*** START OF THE PROJECT GUTENBERG EBOOK 54210 ***
THE
PRINCIPLES OF CHEMISTRY
By D. MENDELÉEFF
TRANSLATED FROM THE RUSSIAN (SIXTH EDITION) BY
GEORGE KAMENSKY, A.R.S.M.
OF THE IMPERIAL MINT, ST PETERSBURG: MEMBER OF THE RUSSIAN
PHYSICO-CHEMICAL SOCIETY
EDITED BY
T. A. LAWSON, B.Sc. PH.D.
EXAMINER IN COAL-TAR PRODUCTS TO THE CITY AND GUILDS OF LONDON
INSTITUTE FELLOW OF THE INSTITUTE OF CHEMISTRY
IN TWO VOLUMES
VOLUME II.
LONGMANS, GREEN, AND CO
39 PATERNOSTER ROW, LONDON
NEW YORK AND BOMBAY
1897
All rights reserved
* * * * *
TABLE III.
_The periodic dependence of the composition of the simplest compounds
and properties of the simple bodies upon the atomic weights of
the elements._
+-------------------------+--------------------------------+
| | |
|Molecular composition of | |
|the higher hydrogen and | Atomic weights of the elements |
|metallo-organic compounds| |
|-------------------------+--------------------------------+
| | |
| | |
|E=CH_{3}, C_{2}H_{5}, &c.| |
| | |
| | |
|[1] [2] [3] [4] | [5] [6] |
| | |
| HH| H 1,005 (mean) |
| | Li 7·02 (Stas) |
| | Be 9·1 (Nilson Pettersson)|
| BE_{3} -- --| B 11·0 (Ramsay Ashton) |
| CH_{4} C_{2}H_{6} | |
| C_{2}H_{4} C_{2}H_{2} | C 12·00 (Roscoe) |
| NH_{3} N_{2}H_{4} --| N 14·04 (Stas) |
| OH_{2} --| O 16 (conventional) |
| FH| F 19·0 (Christiansen) |
| | |
| NaE| Na 23·04 (Stas) |
| MgE_{2} --| Mg 24·3 (Burton) |
| AlE_{3} -- --| Al 27·1 (Mallet) |
|SiH_{4} Si_{2}E_{6} -- --| Si 28·4 (Thorpe Young) |
| PH_{3} P_{2}H_{4} --| P 31·0 (v. d. Plaats) |
| SH_{2} --| S 32·06 (Stas) |
| ClH| Cl 35·45 (Stas) |
| | |
| | K 39·15 (Stas) |
| | Ca 40·0 (Dumas) |
| | Sc 44·0 (Nilson) |
| | Ti 48·1 (Thorpe) |
| | V 51·2 (Roscoe) |
| | Cr 52·1 (Rawson) |
| | Mn 55·1 (Marignac) |
| | Fe 56·0 (Dumas) |
| | Co 58·9 (Zimmermann) |
| | Ni 59·4 (Winkler) |
| | Cu 63.6 (Richards) |
| ZnE_{2} --| Zn 65·3 (Marignac) |
| GaE_{3} -- --| Ga 69·9 (Boisbaudran) |
| GeE_{4} -- -- --| Ge 72·3 (Winkler) |
| AsH_{3} -- --| As 75·0 (Dumas) |
| SeH_{2} --| Se 79·0[A] (Pettersson) |
| BrH| Br 79·95 (Stas) |
| | |
| | Rb 85·5 (Godeffroy) |
| | Sr 87·6 (Dumas) |
| | Y 89 (Clève) |
| | Zr 90·6 (Bailey) |
| | Nb 94 (Marignac) |
| | Mo 96·1 (Maas) |
| | Unknown metal |
| | |
| | Ru 101·7 (Joly) |
| | Rh 102·7 (Seubert) |
| | Pd 106·4 (Keller Smith) |
| | Ag 107·92 (Stas) |
| CdE_{2} --| Cd 112·1 (Lorimer Smith) |
| InE_{3} -- --| In 113·6 (Winkler) |
| SnE_{4} -- -- --| Sn 119·1 (Classen) |
| SbH_{3} -- --| Sb 120·4 (Schneider) |
| TeH_{2} --| Te 125·1 (Brauner) |
| | |
| | Cs 132·7 (Godeffroy) |
| | Ba 137·4 (Richards) |
| | La 138·2 (Brauner) |
| | Ce 140·2 (Brauner) |
| | |
| | Ta 182·7 (Marignac) |
| | W 184·0 (Waddel) |
| | Unknown element. |
| | |
| | Ir 193·3 (Joly) |
| | Pt 196·0 (Dittmar McArthur) |
| | Au 197·5 (Mallet) |
| HgE_{2} --| Hg 200·5 (Erdmann Mar.) |
| TlE_{3} -- --| Tl 204·1 (Crookes) |
| PbE_{4} -- -- --| Pb 206·90 (Stas) |
| BiE_{3} -- --| Bi 208·9 (Classen) |
| | Five unknown elements. |
| | Th 232·4 (Krüss Nilson) |
| | Unknown element. |
| | U 239·3 (Zimmermann) |
+-------------------------+--------------------------------+
+----------------------------------------------------------------------+
| |
| |
| Composition of the saline compounds, X = Cl |
| |
+----------------------------------------------------------------------+
| Br, (NO_{3}), 1/2 O, 1/2 (SO_{4}), OH, (OM) = Z, where M = K, |
| 1/2 Ca, 1/3 Al, &c. |
|Form RX RX_{2} RX_{3} RX_{4} RX_{5} RX_{6} RX_{7} RX_{8}|
|Oxi- R_{2}O RO R_{2}O_{3} RO_{2} R_{2}O_{5} RO_{3} R_{2}O_{7} RO_{4}|
|des |
| [7] [8] [9] [10] [11] [12] [13] [14] |
| |
| X or H_{2}O |
| iX |
| -- BeX_{2} |
| -- -- BX_{3} |
| |
| -- CO -- COZ_{2} |
| N_{2}O NO NOZ NO_2 NO_{2}Z |
| -- OX_{2} |
| FZ |
| |
| NaX |
| -- MgX_{2} |
| -- -- AlX_{3} |
| -- -- -- SiOZ_{2} |
| -- -- PX_{3} -- POZ_{3} |
| -- SX_{2} -- SOZ_{2} -- SO_{2}Z_{2} |
| ClZ -- ClOZ -- ClO_{2}Z -- ClO_{3}Z |
| |
| KX |
| -- CaX_{2} |
| -- -- ScX_{3} |
| -- TiX_{2} TiX_{3} TiX_{4} |
| -- VO VOX -- VOZ_{3} |
| -- CrX_{2} CrX_{3} CrO_{2} -- CrO_{2}Z_{2} |
| -- MnX_{2} MnX_{3} MnO_{2} -- MnO_{2}Z_{2} MnO_{3}Z |
| -- FeX_{2} FeX_{3} -- -- FeO_{2}Z_{2} |
| -- CoX_{2} CoX_{3} CoO_{2} |
| -- NiX_{2} NiX_{3} |
| CuX CuX_{2} |
| -- ZnX_{2} |
| -- -- GaX_{3} |
| -- GeX_{2} -- GeX_{4} |
| -- AsS AsX_{3} AsS_{2} AsO_{2}Z |
| -- -- -- SeOZ_{2} -- SeO_{2}Z_{2} |
| BrZ -- BrOZ -- BrO_{2}Z -- BrO_{3}Z |
| |
| RbX |
| -- SrX_{2} |
| -- -- YX_{3} |
| -- -- -- ZrX_{4} |
| -- -- NbX_{3} -- NbO_{2}Z |
| -- -- MoX_{3} MoX_{4} -- MoO_{2}Z_{2} |
|(eka-manganese, Em = 99). EmO_{3}Z |
| RuO_{4}|
| -- RuX_{2} RuX_{3} RuX_{4} -- RuO_{2}Z_{2} RuO_{3}Z |
| -- RhX_{2} RhX_{3} RhX_{4} -- RhO_{2}Z_{2} |
| PdX PdX_{2} -- PdX_{4} |
| AgX |
| -- CdX_{2} |
| -- InX_{2} InX_{3} |
| -- SnX_{2} -- SnX_{4} |
| -- -- SbX_{3} -- SbO_{2}Z |
| -- -- -- TeOZ_{2} -- TeO_{2}Z_{2} |
| IZ -- IZ_{3} -- IO_{2}Z -- IO_{3}Z |
| |
| CsX |
| -- BaX_{2} |
| -- -- LaX_{3} |
| -- -- CeX_{3} CeX_{4} |
| Little known Di = 142.1 and Yb = 173.2, and over 15 unknown elements.|
| -- -- -- -- TaO_{2}Z |
| -- -- -- WX_{4} -- WO_{2}Z_{2} |
| |
| OsO_{4}|
| -- -- OsX_{3} OsX_{4} -- OsO_{2}Z_{2} -- |
| -- -- IrX_{3} IrX_{4} -- IrO_{2}Z_{2} |
| -- PtX_{2} -- PtX_{4} |
| AuX -- AuX_{3} |
| HgX HgX_{2} |
| TlX -- TlX_{3} |
| -- PbX_{2} -- PbOZ_{2} |
| -- -- BiX_{3} -- BiO_{2}Z |
| |
| -- -- -- ThX_{4} |
| |
| -- -- -- UO_{2} -- UO_{2}X_{2} UO_{4}|
+----------------------------------------------------------------------+
+-------------------------+------------+---------+---------------------+
| | | Lower | Simple bodies |
|Molecular composition of | |hydrogen +-----+-------+-------|
|the higher hydrogen and | Peroxides | com- | Sp. | Sp. |Melting|
|metallo-organic compounds| | pounds | gr | vol. | point |
|-------------------------+------------+---------+-----+-------+-------|
| | | | | | |
| | | | | | |
|E=CH_{3}, C_{2}H_{5}, &c.| | | | | |
| | | | | | |
| | | | | | |
|[1] [2] [3] [4] | [15] | [16] |[17] | [18] | [19] |
| | | | | | |
| HH|H_{2}O_{2} | -- |*0·05| 20 | -250°?|
| | -- | -- | 0·59| 11·9 | 180° |
| | -- | BeH | 1·64| 5·5 | 900°?|
| BE_{3} -- --| -- | -- | 2·5 | 4·4 |1,300°?|
| CH_{4} C_{2}H_{6} | | | | | |
| C_{2}H_{4} C_{2}H_{2} |C_{2}O_{5}* | -- |*1·9 | 6·3 |2,600°?|
| NH_{3} N_{2}H_{4} --|N_{2}O_{6}* | N_{3}H |*0·6 | 23 | -203° |
| OH_{2} --|O_{3} | -- |*0·9 | 18 | -230°?|
| FH| -- | -- |?1·0 | 19 | ? |
| | | | | | |
| NaE|NaO | Na_{2}H | 0·98| 23·5 | 96° |
| MgE_{2} --| -- | MgH | 1·74| 14 | 500° |
| AlE_{3} -- --| -- | -- | 2·6 | 11 | 600° |
|SiH_{4} Si_{2}E_{6} -- --| -- | -- | 2·3 | 12 |1,300°?|
| PH_{3} P_{2}H_{4} --| -- | P_2H | 2·2 | 14 | 44° |
| SH_{2} --|S_{2}O_{7} | -- | 2·07| 15 | 114° |
| ClH| -- | -- |*1·3 | 27 | -75° |
| | | | | | |
| |KO_{2} | K_{2}H | 0·87| 45 | 58° |
| |CaO_{2} | CaH | 1·56| 26 | 800° |
| | -- | -- |?2·5 | ?18 |1,200°?|
| |TiO_{3} | -- | 3·6 | 13 |2,500°?|
| | -- | -- | 5·5 | 9 |3,000°?|
| |Cr_{2}O_{7} | -- | 6·7 | 7·7 |2,000°?|
| | -- | -- | 7·5 | 7·3 |1,500° |
| | -- |Fe_{n}H* | 7·8 | 7·2 |1,450° |
| | -- | -- | 8·6 | 6·8 |1,400° |
| | -- | Ni_{n}H | 8·7 | 6·8 |1,350° |
| |Cu_{2}O_{5}*| CuH | 8·8 | 7·2 |1,054° |
| ZnE_{2} --|ZnO_{2} | -- | 7·1 | 9·2 | 418° |
| GaE_{3} -- --| -- | -- | 5·96| 11·7 | 30° |
| GeE_{4} -- -- --| -- | -- | 5·47| 13·2 | 900° |
| AsH_{3} -- --| -- |As_{4}H* | 5·65| 13·3 | 500° |
| SeH_{2} --| -- | -- | 4·8 | 16 | 217° |
| BrH| -- | -- | 3·1 | 26 | -7° |
| | | | | | |
| |RbO |Rb_{2}H* | 1·5 | 57 | 39° |
| |SrO_{2} | SrH | 2·5 | 35 | 600°?|
| | -- | -- |*3·4 | *26 |1,000°?|
| | -- |Zr_{4n}H*| 4·1 | 22 |1,500°?|
| | -- |Nb_{n}H* | 7·1 | 13 |1,800°?|
| |Mo_{2}O_{7} | -- | 8·6 | 11 |2,200°?|
| | -- | -- | -- | -- | -- |
| | | | | | |
| | -- |Ru_{n}H* |12·2 | 8·4 |2,000°?|
| | -- |Rh_{n}H* |12·1 | 8·6 |1,900°?|
| | -- | Pd_{2}H |11·4 | 8·3 |1,500° |
| |AgO | -- |10·5 | 10·3 | 950° |
| CdE_{2} --|CdO_{2} | -- | 8·6 | 13 | 320° |
| InE_{3} -- --| -- | -- | 7·4 | 14 | 176° |
| SnE_{4} -- -- --|SnO_{3} | -- | 7·2 | 16 | 232° |
| SbH_{3} -- --| -- | -- | 6·7 | 18 | 432° |
| TeH_{2} --| -- | -- | 6·4 | 20 | 455° |
| IH| -- | -- | 4·9 | 26 | 114° |
| | | | | | |
| | -- |Cs_{2}H* | 2·37| 56 | 27° |
| |BaO_{2} | BaH | 3·75| 36 | ? |
| | -- | -- | 6·1 | 23 | ? |
| | -- | -- | 6·6 | 21 | 700°?|
| | | | | | |
| | -- |Ta_{n}H* |10·4 | 18 | ? |
| |W_{2}O_{7} | -- |19·1 | 9·6 |2,600° |
| | | | | | |
| | | | | | |
| | -- | -- |22·5 | 8·5 |2,700°?|
| | -- | Ir_nH* |22·4 | 8·6 |2,000° |
| | -- |Pt_{n}H* |21·4 | 9·2 |1,775° |
| | -- | -- |19·3 | 10 |1,045° |
| HgE_{2} --| -- | -- |13·6 | 15 | -39° |
| TlE_{3} -- --| -- | -- |11·8 | 17 | 294° |
| PbE_{4} -- -- --| -- | -- |11·3 | 18 | 328° |
| BiE_{3} -- --| -- | -- | 9·8 | 21 | 269° |
| | | | | | |
| | -- | -- |11·1 | 21 | ? |
| | | | | | |
| | -- | -- |18·7 | 13 |2,400°?|
+-------------------------+------------+---------+-----+-------+-------+
[A] From analogy there is reason for thinking that the atomic weight
of selenium is really slightly less than 79·0.
Columns 1, 2, 3, and 4 give the molecular composition of the hydrogen
and metallo-organic compounds, exhibiting the most characteristic forms
assumed by the elements. The first column contains only those which
correspond to the form RX_{4}, the second column those of the form
RX_{3}, the third of the form RX_{2}, and the fourth of the form RX, so
that the periodicity stands out clearly (see Column 16).
Column 5 contains the symbols of all the more or less well-known
elements, placed according to the order of the magnitude of their atomic
weights.
Column 6 contains the atomic weights of the elements according to the
most trustworthy determinations. The names of the investigators are given
in parenthesis. The atomic weight of oxygen, taken as 16, forms the basis
upon which these atomic weights were calculated. Some of these have been
recalculated by me on the basis of Stas's most trustworthy data (_see_
Chapter XXIV. and the numbers given by Stas in the table, where they are
taken according to van der Plaats and Thomsen's calculations).
Columns 7-14 contain the composition of the saline compounds of the
elements, placed according to their forms, RX, RX_{2} to RX_{8} (in the
14^{th} column). If the element R has a metallic character like H, Li,
Be, &c., then X represents Cl, NO_{3}, 1/2 SO_{4}, &c., haloid radicles,
or (OH) if a perfect hydrate is formed (alkali, aqueous base), or 1/2 O,
1/2 S, &c. when an anhydrous oxide, sulphide, &c. is formed. For
instance, NaCl, Mg(NO_{3})_{2}, Al_{2}(SO_{4})_{3}, correspond to NaX,
MgX_{2}, and AlX_{3}; so also Na(OH), Mg(OH)_{2}, Al(OH)_{3}, Na_{2}O,
MgO, Al_{2}O_{3}, &c. But if the element, like C or N, be of a metalloid
or acid character, X must be regarded as (OH) in the formation of
hydrates; (OM) in the formation of salts, where M is the equivalent of a
metal, 1/2 O in the formation of an anhydride, and Cl in the formation of
a chloranhydride; and in this case (_i.e._ in the acid compounds) Z is
put in the place of X; for example, the formulæ COZ_{2}, NO_{2}Z,
MNO_{2}Z, FeO_{2}Z_{2}, and IZ_{3} correspond to CO(NaO)_{2} =
Na_{2}CO_{3}, COCl_{2}, CO_{2}, NO_{2}(NaO) = NaNO_{3}, NO_{2}Cl,
NO_{2}(OH) = HNO_{3}; MnO_{3}(OK) = KMnO_{4}, ICl, &c.
The 15th column gives the compositions of the peroxides of the
elements, _taking them as anhydrous_. An asterisk (*) is attached to
those of which the composition has not been well established, and a dash
(--) shows that for a given element no peroxides have yet been obtained.
The peroxides contain more oxygen than the higher saline oxides of the
same elements, are powerfully oxidising, and easily give peroxide of
hydrogen. This latter circumstance necessitates their being referred to
the type of peroxide of hydrogen, if bases and acids are referred to the
type of water (see Chapter XV., Note 7 and 11 bis).
The 16th column gives the composition of the lower hydrogen
compounds like N_{3}H and Na_{2}H. They may often be regarded as alloys
of hydrogen, which is frequently disengaged by them at a comparatively
moderate temperature. They differ greatly in their nature from the
hydrogen compounds given in columns 1-4 (_see_ Note 12).
Column 17 gives the specific gravity of the elements in a solid and
a liquid state. An asterisk (*) is placed by those which can either only
be assumed from analogy (for example, the sp. gr. of fluorine and
hydrogen, which have not been obtained in a liquid state), or which vary
very rapidly with a variation of temperature and pressure (like oxygen
and nitrogen), or physical state (for instance, carbon in passing from
the state of charcoal to graphite and diamond). But as the sp. gr. in
general varies with the temperature, mechanical condition, &c., the
figures given, although chosen from the most trustworthy sources, can
only be regarded as approximate, and not as absolutely true. They clearly
show a certain periodicity; for instance, the sp. gr. diminishes from Al
on both sides (Al, Mg, Na, with decreasing atomic weight; and Al, Si, P,
S, Cl, with increasing atomic weight, it also diminishes on both sides
from Cu, Ru, and Os.)
The same remarks refer to the figures in the 18th column, which
gives the so-called atomic volumes of the simple bodies, or the quotient
of their atomic weight and specific gravity. For Na, K, Rb, and Cs the
atomic volume is greatest among the neighbouring elements. For Ni, Pd,
and Os it is least, and this indicates the periodicity of this property
of the simple bodies.
The last (19th) column gives the melting points of the simple
bodies. Here also a periodicity is seen, i.e. a maximum and minimum value
between which there are intermediate values, as we see, for instance, in
the series Cl, K, Ca, Sc, and Ti, or in the series Cr, Mn, Fe, Co, Ni,
Cu, Zn, Ga, and Ge.
* * * * *
CHAPTER XV
THE GROUPING OF THE ELEMENTS AND THE PERIODIC LAW
It is seen from the examples given in the preceding chapters that the sum
of the data concerning the chemical transformations proper to the
elements (for instance, with respect to the formation of acids, salts,
and other compounds having definite properties) is insufficient for
accurately determining the relationship of the elements, inasmuch as this
may be many-sided. Thus, lithium and barium are in some respects
analogous to sodium and potassium, and in others to magnesium and
calcium. It is evident, therefore, that for a complete judgment it is
necessary to have, not only qualitative, but also quantitative, exact and
measurable, data. When a property can be measured it ceases to be vague,
and becomes quantitative instead of merely qualitative.
Among these measurable properties of the elements, or of their
corresponding compounds, are: (_a_) isomorphism, or the analogy of
crystalline forms; and, connected with it, the power to form crystalline
mixtures which are isomorphous; (_b_) the relation of the volumes of
analogous compounds of the elements; (_c_) the composition of their
saline compounds; and (_d_) the relation of the atomic weights of the
elements. In this chapter we shall briefly consider these four aspects of
the matter, which are exceedingly important for a natural and fruitful
grouping of the elements, facilitating, not only a general acquaintance
with them, but also their detailed study.
Historically the first, and an important and convincing, method for
finding a relationship between the compounds of two different elements is
by _isomorphism_. This conception was introduced into chemistry by
Mitscherlich (in 1820), who demonstrated that the corresponding salts of
arsenic acid, H_{3}AsO_{4}, and phosphoric acid, H_{3}PO_{4}, crystallise
with an equal quantity of water, show an exceedingly close resemblance in
crystalline form (as regards the angles of their faces and axes), and are
able to crystallise together from solutions, forming crystals containing
a mixture of the isomorphous compounds. Isomorphous substances are those
which, with an equal number of atoms in their molecules, present an
analogy in their chemical reactions, a close resemblance in their
properties, and a similar or very nearly similar crystalline form: they
often contain certain elements in common, from which it is to be
concluded that the remaining elements (as in the preceding example of As
and P) are analogous to each other. And inasmuch as crystalline forms are
capable of exact measurement, the external form, or the relation of the
molecules which causes their grouping into a crystalline form, is
evidently as great a help in judging of the internal forces acting
between the atoms as a comparison of reactions, vapour densities, and
other like relations. We have already seen examples of this in the
preceding pages.[1] It will be sufficient to call to mind that the
compounds of the alkali metals with the halogens RX, in a crystalline
form, all belong to the cubic system and crystallise in octahedra or
cubes--for example, sodium chloride, potassium chloride, potassium
iodide, rubidium chloride, &c. The nitrates of rubidium and cæsium appear
in anhydrous crystals of the same form as potassium nitrate. The
carbonates of the metals of the alkaline earths are isomorphous with
calcium carbonate--that is, they either appear in forms like calc spar or
in the rhombic system in crystals analogous to aragonite.[1 bis]
Furthermore, sodium nitrate crystallises in rhombohedra, closely
resembling the rhombohedra of calc spar (calcium carbonate), CaCO_{3},
whilst potassium nitrate appears in the same form as aragonite, CaCO_{3},
and the number of atoms in both kinds of salts is the same: they all
contain one atom of a metal (K, Na, Ca), one atom of a non-metal (C, N),
and three atoms of oxygen. The analogy of form evidently coincides with
an analogy of atomic composition. But, as we have learnt from the
previous description of these salts, there is not any close resemblance
in their properties. It is evident that calcium carbonate approaches more
nearly to magnesium carbonate than to sodium nitrate, although their
crystalline forms are all equally alike. Isomorphous substances which are
perfectly analogous to each other are not only characterised by a close
resemblance of form (homeomorphism), but also by the faculty of entering
into analogous reactions, which is not the case with RNO_{3} and RCO_{3}.
The most important and direct method of recognising perfect
isomorphism--that is, the absolute analogy of two compounds--is given by
that property of analogous compounds of separating from solutions _in
homogeneous crystals, containing the most varied proportions_ of the
analogous substances which enter into their composition. These quantities
do not seem to be in dependence on the molecular or atomic weights, and
if they are governed by any laws they must be analogous to those which
apply to indefinite chemical compounds.[2] This will be clear from the
following examples. Potassium chloride and potassium nitrate are not
isomorphous with each other, and are in an atomic sense composed in a
different manner. If these salts be mixed in a solution and the solution
be evaporated, independent crystals of the two salts will separate, each
in that crystalline form which is proper to it. The crystals will not
contain a mixture of the two salts. But if we mix the solutions of two
isomorphous salts together, then, under certain circumstances, crystals
will be obtained which contain both these substances. However, this
cannot be taken as an absolute rule, for if we take a solution saturated
at a high temperature with a mixture of potassium and sodium chlorides,
then on evaporation sodium chloride only will separate, and on cooling
only potassium chloride. The first will contain very little potassium
chloride, and the latter very little sodium chloride.[3] But if we take,
for example, a mixture of solutions of magnesium sulphate and zinc
sulphate, they cannot be separated from each other by evaporating the
mixture, notwithstanding the rather considerable difference in the
solubility of these salts. Again, the isomorphous salts, magnesium
carbonate, and calcium carbonate are found together--that is, in one
crystal--in nature. The angle of the rhombohedron of these magnesia-lime
spars is intermediate between the angles proper to the two spars
individually (for calcium carbonate, the angle of the rhombohedron is
105° 8´; magnesium carbonate, 107° 30´; CaMg(CO_{3})_{2}, 106° 10´).
Certain of these _isomorphous mixtures_ of calc and magnesia spars appear
in well-formed crystals, and in this case there not unfrequently exists a
simple molecular proportion of strictly definite chemical combination
between the component salts--for instance, CaCO_{3},MgCO_{3}--whilst in
other cases, especially in the absence of distinct crystallisation (in
dolomites), no such simple molecular proportion is observable: this is
also the case in many artificially prepared isomorphous mixtures. The
microscopical and crystallo-optical researches of Professor Inostrantzoff
and others show that in many cases there is really a mechanical, although
microscopically minute, juxtaposition in one whole of the heterogeneous
crystals of calcium carbonate (double refracting) and of the compound
CaMgC_{2}O_{6}. If we suppose the adjacent parts to be microscopically
small (on the basis of the researches of Mallard, Weruboff, and others),
we obtain an idea of isomorphous mixtures. A formula of the following
kind is given to isomorphous mixtures: for instance, for spars, RCO_{3},
where R = Mg, Ca, and where it may be Fe,Mn ..., &c. This means that the
Ca is partially replaced by Mg or another metal. Alums form a common
example of the separation of isomorphous mixtures from solutions. They
are double sulphates (or seleniates) of alumina (or oxides isomorphous
with it) and the alkalis, which crystallise in well-formed crystals. If
aluminium sulphate be mixed with potassium sulphate, an alum separates,
having the composition KAlS_{2}O_{8},12H_{2}O. If sodium sulphate or
ammonium sulphate, or rubidium (or thallium) sulphate be used, we obtain
alums having the composition RAlS_{2}O_{8},12H_{2}O. Not only do they all
crystallise in the cubic system, but they also contain an equal atomic
quantity of water of crystallisation (12H_{2}O). Besides which, if we mix
solutions of the potassium and ammonium (NH_{4}AlS_{2}O_{8},12H_{2}O)
alums together, then the crystals which separate will contain various
proportions of the alkalis taken, and separate crystals of the alums of
one or the other kind will not be obtained, but each separate crystal
will contain both potassium and ammonium. Nor is this all; if we take a
crystal of a potassium alum and immerse it in a solution capable of
yielding ammonia alum, the crystal of the potash alum will continue to
grow and increase in size in this solution--that is, a layer of the
ammonia or other alum will deposit itself upon the planes bounding the
crystal of the potash alum. This is very distinctly seen if a colourless
crystal of a common alum be immersed in a saturated violet solution of
chrome alum, KCrS_{2}O_{8},12H_{2}O, which then deposits itself in a
violet layer over the colourless crystal of the alumina alum, as was
observed even before Mitscherlich noticed it. If this crystal be then
immersed in a solution of an alumina alum, a layer of this salt will form
over the layer of chrome alum, so that one alum is able to incite the
growth of the other. If the deposition proceed simultaneously, the
resultant intermixture may be minute and inseparable, but its nature is
understood from the preceding experiments; the attractive force of
crystallisation of isomorphous substances is so nearly equal that the
attractive power of an isomorphous substance induces a crystalline
superstructure exactly the same as would be produced by the attractive
force of like crystalline particles. From this it is evident that one
isomorphous substance may _induce the crystallisation_[4] of another.
Such a phenomenon explains, on the one hand, the aggregation of different
isomorphous substances in one crystal, whilst, on the other hand, it
serves as a most exact indication of the nearness both of the molecular
composition of isomorphous substances and of those forces which are
proper to the elements which distinguish the isomorphous substances.
Thus, for example, ferrous sulphate or green vitriol crystallises in the
monoclinic system and contains seven molecules of water,
FeSO_{4},7H_{2}O, whilst copper vitriol crystallises with five molecules
of water in the triclinic system, CuSO_{4},5H_{2}O; nevertheless, it may
be easily proved that both salts are perfectly isomorphous; that they are
able to appear in identically the same forms and with an equal molecular
amount of water. For instance, Marignac, by evaporating a mixture of
sulphuric acid and ferrous sulphate under the receiver of an air-pump,
first obtained crystals of the hepta-hydrated salt, and then of the
penta-hydrated salt FeSO_{4},5H_{2}O, which were perfectly similar to the
crystals of copper sulphate. Furthermore, Lecoq de Boisbaudran, by
immersing crystals of FeSO_{4},7H_{2}O in a supersaturated solution of
copper sulphate, caused the latter to deposit in the same form as ferrous
sulphate, in crystals of the monoclinic system, CuSO_{4},7H_{2}O.
[1] For instance the analogy of the sulphates of K, Rb, and Cs (Chapter
XIII., Note 1).
[1 bis] The crystalline forms of aragonite, strontianite, and witherite
belong to the rhombic system; the angle of the prism of CaCO_{3} is
116° 10´, of SrCO_{3} 117° 19´, and of BaCO_{3} 118° 30´. On the
other hand the crystalline forms of calc spar, magnesite, and
calamine, which resemble each other quite as closely, belong to the
rhombohedral system, with the angle of the rhombohedra for CaCO_{3}
105° 8´, MgCO_{3} 107° 10´, and ZnCO_{3} 107° 40´. From this
comparison it is at once evident that zinc is more closely allied
to magnesium than magnesium to calcium.
[2] Solutions furnish the commonest examples of indefinite chemical
compounds. But the isomorphous mixtures which are so common among
the crystalline compounds of silica forming the crust of the earth,
as well as alloys, which are so important in the application of
metals to the arts, are also instances of indefinite compounds. And
if in Chapter I., and in many other portions of this work, it has
been necessary to admit the presence of definite compounds (in a
state of dissociation) in solutions, the same applies with even
greater force to isomorphous mixtures and alloys. For this reason
in many places in this work I refer to facts which compel us to
recognise the existence of definite chemical compounds in all
isomorphous mixtures and alloys. This view of mine (which dates
from the sixties) upon isomorphous mixtures finds a particularly
clear confirmation in B. Roozeboom's researches (1892) upon the
solubility and crystallising capacity of mixtures of the chlorates
of potassium and thallium, KClO_{3} and TlClO_{3}. He showed that
when a solution contains different amounts of these salts, it
deposits crystals containing either an excess of the first salt,
from 98 p.c. to 100 p.c., or an excess of the second salt, from
63·7 to 100 p.c.; that is, in the crystalline form, either the
first salt saturates the second or the second the first, just as in
the solution of ether in water (Chapter I.); moreover, the
solubility of the mixtures containing 36·3 and 98 p.c. KClO_{3} is
similar, just as the vapour tension of a saturated solution of
water in ether is equal to that of a saturated solution of ether in
water (Chapter I., Note 47). But just as there are solutions
miscible in all proportions, so also certain isomorphous bodies can
be present in crystals in all possible proportions of their
component parts. Van 't Hoff calls such systems 'solid solutions.'
These views were subsequently elaborated by Nernst (1892), and Witt
(1891) applied them in explaining the phenomena observed in the
coloration of tissues.
[3] The cause of the difference which is observed in different
compounds of the same type, with respect to their property of
forming isomorphous mixtures, must not be looked for in the
difference of their volumetric composition, as many investigators,
including Kopp, affirm. The molecular volumes (found by dividing
the molecular weight by the density) of those isomorphous
substances which do give intermixtures are not nearer to each other
than the volumes of those which do not give mixtures; for example,
for magnesium carbonate the combining weight is 84, density 3·06,
and volume therefore 27; for calcium carbonate in the form of calc
spar the volume is 37, and in the form of aragonite 33; for
strontium carbonate 41, for barium carbonate 46; that is, the
volume of these closely allied isomorphous substances increases
with the combining weight. The same is observed if we compare
sodium chloride (molecular volume = 27) with potassium chloride
(volume = 37), or sodium sulphate (volume = 55) with potassium
sulphate (volume = 66), or sodium nitrate 39 with potassium nitrate
48, although the latter are less capable of giving isomorphous
mixtures than the former. It is evident that the cause of
isomorphism cannot be explained by an approximation in molecular
volumes. It is more likely that, given a similarity in form and
composition, the faculty to give isomorphous mixtures is connected
with the laws and degree of solubility.
[4] A phenomenon of a similar kind is shown for magnesium sulphate in
Note 27 of the last chapter. In the same example we see what a
complication the phenomena of dimorphism may introduce when the
forms of analogous compounds are compared.
Hence it is evident that isomorphism--that is, the analogy of forms and
the property of inducing crystallisation--may serve as a means for the
discovery of analogies in molecular composition. We will take an example
in order to render this clear. If, instead of aluminium sulphate, we add
magnesium sulphate to potassium sulphate, then, on evaporating the
solution, the double salt K_{2}MgS_{2}O_{8},6H_{2}O (Chapter XIV., Note
28) separates instead of an alum, and the ratio of the component parts
(in alums one atom of potassium per 2SO_{4}, and here two atoms) and the
amount of water of crystallisation (in alums 12, and here 6 equivalents
per 2SO_{4}) are quite different; nor is this double salt in any way
isomorphous with the alums, nor capable of forming an isomorphous
crystalline mixture with them, nor does the one salt provoke the
crystallisation of the other. From this we must conclude that although
alumina and magnesia, or aluminium and magnesium, resemble each other,
they are not isomorphous, and that although they give partially similar
double salts, these salts are not analogous to each other. And this is
expressed in their chemical formulæ by the fact that the number of atoms
in alumina or aluminium oxide, Al_{2}O_{3}, is different from the number
in magnesia, MgO. Aluminium is trivalent and magnesium bivalent. Thus,
having obtained a double salt from a given metal, it is possible to judge
of the analogy of the given metal with aluminium or with magnesium, or of
the absence of such an analogy, from the composition and form of this
salt. Thus zinc, for example, does not form alums, but forms a double
salt with potassium sulphate, which has a composition exactly like that
of the corresponding salt of magnesium. It is often possible to
distinguish the bivalent metals analogous to magnesium or calcium from
the trivalent metals, like aluminium, by such a method. Furthermore, the
specific heat and vapour density serve as guides. There are also indirect
proofs. Thus iron gives ferrous compounds, FeX_{2}, which are isomorphous
with the compounds of magnesium, and ferric compounds, FeX_{3}, which are
isomorphous with the compounds of aluminium; in this instance the
relative composition is directly determined by analysis, because, for a
given amount of iron, FeCl_{2} only contains two-thirds of the amount of
chlorine which occurs in FeCl_{3}, and the composition of the
corresponding oxygen compounds, _i.e._ of ferrous oxide, FeO, and ferric
oxide, Fe_{2}O_{3}, clearly indicates the analogy of the ferrous oxide
with MgO and of the ferric oxide with Al_{2}O_{3}.
Thus in the building up of similar molecules in crystalline forms we
see one of the numerous means for judging of the internal world of
molecules and atoms, and one of the weapons for conquests in the
invisible world of molecular mechanics which forms the main object of
physico-chemical knowledge. This method[5] has more than once been
employed for discovering the analogy of elements and of their compounds;
and as crystals are measurable, and the capacity to form crystalline
mixtures can be experimentally verified, this method is a numerical and
measurable one, and in no sense arbitrary.
[5] The property of solids of occurring in regular crystalline
forms--the occurrence of many substances in the earth's crust in
these forms--and those geometrical and simple laws which govern the
formation of crystals long ago attracted the attention of the
naturalist to crystals. The crystalline form is, without doubt, the
expression of the relation in which the atoms occur in the
molecules, and in which the molecules occur in the mass, of a
substance. Crystallisation is determined by the distribution of the
molecules along the direction of greatest cohesion, and therefore
those forces must take part in the crystalline distribution of
matter which act between the molecules; and, as they depend on the
forces binding the atoms together in the molecules, a very close
connection must exist between the atomic composition and the
distribution of the atoms in the molecule on the one hand, and the
crystalline form of a substance on the other hand; and hence an
insight into the composition may be arrived at from the crystalline
form. Such is the elementary and _a priori_ idea which lies at the
base of all researches into _the connection between composition and
crystalline form_. Haüy in 1811 established the following
fundamental law, which has been worked out by later investigators:
That the fundamental crystalline form for a given chemical compound
is constant (only the combinations vary), and that with a change of
composition the crystalline form also changes, naturally with the
exception of such limiting forms as the cube, regular octahedron,
&c., which may belong to various substances of the regular system.
The fundamental form is determined by the angles of certain
fundamental geometric forms (prisms, pyramids, rhombohedra), or the
ratio of the crystalline axes, and is connected with the optical
and many other properties of crystals. Since the establishment of
this law the description of definite compounds in a solid state is
accompanied by a description (measurement) of its crystals, which
forms an invariable, definite, and measurable character. The most
important epochs in the further history of this question were made
by the following discoveries:--Klaproth, Vauquelin, and others
showed that aragonite has the same composition as calc spar, whilst
the former belongs to the rhombic and the latter to the hexagonal
system. Haüy at first considered that the composition, and after
that the arrangement, of the atoms in the molecules was different.
This is dimorphism (_see_ Chapter XIV., Note 46). Beudant,
Frankenheim, Laurent, and others found that the forms of the two
nitres, KNO_{3} and NaNO_{3}, exactly correspond with the forms of
aragonite and calc spar; that they are able, moreover, to pass from
one form into another; and that the difference of the forms is
accompanied by a small alteration of the angles, for the angle of
the prisms of potassium nitrate and aragonite is 119°, and of
sodium nitrate and calc spar, 120°; and therefore dimorphism, or
the crystallisation of one substance in different forms, does not
necessarily imply a great difference in the distribution of the
molecules, although some difference clearly exists. The researches
of Mitscherlich (1822) on the dimorphism of sulphur confirmed this
conclusion, although it cannot yet be affirmed that in dimorphism
the arrangement of the atoms remains unaltered, and that only the
molecules are distributed differently. Leblanc, Berthier,
Wollaston, and others already knew that many substances of
different composition appear in the same forms, and crystallise
together in one crystal. Gay-Lussac (1816) showed that crystals of
potash alum continue to grow in a solution of ammonia alum. Beudant
(1817) explained this phenomenon as the _assimilation_ of a foreign
substance by a substance having a great force of crystallisation,
which he illustrated by many natural and artificial examples. But
Mitscherlich, and afterwards Berzelius and Henry Rose and others,
showed that such an assimilation only exists with a similarity or
approximate similarity of the forms of the individual substances
and with a certain degree of chemical analogy. Thus was established
the idea of _isomorphism_ as an analogy of forms by reason of a
resemblance of atomic composition, and by it was explained the
variability of the composition of a number of minerals as
isomorphous mixtures. Thus all the garnets are expressed by the
general formula: (RO)_{3}M_{2}O_{3}(SiO_{2})_{3}, where R = Ca, Mg,
Fe, Mn, and M = Fe, Al, and where we may have either R and M
separately, or their equivalent compounds, or their mixtures in all
possible proportions.
But other facts, which render the correlation of form and
composition still more complex, have accumulated side by side with
a mass of data which may be accounted for by admitting the
conceptions of isomorphism and dimorphism. Foremost among the
former stand the phenomena of _homeomorphism_--that is, a nearness
of forms with a difference of composition--and then the cases of
polymorphism and hemimorphism--that is, a nearness of the
fundamental forms or only of certain angles for substances which
are near or analogous in their composition. Instances of
homeomorphism are very numerous. Many of these, however, may be
reduced to a resemblance of atomic composition, although they do
not correspond to an isomorphism of the component elements; for
example, CdS (greenockite) and AgI, CaCO_{3} (aragonite) and
KNO_{3}, CaCO_{3} (calc spar) and NaNO_{3}, BaSO_{4} (heavy spar),
KMnO_{4} (potassium permanganate), and KClO_{4} (potassium
perchlorate), Al_{2}O_{3} (corundum) and FeTiO_{3} (titanic iron
ore), FeS_{2} (marcasite, rhombic system) and FeSAs (arsenical
pyrites), NiS and NiAs, &c. But besides these instances there are
homeomorphous substances with an absolute dissimilarity of
composition. Many such instances were pointed out by Dana.
Cinnabar, HgS, and susannite, PbSO_{4}3PbCO_{3} appear in very
analogous crystalline forms; the acid potassium sulphate
crystallises in the monoclinic system in crystals analogous to
felspar, KAlSi_{3}O_{8}; glauberite, Na_{2}Ca(SO_{4})_{2}, augite,
RSiO_{3} (R = Ca, Mg), sodium carbonate, Na_{2}CO_{3},10H_{2}O,
Glauber's salt, Na_{2}SO_{4},10H_{2}O, and borax,
Na_{2}BrO_{7},10H_{2}O, not only belong to the same system
(monoclinic), but exhibit an analogy of combinations and a nearness
of corresponding angles. These and many other similar cases might
appear to be perfectly arbitrary (especially as a _nearness_ of
angles and fundamental forms is a relative idea) were there not
other cases where a resemblance of properties and a distinct
relation in the variation of composition is connected with a
resemblance of form. Thus, for example, alumina, Al_{2}O_{3}, and
water, H_{2}O, are frequently found in many pyroxenes and
amphiboles which only contain silica and magnesia (MgO, CaO, FeO,
MnO). Scheerer and Hermann, and many others, endeavoured to explain
such instances by _polymetric isomorphism_, stating that MgO may be
replaced by 3H_{2}O (for example, olivine and serpentine), SiO_{2}
by Al_{2}O_{3} (in the amphiboles, talcs), and so on. A certain
number of the instances of this order are subject to doubt, because
many of the natural minerals which served as the basis for the
establishment of polymeric isomorphism in all probability no longer
present their original composition, but one which has been altered
under the influence of solutions which have come into contact with
them; they therefore belong to the class of _pseudomorphs_, or
false crystals. There is, however, no doubt of the existence of a
whole series of natural and artificial homeomorphs, which differ
from each other by atomic amounts of water, silica, and some other
component parts. Thus, Thomsen (1874) showed a very striking
instance. The metallic chlorides, RCl_{2}, often crystallise with
water, and they do not then contain less than one molecule of water
per atom of chlorine. The most familiar representative of the order
RCl_{2},2H_{2}O is BaCl_{2},2H_{2}O, which crystallises in the
rhombic system. Barium bromide, BaBr_{2},2H_{2}O, and copper
chloride, CuCl_{2},2H_{2}O, have nearly the same forms: potassium
iodate, KIO_{4}; potassium chlorate, KClO_{4}; potassium
permanganate, KMnO_{4}; barium sulphate, BaSO_{4}; calcium
sulphate, CaSO_{4}; sodium sulphate, Na_{2}SO_{4}; barium formate,
BaC_{2}H_{2}O_{4}, and others have almost the same crystalline form
(of the rhombic system). Parallel with this series is that of the
metallic chlorides containing RCl_{2},4H_{2}O, of the sulphates of
the composition RSO_{4},2H_{2}O, and the formates
RC_{2}H_{2}O_{4},2H_{2}O. These compounds belong to the monoclinic
system, have a close resemblance of form, and differ from the first
series by containing two more molecules of water. The addition of
two more molecules of water in all the above series also gives
forms of the monoclinic system closely resembling each other; for
example, NiCl_{2},6H_{2}O and MnSO_{4},4H_{2}O. Hence we see that
not only is RCl_{2},2H_{2}O analogous in form to RSO_{4} and
RC_{2}H_{2}O_{4}, but that their compounds with 2H_{2}O and with
4H_{2}O also exhibit closely analogous forms. From these examples
it is evident that the conditions which determine a given form may
be repeated not only in the presence of an isomorphous
exchange--that is, with an equal number of atoms in the
molecule--but also in the presence of an unequal number when there
are peculiar and as yet ungeneralised relations in composition.
Thus ZnO and Al_{2}O_{3} exhibit a close analogy of form. Both
oxides belong to the rhombohedral system, and the angle between the
pyramid and the terminal plane of the first is 118° 7´, and of the
second 118° 49´. Alumina, Al_{2}O_{3}, is also analogous in form to
SiO_{2}, and we shall see that these analogies of form are
conjoined with a certain analogy in properties. It is not
surprising, therefore, that in the complex molecule of a siliceous
compound it is sometimes possible to replace SiO_{2} by means of
Al_{2}O_{3}, as Scheerer admits. The oxides Cu_{2}O, MgO, NiO,
Fe_{3}O_{4}, CeO_{2}, crystallise in the regular system, although
they are of very different atomic structure. Marignac demonstrated
the perfect analogy of the forms of K_{2}ZrF_{6} and CaCO_{3}, and
the former is even dimorphous, like the calcium carbonate. The same
salt is isomorphous with R_{2}NbOF_{5} and R_{2}WO_{2}F_{4}, where
R is an alkali metal. There is an equivalency between CaCO_{3} and
K_{2}ZrF_{6}, because K_{2} is equivalent to Ca, C to Zr, and F_{6}
to O_{3}, and with the isomorphism of the other two salts we find
besides an equal contents of the alkali metal--an equal number of
atoms on the one hand and an analogy to the properties of
K_{2}ZrF_{6} on the other. The long-known isomorphism of the
corresponding compounds of potassium and ammonium, KX and NH_{4}X,
may be taken as the simplest example of the fact that an analogy of
form shows itself with an analogy of chemical reaction even without
an equality in atomic composition. Therefore the ultimate progress
of the entire doctrine of the correlation of composition and
crystalline forms will only be arrived at with the accumulation of
a sufficient number of facts collected on a plan corresponding with
the problems which here present themselves. The first steps have
already been made. The researches of the Geneva _savant_, Marignac,
on the crystalline form and composition of many of the double
fluorides, and the work of Wyruboff on the ferricyanides and other
compounds, are particularly important in this respect. It is
already evident that, with a definite change of composition,
certain angles remain constant, notwithstanding that others are
subject to alteration. Such an instance of the relation of forms
was observed by Laurent, and named by him _hemimorphism_ (an
anomalous term) when the analogy is limited to certain angles, and
_paramorphism_ when the forms in general approach each other, but
belong to different systems. So, for example, the angle of the
planes of a rhombohedron may be greater or less than 90°, and
therefore such acute and obtuse rhombohedra may closely approximate
to the cube. Hausmannite, Mn_{3}O_{4}, belongs to the tetragonal
system, and the planes of its pyramid are inclined at an angle of
about 118°, whilst magnetic iron ore, Fe_{3}O_{4}, which resembles
hausmannite in many respects, appears in regular octahedra--that
is, the pyramidal planes are inclined at an angle of 109° 28´. This
is an example of paramorphism; the systems are different, the
compositions are analogous, and there is a certain resemblance in
form. Hemimorphism has been found in many instances of saline and
other substitutions. Thus, Laurent demonstrated, and Hintze
confirmed (1873), that naphthalene derivatives of analogous
composition are hemimorphous. Nicklès (1849) showed that in
ethylene sulphate the angle of the prism is 125° 26´, and in the
nitrate of the same radicle 126° 95´. The angle of the prism of
methylamine oxalate is 131° 20´, and of fluoride, which is very
different in composition from the former, the angle is 132°. Groth
(1870) endeavoured to indicate in general what kinds of change of
form proceed with the substitution of hydrogen by various other
elements and groups, and he observed a regularity which he termed
_morphotropy_. The following examples show that morphotropy recalls
the hemimorphism of Laurent. Benzene, C_{6}H_{6}, rhombic system,
ratio of the axes 0·891 : 1 : 0·799. Phenol, C_{6}H_{5}(OH), and
resorcinol, C_{6}H_{4}(OH)_{2}, also rhombic system, but the ratio
of one axis is changed--thus, in resorcinol, 0·910 : 1 : 0·540;
that is, a portion of the crystalline structure in one direction is
the same, but in the other direction it is changed, whilst in the
rhombic system dinitrophenol, C_{6}H_{3}(NO_{2})_{2}(OH) =
O·833 : 1 : 0·753; trinitrophenol (picric acid),
C_{6}H_{2}(NO)_{3}(OH) = 0·937 : 1 : 0·974; and the potassium salt
= 0·942 : 1 : 1·354. Here the ratio of the first axis is
preserved--that is, certain angles remain constant, and the
chemical proximity of the composition of these bodies is undoubted.
Laurent compares hemimorphism with architectural style. Thus,
Gothic cathedrals differ in many respects, but there is an analogy
expressed both in the sum total of their common relations and in
certain details--for example, in the windows. It is evident that we
may expect many fruitful results for molecular mechanics (which
forms a problem common to many provinces of natural science) from
the further elaboration of the data concerning those variations
which take place in crystalline form when the composition of a
substance is subjected to a known change, and therefore I consider
it useful to point out to the student of science seeking for matter
for independent scientific research this vast field for work which
is presented by the correlation of form and composition. The
geometrical regularity and varied beauty of crystalline forms offer
no small attraction to research of this kind.
The regularity and simplicity expressed by the exact laws of crystalline
form repeat themselves in the aggregation of the atoms to form molecules.
Here, as there, there are but few forms which are essentially different,
and their apparent diversity reduces itself to a few fundamental
differences of type. There the molecules aggregate themselves into
crystalline forms; here, the atoms aggregate themselves into molecular
forms or into _the types of compounds_. In both cases the fundamental
crystalline or molecular forms are liable to variations, conjunctions,
and combinations. If we know that potassium gives compounds of the
fundamental type KX, where X is a univalent element (which combines with
one atom of hydrogen, and is, according to the law of substitution, able
to replace it), then we know the composition of its compounds: K_{2}O,
KHO, KCl, NH_{2}K, KNO_{3}, K_{2}SO_{4}, KHSO_{4},
K_{2}Mg(SO_{4})_{2},6H_{2}O, &c. All the possible derivative crystalline
forms are not known. So also all the atomic combinations are not known
for every element. Thus in the case of potassium, KCH_{3}, K_{3}P,
K_{2}Pt, and other like compounds which exist for hydrogen or chlorine,
are unknown.
Only a few fundamental types exist for the building up of atoms into
molecules, and the majority of them are already known to us. If X stand
for a univalent element, and R for an element combined with it, then
eight atomic types may be observed:--
RX, RX_{2}, RX_{3}, RX_{4}, RX_{5}, RX_{6}, RX_{7}, RX_{8}.
Let X be chlorine or hydrogen. Then as examples of the first type we
have: H_{2}, Cl_{2}, HCl, KCl, NaCl, &c. The compounds of oxygen or
calcium may serve as examples of the type RX_{2}: OH_{2}, OCl_{2}, OHCl,
CaO, Ca(OH)_{2}, CaCl_{2}, &c. For the third type RX_{3} we know the
representative NH_{3} and the corresponding compounds N_{2}O_{3}, NO(OH),
NO(OK), PCl_{3}, P_{2}O_{3}, PH_{3}, SbH_{3}, Sb_{2}O_{3}, B_{2}O_{3},
BCl_{3}, Al_{2}O_{3}, &c. The type RX_{4} is known among the hydrogen
compounds. Marsh gas, CH_{4}, and its corresponding saturated
hydrocarbons, C_{_n_}H_{2_n_ + 2}, are the best representatives. Also
CH_{3}Cl, CCl_{4}, SiCl_{4}, SnCl_{4}, SnO_{2}, CO_{2}, SiO_{2}, and a
whole series of other compounds come under this class. The type RX_{5} is
also already familiar to us, but there are no purely hydrogen compounds
among its representatives. Sal-ammoniac, NH_{4}Cl, and the corresponding
NH_{4}(OH), NO_{2}(OH), ClO_{2}(OK), as well as PCl_{5}, POCl_{3}, &c.,
are representatives of this type. In the higher types also there are no
hydrogen compounds, but in the type RX_{6} there is the chlorine compound
WCl_{6}. However, there are many oxygen compounds, and among them SO_{3}
is the best known representative. To this class also belong
SO_{2}(OH)_{2}, SO_{2}Cl_{2}, SO_{2}(OH)Cl, CrO_{3}, &c., all of an acid
character. Of the higher types there are in general only oxygen and acid
representatives. The type RX_{7} we know in perchloric acid, ClO_{3}(OH),
and potassium permanganate, MnO_{3}(OK), is also a member. The type
RX_{8} in a free state is very rare; osmic anhydride, OsO_{4}, is the
best known representative of it.[6]
[6] The still more complex combinations--which are so clearly expressed
in the crystallo-hydrates, double salts, and similar
compounds--although they may be regarded as independent, are,
however, most easily understood with our present knowledge as
aggregations of whole molecules to which there are no corresponding
double compounds, containing one atom of an element R and many
atoms of other elements RX_{_n_}. The above types embrace all cases
of direct combinations of atoms, and the formula MgSO_{4},7H_{2}O
cannot, without violating known facts, be directly deduced from the
types MgX_{_n_} or SX_{_n_}, whilst the formula MgSO_{4}
corresponds both with the type of the magnesium compounds MgX_{2}
and with the type of the sulphur compounds SO_{2}X_{2}, or in
general SX_{6}, where X_{2} is replaced by (OH)_{2}, with the
substitution in this case of H_{2} by the atom Mg, which always
replaces H_{2}. However, it must be remarked that the sodium
crystallo-hydrates often contain 10H_{2}O, the magnesium
crystallo-hydrates 6 and 7H_{2}O, and that the type PtM_{2}X_{6} is
proper to the double salts of platinum, &c. With the further
development of our knowledge concerning crystallo-hydrates, double
salts, alloys, solutions, &c., in the _chemical sense_ of feeble
compounds (that is, such as are easily destroyed by feeble chemical
influences) it will probably be possible to arrive at a perfect
generalisation for them. For a long time these subjects were only
studied by the way or by chance; our knowledge of them is
accidental and destitute of system, and therefore it is impossible
to expect as yet any generalisation as to their nature. The days of
Gerhardt are not long past when only three types were recognised:
RX, RX_{2}, and RX_{3}; the type RX_{4} was afterwards added (by
Cooper, Kekulé, Butleroff, and others), mainly for the purpose of
generalising the data respecting the carbon compounds. And indeed
many are still satisfied with these types, and derive the higher
types from them; for instance, RX_{5} from RX_{3}--as, for example,
POCl_{3} from PCl_{3}, considering the oxygen to be bound both to
the chlorine (as in HClO) and to the phosphorus. But the time has
now arrived when it is clearly seen that the forms RX, RX_{2},
RX_{3}, and RX_{4} do not exhaust the whole variety of phenomena.
The revolution became evident when Würtz showed that PCl_{5} is not
a compound of PCl_{3} + Cl_{2} (although it may decompose into
them), but a whole molecule capable of passing into vapour, PCl_{5}
like PF_{5} and SiF_{4}. The time for the recognition of types even
higher than RX_{8} is in my opinion in the future; that it will
come, we can already see in the fact that oxalic acid,
C_{2}H_{2}O_{4}, gives a crystallo-hydrate with 2H_{2}O; but it may
be referred to the type CH_{4}, or rather to the type of ethane,
C_{2}H_{6}, in which all the atoms of hydrogen are replaced by
hydroxyl, C_{2}H_{2}O_{4}2H_{2}O = C_{2}(OH)_{6} (_see_ Chapter
XXII., Note 35).
The four lower types RX, RX_{2}, RX_{3}, and RX_{4} are met with in
compounds of the elements R with chlorine and oxygen, and also in their
compounds with hydrogen, whilst the four higher types only appear for
such acid compounds as are formed by chlorine, oxygen, and similar
elements.
Among the oxygen compounds the _saline oxides_ which are capable of
forming salts either through the function of a base or through the
function of an acid anhydride attract the greatest interest in every
respect. Certain elements, like calcium and magnesium, only give one
saline oxide--for example, MgO, corresponding with the type MgX_{2}. But
the majority of the elements appear in several such forms. Thus copper
gives CuX and CuX_{2}, or Cu_{2}O and CuO. If an element R gives a higher
type RX_{_n_}, then there often also exist, as if by symmetry, lower
types, RX_{_n_-2}, RX_{_n_-4}, and in general such types as differ from
RX_{_n_} by an even number of X. Thus in the case of sulphur the types
SX_{2}, SX_{4}, and SX_{6} are known--for example SH_{2}, SO_{2}, and
SO_{3}. The last type is the highest, SX_{6}. The types SX_{5} and SX_{3}
do not exist. But even and uneven types sometimes appear for one and the
same element. Thus the types RX and RX_{2} are known for copper and
mercury.
Among the _saline_ oxides only the _eight types_ enumerated below are
known to exist. They determine the possible formulæ of the compounds of
the elements, if it be taken into consideration that an element which
gives a certain type of combination may also give lower types. For this
reason the rare type of the _suboxides_ or quaternary oxides R_{4}O (for
instance, Ag_{4}O, Ag_{2}Cl) is not characteristic; it is always
accompanied by one of the higher grades of oxidation, and the compounds
of this type are distinguished by their great chemical instability, and
split up into an element and the higher compound (for instance, Ag_{4}O =
2Ag + Ag_{2}O). Many elements, moreover, form transition oxides whose
composition is intermediate, which are able, like N_{2}O_{4}, to split up
into the lower and higher oxides. Thus iron gives magnetic oxide,
Fe_{3}O_{4}, which is in all respects (by its reactions) a compound of
the suboxide FeO with the oxide Fe_{2}O_{3}. The independent and more or
less stable saline compounds correspond with the following eight
types:--
R_{2}O; salts RX, hydroxides ROH. Generally basic like K_{2}O, Na_{2}O,
Hg_{2}O, Ag_{2}O, Cu_{2}O; if there are acid oxides of this
composition they are very rare, are only formed by distinctly acid
elements, and even then have only feeble acid properties; for
example, Cl_{2}O and N_{2}O.
R_{2}O_{2} or RO; salts RX_{2}, hydroxides R(OH)_{2}. The most simple
basic salts R_{2}OX_{2} or R(OH)X; for instance, the chloride
Zn_{2}OCl_{2}; also an almost exclusively basic type; but the basic
properties are more feebly developed than in the preceding type.
For example, CaO, MgO, BaO, PbO, FeO, MnO, &c.
R_{2}O_{3}; salts RX_{3}, hydroxides R(OH)_{3}, RO(OH), the most simple
basic salts ROX, R(OH)X_{3}. The bases are feeble, like
Al_{2}O_{3}, Fe_{2}O_{3}, Tl_{2}O_{3}, Sb_{2}O_{3}. The acid
properties are also feebly developed; for instance, in B_{2}O_{3};
but with the non-metals the properties of acids are already clear;
for instance, P_{2}O_{3}, P(OH)_{3}.
R_{2}O_{4} or RO_{2}; salts RX_{4} or ROX_{2}, hydroxides R(OH)_{4},
RO(OH)_{2}. Rarely bases (feeble), like ZrO_{2}, PtO_{2}; more
often acid oxides; but the acid properties are in general feeble,
as in CO_{2}, SO_{2}, SnO_{2}. Many intermediate oxides appear in
this and the preceding and following types.
R_{2}O_{5}; salts principally of the types ROX_{3}, RO_{2}X,
RO(OH)_{3}, RO_{2}(OH), rarely RX_{5}. The basic character (X, a
halogen, simple or complex; for instance, NO_{3}, Cl, &c.) is
feeble; the acid character predominates, as is seen in N_{2}O_{5},
P_{2}O_{5}, Cl_{2}O_{5}; then X = OH, OK, &c., for example
NO_{2}(OK).
R_{2}O_{6} or RO_{3}; salts and hydroxides generally of the type
RO_{2}X_{2}, RO_{2}(OH)_{2}. Oxides of an acid character, as
SO_{3}, CrO_{3}, MnO_{3}. Basic properties rare and feebly
developed as in UO_{3}.
R_{2}O_{7}; salts of the form RO_{3}X, RO_{3}(OH), acid oxides; for
instance, Cl_{2}O_{7}, Mn_{2}O_{7}. Basic properties as feebly
developed as the acid properties in the oxides R_{2}O.
R_{2}O_{8} or RO_{4}. A very rare type, and only known in OsO_{4} and
RuO_{4}.
It is evident from the circumstance that in all the higher types the
_acid hydroxides_ (for example, HClO_{4}, H_{2}SO_{4}, H_{3}PO_{4}) and
salts with a single atom of one element contain, like the higher saline
type RO_{4}, _not more than four atoms of oxygen_; that the formation of
the saline oxides is governed by a certain common principle which is best
looked for in the fundamental properties of oxygen, and in general of the
most simple compounds. The hydrate of the oxide RO_{2} is of the higher
type RO_{2}2H_{2}O = RH_{4}O_{4} = R(HO)_{4}. Such, for example, is the
hydrate of silica and the salts (orthosilicates) corresponding with it,
Si(MO)_{4}. The oxide R_{2}O_{5}, corresponds with the hydrate
R_{2}O_{5}3H_{2}O = 2RH_{3}O_{4} = 2RO(OH)_{3}. Such is orthophosphoric
acid, PH_{3}O_{3}. The hydrate of the oxide RO_{3} is RO_{3}H_{2}O =
RH_{2}O_{4} = RO_{2}(OH)_{2}--for instance, sulphuric acid. The hydrate
corresponding to R_{2}O_{7} is evidently RHO = RO_{3}(OH)--for example,
perchloric acid. Here, besides containing O_{4}, it must further be
remarked that _the amount of hydrogen in the hydrate is equal to the
amount of hydrogen in the hydrogen compound_. Thus silicon gives SiH_{4}
and SiH_{4}O_{4}, phosphorus PH_{3} and PH_{3}O_{4}, sulphur SH_{2} and
SH_{2}O_{4}, chlorine ClH and ClHO_{4}. This, if it does not explain, at
least connects in a harmonious and general system the fact that _the
elements are capable of combining with a greater amount of oxygen, the
less the amount of hydrogen which they are able to retain_. In this the
key to the comprehension of all further deductions must be looked for,
and we will therefore formulate this rule in general terms. An element R
gives a hydrogen compound RH_{_n_}, the hydrate of its higher oxide will
be RH_{_n_}O_{4}, and therefore the higher oxide will contain
2RH_{_n_}O_{4} - _n_H_{2}O = R_{2}O_{8 - _n_}. For example, chlorine
gives ClH, hydrate ClHO_{4}, and the higher oxide Cl_{2}O_{7}. Carbon
gives CH_{4} and CO_{2}. So also, SiO_{2} and SiH_{4} are the higher
compounds of silicon with hydrogen and oxygen, like CO_{2} and CH_{4}.
Here the amounts of oxygen and hydrogen are equivalent. Nitrogen combines
with a large amount of oxygen, forming N_{2}O_{5}, but, on the other
hand, with a small quantity of hydrogen in NH_{3}. _The sum of the
equivalents of hydrogen and oxygen_, occurring in combination with an
atom of nitrogen, is, as always in the higher types, equal to _eight_. It
is the same with the other elements which combine with hydrogen and
oxygen. Thus sulphur gives SO_{3}; consequently, six equivalents of
oxygen fall to an atom of sulphur, and in SH_{2} two equivalents of
hydrogen. The sum is again equal to eight. The relation between
Cl_{2}O_{7} and ClH is the same. This shows that the property of elements
of combining with such different elements as oxygen and hydrogen is
subject to one common law, which is also formulated in the system of the
elements presently to be described.[7]
[7] The hydrogen compounds, R_{2}H, in equivalency correspond with the
type of the suboxides, R_{4}O. Palladium, sodium, and potassium
give such hydrogen compounds, and it is worthy of remark that
according to the periodic system these elements stand near to each
other, and that in those groups where the hydrogen compounds R_{2}H
appear, the quaternary oxides R_{4}O are also present.
Not wishing to complicate the explanation, I here only touch on the
general features of the relation between the hydrates and oxides
and of the oxides among themselves. Thus, for instance, the
conception of the ortho-acids and of the normal acids will be
considered in speaking of phosphoric and phosphorous acids.
As in the further explanation of the periodic law only those oxides
which give salts will be considered, I think it will not be
superfluous to mention here the following facts relative to the
peroxides. Of the _peroxides_ corresponding with hydrogen peroxide,
the following are at present known: H_{2}O_{2}, Na_{2}O_{2},
S_{2}O_{7} (as HSO_{4}?), K_{2}O_{4}, K_{2}O_{2}, CaO_{2}, TiO_{3},
Cr_{2}O_{7}, CuO_{2}(?), ZnO_{2}, Rb_{2}O_{2}, SrO_{2},
Ag_{2}O_{2}, CdO_{2}, CsO_{2}, Cs_{2}O_{2}, BaO_{2}, Mo_{2}O_{7},
SnO_{3}, W_{2}O_{7}, UO_{4}. It is probable that the number of
peroxides will increase with further investigation. A periodicity
is seen in those now known, for the elements (excepting Li) of the
first group, which give R_{2}O, form peroxides, and then the
elements of the sixth group seem also to be particularly inclined
to form peroxides, R_{2}O_{7}; but at present it is too early, in
my opinion, to enter upon a generalisation of this subject, not
only because it is a new and but little studied matter (not
investigated for all the elements), but also, and more especially,
because in many instances only the hydrates are known--for
instance, Mo_{2}H_{2}O_{8}--and they perhaps are only compounds of
peroxide of hydrogen--for example, Mo_{2}H_{2}O_{8} = 2MoO_{3} +
H_{2}O_{2}--since Prof. Schöne has shown that H_{2}O_{2} and
BaO_{2} possess the property of combining together and with other
oxides. Nevertheless, I have, in the general table expressing the
periodic properties of the elements, endeavoured to sum up the data
respecting all the known peroxide compounds whose characteristic
property is seen in their capability to form peroxide of hydrogen
under many circumstances.
In the preceding we see not only the regularity and simplicity which
govern the formation and properties of the oxides and of all the
compounds of the elements, but also a fresh and exact means for
recognising the analogy of elements. Analogous elements give compounds of
analogous types, both higher and lower. If CO_{2} and SO_{2} are two
gases which closely resemble each other both in their physical and
chemical properties, the reason of this must be looked for not in an
analogy of sulphur and carbon, but in that identity of the type of
combination, RX_{4}, which both oxides assume, and in that influence
which a large mass of oxygen always exerts on the properties of its
compounds. In fact, there is little resemblance between carbon and
sulphur, as is seen not only from the fact that CO_{2} is the _higher
form_ of oxidation, whilst SO_{2} is able to further oxidise into SO_{3},
but also from the fact that all the other compounds--for example, SH_{2}
and CH_{4}, SCl_{2} and CCl_{4}, &c.--are entirely unlike both in type
and in chemical properties. This absence of analogy in carbon and sulphur
is especially clearly seen in the fact that the highest saline oxides are
of different composition, CO_{2} for carbon, and SO_{3} for sulphur. In
Chapter VIII. we considered the limit to which carbon tends in its
compounds, and in a similar manner there is for every element in its
compounds a tendency to attain a certain highest limit RX_{_n_}. This
view was particularly developed in the middle of the present century by
Frankland in studying the metallo-organic compounds, _i.e._ those in
which X is wholly or partially a hydrocarbon radicle; for instance, X =
CH_{3} or C_{2}H_{5} &c. Thus, for example, antimony, Sb (Chapter XIX.)
gives, with chlorine, compounds SbCl_{3} and SbCl_{5} and corresponding
oxygen compounds Sb_{2}O_{3} and Sb_{2}O_{5}, whilst under the action of
CH_{3}I, C_{2}H_{5}I, or in general EI (where E is a hydrocarbon radicle
of the paraffin series), upon antimony or its alloy with sodium there are
formed SbE_{3} (for example, Sb(CH_{3})_{3}, boiling at about 81°),
which, corresponding to the lower form of combination SbX_{3}, are able
to combine further with EI, or Cl_{2}, or O, and to form compounds of the
limiting type SbX_{5}; for example, SbE_{4}Cl corresponding to NH_{4}Cl
with the substitution of nitrogen by antimony, and of hydrogen by the
hydrocarbon radicle. The elements which are most chemically analogous are
characterised by the fact of their giving compounds of similar form
RX_{_n_}. The halogens which are analogous give both higher and lower
compounds. So also do the metals of the alkalis and of the alkaline
earths. And we saw that this analogy extends to the composition and
properties of the nitrogen and hydrogen compounds of these metals, which
is best seen in the salts. Many such groups of analogous elements have
long been known. Thus there are analogues of oxygen, nitrogen, and
carbon, and we shall meet with many such groups. But an acquaintance with
them inevitably leads to the questions, what is the cause of analogy and
what is the relation of one group to another? If these questions remain
unanswered, it is easy to fall into error in the formation of the groups,
because the notions of the degree of analogy will always be relative, and
will not present any accuracy or distinctness Thus lithium is analogous
in some respects to potassium and in others to magnesium; beryllium is
analogous to both aluminium and magnesium. Thallium, as we shall
afterwards see and as was observed on its discovery, has much kinship
with lead and mercury, but some of its properties appertain to lithium
and potassium. Naturally, where it is impossible to make measurements one
is reluctantly obliged to limit oneself to approximate comparisons,
founded on apparent signs which are not distinct and are wanting in
exactitude. But in the elements there is one accurately measurable
property, which is subject to no doubt--namely, that property which is
expressed in their atomic weights. Its magnitude indicates the relative
mass of the atom, or, if we avoid the conception of the atom, its
magnitude shows the relation between the masses forming the chemical and
independent individuals or elements. And according to the teaching of all
exact data about the phenomena of nature, the mass of a substance is that
property on which all its remaining properties must be dependent, because
they are all determined by similar conditions or by those forces which
act in the weight of a substance, and this is directly proportional to
its mass. Therefore it is most natural to seek for a dependence between
the properties and analogies of the elements on the one hand and their
atomic weights on the other.
This is the fundamental idea which leads _to arranging all the elements
according to their atomic weights_. A periodic repetition of properties
is then immediately observed in the elements. We are already familiar
with examples of this:--
F = 19, Cl = 35·5, Br = 80, I = 127,
Na = 23, K = 39, Rb = 85, Cs = 133,
Mg = 24, Ca = 40, Sr = 87, Ba = 137.
The essence of the matter is seen in these groups. The halogens have
smaller atomic weights than the alkali metals, and the latter than the
metals of the alkaline earths. Therefore, _if all the elements be
arranged in the order of their atomic weights, a periodic repetition of
properties is obtained_. This is expressed by the _law of periodicity_,
_the properties of the elements, as well as the forms and properties of
their compounds, are in periodic dependence or (expressing ourselves
algebraically) form a periodic function of the atomic weights of the
elements_.[8] Table I. of _the periodic system of the elements_, which is
placed at the very beginning of this book, is designed to illustrate this
law. It is arranged in conformity with the eight types of oxides
described in the preceding pages, and those elements which give the
oxides, R_{2}O and consequently salts RX, form the 1st group; the
elements giving R_{2}O_{2} or RO as their highest grade of oxidation
belong to the 2nd group; those giving R_{2}O_{3} as their highest oxides
form the 3rd group, and so on; whilst the elements of all the groups
which are nearest in their atomic weights are arranged in series from 1
to 12. The even and uneven series of the same groups present the same
forms and limits, but differ in their properties, and therefore two
contiguous series, one even and the other uneven--for instance, the 4th
and 5th--form a period. Hence the elements of the 4th, 6th, 8th, 10th,
and 12th, or of the 3rd, 5th, 7th, 9th, and 11th, series form analogues,
like the halogens, the alkali metals, &c. The conjunction of two series,
one even and one contiguous uneven series, thus forms one large _period_.
These periods, beginning with the alkali metals, end with the halogens.
The elements of the first two series have the lowest atomic weights, and
in consequence of this very circumstance, although they bear the general
properties of a group, they still show many peculiar and independent
properties.[9] Thus fluorine, as we know, differs in many points from the
other halogens, and lithium from the other alkali metals, and so on.
These lightest elements may be termed _typical elements_. They include--
H.
Li, Be, B, C, N, O, F.
Na, Mg....
In the annexed table all the remaining elements are arranged, not in
groups and series, but _according to periods_. In order to understand the
essence of the matter, it must be remembered that here the atomic weight
gradually increases along a given line; for instance, in the line
commencing with K = 39 and ending with Br = 80, the intermediate elements
have intermediate atomic weights, as is clearly seen in Table III., where
the elements stand in the order of their atomic weights.
I. II. III. IV. V. VI. VII. I. II. III. IV. V. VI. VII.
{ Even Series. } Mg Al Si P S Cl
K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br
Rb Sr Y Zr Nb Mo -- Ru Rh Pd Ag Cd In Sn Sb Te I
Cs Ba La Ce Di? -- -- -- -- -- -- -- -- -- -- -- --
-- -- Yb -- Ta W -- Os Ir Pt Au Hg Tl Pb Bi -- --
-- -- -- Th -- U { Uneven Series }
The same degree of analogy that we know to exist between potassium,
rubidium, and cæsium; or chlorine, bromine, and iodine; or calcium,
strontium, and barium, also exists between the elements of the other
vertical columns. Thus, for example, zinc, cadmium, and mercury, which
are described in the following chapter, present a very close analogy with
magnesium. For a true comprehension of the matter[10] it is very
important to see that all the aspects of the distribution of the elements
according to their atomic weights essentially express one and the same
fundamental _dependence_--_periodic properties_.[11] The following points
then must be remarked in it.
[8] The periodic law and the periodic system of the elements appeared
in the same form as here given in the first edition of this work,
begun in 1868 and finished in 1871. In laying out the accumulated
information respecting the elements, I had occasion to reflect on
their mutual relations. At the beginning of 1869 I distributed
among many chemists a pamphlet entitled 'An Attempted System of the
Elements, based on their Atomic Weights and Chemical Analogies,'
and at the March meeting of the Russian Chemical Society, 1869, I
communicated a paper 'On the Correlation of the Properties and
Atomic Weights of the Elements.' The substance of this paper is
embraced in the following conclusions: (1) The elements, if
arranged according to their atomic weights, exhibit an evident
_periodicity_ of properties. (2) Elements which are similar as
regards their chemical properties have atomic weights which are
either of nearly the same value (platinum, iridium, osmium) or
which increase regularly (_e.g._ potassium, rubidium, cæsium). (3)
The arrangement of the elements or of groups of elements in the
order of their atomic weights corresponds with their so-called
_valencies_. (4) The elements, which are the most widely
distributed in nature, have _small_ atomic weights, and all the
elements of small atomic weight are characterised by sharply
defined properties. They are therefore typical elements. (5) The
_magnitude_ of the atomic weight determines the character of an
element. (6) The discovery of many yet unknown elements may be
expected. For instance, elements analogous to aluminium and
silicon, whose atomic weights would be between 65 and 75. (7) The
atomic weight of an element may sometimes be corrected by aid of a
knowledge of those of the adjacent elements. Thus the combining
weight of tellurium must lie between 123 and 126, and cannot be
128. (8) Certain characteristic properties of the elements can be
foretold from their atomic weights.
The entire periodic law is included in these lines. In the series
of subsequent papers (1870-72, for example, in the _Transactions_
of the Russian Chemical Society, of the Moscow Meeting of
Naturalists, of the St. Petersburg Academy, and Liebig's _Annalen_)
on the same subject we only find applications of the same
principles, which were afterwards confirmed by the labours of
Roscoe, Carnelley, Thorpe, and others in England; of Rammelsberg
(cerium and uranium), L. Meyer (the specific volumes of the
elements), Zimmermann (uranium), and more especially of C. Winkler
(who discovered germanium, and showed its identity with
ekasilicon), and others in Germany; of Lecoq de Boisbaudran in
France (the discoverer of gallium = ekaaluminium); of Clève (the
atomic weights of the cerium metals), Nilson (discoverer of
scandium = ekaboron), and Nilson and Pettersson (determination of
the vapour density of beryllium chloride) in Sweden; and of Brauner
(who investigated cerium, and determined the combining weight of
tellurium = 125) in Austria, and Piccini in Italy.
I consider it necessary to state that, in arranging the periodic
system of the elements, I made use of the previous researches of
Dumas, Gladstone, Pettenkofer, Kremers, and Lenssen on the atomic
weights of related elements, but I was not acquainted with the
works preceding mine of De Chancourtois (_vis tellurique_, or the
spiral of the elements according to their properties and
equivalents) in France, and of J. Newlands (Law of Octaves--for
instance, H, F, Cl, Co, Br, Pd, I, Pt form the first octave, and O,
S, Fe, Se, Rh, Te, Au, Th the last) in England, although certain
germs of the periodic law are to be seen in these works. With
regard to the work of Prof. Lothar Meyer respecting the periodic
law (Notes 12 and 13), it is evident, judging from the method of
investigation, and from his statement (Liebig's _Annalen, Supt.
Band 7_, 1870, 354), at the very commencement of which he cites my
paper of 1869 above mentioned, that he accepted the periodic law in
the form which I proposed.
In concluding this historical statement I consider it well to
observe that no law of nature, however general, has been
established all at once; its recognition is always preceded by many
hints; the establishment of a law, however, does not take place
when its significance is recognised, but only when it has been
confirmed by experiment, which the man of science must consider as
the only proof of the correctness of his conjectures and opinions.
I therefore, for my part, look upon Roscoe, De Boisbaudran, Nilson,
Winkler, Brauner, Carnelley, Thorpe, and others who verified the
adaptability of the periodic law to chemical facts, as the true
founders of the periodic law, the further development of which
still awaits fresh workers.
[9] This resembles the fact, well known to those having an acquaintance
with organic chemistry, that in a series of homologues (Chapter
VIII.) the first members, in which there is the least carbon,
although showing the general properties of the homologous series,
still present certain distinct peculiarities.
[10] Besides arranging the elements (_a_) in a successive order
according to their atomic weights, with indication of their
analogies by showing some of the properties--for instance, their
power of giving one or another form of combination--both of the
_elements_ and of their compounds (as is done in Table III. and in
the table on p. 36), (_b_) according to periods (as in Table I. at
the commencement of volume I. after the preface), and (_c_)
according to groups and series or small periods (as in the same
tables), I am acquainted with the following methods of expressing
the periodic relations of the elements: (1) By a curve drawn
through points obtained in the following manner: The elements are
arranged along the horizontal axis as abscissæ at distances from
zero proportional to their atomic weights, whilst the values for
all the elements of some property--for example, the specific
volumes or the melting points, are expressed by the ordinates.
This method, although graphic, has the theoretical disadvantage
that it does not in any way indicate the existence of a limited
and definite number of elements in each period. There is nothing,
for instance, in this method of expressing the law of periodicity
to show that between magnesium and aluminium there can be no other
element with an atomic weight of, say, 25, atomic volume 13, and
in general having properties intermediate between those of these
two elements. The actual periodic law does not correspond with a
continuous change of properties, with a continuous variation of
atomic weight--in a word, it does not express an uninterrupted
function--and as the law is purely chemical, starting from the
conception of atoms and molecules which combine in multiple
proportions, with intervals (not continuously), it _above all_
depends on there being but few types of compounds, which are
arithmetically simple, _repeat themselves_, and offer no
uninterrupted transitions, so that each period can only contain a
definite number of members. For this reason there can be no other
elements between magnesium, which gives the chloride MgCl_{2}, and
aluminium, which forms AlX_{3}; there is a break in the
continuity, according to the law of multiple proportions. The
periodic law ought not, therefore, to be expressed by geometrical
figures in which continuity is always understood. Owing to these
considerations I never have and never will express the periodic
relations of the elements by any geometrical figures. (2) _By a
plane spiral._ Radii are traced from a centre, proportional to the
atomic weights; analogous elements lie along one radius, and the
points of intersection are arranged in a spiral. This method,
adopted by De Chancourtois, Baumgauer, E. Huth, and others, has
many of the imperfections of the preceding, although it removes
the indefiniteness as to the number of elements in a period. It is
merely an attempt to reduce the complex relations to a simple
graphic representation, since the equation to the spiral and the
number of radii are not dependent upon anything. (3) _By the lines
of atomicity_, either parallel, as in Reynolds's and the Rev. S.
Haughton's method, or as in Crookes's method, arranged to the
right and left of an axis, along which the magnitudes of the
atomic weights are counted, and the position of the elements
marked off, on the one side the members of the even series
(paramagnetic, like oxygen, potassium, iron), and on the other
side the members of the uneven series (diamagnetic, like sulphur,
chlorine, zinc, and mercury). On joining up these points a
periodic curve is obtained, compared by Crookes to the
oscillations of a pendulum, and, according to Haughton,
representing a cubical curve. This method would be very graphic
did it not require, for instance, that sulphur should be
considered as bivalent and manganese as univalent, although
neither of these elements gives stable derivatives of these
natures, and although the one is taken on the basis of the lowest
possible compound SX_{2}, and the other of the highest, because
manganese can be referred to the univalent elements only by the
analogy of KMnO_{4} to KClO_{4}. Furthermore, Reynolds and Crookes
place hydrogen, iron, nickel, cobalt, and others outside the axis
of atomicity, and consider uranium as bivalent without the least
foundation. (4) Rantsheff endeavoured to classify the elements in
their periodic relations by a system dependent on solid geometry.
He communicated this mode of expression to the Russian Chemical
Society, but his communication, which is apparently not void of
interest, has not yet appeared in print. (5) _By algebraic
formulæ_: for example, E. J. Mills (1886) endeavours to express
all the atomic weights by the logarithmic function A = 15(_n_ -
0·9375_t_), in which the variables _n_ and _t_ are whole numbers.
For instance, for oxygen _n_ = 2, _t_ = 1; hence A = 15·94; for
antimony _n_ = 9, _t_ = 0; whence A = 120, and so on. _n_ varies
from 1 to 16 and _t_ from 0 to 59. The analogues are hardly
distinguishable by this method: thus for chlorine the magnitudes
of _n_ and _t_ are 3 and 7; for bromine 6 and 6; for iodine 9 and
9; for potassium 3 and 14; for rubidium 6 and 18; for cæsium 9 and
20; but a certain regularity seems to be shown. (6) A more natural
method of expressing the dependence of the properties of elements
on their atomic weights is obtained by _trigonometrical
functions_, because this dependence is periodic like the functions
of trigonometrical lines, and therefore Ridberg in Sweden (Lund,
1885) and F. Flavitzky in Russia (Kazan, 1887) have adopted a
similar method of expression, which must be considered as worthy
of being worked out, although it does not express the absence of
intermediate elements--for instance, between magnesium and
aluminium, which is essentially the most important part of the
matter. (7) The investigations of B. N. Tchitchérin (1888,
_Journal of the Russian Physical and Chemical Society_) form the
first effort in the latter direction. He carefully studied the
alkali metals, and discovered the following simple relation
between their atomic volumes: they can all be expressed by A(2 -
0·0428A_n_), where A is the atomic weight and _n_ = 1 for lithium
and sodium, 4/8 for potassium, 3/8 for rubidium, and 2/8 for
cæsium. If _n_ always = 1, then the volume of the atom would
become zero at A = 46-2/3, and would reach its maximum when A =
23-1/3, and the density increases with the growth of A. In order
to explain the variation of _n_, and the relation of the atomic
weights of the alkali metals to those of the other elements, as
also the atomicity itself, Tchitchérin supposes all atoms to be
built up of a primary matter; he considers the relation of the
central to the peripheric mass, and, guided by mechanical
principles, deduces many of the properties of the atoms from the
reaction of the internal and peripheric parts of each atom. This
endeavour offers many interesting points, but it admits the
hypothesis of the building up of all the elements from one primary
matter, and at the present time such an hypothesis has not the
least support either in theory or in fact. Besides which the
starting-point of the theory is the specific gravity of the metals
at a definite temperature (it is not known how the above relation
would appear at other temperatures), and the specific gravity
varies even under mechanical influences. L. Hugo (1884)
endeavoured to represent the atomic weights of Li, Na, K, Rb, and
Cs by geometrical figures--for instance, Li = 7 represents a
central atom = 1 and six atoms on the six terminals of an
octahedron; Na, is obtained by applying two such atoms on each
edge of an octahedron, and so on. It is evident that such methods
can add nothing new to our data respecting the atomic weights of
analogous elements.
[11] Many natural phenomena exhibit a dependence of a periodic
character. Thus the phenomena of day and night and of the seasons
of the year, and vibrations of all kinds, exhibit variations of a
periodic character in dependence on time and space. But in
ordinary periodic functions one variable varies continuously,
whilst the other increases to a limit, then a period of decrease
begins, and having in turn reached its limit a period of increase
again begins. It is otherwise in the periodic function of the
elements. Here the mass of the elements does not increase
continuously, but abruptly, by steps, as from magnesium to
aluminium. So also the valency or atomicity leaps directly from 1
to 2 to 3, &c., without intermediate quantities, and in my opinion
it is these properties which are the most important, and it is
their periodicity which forms the substance of the periodic law.
It expresses _the properties of the real elements_, and not of
what may be termed their manifestations visually known to us. The
external properties of elements and compounds are in periodic
dependence on the atomic weight of the elements only because these
external properties are themselves the result of the properties of
the real elements which unite to form the 'free' elements and the
compounds. To explain and express the periodic law is to explain
and express the cause of the law of multiple proportions, of the
difference of the elements, and the variation of their atomicity,
and at the same time to understand what mass and gravitation are.
In my opinion this is still premature. But just as without knowing
the cause of gravitation it is possible to make use of the law of
gravity, so for the aims of chemistry it is possible to take
advantage of the laws discovered by chemistry without being able
to explain their causes. The above-mentioned peculiarity of the
laws of chemistry respecting definite compounds and the atomic
weights leads one to think that the time has not yet come for
their full explanation, and I do not think that it will come
before the explanation of such a primary law of nature as the law
of gravity.
It will not be out of place here to turn our attention to the
many-sided correlation existing between the undecomposable
_elements and the compound carbon radicles_, which has long been
remarked (Pettenkofer, Dumas, and others), and reconsidered in
recent times by Carnelley (1886), and most originally in
Pelopidas's work (1883) on the principles of the periodic system.
Pelopidas compares the series containing eight hydrocarbon
radicles, C_{_n_}H_{2_n_ + 1}, C_{_n_}H_{2_n_} &c., for instance,
C_{6}H_{13}, C_{6}H_{12}, C_{6}H_{11}, C_{6}H_{10}, C_{6}H_{9},
C_{6}H_{8}, C_{6}H_{7}, and C_{6}H_{6}--with the series of the
elements arranged in eight groups. The analogy is particularly
clear owing to the property of C_{_n_}H_{2_n_+1} to combine with
X, thus reaching saturation, and of the following members with
X_{2}, X_{3} ... X_{8}, and especially because these are followed
by an aromatic radicle--for example, C_{6}H_{5}--in which, as is
well known, many of the properties of the saturated radicle
C_{6}H_{13} are repeated, and in particular the power of forming a
univalent radicle again appears. Pelopidas shows a confirmation of
the parallel in the property of the above radicles of giving
oxygen compounds corresponding with the groups in the periodic
system. Thus the hydrocarbon radicles of the first group--for
instance, C_{6}H_{13} or C_{6}H_{5}--give oxides of the form
R_{2}O and hydroxides RHO, like the metals of the alkalis; and in
the third group they form oxides R_{2}O_{3} and hydrates RO_{2}H.
For example, in the series CH_{3} the corresponding compounds of
the third group will be the oxide (CH)_{2}O_{3} or
C_{2}H_{2}O_{3}--that is, formic anhydride and hydrate, CHO_{2}H,
or formic acid. In the sixth group, with a composition of C_{2},
the oxide RO_{3} will be C_{2}O_{3}, and hydrate
C_{2}H_{2}O_{4}--that is, also a bibasic acid (oxalic) resembling
sulphuric, among the inorganic acids. After applying his views to
a number of organic compounds, Pelopidas dwells more particularly
on the radicles corresponding with ammonium.
With respect to this remarkable parallelism, it must above all be
observed that in the elements the atomic weight increases in
passing to contiguous members of a higher valency, whilst here it
decreases, which should indicate that the periodic variability of
elements and compounds is subject to some higher law whose nature,
and still more whose cause, cannot at present be determined. It is
probably based on the fundamental principles of the internal
mechanics of the atoms and molecules, and as the periodic law has
only been generally recognised for a few years it is not
surprising that any further progress towards its explanation can
only be looked for in the development of facts touching on this
subject.
1. The composition of the higher oxygen compounds is determined by the
groups: the first group gives R_{2}O, the second R_{2}O_{2} or RO, the
third R_{2}O_{3}, &c. There are eight types of oxides and therefore eight
groups. Two groups give a period, and the same type of oxide is met with
twice in a period. For example, in the period beginning with potassium,
oxides of the composition RO are formed by calcium and zinc, and of the
composition RO_{3} by molybdenum and tellurium. The oxides of the even
series, of the same type, have stronger basic properties than the oxides
of the uneven series, and the latter as a rule are endowed with an acid
character. Therefore the elements which exclusively give bases, like the
alkali metals, will be found at the commencement of the period, whilst
such purely acid elements as the halogens will be at the end of the
period. The interval will be occupied by intermediate elements, whose
character and properties we shall afterwards describe. It must be
observed that the acid character is chiefly proper to the elements with
small atomic weights in the uneven series, whilst the basic character is
exhibited by the heavier elements in the even series. Hence elements
which give acids chiefly predominate among the lightest (typical)
elements, especially in the last groups; whilst the heaviest elements,
even in the last groups (for instance, thallium, uranium) have a basic
character. Thus the basic and acid characters of the higher oxides are
determined (_a_) by the type of oxide, (_b_) by the even or uneven
series, and (_c_) by the atomic weight.[11 bis] The groups are indicated
by Roman numerals from I. to VIII.
2. _The hydrogen compounds_ being volatile or gaseous substances which
are prone to reaction--such as HCl, H_{2}O, H_{3}N, and H_{4}C[12]--are
only formed by the elements of the uneven series and higher groups giving
oxides of the forms R_{2}O_{_n_}, RO_{3}, R_{2}O_{5}, and RO_{2}.
3. If an element gives a hydrogen compound, RX_{_m_}, it forms an
_organo-metallic compound_ of the same composition, where X =
C_{_n_}H_{2_n_ + 1}; that is, X is the radicle of a saturated
hydrocarbon. The elements of the uneven series, which are incapable of
giving hydrogen compounds, and give oxides of the forms RX, RX_{2},
R_{X}3, also give organo-metallic compounds of this form proper to the
higher oxides. Thus zinc forms the oxide ZnO, salts ZnX_{2} and zinc
ethyl Zn(C_{2}H_{5})_{2}. The elements of the even series do not seem to
form organo-metallic compounds at all; at least all efforts for their
preparation have as yet been fruitless--for instance, in the case of
titanium, zirconium, or iron.
4. The atomic weights of elements belonging to contiguous periods differ
approximately by 45; for example, K<Rb, Cr<Mo, Br<I. But the elements of
the typical series show much smaller differences. Thus the difference
between the atomic weights of Li, Na, and K, between Ca, Mg, and Be,
between Si and C, between S and O, and between Cl and F, is 16. As a
rule, there is a greater difference between the atomic weights of two
elements of one group and belonging to two neighbouring series (Ti-Si =
V-P = Cr-S = Mn-Cl = Nb-As, &c. = 20); and this difference attains a
maximum with the heaviest elements (for example, Th-Pb = 26, Bi-Ta = 26,
Ba-Cd = 25, &c.). Furthermore, the difference between the atomic weights
of the elements of even and uneven series also increases. In fact, the
differences between Na and K, Mg and Ca, Si and Ti, are less abrupt than
those between Pb and Th, Ta and Bi, Cd and Ba, &c. Thus even in the
magnitude of the differences of the atomic weights of analogous elements
there is observable a certain connection with the gradation of their
properties.[12 bis]
5. According to the periodic system every element occupies a certain
position, determined by the group (indicated in Roman numerals) and
series (Arabic numerals) in which it occurs. These indicate the atomic
weight, the analogues, properties, and type of the higher oxide, and of
the hydrogen and other compounds--in a word, all the chief quantitative
and qualitative features of an element, although there yet remain a whole
series of further details and peculiarities whose cause should perhaps be
looked for in small differences of the atomic weights. If in a certain
group there occur elements, R_{1}, R_{2}, R_{3}, and if in that series
which contains one of these elements, for instance R_{2}, an element
Q_{2} precedes it and an element T_{2} succeeds it, then the properties
of R_{2} are determined by the properties of R_{1}, R_{3}, Q_{2}, and
T_{2}. Thus, for instance, the atomic weight of R_{2} = 1/4(R_{1} +
R_{3} + Q_{2} + T_{2}). For example, selenium occurs in the same group as
sulphur, S = 32, and tellurium, Te = 125, and, in the 7th series As = 75
stands before it and Br = 80 after it. Hence the atomic weight of
selenium should be 1/4(32 + 125 + 75 + 80) = 78, which is near to the
truth. Other properties of selenium may also be determined in this
manner. For example, arsenic forms H_{3}As, bromine gives HBr, and it is
evident that selenium, which stands between them, should form H_{2}Se,
with properties intermediate between those of H_{3}As and HBr. Even the
physical properties of selenium and its compounds, not to speak of their
composition, being determined by the group in which it occurs, may be
foreseen with a close approach to reality from the properties of sulphur,
tellurium, arsenic, and bromine. _In this manner it is possible to
foretell the properties of still unknown elements._ For instance in the
position IV, 5--that is, in the IVth group and 5th series--an element is
still wanting. These unknown elements may be named after the preceding
known element of the same group by adding to the first syllable the
prefix _eka_-, which means _one_ in Sanskrit. The element IV, 5, follows
after IV, 3, and this latter position being occupied by silicon, we call
the unknown element ekasilicon and its symbol Es. The following are the
properties which this element should have on the basis of the known
properties of silicon, tin, zinc, and arsenic. Its atomic weight is
nearly 72, higher oxide EsO_{2}, lower oxide EsO, compounds of the
general form EsX_{4}, and chemically unstable lower compounds of the form
EsX_{2}. Es gives volatile organo-metallic compounds--for instance,
Es(CH_{3})_{4}, Es(CH_{3})_{3}Cl, and Es(C_{2}H_{5})_{4}, which boil at
about 160°, &c.; also a volatile and liquid chloride, EsCl_{4}, boiling
at about 90° and of specific gravity about 1·9. EsO_{2} will be the
anhydride of a feeble colloidal acid, metallic Es will be rather easily
obtainable from the oxides and from K_{2}EsF_{6} by reduction, EsS_{2}
will resemble SnS_{2} and SiS_{2}, and will probably be soluble in
ammonium sulphide; the specific gravity of Es will be about 5·5, EsO_{2}
will have a density of about 4·7, &c. Such a prediction of the properties
of ekasilicon was made by me in 1871, on the basis of the properties of
the elements analogous to it: IV, 3, = Si, IV, 7 = Sn, and also II, 5 =
Zn and V, 5 = As. And now that this element has been discovered by C.
Winkler, of Freiberg, it has been found that its actual properties
entirely correspond with those which were foretold.[13] In this we see a
most important confirmation of the truth of the periodic law. This
element is now called germanium, Ge (_see_ Chapter XVIII.). It is not the
only one that has been predicted by the periodic law.[14] We shall see in
describing the elements of the third group that properties were foretold
of an element ekaaluminium, III, 5, El = 68, and were afterwards verified
when the metal termed 'gallium' was discovered by De Boisbaudran. So also
the properties of scandium corresponded with those predicted for
ekaboron, according to Nilson.[15]
[11 bis] True peroxides (_see_ Note 7), like H_{2}O_{2}, BaO_{2},
S_{2}O_{7} (Chapter XX.), must not be confused with true saline
oxides even if the latter contain much oxygen (for instance,
N_{2}O_{5}, CrO_{3}, &c.) although one and the other easily
oxidise. The difference between them is seen in their fundamental
properties: the saline oxides correspond to water, while the
peroxides correspond in their reactions and origin to peroxide of
hydrogen. This is clearly seen in the difference between Na_{2}O
and Na_{2}O_{2} (Chapter XII.). Therefore the peroxides should
also have their periodicity. An element R, giving a highest degree
of oxidation, R_{2}O_{_n_}, may give both a lower degree of
oxidation, R_{2}O_{_n_ - _m_} (where _m_ is evidently less than
_n_), and peroxides, R_{2}O_{_n_ + 1}, R_{2}O_{_n_ + 2}, or even
more oxygen. This class of oxides, to which attention has only
recently been turned (Berthelot, Piccini, &c.), may perhaps on
further study give the possibility of generalising the capability
of the elements to give unstable complex higher forms of
combination, such as double salts, and in my opinion should in the
near future be the field of new and important discoveries. And in
contemporary chemistry, salts, saline oxides, hydrogen compounds,
and other combinations of the elements corresponding to them
constitute an important and very complex problem for
generalisation, which is satisfied by the periodic law in its
present form, to which it has risen from its first state, in which
it gave the means of foreseeing (_see_ later on) the existence of
unknown elements (Ga, Sc, and Ge), their properties, and many
details respecting their compounds. Until those improvements in
the periodic system which have been proposed by Prof. Flavitzky
(of Kazan) and Prof. Harperath (of Cordoba, in the Argentine
Republic), Ugo Alvisi (Italy), and others give similar practical
results, I think it unnecessary to discuss them further.
[12] The hydrides generalised by the periodic law are those to which
metallo-organic compounds correspond, and they are themselves
either volatile or gaseous. The hydrogen compounds like Na_{2}H,
BaH, &c. are distinguished by other signs. They resemble alloys.
They show (_see_ end of last chapter) a systematic harmony, but
they evidently should not be confused with true hydrides, any more
than peroxides with saline oxides. Moreover, such hydrides have,
like the peroxides, only recently been subjected to research, and
have been but little studied. The best known of these compounds
are given in the 16th column of Table III., and it will be seen
that they already exhibit a periodicity of properties and
composition.
[12 bis] The relation between the atomic weights, and especially the
difference = 16, was observed in the sixth and seventh decades of
this century by Dumas, Pettenkofer, L. Meyer, and others. Thus
Lothar Meyer in 1864, following Dumas and others, grouped together
the tetravalent elements carbon and silicon; the trivalent
elements nitrogen, phosphorus, arsenic, antimony, and bismuth; the
bivalent oxygen, sulphur, selenium, and tellurium; the univalent
fluorine, chlorine, bromine, and iodine; the univalent metals
lithium, sodium, potassium, rubidium, cæsium, and thallium, and
the bivalent metals beryllium, magnesium, strontium and
barium--observing that in the first the difference is, in general
= 16, in the second about = 46, and the last about = 87-90. The
first germs of the periodic law are visible in such observations
as these. Since its establishment this subject has been most fully
worked out by Ridberg (Note 10), who observed a periodicity in the
variation of the differences between the atomic weights of two
contiguous elements, and its relation to their atomicity. A.
Bazaroff (1887) investigated the same subject, taking, not the
arithmetical differences of contiguous and analogous elements, but
the ratio of their atomic weights; and he also observed that this
ratio alternately rises and falls with the rise of the atomic
weights. I will here remark that the relation of the eighth group
to the others will be considered at the end of this work in
Chapter XXII.
[13] The laws of nature admit of no exceptions, and in this they
clearly differ from such rules and maxims as are found in grammar,
and other inventions, methods, and relations of man's creation.
The confirmation of a law is only possible by deducing
consequences from it, such as could not possibly be foreseen
without it, and by verifying those consequences by experiment and
further proofs. Therefore, when I conceived the periodic law, I
(1869-1871, Note 9) deduced such logical consequences from it as
could serve to show whether it were true or not. Among them was
the prediction of the properties of undiscovered elements and the
correction of the atomic weights of many, and at that time little
known, elements. Thus uranium was considered as trivalent, U =
120; but as such it did not correspond with the periodic law. I
therefore proposed to double its atomic weight--U = 240, and the
researches of Roscoe, Zimmermann, and others justified this
alteration (Chapter XXI.). It was the same with cerium (Chapter
XVIII.) whose atomic weight it was necessary to change according
to the periodic law. I therefore determined its specific heat, and
the result I obtained was verified by the new determinations of
Hillebrand. I then corrected certain formulæ of the cerium
compounds, and the researches of Rammelsberg, Brauner, Clève, and
others verified the proposed alteration. It was necessary to do
one or the other--either to consider the periodic law as
completely true, and as forming a new instrument in chemical
research, or to refute it. Acknowledging the method of experiment
to be the only true one, I myself verified what I could, and gave
every one the possibility of proving or confirming the law, and
did not think, like L. Meyer (Liebig's _Annalen, Supt. Band 7_,
1870, 364), when writing about the periodic law that 'it would be
rash to change the accepted atomic weights on the basis of so
uncertain a starting-point.' ('Es würde voreilig sein, auf so
unsichere Anhaltspunkte hin eine Aenderung der bisher angenommenen
Atomgewichte vorzunehmen.') In my opinion, the basis offered by
the periodic law had to be verified or refuted, and experiment in
every case verified it. The starting-point then became general. No
law of nature can be established without such a method of testing
it. Neither De Chancourtois, to whom the French ascribe the
discovery of the periodic law, nor Newlands, who is put forward by
the English, nor L. Meyer, who is now cited by many as its
founder, ventured to foretell the _properties_ of undiscovered
elements, or to alter the 'accepted atomic weights,' or, in
general, to regard the periodic law as a new, strictly established
law of nature, as I did from the very beginning (1869).
[14] When in 1871 I wrote a paper on the application of the periodic
law to the determination of the properties of hitherto
undiscovered elements, I did not think I should live to see the
verification of this consequence of the law, but such was to be
the case. Three elements were described--ekaboron, ekaaluminium,
and ekasilicon--and now, after the lapse of twenty years, I have
had the great pleasure of seeing them discovered and named
Gallium, Scandium, and Germanium, after those three countries
where the rare minerals containing them are found, and where they
were discovered. For my part I regard L. de Boisbaudran, Nilson,
and Winkler, who discovered these elements, as the true
corroborators of the periodic law. Without them it would not have
been accepted to the extent it now is.
[15] Taking indium, which occurs together with zinc, as our example, we
will show the principle of the method employed. The equivalent of
indium to hydrogen in its oxide is 37·7--that is, if we suppose
its composition to be like that of water; then In = 37·7, and the
oxide of indium is In_{2}O. The atomic weight of indium was taken
as double the equivalent--that is, indium was considered to be a
bivalent element--and In = 2 × 37·7 = 75·4. If indium only formed
an oxide, RO, it should be placed in group II. But in this case it
appears that there would be no place for indium in the system of
the elements, because the positions II., 5 = Zn = 65 and II., 6 =
Sr = 87 were already occupied by known elements, and according to
the periodic law an element with an atomic weight 75 could not be
bivalent. As neither the vapour density nor the specific heat, nor
even the isomorphism (the salts of indium crystallise with great
difficulty) of the compounds of indium were known, there was no
reason for considering it to be a bivalent metal, and therefore it
might be regarded as trivalent, quadrivalent, &c. If it be
trivalent, then In = 3 × 37·7 = 113, and the composition of the
oxide is In_{2}O_{3}, and of its salts InX_{3}. In this case it at
once falls into its place in the system, namely, in group III. and
7th series, between Cd = 112 and Sn = 118, as an analogue of
aluminium or dvialuminium (dvi = 2 in Sanskrit). All the
properties observed in indium correspond with this position; for
example, the density, cadmium = 8·6, indium = 7·4, tin = 7·2; the
basic properties of the oxides CdO, In_{2}O_{3}, SnO_{2},
successively vary, so that the properties of In_{2}O_{3} are
intermediate between those of CdO and SnO_{2} or Cd_{2}O_{2} and
Sn_{2}O_{4}. That indium belongs to group III. has been confirmed
by the determination of its specific heat, (0·057 according to
Bunsen, and 0·055 according to me) and also by the fact that
indium forms alums like aluminium, and therefore belongs to the
same group.
The same kind of considerations necessitated taking the atomic
weight of titanium as nearly 48, and not as 52, the figure derived
from many analyses. And both these corrections, made on the basis
of the law, have now been confirmed, for Thorpe found, by a series
of careful experiments, the atomic weight of titanium to be that
foreseen by the periodic law. Notwithstanding that previous
analyses gave Os = 199·7, Ir = 198, and Pt = 187, the periodic law
shows, as I remarked in 1871, that the atomic weights should rise
from osmium to platinum and gold, and not fall. Many recent
researches, and especially those of Seubert, have fully verified
this statement, based on the law. Thus a true law of nature
anticipates facts, foretells magnitudes, gives a hold on nature,
and leads to improvements in the methods of research, &c.
6. As a true law of nature is one to which there are no exceptions, the
periodic dependence of the properties on the atomic weights of the
elements gives a _new means for determining by the equivalent the atomic
weight_ or atomicity of imperfectly investigated but known elements, for
which no other means could as yet be applied for determining the true
atomic weight. At the time (1869) when the periodic law was first
proposed there were several such elements. It thus became possible to
learn their true atomic weights, and these were verified by later
researches. Among the elements thus concerned were indium, uranium,
cerium, yttrium, and others.
7. The periodic variability of the properties of the elements in
dependence on their masses presents a distinction from other kinds of
periodic dependence (as, for example, the sines of angles vary
periodically and successively with the growth of the angles, or the
temperature of the atmosphere with the course of time), in that the
weights of the atoms do not increase gradually, but by leaps; that is,
according to Dalton's law of multiple proportions, there not only are
not, but there cannot be, any transitive or intermediate elements between
two neighbouring ones (for example, between K = 39 and Ca = 40, or Al =
27 and Si = 28, or C = 12 and N = 14, &c.) As in a molecule of a hydrogen
compound there may be either one, as in HF, or two, as in H_{2}O, or
three, as in NH_{3}, &c., atoms of hydrogen; but as there cannot be
molecules containing 2-1/2 atoms of hydrogen to one atom of another
element, so there cannot be any element intermediate between N and O,
with an atomic weight greater than 14 or less than 16, or between K and
Ca. Hence the periodic dependence of the elements cannot be expressed by
any algebraical continuous function in the same way that it is possible,
for instance, to express the variation of the temperature during the
course of a day or year.
8. The essence of the notions giving rise to the periodic law consists
in a general physico-mechanical principle which recognises the
correlation, transmutability, and equivalence of the forces of nature.
Gravitation, attraction at small distances, and many other phenomena are
in direct dependence on the mass of matter. It might therefore have been
expected that chemical forces would also depend on mass. A dependence is
in fact shown, the properties of elements and compounds being determined
by the masses of the atoms of which they are formed. The weight of a
molecule, or its mass, determines, as we have seen, (Chapter VII. and
elsewhere) many of its properties independently of its composition. Thus
carbonic oxide, CO, and nitrogen, N_{2}, are two gases having the same
molecular weight, and many of their properties (density, liquefaction,
specific heat, &c.) are similar or nearly similar. The differences
dependent on the nature of a substance play another part, and form
magnitudes of another order. But the properties of atoms are mainly
determined by their mass or weight, and are in dependence upon it. Only
in this case there is a peculiarity in the dependence of the properties
on the mass, for this _dependence is determined by a periodic law_. As
the mass increases the properties vary, at first successively and
regularly, and then return to their original magnitude and recommence a
fresh period of variation like the first. Nevertheless here as in other
cases a small variation of the mass of the atom generally leads to a
small variation of properties, and determines differences of a second
order. The atomic weights of cobalt and nickel, of rhodium, ruthenium,
and palladium, and of osmium, iridium, and platinum, are very close to
each other, and their properties are also very much alike--the
differences are not very perceptible. And if the properties of atoms are
a function of their weight, many ideas which have more or less rooted
themselves in chemistry must suffer change and be developed and worked
out in the sense of this deduction. Although at first sight it appears
that the chemical elements are perfectly independent and individual,
instead of this idea of the nature of the elements, the notion of the
dependence of their properties upon _their mass_ must now be established;
that is to say, the subjection of the individuality of the elements to a
common higher principle which evinces itself in gravity and in all
physico-chemical phenomena. Many chemical deductions then acquire a new
sense and significance, and a regularity is observed where it would
otherwise escape attention. This is more particularly apparent in the
physical properties, to the consideration of which we shall afterwards
turn, and we will now point out that Gustavson first (Chapter X., Note
28) and subsequently Potilitzin (Chapter XI., Note 66) demonstrated the
direct dependence of the reactive power on the atomic weight and that
fundamental property which is expressed in the forms of their compounds,
whilst in a number of other cases the purely chemical relations of
elements proved to be in connection with their periodic properties. As a
case in point, it may be mentioned that Carnelley remarked a dependence
of the decomposability of the hydrates on the position of the elements in
the periodic system; whilst L. Meyer, Willgerodt, and others established
a connection between the atomic weight or the position of the elements in
the periodic system and their property of serving as media in the
transference of the halogens to the hydrocarbons.[16] Bailey pointed out
a periodicity in the stability (under the action of heat) of the oxides,
namely: (_a_) in the even series (for instance, CrO_{3}, MoO_{3}, WO_{3},
and UO_{3}) the higher oxides of a given group decompose with greater
ease the smaller the atomic weight, while in the uneven series (for
example, CO_{2}, GeO_{2}, SnO_{2}, and PbO_{2}) the contrary is the case;
and (_b_) the stability of the higher saline oxides in the even series
(as in the fourth series from K_{2}O to Mn_{2}O_{7}) decreases in passing
from the lower to the higher groups, while in the uneven series it
increases from the Ist to the IVth group, and then falls from the IVth to
the VIIth; for instance, in the series Ag_{2}O, CdO, In_{2}O_{3},
SnO_{2}, and then SnO_{2}, Sb_{2}O_{5}, TeO_{3}, I_{2}O_{7}. K. Winkler
looked for and actually found (1890) a dependence between the
reducibility of the metals by magnesium and their position in the
periodic system of the elements. The greater the attention paid to this
field the more often is a distinct connection found between the variation
of purely chemical properties of analogous substances and the variation
of the atomic weights of the constituent elements and their position in
the periodic system. Besides, since the periodic system has become more
firmly established, many facts have been gathered, showing that there are
many similarities between Sn and Pb, B and Al, Cd and Hg, &c., which had
not been previously observed, although foreseen in some cases, and a
consequence of the periodic law. Keeping our attention in the same
direction, we see that the most widely distributed elements in nature are
those with small atomic weights, whilst in organisms the lightest
elements exclusively predominate (hydrogen, carbon, nitrogen, oxygen),
whose small mass facilitates those transformations which are proper to
organisms. Poluta (of Kharkoff), C. C. Botkin, Blake, Brenton, and others
even discovered a correlation between the physiological action of salts
and other reagents on organisms and the positions occupied in the
periodic system by the metals contained in them.[17]
[16] Meyer, Willgerodt, and others, guided by the fact that Gustavson
and Friedel had remarked that metalepsis rapidly proceeds in the
presence of aluminium, investigated the action of nearly all the
elements in this respect. For example, they took benzene, added
the metals to be experimented on to it, and passed chlorine
through the liquid in diffused light. When, for instance, sodium,
potassium, barium, &c. are taken, there is no action on the
benzene; that is, hydrochloric acid is not disengaged; but if
aluminium, gold, or, in general, any metal having this power of
aiding chlorination (Halogen-überträger) is employed, then the
action is clearly seen from the volumes of hydrochloric acid
evolved (especially if the metallic chloride formed is soluble in
benzene). Thus, in group I., and in general among the even and
light elements, there are none capable of serving as agents of
metalepsis; but aluminium, gallium, indium, antimony, tellurium,
and iodine, which are contiguous members in the periodic system,
are excellent transmitters (carriers) of the halogens.
[17] The periodic relations enumerated above appertain to the real
elements, and not to the elements in the free state as we know
them; and it is very important to note this, because the periodic
law refers to the real elements, inasmuch as the atomic weight is
proper to the real element, and not to the 'free' element, to
which, as to a compound, a molecular weight is proper. Physical
properties are chiefly determined by the properties of molecules,
and only indirectly depend on the properties of the atoms forming
the molecules. For this reason the periods, which are clearly and
quite distinctly expressed--for instance, in the forms of
combination--become to some extent involved (complicated) in the
physical properties of their members. Thus, for instance, besides
the _maxima_ and _minima_ corresponding with the periods and
groups, new molecules appear; thus, as regards the melting-point
of germanium, a local maximum appears, which was, however,
foreseen by the periodic law when the properties of germanium
(ekasilicon) were forecast.
As, from the necessity of the case, the physical properties must be in
dependence on the composition of a substance, _i.e._ on the quality and
quantity of the elements forming it, so for them also a dependence on the
atomic weight of the component elements must be expected, and
consequently also on their periodic distribution. We shall meet with
repeated proofs of this in the further exposition of our treatise, and
for the present will content ourselves with citing the discovery by
Carnelley in 1879 of the dependence of the magnetic properties of the
elements on the position occupied by them in the periodic system.
Carnelley showed that all the elements of the _even series_ (beginning
with lithium, potassium, rubidium, cæsium) belong to the number of
magnetic (paramagnetic) substances; for example, according to Faraday and
others,[17 bis] C, N, O, K, Ti, Cr, Mn, Fe, Co, Ni, Ce, are magnetic; and
the elements of the _uneven series are diamagnetic_, H, Na, Si, P, S, Cl,
Cu, Zn, As, Se, Br, Ag, Cd, Sn, Sb, I, Au, Hg, Tl, Pb, Bi.
[17 bis] The relation of certain elements (for instance, the analogues
of Pt) among diamagnetic and paramagnetic bodies is sometimes
doubtful (probably partly owing to the imperfect purity of the
reagents under investigation). This subject has been studied in
some detail by Bachmetieff in 1889.
Carnelley also showed that the _melting-point_ of elements varies
periodically, as is seen by the figures in Table III. (nineteenth
column),[18] where all the most trustworthy data are collected, and
predominance is given to those having maximum and minimum values.[19]
[18] It is evident that many of the figures, especially those exceeding
1000°, have been determined with but little exactitude, and some,
placed in Table III. with the sign (?), I have only given on the
basis of rough and comparative determinations, calculated from the
melting-points of silver and platinum, now established by many
observers. In Table III., besides the large periods whose maxima
correspond with carbon, silicon, titanium, ruthenium (?), and
osmium (?), there are also small periods in the melting-points,
and their maxima correspond with sulphur, arsenic, antimony. The
minima correspond with the halogens and metals of the alkalis. A
distinct periodicity is also seen in taking the coefficients of
linear expansion (chiefly according to Fizeau); for instance, in
the vertical series (according to the magnitude of the atomic
weight), Fe, Co, Ni, Cu, the linear expansion in millionths of an
inch = 12, 13, 17, and 29; for Rh, Pd, Ag, Cd, In, Sn, and Sb the
coefficients are 8, 12, 19, 31, 46, 26, and 12, so that a maximum
is reached at In. In the series Ir (7), Pt (5), Au (14), Hg (60),
Tl (31), Pb (29), and Bi (14), the maximum is at Hg and the
minimum at Pt. Raoul Pictet expressed this connection by the fact
that he found the product [alpha](_t_ + 273)[3root](A/_d_) to be
nearly constant for all elements in the free state, and nearly
equal to 0·045, and being the coefficient of linear expansion, _t_
+ 273, the melting-point calculated from the absolute zero
(-273°), and [3root](A/_d_), the mean distance between the atoms,
if A is the atomic weight and _d_ the sp. gr. of an element.
Although the above product is not strictly constant, nevertheless
Pictet's rule gives an idea of the bond between magnitudes which
ought to have a certain connection with each other. De Heen,
Nadeschdin, and others also studied this dependence, but their
deductions do not give a general and exact law.
[19] Carnelley found a similar dependence in comparing the
melting-points of the metallic chlorides, many of which he
redetermined for this purpose. The melting-points (and
boiling-points, in brackets) of the following chlorides are known,
and a certain regularity is seen to exist in them, although the
number (and degree of accuracy) of the data is insufficient for a
generalisation:--
LiCl 598° BeCl_{2} 600° BCl_{3} -20°
NaCl 772° MgCl_{2} 708° AlCl_{3} 187°
KCl 734° CaCl_{2} 719° ScCl_{3} ?
{CuCl 434° ZnCl_{2} 262° GaCl_{3} 76°
{(993°) (680°) (217°)
AgCl 451° CdCl_{2} 541° InCl_{3} ?
{TlCl 427° PbCl_{2} 498° BiCl_{3} 227°
{(713°) (908°)
We will also enumerate the following data given by Carnelley,
which are interesting for comparison: HCl -112° (-102°); RbCl
710°, SrCl_{2} 825°, CsCl 631°, BaCl_{2} 860°, SbCl_{3} 73°
(223°), TeCl_{2} 209° (327°), ICl 27°, HgCl_{2} 276° (303°),
FeCl_{3} 306°, NbCl_{5} 194° (240°), TaCl_{3} 211° (242°), WCl_{6}
190°. The melting-points of the bromides and iodides are higher or
lower than those of the corresponding chlorides, according to the
atomic weight of the element and number of atoms of the halogen,
as is seen from the following examples:--1. KCl 734°, KBr 699°, KI
634°; 2. AgCl 454°, AgBr 427° AgI 527°; 3. PbCl_{2} 498° (900°),
PbBr_{2} 499° (861°), PbI_{2} 383° (906°); 4. SnCl_{4} below -20°
(114°), SnBr_{2} 30° (201°), SnI_{4} 146° (295°) (_see_ Chapter
II. Note 27, and Chapter XI. Note 47^{bis}, &c.)
Laurie (1882) also observed a periodicity in the _quantity of
heat_ developed in the formation of the chlorides, bromides, and
iodides (fig. 79), as is seen from the following figures, where
the heat developed is expressed in thousands of calories, and
referred to a molecule of chlorine, Cl_{2}, so that the heat of
formation of KCl is doubled, and that of SnCl_{4} halved, &c.: Na
195 (Ag 59, Au 12), Mg 151 (Zn 97, Cd 93, Hg 63), Al 117, Si 79
(Sn 64), K 211 (Li 187), Ca 170 (Sr 185, Ba 194), whence it is
seen that the greatest amount of heat is evolved by the metals of
the alkalis, and that in each period it falls from them to the
halogens, which evolve very little heat in combining together.
Richardson, by comparing the heats of formation of the fluorides
also came to the conclusion that they are in periodic dependence
upon the atomic weights of the combined elements.
[Illustration: FIG. 79.--Laurie's diagram for expressing the
periodic variation of the heat of formation of the chlorides. The
abscissæ give the atomic weights from 0 to 210, and the ordinates
the amounts of heat from 0 to 220 thousand calories evolved in the
combination with Cl_{2}, (_i.e._ with 71 parts of chlorine). The
apices of the curve correspond to Li, Na, K, Rb, Cs, and the lower
extremities to F, Cl, Br, and I.]
In this respect it may not be superfluous to remark (1) that
Thomsen, whose results I have employed above, observed a
correlation in the calorific equivalents of analogous elements,
although he did not remark their periodic variation; (2) that the
uniformity of many thermochemical deductions must gain
considerably by the application of the periodic law, which
evidently repeats itself in calorimetric data; and if these data
frequently lead to true forecasts, this is due to the periodicity
of the thermal as well as of many other properties, as Laurie
remarked; and (3) that the heat of formation of the oxides is also
subject to a periodic dependence which differs from that of the
heat of formation of the chlorides, in that the greatest quantity
corresponds with the bivalent metals of the alkaline earths
(magnesium, calcium, strontium, barium), and not with the
univalent metals of the alkalis, as is the case with chlorine,
bromine, and iodine. This circumstance is probably connected with
the fact that chlorine, bromine, and iodine are univalent
elements, and oxygen bivalent (compare, for instance, Chapter XI.,
Note 13, Chapter XXII., Note 40, Chapter XXIV., Note 28^{bis},
&c.)
Keyser (1892), in investigating the spectra of the alkali metals
and metals of the alkaline earths, came to the conclusion that in
this respect also there is a regularity of a periodic character in
dependence upon the atomic weights. Probably a closer and
systematic study of many of the properties of the elements and of
complex and simple bodies formed by them will more and more
frequently lead to similar conclusions, and to extending the range
of application of the periodic law.
There is no doubt that many other physical properties will, when
further studied, also prove to be in periodic dependence on the atomic
weights,[19 bis] but at present only a few are known with any
completeness, and we will only refer to the one which is the most easily
and frequently determined--namely, the _specific gravity_ in a solid and
liquid state, the more especially as its connection with the chemical
properties and relations of substances is shown at every step. Thus, for
instance, of all the metals those of the alkalis, and of all the
non-metals the halogens, are the most energetic in their reactions, and
they have the lowest specific gravity among the adjacent elements, as is
seen in Table III., column 17. Such are sodium, potassium, rubidium,
cæsium among the metals, and chlorine, bromine, and iodine among the
non-metals; and as such less energetic metals as iridium, platinum, and
gold (and even charcoal or the diamond) have the highest specific gravity
among the elements near to them in atomic weight; therefore the degree of
the condensation of matter evidently influences the course of the
transformations proper to a substance, and furthermore this dependence on
the atomic weight, although very complex, is of a clearly periodic
character. In order to account for this to some extent, it may be
imagined that the lightest elements are porous, and, like a sponge, are
easily penetrated by other substances, whilst the heavier elements are
more compressed, and give way with difficulty to the insertion of other
elements. These relations are best understood when, instead of the
specific gravities referring to a unit of volume,[20] the _atomic volumes
of the elements_--that is, the quotient _A_/_d_ of the atomic weight _A_
by the specific gravity _d_--are taken for comparison. As, according to
the entire sense of the atomic theory, the actual matter of a substance
does not fill up its whole cubical contents, but is surrounded by a
medium (ethereal, as is generally imagined), like the stars and planets
which travel in the space of the heavens and fill it, with greater or
less intervals, so the quotient _A_/_d_ only expresses the _mean_ volume
corresponding to the sphere of the atoms, and therefore [3root]_A_/_d_
_is the mean distance between the centres of the atoms_. For compounds
whose molecules weigh _M_, the mean magnitude of the atomic volume is
obtained by dividing the mean molecular volume _M_/_d_ by the number of
atoms _n_ in the molecule.[21] The above relations may easily be
expressed from this point of view by comparing the atomic volumes. Those
comparatively light elements which easily and frequently enter into
reaction have the greatest atomic volumes: sodium 23, potassium 45,
rubidium 57, cæsium 71, and the halogens about 27; whilst with those
elements which enter into reaction with difficulty, the mean atomic
volume is small; for carbon in the form of a diamond it is less than 4,
as charcoal about 6, for nickel and cobalt less than 7, for iridium and
platinum about 9. The remaining elements having atomic weights and
properties intermediate between those elements mentioned above have also
intermediate atomic volumes. Therefore _the specific gravities and
specific volumes of solids and liquids stand in periodic dependence on
the atomic weights_, as is seen in Table III., where both _A_ (the atomic
weight) and _d_ (the specific gravity), and _A_/_d_ (specific volumes of
the atoms) are given (column 18).
[19 bis] Probably, besides thermo-chemical data (Note 19), the
refractive index, cohesion, ductility, and similar properties of
corresponding compounds or of the elements themselves will be
found to exhibit a dependence of the magnitude of the atomic
weight upon the periodic law.
[20] Having occupied myself since the fifties (my dissertation for the
degree of M.A. concerned the specific volumes, and is printed in
part in the _Russian Mining Journal_ for 1856) with the problems
concerning the relations between the specific gravities and
volumes, and the chemical compositions of substances, I am
inclined to think that the direct investigation of specific
gravities gives essentially the same results as the investigation
of specific volumes, only that the latter are more graphic. Table
III. of the periodic properties of the elements clearly
illustrates this. Thus, for those members whose volume is the
greatest among the contiguous elements, the specific gravity is
least--that is, the periodic variation of both properties is
equally evident. In passing, for instance, from silver to iodine
we have a successive decrease of specific gravity and successive
increase of specific volume. The periodic alternation of the rise
and fall of the specific gravity and specific volume of the free
elements was communicated by me in August 1869 to the Moscow
Meeting of Russian Naturalists. In the following year (1870) L.
Meyer's paper appeared, which also dealt with the specific volume
of the elements.
[21] In my opinion the mean volume of the atoms of compounds deserves
more attention than has yet been paid to it. I may point out, for
instance, that for feebly energetic oxides the mean volume of the
atom is generally nearly 7; for example, the oxides SiO_{2},
Sc_{2}O_{3}, TiO_{2}, V_{2}O_{5}, as well as ZnO, Ga_{2}O_{3},
GeO_{2}, ZrO_{2}, In_{2}O_{3}, SnO_{2}, Sb_{2}O_{5}, &c., whilst
the mean volume of the atom of the alkali and acid oxides is
greater than 7. Thus we find in the magnitudes of the mean volumes
of the atom in oxides and salts both a periodic variation and a
connection with their energy of essentially the same character as
occurs in the case of the free elements.
Thus we find that in the large periods beginning with lithium, sodium,
potassium, rubidium, cæsium, and ending with fluorine, chlorine, bromine,
iodine, the extreme members (energetic elements) have a small density and
large volume, whilst the intermediate substances gradually increase in
density and decrease in volume--that is, as the atomic weight increases
the density rises and falls, again rises and falls, and so on.
Furthermore, the energy decreases as the density rises, and the greatest
density is proper to the atomically heaviest and least energetic
elements; for example, Os, Ir, Pt, Au, U.
In order to explain the relation between the volumes of the elements and
of their compounds, the densities (column S) and volumes (column M/_s_)
of some of the higher saline oxides arranged in the same order as in the
case of the elements are given on p. 36. For convenience of comparison
the volumes of the oxides are all calculated per two atoms of an element
combined with oxygen. For example, the density of Al_{2}O_{3} = 4·0,
weight Al_{2}O_{3} = 102, volume Al_{2}O_{3} = 25·5. Whence, knowing the
volume of aluminium to be 11, it is at once seen that in the formation of
aluminium oxide, 22 volumes of it give 25·5 volumes of oxide. A distinct
periodicity may also be observed with respect to the specific gravities
and volumes of the higher saline oxides. Thus in each period, beginning
with the alkali metals, the specific gravity of the oxides first rises,
reaches a maximum, and then falls on passing to the acid oxides, and
again becomes a minimum about the halogens. But it is especially
important to call attention to the fact that the volume of the alkali
oxides is less than that of the metal contained in them, which is also
expressed in the last column, giving this difference for each atom of
oxygen.[22] Thus 2 atoms of sodium, or 46 volumes, give 24 volumes of
Na_{2}O, and about 37 volumes of 2NaHO--that is, the oxygen and hydrogen
in distributing themselves in the medium of sodium have not only not
increased the distance between its atoms, but have brought them nearer
together, have drawn them together by the force of their great affinity,
by reason, it may be presumed, of the small mutual attraction of the
atoms of sodium. Such metals as aluminium and zinc, in combining with
oxygen and forming oxides of feeble salt-forming capacity, hardly vary in
volume, but the common metals and non-metals, and especially those
forming acid oxides, always give an increased volume when oxidised--that
is, the atoms are set further apart in order to make room for the oxygen.
The oxygen in them does not compress the molecule as in the alkalis; it
is therefore comparatively easily disengaged.
[22] The volume of oxygen (judging by the table on p. 36) is evidently
a variable quantity, forming a distinctly periodic function of the
atomic weight and type of the oxide, and therefore the efforts
which were formerly made to find the volume of the atom of oxygen
in the volumes of its compounds may be considered to be futile.
But since a distinct contraction takes place in the formation of
oxides, and the volume of an oxide is frequently less than the
volume in the free state of the element contained in it, it might
be surmised that the volume of oxygen in a free state is about 15,
and therefore the specific gravity of solid oxygen in a free state
would be about O·9.
S M/_s_ Volume of Oxygen
H_{2}O 1·0 18 ?- 22
Li_{2}O 2·0 15 - 9
Be_{2}O_{2} 3·06 16 + 2·6
B_{2}O_{3} 1·8 39 + 10·0
C_{2}O_{4} 1·6 55 + 10·6
N_{2}O_{5} 1·64 66 ?+ 4
Na_{2}O 2·6 24 - 22
Mg_{2}O_{2} 3·5 23 - 4·5
Al_{2}O_{3} 4·0 26 + 1·3
Si_{2}O_{4} 2·65 45 + 5·2
P_{2}O_{5} 2·39 59 + 6·2
S_{2}O_{6} 1·96 82 + 8·7
Cl_{2}O_{7} ?1·92 95 + 6
K_{2}O 2·7 35 - 35
Ca_{2}O_{2} 3·25 34 - 8
Sc_{2}O_{3} 3·86 35 ? 0
Ti_{2}O_{4} 4·2 38 + 3
V_{2}O_{5} 3·49 52 + 6·7
Cr_{2}O_{6} 2·74 73 + 9·5
Cu_{2}O 5·9 24 + 9·6
Zn_{2}O_{2} 5·7 23 + 4·8
Ga_{2}O_{3} ?5·1 36 + 4
Ge_{2}O_{4} 4·7 44 + 4·5
As_{2}O_{5} 4·1 56 + 6·0
Sr_{2}O_{2} 4·7 44 - 13
Y_{2}O_{3} 5·0 45 ?- 2
Zr_{2}O_{4} 5·5 44 0
Nb_{2}O_{5} 4·7 57 + 6
MoO_{6} 4·4 65 + 6·8
Ag_{2}O 7·5 31 + 11
Cd_{2}O_{3} 8·0 32 + 3
In_{2}O_{3} 7·18 38 + 2·7
Sn_{2}O_{4} 7·O 43 + 2·7
Sb_{2}O_{5} 6·5 49 + 2·6
TeO_{6} 5·1 68 + 4·7
Ba_{2}O_{2} 5·7 52 - 10
La_{2}O_{3} 6·5 50 + 1
Ce_{2}O_{4} 6·74 50 + 2
Ta_{2}O_{5} 7·5 59 + 4·6
W_{2}O_{6} 6·8 68 + 8·2
Hg_{2}O_{2} 11·1 39 + 4·5
Pb_{2}O_{4} 8·9 53 + 4·2
Th_{2}O_{4} 9·86 54 + 2
As the volumes of the chlorides, organo-metallic and all other
corresponding compounds, also vary in a like periodic succession with a
change of elements, it is evidently possible to indicate the properties
of substances yet uninvestigated by experimental means, and even those of
yet undiscovered elements. It was possible by following this method to
foretell, on the basis of the periodic law, many of the properties of
scandium, gallium, and germanium, which were verified with great accuracy
after these metals had been discovered.[23] The periodic law, therefore,
has not only embraced the mutual relations of the elements and expressed
their analogy, but has also to a certain extent subjected to law the
doctrine of the types of the compounds formed by the elements: it has
enabled us to see a regularity in the variation of all chemical and
physical properties of elements and compounds, and has rendered it
possible to foretell the properties of elements and compounds yet
uninvestigated by experimental means; thus it has prepared the ground for
the building up of atomic and molecular mechanics.[24]
[23] As an example we will take indium oxide, In_{2}O_{3}. Its sp. gr.
and sp. vol. should be the mean of those of cadmium oxide,
Cd_{2}O_{2}, and stannic oxide, Sn_{2}O_{4}, as indium stands
between cadmium and tin. Thus in the seventies it was already
evident that the volume of indium oxide should be about 38, and
its sp. gr. about 7·2, which was confirmed by the determinations
of Nilson and Pettersson (7·179) made in 1880.
[24] As the distance between, and the volumes of, the molecules and
atoms of solids and liquids certainly enter into the data for the
solution of the problems of molecular mechanics, which as yet have
only been worked out to any extent for the gaseous state, the
study of the specific gravity of solids, and especially of
liquids, has long had an extensive literature. With respect to
solids, however, a great difficulty is met with, owing to the
specific gravity varying not only with a change of isomeric state
(for example, for silica in the form of quartz = 2·65, and in
tridymite = 2·2) but also directly under mechanical pressure (for
example, in a crystalline, cast, and forged metal), and even with
the extent to which they are powdered, &c., which influences are
imperceptible in liquids. Compare Chapter XIV., Note 55^{bis}.
Without going into further details, we may add to what has been
said above that the conception of specific volumes and atomic
distances has formed the subject of a large number of researches,
but as yet it is only possible to lay down a few generalisations
given by Dumas, Kopp, and others, which are mentioned and
amplified by me in my work cited in Note 20, and in my memoirs on
this subject.
1. Analogous compounds and their isomorphs have frequently
approximately the same molecular volumes.
2. Other compounds, analogous in their properties, exhibit
molecular volumes which increase with the molecular weight.
3. When a contraction takes place in combination in a gaseous
state, then contraction is in the majority of instances also to be
observed in the solid or liquid state--that is, the sum of the
volumes of the reacting substances is greater than the volume of
the resultant substance or substances.
4. In decomposition the reverse takes place to that which occurs
in combination.
5. In substitution (when the volumes in a state of vapour do not
vary) a very small change of volume generally takes place--that
is, the sum of the volumes of the reacting substances is almost
equal to the sum of the resultant substances.
6. Hence it is impossible to judge the volume of the component
substances from the volume of a compound, although it is possible
to do so from the product of substitution.
7. The replacement of H_{2} by sodium, Na_{2}, and by barium, Ba,
as well as the replacement of SO_{4} by Cl_{2}, scarcely changes
the volume, but the volume increases with the replacement of Na by
K, and decreases with the replacement of H_{2}, by Li_{2} Cu, and
Mg.
8. There is no need for comparing volumes in a solid and liquid
state at the so-called corresponding temperatures--that is at
temperatures at which vapour tension is equal in each case. The
comparison of volumes at the ordinary temperature is sufficient
for finding a regularity in the relations of volumes (this
deduction was developed with particular detail by me in 1856).
9. Many investigators (Perseau, Schröder, Löwig, Playfair and
Joule, Baudrimont, Einhardt) have sought in vain for a multiple
proportion in the specific volumes of solids and liquids.
10. The truth of the above is seen very clearly in comparing the
volumes of polymeric substances. The volumes of their molecules
are equal in a state of vapour, but are very different in a solid
and liquid state, as is seen from the close resemblance of the
specific gravities of polymeric substances. But as a rule the more
complex polymerides are denser than the simpler.
11. We know that the hydroxides of light metals have generally a
smaller volume than the metals, whilst that of magnesium hydroxide
is considerably greater, which is explained by the stability of
the former and instability of the latter. In proof of this we may
cite, besides the volumes of the true alkali metals, the volume of
barium (36) which is greater than that of its stable hydroxide
(sp. gr. 4·5, sp. vol. 30). The volumes of the salts of magnesium
and calcium are greater than the volume of the metal, with the
single exception of the fluoride of calcium. With the heavy metals
the volume of the compound is always greater than the volume of
the metal, and, moreover, for such compounds as silver iodide, AgI
(_d_ = 5·7), and mercuric iodide, HgI_{2} (_d_ = 6·2, and the
volumes of the compounds 41 and 73), the volume of the compound is
greater than the sum of the volumes of the component elements.
Thus the sum of the volumes Ag + I = 36, and the volume of AgI =
41. This stands out with particular clearness on comparing the
volumes K + I = 71 with the volume of KI, which is equal to 54,
because its density = 3·06.
12. In such combinations, between solids and liquids, as
solutions, alloys, isomorphous mixtures, and similar feeble
chemical compounds, the sum of the reacting substances is always
very nearly that of the resulting substance, but here the volume
is either slightly larger or smaller than the original; speaking
generally, the amount of contraction depends on the force of
affinity acting between the combining substances. I may here
observe that the present data respecting the specific volumes of
solid and liquid bodies deserve a fresh and full elaboration to
explain many contradictory statements which have accumulated on
this subject.
CHAPTER XVI
ZINC, CADMIUM, AND MERCURY
These three metals give, like magnesium, oxides RO, which form feebly
energetic bases, and like magnesium they are volatile. The volatility
increases with the atomic weight. Magnesium can be distilled at a white
heat, zinc at a temperature of about 930°, cadmium about 770°, and
mercury about 351°. Their oxides, RO, are more easily reducible than
magnesia, and mercuric oxide is the most easily reducible. The properties
of their salts RX_{2} are very similar to the properties of MgX_{2}.
Their solubility, power of forming double and basic salts, and many other
qualities are in many respects identical with those of MgX_{2}. The
greater or less ease with which they are oxidised, the instability of
their compounds, the density of the metals and their compounds, their
scarcity in nature, and many other properties gradually change with the
increase of atomic weight, as might be expected from the periodicity of
the elements. Their principal characteristics, as contrasted with
magnesium, find a general expression in the fact that zinc, cadmium, and
mercury are heavy metals.
_Zinc_ stands nearest to magnesium in atomic weight and in properties.
Thus zinc sulphate, or white vitriol, easily crystallises with seven
molecules of water, ZnSO_{4},7H_{2}O. It is isomorphous with Epsom salts,
and parts with difficulty with the last molecule of water; it forms
double salts--for instance, ZnK_{2}(SO_{4})_{2},6H_{2}O--exactly as
magnesium sulphate does.[1] _Zinc oxide_, ZnO, is a white powder, almost
insoluble in water,[2] like magnesia, from which, however, it is
distinguished by its solubility in solutions of sodium and potassium
hydroxides.[3] Zinc chloride[4] is decomposed by water, combines with
ammonium chloride, potassium chloride, &c., just like magnesium chloride,
forms an oxychloride, and also combines with zinc oxide.[4 bis]
[1] Zinc sulphate is often obtained as a by-product--for instance, in
the action of galvanic batteries containing zinc and sulphuric
acid. When the anhydrous salt is heated it forms zinc oxide,
sulphurous anhydride, and oxygen. The solubility in 100 parts of
water at O° = 43, 20° = 53, 40° = 63-1/2, 60° = 74, 80° = 84-1/2,
100° = 95 parts of anhydrous zinc sulphate--that is to say, it is
closely expressed by the formula 43 + 0·52_t_.
An admixture of iron is often found in ordinary sulphate of zinc in
the form of ferrous sulphate, FeSO_{4}, isomorphous with the zinc
sulphate. In order to separate it, chlorine is passed through the
solution of the impure salt (when the ferrous salt is converted
into ferric), the solution is then boiled, and zinc oxide is
afterwards added, which, after some time has elapsed, precipitates
all the ferric oxide. Ferric oxide of the form R_{2}O_{3} is
displaced by zinc oxide of the form RO.
[2] Zinc oxide is obtained both by the combustion and oxidation of
zinc, and by the ignition of some of its salts--for instance, those
of carbonic and nitric acids; it is likewise precipitated by
alkalis from a solution of ZnX_{2} in the form of a gelatinous
hydroxide. The oxide produced by roasting zinc blende (by burning
in the air, when the sulphur is converted into sulphurous
anhydride) contains various impurities. For purification, the oxide
is mixed with water, and the sulphurous anhydride formed by
roasting the blende is passed through it. Zinc bisulphite,
ZnSO_{3},H_{2}SO_{3}, then passes into solution. If a solution of
this salt be evaporated, and the residue ignited, zinc oxide, free
from many of its impurities, will remain. Zinc oxide is a light
white powder, used as a paint instead of _white lead_; the basic
salt, corresponding with magnesia alba, is used for the same
purpose. V. Kouriloff (1890) by boiling the hydrate of the oxide
with a 3 p.c. solution of peroxide of hydrogen obtained
Zn_{2}H_{2}O_{4} or the hydrate of the peroxide
(= ZnO_{2}ZnH_{2}O_{2} or a compound of 2ZnO with H_{2}O_{2}),
which did not part with its oxygen at 100°, but only above 120°.
Cadmium gives a similar compound of a yellow colour. Magnesium,
although it does form such a compound, does so with great
difficulty.
[3] For the solution of one part of the oxide 55,400 parts of water are
required. Nevertheless, even in such a weak solution, zinc oxide
(hydroxide, ZnH_{2}O_{2}) changes the colour of red litmus paper.
Zinc oxide is obtained in the wet way by adding an alkali hydroxide
to a solution of a zinc salt--for instance: ZnSO_{4} + 2HKO =
K_{2}SO_{4} + ZnH_{2}O_{2}. The gelatinous precipitate of zinc
hydroxide is _soluble_ in an excess of alkali, which clearly
distinguishes it from magnesia. This solubility of zinc hydroxide
in alkalis is due to the power of zinc oxide to form a compound,
although an unstable one, with alkalis--that is to say, points to
the fact that zinc oxide already partly belongs to the intermediate
oxides. The oxides of the metals above mentioned (except BeO) do
not show this property. The property which metallic zinc itself has
of dissolving in caustic alkali with the disengagement of hydrogen
(the solution is facilitated by contact with platinum or iron)
depends on the formation of such a compound of the oxides of zinc
and the alkali metals. The solution of zinc hydroxide,
ZnH_{2}O_{2}, in potash (in a strong solution), proceeds when these
hydrates are taken in proportion to ZnH_{2}O_{2} + KHO. If such a
solution be evaporated to dryness, water extracts only caustic
potash from the fused residue. When a solution of zinc hydroxide in
strong alkali is mixed with a large mass of water, nearly all the
oxide of zinc is precipitated; and, therefore, in weak solutions, a
large quantity of the alkali is required to effect solution, which
points to the decomposition of the zinc-alkali compounds by water.
If strong alcohol be added to a solution of zinc oxide in sodium
hydroxide, the crystallo-hydrate, 2Zn(OH)(ONa),7H_{2}O, separates.
[4] _Zinc chloride_, ZnCl_{2}, is generally employed in the arts in the
form of a solution obtained by dissolving zinc in hydrochloric
acid. This solution is used for soldering metals, impregnating
wood, &c. The reason why it is thus employed may be understood from
its properties. When evaporated it first parts with its water of
crystallisation; on being further heated, however, it loses all
traces of water, and forms an oily mass of anhydrous salt which
solidifies on cooling. This substance melts at 250°, commences to
volatilise at about 400°, and boils at 730°. The soldering of
metals--that is, the introduction of an easily fusible metal
between two contiguous metallic objects--is hindered by any film of
oxide upon them; and, as heated metals easily oxidise, they are
naturally difficult to solder. Zinc chloride is used to prevent the
oxidation. It fuses on being heated, and, covering the metal with
an oily coating, prevents contact with the air; but even if any
oxide has formed, the free hydrochloric acid generally existing in
the zinc chloride solution dissolves it, and in this way the
metallic surface of the metals to be soldered is preserved fit for
the adhesion of the liquid solder, which, on cooling, binds the
objects together. Much zinc chloride is used also for steeping wood
(telegraph-posts and railway-sleepers) in order to preserve it from
decaying quickly; this preservative action is in all probability
mainly due to the poisonous character of zinc salts (corrosive
sublimate is still more poisonous, and a still better agent to
preserve wood from decay), since decay is due to the action of
lower organisms.
The specific gravity of solutions containing _p_ per cent. of zinc
chloride, ZnCl_{2}, is as follows:
_p_ = 10 20 30 40 50
15°/4° = 1·093 1·184 1·293 1·411 1·554
_ds/dt_ = -3 -5 -7 -8 -9
The last line shows the change of specific gravity for 1° in
ten-thousandth parts for temperatures near 15°. More accurate
determinations of Cheltzoff, personally communicated by him, led
him to conclude that solutions of zinc chloride follow the same
laws as the solutions of sulphuric acid, which will be considered
in Chapter XX.: (1) from H_{2}O to ZnCl_{2},120H_{2}O _s_ = S_{0} +
92·85_p_ + 0·1748_p_^2; (2) from thence to ZnCl_{2},40H_{2}O _s_ =
S_{0} + 93·96_p_ - 0·0126_p_^2; (3) thence to ZnCl_{2},25H_{2}O _s_
= 11481·5 + 96·45(_p_ - 15·89) + 0·4567(_p_ - 15·89)^2; (4) thence
to ZnCl_{2},10H_{2}O _s_ = 12212·1 + 104·82(_p_ - 23·21) +
0·7992(_p_ - 23·21)^2; (5) thence to _p_ = 65 p.c. _s_ = 14606·3 +
140·96(_p_ - 43·05) + 1·4905(_p_ - 43·05)^2, where _s_ is the
specific gravity of the solution at 15°, containing _p_ p.c. of
ZnCl_{2} by weight, taking water at 4° = 10000, and where S_{0} =
9991·6 (specific gravity of water at 15°). The compound of zinc
chloride with hydrochloric acid has been mentioned in Vol. I.
Chapter X.
Zinc chloride has a great affinity for water; it is not only
soluble in it, but in alcohol, and on being dissolved in water
becomes considerably heated, like magnesium and calcium chlorides.
Zinc chloride is capable of taking up water, not only in a free
state, but also in chemical combination with many substances. Thus,
for instance, it is used in organic researches for removing the
elements of water from many of the organic compounds.
[4 bis] When mixed with zinc oxide it forms, with remarkable ease, a
very hard mass of zinc oxychloride, which is applied in the arts;
for instance, in painting, to resist the action of water, or for
cementing such objects as are destined to remain in water. Zinc
oxychloride, ZnCl_{2},3ZnO,2H_{2}O (= Zn_{2}OCl_{2},2ZnH_{2}O_{2}),
is also formed from a solution of zinc chloride by the action of a
small quantity of ammonia on it after heating the precipitate
obtained with the liquid for a considerable time; the admixture of
ammonium salts with a mixture of a strong solution of zinc chloride
with its oxide makes a similar mass, which does not solidify so
rapidly, and is therefore more useful for some purposes. Moisture
and cold do not change the hardened mass of oxychloride, and it
also resists the action of many acids, and a temperature of 300°,
which makes it a useful cement for many purposes. A solution of
magnesium chloride with magnesium oxide forms a similar
oxychloride. The mass solidifies best when there are equal
quantities by weight of zinc in the chloride and oxide, and
therefore when it has the composition Zn_{2}OCl_{2} In preparing
such a cement, naturally zinc oxide alone may be taken, and the
requisite quantity of hydrochloric acid added to it. The capacity
of ZnCl_{2} to combine with water, ZnO, and HCl (and also with
other metallic chlorides) indicates its property to combine with
molecules of other substances, and therefore its compounds with
NH_{3}, and especially a compound, ZnCl_{2}2NH_{3}, similar to
sal-ammoniac, might be expected (_i.e._ 2NH_{4}Cl, in which H_{2}
is replaced by Zn). And indeed it has long been known that ZnCl_{2}
absorbs ammonia and gives solid substances capable of dissociating
with the disengagement of NH_{3}. Among these compounds Isambert
and V. Kouriloff (1894) obtained ZnCl_{2}6NH_{3}, ZnCl_{2}4NH_{3},
ZnCl_{2}2NH_{3}, and ZnCl_{2}NH_{3}. The dissociation tension of
the two last-mentioned compounds at 218° is equal to 43·6 mm. and
6·7 mm. CdCl_{2} also forms similar compounds with NH_{3}
(Kouriloff, 1894).
Zinc, like many heavy metals, is often _found in nature in combination
with sulphur_, forming the so-called _zinc blende_,[5] ZnS. It sometimes
occurs in large masses, often crystallised in cubes; it is frequently
translucent, and has a metallic lustre, although this is not so clearly
developed as in many other metallic sulphides with which we shall
hereafter become acquainted. The ores of zinc also comprise the
carbonate, calamine, and silicate, _siliceous calamine_.
[5] This mineral has been given the name of 'mock-ore,' on account of
its having the appearance (considerable density, 4·06, &c.) of
ordinary metallic ores; it deceived the first miners, because it
did not, like other ores, give metal when simply roasted in air and
fused with charcoal. The white zinc oxide, formed by burning the
vapours of zinc, was also called 'nihil album,' or 'white nothing,'
on account of its lightness.
[Illustration: FIG. 80.--Distillation of zinc in a crucible placed in a
furnace. _o c_, tube along which the vapour passes and condenses.]
Metallic zinc (spelter) is most frequently obtained from the ores
containing the carbonate[6]--that is, from calamine, which is sometimes
found in thick veins: for instance, in Poland, Galicia, in some places on
the banks of the Rhine, and in considerable masses in Belgium and
England. In Russia beds of zinc ore are met with in Poland and the
Caucasus, but the output is small. In Sweden, as early as the fifteenth
century, calamine was worked up into an alloy of zinc and copper (brass),
and Paracelsus produced zinc from calamine; but the technical production
of the metal itself, long ago practised in China, only commenced in
Europe in 1807--in Belgium, when the Abbé Donnet discovered that zinc was
volatile. From that time the production increased until it is now about
150 million kilograms in Germany alone.
[6] It may be here mentioned that by the word _ore_ is meant a hard,
heavy substance dug out of the earth, which is used in
metallurgical works for obtaining the usual heavy metals long known
and used. The natural compounds of sodium, or magnesium, are not
called ores, because magnesium and sodium have not been long
obtainable in quantity. The heavy metals, those which are easily
reduced and do not easily oxidise, are exclusively those which are
directly applied in manufactures. Ores either contain the metals
themselves (for instance, ores of silver or bismuth), and the
metals are then said to be in a native state, or else their sulphur
compounds (blende, mock-ore, pyrites--as, for example, galena, PbS;
zinc blende, ZnS; copper pyrites, CuFeS) or oxides (as the ores of
iron), or salts (calamine, for instance). Zinc is incomparably
rarer than magnesium, and is only well known because it is
transformed from its ores into a metal which finds direct use in
many branches of industry.
The reduction of metallic zinc from its ores is based on the fact that
zinc oxide[7] is easily reduced by charcoal at a red heat: ZnO + C = Zn +
CO. The zinc thus obtained is in a finely divided state and impure, being
mixed with other metals reduced with it, but the greater portion is
_converted into vapour_, from which it easily passes into a liquid or
solid state. The reduction and distillation are carried on in earthenware
retorts, filled with a mixture of the divided ore and charcoal. The
vapours of zinc and gases formed during the reaction escape by means of a
pipe leading downwards, and are led to a chamber where the vapours are
cooled. By this means they do not come into contact with the air, because
the neck of the retort is filled with gaseous carbonic oxide, and
therefore the zinc does not oxidise; otherwise its vapour would burn in
the air.[7 bis] The vapours of zinc, entering into the cooling chamber,
condense into white zinc powder or zinc dust. When the neck of the retort
is heated the zinc is obtained in a liquid state, and is cast into
plates, in which form it is generally sold.
[7] Ores, when extracted from the earth by the miners, are often
enriched by sorting, washing, and other mechanical operations. The
sulphurous ores (and likewise others) are then generally roasted.
Roasting an ore means heating it to redness in air. The sulphur
then burns, and passes off in the form of sulphurous anhydride,
SO_{2}, and the metal oxidises. The roasting is carried on in order
to obtain an oxide instead of a sulphur compound, the oxide being
reducible by charcoal. These methods, introduced ages ago, are met
with in nearly all metallurgical works for practically all ores.
For this reason the preparatory treatment of zinc blende furnishes
zinc oxide: this is already contained in calamine.
[7 bis] with very impure ores, especially such as contain lead (PbS
often accompanies zinc), the vapour of the reduced zinc is allowed
to pass directly into the air. It burns and gives ZnO, which is
used as a pigment.
Commercial zinc is generally impure, containing a mixture of lead,
particles of carbon, iron, and other metals carried over with the
vapours, although they are not volatile at a temperature approaching
1000°. If it be required to obtain pure zinc from the commercial article,
it is subjected to a further distillation in a crucible with a pipe
passing through the bottom, the vapours formed by the heated zinc only
having exit through the pipe cemented into the bottom of the crucible.
Passing through this pipe, the vapours condense to a liquid, which is
collected in a receiver. Zinc thus purified is generally re-melted and
cast into rods, and in this form is often used for physical and chemical
researches where a pure article is required.[8]
[8] This zinc, although homogeneous, still contains certain impurities,
to remove which it is necessary to prepare some salt of zinc in a
pure state and transform it into carbonate, which latter is then
distilled with charcoal; and, as thin sheets of zinc can only be
obtained from very pure metal, they are frequently made use of in
cases where pure zinc is required. In order to remove the arsenic
from zinc, it was proposed to melt it and mix it with anhydrous
magnesium chloride, by which means vapours of zinc chloride and
arsenic chloride are formed. Perfectly pure zinc is made (V. Meyer
and others) by decomposing, by means of the galvanic current, a
solution of zinc sulphate to which an excess of ammonia has been
added. The zinc used for Marsh's arsenic test (Chapter XIX.) is
purified from As by fusing it with KNO_{3} and then with ZnCl_{2}.
Metallic zinc has a bluish-white colour; its lustre, compared with many
other metals, is insignificant. When cast it exhibits a crystalline
structure. Its specific gravity is about 7--that is, varies from 6·8 to
7·2, according to the degree of compression (by forging, rolling, &c.) to
which it has been subjected. It is very ductile, considering its
hardness. For this reason it chokes up files when being worked. Its
malleability is considerable when pure, but in the ordinary impure
condition in which it is sold, it is impossible to roll it at the
ordinary temperature, as it easily breaks. At a temperature of 100°,
however, it easily undergoes such operations, and can then be drawn into
wire or rolled into sheets. If heated further it again becomes brittle,
and at 200° may be even crushed into powder, so completely does it lose
its molecular cohesion. It melts at 418°, and distils at 930°.
Zinc does not undergo any change in the atmosphere. Even in very damp air
it only becomes slowly coated with a very thin white coating of oxide.
For this reason it is available for all objects which are only in contact
with air. Therefore sheet zinc may be used for roofing and many other
purposes.[9] This great unchangeability of zinc in the air shows its
slight energy with regard to oxygen compared with the metals already
mentioned, which are capable of reducing zinc from solutions. But zinc
plays this part with regard to the remaining metals--for example, it
reduces salts of lead, copper, mercury, &c. Although zinc is an almost
unoxidisable metal at the ordinary temperature, it burns in the air on
being heated, particularly when in the form of shavings or in the
condition of vapour. At the ordinary temperature zinc does not decompose
water--at any rate, if the metal be in a dense mass. But even at a
temperature of 100° zinc begins little by little to decompose water; it
easily displaces the hydrogen of acids at the ordinary temperature, and
of alkalis on being heated.
[9] Cornices and other architectural ornaments, remarkable for their
lightness and beauty, are stamped out of sheet zinc. Zinc-roofing
does not require painting, but it melts during a conflagration, and
even burns at a strong heat. Many iron vessels, &c., are covered
with zinc ('galvanised') in order to prevent them from rusting.
In this respect the action of zinc varies a great deal with the degree
of its purity. Weak sulphuric acid (corresponding with the composition
H_{2}SO_{4},8H_{2}O) at the ordinary temperature does not act at all on
chemically pure zinc, and even a stronger solution acts very slowly. If
the temperature be raised, and particularly if the zinc be previously
slightly heated, so as to cover the surface with a film of oxide,
chemically pure zinc acts on sulphuric acid. Thus, for example, one cubic
centimetre of zinc in sulphuric acid having a composition
H_{2}SO_{4},6H_{2}O at the ordinary temperature in two hours only
dissolves to the extent of 0·018 gram, and at a temperature of 100° about
3·5 grams. If we compare this slow action with that rapid evolution of
hydrogen which occurs in the case of commercial zinc, we see that the
influence of those impurities in the zinc is very great. Every particle
of charcoal or iron introduced into the mass of the zinc, and likewise
the connection of the zinc with a piece of another electro-negative
metal, assists such a dissolution. The slowness of the action of
sulphuric acid on pure zinc (and likewise on amalgamated zinc) may also
be explained by the fact that a layer of hydrogen[10] collects on the
surface of the metal, preventing contact between the acid and the
metal.[10 bis]
[10] Veeren (1891) proved this by simple experiments, finding that in
vacuo the solution proceeds far more rapidly for both pure and
commercial zinc, and still more rapidly in the presence of
oxidising agents (which absorb the hydrogen) like CrO_{3} and
H_{2}O_{2}.
[10 bis] The addition of cupric sulphate, or, better still, a few drops
of platinic chloride (the metals become reduced), to the sulphuric
acid greatly accelerates the evolution of the hydrogen, because in
this case, as with commercial zinc, galvanic couples are formed
locally by the copper or platinum and the zinc, under the
influence of which the zinc rapidly dissolves. The action of acids
on metallic zinc of various degrees of purity has been the subject
of many investigations, particularly important with reference to
the application of zinc in galvanic batteries, whilst some
investigations have direct significance for chemical mechanics,
although from many points of view the matter is not clear. I
consider it useful to mention certain of these investigations.
Calvert and Johnson made the following series of observations on
the action of sulphuric acid of various degrees of concentration
on 2 grams of pure zinc during two hours. In the cold the
concentrated acid, H_{2}SO_{4}, does not act, H_{2}SO_{4},2H_{2}O
dissolves about 0·002 gram, but principally forms hydrogen
sulphide, which is obtained also when the dilution reaches
H_{2}SO_{4},7H_{2}O, when 0·035 gram of zinc is dissolved. When
largely diluted with water, pure hydrogen begins to be disengaged.
H_{2}SO_{4},2H_{2}O at 130° gives a mixture of hydrogen sulphide
and sulphurous anhydride dissolving 0·156 gram of zinc.
Bouchardat showed that if in a vessel made of glass or sulphur
dilute sulphuric acid acting on a piece of zinc liberates one part
of hydrogen, then the same acid with the same piece of zinc in the
same time will liberate 4 parts of hydrogen if the vessel be made
of tin--that is, zinc forms a galvanic couple with tin; in a
leaden vessel 9 parts of hydrogen are set free, with a vessel of
antimony or bismuth 13 parts, silver or platinum 38 parts, copper
50 parts, iron 43 parts. If a salt of platinum be added to the
dilute sulphuric acid (1 part of acid and 12 parts of water),
Millon determined that the rapidity of the action on the zinc is
increased 149 times, and by the addition of copper sulphate is
rendered 45 times greater than the action of pure sulphuric acid.
The salts which are added are reduced to metals by the zinc, their
contact serving to promote the reaction because they form local
galvanic currents.
According to the observations of Cailletet, if, at the ordinary
pressure, sulphuric acid with zinc liberates 100 parts of
hydrogen, then with a pressure of 60 atmospheres 47 parts will be
liberated and 1 part at a pressure of 120 atmospheres. With a
reduced pressure under the receiver of an air-pump 168 parts are
liberated. Helmholtz showed that a reduced pressure also exercises
its influence on galvanic elements.
Debray, Löwel, and others showed that zinc liberates hydrogen and
forms basic salts and zinc oxide with solutions of many salts--for
instance, MCl_{_n_}, aluminium sulphate, and alum, Sodium and
potassium carbonates scarcely act, because they form carbonates.
The salts of ammonia act more strongly than the salts of potassium
and sodium; the zinc remains bright. It is evident that this
action is founded on the formation of double salts and basic
salts.
The variation with concentration in the rate of the action of
sulphuric acid on zinc (containing impurities) under otherwise
uniform conditions is in evident connection with the electrical
conductivity of the solution and its viscosity, although, when
largely diluted, the action is almost proportional to the amount
of acid in a known volume of the solution. Forging, casting the
molten metal, and similar mechanical influences change the density
and hardness of zinc, and also strongly influence its power of
liberating hydrogen from acids. Kayander showed (1881) that when
magnesium is submitted to the action of acids: (_a_) the action
depends, not on the nature of the acid, but on its basicity; (_b_)
the increase of the action is more rapid than the growth of the
concentration; and (_c_) there is a decrease of action with the
increase of the coefficient of internal friction and electrical
conductivity.
Spring and Aubel (1887) measured the volume of hydrogen disengaged
by an alloy of zinc and a small quantity of lead (0·6 p.c.),
because the action of acids is then uniform. In order to deal with
a known surface, spheres were taken (9·5 millimetres diameter) and
cylinders (17 mm. dia.), the sides of which were covered with wax
in order to limit the action to the end surfaces. During the
commencement of the action of a definite quantity of acid the
rapidity increases, attains a maximum, and then declines as the
acid becomes exhausted. The results for 5, 10, and 15 per cent. of
hydrochloric acid are given below. H denotes the number of cubic
centimetres of hydrogen, D the time in seconds elapsing after the
zinc spheres have been plunged into the acid. At 15° were
obtained:
H = 50 100 200 400 600 800 1000
5 p.c. D = 714 1152 1755 2731 3908 6234 15462
10 p.c. D = 301 455 649 995 1573 2746 6748
15 p.c. D = 106 151 233 440 826 1604 4289
At 35°:
5 p.c. D = 462 705 1058 1700 2525 4132 8499
10 p.c. D = 96 148 239 460 835 1594 3735
15 p.c. D = 44 64 112 255 505 1011 2457
At 55°:
5 p.c. D = 178 276 408 699 1164 2105 5093
10 p.c. D = 34 60 113 258 491 970 2457
15 p.c. D = 24 35 58 136 239 610 1593
In consequence of the complex character of the phenomenon, the
authors themselves do not consider their determinations as being
conclusive, and only give them a relative significance; and in
this connection it is remarkable that hydrobromic acid under
similar conditions (with an equivalent strength) gives a greater
(from 2 to 5 times) rapidity of action than hydrochloric acid, but
sulphuric acid a far smaller velocity (nearly 25 times smaller).
It is also remarkable that during the reaction the metal becomes
much more heated than the acid.
It may be mentioned that zinc dust and zinc itself, when heated
with hydrated lime and similar hydrates, disengages hydrogen: this
method has even been proposed for obtaining hydrogen for filling
war balloons.
The action of zinc on acids, and the consequent formation of zinc
salts, interferes with its application in many cases, particularly for
the preservation of liquids either containing or capable of developing
acid. For this reason zinc vessels ought not to be used for the
preparation or preservation of food, as this often contains acids which
form poisonous salts with the zinc. Even ordinary water, containing
carbonic acid, slowly attacks zinc.
Finely divided zinc, or _zinc dust_, obtained in the distillation of the
metal when the receiver is not heated up to the melting point, on account
of its presenting a large surface of contact and containing foreign
matter (particularly zinc oxide), has in the highest degree the property
of decomposing acids, and even water, which it easily decomposes,
particularly if slightly heated. On this account zinc dust is often used
in laboratories and factories as a reducing agent. A similar influence of
the finely divided state is also noticed in other metals--for instance,
copper and silver--which again shows the close connection between
chemical and physico-mechanical phenomena. We must first of all turn to
this close connection for an explanation of the widely spread application
of zinc in galvanic batteries, where the chemical (latent, potential)
energy of the acting substances is transformed into (evident, kinetic)
galvanic energy, and through this latter into heat, light, or mechanical
work.
Hermann and Stromeyer, in 1819, showed that _cadmium_ is almost always
found with zinc, and in many respects resembles it. When distilled the
cadmium volatilises sooner, because it has a lower boiling point.
Sometimes the zinc dust obtained by the first distillation of zinc
contains as much as 5 per cent. of cadmium. When zinc blende, containing
cadmium, is roasted, the zinc passes into the state of oxide, and the
cadmium sulphide in the ore oxidises into cadmium sulphate, CdSO_{4},
which resists tolerably well the action of heat; therefore if roasted
zinc blende be washed with water, a solution of cadmium sulphate will be
obtained, from which it is very easy to prepare metallic cadmium.
Hydrogen sulphide may be used for separating cadmium from its solutions;
it gives _a yellow precipitate of cadmium sulphide_, CdS (according to
the equation CdSO_{4} + H_{2}S = H_{2}SO_{4} + CdS),[11] which, on
account of its characteristic colour, is used as a pigment.[11 bis]
Cadmium sulphide, when strongly heated in air, leaves cadmium oxide, from
which the metal may be obtained in precisely the same way as in the case
of zinc.
[11] It may be here remarked that sulphate of zinc (especially in the
presence of mineral acids) does not give a precipitate of sulphide
of zinc, or is only slightly precipitated by sulphuretted
hydrogen.
[11 bis] Sulphide of cadmium appears in two varieties of a similar
chemical but different physical character: one is of a lemon
colour, and the other bright red. Kloboukoff (1890) studied the
physical properties of these varieties more closely. The sp. gr.
of the former is 3·906, and of the latter 4·513. They belong to
different crystallographic systems. The first variety may be
converted into the second by friction or pressure, but the second
cannot be converted into the first variety by these means.
Cadmium is a white metal, and when freshly cut is almost as white and
lustrous as tin. It is so soft that it may be easily cut with a knife,
and so malleable that it can be easily drawn into wire, rolled into
sheets, &c. Its specific gravity is 8·67, melting point 320°, boiling
point 770°; its vapours burn, forming a brown powder of the oxide.[12]
Next to mercury it is the most volatile metal; hence Deville determined
the density of its vapours compared with hydrogen, and found it to be
equal to 57·1; therefore the molecule contains _one atom_ whose weight =
112. V. Meyer found the like for zinc; the molecule of mercury also
contains one atom.
[12] Amongst the compounds of cadmium very closely allied to the
compounds of zinc, we must mention _cadmium iodide_, CdI_{2},
which is used in medicine and photography. This salt crystallises
very well: it is prepared by the direct action of iodine, mixed
with water, on metallic cadmium. One part of cadmium iodide at 20°
requires for its solution 1·08 part of water. It may be remarked
that cadmium chloride at the same temperature requires 0·71 part
of water to dissolve it, so that the iodine compound of this metal
is less soluble than the chloride, whilst the reverse relation
holds in the case of the corresponding compounds of the alkali or
alkaline earthy metals. Cadmium sulphate crystallises well, and
has the composition 3CdSO_{4},8H_{2}O, thus differing from zinc
sulphate.
Cadmium oxide is soluble, although sparingly, in alkalis, but in
the presence of tartaric and certain other acids the alkaline
solution of cadmium oxide does not change when boiled, whilst a
_diluted_ solution in that case deposits cadmium oxide: this may
also serve for separating zinc compounds from those of cadmium.
Cadmium is precipitated from its salts by zinc, which fact may
also be taken advantage of for separating cadmium; for this
reason, in an alloy of zinc and cadmium, acids first of all
extract the zinc. Cadmium is in all respects less energetic than
zinc. Thus, for instance, it decomposes water with difficulty, and
this only when strongly heated. It even acts but slowly on acids,
but then displaces hydrogen from them. It is necessary here to
call attention to the fact that for alkali and alkaline earthy
metals (of the even series) the highest atomic weight determines
the greatest energy; but cadmium (of the uneven series), whilst
having a larger atomic weight than zinc, is less energetic. The
salts of cadmium are colourless, like those of zinc. De Schulten
obtained a crystalline oxychloride, Cd(OH)Cl by heating marble
with a solution of cadmium chloride in a sealed tube at 200°.
_Mercury_ resembles zinc and cadmium in many respects, but
presents that distinction from them which is always noticed in all
the heaviest metals (with regard to atomic weight and density)
compared with the lighter ones--namely, that it oxidises with more
difficulty, and its compounds are more easily decomposed.[12 bis]
Besides compounds of the usual type RX_{2}, it also gives those of
the lower type, RX, which are unknown for zinc and cadmium.[13]
Mercury therefore gives salts of the composition HgX (mercurous
salts) and HgX_{2} (mercuric salts), the oxides having the formulæ
Hg_{2}O and HgO respectively.
[12 bis] According to its atomic weight, mercury follows gold in the
periodic system, just as cadmium follows silver and zinc follows
copper:--
Ni = 59 Cu = 63 Zn = 65
Pd = 106 Ag = 108 Cd = 112
Pt = 196 Au = 198 Hg = 200
Eventually we shall see the near relation of platinum, palladium,
and nickel, and also of gold, silver, and copper, but we will now
point out the parallelism between these three groups. The relation
between the physical and also chemical properties is here
strikingly similar. Nickel, palladium, and platinum are very
difficult to fuse (far more so than iron, ruthenium, and osmium,
which stand before them). Copper, silver, and gold melt far more
easily in a strong heat than the three preceding metals, and zinc,
cadmium, and mercury melt still more easily. Nickel, palladium,
and platinum are very slightly volatile; copper, silver, and gold
are more volatile; and zinc, cadmium, and mercury are among the
most volatile metals. Zinc oxidises more easily than copper, and
is reduced with more difficulty, and the same is true for mercury
as compared with gold. These properties for cadmium and silver are
intermediate in the respective groups. Relations of this kind
clearly show the nature of the periodic law.
[13] Thus thallium, lead, and bismuth, following mercury according to
their atomic weights, form, besides compounds of the highest
types, TlX_{3}, PbX_{4}, and BiX_{5}, also the lower ones TlX,
PbX_{2}, and BiX_{3}.
Mercury is found _in nature_ almost exclusively in combination with
sulphur (like zinc and cadmium, but is still rarer than them) in the form
known as cinnabar, HgS (Chapter XX., Note 29). It is far more rarely met
with in the native or metallic condition, and this in all probability has
been derived from cinnabar. Mercury ore is found only in a few
places--namely, in Spain (in Almaden), in Idria, Japan, Peru, and
California. About the year 1880 Minenkoff discovered a rich bed of
cinnabar in the Bahmout district (near the station of Nikitovka), in the
Government of Ekaterinoslav, so that now Russia even exports mercury to
other countries. Cinnabar is now being worked in Daghestan in the
Caucasus. Mercury ores are easily reduced to metallic mercury, because
the combination between the metal and the sulphur is one of but little
stability. Oxygen, iron, lime, and many other substances, when heated,
easily destroy the combination. If iron is heated with cinnabar, iron
sulphide is formed; if cinnabar is heated with lime, mercury and calcium
sulphide and sulphate are formed, 4HgS + 4CaO = 4Hg + 3CaS + CaSO_{4}. On
being heated in the air, or roasted, the sulphur burns, oxidises, forming
sulphurous anhydride, and vapours of metallic mercury are formed. Mercury
is more easily distilled than all other metals, its boiling point being
about 351°, and therefore its separation from natural admixtures,
decomposed by one of the above-mentioned methods, is effected at the
expense of a comparatively small amount of heat. The mixture of mercury
vapour, air, and products of combustion obtained is cooled in tubes (by
water or air), and the mercury condenses as liquid metal.[14]
[14] During the condensation of the vapours of mercury in works, a part
forms a black mass of finely-divided particles, which gives
metallic mercury when worked up in centrifugal machines, or on
pressure, or on re-distillation. In mercury we observe a tendency
to easily split up into the finest drops, which are difficult to
unite into a dense mass. It is sufficient to shake up mercury with
nitric and sulphuric acids in order to produce such a mercury
_powder_. The mercury separated (for instance, reduced by
substances like sulphurous anhydride) from solutions, forms such a
powder. According to the experiments of Nernst, this disintegrated
mercury when entering into reactions develops more heat than the
dense liquid metal--that is to say, the work of disintegration
reappears in the form of heat. This example is instructive in
considering thermochemical deductions.
Mercury, as everybody knows, is a liquid metal at the ordinary
temperature. In its lustre and whiteness it resembles silver.[15] At -39°
mercury is transformed into a malleable crystalline metal; at 0° its
specific gravity is 13·596, and in the solid state at -40° it is
14·39.[16] Mercury does not change in the air--that is to say, it does
not oxidise at the ordinary temperature--but at a temperature approaching
the boiling-point, as was stated in the Introduction, it oxidises,
forming mercuric oxide. Both metallic mercury and its compounds in
general produce salivation, trembling of the hands, and other unhealthy
symptoms which are found in the workmen exposed to the influence of
mercurial vapours or the dust of its compounds.
[15] Mercury may sometimes be obtained in a perfectly pure state from
works (in iron bottles holding about 35 kilos), but after being
used in laboratories (for baths, calibration, &c.) it contains
impurities. It may be purified mechanically in the following way:
a paper filter with a fine hole (pricked with a needle) is placed
in a glass funnel and mercury is poured into it, which slowly
trickles through the hole, leaving the impurities upon the filter.
Sometimes it is squeezed through chamois leather or through a
block of wood (as in the well-known experiment with the air-pump).
It may be purified from many metals by contact with dilute nitric
acid, if small drops of mercury are allowed to pass through a long
column of it (from the fine end of a funnel); or by shaking it up
with sulphuric acid in air. Mercury may be purified by the action
of an electric current, if it be covered with a solution of
HgNO_{3}. But the complete purification of mercury for barometers
and thermometers can only be attained by distillation, best in a
vacuum (the vapour-tension of mercury is given in Chapter II.,
Note 27). For this purpose Weinhold's apparatus is most often
used. The principle of this apparatus is very ingenious, the
distillation being effected in a Torricellian vacuum continuously
supplied with fresh mercury, whilst the condensed mercury is
continuously removed. This process of distillation requires very
little attention, and gives about one kilo of pure mercury per
hour.
[16] If the volume of _liquid_ mercury at 0° be taken as 1000000, then,
according to the determinations of Regnault (recalculated by me in
1875), at _t_ it will be 1000000 + 180·1_t_ + 0·02_t_^2.
As many of the compounds of mercury decompose on being heated--for
instance, the oxide or carbonate[17]--and as zinc, cadmium, copper, iron,
and other metals separate mercury from its salts,[18] it is evident that
mercury has less chemical energy than the metals already described, even
than zinc and cadmium. Nitric acid, when acting on _an excess_ of mercury
at the ordinary temperature, gives mercurous nitrate, HgNO_{3}.[19] The
same acid, under the influence of heat and when in excess (nitric oxide
being liberated), forms mercuric nitrate, Hg(NO_{3})_{2}. This,[20] both
in its composition and properties, resembles the salts of zinc and
cadmium. Dilute sulphuric acid does not act on mercury, but strong
sulphuric acid dissolves it, with evolution of _sulphurous anhydride_
(not hydrogen), and on being slightly heated with an excess of mercury it
forms the sparingly soluble mercurous sulphate, Hg_{2}SO_{4}; but if
mercury be strongly heated with an excess of the acid, the mercuric salt,
HgSO_{4},[21] is formed. Alkalis do not act on mercury, but the
non-metals chlorine, bromine, sulphur, and phosphorus easily combine with
it. They form, like the acids, two series of compounds, HgX and HgX_{2}.
The oxygen compound of the first series is the suboxide of mercury, or
mercurous oxide, Hg_{2}O, and of the second order the oxide HgO, mercuric
oxide. The chlorine compound corresponding with the suboxide is HgCl
(calomel), and with the oxide HgCl_{2} (corrosive sublimate or mercuric
chloride). In the compounds HgX, mercury resembles the metals of the
first group, and more especially silver. In the mercuric compounds there
is an evident resemblance to those of magnesium, cadmium, &c. Here the
atom of mercury is bivalent, as in the type RX_{2}.[22] Every soluble
mercurous compound (corresponding with the type of the suboxide of
mercury), HgX, forms a white precipitate of calomel, HgCl, with
hydrochloric acid or a metallic chloride, because HgCl is very slightly
soluble in water, HgX + MCl = HgCl + MX. In soluble mercuric compounds,
HgX_{2}, hydrochloric acid and metallic chlorides do not form a
precipitate, because corrosive sublimate, HgCl_{2}, is soluble in water.
Alkali hydroxides precipitate the yellow mercuric oxide from a solution
of HgX_{2}, and the black mercurous oxide from HgX. Potassium iodide
forms a dirty greenish precipitate, HgI, with mercurous salts, HgX, and a
red precipitate, HgI_{2}, with the mercuric salts, HgX_{2}. These
reactions distinguish the mercuric from the mercurous salts, which latter
represent the transition from the mercuric salts to mercury itself, 2HgX
= Hg + HgX_{2}. The salts, HgX, as well as HgX_{2}, are reduced by
nascent hydrogen (_e.g._ from Zn + H_{2}SO_{4}), by such metals as zinc
and copper, and also by many reducing agents--for example,
hypophosphorous acid, the lowest grade of oxidation of phosphorus, by
sulphurous anhydride, stannous chloride, &c. Under the action of these
reagents the mercuric salts are first transformed into the mercurous
salts, and the latter are then reduced to metallic mercury. This reaction
is so delicate that it serves to detect the smallest quantity of mercury;
for instance, in cases of poisoning, the mercury is detected by immersing
a copper plate in the solution to be tested, the mercury being then
deposited upon it (more readily on passing a galvanic current). The
copper plate, on being rubbed, shows a silvery white colour; on being
heated, it yields vapours of mercury, and then again assumes its original
red colour (if it does not oxidise). The mercurous compounds, HgX, under
the action of oxidising agents, even air, pass into mercuric compounds,
especially in the presence of acids (otherwise a basic salt is produced),
2HgX + 2HX + O = 2HgX_{2} + H_{2}O; but the mercuric compounds, when in
contact with mercury, change more or less readily, and turn into
mercurous compounds, HgX_{2} + Hg = 2HgX. For this reason, in order to
preserve solutions of mercurous salts, a little mercury is generally
added to them.
[17] All salts of mercury, when mixed with sodium carbonate and heated,
give mercurous or mercuric carbonates; these decompose on being
heated, forming carbonic anhydride, oxygen, and vapours of
mercury.
[18] Spring (1888) showed that, solid dry HgCl is gradually decomposed
in contact with metallic copper. According to the determinations
of Thomsen, the formation of a gram of mercurial compounds from
their elements develops the following amounts of heat (in
thousands of units): Hg_{2} + O, 42; Hg + O, 31; Hg + S, 17; Hg +
Cl, 41; Hg + Br, 34; Hg + I, 24; Hg + Cl_{2}, 63; Hg + Br_{2}, 51;
Hg + I_{2}, 34; Hg + C_{2}N_{2}, 19. These numbers are less than
the corresponding ones for potassium, sodium, calcium, barium, and
for zinc and cadmium--for instance, Zn + O, 85; Zn + Cl_{2}, 97;
Zn + Br_{2}, 76; Zn + I_{2}, 49; Cd + Cl_{2}, 93; Cd + Br_{2}, 75;
Cd + I_{2}, 49.
[19] This salt easily forms the crystallo-hydrate HgNO_{3},H_{2}O,
corresponding with ortho-nitric acid, H_{3}NO_{4} (the terms
ortho-, pyro-, and meta-acids are explained in the chapter on
Phosphorus), with the substitution of Hg for H. In an aqueous
solution this salt can only be preserved in the presence of free
mercury, otherwise it forms basic salts, which will be mentioned
hereafter (Chapter VI., Note 59).
[20] Mercuric nitrate, Hg(NO_{3})_{2},8H_{2}O, crystallises from a
concentrated solution of mercury in an excess of boiling nitric
acid. Water decomposes this salt; at the ordinary temperature
crystals of a basic salt of the composition
Hg(NO_{3})_{2},HgO,2H_{2}O are formed, and with an excess of water
the insoluble yellow basic salt Hg(NO_{3})_{2},H_{2}O,2HgO. These
three salts correspond with the type of ortho-nitric acid,
(H_{3}NO_{4})_{2}, in which mercury is substituted for 1, 2 and 3
times H_{2}. As all these salts still contain water, it is
possible that they correspond with the tetrahydrate = N_{2}O_{5} +
4H_{2}O = N_{2}O(OH)_{8} if ortho-nitric acid = N_{2}O_{5} +
3H_{2}O = 2NO(OH)_{3}.
[21] To obtain the mercuric salt a large excess of strong sulphuric
acid must be taken and strongly heated. With a small quantity of
water colourless crystals of HgSO_{4},H_{2}O may be obtained. An
excess of water, especially when heated, forms the basic salt (as
in Note 20), HgSO_{4},2HgO, which corresponds with trihydrated
sulphuric acid, SO_{3} + 3H_{2}O = S(OH)_{6}, with the
substitution of H_{6} by 3Hg, which in mercuric salts is
equivalent to H_{6}. Le Chatelier (1888) gives the following ratio
between the amounts of equivalents per litre:
HgSO_{4} 0·318 0·890 1·80 2·02
SO_{3} 0·752 1·42 2·10 2·40
--that is, the relative amount of free acid decreases as the
strength of the solution increases.
[22] The question of the molecular weight of calomel--that is, whether
the mercury in the salts of the suboxide is monatomic or
diatomic--long occupied the minds of chemists, although it is not
of very great importance. It is only recently (1894) that this
question can be considered as answered, thanks to the researches
of V. Meyer and Harris, in favour of diatomicity--that is, that
calomel is analogous to peroxide of hydrogen and contains
Hg_{2}Cl_{2} (like O_{2}H_{2}) in its molecule if corrosive
sublimate contains HgCl_{2} (like water OH_{2}). As a matter of
fact, direct experiment gives the vapour density of calomel as
about 118--that is, indicates that its molecule contains HgCl,
whilst the molecule of the sublimate, judging also by the vapour
density (nearly 136), contains HgCl_{2}; it might therefore be
concluded that the mercury in the suboxide is not only monovalent
(corresponding to H) but also monatomic, whilst in the oxide it is
divalent and diatomic. Instances of a variable atomicity, as shown
by the vapour density, are known in N_{2}O, NO, and NH_{3}, CO and
CO_{2}, PCl_{3} and PCl_{5}, and it might therefore be supposed
that the present was a similar instance. But there are also
instances of a variable equivalency which do not correspond to a
variation of atomicity--for example, OH_{2} (water) and OH
(peroxide of hydrogen), CH_{4} (methane), C_{2}H_{5} (ethyl), and
CH_{2} (ethylene), &c. Here, according to the law of substitution,
the residues of OH_{2} and CH_{4} combine together and give
molecules; OHOH = O_{2}H_{2} (peroxide of hydrogen) and
CH_{3}CH_{3} = C_{2}H_{6} (ethane), &c. The same may be assumed
also to be the relation of calomel to sublimate; the residue HgCl,
which is combined with Cl in sublimate, corresponds to HgCl_{2},
and in calomel it may be supposed that this residue is combined
with itself, forming the molecule Hg_{2}Cl_{2}. On this view of
the composition of the molecule of calomel it would follow that in
the state of vapour it breaks up into two molecules, HgCl_{2} and
Hg, when the vapour density would be about 118 (because that of
sublimate is about 136 and that of mercury about 100), and that in
cooling this mixture (like a mixture of HCl and NH_{3}) again
gives Hg_{2}Cl_{2}. It was therefore necessary to prove that
calomel is decomposed in the state of vapour. This was not
effected for a long time, although Odling, as far back as the
thirties, showed that gold becomes amalgamated (_i.e._ absorbs
metallic mercury) in the vapour of calomel, but not in the vapour
of sublimate. Recently, however, V. Meyer and Harris (1894) have
shown that a greater amount of the vapour of mercury than of
calomel passes (at about 465°) through a porous clay cell,
containing calomel. This proves that the vapour of calomel
contains a mixture of the vapours of Hg and HgCl_{2}, as would
follow from the second hypothesis. Moreover, on introducing a
heated piece of KHO into the vapour of calomel, Meyer observed the
formation, not of suboxide (black), but of oxide of mercury
(yellow). Therefore the molecular formula of calomel must be taken
as Hg_{2}Cl_{2} (and not HgCl).
The lowest oxygen compound of mercury--that is, _mercurous oxide_,
Hg_{2}O--does not seem to exist, for the substance precipitated in the
form of a black mass by the action of alkalis on a solution of mercurous
salts gradually separates on keeping into the yellow mercuric oxide and
metallic mercury, as does also a simple mechanical mixture of oxide, HgO,
with mercury (Guibourt, Barfoed). The other compound of mercury with
oxygen is already known to us as _mercuric oxide_, HgO, obtained in the
form of a red crystalline substance by the oxidation of mercury in the
air, and precipitated as a yellow powder by the action of sodium
hydroxide on solutions of salts of the type HgX_{2}. In this case it is
amorphous and more amenable to the action of various reagents (Chap. XI.,
Note 32) than when it is in the crystalline state. Indeed, on
trituration, the red oxide is changed into a powder of a yellow colour.
It is very sparingly soluble in water, and forms an alkaline solution
which precipitates magnesia from the solution of its salts.
Mercury combines directly with chlorine, and the first product of
combination is _calomel_ or _mercurous chloride_, Hg_{2}Cl_{2}. This is
obtained, as above stated, in the form of a white precipitate by mixing
solutions of mercurous salts with hydrochloric acid or with metallic
chlorides. A precipitate of calomel is also obtained by reducing a
boiling aqueous solution of corrosive sublimate, HgCl_{2}, with
sulphurous anhydride. It is likewise produced by heating corrosive
sublimate with mercury.[22 bis] Calomel may be distilled (although in so
doing it decomposes and recombines on cooling from a state of vapour);
its vapour density equals 118 compared with hydrogen (= 1) (_see_ Note
23). In the solid state its specific gravity is 7·0; it crystallises in
rhombic prisms, is colourless, but has a yellowish tint, turns brown from
the action of light, and when boiled with hydrochloric acid decomposes
into mercury and corrosive sublimate. It is used as a strong purgative.
_Corrosive sublimate or mercuric chloride_, HgCl_{2}, can be obtained
from or converted into calomel by many methods.[23] An excess of chlorine
(for instance, _aqua regia_) converts calomel and also mercury into
corrosive sublimate. It owes its name corrosive sublimate to its
volatility, and, in medicine up to the present day, it is termed
_Mercurius sublimatus seu corrosivus_. The vapour density, compared with
hydrogen (= 1) is 135; therefore its molecule contains HgCl_{2}. It forms
colourless prismatic crystals of the rhombic system, boils at 307°, and
is soluble in alcohol. It is usually prepared by subliming a mixture of
mercuric sulphate with common salt, HgSO_{4} + 2NaCl = Na_{2}SO_{4} +
HgCl_{2}. Corrosive sublimate combines with mercuric oxide, forming an
oxychloride or basic salt,[23 bis] of the composition HgCl_{2},2HgO
(magnesium and zinc form similar compounds). This compound is obtained by
mixing a solution of corrosive sublimate with mercuric oxide or with a
solution of sodium bicarbonate. In general, with both mercurous and
mercuric salts, there is a marked tendency to form basic salts.[24]
[22 bis] Calomel (in Japanese 'Keyfun') has been prepared in Japan (and
China) for many centuries, by heating mercury in clay crucibles
with sea salt, which contains MgCl_{2} and gives HCl. The vapour
of the mercury reacts with this HCl and the oxygen of the air and
forms calomel: 2Hg + 2HCl + O = Hg_{2}Cl_{2} + H_{2}O. The calomel
collects on the lid of the crucible in the form of a sublimate
(Divers, 1894).
[23] HgCl_{2} is partially converted into calomel even in the act of
dissolving in ordinary water, especially under the action of light.
[23 bis] As feebly energetic bases (for instance, the oxides MgO, ZnO,
PbO, CuO, Al_{2}O_{3}, Bi_{2}O_{3}, &c.), mercuric oxide (_see_
Notes 20, 21) and mercurous oxide easily give basic salts, which
are usually directly formed by the action of water on the normal
salt, according to the general equation (for mercuric compounds,
RX_{2}):
_n_RX_{2} + _m_H_{2}O = 2_m_HX + (_n_-_m_)RX_{2}_m_RO
neutral salt water acid basic salt
or else are produced directly from the normal salt and the oxide
or its hydroxide. Thus mercurous nitrate, when treated with water,
forms basic salts of the composition 6(HgNO_{3}),Hg_{2}O,H_{2}O,
2(HgNO_{3}),Hg_{2}O,H_{2}O, and 3(HgNO_{3}),Hg_{2}O,H_{2}O, the
first two of which crystallise well. Naturally it is possible
either to refer similar salts to the type of hydrates--for
instance, the second salt to the hydrate N_{2}O_{5},4H_{2}O--or to
view it as a compound, HgNO_{3},HgHO, but our present knowledge of
basic salts is not sufficiently complete to admit of
generalisations. However, it is already possible to view the
subject in the following aspects: (1) basic salts are principally
formed from feeble bases; (2) certain metals (mentioned above)
form them with particular ease, so that one of the causes of the
formation of many basic salts must depend on the property of the
metal itself; (3) those bases which readily form basic salts as a
rule also readily form double salts; (4) in the formation of basic
salts, as also everywhere in chemistry, where sufficient facts
have accumulated, we clearly see the conditions of equally
balanced heterogeneous systems, such as we saw, for instance, in
the formation of double salts, crystallo-hydrates, &c.
The mercuric salts often form double salts (confirming the third
thesis), and mercuric chloride easily combines with ammonia,
forming Hg(NH_{4})_{2}Cl_{4}, or in general HgCl_{2}_n_MCl. If a
mixture of mercurous and potassium sulphates be dissolved in
dilute sulphuric acid, the solution easily yields large colourless
crystals of a double salt of the composition
K_{2}SO_{4},3HgSO_{4},2H_{2}O. Boullay obtained crystalline
compounds of mercuric chloride with hydrochloric acid, and
mercuric iodide with hydriodic acid; and Thomsen describes the
compound HgBr_{2},HBr,4H_{2}O as a well-crystallised salt, melting
at 13°, and having, in a molten state, a specific gravity 3·17 and
a high index of refraction. Moreover, the capacity of salts for
forming basic compounds has been considerably cleared up since the
investigation (by Würtz, Lorenz, and others) of glycol,
C_{2}H_{4}(OH)_{2} (and of polyatomic alcohols resembling it),
because the ethers C_{2}H_{4}X_{2}, corresponding with it, are
capable of forming compounds containing
C_{2}H_{4}X_{2}_n_C_{2}H_{4}O.
On the other hand, there is reason to think that the property of
forming basic salts is connected with the polymerisation of bases,
especially colloidal ones (_see_ the chapter on Silica, Lead
Salts, and Tungstic Acid).
[24] Mercuric iodide, HgI_{2}, is obtained first as a yellow, and then
as a red, precipitate on mixing solutions of mercuric salts and
potassium iodide, and is soluble in an excess of the latter (in
consequence of the formation of the double salt, HgKI_{3}); of
ammonium chloride (for a similar reason), &c. It crystallises at
the ordinary temperature in square prisms of a red colour. On
being heated, these change into yellow rhombic crystals,
isomorphous with mercuric chloride. This yellow form of mercuric
iodide is very unstable, and when cooled and triturated easily
again assumes the more stable red form. When fused, a yellow
liquid is obtained. _Mercuric cyanide_, Hg(CN)_{2}, forms one of
the most stable metallic cyanides. It is obtained by dissolving
mercuric oxide in prussic acid, and by boiling Prussian blue with
water and mercuric oxide, ferric oxide being then obtained in the
precipitate. Mercuric cyanide is a colourless crystalline
substance, soluble in water, and distinguished by its great
stability; sulphuric acid does not liberate prussic acid from it,
and even caustic potash does not remove the cyanogen (a complex
salt is probably produced), but the halogen acids disengage HCN.
Like the chloride, it combines with mercuric oxide, forming the
oxycyanide, Hg_{2}O(CN)_{2}, and it shows a very marked tendency
to form double compounds--for example, K_{2}Hg(CN)_{4}. The alkali
chlorides and iodides form similar compounds--for instance, the
salt HgKI(CN)_{2} crystallises very well, and is produced by
directly mixing solutions of potassium iodide and mercuric
cyanide.
Wells (1889) and Vare obtained and investigated many such double
salts, and showed the possibility of the formation, not only of
HgCl_{2}MCl and HgCl_{2}2MCl where M is a metal of the
alkalis--for example, Cs--but also of HgCl_{2}3MCl,2(HgCl_{2})MCl,
and in general _n_HgX_{2}_m_MX, where X stands for various
haloids.
Mercury has a remarkable power of forming very unstable compounds with
ammonia, in which the mercury replaces the hydrogen, and, if a mercuric
compound be taken, its atom occupies the place of two atoms of the
hydrogen in the ammonia. Thus Plantamour and Hirtzel showed that
precipitated mercuric oxide dried at a gentle heat, when continuously
heated (up to 100°-150°) in a stream of dry ammonia, leaves a brown
powder of _mercuric nitride_, N_{2}Hg_{3}, according to the equation 3HgO
+ 2NH_{3} = N_{2}Hg_{3} + 3H_{2}O.[24 bis] This substance, which is
attacked by water, acids, and alkalis (giving a white powder), is very
explosive when struck or rubbed, evolving nitrogen, proving that the bond
between the mercury and the nitrogen is very feeble.[25] By the action of
liquefied ammonia on yellow mercuric oxide Weitz also obtained an
explosive compound, dimercurammonium hydroxide, N_{2}Hg_{4}O, which
corresponds with an ammonium oxide, (NH_{4})_{2}O, in which the whole of
the hydrogen is replaced by mercury. A solution of ammonia reacts with
mercuric oxide, forming the hydroxide, NHg_{2}.OH, to which a whole
series of salts, NHg_{2}X, correspond; these are generally insoluble in
water and capable of decomposing with an explosion. But salts of the same
type, but with one atom of mercury, NH_{2}HgX, are more frequently and
more easily formed; they were principally studied by Kane, although known
much earlier. Thus, if ammonia be added to a solution of corrosive
sublimate (or, still better, in reverse order), a precipitate is obtained
known as white precipitate (_Mercurius præcipitatus albus_) or
_mercurammonium chloride_, NH_{2}HgCl, which may also be regarded not
only as sal-ammoniac with the substitution of H_{2} by mercury, but also
as HgX_{2}, where one X represents Cl and the other X represents the
ammonia radicle, HgCl_{2} + 2NH_{3} = NH_{2}.HgCl + NH_{4}Cl. When
heated, mercurammonium chloride decomposes, yielding mercurous chloride;
when heated with dry hydrochloric acid it forms ammonium chloride and
mercuric chloride. Other simple and double salts of mercurammonium,
NH_{2}HgX, are also known. Pici (1890) showed that all the compounds
HgH_{2}NX may be regarded as compounds of the above-named Hg_{2}NX with
NH_{4}X because their sum equals 2HgH_{2}X.[25 bis]
[24 bis] _See_ Chapter XIX., Note 6 bis: Hg_{3}P_{2}. In studying the
metallic nitrides it is necessary to keep the corresponding
phosphides in mind.
[25] Hg_{3}N_{2} is similar in composition to Mg_{3}N_{2}, &c. (Chapter
XIV.) The readiness with which mercuric nitride explodes shows
that the connection between the nitrogen and the mercury is very
unstable, and explains the circumstance that the so-called
_mercury fulminate_, or _fulminating mercury_, is an exceedingly
explosive substance. This substance is prepared in large
quantities for explosive mixtures; it enters into the composition
of percussion caps, which explode when struck, and ignite
gunpowder. Mercury fulminate was discovered by Howard, and from
that time has been prepared in the following way: one part of
mercury is dissolved in twelve parts of nitric acid, of sp. gr.
1·36, and when the whole of the mercury is dissolved, 5·5 parts of
90 p.c. alcohol are added, and the mass is shaken. A reaction then
commences, accompanied by a rise in temperature due to the
oxidation of the alcohol. As a matter of fact, many oxidation
products are produced during the action of the nitric acid on the
alcohol (glycolic acid, ethers, &c.) When the reaction becomes
tolerably vigorous, the same quantity of alcohol is added as at
the commencement, when a grey precipitate of the fulminate
separates. This salt has the composition C_{2}Hg(NO_{2})N. It
explodes when struck or heated. The mercury in it may be replaced
by other metals--for instance, copper or zinc, and also silver.
The silver salt, C_{2}Ag_{2}(NO_{2})N, is obtained in a precisely
analogous manner, and is even more explosive. Under the action of
alkali chlorides, only half the silver is replaced by the alkali
metal, but if the whole of the silver be replaced by an alkali
metal, then the salt decomposes. This is evidently because
combinations of this kind proceed in virtue of the formation of
substances in which mercury, and metals akin to it, are connected
in an unstable way with nitrogen. Potassium and other light metals
are incapable of entering into such connection and therefore, the
substitution of potassium for mercury entails the splitting-up of
the combination. Investigations of the fulminates were carried on
by Gay-Lussac and Liebig, but only the investigations of L. N.
Shishkoff fully cleared up the composition and relation of these
substances to the other carbon compounds. Shishkoff showed that
fulminates correspond with the nitro-acid, C_{2}H_{2}(NO_{2})N.
The explosiveness of the group depends partly on its containing at
the same time NO_{2} and carbon; we already know that all such
nitrogen compounds are explosive. If we imagine that the NO_{2} is
replaced by hydrogen, we shall have a substance of the composition
C_{2}H_{3}N. This is acetonitrile--that is, acetic acid + NH_{3} -
2H_{2}O, or ethenyl nitrile, as shown in Chapter VI. The formation
of an acetic compound by the action of nitric acid on alcohol is
easily understood, because acetic acid is produced by the
oxidation of alcohol, and the production of the elements of
ammonia, indispensable for the formation of a nitrile, is
accounted for by the fact that nitric acid under the action of
reducing substances in many cases forms ammonia. Moreover a
certain analogy has been found between fulminating acid and
hydroxylamine, but details upon this subject must be looked for in
works on organic chemistry. The explosiveness of fulminating
mercury, the rapidity of its decomposition (gunpowder, and even
guncotton, burn more slowly and explode less violently), and the
force of its explosion, are such that a small quantity (loosely
covered) will shatter massive objects.
The investigations of Abel on the communication of explosion from
one substance to another are remarkable. If guncotton be ignited
in an open space, it burns quietly; but if fulminating mercury be
exploded by the side of it, the decomposition of the guncotton is
effected instantaneously, and it then shatters the objects upon
which it lies, so rapid is the decomposition. Abel explains this
by supposing that the explosion of the fulminating salt brings the
molecules of guncotton into a uniform or as it were harmonious
state of vibration, which causes the rapid decomposition of the
whole mass. This rapid decomposition of explosive substances
defines the distinction between explosion and combustion. Besides
this, Berthelot showed that from that form of powerful molecular
concussion which takes place during the explosion of fulminating
mercury, the state of strain and stability of equilibrium of
substances which are endothermal, or capable of decomposing with
the disengagement of heat--for instance, cyanogen, nitro
compounds, nitrous oxide, &c.--is generally destroyed. Thorpe
showed that carbon bisulphide, CS_{2}, also an endothermal
substance, decomposes into sulphur and charcoal, when fulminating
mercury is exploded in contact with it.
[25 bis] The capacity for replacing hydrogen in chloride of ammonium by
metals also belongs to Zn and Cd. Kvasnik (1892), by the action of
ammonia upon alcoholic solutions of CdCl_{2} and ZnCl_{2},
obtained substances of the general formula M(NH_{3}Cl)_{2}, formed
as it were from two molecules of sal-ammoniac by the substitution
of two atoms of hydrogen by a diatomic metal. These substances
appear as white, finely crystalline powders. Under the action of
heat half the ammonia passes off, and a compound of the
composition MClNH_{3}Cl is formed. The compounds of cadmium and
zinc are distinguished from each other by the former being more
volatile than the latter.
We may further remark that in the series Mg, Zn, Cd, and Hg the
capacity to form double salts of diverse composition increases
with the atomic weight. Thus, according to Wells and Walden's
observations (1893), the ratio _n_ : _m_ for the type
_n_MCl_m_RCl_{2} (M = K, Li, Na ... R = Mg, Zn ...) is for Mg 1 :
1, for Zn 3 : 1, 2 : 1, and 1 : 1; for Cd, besides this, salts are
known with the ratio 4 : 1, and for Hg 3 : 1, 2 : 1, 1 : 1, 2 : 3,
1 : 2, and 1 : 5.
Mercury as a liquid metal is capable of dissolving other metals and
forming metallic solutions. These are generally called 'amalgams.' The
formation of these solutions is often accompanied by the development of a
large amount of heat--for instance, when potassium and sodium are
dissolved (Chapter XII., Note 39); but sometimes heat is absorbed, as,
for instance, when lead is dissolved. It is evident that phenomena of
this kind are exceedingly similar to the phenomena accompanying the
dissolution of salts and other substances in water, but here it is easy
to demonstrate that which is far more difficult to observe in the case of
salts: the solution of metals in mercury is accompanied by the formation
of definite chemical compounds of the mercury with the metals dissolved.
This is shown by the fact that when pressed (best of all in chamois
leather) such solutions leave solid, definite compounds of mercury with
metals. It is, however, very difficult to obtain them in a pure state, on
account of the difficulty of separating the last traces of mercury, which
is mechanically distributed between the crystals of the compounds.
Nevertheless, in many cases such compounds have undoubtedly been
obtained, and their existence is clearly shown by the evident crystalline
structure and characteristic appearance of many amalgams. Thus, for
instance, if about 2-1/2 p.c. of sodium be dissolved in mercury, a hard,
crystalline amalgam is obtained, very friable and little changeable in
air. It contains the compound NaHg_{5} (Chapter XII., Note 39). Water
decomposes it, with the evolution of hydrogen, but more slowly than other
sodium amalgams, and this action of water only shows that the bond
between the sodium and the mercury is weak, just like the connection
between mercury and many other elements--for instance, nitrogen. Mercury
directly and easily dissolves potassium, sodium, zinc, cadmium, tin,
gold, bismuth, lead, &c., and from such solutions or alloys it is in most
cases easy to extract definite compounds--thus, for instance, the
compounds of mercury and silver have the compositions HgAg and
Ag_{2}Hg_{3}. Objects made of copper when rubbed with mercury become
covered with a white coating of that metal, which slowly forms an
amalgam; silver acts in the same way, but more slowly, and platinum
combines with mercury with still greater difficulty. This metal only
readily forms an amalgam when in the form of a fine powder. If salts of
platinum in solution are poured on to an amalgam of sodium, the latter
element reduces the platinum, and the platinum separated is dissolved by
the mercury. Almost all metals readily form amalgams if their solutions
are decomposed by a galvanic current, where mercury forms the negative
pole. In this way an amalgam may even be made with iron, although iron in
a mass does not dissolve in mercury. Some amalgams are found in
nature--for instance, silver amalgams. Amalgams are used in considerable
quantities in the arts. Thus the solubility of silver in mercury is taken
advantage of for extracting that metal from the ore by means of
amalgamation, and for silvering by fire. The same is the case with gold.
Tin amalgam, which is incapable of crystallising and is obtained by
dissolving tin in mercury, composes the brilliant coating of ordinary
looking-glasses, which is made to adhere to the surface of the polished
glass by simply pressing by mechanical means sheets of tin foil bathed in
mercury on to the cleansed surface of the glass.[26] (_See_ 'The Nature
of Amalgams,' by W. L. Dudley; Toronto, 1889.)
[26] I consider it appropriate here to call attention to the want of an
element (ekacadmium) between cadmium and mercury in the periodic
system (Chapter XV.) But as in the ninth series there is not a
single known element, it may be that this series is entirely
composed of elements incapable of existing under present
conditions. However, until this is proved in one way or another,
it may be concluded that the properties of ekacadmium will be
between those of cadmium and mercury. It ought to have an atomic
weight of about 155, to form an oxide EcO, a slightly stable oxide
Ec_{2}O. Both ought to be feeble bases, easily forming double and
basic salts. The volume of the oxide will be nearly 17·5, because
the volume of cadmium oxide is about 16, and that of mercuric
oxide 19. Therefore the density of the oxide will approach 171 ÷
17·5 = 9·7. The metal ought to be easily fusible, oxidising when
heated, of a grey colour, with a specific volume, about 14
(cadmium = 13, mercury = 15), and, therefore, its specific gravity
(155 ÷ 14) will nearly = 11. Such a metal is unknown. But in 1879
Dahl, in Norway, discovered in the island of Oterö, not far from
Kragerö, in a vein of Iceland spar in a nickel mine, traces of a
new metal which he called norwegium, and which presented a certain
resemblance to ekacadmium. Perfect purity of the metal was not
attained, and therefore the properties ascribed to norwegium must
be regarded as approximate, and likely to undergo considerable
alteration on further study. A solution of the roasted mineral in
acid was twice precipitated by sulphuretted hydrogen, and again
ignited; the oxide obtained was easily reduced. When the metal was
dissolved in hydrochloric acid largely diluted with water, and the
solution boiled, the basic salt was precipitated, and thus freed
from the copper which remained in the solution. The reduced metal
had a density 9·44, and easily oxidised. If the composition NgO be
assigned to the oxide, then Ng = 145·9. It fused at 254°; the
hydroxide was soluble in alkalis and potassium carbonate. In any
case, if norwegium is not a mixture of other metals, it belongs to
the uneven series, because the heavy metals of the even series are
not easily reducible. Brauner thinks that norwegium oxide is
Ng_{2}O_{3}, the atom Ng = 219, and places it in Group VI., series
11, but then the feebly acid higher oxide, NgO_{3}, ought to be
formed.
Amongst the metals accompanying zinc which have been named, but
not authentically separated, must be included the _actinium_ of
Phipson (1881). He remarked that certain sorts of zinc give a
white precipitate of zinc sulphide which blackens on exposure to
light and then becomes white in the dark again. Its oxide, closely
resembling in many ways cadmium oxide, is insoluble in alkalis,
and it forms a white metallic sulphide, blackening on exposure to
light. As no further mention has been made of it since 1882, its
existence must be regarded as doubtful.
CHAPTER XVII
BORON, ALUMINIUM, AND THE ANALOGOUS METALS OF THE THIRD GROUP
If the elements of small atomic weight which we have hitherto discussed
be placed in order, it will be clearly seen that, judging by the formulæ
of their higher compounds, one element is wanting between beryllium and
carbon. For lithium gives LiX, beryllium forms BeX_{2}, and then comes
carbon giving CX_{4}. Evidently to complete the series we must look for
an element forming RX_{3}, and having an atomic weight greater than 9 and
less than 12. And _boron_ is such a one; its atomic weight is 11, and its
compounds are expressed by BX_{3}. Lithium and beryllium are metals;
carbon has no metallic properties; boron appears in a free state in
several forms which are intermediate between the metals and non-metals.
Lithium gives an energetic caustic oxide, beryllium forms a very feeble
base; hence one would expect to find that the oxide of boron, B_{2}O_{3},
has still more feeble basic properties and some acid properties, all the
more as CO_{2} and N_{2}O_{5}, which follow after B_{2}O_{3} in their
composition and in the periodic system, are acid oxides. And, indeed, the
only known _oxide of boron_ exhibits a feeble basic character, together
with the properties of a feeble acid oxide. This is even seen from the
fact that a solution of boron oxide reddens blue litmus and acts on
turmeric paper as an alkali, and these reactions may be used for
determining the presence of B_{2}O_{3} in solutions. By themselves the
alkali borates have an alkaline reaction, which clearly indicates the
feeble acid character of boric acid. If they are mixed in solution with
hydrochloric acid, boric acid is liberated, and if a piece of turmeric
paper be immersed in this solution and then dried, the excess of
hydrochloric acid volatilises, while the boric acid remains on the paper
and communicates a _brown coloration_ to it, just like alkalis.
Boron trioxide or boric anhydride enters into the composition of many
minerals, in the majority of cases in small quantities as an isomorphous
admixture, not replacing acids but bases, and most frequently alumina
(Al_{2}O_{3}), for as a rule the amount of alumina decreases as that of
the boric anhydride increases in them. This substitution is explained by
the similarity between the atomic composition of the oxides of aluminium
(alumina) and boron. The subdivision of oxides into basic and acid can in
no way be sharply defined, and here we meet with the most conclusive
proof of the fact, for the oxides of boron and aluminium belong to the
number of intermediate oxides, closely approaching the limit separating
the basic from the acid oxides. Their type (Chapter XV.) R_{2}O_{3} is
intermediate between those of the basic oxides R_{2}O and RO and those of
the acid oxides R_{2}O_{5} and RO_{3}. If we turn our attention to the
chlorides, we remark that lithium chloride is soluble in water, is not
volatile, and is not decomposed by water; the chlorides of beryllium and
magnesium are more volatile, and although not entirely, still are
decomposed by water; whilst the chlorides of boron and aluminium are
still more volatile and are decomposed by water. Thus the position of
boron and aluminium in the series of the other elements is clearly
defined by their atomic weights, and shows us that we must not expect any
new and distinct functions in these elements.
Boron was originally known in the form of sodium borate,
Na_{2}B_{4}O_{7},10H_{2}O, or _borax_, or _tincal_, which was exported
from Asia, where it is met with in solution in certain lakes of Thibet;
it has also been discovered in California and Nevada, U.S.A.[1] Boric
acid was afterwards found in sea-water and in certain mineral springs.[2]
Its presence may be discovered by means of the green coloration which it
communicates to the flame of alcohol, which is capable of dissolving free
boric acid.[3] Many of the boron compounds employed in the arts are
obtained from the impure boric acid which is extracted in Tuscany from
the so-called _suffioni_. In these localities, which present the remains
of volcanic action, steam mixed with nitrogen, hydrogen sulphide, small
quantities of boric acid, ammonia, and other substances, issue from the
earth.[3 bis] The boric acid partially volatilises with the steam, for if
a solution of boric acid be boiled, the distillate will always contain a
certain amount of this substance.[4]
[1] Borax is either directly obtained from lakes (the American lakes
give about 2,000 tons and the lakes of Thibet about 1,000 tons per
annum), or by heating native calcium borate (_see_ Note 2) with
sodium carbonate (about 4,000 tons per annum), or it is obtained
(up to 2,000 tons) from the Tuscan impure boric acid and sodium
carbonate (carbonic anhydride is evolved). Borax gives
supersaturated solutions with comparative ease (Gernez), from which
it crystallises, both at the ordinary and higher temperatures, in
octahedra, containing Na_{2}B_{4}O_{7},5H_{2}O. Its sp. gr. is
1·81. But if the crystallisation proceeds in open vessels, then at
temperatures below 56°, the ordinary prismatic crystallo-hydrate
B_{4}Na_{2}O_{7},10H_{2}O is obtained. Its sp. gr. is 1·71, it
effloresces in dry air at the ordinary temperature, and at 0° 100
parts of water dissolve about 8 parts of this crystallo-hydrate, at
50° 27 parts, and at 100° 201 parts. Borax fuses when heated, loses
its water and gives an anhydrous salt which at a red heat fuses
into a mobile liquid and solidifies into a transparent amorphous
_glass_ (sp. gr. 2·37), which before hardening acquires the pasty
condition peculiar to common molten glass. Molten borax dissolves
many oxides and on solidifying acquires characteristic tints with
the different oxides; thus oxide of cobalt gives a dark blue glass,
nickel a yellow, chromium a green, manganese an amethyst, uranium a
bright yellow, &c. Owing to its fusibility and property of
dissolving oxides, borax is employed in soldering and brazing
metals. Borax frequently enters into the composition of strass and
fusible glasses.
[2] We may mention the following among the minerals which contain
boron: calcium borate, (CaO)_{3}(B_{2}O_{3})(H_{2}O)_{6}, found and
extracted in Asia Minor, near Brusa; _boracite_ (stassfurtite),
(MgO)_{6}(B_{2}O_{3})_{8},MgCl_{2}, at Stassfurt, in the regular
system, large crystals and amorphous masses (specific gravity
2·95), used in the arts; _ereméeffite_ (Damour), AlBO_{3} or
Al_{2}O_{3}B_{2}O_{3}, found in the Adulchalonsk mountains in
colourless, transparent prisms (specific gravity 3·28) resembling
apatite; _datholite_, (CaO)_{2}(SiO_{2})_{2}B_{2}O_{3},H_{2}O; and
ulksite, or the boron-sodium carbonate from which a large quantity
of borax is now extracted in America (Note 1). As much as 10 p.c.
of boric anhydride sometimes enters into the composition of
tourmalin and axinite.
[3] This green coloration is best seen by taking an alcoholic solution
of volatile ethyl borate, which is easily obtained by the action of
boron chloride on alcohol.
[3 bis] P. Chigeffsky showed in 1884 (at Geneva) that in the
evaporation of saline solutions many salts are carried off by the
vapour--for instance, if a solution of potash containing about
17-20 grams of K_{2}CO_{3} per litre be boiled, about 5 milligrams
of salt are carried off for every litre of water evaporated. With
Li_{2}CO_{3} the amount of salt carried over is infinitesimal, and
with Na_{2}CO_{3} it is half that given by K_{2}CO_{3}. The
volatilisation of B_{2}O_{3} under these circumstances is
incomparably greater--for instance, when a solution containing 14
grams of B_{2}O_{3} per litre is boiled, every litre of water
evaporated carries over about 350 milligrams of B_{2}O_{3}. When
Chigeffsky passed steam through a tube containing B_{2}O_{3} at
400°, it carried over so much of this substance that the flame of a
Bunsen's burner into which the steam was led gave a distinct green
coloration; but when, instead of steam, air was passed through the
tube there was no coloration whatever. By placing a tube with a
cold surface in steam containing B_{2}O_{3}, Chigeffsky obtained a
crystalline deposit of the hydrate B(OH)_{3} on the surface of the
tube. Besides this, he found that the amount of B_{2}O_{3} carried
over by steam increases with the temperature, and that crystals of
B(OH)_{3} placed in an atmosphere of steam (although perfectly
still) volatilise, which shows that this is not a matter of
mechanical transfer, but is based on the capacity of B_{2}O_{3} and
B(OH)_{3} to pass into a state of vapour in an atmosphere of steam.
[4] How it is that these vapours containing boric acid are formed in
the interior of the earth is at present unknown. Dumas supposes
that it depends on the presence of _boron sulphide_, B_{2}S_{3}
(others think boron nitride), at a certain depth in the earth. This
substance may be artificially prepared by heating a mixture of
boric acid and charcoal in a stream of carbon bisulphide vapour,
and by the direct combination of boron and the vapour of sulphur at
a white heat. The almost non-crystalline compound B_{2}S_{3}, sp.
gr. 1·55, thus obtained is somewhat volatile, has an unpleasant
smell, and is very easily decomposed by water, forming boric acid
and hydrogen sulphide, B_{2}S_{3} + 3H_{2}O = B_{2}O_{3} + 3H_{2}S.
It is supposed that a bed of boron sulphide lying at a certain
depth below the surface of the earth comes into contact with sea
water which has percolated through the upper strata, becomes very
hot, and gives steam, hydrogen sulphide, and boric acid. This also
explains the presence of ammonia in the vapours, because the sea
water certainly passes through crevices containing a certain amount
of animal matter, which is decomposed by the action of heat and
evolves ammonia. There are several other hypotheses for explaining
the presence of the vapours of boric acid, but owing to the want of
other known localities the comparison of these hypotheses is at
present hardly possible. The amount of boric anhydride in the
vapours which escape from the Tuscan fumerolles and suffioni is
very inconsiderable, less than one-tenth per cent., and therefore
the direct extraction of the acid would be very uneconomical, hence
the heat contained in the discharged vapours is made use of for
evaporating the water. This is done in the following manner.
Reservoirs are constructed over the crevices evolving the vapours,
and the water of some neighbouring spring is passed into them. The
vapours are caused to pass through these reservoirs, and in so
doing they give up all their boric acid to the water and heat it,
so that after about twenty-four hours it even boils; still this
water only forms a very weak solution of boric acid. This solution
is then passed into lower basins and again saturated by the vapours
discharged from the earth, by which means a certain amount of the
water is evaporated and a fresh quantity of boric acid absorbed;
the same process is repeated in another reservoir, and so on until
the water has collected a somewhat considerable amount of boric
acid. The solution is drawn from the last reservoir A into settling
vessels B D, and then into a series of vessels _a_, _b_, _c_. In
these vessels, which are made of lead, the solution is also
evaporated by the vapours escaping from the earth, and attains a
density of 10° to 11° Baumé. It is allowed to settle in the vessel
C, in which it cools and crystallises, yielding (not quite pure)
crystalline boric acid. At temperatures above 100°, for instance,
with superheated steam, boric acid volatilises with steam very
easily.
[Illustration: FIG. 81.--Extraction of boric acid in Tuscany.]
If boric acid be introduced into an excess of a strong hot solution of
sodium hydroxide, then, on slowly cooling, the salt NaBO_{2},4H_{2}O
crystallises out. This salt contains an equivalent of Na_{2}O to one
equivalent B_{2}O_{3}. It might be termed a neutral salt did it not
possess strongly alkaline reactions and easily split up into the alkali
and the more stable borax or biborate of sodium mentioned above, which
contains 2B_{2}O_{3} to Na_{2}O.[5] This salt is prepared by the action
of boric acid on a solution of sodium carbonate. Borax may be perfectly
purified by crystallisation. If a saturated and hot solution of borax be
mixed with strong hydrochloric acid, common salt and a normal crystalline
hydrate of _boric acid_ are formed. The composition of this hydrate is
B(HO)_{3}, according to the form BX_{3}--that is, of the composition
B_{2}O_{3},3H_{2}O. This is the easiest method of obtaining pure boric
acid. The water is easily expelled from this hydrate; it loses half at
100° and the remainder on further heating, and the remaining B_{2}O_{3}
or boric anhydride fuses at 580° (according to Carnelley), forming at
first a ductile (easily drawn out into threads), tenacious mass and then
a colourless liquid solidifying to a transparent glass, which absorbs
moisture from the atmosphere and then becomes cloudy.[6] Only the
alkaline salts of boric acid are soluble in water, but all borates are
soluble in acids, owing to their easy decomposability and the solubility
of boric acid itself. Although boric anhydride, B_{2}O_{3}, absorbs
3H_{2}O from damp air, still in the presence of water it always[7]
combines with a less quantity of bases (borax only contains 1/6).
However, fused boric anhydride forms a crystalline compound with
magnesium of the same type as the hydrate (MgO)_{3}B_{2}O_{3} (Ebelmann),
and even with sodium it forms (Na_{2}O)_{3}B_{2}O_{3} or Na_{3}BO_{3}
(Benedict). As a rule, the salts of boric acid contain less base,
although they are all able to form saline compounds with bases when
fused. Generally, vitreous fluxes are formed by this means,[8] which when
fused recall ordinary aqueous solutions in many respects. Some of them
crystallise on solidifying, and then they have, like salts, a definite
composition. The property of boric anhydride of forming higher grades of
combination with basic oxides when fused explains the power of fused
borax to dissolve metallic oxides, and the experiments of Ebelmann on the
preparation of artificial crystals of the precious stones by means of
boric anhydride. Boric anhydride is, although with difficulty, volatile
at a high temperature, and therefore if it dissolves an oxide, it may be
partially driven off from such a solution by prolonged and powerful
ignition; in which case the oxides previously in solution separate out in
a crystalline form, and frequently in the same forms as those in which
they occur in nature--for example, crystals of alumina, which by itself
fuses with difficulty, have been obtained in this manner. It dissolves in
molten boric anhydride, and separates out in natural rhombohedric
crystals. In this way Ebelmann also obtained _spinel_--that is, a
compound of magnesium and aluminium oxides which occurs in nature.[9]
[5] Metals, like Na, K, Li, give salts of the type of borax, MBO_{2} or
MH_{2}BO_{3}. A solution of borax, Na_{2}B_{4}O_{7}, has an
alkaline reaction, decomposes ammonia salts with the liberation of
ammonia (Bolley), absorbs carbonic anhydride like an alkali,
dissolves iodine like an alkali (Georgiewics), and seems to be
decomposed by water. Thus Rose showed that strong solutions of
borax give a precipitate of silver borate with silver nitrate,
whilst dilute solutions precipitate silver oxide, like an alkali.
Georgiewics even supposes (1888) boric anhydride to be entirely
void of acid properties; for all acids, on acting on a mixture of
solutions of potassium iodide and iodate, evolve iodine, but boric
acid does not do this. With dilute solutions of sodium hydroxide
Berthelot obtained a development of heat equal to 11-1/2 thousand
calories per equivalent of alkali (40 grams sodium hydroxide) when
the ratio Na_{2}O : 2B_{2}O_{3} (as in borax) was taken, and only 4
thousand calories when the ratio was Na_{2}O : B_{2}O_{3}, whence
he concludes that water powerfully decomposes those sodium borates
in which there is more alkali than in borax. Laurent (1849)
obtained a sodium compound, Na_{2}O,4B_{2}O_{3},10H_{2}O,
containing twice as much boric anhydride as borax, by boiling a
mixture of borax with an equivalent quantity of sal-ammoniac until
the evolution of ammonia entirely ceased.
Hence it is evident that feeble acids are as prone to, and as
easily, form acid salts (that is, salts containing much acid oxide)
as feeble bases are to give basic salts. These relations become
still clearer on an acquaintance with such feeble acids as silicic,
molybdic, &c. This variety of the proportions in which bases are
able to form salts recalls exactly the variety of the proportions
in which water combines with crystallo-hydrates. But the want of
sufficient data in the study of these relations does not yet permit
of their being generalised under any common laws.
With respect to the feeble acid energy of boric anhydride I think
it useful to add the following remarks. Carbonic anhydride is
absorbed by a solution of borax, and displaces boric anhydride; but
it is also displaced by it, not only on fusion, but also on
solution, as the preparation of borax itself shows. Sulphuric
anhydride is absorbed by boric acid, forming a compound
B(HSO_{4})_{3}, where HSO_{4} is the radicle of sulphuric acid
(D'Ally). With phosphoric acid, boric acid forms a stable compound,
BPO_{4}, or B_{2}O_{3}P_{2}O_{5}, undecomposable by water, as
Gustavson and others have shown. With respect to tartaric acid,
boric anhydride is able to play the same part as antimonious oxide.
Mannitol, glycerol, and similar polyhydric alcohols also seem able
to form particularly characteristic compounds with boric anhydride.
All these aspects of the subject require still further explanation
by a method of fresh and detailed research.
[6] Ditte determined the sp. gr.:--
0° 12° 80°
B_{2}O_{3} 1·8766 1·8470 1·6988
B(OH)_{3} 1·5463 1·5172 1·3828
Solubility 1·95 2·92 16·82
The last line gives the solubility, in grams, of boric acid,
B(OH)_{3}, per 100 c.c. of water, also according to the
determinations of Ditte.
[7] It is evident that, in the presence of basic oxides, water competes
with them, which fact in all probability determines both the amount
of water in the salts of boric acid as well as their decomposition
by an excess of water. In confirmation of the above-mentioned
competing action between water and bases, I think it useful to
point out that the crystallo-hydrate of borax containing 5H_{2}O
may be represented as B(HO)_{3}, or rather as B_{2}(OH)_{6}, with
the substitution of one atom of hydrogen by sodium, since
Na_{2}B_{4}O_{7},5H_{2}O = 2B_{2}(OH)_{5}(ONa). The composition of
the acid boric salts is very varied, as is seen from the fact that
Reychler (1893) obtained (Cs_{2}O)3B_{2}O_{3}, (Rb_{2}O)2B_{2}O_{3}
(corresponding to borax) and (Li_{2}O)B_{2}O_{3}, and that Le
Chatelier and Ditte obtained, for CaO, MgO, &c., (RO)B_{2}O_{3},
(RO)_{2}3B_{2}O_{3}, (RO)2B_{2}O_{3}, (RO)_{2}B_{2}O_{3}, and even
(RO)_{3}B_{2}O_{3}.
[8] A glass can only be formed by those slightly volatile oxides which
correspond with feeble acids, like silica, phosphoric and boric
anhydrides, &c., which themselves give glassy masses, like quartz,
glacial phosphoric acid, and boric anhydride. They are able, like
aqueous solutions and like metallic alloys, to solidify either in
an amorphous form or to yield (or even be wholly converted into)
definite crystalline compounds. This view illustrates the position
of solutions amongst the other chemical compounds, and allows all
alloys to be regarded from the aspect of the common laws of
chemical reactions. I have therefore frequently recurred to it in
this work, and have since the year 1850 introduced it into various
provinces of chemistry.
[9] If boric acid in its aqueous solutions proves to be exceedingly
feeble, unenergetic, and easily displaced from its salts by other
acids, yet in an anhydrous state, as anhydride, it exhibits the
properties of an energetic acid oxide, and it _displaces_ the
anhydrides of other acids. This of course does not mean that the
acid then acquires new chemical properties, but only depends on the
fact that the anhydrides of the majority of acids are much more
volatile than boric anhydride, and therefore the salts of many
acids--even of sulphuric acid--are decomposed when fused with boric
anhydride.
By itself boric acid is used in the arts in small quantity, chiefly
for the preservation of meat and fish (which must be afterwards
well washed in water) and of milk, and for soaking the wicks of
stearin candles; the latter application is based on the fact that
the wicks, which are made of cotton twist, contain an ash which is
infusible by itself but which fuses when mixed with boric acid.
Free _boron_ was obtained (1809) by Davy, Gay-Lussac, and Thénard when
they obtained the metals of the alkalis, for boric anhydride when fused
with sodium gives up its oxygen to the sodium, and free boron is
liberated as an _amorphous_ powder like charcoal.[10] It is of a brown
colour, specific gravity 2·45 (Moissan), and when dry does not alter in
the air at the ordinary temperature; but it burns when ignited to 700°,
and in so doing combines not only with the oxygen of the air, but also
with the nitrogen. However, the combustion is never complete, because the
boric anhydride formed on the surface covers the remaining mass of the
boron, and so preserves it from the action of the oxygen. Acids, even
sulphuric (forming SO_{2}) and phosphoric (forming phosphorus), easily
oxidise amorphous boron, especially when heated, converting it into boric
acid. Alkalis have the same action on it, only in this case hydrogen is
evolved. Boron decomposes steam at a red heat, also with evolution of
hydrogen.
[10] _Amorphous boron_ is prepared by mixing 100 parts of powdered
boric anhydride with 50 parts of sodium in small lumps; this
mixture is thrown into a powerfully heated cast-iron crucible,
covered with a layer of ignited salt, and the crucible covered.
Reaction proceeds rapidly; the mass is stirred with an iron rod,
and poured directly into water containing hydrochloric acid. The
action is naturally accompanied by the formation of sodium borate,
which is dissolved, together with the salt, by the water, whilst
the boron settles at the bottom of the vessel as an insoluble
powder. It is washed in water, and dried at the ordinary
temperature. Magnesium, and even charcoal and phosphorus, are also
able to reduce boron from its oxide. Boron, in the form of an
amorphous powder, very easily passes through filter-paper, remains
suspended in water, and colours it brown, so that it appears to be
soluble in water. Sulphur precipitated from solutions shows the
same (colloidal) property. When borax is fused with magnesium
powder, it gives a brown powder of a compound of boron and
magnesium, Mg_{2}B (Winkler, 1890), but when a mixture of 1 part
of magnesium and 3 parts of B_{2}O_{3} is heated to redness
(Moissan, 1892), it forms amorphous boron in the form of a
chestnut-coloured powder, which, after being washed with water,
hydrochloric and hydrofluoric acids, is fused again with
B_{2}O_{3} in an atmosphere of hydrogen in order to prevent the
access of the nitrogen of the air, which is easily absorbed by
incandescent amorphous boron.
Sabatier (1891) considers that a certain amount of gaseous hydride
of boron is evolved in the action of hydrochloric acid upon the
alloys of magnesium and boron, because the gas disengaged burns
with a green flame. Still, the existence of hydride of boron
cannot be regarded as certain.
Under the action of the heat of the electric furnace boron forms
with carbon a _carbide_, BC, as Mühlhäuser and Moissan showed in
1893.
Amorphous boron, like charcoal, dissolves in certain molten metals. The
property of fused _aluminium of dissolving boron_ in considerable
quantity is very striking; on cooling such a solution, the boron
partially combined with the aluminium separates out in a crystalline
form, and its properties are then exceedingly remarkable. The crystalline
boron may be obtained by heating (to 1,300°) the pulverulent boron with
aluminium in a well-closed crucible, the access of air being prevented as
far as possible. After cooling, crystals are observed on the surface of
the aluminium, and may easily be separated by dissolving the latter in
hydrochloric acid, which does not act on the crystals. The specific
gravity of the crystals is 2·68; they are partially transparent, but are
for the most part coloured dark brown; they contain about 4 p.c. of
carbon and up to 7 p.c. of aluminium, so that they cannot be considered
as pure boron. Nevertheless, the properties of this _crystalline_
substance, which was obtained by Wöhler and Deville, are very remarkable.
It most closely resembles _the diamond in its properties_--in fact, these
crystals have the lustre and high refracting power proper to the diamond
only, whilst their hardness competes with that of the diamond. Their
powder polishes even the diamond, and like the diamond scratches the
sapphire and corundum. Crystalline boron is much more stable with respect
to chemical reagents than the amorphous variety, and as it resembles the
diamond, so amorphous boron, on the other hand, distinctly recalls
certain of the properties of charcoal; thus a certain resemblance exists
between boron and carbon in a free state, which is further justified by
the proximity of their positions in the periodic system.
Among the other compounds of boron, those with nitrogen and the halogens
are the most remarkable. As already mentioned above, amorphous boron
combines directly with _nitrogen_ at a red heat. If it be heated in a
glass tube in a stream of nitric oxide, perfect combustion takes place,
5B + 3NO = B_{2}O_{3} + 3BN. If the residue be treated with nitric acid,
the boric anhydride dissolves, whilst the _boron nitride_ remains[11] as
an extremely light white powder, which is sometimes partially crystalline
and greasy to the touch, like talc. It is infusible and unchanged, even
at the melting-point of nickel. In general, it is remarkable for its
great stability with respect to chemical reagents. Nitric and
hydrochloric acids, as well as alkaline solutions, and hydrogen and
chlorine at a red heat, have no action on it. When fused with potash, it
evolves ammonia, and when ignited in steam it also yields ammonia: 2BN +
3H_{2}O = B_{2}O_{3} + 2NH_{3}.[12]
[11] At first boron nitride was obtained by heating boric acid with
potassium cyanide or other cyanogen compounds. It may be more
simply prepared by heating anhydrous borax with potassium
ferrocyanide, or by heating borax with ammonium chloride. For this
purpose one part of borax is intimately mixed with two parts of
dry ammonium chloride, and the mixture heated in a platinum
crucible. A porous mass is formed, which after crushing and
treating with water and hydrochloric acid, leaves boron nitride.
_Boron fluoride_, BF, is known, corresponding to BN; this body was
obtained by Besson and Moissan (1891). The action of phosphorus
upon iodide of boron, BI_{3}, forms PBI_{2}, and when heated to
500° in hydrogen it forms BP, which gives PH_{3} with fused KHO.
[12] When fused with potassium carbonate it forms potassium cyanate,
BN + K_{2}CO_{3} = KBO_{2} + KCNO. All this shows that boron
nitride is a nitrile of boric acid, BO(OH) + NH_{3} - 2H_{2}O =
BN. The same is expressed by saying that boron nitride is a
compound of the type of the boron compounds BX_{3}, with the
substitution of X_{3} by nitrogen, as the trivalent radicle of
ammonia, NH_{3}.
No less remarkable is the compound of boron with fluorine--_boron
fluoride_, BF_{3}. It is produced in many instances when compounds of
boron and of fluorine are brought together.[13] The most convenient
method of preparing it is by heating a mixture of calcium fluoride with
boric anhydride and sulphuric acid, 3CaF_{2} + B_{2}O_{3} + 3H_{2}SO_{4}
= 3CaSO_{4} + 3H_{2}O + 2BF_{3}.[14] It is a colourless liquefiable _gas_
(the liquid boils at -100°), which on coming into contact with damp air
forms white fumes, owing to its combining with water. One volume of water
dissolves as much as 1,050 volumes of this gas (Bazaroff), forming a
liquid which disengages boron fluoride when heated, and distils over
unaltered. Boron fluoride chars organic matter, owing to its taking up
the water from it, and in this respect it acts like sulphuric acid. The
behaviour of boron fluoride with water must be understood as a reversible
reaction, since with water it yields hydrofluoric and boric acids, whilst
they, acting on one another, re-form boron fluoride and water. A state of
equilibrium is set up between these four substances (and between two
reversible reactions) which is distinctly dependent on the mass of the
water.[14 bis] When boron fluoride is in great excess, the equilibrated
system, which is capable of distilling over (sp. gr. of the liquid 1·77),
has a composition BF_{3},2H_{2}O (or B_{2}O_{3},H_{2}O,6HF). It has also
its corresponding salts.[15] It is a caustic liquid, having the
properties of a powerful acid; but it does not act on glass, which shows
that there is no free hydrofluoric acid present. Under the action of
water this system changes, with the formation of boric acid and
hydroborofluoric acid (HBF_{4}) according to the equation
4BF_{3}H_{4}O_{2} = 3HBF_{4} + BH_{3}O_{3} + 5H_{2}O.[16] This
hydroborofluoric acid has its corresponding salts--for instance, KBF_{4}.
On evaporating the aqueous solution this free acid decomposes, with the
evolution of hydrofluoric acid, and a stable system is again obtained:
2HBF_{4} + 5H_{2}O = B_{2}F_{6}H_{10}O_{5} + 2HF. The resultant solution
(containing 2BF_{3},5H_{2}O, sp. gr. 1·58), which is identical with that
formed by the evaporation of a solution of boric acid with hydrofluoric
acid, again only contains a compound of boron fluoride with water.
Probably there are various other possible and more or less stable states
of equilibrium and definite compounds of boron fluoride, hydrofluoric
acid, and water.
[13] Boron fluoride is frequently evolved on heating certain compounds
occurring in nature containing both boron and fluorine. If calcium
fluoride is heated with boric anhydride, calcium borate and boron
fluoride are formed, and the latter, as a gas, is volatilised:
2B_{2}O_{3} + 3CaF_{2} = 2BF_{3} + Ca_{3}B_{2}O_{6}. The calcium
borate, however, retains a certain amount of calcium fluoride.
[14] In order to avoid the formation of silicon fluoride the
decomposition should not be carried on in glass vessels, which
contain silica, but in lead or platinum vessels. Boron fluoride by
itself does not corrode glass, but the hydrofluoric acid liberated
in the reaction may bring a part of the silica into reaction.
Boron fluoride should be collected over mercury, as water acts on
it, as we shall see afterwards.
[14 bis] It appears to me that from this point of view it is possible
to understand the apparently contradictory results of different
investigators, especially those of Gay-Lussac (and Thénard), Davy,
Berzelius, and Bazaroff. In the form in which the reaction of
BF_{3} on water is given here, it is evident that the act of
solution in water is accompanied by complex but direct chemical
transformations, and I think that this example should prove the
justness of those observations upon the nature of solutions which
are given in Chapter I.
[15] They are called fluoborates. They may be prepared directly from
fluorides and borates. Such compounds of halogens with oxygen
salts are known in nature (for instance, apatite and boracite),
and may be artificially prepared. The composition of the
fluoborates--for example, K_{4}BF_{3}O_{2}--may be expressed as
that of a double salt, BO(OK),3KF. If an excess of water
decomposes them (Bazaroff), this does not prove that they do not
exist as such, for many double salts are decomposed by water.
[16] Fluoboric acid contains boron fluoride and water, hydrofluoboric
acid, boron fluoride, and hydrofluoric acid. It is evident that on
the one side the competition between water and hydrofluoric acid,
and, on the other hand, their power to combine, are among the
forces which act here. From the fact that hydroborofluoric acid,
HBF_{4}, can only exist in an aqueous solution, it must be assumed
that it forms a somewhat stable system only in the presence of
3H_{2}O.
Nothing of this kind occurs with boron chloride, because hydrochloric
acid does not act on boric acid. However, amorphous boron at 400° burns
in chlorine, and at 410° forms _boron chloride_, BCl_{3}. The boron burns
in the chlorine, forming a gas which, in a freezing mixture, condenses
into a liquid boiling at 17°, and gives up its excess of chlorine, if
there be any, to mercury. The specific gravity of this liquid is 1·42 at
6°. Boron chloride may also be directly obtained from boric anhydride by
the simultaneous action of charcoal and chlorine at a high temperature:
B_{2}O_{3} + 3C + 3Cl_{2} = 2BCl_{3} + 3CO. It is also obtained by the
action of phosphoric chloride on boric anhydride in a closed tube at 200°
It is completely decomposed by water, like the chloranhydride of an acid,
boric acid being formed; hence it fumes in the air: 2BCl_{3} + 6H_{2}O =
2BH_{3}O_{3} + 6HCl. Boron forms with bromine a similar compound,
BBr_{3}, specific gravity at 6° = 2·64, boiling at 90°. The vapour
densities of the fluoride, chloride, and bromide of boron show that they
contain three atoms of the halogen in the molecule--that is, that boron
is a trivalent element forming BX_{3}.[16 bis]
[16 bis] Iodide of boron, BI_{3}, was obtained by Moissan (1891), by
heating a mixture of the vapours of HI and BCl_{3} in a tube, or
by the action of iodine vapour (at 750°) or HI upon amorphous
boron. BI_{3} is a solid substance which dissolves in benzol and
CS_{2}, reacts with water, melts at 43°, boils at 210°, has a
density 3·3 at 50°, and partially decomposes in the light. Besson
(1891) obtained BIBr_{2} (boiling at 125°), and BI_{2}Br (boiling
at 180°) by heating (300-400°) a mixture of the vapours of HI and
BBr_{3}, and showed that NH_{3} combines with BBr_{3} and BI_{3}
in various proportions.
As in the first group lithium is followed by sodium, giving a more basic
oxide, so in the second group beryllium is followed by magnesium, and so
also in the third group there is, besides the lightest element, boron,
whose basic character is scarcely defined, _aluminium_, Al = 27, whose
oxide, alumina, has somewhat distinct basic properties, which, although
not so powerful as in magnesium oxide, are more distinct than in boric
anhydride. Among the elements of the third group, aluminium is the most
widely distributed in nature; it will be sufficient to mention that it
enters into the composition of clay to demonstrate the universal
distribution of aluminium in the earth's crust.
Alumina is so named from its being the metal of alums (_alumen_).
_Clay_, which is so widely distributed and familiar to everybody, is the
insoluble residue obtained after the action of water containing carbonic
acid on many rocks, and especially on the felspars contained in some of
them. Felspar is a compound containing potash or soda, alumina, and
silica. The primary rocks, like granite, contain many similar compounds
(_see_ Chapter XVIII.: Felspars). Felspar is acted on by water containing
carbonic acid, all the alkalis (potash and soda), and a portion of the
silica passing into the water as substances which are soluble and carried
away by it, whilst the alumina and silica left from the felspar remain on
the spot where the solution has taken place. This is the original method
of the formation of clay in its primary deposits among rocks along whose
crevices the atmospheric water has permeated. Such primary deposits often
contain a white pure clay, termed _kaolin_ or _porcelain clay_. But such
clay is a rarity, because the conditions for its formation are rarely met
with. The water, whilst acting chemically on rocks, at the same time
destroys them _mechanically_, and carries off the finely divided residues
of disintegration with it. Clay is most easily subjected to this
mechanical action of water, because it is composed of grains of
exceedingly small size and void of any visible crystalline structure,
which easily remain suspended in water. The cloudy water of running
mountain streams generally contains particles of clay in suspension,
owing to the above-described chemical and mechanical action of the water
on the minerals contained in the mountain rocks. Together with these
minute particles of clay the water carries away the coarser components on
which it is not able to act--for example, splinters of rock, grains of
mica, quartz, &c. They were originally held together by those minerals
which form clay. When the water acts on these binding minerals, a sandy
mass is formed which water bears away. The cloudy water in which the
particles of clay and sand are held in suspension carries them to, and
deposits them at, the estuaries of rivers, lakes, seas, and oceans. The
coarser particles are first deposited and form sand and similar
disintegrated rocky matter, whilst the clay, owing to its finely divided
state, is carried on further, and is only deposited in the still parts of
the rivers, lakes, &c. Such disintegrations of rocks and separations of
clay from sand have been gradually going on during the millions of years
of the earth's existence, and are now proceeding, and have been the cause
of the formation of the immense deposits of sandstone and clay now
forming a part of the earth's strata. Such beds of clay may have been
transferred by currents and streams from one locality to another, so that
we must distinguish between primary and secondary deposits of clay. In
places these beds of clay have, owing to long exposure under water, and
perhaps partially owing to the action of heat, undergone compression, and
have formed the rocky masses known as clay slates and schists, which
sometimes form entire mountains. Roofing slates belong to this class of
rocks.
From what has been said above it will be evident that these deposits can
never consist of a chemically pure and homogeneous substance, but will
contain all kinds of extraneous insoluble finely divided matter, and
especially sand--that is, fragments of rock, chiefly quartz (SiO_{2}). It
is, however, possible to considerably purify clay from these impurities,
owing to the fact that they are the result of mechanical disintegration,
whilst the clay has been formed as a residue of the chemical alteration
of rocky matter, and therefore its particles are incomparably more minute
than the particles of sand and other rock fragments mixed with it. This
difference in the size of the grains causes the clay to remain longer in
suspension when shaken up in water than the coarser grains of sand. If
clay be shaken up in water, and especially if it be previously boiled in
it, and if after the first portion has settled the cloudy water be
decanted, it will give a deposit of a very much purer clay than the
original. This method is employed for purifying kaolin designed for the
manufacture of the best kinds of china, earthenware, &c. A similar method
is also employed in the investigation of earths for determining the
_composition of soils_ chiefly composed of a mixture of sand, clay,
limestone, and mould. The limestone is soluble in dilute acids, but
neither the clay nor sand passes into solution by this means, and
therefore the limestone is easily separated in the investigation of
soils. The clay is separated from the sand by a mechanical method similar
to that described above, and termed _levigation_.[17]
[17] The process of _levigation_ is based on the difference in the
diameters of the particles of clay and sand. In density these
particles differ but little from each other, and therefore a
stream of water of a certain velocity can only carry away the
particles of a certain diameter, whilst the particles of a larger
diameter cannot be borne away by it. This is due to the resistance
to falling offered by the water. This resistance to substances
moving in it increases with the velocity, and therefore a
substance falling into water will only move with an increasing
velocity until its weight equals the resistance offered by the
water, and then the velocity will be uniform. And as the weight of
the minute particles of clay is small, the maximum velocity
attained by them in falling is also small. A detailed account of
the theory of falling bodies in liquid, and of the experiments
bearing on this subject, may be found in my work, _Concerning the
Resistance of Liquids and Aeronautics_, 1880. The minute particles
of clay remain suspended longer in water, and take longer to fall
to the bottom. Heavy particles, although of small dimensions, fall
more quickly, and are borne away by water with greater difficulty
than the lighter. In this way gold and other heavy ores are washed
free from sand and clay, and the coarser portions and heavier
particles are left behind. A current of water of a certain
velocity cannot carry away with it particles of more than a
definite diameter and density, but by increasing the velocity of
the current a point may be arrived at when it will bear away
larger particles. A description of apparatus for the observation
of phenomena of this kind is given by Schöne in his memoir in the
Transactions of the Moscow Society of Natural Sciences for 1867.
In order to be able accurately to vary the velocity of the current
of water, a cylinder is employed in which the earth to be
experimented on is placed, and water is introduced through the
conical bottom of the cylinder. The rate at which the water rises
in the cylinder will vary according to the quantity of water
flowing per unit of time into the vessel, and consequently
particles of various sizes will be carried away by the water
flowing over the upper edges of the vessel. Schöne showed by
direct experiment that a current of water having a velocity of 0·1
mm. per second will carry away particles having a diameter of not
more than 0·0075 mm., that is, only the most minute; with a
velocity _v_ = 0·2 mm. per second, particles having a diameter _d_
= 0·011 mm. are carried away; with _v_ = 0·3 mm., _d_ = 0·0146
mm.; with _v_ = 0·4 mm., _d_ = 0·017 mm.; with _v_ = 0·5 mm., _d_
= 0·02 mm.; with _v_ = 1 mm., _d_ = 0·03 mm.; with _v_ = 4 mm.,
_d_ = 0·07 mm.; with _v_ = 10 mm., _d_ = 0·137 mm.; with _v_ = 12
mm., _d_ = 0·15 mm.; and therefore if the current does not exceed
one of these velocities, it will only carry away or wash away
particles having a diameter less than that indicated. The sand and
other particles mixed with the clay will then remain in the
vessel. The very minute particles obtained after levigation are
all considered as clay, although not only clay but other rock
residue may also exist in it as very fine particles. However, this
is very seldom the case, and the fine mud separated from all clays
has practically the same composition as the purest kinds of
kaolin.
The relation between the amounts of clay and sand in soils used
for the cultivation of plants is very important, because a soil
rich in clay is denser, heavier, shrinks up under the action of
heat, and does not readily yield to the plough in dry or wet
weather, whilst a soil rich in sand is friable, crumbling, easily
parts with its moisture and dries rapidly, but is comparatively
easily worked. Neither crumbling sand nor pure clay can be
regarded as a good _cultivating soil_. The difference in the
amounts of clay and sand in a soil has also a purely chemical
signification. Sand is easily permeated by the air, because its
particles are not closely packed together. Hence the chemical
change of manures proceeds very easily in sandy soils. But on the
other hand such soils do not retain the nutritious principles
contained in the manure, nor the water necessary for the
nourishment of plants by means of their roots. Solutions of
nutritious substances, containing salts of potassium, phosphoric
acid, &c., when passed through sand only leave a portion
moistening the surface of its particles. The sand has only to be
washed with pure water and all the adhering films of solution are
washed away. It is not so with clay. If the above solutions be
passed through a layer of clay the retention of the nutritive
substances of these solutions will be very marked; this is partly
because of the very large surface which the minute particles of
clay expose. The nutritive elements dissolved in water are
retained by the particles of clay in a peculiar manner--that is,
the absorptive power of clay is very great compared to that of
sand--and this has a great significance in the economy of nature
(Chapter XIII., p. 547). It is evident that for cultivation the
most convenient soils in every respect will be those containing a
definite mixture of clay and sand, and indeed the most fertile
soils have this composition. The study of fertile soils, which is
so important for a knowledge of the natural conditions for the
application of fertilisers, belongs, strictly speaking, to the
province of agriculture. In Russia the first foundation of a
scientific fertilisation has been laid by Dokuchaeff. As an
example only, we will give the composition of four soils; (1) The
black earth of the Simbirsk Government; (2) a clay soil from the
Smolensk Government; (3) a more sandy soil from the Moscow
Government; and (4) a peaty soil from near St. Petersburg. These
analyses were made in the laboratory of the St. Petersburg
University about 1860, in connection with experiments on
fertilisation (conducted by me) by the Imperial Free Economical
Society. 10,000 grams of air-dried soil contain the following
quantities (in grams) of substances capable of dissolving in
acids, and of serving for the nourishment of plants.
(1) (2) (3) (4)
Na_{2}O 11 5 4 4
K_{2}O 58 10 7 5
MgO 92 33 19 7
CaO 134 17 14 11
P_{2}O_{5} 7 1 7 3
N 44 11 13 16
S 13 7 7 6
Fe_{2}O_{3} 341 155 111 46
By chemical and mechanical analysis, the chief component parts per
100 parts of air-dried soil are
Clay 46 29 12 10
Sand 40 67 86 84
Organic matter 3·7 1·7 0·6 4·1
Hygroscopic water 6·3 1·3 0·8 1·9
Weight of a litre in grams 1150 1270 1350 960
The black earth excels the other soils in many respects, but
naturally its stores are also exhausted by cultivation if nothing
be returned to it in the form of fertilisers; and the improvement
of a soil (for instance, by the addition of marl or peat, and by
drainage and watering), and its fertilisation, if carried on in
conformity with its composition and with the properties of the
plants to be cultivated, are capable of rendering not only every
soil fit for cultivation, but also of improving its value, so that
in the course of time whole countries (like Holland) may clearly
improve their agricultural position, whilst under the ordinary
_régime_ of continued exhaustion of the soil, entire regions (as,
for instance, many parts of Central Asia) may be rendered unfit
for any agriculture.
By treating clay with strong sulphuric acid, which dissolves the alumina
in it, and then (by means of an alkaline carbonate) dissolving the silica
which was combined with the alumina in the clay (but not that occurring
in the form of sand, &c., which is hardly dissolved by carbonate of soda
solution at all even on boiling), we may form an idea of the proportion
between the component parts of a clay; and by igniting it at a high
temperature, we may determine the amount of water held in it. In the
purer sorts of clay dried at 100° (sp. gr. of pure kaolin is about 2·5)
this proportion is about 2SiO_{2} : 2H_{2}O : Al_{2}O_{3}. In this case
the conversion of felspar into kaolin is expressed by the equation:--
K_{2}O,Al_{2}O_{3},6SiO_{2} = Al_{2}O_{3},2SiO_{2} + K_{2}O,4SiO_{2};
Felspar Kaolin
the compound K_{2}O,4SiO_{2} passes into solution.
But as a rule clays contain from 45 to 60 p.c. of silica, from 20 to 30
p.c. of alumina, and about 12 p.c. of water; and it cannot be supposed
that clays are always homogeneous, because they are an aggregation of
residues (of silico-aluminous compounds) which are unacted on by water.
Nevertheless, clays always contain a hydrous compound of alumina and
silica, which is able to give up the alumina contained by it as a base to
strong sulphuric acid, forming aluminium sulphate, which is soluble in
water. After this treatment the silica remains, and is soluble in a
solution of an alkaline carbonate.[18]
[18] Everyone knows that a mixture of clay and water is endowed with
the property of taking a given form when subjected to a moderate
pressure. This plasticity of clay renders it an invaluable
material for practical purposes. From clay are moulded and
manufactured a variety of objects, beginning with the common brick
and ending with the most delicate china works of art. This
_plasticity of clay_ increases with its purity. When articles made
of clay are dried, the well-known hard mass is obtained; but water
washes it away, and furthermore, the cohesion of its particles is
not sufficiently great for it to resist the impression of blows,
shocks, &c. If such an article be subjected to the action of heat,
its volume first decreases, then it begins to lose water, and it
shrinks still further (in the case of a compact mass approximately
by 1/5 of its linear measurement). On the other hand, a great
coherence of particles is obtained, and thus burnt clay has the
hardness of stone. Pure clay, however, shrinks so considerably
when burnt that the form given to it is destroyed and cracks
easily form; such vessels are also porous, so that they will not
hold water. The addition of sand--that is, silica in fine
particles--or of _chamotte_--that is, already burnt and crushed
clay--renders the mass much more dense and incapable of cracking
in the furnace. Nevertheless, such clay articles (bricks,
earthenware vessels, &c.) are still porous to liquids after being
burnt, because the clay in the furnace is only baked and does not
fuse. In order to obtain articles impervious to water the clay
must either be mixed with substances which form a glassy mass in
the furnace, permeating the clay and filling up its pores, or else
only the surface of the article is covered with such a glassy
fusible substance. In the first case the purest kinds of clay give
what is known as china, in the second case porcelain or 'faïence.'
So, for instance, by covering the surface of clay articles with a
layer of the oxides of lead and tin, the well-known white glaze is
obtained, because the oxides of these metals give a white gloss
when fused with silica and clay. In the preparation of china,
fluor spar and finely ground silica is mixed up into the clay;
these ingredients give a mass which is infusible but softens in
the furnace, so that all the particles of the clay cohere in this
softened mass, which hardens on cooling. A glaze composed of
glassy substances, which only fuse at a high temperature, is also
applied to the surface of china articles.
Clay is the source from which alumina, Al_{2}O_{3}, and the majority of
the compounds of aluminium are prepared. Among these compounds the most
important are the alums--that is, the double sulphates of potassium (and
allied metals) and aluminium, AlK(SO_{4})_{2},12H_{2}O. When clay is
treated with sulphuric acid diluted with a certain amount of water,
aluminium sulphate, Al_{2}(SO_{4})_{3}, is formed; and if potassium
carbonate or sulphate be added to this solution, a double salt or alum is
obtained in solution. The alums crystallise easily, and are prepared on a
very large manufacturing scale owing to their being employed in the
process of dyeing. Alums are soluble in water, and, on the addition of
ammonia to their solutions, they give _hydrous alumina_, or _aluminium
hydroxide_, as a white gelatinous precipitate, which is insoluble in
water but easily soluble in acids, even when dilute, and in aqueous soda
or potash. The solubility of alumina in acids indicates the basic
character of the oxide, and its solubility in alkalis and its power of
forming compounds with them shows the weakness of this basic character.
However, the feeblest acids, even carbonic acid, take up the alkali from
such a solution, and the alumina then separates out in a precipitate as
the hydroxide. It must also be remembered as characteristic of the
salt-forming properties of alumina that it does not combine with such
feeble acids as carbonic, sulphurous, or hypochlorous, &c.--that is, its
compounds with these acids are decomposed by water. It is also important
to observe that the hydroxide is not soluble in aqueous ammonia.
_Alumina_, Al_{2}O_{3}--that is, the anhydrous aluminium oxide--is
met with in nature, sometimes in a somewhat pure state, having
crystallised in transparent crystals, which are often coloured by
impurities (chromic, cobaltic, and ferric compounds). Such are the ruby
and sapphire, the former red and the latter blue. They have a specific
gravity 4·0, are distinguished by their very great hardness, which is
second only to that of the diamond, and they represent the purest form of
alumina. They are found in Ceylon and other islands of the Indian
Archipelago, embedded in a rock matrix.[18 bis] _Corundum_ is the same
crystallised anhydrous alumina coloured brown by a trace of oxide of
iron. A very much larger portion of this impurity occurs in _emery_,
which is found in crystalline masses in Asia Minor and in Massachusetts,
and owing to its extreme hardness is employed for polishing stones and
metals. In this anhydrous and crystalline state the aluminium oxide is a
substance which very powerfully resists the action of reagents, and is
insoluble both in solutions of the alkalis and in strong acids. It is
only capable of passing into solution after being fused with alkalis.[19]
Alumina may be obtained in this form by artificial means if the hydroxide
be ignited and then fused in the oxyhydrogen flame.[20] Alumina also
occurs in nature in combination with water--as, for instance, in the
rather rare minerals _hydrargillite_ (sp. gr. 2·3), Al_{2}O_{3},3H_{2}O =
2Al(HO)_{3}, and _diaspore_, Al_{2}O_3,H_{2}O = 2AlO(HO) (sp. gr. 3·4). A
less pure hydrate, mixed with ferric oxide, sometimes occurs in masses
(at Baux in the south of France) and is termed _bauxite_; it contains
Al_{2}O_{3},2H_{2}O = Al_{2}O(HO)_{4} (sp. gr. 2·6). When bauxite is
ignited with sodium carbonate, carbonic anhydride is liberated and the
alumina then combines with the sodium oxide, forming a saline aluminate
of the oxides of aluminium and sodium. This is taken advantage of in
practice for the preparation of pure alumina compounds on a large scale,
for bauxite is found in large masses (in the South of France, in Austria,
and in Carolina in South America), and the resultant compound of alumina
and sodium is soluble in water and does not contain ferric oxide. This
solution when subjected to the action of carbonic anhydride gives a
precipitate of aluminium hydroxide,[21] which with acids forms aluminium
salts. If aqueous ammonia be added to a solution of aluminium sulphate a
gelatinous precipitate is formed, which at first remains suspended in the
liquid and then on settling forms a gelatinous mass, which itself
indicates the _colloidal property of aluminium hydroxide_. The following
points are characteristic of this colloidal state: (1) in an anhydrous
state such a colloidal substance is insoluble in water, as alumina is;
(2) in the hydrated state, it is gelatinous and insoluble in water; and
(3) it is also capable of existing in solutions, from which it separates
out in a non-crystalline state, forming a substance resembling glue.
These different states of colloids were distinguished by Graham, who gave
them the following very characteristic names. He called the gelatinous
form of the hydrate _hydrogel_, _i.e._ a gelatinous hydrate, and the
soluble form of the aqueous compound, _hydrosol_, from the Latin for a
soluble hydrate. Alumina readily and frequently assumes these states. The
gelatinous hydrate of alumina is its hydrogel. It is, as has been already
mentioned, insoluble in water, and, like all similar hydrogels, shows not
the faintest sign of crystallisation; it is apt to vary in many of its
properties with the amount of water it contains, and loses its water on
ignition, leaving a white powder of the anhydrous oxide. The hydrogel of
alumina is soluble both in acids and alkalis. It may also be obtained by
the evaporation of its solutions in such feebly energetic acids as
volatile acetic acid. These properties are very frequently made use of in
the arts, and especially in _the processes of dyeing_, because the
hydrogel of alumina in precipitating attracts a number of colouring
matters from their solutions, the precipitate being thus coloured by the
dyes attracted.[22] The preparation of fixed dyes and the employment of
aluminous compounds (mordants) in the processes of dyeing are founded on
this fact.[23] When precipitated upon the fibres of tissues (calicoes,
linens, &c.) the aluminium hydroxide renders them impermeable to water;
this may be taken advantage of for the preparation of waterproof tissues.
[18 bis] Frémy (1890) obtained transparent rubies, which crystallised
in rhombohedra, and resembled natural rubies in their hardness,
colour, size, and other properties. He heated together a mixture
of anhydrous alumina containing more or less caustic potash, with
barium fluoride and bichromate of potassium. The latter is added
to give the ruby its colour, and is taken in small quantity (not
more than 4 parts by weight to 100 parts of alumina). The mixture
is put into a clay crucible, and heated (for from 100 hours to 8
days) in a reverberatory furnace at a temperature approaching
1,500°. At the end of the experiment the crucible was found to
contain a crystalline mass, and the walls were covered with
crystals of the ruby of a beautiful rose colour. It was found that
the access of moist air was indispensable for the reaction.
According to Frémy, the formation of the ruby may be here
explained by the formation of fluoride of aluminium which under
the action of the moist air at the high temperature of the furnace
gives the ruby and hydrofluoric acid gas.
[19] The effects of purely mechanical subdivision on the solubility of
alumina are evident from the fact that native anhydrous alumina,
when converted into an exceedingly fine powder by means of
levigation, dissolves in a mixture of strong sulphuric acid and a
small quantity of water, especially when heated in a closed tube
at 200°, or when fused with acid sulphate of potassium (_see_
Chapter XIII., Note 9).
[20] The preparation of crystallised alumina is given on p. 65, and in
Note 18 bis. When alumina, moistened with a solution of cobalt
salt, is ignited, it forms a blue mass called Thénard's salt. This
coloration is taken advantage of not only in the arts, but also
for distinguishing alumina from other earthy substances resembling
it.
[21] The treatment of bauxite is carried on on a large scale, chiefly
in order to obtain alumina from alkaline solutions, free from
ferric oxide, because in dyeing it is necessary to have salts of
aluminium which do not contain iron. But this end, it would seem,
may also be obtained by igniting alumina containing ferric oxide
in a stream of chlorine mixed with hydrocarbon vapours, as ferric
chloride then volatilises. K. Bayer observed that in the treatment
of bauxite with soda, about 4 molecules of sodium hydroxide pass
into solution to 1 molecule of alumina, and that on agitating this
solution (especially in the presence of some already precipitated
aluminium hydroxide), about two-thirds of the alumina is
precipitated, so that only 1 molecule of alumina to 12 molecules
of sodium hydroxide remains in solution. This solution is
evaporated directly, and used again. He therefore treats bauxite
directly with a solution of NaHO at 170° in a closed boiler, and
on cooling adds hydrated alumina to the resultant solution. The
greater part of the dissolved alumina then precipitates on this
hydrated alumina, and the solution is used over again. The
hydroxide which separates from the alkaline solution contains
Al(OH)_{3}. All these properties bear a great resemblance to those
of boric acid. It may be taken for granted that the relation
between sodium hydroxide and alumina in solution varies with the
mass of water.
If lime be added to a solution of alumina in alkali (sodium
aluminate) calcium aluminate is precipitated, from which acids
first extract the lime, leaving aluminium hydroxide, which is
easily soluble in acids (Loewig). When sodium aluminate is mixed
with a solution of sodium bicarbonate, a double carbonate of the
alkali and aluminium is precipitated, which is easily soluble in
acids.
[22] These coloured precipitates of alumina are termed _lakes_, and are
employed in dyeing tissues and in the formation of various
pigments--such as pastels, oil colours, &c. Thus, if organic
colouring matters, such as logwood, madder, &c., are added to a
solution of any aluminium salt, and then an alkali is added, so
that alumina may be precipitated, these pigments, which are by
themselves soluble in water, will come down with the precipitate.
This shows that alumina is able to combine with the colouring
matter, and that this compound is not decomposed by water. The
dyes then become insoluble in water. If a dye be mixed with starch
paste and aluminium acetate, and then, by means of engraved blocks
having a design in relief, we transfer this mixture to a fabric
which is then heated, the aluminium acetate will leave the
hydrogel of alumina which binds the colouring matter, and water
will no longer be able to wash the pigment from the material--that
is, a so-called 'fixed' dye is obtained. In the case of dyeing a
fabric a uniform tint, it is first soaked in a solution of
aluminium acetate and then dried, by which means the acetic acid
is driven off, while the hydrogel of alumina adheres to the fibres
of the material. If the latter be then passed through a solution
of a dye in water, the former will be attracted to the portions
covered with alumina, and closely adhere to them. If certain parts
of the material be protected by the application of an acid, such
as tartaric, C_{4}H_{6}O_{6}, oxalic, citric, &c. (these acids
being non-volatile), the alumina will be dissolved in those parts,
and the pigment will not adhere, so that after washing, a white
design will be obtained on those parts which have been so
protected.
In dye-works the aluminium acetate is generally obtained in
solution by taking a solution of alum, and mixing it with a
solution of lead acetate. In this case lead sulphate is
precipitated and aluminium acetate remains in solution, together
with either acetate or sulphate of potassium, according to the
amount of acetate of lead first taken. The complete decomposition
will be as follows: KAl(SO_{4})_{2} + 2Pb(C_{2}H_{3}O_{2})_{2} =
KC_{2}H_{3}O_{2} + Al(C_{2}H_{3}O_{2})_{3} + 2PbSO_{4}, or the
less complete decomposition, 2KAl(SO_{4})_{2} +
3Pb(C_{2}H_{3}O_{2})_{2} = 2Al(C_{2}H_{3}O_{2})_{3} + K_{2}SO_{4}
+ 3PbSO_{4}. If the resultant solution of aluminium acetate be
evaporated or further boiled, the acetic acid passes off and the
hydrogel of alumina remains.
As the salt of potassium obtained in the solution passes away with
the water used for washing, and the salt of lead precipitated has
no practical use, this method for the preparation of aluminium
acetate cannot be considered economical; it is retained in the
process of dyeing mainly because both the salts employed, alum and
sugar of lead, easily crystallise, and it is easy to judge of
their degree of purity in this form. Indeed, it is very important
to employ pure reagents in dyeing, because if impurity is
present--such as a small quantity of an iron compound--the tint of
the dye changes; thus madders give a red colour with alumina, but
if oxide of iron be present the red changes into a violet tint.
The aluminium hydroxide is soluble in alkalis, whilst ferric oxide
is not. Therefore sodium aluminate--that is, the dissolved
compound of alumina and caustic soda--obtained, as already
described, from bauxite, is sometimes employed in dyeing. Every
aluminium salt gives a solution containing sodium aluminate free
from iron, when it is mixed with excess of caustic soda. This
solution, when mixed with a solution of ammonium chloride, gives a
precipitate of the hydrogel of alumina: Al(OH)_{3} + 3NaHO +
3NH_{4}Cl = Al(OH)_{3} + 3NaCl + 3NH_{4}OH. There was originally
free soda, and on the addition of sal-ammoniac there is free
ammonia, and this does not dissolve alumina, therefore the
hydrogel of the latter is precipitated.
[23] Another direct method for the preparation of pure aluminium
compounds consists in the treatment of _cryolite_ containing
aluminium fluoride together with sodium fluoride, AlNa_{3}F_{6}.
This mineral is exported from Greenland, and is also found in the
Urals. It is crushed and heated in reverberatory furnaces with
lime, and the resultant mass is treated with water; sodium
aluminate is then obtained in solution, and calcium fluoride in
the precipitate AlNa_{3}F_{6} + 3CaO = 3CaF_{2} + AlNa_{3}O_{3}.
_The hydrosol_ of alumina--_i.e._ the soluble aluminium hydroxide--is
more difficult to obtain.[24] In order to obtain this soluble variety of
alumina, Graham took a solution of its hydrogel in hydrochloric
acid--that is, a solution of aluminium chloride, which is able to
dissolve a still further quantity of the hydrogel of alumina, forming a
basic salt having probably one of the compositions Al(HO)Cl_{2} or
Al(HO)_{2}Cl. When such a solution, considerably diluted with water, is
subjected to dialysis--that is, to diffusion through a membrane[25]--the
hydrochloric acid diffuses through the membrane and leaves the alumina in
the form of hydrosol. The resultant solution, even when only containing
two or three per cent. of alumina, passes into the hydrogel state with
such facility that it is sufficient to transfer it from one vessel to
another which has not been previously washed with water, for the entire
mass to solidify into a jelly. But a solution containing not more than
one-half per cent. of alumina may even be boiled without coagulating;
however, after the lapse of several days this solution will of its own
accord yield the hydrogel of alumina.[25 bis]
[24] Crum first prepared a solution of basic acetate of alumina--that
is, a salt containing as large as possible an excess of aluminium
hydroxide with as small as possible a quantity of acetic acid. The
solution must be dilute--that is, not contain more than one part
of alumina per 200 of water--and if this solution be heated in a
closed vessel (so that the acetic acid cannot evaporate) to the
boiling point of water, for one and a half to two days, then the
solution, which apparently remains unaltered, loses its original
astringent taste, proper to solutions of all the salts of alumina,
and has instead the purely acid taste of vinegar. The solution
then no longer contains the salt, but acetic acid and the hydrosol
of alumina in an uncombined state; they may be isolated from each
other by evaporating the acetic acid in shallow vessels at the
ordinary temperature, and with a thin layer of liquid the alumina
does not separate as a precipitate. When the acid vapours cease to
come off there remains a solution of the hydrosol of alumina,
which is tasteless and has no action on litmus paper. When
concentrated, this solution acquires a more and more gluey
consistency, and when completely evaporated over a water-bath it
leaves a non-crystalline glue-like hydrate, whose composition is
Al_{2}H_{4}O_{5} = Al_{2}O_{3},2H_{2}O. The smallest quantity of
alkalis, and of many acids and salts, will convert the hydrosol
into the hydrogel of alumina--that is, convert the aluminium
hydroxide from a soluble into an insoluble form, or, as it is
said, cause the hydrate to coagulate or gelatinise. The smallest
amount of sulphuric acid and its salts will cause the alumina to
gelatinise--that is, cause the hydrogel to separate. Many such
colloidal solutions are known (Vol. I. p. 98, Note 57).
[25] In a dialyser, Vol. I. p. 63, Note 18.
[25 bis] The different states in which the hydrates of alumina occur
and are prepared resemble similar varieties of the hydrates of the
oxides of iron and chromium, of molybdic and tungstic acids, as
well as of phosphoric and silicic acids, of many sulphides,
proteid substances, &c. We shall therefore have occasion to recur
to this subject in the further course of this work.
The most remarkable peculiarity of Graham's solution is that it
solidifies on litmus paper, and leaves a blue ring on it, which
shows the alkaline--that is, basic--character of the alumina in
such a solution. If in the dialysis the basic hydrochloric acid
salt be replaced by a similar acetic acid salt, a hydrosol of
alumina is obtained which does not act upon litmus.
With respect to alumina as a base, it is very important to observe that
it is not only capable of combining with other bases[26] but that it does
not give salts with feeble volatile acids (like carbonic and
hypochlorous); it forms salts which are easily decomposed by water,
especially when heated,[27] as well as double and basic salts,[28] _so
that it forms a clear example of a feeble base_.[29] To these
characteristics of alumina we must add that it not only gives compounds
of the type AlX_{3}, but also the polymeric type Al_{2}X_{6}, even when X
is a simple univalent haloid like chlorine. Deville and Troost showed
(1857) that the vapour density of aluminium chloride (at about 400°) is
9·37 with respect to air--that is, nearly 135 with respect to hydrogen,
and therefore the formula of its molecule is expressed by Al_{2}Cl_{6},
and not AlCl_{3},[30] although in the case of boron, arsenic, and
antimony, which give oxides R_{2}O_{3} of the same composition as
Al_{2}O_{3}, the chlorine compounds form non-polymeric molecules,
BCl_{3}, AsCl_{3}, SbCl_{3}.[31] This duplication (polymerisation) of the
form AlX_{3} is connected with the facility with which the salts of
aluminium combine with other salts to form double salts and with
aluminium hydroxide itself to form basic salts.
[26] Compounds of alumina with bases (aluminates, _see_ Note 21) are
sometimes met with in nature. Such are spinel (_see_ p. 65),
MgO,Al_{2}O_{3} = MgAl_{2}O_{4}, chrysoberyl, BeAl_{2}O_{4}, and
others. Magnetic oxide of iron, FeO,Fe_{2}O_{3} = Fe_{3}O_{4}, and
compounds like it, belong to the same class. Here we evidently
have a case of combination 'by analogy,' as in solutions and
alloys, accompanied by the formation of strictly definite saline
compounds, and such instances form a clear transition from
so-called solutions and certain mixtures to the type of true
salts.
[27] Not only aluminium acetate (Note 24), but also every other
aluminium salt with a volatile acid, parts with its acid on
heating an aqueous solution--that is, is decomposed by water, and
forms either basic salts or a hydrate of alumina. By dissolving
aluminium hydroxide in nitric acid we may easily obtain a
well-crystallising _aluminium nitrate_, Al(NO_{3})_{3},9H_{2}O,
which fuses at 73° without decomposing (Ordway), gives a basic
salt, 2Al_{2}O_{3},6HNO_{3}, at 100°, and at 140° leaves the
aluminium hydroxide perfectly free from the elements of nitric
acid. But the solutions of this salt, like those of the acetate,
are also able to yield aluminium hydroxide. From all this it is
evident that we must suppose that the solutions of this and
similar salts contain an equilibrated dissociated system,
containing the salt, the acid, and the base, and their compounds
with water, as well as partly the molecules of water itself. Such
examples much more clearly confirm those conceptions of solutions
which are given in the first chapter than a general preliminary
acquaintance with the subject can do.
[28] As an example of native basic salts we may cite _alunite_, or
alum-stone (sp. gr. 2·6), which sometimes occurs in crystals, but
more frequently in fibrous masses. It has been found in masses in
the Caucasus (at Zaglik, forty versts distance from Elizabetpol),
and at Tolfa, near Rome. Its composition is
K_{2}O,3Al_{2}O_{3},4SO_{3},6H_{2}O (alunite contains 9H_{2}O). It
is soluble in water but not decomposed by it, but after being
slightly ignited it gives up alum to it. It may be artificially
prepared by heating a mixture of alum with aluminium sulphate in a
closed tube at 230°.
[29] As the colloidal properties are particularly sharply developed in
those oxides (Al_{2}O_{3}, SiO_{2}, MoO_{3}, SnO_{2}, &c.) which
show (like water also) the properties of feeble bases and feeble
acids, there is probably some causal reason for this coincidence,
all the more so since among organic substances--gelatins,
albumins, &c.--the representatives of the colloids also have the
property of feebly combining with bases and acids.
[30] Since Deville's experiments the question of the density of
aluminium chloride has been frequently re-investigated. The
subject has more especially occupied the attention of Nilson,
Pettersson, Friedel and Crafts, and V. Meyer and his
collaborators. In general, it has been found that at low
temperatures (up to 440°) the density is constant, and indicates a
molecule Al_{2}Cl_{6}; whilst depolymerisation probably (although
it is not yet certain) takes place at higher temperatures, and the
molecule AlCl_{3} is obtained. Along with this there has been, and
still is, a difference of opinion as to the vapour density of
aluminium ethyl and methyl--whether for instance, Al(CH_{3})_{3}
or Al_{2}(CH_{3})_{6} expresses the molecule of the latter. The
interest of these researches is intimately connected with the
question of the valency of aluminium, if we hold to the opinion
that elements in their various compounds have a constant and
strictly definite valency. In this case the formula AlCl_{3} or
Al(CH_{3})_{3} would show that Al is trivalent, and that
consequently the compounds of aluminium are Al(OH)_{3}, AlO_{3}Al,
and, in general, AlX_{3}. But if the molecule be Al_{2}Cl_{6}, it
is--for the followers of the doctrine of the invariable valency of
the elements--incompatible with the idea of the trivalency of
aluminium, and they assume it to be quadrivalent like carbon,
likening Al_{2}Cl_{6} to ethane C_{2}H_{6} = CH_{3}CH_{3},
although this does not explain why Al does not form AlCl_{4}, or,
in general, AlX_{4}. In this work another supposition is
introduced; according to this, although aluminium, as an element
of group III., gives compounds of the type AlX_{5}, this does not
exclude the possibility of these molecules combining with others,
and consequently with _each other_--that is, forming Al_{2}X_{6};
just as the molecules of univalent elements exist either as H_{2},
Cl_{2}, &c., or as Na, and the molecules of bivalent elements
either as Zn, or as S_{2}, or even S_{6}. In the first place it
must be recognised that the limiting form does not exhaust all
power of combination, it only exhausts the capacity of the element
for combining with X's, but the saturated substance may afterwards
combine with _whole molecules_, which fact is best proved by the
capacity of substances to form crystalline compounds with water,
ammonia, &c. But in some substances this faculty for further
combinations is less developed (for instance, in carbon
tetrachloride, CCl_{4}), whilst in others it is more so. AlX_{3}
combines with many other molecules. Now if a limiting form, which
does not combine with new X's, nevertheless combines with other
whole molecules, it will naturally in some instances combine with
itself, will polymerise. In this manner the mind clearly grasps
the idea that the same forces which cause S_{2} to unite itself to
Cl_{2}, or C_{2}H_{4} to Cl_{2}, &c., also unite molecules of a
similar kind together; thus _polymerisation_ ceases to be an
isolated fragmentary phenomenon, and chemical combinations 'by
analogy' acquire a particular and important interest. In
conformity with these views the following proposition may be made
concerning the compounds of aluminium. They are of the type
AlX_{3} in the limit, like BX_{3}, but those limiting forms are
still able to combine to form AlX_{3},RZ, and the aluminium
chloride is a compound of this kind--_i.e._ (AlX_{3})_{2}. In
boron, for example, in BCl_{3}, this tendency to form further
compounds is less developed. Hence boron chloride appears as
BCl_{3}, and not (BCl_{3})_{2}. Polymerisation is not only
possible when a substance has not attained the limit (although it
is more probable then), but also when the limiting form has been
reached, if only the latter has the faculty of combining with
other whole molecules. We may therefore conclude that aluminium,
like boron, is trivalent in the same sense that lithium and sodium
are univalent, magnesium bivalent, and carbon tetravalent. In a
word, there is no reason to consider that aluminium is capable of
forming compounds AlX_{4}, and in that way to explain the
existence of the molecule Al_{2}Cl_{6}. Furthermore, there are
many reasons for thinking that AlF_{3}, Al_{2}O_{3}, and other
empirical formulæ do not express the molecular weights of these
compounds, but that they are much higher: Al_{_n_}F_{3_n_},
Al_{2_n_}O_{3_n_}. In recent years convincing proofs of the truth
of the above statements have been obtained, and of the independent
existence of AlX_{3} in a state of vapour; for Comb has determined
the vapour density of the volatile acetyl of aluminium acetate
Al(C_{3}H_{7}O_{2})_{3} (which melts at 193°, boils at 315°, and
distils without a trace of decomposition), and has found that it
exactly corresponds to the above molecular composition. On the
other hand, Louise and Roux (1889) by employing the method of
'freezing point depression' of solutions (Chapter I., Note 49)
found that the molecules Al_{2}(C_{2}H_{5})_{6} and
Al_{2}(C_{5}H_{11})_{6}, &c., correspond to the type Al_{2}X_{6}.
Thus it may now be accepted that the molecular composition of the
compounds of aluminium in their simplest form is AlX_{3}, but that
they may polymerise and give Al_{2}X_{6} or, in general,
Al_{2}X_{3_n_}.
[31] In the case of gallium, as a close analogue of aluminium, Lecoq de
Boisbaudran (1880) showed that probably the molecule gallium
chloride contains Ga_{2}Cl_{6} at low temperatures and high
pressures, and that it dissociates into GaCl_{3} at high
temperatures and low pressures. The molecule of indium chloride
seems to exist only in the simplest form, InCl_{3}.
_Aluminium sulphate_, Al_{2}(SO_{4})_{3}, which is obtained by treating
clay or the hydrates of alumina with sulphuric acid, crystallises in the
cold with 27H_{2}O, or at the ordinary temperature in pearly crystals,
which are greasy to the touch and contain 16H_{2}O.[32] Its solutions act
like sulphuric acid--for instance, they evolve hydrogen with zinc,
forming basic salts, which are sometimes met with in nature (_aluminite_,
Al_{2}O_{3},SO_{3},9H_{2}O, _alumiane_, Al_{2}O_{3},2SO_{3}, and others),
and may be obtained by the decomposition of normal salts and by the
direct solution of the hydroxide in normal salts: these exhibit a varying
composition, (Al_{2}O_{3})_{_n_}(SO_{3})_{_m_}(H_{2}O)_{_q_}, where _m/n_
is less than 3. Aluminium sulphate is now prepared (from the pure hydrate
obtained from bauxite, Note 21) in large quantities for dyeing purposes
(instead of alums) as a mordant. With solutions of the alkali sulphates
(potassium, sodium, ammonium, rubidium, and cæsium sulphates), the normal
salt easily forms double salts, termed _alums_--for example, the ordinary
crystalline alum contains KAl(SO_{4})_{2},12H_{2}O, or
K_{2}SO_{4},Al_{2}(SO_{4})_{3},24H_{2}O. In the ammonium alums (which
leave a residue of alumina when ignited) the potassium is replaced by
ammonium (NH_{4}). Alums are used in large quantities, because there is
scarcely any other salt which crystallises so easily. In this respect the
alums formed by potassium and ammonium are equally convenient to purify,
because they present a considerable difference in their solubility at the
ordinary and higher temperatures. If the crystallisation be conducted
rapidly, the salt separates in minute crystals, but if it be slowly
deposited, especially in large masses, as in factories, then crystals
several centimetres long are sometimes obtained. At a higher temperature
alums are very much more soluble, and crystallise with greater
difficulty, and are therefore less easily freed from impurities; at 0°
100 parts of water dissolve 3 parts, at 30° 22 parts, at 70° 90 parts,
and at 100° 357 parts of potassium alum.[33] The solubility of ammonium
alum is slightly less. The specific gravity of potassium alum is 1·74, of
ammonium alum 1·63, and of sodium alum 1·60. Alums easily part with their
water of crystallisation; thus potash alum partially effloresces when
exposed to the air, and loses 9 mol. H_{2}O under the receiver of an
air-pump. At 100°, dry air passed over alums takes up nearly all their
water. As we have already mentioned (Chapter XV.), the law of isomorphous
substitutions exhibits itself more clearly in the alums than in any other
salts, and all alums not only contain the same amount of water of
crystallisation, MR(SO_{4})_{2},12H_{2}O (where M = K, NH_{4}, Na; R =
Al, Fe, Cr), and appear in crystals whose planes are inclined at equal
angles, but they also give every possible kind of isomorphous mixture.
The aluminium in them is easily replaced by iron, chromium, indium and
sometimes by other metals, whilst the potassium may be substituted by
sodium, rubidium, ammonium, and thallium, and the sulphuric acid may be
replaced by selenic and chromic acids.
[32] The pure salt (16H_{2}O) is not hygroscopic. In the presence of
impurities the amount of water increases to 18H_{2}O, and the salt
becomes hygroscopic.
[33] The common form of crystals of alums is octahedral, but if this
solution contains a certain small excess of alumina above the
ratio 2Al(OH)_{3} to K_{2}SO_{4}, and not more sulphuric acid than
3H_{2}SO_{4} to 2Al(OH)_{3}, then it easily forms combinations of
the cube and octahedron, and these alums are called 'cubic' alums.
They are valued by the dyer because they can contain no iron in
solution, for oxide of iron is precipitated before alumina, and if
the latter be in excess there can be no oxide of iron present.
These alums were long exported from Italy, where they were
prepared from alunite (Note 28).
_Aluminium chloride_, Al_{2}Cl_{6}, is obtained, like other similar
chlorides, (for instance MgCl_{2}) either directly from chlorine and the
metal, or by heating to redness an intimate mixture of the amorphous
anhydrous oxide and charcoal in a stream of dry chlorine.[33 bis] The
resultant sublimate is very volatile,[34] and forms a crystalline, easily
fusible mass, which deliquesces in the air and easily dissolves in water,
with the evolution of a large amount of heat.[34 bis] On evaporating this
solution, hydrochloric acid and aluminium hydroxide are liberated. But if
the solution be heated in a closed tube, with an excess of hydrochloric
acid, then, on cooling, crystals of AlCl_{3},6H_{2}O are obtained--that
is, aluminium chloride both combines with water and is decomposed by it.
And the faculty of the type AlX_{3} for combining with other molecules is
seen in the compounds of AlCl_{3} with many other chlorine compounds.
Thus, for example, a mixture of aluminium chloride with sulphur
tetrachloride gives Al_{2}Cl_{6},SCl_{4}, under the action of chlorine,
whilst with phosphorus pentachloride it forms AlCl_{3},PCl_{5}; it also
combines with NOCl. Thus, the compounds AlCl_{3},NOCl, AlCl_{3},POCl_{3},
AlCl_{3},3NH_{3}, AlCl_{3},KCl, AlCl_{3},NaCl are known.[35] The compound
of aluminium and sodium chlorides, AlNaCl_{4}, is very fusible and much
more stable in the air than aluminium chloride itself. It seems to be of
the same type as the alums. This compound, AlNaCl_{4}, is employed in the
extraction of metallic aluminium, as we shall presently proceed to
describe. Aluminium bromide, which is obtained by the direct combination
of metallic aluminium with bromine, closely resembles the chloride; it
melts at 90°, volatilises at 270°, and its vapour density indicates the
formula Al_{2}Br_{6}. Aluminium iodide is obtained by heating iodine with
finely divided aluminium in a closed tube; it is so easily decomposed by
oxygen that its vapour even explodes when mixed with it.[36]
[33 bis] It is also formed by the action of hydrochloric acid upon
metallic aluminium (Nilson and Pettersson), by heating alumina in
a mixture of the vapours of naphthaline and HCl (Faure, 1889), and
by the action of dry HCl upon an alloy of 14 p.c., or more of Al
and copper (Mobery).
[34] Aluminium chloride fuses at 178°, boils at 183° (pressure 755 mm.,
at 168° under a pressure of 250 mm., and at 213° under 2,278 mm.),
according to Friedel and Crafts, so that it boils immediately
after fusion. According to Seubert and Pallard (1892),
Al_{2}Cl_{6} fuses at 193°. Aluminium bromide fuses at about 92°,
and the iodide at 185° according to Weber, at 125° according to
Deville and Troost.
All these halogen compounds of aluminium are soluble in water.
_Aluminium fluoride_, AlF_{3} (Al_{_n_}F_{3_n_}), is insoluble in
water. It is obtained by dissolving alumina in hydrofluoric acid;
a solution is then formed, but it contains an excess of
hydrofluoric acid. When this solution is evaporated, crystals
containing Al_{2}F_{6},HF,H_{2}O are obtained. They are also
insoluble in water. By saturating the above solution with a large
quantity of alumina, and then evaporating, we obtain crystals
having the composition Al_{2}F_{6},7H_{2}O. All these compounds,
when ignited, leave insoluble anhydrous aluminium fluoride. It
forms colourless rhombohedra, which are non-volatile, of sp. gr.
3·1, and are decomposed by steam into alumina and hydrofluoric
acid. The acid solution apparently contains a compound which has
its corresponding salts; by the addition of a solution of
potassium fluoride, a gelatinous precipitate of AlK_{3}F_{6} is
obtained. A similar compound occurs in nature--namely,
AlNa_{3}F_{6}, or _cryolite_, sp. gr. 3·0.
[34 bis] In this respect aluminium chloride resembles the chloranhydrides
of the acids, and probably in the aqueous solution the elements of
the hydrochloric acid are already separated, at least partially,
from the aluminium hydroxide. The solution may also be obtained by
the action of aluminium hydroxide on hydrochloric acid.
[35] Here we see an instance in confirmation of what has been said in
Note 30--_i.e._ the action of the molecule AlCl_{3}. We will cite
still another instance confirming the power of alumina to enter
into complex combinations. Alumina, moistened with a solution of
calcium chloride, gives, when ignited, an anhydrous crystalline
substance (tetrahedral), which is soluble in acids, and contains
(Al_{2}O_{3})_{6}(CaO)_{10}CaCl_{2}. Even clay forms a similar
stony substance, which might be of practical use.
Among the most complex compounds of aluminium, _ultramarine_, or
_lapis lazuli_, must be mentioned. It occurs in nature near Lake
Baikal, in crystals, some colourless and others of various
tints--green, blue, and violet. When heated it becomes dull and
acquires a very brilliant blue colour. In this form it is used for
ornaments (like malachite), and as a brilliant blue pigment. At
the present time ultramarine is prepared artificially in large
quantities, and this process is one of the most important
conquests of science; for the blue tint of ultramarine has been
the object of many scientific researches, which have culminated in
the manufacture of this native substance. The most characteristic
property of ultramarine is that when placed in sulphuric acid it
evolves hydrogen sulphide and becomes colourless. This shows that
the blue colour of ultramarine is due to the presence of
sulphides. If clay be heated in a furnace with sodium sulphate and
charcoal (forming sodium sulphide) without access of air, a white
mass is obtained, which becomes green when heated in the air, and
when treated with water leaves a colourless substance known as
'white ultramarine.' When ignited in the air it absorbs oxygen and
turns blue. The coloration is ascribed to the presence of metallic
sulphides or polysulphides, but it is most probable that silicon
sulphide, or its oxysulphide, SiOS, is present. At all events the
sulphides play an important part, but the problem is not yet quite
settled. The formula Na_{8}Al_{6}Si_{6}O_{24}S is ascribed to
white ultramarine. The green probably contains more sulphur, and
the blue a still larger quantity. The last is supposed to contain
Na_{8}Al_{6}Si_{6}O_{24}S_{3}. It is more probable (according to
Guckelberger, 1882) that the composition of the blue varies
between Si_{18}Al_{18}Na_{20}S_{6}O_{71} and
Si_{18}Al_{12}Na_{20}S_{6}O_{69}. The latter may be expressed as
(Al_{2}O_{3})_{6}(SiO_{2})_{18}(Na_{2}O)_{10}S_{6}O_{5}, which
would indicate the presence of insufficiently-oxidised sulphur in
ultramarine.
[36] At the ordinary temperature aluminium does not decompose water,
but if a small quantity of iodine, or of hydriodic acid and
iodine, or of aluminium iodide and iodine, is added to the water,
then hydrogen is abundantly evolved. It is evident that here the
reaction proceeds at the expense of the formation of Al_{2}I_{6},
and that this substance, with water, gives aluminium hydroxide and
hydriodic acid, which, with aluminium, evolves hydrogen. Aluminium
probably belongs to those metals having a greater affinity for
oxygen than for the halogens (Note 36 tri).
_Metallic Aluminium_ was first prepared by Wöhler in 1822 as a grey
powder by the action of potassium on aluminium chloride. He afterwards
(in 1845) obtained it as a white compact metal, unoxidisable in the air,
and only slowly attacked by acids. Owing to the vast and wide occurrence
of clay, many efforts have been made in investigating in detail the
methods for the extraction of this metal. These efforts were brought to a
successful issue (1854) by Sainte-Claire Deville, who is also renowned
for his doctrine of dissociation. Experiments on a large scale have
proved that metallic aluminium, although possessed of great lightness,
strength, and durability, is not so generally suitable for technical
purposes as was at first thought. Nitric and many other acids, indeed, do
not act on it, but the alkalis, alkaline substances, and even salts--for
instance, moist table salt--humidity, &c.,[36 bis] tarnish it, and hence
objects made of aluminium suffer at the surfaces, alter, and cannot, as
was hoped, replace the precious metals, from which it differs in its
extreme lightness. But the alloys made with aluminium (especially with
copper, for example aluminium bronze) are very valuable in their
properties and applications.
[36 bis] As an example we may mention that if mercury comes in contact
with metallic aluminium and especially if it be rubbed upon the
surface of aluminium moistened with a dilute acid, the Al becomes
rapidly oxidised (Al_{2}O_{3} being formed). The oxidation is
accompanied by a very curious appearance, as it were of wool (or
fur) formed by threads of oxide of aluminium growing upon the
metal. This was first pointed out by Cass in 1870, and
subsequently by A. Sokoleff in 1892. This interesting and curious
phenomenon has not to my knowledge been further studied.
I think it necessary, however, to add that according to Lubbert
and Rascher's researches (1891), wine, coffee, milk, oil, urine,
earth, &c., have no more action upon aluminium vessels than upon
copper, tin, and other similar articles. In the course of four
months ordinary vinegar dissolved 0·35 grm. of Al per sq.
centimetre, whilst a 5 per cent. solution of common salt dissolved
about 0·05 grm. of aluminium. Ditte (1890) showed that Al is acted
upon by nitric and sulphuric acids, although only slowly (owing to
the formation of a layer of gas, as in Chapter XVI., Note 10) and
that the reaction proceeds much more rapidly in vacuo or in the
presence of oxidising agents. Al is even oxidised by water on the
surface, but the thin coating of alumina formed prevents further
action. In the course of twelve hours nitric acid sp. gr. 1·383
dissolved at 17° about 20 grms. of aluminium (containing only a
small amount of Si, 1-1/4 p.c.) from a sq. metre of surface (Le
Rouart, 1891).
The Deville method for the preparation of metallic aluminium is
based on the decomposition of the above-mentioned compound of
sodium and aluminium chlorides by metallic sodium. The compound is
obtained by passing the vapour of aluminium chloride (evolved from
a mixture of alumina, extracted from bauxite or cryolite, with
charcoal ignited in a stream of chlorine) over red-hot salt, when
the compound AlNaCl_{4}, is itself volatilised, and may in this
manner be obtained pure. A mixture of this compound with salt and
fluor spar, or with cryolite, is heated with a certain excess of
sodium, cut into small lumps. On a large scale this operation is
carried on in special furnaces with a small access of air and at a
high temperature. The decomposition takes place chiefly according
to the equation NaAlCl_{4} + 3Na = 4NaCl + Al. Neither charcoal
nor zinc will reduce the oxygen compounds of aluminium; even
sodium and potassium do not act on alumina. Moreover, metallic
aluminium, like magnesium, is able to reduce even the metals of
the alkalis from their oxygen compounds. This is connected with
the fact that the atom of oxygen evolves more heat in combining
with Al (and Mg) than it does in combining with other metals;
whilst on the other hand, chlorine (and the other halogens) evolve
more heat in combining with the metals of the alkalis.[36 tri]
[36 tri] In addition to the data given in Chapters XI., XIII., and
in Chapter XV., Note 19, the following are the amounts of heat in
thousands of units, evolved in the formation of the oxides and
chlorides from the metals taken in gram-atomic quantities:
Na_{2}O 100; MgO 140*; 1/3Al_{2}O_{3} 120*;
1/3Fe_{2}O_{3} 63*;
Na_{2}Cl_{2} 195; MgCl_{2} 151; 1/3Al_{2}Cl_{6} 107;
1/3Fe_{2}Cl_{6} 64.
The asterisks following the oxides of Mg, Al and Fe call attention
to the fact that the existing data refer to the formation of the
hydrates of these metals, from which the heat of formation of the
anhydrous oxides may easily be assumed, because the heat of
hydration (for example, MgO + H_{2}O) has not yet been determined.
Since the close of the eighties the metallurgy of aluminium has taken a
new direction, based upon the action of an electric current upon cryolite
at a high temperature,[37] and the solution of oxide of aluminium
(obtained from bauxite or in the form of corundum) in it; under these
conditions metallic aluminium is reduced at the negative pole (cathode)
in a sufficiently pure state, and if the cathode be copper, forms alloys
with it. Such are Hall's and Cowle's (both in the United States) and the
Neuhausen process (where the current is obtained from a dynamo worked by
the Falls of the Rhine at Schaffhausen). As an example, we will describe
(in the words of Prof. D. P. Konovaloff, who became acquainted with this
process at the Chicago Exhibition), Hall's process as applied near
Pittsburg, where it gives about 1,500 kilos of Al a day. An iron box
(about 1 metre long and 1/2 metre wide), provided with a well rammed down
charcoal lining, is charged with a mixture of cryolite and Al_{2}O_{3}
(from bauxite), over which salt is strewn, and a current of 5,000 ampères
at 20 volts is passed through the mixture. The anode is composed of a
carbon cylinder (about 9 cm. in diameter), while the charcoal lining
forms the cathode. When the temperature inside the box is raised to a red
heat by the current, the mixture fuses and the Al_{2}O_{3} begins to
decompose. The Al liberated collects at the bottom of the box, whilst the
oxygen evolved burns the charcoal anode. When the decomposition is at an
end, and the resistance of the mass increases, a fresh quantity of
Al_{2}O_{3} is added, and this is continued until the amount of
impurities accumulated in the furnace and passing into the metal becomes
too great.[37 bis]
[37] Cryolite under the action of the current at about 1,000° gives off
the vapour of Na which reduces the Al, but it recombines with the
liberated fluorine and again passes into the fused mass. It is
important to obtain aluminium at as low a temperature as possible,
but the action proceeds far more easily with the solution (alloy)
of oxide of aluminium in cryolite.
[37 bis] The cost of working this process can be brought as low as 20
cents per lb. or about 2-1/2 fcs. per kilo. In England, Castner,
prior to the introduction of the electric method, obtained Al by
taking a mixture of 1,200 parts of the double salt NaAlCl_{4}, 600
parts of cryolite, and 350 parts of Na, and obtained about 120
parts of Al, so that the cost of this process is about 1-1/2 time
that of the electric method.
Buchner found that sulphide of aluminium, Al_{2}S_{3}, is more
suitable for the preparation of Al by the electrolytic method than
Al_{2}O_{3}, but since the formation of Al_{2}S_{3} by heating a
mixture of Al_{2}O_{3}, and charcoal in sulphur vapour proceeds
with difficulty, Gray (1894) proposed to prepare Al_{2}S_{3} by
heating a mixture of charcoal, sulphate of aluminium, and sodium
fluoride. The resultant molten mixture of NaF and Al_{2}S_{3}
gives aluminium directly under the action of an electric current.
Aluminium has a white colour resembling that of tin--that is, it is
greyer than silver and has the feebly dull lustre of tin, but compared to
tin and pure silver, aluminium is very hard. Its density is 2·67--that
is, it is nearly four times lighter than silver and three times lighter
than copper. It melts at an incipient red heat (600°), and in so doing is
but slightly oxidised. At the ordinary temperature it does not alter in
the air, and in a compact mass it burns with great difficulty at a white
heat, but in thin sheets, into which it may be rolled, or as a very fine
wire, it burns with a brilliant white light, since it forms an infusible
and non-volatile oxide. Aluminium itself is non-volatile at a furnace
heat. These properties render Al a very good reducing agent, and N. N.
Beketoff showed that it reduces the oxides of the alkali metals (Chapter
XIII., Note 42 bis). Dilute sulphuric acid has scarcely any action on it,
but the strong acid dissolves it, especially with the aid of heat. Nitric
acid, dilute or strong, has no action whatever on it. On the other hand,
hydrochloric acid dissolves aluminium with great ease, as do also
solutions of caustic soda and potash. In the latter cases hydrogen is
evolved.[38]
[38] Aluminium, when heated to the high temperature of the electric
furnace, dissolves carbon and forms an alloy which, according to
Moissan, when rapidly treated with _cold_ hydrochloric acid leaves
a compound C_{3}Al_{4} in the form of a yellow crystalline
transparent powder, sp. gr. 2·36 (_see_ Chapter VIII. Note 12
bis). This _carbide of aluminium_ C_{3}Al_{4} corresponds to
methane CH_{4}, for Al replaces H_{3} and carbon O_{2} or H_{4},
that is, it is equal to three molecules of CH_{4} with the
substitution of twelve atoms of H in it by four of Al, or, what is
the same thing, it is the duplicated molecule of Al_{2}O_{3} with
the substitution of O_{6} by C_{3}. And indeed C_{3}Al_{4} under
the action of water forms marsh gas and hydrate of alumina:
C_{3}Al_{4} + 12H_{2}O = 3CH_{4} + 4Al(OH)_{3}. This decomposition
gives a new aspect of the synthesis of hydrocarbons, and quite
agrees with what should follow from the action of water upon the
metallic carbides as applied by me for explaining the origin of
naphtha (Chapter VIII., Notes 57, 58, and 59). Frank (1894) by
heating Al with carbon obtained a similar although not quite pure
compound, which (like CaC_{2}) evolves acetylene with hydrochloric
acid _i.e._ probably has the composition AlC_{3}.
Aluminium forms alloys with different metals with great ease. Among them
the copper alloy is of practical use. It is called _aluminium bronze_.
This alloy is prepared by dissolving 11 p.c. by weight of metallic
aluminium in molten copper at a white heat. The formation of the alloy is
accompanied by the development of a considerable quantity of heat, so
that it glows to a bright white heat. This alloy, which corresponds with
the formula AlCu_{3}, presents an exceedingly homogeneous mass,
especially if perfectly pure copper be taken. It is distinguished for its
capacity to fill up the most minute impressions of the mould into which
it may be cast, and by its extraordinary elasticity and toughness, so
that objects cast from it may be hammered, drawn, &c., and at the same
time it is fine-grained and exceedingly hard, takes an excellent polish,
and, what is most important, its surface then remains almost unchangeable
in the air, and has a colour and lustre which may be compared to that of
gold alloys. Hence aluminium bronze is much used in the arts for making
spoons, watches, vessels, forks, knives, and for ornaments, &c. No less
important is the fact that the admixture of one-thousandth part of
aluminium with steel renders its castings homogeneous (free from
cavities) to an extent that could not be arrived at by other means, nor
does the quality of the steel in any respect deteriorate by this
admixture, but rather is it improved. In a pure state, aluminium is only
employed for such objects as require the hardness of metals with
comparative lightness, such as telescopes and various physical apparatus
and small articles.
According to the periodic system of the elements, the analogues of
magnesium are zinc, cadmium, and mercury in the second group. So also in
the third group, to which aluminium belongs, we find its corresponding
analogues _gallium_, _indium_, and _thallium_. They are all three so
rarely and sparingly met with in nature that they could only be
discovered by means of the spectroscope. This fact shows that they are
partially volatile, as should be the case according to the property of
their nearest neighbours, the very volatile zinc, cadmium and mercury. As
with them, in gallium, indium, and thallium the density of the metal,
decomposability of compounds, &c., rises with the atomic weight. But here
we find a peculiarity which does not exist in the second group. In the
latter, the fusibility increases with the atomic weight of magnesium,
zinc, cadmium, and mercury; indeed, the heaviest metal--mercury--is a
liquid. In the third group it is not so. In order to understand this it
is sufficient to turn our attention to the elements of the further groups
of the uneven series--for instance, to group V., containing phosphorus,
arsenic, and antimony, or to group VI., with sulphur, selenium, and
tellurium, and also to group VII., where chlorine, bromine and iodine are
situated. In all these instances the fusibility decreases with a rise of
atomic weight; the members of the higher series, the elements of a high
atomic weight, fuse with greater difficulty than the lighter elements.
The representatives of the uneven series of group III., aluminium,
gallium, indium, thallium, forming, as they do, a transition, all show an
intermediate behaviour. Here the most fusible of all is the medium metal
gallium,[38 bis] which fuses at the heat of the hand; whilst indium,
thallium, and aluminium fuse at much higher temperatures.
[38 bis] The same is the case in group IV. of the uneven series, where
tin is the most fusible. Thus the temperature of fusion rises on
both sides of tin (silicon is very infusible; germanium, 900°;
tin, 230°; lead, 326°); as it also does in group III., starting
from gallium, for indium fuses at 176°, less easily than gallium
but more easily than thallium (294°). Aluminium also fuses with
greater difficulty than gallium.
Zinc (group II.), which has an atomic weight 65, should be followed in
group III. by an element with an atomic weight of about 69. It will be in
the same group as Al and should consequently give R_{2}O_{3}, RCl_{3},
R_{2}(SO_{4})_{3}, alums and similar compounds analogous to those of
aluminium. Its oxide should be more easily reducible to metal than
alumina, just as zinc oxide is more easily reduced than magnesia. The
oxide R_{2}O_{3} should, like alumina, have feeble but clearly expressed
basic properties. The metal reduced from its compounds should have a
greater atomic volume than zinc, because in the fifth series, proceeding
from zinc to bromine, the volume increases. And as the volume of zinc =
9·2, and of arsenic = 18, that of our metal should be near to 12. This is
also evident from the fact that the volume of aluminium = 11, and of
indium = 14, and our metal is situated in group III., between aluminium
and indium. If its volume = 11·5 and its atomic weight be about 69, then
its density will be nearly 5·9. The fact that zinc is more volatile than
magnesium gives reason for thinking that the metal in question will be
more volatile than aluminium, and therefore for expecting its discovery
by the aid of the spectroscope, &c.
These properties were indicated by me for the analogue of aluminium in
1871, and I named it (_see_ Chapter XV.) eka-aluminium. In 1875, Lecoq de
Boisbaudran, who had done much work in spectrum analysis, discovered a
new metal in a zinc blende from the Pyrenees (Pierrefitte). He recognised
its individuality and difference from zinc, cadmium, indium, and the
other companions of zinc by means of the spectroscope; but he only
obtained some fractions of a centigram of it in a free state.
Consequently only a few of its reactions were determined, as, for
instance, that barium carbonate precipitates the new oxide from its salts
(alumina, as is known, is also precipitated). Lecoq de Boisbaudran named
the newly discovered metal _gallium_. As one would expect the same
properties for eka-aluminium as were observed in gallium, I pointed out
this fact at the time in the Memoirs of the Paris Academy of Sciences.
All the subsequent observations of Lecoq de Boisbaudran confirmed the
identity between the properties of gallium and those indicated for
eka-aluminium. Immediately after this the ammonium alum of gallium was
obtained, but the most convincing proof of all was found in the fact that
the density of gallium although first apparently different (4·7) from
that indicated above, afterwards, when the metal was carefully purified
from sodium (which was first used as a reducing agent), proved to be just
that (5·9) which would have been looked for in the analogue of aluminium;
and, what was very important, the equivalent (23·3) and atomic weight
(69·8) determined by the specific heat (0·08) were shown by experiment to
be such as would be expected. These facts confirmed the universality and
applicability of the periodic system of the elements. It must be remarked
that previous to it there was no means of either foretelling the
properties or even the existence of undiscovered elements.[39]
[39] The spectrum of gallium is characterised by a brilliant violet
line of wave-length = 417 millionths of a millimetre. The metal
can be separated from the solution, containing a mixture of the
many metals occurring in the zinc blende, by making use of the
following reactions: it is precipitated by sodium carbonate in the
first portions; it gives a sulphate which, on boiling, easily
decomposes into a basic salt, very slightly soluble in water; and
it is deposited in a metallic state from its solutions by the
action of a galvanic current. It fuses at +30°, and, when once
fused, remains liquid for some time. It oxidises with difficulty,
evolves hydrogen from hydrochloric acid and from potassium
hydroxide, and, like all feeble bases (for instance, alumina and
indium oxide), it easily forms basic salts. The hydroxide is
soluble in a solution of caustic potash, and slightly so in
caustic ammonia. Gallium forms volatile GaCl_{3} and GaCl_{2}
(Nilson and Pettersson).
Much more light has been thrown on that element of the aluminium group
which follows after cadmium (its position in the periodic system is III.,
7, that is, it is in group III. in the 7th series). This is _indium_, In,
which also occurs in small quantities in certain zinc ores. It was
discovered (1863) by Reich and Richter (and more fully investigated by
Winkler) in the Freiberg zinc ores, and was named indium from the fact
that it gives to the flame of a gas-burner a blue coloration, owing to
the indigo blue spectral lines proper to it. The equivalent (_see_
Chapter XV., Note 15), specific heat, and other properties of the metal
confirm the atomic weight In = 113.[40]
[40] The vapour density of indium chloride, InCl_{3} (Note 31),
determined by Nilson and Pettersson, confirms this atomic weight.
Indium is separated from zinc and cadmium, with which it occurs,
by taking advantage of the fact that its hydroxide is insoluble in
ammonia, that the solutions of its salts give indium when treated
with zinc (hence indium is dissolved after zinc by acids) and that
they give a precipitate with hydrogen sulphide even in acid
solutions. Metallic indium is grey, has a sp. gr. of 7·42, fuses
at 176°, and does not oxidise in the air; when ignited, it first
gives a black suboxide, In_{4}O_{3}, then volatilises and gives a
brown oxide, In_{2}O_{3}, whose salts, InX_{3}, are also formed by
the direct action of acids on the metal, hydrogen being evolved.
Caustic alkalis do not act on indium, from which it is evident
that it is less capable of forming alkaline compounds than
aluminium is; however, with potassium and sodium hydroxides,
solutions of indium salts give a colourless precipitate of the
hydroxide, which is soluble in an excess of the alkali, like the
hydroxides of aluminium and zinc. Its salts do not crystallise.
Nilson and Pettersson (1889), by the action of HCl upon In,
obtained volatile crystalline, InCl_{2}, and by treating this
compound with In, InCl also.
Inasmuch as we found among the analogues of magnesium in group II. a
metal, mercury, heavier and more easily reduced than the rest, and giving
two grades of oxidation, so we should expect to find a metal among the
analogues of aluminium in group III. which would be heavy, easily
reduced, and give two grades of oxidation, and would have an atomic
weight greater than 200. Such is _thallium_. It forms compounds of a
lower type, TlX, besides the higher unstable type TlX_{3}, just as
mercury gives HgX_{2} and HgX. In the form of the thallic oxide,
Tl_{2}O_{3}, the base is but feebly energetic, as would be expected by
analogy with the oxides Al_{2}O_{3}, Ga_{2}O_{3}, and In_{2}O_{3}, whilst
in thallous oxide, Tl_{2}O, the basic properties are sharply defined, as
might be expected according to the properties of the type R_{2}O (Chapter
XV.). _Thallium_ was discovered in 1861 by Crookes and by Lamy in certain
pyrites. When pyrites are employed in the manufacture of sulphuric acid,
they are burned, and give besides sulphurous anhydride the vapours of
various substances which accompany the sulphur, and are volatile. Among
these substances arsenic and selenium are found, and together with them,
thallium. These substances accumulate in a more or less considerable
quantity in the tubes through which the vapours formed in the combustion
of the pyrites have to pass. When the methods of spectrum analysis were
discovered (1860), a great number of substances were subjected to
spectroscopic research, and it was observed that those sublimations which
are obtained in the combustion of certain pyrites contained an element
having a very sharply-defined and characteristic spectrum--namely, in the
green portion of the spectra it gave a well-defined band (wave-length 535
millionth millimetres) which did not correspond with any then known
element.[41]
[41] Thallium was afterwards found in certain micas and in the rare
mineral crookesite, containing lead, silver, thallium, and
selenium. Its isolation depends on the fact that in the presence
of acids thallium forms thallous compounds, TlX. Among these
compounds the chloride and sulphate are only slightly soluble, and
give with hydrogen sulphide a black precipitate of the sulphide
Tl_{2}S, which is soluble in an excess of acid, but insoluble in
ammonium sulphide.
Under the action of a galvanic current solutions of thallium salts
deposit the metal in the form of a heavy powder. It is of a grey colour
like tin, is soft like sodium, and has a metallic lustre. Its specific
gravity is 11·8, it melts at 290°, and volatilises at a high temperature.
When heated slightly above its melting point it forms an insoluble (in
water) higher oxide, Tl_{2}O_{3}, as a dark-coloured powder, generally
however accompanied by the lower oxide Tl_{2}O, which is also black but
soluble in water and alcohol. This solution has a distinctly alkaline
reaction. This _thallous oxide_, melts at 300°, and is easily obtained
from the hydroxide TlHO by igniting it without access of air (in the
presence of air the incandescent thallous oxide partly passes into
thallic oxide). _Thallous hydroxide_, TlOH, crystallises with one
molecule H_{2}O in yellow prisms which are very easily soluble in water.
Metallic thallium may be used for its preparation, as the metal in the
presence of water attracts oxygen from the air and forms the hydroxide.
But metallic thallium does not decompose water, although it gives a
hydroxide which is soluble in water.[41 bis] All the other data for the
chemical and physical properties of thallium, of its two grades of
oxidation and of their corresponding salts, are expressed by the position
occupied by this metal in virtue of its atomic weight Tl = 204, between
mercury Hg = 200, and lead Pb = 206.
[41 bis] The best method of preparing thallous hydroxide, TlOH, is by
the decomposition of the requisite quantity of baryta by thallous
sulphate, which is slightly soluble in water; barium sulphate is
then obtained in the precipitate and thallous hydroxide in
solution. This solubility of the hydroxide is exceedingly
characteristic, and forms one of the most important properties of
thallium. These lower (thallous) compounds are of the type TlX,
and recall the salts of the alkalis. The salts TlX are colourless,
do not give a precipitate with the alkalis or ammonia, but are
precipitated by ammonium carbonate, because thallous carbonate,
Tl_{2}CO_{3}, is sparingly soluble in water. Platinic chloride
gives the same kind of precipitate as it does with the salts of
potassium--that is, thallous platinochloride, PtTl_{2}Cl_{6}. All
these facts, together with the isomorphism of the salts TlX with
those of potassium, again point out what an important significance
the types of compounds have in the determination of the character
of a given series of substances. Although thallium has a greater
atomic weight and greater density than potassium, and although it
has a less atomic volume, nevertheless thallous oxide is analogous
to potassium oxide in many respects, for they both give compounds
of the same type, RX. We may further remark that thallous
fluoride, TlF, is easily soluble in water as well as thallous
silicofluoride, SiTl_{2}F_{6}, but that thallous cyanide, TlCN, is
sparingly soluble in water. This, together with the slight
solubility of thallous chloride, TlCl, and sulphate, Tl_{2}SO_{4},
indicates an analogy between TlX and the salts of silver, AgX.
As regards the higher oxide or the _thallic oxide_, Tl_{2}O_{3},
the thallium is trivalent in it--that is, it forms compounds of
the type TlX_{3}. The hydroxide, TlO(OH), is formed by the action
of hydrogen peroxide on thallous oxide, or by the action of
ammonia on a solution of thallic chloride, TlCl_{3}. It is
obtained as a brown precipitate, insoluble in water but easily
soluble in acids, with which it gives thallic salts, TlX_{3}.
Thallic chloride, which is obtained by cautiously heating the
metal in a stream of chlorine, forms an easily fusible white mass,
which is soluble in water and able to part with two-thirds of its
chlorine when heated. An aqueous solution of this salt yields
colourless crystals containing one equivalent of water. It is
evident from the above that all the thallic salts can easily be
reduced to thallous salts by reducing agents such as sulphurous
anhydride, zinc, &c. Besides these salts, thallic sulphate,
Tl_{2}(SO_{4})_{3},7H_{2}O, thallic nitrate
Tl(NO_{3})_{3},4H_{2}O, &c., are known. These salts are decomposed
by water, like the salts of many feeble basic metals--for example,
aluminium.
Gallium, indium, and thallium belong to the uneven series, and there
should be elements of the even series in group III. corresponding with
calcium, strontium, and barium in group II. These elements should in
their oxides R_{2}O_{3} present basic characters of a more energetic kind
than those shown by alumina, just as calcium, strontium, and barium give
more energetic bases than magnesium, zinc, and cadmium. Such are
_yttrium_ and _ytterbium_, which occur in a rare Swedish mineral called
_gadolinite_, and are therefore termed the gadolinite metals. To these
belong also the metal _lanthanum_, which accompanies the two other metals
_cerium_ and _didymium_ in the mineral _cerite_, and it therefore belongs
to the cerite metals. All these metals and certain others accompanying
them, give basic oxides R_{2}O_{3}. At first their formula was supposed
to be RO, but the application of the periodic system required their being
counted as elements of groups III. and IV., which was also confirmed by
the determination of the specific heats of these metals,[42] and better
still by the fact that Nilson and Clève, in their researches on the
gadolinite metals (1879), discovered that they contain a peculiar and
very rare element, _scandium_, which by the magnitude of its atomic
weight, Sc = 44, and in all its properties, exactly corresponds with the
metal (previously foretold on the basis of the periodic system)
_ekaboron_, whose properties were determined by taking the cerite and
gadolinite metals as forming oxides R_{2}O_{3}.[43]
[42] The specific heat of cerium determined (1870) by me, and
afterwards confirmed by Hillebrand, corresponds with that atomic
weight of cerium according to which the composition of two oxides
should be Ce_{2}O_{3} and CeO_{2}. Hillebrand also obtained
metallic lanthanum and didymium by decomposing their salts by a
galvanic current, and he found their specific heats to be near
that of cerium and about 0·04, and it is therefore justifiable to
give them an atomic weight near that of cerium, as was done on the
basis of the periodic law. Up to 1870 yttrium oxide was also given
the formula RO. Having re-determined the equivalent of yttrium
oxide (with respect to water), and found it to be 74·6, I
considered it necessary to also ascribe to it the composition
Y_{2}O_{3}, because then it falls into its proper place in the
periodic system. If the equivalent of the oxide to water be 74·6,
it contains 58·6 of metal per 16 of oxygen, and consequently one
part by weight of hydrogen replaces 29·3 of yttrium, and if it be
regarded as bivalent (oxide RO), it would not, by its atomic
weight 58·6, find a place in the second group. But if it be taken
as trivalent--that is, if the formula of its oxide be R_{2}O_{3}
and salts RX_{3}--then Y = 88, and a position is open for it in
the third group in the sixth series after rubidium and strontium.
These alterations in the atomic weights of the cerite and
gadolinite metals were afterwards accepted by Clève and other
investigators, who now ascribe a formula R_{2}O_{3} to all the
newly discovered oxides of these metals. But still the position in
the periodic system of certain elements--for example of holmium,
thulium, samarium, and others--has not yet been determined for
want of a sufficient knowledge of their properties in a state of
purity.
[43] So, for example, in 1871, in the _Journal of the Russian
Physico-Chemical Society_ (p. 45) and in Liebig's _Annalen_, Supt.
Band viii. 198, I deduced, on the basis of the periodic law, an
atomic weight 44 for ekaboron, and Nilson in 1888 found that of
scandium, which is ekaboron, to be Sc = 44·03, The periodic law
showed that the specific gravity of the ekaboron oxide would be
about 8·5, that it would have decided but feeble basic properties
and that it would give colourless salts. And this proved to be the
case with scandium oxide. In describing scandium, Clève and Nilson
acknowledge that the particular interest attached to this element
is due to its complete identity with the expected element
ekaboron. And this accurate foretelling of properties could only
be arrived at by admitting that alteration of the atomic weights
of the cerite and gadolinite metals which was one of the first
results of the application of the periodic system of the elements
to the interpretation of chemical facts. In my first memoirs,
namely, in the _Bulletin of the St. Petersburg Academy of
Sciences_, vol. viii. (1870), and in Liebig's _Annalen_ (_l. c._
p. 168) and others, I particularly insisted on the necessity of
altering the then accepted atomic weights of cerium, lanthanum,
and didymium. Clève, Höglund, Hillebrand and Norton, and more
especially Brauner, and others accepted the proposed alteration,
and gave fresh proofs in favour of the proposed alterations of
these atomic weights. The study of the fluorides was particularly
important. Placing cerium in the fourth group, the composition of
its highest oxide would then be CeO_{2}, and its compounds CeX_{4}
and the lower oxide, Ce_{2}O_{3} or CeX_{3}. Brauner obtained the
fluoride CeF_{4},H_{2}O corresponding with the first, and a double
crystalline salt, 3KF,2CeF_{4},2H_{2}O, without any admixture of
compound of the lower grade CeX_{3}, which generally occur
together with the majority of salts corresponding with CeX_{4}. It
will be seen from these formulæ and from the tables of the
elements, that cerium and didymium do not belong to the third
group, which is now being described, but we mention them here for
convenience, as all the cerite and gadolinite metals have much in
common. These metals, which are rare in nature, resemble each
other in many respects, always accompany each other, are with
difficulty isolated from each other, and stand together in the
periodic system of the elements; they have acquired a peculiar
interest owing to their having been in 1870 the objects of the
study of Marignac, Delafontaine, Soret, Lecoq de Boisbaudran,
Brauner, Clève, Nilson, the professors of Upsala, and others.
The cerite and gadolinite metals occur in rare siliceous minerals
from Sweden, America, the Urals, and Baikal, such as cerite (in
Sweden), gadolinite, and orthite; and in still rarer minerals
formed by titanic, niobic, and tantalic acids, such as euxenite in
Norway and America, and samarskite in Norway, the Urals and
America, and in a few rare fluorides and phosphates. Among the
latter, monazite is found in somewhat considerable quantities in
Brazil and North Carolina; this contains the phosphate of cerium,
CePO_{4} (= Ce_{2}O_{3}P_{2}O_{3}), together with didymium,
thorium and lanthanum (according to W. Edron and Shapleigh's
analyses), and is now used for preparing that mixture of the
oxides of the rare metals (especially ThO_{2}, Ce_{2}O_{3},
La_{2}O_{3}, &c.), which is employed for incandescent burners
(Auer von Welsbach), as it has been found by experiment that these
oxides when raised to incandescence in a non-luminous gas flame,
give a far more brilliant flame with a smaller consumption of gas,
besides being suitable for such non-luminous gases as water gas.
The insufficiency of material to work upon, and the difficulty of
separating the oxides from each other, are the chief reasons why
the composition of the compounds of these rare metals is so
imperfectly known. Cerite is the most accessible of these
minerals. Besides silica it contains more than 50 p.c. of the
oxides of cerium, lanthanum (from 4 p.c.), and didymium. The
decomposition of its powder by sulphuric acid gives sulphates, all
of which are soluble in water. The other minerals mentioned above
are also decomposed in the same manner. The solution of sulphates
is precipitated with free oxalic acid, which forms salts insoluble
in water and dilute acids with all the cerite and gadolinite
oxides. The oxides themselves are obtained by igniting the
oxalates. When ignited in the air the cerium passes from its
ordinary oxide Ce_{2}O_{3} into the higher oxide CeO_{2}, which is
so feeble a base that its salts are decomposed by water, and it is
insoluble in dilute nitric acid. Therefore it is always possible
to remove all the cerium oxide by repeated ignitions and solutions
in sulphuric acid. The further separation of the metals is mainly
based on four methods employed by many investigators.
(_a_) A solution of the mixed salts is treated with an excess of
solid potassium sulphate. Double salts, such as
Ce_{2}(SO_{4})_{3},3K_{2}SO_{4}, are thus formed. The gadolinite
metals, namely yttrium, ytterbium, and erbium, then remain in
solution--that is, their double salts are soluble in a solution of
potassium sulphate, whilst the cerite metals--namely, cerium,
lanthanum, and didymium--are precipitated, that is, their double
salts are insoluble in a saturated solution of potassium sulphate.
This ordinary method of separation, however, appears from the
researches of Marignac to be so untrustworthy that a considerable
amount of didymium and the other metals remain in the soluble
portion, owing to the fact that, although individually insoluble,
they are dissolved when mixed together. Thus erbium and terbium
occur both in the solution and precipitate. Nevertheless,
beryllium, yttrium, erbium, and ytterbium belong to the soluble,
and scandium, cerium, lanthanum, didymium, and thorium to the
insoluble portion. The insoluble salt of scandium, for example
(_i.e._ insoluble in a solution of potassium sulphate), has a
composition Sc_{2}(SO_{4})_{3},3K_{2}SO_{4}.
(_b_) The oxides obtained by the ignition of the oxalates are
dissolved in nitric acid (the nitrates of the cerite metals easily
form double salts with those of the alkali metals, and as
some--for example, the ammonio-lanthanum salt--crystallise very
well, they should be studied and applied to the analytical
separation of these metals), the solution is then evaporated to
dryness, and the residue fused. All nitrates are destroyed by
heat; those of aluminium and iron, &c., very easily, those of the
cerite and gadolinite metals also easily (although not so easily
as the above) but in different degrees and sequence; so that by
carrying on the decomposition carefully from the beginning it is
possible to destroy the nitrate of only one metal without touching
the others, or leaving them as insoluble basic salts. This method,
like the preceding and the two following, must be repeated as many
as seventy times to attain a really constant product of fixed
properties, that is, one in which the decomposed and undecomposed
portions contain one and the same oxide. This method, due to
Berlin and worked out by Bunsen, has given in the hands of
Marignac and Nilson the best results, especially for the
separation of the gadolinite metals, ytterbium and scandium.
(_c_) A solution of the salts is partially precipitated by
ammonia; that is, the solution is mixed with a small quantity of
ammonia insufficient for the precipitation of the entire quantity
of the bases (fractional precipitation). Thus, the didymium
hydroxide is first precipitated from a mixture of the salts of
didymium and lanthanum. A partial separation may be effected by
repeating the solution of the precipitate and fractional
precipitation, but a perfectly pure product is scarcely
attainable.
(_d_) The formates having different degrees of solubility
(lanthanum formate 420 parts of water per one of salt, didymium
formate 221, cerium formate 360, yttrium and erbium formates
easily soluble) give a possible means of separating certain of the
gadolinite metals from each other by a method of fractional
solution and precipitation, as Bunsen, Bahr, Clève, and others
have pointed out.
(_e_) Crookes (1893) took advantage of the fractional
precipitation of alcoholic solutions of the chlorides by amylene,
and by this means separated, for example, erbium, terbium, and
others.
(_f_) Lastly, oxide of thorium ThO_{2} (Chapter VIII., Note 59) is
separated by means of its solubility in a solution of sodium
carbonate.
A good method of separating these metals is not known, for they
are so like each other. There are also only a few _methods of
distinguishing_ them from each other, and we can only add the
following four to the above.
^a The faculty of oxidising into a higher oxide. This is very
characteristic for cerium, which gives the oxides Ce_{2}O_{3} and
CeO_{2} or Ce_{2}O_{4}. Didymium also gives one colourless oxide,
Di_{2}O_{3}, which is capable of forming salts (of a lilac
colour), and another, according to Brauner, Di_{2}O_{5} which is
dark brown and does not form salts, so far as is known, and (like
ceric oxide) acts as an oxidising agent, like the higher oxides of
tellurium, manganese, lead, and others. Lanthanum, yttrium, and
many others are not capable of such oxidation. The presence of the
higher oxides may be recognised by ignition in a stream of
hydrogen, by which means the higher oxides are reduced to the
lower, which then remain unaltered.
^b The majority of the salts of the gadolinite and cerite metals
are colourless, but those of didymium and erbium are
rose-coloured, the salts of the higher oxide of cerium, CeX_{4},
yellow, of the higher oxide of terbium, yellow, &c. Thus, the
first metals obtained from gadolinite were yttrium, giving
colourless, and erbium, giving rose-coloured, salts. Afterwards it
was found that the salts of erbium of former investigators
contained numerous colourless salts of scandium, ytterbium, &c.,
so that a coloration sometimes indicates the presence of a small
impurity, as was long known to be the case in minerals, and
therefore this point of distinction cannot be considered
trustworthy.
^c In a solid state and in solutions, the salts of didymium,
samarium, holmium, &c., give characteristic absorption spectra, as
we pointed out in Chapter XIII., and this naturally is connected
with the colour of these salts. The most important point is, that
those metals which do not give an absorption spectrum--for
example, lanthanum, yttrium, scandium, and ytterbium--may be
obtained free from didymium, samarium, and the other metals giving
absorption spectra, because the presence of the latter may be
easily recognised by means of the spectroscope, whilst the
presence of the former in the latter cannot be distinguished, and
therefore the purification of the former can be carried further
than that of the latter. We may further remark that the
sensitiveness of the spectrum reaction for didymium is so great
that it is possible with a layer of solution half a metre thick to
recognise the presence of 1 part of didymium oxide (as salt) in
40,000 parts of water. Cossa determined the presence of didymium
(together with cerium and lanthanum) in apatites, limestones,
bones, and the ashes of plants by this method. The main group of
dark lines of didymium correspond with wave-lengths of from 580 to
570 millionths mm.; and the secondary to about 520, 730, 480, &c.
The chief absorption bands of samarium are 472-486, 417, 500, and
559. Besides which, Crookes applied the investigation of the
spectra of the phosphorescent light which is emitted by certain
earths in an almost perfect vacuum, when an electric discharge is
passed through it, to the discovery and characterisation of these
rare metals. But it would seem that the smallest admixture of
other oxides (for example, bismuth, uranium) so powerfully
influences these spectra that the fundamental distinctions of the
oxides cannot be determined by this method. Besides which, the
spectra obtained by the passage of sparks through solutions or
powders of the salts are determined and applied to distinguishing
the elements, but as spectra vary with the temperature and
elasticity (concentration) this method cannot be considered as
trustworthy.
^d The most important point of distinction of individual metallic
oxides is given by the direct _determination of their equivalent
with respect to water_--that is, the amount of the oxide by weight
which combines (like water) with 80 parts by weight of sulphuric
anhydride, SO_{3}, for the formation of a normal salt. For this
purpose the oxide is weighed and dissolved in nitric acid,
sulphuric acid is then added, and the whole is evaporated to
dryness over a water-bath and then heated over a naked flame
sufficiently strongly to drive off the excess of sulphuric acid,
but so as not to decompose the salt (the product would in that
case not be perfectly soluble in water); then, knowing the weight
of the oxide and of the anhydrous sulphate, we can find the
equivalent of the oxide. The following are the most trustworthy
figures in this connection: scandium oxide 45·35 (Nilson), yttrium
oxide 75·7 (Clève; according to my determination, 1871--74·6),
cerous oxide--that is, the lower form of oxidation of cerium,
according to various investigators (Bunsen, Brauner, and others)
from 108 to 111, the higher oxide of cerium from 85 to 87,
lanthanum oxide, according to Brauner, 108, didymium oxide (in
salts of the ordinary lower form of oxidation) about 112
(Marignac, Brauner, Clève), samarium oxide about 116 (Clève),
ytterbium oxide 131·3 (Nilson). It may not be superfluous here to
draw attention to the fact that the equivalent of the oxides of
all the gadolinite and cerite metals for water distribute
themselves into four groups with a somewhat constant difference of
nearly 30. In the first group is scandium oxide with equivalent
45, in the second, yttrium oxide 76, in the third, lanthanum,
cerium, didymium, and samarium oxides with equivalent about 110,
and, in the fourth, erbium, ytterbium, and thorium oxides with
equivalent about 131. The common difference of period is nearly
45. And if we ascribe the type R_{2}O_{3} to all the oxides--that
is, if we triple the weight of the equivalent of the oxide--we
shall obtain a difference of the groups nearly equal to 90, which,
for two atoms of the metal, forms the ordinary periodic difference
of 45. If one and the same type of oxide R_{2}O_{3} be ascribed to
all these elements (as now generally accepted, in many cases there
being insufficiently trustworthy data), then the atomic weights
should be Sc = 44, Y = 89, La = 138, Ce = 140, Di = 144,
(neodymium 140, praseodymium 144), Sm = 150, Yb = 173, also
terbium 147, holmium 162, alphayttrium 157, erbium 166, thulium
170, decipium 171. It should be observed that there may be
instances of basic salts. If, for example, an element with an
atomic weight 90 gave an oxide RO_{2}, but salts ROX_{2}, then by
counting its oxide as R_{2}O_{3} its atomic weight would be 159.
All the points distinguishing many gadolinite and cerite elements
have not been sufficiently well established in certain cases (for
example, with decipium, thulium, holmium, and others). At present
the most certain are yttrium, scandium, cerium, and lanthanum. In
the case of didymium, for example, there is still much that is
doubtful. Didymium, discovered in 1842 by Mosander after
lanthanum, differs from the latter in its absorption spectrum and
the lilac-rose colour of its salts. Delafontaine (1878) separated
samarium from it. Welsbach showed that it contains two particular
elements, neodymium (salts bluish-red) and praseodymium (salts
apple-green), and Becquerel (1887) by investigating the spectra of
crystals, recognised the presence of six individual elements.
Probably, therefore, many of the now recognised elements contain a
mixture of various others, and as yet there is not enough
confirmation of their individuality. As regards yttrium, scandium,
cerium, and lanthanum, which have been established without doubt,
I think that, owing to their great rarity in nature and chemical
art, it would be superfluous to describe them further in so
elementary a work as the present. We may add that Winkler (1891)
obtained a hydrogen compound of lanthanum, whose composition
(according to Brauner) is La_{2}H_{3}, as would be expected from
the composition of Na_{2}H, Mg_{2}H_{2}, &c. C. Winkler (1891), on
reducing CeO_{2} with magnesium, also remarked a rapid absorption
of hydrogen, and showed that a _hydride of cerium_, CeH_{2},
corresponding to CaH, and the other similar hydrides of metals of
the alkaline earths, is formed (Chapter XIV., Note 63).
The brevity of this work and the great rarity of the above-mentioned
elements will give me the right to exclude their description, all the
more as the principles of the periodic system enable many of their
properties to be foreseen, and as their practical uses (cerium oxalate is
used in medicine, and didymium oxide in the manufacture of glass, a
mixture of the oxides of lanthanum and similar metals is employed for
giving a bright light, as this mixture emits a brilliant white light when
brought to incandescence) are very limited, by reason of their great
rarity in nature, and the difficulty of separating them from one another.
CHAPTER XVIII
SILICON AND THE OTHER ELEMENTS OF THE FOURTH GROUP
Carbon, which gives the compounds CH_{4}, and CO_{2}, belongs to the
fourth group of elements. The nearest element to carbon is silicon, which
forms the compounds SiH_{4} and SiO_{2}; its relation to carbon is like
that of aluminium to boron or phosphorus to nitrogen. As carbon composes
the principal and most essential part of animal and vegetable substances,
so is silicon almost an invariable component part of the rocky formations
of the earth's crust. Silicon hydride, SiH_{4}, like CH_{4}, has no acid
properties, but silica, SiO_{2}, shows feeble acid properties like
carbonic anhydride. In a free state silicon is also a non-volatile,
slightly energetic non-metal, like carbon. Therefore the form and nature
of the compounds of carbon and silicon are very similar. In addition to
this resemblance, silicon presents one exceedingly important distinction
from carbon: namely, the nature of the higher degree of oxidation. That
is, silica, silicon dioxide, or silicic anhydride, SiO_{2} is a solid,
non-volatile, and exceedingly infusible substance, very unlike carbonic
anhydride, CO_{2}, which is a gas. This expresses the essential
peculiarity of silicon. The cause of this distinction may be most
probably sought for in the polymeric composition of silica compared with
carbonic anhydride. The molecule of carbonic anhydride contains CO_{2},
as seen by the density of this gas. The molecular weight and vapour
density of silica, were it volatile, would probably correspond with the
formula SiO_{2}, but it might be imagined that it would correspond to a
far higher atomic weight of Si_{_n_}O_{2_n_}, principally from the fact
that SiH_{4} is a gas like CH_{4}, and SiCl_{4} is a liquid and volatile,
boiling at 57°--that is, even lower than CCl_{4}, which boils at 76°. In
general, analogous compounds of silicon and carbon have nearly the same
boiling points if they are liquid and volatile.[1] From this it might be
expected that silicic anhydride, SiO_{2}, would be a gas like carbonic
anhydride, whilst in reality silica is a hard non-volatile substance,[1
bis] and therefore it may with great certainty be considered that in this
condition it is polymeric with SiO_{2}, as on polymerisation--for
instance, when cyanogen passes into paracyanogen, or hydrocyanic acid
into cyanuric acid (Chapter IX.)--very frequently gaseous or volatile
substances change into solid, non-volatile, and physically denser and
more complex substances.[2] We will first make acquaintance with free
silicon and its volatile compounds, as substances in which the analogy of
silicon with carbon is shown, not only in a chemical but also in a
physical sense.[3]
[1] Chloroform, CHCl_{3}, boils at 60°, and silicon chloroform,
SiHCl_{3}, at 34°; silicon ethyl, Si(C_{2}H_{5})_{4}, boils at
about 150°, and its corresponding carbon compound,
C(C_{2}H_{5})_{4}, at about 120°; ethyl orthosilicate,
Si(OC_{2}H_{5})_{4}, boils at 160°, and ethyl orthocarbonate,
C(OC_{2}H_{5})_{4}, at 158°. The specific volumes in a liquid
state--that is, those of the silicon compounds--generally are
slightly greater than those of the carbon compounds; for example,
the volumes of CCl_{4} = 94, SiCl_{4} = 112, CHCl_{3} = 81,
SiHCl_{3} = 82, of C(OC_{2}H_{5})_{4} = 186, and
Si(OC_{2}H_{5})_{4} = 201. The corresponding salts have also nearly
equal specific volumes; for example, CaCO_{3} = 37, CaSiO_{3} = 41.
It is impossible to compare SiO_{2} and CO_{2}, because their
physical states are so widely different.
[1 bis] But silica fuses and volatilises (Moissan) in the heat of the
electric furnace, about 3000°, SiO_{2} is also partially volatile
at the temperature attained in the flame of detonating gas (Cremer,
1892).
[2] A property of intercombination is observable in the atoms of
carbon, and a faculty for intercombination, or polymerisation, is
also seen in the unsaturated hydrocarbons and carbon compounds in
general. In silicon a property of the same nature is found to be
particularly developed in silica, SiO_{2}, which is not the case
with carbonic anhydride. The faculty of the molecules of silica for
combining both with other molecules and among themselves is
exhibited in the formation of most varied compounds with bases, in
the formation of hydrates with a gradually decreasing proportion of
water down to anhydrous silica, in the colloid nature of the
hydrate (the molecules of colloids are always complex), in the
formation of polymeric ethereal salts, and in many other properties
which will be considered in the sequel. Having come to this
conclusion as to the polymeric state of silica since the years
1850-1860, I have found it to be confirmed by all subsequent
researches on the compounds of silica, and, if I mistake not, this
view has now been very generally accepted.
[3] It was only after Gerhardt, and in general subsequently to the
establishment of the true atomic weights of the elements (Chapter
VII.), that a true idea of the atomic weight of silicon and of the
composition of silica was arrived at from the fact that the
molecules of SiCl_{4}, SiF_{4}, Si(OC_{2}H_{5})_{4}, &c., never
contain less than 28 parts of silicon.
The question _of the composition of silica_ was long the subject of
the most contradictory statements in the history of science. In the
last century Pott, Bergmann, and Scheele distinguished silica from
alumina and lime. In the beginning of the present century Smithson
for the first time expressed the opinion that silica was an acid,
and the minerals of rocks salts of this acid. Berzelius determined
the presence of oxygen in silica--namely, that 8 parts of oxygen
were united with 7 of silicon. The composition of silica was first
expressed as SiO (and for the sake of shortness S only was
sometimes written instead). An investigation in the amount of
silica present in crystalline minerals showed that the amount of
oxygen in the bases bears a very varied proportion to the amount of
oxygen in the silica, and that this ratio varies from 2 : 1 to 1 :
3. The ratio 1 : 1 is also met with, but the majority of these
minerals are rare. Other more common minerals contain a larger
proportion of silica, the ratio between the oxygen of the bases and
the oxygen of the silica being equal to 1 : 2, or thereabouts; such
are the augites, labradorites, oligoclase, talc, &c. The higher
ratio 1 : 3 is known for a widely distributed series of natural
silicates--for example, the felspars. Those silicates in which the
amount of oxygen in the bases is equal to that in the silica are
termed _monosilicates_; their general formula will be
(RO)_{2}SiO_{2} or (R_{2}O_{3})_{2}(SiO_{2})_{3}. Those in which
the ratio of the oxygen is equal to 1 : 2 are termed _bisilicates_,
and their general formula will be ROSiO_{2} or
R_{2}O_{3}(SiO_{2})_{3}. Those in which the ratio is 1 : 3 will be
_trisilicates_, and their general formula (RO)_{2}(SiO_{2})_{3} or
(R_{2}O_{3})_{2}(SiO_{2})_{9}.
In these formulæ the now established composition of SiO_{2}--that
is, that in which the atom of Si = 28--is employed. Berzelius, who
made an accurate analysis of the composition of felspar, and
recognised it as a trisilicate formed by the union of potassium
oxide and alumina with silica, in just the same manner as the alums
are formed by sulphuric acid, gave silica the same formula as
sulphuric anhydride--that is, SiO_{3}. In this case the formula of
felspar would be exactly similar to that of the alums--that is,
KAl(SiO_{4})_{2}, like the alums, KAl(SO_{4})_{2}. If the
composition of silica be represented as SiO_{3}, the atom of
silicon must be recognised as equal to 42 (if O = 16; or if O = 8,
as it was before taken to be, Si = 21).
The former formulæ of silica, SiO (Si = 14) and SiO_{3} (Si = 42),
were first changed into the present one, SiO_{2} (Si = 28), on the
basis of the following arguments:--An excess of silica occurs in
nature, and in siliceous rocks free silica is generally found side
by side with the silicates, and one is therefore led to the
conclusion that it has formed acid salts. It would therefore be
incorrect to consider the trisilicates as normal salts of silica,
for they contain the largest proportion of silica; it is much
better to admit another formula with a smaller proportion of oxygen
for silica, and it then appears that the majority of minerals are
normal or slightly basic salts, whilst some of the minerals
predominating in nature contain an excess of silica--that is,
belong to the order of acid salts.
At the present time, when there is a general method (Chapter VII.)
for the determination of atomic weights, the volumes of the
volatile compounds of silica show that its atomic weight Si = 28,
and therefore silica is SiO_{2}. Thus, for example, the vapour
density of silicon chloride with respect to air is, as Dumas showed
(1862), 5·94, and hence with respect to hydrogen it is 85·5, and
consequently its molecular weight will be 171 (instead of 170 as
indicated by theory). This weight contains 28 parts of silicon and
142 parts of chlorine, and as an atom of the latter is equal to
35·5, the molecule of silicon chloride contains SiCl_{4}. As two
atoms of chlorine are equivalent to one of oxygen, the composition
of silica will be SiO_{2}--that is, the same as stannic oxide,
SnO_{2}, or titanic oxide, TiO_{2}, and the like, and also as
carbonic and sulphurous anhydrides, CO_{2} and SO_{2}. But silica
bears but little physical resemblance to the latter compounds,
whilst stannic and titanic oxides resemble silica both physically
and chemically. They are non-volatile, crystalline insoluble, are
colloids, also form feeble acids like silica, &c., and they might
therefore be expected to form analogous compounds, and be
isomorphous with silica, as Marignac (1859) found actually to be
the case. He obtained stannofluorides, for example an easily
soluble strontium salt, SrSnF_{6},2H_{2}O, corresponding with the
already long known silicofluorides, such as SrSiF_{6},2H_{2}O.
These two salts are almost identical in crystalline form
(monoclinic; angle of the prism, 83° for the former and 84° for the
latter; inclination of the axes, 103° 46´ for the latter and 103°
30´ for the former), that is, they are isomorphous. We may here add
that the specific volume of silica in a solid form is 22·6, and of
stannic oxide 21·5.
Free silicon can be obtained in an amorphous or crystalline state.
Amorphous silicon is produced, like aluminium, by decomposing the double
fluoride of sodium and silicon (sodium silicofluoride) by means of
sodium: Na_{2}SiF_{6} + 4Na = 6NaF + Si. By treating the mass thus
obtained with water the sodium fluoride may be extracted and the residue
will consist of brown, powdery silicon. In order to free it from any
silica which might be formed, it is treated with hydrofluoric acid. This
silicon powder is not lustrous; when heated it easily ignites, but does
not completely burn. It fuses when very strongly heated, and has then the
appearance of carbon.[4] Crystalline silicon is obtained in a similar
way, but by substituting an excess of aluminium for the sodium:
3Na_{2}SiF_{6} + 4Al = 6NaF + 4AlF_{3} + 3Si. The part of the aluminium
remaining in the metallic state dissolves the silicon, and the latter
separates from the solution on cooling in a crystalline form. The excess
of aluminium after the fusion is removed by means of hydrochloric and
hydrofluoric acid. The best silicon crystals are obtained from molten
zinc; 15 parts of sodium silicofluoride are mixed with 20 parts of zinc
and 4 parts of sodium, and the mixture is thrown into a strongly heated
crucible, a layer of common salt being used to cover it; when the mass
fuses it is stirred, cooled, treated with hydrochloric acid, and then
washed with nitric acid. Silicon, especially when crystalline, like
graphite and charcoal, does not in any way act on the above-mentioned
acids. It forms black, very brilliant, regular octahedra having a
specific gravity of 2·49; it is a bad conductor of electricity, and does
not burn even in pure oxygen (but it burns in gaseous fluorine). The only
acid which acts on it is a mixture of hydrofluoric and nitric acids; but
caustic alkalis dissolve in it like aluminium, with evolution of
hydrogen, thus showing its acid character. In general silicon strongly
resists the action of reagents, as do also boron and carbon. Crystalline
silicon was obtained in 1855 by Deville, and amorphous silicon in 1826 by
Berzelius.[4 bis]
[4] A similar form of silicon is obtained by fusing SiO_{2} with
magnesium, when an alloy of Si and Mg is also formed (Gattermann).
Warren (1888) by heating magnesium in a stream of SiF_{4} obtained
silicon and its alloy with magnesium. Winkler (1890) found that
Mg_{5}Si_{3} and Mg_{2}Si are formed when SiO_{2} and Mg are heated
together at lower temperatures, whilst at a high temperature Si
only is formed.
[4 bis] It is very remarkable that silicon decomposes carbonic
anhydride at a white heat, forming a white mass which, after being
treated with potassium hydroxide and hydrofluoric acid, leaves a
very stable yellow substance of the formula SiCO, which is formed
according to the equation, 3Si + 2CO_{2} = SiO_{2} + 2SiCO. It is
also slowly formed when silicon is heated with carbonic oxide. It
is not oxidised when heated in oxygen. A mixture of silicon and
carbon when heated in nitrogen gives the compound Si_{2}C_{2}N,
which is also very stable. On this basis Schützenberger recognises
a group, C_{2}Si_{2}, as capable of combining with O_{2} and N,
like C.
We may add that Troost and Hautefeuille, by heating amorphous
silicon in the vapour of SiCl_{4}, obtained crystalline silicon,
and probably at the same time lower compounds of Si and Cl were
temporarily formed. In the vapour of TiCl_{4} under the same
conditions crystalline titanium is formed (Levy, 1892).
Silicon hydride, SiH_{4}, analogous to marsh gas was obtained first of
all in an impure state, mixed with hydrogen, by two methods: by the
action of an alloy of silicon and magnesium on hydrochloric acid,[5] and
by the action of the galvanic current on dilute sulphuric acid, using
electrodes of aluminium, containing silicon. In these cases silicon
hydride is set free, together with hydrogen, and the presence of the
hydride is shown by the fact that the hydrogen separated ignites
spontaneously on coming into contact with the air, forming water and
silica. The formation of silicon hydride by the action of hydrochloric
acid on magnesium silicide is perfectly akin to the formation of
phosphuretted hydrogen by the action of hydrochloric acid on calcium
phosphide, to the formation of hydrogen sulphide by the action of acids
on many metallic sulphides, and to the formation of hydrocarbons by the
action of hydrochloric acid on white cast iron. On heating silicon
hydride--that is, on passing it through an incandescent tube, it is
decomposed into silicon and hydrogen, just like the hydrocarbons, but the
caustic alkalis, although without action on the latter, react with
silicon hydride according to the equation: SiH_{4} + 2KHO + H_{2}O =
SiK_{2}O_{3} + 4H_{2}.
[5] This alloy, as Beketoff and Cherikoff showed, is easily obtained by
directly heating finely divided silica (the experiment may be
conducted in a test tube) with magnesium powder (Chapter XIV.,
Notes 17, 18). The substance formed, when thrown into a solution of
hydrochloric acid, evolves spontaneously inflammable and impure
silicon hydride, so that the self-inflammability of the gas is
easily demonstrated by this means.
In 1850-60 Wöhler and Buff obtained an alloy of silicon and
magnesium by the action of sodium on a molten mixture of magnesium
chloride, sodium silicofluoride, and sodium chloride. The sodium
then simultaneously reduces the silicon and magnesium.
Friedel and Ladenburg subsequently prepared silicon hydride in a
pure state, and showed that it is not spontaneously inflammable in
air, at the ordinary pressure, but that, like PH_{3}, and like the
mixture prepared by the above methods, it easily takes fire in air
under a lower pressure or when mixed with hydrogen. They prepared
the pure compound in the following manner: Wöhler showed that when
dry hydrochloric acid gas is passed through a slightly heated tube
containing silicon it forms a very volatile colourless liquid,
which fumes strongly in air; this is a mixture of silicon chloride,
SiCl_{4}, and _silicon chloroform_, SiHCl_{3}, which corresponds
with ordinary chloroform, CHCl_{3}. This mixture is easily
separated by distillation, because silicon chloride boils at 57°,
and silicon chloroform at 36°. The formation of the latter will be
understood from the equation Si + 3HCl = H_{2} + SiHCl_{3}. It is
an anhydrous inflammable liquid of specific gravity 1·6. It forms a
transition product between SiH_{4} and SiCl_{4}, and may be
obtained from silicon hydride by the action of chlorine and
SbCl_{5}, and is itself also transformed into silicon chloride by
the action of chlorine. Gattermann obtained SiHCl_{3} by heating
the mass obtained after the action (Note 4) of Mg upon SiO_{2}, in
a stream of chlorine (with HCl) at about 470°. Friedel and
Ladenburg, by acting on anhydrous alcohol with silicon chloroform,
obtained an ethereal compound having the composition
SiH(OC_{2}H_{5})_{3}. This ether boils at 136°, and when acted on
by sodium disengages silicon hydride, and is converted into ethyl
orthosilicate, Si(OC_{2}H_{5})_{4}, according to the equation
4SiH(OC_{2}H_{5})_{3} = SiH_{4} + 3Si(OC_{2}H_{5})_{4} (the sodium
seems to be unchanged), which is exactly similar to the
decomposition of the lower oxides of phosphorus, with the evolution
of phosphuretted hydrogen. If we designate the group C_{2}H_{5},
contained in the silicon ethers by Et, the parallel is found to be
exact:
4PHO(OH)_{2} = PH_{3} + 3PO(OH)_{3};
4SiH(OEt)_{3} = SiH_{4} + 3Si(OEt)_{4}.
_Silicon chloride_, SiCl_{4}, is obtained from amorphous anhydrous
silica (made by igniting the hydrate) mixed with charcoal,[6] heated to a
white heat in a stream of dry chlorine--that is, by that general method
by which many other chloranhydrides having acid properties are obtained.
Silicon chloride is purified from free chlorine by distillation over
metallic mercury. Free silicon forms the same substance when treated with
dry chlorine. It is a volatile colourless liquid, which boils at 59° and
has a specific gravity of 1·52. It fumes strongly in air, has a pungent
smell, and in general has the characteristic properties of the acid
chloranhydrides. It is completely decomposed by water, forming
hydrochloric acid and silicic acid, according to the equation: SiCl_{4} +
4H_{2}O = Si(OH)_{4} + 4HCl.[7]
[6] The amorphous silica is mixed with starch, dried, and then charred
by heating the mixture in a closed crucible. A very intimate
mixture of silica and charcoal is thus formed. In Chapter XI., Note
13, we saw that elements like silicon disengage more heat with
oxygen than with chlorine, and therefore their oxygen compounds
cannot be directly decomposed by chlorine, but that this can be
effected when the affinity of carbon for oxygen is utilised to aid
the action. When the mass obtained by the action of Mg upon SiO_{2}
is heated to 300° in a current of chlorine, it easily forms
SiCl_{4} (Gattermann): besides which two other compounds,
corresponding to SiCl_{4}, are formed, namely: Si_{2}Cl_{6}, which
boils at 145° and solidifies at -1°, and Si_{3}Cl_{8}, which boils
at about 212°. These substances, which answer to corresponding
carbon compounds (C_{2}H_{6} and C_{3}H_{8}), act upon water and
form corresponding oxygen compounds; for instance, Si_{2}Cl_{6} +
4H_{2}O = (SiO_{2}H)_{2} + 6HCl gives the analogue of oxalic acid
(CO_{2}H)_{2}. This substance is insoluble in water, decomposes
under the action of friction and heat with an explosion, and should
be called _silico-oxalic acid_, Si_{2}H_{2}O_{4} (_see_ later, Note
11 ^{bis}).
[7] Silicon chloride shows a similar behaviour with alcohol. This is
accompanied by a very characteristic phenomenon; on pouring silicon
chloride into anhydrous alcohol a momentary evolution of heat is
observed, owing to a reaction of double decomposition, but this is
immediately followed by a powerful cooling effect, due to the
disengagement of a large amount of hydrochloric acid--that is,
there is an absorption of heat from the formation of gaseous
hydrochloric acid. This is a very instructive example in this
respect; here two processes occurring simultaneously--one chemical
and the other physical--are divided from each other by time, the
latter process showing itself by a distinct fall in temperature. In
the majority of cases the two processes proceed simultaneously, and
we only observe the difference between the heat developed and
absorbed. In acting on alcohol, silicon chloride forms ethyl
orthosilicate, SiCl_{4} + 4HOC_{2}H_{5} = 4HCl +
Si(OC_{2}H_{5})_{4}. This substance boils at 160°, and has a
specific gravity 0·94. Another salt, ethyl metasilicate,
SiO(OC_{2}H_{5})_{2}, is also formed by the action of silicon
chloride on anhydrous alcohol; it volatilises above 300°, having a
sp. gr. 1·08. It is exceedingly interesting that these two ethereal
salts are both volatile, and both correspond with silica, SiO_{2}:
the first ether corresponds to the hydrate Si(OH)_{4}, orthosilic
acid, and the second to the hydrate SiO(OH)_{2}, metasilicic acid.
As the nature of hydrates may be judged from the composition of
salts, so also, with equal right, can ethereal salts serve the same
purpose. The composition of an ethereal salt corresponds with that
of an acid in which the hydrogen is replaced by a hydrocarbon
radicle--for instance, by C_{2}H_{5}. And, therefore, it may be
truly said that there exist at least the two silicic acids above
mentioned. We shall afterwards see that there are really several
such hydrates; that these ethereal salts actually correspond with
hydrates of silica is clearly shown from the fact that they are
decomposed by water, and that in moist air they give alcohol and
the corresponding hydrate, although the hydrate which is obtained
in the residue always corresponds with the second ethereal salt
only--that is, it has the composition SiO(OH)_{2}; this form
corresponds also to carbonic acid in its ordinary salts. This
hydrate is formed as a vitreous mass when the ethyl silicates are
exposed to air, owing to the action of the atmospheric moisture on
them. Its specific gravity is 1·77.
_Silicon bromide_, SiBr_{4}, as well as silicon bromoform,
SiHBr_{3}, are substances closely resembling the chlorine compounds
in their reactions, and they are obtained in the same manner.
Silicon iodoform, SiHI_{3}, boils at about 220°, has a specific
gravity of 3·4, reacts in the same manner as silicon chloroform,
and is formed, together with silicon iodide, SiI_{4}, by the action
of a mixture of hydrogen and hydriodic acid on heated silicon.
Silicon iodide is a solid at the ordinary temperature, fusing at
about 120°; it may be distilled in a stream of carbonic anhydride,
but easily takes fire in air, and behaves with water and other
reagents just like silicon chloride. It may be obtained by the
direct action of the vapour of iodine on heated silicon. Besson
(1891) also obtained SiCl_{3}I (boils at 113°), SiCl_{2}I_{2}
(172°), and SiClI_{3} (220°), and the corresponding bromine
compounds. All the halogen compounds of Si are capable of absorbing
6NH_{3} and more. Besides which Besson obtained SiSCl_{2} by
heating Si in the vapour of chloride of sulphur; this compound
melts at 74°, boils at 185°, and gives with water the hydrate of
SiO_{2}, HCl, and H_{2}S.
The most remarkable of the haloid compounds of silicon is _silicon
fluoride_, SiF_{4}. It is a gaseous substance only liquefied by intense
cold, -100°, and is obtained (Chapter XI.) directly by the action of
hydrofluoric acid on silica and its compounds (SiO_{2} + 4HF = 2H_{2}O +
SiF_{4}), and also by heating fluorspar with silica (2CaF_{2} + 3SiO_{2}
= 2CaSiO_{3} + SiF_{4}).[8] In order to prepare silicon fluoride, sand or
broken glass is mixed with an equal quantity by weight of fluorspar and 6
parts by weight of strong sulphuric acid, and the mixture is gently
heated. It fumes strongly in air, reacting with the aqueous vapours,
although it is produced from silica and hydrofluoric acid with the
separation of water. It is evident that a reverse reaction occurs here;
that is to say, the water reacts with the silicon fluoride, but the
reaction is not complete. This phenomenon is similar to that which occurs
when water decomposes aluminium chloride, but at the same time
hydrochloric acid dissolves aluminium hydroxide and forms the same
aluminium chloride. The relative amount of water present (together with
the temperature) determines the limit and direction of the reaction. The
faculty which silicon fluoride has of reacting with water is so great
that it takes up the elements of water from many substances--for
instance, like sulphuric acid, it chars paper. Water dissolves about 300
volumes of this gas, but in this case it is not a common dissolution
which takes place, but a reaction. During the first absorption of silicon
fluoride by water, silicic acid is separated in the form of a jelly, but
a certain quantity of the silicon fluoride also remains in the liquid,
because the hydrofluoric acid formed dissolves the other part of the
silica[9] and forms the so-called _hydrofluosilicic acid_: H_{2}SiF_{6} =
SiF_{4} + 2HF = SiH_{2}O_{3} + 6HF - 3H_{2}O. That is to say, a
metasilicic acid, SiH_{2}O_{3}, in which O_{3} is replaced by F_{6}. This
view of the composition of hydrofluosilicic acid may be admitted, because
it forms a whole series of crystallisable and well defined salts. In
general, the whole reaction of water on silicon fluoride may be expressed
by the equation: 3SiF_{4} + 3H_{2}O = SiO(OH)_{2} + 2SiH_{2}F_{6}.
Hydrofluosilicic acid and silicic acid resemble each other as much, and
differ as much, in their chemical character as water and hydrofluoric
acid. For this reason silicic acid is a feebler acid than
hydrofluosilicic acid, and in addition to this the former is insoluble,
and the latter soluble, in water.[10] Hydrofluosilicic acid is also
formed if silicic acid be dissolved in a solution of hydrofluoric acid.
It is incapable of volatilising without decomposition, and on heating the
concentrated acid silicon fluoride is evolved, leaving an aqueous
solution of hydrofluoric acid. This is the reason why solutions of
hydrofluosilicic acid corrode glass. This decomposition may be further
accelerated by the addition of sulphuric acid, or even of other acids.
Hydrofluosilicic acid, when acting on potassium and barium salts, gives
precipitates, because the salts of these metals are but sparingly soluble
in water: thus 2KX + H_{2}SiF_{6} = 2HX + K_{2}SiF_{6}. The potassium
salt is obtained in the form of very fine octahedra, but the precipitate
does not form quickly, and at first appears as a jelly. Nevertheless, the
decomposition is complete, and it is taken advantage of for obtaining
their corresponding acids from salts of potassium.[10 bis]
[8] This property of calcium fluoride of converting silica into a gas
and a vitreous fusible slag of calcium silicate is frequently taken
advantage of in the laboratory and in practice in order to remove
silica. The same reaction is employed for preparing silicon
fluoride on a large scale in the manufacture of hydrofluosilicic
acid (see sequel).
[9] The amount of heat developed by the solution of silicic acid,
SiO_{2}_n_H_{2}O, in aqueous hydrofluoric acid, _x_HF_n_H_{2}O,
increases with the magnitude of _x_ and normally equals _x_5,600
heat units, where _x_ varies between 1 and 8. However, when _x_ =
10 the maximum amount of heat is developed (= 49,500 units), and
beyond that the amount decreases (Thomsen).
[10] In reality, however, it would seem that the reaction is still more
complex, because the aqueous solution of silicon fluoride does not
yield a hydrate of silica, but a fluo-hydrate (Schiff),
Si_{2}O_{3}(OH)F, corresponding to the (pyro) hydrate
Si_{2}O_{3}(OH)_{2}, equal to SiO(OH)_{2}SiO_{2}, so that the
reaction of silicon fluoride on water is expressed by the
equation: 5SiF_{4} + 4H_{2}O = 3SiH_{2}F_{6} + Si_{2}O_{3}(OH)F +
HF. However, Berzelius states that the hydrate, when well washed
with water, contains no fluorine, which is probably due to the
fact that an excess of water decomposes Si_{2}O_{3}(OH)F, forming
hydrofluoric acid and the compound Si_{2}O_{3}(OH)_{2}. Water
saturated with silicon fluoride disengages silicon fluoride and
hydrofluoric acid when treated with hydrochloric acid, the
gelatinous precipitate being simultaneously dissolved. It may be
further remarked that hydrofluosilicic acid has been frequently
regarded as SiO_{2},6HF, because it is formed by the solution of
silica in hydrofluoric acid, but only two of these six hydrogens
are replaced by metals. On concentration, solutions of the acid
begin to decompose when they reach a strength of 6H_{2}O per
H_{2}SiF_{6}, and therefore the acid may be regarded as
Si(OH)_{4},2H_{2}O,6HF, but the corresponding salts contain less
water, and there are even anhydrous salts, R_{2}SiF_{6}, so that
the acid itself is most simply represented as H_{2}SiF_{6}.
If gaseous silicon fluoride be passed directly into water, the
gas-conducting tube becomes clogged with the precipitated silicic
acid. This is best prevented by immersing the end of the tube
under mercury, and then pouring water over the mercury; the
silicon fluoride then passes through the mercury, and only comes
into contact with the water at its surface, and consequently the
gas-conducting tube remains unobstructed. The silicic acid thus
obtained soon settles, and a colourless solution with a pleasant
but distinctly acid taste is procured.
Mackintosh, by taking 9 p.c. of hydrofluoric acid, observed that
in the course of an hour its action on opal attained 77 p.c. of
the possible, and did not exceed 1-1/2 p.c. of its possible action
on quartz during the same time. This shows the difference of the
structure of these two modifications of silica, which will be more
fully described in the sequel.
[10 bis] The sodium salt is far more soluble in water, and crystallises
in the hexagonal system. The magnesium salt, MgSiF_{6}, and
calcium salt are soluble in water. The salts of hydrofluosilicic
acid may be obtained not only by the action of the acid on bases
or by double decompositions, but also by the action of
hydrofluoric acid on metallic silicates. Sulphuric acid decomposes
them, with evolution of hydrofluoric acid and silicon fluoride,
and the salts when heated evolve silicon fluoride, leaving a
residue of metallic fluoride, R_{2}F_{2}.
Silicon, having so much in common with carbon, is also able to combine
with it in the proportion given by the law of substitution, that is, it
forms a carbide of silicon CSi, called _carborundum_ and obtained by
Mühlhäuser and Acheson in the United States, and by Moissan in France
(1891), and others, by reducing silica with carbon in the electrical
furnace at a temperature of about 2500°[11], _i.e._ by the action of an
electrical current upon a mixture of carbon and SiO_{2} with NaCl. After
treating the resultant mass with acids and washing with water,
carborundum is obtained in transparent, lustrous grains of a greenish
color, possessing great hardness (greater than corundum) and therefore
used for polishing the hardest kinds of steel and stones. The specific
gravity is about 3·1. Carborundum does not alter at a red heat, does not
burn, and apparently approaches the diamond in its properties. (Moissan
obtained, 1894, a similar very hard compound for boron, B_{6}C, sp. gr.
2·5.)
[11] _See_ Note 4 bis. Probably Schützenberger had already obtained CSi
in his researches together with other silicon compounds. An
amorphous, less hard compound of the same alloy is also obtained
together with the hard crystalline CSi.
According to the principle of substitution, if silicon forms SiH_{4}, a
series of hydrates, or hydroxyl derivatives, ought to exist corresponding
to it. The first hydrate of an alcoholic character ought to have the
composition SiH_{3}(OH); the second hydrate SiH_{2}(OH)_{2}; the third,
SiH(OH)_{3};[11 bis] and the last, Si(OH)_{4}. The last is a hydrate of
silica, because it is equal to SiO_{2} + 2H_{2}O); and it is formed by
the action of water on silicon chloride, when all four atoms of chlorine
are replaced by four hydroxyl groups. It does not, however, remain in
this state, but easily loses part of its water.
[11 bis] The following consideration is very important in explaining
the nature of the lower hydrates which are known for silicon. If
we suppose water to be taken up from the first hydrates (just as
formic acid is CH(OH)_{3}, _minus_ water), we shall obtain the
various lower hydrates corresponding with silicon hydride. When
ignited they should, like phosphorous and hypophosphorous acids,
disengage silicon hydride, and leave a residue of silica
behind--_i.e._ of the oxide corresponding to the highest
hydrate--just as organic hydrates (for example, formic acid with
an alkali) form carbonic anhydride as the highest oxygen compound.
Such imperfect hydrates of silicon, or, more correctly speaking,
of silicon hydride, were first obtained by Wöhler (1863) and
studied by Geuther (1865), and were named after their
characteristic colours. (_See_ Note 6).
_Leucone_ is a white hydrate of the composition SiH(OH)_{3}. It is
obtained by slowly passing the vapour of silicon chloroform into
cold water: SiHCl_{3} + 3H_{2}O = SiH(OH)_{3} + 3HCl. But this
hydrate, like the corresponding hydrate of phosphorus or carbon,
does not remain in this state of hydration, but loses a portion of
its water. The carbon hydrate of this nature, CH(OH)_{3}, loses
water and forms formic acid, CHO(OH); but the silicon hydrate
loses a still greater proportion of water, 2SiH(OH)_{3}, parting
with 3H_{2}O, and consequently leaving Si_{2}H_{2}O_{3}. This
substance must be an anhydride; all the hydrogen previously in the
form of hydroxyl has been disengaged, two remaining hydrogens
being left from SiH_{4}. The other similar hydrate is also white,
and has the composition Si_{3}H_{2}O (nearly). It may be regarded
as the above white hydrate + SiO_{2}. A yellow hydrate, known as
_chryseone_ (silicone), is obtained by the action of hydrochloric
acid on an alloy of silicon and calcium; its composition is about
Si_{6}H_{4}O_{3}. Most probably, however, chryseone has a more
complex composition, and stands in the same relation to the
hydrate SiH_{2}(OH)_{3} as leucone does to the hydrate
SiH(OH)_{3}, because this very simply expresses the transition of
the first compound into the second with the loss of water,
SiH_{2}(OH)_{3} - H_{2} + H_{2}O = SiH(OH)_{3}. When these lower
hydrates are ignited without access of air, they are decomposed
into hydrogen, silicon, and silica--that is, it may be supposed
that they form silicon hydride (which decomposes into silicon and
hydrogen) and silica (just as phosphorous and hypophosphorous
acids give phosphoric acid and phosphuretted hydrogen). When
ignited in air, they burn, forming silica. They are none of them
acted on by acids, but when treated with alkalis they evolve
hydrogen and give silicates; for example, leucone: SiH_{2}O_{3} +
4KHO = 2SiK_{2}O_{3} + H_{2}O + 2H_{2}. They have no acid
properties.
Silica or silicic anhydride, both in the free state and in combination
with other oxides, enters into the composition of most of the rocky
formations of the earth's crust. These silicious compounds are substances
varying so much in their properties, crystalline forms, and relations to
one another that they are comprised in a special branch of natural
science (like the carbon compounds), and are treated of in works on
mineralogy; so that, in dealing with them further, we shall only give a
short description of these various compounds. It is first of all
necessary to turn to the description of silica itself, especially as it
is not unfrequently met with in nature in a separate state, and often
forms whole masses of rocky formations, called 'quartz.' In an anhydrous
condition silica appears in the greatest variety of natural
forms--sometimes in well-formed crystals, hexagonal prisms, terminated by
hexagonal pyramids. If the crystals are colourless and transparent, they
are called _rock crystal_. This is the purest form of silica. Prismatic
crystals of rock crystal sometimes attain considerable size, and as they
are remarkable for their unchangeability, great hardness, and high index
of refraction, they are used for ornaments, for seals, making necklaces,
&c.[12] Rock crystal coloured with organic matter in contact with which
it has been produced has a brown or greyish colour, and then bears the
name of _cairngorm_ or _smoky quartz_. In this form it has the same uses
as rock crystal, especially as it is often found in large masses. The
same mineral, frequently occurs, coloured red or pink by manganese or
iron oxides, especially in aqueous formations, and is then known as
_amethyst_. When finely coloured the amethyst is used as a precious
stone, but amethysts most frequently occur as small crystals in the
cavities formed in other rocky formations, and especially in those formed
in silica itself. A similar anhydrous silica is often found in
transparent non-crystalline masses, having the same specific gravity as
rock crystal itself (2·66). In this case it is called _quartz_. Sometimes
it forms complete rocky formations, but more often penetrates or is
interspersed through other rocky formations, together with other
siliceous compounds. Thus, in granite, quartz is mixed with felspar and
similar substances. Sometimes the colouring of quartz is so considerable
that it is hardly transparent in thin sheets, but it is often found in
transparent masses slightly coloured with various tints. The existence in
nature of enormous masses of quartz proves that it resists the action of
water. When water destroys rocky formations, the siliceous minerals which
they contain are partly dissolved and partly transformed into clay, &c.
But the quartz remains untouched, in the form of grains in which it
existed in the rocky formation; sometimes, when crushed, it is carried
away by the water and deposited. This is the nature of _sand_. Naturally,
sometimes other rocky substances which are not changed by water, or only
slightly acted on by it, are found in sand; but as these latter are more
or less changed by the continuous action of water, it is not unusual to
find sand which consists almost entirely of pure quartz. Common sand is
generally coloured yellow or reddish-brown by foreign mineral matter,
consisting principally of ferruginous minerals and clays. The purest or
so-called quartz sand is, however, rarely found, and is recognised by the
absence of colour, and also by the test that when shaken in water it does
not form any turbidity: this shows the absence of clay; when fused with
bases it forms a colourless glass, and on this account is a valuable
material for the manufacture of glass. Sands were formed at all periods
of the earth's existence; the ancient ones, compressed by strata of more
recent formation and permeated with various substances (deposited from
the infiltrating water), are sometimes solidified into rock, called
_sandstone_, composing, in some places, whole mountain chains, and
serviceable as a most excellent building material, on account of the
slight change it undergoes under the influence of atmospheric agencies,
and on account of the facility with which it may be wrought from rocky
formations into immense regularly-shaped flags--the latter property is
due to the primary laminar structure of the sand formations deposited, as
above-mentioned, by water. Many grindstones and whetstones are made from
such rocks.
[12] Two modifications of rock crystal are known. They are very easily
distinguished from each other by their relation to polarised
light; one rotates the plane of polarisation to the right and the
other to the left--in the one the hemihedral faces are right and
in the other they are left; this opposite rotatory power is taken
advantage of in the construction of polarisers. But, with this
physical difference--which is naturally dependent on a certain
difference in the distribution of the molecules--there is not only
no observable difference in the chemical properties, but not even
in the density of the mass. Perfectly pure rock crystal is a
substance which is most invariable with respect to its specific
gravity. The numerous and accurate determinations made by
Steinheil on the specific gravity of rock crystal show that (if
the crystal be free from flaws) it is very constant and is equal
to 2·66.
Perfectly pure anhydrous silica is not only known in the condition of
rock crystal and quartz having a specific gravity of 2·6, but also in
another special form, having other chemical and physical properties. This
variety of silica has a specific gravity of 2·2, and is formed by fusing
rock crystal or heating silicic acid.[12 bis] Silicic acid, when heated
to a dull red heat, parts entirely with the water it contains, and leaves
an exceedingly fine amorphous mass of silica (easily levigated, but
difficult to moisten); it is characterised by such excessive friability
that, when lightly blown on, a large mass of it rises into the air like a
cloud of dust. A mass of anhydrous silica maybe poured in this way from
one vessel to another like a liquid, and like the latter it takes a
horizontal position in the vessel containing it.[13] Anhydrous silica,
like quartz, does not fuse in the heat of a furnace, but it fuses in the
oxyhydrogen flame to a colourless glassy mass exactly similar to that
formed in the same way from rock crystal. In this condition silica has a
specific gravity of 2·2.[13 bis] Both forms of silica are insoluble in
ordinary acids, and even when they are in the state of powder, alkalis in
solution act very slowly and feebly on them; rock crystal offers much
greater resistance to the action of alkalis than the powder obtained by
heating the hydrate. The latter is quite soluble, although but slowly, in
hot alkaline solutions. This last property appertains in a greater degree
to anhydrous silica having a specific gravity of 2·2 than to that which
has a specific gravity of 2·6. Hydrofluoric acid more easily transforms
the former into silicon fluoride than it does the latter. Both varieties
of silica, when taken in the form of powder, easily combine with bases,
forming, on being fused with an alkali, a vitreous slag, which is a salt
corresponding with silica. Glass is such a salt, formed of alkalis and
alkaline earthy bases; if the glass does not contain any of the
latter--that is, if only alkaline glass be taken--a mass soluble in water
is obtained. In order to obtain such _soluble glass_, potassium or sodium
carbonates, or better a mixture of the two (fusion mixture), is fused
with fine sand. A still better and further saturation of the alkalis with
silica is effected by the action of alkaline solutions on the silicon
hydrate met with in nature; for instance, an alkaline solution is often
made use of to act on the so-called _tripoli_, or collection of siliceous
skeletons of the lowest microscopical infusoria, which is sometimes found
in considerable layers in the form of a sandy mass. Tripoli is used for
polishing, not only on account of the considerable hardness of the
silica, but also because the microscopic bodies of the infusoria have a
pointed shape, which, however, is not angular, so that they do not
scratch metals like sand.[14] The alkaline solutions of silica obtained
by boiling tripoli with caustic soda under pressure contain various
proportions of silica and alkali.[14 bis] In order that it may contain
the greatest amount of silica, silicic acid should be added to the heated
solution. Silicic acid is formed by taking any solution containing silica
and alkali, and adding to it, by degrees, some acid--for instance,
sulphuric or hydrochloric; if the experiment be carried on carefully and
the solution be concentrated, the whole mass thickens to a jelly, due to
the gelatinous form of the _silicic acid_ separated from the salt by the
action of the acid. The decomposition may be expressed by the following
equation: Si(ONa)_{4} + 4HCl = 4NaCl + Si(OH)_{4}. The hydrate separated,
Si(OH)_{4}, easily loses part of the water and forms a jelly, the whole
mass gelatinising if the solution be strong enough.[15]
[12 bis] Several other modifications are known as minute crystals. For
example, there is a particular mineral first found in Styria and
known as _tridymite_. Its specific gravity 2·3 and form of
crystals clearly distinguish it from rock crystal; its hardness is
the same as that of quartz--that is, slightly below that of the
ruby and diamond.
[13] There is a distinct rise of temperature (about 4°) when amorphous
silica is moistened with water. Benzene and amyl alcohol also give
an observable rise of temperature. Charcoal and sand give the same
result, although to a less extent.
[13 bis] Silica also occurs in nature in two modifications. The opal
and tripoli (infusorial earth) have a specific gravity of about
2·2, and are comparatively easily soluble in alkalis and
hydrofluoric acid. Chalcedony and flint (tinted quartzose
concretions of aqueous origin), agate and similar forms of silica
of undoubted aqueous origin, although still containing a certain
amount of water, have a specific gravity of 2·6, and correspond
with quartz in the difficulty with which they dissolve. This form
of silica sometimes permeates the cellulose of wood, forming one
of the ordinary kinds of petrified wood. The silica may be
extracted from it by the action of hydrofluoric acid, and the
cellulose remains behind, which clearly shows that silica in a
soluble form (see sequel) has permeated into the cells, where it
has deposited the hydrate, which has lost water, and given a
silica of sp. gr. 2·6. The quartzose stalactites found in certain
caves are also evidently of a similar aqueous origin; their sp.
gr. is also 2·6. As crystals of amethyst are frequently found
among chalcedonies, and as Friedau and Sarrau (1879) obtained
crystals of rock crystal by heating soluble glass with an excess
of hydrate of silica in a closed vessel, there is no doubt but
that rock crystal itself is formed in the wet way from the
gelatinous hydrate. Chroustchoff obtained it directly from soluble
silica. Thus this hydrate is able to form not only the variety
having the specific gravity 2·2 but also the more stable variety
of sp. gr. 2·6; and both exist with a small proportion of water
and in a perfectly anhydrous state in an amorphous and crystalline
form. All these facts are expressed by recognising silica as
dimorphous, and their cause must be looked for in a difference in
the degree of polymerisation.
[14] Deposits of perfectly white tripoli have been discovered near
Batoum, and might prove of some commercial importance.
[14 bis] Alkaline solutions, saturated with silica and known as _soluble
glass_, are prepared on a large scale for technical purposes by
the action of potassium (or sodium) hydroxide in a steam boiler on
tripoli or infusorial earth, which contains a large proportion of
amorphous silica. All solutions of the alkaline silicates have an
alkaline reaction, and are even decomposed by carbonic acid. They
are chiefly used by the dyer, for the same purposes as sodium
aluminate, and also for giving a hardness and polish to stucco and
other cements, and in general to substances which contain lime. A
lump of chalk when immersed in soluble glass, or better still when
moistened with a solution and afterwards washed in water (or
better in hydrofluosilicic acid, in order to bind together the
free alkali and make it insoluble), becomes exceedingly hard,
loses its friability, is rendered cohesive, and cannot be
levigated in water. This transformation is due to the fact that
the hydrate of silica present in the solution acts upon the lime,
forming a stony mass of calcium silicate, whilst the carbonic acid
previously in combination with the lime enters into combination
with the alkali and is washed away by the water.
[15] The equation given above does not express the actual reaction, for
in the first place silica has the faculty of forming compounds
with bases, and therefore the formula SiNa_{4}O_{4} is not rightly
deduced, if one may so express oneself. And, in the second place,
silica gives several hydrates. In consequence of this, the hydrate
precipitated does not actually contain so high a proportion of
water as Si(OH)_{4}, but always less. The insoluble gelatinous
hydrate which separates out is able (before, but not after, having
been dried) to dissolve in a solution of sodium carbonate. When
dried in air its composition corresponds with the ordinary salts
of carbonic acid--that is, SiH_{2}O_{3}, or SiO(OH)_{2}. If
gradually heated it loses water by degrees, and, in so doing,
gives various degrees of combination with it. The existence of
these degrees of hydration, having the composition
SiH_{2}O_{3}_n_SiO_{2}, or, in general, _n_SiO_{2}_m_H_{2}O, where
_m_ < _n_, must be recognised, because most varied degrees of
combination of silica with bases are known. The hydrate of silica,
when not dried above 30°, has a composition of nearly
H_{4}Si_{3}O_{8} = (H_{2}SiO_{3})_{2}SiO_{2}, but at 60° contains
a greater proportion of silica--that is, it loses still more
water; and at 100° a hydrate of the composition
SiH_{2}O_{3}2SiO_{2}, and at 250° a hydrate having approximately a
composition SiH_{2}O_{3}7SiO_{2} is obtained.
These data show the complexity of the molecules of anhydrous
silica. The hydrates of silica easily lose water and give the
hydrates (SiO_{2})_{_n_}(H_{2}O)_{_m_}, where _m_ becomes smaller
and smaller than _n_. In the natural hydrates, this decrement of
water proceeds quite consecutively, and, so to say, imperceptibly,
until _n_ becomes incomparably greater than _m_, and when the
ratio becomes very large, anhydrous silica of the two
modifications 2·6 and 2·2 is obtained. The composition
(SiO_{2})_{10},H_{2}O still corresponds with 2·9 p.c. of water,
and natural hydrates often contain still less water than this.
Thus some opals are known which contain only 1 p.c. of water,
whilst others contain 7 and even 10 p.c. As the artificially
prepared gelatinous hydrate of silica when dried has many of the
properties of native opals, and as this hydrate always loses water
easily and continually, there can be no doubt that the transition
of (SiO_{2})_{_n_}(H_{2}O)_{_m_} into anhydrous silica, both
amorphous and crystalline (in nature, chalcedony), is accomplished
gradually. This can only be the case if the magnitude of _n_ be
considerable, and therefore the molecule of silica in the hydrate
is undoubtedly complex, and hence the anhydrous silica of sp. gr.
2·2 and 2·6 does not contain SiO_{2}, but a complex molecule,
Si_{_n_}O_{2_n_}--that is, the structure of silica is polymeric
and complex, and not simple as represented above by the formula
SiO_{2}.
Neither of the two varieties of anhydrous silica, nor the various
natural gelatinous hydrates, are directly soluble in water. There is,
however, a condition of silica known which is soluble in water, _soluble
silica_, and silica is found in this state in nature. Small quantities of
soluble silica are met with in all waters. Certain mineral springs, and
especially hot springs--of which the best known are the Geysers of
Iceland and those in the North American National Park (Yellowstone
Valley)--contain a considerable amount of silica in solution. Such water,
permeating the objects it meets with--for instance, wood--penetrates into
them and deposits silica inside them, that is, transforms them into a
petrified condition. Siliceous stalactites, and also many (if not all)
forms of silica are formed by such water. The absorption of silica by
plants by means of their roots, and also by the lower organisms having
siliceous bodies, is due also to their nourishing themselves with the
solutions containing silica continually formed in nature. Thus, in
plants, in the straws of the grasses, in hard shave-grass, and especially
in the knots of bamboo and other straw-like plants, a considerable
quantity of silica is deposited, which must previously have been absorbed
by the plants.
Silicic acid is a colloid. The gelatinous silicon hydrate is its
hydrogel, the soluble hydrate is the hydrosol (Chapter XII.) Both
varieties may be easily obtained from the alkaline silicates and from
water-glass. The very same substances--that is, aqueous solutions of
soluble glass and acid--taken in the same proportion, may produce either
the gelatinous or the soluble silica, according to the way these
solutions are mixed together. If the acid be added little by little to
the _alkaline silicate_, with continuous stirring, a moment arrives when
the whole mass thickens to a jelly, hydrogel; in this case the silicic
acid is formed in the midst of the alkaline solution and becomes
insoluble. But if the mixing be done in the reverse order--that is, if
the soluble glass be added to the acid, or if a quantity of acid be
rapidly poured into the solution of the salt--then the separation of the
silica takes place in the midst of the acid liquid, and it is obtained in
the form of the soluble hydrate, the hydrosol.[16]
[16] The presence of an excess of acid aids the retention of the silica
in the solution, because the gelatinous silica obtained in the
above manner, but not heated to 60°--that is, containing more
water than the hydrate H_{2}SiO_{3}--is more soluble in water
containing acid than in pure water. This would seem to indicate a
feeble tendency of silica to combine with acids, and it might even
have been imagined that in such a solution the hydrate of silica
is held in combination by an excess of acid, had Graham not
obtained soluble silica perfectly free from acid, and if there
were not solutions of silica free from any acid in nature. At all
events a tolerably strong solution of free silica or silicic acid
may be obtained from soluble glass diluted with water. The
solution, besides silica, will contain sodium chloride and an
excess of the acid taken. If this solution remains for some time
exposed to the air, or in a closed vessel, and under various other
conditions, it is found that, after a time, insoluble gelatinous
silica separates out--that is, the soluble form of silica is
unstable, like the soluble form of alumina. The analogous forms of
molybdic or tungstic acids may be heated, evaporated, and kept for
a long period of time without the soluble form being converted
into the insoluble.
The hydrosol of silica prepared by mixing an excess of hydrochloric acid
with a solution of sodium silicate, may be freed from the admixtures both
of hydrochloric acid and salt, sodium chloride, _by means of
dialysis_,[17] as Graham showed (in 1861) in enquiring into the nature of
colloids (Chapter I.), and making many other important chemical
investigations. The solution, containing the acid, salt, and silica, all
dissolved in water, is poured into a dialyser--that is, a vessel with a
porous diaphragm surrounded by water. Certain substances pass more easily
through the diaphragm than others. This may be represented thus: the
passage through the diaphragm proceeds in both directions, and if the
solutions on each side of the diaphragm be equally strong, there will be
equal numbers of molecules of the soluble substance passing into either
side in a given time, some passing quickly and others slowly. The
metallic chlorides and hydrochloric acid belong to the series of
crystalloids which easily pass through a diaphragm, and therefore the
hydrochloric acid and sodium chloride contained in the above-mentioned
dialyser pass from the solution through the diaphragm into the water of
the external vessel with considerable rapidity. The aqueous solution of
colloidal silica also penetrates through the diaphragm, but very much
more slowly. But if the amount of the substance dissolved is not equal on
either side of the diaphragm, the whole system strives to attain a state
of equilibrium; that is, the given substance penetrates through the
diaphragm from the side where it is in excess to the part where there is
a smaller quantity of it. All substances which are soluble in water have
the faculty of penetrating through a membrane swollen in water, but the
velocity of penetration is not equal, and in this respect the dialyser
separates substances like a sieve. The silica passes less rapidly through
the diaphragm than the sodium chloride and hydrochloric acid, so that by
repeatedly changing the external water it is easy to effect the
extraction of the chlorine compounds from the dialyser, which will
finally only contain a solution of silica. This extraction (of HCl and
NaCl) may be so complete that the liquid taken from the dialyser will not
give any precipitate with a solution of silver nitrate. Graham obtained
in this way soluble silica having a distinctly acid reaction, which,
however, disappeared on the addition of a very minute quantity of alkali;
for ten parts of silica in the solution it was sufficient to take one
part of alkali in order to give the liquid an alkaline reaction, so
slightly energetic are the acid properties of silicic acid. The solution
of silica obtained by this method becomes gelatinous on standing, on
being heated, or on evaporation under the receiver of an air-pump, &c.
The hydrosol is transformed into the hydrogel, the soluble hydrate into
the gelatinous.
[17] _See_ Chapter I., Note 18. A solution of water-glass mixed with an
excess of hydrochloric acid is poured into the dialyser, and the
outer vessel is filled with water, which is continually renewed.
The water carries off the sodium chloride and hydrochloric acid,
and the hydrosol remains in the dialyser.
Thus in addition to the gelatinous form of the silicic acid, there
exists also a variety of this substance, soluble in water, as is the case
with alumina. Such variation in properties and exactly the same relations
with regard to water characterise an immense series of other substances
having a great significance in nature. The number of such substances is
especially great among organic compounds, and particularly in those
classes of them which compose the principal material of the bodies of
animals and plants. It is sufficient to mention, for instance, the
gelatin which is familiar to all as carpenter's and other glues, and in
the form of size and jelly. The same substance is also known in the
solution which is used to join objects together. In a peculiar insoluble
condition it enters into the composition of hides and bones. These
various forms of gelatin differ in the same way as the different
varieties of silica. The property of forming a jelly is exactly the same
as in silica, and the adhesiveness of the solutions of both substances is
identical; soluble silica adheres like a solution of gelatin. The same
properties are again shown by starch, rosin, and albumin, and by a series
of similar substances. The diaphragms used in dialysis are also
insoluble, gelatinous, forms of colloids. The bodies of animals and
plants consist largely of similar matter, insoluble in water,
corresponding with the gelatinous or insoluble silicon hydrate, or with
glue. The albumin which coagulates when eggs are boiled is a typical form
of the gelatinous condition of such substances in the body. These slight
indications are sufficient in order to show how great is the significance
of those transformations which are so well marked in silica. The facts
discovered by _Graham_ in 1861-1864 comprise the most essential
acquisitions in the general association of these phenomena of nature in
the history of organic forms. The facility of transit from hydrogel to
hydrosol is the first condition of the possibility of the development of
organisms. The blood contains hydrosols, and the hydrogels of the same
substances are contained in the muscles and tissues, and especially on
the surface, of the body. All tissues are formed from the blood, and in
that case the hydrosols are converted into hydrogels.[18] The absence of
crystallisation, the property, apparently under the influence of feeble
agencies, of passing from the soluble condition to the insoluble, to the
gelatinous condition of the hydrogel, constitute the fundamental
properties of all colloids.[19]
[18] A similar process occurs in plants--for example, when they secrete
a store of material for the following year in their bulbs, roots,
&c. (for instance, the potato in its tubers), the solutions from
the leaves and stems penetrate into the roots and other parts in
the form of hydrosols, where they are converted into
hydrogels--that is, into an insoluble form, which is acted on with
difficulty and is easily kept unaltered until the period of
growth--for example, until the following spring--when they are
reconverted into hydrosols, and the insoluble substance re-enters
into the sap, and serves as a source of the hydrogels in the
leaves and other portions of plants.
[19] As regards their chemical composition the colloids are very
complex--that is, they have a high molecular weight and a large
molecular volume--in consequence of which they do not penetrate
through membranes, and are easily subject to variation in their
physical and chemical properties (owing to their complex structure
and polymerism?) They have but little chemical energy, and are
generally feeble acids, if belonging to the order of oxides or
hydrates, such as the hydrates of molybdic and tungstic acids
(Chapter XXI.). But now the number of substances capable, like
colloids, of passing into aqueous solutions and of easily
separating out from them, as well as of appearing in an insoluble
form, must be supplemented by various other substances, among
which soluble gold and silver (Chapter XXIV.) are of particular
interest. So that now it may be said that the capacity of forming
colloid solutions is not limited to a definite class of compounds,
but is, if not a general, at all events, an exceedingly widely
distributed phenomenon.
Silica, as regards its _salt forming properties_, stands in the series of
oxides on the boundary line on the side of the acids in just such a place
as alumina occupies on the side of the bases--that is, aluminium
hydroxide is the representative of the feeblest bases and silicic acid is
the least energetic of acids (at least in the presence of water--that is,
in aqueous solutions); in alumina, however, the basic properties are
distinctly expressed, while in silica the acid properties preponderate.
Like all feeble acid oxides it is capable of forming, with other acids,
saline compounds which are but slightly stable and are very easily
decomposed in the presence of water. The chief peculiarity of the
silicates consists in the number of their types. The salts formed with
nitric or sulphuric acid exist in one, two, and three tolerably stable
forms, but for acids like silicic acid the number of forms is very great,
almost unlimited. The natural silicates in particular furnish proof of
this fact; they contain various bases in combination with silica, and for
one and the same base there often exist various degrees of combination.
As feeble bases are capable of forming basic salts in addition to normal
salts--that is, a compound of a normal salt with a feeble base (either
the hydroxide or the oxide)--so the feeble acid oxides (although not all)
form, in addition to normal salts, highly acid salts--that is, normal
salts _plus_ acid (hydrate or anhydride). Such acids are boric,
phosphoric, molybdic, chromic, and especially silicic, acid.
In order to explain these relations it is necessary first to recollect
the existence of the various hydrates of silica, or silicic acids,[20]
and then to turn our attention to the similarity between silicon
compounds and metallic alloys. Silica is an oxide having the appearance
of, and in many respects the same properties as, those oxides which
combine with it, and if two metals are capable of forming homogeneous
alloys in which there exist definite or indefinite compounds, it is
permissible to assume a similar power of forming alloys in the case of
analogous oxides. Such alloys are found in indefinite, amorphous masses
in the form of glass, lava, slags, and a number of similar siliceous
compounds which do not contain any definite types of combination, but
nevertheless are homogeneous throughout their mass. By slow cooling, or
under other circumstances, definite crystalline compounds may--and
sometimes do--separate from this homogeneous mass, as also sometimes
definite crystalline alloys separate from metallic alloys.
[20] This is in accordance with the generally-accepted representation
of the relations between salts and the hydrates of acids, but it
is of little help in the study of siliceous compounds. Generally
speaking, it becomes necessary to explain the property of
(SiO_{2})_{_n_} to combine with (RO)_{_m_}, where _n_ may be
greater than _m_, and where R may be H_{2}, Ca, &c. Here we are
aided by those facts which have been attained by the investigation
of carbon compounds, especially with respect to glycol. Glycol is
a compound having the composition C_{2}H_{6}O_{2}, only differing
from alcohol, C_{2}H_{6}O, by an extra atom of oxygen. This
hydrate contains two hydroxyl groups, which may be successively
replaced by chlorine, &c. Hence the composition of glycol should
be represented as C_{2}H_{4}(OH)_{2}. It has been found that
glycol forms so-called polyglycols. Their origin will be
understood from the fact that glycol as a hydrate has a
corresponding anhydride of the composition C_{2}H_{4}O, known as
ethylene oxide. This substance is ethane, C_{2}H_{6}, in which two
hydrogens are replaced by one atom of oxygen. Ethylene oxide is
not the only anhydride of glycol, although it is the simplest one,
because C_{2}H_{4}O = C_{2}H_{4}(OH)_{2} - H_{2}O. Various other
anhydrides of glycol are possible, and have actually been
obtained, of the composition _n_C_{2}H_{4}(OH)_{2} - (_n_ -
1)H_{2}O = (C_{2}H_{4})_{_n_}O_{_n_ - 1}(OH)_{2}. These imperfect
anhydrides of glycol, or _polyglycols_, still contain hydroxyls
like glycol itself, and therefore are of an alcoholic character in
the same sense as glycol itself. They are obtained by various
methods, and, amongst others, by the direct combination of
ethylene oxide with glycol, because C_{2}H_{4}(OH)_{2} + (_n_ -
1)C_{2}H_{4}O = (C_{2}H_{4})_{_n_}O_{_n_ - 1}(OH)_{2}. The most
important circumstance, from a theoretical point of view, is that
these polyglycols may be distilled without undergoing
decomposition, and that the general formula given above expresses
their actual molecular composition. Hence we have here a direct
combination of the anhydride with the hydrate, and, moreover, a
repeated one. The formula A_{_n_}H_{2}O may be used to express the
composition of glycol and polyglycols with respect to ethylene
oxide in the most simple manner, if A stand for ethylene oxide.
When _n_ = 1 we have glycol, when _n_ is greater than 1 a
polyglycol. Such also is the relationship of the salts of hydrate
of silica, if A stand for silica, and if we imagine that H_{2}O
may also be taken _m_ times. Such a representation of the
_polysilicic acids_ corresponds with the representation of the
polymerism of silica. Laurent supposed the existence of several
polymeric forms, Si_{2}O_{4}, Si_{3}O_{6}, &c., besides silica,
SiO_{2}.
The formation of crystalline rocks in nature is partly of such a
nature. By aqueous or igneous agency, but in any case in a liquid
condition, those oxides which form the earth's crust and her crystalline
minerals came into mutual contact. First of all they formed a shapeless
mass, of which lava, glass, slags and solutions are examples, but little
by little, or else suddenly, some definite compounds of certain oxides
existing in this alloy or in the shapeless mass were formed. This is
entirely similar to two metals forming a homogeneous alloy,[21] and under
known circumstances (for instance, on cooling the alloy, or in the case
of aqueous solution when the two metals are simultaneously liberated from
the solution), definite crystalline compounds are separated. In any case
there is no doubt that there is less distinction between silica and
bases, than between bases and such anhydrides as, for instance, sulphuric
or nitric, or even carbonic, as is seen on comparing the physical and
chemical properties of silica and various kinds of oxides. Alumina,
especially, is exceedingly near akin to silica; not only in the hydrated
state, but also in the anhydrous condition, there exists a certain
similarity between the crystalline forms of alumina and silica, in the
uncombined state. Both are very hard, transparent, inactive,
non-volatile, infusible, and crystallise in the hexagonal system--in a
word, they are remarkably similar, and for this reason they are capable,
like two kindred metals, of entering into many different degrees of
combination. Isomorphous mixtures--that is, differing by the substitution
of oxides akin both in their physical and chemical characters--are very
frequently met with among minerals, and the study of the latter gave the
principal impetus to the study of isomorphism. Thus, in a whole series of
minerals, lime and magnesia are found in variable and interchangeable
proportions. Exactly the same may be said of potassium and sodium, of
alumina and ferric oxide, of manganous, ferrous, magnesium oxides, &c.
Such isomorphism does not, however, extend without change of form and
properties beyond certain rather narrow limits.[22] What I mean by this
is that lime is not always replaced totally, but often only in small
quantities, by magnesia, or by the manganous and ferrous oxides, without
changing the crystalline form. The same may be observed with regard to
potassium and lithium, which may be in part, but not completely, replaced
by sodium. On the total substitution of one metal for another, often
(although not invariably) the entire nature of the substance is changed;
for instance, _enstatite_ (or bronzite) is a magnesium bisilicate with a
small isomorphous substitution of calcium for magnesium; its composition
is expressed by the formula MgSiO_{3}, it belongs to the rhombic system.
On the entire substitution of calcium, _wollastonite_, CaSiO_{3}, of the
monoclinic system, is obtained; when manganese is substituted,
_rhodonite_, of the triclinic system, is produced; but in all of them the
angles of the prism are 86° to 88°.[23]
[21] For us the latter have not a saline character, only because they
are not regarded from this point of view, but an alloy of sodium
and zinc is, in a broad sense, a salt in many of its reactions,
for it is subject to the same double decompositions as sodium
phosphide or sulphide, which clearly have saline properties. The
latter (sodium phosphide), when heated with ethyl iodide, forms
ethyl phosphide, and the former--_i.e._ the alloy of zinc and
sodium--gives zinc ethyl; that is, the element (P, S, Zn) which
was united with the sodium passes into combination with the ethyl:
RNa + EtI = REt + NaI. By combining sodium successively with
chlorine, sulphur, phosphorus, arsenic, antimony, tin, and zinc,
we obtain substances having less and less the ordinary appearance
of salts, but if the alloy of sodium and zinc cannot be termed a
salt, then perhaps this name cannot be given to sodium sulphide,
and the compounds of sodium with phosphorus. The following
circumstance may also be observed: with chlorine, sodium gives
only one compound (with oxygen, at the most three), with sulphur
five, with phosphorus probably still more, with antimony naturally
still more, and the more analogous an element is to sodium, the
more varied are the proportions in which it is able to combine
with it, the less are the alterations in the properties which take
place by this combination, and the nearer does the compound formed
approach to the class of compounds known as indefinite chemical
compounds. In this sense a siliceous alloy, containing silica and
other acids, is a salt. The oxide to a certain extent plays the
same part as the sodium, whilst the silica plays the part of the
acid element which was taken up successively by zinc, phosphorus,
sulphur, &c., in the above examples. Such a comparison of the
silica compounds with alloys presents the great advantage of
including under one category the definite and indefinite silica
compounds which are so analogous in composition--that is, brings
under one head such crystalline substances as certain minerals,
and such amorphous substances as are frequently met with in
nature, and are artificially prepared, as glass, slags, enamels,
&c.
If the compounds of silica are substances like the metallic
alloys, then (1) the chemical union between the oxides of which
they are composed must be a feeble one, as it is in all compounds
formed between analogous substances. In reality such feeble
agencies as water and carbonic acid are able, although slowly, to
act on and destroy the majority of the complex silica compounds in
rocks, as we saw in the preceding chapter; (2) their formation,
like that of alloys, should not be accompanied by a considerable
alteration of volume; and this is actually the case. For example,
felspar has a specific gravity of about 2·6, and therefore, taking
its composition to be K_{2}O,Al_{2}O_{3},6SiO_{2}, we find its
volume, corresponding with this formula, to be 556·8. 2·6 = 214,
the volume of K_{2}O = 35, of Al_{2}O_{3} = 26, and of SiO_{2} =
22·6. Hence the sum of the volumes of the component oxides, 35 +
26 + 6 × 22·6 = 196, which is very nearly equal to that of the
felspar; that is, its formation is attended by a slight expansion,
and not by contraction, as is the case in the majority of other
cases when combinations determined by strong affinities are
accomplished. In the case in question the same phenomenon is
observed as in solutions and alloys--that is, as in cases of
feeble affinities. So also the specific gravity of glass is
directly dependent on the amount of those oxides which enter into
its composition. If in the preceding example we take the sp. gr.
of silica to be, not 2·65, but 2·2, its volume = 27·3, and the sum
of the volumes will be = 224--that is, greater than that of
orthoclase.
[22] It is, however, easy to imagine, and experience confirms the
supposition, that in a complex siliceous compound containing for
instance sodium and calcium, the whole of the sodium may be
replaced by potassium, and _at the same time_ the whole of the
calcium by magnesium, because then the substitution of potassium
for the sodium will produce a change in the nature of the
substance contrary to that which will occur from the calcium being
replaced by magnesium. That increase in weight, decrease in
density, increase of chemical energy, which accompanies the
exchange of sodium for potassium will, so to speak, be compensated
by the exchange of calcium for magnesium, because both in weight
and in properties the sum of Na + Ca is very near to the sum of K
+ Mg. _Pyroxene_ or _augite_ can be taken as an example; its
composition may be expressed by the formula CaMgSi_{2}O_{6}; that
is, it corresponds with the acid H_{2}SiO_{3}; it is a bisilicate.
In many respects it closely resembles another mineral called
'_spodumene_' (they are both monoclinic). This latter has the
composition Li_{6}Al_{8}Si_{15}O_{45}. On reducing both formulæ to
an equal contents of silica the following distinction will be
observed between them: spodumene
(Li_{2}O)_{6}(Al_{2}O_{3})_{8}30SiO_{2}; augite
(CaO)_{15}(MgO)_{15}30SiO_{2}. That is, the difference between
them consists in the sum of the magnesia and lime (MgO)_{15} +
(CaO)_{15} replacing the sum of the lithium oxide and alumina
(Li_{2}O)_{6} + (Al_{2}O_{3})_{8}; and in the chemical relation
these sums are near to one another, because magnesium and calcium,
both in forms of oxidation and in energy (as bases), in all
respects occupy a position intermediate between lithium and
aluminium, and therefore the sum of the first may be replaced by
the sum of the second.
If we take the composition of spodumene, as it is often
represented to be, Li_{2}O,Al_{2}O_{3},4SiO_{2}, the corresponding
formula of augite will be (CaO)_{2},(MgO)_{2},4SiO_{2}, and also
the amount of oxygen in the sum of Li_{2}OAl_{2}O_{3} will be the
same as in (CaO)_{2}(MgO)_{2}. I may remark, for the sake of
clearness, that lithium belongs to the first, aluminium to the
third group, and calcium and magnesium to the intermediate second
group; lithium, like calcium, belongs to the even series, and
magnesium and aluminium to the uneven.
The representation of the substitutions of analogous compounds
here introduced was first deduced by me in 1856. It finds much
confirmation in facts which have been subsequently discovered--for
example, with respect to tourmalin. Wülfing (1888), on the basis
of a number of analyses (especially of those by Röggs), states
that all varieties contain an isomorphous mixture of alkali and
magnesia tourmalin; into the composition of the former there
enters 12SiO_{2},3B_{2}O_{3},8Al_{2}O_{3},2Na_{2}O,4H_{2}O, and of
the latter 12SiO_{2},3B_{2}O_{3},5Al_{2}O_{3},12MgO,3H_{2}O. Hence
it is seen that the former contains in addition the sum of
3Al_{2}O_{3},2Na_{2}O,H_{2}O, whilst in the latter this sum of
oxides is replaced by 12MgO, in which there is as much oxygen as
in the sum of the more clearly-defined base 2Na_{2}O and less
basic 3Al_{2}O_{3}H_{2}O--that is, the relation is just the same
here as between augite and spodumene.
[23] With respect to the silica compounds of the various oxides, it
must be observed that only the _alkali salts_ are known in a
soluble form; all the others only exist in an insoluble form, so
that a solution of the alkali compounds of silica, or soluble
glass, gives a precipitate with a solution of the salts of the
majority of other metals, and this precipitate will contain the
silica compounds of the other bases. The maximum amount of the
gelatinous hydrate of silica, which dissolves in caustic potash,
corresponds with the formation of a compound, 2K_{2}O,9SiO_{2}.
But this compound is partially decomposed, with the precipitation
of hydrate of silica, on cooling the solution. Solutions
containing a smaller amount of silica may be kept for an
indefinite time without decomposing, and silica does not separate
out from the solution; but such compounds crystallise from the
solutions with difficulty. However, a crystalline bisilicate (with
water) has been obtained for sodium having the composition
Na_{2}O,SiO_{2}--_i.e._ corresponding to sodium carbonate. The
whole of the carbonic acid is evolved, and a similar soluble
sodium metasilicate is obtained on fusing 3·5 parts of sodium
carbonate with 2 parts of silica. If less silica be taken a
portion of the sodium carbonate remains undecomposed; however, a
substance may then be obtained of the composition Si(ONa)_{4},
corresponding with orthosilicic acid. It contains the maximum
amount of sodium oxide capable of combining with silica under
fusion. It is a sodium orthosilicate, (Na_{2}O)_{2},SiO_{2}.
Calcium carbonate, and the carbonates of the alkaline earths in
general, also evolve all their carbonic acid when heated with
silica, and in some instances even form somewhat fusible
compounds. Lime forms a fusible slag of _calcium silicate_, of the
composition CaO,SiO_{2} and 2CaO,3SiO_{2}. With a larger
proportion of silica the slags are infusible in a furnace. The
magnesium _slags_ are less fusible than those with lime, and are
often formed in smelting metals. Many compounds of the metals of
the alkaline earths with silica are also met with in nature. For
instance, among the magnesium compounds there is _olivine_,
(MgO)_{2},SiO_{2}, sp. gr. 3·4, which occurs in meteorites, and
sometimes forms a precious stone (peridote), and occurs in slags
and basalts. It is decomposed by acids, is infusible before the
blow-pipe, and crystallises in the rhombic system. _Serpentine_
has the composition 3MgO,2SiO_{2},2H_{2}O; it sometimes forms
whole mountains, and is distinguished for its great cohesiveness,
and is therefore used in the arts. It is generally tinted green;
its specific gravity is 2·5; it is exceedingly infusible, even
before the blowpipe. It is acted on by acids. Among the magnesium
compounds of silica, _talc_ is very widely used. It is frequently
met with in rocks which are widely distributed in nature, and
sometimes in compact masses; it can be used for writing like a
slate pencil or chalk, and being greasy to the touch, is also
known as _steatite_. It crystallises in the rhombic system, and
resembles mica in many respects; like it, it is divisible into
laminæ, greasy to the touch, and having a sp. gr. 2·7. These
laminæ are very soft, lustrous, and transparent, and are infusible
and insoluble in acids. The composition of talc approaches nearly
to 6MgO,5SiO_{2},2H_{2}O.
Among the crystalline silicates the following minerals are
known:--_Wollastonite_ (tabular-spar), crystallises in the
monoclinic system; sp. gr. 2·8; it is semi-transparent,
difficultly fusible, decomposed by acids, and has the composition
of a metasilicate, CaOSiO_{2}. But isomorphous mixtures of calcium
and magnesium silicates occur with particular frequency in nature.
The _augites_ (sp. gr. 3·3), diallages, hypersthenes, hornblendes
(sp. gr. 3·1), amphiboles, common asbestos, and many similar
minerals, sometimes forming the essential parts of entire rock
formations, contain various relative proportions of the
bisilicates of calcium and magnesium partially mixed with other
metallic silicates, and generally anhydrous, or only containing a
small amount of water. In the pyroxenes, as a rule, lime
predominates, and in the amphiboles (also of the monoclinic
system) magnesia predominates. Details upon this subject must be
looked for in works upon mineralogy.
The most remarkable complex siliceous compounds are the _felspars_, which
enter into nearly all the primary rocks like porphyry, granite, gneiss,
&c. These felspars always contain, in addition to silica and alumina,
oxides presenting more marked basic properties, such as potash, soda, and
lime. Thus the _orthoclase_ (adularia), or ordinary felspar (monoclinic)
of the granites, contains K_{2}O,Al_{2}O_{3},6SiO_{2}; _albite_ contains
the same substances, only with Na_{2}O instead of K_{2}O (it already
appertains to the triclinic system); _anorthite_ contains lime, and its
composition is CaO,Al_{2}O_{3},2SiO_{2}. On expressing the two last as
containing equal quantities of oxygen, we have:--
Albite Na_{2} Al_{2} Si_{6} O_{16}
Anorthite Ca_{2} Al_{4} Si_{4} O_{16}
It is then evident that on the conversion of albite into anorthite,
Na_{2}Si_{2} is replaced by Ca_{2}Al_{2}, and this sum, both in chemical
energy and in the form of oxide, may be considered as corresponding with
the first, because sodium and silicon are extreme elements in chemical
character (from groups I. and IV.), and calcium and aluminium are means
between them (from groups II. and III.), and actually both these felspar
minerals are not only of one (triclinic) system, but form (Tchermak,
Schuster) all possible kinds of definite compounds (isomorphous mixtures)
between themselves, as indicated by their composition and all their
properties. Thus oligoclase, andesine, labradorite, &c. (plagioclases),
are nothing more than mutual combinations of albite and anorthite.
Labradorite consists of albite, in combination with 1 to 2 molecules of
anorthite. The class of _zeolites_ corresponds to the felspars; they are
hydrated compounds of a similar composition to the felspars. Thus
_natrolite_ contains Na_{2}O,Al_{2}O_{3},3SiO_{2},2H_{2}O, and _analcime_
presents the same composition, but contains 4SiO_{2} instead of 3SiO_{2}.
In general, the felspars and zeolites contain RO,Al_{2}O_{3},_n_SiO_{2},
where _n_ varies considerably.[24]
[24] The majority of the siliceous minerals have now been obtained
artificially under various conditions. Thus N. N. Sokoloff showed
that slags very frequently contain peridote. Hautefeuille,
Chroustchoff, Friedel, and Sarasin obtained felspar identical in
all respects with the natural minerals. The details of the methods
here employed must be looked for in special works on mineralogy;
but, as an example, we will describe the method of the preparation
of felspar employed by Friedel and Sarasin (1881). From the fact
that felspar gives up potassium silicate to water even at the
ordinary temperature (Debray's experiments), they concluded that
the felspar in granites had an aqueous origin (and this may be
supposed to be the case from geological data); then, in the first
place, its formation could not be accomplished unless in the
presence of an excess of a solution of potassium silicate. In
order to render this argument clear I may mention, as an example,
that carnallite is decomposed by water into easily soluble
magnesium chloride and potassium chloride, and therefore if it is
of aqueous origin it could not be formed otherwise than from a
solution containing an excess of magnesium chloride, and, in the
second place, from a strongly-heated solution; again, felspar
itself and its fellow-components in granites are anhydrous. On
these facts were based experiments of heating hydrates of silica
with alumina and a solution of potassium silicate in a closed
vessel. The mixture was placed in a sealed platinum tube, which
was enclosed in a steel tube and heated to dull redness. When the
mixture contained an excess of silica the residue contained many
crystals of rock crystal and tridymite, together with a powder of
felspar, which formed the main product of the reaction when the
proportion of hydrate of silica was decreased, and a mixture of a
solution of potassium silicate with alumina precipitated together
with the silica by mixing soluble glass with aluminium chloride
was employed. The composition, properties, and forms of the
resultant felspar proved it to be identical with that found in
nature. The experiments approach very nearly to the natural
conditions, all the more as felspar and quartz are obtained
together in one mixture, as they so often occur in nature.
Such complex silicates are generally insoluble in water,[25] and if
they undergo change in it, it is but very slow, and more often only in
the presence of carbonic acid. Some of the silicates which are insoluble
in water are easily and directly decomposed by acids; for instance, the
zeolites and those fused silicates which contain a large quantity of
energetic bases--such as lime. Many of the silicates, like glass,[26] are
hardly changed by acids, particularly if they contain much silica, whilst
fusion with alkalis leads to the formation of compounds rich in bases,
after which acids decompose the alloys formed.[27]
[25] The application of _cements_ is based on this principle; they are
those sorts of 'hydraulic' lime which generally form a stony mass,
which hardens even under water, when mixed with sand and water.
The hydraulic properties of cements are due to their containing
calcareous and silico-aluminous compounds which are able to
combine with water and form hydrates, which are then unacted on by
water. This is best proved, in the first place, by the fact that
certain slags containing lime and silica, and obtained by fusion
(for example, in blast-furnaces), solidify like cements when
finely ground and mixed with water; and, in the second place, by
the method now employed for the manufacture of artificial cements
(formerly only peculiar and comparatively rare natural products
were used). For this purpose a mixture of lime and clay is taken,
containing about 25 p.c. of the latter; this mixture is then
heated, not to fusion, but until both the carbonic anhydride and
water contained in the clay are expelled. This mass when finely
ground forms Portland cement, which hardens under water. The
process of hardening is based on the formation of chemical
compounds between the lime, silica, alumina, and water. These
substances are also found combined together in various natural
minerals--for example, in the zeolites, as we saw above. In all
cases cement which has set contains a considerable amount of
water, and its hardening is naturally due to hydration--that is,
to the formation of compounds with water. Well-prepared and very
finely-ground cement hardens comparatively quickly (in several
days, especially after being rammed down), with 3 parts (and even
more) of coarse sand and with water, into a stony mass which is as
hard and durable as many stones, and more so than bricks and
limestone. Hence not only all maritime constructions (docks,
ports, bridges, &c.), but also ordinary buildings, are made of
Portland cement, and are distinguished for their great durability.
A combination of ironwork (ties, girders) and cement is
particularly suitable for the construction of aqueducts, arches,
reservoirs, &c. Arches and walls made of such cements may be much
less thick than those built up of ordinary stone. Hence the
production and use of cement rapidly increases from year to year.
The origin of accurate data respecting cements is chiefly due to
Vicat. In Russia Professor Schuliachenko has greatly aided the
extension of accurate data concerning Portland cement. Many works
for the manufacture of cement have already been established in
various parts of Russia, and this industry promises a great future
in the arts of construction.
[26] _Glass_ presents a similar complex composition, like that of many
minerals. The ordinary sorts of white glass contain about 75 p.c.
of silica, 13 p.c. of sodium oxide, and 12 p.c of lime; but the
inferior sorts of glass sometimes contain up to 10 p.c. of
alumina. The mixtures which are used for the manufacture of glass
are also most varied. For example, about 300 parts of pure sand,
about 100 parts of sodium carbonate, and 50 of limestone are
taken, and sometimes double the proportion of the latter. Ordinary
_soda-glass_ contains sodium oxide, lime, and silica as the chief
component parts. It is generally prepared from sodium sulphate
mixed with charcoal, silica, and lime (Chapter XII.), in which
case the following reaction takes place at a high temperature:
Na_{2}SO_{4} + C + SiO_{2} = Na_{2}SiO_{3} + SO_{2} + CO.
Sometimes potassium carbonate is taken for the preparation of the
better qualities of glass. In this case a glass, _potash-glass_,
is obtained containing potassium oxide instead of sodium oxide.
The best-known of these glasses is the so-called Bohemian glass or
crystal, which is prepared by the fusion of 50 parts of potassium
carbonate, 15 parts of lime, and 100 parts of quartz. The
preceding kinds of glass contain lime, whilst crystal glass
contains lead oxide instead. Flint glass--that is, the lead glass
used for optical instruments--is prepared in this manner,
naturally from the purest possible materials.
_Crystal-glass_--_i.e._ glass containing lead oxide--is softer
than ordinary glass, more fusible and has a higher index of
refraction. However, although the materials for the preparation of
glass be most carefully sorted, a certain amount of iron oxides
falls into the glass and renders it greenish. This coloration may
be destroyed by adding a number of substances to the vitreous
mass, which are able to convert the ferrous oxide into ferric
oxide; for example, manganese peroxide (because the peroxide is
deoxidised to manganous oxide, which only gives a pale violet tint
to the glass) and arsenious anhydride, which is deoxidised to
arsenic, and this is volatilised. The manufacture of glass is
carried on in furnaces giving a very high temperature (often in
regenerative furnaces, Chapter IX.). Large clay crucibles are
placed in these furnaces, and the mixture destined for the
preparation of the glass, having been first roasted, is charged
into the crucibles. The temperature of the furnace is then
gradually raised. The process takes place in three separate
stages. At first the mass intermixes and begins to react; then it
fuses, evolves carbonic acid gas, and forms a molten mass; and,
lastly, at the highest temperature, it becomes homogeneous and
quite liquid, which is necessary for the ultimate elimination of
the carbonic anhydride and solid impurities, which latter collect
at the bottom of the crucible. The temperature is then somewhat
lowered, and the glass is taken out on tubes and blown into
objects of various shapes. In the manufacture of window-glass it
is blown into large cylinders, which are then cut at the ends and
across, and afterwards bent back in a furnace into the ordinary
sheets. After being worked up, all glass objects have to be
subjected to a slow cooling (_annealing_) in special furnaces,
otherwise they are very brittle, as is seen in the so-called
'Rupert's drops,' formed by dropping molten glass into water;
although these drops preserve their form, they are so brittle that
they break up into a fine powder if a small piece be knocked off
them. Glass objects have frequently to be polished and chased. In
the manufacture of mirrors and many massive objects the glass is
cast and then ground and polished. Coloured glasses are either
made by directly introducing into the glass itself various oxides,
which give their characteristic tints, or else a thin layer of a
coloured glass is laid on the surface of ordinary glass. Green
glasses are formed by the oxides of chromium and copper, blue by
cobalt oxide, violet by manganese oxide, and red glass by cuprous
oxide and by the so-called purple of Cassius--_i.e._ a compound of
gold and tin--which will be described later. A yellow coloration
is obtained by means of the oxides of iron, silver, or antimony,
and also by means of carbon, especially for the brown tints for
certain kinds of bottle-glass.
From what has been said about glass it will be understood that it
is impossible to give a definite formula for it, because it is a
non-crystalline or amorphous alloy of silicates; but such an alloy
can only be formed within certain limits in the proportions
between the component oxides. With a large proportion of silica
the glass very easily becomes clouded when heated; with a
considerable proportion of alkalis it is easily acted on by
moisture, and becomes cloudy in time on exposure to the air; with
a large proportion of lime it becomes infusible and opaque, owing
to the formation of crystalline compounds in it; in a word, a
certain proportion is practically attained among the component
oxides in order that the glass formed may have suitable
properties. Nevertheless, it may be well to remark that the
composition of common glass approaches to the formula
Na_{2}O,CaO,4SiO_{2}.
The coefficient of cubical expansion of glass is nearly equal to
that of platinum and iron, being approximately 0·000027. The
specific heat of glass is nearly 0·18, and the specific gravity of
common soda glass is nearly 2·5, of Bohemian glass 2·4, and of
bottle glass 2·7. Flint glass is much heavier than common glass,
because it contains the heavier oxide of lead, its specific
gravity being 2·9 to 3·2.
[27] It must be recollected that although acids seem to act only feebly
on the majority of silicates, nevertheless a finely-levigated
powder of siliceous compounds is acted on by strong acids,
especially with the aid of heat, the basic oxides being taken up
and gelatinous silica left behind. In this respect sulphuric acid
heated to 200° with finely-divided siliceous compounds in a closed
tube acts very energetically.
According to the periodic law, the nearest analogues of silicon ought to
be elements of the uneven series, because silicon, like sodium,
magnesium, and aluminium, belongs to the uneven series.[28] Immediately
after silicon follows ekasilicon or _germanium_, Ge = 72, whose
properties were predicted (1871) before Winkler (1886) in Freiberg,
Saxony (Chapter XV. § 5), discovered this element in a peculiar silver
ore called _argyrodite_, Ag_{6}GeS_{5}.[29] Easily reduced from the oxide
by heating with hydrogen and charcoal, and separated from its solutions
by zinc, metallic germanium proved to be greyish white, easily
crystallisable (in octahedra), brittle, fusible (under a coating of fused
borax) at about 900°, and easily oxidisable; the specific gravity =
5·469, the atomic weight = 72·3, and the specific heat = 0·076,[30] as
might be expected for this element according to the periodic law. The
corresponding _germanium dioxide_, GeO_{2}, is a white powder having a
specific gravity of 4·703; water, especially when boiling, dissolves this
dioxide (1 part of GeO_{2} requires for solution 247 parts of water at
20°, 95 parts at 100°). It forms soluble salts with alkalis and is but
sparingly soluble in acids.[31] In a stream of chlorine the metal forms
_germanium chloride_, GeCl_{4}, which boils at 86°, and has a specific
gravity of 1·887 at 18°; water decomposes it, forming the oxide. All
these properties[32] of germanium, showing its analogy to silicon and
tin, form a most beautiful demonstration of the truth of the periodic
law.[33]
[28] Such elements as silicon, tin, and lead were only brought together
under one common group by means of the periodic law, although the
quadrivalency of tin and lead was known much earlier. Generally
silicon was placed among the non-metals, and tin and lead among
the metals.
[29] At first (February 1886) the want of material to work on, the
absence of a spectrum in the Bunsen's flame, and the solubility of
many of the compounds of germanium, presented difficulties in the
researches of Professor Winkler, who, on analysing argyrodite by
the usual method, obtained a constant loss of 7 p.c., and was thus
led to search for a new element. The presence of arsenic and
antimony in the accompanying minerals also impeded the separation
of the new metal. After fusion with sulphur and sodium carbonate,
argyrodite gives a solution of a sulphide which is precipitated by
an _excess_ of hydrochloric acid; germanium sulphide is soluble in
ammonia and then precipitated by hydrochloric acid, as a _white_
precipitate, which is dissolved (or decomposed) by water. After
being oxidised by nitric acid, dried and ignited germanium
sulphide leaves the oxide GeO_{2}, which is reduced to the metal
when ignited in a stream of hydrogen.
[30] G. Kobb determined the spectrum of germanium, when the metal was
taken as one of the electrodes of a powerful Ruhmkorff's coil. The
wave-lengths of the most distinct lines are 602, 583, 518, 513,
481, 474, millionths of a millimetre.
[31] If germanium or germanium sulphide be heated in a stream of
hydrochloric acid, it forms a volatile liquid, boiling at 72°,
which Winkler regarded as germanium chloride, GeCl_{2}, or
germanium chloroform, GeHCl_{3}. It is decomposed by water,
forming a white substance, which may perhaps be the hydrate of
germanious oxide, GeO, and acts as a powerfully reducing agent in
a hydrochloric acid solution.
[32] Under certain circumstances germanium gives a blue coloration like
that of ultramarine, as Winkler showed, which might have been
expected from the analogy of germanium with silicon.
[33] Winkler expressed this in the following words (_Jour. f. pract.
Chemie_, 1886 [2], 34, 182-183): '... es kann keinem Zweifel mehr
unterliegen, dass das neue Element nichts Anderes, als das vor
fünfzehn Jahren von _Mendeléeff_ prognosticirte _Ekasilicium_
ist.'
'Denn einen schlagenderen Beweis für die Richtigkeit der Lehre von
der Periodicität der Elemente, als den, welchen die Verkörperung
des bisher hypothetischen "Ekasilicium" in sich schliesst, kann es
kaum geben, und er bildet in Wahrheit mehr, als die blosse
Bestätigung einer kühn aufgestellten Theorie, er bedeutet eine
eminente Erweiterung des chemischen Gesichtfeldes, einen mächtigen
Schritt in's Reich der Erkenntniss.'
The increase of atomic weight from silicon 28 to germanium 72 is 44--that
is, about the same difference as there is in the atomic weights of
chlorine and bromine; between germanium and its next analogue, _tin_ (Sn
= 118), the difference is 46--that is, almost as much as the amount by
which the atomic weight of iodine exceeds that of bromine.
Metallic tin is rarely met with in _nature_; it occurs in the veins of
ancient formations, almost exclusively in the form of oxide, SnO_{2},
called _tin-stone_. The best known tin deposits are in Cornwall and in
Malacca. In Russia, tin ores have been found in small quantities on the
shores of Lake Ladoga, in Pitkarand. The crushed ore may easily be
separated from the earthy matter accompanying it by washing on inclined
tables, as the tin-stone has a specific gravity of 6·9, whilst the
impurities are much lighter. _Tin oxide is very easily reduced_ to
metallic tin by heating with charcoal. For this reason tin was known in
ancient times, and the Ph[oe]nicians brought it from England. Metallic
tin is cast into ingots of considerable weight or into thin sticks or
rods. Tin has a white colour, rather duller than that of silver. It fuses
easily at 232°, and crystallises on cooling. Its specific gravity is
7·29. The crystalline structure of ordinary tin is noticed in bending tin
rods, when a peculiar sound is heard, produced by the fracture of the
particles of tin along the surfaces of crystalline structure.
When pure tin is cooled to a low temperature it splits up into separate
crystals, the bond between the particles is lost, the tin assumes a grey
colour, becomes less brilliant--in a word, its properties become changed,
as Fritzsche showed. This depends on the peculiar structure which the tin
then acquires, and is particularly remarkable because it is effected by
cold in a solid.[33 bis] If such tin be fused, or even simply heated, it
becomes like ordinary tin, but is again changed when cooled. When in this
condition tin has a specific gravity of 7·19. Similarly, tin is obtained
by the action of the galvanic current on a solution of tin chloride; it
then appears in crystals of the cubic system, and has a specific gravity
of 7·18--that is, the same as when cooled.[34]
[33 bis] Emilianoff (1890) states that in the cold of the Russian
winter 30 out of 200 tin moulds for candles were spoilt through
becoming quite brittle.
[34] The tin deposited by an electric current from a neutral solution
of SnCl_{2} easily oxidises and becomes coated with SnO (Vignon,
1889).
Tin is softer than silver and gold, and is only surpassed by lead in
this respect. In addition to this it is very ductile, but its tenacity is
very slight, so that wire made from it will bear but little strain. In
consequence of its ductility it is easily worked, by forging and rolling
into very thin sheets (tin foil), which are used for wrapping many
articles to preserve them from moisture, &c. In this case, however, and
in many others, lead is mixed with the tin, which, within certain limits,
does not alter the ductility. Whilst so soft at the ordinary temperatures
tin becomes brittle at 200°, before fusing. Tin powder may be easily
obtained if the metal be fused and then stirred whilst cooling. At a
white heat tin may be distilled, but with more difficulty than zinc. If
molten tin comes into contact with oxygen, it oxidises, forming stannic
oxide, SnO_{2}, _and its vapour burns_ with a white flame. _At ordinary
temperatures tin does not oxidise_, and this very important property of
tin allows it to be applied in many cases for covering other metals to
prevent their oxidising. This is termed _tinning_. Iron and copper are
frequently tinned. Iron and steel sheets, coated with tin, bear the name
of tin plate (for the most part made in England), and are used for
numerous purposes. Tin plate is prepared by immersing iron sheets,
previously thoroughly cleansed by acid and mechanical means, into molten
tin.[34 bis]
[34 bis] If after this the coating of tin be rapidly cooled--for
instance, by dashing water over it--it crystallises into diverse
star-shaped figures, which become visible when the sheets are
first immersed in dilute aqua regia and then in a solution of
caustic soda.
The coating of iron by tin, guards it against the direct access of
air, but it only preserves the iron from oxidation so long as it
forms a perfectly continuous coating. If the iron is left bare in
certain places, it will be powerfully oxidised at these spots,
because the tin is electro-negative with respect to the iron, and
thus the oxidation is confined entirely to the iron in the
presence of tin. Hence a coating of tin over iron objects only
partially preserves them from rusting. In this respect a coating
of zinc is more effectual. However, a dense and invariable alloy
is formed over the surface of contact of the iron and tin, which
binds the coating of tin to the remaining mass of the iron. Tin
may be fused with cast iron, and gives a greyish-white alloy,
which is very easily cast, and is used for casting many objects
for which iron by itself would be unsuitable owing to its ready
oxidisability and porosity. The coating of copper objects by tin
is generally done to preserve the copper from the action of acid
liquids, which would attack the copper in the presence of air and
convert it into soluble salts. Tin is not acted on in this manner,
and therefore copper vessels for the preparation of food should be
tinned.
Tin with copper forms _bronze_, an alloy which is most extensively used
in the arts. Bronze has various colours and a variety of physical
properties, according to the relative amount of copper and tin which it
contains. With an excess of copper the alloy has a yellow colour; the
admixture of tin imparts considerable hardness and elasticity to the
copper. An alloy containing 78 parts of copper and about 22 per cent. of
tin is so elastic that it is used for casting bells, which naturally
require a very elastic and hard alloy.[35] For casting statues and
various large or small ornamental articles alloys containing 2 to 5 p.c.
of tin, 10 to 30 p.c. of zinc, and 65 to 85 p.c. of copper are used.[36]
Tin is also often used alloyed with lead, for making various objects--for
instance, drinking vessels.
[35] The ancient Chinese alloys, containing about 20 p.c. of tin
(specific gravity of alloys about 8·9), which have been rapidly
cooled, are distinguished for their resonance and elasticity.
These alloys were formerly manufactured in large quantities in
China for the musical instruments known as _tom-toms_. Owing to
their hardness, alloys of this nature are also employed for
casting guns, bearings, &c., and an alloy containing about 11 p.c.
of tin (corresponding with the ratio Cu_{15}Sn) is known as
gun-metal. The addition of a small quantity of phosphorus, up to 2
p.c., renders bronze still harder and more elastic, and the alloy
so formed is now used under the name of phosphor-bronze.
The alloy SnCu_{3} is brittle, of a bluish colour, and has nothing
in common with either copper or tin in its appearance or
properties. It remains perfectly homogeneous on cooling, and
acquires a crystalline structure (Riche). All these signs clearly
indicate that the alloy SnCu_{3} is a product of chemical
combination, which is also seen to be the case from its density,
8·91. Had there been no contraction, the density of the alloy
would be 8·21. It is the heaviest of all the alloys of tin and
copper, because the density of tin is 7·29 and of copper 8·8. The
alloy SnCu_{4}, specific gravity 8·77, has similar properties. All
the alloys except SnCu_{3} and SnCu_{4} split up on cooling; a
portion richer in copper solidifies first (this phenomenon is
termed the _liquation_ of an alloy), but the above two alloys do
not split up on cooling. In these and many similar facts we can
clearly distinguish a _chemical union between the metals_ forming
an alloy. The alloys of tin and copper were known in very remote
ages, before iron was used. The alloys of zinc and tin are less
used, but alloys composed of zinc, tin, and copper frequently
replace the more costly bronze. Concerning the alloys of lead
_see_ Note 46.
[36] An excellent proof of the fact that alloys and solutions are
subject to law is given, amongst others, by the application of
Raoult's method (Chapter I., Note 49) to solutions of different
metals in tin. Thus Heycock and Neville (1889) showed that the
temperature of solidification of molten tin (226°·4) is lowered by
the presence of a small quantity of other metals in proportion to
the concentration of the solution. The following were the
reductions of the temperature of solidification of tin obtained by
dissolving in it atomic proportions of different metals (for
example, 65 parts of zinc in 11,800 parts of tin); Zn 2°·53, Cu
2°·47, Ag 2°·67, Cd 2°·16, Pb 2°·22, Hg 2°·3, Sb 2° [rise], Al
1°·34. As Raoult's method (Chapter VII.) enables the molecular
weight to be determined, the almost perfect identity of the
resultant figures (except for aluminium) shows that the molecules
of copper, silver, lead, and antimony contain _one atom in the
molecule_, like zinc, mercury, and cadmium. They obtained the same
result (1890) for Mg, Na, Ni, Au, Pd, Bi and In. It should here be
mentioned that Ramsay (1889) for the same purpose (the
determination of the molecular weight of metals on the basis of
their mutual solution) took advantage of the variation of the
vapour tension of mercury (_see_ Vol. I., p. 134), containing
various metals in solution, and he also found that the
above-mentioned metals contain but one atom in the molecule.
Tin decomposes the vapour of water when heated with it, liberating the
hydrogen and forming stannic oxide. Sulphuric acid, diluted with a
considerable quantity of water, does not act, or at all events acts very
slightly, on tin, but tin reduces hot strong sulphuric acid, when not
only sulphurous anhydride but also sulphuretted hydrogen is evolved.
Hydrochloric acid acts very easily on tin, with evolution of hydrogen and
formation of stannous chloride, SnCl_{2}, in solution, which, with an
excess of hydrochloric acid and access of air, is converted into stannic
chloride: SnCl_{2} + 2HCl + O = SnCl_{4} + H_{2}O.[36 bis] Nitric acid
diluted with a considerable quantity of water dissolves tin at the
ordinary temperature, whilst the nitric acid itself is reduced, forming,
amongst other products, ammonia and hydroxylamine. Here the tin passes
into solution in the form of stannous nitrate. Stronger nitric acid (also
more dilute, when heated) transforms the tin into its highest grade of
oxidation, SnO_{2}, but the latter then appears as the so-called
metastannic acid, which does not dissolve in nitric acid, and therefore
the tin does not pass into solution. Feeble acids--for instance, carbonic
and organic acids--do not act on tin even in the presence of oxygen,
because tin does not form any powerful bases.
[36 bis] The action of a mixture of hydrochloric acid and tin forms an
excellent means of reducing, wherein both the hydrogen liberated
by the mixture (at the moment of separation) and the stannous
chloride act as powerful reducing and deoxidising agents. Thus,
for instance, by this mixture nitro-compounds are transformed into
amido-compounds--that is, the elements of the group NO_{2} are
reduced to NH_{2}.
It is important to remark as a characteristic of tin that it is reduced
from its solutions by many metals which are more easily oxidised, as, for
instance, by zinc.
_In combination_, _tin_ appears in the two types, SnX_{4} and
SnX_{2},[37] compounds of the intermediate type, Sn_{2}X_{6}, being also
known, but these latter pass with remarkable facility in most cases into
compounds of the higher and lower types, and therefore the form SnX_{3}
cannot be considered as independent.
[37] Many volatile compounds of tin are known, whose molecular weights
can therefore be established from their vapour densities. Among
these may be mentioned stannic chloride, SnCl_{4}, and stannic
ethide, Sn(C_{2}H_{5})_{4} (the latter boils at about 150°). But
V. Meyer found the vapour density of stannous chloride, SnCl_{2},
to be variable between its boiling point (606°) and 1100°, owing,
it would seem, to the fact that the molecule then varies from
Sn_{2}Cl_{4} to SnCl_{2}, but the vapour density proved to be less
than that indicated by the first and greater than that shown by
the second formula, although it approaches to the latter as the
temperature rises--that is, it presents a similar phenomenon to
that observed in the passage of N_{2}O_{4} into NO_{2}.
_Stannous oxide_, SnO, in an anhydrous condition is obtained by boiling
solutions of stannous salts with alkalis, the first action of the alkali
being to precipitate a white hydrate of stannous oxide, Sn(OH)_{2}SnO.
The latter when heated parts with water as easily as the hydrate of
copper oxide. In this form stannous oxide is a black crystalline powder
(specific gravity 6·7) capable of further oxidation when heated. The
hydrate is freely soluble in acids, and also in potassium and sodium
hydroxides, but not in aqueous ammonia.[38] This property indicates the
feeble basic properties of this lower oxide, which acts in many cases as
a reducing agent.[39] Among the compounds corresponding with stannous
oxide the most remarkable and the one most frequently used is stannous
chloride or _chloride of tin_, SnCl_{2}, also called proto-chloride of
tin (because it is the lowest chloride, containing half as much Cl as
SnCl_{4}). It is a transparent, colourless, crystalline substance,
melting at 250° and boiling at 606°. Water dissolves it, without visible
change (in reality partial decomposition occurs, as we shall see
presently). It is also soluble in alcohol. It is obtained by heating tin
in dry hydrochloric acid gas, the hydrogen being then liberated, or by
dissolving metallic tin in hot strong hydrochloric acid and then
evaporating quickly. On cooling, crystals of the monoclinic system are
obtained having the composition SnCl_{2},2H_{2}O. An aqueous solution of
this substance absorbs oxygen from the atmosphere, and gives a
precipitate containing stannic oxide. From this it follows that a
solution of stannous chloride will act as a reducing agent, a fact
frequently made use of in chemical investigations--for example, for
reducing metals from their solutions--since even mercury may be reduced
to a metallic state from its salts by means of stannous chloride. This
reducing property is also employed in the arts, especially in the dyeing
industry, where this substance in the form of a crystalline salt finds an
extensive application, and is known as _tin salt_ or tin crystals.
[38] When rapidly boiled, an alkaline solution of stannous oxide
deposits tin and forms stannic oxide, 2SnO = Sn + SnO_{4}, which
remains in the alkaline solution.
[39] Weber (1882) by precipitating a solution of stannous chloride with
sodium sulphite (this salt as a reducing agent prevents the
oxidation of the stannous compound) and dissolving the washed
precipitate in nitric acid, obtained crystals of _stannous
nitrate_, Sn(NO_{3})_{2},20H_{2}O, on refrigerating the solution.
This crystallo-hydrate easily melts, and is deliquescent. Besides
this, a more stable anhydrous basic salt, Sn(NO_{3})_{2},SnO, is
easily formed. In general, stannous oxide as a feeble base easily
forms basic salts, just as cupric and lead oxides do. For the same
reason SnX_{2} easily forms double salts. Thus a potassium salt,
SnK_{2}Cl_{4},H_{2}O, and especially an ammonium salt,
Sn(NH_{4})_{2}Cl_{4},H_{2}O, called _pink salt_, are known. Some
of these salts are used in the arts, owing to their being more
stable than tin salts alone. Stannous bromide and iodide, SnBr_{2}
and SnI_{2}, resemble the chloride in many respects.
Among other stannous salts a sulphate, SnSO_{4}, is known. It is
formed as a crystalline powder when a solution of stannous oxide
in sulphuric acid is evaporated under the receiver of an air-pump.
The feeble basic character of the stannous oxide is clearly seen
in this salt. It decomposes with extreme facility, when heated,
into stannic oxide and sulphurous anhydride, but it easily forms
double salts with the salts of the alkali metals.
In gaseous hydrochloric acid, stannous chloride, SnCl_{2},2H_{2}O,
forms a liquid having the composition SnCl_{2},HCl,3H_{2}O (sp gr.
2·2, freezes at -27°), and a solid salt, SnCl_{2},H_{2}O (Engel).
_Stannic oxide_, SnO_{2}, occurring in nature as _tinstone_, or
_cassiterite_, is formed during the oxidation or combustion of heated tin
in air as a white or yellowish powder which fuses with difficulty. It is
prepared in large quantities, being used as a white vitreous mixture for
coating ordinary tiles and similar earthenware objects with a layer of
easily fusible glass or enamel. Acid solutions of stannic oxide treated
with alkalis, and alkaline solutions treated with acids, give a
precipitate of stannic hydroxide, Sn(OH)_{4}, also known as stannic acid,
which, when heated, gives up water and leaves the anhydride, SnO_{2},
which is insoluble in acids, clearly showing the feebleness of its basic
character. When fused with alkali hydroxides (not with their carbonates
or acid sulphates), an alkaline compound is obtained which is soluble in
water. Stannic hydroxide, like the hydrates of silica, is a colloidal
substance, and presents several different modifications, depending on the
method of preparation, but having an identical composition; the various
hydroxides have also a different appearance, and act differently with
reagents. For instance, a distinction is made between ordinary stannic
acid and metastannic acid. _Stannic acid_ is produced by precipitation by
soda or ammonia from a freshly-prepared solution of stannic chloride,
SnCl_{4}, in water; on drying the precipitate thus obtained, a
non-crystalline mass is formed, which is freely soluble in strong
hydrochloric or nitric acids, and also in potassium and sodium
hydroxides. This ordinary stannic acid may be still better obtained from
sodium stannate by the action of acids. _Metastannic acid_ is insoluble
in sulphuric and nitric acids. It is obtained in the form of a heavy
white powder by treating tin with nitric acid; hydrochloric acid does not
dissolve it immediately, but changes it to such an extent that, after
pouring off the acid, water extracts the stannic chloride, SnCl_{4},
already formed. Dilute alkalis not only dissolve metastannic acid, but
also transform it into salts, which, slowly, yet completely, dissolve in
_pure water_, but are insoluble even in dilute alkali hydroxides. Dilute
hydrochloric acid, especially when boiling, changes the ordinary hydrate
into metastannic acid. On this depends, by the way, the formation of a
white precipitate, stannic hydroxide, from solutions of stannous and
stannic chlorides diluted with water. The stannic oxide first dissolved
changes under the influence of hydrochloric acid into metastannic acid,
which is insoluble in water in the presence of hydrochloric acid.
Solutions of metastannic acid differ from solutions of ordinary stannic
acid, and in the presence of alkali they change into solutions of
ordinary acid, so that metastannic acid corresponds principally with the
acid compounds of stannic oxide, and ordinary stannic acid with the
alkaline compounds.[40] Graham obtained a soluble colloidal hydroxide; it
is subject to the same transformations that are in general peculiar to
colloids.
[40] Frémy supposes the cause of the difference to consist in a
difference of polymerisation, and considers that the ordinary acid
corresponds with the oxide SnO_{2}, and the meta-acid with the
oxide Sn_{5}O_{10}, but it is more probable that both are
polymeric but in a different degree. Stannic acid with sodium
carbonate gives a salt of the composition Na_{2}SnO_{3}. The same
salt is also obtained by fusing metastannic acid with sodium
hydroxide, whilst metastannic acid gives a salt,
Na_{2}SnO_{3},4SnO_{2} (Frémy), when treated with a dilute
solution of alkali; moreover, stannic acid is also soluble in the
ordinary stannate, Na_{2}SnO_{3} (Weber), so that both stannic
acids (like both forms of silica) are capable of polymerisation,
and probably only differ in its degree. In general, there is here
a great resemblance to silica, and Graham obtained a solution of
stannic acid by the direct dialysis of its alkaline solution. The
main difference between these acids is that the meta-acid is
soluble in hydrochloric acid, and gives a precipitate with
sulphuric acid and stannous chloride, which do not precipitate the
ordinary acid. Vignon (1889) found that more heat is evolved in
dissolving stannic acid in KHO than metastannic.
Stannic oxide shows the properties of a slightly energetic and
intermediate oxide (like water, silica, &c.); that is to say, it forms
saline compounds both with bases and with acids, but both are easily
decomposed, and are but slightly stable. But still the acid character is
more clearly developed than the basic, as in silica, germanic oxide, and
lead dioxide. This determines the character of the compounds SnX_{4},
corresponding to stannic chloride, SnCl_{4} (also called tetrachloride of
tin). It is obtained in an anhydrous condition by the direct action of
chlorine on tin, and is then easily purified, because it is a liquid
boiling at 114°, and therefore can be easily distilled. Its specific
gravity is 2·28 (at 0°), and it fumes in the open air (spiritus fumans
libavii), reacting on the moisture of the air, thus showing the
properties of a chloranhydride. Water however does not at first decompose
it, but dissolves it, and on evaporation gives the crystallo-hydrate
SnCl_{4},5H_{2}O. If but little water be taken, crystals containing
SnCl_{4},3H_{2}O are formed, which part with one-third of the water when
placed under the receiver of the air-pump. A large quantity of water
however, especially on heating, causes a precipitate of metastannic
acid[41] and formation of HCl.
[41] The formation of the compound SnCl_{4},3H_{2}O is accompanied by
so great a contraction that these crystals, although they contain
water, are heavier than the anhydrous chloride SnCl_{4}. The
penta-hydrated crystallo-hydrate absorbs dry hydrochloric acid,
and gives a liquid of specific gravity 1·971, which at 0° yields
crystals of the compound SnCl_{4},2HCl,6H_{2}O (it corresponds
with the similar platinum compound), which melt at 20° into a
liquid of specific gravity 1·925 (Engel).
Stannic chloride combines with ammonia (SnCl_{4},4NH_{3}),
hydrocyanic acid, phosphoretted hydrogen, phosphorus pentachloride
(SnCl_{4},PCl_{5}), nitrous anhydride and its chloranhydride
(SnCl_{4},N_{2}O_{3} and SnCl_{4},2NOCl), and with metallic
chlorides (for example, K_{2}SnCl_{6}, (NH_{4})_{2}SnCl_{6}, &c.)
In general, a highly-developed faculty for combination is observed
in it.
Tin does not combine directly with iodine, but if its filings be
heated in a closed tube with a solution of iodine in carbon
bisulphide, it forms stannic iodide, SnI_{4}, in the form of red
octahedra which fuse at 142° and volatilise at 295°. The fluorine
compounds of tin have a special interest in the history of
chemistry, because they give a series of double salts which are
isomorphous with the salts of hydrofluosilicic acid, SiR_{2}F_{6},
and this fact served to confirm the formula SiO_{2} for silica, as
the formula SnO_{2} was indubitable. Although _stannic fluoride_,
SnF_{4}, is almost unknown in the free state, its corresponding
double salts are very easily formed by the action of hydrofluoric
acid on alkaline solutions of stannic oxide; thus, for example, a
crystalline salt of the composition SnK_{2}F_{6},H_{2}O is
obtained by dissolving stannic oxide in potassium hydroxide and
then adding hydrofluoric acid to the solution. The barium salt,
SnBaF_{6},3H_{2}O, is sparingly soluble like its corresponding
silicofluoride. The more soluble salt of strontium,
SnSrF_{6},2H_{2}O, crystallises very well, and is therefore more
important for the purposes of research; it is isomorphous with the
corresponding salt of silicon (and titanium); the magnesium salt
contains 6H_{2}O.
Stannic sulphide, SnS_{2}, is formed, as a yellow precipitate, by
the action of sulphuretted hydrogen on acid solutions of stannic
salts; it is easily soluble in ammonium and potassium sulphides,
because it has an acid character, and then forms thiostannates
(see Chapter XX.). In an anhydrous state it has the form of
brilliant golden yellow plates, which may be obtained by heating a
mixture of finely-divided tin, sulphur, and sal-ammoniac for a
considerable time. It is sometimes used in this form under the
name of mosaic gold, as a cheap substitute for gold-leaf in
gilding wood articles. On ignition it parts with a portion of its
sulphur, and is converted into stannous sulphide SnS. It is
soluble in caustic alkalis. Hydrochloric acid does not dissolve
the anhydrous crystalline compound, but the precipitated powdery
sulphide is soluble in boiling strong hydrochloric acid, with the
evolution of hydrogen sulphide.
_The alkali compounds of stannic oxide_--that is, the compounds in which
it plays the part of an acid, corresponding in this respect to the
compounds of silica and other anhydrides of the composition RO_{2}--are
very easily formed and are used in the arts. Their composition in most
cases corresponds with the formula SnM_{2}O_{3}--that is, SnO(MO)_{2},
similar to CO(MO)_{2}, where M = K, Na. Acids, even feeble acids like
carbonic, decompose the salts, like the corresponding compounds of
alumina or silica. In order to obtain _potassium stannate_, which
crystallises in rhombohedra, and has the composition
SnK_{2}O_{3},3H_{2}O, potassium hydroxide (8 parts) is fused, and
metastannic acid (3 parts) gradually added. _Sodium stannate_ is prepared
in practice in large quantities by heating a solution of caustic soda
with lead oxide and metallic tin. In this last case an alkaline solution
of lead oxide is formed, and the tin acts on the solution in such a way
as to reduce the lead to the metallic state, and itself passes into
solution. It is very remarkable that lead displaces tin when in
combination with acids, whilst tin, on the contrary, displaces lead from
its alkali compounds. By dissolving the mass obtained in water, and
adding alcohol, sodium stannate is precipitated, which may then be
dissolved in water and purified by re-crystallisation. In this case it
has the composition SnNa_{2}O_{3},3H_{2}O if separated from strong
solutions, and SnNa_{2}O_{3},10H_{2}O when crystallised at a low
temperature from dilute solutions. In the arts this salt is used as a
mordant in dyeing operations. With a cold solution of sodium hydroxide
metastannic acid forms a salt of the composition
(NaHO)_{2},5SnO_{2},3H_{2}O, from which Frémy drew his conclusions
concerning the polymerism of metastannic acid. Tin, like other metals and
many metalloids, gives a peroxide form of combination or _perstannic
oxide_. This substance was obtained by Spring (1889) in the form of a
hydrate, H_{2}Sn_{2}O_{7} = 2(SnO_{3})H_{2}O, by mixing a solution of
SnCl_{2}, containing an excess of HCl, with freshly prepared peroxide of
barium. A cloudy liquid is then obtained, and this after being subjected
to dialysis leaves a gelatinous mass which on drying is found to have the
composition Sn_{2}H_{2}O_{7}. Above 100° this substance gives off oxygen
and leaves SnO_{2}. It is evident that SnO_{3} bears the same relation to
SnO_{2} as H_{2}O_{2} to H_{2}O or ZnO_{2} to ZnO, &c.
Tin occupies the same position amongst the analogues of silicon as
cadmium and indium amongst the analogues of magnesium and aluminium
respectively, and as in each of these cases the heavier analogues with a
high atomic weight and a special combination of properties--namely,
mercury and thallium--are known, so also for silicon we have _lead_ as
the heaviest analogue (Pb = 206), with a series of both kindred and
special properties. The higher type, PbX_{4}--for instance, PbO_{2}--is
in a chemical sense far less stable than the lower type, PbX. The
ordinary compounds of lead correspond with the latter, and in addition to
this, PbO, although not particularly energetic, is still a decided base
easily forming basic salts, PbX_{2}(PbO)_{n}. Although the compounds
PbX_{4}, are unstable they offer many points of analogy with the
corresponding compounds of tin SnO_{2}; this is seen, for instance, in
the fact that PbO_{2} is a feeble acid, giving the salt PbK_{2}O_{3},
that PbCl_{4} is a liquid like SnCl_{4} which is not affected by
sulphuric acid, and that PbF_{4} gives double salts, like SnF_{4} or
SiF_{4} (Brauner 1894. See Chapter II., Note 49 bis); Pb(C_{2}H_{5})_{4}
also resembles Sn(C_{2}H_{5})_{4} &c. All this shows that lead is a true
analogue of tin, as Hg is of cadmium.[41 bis]
[41 bis] Although this has long been generally recognised from the
resemblance between the two metals, still from a chemical point of
view it has only been demonstrated by means of the periodic law.
_Lead_ is found in nature in considerable masses, in the form of galena,
_lead sulphide_, PbS.[42] The specific gravity of galena is 7·58, colour
grey; it crystallises in the regular system, and has a fine metallic
lustre. Both the native and artificial sulphides are insoluble in acids
(hydrogen sulphide gives a black precipitate with the salts PbX_{2}).[42
bis] When heated, lead melts, and in the open air is either totally or
partially transformed into white lead sulphate, PbSO_{4}, as it also is
by many oxidising agents (hydrogen peroxide, potassium nitrate). Lead
sulphate is also insoluble in water,[43] and lead is but rarely met with
in this form in nature. The chromates, vanadates, phosphates, and similar
salts of lead are also somewhat rare. The carbonate, PbCO_{2}, is
sometimes found in large masses, especially in the Altai region. Lead
sulphide is often worked for extracting the silver which it contains; and
as the lead itself also finds manifold industrial applications, this work
is carried out on an exceedingly large scale. Many methods are employed.
Sometimes the lead sulphide is decomposed by heating it with cast iron.
The iron takes up the sulphur from the lead and forms easily-fusible iron
sulphide, which does not mix with the heavier reduced lead. But another
process is more frequently used: the lead ore (it must be clean; that is,
free from earthy matter, which may be easily removed by washing) is
heated in a reverberatory furnace to a moderate temperature with a free
access of air. During this operation part of the lead sulphide oxidises
and forms lead sulphate, PbSO_{4}, and lead oxide. When the oxidation of
part of the lead has been attained, it is necessary to shut off the air
supply and increase the temperature, then the oxidised compounds of the
lead enter into reaction with the remaining lead sulphide, with formation
of sulphurous anhydride and metallic lead. At first from PbS + O_{3}, PbO
+ SO_{2} are formed, and also from PbS + O_{4} lead sulphate PbSO_{4},
and then PbO and PbSO_{4} react with the remaining PbS, according to the
equations 2PbO + PbS = 3Pb + SO_{2} and also PbSO_{4} + PbS = 2Pb +
2SO_{2}.[44]
[42] Mixed ores of copper compounds together with PbS and ZnS are
frequently found in the most ancient primary rocks. As the
separation of the metals themselves is difficult, the ores are
separated by a method of selection or mechanical sorting. Such
mixed ores occur in Russia, in many parts of the Caucasus, and in
the Donetz district (at Nagolchik).
[42 bis] Lead sulphide in the presence of zinc and hydrochloric acid is
completely reduced to metallic lead, all the sulphur being given
off as hydrogen sulphide.
[43] Lead sulphate, PbSO_{4}, occurs in nature (_anglesite_) in
transparent brilliant crystals which are isomorphous with barium
sulphate, and have a specific gravity of 6·3. The same salt is
formed on mixing sulphuric acid or its soluble salts with
solutions of lead salts, as a heavy white precipitate, which is
insoluble in water and acids, but dissolves in a solution of
ammonium tartrate in the presence of an excess of ammonia. This
test serves to distinguish this salt from the similar salts of
strontium and barium.
[44] According to J. B. Hannay (1894) the last named decomposition
(PbS + PbSO_{4} = 2Pb + 2SO_{2}) is really much more complicated,
and in fact a portion of the PbS is dissolved in the Pb, forming a
slag containing PbO, PbS and PbSO_{4}, whilst a portion of the
lead volatilises with the SO_{2} in the form of a compound
PbS_{2}O_{2}, which is also formed in other cases, but has not yet
been thoroughly studied.
Besides these methods for extracting lead from PBS in its ores,
roasting (the removal of the S in the form of SO_{2}) and smelting
with charcoal with a blast in the same manner as in the
manufacture of pig iron (Chapter XXII.) are also employed.
We may add that PbS in contact with Zn and hydrochloric acid
(which has no action upon PbS alone) entirely decomposes, forming
H_{2}S and metallic lead: PbS + Zn + 2HCl = Pb + ZnCl_{2} +
H_{2}S.
As lead is easily reduced from its ores, and the ore itself has a
metallic appearance, it is not surprising that it was known to the
ancients, and that its properties were familiar to the alchemists,
who called it 'Saturn.' Hence metallic lead, reduced from its
salts in solution by zinc, having the appearance of a tree-like
mass of crystals, is called 'arbor saturni,' &c.
The appearance of lead is well known; its specific gravity is 11·3; the
bluish colour and well-marked metallic lustre of freshly-cut lead quickly
disappear when exposed to the air, because it becomes coated with a
layer--although a very thin layer--of oxide and salts formed by the
moisture and acids in the atmosphere. It melts at 320°, and crystallises
in octahedra on cooling. Its softness is apparent from the flexibility of
lead pipes and sheets, and also from the fact that it may be cut with a
knife, and also that it leaves a grey streak when rubbed on paper. On
account of its being so soft, lead naturally cannot be applied in many
cases where most metals may be used; but on the other hand it is a metal
which is not easily changed by chemical reagents, and as it is capable of
being soldered and drawn into sheets, &c., lead is most valuable for many
technical uses. Lead pipes are used for conveying water[45] and many
other liquids, and sheet lead is used for lining all kinds of vessels
containing liquids--(acids, for instance) which act on other metals. This
particularly refers to sulphuric and hydrochloric acids, because at a low
temperature they do not act on lead, and if they form lead sulphate,
PbSO_{4}, and chloride, PbCl_{2}, these salts being insoluble in water
and in acids, cover the lead and protect it from further corrosion.[46]
All soluble preparations of lead are poisonous. At a white heat lead may
be partially distilled; the vapours oxidise and burn. Lead may also be
easily oxidised at low temperatures. Lead only decomposes water at a
white heat, and does not liberate hydrogen from acids, with the exception
only of very strong hydrochloric acid and then only when boiling.
Sulphuric acid diluted with water does not act on it, or only acts very
feebly at the surface; but strong sulphuric acid, when heated, is
decomposed by it, with the evolution of sulphurous anhydride. The best
solvent for lead is nitric acid, which transforms it into a soluble salt,
Pb(NO_{3})_{2}.
[45] Freshly laid new lead pipes contaminate the water with a certain
amount of lead salts, arising from the presence of oxygen,
carbonic acid, &c., in the water. But the lead pipes under the
action of running water soon become coated with a film of
salts--lead sulphate, carbonate, chloride, &c.--which are
insoluble in water, and the water pipes then become harmless.
[46] Lead is used in the arts, and owing to its considerable density,
it is cast, mixed with small quantities of other metals, into
shot. A considerable amount is employed (together with mercury) in
extracting gold and silver from poor ores, and in the manufacture
of chemical reagents, and especially of lead chromate. _Lead
chromate_, PbCrO_{4}, is distinguished for its brilliant yellow
colour, owing to which it is employed in considerable quantities
as a dye, mainly for dyeing cotton tissues yellow. It is formed on
the tissue itself, by causing a soluble salt of lead to react on
potassium chromate. Lead chromate is met with in nature as 'red
lead ore.' It is insoluble in water and acetic acid, hut it
dissolves in aqueous potash. The so-called pewter vessels often
consist of an alloy of 5 parts of tin and 1 part of lead, and
solder is composed of 1 to 2 parts of tin with 1/2 part of lead.
Amongst the alloys of lead and tin, Rudberg states that the alloy
PbSn_{3} stands out from the rest, since, according to his
observations, the temperature of solidification of the alloy is
187°.
Although acids thus have directly but little effect on lead, and this is
one of its most important practical properties, _yet when air has free
access, lead (like copper) very easily reacts with many acids_, even with
those which are comparatively feeble. The action of acetic acid on lead
is particularly striking and often applied in practice. If lead be
plunged into acetic acid it does not change at all and does not pass into
solution, but if part of the lead be immersed in the acid, and the other
part remain in contact with the air, or if lead be merely covered with a
thin layer of acetic acid in such a way that the air is practically in
contact with the metal, then it unites with the oxygen of the air to form
oxide, which combines with the acetic acid and forms lead acetate,
soluble in water. The formation of lead oxide is especially marked from
the fact that with a sufficient quantity of air not only is the normal
lead acetate formed but also the basic salts.[47]
[47] The normal lead acetate, known in trade as _sugar of lead_, owing
to its having a sweetish taste, has the formula
Pb(C_{2}H_{3}O_{2})_{2},3H_{2}O. This salt only crystallises from
acid solutions. It is capable of dissolving a further quantity of
lead oxide or of metallic lead in the presence of air. A basic
salt of the composition Pb(C_{2}H_{3}O_{2})_{2},PbH_{2}O_{2} is
then formed which is soluble in water and alcohol. As in this salt
the number of atoms is even and the same as in the hydrate of
acetic acid, C_{2}H_{4}O_{2},H_{2}O = C_{2}H_{3}(OH)_{3}, it may
be represented as this hydrate in which two of hydrogen are
replaced by lead--that is, as C_{2}H_{3}(OH)(O_{2}Pb). This basic
salt is used in medicine as a remedy for inflammation, for
bandaging wounds, &c., and also in the manufacture of white lead.
Other basic acetates of lead, containing a still greater amount of
lead oxide, are known. According to the above representation of
the composition of the preceding lead acetate, a basic salt of the
composition (C_{2}H_{3})_{2}(O_{2}Pb)_{3} would be also possible,
but what appear to be still more basic salts are known. As the
character of a salt also depends on the property of the base from
which it is formed, it would seem that lead forms a hydroxide of
the composition HOPbOH, containing two water residues, one or both
of which may be replaced by the acid residues. If both water
residues are replaced, a normal salt, XPbX, is obtained, whilst if
only one is replaced a basic salt, XPbOH, is formed. But lead does
not only give this normal hydroxide, but also polyhydroxides,
Pb(OH),_n_PbO, and if we may imagine that in these polyhydroxides
there is a substitution of both the water residues by acid
residues, then the power of lead for forming basic salts is
explained by the properties of the base which enters into their
composition.
When oxidising in the presence of air,[48] when heated or in the
presence of an acid at the ordinary temperature, lead forms compounds of
the type PbX_{2}. _Lead oxide_, PbO, known in industry as _litharge_,
silberglätte (this name is due to the fact that silver is extracted from
the lead ores of this kind) and massicot. If the lead is oxidised in air
at a high temperature, the oxide which is formed fuses, and on cooling is
easily obtained in fused masses which split up into scales of a yellowish
colour, having a specific gravity of 9·3; in this form it bears the name
of litharge. Litharge is principally used for making lead salts, for the
extraction of metallic lead, and also for the preparation of drying
oils--for instance, from linseed oil.[49] When oxidised carefully and
slightly heated, lead forms a powdery (not fused) oxide known under the
name of _massicot_. It is best prepared in the laboratory by heating lead
nitrate, or lead hydroxide. It has a yellow colour, and differs from
litharge in the greater difficulty with which it forms lead salts with
acids. Thus, for instance, when massicot is moistened with water it does
not attract the carbonic acid of the air so easily as litharge does. It
may, however, be imagined that the cause of the difference depends only
on the formation of dioxide on the surface of the lead oxide, on which
the acids do not act. In any case lead oxide is comparatively easily
soluble in nitric and acetic acids. It is but slightly soluble in water,
but communicates an alkaline reaction to it, since it forms the
hydroxide. This hydroxide is obtained in the shape of a white precipitate
by the action of a small quantity of an alkali hydroxide on a solution of
a lead salt. An excess of alkali dissolves the hydroxide separated, which
fact demonstrates the comparatively indistinct basic properties of lead
oxide. The normal lead hydroxide, which should have the composition
Pb(OH)_{2}, is unknown in a separate state, but it is known in
combination with lead oxide as Pb(OH)_{2},2PbO or Pb_{3}O_{2}(OH)_{2}.
The latter is obtained in the form of brilliant, white, octahedral
crystals when basic lead acetate is mixed with ammonia and gently heated.
The basic qualities of this hydroxide are shown distinctly by its
absorbing the carbonic anhydride of the air. When an alkaline solution of
the hydroxide is boiled, it deposits lead oxide in the form of a
crystalline powder.
[48] Few compounds are known of the lower type PbX, and still fewer of
the intermediate type PbX_{3}. To the first type belongs the
so-called lead suboxide, Pb_{2}O, obtained by the ignition of lead
oxalate, C_{2}PbO_{4}, without access of air. It is a black
powder, which easily breaks up under the action of acids, and even
by the simple action of heat, into metallic lead and lead oxide.
This is the character of all suboxides. They cannot be regarded as
independent salt-forming oxides, neither can those forms of
oxidation of lead which contain more oxygen than the oxide of
lead, PbO, and less than the dioxide, PbO_{2}. As we shall see, at
least two such compounds are formed. Thus, for example, an oxide
having the composition Pb_{2}O_{3} is known, but it is decomposed
by the action of acids into lead oxide, which passes into
solution, and lead dioxide, which remains behind. Such is red
lead. (See further on.)
[49] In the boiling of drying oils, the lead oxide partially passes
into solution, forming a saponified compound capable of attracting
oxygen and solidifying into a tar-like mass, which forms the oil
paint. Perhaps, however, glycerine partially acts in the process.
Ossovetsky by saturating drying oil with the salts of certain
metals obtained oil colours of great durability.
A mixture of very finely-divided litharge with glycerine (50 parts
of litharge to 5 c.c. of anhydrous glycerine) forms a very quick
(two minutes) setting cement, which is insoluble in water and
oils, and is very useful in setting up chemical apparatus. The
hardening is based on the reaction of the lead oxide with
glycerine (Moraffsky).
Lead oxide forms but few soluble salts--for instance, the nitrate and the
acetate. The majority of its salts (sulphate, PbSO_{4}; carbonate,
PbCO_{3}; iodide, PbI_{2}, &c.) are insoluble in water. These salts are
colourless or light yellow if the acid be colourless. In lead oxide _the
faculty of forming basic salts_, PbX_{2}_n_PbO or PbX_{2}_n_PbH_{2}O_{2},
is strongly developed. A similar property was observed in magnesium and
also in the salts of mercury, but lead oxide forms basic salts with still
greater facility, although double salts are in this case more rarely
formed.[50]
[50] It is very instructive to observe that lead not only easily forms
basic salts, but also salts containing several acid groups. Thus,
for example, lead carbonate occurs in nature and forms compounds
with lead chloride and sulphate. The first compound, known as
_corneous lead_, _phosgenite_, has the composition
PbCO_{3},PbCl_{2}; it occurs in nature in bright cubical crystals,
and is prepared artificially by simply boiling lead chloride with
lead carbonate. A similar compound of normal salts,
PbSO_{4},PbCO_{3}, occurs in nature as _lanarkite_ in monoclinic
crystals. _Leadhillite_ contains PbSO_{4},3PbCO_{3}, and also
occurs in yellowish, monoclinic, tabular crystals. We will turn
our attention to these salts of lead, because it is very probable
that their formation is allied to the formation of the basic
salts, and the following considerations may lead to the
explanation of the existence of both. In describing silica we
carefully developed the conception of polymerisation, which it is
_also indispensable to recognise in the composition of many other
oxides_. Thus it may be supposed that PbO_{2} is a similar
polymerised compound to SiO_{2}--_i.e._ that the composition of
lead peroxide will be Pb_{_n_}O_{2_n_}, because lead methyl,
PbMe_{4}, and lead ethyl, PbEt_{4}, are volatile compounds, whilst
PbO_{2} is non-volatile, and is very like silica in this respect,
and not in the least like carbonic anhydride. Still more should a
polymeric structure, Pb_{_n_}O_{_n_}, be ascribed to lead oxide,
since it differs as little from lead dioxide in its physical
properties as carbonic oxide does from carbonic anhydride, and
being an unsaturated compound is more likely to be capable of
intercombination (polymerisation) than lead dioxide. These
considerations respecting the complexity of lead oxide could have
no real significance, and could not be accepted, were it not for
the existence of the above-mentioned basic and mixed salts. The
oxide apparently corresponds with the composition
Pb_{_n_}X_{2_n_}, and since, according to this representation, the
number of X's in the salts of lead is considerable, it is obvious
that they may be diverse. When a part of these X's is replaced by
the water residue (OH) or by oxygen, X_{2} = O, and the other
parts by an _acid residue_, X, then basic salts are obtained, but
if a part of the X's is replaced by acid residues of one kind, and
the other part by acid residues of another kind, then those mixed
salts about which we are now speaking are formed. Thus, for
example, we may suppose, for a comparison of the composition of
the majority of the salts of lead, that _n_ = 12, and then the
above-mentioned compounds will present themselves in the following
form:--Lead oxide, Pb_{12}O_{12}, its crystalline hydrate,
Pb_{12}O_{8}(OH)_{8}, lead chloride, Pb_{12}Cl_{24}, lead
oxychloride, Pb_{12}Cl_{12}O_{6}, the other oxychloride,
Pb_{12}(OH)_{8}Cl_{6}O_{6}, mendipite (_see_ Note 51),
Pb_{12}Cl_{8}O_{8}, normal lead carbonate, Pb_{12}(CO_{3})_{12},
crystalline basic salt, Pb_{12}(OH)_{6}(CO_{3})_{6}, white lead,
Pb_{12}(CO_{3})_{8}(HO)_{8}, corneous lead,
Pb_{12}Cl_{12}(CO_{3})_{6}, lanarkite,
Pb_{12}(CO_{3})_{6}(SO_{4})_{6}, leadhillite,
Pb_{12}(CO_{3})_{9}(SO_{4})_{3}, &c. The number 12 is only taken
to avoid fractional quantities. Possibly the polymerisation is
much higher than this. The theory of the polymerisation of oxides
introduced by me in the first edition of this work (1869) is now
beginning to be generally accepted.
Amongst the soluble lead salts, that best known and most often applied in
practical chemistry is _lead nitrate_, obtained directly by dissolving
lead or its oxide in nitric acid. The normal salt, Pb(NO_{3})_{2},
crystallises in octahedra, dissolves in water, and has a specific gravity
of 4·5. When a solution of this salt acts on white lead or is boiled with
litharge, the basic salt, having a composition Pb(OH)(NO_{3}), is formed
in crystalline needles, sparingly soluble in cold water but easily
dissolved in hot water, and therefore in many respects resembling lead
chloride. When the nitrate is heated, either lead oxide is obtained or
else the oxide in combination with peroxide.
_Lead chloride_, PbCl_{2}, is precipitated from the soluble salts of
lead when a strong solution is treated with hydrochloric acid or a
metallic chloride. It is soluble in considerable quantities in hot water,
and therefore if the solutions be dilute or hot, the precipitation of
lead chloride does not occur, and if a hot solution be cooled, the salt
separates in brilliant prismatic crystals. It fuses when heated (like
silver chloride), but is insoluble in ammonia. This salt is sometimes met
with in nature, and when heated in air is capable of exchanging half its
chlorine for oxygen, forming the basic salt or lead oxychloride,
PbCl_{2}PbO, which may also be obtained by fusing PbCl_{2} and PbO
together. The reaction of lead chloride with water vapour leads to the
same conclusion, showing the feeble basic character of lead 2PbCl_{2} +
H_{2}O = PbCl_{2},PbO + 2HCl. When ammonia is added to an aqueous
solution of lead chloride a white precipitate is formed, which parts with
water on being heated, and has the composition Pb(OH)Cl,PbO. This
compound is also formed by the action of metallic chlorides on other
soluble basic salts of lead.[51]
[51] A similar basic salt having a white colour, and therefore used as
a substitute for white lead, is also obtained by mixing a solution
of basic lead acetate with a solution of lead chloride. Its
formation is expressed by the equation: 2PbX(OH),PbO + PbCl_{2} =
2Pb(OH)Cl,PbO + PbX_{2}. Similar basic compounds of lead are met
with in nature--for instance, _mendipite_, PbCl,2PbO, which
appears in brilliant yellowish-white masses. The ignition of red
lead with sal-ammoniac results in similar polybasic compounds of
lead chloride, forming the _Cassel's_, or _mineral yellow_ of the
composition PbCl_{2}_n_PbO. _Lead iodide_, PbI_{2}, is still less
soluble than the chloride, and is therefore obtained by mixing
potassium iodide with a solution of a lead salt. It separates as a
yellow powder, which may be dissolved in boiling water, and on
cooling separates in very brilliant crystalline scales of a golden
yellow colour. The salts PbBr_{2}, PbF_{2}, Pb(CN)_{2},
Pb_{2}Fe(CN)_{6} are also insoluble in water, and form white
precipitates.
Lead carbonate, or _white lead_, is the most extensively used basic lead
salt. It has the valuable property of 'covering,' which only to a certain
extent appertains to lead sulphate and other white powdery substances
used as pigments. This faculty of 'covering' consists in the fact that a
small quantity of white lead mixed with oil spreads uniformly, and if
such a mixture be spread over a surface (for instance, of wood or metal)
the surface is quickly covered--that is, light does not penetrate through
even a very thin layer of superposed white lead; thus, for example, the
grain of the wood remains invisible.[52] White lead, or _basic lead
carbonate_, after being dried at 120°, has a composition
Pb(OH)_{2},2PbCO_{3}.[53] It may be obtained by adding a solution of
sodium carbonate to a solution of one of the basic salts of lead--for
instance, the basic acetate--and likewise by treating this latter with
carbonic acid. For this purpose the solution of basic acetate is poured
into the vessel _f_; it is prepared in the vat A, containing litharge,
into which the pump P delivers the solution of the acetate, which remains
after the action of carbonic anhydride on the basic salt. In A a basic
salt is formed having a composition approaching to
Pb_{4}(OH)_{6}(C_{2}H_{3}O_{2})_{2}; carbonic anhydride, 2CO_{2}, is
passed through this solution and precipitates white lead,
Pb_{2}(OH)_{2}(CO_{3})_{2}, and normal lead acetate,
Pb(C_{2}H_{3}O_{2})_{2}, remains in the solution, and is pumped back into
the vat A containing lead oxide, where the normal salt is again (on being
agitated) converted into the basic salt. This is run into the vessel E,
and thence into _f_. Into the latter carbonic anhydride is delivered from
the generator D, and forms a precipitate of white lead.[53 bis]
[52] It is remarkable that a peculiar kind of attraction exists between
boiled linseed oil and white lead, as is seen from the following
experiments. White lead is triturated in water. Although it is
heavier than water, it remains in suspension in it for some time
and is thoroughly moistened by it, so that the trituration may be
made perfect; boiled linseed oil is then added, and shaken up with
it. A mixture of the oil and white lead is then found to settle at
the bottom of the vessel. Although the oil is much lighter than
the water it does not float on the top, but is retained by the
white lead and sinks under the water together with it. There is
not, however, any more perfect combination nor even any solution.
If the resultant mass be then treated with ether or any other
liquid capable of dissolving the oil, the latter passes into
solution and leaves the white lead unaltered.
[53] It may be regarded as a salt corresponding with the normal hydrate
of carbonic acid, C(OH)_{4}, in which three-quarters of the
hydrogen is replaced by lead. A salt is also known in which all
the hydrogen of this hydrate of carbonic acid is replaced by
lead--namely, the salt containing CO_{4}Pb_{2}. This salt is
obtained as a white crystalline substance by the action of water
and carbonic acid on lead. The normal salt, PbCO_{3}, occurs in
nature under the name of white lead ore (sp. gr. 6·47), in
crystals, isomorphous with aragonite, and is formed by the double
decomposition of lead nitrate with sodium carbonate, as a heavy
white precipitate. Thus both these salts resemble white lead, but
the first-named salt is exclusively used in practice, owing to its
being very conveniently prepared, and being characterised by its
great covering capacity, or 'body,' due to its fine state of
division.
[53 bis] One of the many methods by which white lead is prepared
consists in mixing massicot with acetic acid or sugar of lead, and
leaving the mixture exposed to air (and re-mixing from time to
time), containing carbonic acid, which is absorbed from the
surface by the basic salt formed. After repeated mixings (with the
addition of water), the entire mass is converted into white lead,
which is thus obtained very finely divided.
[Illustration: FIG. 82.--Manufacture of white lead.]
In order to mark the transition from lead oxide, PbO, into lead dioxide
PbO_{2} (plumbic anhydride), it is necessary to direct our attention to
the intermediate oxide, or _red lead_, Pb_{3}O_{4}.[54] In the arts it is
used in considerable quantities, because it forms a very durable
yellowish-red paint used for colouring the resins (shellac, colophony,
&c.) composing sealing wax. It also forms a very good cheap oil paint,
used especially for painting metals, more particularly because drying
oils--for instance, hemp seed, linseed oils--very quickly dry with red
lead and with lead salts. Red lead is prepared by slightly heating
massicot, for which purpose two-storied stoves are used. In the lower
story the lead is turned into massicot, and in the higher one, having the
lower temperature (about 300°), the massicot is transformed into red
lead. Frémy and others showed the instability of red lead prepared by
various methods, and its decomposition by acids, with formation of lead
dioxide, which is insoluble in acids, and a solution of the salts of lead
oxide. The artificial production (synthesis) of red lead by double
decomposition was most important. For this purpose Frémy mixed an
alkaline solution of potassium plumbate, K_{2}PbO_{3} (prepared by
dissolving the dioxide in fused potash),[54 bis] with an alkaline
solution of lead oxide. In this way a yellow precipitate of minium
hydrate is formed, which, when slightly heated, loses water and turns
into bright red anhydrous minium Pb_{3}O_{4}.
[54] If lead hydroxide be dissolved in potash and sodium hypochlorite
be added to the solution, the oxygen of the latter acts on the
dissolved lead oxide, and partially converts it into dioxide, so
that the so-called lead sesquioxide is obtained; its empirical
formula is Pb_{2}O_{3}. Probably it is nothing but a lead
salt--_i.e._ is referable to the type of dioxide of lead, or its
hydroxide, PbO(OH)_{2}, in which two atoms of hydrogen are
replaced by lead, PbO(O_{2}Pb). The brown compound precipitated by
the action of dilute acids--for example, nitric--splits up, even
at the ordinary temperature, into insoluble lead dioxide and a
solution of a lead salt. This compound evolves oxygen when it is
heated. It dissolves in hydrochloric acid, forming a yellow
liquid, which probably contains compounds of the composition
PbCl_{2} and PbCl_{4}, but even at the ordinary temperature the
latter soon loses the excess of chlorine, and then only lead
chloride, PbCl_{2}, remains. In order to see the relation between
red lead and lead sesquioxide, it must be observed that they only
differ by an extra quantity of lead oxide--that is, red lead is a
basic salt of the preceding compound, and if the compound
Pb_{2}O_{3} may be regarded as PbO_{3}Pb, then red lead should be
looked on as PbO_{3}Pb,PbO--that is, as basic lead plumbate.
[54 bis] Frémy obtained potassium plumbate in the following manner.
Pure lead dioxide is placed in a silver crucible, and a strong
solution of pure caustic potash is poured over it. The mixture is
heated and small quantities are removed from time to time for
testing, which consists in dissolving in a small quantity of water
and decomposing the resultant solution with nitric acid. There is
a certain moment during the heating when a considerable amount of
insoluble lead dioxide is precipitated on the addition of the
nitric acid; the solution then contains the salt in question, and
the heating must be stopped, and a small amount of water added to
dissolve the potassium plumbate formed. On cooling the salt
separates in somewhat large crystals, which have the same
composition as the stannate--that is, PbO(KO)_{2},3H_{2}O.
Minium is the first and most ordinary means of producing _lead
dioxide_, or plumbic anhydride, PbO_{2},[55] because when red lead is
treated with dilute nitric acid it gives up lead oxide, and PbO_{2}
remains, on which dilute nitric acid does not act. The composition of
minium is Pb_{3}O_{4}, and therefore the action of nitric acid on it is
expressed by the equation: Pb_{3}O_{4} + 4HNO_{3} = PbO_{2} +
2Pb(NO_{3})_{2} + 2H_{2}O. The dioxide may also be obtained by treating
lead hydroxide suspended in water with a stream of chlorine. Under these
conditions the chlorine takes up the hydrogen from the water, and the
oxygen passes over to the lead oxide.[56] When a strong solution of lead
nitrate is decomposed by the electric current, the appearance of
crystalline lead dioxide is also observed upon the positive pole; it is
also found in nature in the form of a black crystalline substance having
a specific gravity of 9·4. When artificially produced it is a fine dark
powder, resisting the action of acids, but nevertheless when treated with
strong sulphuric acid it evolves oxygen and forms lead sulphate, and with
hydrochloric acid it evolves chlorine. The oxidising property of lead
dioxide depends of course on the facility of its transition into the more
stable lead oxide, which is easily understood from the whole history of
lead compounds. In the presence of alkalis it transforms chromium oxide
into chromic acid, whilst lead chromate, PbCrO_{4}, is formed, remaining,
however, in solution, on account of its being soluble in caustic alkalis.
The oxidising action of lead dioxide on sulphurous anhydride is most
striking, as it immediately absorbs it, with formation of lead sulphate.
This is accompanied by a change of colour and development of heat,
PbO_{2} + SO_{2} = PbSO_{4}. When triturated with sulphur the mixture
explodes, the sulphur burning at the expense of the oxygen of the lead
dioxide. _Tetrachloride of lead_, PbCl_{4}, belongs to the same class of
lead compounds as PbO_{2}. This chloride is formed by the action of
strong hydrochloric acid upon PbO_{2}, or, in the cold, by passing a
stream of chlorine through water containing PbCl_{2} in suspension. The
resultant yellow solution gives off chlorine when heated. With a solution
of sal ammoniac (Nicolukin, 1885) it gives a precipitate of a double
salt, (NH_{4})_{2}PbCl_{6} (very slightly soluble in a solution of sal
ammoniac), which when treated with strong sulphuric acid (Friedrich,
1890) gives PbCl_{4} as a yellow liquid sp. gr. 3·18, which solidifies at
-18°, and when heated gives PbCl_{2} + Cl_{2}. It is not acted upon by
H_{2}SO_{4} like SnCl_{4}. Tetrafluoride of lead (Brauner) belongs to the
same class of compounds, it easily forms double salts and decomposes with
the evolution of fluorine (Chapter II., Note 49 bis).[56 bis]
[55] Lead dioxide is often called lead peroxide, but this name leads to
error, because PbO_{2} does not show the properties of true
peroxides, like hydrogen or barium peroxides, but is endowed with
acid properties--that is, it is able to form true salts with
bases, which is not the case with true peroxides. Lead dioxide is
a normal salt-forming compound of lead, as Bi_{2}O_{5} is for
bismuth, CeO_{2} for cerium, and TeO_{3} for tellurium, &c. They
all evolve chlorine when treated with hydrochloric acid, whilst
true peroxides form hydrogen peroxide. The true lead peroxide, if
it were obtained, would probably have the composition Pb_{2}O_{5},
or, in combination with peroxide of hydrogen, H_{2}Pb_{2}O_{7} =
H_{2}O_{2} + Pb_{2}O_{5}, judging from the peroxides corresponding
with sulphuric, chromic, and other acids, which we shall
afterwards consider.
As a proof of the fact, that the form PbO_{2}, or PbX_{4}, is the
highest normal form of any combination of lead, it is most
important to remark that it might be expected that the action of
lead chloride, PbCl_{2}, on zinc-ethyl, ZnEt_{2}, would result in
the formation of zinc chloride, ZnCl_{2}, and lead-ethyl,
PbEt_{2}, but that in reality the reaction proceeds otherwise.
Half of the lead is set free, and lead tetrethyl, PbEt_{4}, is
formed as a colourless liquid, boiling at about 200° (Butleroff,
Frankland, Buckton, Cahours, and others). The type PbX_{4} is not
only expressed in PbEt_{4} and PbO_{2}, but also in PbF_{4},
obtained by Brauner.
[56] According to Carnelley and Walker, the hydrate
(PbO_{2})_{3},H_{2}O is then formed; it loses water at 230°. The
anhydrous dioxide remains unchanged up to 280°, and is then
converted into the sesquioxide, Pb_{2}O_{3}, which again loses
oxygen at about 400°, and forms red lead, Pb_{3}O_{4}. Red lead
also loses oxygen at about 550°, forming lead oxide, PbO, which
fuses without change at about 600°, and remains constant as far as
the limit of the observations made (about 800°).
The best method for preparing pure lead dioxide consists in mixing
a hot solution of lead chloride with a solution of bleaching
powder (Fehrman).
[56 bis] The plumbates of Ca and other similar metals, mentioned in
Chapter III., Note 7, also belong to the form PbX_{4}.
Amongst the elements of the second and third groups it was observed that
the elements were more basic in the even than in the uneven series. It is
sufficient to remember calcium, strontium, and barium in the even, and
magnesium, zinc, and cadmium in the uneven series. In addition to this,
in the even series, as the atomic weight increases, in the same type of
oxidation the basic properties increase (the acid properties decrease);
for example, in the second group, calcium, strontium, barium. The same
also appears in the fourth and all the following groups. In the even
series of the fourth group titanium, zirconium, cerium, and thorium are
found. All their highest oxides, RO_{2}, even the lightest, titanic
oxide, TiO_{2}, have more highly developed basic properties than silica,
SiO_{2}, and in addition to this the basic properties are more distinctly
seen in zirconium dioxide, ZrO_{2}, than in titanic oxide, TiO_{2},
although the acid property of combining with bases still remains. In the
heaviest oxides, cerium dioxide, CeO_{2}, and thorium dioxide, ThO_{2},
no acid properties are observed, these being both purely basic oxides. In
Chapter XVII. (Note 43) we already pointed out this higher oxide of
cerium. As the above-mentioned elements are rather rare in nature, have
but little practical application, and do not present any new forms of
combination, it is unadvisable to dwell on them in this treatise.
_Titanium_ is found in nature in the form of its anhydride or oxide,
TiO_{2}, mixed with silicon in many minerals, but the oxide is also found
separately in the form of semi-metallic _rutile_ (sp. gr. 4·2). Another
titanic mineral is found as a mixture in other ores, known as _titanic
iron ore_ (in the Thuensky mountains of the southern Ural; it is known as
_thuenite_), FeTiO_{3}. This is a salt of ferrous oxide and titanic
anhydride. It crystallises in the rhombohedric system, has a metallic
lustre, grey colour, sp. gr. 4·5. The third mineral in which titanium is
found in considerable quantities in nature is _sphene_ or _titanite_,
CaTiSiO_{5} = CaO,SiO_{2},TiO_{2}, sp. gr. 3·5, colour yellow, green, or
the like, crystallises in tablets. The fourth, but rare, titanic mineral
is _peroffskite_, calcium titanate, CaTiO_{3}; it forms blackish-grey or
brown cubic crystals, sp. gr. 4·02, and occurs in the Ural and other
localities. It may be prepared artificially by fusing sphene in an
atmosphere of water vapour and carbonic anhydride. At the end of the last
century Klaproth showed the distinction between titanic compounds and all
others then known.[57]
[57] The compounds of titanium are generally obtained from rutile; the
finely-ground ore is fused with a considerable amount of acid
potassium sulphate, until the titanic anhydride, as a feeble base,
passes into solution. After cooling, the resultant mass is ground
up, dissolved in cold water, and treated with ammonium
hydrosulphide; a black precipitate then separates out from the
solution. This precipitate contains TiO_{2} (as hydrate) and
various metallic sulphides--for example, iron sulphide. It is
first washed with water and then with a solution of sulphurous
anhydride until it becomes colourless. This is due to the iron
sulphide contained in the precipitate, and rendering it black,
being converted into dithionate by the action of the sulphurous
acid. The titanic acid left behind is nearly pure. The
considerable volatility of titanium chloride may also be taken
advantage of in preparing the compounds of titanium from rutile.
It is formed by heating a mixture of rutile and charcoal in dry
chlorine; the distillate then contains _titanium chloride_,
TiCl_{4}. It may be easily purified, owing to its having a
constant boiling point of 136°. Its specific gravity is 1·76; it
is a colourless liquid, which fumes in the air, and is perfectly
soluble in water if it be not heated. When hot water acts on
titanic chloride, a large proportion of titanic acid separates out
from the solution and passes into metatitanic acid. A similar
decomposition of acid solutions of titanic acid is accomplished
whenever they are heated, and especially in the presence of
sulphuric acid, just as with metastannic acid, which titanic acid
resembles in many respects. On igniting the titanic acid a
colourless powder of the anhydride, TiO_{2}, is obtained. In this
form it is no longer soluble in acids or alkalis, and only fuses
in the oxyhydrogen flame; but, like silica, it dissolves when
fused with alkalis and their carbonates; as already mentioned, it
dissolves when fused with a considerable excess of acid potassium
sulphate--that is, it then reacts as a feeble base. This shows the
basic character of titanic anhydride; it has at once, although
feebly developed, both basic and acid properties. The fused mass,
obtained from titanic anhydride and alkali when treated with
water, parts with its alkali, and a residue is obtained of a
sparingly-soluble poly-titanate, K_{2}TiO_{3}_n_TiO_{2}. The
hydrate, which is precipitated by ammonia from the solutions
obtained by the fusion of TiO_{2} with acid potassium sulphate,
when dried forms an amorphous mass of the composition Ti(OH)_{4}.
But it loses water over sulphuric acid, gradually passing into a
hydrate of the composition TiO(OH)_{2}, and when heated it parts
with a still larger proportion of water; at 100° the hydrate
Ti_{2}O_{3}(OH)_{2} is obtained, and at 300° the anhydride itself.
The higher hydrate, Ti(OH)_{4}, is soluble in dilute acid, and the
solution may be diluted with water; but on boiling the sulphuric
acid solution (though not the solution in hydrochloric acid), all
the titanic acid separates in a modified form, which is, however,
not only insoluble in dilute acids, but even in strong sulphuric
acid. This hydrate has the composition Ti_{2}O_{3}(OH)_{2}, but
shows different properties from those of the hydrate of the same
composition described above, and therefore this modified hydrate
is called _metatitanic acid_. It is most important to note the
property of the ordinary gelatinous hydrate (that precipitated
from acid solutions by ammonia) of dissolving in acids, the more
so since silica does not show this property. In this property a
transition apparently appears between the cases of common solution
(based on a capacity for unstable combination) and the case of the
formation of a hydrosol (the solubility of germanium oxide,
GeO_{2}, perhaps presents another such instance). If titanium
chloride be added drop by drop to a dilute solution of alcohol and
hydrogen peroxide, and then ammonia be added to the resultant
solution, a yellow precipitate of _titanium trioxide_,
TiO_{3}H_{2}O, separates out, as Piccini, Weller, and Classen
showed. This substance apparently belongs to the category of true
peroxides.
Titanium chloride absorbs ammonia and forms a compound,
TiCl_{4},4NH_{3}, as a red-brown powder which attracts moisture
from the air and when ignited forms _titanium nitride_,
Ti_{3}N_{4}. Phosphuretted hydrogen, hydrocyanic acid, and many
similar compounds are also absorbed by titanium chloride, with the
evolution of a considerable amount of heat. Thus, for example, a
yellow crystalline powder of the composition TiCl_{4},2HCN is
obtained by passing dry hydrocyanic acid vapour into cold titanium
chloride. Titanium chloride combines in a similar manner with
cyanogen chloride, phosphorus pentachloride, and phosphorus
oxychloride, forming molecular compounds, for example
TiCl_{4},POCl_{3}. This faculty for further combination probably
stands in connection, on the one hand, with the capacity of
titanium oxide to give polytitanates, TiO(MO)_{2},_n_TiO_{2}; on
the other hand, it corresponds with the kindred faculty of stannic
chloride for the formation of poly-compounds (Note 41), and lastly
it is probably related to the remarkable behaviour of titanium
towards nitrogen. Metallic titanium, obtained as a grey powder by
reducing potassium titanofluoride, K_{2}TiF_{6}, (sp. gr. 3·55 K.
Hofman 1893), with iron in a charcoal crucible, combines directly
with nitrogen at a red heat. If titanic anhydride be ignited in a
stream of ammonia, all the oxygen of the titanic oxide is
disengaged, and the compound TiN_{2} is formed as a dark violet
substance having a copper-red lustre. A compound Ti_{5}N_{6} is
also known; it is obtained by igniting the compound Ti_{3}N_{4} in
a stream of hydrogen, and is of a golden-yellow colour with a
metallic lustre. To this order of compounds also belongs the
well-known and chemically historical compound known as _titanium
nitrocyanide_; its composition is Ti_{5}CN_{4}. This substance
appears as infusible, sometimes well-formed, cubical crystals of
sp. gr, 4·3, and having a red copper colour and metallic lustre;
it is found in blast furnace slag. It is insoluble in acids but is
acted on by chlorine at a red heat, forming titanium chloride. It
was at first regarded as metallic titanium; it is formed in the
blast furnace at the expense of those cyanogen compounds
(potassium cyanide and others) which are always present, and at
the expense of the titanium compounds which accompany the ores of
iron. Wöhler, who investigated this compound, obtained it
artificially by heating a mixture of titanic oxide with a small
quantity of charcoal, in a stream of nitrogen, and thus proved the
direct power for combination between nitrogen and titanium. When
fused with caustic potash, all the nitrogen compounds of titanium
evolve ammonia and form potassium titanate. Like metals they are
able to reduce many oxides--for example, oxides of copper--at a
red heat. Among the alloys of titanium, the crystalline compound
Al_{4}Ti is remarkable. It is obtained by directly dissolving
titanium in fused aluminium; its specific gravity is 3·11. The
crystals are very stable, and are only soluble in aqua regia and
alkalis.
The comparatively rare element _zirconium_, Zr = 90, is very similar to
titanium, but has a more basic character. It is rarer in nature than
titanium, and is found principally in a mineral called _zircon_,
ZrSiO_{4} = ZrO_{2}.SiO_{2}, crystallising in square prisms, sp. gr. 4·5.
It has considerable hardness and a characteristic brownish-yellow colour,
and is occasionally found in the form of transparent crystals, as a
precious stone called hyacinth.[58] Metallic zirconium was obtained, by
Berzelius and Troost, by the action of aluminium on potassium
zirconofluoride in the same way that silicon is prepared; it forms a
crystalline powder, similar in appearance to graphite and antimony, but
having a very considerable hardness, not much lustre, sp. gr. 4·15. In
many respects it resembles silicon; it does not fuse when heated, and
even oxidises with difficulty, but liberates hydrogen when fused with
potash. When fused with silica it liberates silicon. With carbon in the
electrical furnace it forms ZrC_{2}, with hydrogen it gives ZrH_{2} (like
CaH_{2}, Winkler, Vol. I., p. 621); hydrochloric and nitric acids act
feebly on it, but aqua regia easily dissolves it. It is distinguished
from silicon by the fact that hydrofluoric acid acts on it with great
facility, even in the cold and when diluted, whilst this acid does not
act on silicon at all.
[58] The formula ZrO was first given to the oxide of zirconium as a
base, in this case Zr = 45 whilst the present atomic weight is Zr
= 90--that is, the formula of the oxide is now recognised as being
ZrO_{2}. The reasons for ascribing this formula to the compounds
of zirconium are as follows. In the first place, the investigation
of the crystalline forms of the zirconofluorides--for example,
K_{2}ZrF_{6}, MgZrF_{6},5H_{2}O--which proved to be analogous in
composition and crystalline form with the corresponding compounds
of titanium, tin, and silicon. In the second place, the specific
heat of Zr is 0·067, which corresponds with the combining weight
90. The third and most important reason for doubling the combining
weight of zirconium was given by Deville's determination of the
vapour density of _zirconium chloride_, ZrCl_{4}. This substance
is obtained by igniting zirconium oxide mixed with charcoal in a
stream of dry chlorine, and is a colourless, saline substance
which is easily volatile at 440°. Its density referred to air was
found to be 8·15, that is 117 in relation to hydrogen, as it
should be according to the molecular formula of this substance
above-cited. It exhibits, however, in many respects, a saline
character and that of an acid chloranhydride, for zirconium oxide
itself presents very feebly developed acid properties but clearly
marked basic properties. Thus zirconium chloride dissolves in
water, and on evaporation the solution only partially disengages
hydrochloric acid--resembling magnesium chloride, for example.
Zirconium was discovered and characterised as an individual
element by Klaproth.
Pure compounds of zirconium are generally prepared from zircon,
which is finely ground, but as it is very hard it is first heated
and thrown into cold water, by which means it is disintegrated.
Zircon is decomposed or dissolved when fused with acid potassium
sulphate, or still more easily when fused with acid potassium
fluoride (a double soluble salt, K_{2}ZrF_{6}, is then formed);
however, zirconium compounds are generally prepared from powdered
zircon by fusing it with sodium carbonate and then boiling in
water. An insoluble white residue is obtained consisting of a
compound of the oxides of sodium and zirconium, which is then
treated with hydrochloric acid and the solution evaporated to
dryness. The silica is thus converted into an insoluble form, and
zirconium chloride obtained in solution. Ammonia precipitates
_zirconium hydroxide_ from this solution, as a white gelatinous
precipitate, ZrO(OH)_{2}. When ignited this hydroxide loses water
and in so doing undergoes a spontaneous recalescence and leaves a
white infusible and exceedingly hard mass of _zirconium oxide_,
ZrO_{2}, having a specific gravity of 5·4 (in the electrical
furnace ZrO_{2} fuses and volatilises like SiO_{2}, Moissan).
Owing to its infusibility, zirconium oxide is used as a substitute
for lime and magnesia in the Drummond light. This oxide, in
contradistinction to titanium oxide, is soluble, even after
prolonged ignition, in hot strong sulphuric acid. The hydroxide is
easily soluble in acids. The composition of the salts is ZrX_{4},
or ZrOX_{2}, or ZrOX_{2},ZrO_{2}, just as with those of its
analogues. But although zirconium oxide forms salts in the same
way with acids, it also gives salts with bases. Thus it liberates
carbonic anhydride when fused with sodium carbonate, forming the
salts Zr(NaO)_{4}, ZrO(NaO)_{2}, &c. Water, however, destroys
these salts and extracts the soda.
The very similar element _thorium_ (Th = 232) was distinguished by
Berzelius from zirconium. It is very rarely met with, in _thorite_ and
_orangeite_, ThSiO_{4},2H_{2}O. The latter is isomorphous with zircon
(sp. gr. 4·8).[59]
[59] Thorium has also been found in the form of oxide in certain
pyrochlores, euxenites, monazites, and other rare minerals
containing salts of niobium and phosphates. The compounds of
thorium are prepared by decomposing thorite or orangeite with
strong sulphuric acid at its boiling point; this renders the
silica insoluble, and the thorium oxide passes into solution when
the residue is treated with cold water, after having been
previously boiled with water (boiling water does not dissolve the
oxide of thorium). Lead and other impurities are separated by
passing sulphuretted hydrogen through the solution, and the
thorium hydroxide is then precipitated by ammonia. If this
hydroxide be dissolved in the smallest possible amount of
hydrochloric acid, and oxalic acid be then added, thorium oxalate
is obtained as a white precipitate, which is insoluble in an
excess of oxalic acid; this reaction is taken advantage of for
separating this metal from many others. It, however, resembles the
cerite metals (Chapter XVII., Note 43) in this and many other
respects. The thorium hydroxide is gelatinous; on ignition it
leaves an infusible oxide, ThO_{2}, which, when fused with borax,
gives crystals of the same form as stannic oxide or titanic
anhydride; sp. gr. 9·86. But the basic properties are much more
developed in thorium oxide than in the preceding oxides, and it
does not even disengage carbonic acid when fused with sodium
carbonate--that is, it is a much more energetic base than
zirconium oxide. The hydrate, ThO_{2}, however, is soluble in a
solution of Na_{2}CO_{3} (Chapter XVII., Note 43). Thorium
chloride, ThCl_{4} is obtained as a distinctly crystalline
sublimate when thorium oxide, mixed with charcoal, is ignited in a
stream of dry chlorine. When heated with potassium, thorium
chloride gives a metallic powder of thorium having a sp. gr. 11·1.
It burns in air, and is but slightly soluble in dilute acids. The
atomic weight of thorium was established by Chydenius and
Delafontaine on the basis of the ismorphism of the double
fluorides.
CHAPTER XIX
PHOSPHORUS AND THE OTHER ELEMENTS OF THE FIFTH GROUP
Nitrogen is the lightest and most widely distributed representative of
the elements of the fifth group, which form a higher saline oxide of the
form R_{2}O_{5}, and a hydrogen compound of the form RH_{3}. Phosphorus,
arsenic, bismuth, and antimony belong to the uneven series of this group.
_Phosphorus_ is the most widely distributed of these elements. There is
hardly any mineral substance composing the mass of the earth's crust
which does not contain some--it may be a small--amount of phosphorus
compounds in the form of the salts of phosphoric acid. The soil and
earthy substances in general usually contain from one to ten parts of
phosphoric acid in 10,000 parts. This amount, which appears so small,
has, however, a very important significance in nature. No plant can
attain its natural growth if it be planted in an artificial soil
completely free from phosphoric acid. Plants equally require the presence
of potash, magnesia, lime, and ferric oxide, among basic, and of
carbonic, sulphuric, nitric, and phosphoric anhydrides, among acid
oxides. In order to increase the fertility of a more or less poor soil,
the above-named nutritive elements are introduced into it by means of
fertilisers. Direct experiment has proved that these substances are
undoubtedly necessary to plants, but that they must be all present
simultaneously and in small quantities, and that an excess, like an
insufficiency, of one of these elements is necessarily followed by a bad
harvest, or an imperfect growth, even if all the other conditions (light,
heat, water, air) are normal. The phosphoric compounds of the soil
accumulated by plants pass into the organism of animals, in which these
substances are assimilated in many instances in large quantities. Thus
the chief component part of bones is calcium phosphate, Ca_{3}P_{2}O_{8},
and it is on this that their hardness depends.[1]
[1] Dry bones contain about one-third of gelatinous matter and about
two-thirds of ash, chiefly calcium phosphate. The salts of
phosphoric acid are also found in the mass of the earth as separate
minerals; for example, the _apatites_ contain this salt in a
crystalline form, combined with calcium chloride or fluoride,
CaR_{2},3Ca_{3}(PO_{4})_{2}, where R = F or Cl, sometimes in a
state of isomorphous mixture. This mineral often crystallises in
fine hexagonal prisms; sp. gr. 3·17 to 3·22. Vivianite is a
hydrated ferrous phosphate, Fe_{3}(PO_{4})_{2},8H_{2}O. Phosphates
of copper are frequently found in copper mines; for example,
_tagilite_, Cu_{3}(PO_{4})_{2},Cu(OH)_{2},2H_{2}O. Lead and
aluminium form similar salts. They are nearly all insoluble in
water. The turquoise, for instance, is hydrated phosphate of
alumina, (Al_{2}O_{3})_{2},P_{2}O_{5}5H_{2}O, coloured with a salt
of copper. Sea and other waters always contain a small amount of
phosphates. The ash of sea-plants, as well as of land-plants,
always contains phosphates. Deposits of calcium phosphate are often
met with; they are termed _phosphorites_ and _osteolites_, and are
composed of the fossil remains of the bones of animals; they are
used for manure. Of the same nature are the so-called guano
deposits from Baker's Island, and entire strata in Spain, France,
and in the Governments of Orloff and Kursk in Russia. It is evident
that if a soil destined for cultivation contain very little
phosphoric acid, the fertilisation by means of these minerals will
be beneficial, but, naturally, only if the other elements necessary
to plants be present in the soil.
Phosphorus was first extracted by Brand in 1669, by the ignition of
evaporated urine. After the lapse of a century Scheele, who knew of the
existence of a more abundant source of phosphorus in bones, pointed out
the method which is now employed for the extraction of this element.
Calcium phosphate in bones permeates a nitrogenous organic substance,
which is called ossein, and forms a gelatin. When bones are treated
exclusively for the extraction of phosphorus, neglecting the gelatin,
they are burnt, in which case all the ossein is burnt away. When,
however, it is desired to preserve the gelatin, the bones are immersed in
cold dilute hydrochloric acid, which dissolves the calcium phosphate and
leaves the gelatin untouched; calcium chloride and acid calcium
phosphate, CaH_{4}(PO_{4})_{2}, are then obtained in the solution. When
the bones are directly burnt in an open fire their mineral components
only are left as an ash, containing about 90 per cent. of calcium
phosphate, Ca_{3}(PO_{4})_{2}, mixed with a small amount of calcium
carbonate and other salts. This mass is treated with sulphuric acid, and
then the same substance is obtained in the solution as was obtained from
the unburnt bones immersed in hydrochloric acid--_i.e._ the acid calcium
phosphate soluble in water, in which reaction naturally the chief part of
the sulphuric acid is converted into calcium sulphate:
Ca_{3}(PO_{4})_{2} + 2H_{2}SO_{4} = 2CaSO_{4} + CaH_{4}(PO_{4})_{2}.
Ca_{3}(PO_{4})_{2} + 4HCl = 2CaCl_{2} + CaH_{4}(PO_{4})_{2}.
On evaporating the solution, crystallisable acid calcium phosphate is
obtained. The extraction of the phosphorus from this salt consists in
_heating it with charcoal to a white heat_. When heated, the acid
phosphate, CaH_{4}(PO_{4})_{2}, first parts with water, and forms the
metaphosphate, Ca(PO_{3})_{2}, which for the sake of simplicity may be
regarded, like the acid salt, as composed of pyrophosphate and phosphoric
anhydride, 2Ca(PO_{3})_{2} = Ca_{2}P_{2}O_{7} + P_{2}O_{5}. The latter,
with charcoal, gives phosphorus and carbonic oxide, P_{2}O_{5} + 5C =
P_{2} + 5CO. So that in reality a somewhat complicated process takes
place here, yielding ultimately products according to the following
equation:
2CaH_{4}(PO_{4})_{2} + 5C = 4H_{2}O + Ca_{2}P_{2}O_{7} + P_{2} + 5CO.
After the steam has come over, phosphorus and carbonic oxide distil over
from the retort and calcium pyrophosphate remains behind.[1 bis]
[1 bis] By subjecting the pyrophosphate to the action of sulphuric or
hydrochloric acid it is possible to obtain a fresh quantity of the
acid salt from the residue, and in this manner to extract all the
phosphorus. It is usual to take burnt bones, but mineral
phosphorites, osteolites, and apatites may also be employed as
materials for the extraction of phosphorus. Its extraction for the
manufacture of matches is everywhere extending, and in Russia, in
the Urals, in the Government of Perm, it has attained such
proportions that the district is able to supply other countries
with phosphorus. A great many methods have been proposed for
facilitating the extraction of phosphorus, but none of them differ
essentially from the usual one, because the problem is dependent on
the liberation of phosphoric acid by the action of acids, and on
its ultimate reduction by charcoal. Thus the calcium phosphate may
be mixed directly with charcoal and sand, and phosphorus will be
liberated on heating the mixture, because the silica displaces the
phosphoric anhydride, which gives carbonic oxide and phosphorus
with the charcoal. It has also been proposed to pass hydrochloric
acid over an incandescent mixture of calcium phosphate and
charcoal; the acid then acts just as the silica does, liberating
phosphoric anhydride, which is reduced by the charcoal. It is
necessary to prevent the access of air in the condensation of the
vapours of phosphorus, because they take fire very easily; hence
they are condensed under water by causing the gaseous products to
pass through a vessel full of water. For this purpose the condenser
shown in fig. 83 is usually employed.
[Illustration: FIG. 83.--Preparation of phosphorus. The mixture is
calcined in the retort _c_. The vapours of phosphorus pass through _a_
into water without coming into contact with air. The phosphorus condenses
in the water, and the gases accompanying it escape through _i_.]
As phosphorus melts at about 40°, it condenses at the bottom of the
receiver in a molten liquid mass, which is cast under water in tubes, and
is sold in the form of sticks. This is common or _yellow phosphorus_. It
is a transparent, yellowish, waxy substance, which is not brittle, almost
insoluble in water, and easily undergoes change in its external
appearance and properties under the action of light, heat, and of various
substances. It crystallises (by sublimation or from its solution in
carbon bisulphide) in the regular system, and[2] (in contradistinction to
the other varieties) is easily soluble in carbon bisulphide, and also
partially in other oily liquids. In this it recalls common sulphur. Its
specific gravity is 1·84. It fuses at 44°, and passes into vapour at
290°; it is easily inflammable, and must therefore be handled with great
caution; careless rubbing is enough to cause phosphorus to ignite. Its
application in the manufacture of matches is based on this.[2 bis] It
emits light in the air owing to its slow[3] oxidation, and is therefore
kept under water (such water is phosphorescent in the dark, like
phosphorus itself). It is also very easily oxidised by various oxidising
agents and takes up the oxygen from many substances.[3 bis] Phosphorus
enters into direct combination with many metals and with sulphur,
chlorine, &c., with development of a considerable amount of heat. It is
very poisonous although not soluble in water.
[2] Vernon (1891) observed that ordinary (yellow) phosphorus is
dimorphous. If it be melted and by careful cooling be brought in a
liquid form to as low a temperature as possible, it gives a variety
which melts at 45°·3 (the ordinary variety fuses at 44°·3), sp. gr.
1·827 (that of the ordinary variety is 1·818) at 13°, crystallises
in rhombic prisms (instead of in forms belonging to the cubical
system). This is similar to the relation between octahedral and
prismatic sulphur (Chapter XX.).
[2 bis] According to Herr Irinyi (an Hungarian student), the first
phosphorus matches were made in Austria at Roemer's works in 1835.
[3] The absorption of the oxygen of the atmosphere at a constant
ordinary temperature by a large surface of phosphorus proceeds so
uniformly, regularly, and rapidly, that it may serve, as Ikeda
(Tokio, 1893) has shown, for demonstrating the law of the velocity
(rate) of reaction, which is considered in theoretical chemistry,
and shows that the rate of reaction is proportional to the active
mass of a substance--_i.e._ _dx_/_dt_ = _k_(A - _x_) where _t_ is
the time, A the initial mass of the reacting substance--in this
case oxygen--_x_ the amount of it which has entered into reaction,
and _k_ the coefficient of proportionality. Ikeda took a test-tube
(diameter about 10 mm.), and covered its outer surface with a
coating of phosphorus (by melting it in a test-tube of large
diameter, inserting the smaller test-tube, and, when the phosphorus
had solidified, breaking away the outer test-tube), and introduced
it into a definite volume of air, contained in a Woulfe's bottle
(immersed in a water bath to maintain a constant temperature), one
of whose orifices was connected with a mercury manometer showing
the fall of pressure, _x_. Knowing that the initial pressure of the
oxygen (in air nearly 750 × ·0209) was about 155 mm. = A, the
coefficient of the rate of reaction _k_ is given, by the law of the
variation of the rate of reaction with the mass of the reacting
substance, by the equation: _k_ = (1/_t_)log(A/(A - _x_)), where
_t_ is the time, counting from the commencement, of the experiment
in minutes. When the surface of the phosphorus was about 11 sq.
cm., the following results were actually obtained.
_t_ = 10 20 30 40 50 60 minutes
_x_ = 10·5 21·5 31·1 40·7 49·1 57·3 mm
10,000 _k_ = 32 32 32 33 33 33
The constancy of _k_ is well shown in this case. The determination
takes a comparatively short time, so that it may serve as a lecture
experiment, and demonstrates one of the most important laws of
chemical mechanics.
[3 bis] Not only do oxidising agents like nitric, chromic, and similar
acids act upon phosphorus, but even the alkalis are attacked--that
is, phosphorus acts as a reducing agent. In fact it reduces many
substances, for instance, copper from its salts. When phosphorus is
heated with sodium carbonate, the latter is partially reduced to
carbon. If phosphorus be placed under water slightly warmed, and a
stream of oxygen be passed over it, it will burn under the water.
Besides this, there is a red variety of phosphorus, which differs
considerably from the above. _Red phosphorus_ (sometimes wrongly called
_amorphous phosphorus_) is partially formed when ordinary phosphorus
remains exposed to the action of light for a long time. It is also formed
in many reactions; for example, when ordinary phosphorus combines with
chlorine, bromine, iodine, or oxygen, a portion of it is converted into
red phosphorus. Schrötter, in Vienna, investigated this variety of
phosphorus, and pointed out by what methods it may be produced in
considerable quantities. Red phosphorus is a powdery red-brown opaque
substance of specific gravity 2·14. It does not combine so energetically
with oxygen and other substances as yellow phosphorus, and evolves less
heat in combining with them.[4] Common phosphorus easily oxidises in the
air; red phosphorus does not oxidise at all at the ordinary temperature;
hence it does not phosphoresce in the air, and may be very conveniently
kept in the form of powder. It does not, like yellow phosphorus, fuse at
44°. After being converted into vapour at 290° or 300°, it again passes
into the ordinary variety when slowly cooled. Red phosphorus is not
soluble in carbon bisulphide and other oily liquids, which permits of its
being freed from any admixture of the ordinary phosphorus. It is not
poisonous, and is used in many cases for which the ordinary phosphorus is
unsuitable or dangerous; for example, in the manufacture of matches,
which are then not poisonous or inflammable by accidental friction, and
therefore the red variety has now replaced the ordinary
phosphorus.[4 bis]
[4] The thermochemical determinations for phosphorus and its compounds
date from the last century, when Lavoisier and Laplace burnt
phosphorus in oxygen in an ice calorimeter. Andrews, Despretz,
Favre, and others have studied the same subject. The most accurate
and complete data are due to Thomsen. To determine the heat of
combustion of yellow phosphorus, Thomsen oxidised it in a
calorimeter with iodic acid in the presence of water, and a mixture
of phosphorous and phosphoric acids was thus formed (was not any
hypophosphoric acid formed?--Salzer), and the iodic acid converted
into hydriodic acid. It was first necessary to introduce two
corrections into the calorimetric result obtained, one for the
oxidation of the phosphorous into phosphoric acid, knowing their
relative amounts by analysis, and the other for the deoxidation of
the iodic acid. The result then obtained expresses the conversion
of phosphorous into hydrated phosphoric acid. This must be
corrected for the heat of solution of the hydrate in water, and for
the heat of combination of the anhydride with water, before we can
obtain the heat evolved in the reaction of P_{2} with O_{5} in the
proportion for the formation of P_{2}O_{5}. It is natural that with
so complex a method there is a possibility of many small errors,
and the resultant figures will only present a certain degree of
accuracy after repeated corrections by various methods. Of such a
kind are the following figures determined by Thomsen, which we
express in thousands of calories:--P_{2} + O_{5} = 370; P_{2} +
O_{3} + 3H_{2}O = 400; P_{2} + O_{5} + a mass of water = 405. Hence
we see that P_{2}O_{5} + 3H_{2}O = 30; 2PH_{3}O_{4} + an excess of
water = 5. Experiment further showed that crystallised PH_{3}O_{4},
in dissolving in water, evolves 2·7 thousand calories, and that
fused (39°) PH_{3}O_{4} evolves 5·2 thousand calories, whence the
heat of fusion of H_{3}PO_{4} = 2·5 thousand calories. For
phosphorous acid, H_{3}PO_{3}, Thomsen obtained P_{2} + O_{3} +
3H_{2}O = 250, and the solution of crystallised H_{3}PO_{3} in
water = -0·13, and of fused H_{3}PO_{3} = +2·9. For hypophosphorous
acid, H_{3}PO_{2}, the heats of solution are nearly the same (-0·17
and +2·1), and the heat of formation P_{2} + O + 3H_{2}O = 75;
hence its conversion into 2H_{3}PO_{3} evolves 175 thousand
calories, and the conversion of 2H_{3}PO_{3} into 2H_{3}PO_{4} =
150 thousand calories. For the sake of comparison we will take the
combination of chlorine with phosphorus, also according to Thomsen,
per 2 atoms of phosphorus, P_{2} + 3Cl_{2} = 151, P_{2} + 5Cl_{2} =
210 thousand calories. In their reaction on an excess of water
(with the formation of a solution), 2PCl_{3} = 130, 2PCl_{5} = 247,
and 2POCl_{3} = 142 thousand calories.
Besides which we will cite the following data given by various
observers: heat of fusion for P (that is, for 31 parts of
phosphorus by weight) -0·15 thousand calories; the conversion of
yellow into red phosphorus for P, from +19 to +27 thousand
calories; P + H_{3} = 4·3, HI + PH_{3} = 24, PH_{3} + HBr = 22
thousand calories.
At the ordinary temperature (20° C.) phosphorus is not oxidised by
pure oxygen; oxidation only takes place with a slight rise of
temperature, or the dilution of the oxygen with other gases
(especially nitrogen or hydrogen), or a decrease of pressure.
[4 bis] Ordinary phosphorus takes fire at a temperature (60°) at which
no other known substance will burn. Its application to the
manufacture of matches is based on this property. In order to
illustrate the easy inflammability of common (yellow) phosphorus,
its solution in carbon bisulphide may be poured over paper; this
solvent quickly evaporates, and the free phosphorus spread over a
large surface takes fire spontaneously, notwithstanding the cooling
effect produced by the evaporation of the bisulphide. The majority
of _phosphorus matches_ are composed of common phosphorus mixed
with some oxidising substance which easily gives up oxygen, such as
lead dioxide, potassium chlorate, nitre, &c. For this purpose
common phosphorus is carefully triturated under warm water
containing a little gum; lead dioxide and potassium nitrate are
then added to the resultant emulsion, and the match ends,
previously coated with sulphur or paraffin, are dipped into this
preparation. After this the matches are dipped into a solution of
gum and shellac, in order to preserve the phosphorus from the
action of the air. When such a match containing particles of yellow
phosphorus is rubbed over a rough surface, it becomes (especially
at the point of rupture of the brittle gummy coating) slightly
heated, and this is sufficient to cause the phosphorus to take fire
and burn at the expense of the oxygen of the other ingredients.
The heads of the 'safety' matches do not contain any phosphorus, but
only substances capable of burning and of supporting combustion. Red
phosphorus is spread over a surface on the box, and it is the friction
against this phosphorus which ignites the matches. There is no danger of
the matches taking fire accidentally, nor are they poisonous.[5] This red
phosphorus is prepared by heating the ordinary phosphorus at 230° to
270°; it is evident that this must be done in an atmosphere incapable of
supporting combustion--for example, in nitrogen, carbonic anhydride,
steam, &c. On a large scale, ordinary phosphorus is placed in closed iron
vessels,[5 bis] and immersed in a bath of different proportions of tin
and lead, by which means the temperature of 250° necessary for the
conversion is easily attained. It is kept at this temperature for some
time. The temperature is at first cautiously raised, and the air is thus
partially expelled by the heat, and also by the evolution of steam (the
phosphorus is damp when put in), whilst the remaining oxygen is also
partially absorbed by the phosphorus, so that an atmosphere of nitrogen
is produced in the iron vessel. Red phosphorus enters into all the
reactions proper to yellow phosphorus, only with greater difficulty and
more slowly;[6] and, as its vapour tension (volatility) is less than that
of the yellow variety, it may be supposed that a polymerisation takes
place in the passage of the yellow into the red modification, just as in
the passage of cyanogen into paracyanogen, or of cyanic acid into
cyanuric acid (Chapter IX. Notes 39 bis and 48).
[5] In the so-called 'safety' or Swedish matches (which are not
poisonous, and do not take fire from accidental friction) a mixture
of red phosphorus and glass forms the surface on which the matches
are struck, and the matches themselves do not contain any
phosphorus at all, but a mixture of antimonious sulphide,
Sb_{2}S_{3} (or similar combustible substances) and potassium
chlorate (or other oxidising agents). The combustion, when once
started by contact with the red phosphorus, proceeds by itself at
the expense of the inflammatory and combustible elements contained
in the tip of the match. The mixture applied on the match itself
must not be liable to take fire from a blow or friction. The
mixture forming the heads of the 'safety' matches has the following
approximate composition: 55-60 parts of chlorate of potassium, 5-10
parts of peroxide of manganese (or of K_{2}Cr_{2}O_{7}), about 1
part of sulphur or charcoal, about 1 part of pentasulphide of
antimony, Sb_{2}S_{5}, and 30-40 parts of rouge and powdered glass.
This mixture is stirred up in gum or glue, and the matches are
dipped into it. The paper on which the matches are struck is coated
with a mixture of red phosphorus and trisulphide of antimony,
Sb_{2}S_{3}, stirred up in dextrine.
[5 bis] Phosphorus only acts on iron at a red heat. The boiler is
provided with a safety valve and gas-conducting tube, which is
immersed in mercury or other liquid to prevent the admission of air
into the boiler.
[6] The specific heat of the yellow variety is 0·189--that is, greater
than that of the red variety, which is 0·170. The sp. gr. of the
yellow is 1·84, and of the red prepared at 260° 2·15, and of that
prepared at 580° and above (_i.e._ 'metallic' phosphorus, _see_
below) = 2·34. At 230° the pressure of the vapour of ordinary
phosphorus = 514 millimetres of mercury, and of the red = 0--that
is to say, the red phosphorus does not form any vapour at this
temperature; at 447° the vapour tension of ordinary phosphorus is
at first = 5500 mm., but it gradually diminishes, whilst that of
red phosphorus is equal to 1636 mm.
Hittorf, by heating the lower portion of a closed tube containing
red phosphorus to 530° and the upper portion to 447°, obtained
crystals of the so-called 'metallic' phosphorus at the upper
extremity. As the vapour tensions (according to Hittorf, at 530°
the vapour tension of yellow phosphorus = 8040 mm., of red = 6139
mm., and of metallic = 4130 mm.) and reactions are different,
_metallic phosphorus_ may be regarded as a distinct variety. It is
still less energetic in its chemical reaction than red phosphorus,
and it is denser than the two preceding varieties: sp. gr. = 2·34.
It does not oxidise in the air; is crystalline, and has a metallic
lustre. It is obtained when ordinary phosphorus is heated with lead
for several hours at 400° in a closed vessel, from which the air
has been exhausted. The resultant mass is then treated with dilute
nitric acid, which first dissolves the lead (phosphorus is
electro-negative to lead, and does not, therefore, act on the
nitric acid at first) and leaves brilliant rhombohedral crystals of
phosphorus of a dark violet colour with a slight metallic lustre,
which conduct an electric current incomparably better than the
yellow variety; this also is characteristic of the metallic state
of phosphorus.
The researches of Lemoine partially explain the passage of yellow
(ordinary) phosphorus into its other varieties. He heated a closed
glass globe containing either ordinary or red phosphorus, in the
vapour of sulphur (440°), and then determined the amount of the red
and yellow varieties after various periods of time, by treating the
mixture with carbon bisulphide. It appeared that after the lapse of
a certain time a mixture of definite and equal composition is
obtained from both--that is, between the red and yellow varieties a
state of equilibrium sets in like that of dissociation, or that
observed in double decompositions. But at the same time, the
progress of the transformation appeared to be dependent on the
relative quantity of phosphorus taken per volume of the globe
(_i.e._ upon the pressure). Neglecting the latter, we will cite as
an example the amounts of the red phosphorus transformed into the
ordinary, and of the ordinary not converted into red, per 30 grams
of red or yellow taken per litre capacity of the globe, heated to
440°. When red phosphorus was taken, 4·75 grams of yellow
phosphorus were formed after two hours, four grams after eight
hours, three grams after twenty-four hours, and the last limit
remained constant on further heating. When thirty grams of yellow
phosphorus were taken, five grams remained unaltered after two
hours, four grams after eight hours, and after twenty-four hours
and more three grams as before. Troost and Hautefeuille showed that
liquid phosphorus in general changes more easily into the red than
does phosphorus vapour, which, however, is able, although slowly,
to deposit red phosphorus.
The question presents itself as to whether phosphorus in a state of
vapour is the ordinary or some other variety? Hittorf (1865)
collected many data for the solution of this problem, which leave
no doubt that (as experimental figures show) the density of the
vapour of phosphorus is always the same, although the vapour
tension of the different varieties and their mixtures is very
variable. This shows that the different varieties of phosphorus
only occur in a liquid and solid state, as indeed is implied in the
idea of polymerisation. Strictly speaking, the vapour of phosphorus
is a particular state of this substance, and the molecular formula
P_{4} refers only to it, and not to any other definite state of
phosphorus. But Raoult's solution method showed that in a benzene
solution the fall of the freezing point indicates for ordinary
phosphorus a molecule P_{4}, judging by the determinations of
Paterno and Nasini (1888), Hirtz (1890), and Beckmann (1891), who
obtained for sulphur by the same method a molecular weight = S_{6},
in conformity with the vapour density. Further research in this
direction will perhaps show the possibility of finding the
molecular weight of red phosphorus, if a means be discovered for
dissolving it without converting it into the yellow variety.
I think it will not be out of place here to draw the reader's
attention to the fact that red phosphorus, which we must recognise
as polymeric with the yellow, stands nearer to nitrogen, whose
molecule is N_{2}, in its small inclination towards chemical
reactions, although judging by its small vapour tension it must be
more complex than ordinary (yellow and white) phosphorus.
The vapour of phosphorus is colourless; its density remains constant
between 300° and 1000° (Dumas, 1833; Mitscherlich, Deville, and Troost,
1859, and others). The density with respect to air has been determined as
from 4·3 to 4·5. Hence, referred to hydrogen, it is 4·4 × 14·4 = 63,
corresponding with a molecular weight 124, _i.e._ the molecule of
phosphorus in a state of vapour contains P_{4}. The reader will remember
that the molecule of nitrogen contains N_{2}, of sulphur S_{6} or S_{2},
and of oxygen O_{2} or O_{3}.
The chemical energy of phosphorus in a free state more nearly approaches
that of sulphur than nitrogen. Phosphorus is combustible and inflames at
60°; but having in the act of combination parted with a portion of its
energy in the form of heat it becomes analogous to nitrogen, so long as
there is no question of its reduction back again into phosphorus. Nitric
acid is easily reduced to nitrogen, whilst phosphoric acid is reduced
with very much greater difficulty. All the compounds of phosphorus are
less volatile than those of nitrogen. Nitric acid, HNO_{3}, is easily
distilled; metaphosphoric acid, HPO_{3}, is generally said to be
non-volatile; triethylamine, N(C_{2}H_{5})_{3}, boils at 90°, and
triethylphosphine, P(C_{2}H_{5})_{3}, at 127°.
Phosphorus not only combines easily and directly with oxygen, but also
with chlorine, bromine, iodine, sulphur, and with certain metals, and red
phosphorus when heated combines with hydrogen also.[6 bis] So, for
instance, when fused with sodium under naphtha, phosphorus gives the
compound Na_{3}P_{2}. Zinc, absorbing the vapour of phosphorus, gives the
phosphide Zn_{3}P_{2} (sp. gr. 4·76); tin, SnP; copper, Cu_{2}P; even
platinum combines with phosphorus (PtP_{2}, sp. gr. 8·77).[6 tri] Iron,
when combined even with a small quantity of phosphorus, becomes
brittle.[7] Some of these compounds of phosphorus are obtained by the
action of phosphorus on the solutions of metallic salts, and by the
ignition of metallic oxides in the vapour of phosphorus, or by heating
mixtures of phosphates with charcoal and metals. Phosphides do not
exhibit the external properties of salts, which are so clearly seen in
the chlorides and still distinctly observable in the sulphides. _The
phosphides of the metals_ of the alkalis and of the alkaline earths are
even immediately and very easily decomposed by water, whereas this is
found to be the case with only a very few sulphides, and still more
rarely and indistinctly with the chlorides. We may take calcium phosphide
as an example.[7 bis] Phosphorus is laid in a deep crucible, and covered
with a clay plug, over which lime is strewn. At a red heat the vapours of
phosphorus combine with the oxygen of the lime and form phosphoric
anhydride, which forms a salt with another portion of the lime, whilst
the liberated calcium combines with the phosphorus and forms calcium
phosphide. Its composition is not quite certain; it may be CaP
(corresponding with liquid phosphuretted hydrogen). This substance is
remarkable for the following reaction: if we take water--or, better
still, a dilute solution of hydrochloric acid--and throw calcium
phosphide into it, bubbles of gas are evolved, which take fire
spontaneously in the air and form white rings. This is owing to the fact
that the liquid hydrogen phosphide, PH_{2}, is first formed, thus, CaP +
2HCl = CaCl_{2} + PH_{2}, which, owing to its instability, very easily
splits up into the solid phosphide, P_{2}H, and gaseous phosphide,
PH_{3}; 5PH_{2} = P_{2}H + 3PH_{3}; the latter corresponds with ammonia.
The mixture of the gaseous and liquid phosphides takes fire spontaneously
in the air, forming phosphoric acid. The same hydrogen phosphides are
formed when water acts on sodium phosphide (P_{2}Na_{3}). A similar
mixture of gaseous liquid and solid phosphuretted hydrogen (Retgers 1894)
is formed by heating (in a glass tube) red phosphorus in a stream of dry
hydrogen. Hence we see that there are _three compounds of phosphorus with
hydrogen_. (1) The first or solid yellow phosphide, P_{2}H (more probably
P_{4}H_{2}), is obtained by the action of strong hydrochloric acid on
sodium phosphide; it takes fire when struck or at 175°. (2) The liquid,
PH_{2}, or more correctly expressed as the molecule, P_{2}H_{4}, is a
colourless liquid which takes fire spontaneously in the air, boils at
30°, is very unstable, and is easily decomposed (by light or hydrochloric
acid) into the two other phosphides of hydrogen. It is prepared by
passing the gases evolved by the action of water on calcium phosphide
through a freezing mixture.[8] And, lastly, (3), gaseous hydrogen
phosphide, _phosphine_, PH_{3}, which is distinguished as being the most
stable. It is a colourless gas, which does not take fire in the air. It
has an odour of garlic, and is very poisonous. It resembles ammonia in
many of its properties.[8 bis] It is easily decomposed by heat, like
ammonia, forming phosphorus and hydrogen; but it is very slightly soluble
in water, and does not saturate acids, although it forms compounds with
some of them which resemble ammonium salts in their form and properties.
Among them the _compound with hydriodic acid_, PH_{4}I, analogous to
ammonium iodide, is remarkable. This compound crystallises on sublimation
in well-formed cubes, like sal-ammoniac, which it resembles in many
respects. However, this compound does not enter into those reactions of
double decomposition which are proper to sal-ammoniac, because its saline
properties are very feebly developed. Phosphuretted hydrogen also
combines, like ammonia, with certain chloranhydrides; but they are
decomposed by water, with the evolution of phosphine. Ogier (1880) showed
that hydrochloric acid also combines with phosphine under a pressure of
20 atmospheres at +18°, and under the ordinary pressure at -35°, forming
the crystalline phosphonium chloride PH_{4}Cl, corresponding to
sal-ammoniac. Hydrobromic acid does the same with greater ease, and
hydriodic acid with still greater facility, forming phosphonium iodide,
PH_{4}I.[9]
[6 bis] Retgers (see further on) showed this in 1894, and observed that
As when heated also combines with hydrogen.
[6 tri] The capacity of mercury (Chapter XVI., Note 25 bis) to give
unstable compounds with nitrogen gives rise to the supposition that
similar compounds exist with phosphorus also. Such a compound was
obtained by Granger (1892) by heating mercury with iodide of
phosphorus in a closed tube at 275°-300°. After removing the iodide
of mercury formed, there remain fine rhombic crystals having a
metallic lustre, and composition Hg_{3}P_{2}. This compound is
stable, does not alter at the ordinary temperature and only
decomposes at a red heat; when heated in air it burns with a flame.
Nitric and hydrochloric acids do not act upon it, but it is easily
decomposed by aqua regia. A phosphide of copper, Cu_{2}P_{2}, was
obtained by Granger (1893) by heating a mixture of water, finely
divided copper and red phosphorus in a sealed tube to 130°. The
excess of copper was afterwards washed away by a solution of NH_{3}
in the presence of air.
[7] The metallic compounds of phosphorus possess a great chemical
interest, because they show a transition from metallic alloys (for
instance, of Sb, As) to the sulphides, halogen salts, and oxides,
and on the other hand to the nitrides. Although there are already
many fragmentary data on the subject, the interesting province of
the metallic phosphides cannot yet be regarded as in any way
generalised. The varied applications (phosphor-iron,
phosphor-bronze, &c.), which the phosphides have recently acquired
should give a strong incentive to the complete and detailed study
of this subject, which would, in my opinion, help to the
explanation of chemical relations beginning with alloys (solutions)
and ending with salts and the compounds of hydrogen (hydrides),
because the phosphor-metals, as is proved by direct experiment,
stand in the same relation to phosphuretted hydrogen as the
sulphides do towards sulphuretted hydrogen, or as the metallic
chlorides to hydrochloric acid.
[7 bis] Many other compounds of phosphorus are also capable of forming
phosphuretted hydrogen. Thus BP also gives PH_{3} (_see_ Chapter
XVII., Note 12). According to Lüpke (1890) phosphuretted hydrogen
is formed by phosphide of tin. The latter is prepared by treating
molten tin covered with a layer of carbonate of ammonium, with red
phosphorus; 200-300 c.c. of water are then poured into a flask, 3-5
grams of this phosphide of tin dropped in, and after driving out
the air by a stream of carbonic acid, hydrochloric acid (sp. gr.
1·104) is poured in. The disengagement of phosphuretted hydrogen
takes place on heating the flask in a water bath. The following is
another easy method for preparing PH_{3}. A mixture of 1 part of
zinc dust (fume) and 2 parts of red phosphorus are heated in an
atmosphere of hydrogen (the mixture burns in air). Combination
takes place accompanied by a flash, and a grey mass of Zn_{3}P_{2}
is formed which gives PH_{3} when treated with dilute H_{2}SO_{4}.
[8] The spontaneous inflammability of the hydride PH_{2} in air is very
remarkable, and it is particularly interesting that its analogues
in composition, P(C_{2}H_{5})_{2} (the formula must be doubled) and
Zn(C_{2}H_{5})_{2}, also take fire spontaneously in air.
[8 bis] The analogy between PH_{3} and NH_{3} is particularly clear in
the hydrocarbon derivatives. Just as NH_{2}R, NHR_{2}, and NR_{3},
where R is CH_{3}, and other hydrocarbon radicles, correspond to
NH_{3}, so there are actually similar compounds corresponding to
PH_{3}. These compounds form a branch of organic chemistry.
[9] The periodic law and direct experiment (the molecular weight) show
that PH_{3} is the normal compound of P and H and that it is more
simple than PH_{2} or P_{2}H_{4}, just as methane, CH_{4}, is more
simple than ethane, C_{2}H_{6}, whose empirical composition is
CH_{3}. The formation of liquid phosphuretted hydrogen may be
understood from the law of substitution. The univalent radicle of
PH_{3} is PH_{2}, and if it is combined with H in PH_{3} it
replaces H in liquid phosphuretted hydrogen, which thus gives
P_{2}H_{4}. This substance corresponds with free amidogen
(hydrazine), N_{2}H_{4} (Chapter VI.) Probably P_{2}H_{4} is able
to combine with HI, and perhaps also with 2HI, or other
molecules--that is, to give a substance corresponding to
phosphonium iodide.
_Phosphonium iodide_, PH_{4}I, may be prepared, according to
Baeyer, in large quantities in the following manner:--100 parts of
phosphorus are dissolved in dry carbon bisulphide in a tubulated
retort: when the mixture has cooled, 175 parts of iodide are added
little by little, and the carbon bisulphide is then distilled off,
this being done towards the end of the operation in a current of
dry carbonic anhydride at a moderate temperature. The neck of the
retort is then connected with a wide glass tube, and the tubulure
with a funnel furnished with a stopcock, and containing 50 parts of
water. This water is added drop by drop to the phosphorous iodide,
and a violent reaction takes place, with the evolution of hydriodic
acid and phosphonium iodide. The latter collects as crystals in the
glass tube and the retort itself. It is purified by further
distillations; more than 100 parts may be obtained. Baeyer
expresses the reaction by the equation P_{2}I + 2H_{2}O = PH_{4}I +
PO_{2}; and the compound PO_{2} may be represented as phosphorous
phosphoric anhydride: P_{2}O_{5} + P_{2}O_{3} = 4PO_{2}. As a
better proportion we may take 400 grams of phosphorus, 680 grams of
iodine, and 240 grams of water, and express the formation thus: 13P
+ 9I + 21H_{2}O = 3H_{4}P_{2}O_{7} + 7PH_{4}I + 2HI (Chapter XI.,
Note 77).
Phosphonium iodide and even phosphine act as reducing agents in
solutions of many metallic salts. Cavazzi showed that with a
solution of sulphurous anhydride phosphine gives sulphur and
phosphoric acid.
_Phosphuretted hydrogen, or phosphine_, PH_{3}, is generally prepared by
the action of caustic potash on phosphorus.[10] Small pieces of
phosphorus are dropped into a flask containing a strong solution of
caustic potash and heated. Potassium hypophosphite, H_{2}KPO_{2}, is then
obtained in solution; gaseous phosphuretted hydrogen is evolved:
P_{4} + 3KHO + 3H_{2}O = 3(KH_{2}PO_{2}) + PH_{3}.
Liquid phosphuretted hydrogen (and free hydrogen) is also formed,
together with the phosphine, so that the gaseous product, on escaping
from the water into the air, takes fire spontaneously, forming beautiful
white rings of phosphoric acid. In this experiment, as in that with
calcium phosphide, it is the liquid, P_{2}H_{4}, that takes fire; but the
phosphine set light to by it also burns, PH_{3} + O_{4} = PH_{3}O_{4}.
The same phosphuretted hydrogen, PH_{3}, may be obtained pure, and not
spontaneously combustible, by igniting the hydrates of phosphorous acid
(4PH_{3}O_{3} = PH_{3} + 3PH_{3}O_{4}) and hypophosphorous acid
(2PH_{3}O_{2} = PH_{3} + PH_{3}O_{4}); or, more simply, by the
decomposition of calcium phosphide by hydrochloric acid, because then all
the liquid phosphide, P_{2}H_{4}, is decomposed into non-volatile P_{2}H
and gaseous PH_{3}. Pure phosphine liquefies when cooled to -90°, boils
at -85°, and solidifies at -135° (Olszewski). When phosphorus burns in an
excess[10 bis] of _dry_ oxygen, then only _phosphoric anhydride_,
P_{2}O_{5} is formed. It is prepared by dropping pieces of phosphorus
through a wide tube, fixed into the upper neck of a large glass globe, on
to a cup suspended in the centre of the globe. These lumps are set alight
by touching them with a hot wire, and the phosphorus burns into
P_{2}O_{5}. The dry air necessary for its combustion is forced into the
globe through a lateral neck, and the white flakes of phosphoric
anhydride formed are carried by the current of air through a second
lateral neck into a series of Woulfe's bottles, where they settle as
friable white flakes. Phosphoric anhydride may also be formed by passing
dry air through a solution of phosphorus in carbon bisulphide. All the
materials for the preparation of this substance must be carefully dried,
because it _combines_ with great eagerness _with water_, at the same time
developing a large amount of heat and forming metaphosphoric acid,
HPO_{3}, from which the water cannot be separated by heat. Phosphoric
anhydride is a colourless snow-like substance, which attracts moisture
from the air with the utmost avidity. It fuses at a red heat, and then
_volatilises_. Its affinity for water is so great that it takes it up
from many substances. Thus it converts sulphuric acid into sulphuric
anhydride, and carbohydrates (wood, paper) are carbonised, and give up
the elements of water when brought into contact with it.
[10] The air must first be expelled from the flask by hydrogen, or some
other gas which will not support combustion, as otherwise an
explosion might take place owing to the spontaneous inflammability
of the phosphuretted hydrogen.
The combustion of phosphuretted hydrogen in oxygen also takes
place under water when the bubbles of both gases meet, and it is
very brilliant. The phosphuretted hydrogen obtained by the action
of phosphorus on caustic potash always contains free hydrogen, and
often even the greater part of the gas evolved consists of
hydrogen.
_Pure phosphuretted hydrogen_ (not containing hydrogen or liquid
or solid phosphides) is obtained by the action of a solution of
potash on phosphonium iodide: PH_{4}I + KHO = PH_{3} + KI + H_{2}O
(in just the same way as ammonia is liberated from ammonium
chloride). The reaction proceeds easily, and the purity of the gas
is seen from the fact that it is entirely absorbed by bleaching
powder and is not spontaneously inflammable. Its mixture with
oxygen explodes when the pressure is diminished (Chapter XVIII.,
Note 8). The vapours of bromine, nitric acid, &c., cause it to
again acquire the property of inflaming in the air; that is, they
partially decompose it, forming the liquid hydride, P_{2}H_{4}.
Oppenheim showed that when red phosphorus is heated at 200° with
hydrochloric acid in a closed tube it forms the compound
PCl_{3}(H_{3}PO_{3}), together with phosphine.
[10 bis] If there be a deficiency of oxygen, _phosphorous anhydride_
P_{2}O_{3} is formed. It was obtained by Thorpe and Tutton (1890)
and is easily volatilised, melts at 22°·5, boils without change
(in an atmosphere of N_{2} or CO_{2}) at 173°, and is therefore
easily separated from P_{2}O_{3}, which volatilises with
difficulty. The vapour density shows that the molecular weight is
double, _i.e._ P_{4}O_{6} (like As_{2}O_{3}). Although colourless,
phosphorous anhydride (its density in a state of fusion at 24° =
1·936) turns yellow and reddens in sun-light (possibly red
phosphorus separates out ?), and decomposes at 400° forming
hypophosphorous anhydride P_{2}O_{4} (Note 11) and phosphorus. It
passes into P_{2}O_{5} in air and oxygen, and when slightly heated
in oxygen becomes luminous, and ultimately takes fire. Cold water
slowly transforms P_{2}O_{3} into phosphoric acid, but hot water
gives an explosion and leads to the formation of PH_{3},
(P_{4}O_{6} + 6H_{2}O = PH_{3} + 3PH_{3}O_{4}). Alkalis act in the
same manner. It takes fire in chlorine and forms POCl_{3} and
PO_{2}Cl, and combines with sulphur at 160°, forming
P_{2}S_{2}O_{3} (the molecular formula is double this) a substance
which volatilises in vacuo and is decomposed by water into H_{2}S
and phosphoric acid, _i.e._ it may be regarded as P_{2}O_{5}, in
which O_{2} has been replaced by two atoms of sulphur. Judging
from the above, the mixture of P_{2}O_{3} and P_{2}O_{5} formed in
the combustion of phosphorus in air is transformed into P_{2}O_{5}
in an excess of oxygen.
When moist phosphorus slowly oxidises in the air, it not only forms
phosphorous and phosphoric acids, but also _hypophosphoric acid_,
H_{4}P_{2}O_{6}, which when in a dry state easily splits up at 60° into
phosphorous and metaphosphoric acids (H_{4}P_{2}O_{6} = H_{3}PO_{3} +
HPO_{3}), but differs from a mixture of these acids in that it forms
well-characterised salts, of which the sodium salt,
H_{2}Na_{2}P_{2}O_{6}, is but slightly soluble in water (the sodium salts
of phosphoric and phosphorous acids are easily soluble), and that it does
not act as a reducing agent, like mixtures containing phosphorous
acid.[11]
[11] Salzer proved the existence of hypophosphoric acid (it is also
called subphosphoric acid), in which many chemists did not
believe. Drawe (1888) and Rammelsberg (1892) investigated its
salts. It may be obtained in a free state by the following method.
The solution of acid produced by the slow oxidation of moist
phosphorus is mixed with a solution (25 p.c.) of sodium acetate. A
salt, Na_{2}H_{2}P_{2}O_{6},6H_{2}O, crystallises out on cooling;
it is soluble in 45 parts of water, and gives a precipitate of
Pb_{2}P_{2}O_{6} with lead salts (Ag_{4}P_{2}O_{6} with salts of
silver). The lead salt is decomposed by a current of hydrogen
sulphide, when lead sulphide is precipitated, while the solution,
evaporated under the receiver of an air-pump, gives crystals of
H_{4}P_{2}O_{6},2H_{2}O, which easily lose water and give
H_{4}P_{2}O_{6}. The salts in which the H_{4} is replaced by
Ni_{2}, or NiNa_{2}, or CdNa_{2}, &c., are insoluble in water.
In order to see the relation between phosphoric acid and
hypophosphoric acid which does not contain the elements of
phosphorous acid (because it does not reduce either gold or
mercury from their solutions), but which nevertheless is capable
of being oxidised (for example, by potassium permanganate) into
phosphoric acid, it is simplest to apply the law of substitution.
This clearly indicates the relation between oxalic acid,
(COOH)_{2}, and carbonic acid, OH(COOH). The relation between the
above acids is exactly the same if we express phosphoric acid as
OH(POO_{2}H_{2}), because in this case P_{2}H_{4}O_{6}, or
(POO_{2}H_{2})_{3}, will correspond with it just as oxalic does
with carbonic acid. A similar relationship exists between
hyposulphuric or dithionic acid, (SO_{2}OH)_{2}, and sulphuric
acid, OH(SO_{2}OH), as we shall find in the following chapter.
Dithionic acid corresponds with the anhydride S_{2}O_{5},
intermediate between SO_{2} and SO_{3}; oxalic acid with
C_{2}O_{3}, intermediate between CO and CO_{2}; hypophosphoric
acid corresponds with the anhydride P_{2}O_{4}, intermediate
between P_{2}O_{3} and P_{2}O_{5}, and the analogue of N_{2}O_{4}.
Judging by the general law of the formation of acids (Chapter XV.), the
series of phosphorus compounds should include the following _ortho-acids_
and their corresponding anhydrides, answering to phosphuretted hydrogen,
H_{3}P:--
H_{3}PO_{4}, phosphoric acid, and P_{2}O_{5}, anhydride,
H_{3}PO_{3}, phosphorous acid, and P_{2}O_{3}, anhydride,
H_{3}PO_{2}, hypophosphorous acid, and P_{2}O, anhydride.[12]
The last of these (the analogue of N_{2}O) is almost unknown. Phosphoric
anhydride (P_{2}O_{5}) with a small quantity of water does not at first
give orthophosphoric acid, PH_{3}O_{4}, but a compound P_{2}O_{5},H_{2}O,
or PHO_{3}, whose composition corresponds with that of nitric acid; this
is _metaphosphoric acid_. Even with an excess of water, combining with
phosphoric anhydride, this metaphosphoric acid, and not the ortho-,
passes at first into solution. Metaphosphoric acid in solution only
passes into orthophosphoric acid when the solution is heated or after a
lapse of time.
[12] Besides the hydrates enumerated, a compound, PH_{3}O, should
correspond with PH_{3}. This hydrate, which is analogous to
hydroxylamine, is not known in a free state, but it is known as
triethylphosphine oxide, P(C_{2}H_{5})_{3}O, which is obtained by
the oxidation of triethylphosphine, P(C_{2}H_{5})_{3}. It must be
observed that there may also be lower oxides of phosphorus
corresponding with PH_{3}, like N_{2}O and NO, and there are even
indications of the formation of such compounds, but the data
concerning them cannot be considered as firmly established.
_Orthophosphoric acid_[13] is obtained by oxidising phosphorus with
nitric acid until the phosphorus entirely passes into solution and the
lower oxides of nitrogen cease to be evolved. The reaction takes place
best with dilute nitric acid, and when aided by heat. The resultant
solution is evaporated to a syrup. If a weighed quantity of phosphorus
(dried in a current of dry carbonic anhydride) be taken, a crystalline
mass of the acid can be obtained by evaporating the solution until it
consists only of the quantity[14] of phosphoric acid corresponding with
the amount of phosphorus taken (from 31 parts of P, 98 parts of
solution). The acid fuses at +39°; specific gravity of the liquid 1·88.
Phosphorus pentachloride, PCl_{5}, and oxychloride, POCl_{3} (see further
on), give orthophosphoric acid and hydrochloric acid with water. The two
other varieties of phosphoric acid, with which we shall presently become
acquainted, give the same ortho-acid when under the influence of acids,
with particular ease when boiled and more slowly in the cold. By itself
orthophosphoric acid (either in solution or when dry) does not pass into
the other varieties; it does not oxidise, and therefore presents the
limiting and stable form. When heated to 300°, it loses water and passes
into pyrophosphoric acid, 2H_{3}PO_{4} = H_{2}O + H_{4}P_{2}O_{7}, whilst
at a red heat it loses twice as much water and is converted into
metaphosphoric acid, H_{3}PO_{4} = H_{2}O + HPO_{3}. In aqueous solution
orthophosphoric acid differs clearly from pyro- or metaphosphoric acids,
because the solutions of these latter acids give different reactions:
thus orthophosphoric acid does not precipitate albumin, does not give a
precipitate with barium chloride, and forms a yellow precipitate of
silver orthophosphate, Ag_{3}PO_{4}, with silver nitrate (in the presence
of alkalis, but not otherwise); whilst a solution of pyrophosphoric acid,
H_{4}P_{2}O_{7}, although it does not precipitate albumin or barium
chloride, gives a white precipitate of silver pyrophosphate,
Ag_{4}P_{2}O_{7}, with silver nitrate; and a solution of metaphosphoric
acid, HPO_{3}, precipitates both albumin and barium chloride, and gives a
white precipitate of silver metaphosphate, AgPO_{3}, with silver nitrate.
These points of distinction were studied by Graham, and are exceedingly
instructive. They show that the solution of a substance does not
determine the maximum of chemical combination with water, that solutions
may contain various degrees of combination with water, and that there is
a clear difference between the water serving for solution and that
entering into chemical combination. Graham's experiments also showed that
the water whose removal or combination determines the conversion of
ortho- into meta- and pyrophosphoric acids differs distinctly from water
of crystallisation, for he obtained the salts of ortho-, meta-, and
pyrophosphoric acids with water of crystallisation, and they differed in
their reactions, like the acids themselves. This water of crystallisation
was expelled with greater ease than the water of constitution of the
hydrates in question.[14 bis]
[13] Phosphoric acid, being a soluble and almost non-volatile
substance, cannot be prepared like hydrochloric and nitric acids
by the action of sulphuric acid on the alkali phosphates, although
it is partially liberated in the process. For this purpose the
salts of barium or lead may be taken, because they give insoluble
salts, thus Ba_{3}(PO_{4})_{2} + 3H_{2}SO_{4} = 3BaSO_{4} +
2H_{3}PO_{4}. Bone ash contains, besides calcium phosphate, sodium
and magnesium phosphates, and fluorides and other salts, so that
it cannot give directly a pure phosphoric acid.
[14] If this is not done the orthophosphoric acid, PH_{3}O_{4}, loses a
portion of its water, and then, as with an excess of water, it
does not crystallise.
[14 bis] The difference between the reactions of ortho-, meta- and
pyrophosphoric acids, established by Graham (_see_ p. 163), is of
such importance for the theory of hydrates and for explaining the
nature of solutions, that in my opinion its influence upon
chemical thought has been far from exhausted. At the present time
many such instances are known both in organic (for instance, the
difference between the reactions of the solutions of certain
anhydrides and hydrates of acids), and inorganic chemistry (for
example, the difference between the rose and purple cobalt
compounds, Chapter XXII. &c.) They essentially recall the long
known and generalised difference between C_{2}H_{4} (ethylene),
C_{2}H_{6}O (ethyl alcohol = ethylene + water), and C_{4}H_{10}O
(ethyl ether = 2 ethylene + water = 2 alcohol - water); but to the
present day the numerous analogous phenomena existing among
inorganic substances are only considered as a simple difference in
degrees of affinity, distinguishing the water of constitution
(hydration), crystallisation, and solution without penetrating
into the difference of the structure or distribution of the
elements, which exists here and gives rise to a distinct isomerism
of solutions. In my opinion the progress of chemistry, especially
with regard to solutions, should make rapid strides when the cause
of the isomerism of solutions, for instance, of ortho- and
pyrophosphoric acids, has become as clear to us as the cause of
many well-studied instances of the isomerism, polymerism, and
metamerism of organic compounds. Here it forms one of those many
important problems which remain for the chemistry of the future in
a state of only indistinct presentiments and in the form of facts
empirically known but insufficiently comprehended.
Orthophosphoric acid has a pleasant acid taste and a distinctly acid
reaction; it is used as a medicine, and is not poisonous (phosphorous
acid is poisonous). Alkalis, like sodium, potassium, and ammonium
hydroxides, saturate the acid properties of phosphoric acid when taken in
the ratio 2NaHO : H_{3}PO_{4}--that is, when salts of the composition
HNa_{2}PO_{4} are formed. When taken in the ratio NaHO : H_{3}PO_{4}, a
solution having an acid reaction is obtained, and when 3NaHO :
H_{3}PO_{4}--that is, when the salt Na_{3}PO_{4} is formed--an alkaline
reaction is obtained. Hence many chemists (Berzelius) even regarded the
salts of composition R_{2}HPO_{4} as normal, and considered phosphoric
acid to be bibasic. But the salt Na_{2}HPO_{4} also shows a feeble
alkaline reaction, so that it is impossible to judge the characteristic
peculiarities of acids by the reactions on litmus paper, as we already
know from many examples. Orthophosphoric acid is tribasic, because it
contains three equivalents of hydrogen replaceable by metals, forming
salts, such as NaH_{2}PO_{4}, Na_{2}HPO_{4}, and Na_{3}PO_{4}. It is also
tribasic, because with silver nitrate its soluble salts always give
Ag_{3}PO_{4},[15] a salt with three equivalents of silver, and because by
double decomposition with barium chloride it forms a salt of the
composition Ba_{3}(PO_{4})_{2}, and silver and barium hardly ever give
basic salts. With the metals of the alkalis, phosphoric acid forms
soluble salts, but the normal salts of the metals of the alkaline earths,
R_{3}(PO_{4})_{2} and even R_{2}H_{2}(PO_{4}), are insoluble in water,
but dissolve in feeble acids, such as phosphoric and acetic, because they
then form soluble acid salts, especially RH_{4}(PO_{4})_{2}.[16]
[15] Silver orthophosphate, Ag_{3}PO_{4}, is yellow, sp. gr. 7·32, and
insoluble in water. When heated it fuses like silver chloride, and
if kept fused for some length of time it gives a white
pyrophosphate (the decomposition which causes this is not known).
It is soluble in aqueous solutions of phosphoric, nitric, and even
acetic acids, of ammonia, and many of its salts. If silver nitrate
acts on a dimetallic orthophosphate--for instance,
Na_{2}HPO_{4}--it still gives Ag_{3}PO_{4}, nitric acid being
disengaged: Na_{2}HPO_{4} + 3AgNO_{3} = Ag_{3}PO_{4} + 2NaNO_{3} +
HNO_{3}. When alcohol is added to silver orthophosphate,
Ag_{3}PO_{4}, dissolved in syrupy phosphoric acid, it precipitates
a white salt (the alcohol takes up the free phosphoric acid)
having the composition Ag_{2}HPO_{4}, which is immediately
decomposed by water into the normal salt and phosphoric acid.
[16] The researches of Thomsen showed that in very dilute aqueous
solutions the majority of monobasic acids--nitric, acetic,
hydrochloric, &c. (but hydrofluoric acid more and hydrocyanic
less)--HX evolve the following amounts of heat (in thousands of
calories) with caustic soda: NaHO + 2HX = 14; NaHO + HX = 14;
2NaHO + HX = 14; that is, if _n_ be a whole number _n_NaHO + HX =
14 and NaHO + _n_HX = 14. Hence reaction here only takes place
between one molecule of NaHO and one molecule of acid, and the
remaining quantity of acid or alkali does not enter into the
reaction. In the case of bibasic acids, H_{2}R´´ (sulphuric,
dithionic, oxalic, sulphuretted hydrogen, &c.), NaHO + 2H_{2}R´´ =
14; NaHO + H_{2}R´´ = 14; 2NaHO + H_{2}R´´ = 28; _n_NaHO +
H_{2}R´´ = 28; that is, with an excess of acid (NaHO + 2H´_{2}R´´)
14 thousand units of heat are developed, and with an excess of
alkali 28. When phosphoric acid is taken (but not all tribasic
acids--for instance, not citric) the general character of the
phenomenon is similar to the preceding, namely, NaHO +
2H_{3}PO_{4} = 14·7; NaHO + H_{3}PO_{4} = 14·8; 2NaHO +
H_{3}PO_{4} = 27·1; 3NaHO + H_{3}PO_{4} = 34·0; 6NaHO +
H_{3}PO_{4} = 35·3; or, in general terms, NaHO + _n_H_{3}PO_{4} =
14 (approximately) and _n_NaHO + H_{3}PO_{4} = 35 and not 42,
which shows a peculiarity of phosphoric acid. In the case of
energetic acids, when one equivalent (23 grams) of sodium (in the
form of hydroxide) replaces one equivalent (1 gram) of hydrogen
(with the formation of water and in dilute solutions), 14,000 heat
units are evolved; and this is true for phosphoric acid when in
H_{3}PO_{4}, Na or Na_{2} replaces H or H_{2}, but when Na_{3}
replaces H_{3} less heat is developed. This will be seen from the
following scheme based on the preceding figures: H_{3}PO_{4} +
NaHO = 14·8; NaH_{2}PO_{4} + NaHO = 12·3; Na_{2}HPO_{4} + NaHO =
5·9; with Na_{3}PO_{4} + NaHO, a very small amount of heat is
evolved, as may be judged from the fact that Na_{3}PO_{4} + 3NaHO
= 1·3, but still heat is evolved. It must be supposed that in
acting on phosphoric acid in the presence of a large quantity of
water, a certain portion of the sodium hydroxide remains as alkali
uncombined with the acid. Thus, on increasing the mass of the
alkali, heat is still evolved, and a fresh interchange between Na
and H takes place. Hence water shows a decomposing action on the
alkali phosphates. The same decomposing action of water is seen,
but to a less extent, with Na_{2}HPO_{4}, as may be judged both
from the reactions of this salt and from the amount of heat
developed by NaH_{2}PO_{4} with NaHO. Such an explanation is in
accordance with many facts concerning the decomposition of salts
by water already known to us. Recent researches made by Berthelot
and Louguinine have confirmed the above deductions made by me in
the first edition (1871) of this work. At the present time views
of this nature are somewhat generally accepted, although they are
not sufficiently strictly applied in other cases. As regards
PH_{3}O_{4} it may be said that: on the substitution of the first
hydrogen this acid acts as a powerful acid (like HCl, HNO_{3},
H_{2}SO_{4}); on the substitution of the second hydrogen as a
weaker acid (like an organic acid); and on the substitution of the
third, as an alcohol, for instance phenol, having the properties
of a feeble acid.
Phosphoric anhydride, or any of its hydrates, when ignited with an excess
of sodium hydroxide, carbonate, &c., forms normal or _trisodium
orthophosphate_, Na_{3}PO_{4}, but when a solution of sodium carbonate is
decomposed by orthophosphoric acid, only the salt Na_{2}HPO_{4} is
formed; and when an excess of sodium chloride is ignited with
orthophosphoric acid, hydrochloric acid is evolved, and the acid salt
H_{2}NaPO_{4} alone is formed. These facts clearly indicate the small
energy of phosphoric acid with respect to the formation of the
tri-metallic salt, which is seen further from the fact that the salt
Na_{3}PO_{4} has an alkaline reaction, decomposes in the presence of
water and carbonic acid, forming Na_{2}HPO_{4}, corrodes glass vessels in
which it is boiled or evaporated, just like solutions of the alkalis,
disengages, like them, ammonia from ammonium chloride, and crystallises
from solutions, as Na_{3}PO_{4},12H_{2}O, only in the presence of an
excess of alkali. At 15° the crystals of this salt require five parts of
water for solution; they fuse at 77°.
_Disodium orthophosphate_, or common sodium phosphate, Na_{2}HPO_{4}, is
more stable both in solution and in the solid state. As it is used in
medicine and in dyeing, it is prepared in considerable quantities, most
frequently from the impure phosphoric acid obtained by the action of
sulphuric acid on bone ash. The solution thus formed--which contains,
besides phosphoric and sulphuric acids, salts of sodium, calcium, and
magnesium--is heated, and sodium carbonate added so long as carbonic
anhydride is disengaged. A precipitate is formed containing the insoluble
salts of magnesium and calcium, whilst the solution contains sodium
phosphate, Na_{2}HPO_{4}, with a small quantity of other salts, from
which it may be easily purified by crystallisation. At the ordinary
temperature its solutions, especially in the presence of a small amount
of sodium carbonate, give finely-formed inclined prismatic crystals,
Na_{2}HPO_{4},12H_{2}O; when the crystallisation takes place above 30°
they only contain 7H_{2}O. The former crystals even lose a portion of
their water of crystallisation at the ordinary temperature (the salt
effloresces), and form the second salt with 7H_{2}O; whilst under the
receiver of an air-pump and over sulphuric acid they also part with this
water.[17] When ignited they lose the last molecule of water of
constitution, and give sodium pyrophosphate, Na_{4}P_{2}O_{7}.
[17] Na_{2}HPO_{4},12H_{2}O has a sp. gr. 1·53. Poggiale determined the
solubility in 100 parts of water (1) of the anhydrous ortho-salt
Na_{2}HPO_{4}, and (2) of the corresponding pyro-salt
Na_{4}P_{2}O_{7}:--
0° 20° 40° 80° 100°
I. 1·5 11·1 30·9 81 108
II. 3·2 6·2 13·5 30 40
At temperatures of 20° to 100° the ortho-salt is so very much less
soluble that this difference alone already indicates the
deeply-seated alteration in constitution which takes place in the
passage from the ortho- to the pyro-salts.
_Monosodium orthophosphate_, NaH_{2}PO_{4}, crystallises with one
equivalent of water; its solution has an acid reaction. At 100° the salt
only loses this water of crystallisation, and at about 200° it parts with
all its water, forming the metaphosphate NaPO_{3}. It is prepared from
ordinary sodium phosphate by adding phosphoric acid until the solution
does not give a precipitate with barium chloride, and then evaporating
and crystallising the solution. The solution of this salt does not absorb
carbonic anhydride, and does not give a precipitate with salts of
calcium, barium, &c.[18]
[18] The _ammonium orthophosphates_ resemble the sodium salts in many
respects, but the instability of the di- and tri-metallic salts is
seen in them still more clearly than in the sodium salts; thus
(NH_{4})_{3}PO_{4}, and even (NH_{4})_{2}HPO_{4}, lose ammonia in
the air (especially when heated, even in solutions);
NH_{4}H_{2}PO_{4} alone does not disengage ammonia and has an acid
reaction. The crystals of the first salt contain 3H_{2}O, and are
only formed in the presence of an excess of ammonia; both the
others are anhydrous, and may be obtained like the sodium salts.
When ignited these salts leave metaphosphoric acid behind; for
example, (NH_{4})_{2}HPO_{4} = 2NH_{3 + H_{2}O + HPO_{3}. Ammonia
also enters into the composition of many double phosphates.
Ammonium sodium orthophosphate, or simply phosphate,
NH_{4}NaHPO_{4},4H_{2}O, crystallises in large transparent
crystals from a mixture of the solutions of disodium phosphate and
ammonium chloride (in which case sodium chloride is obtained in
the mother liquid), or, better still, from a solution of
monosodium phosphate saturated with ammonia. It is also formed
from the phosphates in urine when it ferments. This salt is
frequently used in testing metallic compounds by the blow-pipe,
because when ignited it leaves a vitreous metaphosphate, NaPO_{3},
which, like borax, dissolves metallic oxides, forming
characteristic tinted glasses.
When a solution of trisodium phosphate is added to a solution of a
magnesium salt it gives a white precipitate of the normal
orthophosphate Mg_{2}(PO_{4})_{2},7H_{2}O. If the trisodium salt
be replaced by the ordinary salt, Na_{2}HPO_{4}, a precipitate is
also formed, and MgHPO_{4},7H_{2}O is obtained. It might be
thought that the normal salt Mg_{3}(PO_{4})_{2} would be
precipitated if disodium phosphate was added to ammonia and a salt
of magnesium, but in reality _ammonium magnesium orthophosphate_,
MgNH_{4}PO_{4},6H_{2}O, is precipitated as a crystalline powder,
which loses ammonia and water when ignited, and gives a
pyrophosphate, Mg_{2}P_{2}O_{7}. This salt occurs in nature as the
mineral struvite, and in various products of the changes of animal
matter. If we consider that the above salt parts with ammonia with
difficulty, and that the corresponding salt of sodium is not
formed under the same conditions (MgNaPO_{4},9H_{2}O is obtained
by the action of magnesia on disodium phosphate), if we turn our
attention to the fact that the salts of calcium and barium do not
form double salts as easily as magnesium, and remember that the
salts of magnesium in general easily form double ammonium salts,
we are led to think that this salt is not really a normal, but an
acid salt, corresponding with Na_{2}HPO_{4}, in which Na_{2} is
replaced by the equivalent group NH_{3}Mg.
The common normal _calcium phosphate_, Ca_{3}(PO_{4})_{2}, occurs
in minerals, in animals, especially in bones, and also probably in
plants, although the ash of many portions of plants, as a rule,
contains less lime than the formation of the normal salt requires.
Thus 100 parts of the ash (from 5,000 parts of grain) of rye grain
contain 47·5 of phosphoric anhydride and only 2·7 of lime, and
even the ash of the whole of the rye (including the straw)
contains twice as much phosphoric anhydride as lime, and the
normal salt contains almost equal weights of these substances.
Only the ash of grasses, and especially of clover, and of trees,
contains in the majority of cases more lime than is required for
the formation of Ca_{3}P_{2}O_{8}. This salt, which is insoluble
in water, dissolves even in such feeble acids as acetic and
sulphurous, and even in water containing carbonic acid. The latter
fact is of immense importance in nature, since by reason of it
rain water is able to transfer the calcium phosphates in the soil
into solutions which are absorbed by plants. The solubility of the
normal salt in acids takes place by virtue of the formation of an
acid salt, which is evident from the quantity of acid required for
its solution, and more especially from the fact that the acid
solutions when evaporated give crystalline scales of the acid
calcium phosphate, CaH_{4}(PO_{4})_{2}, soluble in water. This
solubility of the acid salt forms the basis of the treatment by
acids of bones, phosphorites, guano, and other natural products
containing the normal salt and employed for fertilising the soil.
The perfect decomposition requires at least 2H_{2}SO_{4} to
Ca_{3}(PO_{4})_{2}, but in reality less is taken, so that only a
portion of the normal salt is converted into the acid salt.
Hydrochloric acid is sometimes used. (In practice such mixtures
are known as _superphosphates_). Certain experiments, however,
show that a thorough grinding, the presence of organic, and
especially of nitrogenous, substances, and the porous structure of
some calcium phosphates (for example, in burnt bones), render the
treatment of phosphoric manures by acids superfluous--that is, the
crop is not improved by it.
As a hydrate, orthophosphoric acid should be expressed, after the fashion
of other hydrates, as containing three water residues (hydroxyl groups),
_i.e._ as PO(OH)_{3}. This method of expression indicates that the type
PX_{5}, seen in PH_{4}I, is here preserved, with the substitution of
X_{2} by oxygen and X_{3} by three hydroxyl groups. The same type appears
in POCl_{3}, PCl_{5}, PF_{5}, &c. And if we recognise phosphoric acid as
PO(OH)_{3}, we should expect to find three anhydrides corresponding with
it: (1) [PO(OH)_{2}]_{2}O, in which two of the three hydroxyls are
preserved; this is pyrophosphoric acid, H_{4}P_{2}O_{7}. (2) PO(OH)O,
where only one hydroxyl is preserved. This is metaphosphoric acid. (3)
(PO)_{2}O_{3} or P_{2}O_{5}, that is, perfect phosphoric anhydride.
Therefore, _pyro- and metaphosphoric acids are imperfect anhydrides_ (or
anhydro-acids) _of orthophosphoric acid_.[19]
[19] In this sense the ortho-acid itself might be regarded as an
anhydro-acid, counting P(HO)_{5} as the perfect hydrate, if PH_{5}
existed; but as in general the normal hydrates correspond with the
existing hydrogen compounds with the addition of up to 4 atoms of
oxygen, therefore PH_{3}O_{4} is the normal acid, just as
SH_{2}O_{4} and ClHO_{4}; while NHO_{3}, CH_{2}O_{3} are
meta-acids, or higher normal acids (NH_{3}O_{4} and CH_{4}O_{4})
with the loss of a molecule of water.
In order to see the relation between the ortho-, pyro-, and
metaphosphoric acids, the first thing to remark in them is that
the anhydride P_{2}O_{5} is combined with 3, 2, and 1 molecules of
water. In the absence of data for the molecular weight of ortho-
and pyrophosphoric acids it is necessary to mention that all
existing data for metaphosphoric acid indicate (Note 21) that its
molecule is much more complex and contains at least
H_{3}P_{3}O_{9} or H_{6}P_{6}O_{18}. The explanation of the
problems which here present themselves can, it seems to me, be
only looked for after a detailed study of the phenomena of the
polymerisations of mineral substances, and of those complex acids,
such as phosphomolybdic, which we shall hereafter describe
(Chapter XXI.) A similar instance is exhibited in the solubility
of hydrate of silica (produced by the action of silicon fluoride
on water) in fused metaphosphoric acid, with the formation, on
cooling, of an octahedral compound (sp. gr., 3·1) containing
SiO_{2},P_{2}O_{5}. A certain indication (but no proof) that
ordinary orthophosphoric acid is polymerised is given by
Staudenmaier (1893), who obtained a salt, K_{5}H_{4}P_{3}O_{12},
by the action of a solution of KH_{2}PO_{4} upon K_{2}CO_{3}; and
a compound, KH_{3}P_{2}O_{8}, corresponding to the doubled
molecule of H_{3}PO_{4}, by the action of KH_{2}PO_{4} upon
H_{3}PO_{4} itself.
_Pyrophosphoric acid_, H_{4}P_{2}O_{7}, is formed by heating
orthophosphoric acid to 250° when it loses water.[19 bis] Its normal
salts are formed by igniting the dimetallic salts of orthophosphoric acid
of the types HM_{2}PO_{4}. Thus from the disodium salt we obtain sodium
pyrophosphate, Na_{4}P_{2}O_{7} (it crystallises from water with
10H_{2}O, is very stable, fuses when heated, has an alkaline reaction,
and does not form ortho-salts when its solution is boiled): and from the
monosodium salt NaH_{2}PO_{4} the acid salt Na_{2}H_{2}P_{2}O_{7} (easily
soluble in water) is formed; this has an acid reaction, and when ignited
further gives the meta-salt.[20]
[19 bis] According to Watson (1893) the ortho-acid is partially
transformed into the pyro-acid at 230°, whilst at 260° the latter
begins to volatilise. At 300° the meta-acid only is formed.
[20] The method of preparation of the acid itself consists in
converting the sodium salt, Na_{4}P_{2}O_{7}, by double
decomposition with water and a salt of lead, into insoluble lead
pyrophosphate, Pb_{2}P_{2}O_{7}, which is then suspended in water
and decomposed by sulphuretted hydrogen; lead sulphide is thus
precipitated, and pyrophosphoric acid remains in solution. This
solution cannot be heated, or the pyro-acid will pass into the
ortho-, but must be evaporated under the receiver of an air-pump.
It concentrates to a syrup and crystallises, and when ignited in
this form loses water, and forms metaphosphoric acid. It resembles
orthophosphoric acid in many respects; its salts with the alkalis
are also soluble, and the others insoluble in water but soluble in
acids. When heated in solution with acid it gives orthophosphoric
acid, as well as when fused with an excess of alkali.
Witt heated ammonium chloride with phosphoric acid (hydrochloric
acid was evolved), ignited the residue to drive off ammonia, and
obtained pyrophosphoric acid in the residue.
_Metaphosphoric acid_, HPO_{3} (the analogue of nitric acid), is formed
by the ignition of the pyro- and ortho-acids (or, better, of their
ammonium salts), as a vitreous, hygroscopic, fused mass (glacial
phosphoric acid, _acidum phosphoricum glaciale_), soluble in water and
volatilising without decomposition. It is also formed in the first slow
action of cold water on the anhydride, but metaphosphoric acid gradually
changes into the ortho-acid when its solution is boiled, or when it is
kept for any length of time, especially in the presence of acids.[21]
[21] As when using phenolphthalein as an indicator in neutralising by
an alkali metaphosphoric acid is monobasic, and orthophosphoric
acid is bibasic, it is possible by means of this difference to
follow the transition of meta- into orthophosphoric acid. Sabatier
(1888) carried on an investigation of this nature, and found that
the rate of transformation is dependent on the temperature, and is
subject to the general laws of the rate of chemical
transformations which belongs to physical chemistry.
Metaphosphoric acid has a particular interest in respect to the
variations to which its salts are subject. The metaphosphates are
formed by the ignition of the acid orthophosphates, MH_{2}PO_{4},
or MNH_{4}HPO_{4}, or of the acid pyrophosphates,
M_{2}H_{2}P_{2}O_{7}, or M_{2}(NH_{4})_{2}P_{2}O_{7}, water and
ammonia being given off in the process. The properties of the
metaphosphates, which have a similar composition to nitrates--for
instance, NaPO_{3}, or Ba(PO_{3})_{2}--vary according to the
duration of the ignition to which the ortho-, or pyrophosphates
from which they are prepared have been subjected. When the salts
NaH_{2}PO_{4} or NH_{4}NaHPO_{4} are strongly ignited, a salt
NaPO_{3} is formed, which deliquesces in the air, and gives a
gelatinous precipitate with salts of the alkaline earths. But, as
Graham (in 1830-40), and many others, especially Fleitmann and
Henneberg (in 1840-50), and Tamman (in the nineties), observed,
under other conditions the salts of the same composition acquire
other properties. The above chemists recognise five polymeric
forms of metaphosphates, (HPO_{3})_{_n_}. We will follow the
nomenclature and researches of Fleitmann.
_Monometaphosphoric acid._ The salts are distinguished for their
insolubility in water; even the salts NaPO_{3}, KPO_{3}, are
insoluble. They are obtained by igniting the monometallic
orthophosphates--for example, RH_{2}PO_{4}--up to the temperature
at which all water is evolved (316°), but not to fusion. No double
salts are known.
_Dimetaphosphoric acid_, on the contrary, easily forms double
salts--for example, KNaP_{2}O_{6}, and also the copper potassium
salt, &c. The copper salt is obtained by evaporating a solution of
copper oxide in orthophosphoric acid. A blue ortho-salt,
CuRHO_{4}, first separates from the solution, then a light-blue
pyro-salt, Cu_{2}P_{2}O_{7}; and above 350°, when metaphosphoric
acid itself begins to volatilise, the dimetaphosphate,
CuP_{2}O_{6}, is formed. The residue is washed with water, and
decomposed with a hot solution of sodium sulphide, when the sodium
salt, Na_{2}P_{2}O_{6}, is obtained in solution. This salt, when
evaporated with alcohol, gives crystals containing 2 mol. H_{2}O,
which, however, retain their solubility (in 7 parts of water)
after the water is driven off at 100°. When fused, these crystals
give a deliquescent salt (hexa-metaphosphate). The solution of the
salt has a neutral reaction, which only after prolonged boiling
becomes acid, owing to the formation of orthophosphate,
NaH_{2}PO_{4}. The soluble salts of dimetaphosphoric acid give the
insoluble silver salt, Ag_{2}P_{2}O_{6}, with silver nitrate, and
a precipitate of BaP_{2}O_{6}2H_{2}O with barium chloride.
_Trimetaphosphoric acid_ is obtained as the sodium salt
Na_{3}P_{3}O_{9} when any other metaphosphate of sodium is fused
and _slowly_ cooled, then dissolved in a slight excess of warm
water, and the resultant solution evaporated. The crystals contain
6 mol. H_{2}O, and dissolve in four parts of water. An acid
reaction is only obtained, as with the preceding salt, after
prolonged boiling with water. The acid is a true analogue of
nitric acid, because _all its metallic salts are soluble_.
_Hexametaphosphoric acid._ Fleitmann so named the ordinary
metaphosphoric acid (glacial) which attracts moisture. The
deliquescent sodium salt is obtained, like the trimetaphosphate,
only by _rapid_ cooling. It is also formed by fusing silver oxide
with an excess of phosphoric acid. The sodium salt is soluble in
water, and gives viscous, elastic precipitates with salts of Ba,
Ca, and Mg. Lubert (1893) obtained salts of Ag, Pb, &c.
Jawein and Thillot (1889), who investigated the sodium salts of
metaphosphoric acid by Raoult's method, came to the conclusion
that the salts of di- and tri-metaphosphoric acid behave in such a
manner that their molecule must be represented as non-polymerised
NaPO_{3}, whilst those of hexametaphosphoric acid behave as
(NaPO_{3})_{4}. At all events, the series of salts which Fleitmann
and Henneberg regard as monometaphosphates--_i.e._ as
non-polymerised--are most probably the most polymerised, because
they are insoluble.
According to Tamman's researches, vitreous metaphosphoric acid
contains a mixture consisting chiefly of two varieties, differing
in the solubility and degree of stability of their salts. The
least stable corresponds to Fleitmann's hexa-acid, and gives three
isomeric salts. Tamman came to the conclusion that there exist
polymers also in the form of penta-, ortho-, and
deca-metaphosphoric acids. Without going into details upon this
subject, I do not think it superfluous to point out that the
undoubted capability of metaphosphoric acid to polymerise should
be connected with its faculty of combining with water, whilst the
degree of polymerisation and the number of polymeric forms cannot
yet be considered as sufficiently explained.
In order to see the relation between phosphoric acid and the lower acids
of phosphorus, it is simplest to imagine the substitution of hydroxyl in
H_{3}PO_{4} or PO(OH)_{3} by hydrogen. Then from orthophosphoric acid,
PO(OH)_{3}, we shall obtain phosphorous acid, POH(OH)_{2}, and
hypophosphorous acid, POH(OH); and, furthermore, phosphorous acid should
be bibasic if orthophosphoric acid was tribasic, and hypophosphorous acid
should be monobasic. This conclusion[21 bis] is, in fact, true, and hence
all the acids of phosphorus may be referred to one common type, PX_{5},
whose representatives are PH_{4}I and PCl_{5}, POCl_{3}, PCl_{2}F_{3},
&c.
[21 bis] The bibasity of H_{3}PO_{3}, established by Würtz, has been
proved by many direct experiments (see, for instance, Note 22),
among which we may mention that Amat (1892) took a mixture of the
aqueous solutions of Na_{2}HPO_{3} and NaHO and added absolute
alcohol to it. Two layers were formed; the upper, alcoholic,
contained all the excess of NaHO, whilst the lower only contained
the salt Na_{2}HPO_{3}, which was therefore unable to react with
the excess of NaHO. Amat also obtained NaH_{2}PO_{3} by saturating
H_{3}PO_{3} with soda until he obtained a neutral reaction with
methyl-orange. The replacement of one atom of H by sodium here, as
in phosphoric acid (Note 16), gives more heat than the replacement
of the second atom. For the third atom there is no formation of a
salt, and therefore no evolution of heat. The monometallic
salts--for example, NaH_{2}PO_{3}--or the ammonia salts, when
heated to 160°, give, as Amat had previously shown, a salt of
bibasic pyrophosphorous acid, Na_{2}H_{2}P_{2}O_{5}.
_Phosphorous acid_, PH_{3}O_{3}, is generally obtained from phosphorus
trichloride, PCl_{3}, by the action of water: PCl_{3} + 3H_{2}O = 3HCl +
PH_{3}O_{3}. Both acids formed are soluble in water, but are easily
separated, because hydrochloric acid is volatile whilst phosphorous acid
volatilises with difficulty, and if a small amount of water be originally
taken the hydrochloric acid nearly all passes off directly. Concentrated
solutions of phosphorous acid give crystals of H_{3}PO_{3}, which fuse at
70°, attract moisture from the air, and deliquesce when ignited, giving
phosphine and phosphoric acid,[22] and are oxidised into orthophosphoric
acid by many oxidising agents. In its salts only two hydrogen atoms are
replaced by metals (Würtz); the salts of the alkaline metals are soluble,
and give precipitates with salts of the majority of other metals.
[22] Phosphorous acid, when subjected to the action of nascent hydrogen
(zinc and sulphuric acid), evolves phosphine, and when boiled with
an excess of alkali it evolves hydrogen (PH_{3}O_{3} + 3KHO =
PK_{3}O_{4} + 2H_{2}O + H_{2}); owing to its liability to
oxidation, it is a reducing agent--for instance, it reduces cupric
chloride to cuprous chloride, and precipitates silver from the
nitrate and mercury from its salts.
These reactions are perhaps connected with the fact that in this
acid one atom of hydrogen should be considered as in the same
condition as in phosphuretted hydrogen, which is expressed by the
formula PHO(OH)_{2}, if we represent it as PH_{4}X, with the
substitution of two of the hydrogen atoms by oxygen and of HX by
two of hydroxyl. The direct passage of phosphorous chloride into
phosphorous acid would, however, indicate that all the three atoms
of hydrogen in it occur in the form of hydroxyl, because no
difference is known between the three atoms of chlorine in
PCl_{3}--they all react alike, as a rule. However, Menschutkin, by
acting on alcohol, C_{2}H_{5}OH, with phosphorous chloride,
obtained hydrochloric acid and a substance P(C_{2}H_{5}O)Cl_{2},
and from it by the action of bromine he obtained ethyl bromide,
C_{2}H_{5}Br, and a compound PBrOCl_{2}, which proves, to a
certain extent, the existence of a difference between the three
atoms of chlorine in phosphorous chloride. If we turn our
attention to the formation of phosphine by the ignition of
phosphorous acid, we see that 4PH_{3}O_{3} only evolve 3H in the
form of PH_{3}, and therefore the residue--that is,
3PH_{3}O_{4}--will still contain one hydrogen of the same nature
as in phosphine, because in 4PH_{3}O_{3} we should recognise four
such hydrogens as in phosphine. We arrive at the same conclusion
by examining the decomposition of hypophosphorous acid,
2PH_{3}O_{2} = PH_{3} + PH_{3}O_{4}. In the two molecules of the
monobasic hypophosphorous acid taken, there are only two atoms of
hydrogen replaceable by metals, whilst in the molecule of the
resultant phosphoric acid there are three. Perhaps relations of
this nature determine the relative stability of the dimetallic
salts of orthophosphoric acid.
The monobasic _hypophosphorous acid_, PH_{3}O_{2}, gives salts
PH_{2}O_{2}Na, (PH_{2}O_{2})_{2}Ba, &c.; the two remaining atoms of
hydrogen (which exist in the same form as in phosphine, PH_{3}) are not
replaceable by metals, and this determines the property of these salts of
evolving phosphuretted hydrogen when heated (especially with alkalis). In
acting on substances liable to reduction it is this hydrogen which acts,
and, for example, _reduces_ gold and mercury from the solutions of their
salts, or converts cupric into cuprous salts. In all these instances the
hypophosphorous acid is converted into phosphoric acid. Under the action
of zinc and sulphuric acid it gives phosphine, PH_{3}. Nevertheless,
neither hypophosphorous acid nor its dry salts absorb oxygen from the
air. The salts of hypophosphorous acid are more soluble than those of the
preceding acids of phosphorus. Thus the sodium salt PNaH_{2}O_{2} does
not give a precipitate with barium chloride, and the salts of calcium,
barium, and many other metals are soluble.[23] The hypophosphites are
prepared by boiling an alkali with phosphorus so long as phosphuretted
hydrogen is evolved. The acid itself is obtained from barium
hypophosphite (prepared in the same manner by boiling phosphorus in
baryta water), by decomposing its solution with sulphuric acid. By
concentration of the solution of hypophosphorous acid (it must not be
heated above 130°, at which temperature it decomposes) a syrup is formed
which is able to crystallise. In the solid state hypophosphorous acid
fuses at +17°, and has the properties of a clearly defined acid.
[23] Calcium hypophosphite is used in medicine. According to Cavazzi, a
mixture of sodium hypophosphite, NaH_{2}PO_{2}, and sodium nitrate
explodes violently.
The types PX_{3} and PX_{5}, which are evident for the hydrogen and
oxygen compounds of phosphorus, are most clearly seen in its halogen
compounds,[24] to the consideration of which we will proceed, fixing our
attention more especially on the chlorine compounds, as being the most
important from the historical, theoretical, and practical point of view.
[24] Fluorine and bromine give PX_{3} and PX_{5}, like chlorine. With
respect to iodine PI_{5} is, in a chemical sense, a very unstable
substance, and generally _phosphorus tri-iodide_ only is formed
(from yellow or red phosphorus and iodine in the requisite
proportions). It is a red crystalline substance, fuses at 55°, is
easily decomposed by water, forming phosphorous and hydriodic
acids, and when heated it evolves iodine vapours and forms
phosphorus di-iodide, PI_{2}. This substance may be obtained in
the same manner as the preceding by taking a smaller proportion of
iodine (8 parts of iodine to 1 part of phosphorus, whilst the
tri-iodide requires 12·3); it also forms red crystals, which melt
at 110°. When decomposed by water it not only gives phosphorous
and hydriodic acids, but also phosphine and a yellow substance (a
lower oxide of phosphorus). In its composition di-iodide of
phosphorus corresponds with liquid phosphuretted hydrogen, PH_{2},
and probably its molecular weight is much higher: P_{2}I_{4} or
P_{3}I_{6}, &c. As the iodine compounds of phosphorus give
hydriodic and phosphorous acids with water, and as both these
substances are reducing agents in the presence of water (and
hydrates), iodide of phosphorus also acts as a reducing agent.
Phosphorus burns in chlorine, forming phosphorous chloride, PCl_{3}, and
with an excess of chlorine, phosphoric chloride, PCl_{5}. The
oxychloride, POCl_{3}, as the simplest chloranhydride according to the
type PX_{5}, and also phosphoric chloride, correspond with
orthophosphoric acid, PO(OH)_{3}, while phosphorous chloride, PCl_{3},
corresponds with phosphorous acid and the type PX_{3}. Phosphoric
oxychloride, POCl_{3}, is a colourless liquid, boiling at 110°.
Phosphorus trichloride is also a colourless liquid, boiling at 76°,[25]
whilst phosphoric chloride is a solid yellowish substance, which
volatilises without melting at about 168°. They are all heavier than
water, and form types of the _chloranhydrides_ or chlorine compounds of
the non-metallic elements whose hydrates are acids, just as NaCl or
BaCl_{2} are types of halogen metallic salts.
[25] In a liquid state the density of phosphorous chloride at 10° =
1·597, and therefore its molecular volume = 137·5/1·597 = 86·0,
and that of phosphorus oxychloride is equal to 153·5/1·693 = 90·7;
hence the addition of oxygen has produced considerable increase in
volume, just as in the conversion of sulphur dichloride, SCl_{2},
into sulphuryl chloride, SOCl_{2}, the volume changes from 64 to
71. It is the same with the boiling-points; phosphorus trichloride
boils at 70°, the oxychloride at 100°, sulphur dichloride at 64°,
and sulphuryl chloride at 78°--that is, the addition of oxygen
raises the boiling points.
_The vapour density_ of phosphorus trichloride and oxychloride
corresponds with their formulæ (Cahours, Würtz)--namely, is equal
to half the molecular weight referred to hydrogen. But it is not
so with phosphorus pentachloride. Cahours showed that the vapour
density of phosphorus pentachloride referred to air = 3·65, to
hydrogen = 52·6, whilst according to the formula PCl_{5} it should
be = 104·2. Hence this formula corresponds with four, and not with
two, molecules. This shows that the vapour of phosphoric chloride
contains two and not one molecule, that in a state of vapour it
splits up, like sal-ammoniac, sulphuric acid, &c. The products of
disruption must here be phosphorous chloride, PCl_{3}, and
chlorine, Cl_{2}, bodies which easily re-form phosphoric chloride,
PCl_{5}, at a lower temperature. This decomposition of phosphoric
chloride in its conversion into vapour is confirmed by the fact
that the vapour of this almost colourless substance shows the
greenish-yellow colour proper to chlorine. This dissociation of
phosphoric chloride has been considered by some chemists as a sign
that phosphorus, like nitrogen, does not give volatile compounds
of the type PX_{5}, and that such substances are only obtained as
unstable molecular compounds which break up when distilled; for
example, PH_{3},HI, PCl_{3},Cl_{2}, NH_{3},HCl, &c. To prove that
the molecule PCl_{5} actually exists, Würtz in 1870 observed that
when mixed with the vapour of phosphorous chloride the vapour of
phosphoric chloride distils over (from 160° to 190°) perfectly
colourless, and has a density which is really near to the
formula--namely, to 104--and the same density was determined for
the pentachloride in an atmosphere of chlorine. Hence at low
temperatures and in admixture with one of the products of
dissociation, there is no longer that decomposition which occurs
at higher temperatures--that is, we have here a case of
dissociation proceeding at moderate temperatures.
An important proof in favour of the type PX_{5} is exhibited by
phosphorus pentafluoride PF_{5}, obtained by Thorpe as a
colourless gas which only corrodes glass after the lapse of time;
it may be kept over mercury, and has a normal density. It is
formed when liquid arsenic trifluoride, AsF_{3}, is added to
phosphoric chloride surrounded by a freezing mixture: 3PCl_{5} +
5AsF_{3} = 3PF_{5} + 5AsCl_{3}.
In general, fluorine and phosphorus give stable compounds: PF_{3},
POF_{3}, and PF_{5}, as would be expected from the fact that in
passing from Cl to I (_i.e._ as the atomic weight of the halogen
increases) the stability of the compounds with P and the tendency
to give PX_{5} (Note 24) decreases. _Phosphorus trifluoride_ is
obtained by heating a mixture of ZnF_{2} and PBr_{3}, by the
action of AsF_{3} upon PCl_{3}, by heating phosphide of copper
with PbF_{2}, &c. It is a strong-smelling gas, which liquefies at
-10° under a pressure of 40 atmospheres, giving a colourless
liquid. It dissolves easily in (is absorbed by, reacts with)
water, and acts upon glass; when mixed with Cl_{2} it combines
with it (Poulenc, 1891), forming PCl_{2}F_{3}, a colourless gas of
normal density, which is transformed into a liquid at 8°,
decomposes into PF_{3} + Cl_{2} at 250°, and, with a small amount
of water, gives _oxy-fluoride_ of phosphorus, POF_{3} (with a
large amount of water it gives PH_{3}O_{4}), which Moissan (1891)
obtained by the action of dry HF upon P_{2}O_{5}, and Thorpe and
Tutton (1890) by heating a mixture of cryolite and P_{2}O_{5}. It
is a gas of normal density, like PF_{3}, and was obtained by
Moissan by the action of fluorine upon PF_{3} (PSF_{3}, _see_
Chapter XX., Note 20). Thus the forms PX_{3} and PX_{5} not only
exist in many solid and non-volatile substances, but also as
vapours.
If a piece of phosphorus be dropped into a flask containing chlorine, it
burns when touched with a red-hot wire, and combines with the chlorine.
If the phosphorus be in excess, liquid _phosphorus trichloride_, PCl_{3},
is always formed, but if the chlorine be in excess the solid
pentachloride is obtained. The trichloride is generally prepared in the
following manner. Dry chlorine (passed through a series of Woulfe's
bottles containing sulphuric acid) is led into a retort containing sand
and phosphorus. The retort is heated, the phosphorus melts, spreads
through the sand, and gradually forms the trichloride, which distils over
into a receiver, where it condenses. _Phosphoric chloride_ or _phosphorus
pentachloride_, PCl_{5}, is prepared by passing dry chlorine into a
vessel containing phosphorus trichloride (purified by distillation).
Phosphorous chloride combines directly with oxygen, but more rapidly with
ozone or with the oxygen of potassium chlorate (3PCl_{3} + KClO_{3} =
3POCl_{3} + KCl), forming _phosphorus oxychloride_, POCl_{3} (Brodie).
This compound is also formed by the first action of water on phosphoric
chloride; for example, if two vessels, one containing phosphoric chloride
and the other water, are placed under a bell jar, after a certain time
the crystals of the chloride disappear and hydrochloric acid passes into
the water. The aqueous vapour acts on the pentachloride, and the
following reaction occurs: PCl_{5} + H_{2}O = POCl_{3} + 2HCl, the result
being that liquid phosphorus oxychloride is found in one vessel, and a
solution of hydrochloric acid in the other. However, an excess of water
directly transforms phosphoric chloride into orthophosphoric acid,
PCl_{5} + 4H_{2}O = PH_{3}O_{4} + 5HCl,[26] since POCl_{3} reacts with
water (3H_{2}O), forming 3HCl and phosphoric acid PO(OH)_{3}.
[26] Phosphorus oxychloride is obtained by the action of phosphoric
chloride on hydrates of acids (because alkalis decompose
phosphorus oxychloride), according to the equation PCl_{5} + RHO =
POCl_{3} + RCl + HCl, where RHO is an acid. The reaction only
proceeds according to this equation with monobasic acids, but then
RCl is volatile, and therefore a mixture is obtained of two
volatile substances, the acid chloride and phosphorus oxychloride,
which are sometimes difficult to separate; whilst if the hydrate
be polybasic the reaction frequently proceeds so that an anhydride
is formed: RH_{2}O_{2} + PCl_{5} = RO + POCl_{3} + 2HCl. If the
anhydride be non-volatile (like boric), or easily decomposed (like
oxalic), it is easy to obtain pure oxychloride. Thus phosphorus
oxychloride is often prepared by acting on boric or oxalic acid
with phosphoric chloride. It is also formed when the vapour of
phosphoric chloride is passed over phosphoric anhydride,
P_{2}O_{5} + 3PCl_{5} = 5POCl_{3}. This forms an excellent example
in proof of the fact that the formation of one substance from two
does not necessarily show that the resultant compound contains the
molecules of these substances in its molecule. But other
oxychlorides of phosphorus are also formed by the interaction of
phosphoric anhydride and chloride; thus at 200° the
chloranhydride, PO_{2}Cl, or chloranhydride of metaphosphoric
acid, is formed (Gustavson). The chloranhydride of pyrophosphoric
acid, P_{2}O_{3}Cl_{4}, was obtained (Hayter and Michaelis),
together with NOCl, &c., by the action of NO upon cold PCl_{3}, as
a fuming liquid boiling at 210°.
The above chlorine compounds serve not only as a type of the
chloranhydrides, but also as a means for the preparation of other _acid
chloranhydrides_. Thus the conversion of acids XHO into chloranhydrides,
XCl, is generally accomplished by means of _phosphorus pentachloride_.
This fact was discovered by Chancel, and adopted by Gerhardt as an
important method for studying organic acids. By this means organic acids,
containing, as we know, RCOOH (where R is a hydrocarbon group, and where
carboxyl may repeat itself several times by replacing the hydrogen of
hydrocarbon compounds), are converted into their chloranhydrides, RCOCl.
With water they again form the acid, and resemble the chloranhydrides of
mineral acids in their general properties.
Since carbonic acid, CO(OH)_{2}, contains two hydroxyl groups, its
perfect chloranhydride, COCl_{2}, _carbonic oxychloride_, _carbonyl
chloride_ or _phosgene gas_, contains two atoms of chlorine, and differs
from the chloranhydrides of organic acids in that in them one atom of
chlorine is replaced by the hydrocarbon radicle RCOCl, if R be a
monatomic radicle giving a hydrocarbon RH. It is evident, on the one
hand, that in RCOCl the hydrogen is replaced by the radicle COCl, which
is also able to replace several atoms of hydrogen (for example,
C_{2}H_{4}(COCl)_{2} corresponds with the bibasic succinic acid); and, on
the other hand, that the reactions of the chloranhydrides of organic
acids will answer to the reactions of carbonyl chloride, as the reactions
of the acids themselves answer to those of carbonic acid. Carbonyl
chloride is obtained directly from dry carbon monoxide and chlorine[27]
exposed to the action of light, and forms a colourless gas, which easily
condenses into a liquid, boiling at +8°, specific gravity 1·43, and
having the suffocating odour belonging to all chloranhydrides. Like all
chloranhydrides, it is immediately decomposed by water, forming carbonic
anhydride, according to the equation COCl_{2} + H_{2}O = CO_{2} + 2HCl,
and thus expresses the type proper to all chloranhydrides of both mineral
and organic acids.[28]
[27] The direct action of the sun's rays, or of magnesium light, is
necessary to start the reaction between carbonic oxide and
chlorine, but when once started it will proceed rapidly in
diffused light. An excess of chlorine (which gives its coloration
to the colourless phosgene) aids the completion of the reaction,
and may afterwards be removed by metallic antimony. Porous
substances, like charcoal, aid the reaction. Phosgene may be
prepared by passing a mixture of carbonic anhydride and chlorine
over incandescent charcoal. Lead or silver chloride, when heated
in a current of carbonic oxide, also partially form phosgene gas.
Carbon tetrachloride, CCl_{4}, also forms it when heated with
carbonic anhydride (at 400°), with phosphoric anhydride (200°),
and most easily of all with sulphuric anhydride (2SO_{3} + CCl_{4}
= COCl_{2} + S_{2}O_{5}Cl_{2}, this is pyrosulphuryl chloride).
Chloroform, CHCl_{3}, is converted into carbonyl chloride when
heated with SO_{2}(OH)Cl (the first chloranhydride of sulphuric
acid); CHCl_{3} + SO_{3}HCl = COCl_{2} + SO_{2} + 2HCl (Dewar),
and when oxidised by chromic acid.
Among the reactions of phosgene we may mention the formation of
urea with ammonia, and of carbonic oxide when heated with metals.
[28] We are already acquainted with some of the chloranhydrides of the
inorganic acids--for instance, BCl_{3}, and SiCl_{4}--and here we
shall describe those which correspond with sulphuric acid in the
following chapter. It may be mentioned here that when hydrochloric
acts on nitric acid (aqua regia, Vol. I. p. 467) there is formed,
besides chlorine, the oxychlorides NOCl and NO_{2}Cl, which may be
regarded as chloranhydrides of nitric and nitrous acids (nitrogen
chloride, Vol. I. p. 476). The former boils at -5°, the latter at
+5°, the specific gravity of the first at -12° = 1·416, and at
-18° = 1·433 (Geuther), and of the second = 1·3; the first is
obtained from nitric oxide and chlorine, the second from nitric
peroxide and chlorine, and also by the action of phosphoric
chloride on nitric acid. If the gases evolved by aqua regia be
passed into cold and strong sulphuric acid, they form crystals of
the composition NHSO_{3} (like chamber crystals), which melt at
86°, and with sodium chloride form acid sodium sulphate and the
oxychloride NOCl. This chloranhydride of nitric acid is termed
_nitrosyl chloride_.
_Cyanogen chloride_, CNCl, is the gaseous chloranhydride of cyanic
acid; it is formed by the action of chlorine on aqueous mercury
cyanide, Hg(CN)_{2} + 2Cl_{2} = HgCl_{2} + 2CNCl. When chlorine
acts on cyanic acid, it forms not only this cyanogen chloride, but
also polymerides of it--a liquid, boiling at 18°, and a solid,
boiling at 190°. The latter corresponds with cyanuric acid, and
consequently contains C_{3}N_{3}Cl_{3}. Details concerning these
substances must be looked for in works on organic chemistry.
In order to show the general method for the preparation of acid
chloranhydrides, we will take that of acetic acid, CH_{3}·COOH, as an
example. Phosphorus pentachloride is placed in a glass retort, and acetic
acid poured over it; hydrochloric acid is then evolved, and the substance
distilling over directly after is a very volatile liquid, boiling at 50°,
and having all the properties of the chloranhydrides. With water it forms
hydrochloric and acetic acids. The reaction here taking place may be
explained thus: the substitution of the oxygen taken from the acetic acid
(from its carboxyl) by two atoms of chlorine from the PCl_{5} should be
as follows: CH_{3}·COOH + PCl_{5} = CH_{3}·COHCl_{2} + POCl_{3}. But the
compound CH_{3}·COHCl_{2} does not exist in a free state (because it
would indicate the possibility of the formation of compounds of the type
CX_{6}, and carbon only gives those of the type CX_{4}); it therefore
splits up into HCl and the chloranhydride CH_{3}·COCl. The general scheme
for the reaction of phosphorus pentachloride with hydrates ROH is exactly
the same as with water; namely, ROH with PCl_{5}, gives POCl_{3} + HCl +
RCl--that is a chloranhydride.[28 bis]
[28 bis] This reaction indeed proceeds very easily and completely with
a number of hydroxides, if they do not react on hydrochloric acid
and phosphorus oxychloride, which is the case when they have
alkaline properties. When the hydroxide is bibasic and is present
in excess, it not unfrequently happens that the elements of water
are taken up: R(OH)_{2} + PCl_{5} = RO + 2HCl + POCl_{3}. The
anhydride RO may then be converted into chloranhydride, RO +
PCl_{5} = RCl_{2} + POCl_{3}--that is, phosphorus pentachloride
brings about the substitution of O by Cl_{2}. Thus carbonyl
chloride, COCl_{2}, boron chloride, 2BCl_{3}, and succinic
chloride, C_{4}H_{4}O_{2}Cl_{2}, &c., are respectively obtained by
the action of phosphoric chloride on carbonic, boric, and succinic
anhydrides. Phosphorus pentachloride reacts in a similar manner on
the aldehydes, RCHO, forming RCHCl_{2}, and on the chloranhydrides
themselves--for example, with acetic chloride, CH_{3}.COCl (when
heated in a closed tube), it forms a substance having the
composition CH_{3}CCl_{3}.
Phosphorus trichloride and oxychloride act in a similar manner to
phosphoric chloride. When phosphorus trichloride acts on an acid,
3RHO + PCl_{3} = 3RCl + P(HO)_{3}. If a salt is taken, then by the
action of phosphorus oxychloride a corresponding chloranhydride
and salt of orthophosphoric acid are easily formed: 3R(KO) +
POCl_{3} = 3RCl + PO(KO)_{3}. The chloranhydride RCl is always
more volatile than its corresponding acid, and distils over before
the hydrate RHO. Thus acetic acid boils at 117°, and its
chloranhydride at 50°. Phosphoric and phosphorous acids are very
slightly volatile, whilst their chloranhydrides are comparatively
easily converted into vapour. The faculty of the chloranhydrides
to react at the expense of their own chlorine determines their
great importance in chemistry. For instance, suppose we require to
know the molecular formula of some hydrate which does not pass
into a state of vapour and does not give a chloranhydride with
hydrochloric acid--that is, which has not any basic or alkaline
properties; we must then endeavour to obtain this chloranhydride
by means of phosphoric chloride, and it frequently happens that
the corresponding chloranhydride is volatile. The resultant
chloranhydride is then converted into vapour, and its composition
is determined; and if we know its composition we are able to
decide that of its corresponding hydrate. Thus, for example, from
the formula of silicon chloride, SiCl_{4}, or of boron chloride,
BCl_{3}, we can judge the composition of their corresponding
hydrates, Si(HO)_{4}, B(HO)_{3}. Having obtained the
chloranhydride RCl or RCl_{_n_}, it is possible by its means to
obtain many other compounds of the same radicle R according to the
equation MX + RCl = MCl + RX. M may be = H, K, Ag, or other metal.
The reaction proceeds thus if M forms a stable compound with
chlorine--for example, silver chloride, hydrochloric acid, and R,
an unstable substance. Hence, a chloranhydride is frequently
employed for the formation of other compounds of a given radicle;
for instance, with ammonia they form amides RNH_{2}, and with
salts ROK, with anhydrides R_{2}O, &c.
Containing, as they do, chlorine, which easily reacts with hydrogen,
phosphorus pentachloride, trichloride, and oxychloride enter into
reaction with ammonia, and give a series of amide and nitrile compounds
of phosphorus. Thus, for example, when ammonia acts on the oxychloride we
obtain sal-ammoniac (which is afterwards removed by water) and an
orthophosphoric triamide, PO(NH_{2})_{3}, as a white insoluble powder on
which dilute acids and alkalis do not act, but which, when fused with
potassium hydroxide, gives potassium phosphate and ammonia like other
amides. When ignited, the triamide liberates ammonia and forms the
nitrile PON, just as urea, CO(NH_{2})_{2}, gives off ammonia and forms
the nitrile CONH. This nitrile, called _monophosphamide_, PON, naturally
corresponds with metaphosphoric acid, namely, with its ammonium salt.
NH_{4}PO_{3} - H_{2}O = PO_{2}·NH_{2}, an as yet unknown amide, and
PO_{2}·NH_{2} - H_{2}O gives the nitrile PON. This relation is confirmed
by the fact that PON, moistened with water, gives metaphosphoric acid
when ignited. It is the analogue of nitrous oxide, NON. It is a very
stable compound, more so than the preceding.[29]
[29] The reaction of ammonia on phosphorus pentachloride is more
complex than the preceding. This is readily understood: to the
oxychloride, POCl_{3}, tere corresponds a hydrate PO(OH)_{3}, and
a salt PO(NH_{4}O)_{3}, and consequently also an amide
PO(NH_{2})_{3}, whilst the pentachloride, PCl_{5}, has no
corresponding hydrate P(OH)_{5}, and therefore there is no amide
P(NH_{2})_{5}. The reaction with ammonia will be of two kinds:
either instead of 5 mol. NH_{3}, only 3 mol. NH_{3} or still less
will act; _i.e._ PCl_{2}(NH_{2})_{3}, PCl_{3}(NH_{2})_{2}, &c. are
formed; or else the pentachloride will act like a mixture of
chlorine with the trichloride, and then as the result there will
be obtained the products of the action of chlorine on those amides
which are formed from phosphorus trichloride and ammonia. It would
appear that both kinds of reaction proceed simultaneously, but
both kinds of products are unstable, at all events complex, and in
the result there is obtained a mixture containing sal-ammoniac,
&c. The products of the first kind should react with water, and we
should obtain, for example, PCl_{3}(NH_{2})_{2} + 2H_{2}O = 3HCl
and PO(HO)(NH_{2})_{2}. This substance has not actually been
obtained, but the compound PONH(NH_{2}) derived from it by
elimination of the elements of water is known, and is termed
_diphosphamide_; it is, however, more probable that it is a
nitrile than an amide, because only amides contain the group
NH_{2}. It is a colourless, stable, insoluble powder, which
possibly corresponds with pyrophosphoric acid, more especially
since when heated it evolves ammonia and gives and leaves
phosphoryl nitride, PON--that is, the nitrile of metaphosphoric
acid. The amide corresponding with the pyrophosphate
P_{2}O_{3}(NH_{4}O)_{4} should be P_{2}O_{3}(NH_{2})_{4}, and the
nitriles corresponding to the latter would be
P_{2}O_{2}N(NH_{2})_{3}, P_{2}ON_{2}(NH_{2})_{2}, and
P_{2}N_{3}(NH_{2}). The composition of the first is the same as
that of the above diphosphamide. The third pyrophosphoric nitrile
has a formula P_{2}N_{4}H_{2}, and this is the composition of the
body known as _phospham_, PHN_{2} (in a certain sense this is the
analogue of N_{3}H polymerised, Chapter VI.) Indeed, phospham has
been obtained by heating the products of the action of ammonia on
phosphoric chloride, as an insoluble and alkaline powder, which
gives ammonia and phosphoric acid when subjected to the action of
water. The same substance is obtained by the action of ammonium
chloride on phosphoric chloride (PNCl_{2} is first formed, and
reacts further with ammonia, forming phospham), and by igniting
the mass which is formed by the action of ammonia on phosphorus
trichloride. Formerly the composition of phospham was supposed to
be PHN_{2}, now there is reason to think that its molecular weight
is P_{3}H_{3}N_{6}.
The above compounds correspond with normal salts, but nitriles and
amides corresponding to acid salts are also possible, and they
will be acids. For example, the amide PO(HO)_{2}(NH_{2}), and its
nitrile, will be either PN(HO)_{2} or PO(HO)(NH), but at all
events of the composition PNH_{2}O_{2}, and having acid
properties. The ammonium salt of this _phosphonitrilic acid_ (it
is called phosphamic acid), PNH(NH_{4})O_{2}, is obtained by the
action of ammonia on phosphoric anhydride, P_{2}O_{5} + 4NH_3 =
H_{2}O + 2PNH(NH_{4})O_{2}. A non-crystalline soluble mass is thus
formed, which is dissolved in a dilute solution of ammonia and
precipitated with barium chloride, and the resultant barium salt
is then decomposed with sulphuric acid, and thus a solution of the
acid of the above composition is obtained.
It is evident from the theory of the formation of amides and
nitriles (Chapter IX.) that very many compounds of this kind can
correspond with the acids of phosphorus; but as yet only a few are
known. The easy transitions of the ortho-, meta-, and
pyrophosphoric acids, by means of the hydrogen of ammonia, into
the lower acids, and conversely, tend to complicate the study of
this very large class of compounds, and it is rarely that the
nature of a product thus obtained can be judged from its
composition; and this all the more that instances of isomerism and
polymerism, of mixture between water of crystallisation and of
constitution, &c., are here possible. Many data are yet needed to
enable us to form a true judgment as to the composition and
structure of such compounds. As the best proof of this we will
describe the very interesting and most fully investigated compound
of this class, PNCl_{2}, called _chlorophosphamide_, or nitrogen
chlorophosphorite. It is formed in small quantities when the
vapour of phosphoric chloride is passed over ignited sal-ammoniac.
Besson (1892) heated the compound PCl_{5}8NH_{3} (which is easily
and directly formed from PCl_{5} and NH_{3}) under a pressure of
about 50 mm. (of mercury) to 200°, and obtained brilliant crystals
of PNCl_{2}, which melted at 106° (in the residue after the
distillation of sal-ammoniacal phospham). The chlorine in it is
very stable--quite different from that in phosphoric chloride.
Indeed, the resultant substance is not only insoluble in water
(though soluble in alcohol and ether), but it is not even
moistened by it, and distils over, together with steam, without
being decomposed. In a free state it easily crystallises in
colourless prisms, fuses at 114°, boils at 250° (Gladstone,
Wichelhaus), and when fused with potash gives potassium chloride
and the amidonitrile of phosphoric acid. Judging from its formula
and the simplicity of its composition and reactions, it might be
thought that the molecular weight of this substance would be
expressed by the formula PCl_{2}N, that it corresponds with PON
and with PCl_{5} (like POCl_3), with the substitution of Cl_3 by
N, just as in POCl_3 two atoms of chlorine are replaced by oxygen;
but all these surmises are incorrect, because its vapour density
(referred to hydrogen--Gladstone, Wichelhaus) = 182--that is, the
molecular formula must be three times greater, P_{3}N_{3}Cl_{6}.
The polymerisation (tripling) is here of exactly the same kind as
with the nitriles.
The most important analogue of phosphorus is _arsenic_, the metallic
aspect of which and the general character of its compounds of the types
AsX_{3} and AsX_{5} at once recall the metals. The hydrate of its highest
oxide, arsenic acid (ortho-arsenic acid), H_{3}AsO_{4}, is an oxidising
agent, and gives up a portion of its oxygen to many other substances;
but, nevertheless, it is very like phosphoric acid. Mitscherlich
established the conception of isomorphism by comparing the salts of these
acids.[30]
[30] It is necessary to remark that, although arsenic is so closely
analogous to phosphorus (especially in the higher forms of
combination, RX_{3} and RX_{5}), at the same time it exhibits a
certain resemblance and even isomorphism with the corresponding
compounds of sulphur (especially the metallic compounds of the
type MAs, corresponding with MS). Thus compounds containing
metals, arsenic, and sulphur are very frequently met with in
nature. Sometimes the relative amounts of arsenic and sulphur
vary, so that an isomorphous substitution between the arsenides
and sulphides must be recognised. Besides FeS_{2} (ordinary
pyrites), and FeAs_{2}, iron forms an arsenical pyrites containing
both sulphur and arsenic, which from its composition, FeAsS or
FeS_{2}FeAs_{2}, resembles the two preceding.
Arsenic occurs _in nature_, not only combined with metals, but also,
although rarely, native and also in combination with sulphur in two
minerals--one red, _realgar_, As_{2}S_{2}, and the other yellow,
_orpiment_, As_{2}S_{3} (Chapter XX., Note 29). Arsenic occurs, but more
rarely, in the form of salts of arsenic acid--for instance, the so-called
cobalt and nickel blooms, two minerals which are found accompanying other
cobalt ores, are the arsenates of these metals. Arsenic is also found in
certain clays (ochres) and has been discovered in small quantities in
some mineral springs, but it is in general of rarer occurrence in nature
than phosphorus. Arsenic is most frequently extracted from arsenical
pyrites, FeSAs, which, when roasted without access of air, evolves the
vapour of arsenic, ferrous sulphide being left behind. It is also
obtained by heating arsenious anhydride with charcoal, in which case
carbonic oxide is evolved. In general, the oxides and other compounds are
very easily reduced. Solid _arsenic_ is a steel-grey brittle _metal_,
having a bright lustre and scaly structure. Its specific gravity is 5·7.
It is opaque and infusible, but volatilises as a yellow vapour which on
cooling deposits rhombohedral crystals.[30 bis] The vapour density of
arsenic is 150 times greater than that of hydrogen--that is, its
molecule, like that of phosphorus, contains 4 atoms, As_{4}. When heated
in the air, arsenic easily oxidises into white arsenious anhydride,
As_{2}O_{3}, but even at the ordinary temperature it loses its lustre
(becomes dull), owing to the formation of a coating of a lower oxide. The
latter appears to be as volatile as arsenious anhydride, and it is
probable that it is owing to the presence of this compound that the
vapours of arsenious compounds, when heated with charcoal (for example,
in the reducing flame of a blow-pipe), have the characteristic smell of
garlic, because the vapour of arsenic itself has not this odour.
[30 bis] According to Retgers (1893) the arsenic mirror (see further
on) is an unstable variety of metallic arsenic, whilst the brown
product which is formed together with it in Marsh's apparatus is a
lower hydride AsH. Schuller and McLeod (1894), however, recognise
a peculiar yellow variety of arsenic.
Arsenic easily combines with bromine and chlorine;[31] nitric acid and
aqua regia also oxidise it into the higher oxide, or rather its hydrate,
arsenic acid.[32] As far as is known, it does not decompose steam, and it
acts exceedingly slowly on those acids, like hydrochloric, which are not
capable of oxidising.
[31] Hydrochloric acid dissolves arsenious anhydride in considerable
quantities, and this is probably owing to the formation of
unstable compounds in which the arsenious anhydride plays the part
of a base. A compound called _arsenious oxychloride_, having the
composition AsOCl, is even known. It is formed when arsenious
anhydride is added little by little to boiling arsenic
trichloride, As_{2}O_{3} + AsCl_{3} = 3AsOCl. It is a transparent
substance, which fumes in air, and combines with water to form a
crystalline mass having the composition As_{2}(OH)_{4}Cl_{2}. When
heated it decomposes into arsenious chloride and a fresh
oxychloride of a more complex composition, As_{6}O_{8}Cl_{2}·
Arsenic trichloride, when treated with a small quantity of water,
forms the crystalline compound, As_{2}(HO)_{4}Cl_{2}, mentioned
above. These compounds resemble the basic salts of bismuth and
aluminium. The existence of these compounds shows that arsenic is
of a more metallic or basic character than phosphorus.
Nevertheless _arsenic trichloride_, AsCl_{3}, resembles phosphorus
trichloride in many respects. It is obtained by the direct action
of chlorine on arsenic, or by distilling a mixture of common salt,
sulphuric acid, and arsenious anhydride. The latter mode of
preparation already indicates the basic properties of the oxide.
Arsenious chloride is a colourless oily liquid, boiling at 130°,
and having a sp. gr. of 2·20. It fumes in air like other
chloranhydrides, but it is much more slowly and imperfectly
decomposed by water than phosphorus trichloride. A considerable
quantity of water is required for its complete decomposition into
hydrochloric acid and arsenious anhydride. It forms an excellent
example of the transition from true metallic chlorides to true
chloranhydrides of the acids. It hardly combines with chlorine,
_i.e._ if AsCl_{5} is formed it is very unstable. _Arsenic
tribromide_, AsBr_{3}, is formed as a crystalline substance,
fusing at 20° and boiling at 220°, by the direct action of
metallic arsenic on a solution of bromine in carbon bisulphide,
the latter being then evaporated. The specific gravity of arsenic
tribromide is 3·36. Crystalline arsenic tri-iodide, AsI_{3},
having a sp. gr. 4·39, may be obtained in a like manner; it may be
dissolved in water, and on evaporation separates out from the
solution in an anhydrous state--that is, it is not decomposed--and
consequently behaves like metallic salts. _Arsenic trifluoride_,
AsF_{3}, is obtained by heating fluor spar and arsenious anhydride
with sulphuric acid. It is a fuming, colourless, and very
poisonous liquid, which boils at 63° and has a sp. gr. of 2·73. It
is decomposed by water. It is very remarkable that fluorine forms
a pentafluoride of arsenic also, although this compound has not
yet been obtained in a separate state, but only in combination
with potassium fluoride. This compound, K_{3}AsF_{8}, is formed as
prismatic crystals when potassium arsenate, K_{3}AsO_{4}, is
dissolved in hydrofluoric acid.
[32] _Arsenic acid_, H_{3}AsO_{4}, corresponding with orthophosphoric
acid, is formed by oxidising arsenious anhydride with nitric acid,
and evaporating the resultant solution until it attains a sp. gr.
of 2·2; on cooling it separates in crystals having the above
composition. This hydrate corresponds with the normal salts of
arsenic acid; but on dissolving in water (without heating), and on
cooling a strong solution, crystals containing a greater amount of
water, namely, (AsH_{3}O_{4})_{2},H_{2}O, separate. This water,
like water of crystallisation, is very easily expelled at 100°. At
120° crystals having a composition identical with that of
pyrophosphoric acid, As_{2}H_{4}O_{7}, separate, but water, on
dissolving this hydrate with the development of heat, forms a
solution in no way differing from a solution of ordinary arsenic
acid, so that it is not an independent pyroarsenic acid that is
formed. Neither is there any true analogue of metaphosphoric acid,
although the compound AsHO_{3} is formed at 200°, and on
solidifying forms a mass having a pearly lustre and sparingly
soluble in cold water; but on coming into contact with warm water
it becomes very hot, and gives ordinary orthoarsenic acid in
solution. Arsenic acid forms three series of salts, which are
perfectly analogous to the three series of orthophosphates. Thus
the normal salt, K_{3}AsO_{4}, is formed by fusing the other
potassium arsenates with potassium carbonate; it is soluble in
water and crystallises in needles which do not contain water.
Di-potassium arsenate, K_{2}HAsO_{4}, is formed in solution by
mixing potassium carbonate and arsenic acid until carbonic
anhydride ceases to be evolved; it does not crystallise, and has
an alkaline reaction; hence it corresponds perfectly with the
sodium phosphate. As was mentioned above, arsenic acid itself acts
as an oxidising agent; for example, it is used in the manufacture
of aniline dyes for oxidising the aniline, and it is prepared in
large quantities for this purpose. When sulphuretted hydrogen is
passed through its solution, sulphuric acid and arsenious
anhydride are obtained in solution. Arsenic acid is very easily
soluble in water, and its solution has an exceedingly acid
reaction, and when boiled with hydrochloric acid evolves chlorine,
like selenic, chromic, manganic, and certain other higher metallic
acids.
_Arsenic anhydride_, As_{2}O_{5}, is produced when arsenic acid is
heated to redness. It must be carefully heated, as at a bright red
heat it decomposes into oxygen and arsenious anhydride. Arsenic
anhydride is an amorphous substance almost entirely insoluble in
water, but it attracts moisture from the air, deliquesces, and
passes into the acid. Hot water produces this transformation with
great ease.
_Arseniuretted hydrogen_, _arsine_, AsH_{3}, resembles phosphuretted
hydrogen in many respects. This colourless gas, which liquefies into a
mobile liquid at -40°, has a disagreeable garlic-like odour, is only
slightly soluble in water, and is exceedingly poisonous. Even in a small
quantity it causes great suffering, and if present to any considerable
amount in air it even causes death. The other compounds of arsenic are
also poisonous, with the exception of the insoluble sulphur compound and
some compounds of arsenic acid. Arseniuretted hydrogen, AsH_{3}, is
obtained by the action of water on the alloy of arsenic and sodium,
sodium hydroxide and arseniuretted hydrogen being formed. It is also
formed by the action of sulphuric acid on the alloy of arsenic and zinc:
Zn_{3}As_{2} + 3H_{2}SO_{4} = 2AsH_{3} + 3ZnSO_{4}.[33] The oxygen
compounds of arsenic are very easily reduced by the action of hydrogen at
the moment of its evolution from acids, and the reduced arsenic then
combines with the hydrogen; hence, if a certain amount of an oxygen
compound of arsenic be put into an apparatus containing zinc and
sulphuric acid (and thus serving for the evolution of hydrogen), the
hydrogen evolved will contain arseniuretted hydrogen. In this case it is
diluted with a considerable amount of hydrogen. But its presence in the
most minute quantities may be easily recognised from the fact that it is
_easily decomposed_ by heat (200° according to Brunn) into metallic
arsenic and hydrogen, and therefore if such impure hydrogen he passed
through a moderately-heated tube metallic arsenic will be deposited as a
bright layer on the part of the tube which was heated (_see_ Note 30
bis). This reaction is so sensitive that it enables the most minute
traces of arsenic to be discovered; hence it is employed in medical
jurisprudence, as a test in poisoning cases. It is easy to discover the
presence of arsenic in common zinc, copper, sulphuric and hydrochloric
acids, _&c._ by this method. It is obvious that in testing for poison by
Marsh's apparatus it is necessary to take zinc and sulphuric acid quite
free from arsenic. The arsenic deposited in the tube may be driven as a
volatile metal from one place to another in the current of hydrogen
evolved, owing to its volatility. This forms a distinction between
arseniuretted and antimoniuretted hydrogen, which is decomposed by heat
in just the same way as arseniuretted hydrogen, but the mirror given by
Sb is not so volatile as that formed by As.
[Illustration: FIG. 84.--Formation and decomposition of arseniuretted
hydrogen. Hydrogen is evolved in the Woulfe's bottle, and when the gas
comes off, a solution containing arsenic is poured through the funnel.
The presence of AsH_{3} is recognised from the deposition of a mirror of
arsenic when the gas-conducting tube is heated. If the escaping hydrogen
be lighted, and a porcelain dish be held in the flame, a film of arsenic
is deposited on it. The gas is dried by passing through the tube
containing calcium chloride. This apparatus is used for the detection of
arsenic by Marsh's test.]
[33] The formation of arseniuretted hydrogen is accompanied by the
absorption of 37,000 heat units, while phosphine evolves 18,000
(Ogier), and ammonia 27,000. Sodium (0·6 p.c.) amalgam, with a
strong solution of As_{2}O_{3}, gives a gas containing 86 vols. of
arsenic and 14 vols. of hydrogen (Cavazzi).
If hydrogen contains arseniuretted hydrogen, it also gives metallic
arsenic when it burns, because in the reducing flame of hydrogen the
oxygen attracted combines entirely with the hydrogen and not with the
arsenic, so that if a cold object, such as a piece of china, be held in
the hydrogen flame the arsenic will be deposited upon it as a metallic
spot.[34]
[34] This spot, or the metallic ring which is deposited on the heated
tube, may easily be tested as to whether it is really due to
arsenic or proceeds from some other substance reduced in the
hydrogen flame--for instance, carbon or antimony. The necessity
for distinguishing arsenic from antimony is all the more
frequently encountered in medical jurisprudence, from the fact
that preparations of antimony are very frequently used as
medicine, and antimony behaves in the hydrogen apparatus just like
arsenic, and therefore in making an investigation for poisoning by
arsenic it is easy to mistake it for antimony. The best method to
distinguish between the metallic spots of arsenic and antimony is
to test them with a solution of sodium hypochlorite, free from
chlorine, because this will dissolve arsenic and not antimony.
Such a solution is easily obtained by the double decomposition of
solutions of sodium carbonate and bleaching powder. A solution of
potassium chlorate acts in the same manner, only more slowly.
Further particulars must be looked for in analytical works.
Arseniuretted hydrogen, like phosphuretted hydrogen, is only
slightly soluble in water, has no alkaline properties--that is, it
does not combine with acids--and acts as a reducing agent. When
passed into a solution of silver nitrate it gives a blackish brown
precipitate of metallic silver, the arsenic being oxidised. If
acting on copper sulphate and similar salts, arseniuretted
hydrogen sometimes forms arsenides--_i.e._ it reduces the metallic
salt with its hydrogen, and is itself reduced to arsenic.
Sulphuric, and even hydrochloric, acid reduces arseniuretted
hydrogen to arsenic, and it is still more easily decomposed by
arsenious chloride, and with phosphorous chloride it gives the
compound PAs. Arseniuretted hydrogen gives metallic arsenic with
an acid solution of arsenious anhydride (Tivoli).
The most common compound of arsenic is the solid and volatile _arsenious
anhydride_, As_{2}O_{3}, which corresponds with phosphorous and nitrous
anhydrides. This very poisonous, colourless, and sweet-tasting substance
is generally known under the name of arsenic, or _white arsenic_. The
corresponding hydrate is as yet unknown; its solutions, when evaporated,
yield crystals of arsenious anhydride. It is chiefly prepared for the
dyer, and is also used as a vermin killer, and sometimes in medicine; it
is a product from which all other compounds of arsenic can be prepared.
It is obtained as a by-product in roasting cobalt and other ores
containing arsenic. Arsenical pyrites are sometimes purposely roasted for
the extraction of arsenious anhydride. When arsenical ores are burnt in
the air, the sulphur and arsenic are converted into the oxides
As_{2}O_{3} and SO_{2}. The former is a solid at the ordinary
temperature, and the latter gaseous, and therefore the arsenious
anhydride is deposited as a sublimate in the cooler portion of the flues
through which the vapours escape from the furnace. It collects in
condensing chambers especially constructed in the flues. The deposit is
collected, and after being distilled gives arsenious anhydride in the
form of a vitreous non-crystalline mass. This is one of the varieties of
arsenious anhydride, which is also known in two crystalline forms. When
sublimed--_i.e._ when it rapidly passes from the state of vapour to the
solid state--it appears in the regular system in the form of
octahedra.[35] It is obtained in the same form when it is crystallised
from acid solutions. The specific gravity of the crystals is 3·7. The
other crystalline form (in prisms) belongs to the rhombohedral system,
and is also formed by sublimation when the crystals are deposited on a
heated surface, or when it is crystallised from alkaline solutions.[36]
[35] According to Mitscherlich's determination, the vapour density of
arsenious anhydride is 199 (H = 1)--that is, it answers to the
molecular formula As_{4}O_{6}. Probably this is connected with the
fact that the molecule of free arsenic contains As_{4}. V. Meyer
and Biltz, however, showed (1889) that at a temperature of about
1,700° the vapour density of arsenic corresponds with the molecule
As_{2}, and not As_{4}, as at lower temperatures.
[36] Arsenious anhydride is obtained in an amorphous form after
prolonged heating at a temperature near to that at which it
volatilises, or, better still, by heating it in a closed vessel.
It then fuses to a colourless liquid, which on cooling forms a
transparent vitreous mass, whose specific gravity is only slightly
less than that of the crystalline anhydride. On cooling, this
vitreous mass undergoes an internal change, in which it
crystallises and becomes opaque, and acquires the appearance of
porcelain. The following difference between the vitreous and
opaque varieties is very remarkable: when the vitreous variety is
dissolved in strong and hot hydrochloric acid it gives crystals of
the anhydride on cooling, and this crystallisation _is accompanied
by the emission of light_ (which is visible in the dark), and the
entire liquid glows as the crystals begin to separate. The opaque
variety does not emit light when the crystals separate from its
hydrochloric acid solution. It is also remarkable that the
vitreous variety passes into the opaque form when it is
pounded--that is, under the action of a series of blows. Thus,
several varieties of arsenious anhydride are known, but as yet
they are not characterised by any special chemical distinctions,
and even differ but little in their specific gravities, so that it
cannot be said that the above differences are due to any isomeric
transformation--that is, to an arrangement of the atoms in the
molecule--but probably only depend on a difference in the
distribution of the molecules, or, in other terms, are physical
and not chemical variations. One part of the vitreous anhydride
requires twelve parts of boiling water for its solution, or
twenty-five parts at the ordinary temperature. The opaque variety
is less soluble, and at the ordinary temperature requires about
seventy parts of water for its solution.
Solutions of arsenious anhydride have a sweet metallic taste, and give
_a feeble acid reaction_. Its solubility increases with the admixture of
acids and alkalis. This shows the property of arsenious anhydride of
forming salts with acids and alkalis. And in fact compounds of it with
hydrochloric acid (Note 31), sulphuric anhydride (_see_ further on), and
with the alkali oxides are known.[37] If silver nitrate be added to a
solution of arsenious anhydride, it does not give any precipitate unless
a certain amount of the arsenious anhydride is saturated with an
alkali--for instance, ammonia. It then gives a precipitate of silver
arsenite, Ag_{3}AsO_{3}. This is yellow, soluble in an excess of ammonia,
and anhydrous; it distinctly shows that arsenious acid is tribasic, and
that it differs in this respect from phosphorous acid, in which only two
atoms of hydrogen can be replaced by metals.[38] The feeble acid
character of arsenious anhydride is confirmed by the formation of saline
compounds with acids. In this respect the most remarkable example is the
anhydrous compound with sulphuric acid, having the composition
As_{2}O_{3},SO_{3}. It is formed in the roasting of arsenical pyrites in
those spaces where the arsenious anhydride condenses, a portion of the
sulphurous anhydride being converted into sulphuric anhydride, SO_{3}, at
the expense of the oxygen of the air. The compound in question forms
colourless tabular crystals, which are decomposed by water with formation
of sulphuric acid and arsenious anhydride.[39]
[37] Arsenious anhydride does not oxidise in air, either in a dry state
or in solution, but in the presence of alkalis it absorbs oxygen
from the air, and acts as an excellent reducing agent. This
probably is connected with the fact that arsenic acid is much more
energetic than arsenious acid, and that it is arsenic acid which
is formed by the oxidation of the latter in the presence of
alkalis. Arsenious anhydride is easily reduced to arsenic by many
metals, even by copper.
[38] The feebleness of the acid properties of arsenious anhydride is
seen in the fact that if it be dissolved in ammonia water, and
then a still stronger solution of ammonia be added, prismatic
crystals separate having the composition of ammonium metarsenite,
NH_{4}AsO_{3}. This ammonium salt deliquesces in air, and loses
all its ammonia. The magnesium salt is tri-metallic,
Mg_{3}(AsO_{3})_{2}; it is insoluble in water, and is formed by
mixing an ammoniacal solution of arsenious anhydride with an
ammoniacal solution of a magnesium salt. It is insoluble even in
ammonia, although it dissolves in an excess of acids. Magnesium
hydroxide gives the same salt with arsenious solutions, and hence
magnesia is one of the best antidotes for arsenic poisoning. _The
arsenites of copper_ are much used in the manufacture of colours,
more especially of pigments. They are distinguished by their
insolubility in water and by their remarkably vivid green colour,
but at the same time by their poisonous character. Not only do
such pigments applied to wall papers or other materials easily
dust off from them, but they give exhalations containing AsH_{3}.
The cupric salts, CuX_{2}, when mixed with an alkaline solution of
arsenious acid, give a green precipitate of a copper salt called
_Scheele's green_. Its composition is probably CuHAsO_{3}. Ammonia
dissolves it, and gives a colourless solution, containing cuprous
arsenate--that is, the cupric compound is reduced and the arsenic
subjected to a further oxidation. The so-called _Schweinfurt
green_ was still more used, especially in former times; it is an
insoluble green cupric salt, which resembles the preceding in many
respects, but has a different tint. It is prepared by mixing
boiling solutions of arsenious acid and cupric acetate. Arsenious
acid forms an insoluble compound with ferric hydroxide, resembling
the phosphate; and this is the reason why freshly precipitated
oxide of iron is employed as an _antidote for arsenic_. The
freshly precipitated oxide of iron, taken immediately after
poisoning by arsenic, converts the arsenious acid into an
insoluble state, by forming a compound on which the acids of the
stomach have no action, so that the poisoning cannot proceed. It
is remarkable that the inhabitants of certain mountainous
countries accustom themselves to taking arsenic, as a means which,
according to their experience, helps to overcome the fatigue of
mountain ascents. Arsenious anhydride and certain of its salts are
also used in medicine, naturally only in small quantities. When
taken internally arsenic passes into the blood, and is mainly
excreted by the urine.
[39] Adie (1889) obtained compounds of As_{2}O_{3} with 1, 2, 4, and 8
SO_{3} by the direct action of ordinary and Nordhausen sulphuric
acid upon As_{2}O_{3}. Weber had previously obtained
As_{2}O_{3}SO_{3} (which disengages SO_{3} at 225°), and also
other As_{2}O_{3}_n_SO_{3} (where _n_ = 3, 6, and 8), by the
action of the vapours of SO_{3} upon As_{2}O_{3} at a definite
temperature. The compound As_{2}O_{3},8SO_{3} loses SO_{3} at
100°. Oxide of antimony, Sb_{2}O_{3}, gives similar compounds.
Adie (1891) also obtained (by the action of SO_{3} upon
H_{3}PO_{4}) a compound H_{3}PO_{4}3SO_{3} in the form of a
viscous liquid decomposed by water.
_Antimony_ (stibium), Sb = 120, is another analogue of phosphorus. In
its external appearance and the properties of its compounds it resembles
the metals still more closely than arsenic. In fact, antimony has the
appearance, lustre, and many of the characteristic properties of the
metals. Its oxide, Sb_{2}O_{3}, exhibits the earthy appearance of rust or
of lime, and has distinctly basic properties, although it corresponds
with nitrous and phosphorous anhydride, and is able, like them, to give
saline compounds with bases. At the same time antimony presents, in the
majority of its compounds, an entire analogy with phosphorus and arsenic.
Its compounds belong to the type SbX_{3} and SbX_{5}. It is found in
nature chiefly in the form of sulphide, Sb_{2}S_{3}. This substance
sometimes occurs in large masses in mineral veins and is known in
mineralogy under the name of antimony glance or _stibnite_, and
commercially as _antimony_ (Chapter XX., Note 29). The most abundant
deposits of antimony ore occur in Portugal (near Oporto on the Douro).
Besides which antimony partially or totally replaces arsenic in some
minerals; thus, for example, a compound of antimony sulphide and arsenic
sulphide with silver sulphide is found in red silver ore. But in every
case antimony is a rather rare metal found in few localities. In Russia
it is known to occur in Daghestan in the Caucasus. It is extracted
chiefly for the preparation of alloys with lead and tin, which are used
for casting printing type.[40] Some of its compounds are also used in
medicine, the most important in this respect being antimony
pentasulphide, Sb_{2}S_{5} (_sulfur auratum antimonii_), and tartar
emetic, which is a double salt derived from tartaric acid and has the
composition C_{4}H_{4}K(SbO)O_{6}. Even the native antimony sulphide is
used in large quantities as a purgative for horses and dogs. Metallic
antimony is extracted from the glance, Sb_{2}S_{2}, by roasting, when the
sulphur burns away and the antimony oxidises, forming the oxide
Sb_{2}O_{3}, which is then heated with charcoal, and thus reduced to a
_metallic state_. The reduction may be carried on in the laboratory on a
small scale by fusing the sulphide with iron which takes up the
sulphur.[40 bis]
[40] Printers' type consists of an alloy known as 'type-metal,'
containing usually about 15 parts of antimony to 85 parts of lead;
sometimes (for example, for stereotypes) from 10 to 15 per cent.
Bi or 8 per cent. Sn and even Cu is added. The hardness of the
alloy, which is essential for printing, evidently depends upon the
presence of antimony, but an excess must be avoided, since this
renders the alloy brittle, and the type after a time loses its
sharpness.
[40 bis] Antimony is prepared in a state of greater purity by heating
with charcoal the oxide obtained by the action of nitric acid on
the impure commercial metallic antimony. This is based on the fact
that by the action of the acid, antimony forms the oxide
Sb_{2}O_{3}, which is but slightly soluble in water. The arsenic,
which is nearly always present, forms soluble arsenious and
arsenic acids, and remains in solution. The purest antimony is
easily obtained from tartar emetic, by heating it with a small
quantity of nitre. Metallic antimony also occurs, although rarely,
native; and as it is very easily obtained, it was known to the
alchemists of the fifteenth century. Very pure metallic antimony
may be deposited by the electric current from a solution of
antimonious sulphide in sodium sulphide after the addition of
sodium chloride to the solution.
Metallic antimony has a white colour and a brilliant lustre; it remains
untarnished in the air, for the metal does not oxidise at the ordinary
temperature. It crystallises in rhombohedra, and always shows a
distinctly crystalline structure which gives it quite a different aspect
from the majority of the metals yet known. It is most like tellurium in
this respect. Antimony is brittle, so that it is very easily powdered;
its specific gravity is 6·7, it melts at about 432°, but only volatilises
at a bright red heat. When heated in the air--for instance, before the
blow-pipe--it burns and gives white odourless fumes, consisting of the
oxide. This oxide is termed antimonious oxide, although it might as well
be termed antimonious anhydride. It is given the first name because in
the majority of cases its compounds with acids are used, but it forms
compounds with the alkalis just as easily.
Antimonious oxide, like arsenious anhydride, crystallises either in
regular octahedra or in rhombic prisms; its specific gravity is 5·56;
when heated it becomes yellow and then fuses, and when further heated in
air it oxidises, forming an oxide of the composition Sb_{2}O_{4}.
Antimonious oxide is insoluble in water and in nitric acid, but it easily
dissolves in strong hydrochloric acid and in alkalis, as well as in
tartaric acid or solutions of its acid salts. When dissolved in the
latter it forms tartar emetic. It is precipitated from its solutions in
alkalis and acids (by the action of acids on the former and alkalis on
the latter). It occurs native but rarely. As a base it gives salts of the
type SbOX (as if the basic salts = SbX_{3}, Sb_{2}O_{3}) and hardly ever
forms salts, SbX_{3}. In the antimonyl salts, SbOX, the group SbO is
univalent, like potassium or silver. The oxide itself is (SbO)_{2}O, the
hydroxide, SbO(OH), &c.; tartar emetic is a salt in which one hydrogen of
tartaric acid is replaced by potassium and the other by antimonyl, SbO.
Antimonious oxide is very easily separated from its salts by any base,
but it must be observed that this separation does not take place in the
presence of tartaric acid, owing to the property of tartaric acid of
forming a soluble double salt--_i.e._ tartar emetic.[41]
[41] As antimonious oxide answers to the type SbX_{3}, it is evident
that compounds may exist in which antimony will replace three
atoms of hydrogen; such compounds have been to some extent
obtained, but they are easily converted by water into substances
corresponding with the ordinary formulæ of the compounds of
antimony. Thus tartar emetic, C_{4}H_{4}(SbO)KO_{6}, loses water
when heated, and forms C_{4}H_{2}SbKO_{6}--that is, tartaric acid,
C_{2}H_{6}O_{6}, in which one atom of hydrogen is replaced by
potassium and three by antimony. But this substance is reconverted
into tartar emetic by the action of water.
A similar compound is seen in that _intermediate oxide of
antimony_ which is formed when antimonious oxide is heated in air:
its composition is SbO_{2} or Sb_{2}O_{4}. This oxide may be
regarded as orthantimonic acid, SbO(HO)_{3}, in which three atoms
of hydrogen are replaced by antimony in that state in which it
occurs in oxide of antimony--_i.e._ SbO(SbO_{3}) = Sb_{2}O_{4}.
Oxide of antimony is also formed when antimonic acid is ignited;
it then loses water and oxygen, and gives this intermediate oxide
as a white infusible powder, of sp. gr. 6·7. It is somewhat
soluble in water, and gives a solution which turns litmus paper
red.
If metallic antimony, or antimonious oxide, be oxidised by an excess of
nitric acid and the resultant mass be carefully evaporated to dryness,
_metantimonic acid_, SbHO_{3}, is formed. Its corresponding potassium
salt, 2SbKO_{3},5H_{2}O, is prepared by fusing metallic antimony with
one-fourth its weight of nitre and washing the resultant mass with cold
water. This potassium salt is only slightly soluble in water (in 50
parts) and the sodium salt is still less so. An ortho-acid, SbH_{3}O_{4},
also appears to exist;[41 bis] it is obtained by the action of water on
antimony pentachloride, but it is very unstable, like the pentachloride,
SbCl_{5}, itself, which easily gives up Cl_{2}, leaving antimony
trichloride, SbCl_{3}, and this is decomposed by water, forming an
oxychloride--SbOCl, only slightly soluble in water. When antimonic acid
is heated to an incipient red heat, it parts with water and forms the
anhydride, Sb_{2}O_{5}, of a yellow colour and specific gravity 6·5.[42]
[41 bis] Beilstein and Blaese (1889), after preparing many salts of
antimonic acid, came to the conclusion that it is monobasic, but
all the salts still contain water, so that their general type is
mostly: MSbO_{3}3H_{2}O, for example, M = Li, Hg (salts of the
suboxide), 1/2 Pb, &c. The type of the ortho-salts, M_{2}SbO_{4},
is quite unknown, although it is reproduced in the thio-compounds,
for instance, Schlippe's salt, Na_{2}SbS_{4}, but this salt also
contains water of crystallisation, 9H_{2}O (Chapter XX., Note 29).
[42] Among the other compounds of antimony, _antimoniuretted hydrogen_,
SbH_{3}, resembles arseniuretted hydrogen in its mode of formation
and properties (it splits up at 150°, Brunn 1890; when liquified,
it boils at -65° and solidifies at -92°), whilst the halogen
compounds differ in many respects from those of arsenic. When
chlorine is passed over an excess of antimony powder, it forms
_antimony trichloride_, SbCl_{3}, but if the chlorine be in excess
it forms the _pentachloride_, SbCl_{5}. The trichloride is a
crystalline substance which melts at 72° and distils at 230°,
whilst the pentachloride is a yellow liquid, which splits up into
chlorine and the trichloride when heated; at 140° it begins to
give off chlorine abundantly, carrying away the vapour of the
trichloride with it; and at 200° the decomposition is complete,
and pure antimonious chloride only passes over. This property of
antimony pentachloride has caused it to be applied in many cases
for the transference of chlorine; all the more that when it has
given up its chlorine, it leaves the trichloride, which is able to
absorb a fresh amount of chlorine; and therefore many substances
which are unable to react directly with gaseous chlorine do so
with antimony pentachloride, and in the presence of a small
quantity of it chlorine will act on them, just as oxygen is able,
in the presence of nitrogen oxides, to oxidise substances which
could not be oxidised by means of free oxygen. Thus carbon
bisulphide is not acted on by chlorine at low temperatures--this
reaction requires a high temperature--but in the presence of
antimony pentachloride its conversion into carbon tetrachloride
takes place at low temperatures. Antimony tri- and pentachloride,
having the character of chloranhydrides, fume in air, attract
moisture, and are decomposed by water, forming antimonious and
antimonic acids. But in the first action of water the trichloride
does not evolve all its chlorine as hydrochloric acid, which is
intelligible in view of the fact that antimonious anhydride is
also a base, and is therefore able to react with acids; indeed
antimony sulphide dissolved in an excess of hydrochloric acid
(hydrogen sulphide is evolved) gives an aqueous solution of
antimony trichloride, which, when carefully distilled, even gives
the anhydrous compound. Antimony trichloride is only decomposed by
an excess of water, and then not completely, for with a large
quantity of water it forms _powder of algaroth_--_i.e._ antimony
oxychloride. The first action of water consists in the formation
of _oxychloride_, SbOCl--that is, a salt corresponding to oxide of
antimony as a base. If antimony oxide or antimony chloride be
dissolved in an excess of hydrochloric acid, and the solution
diluted with a considerable amount of water, then this same powder
of algaroth is precipitated. The composition varies with the
relative amount of water; namely, between the limits SbOCl and
Sb_{4}O_{5}Cl_{2}. The latter compound is, as it were, a basic
salt of the former, because its composition = 2(SbOCl)Sb_{2}O_{3}.
With bromine and iodine, antimony forms compounds similar to those
with chlorine. Antimonious bromide, SbBr_{3}, crystallises in
colourless prisms, melts at 94°, and boils at 270°; antimonious
iodide, SbI_{3}, forms red crystals of sp. gr. 5·0; antimony
trifluoride, SbF_{3} separates from a solution of antimonious
oxide in hydrofluoric acid, and SbF_{5} is formed by a similar
treatment of antimonic acid. The latter gives easily-soluble
double salts with the fluorides of the metals of the alkalis.
De Haën (1887) obtained very stable double soluble salts,
SbF_{3},KCl (100 parts of water dissolve 57 parts of salt),
SbF_{3},K_{2}SO_{4}, &c., which he proposed to make use of in the
arts as very easily crystallisable and soluble salts of antimony.
Engel, by passing hydrochloric acid gas into a saturated solution
of antimonious chloride at 0°, obtained a compound
HCl,2SbCl_{3},2H_{2}O, and with the pentachloride a compound
SbCl_{5},5HCl,10H_{2}O. Bismuth trichloride, BiCl_{3}, gives a
similar compound.
Saunders (1892) obtained 5RbCl,3SbCl_{3} and RbCl,SbCl_{3}. Ditte
and Metzner (1892) showed that Sb and Bi dissolve in hydrochloric
acid only owing to the participation of the oxygen of the air or
of that dissolved in the acid.
The heaviest analogue of nitrogen and phosphorus is _bismuth_, Bi = 208.
Here, as in the other groups, the basic, metallic, properties increase
with the atomic weight. Bismuth does not give any hydrogen compound and
the highest oxide, Bi_{2}O_{5}, is a very feeble acid oxide. Bismuthous
oxide, Bi_{2}O_{3}, is a base, and bismuth itself a perfect metal. To
explain the other properties of bismuth it must further be remarked that
in the eleventh series it follows mercury, thallium and lead, whose
atomic weights are near to that of bismuth, and that therefore it
resembles them and more especially its nearest neighbour, lead. Although
PbO and PbO_{2}, represent types different from Bi_{2}O_{3} and
Bi_{2}O_{5}, they resemble them in many respects, even in their external
appearance, moreover the lower oxides both of Pb and Bi are basic and the
higher acid, which easily evolve oxygen. But judging by the formula,
Bi_{2}O_{3} is a more feeble base than PbO. They both easily give basic
salts.
Bismuth forms compounds of two types, BiX_{3} and BiX_{5},[43] which
entirely recall the two types we have already established for the
compounds of lead. Just as in the case of lead, the type PbX_{2}, is
basic, stable, easily formed, and passes with difficulty into the higher
and lower types, which are unstable, so also in the case of bismuth the
type of combination BiX_{3} is the usual basic form. The higher type of
combination, BiX_{5},[44] in fact behaves toward this stable type,
BiX_{3}, in exactly the same manner as lead dioxide does to the monoxide;
and bismuthic acid is obtained by the action of chlorine on bismuth oxide
suspended in water, in exactly the same way as lead dioxide is obtained
from lead oxide. It is an oxidising agent like lead dioxide, and even the
acid character in bismuthic acid is only slightly more developed than in
lead dioxide. Here, as in the case of lead (minium), intermediate
compounds are easily formed in which the bismuth of the lower oxide plays
the part of a base combined with the acid which is formed by the higher
form of the oxidation of bismuth.
[43] Metallic bismuth is very easily obtained when the compounds of the
oxide are reduced by powerful reducing agents, but when less
powerful reducing agents--for example, stannous oxide--are taken,
bismuth suboxide is formed as a black crystalline powder. It is a
compound of the type BiX_{2}, its composition being BiO; it is
decomposed by acids into the metal and oxide, which passes into
solution.
[44] The type BiX_{5} is represented by the pentoxide, Bi_{2}O_{5}, its
metahydrate, Bi_{2}O_{5},H_{2}O, or BiHO_{3}, known as bismuthic
acid, and the pyrohydrate, Bi_{2}H_{4}O_{7}. _Bismuth pentoxide_
is obtained by the prolonged passage of chlorine through a boiling
solution of potassium hydroxide (sp. gr. 1·38), containing bismuth
oxide in suspension; the precipitate is washed with water, with
boiling nitric acid (but not for long, as otherwise the bismuthic
acid is decomposed), then again with water, and finally the
resultant bright red powder of the hydrate BiHO_{3} is dried at
125°. The prolonged action of nitric acid on bismuthic anhydride,
Bi_{2}O_{5}, results in the formation of the compound
Bi_{2}O_{4},H_{2}O, which decomposes in moist air, forming
Bi_{2}O_{3}. The density of bismuthic anhydride is 5·10, of the
tetroxide, Bi_{2}O_{4}, 3·60, and of bismuthic acid, BiHO_{3},
5·75. _Pyrobismuthic acid_, Bi_{2}H_{4}O_{7}, forms a brown
powder, which loses a portion of its water at 150°, and decomposes
on further heating, with the evolution of oxygen and water. It is
obtained by the action of potassium cyanide on a solution of
bismuth nitrate. The meta-salts of bismuthic acid are known, for
example KBiO_{3}. They generally occur, however, in combinations
with metabismuthic acid itself. Thus André (1891) took a solution
of the double salt of BiBr_{3} and KBr, treated it with bromine
after adding ammonia, and obtained a red-brown precipitate, which
after being washed (for several weeks) had the composition
KBiO_{3},HBiO_{3} When washed with dilute nitric acid this salt
gave bismuthic acid.
[Illustration: FIG. 85.--Furnace used for the extraction of bismuth from
its ores.]
In nature, bismuth occurs in only a few localities and in small
quantities, most frequently in a native state, and more rarely as oxide
and as a compound of bismuth sulphide with the sulphides of other metals,
and sometimes in gold ores. It is extracted from its native ores by
simple fusion in the furnace shown in fig. 85. This furnace contains an
inclined iron retort, into the upper extremity of which the ore is
charged, and the molten _metal_ flows from the lower extremity. It is
refined by re-melting, and the pure metal may be obtained by dissolving
in nitric acid, decomposing the resultant salt with water, and reducing
the precipitate by heating it with charcoal. Bismuth is a metal which
crystallises very well from a molten state. Its specific gravity is 9·8;
it melts at 269°, and if it be melted in a crucible, allowed to cool
slowly, and the crust broken and the remaining molten liquid poured out,
perfect rhombohedral crystals of bismuth are obtained on the sides of the
crucible.[44 bis] It is brittle, has a grey-coloured fracture with a
reddish lustre, is not hard, and is but very slightly ductile and
malleable; it volatilises at a white heat and easily oxidises. It recalls
antimony and lead in many of its properties. When oxidised in air, or
when the nitrate is ignited, bismuth forms the _oxide_, Bi_{2}O_{3}, as a
white powder which fuses when heated and resembles massicot. The addition
of an excess of caustic potash to a solution of a bismuthous salt gives a
white precipitate of the hydroxide, BiO(OH), which loses its water and
gives the anhydrous oxide when boiled with a solution of caustic potash.
Both the hydroxide and oxide easily dissolve in acids and form bismuthous
salts.
[44 bis] Hérard (1889) obtained a peculiar variety of bismuth by
heating pure crystalline bismuth to a bright red heat in a stream
of nitrogen. A greenish vapour was deposited in the cooler
portions of the apparatus in the form of a grey powder, which
under the microscope had the appearance of minute globules. An
atmosphere of nitrogen is necessary for this transformation, other
gases such as hydrogen and carbonic oxide do not favour the
transition. The resultant amorphous bismuth fuses at 410° (the
crystalline variety at 269°), sp. gr. 9·483. (Does it not contain
a nitride?)
_Bismuthous oxide_, Bi_{2}O_{3}, is a feeble and unenergetic base. The
normal hydroxide of the oxide Bi_{2}O_{3} is Bi(OH)_{3}; it parts with
water and forms a metahydroxide (bismuthyl hydroxide), BiO(OH). Both of
these hydroxides have their corresponding saline compounds of the
composition BiX_{3} and BiOX. And the form BiOX is nothing else but the
type of the basic salt, because 3ROX = RX + R_{2}O_{3}. It is evident
that in the type BiX_{3} the bismuth replaces three atoms of hydrogen.
And indeed with phosphoric acid solutions of the bismuthous salts give a
precipitate of the composition BiPO_{4}. On the other hand, in the form
of compounds BiOX or Bi(OH)_{2}X, the univalent group (BiO) or
(BiH_{2}O_{2}) is combined with X. Many bismuth salts are formed
according to the type BiOX. For instance the carbonate, (BiO)_{2}CO_{3},
which corresponds with the other carbonates M_{2}CO_{3}. It is obtained
as a white precipitate when a solution of sodium carbonate is added to a
solution of a bismuth salt.[45] The compound radicle BiO is not a special
natural grouping, as it was formerly represented to be; it is simply a
mode of expression for showing the relation between the compound in
question and the compounds of other oxides.
[45] Basic bismuth carbonate is employed for whitening the skin
(veloutine, &c.)
Three _salts of nitric acid_ are known containing bismuthous oxide. If
metallic bismuth or its oxide be dissolved in nitric acid, it forms a
colourless transparent solution containing a salt which separates in
large transparent crystals containing Bi(NO_{3})_{3},5H_{2}O. When heated
at 80° these crystals melt in their water of crystallisation, and in so
doing lose a portion of their nitric acid together with water, forming a
salt whose empirical formula is Bi_{2}N_{2}H_{2}O_{9}. If the preceding
salt belongs to the type BiX_{3}, this one should belong to the form
BiOX, because it = BiO(NO_{3}) + Bi(H_{2}O_{2})(NO_{3}). This salt may be
heated to 150° without change. When the first colourless crystalline salt
dissolves in water _it is decomposed_. There is no decomposition if an
excess of acid be added to the water--that is to say, the salt is able to
exist in an acid solution without decomposing, without separation of the
so-called basic salt--but by itself it cannot be kept in solution; water
decomposes this salt, acting on it like an alkali. In other words the
basic properties of bismuthic oxide are so feeble that even water acts by
taking up a portion of the acid from it. Here we see one of the most
striking facts, long since observed, confirming that action of water on
salts about which we have spoken in Chapter X. and elsewhere. This action
on water may be expressed thus:--BiX_{3} + 2H_{2}O = Bi(OH)_{2}X + 2XH. A
salt of the type Bi(OH)_{2}X is obtained in the precipitate. But if the
quantity of acid, HX, be increased, the salt BiX_{3} is again formed and
passes into solution. The quantity of the salt BiOX which passes into
solution on the addition of a given quantity of acid depends indisputably
on the amount (mass) of water (Muir). The solution, which is perfectly
transparent with a small amount of water, becomes cloudy and deposits the
salt of the type BiOX, when diluted. The white flaky precipitate of
Bi(OH)_{2}NO_{3} formed from the normal salt Bi(NO_{3})_{3} by mixing it
with five parts of water, and in general with a small amount of water, is
used in medicine under the name of magistery of bismuth.[46]
[46] With an excess of water a further quantity of acid is separated
and a still more basic salt formed. The ultimate product, on which
an excess of water has apparently no action whatever, is a
substance having the composition BiO(NO_{3}).BiO(OH). In the
latter salt we see the limit of change, and this limit appears to
show that the type of the saline compounds of bismuthic oxide is
of the form Bi_{2}X_{6}, and not BiX_{3}; but it is very probable,
on the basis of the examples which we considered in the case of
lead, that this type should be still further polymerised in order
to give a correct idea of the type of the bismuthous compounds. If
we refer all the bismuthous compounds to this type, Bi_{2}X_{6},
we shall obtain the following expression for the composition of
the nitrates: normal salt, Bi_{2}(NO_{3})_{6}, first basic salt,
Bi_{2}O(OH)_{2}(NO_{3})_{2}, magistery of bismuth,
Bi_{2}(OH)_{4}(NO_{3})_{2}, and the limiting form
Bi_{2}O_{2}(OH)(NO_{3}).
The general character of bismuthous oxide in its compounds is well
exemplified in the nitrate; bismuthous chloride, BiCl_{3}, which
is obtained by heating bismuth in chlorine, or by dissolving it in
aqua regia, and then distilling without access of air, is also
decomposed by water in exactly the same manner, and forms basic
salts--for instance, first, BiOCl, like the above salt of nitric
acid. Bismuth chloride boils at 447° and probably its formula is
BiCl_{3}. Polymerisation may take place in some compounds and not
in others. A volatile compound of the composition
Bi(C_{2}H_{5})_{3} is also known as a liquid which is insoluble in
water and decomposes with explosion when heated at 130°. Double
salts containing chloride of bismuth are: 2(KCl)BiCl_{3}2H_{2}O
(from a solution of Bi_{2}O_{3} and KCl in hydrochloric acid) and
KClBiCl_{3}H_{2}O. Bigham (1892) also obtained KBr(SO_{4})_{2} in
tabular crystals by treating the above-named double salt with
strong sulphuric acid. The composition of this salt recalls that
of alum.
Metallic bismuth is used in the preparation of fusible alloys. The
addition of bismuth to many metals renders them very hard, and at the
same time generally lowers their melting point to a considerable extent.
Thus Wood's metal, which contains one part of cadmium, one part of tin,
two parts of lead, and four parts of bismuth, fuses at about 60°, and in
general many alloys composed of bismuth, tin, lead, and antimony melt
below or about the boiling point of water.[47]
[47] As the metals contained in alloys like the above (bismuth, lead,
tin, cadmium) are difficultly volatile and their alloys are
fusible, they may be employed in the place of mercury in many
physical experiments conducted at or above 70°, and they offer the
advantage that they do not give any vapour having an appreciable
tension (mercury at 100°, 0·75 mm.) Bismuth expands in passing
into a molten state, but it has a temperature of maximum density.
According to Luedeking the mean coefficient of expansion of liquid
bismuth is 0·0000442 (between 270° and 303°), and of solid bismuth
0·0000411.
Just as in group II., side by side with the elements zinc, cadmium, and
mercury in the uneven series, we found calcium, strontium, and barium in
the even series; and as in group IV., parallel to silicon, germanium,
tin, and lead, we noticed thallium, zirconium, cerium, and thorium; so
also in group V. we find, beside those elements of the uneven series just
considered by us, a series of analogues in the even series, which, with a
certain degree of similarity (mainly quantitative, or relative to the
atomic weights), also present a series of particular (qualitative)
independent points of distinction. In the even series are known
_vanadium_, which stands between titanium and chromium, _niobium_,
between zirconium and molybdenum, and _tantalum_, situated near tungsten
(an element of group VI. like chromium and molybdenum). Just as bismuth
is similar in many respects to its neighbour lead, so also do these
neighbouring elements resemble each other, even in their external
appearance, not to mention the quality of their compounds, naturally
taking into account the differences of type corresponding with the
different groups. The occurrence in group V. determines the type of the
oxides, R_{2}O_{3} and R_{2}O_{5}, and the development of an acid
character in the higher oxides. The occurrence in the even series
determines the absence of volatile compounds, RH_{3}, for these metals,
and a more basic character of the oxides of a given composition than in
the uneven series, &c.[48] Vanadium, niobium, and tantalum belong to the
category of rare metals, and are exceedingly difficult to obtain pure,
more especially owing to their similarity to, and occurrence with,
chromium, tungsten and other metals, and also in combination among
themselves; therefore it is natural that they have been far from
completely studied, although since 1860 chemists have devoted not a
little time to their investigation. The researches carried out by
Marignac, at Geneva, on niobium, and by Sir Henry Roscoe, at Manchester,
on vanadium deserve special attention. The undoubted external resemblance
of the compounds of chromium and vanadium, as well as the want of
completeness in the knowledge of the compounds of vanadium, long caused
its oxides to be considered analogous in atomic composition to those
formed by chromium. The higher oxide of vanadium was therefore supposed
to have the formula VO_{3}. But the fact of the matter is, that the
chemical analogy of the elements does not hold in one direction only;
vanadium is at one and the same time the analogue of chromium, and
consequently of the elements like sulphur of group VI, and also the
analogue of phosphorus, arsenic, and antimony; just as bismuth stands in
respect to lead and antimony. Investigation has shown that the compounds
of vanadium are always accompanied by those of phosphorus as well as of
iron, and that it is even more difficult to separate it from the
compounds of phosphorus than from those of iron and tungsten. We should
have to extend our description considerably if we wished to give the
complete history, even of vanadium alone, not to mention niobium and
tantalum, all the more as questions would not unfrequently arise
concerning the compounds of these elements which have not yet been fully
elucidated. We shall therefore limit ourselves to pointing out the most
important features in the history of these elements, the more so since
the minerals themselves in which they occur are exceedingly rare and only
accessible to a few investigators.
[48] Although, guided by Brauner, who showed that didymium gives a
higher oxide, Di_{2}O_{5}, I place this element in the fifth
group, still I am not certain as to its position, because I
consider that the questions relating to this metal are still far
from being definitely answered.
An important point in the history of the members of this group is the
circumstance that they form volatile compounds with chlorine, similar to
the compounds of the elements of the phosphorus group, namely, of the
type RX_{5}. The vapour densities of the compounds of this kind were
determined, and served as the most important basis for the explanation of
the atomic composition of these molecules. In this we see the power of
general and fundamental laws, like the law of Avogadro-Gerhardt. An
oxychloride, VOCl_{3}, is known for vanadium, which is the perfect
analogue of phosphorus oxychloride. It was formerly considered to be
vanadium chloride, for just as in the case of uranium (Chapter XXI.), its
lower oxide, VO, was considered to be the metal, because it is
exceedingly difficultly reduced--even potassium does not remove all the
oxygen, besides which it has a metallic appearance, and decomposes acids
like a metal; in a word, it simulates a metal in every respect. _Vanadium
oxychloride_ is obtained by heating the trioxide, V_{2}O_{3}, mixed with
charcoal, in a current of hydrogen; the lower oxide of vanadium is then
formed, and this, when heated in a current of dry chlorine, gives the
oxychloride VOCl_{3} as a reddish liquid which does not act on sodium and
may be purified by distillation over this metal. It fumes in the air,
giving reddish vapours; it reacts on water, forming hydrochloric and
vanadic acids; hence, on the one hand it is very similar to phosphorus
oxychloride, and on the other hand to chromium oxychloride, CrO_{2}Cl_{2}
(Chapter XXI.). It is of a yellow colour, its specific gravity is 1·83,
it boils at 120°, and its vapour density is 86 with respect to hydrogen;
therefore the above formula expresses its molecular weight.[49]
[49] When the vapours of vanadium oxychloride are heated with zinc in a
closed tube at 400°, they lose a portion of their chlorine and
form a green crystalline mass of sp. gr. 2·88, which is
deliquescent in air and has the composition VOCl_{2}. Only its
vapour density is unknown, and it would be extremely important to
determine whether its molecular composition is that given above,
or whether it corresponds with the formula V_{2}O_{2}Cl_{4}.
Another less volatile oxychloride, VOCl, is formed with it as a
brown insoluble substance, which is, however, soluble in nitric
acid like the preceding. Roscoe obtained a still less chlorinated
substance, namely, (VO)_{2}Cl; but it may only consist of a
mixture of VO and VOCl. At all events, we here find a graduated
series such as is met with in the compounds of very few other
elements.
_Vanadic anhydride_, V_{2}O_{5}, is obtained either in small quantities
from certain clays where it accompanies the oxides of iron (hence some
sorts of iron contain vanadium) and phosphoric acid, or from the rare
minerals: _volborthite_, CuHVO_{4}, or basic vanadate of copper;
_vanadinite_, PbCl_{2}3Pb_{3}(VO_{4})_{2}; lead vanadate,
Pb_{3}(VO_{4})_{2}, &c. The latter salts are carefully ignited for some
time with one-third of their weight of nitre; the fused mass thus formed
is powdered and boiled in water: the yellow solution obtained contains
potassium vanadate. The solution is neutralised with acid, and barium
chloride added; a meta-salt, Ba(VO_{3})_{2}, is then precipitated as an
almost insoluble white powder, which gives a solution of vanadic acid
when boiled with sulphuric acid. (The precipitate is at first yellow, as
long as it remains amorphous, but it afterwards becomes crystalline and
white.) The solution thus obtained is neutralised with ammonia, which
thus forms ammonium (meta) vanadate, NH_{4}VO_{3}, which, when
evaporated, gives colourless crystals, insoluble in water containing
sal-ammoniac; hence this salt is precipitated by adding solid
sal-ammoniac to the solution. Ammonium vanadate, when ignited, leaves
vanadic acid behind. In this it differs from the corresponding chromium
salt, which is deoxidised into chromium oxide when ignited. In general,
vanadic acid has but a small oxidising action. It is reduced with
difficulty, like phosphoric or sulphuric acid, and in this differs from
arsenic and chromic acids. Vanadic acid, like chromic acid, separates
from its solution as the anhydride V_{2}O_{5}, and not in a hydrous
state. Vanadic anhydride, V_{2}O_{5}, forms a reddish-brown mass, which
easily fuses and re-solidifies into transparent crystals having a violet
lustre (another point of resemblance to chromic acid); it dissolves in
water, forming a yellow solution with a slightly acid reaction.[50]
[50] Strong acids and alkalis dissolve vanadic anhydride in
considerable quantities, forming yellow solutions. When it is
ignited, especially in a current of hydrogen, it evolves oxygen
and forms the lower oxides; V_{2}O_{4} (acid solutions of a green
colour, like the salts of chromic oxide), V_{2}O_{3}, and the
lowest oxide, VO. The latter is the metallic powder which is
obtained when the vanadium oxychloride is heated in an excess of
hydrogen, and was formerly mistaken for metallic vanadium. When a
solution of vanadic acid is treated with metallic zinc it forms a
blue solution, which seems to contain this oxide. It acts as a
reducing agent (and forms a close analogue to chromous oxide,
CrO). Metallic _vanadium_ can only be obtained from vanadium
chloride which is quite free from oxygen. Moissan (1893) obtained
it by reducing the oxide with carbon in the electric furnace, and
considered it to be most infusible of the metals in the series Pt,
Cr, Mo, U, W, and V (he also obtained a compound of vanadium and
carbon). The specific gravity of this metal is 5·5. It is of a
grey-white colour, is not decomposed by water, and is not oxidised
in air, but burns when strongly heated, and can be fused in a
current of hydrogen (forming perhaps a compound with hydrogen). It
is insoluble in hydrochloric acid, but easily dissolves in nitric
acid, and when fused with caustic soda it forms sodium vanadate.
As regards the salts of vanadic acid, three different classes are
known; the first correspond with metavanadic acid, VMO_{3} =
M_{2}OV_{2}O_{5}, the second correspond with the dichromates--that
is, have the composition V_{4}M_{2}O_{11}, which is equal to
M_{2}O + 2V_{2}O_{5}--and the third correspond with orthovanadic
acid, VM_{3}O_{4} or 3M_{2}O + V_{2}O_{5}. The latter are formed
when vanadic anhydride is fused with an excess of an alkaline
carbonate.
Vanadic acid gives the so-called 'complex' acids (which are
considered more fully in Chapter XXI. in speaking of Mo and
W)--_i.e._ acids formed of two acids assimilated into one. Thus
Friedheim (1890) obtained phosphor-vanadic acid, and
Schmitz-Dumont (1890) a similar arseno-vanadic acid. The former is
obtained by heating V_{2}O_{5} with sirupy phosphoric acid. The
resultant golden-yellow tabular crystals have the composition
H_{2}OV_{2}O_{5}P_{2}O_{5}9H_{2}O, and there are corresponding
salts--for example, (NH_{4})_{2}V_{2}O_{5}P_{2}O_{5} with 3 and
7H_{2}O, &c. These salts cannot be separated by crystallisation,
so that there are 'complexes' of these acids in a whole series of
salts (and also in nature). It may be supposed (Friedheim) that
V_{2}O_{5} here, as it were, plays the part of a base, or that
those acids may be looked upon as double salts. Among the true
double salts of vanadium (Nb and Ta) very many are known among the
fluorides, such as VF_{3}2NH_{4}F, VOF_{2}2NH_{4}F,
VO_{2}F,3NH_{4}F, &c. (Pettersson, Piccini, and Georgi, 1890-92).
Vanadium was discovered at the beginning of this century by
Del-Rio, and afterwards investigated by Sefström, but it was only
in 1868 that Roscoe established the above formulæ of the vanadic
compounds.
_Niobium and tantalum_[51] occur as acids in rare minerals, and are
mainly extracted from _tantalite_ and _columbite_, which are found in
Bavaria, Finland, North America, and in the Urals. These minerals are
composed of the ferrous salts of niobic and tantalic acids; they contain
about 15 per cent. of ferrous oxide in isomorphous mixture with manganous
oxide, in combination with various proportions of tantalic and niobic
anhydrides. These minerals are first fused with a considerable amount of
potassium bisulphate, and the fused mass is boiled in water, which
dissolves the ferrous and potassium salts and leaves an insoluble residue
of impure niobic and tantalic acids. This raw product is then treated
with ammonium sulphide, in order to extract the tin and tungsten, which
pass into solution. The residue containing the acids (according to
Marignac) is then treated with hydrofluoric acid, in which it entirely
dissolves, and potassium fluoride is added to the resultant hot solution;
on cooling, a sparingly soluble double fluoride of potassium and tantalum
separates out in fine crystals, while the much more soluble niobium salt
remains in solution. The difference in the solubility of these double
salts in water acidified with hydrofluoric acid (in pure water the
solution becomes cloudy after a certain time) is so great that the
tantalum compound requires 150 parts of water for its solution, and the
niobium compound only 13 parts. The Greenland columbite (specific gravity
5·36) only contains niobic acid, and that from Bodenmais, Bavaria
(specific gravity 6·06) almost equal quantities of tantalic and niobic
acids. Having isolated tantalic and niobic salts, Marignac found that the
relation between the potassium and fluorine in them is very
variable--that is, that there exist various double salts of fluoride of
potassium, and of the fluorides of the metals of this group, but that
with an excess of hydrofluoric acid both the tantalum and niobium
compounds contain seven atoms of fluorine to two of potassium, whence it
must be concluded that the simplest formula for these double salts will
be K_{2}RF_{7} = RF_{5},2KF; that is, that the type of the higher
compounds of niobium and tantalum is RX_{5}, and hence is similar to
phosphoric acid. A chloride, TaCl_{5}, may be obtained from pure tantalic
acid by heating it with charcoal in a current of chlorine. This is a
yellow crystalline substance, which melts at 211°, and boils at 241°; its
vapour density with respect to hydrogen is 180, as would follow from the
formula TaCl_{5}. It is completely decomposed by water into tantalic and
hydrochloric acids. _Niobium pentachloride_ may be prepared in the same
manner; it fuses at 194°, and boils at 240°. When treated with water this
substance gives a solution containing niobic acid, which only separates
out on boiling the solution. Delafontaine and Deville found its vapour
density to be 9·3 (air = 1), as is shown by its formula NbCl_{5}.[52]
[51] The researches made by Roscoe were preceded by those of Marignac
in 1865, on the _compounds_ of _niobium_ and _tantalum_, to which
were also ascribed different formulæ from those now recognised.
Tantalum was discovered simultaneously with vanadium by Hatchett
and Ekeberg, and was afterwards studied by Rose, who in 1844
discovered niobium in it. Notwithstanding the numerous researches
of Hermann (in Moscow), Kobell, Rose, and Marignac, still there is
not yet any certainty as to the purity of, and the properties
ascribed to, the compounds of these elements. They are difficult
to separate from each other, and especially from the cerite metals
and titanium, &c., which accompany them. Before the investigations
of Rose the highest oxide of tantalum was supposed to belong to
the type TaX_{6}--that is, its composition was taken as TaO_{3},
and to the lower oxide was ascribed a formula TaO_{2}. Rose gave
the formula TaO_{2} to the higher oxide, and discovered a new
element called niobium in the substance previously supposed to be
the lower oxide. He even admitted the existence of a third element
occurring together with tantalum and niobium, which he named
pelopium, but he afterwards found that pelopic acid was only
another oxide of niobium, and he considered it probable that the
higher oxide of this element is NbO_{2}, and the lower
Nb_{2}O_{3}. Hermann found that niobic acid which was considered
pure contained a considerable quantity of tantalic acid, and
besides this he admitted the existence of another special metallic
acid, which he called ilmenic acid, after the locality (the Ilmen
mountains of the Urals) of the mineral from which he obtained it.
V. Kobell recognised still another acid, which he called dianic
acid, and these diverse statements were only brought into
agreement in the sixties by Marignac. He first of all indicated an
accurate method for the separation of tantalic and niobic
compounds, which are always obtained in admixture.
[52] If niobic acid be mixed with a small quantity of charcoal and
ignited in a stream of chlorine, a difficultly-fusible and
difficultly-volatile oxychloride, NbOCl_{3} separates. The vapour
density of this compound with respect to air is 7·5, and this
vapour density perfectly confirms the accuracy of the formulæ
given by Marignac, and indicates the quantitative analogy between
the compounds of niobium and tantalum, and those of phosphorus and
arsenic, and consequently also of vanadium. In their qualitative
relations (as is evident also from the correspondence of the
atomic weights), the compounds of tantalum and niobium exhibit a
great analogy with the compounds of molybdenum and tungsten. Thus
zinc, when acting on acid solutions of tantalic and niobic
compounds, gives a blue coloration, exactly as it does with those
of tungsten and molybdenum (also titanium). These acids form the
same large number of salts as those of tungsten and molybdenum.
The anhydrides of the acids are also insoluble in water, but as
colloids are sometimes held in solution, just like those of
titanic and molybdic acids. Furthermore, niobium is in every
respect the nearest analogue of molybdenum, and tantalum of
tungsten. _Niobium_ is obtained by reducing the double fluoride of
niobium and sodium, with sodium. It is difficult to obtain in a
pure state. It is a metal on which hydrochloric acid acts with
some energy, as also does hydrofluoric acid mixed with nitric
acid, and also a boiling solution of caustic potash. _Tantalum_,
which is obtained in exactly the same way, is a much heavier
metal. It is infusible, and is only acted on by a mixture of
hydrofluoric and nitric acids. Rose in 1868 showed that in the
reduction of the double fluoride, NbF_{5},2KF, by sodium, a
greyish powder is obtained after treating with water. The specific
gravity of this powder is 6·8, and he considers it to be niobium
hydride, NbH. Neither did he obtain metallic niobium when he
reduced with magnesium and aluminium, but an alloy, Al_{3}Nb,
having a sp. gr. of 4·5.
Niobium, so far as is known, unites in three proportions with
oxygen. NbO, which is formed when NbOF_{3},2KF is reduced by
sodium; NbO_{2}, which is formed by igniting niobic acid in a
stream of hydrogen, and niobic anhydride, Nb_{2}O_{5}, a white
infusible substance, which is insoluble in acids, and has a
specific gravity of 4·5. Tantalic anhydride closely resembles
niobic anhydride, and has a specific gravity of 7·2. _The
tantalates and niobates_ present the type of ortho-salts--for
example, Na_{2}HNbO_{4},6H_{2}O, and also of pyro-salts, such as
K_{3}HNb_{2}O_{7},6H_{2}O, and of meta-salts--for example,
KNbO_{3},2H_{2}O. And, besides these, they give salts of a more
complex type, containing a larger amount of the elements of the
anhydride; thus, for instance, when niobic anhydride is fused with
caustic potash it forms a salt which is soluble in water, and
crystallises in monoclinic prisms, having the composition
K_{8}Nb_{6}O_{19},16H_{2}O. There is a perfectly similar
isomorphous salt of tantalic acid. Tantalite is a salt of the type
of metatantalic acid, Fe(TaO_{3})_{2}. The composition of
Yttrotantalite appears to correspond with orthotantalic acid.
CHAPTER XX
SULPHUR, SELENIUM, AND TELLURIUM
The acid character of the higher oxides RO_{3} of the elements of group
VI. is still more clearly defined than that of the higher oxides of the
preceding groups, whilst feeble basic properties only appear in the
oxides RO_{3} of the elements of the even series, and then only for those
elements having a high atomic weight--that is, under those two conditions
in which, as a rule, the basic characters increase. Even the lower types
RO_{2} and R_{2}O_{3}, &c., formed by the elements of group VI., are acid
anhydrides in the uneven series, and only those of the elements of the
even series have the properties of peroxides or even of bases.
_Sulphur_ is the typical representative of group VI., both on account of
the fact that the acid properties of the group are clearly defined in it,
and also because it is more widely distributed in nature than any of the
other elements belonging to this group. As an element of the uneven
series of group VI., sulphur gives H_{2}S, sulphuretted hydrogen, SO_{3},
sulphuric anhydride, and SO_{2}, sulphurous anhydride. And in all of them
we find acid properties--SO_{3} and SO_{2} are anhydrides of acids, and
H_{2}S is an acid, although a feeble one. As an element sulphur has all
the properties of a true non-metal; it has not a metallic lustre, does
not conduct electricity, is a bad conductor of heat, is transparent, and
combines directly with metals--in short it has all the properties of the
non-metals, like oxygen and chlorine. Furthermore, sulphur exhibits a
great qualitative and quantitative _resemblance to oxygen_, especially in
the fact that, like oxygen, it combines _with two atoms of hydrogen_, and
forms compounds resembling oxides with metals and non-metals. From this
point of view sulphur is bivalent, if the halogens are univalent.[1] The
chemical character of sulphur is expressed by the fact that it forms a
very slightly stable and feebly energetic acid with hydrogen. The salts
corresponding with this acid are the sulphides, just as the oxides
correspond to water and the chlorides to hydrochloric acid. However, as
we shall afterwards see more fully, the sulphides are more analogous to
the former than to the latter. But although combining with metals, like
oxygen, sulphur also forms chemically stable compounds with oxygen, and
this fact impresses a peculiar character on all the relations of this
element.[2]
[1] The character of sulphur is very clearly defined in the
organo-metallic compounds. Not to dwell on this vast subject, which
belongs to the province of organic chemistry, I think it will be
sufficient for our purpose to compare the physical properties of
the ethyl compounds of mercury, zinc, sulphur and oxygen. The
composition of all of them is expressed by the general formula
(C_{2}H_{5})_{2}R, where R = Hg, Zn, S, or O. They are all
volatile: mercury ethyl, Hg(C_{2}H_{5})_{2}, boils at 159°, its sp.
gr. is 2·444, molecular volume = 106; zinc ethyl boils at 118°, sp.
gr. 1·882, volume 101; ethyl sulphide, S(C_{2}H_{5})_{2}, boils at
90°, sp. gr. 0·825, volume 107; common ether, or ethyl oxide,
O(C_{2}H_{5})_{2}, boils at 35°, sp. gr. 0·736, volume 101, in
addition to which diethyl itself, (C_{2}H_{5})_{2} = C_{4}H_{10},
boils about 0°, sp. gr. about 0·62, volume about 94. Thus the
substitution of Hg, S, and O scarcely changes the volume,
notwithstanding the difference of the weights; the physical
influence, if one may so express oneself, of these elements, which
are so very different in their atomic weights, is almost alike.
[2] Therefore in former times sulphur was known as an amphid element.
Although the analogy between the compounds of sulphur and oxygen
has been recognised from the very birth of modern chemistry (owing,
amongst other things, to the fact that the oxides and sulphides are
the most widely spread metallic ores in nature), still it has only
been clearly expressed by the periodic system, which places both
these elements in group VI. Here, moreover, stands out that
parallelism which exists between SO_{2} and ozone OO_{2}, between
K_{2}SO_{3} and peroxide of potassium K_{2}O_{4} (Volkovitch in
1893 again drew attention to this parallelism).
Sulphur belongs to the number of those elements which _are very widely
distributed in nature_, and occurs both free and combined in various
forms. The atmosphere, however, is almost entirely free from compounds of
sulphur, although a certain amount of them should be present, if only
from the fact that sulphurous anhydride is emitted from the earth in
volcanic eruptions, and in the air of cities, where much coal is burnt,
since this always contains FeS_{2}. Sea and river water generally contain
more or less sulphur in the form of sulphates. The beds of gypsum, sodium
sulphate, magnesium sulphate, and the like are formations of undoubtedly
aqueous origin. The sulphates contained in the soil are the source of the
sulphur found in plants, and are indispensable to their growth. Among
vegetable substances, the proteïds always contain from one to two per
cent. of sulphur. From plants the albuminous substances, together with
their sulphur, pass into the animal organism, and therefore the
decomposition of animal matter is accompanied by the odour of
sulphuretted hydrogen, as the product into which the sulphur passes in
the decomposition of the albuminous substances. Thus a rotten egg emits
sulphuretted hydrogen. Sulphur occurs largely in nature, as the various
insoluble sulphides of the metals. Iron, copper, zinc, lead, antimony,
arsenic, &c., occur in nature combined with sulphur. These _sulphides_
frequently have a metallic lustre, and in the majority of cases occur
crystallised, and also very often several sulphides occur combined or
mixed together in these crystalline compounds. If they are yellow and
have a metallic lustre they are called pyrites. Such are, for example,
copper pyrites, CuFeS_{2}, and iron pyrites, FeS_{2}, which is the
commonest of all. They are all also known as glances or blendes if they
are greyish and have a metallic lustre--for example, zinc blende, lead
glance, PbS, antimony glance, Sb_{2}S_{3}, &c. And, lastly, sulphur
occurs _native_. It occurs in this form in the most recent geological
formations in admixture with limestone and gypsum, and most frequently in
the vicinity of active or extinct volcanoes. As the gases of volcanoes
contain sulphur compounds--namely, sulphuretted hydrogen and sulphurous
anhydride, which by reacting on one another may produce sulphur, which
also frequently appears in the craters of volcanoes as a sublimate--it
might be imagined that the sulphur was of volcanic origin. But on a
nearer acquaintance with its mode of occurrence, and more especially
considering its relation to gypsum, CaSO_{4}, and limestone, the present
general opinion leads to the conclusion that the 'native' sulphur has
been formed by the reduction of the gypsum by organic matter and that its
occurrence is only indirectly connected with volcanic agencies. Near
Tetush, on the Volga, there are beds containing gypsum, sulphur, and
asphalt (mineral tar). In Europe the most important deposits of sulphur
are in the south of Sicily from Catania to Girgenti.[3] There are very
rich deposits of sulphur in Daghestan near Cherkai and Cherkat in Khyut,
near Mount Kanabour-bam, near Petrovsk, and in the Kira Koumski steppes
in the Trans-Caspian provinces, which are able to supply the whole of
Russia with this mineral. Abundant deposits of sulphur have also been
found in Kamtchatka in the neighbourhood of the volcanoes. The method of
separation of the sulphur from its earthy impurities is based on the fact
that sulphur melts when it is heated. The fusion is carried on at the
expense of a portion of the sulphur, which is burnt, so that the
remainder may melt and run from the mass of the earth. This is carried on
in special furnaces called calcaroni, built up of unhewn stone in the
neighbourhood of the mines.[4]
[3] When in Sicily, I found, near Caltanisetta, a specimen of sulphur
with mineral tar. In the same neighbourhood there are naphtha
springs and mud volcanoes. It may be that these substances have
reduced the sulphur from gypsum.
The chief proof in favour of the origin of sulphur from gypsum is
that in treating the deposits for the extraction of the sulphur it
is found that the proportion of sulphur to calcium carbonate never
exceeds that which it would be had they both been derived from
calcium sulphate.
[4] Naturally only those ores of sulphur which contain a considerable
amount of sulphur can be treated by this method. With poor ores it
is necessary to have recourse to distillation or mechanical
treatment in order to separate the sulphur, but its price is so low
that this method in most cases is not profitable.
The sulphur obtained by the above-described method still contains
some impurities, but it is frequently made use of in this form for
many purposes, and especially in considerable quantities for the
manufacture of sulphuric acid, and for strewing over grapes. For
other purposes, and especially in the preparation of gunpowder, a
purer sulphur is required. Sulphur may be purified by distillation.
The crude sulphur is called _rough_, and the distilled sulphur
_refined_. The arrangement given in fig. 86 is employed for
refining sulphur. The rough sulphur is melted in the boiler _d_,
and as it melts it is run through the tube F into an iron retort B
heated by the naked flame of the furnace. Here the sulphur is
converted into vapour, which passes through a wide tube into the
chamber G, surrounded by stone walls and furnished with a
safety-valve S.
[Illustration: FIG. 86.--Refining sulphur by sublimation.]
Sulphur is purified by distillation in special retorts (see fig. 86) by
passing the vapour into a chamber G built of stone. The first portions of
the vapour entering into the condensing chamber are condensed straightway
from the vapour into a solid state, and form a fine powder known as
_flowers of sulphur_.[5] But when the temperature of the receiver attains
the melting point of sulphur, it passes into a liquid state and is cast
into moulds (like sealing wax), and is then known under the name of _roll
sulphur_.[6]
[5] Flowers of sulphur always contain a certain amount of the oxides of
sulphur.
[6] Sulphur may be extracted by various other means. It may be
extracted from iron pyrites, FeS_{2}, which is very widely
distributed in nature. From 100 parts of iron pyrites about half
the sulphur contained, namely, about 25 parts, may be extracted by
heating without the access of air, a lower sulphide of iron, which
is more stable under the action of heat, being left behind. Alkali
waste (Chapter XII.), containing calcium sulphide and gypsum,
CaSO_{4}, may be used for the same purpose, but native sulphur is
so cheap that recourse can only be had to these sources when the
calcium sulphide appears as a worthless by-product. The most simple
process for the extraction of sulphur from alkali waste, in a
chemical sense, consists in evolving sulphuretted hydrogen from the
calcium sulphide by the action of hydrochloric acid. The
sulphuretted hydrogen when burnt gives water and sulphurous
anhydride, which reacts on fresh sulphuretted hydrogen with the
separation of sulphur. The combustion of the sulphuretted hydrogen
may be so conducted that a mixture of 2H_{2}S and SO_{2} is
straightway formed, and this mixture will deposit sulphur (Chapter
XII., Note 14). Gossage and Chance treat alkali waste with carbonic
anhydride, and subject the sulphuretted hydrogen evolved to
incomplete combustion (this is best done by passing a mixture of
sulphuretted hydrogen and air, taken in the requisite proportions,
over red-hot ferric oxide), by which means water and the vapour of
sulphur are formed: H_{2}S + O = H_{2}O + S.
In an uncombined state sulphur exists in _several modifications_, and
forms a good example of the facility with which an alteration of
properties can take place without a change of composition--that is, as
regards the material of a substance. Common sulphur has the well-known
yellow colour. This colour fades as the temperature falls, and at -50°
sulphur is almost colourless. It is very brittle, so that it may be
easily converted into a powder, and it presents a crystalline structure,
which, by the way, shows itself in the unequal expansion of lumps of
sulphur by heat. Hence when a piece of sulphur is heated by the warmth of
the hand, it emits sounds and sometimes cracks, which probably also
depends on the bad heat-conducting power of this substance. It is easily
obtained in a crystalline form by artificial means, because although
insoluble in water it dissolves in carbon bisulphide, and in certain
oils.[7] Solutions of sulphur in carbon bisulphide when evaporated at the
ordinary temperature yield well-formed transparent crystals of sulphur in
the form of rhombic octahedra, in which form it occurs native. The
specific gravity of these crystals is 2·045. Fused sulphur, cast into
moulds and cooled, has, after being kept a long time, a specific gravity
2·066; almost the same as that of the crystalline sulphur of the above
form, which shows that common sulphur is the same as that which
crystallises in octahedra. The specific heat of octahedral sulphur is
0·17; it melts at 114°, and forms a bright yellow mobile liquid. On
further heating, the fused sulphur undergoes an alteration, which we
shall presently describe, first observing that the above octahedral state
of sulphur is its most stable form. Sulphur may be kept at the ordinary
temperature in this form for an indefinite length of time, and many other
modifications of sulphur pass into this form after being left for a
certain time at ordinary temperature.
[7] One hundred parts of liquid carbon bisulphide, CS_{2}, dissolve
16·5 parts of sulphur at -11°, 24 parts at 0°, 37 parts at 15°, 46
parts at 22°, and 181 parts at 55°. The saturated solution boils at
55°, whilst pure carbon bisulphide boils at 47°. The solution of
sulphur in carbon bisulphide reduces the temperature, just as in
the solution of salts in water. Thus the solution of 20 parts of
sulphur in 50 parts of carbon bisulphide at 22° lowers the
temperature by 5°; 100 parts of benzene, C_{6}H_{6}, dissolves
0·965 part of sulphur at 26°, and 4·377 parts at 71°; chloroform,
CHCl_{3}, dissolves 1·2 part of sulphur at 22°, and 16·35 parts at
174°.
If sulphur be melted and then slightly cooled, so that it forms a crust
on the surface and over the sides of the crucible, while the internal
mass remains liquid, then the sulphur takes another crystalline form as
it solidifies. This may be seen by breaking the crust, and pouring out
the remaining molten sulphur.[8] It is then found that the sides of the
crucible are covered with _prismatic crystals_ of the monoclinic system;
they have a totally different appearance from the above-described
crystals of rhombic sulphur. The prismatic crystals are brown,
transparent, and less dense than the crystals of rhombic sulphur, their
specific gravity being only 1·93, and their melting point higher--about
120°. These crystals of sulphur cannot be kept at the ordinary
temperature, which is indeed evident from the fact that in time they turn
yellow; the specific gravity also changes, and they pass completely into
the ordinary modification. This is accompanied by a considerable
development of heat, so that the temperature of the mass may rise 12°.
Thus sulphur is _dimorphous_--that is, it exists in two crystalline
forms, and in both forms it has independent physical properties. However,
no chemical reactions are known which distinguish the two modifications
of sulphur, just as there are none distinguishing aragonite from
calcspar.[9]
[8] If the experiment be made in a vessel with a narrow capillary tube,
the sulphur fuses at a lower temperature (occurs, as it were, in a
supersaturated state), and solidifying at 90°, appears in a rhombic
form (Schützenberger).
[9] If sulphur be cautiously melted in a U tube immersed in a salt
bath, and then gradually cooled, it is possible for all the sulphur
to remain liquid at 100°. It will now be in a state of superfusion;
thus also by careful refrigeration water may be obtained in a
liquid state at -10°, and a lump of ice then causes such water to
form ice, and the temperature rises to 0°. If a prismatic crystal
of sulphur be thrown into one branch of the U tube containing the
liquid sulphur at 100°, and an octahedral crystal be thrown into
the other branch, then, as Gernez showed, the sulphur in each
branch will crystallise in the corresponding form, and both forms
are obtained at the same temperature; therefore it is not the
influence of temperature only which causes the molecules of sulphur
to distribute themselves in one or another form, but also the
influence of the crystalline parts already formed. This phenomenon
is essentially analogous to the phenomena of supersaturated
solutions.
If molten sulphur be heated to 158° it loses its mobility and becomes
thick and very dark-coloured, so that the crucible in which it is heated
may be inverted without the sulphur running out. When heated above this
temperature the sulphur again becomes liquid, and at 250° it is very
mobile, although it does not acquire its original colour, and at 440° it
boils. These modifications in the properties of sulphur depend not only
on the variations of temperature, but also on a change of structure. If
sulphur, heated to about 350°, be poured in a thin stream into cold
water, it does not solidify into a solid mass, but retains its brown
colour and _remains soft_, may be stretched out into threads, and is
elastic, like guttapercha. But in this soft and ductile state, also, it
does not remain for a long time. After the lapse of a certain period this
soft transparent sulphur hardens, becomes opaque, passes into the
ordinary yellow modification of sulphur, and in so doing develops heat,
just as in the conversion of the prismatic into the octahedral variety.
The soft sulphur is characterised by the fact that a certain portion of
it is insoluble in carbon bisulphide. When soft sulphur is immersed in
this liquid, only a portion of common sulphur passes into solution,
whilst a certain portion is quite insoluble and remains so for a long
time. The maximum proportion of insoluble sulphur is obtained by heating
slightly above 170°. It melts at 114°. An exactly similar _insoluble
amorphous sulphur_ is obtained in certain reactions in the wet way, when
sulphur separates out from solutions. Thus sodium thiosulphate,
Na_{2}S_{2}O_{3}, when treated with acids, gives a precipitate of
sulphur, which is insoluble in carbon bisulphide. The action of water on
sulphur chloride also gives a similar modification of sulphur. Certain
sulphides, when treated with nitric acid, also yield sulphur in this
form.[10]
[10] A certain amount of insoluble sulphur remains for a long time in
the mass of soft sulphur, changing into the ordinary variety.
Freshly-cooled soft sulphur contains about one-third of insoluble
sulphur, and after the lapse of two years it still contains about
15 p.c. Flowers of sulphur, obtained by the rapid condensation of
sulphur from a state of vapour, also contains a certain amount of
insoluble sulphur. _Rapidly distilled and condensed sulphur_ also
contains some insoluble sulphur. Hence a certain amount of
insoluble sulphur is frequently found in roll sulphur. The action
of light on a solution of sulphur converts a certain portion into
the insoluble modification. Insoluble sulphur is of a lighter
colour than the ordinary variety. It is best prepared by
vaporising sulphur in a stream of carbonic anhydride, hydrochloric
acid, &c., and collecting the vapour in cold water. When condensed
in this manner it is nearly all insoluble in carbon bisulphide. It
then has the form of hollow spheroids, and is therefore lighter
than the common variety: sp. gr. 1·82. An idea of the
modifications taking place in sulphur between 110° and 250° may be
formed from the fact that at 150° liquid sulphur has a coefficient
of expansion of about 0·0005, whilst between 150° and 250° it is
less than 0·0003.
Engel (1891), by decomposing a saturated solution of hyposulphite
of sodium (Note 42) with HCl in the cold (the sulphur is not
precipitated directly in this case), obtained, after shaking up
with chloroform and evaporation, crystals of sulphur (sp. gr.
2·135), which, after several hours, passed into the insoluble (in
CS_{2}) state, and in so doing became opaque, and increased in
volume. But if a mixture of solution of Na_{2}S_{2}O_{3} and HCl
be allowed to stand, it deposits sulphur, which, after sufficient
washing, is able to dissolve in water (like the colloid varieties
of the metallic sulphides, alumina, boron, and silver), but this
colloid _solution of sulphur_ soon deposits sulphur insoluble in
CS_{2}.
When a solution of sulphuretted hydrogen in water is decomposed by
an electric current the sulphur is deposited on the positive pole,
and has therefore an electro-negative character, and this sulphur
is soluble in carbon bisulphide. When a solution of sulphurous
acid is decomposed in the same manner, the sulphur is deposited on
the negative pole, and is therefore electro-positive, and the
sulphur so deposited is insoluble in carbon bisulphide. The
sulphur which is combined with metals must have the properties of
the sulphur contained in sulphuretted hydrogen, whilst the sulphur
combined with chlorine is like that which is combined with oxygen
in sulphurous anhydride. Hence Berthelot recognises the presence
of soluble sulphur in metallic sulphides, and of the insoluble
modification of amorphous sulphur in sulphur chloride. Cloez
showed that the sulphur precipitated from solutions is either
soluble or insoluble, according to whether it separates from an
alkaline or acid solution. If sulphur be melted with a small
quantity of iodine or bromine, then on pouring out the molten mass
it forms amorphous sulphur, which keeps so for a very long time,
and is insoluble, or nearly so, in carbon bisulphide. This is
taken advantage of in casting certain articles in sulphur, which
by this means retain their tenacity for a long time; for example,
the discs of electrical machines.
At temperatures of 440° to 700° the vapour density of sulphur is 6·6
referred to air--_i.e._ about 96 referred to hydrogen. Hence, at these
temperatures _the molecule of sulphur contains six atoms_, it has the
composition S_{6}. The agreement between the observations of Dumas,
Mitscherlich, Bineau, and Deville confirms the accuracy of this result.
But in this respect the properties of sulphur were found to be variable.
When heated to higher temperatures, that is to say, _above_ 800°, the
vapour density of sulphur is found to be one-third of this quantity,
_i.e._ about 32 referred to hydrogen. At this temperature _the molecule
of sulphur_, like that of hydrogen, oxygen, nitrogen, and chlorine,
_contains two atoms_; hence the molecular formula is then S_{2}. This
variation in the vapour density of sulphur evidently corresponds with a
polymeric modification, and may be likened to the transformation of
ozone, O_{3}, into oxygen, O_{2}, or better still, of benzene,
C_{6}H_{6}, into acetylene, C_{2}H_{2}.[11]
[11] Here, however, it is very important to remark that both benzene
and acetylene can exist at the ordinary temperature, whilst the
sulphur molecule S_{2} only exists at high temperatures; and if
this sulphur be allowed to cool, it passes first into S_{6} and
then into a liquid state. Were it possible to have sulphur at the
ordinary temperature in both the above modifications, then in all
probability the sulphur in the state S_{2} would present totally
different properties from those which it has in the form S_{6},
just as the properties of gaseous acetylene are far from being
similar to those of liquid benzene. Sulphur, in the form of S_{2},
is probably a substance which boils at a much lower temperature
than the variety with which we are now dealing. Paterno and Nasini
(1888), following the method of depression or fall of the
freezing-point in a benzene solution, found that the molecule of
sulphur in solution contains S_{6}.
One must here call attention to the fact that sulphur, with all
its analogy to oxygen (which also shows itself in its faculty to
give the modification S_{2}), is also able to give a series of
compounds containing more atoms of sulphur than the analogous
oxygen compounds do of oxygen. Thus, for instance, compounds of 5
atoms of sulphur with 1 atom of barium, BaS_{5}, are known,
whereas with oxygen only BaO_{2} is known. On every side one
cannot but see in sulphur a faculty for the union of a greater
number of atoms than with oxygen. With oxygen the form of ozone,
O_{3}, is very unstable, the stable form is O_{2}; whilst with
sulphur S_{6} is the stable form, and S_{2} is exceedingly
unstable. Furthermore, it is remarkable that sulphur gives a
higher degree of oxidation, H_{2}SO_{4}, corresponding, as it
were, with its complex composition, if we suppose that in S_{6}
four atoms of sulphur are replaced by oxygen and one by two atoms
of hydrogen. The formulæ of its compounds, K_{2}SO_{4},
K_{2}S_{2}O_{3}, K_{2}S_{5}, BaS_{5}, and many others, have no
analogues among the compounds of oxygen. They all correspond with
the form S_{6} (one portion of the sulphur being replaced by
oxygen and another by metals), which is not attained by oxygen. In
this faculty of sulphur to hold many atoms of other substances the
same forces appear which cause many atoms of sulphur to form one
complex molecule.
_In its faculty for combination_, sulphur most closely resembles oxygen
and chlorine; like them, it combines with nearly all elements, with the
development of heat and light, forming sulphur compounds, but as a rule
this only takes place at a high temperature. At the ordinary temperature
it does not enter into reactions, owing, amongst other things, to the
fact that it is a solid. In a molten state it acts on most metals and on
the halogens. It burns in air at about 300°, and with carbon at a red
heat, but it does not combine with nitrogen.
Fine wires, or the powders of the greater number of metals, burn in the
vapour of sulphur. The direct combination of hydrogen with sulphur is
restricted by a limit--that is, at a given temperature and under other
given conditions it does not proceed unrestrictedly; there is no
explosion or recalescence. Sulphuretted hydrogen, H_{2}S, decomposes at
its temperature of combination--that is, it is easily dissociated.[12]
The same phenomenon is repeated here as with water, except that the
temperatures at which the attraction of hydrogen for sulphur begins and
ceases are much lower than in the case of oxygen and hydrogen. The
temperature at which combination takes place is here, as in many other
instances, nearly the same as that at which dissociation begins. Hence
_sulphuretted hydrogen_ is formed in a small quantity by the direct
ignition of a mixture of the vapour of sulphur and hydrogen. However, the
temperature must not be high, because otherwise the whole of the
sulphuretted hydrogen is decomposed; but at lower temperatures a small
amount of sulphuretted hydrogen is formed by direct combination.[13]
Sulphuretted hydrogen however, like all other hydrogen compounds, may be
easily obtained by the double decomposition of its corresponding metallic
compounds, the replacement of the metal by hydrogen being effected by the
action of acids on the sulphides. The metallic sulphides are, as a rule,
easily formed. A sulphide, when mixed with a non-volatile acid, may give,
by double decomposition, a salt of the acid taken and sulphuretted
hydrogen, M_{2}S + H_{2}SO_{4} = H_{2}S + M_{2}SO_{4}. However, it is not
all sulphides nor solutions of all acids that will evolve sulphuretted
hydrogen, which fact is exceedingly characteristic, because, for example,
all carbonates evolve carbonic anhydride when treated with any acid.
Sulphuric acid will only evolve sulphuretted hydrogen from those
sulphides which contain a metal capable of decomposing the acid with the
evolution of hydrogen. Thus zinc, iron, calcium, magnesium, manganese,
potassium, sodium, &c., form sulphides which evolve sulphuretted hydrogen
when treated with sulphuric acid, and the metals themselves evolve
hydrogen with acids.[14] The sulphides of those metals which do not
liberate hydrogen from acids do not generally act on acids--that is, do
not form sulphuretted hydrogen with them; such are, for example, the
sulphides of lead, silver, copper, mercury, tin, &c. Therefore, the
_modus operandi_ of the formation of sulphuretted hydrogen by the action
of acids on metallic sulphides may be looked on as a phenomenon of the
combination of hydrogen, at the moment of its evolution, with the
sulphur, which is combined with the metal. Such a representation is all
the more simple as all the circumstances under which sulphuretted
hydrogen is formed are exactly similar to the conditions of the formation
of hydrogen itself. Thus the usual mode of preparing sulphuretted
hydrogen is by the action of _sulphuric acid on ferrous sulphide_, in
which the same apparatus and method are employed as in the preparation of
hydrogen, only replacing the metallic iron or zinc by ferrous sulphide or
zinc sulphide. The reaction between sulphide of iron and sulphuric acid
takes place at the ordinary temperature, and is accompanied by just as
small a development of heat as in the liberation of hydrogen itself, FeS
+ H_{2}SO_{4} = FeSO_{4} + H_{2}S.[15]
[12] In the formation of potassium sulphide, K_{2}S (that is, in the
combination of 32 parts of sulphur with 78 parts of potassium),
about 100 thousand heat units are developed. Nearly as much heat
is developed in the combination of an equivalent quantity of
sodium; about 90 thousand heat units in the formation of calcium
or strontium sulphide; about 40 thousand for zinc or cadmium
sulphide, and about 20 thousand for iron, cobalt, or nickel
sulphide. Less heat is evolved in the combination of sulphur with
copper, lead, and silver. According to Thomsen, sulphur develops
heat with hydrogen in solutions. The reaction I_{2},Aq,H_{2}S =
21,830 calories. But, as the reaction I_{2} + H_{2} + Aq develops
26,842 calories, it follows that the reaction H_{2} + S develops
4,512 calories.
[13] If sulphur be melted in a flask and heated nearly to its boiling
point, as Lidoff showed, the addition, drop by drop (from a funnel
with a stopcock) of heavy (0·9) naphtha oil (of lubricating
oleonaphtha), &c., is followed by a regular evolution of
sulphuretted hydrogen. This is analogous to the action of bromine
or iodine on paraffin and other oils, because hydrobromic or
hydriodic acid is then formed (Chapter XI.) A certain amount of
hydrogen sulphide is even formed when sulphur is boiled with
water.
[14] However, the matter is really much more complicated. Thus zinc
sulphide evolves sulphuretted hydrogen with sulphuric or
hydrochloric acids, but does not react with acetic acid and is
oxidised by nitric acid. Ferrous sulphide evolves sulphuretted
hydrogen with acids, whilst the bisulphide, FeS_{2}, does not
react with acids of ordinary strength. This absence of action
depends, among other things, on the form in which the native iron
pyrites occurs; it is a crystalline, compact, and very dense
substance; and acids in general react with great difficulty on
such metallic sulphides. This is seen very clearly in the case of
zinc sulphide; if this substance is obtained by double
decomposition, it separates as a white precipitate, which evolves
sulphuretted hydrogen with great ease when treated with acids.
Zinc sulphide is obtained in the same form when zinc is fused with
sulphur, but native zinc sulphide--which occurs in compact masses
of zinc blende, and has a metallic lustre--is not decomposed or
scarcely decomposed by sulphuric acid.
Another source of complication in the behaviour of the metallic
sulphides towards acids depends on the action of water, and is
shown in the fact that the action varies with different degrees of
dilution or proportion of water present. The best known example of
this is antimonious sulphide, Sb_{2}S_{3}, for strong hydrochloric
acid, containing not more water than corresponds with HCl,6H_{2}O,
even decomposes native antimony glance, with evolution of
sulphuretted hydrogen, whilst dilute acid has no action, and in
the presence of an excess of water the reaction 2SbCl_{3} +
3H_{2}S = Sb_{2}S_{3} + 6HCl occurs, whilst in the presence of a
small amount of water the reaction proceeds in exactly the
opposite direction. Here the participation of water in the
reaction and its affinity are evident.
The facts that lead sulphide is insoluble in acids, that zinc
sulphide is soluble in hydrochloric acid but insoluble in acetic
acid, that calcium sulphide is even decomposed by carbonic acid,
&c.--all these peculiarities of the sulphides are in correlation
with the amount of heat evolved in the reaction of the oxides with
hydrogen sulphide and with acids, as is seen from the observations
of Favre and Silberman, and from the comparisons made by Berthelot
in the Proceedings of the Paris Academy of Sciences, 1870, to
which we refer the reader for further details.
[15] _Ferrous sulphide_ is formed by heating a piece of iron to an
incipient white heat, and then removing it from the furnace and
bringing it into contact with a piece of sulphur. Combination then
proceeds, accompanied by the development of heat, and the ferrous
sulphide formed fuses. The sulphide of iron thus formed is a
black, easily-fusible substance, insoluble in water. When damp it
attracts oxygen from the air, and is converted into green vitriol,
FeSO_{4}. If all the iron does not combine with the sulphur in the
method described above, the action of sulphuric acid will evolve
hydrogen as well as hydrogen sulphide.
We will not describe the details of the preparation of
sulphuretted hydrogen employed as a reagent in the laboratory,
because, in the first place, the methods are essentially the same
as in the preparation of hydrogen, and, in the second place,
because the apparatus and methods employed are always described in
text-books of analytical chemistry. Ferrous sulphide may be
advantageously replaced by calcium sulphide or a mixture of
calcium and magnesium sulphides. A solution of magnesium
hydrosulphide, MgS,H_{2}S, is very convenient, as at 60° it
evolves a stream of pure hydrogen sulphide. A paste, consisting of
CuS with crystals of MgCl_{2} and water, may also be employed,
since it only evolves H_{2}S when heated (Habermann).
_In nature_ sulphuretted hydrogen is formed in many ways. The most usual
mode of its formation is by the decomposition of albuminous substances
containing sulphur, as mentioned above. Another method is by the reducing
action of organic matter on sulphates, and by the action of water and
carbonic acid on the sulphides formed by this reduction. Volcanic
eruptions are a third source of sulphuretted hydrogen in nature. Although
sulphuretted hydrogen is formed in small quantities everywhere, it
nevertheless soon disappears from the atmosphere, owing to its being
easily decomposed by oxidising agencies. Many mineral waters contain
sulphuretted hydrogen, and smell of it; they are called 'sulphur waters.'
Sulphuretted hydrogen, at the ordinary temperature, is a colourless gas,
having a very unpleasant odour. It has, as its composition H_{2}S shows,
a specific gravity seventeen times greater than hydrogen, and therefore
it is somewhat heavier than air. Sulphuretted hydrogen _liquefies_ at
about -74°, and at the ordinary temperature when subjected to a pressure
of 10 to 15 atmospheres; at -85° it is converted into a solid crystalline
mass.[15 bis] The easy liquefaction of sulphuretted hydrogen is evidently
allied to its solubility. One volume of water at 0° dissolves 4·37
volumes of sulphuretted hydrogen, at 10° 3·58 volumes, and at 20° 2·9
volumes.[16] The solutions impart a very feeble red coloration to litmus
paper. This gas is poisonous. One part in fifteen hundred parts of air
will kill birds. Mammalia die in an atmosphere containing 1/200 of this
gas.
[15 bis] Liquid sulphuretted hydrogen is most easily obtained by the
decomposition of hydrogen polysulphide, which we shall presently
describe, by the action of heat, and in the presence of a small
amount of water. If poured into a bent tube, like that described
for the liquefaction of ammonia (Chapter VI.), the hydrogen
polysulphide is decomposed by heat, in the presence of water, into
sulphur and sulphuretted hydrogen, which condenses in the cold end
of the tube into a colourless liquid.
[16] Sulphuretted hydrogen is still more soluble in alcohol than in
water; one volume at the ordinary temperature dissolves as much as
eight volumes of the gas. The solutions in water and alcohol
undergo change, especially in open vessels, owing to the fact that
the water and alcohol dissolve oxygen from the atmosphere, which,
acting on the sulphuretted hydrogen, forms water and sulphur. The
solution may be so altered in this manner that every trace of
sulphuretted hydrogen disappears. Solutions of sulphuretted
hydrogen in glycerine change much more slowly, and may therefore
be kept for a long time as reagents. De Forcrand obtained a
hydrate, H_{2}S,16H_{2}O, resembling the hydrates given by many
gases.
Sulphuretted hydrogen is very easily _decomposed_ into its component
parts by the action of heat or a series of electric sparks. Hence it is
not surprising that sulphuretted hydrogen undergoes change under the
action of many substances having a considerable affinity for hydrogen and
oxygen. Very many metals[17] evolve hydrogen with sulphuretted hydrogen,
so that in this respect it presents the property of an acid; for
instance, 2H_{2}S + Sn = 2H_{2} + SnS_{2}. This may be taken advantage of
for determining the composition of sulphuretted hydrogen, because a given
volume then leaves the same volume of hydrogen. On the other hand,
oxygen,[18] chlorine,[19] and even iodine decompose sulphuretted
hydrogen, removing the hydrogen from it and leaving free sulphur, so that
in this reaction the sulphur is replaced by the above-named elements; for
example, H_{2}S + Br_{2} = 2HBr + S. In no other hydrogen compound is it
so easy to show the _substitution_, both of hydrogen and of the element
combined with it, as in hydrogen sulphide. This clearly proves the feeble
union between the elements forming this gas. Compounds containing a
considerable amount of oxygen, with which they easily part, can
accomplish the separation of the sulphur very easily. Such are, for
instance, nitrous acid, chromic acid, and even ferric oxide and the
higher oxides like it. Thus, if sulphuretted hydrogen be passed into a
solution of chromic acid or an acid solution of ferric oxide, water is
formed, _and the sulphur is separated in a free state_. Thus,
sulphuretted hydrogen acts as a _reducing agent_, in virtue of the
hydrogen it contains. Salts of iodic, chlorous, chloric, and other acids
are reduced by sulphuretted hydrogen, their oxygen acting mainly on its
hydrogen; but in the presence of an excess of a powerful oxidising agent
a portion of the sulphur may also be oxidised to sulphurous anhydride.
The reducing action of sulphuretted hydrogen is frequently applied in
chemical manipulations for the preparation of lower oxides, and for the
conversion of certain oxygen compounds into hydrogen compounds: thus, the
higher oxides of nitrogen are converted into ammonia by it, and in the
presence of alkalis the nitro-compounds are converted into ammonia
derivatives. The reaction of sulphuretted hydrogen on sulphurous
anhydride belongs to this class of phenomena, the chief products of which
are sulphur and water, 2H_{2}S + SO_{2} = 2H_{2}O + S_{3}.
[17] Some metals evolve hydrogen from sulphuretted hydrogen at the
ordinary temperature. For example, the light metals, and copper
and silver (especially with the access of air?) among the heavy
metals. Hence articles made of silver turn black in the presence
of vapours containing sulphuretted hydrogen, because silver
sulphide is black. Zinc and cadmium act at a red heat, but not
completely.
[18] If sulphuretted hydrogen escapes from a fine orifice into the air,
it will burn when lighted, and be transformed into sulphurous
anhydride and water. But if it burns in a limited supply of
air--for instance, when a cylinder is filled with it and
lighted--then only the hydrogen burns, which has, judging from the
amount of heat developed in its combustion and from all its
properties, a greater affinity for oxygen than sulphur. In this
respect the combustion of sulphuretted hydrogen resembles that of
hydrocarbons.
[19] Hence bleaching powder and chlorine destroy the disagreeable smell
of sulphuretted hydrogen. (For the reaction of hydrogen sulphide
and iodine, _see_ Chapter XI. p. 504.)
The acid character of sulphuretted hydrogen is clearly seen in its
action on alkalis and salts.[19 bis] Thus lead oxide and its salts in the
presence of sulphuretted hydrogen form water or an acid, and sulphide of
lead: PbX_{2} + H_{2}S = PbS + 2HX. This reaction takes place even in the
presence of powerful acids, because lead sulphide is one of those
sulphides which are unacted on by acids, and in solutions the reaction is
a complete one. This reaction is taken advantage of for the preparation
of many acids, by first converting into a lead salt, and then submitting
this salt to the action of sulphuretted hydrogen. For example, lead
formate with sulphuretted hydrogen gives formic acid. Sulphuretted
hydrogen in acting on a number of metallic acid substances in solution or
in an anhydrous state also forms corresponding sulphates: (1) if it does
not reduce the acid; (2) if the sulphur compound corresponding with the
anhydride of the acid be insoluble in water, the reaction proceeds in
solutions; (3) if the sulphuretted hydrogen and the acid taken do not
come in contact with an alkali, on which they would be able to act first;
and (4) if the sulphur compound be not decomposed by water. Thus
solutions of arsenious acid give a precipitate of arsenious sulphide,
As_{2}S_{3}, with sulphuretted hydrogen. This reaction proceeds not only
in the presence of water, but also of acids, because the latter do not
decompose the resultant sulphur compounds. The type of the decomposition
is the same as with bases--that is, the sulphur and oxygen change places:
RO_{_n_} + _n_H_{2}S = RS_{_n_} + _n_H_{2}O. Some sulphides corresponding
with acid anhydrides are decomposed by water, and therefore are not
formed in the presence of water. Such, for example, are the sulphides of
phosphorus.[20]
[19 bis] Perfectly dry H_{2}S (Hughes 1892) has no action upon
perfectly dry salts, just as dry HCl does not react with dry
NH_{3} or metals (Chapter IX., Note 29).
[20] The sulphide P_{4}S is obtained by cautiously fusing the requisite
proportions of common phosphorus and sulphur under water; it is a
liquid which solidifies at 0°, and may be distilled without
undergoing change, but it fumes in air and easily takes fire. The
higher sulphide, P_{2}S, has similar properties. But little heat
is evolved in the formation of these compounds, and it may be
supposed that they are formed by the direct conjunction of whole
molecules of phosphorus and sulphur; but if the proportion of
sulphur be increased, the reaction is accompanied by so
considerable a rise of temperature that an explosion takes place,
and for the sake of safety red phosphorus must be used, mixed as
intimately as possible with powdered sulphur and heated in an
atmosphere of carbonic anhydride. The higher compounds are
decomposed by water. By increasing the proportion of sulphur, the
following compounds have been obtained: P_{4}S_{3} as prisms
(fuses at 165°, Rebs), soluble in carbon bisulphide, and unaltered
by air and water; _phosphorus trisulphide_, P_{2}S_{3}, is the
analogue of P_{2}O_{3}; it is a light yellow crystalline compound
only slightly soluble in carbon bisulphide, fusible and volatile,
decomposed into hydrogen sulphide and phosphorous acid by water,
and, like the highest compound of sulphur and phosphorus,
P_{2}S_{5}, it forms thio-salts with potassium sulphide, &c. This
_phosphorus pentasulphide_ corresponds with phosphoric anhydride;
like the trisulphide it gives hydrogen sulphide and phosphoric
acid with an excess of water. It reacts in many respects like
phosphoric chloride. The sulphide PS_{2} is also known; the vapour
density of this compound seems to indicate a molecule P_{3}S_{6}.
_Phosphorus sulphochloride_, PSCl_{3}, corresponds with phosphorus
oxychloride. It is a colourless, pleasant-smelling liquid, boiling
at 124°, and of sp. gr. 1·63; it fumes in air and is decomposed by
water: PSCl_{3} + 4H_{2}O = PH_{3}O_{4} + H_{2}S + 3HCl. It is
obtained when phosphoric chloride is treated with hydrogen
sulphide, hydrochloric acid being also formed; it is also produced
by the action of phosphoric chloride on certain sulphides--for
example, on antimonious sulphide, also by the (cautious) action of
phosphorus on sulphur chloride: 2P + 3S_{2}Cl_{2} = 2PSCl_{3} +
4S, by the action of PCl_{5} upon certain sulphides, for example,
Sb_{2}S_{3}, by the reaction: 3MCl + P_{2}S_{5} = PSCl_{3} +
M_{3}PS_{4} (Glatzel, 1893), and in the reaction 3PCl_{3} +
SOCl_{2} = PCl_{5} + POCl_{3} + PSCl_{3}, showing the reducing
action of phosphorus trichloride, which is especially clear in the
reaction SO_{3} + PCl_{3} = SO_{2} + POCl_{3}. Thorpe and Rodger
(1889), by heating 3PbF_{2} or BiF_{3} with phosphorus
pentasulphide (and also by heating AsF_{3} and PSCl_{3} to 150°),
obtained thiophosphoryl fluoride as a colourless, spontaneously
inflammable gas (see further on, Note 74 bis, and Chapter XIX.,
Note 25). The action of PSCl_{3} upon NaHO gives a salt of
monothiophosphoric acid (Würtz, Kubierschky), H_{3}PSO_{3}, which
gives soluble salts of the alkalis.
The metallic sulphides corresponding with the metallic oxides have
either a feeble alkaline or a feeble acid character, according to the
character of the corresponding oxide, and therefore by combining together
they are able to form saline substances--that is, salts in which the
oxygen is replaced by sulphur. Thus sulphuretted hydrogen having the
properties of a feeble acid[21] has, at the same time, the properties of
water, and forms the type of the sulphur derivatives, which may also be
formed by means of sulphuretted hydrogen, just as the oxides may be
formed by the aid of water. But as sulphuretted hydrogen has acid
properties, it combines more easily with the basic metallic sulphides.
Hence, for instance, there exists a compound of sulphuretted hydrogen
with potassium sulphide, potassium hydrosulphide, 2KHS = K_{2}S + H_{2}S,
just as there are potassium hydroxides; but there are scarcely any
compounds of sulphuretted hydrogen with the sulphides corresponding with
acids. Thus the sulphides of the metals may be regarded either as salts
of sulphuretted hydrogen or as oxides of the metals in which the oxygen
is replaced by sulphur. In general terms the sulphides exhibit the same
degrees of difference with respect to their solubility in water as do the
oxides. Thus the oxides of the alkali metals, and of some of the metals
of the alkaline earths, are soluble in water, whilst those of nearly all
the other metals are insoluble. The same may be said as to the sulphides;
the sulphides of the metals of the alkalis and certain of the alkaline
earths are soluble in water, whilst those of the other metals are
insoluble. Those metals, like aluminium, whose oxides--for example,
Al_{2}O_{3}--have intermediate properties and do not form compounds with
feeble acids, at least in a wet way, also do not form sulphides by this
method, although these may be obtained indirectly. And in general the
sulphides of the metals are easily formed in a wet way, and with
particular ease if they are insoluble in water. In this case their salts
enter into double decomposition with sulphuretted hydrogen, or with
soluble sulphides, and give an insoluble sulphide--for instance, a salt
of lead gives lead sulphide with sulphuretted hydrogen. By the action of
sulphuretted hydrogen on a salt of a metal, a free acid must be formed
besides the metallic sulphide. Thus if a metal M be in a state of
combination MX_{2}, then by the action of sulphuretted hydrogen there
will be formed, besides MS,[22] an acid 2HX. It is evident that
sulphuretted hydrogen will not precipitate an insoluble sulphide from the
salts of those metals whose sulphides react with free acid, such as zinc,
iron, manganese, &c. The reaction FeCl_{2} + H_{2}S = FeS + 2HCl, and the
like, do not take place because the acid acts on the ferrous sulphide.
Antimonious sulphide is not acted on by dilute hydrochloric acid, but it
is decomposed by strong acid, and therefore in presence of an excess of
hydrochloric acid antimonious chloride does not entirely react with
hydrogen sulphide, whilst the reaction 2SbCl_{3} + 3H_{2}S = Sb_{2}S_{3}
+ 6HCl is a complete one in a dilute solution and with a small quantity
of acid. Those metallic sulphides which are decomposed by acids may be
obtained in a wet way by the double decomposition of the salts of the
metals, not with hydrogen sulphide, but with soluble metallic sulphides,
such as sulphide of ammonium or of potassium, because then no free acid
is formed, but a salt of the metal (potassium or ammonium) which was
taken as a soluble sulphide. So, for example, FeCl_{2} + K_{2}S = FeS +
2KCl.[23]
[21] Sulphuretted hydrogen does not saturate the alkaline properties of
alkali hydroxides, so that a solution of potassium hydroxide will
not under any circumstances give a neutral liquid with
sulphuretted hydrogen. In this case the sulphuretted hydrogen
forms in solution only an acid salt with the potassium: KHO +
H_{2}S = KHS + H_{2}O. It must be supposed that the normal salt is
not formed in the solution--that is, that the reaction 2KHO +
H_{2}S = K_{2}S + 2H_{2}O does not take place. This is seen from
the fact that a development of heat, depending on the formation of
potassium hydrosulphide, KHS, is remarked when as much hydrogen
sulphide is passed into a solution of potassium hydroxide as it
will absorb. But if a further quantity of potassium hydroxide be
added to the resultant solution, heat is not developed, whilst if
alkali be added to potassium acid sulphate or sodium acid
carbonate, heat is developed. It must not be concluded from this
that H_{2}S is a monobasic acid, for here there is a question of
the decomposing action of water upon K_{2}S; K_{2}S and H_{2}O in
reacting on each other should absorb heat if the reaction of KHS
upon KHO evolves heat. Furthermore, it must be taken into account
that potassium oxide, K_{2}O, and the anhydrous oxides like it,
also do not exist in solutions, for whenever they are formed they
immediately react with the water, forming caustic potash, KHO, &c.
In the same way, directly potassium sulphide, K_{2}S, is formed in
water it is decomposed into potassium hydroxide and hydrosulphide:
K_{2}S + H_{2}O = KHO + KHS. Potassium sulphide, K_{2}S, in a
solid state corresponds with K_{2}O, although neither can exist in
solution.
[22] During recent years (beginning with Schulze, 1882) it has been
found that many metallic sulphides which were considered totally
insoluble do, under certain circumstances, form very unstable
solutions in water, as already mentioned in Chapter I., Note 57.
Arsenic sulphide is very easily obtained in the form of a solution
(hydrosol). Solutions of copper and cadmium sulphides may also be
easily obtained by precipitating their salts CuX_{2}, or CdX_{2},
with ammonium sulphide, and washing the precipitate; but they are
re-precipitated by the addition of foreign salts.
[23] In reality the preceding reaction should be expressed thus:
FeCl_{2} + 2KHS = FeS + 2KCl + H_{2}S (Note 21), because in the
presence of water not K_{2}S but KHS reacts. But as the
sulphuretted hydrogen takes no part in the reaction, it is usual
to express the formation of such sulphides without taking the
hydrogen sulphide proceeding from the potassium or ammonium
hydrosulphides into account. It is not usual to employ potassium
sulphide but ammonium sulphide--or, to speak more accurately,
ammonium hydrosulphide--in order to avoid the formation of a
non-volatile salt of potassium and to have, together with the
formation of the sulphide, a salt of ammonium which can always be
driven off by evaporating the solution and igniting the
residue--for instance: FeCl_{2} + (NH_{4})_{2}S = FeS + 2NH_{4}Cl.
Thus the metallic sulphides may be divided into three chief
classes: (1) _those soluble in water_, (2) _those insoluble in
water but reacting with acids_, and (3) _those insoluble both in
water and acids_. The third class may be easily subdivided into
two groups; to the first group belong those sulphides which
correspond with bases or basic oxides, and are therefore unable to
play the part of an acid with the sulphides of the alkalis, and
are insoluble in NH_{4}HS, whilst the sulphides of the second
group are of an acid character, and give soluble thio-salts with
the sulphides of the alkaline metals, in which they play the part
of an acid. To this group belong those metals whose corresponding
oxides have acid properties. It must be observed, however, that
not all metallic acids have corresponding sulphides, partly owing
to the fact that certain acids are reducible by sulphuretted
hydrogen, especially when their lower degrees of oxidation are of
a basic character. Such are, for instance, the acids of chromium,
manganese, &c. Sulphuretted hydrogen converts them into lower
oxides, having the properties of bases. Those bases which do not
combine with feeble acids, such as carbonic acid and hydrogen
sulphide, give a precipitate of hydroxide with ammonium
sulphide--for example, aluminium salts react in this manner. This
difference of the metals in their behaviour towards sulphuretted
hydrogen gives a very valuable means of separating them from each
other, and _is taken advantage of in analytical chemistry_. If,
for instance, the metals of the first and third groups occur
together, it is only necessary to convert them into soluble salts,
and to act on the solution of the salts with sulphuretted
hydrogen; this will precipitate the metals of the third group in
the form of sulphides, whilst the metals of the first group will
not be in the least acted on. Such a method of separating the
metals is considered more fully in analytical chemistry, and we
will therefore limit ourselves here to pointing out to which
groups the most common metals belong, and the colour which is
proper to the sulphide precipitated.
_Metals which are precipitated by sulphuretted hydrogen_, as
sulphides from a solution of their salts, even in the presence of
free acid:
The precipitate is soluble in ammonium sulphide:
_Platinum_ (dark brown) | _Antimony_ (orange)
_Gold_ (dark brown) | _Arsenic_ (yellow)
_Tin_ (yellow and brown) |
The precipitate is insoluble in ammonium sulphide:
_Copper_ (black) | _Mercury_ (black)
_Silver_ (black) | _Lead_ (black)
_Cadmium_ (yellow) |
_Metals which are precipitated by ammonium sulphide_ from neutral
solutions, but not precipitated from acid solutions by
sulphuretted hydrogen:
The sulphide precipitated is soluble in hydrochloric acid:
_Zinc_ (white) | _Manganese_ (rose colour) | _Iron_ (black)
The sulphide precipitated is not soluble in dilute hydrochloric
acid:
_Nickel_ (black) | _Cobalt_ (black)
A hydroxide, and not a sulphide, is precipitated:
_Chromium_ (green) | _Aluminium_ (white)
The metals of the alkalis and of the alkaline earths are not
precipitated either by sulphuretted hydrogen or ammonium sulphide.
The metals of the alkaline earths when in acid solutions in the
form of phosphates and many other salts are precipitated by
ammonium sulphide, because the latter neutralises the free acid,
with formation of an ammonium salt of the acid and evolution of
sulphuretted hydrogen.
Metallic sulphides may be obtained by many other means besides the
action of sulphuretted hydrogen on salts and oxides, or by the simple
combination of metals with sulphur when heated or fused. Thus they may
also be formed by the reduction of sulphates by heating them with
charcoal or other means. Charcoal takes up the oxygen from many
sulphates, leaving corresponding sulphides. Thus sodium sulphate,
Na_{2}SO_{4}, when heated with charcoal, forms sodium sulphide, Na_{2}S.
Besides which metallic sulphides are also obtained by heating metals or
their oxides in the vapours of many sulphur compounds--for example, in
the vapour of carbon bisulphide, CS_{2}, when the carbon takes up the
oxygen and the sulphur combines with the metal. The sulphides formed in
this manner are often crystalline, and often appear with those properties
and in that crystalline form in which they occur in nature. Besides which
we must mention that many of the sulphides of the metals are oxidised in
air at the ordinary, and especially at a higher, temperature, forming
either SO_{2} and the oxide of the metal or sulphates. This oxidation
proceeds with particular ease, even at the ordinary temperature, when a
metallic sulphide is precipitated from its solutions, as a fine powder
containing water. The sulphides of iron and manganese, &c., are very
easily oxidised in this manner. But if these hydrates be ignited, they
lose their water (the ignition must be carried on in a stream of hydrogen
to prevent their oxidation during the process), become denser, and are no
longer oxidised at the ordinary temperature. Those sulphides whose
corresponding sulphates are decomposed by heat part with their sulphur in
the form of sulphurous anhydride when they are ignited in air, and the
metal, as a rule, remains behind as oxide. This is taken advantage of in
the treatment of sulphurous ores. The process is called _roasting_.
Hydrogen not only forms sulphuretted hydrogen with sulphur, but it also
combines with it in several other proportions, just as it combines with
oxygen, forming not only water but also hydrogen peroxide. Moreover these
_polysulphides of hydrogen_ are also unstable, like hydrogen peroxide,
and are also obtained from the corresponding polysulphides of the metals
of the alkaline earths, just as hydrogen peroxide is obtained from barium
peroxide. Thus calcium forms not only calcium sulphide, CaS, but also as
bi-, tri-, and pentasulphide, CaS_{5}, and all these compounds are
soluble in water. Sodium also combines with sulphur in the same
proportions, forming sulphides from Na_{2}S to Na_{2}S_{5}. If an acid be
added to a solution of a polysulphide, it gives sulphur, sulphuretted
hydrogen, and a salt of the metal. For instance, MS_{5}, + 2HCl = MCl_{2}
+ H_{2}S + 4S. If we reverse the operation, and pour a solution of a
polysulphide into an acid, sulphur is not precipitated, but an oily
liquid is formed which is heavier than water and insoluble in it. This is
the polysulphide of hydrogen: MS_{5} + 2HCl = MCl_{2} + H_{2}S_{5}. As
Rebs showed (1888), whatever polysulphide be taken--of sodium, for
instance--it always gives one and the same _hydrogen pentasulphide_,[24]
of specific gravity 1·71 (15°). It can only be preserved in the absence
of water and at low temperatures, and then not for long: for, especially
in the presence of alkalis and when slightly warmed, it splits up very
easily into sulphuretted hydrogen and sulphur.[25]
[24] Rebs took di-, tri-, tetra-, and pentasulphides of sodium,
potassium, and barium, which he prepared by dissolving sulphur in
solutions of the normal sulphides; on adding hydrochloric acid he
always obtained hydrogen pentasulphide, whence it is evident that
4H_{2}S_{n} = (_n_ - 1)H_{2}S_{5} + (5 - _n_)H_{2}S. For example,
if H_{2}S_{2} were formed, it would decompose according to the
equation 4H_{2}S_{2} = H_{2}S_{5} + 3H_{2}S. The hydrogen
pentasulphide formed breaks up into hydrogen sulphide and sulphur
when brought into contact with water. Previous to Rebs' researches
many chemists stated that all polysulphides gave the bisulphide
H_{2}S_{2}, and Hofmann recognised only hydrogen trisulphide,
H_{2}S_{3}.
[25] The formation of the polysulphides of hydrogen, H_{2}S_{n} is
easily understood from the law of substitution, like that of the
saturated hydrocarbons, C_{n}H_{2n + 2}, knowing that sulphur
gives H_{2}S, because the molecule of sulphuretted hydrogen may be
divided into H and HS. This radicle, HS, is equivalent to H. By
substituting this radicle for hydrogen in H_{2}S we obtain (HS)HS
= H_{2}S_{2}, (HS)(HS)S = H_{2}S_{3}, &c., in general H_{2}S_{n}.
The homologues of CH_{4}, C_{n}H_{2n + 2} are formed in this
manner from CH_{4}, and consequently the polysulphides H_{2}S_{n}
are the homologues of H_{2}S. The question arises why in
H_{2}S_{n} the apparent limit of _n_ is 5--that is, why does the
substitution end with the formation of H_{2}S_{5}? The answer
appears to me to be clearly because in the molecule of sulphur,
S_{6}, there are six atoms of sulphur (Note 11). The forces in one
and the other case are the same. In the one case they hold S_{6}
together, in the other S_{5} and H_{2}; and, judging from H_{2}S,
the two atoms of hydrogen are equal in power and significance to
the atom of sulphur. Just as hydrogen peroxide, H_{2}O_{2},
expresses the composition of ozone, O_{3}, in which O is replaced
by H_{2}, so also H_{2}S_{5} corresponds with S_{6}.
The soluble sulphides and polysulphides of the metals of the alkalis
and alkaline earths--for example, of ammonium,[26] potassium,[27] and
calcium,[28]--have the appearance and properties of salts, just as the
hydrated oxides have, whilst the sulphides of the metals of the higher
groups resemble their oxides and have not at all the appearance
of salts, and this is more especially the case with regard to the
crystalline forms in which they frequently occur in nature.[29]
[26] _Ammonium sulphide_, (NH_{4})_{2}S, may be prepared by passing
sulphuretted hydrogen into a vessel full of dry ammonia, or by
passing both dry gases together into a very cold receiver. In the
latter case it is necessary to prevent the access of air, and to
have an excess of ammonia. Under these circumstances, two volumes
of ammonia combine with one volume of sulphuretted hydrogen, and
form a colourless, very volatile, crystalline substance, having a
very unpleasant odour, which is very poisonous and exceedingly
unstable. When exposed to the air it absorbs oxygen and acquires a
yellow colour, and then contains oxygen and polysulphide compounds
(because a portion of the hydrogen sulphide gives water and
sulphur). It is soluble in water and forms a colourless solution,
which, however, in all probability contains free ammonia and the
acid salt--that is, ammonium hydrosulphide, NH_{4}HS, or
(NH_{4})_{2},S,H_{2}S. This salt is formed when dry ammonia is
mixed with an excess of dry sulphuretted hydrogen. The compound
contains equal volumes of the components NH_{3} + H_{2}S =
(NH_{4})HS. It crystallises in an anhydrous state in colourless
plates, and may be easily volatilised (dissociating like ammonium
chloride), even at the ordinary temperature; it has an alkaline
reaction, absorbs oxygen from the air, is soluble in water, and
its solution is usually prepared by saturating an aqueous solution
of ammonia with sulphuretted hydrogen. According to the ordinary
rule, these salts, like other ammonium salts, split up into
ammonia and sulphuretted hydrogen when they are distilled.
A solution of ammonium sulphide is able to dissolve sulphur, and
it then contains compounds of hydrogen polysulphide and ammonia.
Some of these compounds may be obtained in a crystalline form.
Thus Fritzsche obtained a compound of ammonia with hydrogen
pentasulphide, or ammonium pentasulphide, (NH_{4})_{2}S_{5}, in
the following manner: He saturated an aqueous solution of ammonia
with sulphuretted hydrogen, added powdered sulphur to it, and
passed ammonia gas into the solution, which then absorbed a fresh
amount. After this he again passed sulphuretted hydrogen into the
solution, and then added sulphur, and then again ammonia. After
repeating this several times, orange-yellow crystals of
(NH_{4})_{2}S_{5} separated out from the liquid. These crystals
melted at 40° to 50°, and were very unstable.
When a solution of ammonium hydrosulphide, prepared by saturating
a solution of ammonia with sulphuretted hydrogen, is exposed to
the air, it turns yellow, owing to the presence of an ammonium
polysulphide, whose formation is due to the sulphuretted hydrogen
being oxidised by the air and converted into water and sulphur,
which is dissolved by the ammonium sulphide. In certain analytical
reactions it is usual to employ a solution of ammonium sulphide
which has been kept for some time and acquired a yellow colour.
This yellow sulphide of ammonium deposits sulphur when saturated
with acids, whilst a freshly-prepared solution only evolves
sulphuretted hydrogen. The yellow solution furthermore contains
ammonium thiosulphate, which is derived not only from the
oxidation of the ammonium sulphide, but also from the action of
the liberated sulphur on the ammonia, just as an alkaline salt of
thiosulphuric acid and a sulphide are formed by the action of
sulphur on a solution of a caustic alkali.
[27] _Potassium sulphide_, K_{2}S, is obtained by heating a mixture of
potassium sulphate and charcoal to a bright-red heat. It may be
prepared in solution by taking a solution of potassium hydroxide,
dividing it into two equal parts, and saturating one portion with
sulphuretted hydrogen so long as it is absorbed. This portion will
then contain the acid salt KHS (Note 21). The two portions are
then mixed together, and potassium sulphide will then be obtained
in the solution. This solution has a strongly alkaline reaction,
and is colourless when freshly prepared, but it very easily
undergoes change when exposed to the air, forming potassium
thiosulphate and polysulphides. When the solution is evaporated at
low temperatures under the receiver of an air-pump, it yields
crystals containing K_{2}S,5H_{2}O (heated at 150°, they part with
3 mol. H_{2}O, and at higher temperatures they lose nearly all
their water without evolving sulphuretted hydrogen). When they are
ignited in glass vessels they corrode the glass. When a solution
of caustic potash, completely saturated with sulphuretted
hydrogen, is evaporated under the receiver of an air-pump it forms
colourless rhombohedra of _potassium hydrosulphide_,
2(KHS),H_{2}O,K_{2}S,H_{2}S,H_{2}O. These crystals are
deliquescent in the air, but do not change in a vacuum when heated
up to 170°, and at higher temperatures they lose water but do not
evolve sulphuretted hydrogen. The anhydrous compound, KHS, fuses
at a dark-red heat into a very mobile yellow liquid, which
gradually becomes darker in colour and solidifies to a red mass.
It is remarkable that when a solution of the compound KHS is
boiled it somewhat easily evolves half its sulphuretted hydrogen,
leaving potassium sulphide, K_{2}S, in solution; and a solution of
the latter in water is also able to evolve sulphuretted hydrogen
on prolonged boiling, but the evolution cannot be rendered
complete, and, therefore, at a certain temperature, a solution of
potassium sulphide will not be capable of absorbing sulphuretted
hydrogen at all. From this we must conclude that potassium
hydroxide, water, and sulphuretted hydrogen form a system whose
complex equilibrium is subject to the laws of dissociation,
depends on the relative mass of each substance, on the
temperature, and the dissociation pressure of the component
elements. Potassium sulphide is not only soluble in water, but
also in alcohol.
Berzelius showed that in addition to potassium sulphide there also
exist potassium bisulphide, K_{2}S_{2}; trisulphide, K_{2}S_{3};
tetrasulphide, K_{2}S_{4}; and pentasulphide, K_{2}S_{5}.
According to the researches of Schöne, the last three are the most
stable. These different compounds of potassium and sulphur may be
prepared by fusing potassium hydroxide or carbonate with an excess
of sulphur in a porcelain crucible in a stream of carbonic
anhydride. At about 600° potassium pentasulphide is formed; this
is the highest sulphur compound of potassium. When heated to 800°
it loses one-fifth of its sulphur and gives the tetrasulphide,
which at this temperature is stable. At a bright-red heat--namely,
at about 900°--the trisulphide is formed. This compound may be
also formed by igniting potassium carbonate in a stream of carbon
bisulphide, in which case a compound, K_{2}CS_{3}, is first formed
corresponding to potassium carbonate, and carbonic anhydride is
evolved. On further ignition this compound splits up into carbon
and potassium trisulphide, K_{2}S_{3}. The tetrasulphide may also
be obtained in solution if a solution of potassium sulphide be
boiled with the requisite amount of sulphur without access of air.
This solution yields red crystals of the composition
K_{2}S_{4},2H_{2}O when it is evaporated in a vacuum. These
crystals are very hygroscopic, easily soluble in water, but very
sparingly in alcohol; when ignited they give off water,
sulphuretted hydrogen, and sulphur. If a solution of potassium
sulphide be boiled with an excess of sulphur it forms the
pentasulphide, which, however, is decomposed on prolonged boiling
into sulphuretted hydrogen and potassium thiosulphate: K_{2}S_{5}
+ 3H_{2}O = K_{2}S_{2}O_{3} + 3H_{2}S. A substance called _liver
of sulphur_ was formerly frequently used in chemistry and
medicine. Under this name is known the substance which is formed
by boiling a solution of caustic potash with an excess of flowers
of sulphur. This solution contains a mixture of potassium
pentasulphide and thiosulphate, 6KHO + 12S = 2K_{2}S_{5} +
K_{2}S_{2}O_{3} + 3H_{2}O. The substance obtained by fusing
potassium carbonate with an excess of sulphur was also known as
liver of sulphur. If this mixture be heated to an incipient
dark-red heat it will contain potassium thiosulphate, but at
higher temperatures potassium sulphate is formed. In either case a
polysulphide of potassium is also present. The sulphides of
sodium, for example Na_{2}S, NaHS, &c., in many respects closely
resemble the corresponding potassium compounds.
[28] The metals of the alkaline earths, like those of the alkalis, form
several compounds with sulphur; thus calcium forms compounds with
one and with five atoms of sulphur. There are doubtless also
intermediate sulphides. If sulphuretted hydrogen be passed over
ignited lime it forms water and _calcium sulphide_, which may also
be formed by heating calcium sulphate with charcoal, whilst if
sulphur be heated with lime or with calcium carbonate, then
naturally oxygen compounds (calcium thiosulphate and sulphate) are
formed at the same time as calcium sulphide. The prolonged action
of the vapour of carbon bisulphide, especially when mixed with
carbonic anhydride, on strongly ignited calcium carbonate entirely
converts it into sulphide. Calcium sulphide is generally obtained
as an almost colourless, opaque, brittle mass, which is infusible
at a white heat, and is soluble in water. The act of solution (as
with K_{2}S, Note 21) is partly accompanied by a double
decomposition with the water. When heated, dry calcium sulphide
does not absorb oxygen from the air. An excess of water decomposes
it, like many other metallic sulphides, precipitating lime (as a
product of the decomposition the lime hinders the action of the
water upon the CaS; see soda refuse, Chapter XII., Note 12), and
forming a hydrosulphide, CaH_{2}S_{2}, in solution. This compound
is also formed by passing sulphuretted hydrogen through an aqueous
solution of calcium sulphide or lime. Its solution, like that of
calcium sulphide, has an alkaline reaction. It decomposes when
evaporated, and absorbs oxygen from the air. _Calcium
pentasulphide_, CaS_{5}, is not known in a pure state, but may be
obtained in admixture with calcium thiosulphate by boiling a
solution of lime or calcium sulphide with sulphur: 3CaH_{2}O_{2} +
12S = 2CaS_{5} + CaS_{2}O_{3} + 3H_{2}O. A similar compound in an
impure form is formed by the action of air on alkali waste, and is
used for the preparation of thiosulphates.
Many of the sulphides of the metals of the alkaline earths are
phosphorescent--that is, they have the faculty of _emitting
light_, after having been subjected to the action of sunlight, or
of any bright source of light (Canton phosphorus, &c.). The
luminosity lasts some time, but it is not permanent, and gradually
disappears. This phosphorescent property is inherent, in a greater
or less degree, to nearly all substances (Becquerel), but for a
very short time, whilst with calcium sulphide it is comparatively
durable, lasting for several hours, and Dewar (1894) showed that
it is far more intense at very low temperatures (for instance, in
bodies cooled in liquid oxygen to -182°). It is due to the
excitation of the surfaces of substances by the action of light,
and is determined by those rays which exhibit a chemical action.
Hence daylight or the light of burning magnesium, &c., acts more
powerfully than the light of a lamp, &c. Warnerke has shown that a
small quantity of magnesium lighted near the surface of a
phosphorescent substance rapidly excites the greatest possible
intensity of luminosity; this enabled him to found a method of
measuring the intensity of light--_i.e._ to obtain a constant unit
of light--and to apply it to photography. The nature of the change
which is accomplished on the surface of the luminous substance is
at present unknown, but in any case it is a renewable one, because
the experiment may be repeated for an infinite number of times and
takes place in a vacuum. The intensity and tint of the light
emitted depend on the method of preparation of the calcium
sulphide, and on the degree of ignition and purity of the calcium
carbonate taken. According to the observations of Becquerel, the
presence of compounds of manganese, bismuth, &c., sodium sulphide
(but not potassium sulphide), &c., although in minute traces, is
perfectly indispensable. This gives reason for thinking that the
formation (in the dark) and decomposition (in light) of double
salts like MnS,Na_{2}S perhaps form the chemical cause of the
phenomena. Compounds of strontium and barium have this property to
even a greater extent than calcium sulphide. These compounds may
be prepared as in the following example: A mixture of sodium
thiosulphate and strontium chloride is prepared; a double
decomposition takes place between the salts, and, on the addition
of alcohol, strontium thiosulphate, SrS_{2}O_{3}, is precipitated,
which, when ignited, leaves strontium sulphide behind. The
strontium sulphide thus prepared emits (when dry) a greenish
yellow light. It contains a certain amount of sulphur, sodium
sulphide, and strontium sulphate. By ignition at various
temperatures, and by different methods of preparation, it is
possible to obtain mixtures which emit different coloured lights.
[29] As examples, we will describe the sulphides of arsenic, antimony,
and mercury. Arsenic trisulphide, or _orpiment_, As_{2}S_{3},
occurs native, and is obtained pure when a solution of arsenious
anhydride in the presence of hydrochloric acid comes into contact
with sulphuretted hydrogen (there is no precipitate in the absence
of free acid). A beautiful yellow precipitate is then obtained:
As_{2}O_{3} + 3H_{2}S = 3H_{2}O + As_{2}S_{3}; it fuses when
heated, and volatilises without decomposition. As_{2}S_{3} is
easily obtained in a colloid form (Chapter I., Note 57). When
fused it forms a semi-transparent, yellow mass, and it is thus
that it enters the market. The specific gravity of native orpiment
is 3·4, and that of the artificially-fused mass is 2·7. It is used
as a yellow pigment, and owing to its insolubility in water and
acids it is less injurious than the other compounds corresponding
to arsenious acid. According to the type AsX_{2}, realgar, AsS, is
known, but it is probable that the true composition of this
compound is As_{4}S_{4}--that is, it presents the same relation to
orpiment as liquid phosphuretted hydrogen does to gaseous.
_Realgar_ (_Sandaraca_) occurs native as brilliant red crystals of
specific gravity 3·59, and may be prepared artificially by fusing
arsenic and sulphur in the proportions indicated by its formulæ.
It is prepared in large quantities by distilling a mixture of
sulphur and arsenical pyrites. Like orpiment it dissolves in
calcium sulphide, and even in caustic potash. It is used for
signal lights and fireworks, because it deflagrates and gives a
large and very brilliant white flame with nitre.
With antimony, sulphur gives a tri- and a pentasulphide. The
former, Sb_{2}S_{3}, which corresponds with antimonious oxide,
occurs native (Chapter XIX.) in a crystalline form; its sp. gr. is
then 4·9, and it presents brilliant rhombic crystals of a grey
colour, which fuse when heated. A substance of the same
composition is obtained as an amorphous orange powder by passing
sulphuretted hydrogen into an acid solution of antimonious oxide.
In this respect antimonious oxide again reacts like arsenious
acid, and the sulphides of both are soluble in ammonium and
potassium sulphides, and, especially in the case of arsenious
sulphide, are easily obtained in colloidal solutions. By prolonged
boiling with water, antimonious sulphide may be entirely converted
into the oxide, hydrogen sulphide being evolved (Elbers). Native
antimony sulphide, or the orange precipitated trisulphide when
fused with dry, or boiled with dissolved, alkalis, forms a
dark-coloured mass (Kermes mineral) formerly much used in
medicine, which contains a mixture of antimonious sulphide and
oxide. There are also compounds of these substances. A so-called
antimony vermilion is much used as a dye; it is prepared by
boiling sodium thiosulphate (six parts) with antimony trichloride
(five parts) and water (fifty parts). This substance probably
contains an oxysulphide of antimony--that is, a portion of the
oxygen in the oxide of antimony in it is replaced by sulphur. Red
antimony ore, and antimony glass, which is obtained by fusing the
trisulphide with antimonious oxide, have a similar composition,
Sb_{2}OS_{2}. In the arts, the _antimony pentasulphide_,
Sb_{2}S_{5}, is the most frequently used of the sulphur compounds
of antimony. It is formed by the action of acids on the so-called
Schlippe's salt, which is a _sodium thiorthantimonate_,
SbS(NaS)_{3}, corresponding with (Chapter XIX., Note 41 bis)
orthantimonic acid, SbO(OH)_{3}, with the replacement of oxygen by
sulphur. It is obtained by boiling finely-powdered native antimony
trisulphide with twice its weight of sodium carbonate, and half
its weight of sulphur and lime, in the presence of a considerable
quantity of water. The processes taking place are as follows:--The
sodium carbonate is converted into hydroxide by the lime, and then
forms sodium sulphide with the sulphur; the sodium sulphide then
dissolves the antimony sulphide, which in this form already
combines with the greatest amount of sulphur, so that a compound
is formed corresponding with antimony pentasulphide dissolved in
sodium sulphide. The solution is filtered and crystallised, care
being taken to prevent access of air, which oxidises the sodium
sulphide. This salt crystallises in large, yellowish crystals,
which are easily soluble in water and have the composition
Na_{3}SbS_{4},9H_{2}O. When heated they lose their water of
crystallisation and then fuse without alteration; but when in
solution, and even in crystalline form, this salt turns brown in
air, owing to the oxidation of the sulphur and the breaking up of
the compound. As it is used in medicine, especially in the
preparation of antimony pentasulphide, it is kept under a layer of
alcohol, in which it is insoluble. Acids precipitate antimony
pentasulphide from a solution of this salt, as an orange powder,
insoluble in acids and very frequently used in medicine (_sulfur
auratum antimonii_). This substance when heated evolves vapours of
sulphur, and leaves antimony trisulphide behind.
Mercury forms compounds with sulphur of the same types as it does
with oxygen. Mercurous sulphide, Hg_{2}S, easily splits up into
mercury and mercuric sulphide. It is obtained by the action of
potassium sulphide on mercurous chloride, and also by the action
of sulphuretted hydrogen on solutions of salts of the type HgX.
Mercuric sulphide, HgS, corresponding with the oxide, is cinnabar;
it is obtained as a black precipitate by the action of an excess
of sulphuretted hydrogen on solutions of mercuric salts. It is
insoluble in acids, and is therefore precipitated in their
presence. If a certain amount of water containing sulphuretted
hydrogen be added to a solution of mercuric chloride, it first
gives a white precipitate of the composition
Hg_{3}S_{2}Cl_{2}--that is, a compound HgCl,2HgS, a sulphochloride
of mercury like the oxychloride. But in the presence of an excess
of sulphuretted hydrogen, the black precipitate of mercuric
sulphide is formed. In this state it is not crystalline (the red
variety is formed by the prolonged action of polysulphides of
ammonium upon the black HgS), but if it be heated to its
temperature of volatilisation it forms a red crystalline sublimate
which is identical with native cinnabar. In this form its specific
gravity is 8·0, and it forms a red powder, owing to which it is
used as a red pigment (vermilion) in oil, pastel, and other
paints. It is so little attacked by reagents that even nitric acid
has no action on it, and the gastric juices do not dissolve it, so
that it is not poisonous. When heated in air, the sulphur burns
away and leaves metallic mercury. On a large scale cinnabar is
usually prepared in the following manner: 300 parts of mercury and
115 parts of sulphur are mixed together as intimately as possible
and poured into a solution of 75 parts of caustic potash in 425
parts of water, and the mixture is heated at 50° for several
hours. Red mercury sulphide is thus formed, and separates out from
the solution. The reaction which takes place is as follows: A
soluble compound, K_{2}HgS_{2}, is first formed; this compound is
able to separate in colourless silky needles, which are soluble in
the caustic potash, but are decomposed by water, and at 50°; this
solution (perhaps by attracting oxygen from the air) slowly
deposits HgS in a crystalline form.
Spring conducted an interesting research (at Liège, 1894) upon the
conversion of the black amorphous sulphide of mercury, HgS, into
red crystalline cinnabar. This research formed a sequel to
Spring's classical researches on the influence of high pressures
upon the properties of solids and their capacity for mutual
combination. He showed, among other things, that ordinary solids
and even metals (for instance, Pb), after being considerably
compressed under a pressure of 20,000 atmospheres, return on
removal of the pressure to their original density like gases. But
this is only true when the compressed solid is not liable to an
allotropic variation, and does not give a denser variety. Thus
prismatic sulphur (sp. gr. 1·9) passes under pressure into the
octahedral (sp. gr. 2·05) variety. Black HgS (precipitated from
solution) has a sp. gr. 7·6, while that of the red variety is 8·2,
and therefore it might be expected that the former would pass into
the latter under pressure, but experiments both at the ordinary
and a higher temperature did not give the looked-for result,
because even at a pressure of 20,000 atmospheres the black
sulphide was not compressed to the density of cinnabar (a pressure
of as much as 35,000 atmospheres was necessary, which could not be
attained in the experiment). But Spring prepared a black HgS,
which had a sp. gr. of 8·0, and this, under a pressure of 2,500
atmospheres, passed into cinnabar. He obtained this peculiar black
variety of HgS (sp. gr. 8·0) by distilling cinnabar in an
atmosphere of CO_{2}, when the greater portion of the HgS is
redeposited in the form of cinnabar. Under the action of a
solution of polysulphide of ammonium, this variety of HgS passes
more slowly into the red variety than the precipitated variety
does, while under pressure the conversion is comparatively easy.
It is worthy of remark, that Linder and Picton obtained complex
compounds of many of the sulphides of the heavy metals (Ca, Hg,
Sb, Zn, Cd, Ag, Au) with H_{2}S, for example H_{2}S,7CuS (by the
action of H_{2}S upon the hydrate of oxide of copper), H_{2}S,9CuS
(in the presence of acetic acid and with an excess of H_{2}S), &c.
Probably we have here a sort of 'solid' solution of H_{2}S in the
metallic sulphides.
As the acids derived from chlorine, phosphorus, and carbon are the
oxidised hydrogen compounds of these elements, so also we can form an
idea of the acid hydrates of sulphur, or of _the normal acids of
sulphur_, by representing them as the oxidised products of sulphuretted
hydrogen--
HCl H_{2}S H_{3}P H_{4}C
HClO H_{2}SO(?) H_{3}PO(?) H_{4}CO
HClO_{2} H_{2}SO_{2}(?) H_{3}PO_{2} H_{4}CO_{2}
HClO_{3} H_{2}SO_{3} H_{3}PO_{3} H_{4}CO_{3}
HClO_{4} H_{2}SO_{4} H_{3}PO_{4} H_{4}CO_{4}[30]
In the case of chlorine, if not all the hydrates, at all events salts of
all the normal hydrates are known, whilst in the case of sulphur only the
acids H_{2}S, H_{2}SO_{3} and H_{2}SO_{4} are known. But, on the other
hand, the latter are obtained not only as hydrates but also as stable
anhydrides, SO_{2} and SO_{3}, which are formed with the evolution of
heat from sulphur and oxygen; 32 parts of sulphur in combining with 32
parts of oxygen--that is, in forming SO_{2}--evolve 71,000 heat
units,[31] and if the oxidation proceeds to the formation of SO_{3},
103,000 heat units are evolved. These figures may be compared with those
which correspond with the passage of carbon into CO and CO_{2}, when
29,000 and 97,000 units of heat are evolved. This determines the
stability of the higher oxides of sulphur, and also expresses the
peculiarity of sulphur as an element which, although an analogue of
oxygen, forms stable compounds with it, and thus fundamentally differs
from chlorine. The higher and lower oxides of chlorine are powerful
oxidising agents, whilst the higher oxide of sulphur, SO_{3}, has but
feeble oxidising powers, and the lower oxide, SO_{2}, frequently acts as
a reducing agent, and is formed by the direct combustion of sulphur, just
as carbonic anhydride, CO_{2}, proceeds from the combustion of carbon. In
the combustion of sulphur, and also in the oxidation (roasting) of the
sulphides and polysulphides by their ignition in air, _sulphurous oxide_,
or _sulphurous anhydride_, or _sulphur dioxide_, SO_{2},[31 bis] is
exclusively formed. It is prepared on a large scale by burning sulphur or
roasting iron pyrites or other sulphides[32] for the manufacture of
sulphuric acid (Chapter VI.), and for direct application in the
manufacture of wine or for bleaching tissues and other purposes. In the
latter instances its application is based on the fact that sulphurous
anhydride acts on certain vegetable matters, and has the property of a
reducing and feeble acid.[32 bis]
[30] CH_{4} gives CH_{4}O or CH_{3}(OH), wood spirit; CH_{4}O_{2} or
CH_{2}(OH)_{2}, which decomposes into water and CH_{2}O--that is,
methylene oxide or formaldehyde; CH_{4}O_{3} = CH(OH)_{3} = H_{2}O
+ CHO(OH), or formic acid; and CH_{4}O_{4} = C(OH)_{4} = 2H_{2}O +
CO_{2}. There are four typical hydrogen compounds, RH, RH_{2},
RH_{3}, and RH_{4}, and each of them has its typical oxide. Beyond
H_{4} and O_{4} combination does not proceed.
[31] Rhombic sulphur, 71,080 heat units; monoclinic sulphur, 71,720
units, according to Thomsen.
[31 bis] However, when sulphur or metallic sulphides burn in an excess
of air, there is always formed a certain, although small, amount
of SO_{3}, which gives sulphuric acid with the moisture of the
air.
[32] The enormous amount of sulphuric acid now manufactured is chiefly
prepared by roasting native pyrites, but a considerable amount of
the SO_{2} for this purpose is obtained by roasting zinc blende
(ZnS) and copper and lead sulphides. A certain amount is also
procured from soda refuse (Note 6) and the residues obtained from
the purification of coal gas.
[32 bis] Sulphurous anhydride is also obtained by the decomposition of
many sulphates, especially of the heavy metals, by the action of
heat; but this requires a very powerful heat. This formation of
sulphurous anhydride from sulphates is based on the decomposition
proper to sulphuric acid itself. When sulphuric acid is strongly
heated (for instance, by dropping it upon an incandescent surface)
it is decomposed into water, oxygen, and sulphurous
anhydride--that is, into those compounds from which it is formed.
A similar decomposition proceeds during the ignition of many
sulphates. Even so stable a sulphate as gypsum does not resist the
action of very high temperatures, but is decomposed in the same
manner, lime being left behind. The decomposition of sulphates by
heat is accomplished with still greater facility in the presence
of sulphur, because in this case the liberated oxygen combines
with the sulphur and the metal is able to form a sulphide. Thus
when ferrous sulphate (green vitriol) is ignited with sulphur, it
gives ferrous sulphide and sulphurous anhydride: FeSO_{4} + 2S =
FeS + 2SO_{2}, and this reaction may even be used for the
preparation of this gas. At 400° sulphuric acid and sulphur give
an extremely uniform stream of pure sulphurous anhydride, so that
it is best prepared on a manufacturing scale by this method. Iron
pyrites, FeS_{2}, when heated to 150° with sulphuric acid (sp. gr.
1·75) in cast-iron vessels also gives an abundant and uniform
supply of sulphurous anhydride.
In the laboratory--that is, on a small scale--sulphurous anhydride is
best prepared by deoxidising sulphuric acid by heating it with charcoal,
or copper, sulphur, mercury, &c. Charcoal produces this decomposition of
sulphuric acid at but moderately high temperatures; it is itself
converted into carbonic anhydride,[32 tri] and therefore when sulphuric
acid is heated with charcoal it evolves a mixture of sulphurous and
carbonic anhydrides: C + 2H_{2}SO_{4} = CO_{2} + 2SO_{2} + 2H_{2}O. The
metals which are unable to decompose water, and which do not, therefore,
expel hydrogen from sulphuric acid, are frequently capable of decomposing
sulphuric acid, with the evolution of sulphurous anhydride, just as they
decompose nitric acid, forming the lower oxides of nitrogen. These metals
are silver, mercury, copper, lead, and others. Thus, for example, the
action of copper on sulphuric acid may be expressed by the following
equation: Cu + 2H_{2}SO_{4} = CuSO_{4} + SO_{2} + 2H_{2}O. In the
laboratory this reaction is carried on in a flask with a gas-conducting
tube, and does not take place unless aided by heat.[33]
[32 tri] Mellitic acid is formed at the same time (Verneuille).
[33] The thermochemical data connected with this reaction are as
follows: A molecule of hydrogen H_{2}, in combining with oxygen (O
= 16) develops about 69,000 heat units, whilst the molecule of
SO_{2}, in combining with oxygen only develops about 32,000 heat
units--that is, about half as much--and therefore those metals
which cannot decompose water may still be able to deoxidise
sulphuric into sulphurous acid. Those metals which decompose water
and sulphuric acid with the evolution of hydrogen, evolve in
combining with sixteen parts by weight of oxygen more heat than
hydrogen does--for example, K_{2}, Na_{2}, Ca develop about or
more than 100,000 heat units; Fe, Zn, Mn about 70,000 to 80,000
heat units; whilst those metals which neither decompose water nor
evolve hydrogen from sulphuric acid, but are still capable of
evolving sulphurous anhydride from it, develop less heat with
oxygen than hydrogen, but nearly the same amount, if not more
than, sulphurous anhydride develops--for example, Cu and Hg
develop about 40,000 and Pb about 50,000 heat units.
In its physical and chemical properties sulphurous anhydride presents a
great _resemblance to carbonic anhydride_. It is a heavy gas, somewhat
considerably soluble in water, very easily condensed into a liquid; it
forms normal and acid salts, does not evolve oxygen under the direct
action of heat,[34] although such metals as sodium and magnesium burn in
it, just as in carbonic anhydride. It has a suffocating odour, which is
well known owing to its being evolved when sulphur or sulphur matches are
burnt. In characterising the properties of sulphurous anhydride, it is
very important to remember (Chapter II.) also that it is more easily
liquefied (at -10°, or at 0° under two atmospheres pressure) than
carbonic anhydride (thirty-six atmospheres at 0°),[35] that it is more
soluble than carbonic anhydride (Vol. I. p. 79); at 0°, 100 vols. of
water dissolve 180 vols. of carbonic anhydride and 688 vols. of sulphuric
anhydride), that the molecular weight of SO_{2} = 64 and of CO_{2} = 44,
and that the density of liquid sulphurous anhydride at 0° = 1·43
(molecular volume = 45) and of carbonic anhydride = 0·95 (molecular
volume = 49). Although sulphur dioxide is the anhydride of an acid,
nevertheless, like carbonic anhydride, it does not form any stable
compounds with water, but gives a solution from which it may be entirely
expelled by the action of heat.[36] The acid character of sulphurous
anhydride is clearly expressed by the fact that it is entirely absorbed
by alkalis, with which it forms acid and normal salts easily soluble in
water. With salts of barium, calcium, and the heavy metals, the normal
salts of the alkalis, M_{2}SO_{3}, give precipitates exactly like those
formed by the carbonates. In general, the salts of sulphurous acid are
closely analogous to the corresponding carbonates.
[34] That is, it only dissociates and re-forms the original product on
cooling.
[35] At a given temperature the pressure of this gas evolved from any
salt will be less than that of carbonic anhydride, if we compare
the separation of a gas from its salts with the phenomenon of
evaporation, as was done in discussing the decomposition of
calcium carbonate.
Liquid sulphurous anhydride is used on a large scale (Pictet) for
the production of cold.
[36] De la Rive, Pierre, and more especially Roozeboom, have
investigated the crystallo-hydrate which is formed by sulphurous
anhydride and water at temperatures below 7° under the ordinary
pressure, and in closed vessels (at temperatures below 12°). Its
composition is SO_{2},7H_{2}O, and density 1·2. This hydrate
corresponds with the similar hydrate CO_{2},8H_{2}O obtained by
Wroblewsky.
_Acid sodium sulphite_, NaHSO_{3}, may be obtained by passing sulphurous
anhydride into a solution of sodium hydroxide. It is also formed by
saturating a solution of sodium carbonate with the gas (carbonic
anhydride is then given off), and as the solubility of the acid sulphite
is much greater than that of the carbonate, a further quantity of the
latter may be dissolved after the passage of the sulphurous anhydride, so
that ultimately a very strong solution of the sulphite may be formed in
this manner, from which it may be obtained in a crystalline form, either
by cooling and evaporating (without heating, for then the salt would give
off sulphurous anhydride) or by adding alcohol to the solution. When
exposed to the air this salt loses sulphurous anhydride and attracts
oxygen, which converts it into sodium sulphate. The acid sulphites of the
alkali metals are able to combine not only with oxygen, but also with
many other substances--for example, a solution of the sodium salt
dissolves sulphur, forming sodium thiosulphate, gives crystalline
compounds with the aldehydes and ketones, and dissolves many bases,
converting them into double sulphites. Having the faculty of attracting
or absorbing oxygen, acid sodium sulphite is also able to absorb
chlorine, and is therefore employed, like sodium thiosulphate, for the
removal of chloride (as an antichlor), especially in the bleaching of
fabrics, when it is necessary to remove the last traces of the chlorine
held in the tissues, which might otherwise have an injurious effect on
them. If a solution of an alkali hydroxide be divided into two parts, and
one half is saturated with sulphurous anhydride, and then the other half
added to it, a normal salt will be obtained in the solution, having an
alkaline reaction, like a solution of sodium carbonate. The acid salt has
a neutral reaction.[36 bis] Like sodium carbonate, _normal sodium
sulphite_ has the composition Na_{2}SO_{3},10H_{2}O, and its maximum
solubility is at 33°--in a word, it very closely resembles sodium
carbonate. Although this salt does not give off sulphurous anhydride from
its solution, it is able, like the acid salt, to absorb oxygen from the
air, and is then converted into sodium sulphate.[37]
[36 bis] Schwicker (1889) by saturating NaHSO_{3} with potash, or
KHSO_{3} with soda, obtained NaKSO_{3}, in the first instance with
H_{2}O, and in the second instance with 2H_{2}O, probably owing to
the different media in which the crystals are formed. In general
sulphurous acid easily forms double salts.
[37] The normal salts of calcium and magnesium are slightly, and the
acid salts easily, soluble in water. These acid sulphites are much
used in practice; thus calcium bisulphite is employed in the
manufacture of cellulose from sawdust, for mixing with fibrous
matter in the manufacture of paper.
Besides the acid character we must also point out the reducing character
of sulphurous anhydride. The reducing action of sulphurous acid, its
anhydride and salts, is due to their faculty of passing into sulphuric
acid and sulphates. The reducing action of the sulphites is particularly
energetic, so that they even convert nitric oxide into nitrous oxide:
K_{2}SO_{3} + 2NO = K_{2}SO_{4} + N_{2}O. The salts of many of the higher
oxides are converted into those of the lower--for example, FeX_{3} into
FeX_{2}, CuX_{2} into CuX, HgX_{2} into HgX; thus 2FeX_{3} + SO_{2} +
2H_{2}O = 2FeX_{2} + H_{2}SO_{4} + 2HX. In the presence of water,
sulphurous anhydride is oxidised by chlorine (SO_{2} + 2H_{2}O + Cl_{2} =
H_{2}SO_{4} + 2HCl), iodine, nitrous acid, hydrogen peroxide,
hypochlorous acid, chloric acid, and other oxygen compounds of the
halogens, chromic, manganic, and many other metallic acids and higher
oxides, as well as all peroxides. Free oxygen in the presence of spongy
platinum is able to oxidise sulphurous anhydride even in the absence of
water, in which case sulphuric anhydride SO_{3} is formed, so that the
latter may be prepared by passing a mixture of sulphurous anhydride and
oxygen over incandescent spongy platinum, or, as it is now prepared on a
large scale in chemical works, by passing this mixture over asbestos or
pumice stone moistened with a solution of platinum salt and ignited.
Sulphurous anhydride is completely absorbed by certain higher oxides--for
instance, by barium peroxide and lead dioxide (PbO_{2} + SO_{2} =
PbSO_{4}).[38]
[38] This reaction is taken advantage of in removing sulphurous
anhydride from a mixture of gases. Lead dioxide, PbO_{2}, is
brown, and when combined with sulphurous anhydride it forms lead
sulphate, PbSO_{4}, which is white, so that the reaction is
evident both from the change in colour and development of heat.
Sulphurous anhydride is slowly decomposed by the action of light,
with the separation of sulphur and formation of sulphuric
anhydride. This explains the fact that sulphurous anhydride
prepared in the dark gives a white precipitate of silver sulphite,
Ag_{2}SO_{3}, with silver chlorate, AgClO_{4}, but when prepared
in the light, even in diffused light, it gives a dark precipitate.
This naturally depends on the fact that the sulphur liberated then
forms silver sulphide, which is black.
There are, however, cases where sulphurous anhydride acts as an oxidising
agent--that is, it is _deoxidised_ in the presence of substances which
are capable of absorbing oxygen with still greater energy than the
sulphurous anhydride itself. This oxidising action proceeds with the
formation of sulphuretted hydrogen or of sulphides, while the reducing
agent is oxidised at the expense of the oxygen of the sulphurous
anhydride. In this respect, the action of stannous salts is particularly
remarkable. Stannous chloride, SnCl_{2}, in an aqueous solution gives a
precipitate of stannic sulphide, SnS_{2}, with sulphurous anhydride--that
is, the latter is deoxidised to sulphuretted hydrogen, while SnX_{2} is
oxidised into SnX_{4}. A solution of sulphurous anhydride has also an
oxidising action on zinc. The zinc passes into solution, but no hydrogen
is evolved,[39] because a salt of _hydrosulphurous acid_, ZnS_{2}O_{4},
is formed. The free acid is still less stable than the salt.
[39] Schönebein observed that the liquid turns yellow, and acquires the
faculty of decolorising litmus and indigo. Schützenberger showed
that this depends on the formation of a zinc salt of a peculiar
and very powerfully-reducing acid, for with cupric salts the
yellow solution gives a red precipitate of cuprous hydrate or
metallic copper, and it reduces salts of silver and mercury
entirely. An exactly similar solution is obtained by the action of
zinc on sodium bisulphite without access of air and in the cold.
The yellow liquid absorbs oxygen from the air with great avidity,
and forms a sulphate. If the solution be mixed with alcohol, it
deposits a double sulphite of zinc and sodium,
ZnNa_{2}(SO_{3})_{2}, which does not decolorise litmus or indigo.
The remaining alcoholic solution deposits colourless crystals in
the cold, which absorb oxygen with great energy in the presence of
water, but are tolerably stable when dried under the receiver of
an air-pump. The solution of these crystals has the
above-mentioned decolorising and reducing properties. These
crystals contain a sodium salt of a lower acid; their composition
was at first supposed to be HNaSO_{2}, but it was afterwards
proved that they do not contain hydrogen, and present the
composition Na_{2}S_{2}O_{4} (Bernthsen). The same salt is formed
by the action of a galvanic current on a solution of sodium
bisulphite, owing to the action of the hydrogen at the moment of
its liberation. If SO_{2} resembles CO_{2} in its composition,
then hyposulphurous acid H_{2}S_{2}O_{4} resembles oxalic acid
H_{2}C_{2}O_{4}. Perhaps an analogue of formic acid SH_{2}O_{2}
will be discovered.
The faculty of sulphurous anhydride of combining with various substances
is evident from the above-cited reactions, where it combines with
hydrogen and with oxygen, and this faculty also appears in the fact that,
like carbonic oxide, it combines with chlorine, forming a chloranhydride
of sulphuric acid, SO_{2}Cl_{2}, to which we shall afterwards return. The
same faculty for combination also appears in the salts of sulphurous
acid, in their liability to oxidation and in the exceedingly
characteristic formation of a peculiar series of salts obtained by
Pelouze and Frémy. At a temperature of -10° or below, nitric oxide NO is
absorbed by alkaline solutions of the alkali sulphites, forming a
peculiar series of _nitrosulphates_. At a higher temperature these salts
are not formed but the nitric oxide is reduced to nitrous oxide. But in
the cold the liquid saturated with nitric oxide after a certain time
gives prismatic crystals resembling those of nitre. The composition of
the potassium salt is K_{2}SN_{2}O_{3}--that is, the salt contains the
elements of potassium sulphite and of nitric oxide.[40]
[40] The instability of this salt is very great, and may be compared to
that of the compound of ferrous sulphate with nitric oxide, for
when heated under the contact influence of spongy platinum,
charcoal, &c., it splits up into potassium sulphate and nitrous
oxide. At 130° the dry salt gives off nitric oxide, and re-forms
potassium sulphite. The free acid has not yet been obtained. These
salts resemble the series of _sulphonitrites_ discovered by Frémy
in 1845. They are obtained by passing sulphurous anhydride through
a concentrated and strongly alkaline aqueous solution of potassium
nitrite. They are soluble in water, but are precipitated by an
excess of alkali. The first product of the action has the
composition K_{3}NS_{3}HO_{9}. It is then converted by the further
action of sulphurous anhydride, cold water, and other reagents
into a series of similar complex salts, many of which give
well-formed crystals. One must suppose that the chief cause of the
formation of these very complex compounds is that they contain
unsaturated compounds, NO, KNO_{2}, and KHSO_{3}, all of which are
subject to oxidation and further combination, and therefore easily
combine among each other. The decomposition of these compounds,
with the evolution of ammonia, when their solutions are heated is
due to the fact that the molecule contains the deoxidant,
sulphurous anhydride, which reduces the nitrous acid, NO(OH), to
ammonia. In my opinion the composition of the sulphonitrites may
be very simply referred to the composition of ammonia, in which
the hydrogen is partly replaced by the radicle of the sulphates.
If we represent the composition of potassium sulphate as
KO.KSO_{3}, the group KSO_{3} will be equivalent (according to the
law of substitution) to HO and to hydrogen. It combines with
hydrogen, forming the potassium acid sulphite, KHSO_{3}. Hence the
group KSO_{3} may also replace the hydrogen in ammonia. Judging by
my analysis (1870) the extreme limit of this substitution,
N(HSO_{3})_{3}, agrees with that of the sulphonitrite, which is
easily formed, simultaneously with alkali, by the action of
potassium sulphite on potassium nitrite, according to the equation
3K(KSO_{3}) + KNO_{2} + 2H_{2}O = N(KSO_{3})_{3} + 4HKO. The
researches of Berglund, and especially of Raschig (1887), fully
verified my conclusions, and showed that we must distinguish the
following types of salts, corresponding with ammonia, where X
stands for the sulphonic group, HSO_{3}, in which the hydrogen is
replaced by potassium; hence X = KSO_{3}: (1) NH_{2}X, (2)
NHX_{2}, (3) NH_{3}, (4) N(OH)XH, (5) N(OH)X_{2}, (6) N(OH)_{2}X,
just as NH_{2}(OH) is hydroxylamine, NH(OH)_{2}, is the hydrate of
nitrous oxide, and N(OH)_{3} is orthonitrous acid, as follows from
the law of substitution. This class of compounds is in most
intimate relation with the series of sulphonitrous compounds,
corresponding with 'chamber crystals' and their acids, which we
shall consider later.
There are also several other substances, formed by the oxides of
nitrogen and sulphur, which belong to this class of complex and, under
some circumstances, unstable compounds. In the manufacture of sulphuric
acid, both these classes of oxides come into contact with each other in
the lead chambers, and if there be insufficient water for the formation
of sulphuric acid they give crystalline compounds, termed _chamber
crystals_. As a rule, the composition of the crystals is expressed by the
formula NHSO_{3}. This is a compound of the radicles NO_{2} of nitric
acid, and HSO_{3} of sulphuric acid, or nitro-sulphuric acid,
NO_{2}.SHO_{3}, if sulphuric acid be expressed as OH.SHO_{3} and nitric
by NO_{2}.OH. The tabular crystals of this substance fuse at about 70°,
are formed both by the direct action of nitrous anhydride or nitric
peroxide (but not NO, which is not absorbed by sulphuric acid) on
sulphuric acid (Weltzien and others), and especially on sulphuric acid
containing an anhydride and the lower oxides of sulphur and nitric
acid.[41]
[41] In the sulphuric acid chambers the lower oxides of nitrogen and
sulphur take part in the reaction. They are oxidised by the oxygen
of the air, and form nitro-sulphuric acid--for example, 2SO_{2} +
N_{2}O_{3} + O_{2} + H_{2}O = 2NHSO_{5}. This compound dissolves
in strong sulphuric acid without changing, and when this solution
is diluted (when the sp. gr. falls to 1·5), it splits up into
sulphuric acid and nitrous anhydride, and by the action of
sulphurous anhydride is converted into nitric oxide, which by
itself (in the absence of nitric acid or oxygen) is insoluble in
sulphuric acid. These reactions are taken advantage of in
retaining the oxides of nitrogen in the Gay-Lussac coke-towers,
and for extracting the absorbed oxides of nitrogen from the
resultant solution in the Glover tower. Although nitric oxide is
not absorbed by sulphuric acid, it reacts (Rose, Brüning) on its
anhydride, and forms sulphurous anhydride and a crystalline
substance, N_{2}S_{2}O_{9} = 2NO + 3SO_{3} - SO_{2} =
N_{2}O_{3}2SO_{3}. This may be regarded as the anhydride of
nitro-sulphuric acid, because N_{2}S_{2}O_{9} = 2NHSO_{5} -
H_{2}O; like nitro-sulphuric acid, it is decomposed by water into
nitro-sulphuric acid and nitrous anhydride. Since boric and
arsenious anhydrides, alumina and other oxides of the form
R_{2}O_{3} are able to combine with sulphuric anhydride to form
similar compounds decomposable by water, the above compound does
not present any exceptional phenomenon. The substance NOClSO_{3}
obtained by Weber by the action of nitrosyl chloride upon
sulphuric anhydride belongs to this class of compounds.
_Thiosulphuric acid_, H_{2}S_{2}O_{3}--that is, a compound of sulphurous
acid and sulphur--also belongs to the products of combination of
sulphurous acid. In the same way that sulphurous acid, H_{2}SO_{3}, gives
H_{2}SO_{4} with oxygen, so it gives H_{2}S_{2}O_{3} with sulphur. In a
free state it is very unstable, and it is only known in the form of its
salts proceeding from the direct action of sulphur on the normal
sulphites; if endeavours be made to separate it in a free state, it
immediately splits up into those elements from which it might be
formed--that is, into sulphur and sulphurous acid. The most important of
its salts is the _sodium thiosulphate_ (known as hyposulphite),
Na_{2}S_{2}O_{3},5H_{2}O, which occurs in colourless crystals, and is
unacted on by atmospheric oxygen either when in a dry state or in
solution. Many other salts of this acid are easily formed by means of
this salt,[41 bis] although this cannot be done with all bases, for such
bases as alumina, ferric oxide, chromium oxide, and others do not give
compounds with thiosulphuric acid, just as they do not form stable
compounds with carbonic acid. Whenever these salts might be formed, they
(like the acid) split up into sulphurous acid and sulphur, and
furthermore the elements of thiosulphuric acid in many cases act in a
reducing manner, forming sulphuric acid and taking up the oxygen from
reducible oxides. Thus when treated with a thiosulphate the soluble
ferric salts give a precipitate of sulphur and form ferrous salts. The
thiosulphates of the metals of the alkalis are obtained directly by
boiling a solution of their sulphites with sulphur: Na_{2}SO_{3} + S =
Na_{2}S_{2}O_{3}. The same salts are formed by the action of sulphurous
anhydride on solutions of the sulphides; thus sodium sulphide dissolved
in water gives sulphur and sodium thiosulphate when a stream of
sulphurous anhydride is passed through it: 2Na_{2}S + 3SO_{2} =
2Na_{2}S_{2}O_{3} + S. The polysulphides of the alkali metals when left
exposed to the air attract oxygen and also form thiosulphates.[42]
[41 bis] Many double salts of thiosulphuric acid are known, for
instance, PbS_{2}O_{3},3Na_{2}S_{2}O_{3},12H_{2}O;
CaS_{2}O_{3},3K_{2}S_{2}O_{3},5H_{2}O, &c. (Fortman, Schwicker,
Fock, and others).
[42] Thus when alkali waste, which contains calcium sulphide, undergoes
oxidation in the air it first forms a calcium polysulphide, and
then calcium thiosulphate, CaS_{2}O_{3}. When iron or zinc acts on
a solution of sulphurous acid, besides the hyposulphurous acid
first formed, a mixture of sulphite and thiosulphate is obtained
(Note 39), 3SO_{2} + Zn_{2} = ZnSO_{3} + ZnS_{2}O_{3}. In this
case, as in the formation of hyposulphurous acid, there is no
hydrogen liberated. One of the most common methods for preparing
thiosulphates consists in the _action of sulphur on the alkalis_.
The reaction is accomplished by the formation of sulphides and
thiosulphates, just as the reaction of chlorine on alkalis is
accompanied by the formation of hypochlorites and chlorides; hence
in this respect the thiosulphates hold the same position in the
order of the compounds of sulphur as the hypochlorites do among
the chlorine compounds. The reaction of caustic soda on an excess
of sulphur may be expressed thus: 6NaHO + 12S = 2Na_{2}S_{5} +
Na_{2}S_{2}O_{3} + 3H_{2}O. Thus sulphur is soluble in alkalis. On
a large scale sodium thiosulphate, Na_{2}S_{2}O_{3}, is prepared
by first heating sodium sulphate with charcoal, to form sodium
sulphide, which is then dissolved in water and treated with
sulphurous anhydride. The reaction is complete when the solution
has become slightly acid. A certain amount of caustic alkali is
added to the slightly acid solution; a portion of the sulphur is
thus precipitated, and the solution is then boiled and evaporated
when the salt crystallises out. The saturation of the solution of
sodium sulphide by sulphurous anhydride is carried on in different
ways--for example, by means of coke-towers, by causing the
solution of sulphide to trickle over the coke, and the sulphurous
anhydride, obtained by burning sulphur, to pass up the coke-tower
from below. An excess of sulphurous anhydride must be avoided, as
otherwise sodium trithionate is formed. Sodium thiophosphate is
also prepared by the double decomposition of the soluble calcium
thiosulphate with sodium sulphate or carbonate, in which case
calcium sulphate or carbonate is precipitated. The calcium
thiosulphate is prepared by the action of sulphurous anhydride on
either calcium sulphide or alkali waste. A dilute solution of
calcium thiosulphate may be obtained by treating alkali waste
which has been exposed to the action of air with water. On
evaporation, this solution gives crystals of the salt containing
CaS_{2}O_{3},5H_{2}O. A solution of calcium thiosulphate must be
evaporated with great care, because otherwise the salt breaks up
into sulphur and calcium sulphide. Even the crystallised salt
sometimes undergoes this change.
The crystals of sodium thiosulphate are stable, do not effloresce
and at 0° dissolve in one part of water, and at 20° in 0·6 part.
The solution of this salt does not undergo any change when boiled
for a short time, but after prolonged boiling it deposits sulphur.
The crystals fuse at 56°, and lose all their water at 100°. When
the dry salt is ignited it gives sodium sulphide and sulphate.
With acids, a solution of the thiosulphate soon becomes cloudy and
deposits an exceedingly fine powder of sulphur (Note 10). If the
amount of acid added be considerable, it also evolves sulphurous
anhydride: H_{2}S_{2}O_{3} = H_{2}O + S + SO_{2}. Sodium
thiosulphate has many practical uses; it is used in photography
for dissolving silver chloride and bromide. Its solvent action on
silver chloride may be taken advantage of in extracting this metal
as chloride from its ores. In dissolving, it forms a double salt
of silver and sodium: AgCl + Na_{2}S_{2}O_{3} = NaCl +
AgNaS_{2}O_{3}. Sodium thiosulphate is an _antichlor_--that is, a
substance which hinders the destructive action of free chlorine
owing to its being very easily oxidised by chlorine into sulphuric
acid and sodium chloride. The reaction with iodine is different,
and is remarkable for the accuracy with which it proceeds. The
iodine takes up half the sodium from the salt and converts it into
a tetrathionate; 2Na_{2}S_{2}O_{3} + I_{2} = 2NaI +
Na_{2}S_{4}O_{6}, and hence this reaction is employed for the
determination of free iodine. As iodine is expelled from potassium
iodide by chlorine, it is possible also to determine the amount of
chlorine by this method if potassium iodide be added to a solution
containing chlorine. And as many of the higher oxides are able to
evolve iodine from potassium iodide, or chlorine from hydrochloric
acid (for example, the higher oxides of manganese, chromium, &c.),
it is also possible to determine the amounts of these higher
oxides by means of sodium thiosulphate and liberated iodine. This
forms the basis of the iodometric method of volumetric analysis.
The details of these methods will be found in works on analytical
chemistry.
On adding a solution of a _lead salt_ gradually to a solution of
sodium thiosulphate a white precipitate of lead thiosulphate,
PbS_{2}O_{3}, is formed (a soluble double salt is first formed,
and if the action be rapid, lead sulphide). When this substance is
heated at 200°, it undergoes a change and takes fire. Sodium
thiosulphate in solution rapidly reduces cupric salts to cuprous
salts by means of the sulphurous acid contained in the
thiosulphate, but the resultant cuprous oxide is not precipitated,
because it passes into the state of a thiosulphate and forms a
double salt. These double cuprous salts are excellent reducing
agents. The solution when heated gives a black precipitate of
copper sulphide.
The following formulæ sufficiently explain the position held by
thiosulphuric acid among the other acids of sulphur:
Sulphurous acid SO_{2}H(OH)
Sulphuric acid SO_{2}OH(OH)
Thiosulphuric acid SO_{2}SH(OH)
Hyposulphurous acid SO_{2}H(SO_{2}H)
Dithionic acid SO_{2}OH(SO_{2}OH)
At one time it was thought that all the salts of thiosulphuric
acid only existed in combination with water, and it was then
supposed that their composition was H_{4}S_{2}O_{4}, or
H_{2}SO_{2}, but Popp obtained the anhydrous salts.
Although sulphur, oxidising at a high temperature, only forms a small
quantity of sulphuric anhydride, SO_{3}, and nearly all passes into
sulphurous anhydride, still the latter may be converted into the higher
oxide, or _sulphuric anhydride_, SO_{3}, by many methods. Sulphuric
anhydride is a solid crystalline substance at the ordinary temperature;
it is easily fusible (15°), and volatile (46°), and rapidly attracts
moisture. Although it is formed by the combination of sulphurous
anhydride with oxygen, it is capable of further combination. Thus it
combines with water, hydrochloric acid, ammonia, with many hydrocarbons,
and even with sulphuric acid, boric and nitrous anhydrides, &c., and also
with bases which burn directly in its vapour, forming sulphates in the
presence of traces of moisture (_see_ Chapter IX., Note 29). The
oxidation of sulphurous anhydride, SO_{2}, into sulphuric anhydride,
SO_{3}, is effected by passing a mixture of the former and dry oxygen or
air over incandescent spongy platinum. An increase of pressure
accelerates the reaction (Hanisch). If the product be passed into a cold
vessel, crystalline sulphuric anhydride is deposited upon the sides of
the vessel, but as it is difficult to avoid all traces of moisture it
always contains compounds of its hydrates: H_{2}S_{2}O_{7} and
H_{2}S_{4}O_{13}, whose presence so modifies the properties of the
anhydride (Weber) that formerly two modifications of the anhydride were
recognised. The same sulphuric anhydride may be obtained from certain
anhydrous sulphates, or those which are almost so, which are decomposed
by heat, whilst an impure but perfectly anhydrous anhydride is formed by
distillation over phosphoric anhydride. For instance, acid sodium
sulphate, NaHSO_{4}, and the pyro- or di-sulphate, Na_{2}S_{2}O_{7}
(Chapter XII.) formed from it, when ignited evolve sulphuric anhydride.
Green vitriol--that is, ferrous sulphate, FeSO_{4}--belongs to the number
of those sulphates which easily give off sulphuric anhydride under the
action of heat. It contains water of crystallisation and parts with it
when it is heated, but the last equivalent of water is driven off with
difficulty, just as is the case with magnesium sulphate, MgSO_{4}7H_{2}O;
however, when thoroughly heated, this evolution of sulphuric anhydride
does take place, although not completely, because at a high temperature a
portion of it is decomposed by the ferrous oxide (SO_{3} + 2FeO), which
is converted into ferric oxide, Fe_{2}O_{3}, and in consequence part of
the sulphuric anhydride is converted into sulphurous anhydride. Thus the
products of the decomposition of ferrous sulphate will be: ferric oxide,
Fe_{2}O_{3}, sulphurous anhydride, SO_{2}, and sulphuric anhydride,
SO_{3}, according to the equation: 2FeSO_{4} = Fe_{2}O_{3} + SO_{2} +
SO_{3}. As water still remains with the ferrous sulphate when it is
heated, the result will partially consist of the hydrate H_{2}SO_{4},
with anhydride, SO_{3}, dissolved in it. Sulphuric acid was for a long
time prepared in this manner; the process was formerly carried on on a
large scale in the neighbourhood of Nordhausen, and hence the sulphuric
acid prepared from ferrous sulphate is called _fuming Nordhausen acid_.
At the present time the fuming acid is prepared by passing the volatile
products of the decomposition of ferrous sulphate through strong
sulphuric acid prepared by the ordinary method. The sulphurous anhydride
is insoluble in it, but it absorbs the sulphuric anhydride. Sulphuric
anhydride may be prepared not only by igniting FeSO_{4} or sodium
pyrosulphate, Na_{2}S_{2}O_{7} (the decomposition proceeds at 600°), but
also by heating a mixture of the latter and MgSO_{4} (Walters); in the
former case a stable double salt MgNa_{2}(SO_{4})_{2} finally remains. It
is also obtained by the direct combination of SO_{2} and O under the
action of spongy platinum or asbestos coated with platinum black (C.
Winkler's process). Nordhausen sulphuric acid fumes in air, owing to its
containing and easily giving off sulphuric anhydride, and it is therefore
also called _fuming sulphuric acid_; these fumes are nothing but the
vapour of sulphuric anhydride combining with the moisture in the air and
forming non-volatile sulphuric acid (hydrate).[43]
[43] Nordhausen sulphuric acid may serve as a very simple means for the
preparation of sulphuric anhydride. For this purpose the
Nordhausen acid is heated in a glass retort, whose neck is firmly
fixed in the mouth of a well-cooled flask. The access of moisture
is prevented by connecting the receiver with a drying-tube. On
heating the retort the vapours of sulphuric anhydride will pass
over into the receiver, where they condense; the crystals of
anhydride thus prepared will, however, contain traces of sulphuric
acid--that is, of the hydrate. By repeatedly distilling over
phosphoric anhydride, it is possible to obtain the pure anhydride,
SO_{3}, especially if the process be carried on without access of
air in a closed vessel.
The ordinary sulphuric anhydride, which is imperfectly freed from
the hydrate, is a snow-white, exceedingly volatile substance,
which crystallises (generally by sublimation) in long silky
prisms, and only gives the pure anhydride when carefully distilled
over P_{2}O_{5}. Freshly prepared crystals of almost pure
anhydride fuse at 16° into a colourless liquid having a specific
gravity at 26° = 1·91, and at 47° = 1·81; it volatilises at 46°.
After being kept for some time the anhydride, even containing only
small traces of water, undergoes a change of the following nature:
A small quantity of sulphuric acid combines by degrees with a
large proportion of the anhydride, forming polysulphuric acids,
H_{2}SO_{4},_n_SO_{3}, which fuse with difficulty (even at 100°,
Marignac), but decompose when heated. In the entire absence of
water this rise in the fusing point does not occur (Weber), and
then the anhydride long remains liquid, and solidifies at about
+15°, volatilises at 40°, and has a specific gravity 1·94 at 16°.
We may add that Weber (1881), by treating sulphuric anhydride with
sulphur, obtained a blue lower oxide of sulphur, S_{2}O_{3}.
Selenium and tellurium also give similar products with SO_{3},
SeSO_{3}, and TeSO_{3}. Water does not act upon them.
Nordhausen sulphuric acid contains a peculiar compound of SO_{3} and
H_{2}SO_{4}, or _pyrosulphuric acid_; an imperfect anhydride of sulphuric
acid, H_{2}S_{2}O_{7}, analogous in composition with the salts
Na_{2}S_{2}O_{7}, K_{2}Cr_{2}O_{7}, and bearing the same relation to
H_{2}SO_{4} that pyrophosphoric acid does to H_{3}PO_{4}. The bond
holding the sulphuric acid and anhydride together is unstable. This is
obvious from the fact that the anhydride may easily be separated from
this compound, by the action of heat. In order to obtain the definite
compound, the Nordhausen acid is cooled to 5°, or, better still, a
portion of it is distilled until all the anhydride and a certain amount
of sulphuric acid have passed over into the distillate, which will then
solidify at the ordinary temperature, because the compound
H_{2}SO_{4},SO_{3} fuses at 35°. Although this substance reacts on water,
bases, &c., like a mixture of SO_{3} + H_{2}SO_{4}, still since a
definite compound, H_{2}S_{2}O_{7}, exists in a free state and gives
salts and a chloranhydride, S_{2}O_{5}Cl_{2},[44] we must admit the
existence of a definite pyrosulphuric acid, like pyrophosphoric acid,
only that the latter has a far greater stability and is not even
converted into a perfect hydrate by water. Further, the salts
M_{2}S_{2}O_{7} dissolved in water react in the same manner as the acid
salts MHSO_{4}, whilst the imperfect hydrates of phosphoric acid (for
example, PHO_{3}, H_{4}P_{2}O_{7}) have independent reactions even in an
aqueous solution which distinguish them and their salts from the perfect
hydrates.
[44] Pyrosulphuric chloranhydride, or _pyrosulphuryl chloride_,
S_{2}O_{5}Cl_{2}, corresponds to pyrosulphuric acid, in the same
way that sulphuryl chloride, SO_{2}Cl_{2}, corresponds to
sulphuric acid. The composition S_{2}O_{5}Cl_{2} = SO_{2}Cl_{2} +
SO_{3}. It is obtained by the action of the vapour of sulphuric
anhydride on sulphur chloride: S_{2}Cl_{2} + 5SO_{3} = 5SO_{2} +
S_{2}O_{5}Cl_{2}. It is also formed (and not sulphuryl chloride,
SO_{2}Cl_{2}, Michaelis) by the action of phosphorus pentachloride
in excess on sulphuric acid (or its first chloranhydride,
SHO_{3}Cl). It is an oily liquid, boiling at about 150°, and of
sp. gr. 1·8. According to Konovaloff (Chapter VII.), its vapour
density is normal. It should be noticed that the same substance is
obtained by the action of sulphuric anhydride on sulphur
tetrachloride, and also on carbon tetrachloride, and this
substance is the last product of the metalepsis of CH_{4}, and
therefore the comparison of SCl_{2} and S_{2}Cl_{2} with products
of metalepsis (_see_ later) also finds confirmation in particular
reactions. Rose, who obtained pyrosulphuryl chloride,
S_{2}O_{5}Cl_{2}, regarded it as SCl_{6},5SO_{3}, for at that time
an endeavour was always made to find two component parts of
opposite polarity, and this substance was cited as a proof of the
existence of a hexachloride, SCl_{6}. Pyrosulphuryl chloride is
decomposed by cold water, but more slowly than chlorosulphuric
acid and the other chloranhydrides.
The relation between pyrosulphuric acid and the normal acid will
be obvious if we express the latter by the formula OH(SO_{3}H),
because the sulphonic group (SO_{3}H) is then evidently equivalent
to OH, and consequently to H, and if we replace both the hydrogens
in water by this radicle we shall obtain (SO_{3}H)_{2}O--that is,
pyrosulphuric acid.
[Illustration: FIG. 87.--Concentration of sulphuric acid in glass
retorts. The neck of each retort is attached to a bent glass tube, whose
vertical arm is lowered into a glass or earthenware vessel acting as a
receiver for the steam which comes over from the acid, as the former
still contains a certain amount of acid.]
_Sulphuric acid_, H_{2}SO_{4}, is formed by the combination of its
anhydride, SO_{3}, and water, with the evolution of a large amount of
heat; the reaction SO_{3} + H_{2}O develops 21,300 heat units. The method
of its preparation on a large scale, and most of the methods employed for
its formation, are dependent on the oxidation of sulphurous anhydride,
and the formation of sulphuric anhydride, which forms sulphuric acid
under the action of water. The technical method of its manufacture has
been described in Chapter VI. The acid obtained _from the lead chambers_
contains a considerable amount of water, and is also impure owing to the
presence of oxides of nitrogen, lead compounds, and certain impurities
from the burnt sulphur which have come over in a gaseous and vaporous
state (for example, arsenic compounds). For practical purposes, hardly
any notice is taken of the majority of these impurities, because they do
not interfere with its general qualities. Most frequently endeavours are
only made to remove, as far as possible, all the water which can be
expelled.[45] That is, the object is to obtain the hydrate, H_{2}SO_{4},
from the dilute acid (60 per cent.), and this is effected by evaporation
by means of heat. Every given mixture of water and sulphuric acid begins
to part with a certain amount of aqueous vapour when heated to a certain
definite temperature. At a low temperature either there is no evaporation
of water, or there can even be an absorption of moisture from the air. As
the removal of the water proceeds, the vapour tension of the residue
decreases for the same temperature, and therefore the more dilute the
acid the lower the temperature at which it gives up a portion of its
water. In consequence of this, the removal of water from dilute solutions
of sulphuric acid may be easily carried on (up to 75 p.c. H_{2}SO_{4}) in
lead vessels, because at low temperatures dilute sulphuric acid does not
attack lead. But as the acid becomes more concentrated the temperature at
which the water comes over becomes higher and higher, and then the acid
begins to act on lead (with the evolution of sulphuretted hydrogen and
conversion of the lead into sulphate), and therefore lead vessels cannot
be employed for the complete removal of the water. For this purpose the
evaporation is generally carried on in glass or platinum retorts, like
those depicted in figs. 87 and 88.
[45] The removal of the water, or concentration to almost the real
acid, H_{2}SO_{4}, is effected for two reasons: in the first place
to avoid the expense of transit (it is cheaper to remove the water
than to pay for its transit), and in the second place because many
processes--for instance, the refining of petroleum--require a
strong acid free from an excess of water, the weak acid having no
action. When in the manufacture of chamber acid, both the
Gay-Lussac tower (cold, situated at the end of the chambers) and
the Glover tower (hot, situated at the beginning of the plant,
between the chambers and ovens for the production of SO_{2}) are
employed, a mixture of nitrose (_i.e._ the product of the
Gay-Lussac tower) and chamber acid containing about 60 p.c.
H_{2}SO_{4}, is poured into the Glover tower, where under the
action of the hot furnace gases containing SO_{2}, and the water
held in the chamber acid (1) N_{2}O_{3} is evolved from the
nitrose; (2) water is expelled from the chamber acid; (3) a
portion of the SO_{2} is converted into H_{2}SO_{4}; and (4) the
furnace gases are cooled. Thus, amongst other things, the Glover
tower facilitates the concentration of the chamber acid (removal
of H_{2}O), but the product generally contains many impurities.
[Illustration: FIG. 88.--Concentration of sulphuric acid in platinum
retorts.]
_The concentration of sulphuric acid_ in glass retorts is not a
continuous process, and consists of heating the dilute 75 per cent. acid
until it ceases to give off aqueous vapour, and until acid containing
93-98 per cent. H_{2}SO_{4} (66° Baumé) is obtained--and this takes place
when the temperature reaches 320° and the density of the residue reaches
1·847 (66° Baumé).[46] The platinum vessels designed for the continuous
concentration of sulphuric acid consist of a still _B_, furnished with a
still head _E_, a connecting pipe _E F_, and a syphon tube _H R_, which
draws off the sulphuric acid concentrated in the boiler. A stream of
sulphuric acid previously concentrated in lead retorts to a density of
about 60° Baumé--_i.e._ to 75 per cent. or a sp. gr. of 1·7--runs
continuously into the retort through a syphon funnel _E´_. The apparatus
is fed from above, because the acid freshly supplied is lighter than that
which has already lost water, and also because the water is more easily
evaporated from the freshly supplied acid at the surface. The platinum
retort is heated, and the steam coming off[47] is condensed in a worm _F
G_, whilst as fresh dilute acid is supplied to the boiler the acid
already concentrated is drawn off through the syphon tube _H B_, which is
furnished with a regulating cock by means of which the outflow of the
concentrated acid from the bottom of the retort can be so regulated that
it will always present one and the same specific gravity, corresponding
with the strength required. For this purpose the acid flowing from the
syphon is collected in a receiver _R_, in which a hydrometer, indicating
its density, floats; if its density be less than 66° Baumé, the
regulating cock is closed sufficiently to retard the outflow of sulphuric
acid, so as to lengthen the time of its evaporation in the retort.[48]
[46] The difficulty with which the last portions of water are removed
is seen from the fact that the boiling becomes very irregular,
totally ceasing at one moment, then suddenly starting again, with
the rapid formation of a considerable amount of steam, and at the
same time bumping and even overturning the vessel in which it is
held. Hence it is not a rare occurrence for the glass retorts to
break during the distillation; this causes platinum retorts to be
preferred, as the boiling then proceeds quite uniformly.
[47] According to Regnault, the vapour tensions (in millimetres of
mercury) of the water given off by the hydrates of sulphuric acid,
H_{2}SO_{4},_n_H_{2}O, are--
_t_=5° 15° 30°
_n_ = 1 0·1 0·1 0·2
2 0·4 0·7 1·5
3 0·9 1·6 4·1
4 1·3 2·8 7·0
5 2·1 4·2 10·7
7 3·2 6·2 15·6
9 4·1 8·0 19·6
11 4·4 9·0 22·2
17 5·5 10·6 26·1
According to Lunge, the vapour tension of the aqueous vapour given
off from solutions of sulphuric acid containing _p_ per cent.
H_{2}SO_{4}, at _t_°, equals the barometric pressure 720 to 730
mm.
_p_= 10 20 30 40 50 60 70 80 85 90 95
_t_= 102° 105° 108° 114° 124° 141° 170° 207° 233° 262° 295°
The latter figures give the temperature at which water is easily
expelled from solutions of sulphuric acid of different strengths.
But the evaporation begins sooner, and concentration may be
carried on at lower temperatures if a stream of air be passed
through the acid. Kessler's process is based upon this (Note 48).
[48] The greatest part of the sulphuric acid is used in the soda
manufacture, in the conversion of the common salt into sulphate.
For this purpose an acid having a density of 60° Baumé is amply
sufficient. Chamber acid has a density up to 1·57 = 50° to 51°
Baumé; it contains about 35 per cent. of water. About 15 per cent.
of this water can be removed in leaden stills, and nearly all the
remainder may be expelled in glass or platinum vessels. Acid of
66° Baumé, = 1·847, contains about 96 per cent. of the hydrate
H_{2}SO_{4}. The density falls with a greater or less proportion
of water, the maximum density corresponding with 97-1/2 per cent.
of the hydrate H_{2}SO_{4}. The concentration of H_{2}SO_{4} in
platinum retorts has the disadvantage that sulphuric acid, upwards
of 90 per cent. in strength, does corrode platinum, although but
slightly (a few grams per tens of tons of acid). The retorts
therefore require repairing, and the cost of the platinum exceeds
the price obtained for concentrating the acid from 90 per cent. to
98 per cent. (in factories the acid is not concentrated beyond
this by evaporation in the air). This inconvenience has lately
(1891, by Mathey) been eliminated by coating the inside of the
platinum retorts with a thin (0·1 to 0·02 mm.) layer of gold which
is 40 times less corroded by sulphuric acid than platinum. Négrier
(1890) carries on the distillation in porcelain dishes, Blond by
heating a thin platinum wire immersed in the acid by means of an
electric current, but the most promising method is that of Kessler
(1891), which consists in passing hot air over sulphuric acid
flowing in a thin stream in stone vessels, so that there is no
boiling but only evaporation at moderate temperatures: the
transference of the heat is direct (and not through the sides of
the vessels), which economises the fuel and prevents the
distilling vessels being damaged.
When, by evaporation of the water, sulphuric acid attains a
density of 66° Baumé (sp. gr. 1·84), it is impossible to
concentrate it further, because it then distils over unchanged.
_The distillation of sulphuric acid_ is not generally carried on
on a large scale, but forms a laboratory process, employed when
particularly pure acid is required. The distillation is effected
either in platinum retorts furnished with corresponding condensers
and receivers, or in glass retorts. In the latter case, great
caution is necessary, because the boiling of sulphuric acid itself
is accompanied by still more violent jerks and greater
irregularity than even the evaporation of the last portions of
water contained in the acid. If the glass retort which holds the
strong sulphuric acid to be distilled be heated directly from
below, it frequently jerks and breaks. For greater safety the
heating is not effected from below, but at the sides of the
retort. The evaporation then does not proceed in the whole mass,
but only from the upper portions of the liquid, and therefore goes
on much more quietly. The acid may be made to boil quietly also by
surrounding the retort with good conductors of heat--for example,
iron filings, or by immersing a bunch of platinum wires in the
acid, as the bubbles of sulphuric acid vapour then form on the
extremities of the wires.
Strictly speaking, _sulphuric acid is not volatile_, and at its
so-called boiling-point it really decomposes into its anhydride and
water; its boiling-point (338°) being nothing else but its temperature of
decomposition. The products of this decomposition are substances boiling
much below the temperature of the decomposition of sulphuric acid. This
conclusion with regard to the process of the distillation of sulphuric
acid may be deduced from Bineau's observations on the vapour-density of
sulphuric acid. This density referred to hydrogen proved to be half that
which sulphuric acid should have according to its molecular weight,
H_{2}SO_{4}, in which case it should be 49, whilst the observed density
was equal to 24·5. Besides which, Marignac showed that the first portions
of the sulphuric acid distilling over contain less of the elements of
water than the portion which remains behind, or which distils over
towards the end. This is explained by the fact that on distillation the
sulphuric acid is decomposed, but a portion of the water proceeding from
its decomposition is retained by the remaining mass of sulphuric acid,
and therefore at first a mixture of sulphuric acid and sulphuric
anhydride--_i.e._ fuming sulphuric acid--is obtained in the distillate.
It is possible by repeating the distillation several times and only
collecting the first portions of the distillate, to obtain a distinctly
fuming acid. To obtain the definite hydrate H_{2}SO_{4} it is necessary
to refrigerate a highly concentrated acid, of as great a purity as
possible, to which a small quantity of sulphuric anhydride has been
previously added. Sulphuric acid containing a small quantity (a fraction
of a per cent. by weight) of water only freezes at a very low
temperature, while the pure normal acid, H_{2}SO_{4}, solidifies when it
is cooled below 0°, and therefore the normal acid first crystallises out
from the concentrated sulphuric acid. By repeating the refrigeration
several times, and pouring off the unsolidified portion, it is possible
to obtain a pure _normal hydrate_, H_{2}SO_{4}, which melts at 10°·4.
Even at 40° it gives off distinct fumes--that is, it begins to evolve
sulphuric anhydride, which volatilises, and therefore even in a dry
atmosphere the hydrate H_{2}SO_{4} becomes weaker, until it contains
1-1/2 p.c. of water.[49]
[49] Thus it appears that so common, and apparently so stable, a
compound as sulphuric acid decomposes even at a low temperature
with separation of the anhydride, but this decomposition is
restricted by a limit, corresponding to the presence of about
1-1/2 p.c. of water, or to a composition of nearly
H_{2}O,12H_{2}SO_{4}.
Now there is no reason for thinking that this substance is a
definite compound; it is an equilibrated system which does not
decompose under ordinary circumstances below 338°. Dittmar carried
on the distillation under pressures varying between 30 and 2,140
millimetres (of mercury), and he found that the composition of the
residue hardly varies, and contains from 99·2 to 98·2 per cent. of
the normal hydrate, although at 30 mm. the temperature of
distillation is about 210° and at 2,140 mm. it is 382°.
Furthermore, it is a fact of practical importance that under a
pressure of two atmospheres the distillation of sulphuric acid
proceeds very quietly.
Sulphuric acid may be _purified_ from the majority of its
impurities by distillation, if the first and last portions of the
distillate be rejected. The first portions will contain the oxides
of nitrogen, hydrochloric acid, &c., and the last portions the
less volatile impurities. The oxides of nitrogen may be removed by
heating the acid with charcoal, which converts them into volatile
gases. Sulphuric acid may be freed from arsenic by heating it with
manganese dioxide and then distilling. This oxidises all the
arsenic into non-volatile arsenic acid. Without a preliminary
oxidation it would partially remain as volatile arsenious acid,
and might pass over into the distillate. The arsenic may also be
driven off by first reducing it to arsenious acid, and then
passing hydrochloric acid gas through the heated acid. It is then
converted into arsenious chloride, which volatilises.
In a concentrated form sulphuric acid is commercially known as _oil of
vitriol_, because for a long time it was obtained from green vitriol and
because it has an oily appearance and flows from one vessel into another
in a thick and somewhat sluggish stream, like the majority of oily
substances, and in this clearly differs from such liquids as water,
spirit, ether, and the like, which exhibit a far greater mobility. Among
its reactions the first to be remarked is its faculty for the formation
of many compounds. We already know that it combines with its anhydride,
and with the sulphates of the alkali metals; that it is soluble in water,
with which it forms more or less stable compounds. Sulphuric acid, when
mixed with water, develops a very considerable amount of heat.[50]
[50] The amount of heat developed by the mixture of sulphuric acid with
water is expressed in the diagram on p. 77, Volume I., by the
middle curve, whose abscissæ are the percentage amounts of acid
(H_{2}SO_{4}) in the resultant solution, and ordinates the number
of units of heat corresponding with the formation of 100 cubic
centimetres of the solution (at 18°). The calculations on which
the curve is designed are based on Thomsen's determinations, which
show that 98 grams or a molecular amount of sulphuric acid, in
combining with _m_ molecules of water (that is, with _m_=18 grams
of water), develop the following number of units of heat, R:--
_m_ = 1 2 3 5 9
R = 6379 9418 11137 13108 14952
_c_ = 0·432 0·470 0·500 0·576 0·701
T = 127° 149° 146° 121° 82°
_m_ = 19 49 100 200
R = 16256 16684 16859 17066
_c_ = 0·821 0·914 0·954 0·975
T = 145° 19° 9° 5°
_c_ stands for the specific heat of H_{2}SO_{4}_m_H_{2}O
(according to Marignac and Pfaundler), and T for the rise in
temperature which proceeds from the mixture of H_{2}SO_{4} with
_m_H_{2}O. The diagram shows that contraction and rise of
temperature proceed almost parallel with each other.
Besides the normal hydrate H_{2}SO_{4}, _another definite
hydrate_, H_{2}SO_{4},H_{2}O (84·48 per cent. of the normal
hydrate, and 15·52 per cent. of water) is known; it
crystallises[50 bis] extremely easily in large six-sided prisms,
which form above 0°--namely, at about +8°·5; when heated to 210°
it loses water.[51] If the hydrates H_{2}SO_{4} and
H_{2}SO_{4},H_{2}O exist at low temperatures as definite
crystalline compounds, and if pyrosulphuric acid,
H_{2}SO_{4}SO_{3}, has the same property, and if they all
decompose with more or less ease on a rise of temperature, with
the disengagement of either SO_{3} or H_{2}O, and in their
ordinary form present all the properties of simple solutions, it
follows that between sulphuric anhydride, SO_{3}, and water,
H_{2}O, there exists a consecutive series of homogeneous liquids
or solutions, among which we must distinguish _definite
compounds_, and therefore it is quite justifiable to look for
other definite compounds between SO_{3} and H_{2}O, beyond the
conditions for a change of state. In this respect we may be guided
by the variation of properties of any kind, proceeding
concurrently with a variation in the composition of a solution.
[50 bis] Pickering (1890) showed (_a_) that dilute solutions of
sulphuric acid containing up to H_{2}SO_{4} + 10H_{2}O deposit ice
(at -0°·12 when there is 2,000H_{2}O per H_{2}SO_{4}, at -0°·23
when there is 1,000H_{2}O, at -1°·04 when there is 200H_{2}O, at
-2°·12 when there is 100H_{2}O, at -4°·5 when there is 50H_{2}O,
at -15°·7 when there is 20H_{2}O, and at -61° when the composition
of the solution is H_{2}SO_{4} + 10H_{2}O); (_b_) that for higher
concentrations crystals separate out at a considerable degree of
cold, having the composition H_{2}SO_{4}4H_{2}O, which melt at
-24°·5, and if either water or H_{2}SO_{4} be added to this
compound the temperature of crystallisation falls, so that a
solution of the composition 12H_{2}SO_{4} + 100H_{2}O gives
crystals of the above hydrate at -70°, 15H_{2}SO_{4} + 100H_{2}O
at -47°, 30H_{2}SO_{4} + 100H_{2}O at -32°, 40H_{2}SO_{4} +
100H_{2}O at -52°; (_c_) that if the amount of H_{2}SO_{4} be
still greater, then a hydrate H_{2}SO_{4}H_{2}O separates out and
melts at +8°·5, while the addition of water or sulphuric acid to
it lowers the temperature of crystallisation so that the
crystallisation of H_{2}SO_{4}H_{2}O from a solution of the
composition H_{2}SO_{4} + 1·73H_{2}O takes place at -22°,
H_{2}SO_{4} + 1·5H_{2}O at -6°·5, H_{2}SO_{4} + 1·2H_{2}O at
+3°·7, H_{2}SO_{4} + 0·75H_{2}O at +2°·8, H_{2}SO_{4} + 0·5H_{2}O
at -16°; (_d_) that when there is less than 40H_{2}O per
100H_{2}SO_{4}, refrigeration separates out the normal hydrate
H_{2}SO_{4}, which melts at +10°·35, and that a solution of the
composition H_{2}SO_{4} + 0·35H_{2}O deposits crystals of this
hydrate at -34°, H_{2}SO_{4} + 0·1H_{2}O at -4°·1, H_{2}SO_{4} +
0·05H_{2}O at +4°·9, while fuming acid of the composition
H_{2}SO_{4} + 0·05SO_{3} deposits H_{2}SO_{4} at about +7°. Thus
the temperature of the separation of crystals clearly
distinguishes the above four regions of solutions, and in the
space between H_{2}SO_{4} + H_{2}O and +25H_{2}O a particular
hydrate H_{2}SO_{4}4H_{2}O separates out, discovered by Pickering,
the isolation of which deserves full attention and further
research. I may add here that the existence of a hydrate
H_{2}SO_{4}4H_{2}O was pointed out in my work, _The Investigation
of Aqueous Solutions_, p. 120 (1887), upon the basis that it has
at all temperatures a smaller value for the coefficient of
expansion _k_ in the formula S_{_t_} = S_{0}/(1 - _kt_) than the
adjacent (in composition) solutions of sulphuric acid. And for
solutions approximating to H_{2}SO_{4}10H_{2}O in their
composition, _k_ is constant at all temperatures (for more dilute
solutions the value of _k_ increases with _t_ and for more
concentrated solutions it decreases). This solution (with
10H_{2}O) forms the point of transition between more dilute
solutions which deposit ice (water) when refrigerated and those
which give crystals of H_{2}SO_{4}4H_{2}O. According to R. Pictet
(1894) the solution H_{2}SO_{4}10H_{2}O freezes at -88° (but no
reference is made as to what separates out), _i.e._ at a lower
temperature than all the other solutions of sulphuric acid.
However, in respect to these last researches of R. Pictet (for
88·88 p.c. H_{2}SO_{4} -55°, for H_{2}SO_{4}H_{2}O +3·5°, for
H_{2}SO_{4}2H_{2}O -70°, for H_{2}SO_{4}4H_{2}O -40°, &c.) it
should be remarked that they offer some quite improbable data; for
example, for H_{2}SO_{4}75H_{2}O they give the freezing point as
0°, for H_{2}SO_{4}300H_{2}O +4°·5, and even for
H_{2}SO_{4}1000H_{2}O +0°·5, although it is well known that a
small amount of sulphuric acid lowers the temperature of the
formation of ice. I have found by direct experiment that a frozen
solidified solution of H_{2}SO_{4} + 300H_{2}O melted completely
at 0°.
[51] With an excess of snow, the hydrate H_{2}SO_{4},H_{2}O, like the
normal hydrate, gives a freezing mixture, owing to the absorption
of a large amount of heat (the latent heat of fusion). In melting,
the molecule H_{2}SO_{4} absorbs 960 heat units, and the molecule
H_{2}SO_{4}H_{2}O 3,680 heat units. If therefore we mix one gram
molecule of this hydrate with seventeen gram molecules of snow,
there is an absorption of 18,080 heat units, because 17H_{2}O
absorbs 17 × 1,430 heat units, and the combination of the
monohydrate with water evolves 9,800 heat units. As the specific
heat of the resultant compound H_{2}SO_{4},18H_{2}O = 0·813, the
fall of temperature will be -52°·6. And, in fact, a very low
temperature may be obtained by means of sulphuric acid.
But only a few properties have been determined with sufficient accuracy.
In those properties which have been determined for many solutions of
sulphuric acid, it is actually seen that the above-mentioned definite
compounds are distinguished by distinctive marks of change. As an example
we may cite the variation of the specific gravity with a variation of
temperature (namely K = _ds/dt_, if _s_ be the sp. gr. and _t_ the
temperature). For the normal hydrate, H_{2}SO_{4}, this factor is easily
determined from the fact that--
_s_ = 18528 - 10·65_t_ + 0·013_t_^2,
where _s_ is the specific gravity at _t_ (degrees Celsius) if the sp.
gr. of water at 4° = 10,000. Therefore K = 10·65 - 0·026_t_. This means
that at 0° the sp. gr. of the acid H_{2}SO_{4} decreases by 10·65 for
every rise of a degree of temperature, at 10° by 10·39, at 20° by 10·13,
at 30° by 9·87.[52] And for solutions containing slightly more anhydride
than the acid H_{2}SO_{4} (_i.e._ for fuming sulphuric acid), as well as
for solutions containing more water, K is greater than for the acid
H_{2}SO_{4}. Thus for the solution SO_{3},2H_{2}SO_{4}, at 10° K = 11·0.
On diluting the acid H_{2}SO_{4} K again increases until the formation of
the solution H_{2}SO_{4},H_{2}O (K = 11·1 at 10°), and then, on further
dilution with water, it again decreases. Consequently both hydrates
H_{2}SO_{4} and H_{2}SO_{4},H_{2}O are here expressed by an alteration of
the magnitude of K.
[52] For example, if it be taken that at 19° the sp. gr. of pure
sulphuric acid is 1·8330, then at 20° it is 1·8330 - (20 -
19)10·13 = 1·8320.
[Illustration: FIG. 89.--Diagram showing the variation of the factor
(_ds/dp_) of the specific gravity of solutions of sulphuric acid. The
percentage quantities of the acid, H_{2}SO_{4}, are laid out on the axes
of abscissæ. The ordinates are the factors or rises in sp. gr. (water
at 4 = 10,000) with the increase in the quantity of H_{2}SO_{4}.]
This shows that in liquid solutions it is possible by studying the
variation of their properties (without a change of physical state) to
recognise the presence or formation of definite hydrate compounds, and
therefore an exact investigation of the properties of solutions, of their
specific gravity for instance, should give direct indications of such
compounds.[53] The mean result of the most trustworthy determinations of
this nature is given in the following tables. The first of these tables
gives the specific gravities (in vacuo, taking the sp. gr. of water at 4°
= 1), at 0° (column 3), 15° (column 4), and 30° (column 5),[53 bis] for
solutions having the composition H_{2}SO_{4} + _n_H_{2}O (the value of
_n_ is given in the first column), and containing _p_ (column 2) per
cent. (by weight in vacuo) of H_{2}SO_{4}.[53 tri]
_n_ _p_ 0° 15° 30°
100 5·16 1·0374 1·0341 1·0292
50 9·82 1·0717 1·0666 1·0603
25 17·88 1·1337 1·1257 1·1173
15 26·63 1·2040 1·1939 1·1837
10 35·25 1·2758 1·2649 1·2540
8 40·50 1·3223 1·3110 1·2998
6 47·57 1·3865 1·3748 1·3622
5 52·13 1·4301 1·4180 1·4062
4 57·65 1·4881 1·4755 1·4631
3 64·47 1·5635 1·5501 1·5370
2 73·13 1·6648 1·6500 1·6359
1 84·48 1·7940 1·7772 1·7608
0·5 91·59 1·8445 1·8284 1·8128
H_{2}SO_{4} 100 1·8529 1·8372 1·8221
[53] Unfortunately, notwithstanding the great number of fragmentary and
systematic researches which have been made (by Parks, Ure, Bineau,
Kolbe, Lunge, Marignac, Kremers, Thomsen, Perkin, and others) for
determining the relation between the sp. gr. and composition of
solutions of sulphuric acid, they contain discrepancies which
amount to, and even exceed, 0·002 in the sp. gr. For instance, at
15°·4 the solution of composition H_{2}SO_{4}3H_{2}O has a sp. gr.
1·5493 according to Perkin (1886), 1·5501 according to Pickering
(1890), and 1·5525 according to Lunge (1890). The cause of these
discrepancies must be looked for in the methods employed for
determining the composition of the solutions--_i.e._ in the
inaccuracy with which the percentage amount of H_{2}SO_{4} is
determined, for a difference of 1 p.c. corresponds to a difference
of from 0·0070 (for very weak solutions) to 0·0118 (for a solution
containing about 73 p.c.) in the specific gravity (that is the
factor _ds/dp_) at 15°. As it is possible to determine the
specific gravity with an accuracy even exceeding 0·0002, the
specific gravities given in the adjoining tables are only averages
and most probable data in which the error, especially for the
30-80 p.c. solutions cannot be less than 0·0010 (taking water at
4° as 1).
[53 bis] Judging from the best existing determinations (of Marignac,
Kremers, and Pickering) for solutions of sulphuric acid
(especially those containing more than 5 p.c. H_{2}SO_{4}) within
the limits of 0° and 30° (and even to 40°), the variation of the
sp. gr. with the temperature _t_ may (within the accuracy of the
existing determinations) be perfectly expressed by the equation
S_{_t_} = S_{_0_} + A_t_ + B_t_^2. It must be added that (1) three
specific gravities fully determine the variation of the density
with _t_; (2) _ds/dt_ = A + 2B_t_--_i.e._ the factor of the
temperature is expressed by a straight line; (3) the value of A
(if _p_ be greater than 5 p.c.) is negative, and numerically much
greater than B; (4) the value of B for dilute solutions containing
less than 25 p.c. is negative; for solutions approximating to
H_{2}SO_{4}3H_{2}O in their composition it is equal to 0, and for
solutions of greater concentration B is positive; (5) the factor
_ds/dp_ for all temperatures attains a maximum value about
H_{2}SO_{4}H_{2}O; (6) on dividing _ds/dt_ by S_{_0_}, and so
obtaining the coefficient of expansion _k_ (_see_ Note 53), a
minimum is obtained near H_{2}SO_{4} and H_{2}SO_{4}4H_{2}O, and a
maximum at H_{2}SO_{4}H_{2}O for all temperatures.
[53 tri] These data (as well as those in the following table) have been
recalculated by me chiefly upon the basis of Kremer's,
Pickering's, Perkin's, and my own determinations; all the
requisite corrections have been introduced, and I have reason for
thinking that in each of them the probable error (or difference
from the true figures, now unknown) of the specific gravity does
not exceed ±0·0007 (if water at 4° = 1) for the 25-80 p.c.
solutions, and ±0·0002 for the more dilute or concentrated
solutions.
In the second table the first column gives the percentage amount _p_ (by
weight) of H_{2}SO_{4}, the second column the weight in grams (S_{15}) of
a litre of the solution at 15° (at 4° the weight of a litre of water =
1,000 grams), the third column, the variation (_d_S/_dt_) of this weight
for a rise of 1°, the fourth column, the variation _d_S/_dp_ of this
weight (at 15°) for a rise of 1 per cent. of H_{2}SO_{4}, the fifth
column, the difference between the weight of a litre at 0° and 15° (S_{0}
- S_{15}), and the sixth column, the difference between the weight of a
litre at 15° and 30° (S_{15} - S_{30}).
_p_ _S__{15} _dS__{15}/_dt_ _dS__{15}/_dp_ _S__{0}- _S__{15}-
_S__{15} _S__{30}
0 999·15 0·148 7·0 0·7 3·4
5 1033·0 0·27 6·8 3·1 5·0
10 1067·7 0·38 7·1 5·2 6·4
20 1141·9 0·58 7·7 8·6 8·9
30 1221·3 0·69 8·2 10·4 10·4
40 1306·6 0·75 8·8 11·3 11·2
50 1397·9 0·79 9·9 11·9 11·8
60 1501·2 0·86 10·8 13·0 12·7
70 1613·1 0·93 11·6 14·1 13·8
80 1731·4 1·04 11·0 15·8 15·4
90 1819·9 1·08 5·4 16·4 16·0
95 1837·6 1·03 +1·7 15·8 15·1
100 1837·2 1·03 -1·9[54] 15·7 15·1
The figures in these tables give the means of finding the amount of
H_{2}SO_{4} contained in a solution from its specific gravity,[55] and
also show that 'special points' in the lines of variation of the specific
gravity with the temperature and percentage composition correspond to
certain definite compounds of H_{2}SO_{4} with OH_{2}. This is best seen
in the variation of the factors (_d_S/_dt_ and _d_S/_dp_) with the
temperature and composition (columns 3, 4, second table). We have already
mentioned how the factor of temperature points to the existence of
hydrates, H_{2}SO_{4} and H_{2}SO_{4},H_{2}O. As regards the factor
_d_S/_dp_ (giving the increase of sp. gr. with an increase of 1 per cent.
H_{2}SO_{4}) the following are the three most salient points: (1) In
passing from 98 per cent. to 100 per cent. the factor is negative, and at
100 per cent. about -0·0019 (_i.e._ at 99 per cent. the sp. gr. is about
1·8391, and at 100 per cent. about 1·8372, at 15°, the amount of
H_{2}SO_{4} has increased whilst the sp. gr. has decreased), but as soon
as a certain amount of SO_{3} is added to the definite compound
H_{2}SO_{4} (and 'fuming' acid formed) the specific gravity rises (for
example, for H_{2}SO_{4} 0·136 SO_{3} the sp. gr. at 15° = 1·866), that
is the factor becomes positive (and, in fact, greater by +0·01), so that
the formation of the definite hydrate H_{2}SO_{4} is accompanied by a
distinct and considerable break in the continuity of the factor[55 bis];
(2) The factor (_d_S/_dp_) in increasing in its passage from dilute to
concentrated solutions, attains a maximum value (at 15° about 0·012)
about H_{2}SO_{4}2H_{2}O, _i.e._ at about the hydrate corresponding to
the form SX_{6}; proper to the compounds of sulphur, for S(OH)_{6} =
H_{2}SO_{4}2H_{2}O; the same hydrate corresponds to the composition of
gypsum CaSO_{4}2H_{2}O, and to it also corresponds the greatest
contraction and rise of temperature in mixing H_{2}SO_{4} with H_{2}O
(_see_ Chapter I., Note 28); (3) The variation of the factor (_d_S/_dp_)
under certain variations in the composition proceeds so uniformly and
regularly, and is so different from the variation given under other
proportions of H_{2}SO_{4} and H_{2}O, that the sum of the variations of
_d_S/_dp_ is expressed by a series of straight lines, if the values of
_p_ be laid along the axis of abscissæ and those of _d_S/_dp_ along the
ordinates.[56] Thus, for instance, for 15°, at 10 per cent. _d_S/_dp_ =
0·0071, at 20 per cent. = 0·0077, at 30 per cent. = 0·0082, at 40 per
cent. = 0·0088, that is, for each 10 per cent. the factor increases by
about 0·0006 for the whole of the above range, but beyond this it becomes
larger, and then, after passing H_{2}SO_{4}2H_{2}O, it begins to fall
rapidly. Such changes in the variation of the factor take place
apparently about definite hydrates,[56 bis] and especially about
H_{2}SO_{4}4H_{2}O, H_{2}SO_{4}2H_{2}O and H_{2}SO_{4}H_{2}O. All this
indicating as it does the special chemical affinity of sulphuric acid for
water, although of no small significance for comprehending the nature of
solutions (_see_ Chapter I. and Chapter VII.), contains many special
points which require detailed investigation, the chief difficulty being
that it requires great accuracy in a large number of experimental data.
[54] The factor _d_S/_dp_ passes through 0, that is, the specific
gravity attains a maximum value at about 98 p.c. This was
discovered by Kohlrausch, and confirmed by Chertel, Pickering, and
others.
[55] Naturally under the condition that there is no other ingredient
besides water, which is sufficiently true. For commercial acid,
whose specific gravity is usually expressed in degrees of Baumé's
hydrometer, we may add that at 15°
Specific gravity 1 1·1 1·2 1·3 1·4 1·5 1·6 1·7 1·8
Degree Baumé 0 13 24 33·3 41·2 48·1 54·1 59·5 64·2
66° Baumé (the strongest commercial acid or oil of vitriol)
corresponds to a sp. gr. 1·84.
By employing the second table (by the method of interpolation) the
specific gravity, at a given temperature (from 0° to 30°) can be
found for any percentage amount of H_{2}SO_{4}, and therefore
conversely the percentage of H_{2}SO_{4} can be found from the
specific gravity.
[55 bis] Whether similar (even small) breaks in the continuity of the
factor _dS_/_dp_ exist or not, for other hydrates (for instance,
for H_{2}SO_{4}H_{2}O and H_{2}SO_{4}4H_{2}O) cannot as yet be
affirmed owing to the want of accurate data (Note 53). In my
investigation of this subject (1887) I admit their possibility,
but only conditionally; and now, without insisting upon a similar
opinion, I only hold to the existence of a distinct break in the
factor at H_{2}SO_{4}, being guided by C. Winkler's observations
ond the specific gravities of fuming sulphuric acid.
[56] In 1887, on considering all the existent observations for a
temperature 0°, I gave the accompanying scheme (p. 243) of the
variation of the factor _ds_/_dp_ at 0°.
I did not then (1887) give this scheme an absolute value, and now
after the appearance of two series of new determinations (Lunge
and Pickering in 1890), which disagree in many points, I think it
well to state quite clearly: (1) that Lunge's and Pickering's new
determinations have not added to the accuracy of our data
respecting the variation of the specific gravity of solutions of
sulphuric acid; (2) that the sum total of existing data does not
negative (within the limit of experimental accuracy) the
possibility of a rectilinear and broken form for the factors
_ds_/_dp_; (3) that the supposition of 'special points' in
_ds_/_dp_, indicating definite hydrates, finds confirmation in all
the latest determinations; (4) that the supposition respecting the
existence of hydrates determining a break of the factor _ds_/_dp_
is in in way altered if, instead of a series of broken straight
lines, there be a continuous series of curves, nearly approaching
straight lines; and (5) that this subject deserves (as I mentioned
in 1887) new and careful elaboration, because it concerns that
foremost problem in our science--solutions--and introduces a
special method into it--that is, the study of differential
variations in a property which is so easily observed as the
specific gravity of a liquid.
[56 bis] These hydrates are: (_a_) H_{2}SO_{4} = SO_{3}H_{2}O (melts
at + 10°·4); (_b_) H_{2}SO_{4}H_{2}O = SO_{3}2H_{2}O
(crystallo-hydrate, melts at +8°·5); (_c_) H_{2}SO_{4}2H_{2}O (is
apparently not crystallisable); (_d_) one of the hydrates between
H_{2}SO_{4}6H_{2}O and H_{2}SO_{4}3H_{2}O, most probably
H_{2}SO_{4}4H_{2}O = SO_{3}5H_{2}O, for it crystallises at -24°·5
(Note 50 bis); and (_e_) a certain hydrate with a large proportion
of water, about H_{2}SO_{4}150H_{2}O. The existence of the last is
inferred from the fact that the factor _ds_/_dp_ first falls,
starting from water, and then rises, and this change takes place
when _p_ is less than 5 p.c. Certainly a change in the variation
of _ds_/_dp_ or _ds_/_dt_ does take place in the neighbourhood of
these five hydrates (Pickering, 1890, recognised a far greater
number of hydrates). I think it well to add that if the
composition of the solutions be expressed by the percentage amount
of molecules--_r__{1}SO_{3} + (100 - _r__{1})H_{2}O we find that
for H_{2}SO_{4}, _r__{1} = 50, for H_{2}SO_{4}2H_{2}O _r__{1} = 25
= 50/2, for H_{2}SO_{4}H_{2}O, _r__{1} = 33·333 = 50 · 2/3, while
for H_{2}SO_{4}4H_{2}O, _r__{1} = 16·666 = 50 · 1/3--_i.e._ that
the chief hydrates are distributed symmetrically between H_{2}O
and H_{2}SO_{4}. Besides which I may mention that my researches
(1887) upon the abrupt changes in the factor for solutions of
sulphuric acid, and upon the correspondence of the breaks of
_ds_/_dp_ with definite hydrates, received an indirect
confirmation not only in the solutions of HNO_{3}, HCl,
C_{2}H_{6}O, C_{3}H_{8}O, &c., which I investigated (in my work
cited in Chapter I., Note 19), but also in the careful
observations made by Professor Cheltzoff on the solutions of
FeCl_{3} and ZnCl_{2} (Chapter XVI., Note 4) which showed the
existence in these solutions of an almost similar change in
_ds_/_dp_ as is found in sulphuric acid. The detailed researches
(1893) made by Tourbaba on the solutions of many organic
substances are of a similar nature. Besides which, H. Crompton
(1888), in his researches on the electrical conductivity of
solutions of sulphuric acid, and Tammann, in his observations on
their vapour tension, found a correlation with the hydrates
indicated as above by the investigation of their specific
gravities. The influence of mixtures of a definite composition
upon the chemical relations of solutions is even exhibited in such
a complex process as electrolysis. V. Kouriloff (1891) showed that
mixtures containing about 3 p.c., 47 p.c. and 73 p.c. of sulphuric
acid--_i.e._ whose composition approaches that of the hydrates
H_{2}SO_{4}150H_{2}O, H_{2}SO_{4}6H_{2}O and
H_{2}SO_{4}2H_{2}O--exhibit certain peculiarities in respect to
the amount of peroxide of hydrogen formed during electrolysis.
Thus a 3 p.c. solution gives a maximum amount of peroxide of
hydrogen at the negative pole, as compared with that given by
other neighbouring concentrations. Starting from 3 p.c., the
formation of peroxide of hydrogen ceases until a concentration of
47 p.c. is reached.
The great affinity of sulphuric acid for water is also seen from the
fact that when the strong acid acts on the majority of _organic
substances_ containing hydrogen and oxygen (especially on heating) it
very frequently _takes up these elements in the form of water_. Thus
strong sulphuric acid acting on alcohol, C_{2}H_{6}O, removes the
elements of water from it, and converts it into olefiant gas, C_{2}H_{4}.
It acts in a similar manner on wood and other vegetable tissues, which it
chars. If a piece of wood be immersed in strong sulphuric acid it turns
black. This is owing to the fact that the wood contains carbohydrates
which give up hydrogen and oxygen as water to the sulphuric acid, leaving
charcoal, or a black mass very rich in it. For example, cellulose,
C_{6}H_{10}O_{5}, acts in this manner.[57]
[57] Cellulose, for instance unsized paper or calico, is dissolved by
strong sulphuric acid. Acid diluted with about half its volume of
water converts it (if the action be of short duration) into
vegetable parchment (Chapter I., Note 18). The action of dilute
solutions of sulphuric acid converts it into hydro-cellulose, and
the fibre loses its coherent quality and becomes brittle. The
prolonged action of strong sulphuric acid chars the cellulose
while dilute acid converts it into glucose. If sulphuric acid be
kept in an open vessel, the organic matter of the dust held in the
atmosphere falls into it and blackens the acid. The same thing
happens if sulphuric acid be kept in a bottle closed by a cork;
the cork becomes charred, and the acid turns black. However, the
chemical properties of the acid undergo only a very slight change
when it turns black. Sulphuric acid which is considerably diluted
with water does not produce the above effects, which clearly shows
their dependence on the affinity of the sulphuric acid for water.
It is evident from the preceding that strong sulphuric acid will
act as a powerful poison; whilst, on the other hand, when very
dilute it is employed in certain medicines and as a fertiliser for
plants.
We have already had frequent occasion to notice the very _energetic acid
properties_ of sulphuric acid, and therefore we will now only consider a
few of their aspects. First of all we must remember that, with calcium,
strontium, and especially with barium and lead, sulphuric acid forms very
slightly soluble salts, whilst with the majority of other metals it gives
more easily soluble salts, which in the majority of cases are able, like
sulphuric acid itself, to combine with water to form crystallo-hydrates.
Normal sulphuric acid, containing two atoms of hydrogen in its molecule,
is able for this reason alone to form two classes of salts, _normal_ and
_acid_, which it does with great facility _with the alkali metals_. The
metals of the alkaline earths and the majority of other metals, if they
do form acid sulphates, do so under exceptional conditions (with an
excess of strong sulphuric acid), and these salts when formed are
decomposable by water--that is, although having a certain degree of
physical stability they have no chemical stability. Besides the acid
salts RHSO_{4}, sulphuric acid also gives other forms of acid salts. An
entire series of salts having the composition RHSO_{4},H_{2}SO_{4}, or
for bivalent metals RSO_{4},3H_{2}SO_{4},[58] has been prepared. Such
salts have been obtained for potassium, sodium, nickel, calcium, silver,
magnesium, manganese. They are prepared by dissolving the sulphates in an
excess of sulphuric acid and heating the solution until the excess of
sulphuric acid is driven off; on cooling, the mass solidifies to a
crystalline salt. Besides which, Rose obtained a salt having the
composition Na_{2}SO_{4},NaHSO_{4}, and if HNaSO_{4} be heated it easily
forms a salt Na_{2}S_{2}O_{7} = Na_{2}SO_{4},SO_{3}; hence it is clear
that sulphuric anhydride combines with various proportions of bases, just
as it combines with various proportions of water.
[58] Weber (1884) obtained a series of salts R_{2}O,8SO_{3}_n_H_{2}O
for K, Rb, Cs, and Tl.
We have already learned that sulphuric acid displaces the acid from the
salts of nitric, carbonic, and many other volatile acids. Berthollet's
laws (Chapter X.) explain this by the small volatility of sulphuric acid;
and, indeed, in an aqueous solution sulphuric acid displaces the much
less soluble boric acid from its compounds--for instance, from borax, and
it also displaces silica from its compounds with bases; but both boric
anhydride and silica, when fused with sulphates, decompose them,
displacing sulphuric anhydride, SO_{3}, because they are less volatile
than sulphuric anhydride. It is also well known that with metals,
sulphuric acid forms salts giving off hydrogen (Fe, Zn, &c.), or sulphur
dioxide (Cu, Hg, &c.).[58 bis]
[58 bis] Ditte (1890) divides all the metals into two groups with
respect to sulphuric acid; the first group includes silver,
mercury, copper, lead, and bismuth, which are only acted upon by
hot concentrated acid. In this case sulphurous anhydride is
evolved without any by-reactions. The second group contains
manganese, nickel, cobalt, iron, zinc, cadmium, aluminium, tin,
thallium, and the alkali metals. They react with sulphuric acid of
any concentration at any temperature. At a low temperature
hydrogen is disengaged, and at higher temperatures (and with very
concentrated acid) hydrogen and sulphurous anhydride are
simultaneously evolved.
The reactions of sulphuric acid _with respect to organic substances_ are
generally determined by its acid character, when the direct extraction of
water, or oxidation at the expense of the oxygen of the sulphuric
acid,[59] or disintegration does not take place. Thus the majority of the
saturated hydrocarbons, C_{_n_}H_{2_m_}, form with sulphuric acid a
special class of _sulphonic acids_, C_{_n_}H_{2_m_-1}(HSO_{3}); for
example, benzene, C_{6}H_{6}, forms benzenesulphonic acid,
C_{6}H_{5}.SO_{3}H, water being separated, for the formation of which
oxygen is taken up from the sulphuric acid, for the product contains less
oxygen than the sulphuric acid. It is evident from the existence of these
acids that the hydrogen in organic compounds is replaceable by the group
SO_{3}H, just as it may be replaced by the radicles Cl, NO_{2}, CO_{2}H
and others. As the radicle of sulphuric acid or _sulphoxyl_, SO_{2}OH or
SHO_{3}, contains, like carboxyl (Vol. I., p. 395), one hydrogen
(hydroxyl) of sulphuric acid, the resultant substances are acids whose
basicity is equal to the number of hydrogens replaced by sulphoxyl. Since
also sulphoxyl takes the place of hydrogen, and itself contains hydrogen,
the sulpho-acids are equal to a hydrocarbon + SO_{3}, just as every
organic (carboxylic) acid is equal to a hydrocarbon + CO_{2}. Moreover,
here this relation corresponds with actual fact, because many sulphonic
acids are obtained by the direct combination of sulphuric anhydride:
C_{6}H_{5},(SO_{3}H) = C_{6}H_{6} + SO_{3}. The sulphonic acids give
soluble barium salts, and are therefore easily distinguished from
sulphuric acid. They are soluble in water, are not volatile, and when
distilled give sulphurous anhydride (whilst the hydroxyl previously in
combination with the sulphurous anhydride remains in the hydrocarbon
group; thus phenol, C_{6}H_{5}.OH, is obtained from benzenesulphonic
acid), and they are very energetic, because the hydrogen acting in them
is of the same nature as in sulphuric acid itself.[60]
[59] For example, the action of hot sulphuric acid on nitrogenous
compounds, as applied in Kjeldahl's method for the estimation of
nitrogen (Volume I. p. 249). It is obvious that when sulphuric
acid acts as an oxidising agent it forms sulphurous anhydride.
The action of sulphuric acid on the alcohols is exactly similar to
its action on alkalis, because the alcohols, like alkalis, react
on acids; a molecule of alcohol with a molecule of sulphuric acid
separates water and forms an _acid_ ethereal salt--that is there
is produced an ethereal compound corresponding with acid salts.
Thus, for example, the action of sulphuric acid, H_{2}SO_{4}, on
ordinary alcohol, C_{2}H_{5}OH, gives water and sulphovinic acid,
C_{2}H_{5}HSO_{4}--that is, sulphuric acid in which one atom of
hydrogen is replaced by the radicle C_{2}H_{5} of ethyl alcohol,
SO_{2}(OH)(OC_{2}H_{5}), or, what is the same thing, the hydrogen
in alcohol is replaced by the radicle (sulphoxyl) of sulphuric
acid, C_{2}H_{5}O.SO_{2}(OH).
[60] We will mention the following difference between the sulphonic
acids and the ethereal acid sulphates (Note 59): the former
re-form sulphuric acid with difficulty and the latter easily. Thus
sulphovinic acid when heated with an excess of water is
reconverted into alcohol and sulphuric acid. This is explained in
the following manner. Both these classes of acids are produced by
the substitution of hydrogen by SO_{3}H, or the univalent radicle
of sulphuric acid, but in the formation of ethereal acid sulphates
the SO_{3}H replaces the hydrogen of the hydroxyl in the alcohol,
whilst in the formation of the sulphonic acids the SO_{3}H
replaces the hydrogen of a hydrocarbon. This difference is clearly
evidenced in the existence of two acids of the composition
SO_{4}C_{2}H_{6}. The one, mentioned above, is sulphovinic acid or
alcohol, C_{2}H_{5}.OH, in which the hydrogen of the hydroxyl is
replaced by sulphoxyl = C_{2}H_{5}.OSO_{3}H, whilst the other is
alcohol, in which one atom of the hydrogen in ethyl, C_{2}H_{5},
is replaced by the sulphonic group--that is =
(C_{2}H_{4})SO_{3}H·OH. The latter is called isethionic acid. It
is more stable than sulphovinic acid. The details as to these
interesting compounds must be looked for in works on organic
chemistry, but I think it necessary to note one of the general
methods of formation of these acids. The sulphites of the
alkalis--for example, K_{2}SO_{3}--when heated with the halogen
products of metalepsis, give a halogen salt and a salt of a
sulphonic acid. Thus methyl iodide, CH_{3}I, derived from marsh
gas, CH_{4}, when heated to 100° with a solution of potassium
sulphite, K_{2}SO_{3}, gives potassium iodide, KI, and potassium
methylsulphonate, CH_{3}SO_{3}K--that is a salt of the sulphonic
acid. This shows that the sulphonic acid may be referred to
sulphurous acid, and that there is a resemblance between sulphuric
and sulphurous acid, which clearly reveals itself here in the
formation of one product from them both.
Sulphuric acid, as containing a large proportion of oxygen, is a
substance which frequently acts as an oxidising agent: in which case it
is _deoxidised, forming sulphurous anhydride_ and water (or even,
although more rarely, sulphuretted hydrogen and sulphur). Sulphuric acid
acts in this manner on charcoal, copper, mercury, silver, organic and
other substances, which are unable to evolve hydrogen from it directly,
as we saw in describing sulphurous anhydride.
Although the hydrate of a higher saline form of oxidation (Chapter XV.),
sulphuric anhydride is capable of further oxidation, and forms a kind of
peroxide, just as hydrogen gives hydrogen peroxide in addition to water,
or as sodium and potassium, besides the oxides Na_{2}O and K_{2}O, give
their peroxides, compounds which are in a chemical sense unstable,
powerfully oxidising, and not directly able to enter into saline
combinations. If the oxides of potassium, barium, &c., be compared to
water, then their peroxides must in like manner correspond to hydrogen
peroxide,[61] not only because the oxygen contained in them is very
mobile and easily liberated, and because their reactions are similar, but
also because they can be mutually transformed into each other, and are
able to form compounds with each other, with bases and with water, and
indeed form a kind of peroxide salts.[62] This is also the character of
_persulphuric acid_, discovered in 1878 by Berthelot, and its
corresponding anhydride or peroxide of sulphur S_{2}O_{7}. It is formed
from 2SO_{3} + O with the absorption of heat (-27 thousand heat units),
like ozone from O_{2} + O (-29 thousand units of heat), or hydrogen
peroxide from H_{2}O + O (-21 thousand heat units).
[61] The reaction BaO + O develops 12,000 heat units, whilst the
reaction H_{2}O + O absorbs 21,000 heat units.
[62] Schöne obtained a compound of peroxide of barium with peroxide of
hydrogen. If barium peroxide be dissolved in hydrochloric (or
acetic) acid, or if a solution of hydrogen peroxide be diluted
with a solution of barium hydroxide, a pure hydrate is
precipitated having the composition BaO_{2},8H_{2}O (sometimes the
composition is taken as BaO_{2},6H_{2}O). This fact was already
known to Thénard. Schöne showed that if hydrogen peroxide be in
excess, a crystalline compound of the two peroxides,
BaO_{2}H_{2}O_{2}, is precipitated. Schöne also obtained small
well-formed crystals of the same composition by adding a solution
of ammonia to an acid solution of barium peroxide (containing a
barium salt and hydrogen peroxide or a compound of BaO_{2} with
the acid). Thus barium peroxide combines with both water and
hydrogen peroxide. This is a very important fact for the
comprehension of the composition of other peroxides. Moreover, if
the peroxides are able to give hydrates they can also form
corresponding salts, _i.e._ they can combine with bases and acids,
as was afterwards found to be the case on further research into
this subject.
Peroxide of sulphur is produced by the action of a silent discharge
upon a mixture of oxygen and sulphurous anhydride.[63] With water
S_{2}O_{7} gives persulphuric acid, H_{2}S_{2}O_{8}. The latter is
obtained more simply by mixing strong sulphuric acid (not weaker than
H_{2}SO_{4},2H_{2}O) directly with hydrogen peroxide, or by the action of
a galvanic current on sulphuric acid mixed with a certain amount of
water, and cooled, the electrodes being platinum wires, when persulphuric
acid naturally appears at the positive pole.[64] When an acid of the
strength H_{2}SO_{4},6H_{2}O is taken, at first the hydrate of the
sulphuric peroxide, S_{2}O_{7},H_{2}O only is formed; but when the
concentration about the positive pole reaches H_{2}SO_{4},3H_{2}O, a
mixture of hydrogen peroxide and the hydrate of sulphuric peroxide begins
to be formed. Dilute solutions of sulphuric peroxide can be kept better
than more concentrated solutions, but the latter may be obtained
containing as much as 123 grams of the peroxide to a litre. It is a very
instructive fact that hydrogen peroxide is always formed when strong
solutions of persulphuric acid break up on keeping. So that the bond
between the two peroxides is established both by analysis and synthesis:
hydrogen peroxide is able to produce S_{2}H_{2}O_{8}, and the latter to
produce hydrogen peroxide. A mixture of sulphuric peroxide with sulphuric
acid or water is immediately decomposed, with the evolution of oxygen,
either when heated or under the action of spongy platinum. The same thing
takes place with a solution of baryta, although at first no precipitate
is formed and the decomposition of the barium salt, BaS_{2}O_{8}, with
the formation of BaSO_{4}, only proceeds slowly, so that the solution may
be filtered (the barium salt of persulphuric acid is soluble in water).
Mercury, ferrous oxide, and the stannous salts, are oxidised by
S_{2}H_{2}O_{8}. These are all distinct signs of true peroxides. The same
common properties (capacity for oxidising, property of forming peroxide
of hydrogen, &c.) are possessed by the alkali salts of persulphuric acid,
which are obtained by the action of an electric current upon certain
sulphates, for instance ammonium or potassium sulphate. The ammonium salt
of persulphuric acid, (NH_{4})_{2}S_{2}O_{8}, is especially easily formed
by this means, and is now prepared on a large scale and used (like
Na_{2}O_{2} and H_{2}O_{2}) for bleaching tissues and fibres.[65]
[63] Anhydrous _sulphuric peroxide_, S_{2}O_{7}, is obtained by the
prolonged (8 to 10 hours) action of a silent discharge of
considerable intensity on a mixture of oxygen and sulphurous
anhydride; the vapour of sulphuric peroxide, S_{2}O_{7}, condenses
as liquid drops, or after being cooled to 0° in the form of long
prismatic crystals, resembling those of sulphuric anhydride. The
anhydrous compound S_{2}O_{7} (and also the hydrated compound)
cannot be preserved long, as it splits up into oxygen and
sulphuric anhydride. Direct experiment shows that a mixture of
equal volumes of sulphurous anhydride and oxygen leaves a residue
of a quarter of the oxygen taken, or half of the whole volume,
which indicates the formula S_{2}O_{7}. This substance is soluble
in water, and it then gives a hydrate, probably having the
composition S_{2}O_{7},H_{2}O = 2SHO_{4}. This solution oxidises
the salts SnX_{2}, potassium iodide, and others, which renders it
possible to prove that the solution actually contains one atom of
oxygen capable of effecting oxidation to two molecules of
sulphuric anhydride.
In order to fully demonstrate the reality of a peroxide form for
acids, it should be mentioned that some years ago Brodie obtained
the so-called _acetic peroxide_, (C_{2}H_{2}O)_{2}O_{2}, by the
action of barium peroxide on acetic anhydride, (C_{2}H_{3}O)_{2}O.
Its corresponding hydrate is also known. This shows that true
peroxides and their hydrates, with reactions similar to those of
hydrogen peroxide, are possible for acids. A similar higher oxide
has long been known for chromium, and Berthelot obtained a like
compound for nitric acid (Chapter VI., Note 26).
[64] When an acid of the strength H_{2}SO_{4}6H_{2}O is taken, at first
only the hydrate of the sulphuric peroxide, S_{2}O_{7}H_{2}O, is
formed, but when the concentration at the positive pole reaches
H_{2}SO_{4}3H_{2}O, a mixture of hydrogen peroxide and the hydrate
of sulphuric peroxide begins to be formed. A state of equilibrium
is ultimately arrived at when the amounts of these substances
correspond to the proportion S_{2}O_{7} : 2H_{2}O_{2}, which, as
it were, answers to a new hydrate, S_{2}O_{9}2H_{2}O. But its
existence cannot be admitted because the sulphuric peroxide can be
easily distinguished from the hydrogen peroxide in the solution
owing to the fact that the former does not act on an acid solution
of potassium permanganate, whilst the hydrogen peroxide disengages
both its own oxygen and that of the permanganic acid, converting
it into manganous oxide. Their common property of liberating
iodine from an acid solution of the potassium iodide enables the
sum of the active oxygen in them both to be determined.
[65] If a solution of sulphuric acid which has been first subjected to
electrolysis be neutralised with potash or baryta, the salt which
is formed begins to decompose rapidly with the evolution of oxygen
(Berthelot, 1890). On saturating with caustic baryta, the solution
of the salt formed may be separated from the sulphate of barium,
and then the composition of the resultant compound, BaS_{2}O_{8},
may be determined from the amount of oxygen disengaged. Marshall
(1891) studied the formation of this class of compounds more
fully; he subjected a saturated solution of bisulphate of
potassium to electrolysis with a current of 3-3-1/2 ampères;
before electrolysis dilute sulphuric acid is added to the liquid
surrounding the negative pole, and during electrolysis the
solution at the anode is cooled. The electrolysis is continued
without interruption for two days, and a white crystalline deposit
separates at the anode. To avoid decomposition, the latter is not
filtered through paper, but through a perforated platinum plate,
and dried on a porous tile. The mother liquor, with the addition
of a fresh solution of bisulphate of potassium, is again subjected
to electrolysis and the crystals formed at the anode are again
collected, &c. The salt so obtained may be recrystallised by
dissolving it in hot water and rapidly cooling the solution after
filtration; a small proportion of the salt is decomposed by this
treatment. Rapid cooling is followed by the formation of small
columnar crystals; slow cooling gives large prismatic crystals.
The composition of the salt is determined either by igniting it,
when it forms sulphate of potassium, or else by titrating the
active oxygen with permanganate: its composition was found to
correspond to the salt of persulphuric acid, K_{2}S_{2}O_{8}. The
solution of the salt has a neutral reaction, and does not give a
precipitate with salts of other metals. K_{2}S_{2}O_{8} is the
most insoluble of the salts of persulphuric acid. With nitrate of
silver it forms persulphate of silver, which gives peroxide of
silver under the action of water according to the equation
Ag_{2}S_{2}O_{8} + 2H_{2}O = Ag_{2}O_{2} + 2H_{2}SO_{4}. With an
alkaline solution of a cupric salt (Fehling's solution) it forms a
red precipitate of peroxide of copper. Manganese and cobalt salts
give precipitates of MnO_{2} and Co_{2}O_{3}. Ferrous salts are
rapidly oxidised, potassium iodide slowly disengages iodine at the
ordinary temperature. All these reactions indicate the powerful
oxidising properties of K_{2}S_{2}O_{8}. In oxidising in the
presence of water it gives a residue of KHSO_{4}. The
decomposition of the dry salt begins at 100° but is not complete
even at 250°. The freshly prepared salt is inodorous, but after
being kept in a closed vessel it evolves a peculiar smell
different from that of ozone. The ammonium salt of persulphuric
acid, (NH_{4})_{2}S_{2}O_{8}, is obtained in a similar manner. It
is soluble to the extent of 58 parts per 100 parts by weight of
water. The decomposition of the ammonium salt by the hydrated
oxide of barium gives the barium salt, BaS_{2}O_{8}4H_{2}O, which
is soluble to the extent of 52·2 parts in 100 parts of water at
0°. The crystals do not deliquesce in the air and decompose in the
course of several days; they decompose most rapidly in perfectly
dry air. Solutions of the pure salt decompose slowly at the
ordinary temperature; on boiling barium sulphate is gradually
precipitated, oxygen being liberated simultaneously. To completely
decompose this salt it is necessary to boil its solution for a
long time. Alcohol dissolves the solid salt; the anhydrous salt
does not separate from the alcoholic solution, but a hydrate
containing one molecule of water, BaS_{2}O_{8}H_{2}O, which is
soluble in water but insoluble in absolute alcohol. Solid barium
persulphate decomposes even when slightly heated. The free acid,
which may serve for the preparation of other salts, is obtained by
treating the barium salt with sulphuric acid. The lead salt,
PbS_{2}O_{8}, has been obtained from the free acid; it
crystallises with two or three molecules of water. It is soluble
in water, deliquesces in the air, and with alkalis gives a
precipitate of the hydrated oxide which rapidly oxidises into the
binoxide.
Traube, before Marshall's researches, thought that the
electrolysis of solutions of sulphuric acid did not give
persulphuric acid but a persulphuric oxide having the composition
SO_{4}. On repeating his former researches (1892) Traube obtained
a persulphuric oxide by the electrolysis of a 70 per cent.
solution of sulphuric acid, and he separated it from the solution
by means of barium phosphate. Analysis showed that this substance
corresponded to the above composition SO_{4}, and therefore Traube
considers it very likely that the salts obtained by Marshall
corresponded to an acid H_{2}SO_{4} + SO_{4}, _i.e._ that the
indifferent oxide, SO_{4}, can combine with sulphuric acid and
form peculiar saline compounds.
In order to understand the relation of sulphuric peroxide to sulphuric
acid we must first remark that hydrogen peroxide is to be considered, in
accordance with the law of substitution, as water, H(OH), in which H is
replaced by (OH). Now the relation of H_{2}S_{2}O_{8} to H_{2}SO_{4} is
exactly similar. The radicle of sulphuric acid, equivalent to hydrogen,
is HSO_{4};[65 bis] it corresponds with the (OH) of water, and therefore
sulphuric acid, H(SHO_{4}), gives (SHO_{4})_{2} or S_{2}H_{2}O_{8}, in
exactly the same manner as water gives (HO)_{2}--_i.e._ H_{2}O_{2}.[66]
[65 bis] Or one of those supposed ions which appear at the positive
pole in the decomposition of sulphuric acid by the action of a
galvanic current.
[66] If this be true one would expect the following peroxide hydrates:
for phosphoric acid, (H_{2}PO_{4})_{2} = H_{4}P_{2}O_{8} = 2H_{2}O
+ 2PO_{3}; for carbonic acid, (HCO_{3})_{2} = H_{2}C_{2}O_{6} =
H_{2}O + C_{2}O_{5}; and for lead the true peroxide will be also
Pb_{2}O_{5}, &c. Judging from the example of barium peroxide (Note
62), these peroxide forms will probably combine together. It seems
to me that the compounds obtained by Fairley for uranium are very
instructive as elucidating the peroxides. In the action of
hydrogen peroxide in an acid solution on uranium oxide, UO_{3},
there is formed a uranium peroxide, UO_{4},4H_{2}O (U = 240), but
hydrogen peroxide acts on uranium oxide in the presence of caustic
soda; on the addition of alcohol a crystalline compound containing
Na_{4}UO_{8},4H_{2}O is precipitated, which is doubtless a
compound of the peroxides of sodium, Na_{2}O_{2}, and uranium,
UO_{4}. It is very possible that the first peroxide,
UO_{4},4H_{2}O, contains the elements of hydrogen peroxide and
uranium peroxide, U_{2}O_{7}, or even U(OH)_{6},H_{2}O_{2}, just
as the peroxide form lately discovered by Spring for tin perhaps
contains Sn_{2}O_{3},H_{2}O_{2}.
The largest part _of the sulphuric acid made_ is used for reacting on
sodium chloride in the manufacture of sodium carbonate; for the
manufacture of the volatile acids, like nitric, hydrochloric, &c., from
their corresponding salts; for the preparation of ammonium sulphate,
alums, vitriols (copper and iron), artificial manures, superphosphate
(Chapter XIX., Note 18) and other salts of sulphuric acid; in the
treatment of bone ash for the preparation of phosphorus, and for the
solution of metals--for example, of silver in its separation from
gold--for cleaning metals from rust, &c. A large amount of oil of vitriol
is also used in treatment of organic substances; it is used for the
extraction of stearin, or stearic acid, from tallow, for refining
petroleum and various vegetable oils, in the preparation of
nitro-glycerine (Chapter VI., Notes 37 and 37 bis), for dissolving indigo
and other colouring matters, for the conversion of paper into vegetable
parchment, for the preparation of ether from alcohol, for the preparation
of various artificial scents from fusel oil, for the preparation of
vegetable acids, such as oxalic, tartaric, citric, for the conversion of
non-fermentable starchy substances into fermentable glucose, and in a
number of other processes. It would be difficult to find another
artificially-prepared substance which is so frequently applied in the
arts as sulphuric acid. Where there are not works for its manufacture,
the economical production of many other substances of great technical
importance is impossible. In those localities which have arrived at a
high technical activity the amount of sulphuric acid consumed is
proportionally large; sulphuric acid, sodium carbonate, and lime are the
most important of the artificially-prepared agents employed in factories.
Besides the normal acids of sulphur, H_{2}SO_{3}, H_{2}SO_{3}S, and
H_{2}SO_{4}, corresponding with sulphuretted hydrogen, H_{2}S, in the
same way that the oxy-acids of chlorine correspond with hydrochloric
acid, HCl, there exists a peculiar series of acids which are termed
_thionic acids_. Their general composition is S_{_n_}H_{2}O_{6}, where
_n_ varies from 2 to 5. If _n_ = 2, the acid is called dithionic acid.
The others are distinguished as trithionic, tetrathionic, and
pentathionic acids. Their composition, existence, and reactions are very
easily understood if they be referred to the class of the sulphonic
acids--that is, if their relation to sulphuric acid be expressed in just
the same manner as the relation of the organic acids to carbonic acid.
The organic acids, as we saw (Chapter IX.), proceed from the hydrocarbons
by the substitution of their hydrogen by carboxyl--that is, by the
radicle of carbonic acid, CH_{2}O_{3} - HO = CHO_{2}. The formation of
the acids of sulphur by means of sulphoxyl may be represented in the same
manner, HSO_{3} = H_{2}SO_{4} - HO. Therefore to hydrogen H_{2}, there
should correspond the acids H.SHO_{3}, sulphurous, and SHO_{3}.SHO_{3} =
S_{2}H_{2}O_{6}, or dithionic; to SH_{2} there should correspond the
acids SH(SHO_{3}) = H_{2}S_{2}O_{3} (thiosulphuric), and S(SHO_{3})_{2} =
H_{2}S_{3}O_{6} (trithionic); to S_{2}H_{2} the acids S_{2}H(SHO_{3}) =
H_{2}S_{3}O_{2} (unknown), and S_{2}(SHO_{3})_{2} = H_{2}S_{4}O_{6}
(tetrathionic); to S_{3}H_{2} the acids S_{3}H(SHO_{3}) and
S_{3}(SHO_{3})_{2} = H_{2}S_{5}O_{6} (pentathionic). We know that iodine
reacts directly with the hydrogen of sulphuretted hydrogen and combines
with it, and if thiosulphuric acid contains the radicle of sulphuretted
hydrogen (or hydrogen united with sulphur) of the same nature as in
sulphuretted hydrogen, it is not surprising that iodine reacts with
sodium thiosulphate and forms sodium tetrathionate. Thus, thiosulphuric
acid, HS(SHO_{3}), when deprived of H, gives a radicle which immediately
combines with another similar radicle, forming the tetrathionate
S_{2}(SO_{2}HO)_{2}. On this view[67] of the structure of the thionic
acids and salts, it is also clear how all the thionic acids, like
thiosulphuric acid, easily give sulphur and sulphides, with the exception
only of dithionic acid, H_{2}S_{2}O_{6}, which, judging from the above,
stands apart from the series of the other thionic acids. Dithionic acid
stands in the same relation to sulphuric acid as oxalic acid does to
carbonic acid. Oxalic acid is dicarboxyl, (CHO_{2})_{2} =
C_{2}H_{2}O_{4}, and so also dithionic acid is disulphoxyl, (SHO_{3})_{2}
= S_{2}H_{2}O_{6}. Oxalic acid when ignited decomposes into carbonic
anhydride and carbonic oxide, CO, and dithionic acid when heated
decomposes into sulphuric anhydride and sulphurous anhydride, SO_{2}, and
SO_{2} stands in the same relation to SO_{3} as CO to CO_{2}. This also
explains the peculiarity of the calcium, barium, and lead, &c. salts of
the thionic acids being easily soluble (although the corresponding salts
of H_{2}SO_{3}, H_{2}SO_{4}, and H_{2}S dissolve with difficulty),
because the former are similar to the salts of the sulphonic acids, which
are also soluble in water. Thus the thionic acids are _disulphonic
acids_, just as many dicarboxylic acids are known--for example,
CH_{2}(CO_{2}H)_{2}, C_{6}H_{4}(CO_{2}H)_{2}.[68]
[67] This view was communicated by me in 1870 to the Russian Chemical
Society.
[68] _Dithionic acid_, H_{2}S_{2}O_{6}, is distinguished among the
thionic acids as containing the least proportion of sulphur. It is
also called hyposulphuric acid, because its supposed anhydride,
S_{2}O_{5}, contains more O than sulphurous oxide, SO_{2} or
S_{2}O_{4}, and less than sulphuric anhydride, SO_{3} or
S_{2}O_{6}. Dithionic acid, discovered by Gay-Lussac and Welter,
is known as a hydrate and as salts, but not as anhydride. The
method for preparing dithionic acid usually employed is by the
action of finely-powdered manganese dioxide on a solution of
sulphurous anhydride. On shaking, the smell of the latter
disappears, and the manganese salt of the acid in question passes
into solution; MnO_{2} + 2SO_{2} = MnS_{2}O_{6}. If the
temperature be raised, the dithionate splits up into sulphurous
anhydride and manganese sulphate, MnSO_{4}. Generally owing to
this a mixture of manganese sulphate and dithionate is obtained in
the solution. They may be separated by mixing the solution of the
manganese salts with a solution of barium hydroxide, when a
precipitate of manganese hydroxide and barium sulphate is
obtained. In this manner barium dithionate only is obtained in
solution. It is purified by crystallisation, and separates as
BaS_{2}O_{6},2H_{2}O; this is then dissolved in water, and
decomposed with the requisite amount of sulphuric acid. Dithionic
acid, H_{2}S_{2}O_{6}, then remains in solution. By concentrating
the resultant solution under the receiver of an air-pump it is
possible to obtain a liquid of sp. gr. 1·347, but it still
contains water, and on further evaporation the acid decomposes
into sulphuric acid and sulphurous anhydride: H_{2}S_{2}O_{6} =
H_{2}SO_{4} + SO_{2}. The same decomposition takes place if the
solution be slightly heated. Like all the thionic acids, dithionic
acid is readily attacked by oxidising agents, and passes into
sulphurous acid. No dithionate is able to withstand the action of
heat, even when very slight, without giving off sulphurous
anhydride: K_{2}S_{2}O_{6} = K_{2}SO_{4} + SO_{2}. The alkali
dithionates have a neutral reaction (which indicates the energetic
nature of the acid) are soluble in water, and in this respect
present a certain resemblance to the salts of nitric acid (their
anhydrides are: N_{2}O_{5} and S_{2}O_{5}). Klüss (1888) described
many of the salts of dithionic acid.
Langlois, about 1840, obtained a peculiar thionic acid by heating
a strong solution of acid potassium sulphite with flowers of
sulphur to about 60°, until the disappearance of the yellow
coloration first produced by the solution of the sulphur. On
cooling, a portion of the sulphur was precipitated, and crystals
of a salt of _trithionic acid_, K_{2}S_{3}O_{6} (partly mixed with
potassium sulphate), separated out. Plessy afterwards showed that
the action of sulphurous acid on a thiosulphate also gives sulphur
and trithionic acid: 2K_{2}S_{2}O_{3} + 3SO_{2} = 2K_{2}S_{3}O_{6}
+ S. A mixture of potassium acid sulphite and thiosulphate also
gives a trithionate. It is very possible that a reaction of the
same kind occurs in the formation of trithionic acid by Langloid's
method, because potassium sulphite and sulphur yield potassium
thiosulphate. The potassium thiosulphate may also be replaced by
potassium sulphide, and on passing sulphurous anhydride through
the solution thiosulphate is first formed and then trithionate:
4KHSO_{3} + K_{2}S + 4SO_{2} = 3K_{2}S_{3}O_{6} + 2H_{2}O. The
sodium salt is not formed under the same circumstances as the
corresponding potassium salt. The sodium salt does not crystallise
and is very unstable: the barium salt is, however, more stable.
The barium and potassium salts are anhydrous, they give neutral
solutions and decompose when ignited, with the evolution of
sulphur and sulphurous anhydride, a sulphate being left behind,
K_{2}S_{3}O_{6} = K_{2}SO_{4} + SO_{2} + S. If a solution of the
potassium salt be decomposed by means of hydrofluosilicic or
chloric acid, the insoluble salts of these acids are precipitated
and trithionic acid is obtained in solution, which however very
easily breaks up on concentration. The addition of salts of
copper, mercury, silver, &c., to a solution of a trithionate is
followed, either immediately or after a certain time, by the
formation of a black precipitate of the sulphides whose formation
is due to the decomposition of the trithionic acid with the
transference of its sulphur to the metal.
_Tetrathionic acid_, H_{2}S_{4}O_{6}, in contradistinction to the
preceding acids, is much more stable in the free state than in the
form of salts. In the latter form it is easily converted into
trithionate, with liberation of sulphur. Sodium tetrathionate was
obtained by Fordos and Gélis, by the action of iodine on a
solution of sodium thiosulphate. The reaction essentially consists
in the iodine taking up half the sodium of the thiosulphate,
inasmuch as the latter contains Na_{2}S_{2}O_{3}, whilst the
tetrathionate contains NaS_{2}O_{3} or Na_{2}S_{4}O_{6}, so that
the reaction is as follows: 2Na_{2}S_{2}O_{3} + I_{2} = 2NaI +
Na_{2}S_{4}O_{6}. It is evident that tetrathionic acid stands to
thiosulphuric acid in exactly the same relation as dithionic acid
does to sulphurous acid; for the same amount of the other elements
in dithionate, KSO_{3}, and tetrathionate, KS_{2}O_{3}, there is
half as much metal as in sulphite, K_{2}SO_{3}, and thiosulphate,
K_{2}S_{2}O_{3}. If in the above reaction the sodium thiosulphate
be replaced by the lead salt PbS_{2}O_{3}, the sparingly-soluble
lead iodide PbI_{2} and the soluble salt PbS_{4}O_{6} are
obtained. Moreover the lead salt easily gives tetrathionic acid
itself (PbSO_{4} is precipitated). The solution of tetrathionic
acid may be evaporated over a water-bath, and afterwards in a
vacuum, when it gives a colourless liquid, which has no smell and
a very acid reaction. When dilute it may be heated to its
boiling-point, but in a concentrated form it decomposes into
sulphuric acid, sulphurous anhydride, and sulphur: H_{2}S_{4}O_{6}
= H_{2}SO_{4} + SO_{2} + S_{2}.
_Pentathionic acid_, H_{2}S_{5}O_{6}, also belongs to this series
of acids. But little is known concerning it, either as hydrate or
in salts. It is formed, together with tetrathionic acid, by the
direct action of sulphurous acid on sulphuretted hydrogen in an
aqueous solution; a large proportion of sulphur being precipitated
at the same time: 5SO_{2} + 5H_{2}S = H_{2}S_{5}O_{6} + 5S +
4H_{2}O.
If, as was shown above, the thionic acids are disulphonic acids,
they may be obtained, like other sulphonic acids, by means of
potassium sulphite and sulphur chloride. Thus Spring demonstrated
the formation of potassium trithionate by the action of sulphur
dichloride on a strong solution of potassium sulphite: 2KSO_{3}K +
SCl_{2} = S(SO_{3}K)_{2} + 2KCl. If sulphur chloride be taken,
sulphur also is precipitated. The same trithionate is formed by
heating a solution of double thiosulphates; for example, of
AgKS_{2}O_{3}. Two molecules of the salts then form silver
sulphide and potassium trithionate. If the thiosulphate be the
potassium silver salt SO_{3}K(AgS), then the structure of the
trithionate must necessarily be (SO_{3}K)_{2}S. Previous to
Spring's researches, the action of iodine on sodium thiosulphate
was an isolated accidentally discovered reaction; he, however,
showed its general significance by testing the action of iodine on
mixtures of different sulphur compounds. Thus with iodine, I_{2},
the mixture Na_{2}S + Na_{2}SO_{3} forms 2NaI + Na_{2}S_{2}O_{3},
whilst the mixture Na_{2}S_{2}O_{3} + Na_{2}SO_{3} + I_{2} gives
2NaI + Na_{2}S_{3}O_{6}--that is, trithionic acid stands in the
same relation to thiosulphuric acid as the latter does to
sulphuretted hydrogen. We adopt the same mode of representation:
by replacing one hydrogen in H_{2}S by sulphuryl we obtain
thiosulphuric acid, HSO_{3}.HS, and by replacing a second hydrogen
in the latter again by sulphuryl we obtain trithionic acid,
(HSO_{3})_{2}S. Furthermore, Spring showed that the action of
sodium amalgam on the thionic acids causes reverse reactions to
those above indicated for iodine. Thus sodium thiosulphate with
Na_{2} gives Na_{2}S + Na_{2}SO_{3}, and Spring showed that the
sodium here is not a simple element taking up sulphur, but itself
enters into double decomposition, replacing sulphur; for on taking
a potassium salt and acting on it with sodium, KSO_{3}(SK) + NaNa
= KSO_{3}Na + (SK)Na. In a similar way sodium dithionate with
sodium gives sodium sulphite: (NaSO_{3})_{2} + Na_{2} =
2NaSO_{3}Na; sodium trithionate forms NaSO_{3}Na and NaSO_{3}.SNa,
and tetrathionate forms sodium thiosulphate,
(NaSO_{3})S_{2}(NaSO_{3}) + Na_{2} = 2(NaSO_{3})(NaS).
In all the oxidised compounds of sulphur we may note the presence
of the elements of sulphurous anhydride, SO_{2}, the only product
of the combustion of sulphur, and in this sense the compounds of
sulphur containing one SO_{2} are--
H HO C_{6}H_{5} HS
SO_{2} SO_{2} SO_{2} SO_{2}
HO HO HO HO
Sulphurous Sulphuric Benzene sulphonic Thiosulphuric
acid acid acid acid
while, according to this mode of representation, the thionic acids
are--
HO HO HO HO
SO_{2} SO_{2} SO_{2} SO_{2}
S S_{2} S_{3}
SO_{2} SO_{2} SO_{2} SO_{2}
HO HO HO HO
Dithionic Trithionic Tetrathionic Pentathionic
Hence it is evident that SO_{2} has (whilst CO_{2} has not) the
faculty for combination, and aims at forming SO_{2}X_{2}. These
X_{2} can = O, and the question naturally suggests itself as to
whether the O_{2} which occurs in SO_{2} is not of the same nature
as this oxygen which adds itself to SO_{2}--that is, whether
SO_{2} does not correspond with the more general type SX_{4}, and
its compounds with the type SX_{6}? To this we may answer 'Yes'
and 'No'--'Yes' in the general sense which proceeds from the
investigation of the majority of compounds, especially metals,
where RO corresponds with RCl_{2}, RX_{2}; 'No' in the sense that
sulphur does not give either SH_{4}, SH_{6}, or SCl_{6}, and
therefore the stages SX_{4} and SX_{6} are only observable in
oxygen compounds. With reference to the type SX_{6} a hydrate,
S(HO)_{6}, might be expected, if not SCl_{6}. And we must
recognise this hydrate from a study of the compounds of sulphuric
acid with water. In addition to what has been already said
respecting the complex acids formed by sulphur, I think it well to
mention that, according to the above view, still more complex
oxygen acids and salts of sulphur may be looked for. For instance,
the salt Na_{2}S_{4}O_{8} obtained by Villiers (1888) is of this
kind. It is formed together with sodium trithionate and sulphur,
when SO_{2} is passed through a cold solution of Na_{2}S_{2}O_{3},
which is then allowed to stand for several days at the ordinary
temperature: 2Na_{2}S_{2}O_{3} + 4SO_{2} = Na_{2}S_{4}O_{8} +
Na_{2}S_{3}O_{6} + S. It may be assumed here, as in the thionic
acids, that there are two sulphoxyls, bound together not only by
S, but also by SO_{2}, or what is almost the same thing, that the
sulphoxyl is combined with the residue of trithionic acid, _i.e._
replaces one aqueous residue in trithionic acid.
Sulphur exhibits an acid character, not only in its compounds with
hydrogen and oxygen, but also in those with other elements. The compound
of sulphur and carbon has been particularly well investigated. It
presents a great analogy to carbonic anhydride, both in its elementary
composition and chemical character. This substance is the so-called
carbon bisulphide, CS_{2}, and corresponds with CO_{2}.
The first endeavours to obtain a compound of sulphur with carbon were
unsuccessful, for although sulphur does combine directly with carbon, yet
the formation of this compound requires distinctly definite conditions.
If sulphur be mixed with charcoal and heated, it is simply driven off
from the latter, and not the smallest trace of carbon bisulphide is
obtained. The formation of this compound requires that the charcoal
should be first heated to a red heat, but not above, and then either the
vapour of sulphur passed over it or lumps of sulphur thrown on to the
red-hot charcoal, but in small quantities, so as not to lower the
temperature of the latter. If the charcoal be heated to a white heat, the
amount of carbon bisulphide formed is less. This depends, in the first
place, on the carbon bisulphide dissociating at a high temperature.[69]
In the second place, Favre and Silberman showed that in the combustion of
one gram of carbon bisulphide (the products will be CO_{2} + 2SO_{2})
3,400 heat units are evolved--that is, the combustion of a molecular
quantity of carbon bisulphide evolves 258,400 heat units (according to
Berthelot, 246,000). From a molecule of carbon bisulphide in grams we may
obtain 12 grams of carbon, whose combustion evolves 96,000 heat units,
and 64 grams of sulphur, evolving by combustion (into SO_{2}) 140,800
heat units. Hence we see that the component elements separately evolve
less heat by their combustion (237,000 heat units) than carbon bisulphide
itself--that is, that heat should be evolved (at the ordinary
temperature) and not absorbed in its decomposition, and therefore that
the formation of carbon bisulphide from charcoal and sulphur is in all
probability accompanied by an absorption of heat.[70] It is therefore not
surprising that, like other compounds produced with an absorption of heat
(ozone, nitrous oxide, hydrogen peroxide, &c.), carbon bisulphide is
unstable and easily converted into the original substances from which it
is obtained. And indeed if the vapour of carbon bisulphide be passed
through a red-hot tube, it is decomposed--that is, it dissociates--into
sulphur and carbon. And this takes place at the temperature at which this
substance is formed, just as water decomposes into hydrogen and oxygen at
the temperature of its formation. In this absorption of heat in the
formation of carbon bisulphide is explained the facility with which it
suffers reactions of decomposition, which we shall see in the sequel, and
its main difference from the closely analogous carbonic anhydride.
[Illustration: FIG. 90.--Apparatus for the manufacture of carbon
bisulphide.
[69] Even light decomposes carbon bisulphide, but not to the extent of
separating carbon; under the action of the sun's rays it is
decomposed into sulphur and solid substance which is considered to
be carbon monosulphide; it is of a red colour, and its sp. gr. is
1·66. (The formation of a red liquid compound C_{3}S_{2} has also
been remarked.) Thorpe (1889) observed a complete decomposition of
carbon bisulphide under the action of a liquid alloy of potassium
and sodium; it is accompanied by an explosion and the deposition
of carbon and sulphur. A similar complete decomposition of carbon
bisulphide is also accomplished by the action of mercury fulminate
(Chapter XVI., Note 26), and is due to the fact that _at the
ordinary temperature_ (at which carbon bisulphide is not produced)
_the decomposition_ of carbon bisulphide takes place with the
development of heat--that is, it presents an exothermal reaction,
like the decomposition of all explosives. It is very possible that
at a higher temperature, when carbon bisulphide is formed, the
_combination_ of carbon with sulphur is also an exothermal
reaction--that is, heat is developed. If this should be the case,
carbon bisulphide would present a most instructive example in
thermochemistry.
[70] The fact should not be lost sight of that sulphur and charcoal are
solids at the ordinary temperature, whilst carbon bisulphide is a
very volatile liquid, and consequently, in the act of combination,
referred to the ordinary temperature (Note 69), there is, as it
were, a passage into a liquid state, and this requires the
absorption of heat. And furthermore, the molecule of sulphur
contains at least six atoms, and the molecule of carbon in all
probability (Chapter VIII.) a very considerable number of atoms;
thus the reaction of sulphur on charcoal may be expressed in the
following manner: 3C_{_n_} + _n_S_{6} = 3_n_CS_{2}--that is, from
_n_ + 3 molecules there proceed 3_n_ molecules, and as _n_ must be
very considerable, 3_n_ must be greater than 3 + _n_, which
indicates a decomposition in the formation of carbon bisulphide,
although the reaction at first sight appears as one of
combination. This decomposition is seen also from the volumes in
the solid and liquid states. Carbon bisulphide has a sp. gr. of
1·29; hence its molecular volume is 59. But the volume of carbon,
even in the form of charcoal, is not more than 6, and the volume
of S_{2} is 30; hence 36 volumes after combination give 59
volumes--an expansion takes place, as in decompositions.
In the laboratory carbon bisulphide is prepared as follows: A porcelain
tube is luted into a furnace in an inclined position, the upper extremity
of the tube being closed by a cork, and the lower end connected with a
condenser. The tube contains charcoal, which is raised to a red heat, and
then pieces of sulphur are placed in the upper end. The sulphur melts,
and its vapour comes into contact with the red-hot charcoal, when
combination takes place; the vapours condense in the condenser, carbon
bisulphide being a liquid boiling at 48°. On a large scale the apparatus
depicted in fig. 90 is employed. A cast-iron cylinder rests on a stand in
a furnace. Wood charcoal is charged into the cylinder through the upper
tube closed by a clay stopper, whilst the sulphur is introduced through a
tube reaching to the bottom of the cylinder. Pieces of sulphur thrown
into this tube fall on to the bottom of the cylinder, and are converted
into vapour, which passes through the entire layer of charcoal in the
cylinder. The vapour of carbon bisulphide thus formed passes through the
exit tube first into a Woulfe's bottle (where the sulphur which has not
entered into the reaction is condensed), and then into a strongly-cooled
condenser or worm.[71]
[71] Carbon bisulphide, as prepared on a large scale, is generally very
impure, and contains not only sulphur, but, more especially, other
impurities which give it a very disagreeable odour. The best
method of purifying this malodorous carbon bisulphide is to shake
it up with a certain amount of mercuric chloride, or even simply
with mercury, until the surface of the metal ceases to turn black.
After this the carbon bisulphide must be poured off and distilled
over a water-bath, after mixing with some oil to retain the
impurities.
Pure carbon bisulphide is a colourless liquid, which refracts light
strongly, and has a pure ethereal smell; at 0° its specific gravity is
1·293, and at 15° 1·271. If kept for a long time it seems to undergo a
change, especially when it is kept under water, in which it is insoluble.
It boils at 48°, and the tension of its vapour is so great that it
evaporates very easily, producing cold,[72] and therefore it has to be
kept in well-stoppered vessels; it is generally kept under a layer of
water, which hinders its evaporation and does not dissolve it.[73]
[72] If carbon bisulphide be evaporated under the receiver of an
air-pump, or by means of a current of air, it is possible to
obtain a temperature as low as -60°, and the carbon bisulphide
does not solidify at this temperature. However, if a series of
air-bubbles be passed through it by means of bellows, a
crystalline white substance remains which volatilises below 0°:
this a hydrate, H_{2}O,2CS_{2}; it easily decomposes into water
and carbon bisulphide. It is formed in the above experiment by the
moisture held in the air passed through the carbon bisulphide, and
the fall of temperature.
[73] Strong alcohol is miscible in all proportions with carbon
bisulphide, but dilute alcohol only in a definite amount, owing to
its diminished solubility from the presence of the water in it.
Ether, hydrocarbons, fatty oils, and many other organic substances
are soluble with great ease in carbon bisulphide. This is taken
advantage of in practice for extracting the fatty oils from
vegetable seeds, such as linseed, palm-nuts, or from bones, &c.
The preparation of vegetable oils is usually done by pressing the
seeds under a press, but the residue always contains a certain
amount of oil. These traces of oil can, however, be removed by
treatment with carbon bisulphide. In this manner a solution is
obtained which when heated easily parts with all the carbon
bisulphide, leaving the non-volatile fatty oil behind, so that the
same carbon bisulphide may be condensed and used over again for
the same purpose. It also dissolves iodine, bromine, indiarubber,
sulphur, and tars.
Carbon bisulphide, especially at high temperatures, very often
acts by its elements in a manner in which carbon and sulphur alone
are not able to react, which will be understood from what has been
said above respecting its endothermal origin. If it be passed over
red-hot metals--even over copper, for instance, not to mention
sodium, &c.--it forms a sulphide of the metal and deposits
charcoal, and if the vapour be passed over incandescent metallic
oxides it forms metallic sulphides and carbonic anhydride (and
sometimes a certain amount of sulphurous anhydride). Lime and
similar oxides give under these circumstances a carbonate and a
sulphide--for example, CS_{2} +3CaO = 2CaS + CaCO_{3}. The
sulphides obtained by this means are often well crystallised, like
those found in nature--for example, lead and antimony sulphides.
Carbon bisulphide enters into many combinations, which are frequently
closely analogous to the compounds of carbonic anhydride. In this respect
it is a _thio-anhydride_--_i.e._ it has the character of the acid
anhydrides,[73 bis] like carbonic anhydride, with the difference that the
oxygen of the latter is replaced by sulphur. By thio-compounds in general
are understood those compounds of sulphur which differ from the compounds
of oxygen as carbon bisulphide does from carbonic anhydride--that is,
which correspond with the oxygen compounds, but with substitution of
sulphur for oxygen. Thus thiosulphuric acid is monothiosulphuric
acid--that is, sulphuric acid in which one atom of sulphur replaces one
atom of oxygen. With the sulphides of the alkalis and alkaline earths, it
forms saline substances corresponding with the carbonates, and these
compounds may be termed _thiocarbonates_. For example, the composition of
the sodium salt Na_{2}CS_{3} is exactly like that of sodium carbonate.
They are formed by the direct solution of carbon bisulphide in aqueous
solutions of the sulphides; but they are difficult to obtain in a
crystalline form, because they are easily decomposable. When the
solutions of these salts are highly concentrated they begin to decompose,
with the evolution of sulphuretted hydrogen and the formation of a
carbonate, water taking part in the reaction--for example, K_{2}CS_{3} +
3H_{2}O = K_{2}CO_{3} + 3H_{2}S.[74]
[73 bis] And just as COCl_{2} corresponds to CO_{2}, so also the
chloranhydride, CSCl_{2}, or _thiophosgene_, corresponds to CS_{2}.
[74] If instead of a sulphide we take an alkali hydroxide, a
thiocarbonate is also formed, together with a carbonate--thus,
3BaH_{2}O_{2} + 3CS_{2} = 2BaCS_{3} + BaCO_{3} + 3H_{2}O. From the
instability of the thiocarbonates of the alkaline metals we can
clearly see the reason of the difficulty with which the salts of
the heavier metals are formed, whose basic properties are
incomparably weaker than those of the alkali metals. However,
these salts may be obtained by double decomposition. Ammonia in
reacting on carbon bisulphide gives, besides products like those
formed by other alkalis, a whole series of products of as complex
a structure as those substances which are produced by the action
of carbonic anhydride on ammonia. In the ninth chapter we examined
the formation of the ammonium carbonates, and saw the transition
from them into the cyanides. It is not surprising after this that
the action of carbon bisulphide on ammonia not only produces the
above-mentioned salts, but also amidic compounds corresponding
with them, in which the oxygen is wholly or partially replaced by
sulphur. Thus ammonium dithiocarbamate is very easily obtained if
carbon bisulphide be added to an alcoholic solution of ammonia,
and the mixture cooled in a closed vessel. The salt then separates
out in minute yellow crystals, CN_{2}H_{6}S_{2}.
Carbon bisulphide not only forms compounds with the metallic
sulphides, but also with sulphuretted hydrogen--that is, it forms
_thiocarbonic acid_, H_{2}CS_{3}. This is obtained by carefully
mixing solutions of thiocarbonates with dilute hydrochloric acid.
It then separates in an oily layer, which easily decomposes in the
presence of water into sulphuretted hydrogen and carbon
bisulphide, just as the corresponding carbonic acid (hydrate)
decomposes into water and carbonic anhydride. Carbon bisulphide
combines not only with sodium sulphide, but also with the
bisulphide, Na_{2}S_{2}, not, however, with the trisulphide,
Na_{2}S_{3}.
The relation of carbon bisulphide to the other carbon compounds
presents many most interesting features which are considered in
organic chemistry. We will here only turn our attention to one of
the compounds of this class. Ethyl sulphide, (C_{2}H_{5})_{2}S,
combines with ethyl iodide, C_{2}H_{5}I, forming a new molecule,
S(C_{2}H_{5})_{3}I. If we designate the hydrocarbon group, for
instance ethyl, C_{2}H_{5}, by Et, the reaction would be expressed
by the following equation : Et_{2}S + EtI = SEt_{3}I. This
compound is of a saline character, corresponds with salts of the
alkalis, and is closely analogous to ammonium chloride. It is
soluble in water; when heated it again splits up into its
components EtI and Et_{2}S, and with silver hydroxide gives a
hydroxide, Et_{3}S·OH, having the property of a distinct and
energetic alkali, resembling caustic ammonia. Thus the compound
group SEt_{3} combines, like potassium or ammonium, with iodine,
hydroxyl, chlorine, &c. The hydroxide SEt_{3}·OH is soluble in
water, precipitates metallic salts, saturates acids, &c. Hence
sulphur here enters into a relation towards other elements similar
to that of nitrogen in ammonia and ammonium salts, with only this
difference, that nitrogen retains, besides iodine, hydroxyl, and
other groups, also H_{4} or Et_{4} (for example, NH_{4}Cl,
NEt_{3}HI, NEt_{4}I), whilst sulphur only retains Et_{3}.
Compounds of the formula SH_{3}X are however unknown, only the
products of substitution SEt_{3}X, &c. are known. The distinctly
alkaline properties of the hydroxide, triethylsulphine hydroxide,
SEt_{3}OH, and also the sharply-defined properties of the
corresponding hydroxide, tetraethylammonium hydroxide, NEt_{4}OH,
depend naturally not only on the properties of the nitrogen and
sulphur entering into their composition, but also on the large
proportion of hydrocarbon groups they contain. Judging from the
existence of the ethylsulphine compounds, it might be imagined
that sulphur forms a compound, SH_{4}, with hydrogen; but no such
compound is known, just as NH_{5} is unknown, although NH_{4}Cl
exists.
A remarkable example[74 bis] of the thio-compounds is found in
_thiocyanic acid_--_i.e._ cyanic acid in which the oxygen is replaced by
sulphur, HCNS. We know (Chapter IX.) that with oxygen the cyanides of the
alkaline metals RCN give cyanates RCNO; but they also combine with
sulphur, and therefore if yellow prussiate of potash be treated as in the
preparation of potassium cyanide, and sulphur be added to the mass,
potassium thiocyanate, KNCS, is obtained in solution. This salt is much
more stable than potassium cyanate; it dissolves without change in water
and alcohol, forming colourless solutions from which it easily
crystallises on evaporation. It may be kept exposed to air even when in
solution; in dissolving in water it absorbs a considerable amount of
heat, and forms a starting-point for the preparation of all the
thiocyanates, RCNS, and organic compounds in which the metals are
replaced by hydrocarbon groups. Such, for example, is volatile mustard
oil, C_{3}H_{5}CSN (allyl thiocyanate),[75] which gives to mustard its
caustic properties. With ferric salts the thiocyanates give an
exceedingly brilliant red coloration, which serves for detecting the
smallest traces of ferric salts in solution. Thiocyanic acid, HCNS, may
be obtained by a method of double decomposition, by distilling potassium
thiocyanate with dilute sulphuric acid. It is a volatile colourless
liquid, having a smell recalling that of vinegar, is soluble in water,
and may be kept in solution without change.[75 bis]
[74 bis] Thorpe and Rodger (1889), by heating a mixture of lead
fluoride and phosphorus pentasulphide to 250° in an atmosphere of
dry nitrogen, obtained gaseous _phosphorus fluosulphide_, or
_thiophosphoryl fluoride_, PSF_{3}, corresponding with POCl_{3}.
This colourless gas is converted into a colourless liquid by a
pressure of eleven atmospheres; it does not act on dry mercury,
and takes fire spontaneously in air or oxygen, forming phosphorus
pentafluoride, phosphoric anhydride, and sulphurous anhydride. It
is soluble in ether, but is decomposed by water: PSF_{3} + 4H_{2}O
= H_{2}S + H_{3}PO_{4} + 3HF (Note 20).
[75] Although mustard oil may be obtained from the thiocyanates, it is
only an isomer of allyl thiocyanate proper, as is explained in
Organic Chemistry.
[75 bis] Sulphur can only replace half the oxygen in CO_{2}, as is seen
in _carbon oxysulphide_, or monothiocarbonic anhydride COS. This
substance was obtained by Than, and is formed in many reactions. A
certain amount is obtained if a mixture of carbonic oxide and the
vapour of sulphur be passed through a red-hot tube. When carbon
tetrachloride is heated with sulphurous anhydride, this substance
is also formed; but it is best obtained in a pure form by
decomposing potassium thiocyanate with a mixture of equal volumes
of water and sulphuric acid. A gas is then evolved containing a
certain amount of hydrocyanic acid, from which it may be freed by
passing it over wool containing moistened mercuric oxide, which
retains the hydrocyanic acid. The reaction is expressed by the
equation: 2KCNS + 2H_{2}SO_{4} + 2H_{2}O = K_{2}SO_{4} +
(NH_{4})_{2}SO_{4} + 2COS. It is also formed by passing the vapour
of carbon bisulphide over alumina or clay heated to redness
(Gautier; silicon sulphide is then formed). COS is also formed by
passing phosgene over a long layer of asbestos mixed with cadmium
sulphide at 270°; CdS + COCl_{3} = CdCl_{2} + COS (Nuricsán,
1892). The pure gas has an aromatic odour, is soluble in an equal
volume of water, which, however, acts on it, so that it must be
collected over mercury. When slightly heated, carbon oxysulphide
decomposes into sulphur and carbonic oxide. It burns in air with a
pale blue flame, explodes with oxygen, and yields potassium
sulphide and carbonate with potassium hydroxide: COS + 4KHO =
K_{2}CO_{3} + K_{2}S + 2H_{2}O.
The sulphur compounds of chlorine Cl_{2}S and Cl_{2}S_{2} may be
regarded on the one hand as products of the metalepsis of the sulphides
of hydrogen, H_{2}S and H_{2}S_{2}; and on the other hand of the oxygen
compounds of chlorine, because chloride of sulphur, Cl_{2}S, resembles
chlorine oxide, Cl_{2}O, whilst Cl_{2}S_{2} corresponds with the higher
oxide of chlorine; or thirdly, we may see in these compounds the type of
the acid chloranhydrides, because they are all decomposed by water,
forming hydrochloric acid, and sulphur tetrachloride, SCl_{4}, is
decomposed with the formation of sulphurous anhydride.[76]
[76] There is no reason for seeing any contradiction or mutual
incompatibility in these three views, because every analogy is
more or less modified by a change of elements. Thus, for instance,
it cannot be expected that the product of the metalepsis of
hydrogen sulphide would resemble the corresponding products of
water in all respects, because water has not the acid properties
of hydrogen sulphide. In the days of dualism and electrical
polarity it was supposed that the sulphur varied in its nature: in
hydrogen sulphide or potassium sulphide it was considered to be
negative, and in sulphurous anhydride or sulphur dichloride
positive. It then appeared evident that sulphur dichloride would
have no point of analogy with potassium sulphide. But metalepsis,
or its expression in the law of substitution, necessitates such
opinions being laid aside. If we can compare CO_{2}, CH_{4},
CCl_{4}, CHCl_{3}, CH_{3}(OH) with each other, we cannot recognise
any difference in the sulphur in SH_{2}, SCl_{2}, SK_{2}, or in
general SX_{2}, for otherwise we should have to acknowledge as
many different states of sulphur, carbon, or hydrogen as there are
compounds of sulphur, carbon, or hydrogen. The essential truth of
the matter is that all the elements in a molecule play their part
in the reactions into which it enters. Often this appears to be
contradicted in the result--for example, hydrogen alone may be
replaced; but it is not this hydrogen alone that has determined
the reaction; all the elements present have participated in it.
This may be made clearer by the following rough illustration.
Supposing two regiments of soldiers were fighting against each
other, and that several men were lost by one of the regiments; no
one could say that it was only these men who took part in the
engagement. The other men fired and the bullets flew over the
heads of their opponents. It was not only those who fell who
fought, although they only were removed from the field of battle;
the fighting proceeded among the masses, but only those few were
disabled who went forward and were more conspicuous &c.; not that
the remainder did not take part in the action; they also fought
and were an object of attack, only they remained sound and unhurt.
Hydrogen is lighter than other elements and its atoms more mobile;
it subjects itself more frequently and easily to reactions; but it
is not it alone which reacts, it is even less liable to attack
than other elements. It participates in exceedingly diverse
reactions, not indeed because the hydrogen itself varies, but
because one atom of it puts itself forward, another is hidden, one
is united with carbon, another feebly held by sulphur, one stands
or moves in the neighbourhood of oxygen, another is joined to a
hydrocarbon. All hydrogen atoms are equal, and equally serve as an
object of attack for the atoms of molecules encountering them, but
those only are removed from the sphere of action which are nearer
the surface of a molecule, which are more mobile, or held by a
less sum of forces. So also sulphur is one and the same in sulphur
dichloride, in sulphurous or sulphuric anhydride, in hydrogen
sulphide, in potassium sulphide, but it reacts differently, and
those elements which are with it also vary in their reactions
because they are with it, and it varies its reactions because it
is with them. It is possible to seize on a character common to
substances quantitatively and qualitatively analogous to each
other. It may be admitted that an element in certain forms is not
able to enter into reactions into which in other forms it enters
willingly, if only the requisite conditions are encountered; but
it must not therefore be concluded that an element changes its
essential quality in these different cases. The preceding remarks
touch on questions which are subject to much argument among
chemists, and I mention them here in order to show the treatment
of those most important problems of chemistry which lie at the
basis of this treatise.
[Illustration: FIG. 91.--Apparatus for the preparation of sulphur
chloride, and similar volatile compounds prepared by combustion in a
stream of chlorine.]
The compounds of sulphur with chlorine are prepared in the apparatus
depicted in fig. 91. As sulphur chloride is decomposed by water, the
chlorine evolved in the flask C must be dried before coming into contact
with the sulphur. It is therefore first passed through a Woulfe's bottle,
B, containing sulphuric acid, and then through the cylinder D containing
pumice stone moistened with sulphuric acid, and then led into the retort
E, in which the sulphur is heated. The compound which is formed distils
over into the receiver R. A certain amount of sulphur passes over with
the sulphur chloride, but if the resultant distillate be re-saturated
with chlorine and distilled no free sulphur remains, the boiling-point
rises to 144°, and pure sulphur chloride, S_{2}Cl_{2}, is obtained. It
has this formula because its vapour density referred to hydrogen is 68.
It is also obtained by heating certain metallic chlorides (stannous,
mercuric) with sulphur; both the metal and chlorine then combine with the
sulphur. Sulphur chloride is a yellowish-brown liquid, which boils at
144°, and has a specific gravity of 1·70 at 0°. It fumes strongly in the
air, reacting on the moisture contained therein, and has a heavy
chloranhydrous odour. It dissolves sulphur, is miscible with carbon
bisulphide, and falls to the bottom of a vessel containing water, by
which it is decomposed, forming sulphurous anhydride and hydrochloric
acid; but it first forms various lower stages of oxidation of sulphur,
because the addition of silver nitrate to the solution gives a black
precipitate. With hydrogen sulphide it gives sulphur and hydrochloric
acid, and it reacts directly with metals--especially arsenic, antimony,
and tin--forming sulphides and chlorides. In the cold, it absorbs
chlorine and gives _sulphur dichloride_, SCl_{2}. The entire conversion
into this substance requires the prolonged passage of dry chlorine
through sulphur chloride surrounded by a freezing mixture. The
distillation of the dichloride must be conducted in a stream of chlorine,
as otherwise it partially decomposes into sulphur chloride and chlorine.
Pure sulphur dichloride is a reddish-brown liquid, which resembles the
lower chloride in many respects; its specific gravity is 1·62; its odour
is more suffocating than that of sulphur chloride; it volatilises at
64°.[77]
[77] The observed vapour density of sulphur dichloride referred to
hydrogen is 53·3, and that given by the formula is 51·5. The
smaller molecular weight explains its boiling point being lower
than that of sulphur chloride, S_{2}Cl_{2}. The reactions of both
these compounds are very similar. Sulphur converts the dichloride,
SCl_{2}, into the monochloride, S_{2}Cl_{2}. In one point the
dichloride differs distinctly from the monochloride--that is, in
its capacity for easily giving up chlorine and decomposing. Even
light decomposes it into chlorine and the monochloride. Hence it
acts on many substances in the same manner as chlorine, or
substances which easily part with the latter, such as phosphoric
or antimonic chloride. In distinction to these, however, sulphur
dichloride would appear to distil without any considerable
decomposition, judging by the vapour density. But this is not a
valid conclusion, for if there be a decomposition, then 2SCl_{2} =
S_{2}Cl_{2} + Cl_{2}; now the density of sulphur chloride = 67·5,
and of chlorine = 35·5, and consequently a mixture of equal
volumes of the two = 51·5, just the same as an equal volume of
sulphur dichloride. _Therefore the distillation of sulphur
dichloride is probably nothing but its decomposition._ Hence the
compound SCl_{2}, which is stable at the ordinary temperature,
decomposes at 64°. In the cold it absorbs a further amount of
chlorine, corresponding to SCl_{4}, but even at -10° a portion of
the absorbed chlorine is given off--that is, dissociation takes
place. Thus the tetrachloride is even less stable than the
dichloride.
_Thionyl chloride_, SOCl_{2}, may be regarded as oxidised sulphur
dichloride; it corresponds with sulphur chloride, S_{2}Cl_{2}, in which
one atom of sulphur is replaced by oxygen. At the same time it is
chlorine oxide (hypochlorous anhydride, Cl_{2}O) combined with sulphur,
and also the chloranhydride of sulphurous acid--that is, SO(HO)_{2}, in
which the two hydroxyl groups are replaced by two atoms of chlorine, or
sulphurous anhydride, SO_{2}, in which one atom of oxygen is replaced by
two atoms of chlorine. All these representations are confirmed by
reactions of formation, or decompositions; they all agree with our
notions of the other compounds of sulphur, oxygen, and chlorine; hence
these definitions are not contradictory to each other. Thus, for
instance, thionyl chloride was first obtained by Schiff, by the action of
dry sulphurous anhydride on phosphorus pentachloride. On distilling the
resultant liquid, thionyl chloride comes over first at 80°, and on
continuing the distillation phosphorus oxychloride distils over at above
100°, PCl_{5} + SO_{2} = POCl_{3} + SOCl_{2}. This mode of preparation is
direct evidence of the oxychloride character of SOCl_{2}. Würtz obtained
the same substance by passing a stream of chlorine oxide through a cold
solution of sulphur in sulphur chloride; the chlorine oxide then combined
directly with the sulphur, S + Cl_{2}O = SOCl_{2}, whilst the sulphur
chloride remained unchanged (sulphur cannot be combined directly with
chlorine oxide, as an explosion takes place). Thionyl chloride is a
colourless liquid, with a suffocating acrid smell; it has a specific
gravity at 0° of 1·675, and boils at 78°. It sinks in water, by which it
is immediately decomposed, like all chloranhydrides--for example, like
carbonyl chloride, which corresponds with it: SOCl_{2} + H_{2}O = SO_{2}
+ 2HCl.[77 bis]
[77 bis] Hartog and Sims (1893) obtained thionyl bromide, SOBr_{2}, by
treating SOCl_{2} with sodium bromide; it is a red liquid, sp. gr.
2·62, and decomposes at 150°.
Normal _sulphuric acid has two corresponding chloranhydrides_; the first,
SO_{2}(OH)Cl, is sulphuric acid, SO_{2}(HO)_{2}, in which one equivalent
of HO is replaced by chlorine; the second has the composition
SO_{2}Cl_{2}--that is, two HO groups are substituted by two of chlorine.
The second chloranhydride, or the compound SO_{2}Cl_{2}, is called
sulphuryl chloride, and the first chloranhydride, SO_{2}HOCl, may be
called chlorosulphonic acid, because it is really an acid; it still
retains one hydroxyl of sulphuric acid, and its corresponding salts are
known. Thus, potassium chloride absorbs the vapour of sulphuric
anhydride, forming a salt, SO_{3}KCl, corresponding with SO_{3}HCl as
acid. In acting on sodium chloride it forms hydrochloric acid and the
salt NaSO_{3}Cl. This first chloranhydride of sulphuric acid, SO_{2}HOCl,
discovered by Williamson, is obtained either by the action of phosphorus
pentachloride on sulphuric acid (PCl_{5} + H_{2}SO_{4} = POCl_{3} + HCl +
HSO_{3}Cl), or directly by the action of dry hydrochloric acid on
sulphuric anhydride, SO_{3} + HCl = HSO_{3}Cl. The most easy and rapid
method of its formation is by direct saturation of cold Nordhausen acid
with dry hydrochloric acid gas (SO_{3} + HCl = HSO_{3}Cl), and
distillation of the resultant solution; the distillate then contains
HSO_{3}Cl. It is a colourless fuming liquid, having an acrid odour; it
boils at 153° (according to my determination, confirmed by Konovaloff),
and its specific gravity at 19° is 1·776. It is immediately decomposed by
water, forming hydrochloric and sulphuric acids, as should be the case
with a true chloranhydride. In the reactions of this chloranhydride we
find the easiest means of introducing the sulphonic group HSO_{3} into
other compounds, because it is here combined with chlorine. The second
chloranhydride of sulphuric acid, or _sulphuryl chloride_, SO_{2}Cl_{2},
was obtained by Regnault by the direct action of the sun's ray on a
mixture of equal volumes of chlorine and sulphurous oxide. The gases
gradually condense into a liquid, combining together as carbonic oxide
does with chlorine. It is also obtained when a mixture of the two gases
in acetic acid is allowed to stand for some time. The first
chloranhydride, SO_{3}HCl, decomposes when heated at 200° in a closed
tube into sulphuric acid and sulphuryl chloride. It boils at 70°, its
specific gravity is 1·7, it gives hydrochloric and sulphuric acids with
water, fumes in the air, and, judging by its vapour density, does not
decompose when distilled.[78]
[78] Pyrosulphuryl chloride, S_{2}O_{5}Cl_{2}. See Note 44. Thorpe and
Kirman, by treating SO_{3} with HF, obtained SO_{2}(OH)F, as a
liquid boiling at 163°, but which decomposed with greater facility
and then gave SO_{2}F_{2}.
The acids of sulphur naturally have their corresponding ammonium
salts, and the latter their amides and nitriles. It will be
readily understood how vast a field for research is presented by
the series of compounds of sulphur and nitrogen, if we only
remember that to carbonic and formic acids there corresponds, as
we saw (Chapter IX.), a vast series of derivatives corresponding
with their ammonium salts. To sulphuric acid there correspond two
ammonium salts, SO_{2}(HO)(NH_{4}O) and SO_{2}(NH_{4}O)_{2}; three
amides: the acid amide SO_{2}(HO)(NH_{2}), or sulphamic acid, the
normal saline compound SO_{2}(NH_{4}O)(NH_{2}), or ammonium
sulphamate, and the normal amide SO_{2}(NH_{2})_{2}, or sulphamide
(the analogue of urea); then the acid nitrile, SON(HO), and two
neutral nitriles, SON(NH_{2}) and SN_{2}. There are similar
compounds corresponding with sulphurous acid, and therefore its
nitriles will be, an acid, SN(HO), its salt, and the normal
compound, SN(NH_{2}). Dithionic and the other acids of sulphur
should also have their corresponding amides and nitriles. Only a
few examples are known, which we will briefly describe. Sulphuric
acid forms salts of very great stability with ammonia, and
ammonium sulphate is one of the commonest ammoniacal compounds. It
is obtained by the direct action of ammonia on sulphuric acid, or
by the action of the latter on ammonium carbonate; it separates
from its solutions in an anhydrous state, like potassium sulphate,
with which it is isomorphous. Hence, the composition of crystals
of ammonium sulphate is (NH_{4})_{2}SO_{4}. This salt fuses at
140°, and does not undergo any change when heated up to 180°. At
higher temperatures it does not lose water, but parts with half
its ammonia, and is converted into the acid salt, HNH_{4}SO_{4};
and this acid salt, on further heating, undergoes a further
decomposition, and splits up into nitrogen, water, and acid
ammonium sulphite, HNH_{4}SO_{3}. At the ordinary temperature the
normal salt is soluble in twice its weight of water and at the
boiling-point of water in an equal weight. In its faculty for
combinations this salt exhibits a great resemblance to potassium
sulphate, and, like it, easily forms a number of double salts; the
most remarkable of which are the ammonia alums,
NH_{4}AlS_{2}O_{8},12H_{2}O, and the double salts formed by the
metals of the magnesium group, having, for example, the
composition (NH_{4})_{2}MgS_{2}O_{8},6H_{2}O. Ammonium sulphate
does not give an amide when heated, perhaps owing to the faculty
of sulphuric anhydride to retain the water combined with it with
great force. But the amides of sulphuric acid may be very
conveniently prepared from sulphuric anhydride. Their formation by
this method is very easily understood because an amide is equal to
an ammonium salt less water, and if the anhydride be taken it will
give an amide directly with ammonia. Thus, if dry ammonia be
passed into a vessel surrounded by a freezing mixture and
containing sulphuric anhydride, it forms a white powdery mass
called sulphatammon, having the composition SO_{3},2H_{3}N, and
resembling the similar compound of carbonic acid, CO_{2},2NH_{3}.
This substance is naturally the ammonium salt of sulphamic acid,
SO_{2}(NH_{4}O)NH_{2}. It is slowly acted on by water, and may
therefore be obtained in solution, in which it slowly reacts with
barium chloride, which proves that with water it still forms
ammonium sulphate. If this substance be carefully dissolved in
water and evaporated, it yields well-formed crystals, whose
solution no longer gives a precipitate with barium chloride. This
is not due to the presence of impurities, but to a change in the
nature of the substance, and therefore Rose calls the crystalline
modification _parasulphatammon_. Platinum chloride only
precipitates half the nitrogen as platinochloride from solutions
of sulphat- and parasulphatammon, which shows that they are
ammonium salts, SO_{2}(NH_{4}O)(NH_{2}). It may be that the reason
of the difference in the two modifications is connected with the
fact that two different substances of the composition
N_{2}H_{4}SO_{2} are possible: one is the amide SO_{2}(NH_{2})_{2}
corresponding with the normal salt, and the other is the salt of
the nitrile acid corresponding with acid ammonium sulphate--that
is, SON(ONH_{4}) corresponds with the acid SON(OH) =
SO_{2}(NH_{4}O)OH - 2H_{2}O. Hence there may here be a difference
of the same nature as between urea and ammonium cyanate. Up to the
present, the isomerism indicated above has been but little
investigated, and might be the subject of interesting researches.
If in the preceding experiment the ammonia, and not the sulphuric
anhydride, be taken in excess, a soluble substance of the
composition 2SO_{2},3NH_{3} is formed. This compound, obtained by
Jacqueline and investigated by Voronin, doubtless also contains a
salt of sulphamic acid--that is, of the amide corresponding with
the acid ammonium sulphate = HNH_{4}SO_{4} - H_{2}O =
(NH_{2})SO_{2}(OH). Probably it is a compound of sulphatammon with
sulphamic acid. Thus it has an acid reaction, and does not give a
precipitate with barium chloride.
With normal sulphate of ammonium, an amide of the composition
N_{2}H_{4}SO_{2} should correspond, which should bear the same
relation to sulphuric acid as urea bears to carbonic acid. This
amide, known as _sulphamide_, is obtained by the action of dry
ammonia on the sulphuryl chloride, SO_{2}Cl_{2}, just as urea is
obtained by the action of ammonia on carbonyl chloride,
SO_{2}Cl_{2} + 4NH_{3} = N_{2}H_{4}SO_{2} + 2NH_{4}Cl. The
ammonium chloride is separated from the resultant sulphamide with
great difficulty. Cold water, acting on the mixture, dissolves
them both; the cold solution does not gives precipitate with
barium chloride. Alkalis act on it slowly, as they do on urea; but
on boiling, especially in the presence of alkalis or acids, it
easily recombines with water, and gives an ammonium salt. V.
Traube (1892) obtained sulphamide by the reaction of sulphuryl,
dissolved in chloroform, upon ammonia. The resultant precipitate
dissolves when shaken up with water, and the solution (after
boiling with the oxides or lead or silver) is evaporated, when a
syrupy liquid remains. With nitrate of silver the latter gives a
solid compound, which, when decomposed by hydrochloric acid, gives
free sulphamide in large colourless crystals, having the
composition SO_{2}(NH_{2})_{2}. This substance fuses at 81°,
begins to decompose below 100°, and is entirely decomposed above
250°; it is soluble in water, and the solution has a neutral
reaction and bitter taste. When heated with acids, sulphamide
gradually decomposes, forming sulphuric acid and ammonia. If the
silver compound obtained by the action of sulphamide on nitrate of
silver be heated at 170°-180° until ammonia is no longer evolved,
and the residue be extracted with water acidulated with nitric
acid, a salt separates out from the solution, answering in its
composition to sulphamide, SO_{2}NAg, which = the amide - NH_{3} =
SO_{2}N_{2}H_{4} - NH_{3} = SO_{2}NH. The action of sulphuryl
chloride (and of the other chloranhydrides of sulphur) on ammonium
carbonate always, as Mente showed (1888), results in the formation
of the salt NH(SO_{3}NH_{4})_{2}.
The nitriles corresponding with sulphuric acid are not as yet
known with any certainty. The most simple nitrile corresponding
with sulphuric acid should have the composition N_{2}H_{8}SO_{4} -
4H_{2}O = N_{2}S. This would be a kind of cyanogen corresponding
with sulphuric acid. On comparing sulphurous acid with carbonic
acid, we saw that they present a great analogy in many respects,
and therefore it might be expected that nitrile compounds having
the composition NHS and N_{2}S_{2} would be found. The latter of
these compounds is well known, and was obtained by Soubeiron, by
the action of dry ammonia on sulphur chloride. This substance
corresponds with cyanogen (paracyanogen), and is known as
_nitrogen sulphide_, N_{2}S_{2}. It is formed according to the
equation 3SCl_{2} + 8NH_{3} = N_{2}S_{2} + S + 6NH_{4}Cl. The free
sulphur and nitrogen sulphide are dissolved by acting on the
product with carbon bisulphide, the nitrogen sulphide being much
less soluble than the sulphur. It is a yellow substance, which is
excessively irritating to the eyes and nostrils. It explodes when
rubbed with a hard substance, being naturally decomposed with the
evolution of nitrogen; but when heated it fuses without
decomposing, and only decomposes with explosion at 157°. It is
insoluble in water, and only slightly so in alcohol, ether, and
carbon bisulphide; 100 parts of the latter dissolve 1·5 part of
nitrogen sulphide at the boiling point. This solution on cooling
deposits it in minute transparent prisms of a golden yellow
colour.
In the group of the halogens we saw four closely analogous
elements--fluorine, chlorine, bromine, and iodine--and we meet with the
same number of closely allied analogues in the oxygen group; for besides
sulphur this group also includes _selenium_ and _tellurium_: O, S, Se,
Te. These two groups are very closely allied, both in respect to the
magnitudes of their atomic weights and also in the faculty of the
elements of both groups for combining with metals. The distinct analogy
and definite degree of variance known to us for the halogens, also repeat
themselves in the same degree for the elements of the oxygen group.
Amongst the halogens fluorine has many peculiarities compared to Cl, Br
and I which are more closely analogous, whilst oxygen differs in many
respects from S, Se, Te, which possess greater similarities. The analogy
in a quantitative respect is perfect in both cases. Thus the halogens
combine with H, and the elements of the oxygen group with H_{2}, forming
H_{2}O, H_{2}S, H_{2}Se, H_{2}Te. The hydrogen compounds of selenium and
tellurium are acids like hydrogen sulphide. Selenium, by simple heating
in a stream of hydrogen, partially combines with it directly, but
seleniuretted hydrogen is more readily decomposable by heat than
sulphuretted hydrogen, and this property is still more developed in
telluretted hydrogen. Hydrogen selenide and telluride are gases like
sulphuretted hydrogen, and, like it, are soluble in water, form saline
compounds with alkalis, precipitate metallic salts, are obtained by the
action of acids on their compounds with metals, &c. Selenium and
tellurium, like sulphur, give two normal grades of combination with
oxygen, both of an acid character, of which only the forms corresponding
to sulphurous anhydride--namely, selenious anhydride, SeO_{2}, and
tellurous anhydride, TeO_{2}[79]--are formed directly. These are both
solids, obtained by the combustion of the elements themselves and by the
action of oxidising agents on them. They form feebly energetic acids,
having distinct bibasic properties; however, a characteristic difference
from SO_{2} is observable both in the physical properties of these
compounds and in their stability and capacity for further oxidation, just
as in the series of the halogens already known to us, only in an inverse
order; in the latter we saw that iodine combines more easily than bromine
or chlorine with oxygen, forming more stable oxygen compounds, whereas
here, on the contrary, sulphurous anhydride, as we know, is difficultly
decomposed, parts with its sulphur with difficulty, and is easily
oxidised and especially in its salts, while selenious and tellurous
anhydrides are oxidised with difficulty and easily reduced, even by means
of sulphurous acid.
[79] _Selenious anhydride_, SeO_{2}, is a volatile solid, which
crystallises in prisms soluble in water. It is best procured by
the action of nitric acid on selenium. The well-known researches
of Nilson (1874) showed that the salts of selenious acid easily
form acid salts, and are so characteristic in many respects that
they may even serve for judging the analogy of types of oxides.
Thus the oxides of the composition RO give normal salts of the
composition RSeO_{3},2H_{2}O, where R = Mn, Co, Ni, Cu, Zn. The
salts of magnesium, barium, and calcium contain a different
quantity of water, as do also the salts of the oxides R_{2}O_{3}.
We here turn attention to the fact that beryllium gives a normal
salt, BeSeO_{3},2H_{2}O, and not a salt analogous to those of
aluminium, scandium, Sc_{2}(SeO_{3})_{3},H_{2}O, yttrium,
Y_{2}(SeO_{3})_{2},12H_{2}O, and other oxides of the form
R_{2}O_{3}, which speaks in favour of the formula BeO.
_Tellurous anhydride_ is also a colourless solid, which
crystallises in octahedra; it also, when heated, first fuses and
then volatilises. It is insoluble in water, and the decomposition
of its salts gives a hydrate, H_{2}TeO_{3}, which is insoluble.
It is a very characteristic circumstance that selenious and
tellurous anhydrides are very easily _reduced_ to selenium and
tellurium. This is not only effected by metals like zinc, or by
sulphuretted hydrogen, which are powerful deoxidisers, but even by
sulphurous anhydride, which is able to precipitate selenium and
tellurium from solutions of the selenites and tellurites, and even
of the acids themselves, which is taken advantage of in obtaining
these elements and separating them from sulphur.
Sulphuric acid, as we know, rarely acts as an oxidising agent. It
is otherwise with selenic and telluric acids, H_{2}SeO_{4} and
H_{2}TeO_{4}, which are powerful oxidising agents--that is, are
easily reduced in many circumstances either into the lower oxide
or even to selenium and tellurium. A powerful oxidising agent is
required in order to convert selenious and tellurous anhydrides
into selenic and telluric anhydrides, and, moreover, it must be
employed in excess. If chlorine be passed through a solution of
potassium selenide, K_{2}Se, telluride, K_{2}Te, selenite,
K_{2}SeO_{3}, or tellurite, K_{2}TeO_{3}, it acts as an oxidiser
in the presence of the water, forming potassium selenate,
K_{2}SeO_{4}, or tellurate, K_{2}TeO_{4}. The same salts are
formed by fusing the lower oxides with nitre. These salts are
isomorphous with the corresponding sulphates, and cannot therefore
be separated from them by crystallisation. The salts of potassium,
sodium, magnesium, copper, cadmium, &c. are soluble like the
sulphates, but those of barium and calcium are insoluble, in
perfect analogy with the sulphates. When copper selenate,
CuSeO_{4}, is treated with sulphuretted hydrogen (CuS is
precipitated), _selenic acid_ remains in solution. On evaporation
and drying in vacuo at 180° it gives a syrupy liquid, which may be
concentrated to almost the pure acid, H_{2}SeO_{4}, having a
specific gravity of 2·6. Cameron and Macallan (1891) showed that
pure H_{2}SeO_{4} only remains liquid in a state of superfusion
whilst the solidified acid melts at +58°, the solid acid
crystallises well, its sp. gr. is then 2·95. The hydrate
H_{2}SeO_{4},H_{2}O melts at +25°. The acid in a superfused state
has a sp. gr. 2·36 and the solid 2·63. Like sulphuric acid strong
selenic acid attracts moisture from the atmosphere; it is not
decomposed by sulphurous acid, but oxidises hydrochloric acid
(like nitric, chromic, and manganic acids), evolving chlorine and
forming selenious acid, H_{2}SeO_{4} + 2HCl = H_{2}SeO_{3} +
H_{2}O + Cl_{2}. _Telluric acid_, H_{2}TeO_{4}, is obtained by
fusing tellurous anhydride with potassium hydroxide and chlorate;
the solution, containing potassium tellurate, is then precipitated
with barium chloride, and the barium tellurate, BaTeO_{4} obtained
in the precipitate is decomposed by sulphuric acid. A solution of
telluric acid is thus obtained, which on evaporation yields
colourless prisms, soluble in water, and containing
TeH_{2}O_{4},2H_{2}O. Two equivalents of water are driven off at
160°; on further heating the last equivalent of water is expelled,
and then oxygen is given off. It also gives chlorine with
hydrochloric acid, like selenic acid. Its salts also correspond
with those of sulphuric acid. It must, however, be remarked that
telluric and selenic acids are able to give poly-acid salts with
much greater ease than sulphuric acid. Thus, for example, there
are known for telluric acid not only K_{2}TeO_{4},5H_{2}O and
KHTeO_{4},3H_{2}O, but also KHTeO_{4},H_{2}TeO_{4},H_{2}O =
K_{2}TeO_{4},3H_{2}TeO_{4},2H_{2}O. This salt is easily obtained
from acid solutions of the preceding salts and is less soluble in
water. As selenious anhydride is volatile and gives similar
poly-salts, it may be surmised that selenious, tellurous, selenic,
and telluric anhydrides are polymeric as compared with sulphurous
and sulphuric anhydrides, for which reason it would be desirable
to determine the vapour density of selenious anhydride. It would
probably correspond with Se_{2}O_{4} or Se_{3}O_{6}.
In order to show the very close analogy of selenium to sulphur, I
will quote two examples. Potassium cyanide dissolves selenium, as
it does sulphur, forming potassium selenocyanate, KCNSe,
corresponding with potassium thiocyanate. Acids precipitate
selenium from this solution, because selenocyanic acid, H_{2}CNSe,
when in a free state is immediately decomposed. A boiling solution
of sodium sulphite dissolves selenium, just as it would sulphur,
forming a salt analogous to thiosulphate of sodium, namely, sodium
selenosulphate, Na_{2}SSeO_{3}. Selenium is separated from a
solution of this salt by the action of acid.
_Selenium_ was obtained in 1817 by Berzelius from the sublimate which
collects in the first chamber in the preparation of sulphuric acid from
Fahlun pyrites. Certain other pyrites also contain small quantities of
selenium. Some native selenides, especially those of lead, mercury, and
copper, have been found in the Hartz Mountains, but only in small
quantities. Pyrites and blendes, in which the sulphur is partially
replaced by selenium, still remain the chief source for its extraction.
When these pyrites are roasted they evolve selenious anhydride, which
condenses in the cooler portions of the apparatus in which the pyrites
are roasted, and is partially or wholly reduced by the sulphurous
anhydride simultaneously formed. The presence of selenium in ores and
sublimates is most simply tested by heating them before the blowpipe,
when they evolve the characteristic odour of garlic. Selenium exhibits
two modifications, like sulphur: one amorphous and insoluble in carbon
bisulphide, the other crystalline and slightly soluble in carbon
bisulphide (in 1,000 parts at 45° and 6,000 at 0°), and separating from
its solutions in monoclinic prisms. If the red precipitate obtained by
the action of sulphurous anhydride on selenious anhydride be dried, it
gives a brown powder, having a specific gravity of 4·26, which when
heated changes colour and fuses to a metallic mass, which gains lustre as
it cools. The selenium acquires different properties according to the
rate at which it is cooled from a fused state; if rapidly cooled, it
remains amorphous and has the same specific gravity (4·28) as the powder,
but if slowly cooled it becomes crystalline and opaque, soluble in carbon
bisulphide, and has a specific gravity of 4·80. In this form it fuses at
214° and remains unchanged, whilst the amorphous form, especially above
80°, gradually passes into the crystalline variety. The transition is
accompanied by the evolution of heat, as in the case of sulphur; thus the
analogy between sulphur and selenium is clearly shown here. In the fused
amorphous form selenium presents a brown mass, slightly translucent, with
a vitreous fracture, whilst in the crystalline form it has the appearance
of a grey metal, with a feeble lustre and a crystalline fracture.[79 bis]
Selenium boils at 700°, forming a vapour whose density is only constant
at a temperature of about 1,400°, when it is equal to 79·4 (referred to
hydrogen)--that is, the molecular formula is then Se_{2}, like sulphur at
an equally high temperature.
[79 bis] Muthmann, in his researches upon the allotropic forms of
selenium, pointed out (1889) a peculiar modification, which
appears, as it were, as a transition between crystalline and
amorphous selenium. It is obtained together with the crystalline
variety by slowly evaporating a solution of selenium in bisulphide
of carbon, and differs from the crystalline variety in the form of
its crystals; it passes into the latter modification when heated.
Schultz also obtained selenium (like Ag, _see_ Chapter XXIV.) in a
soluble form, but these researches are not so conclusive as those
upon soluble silver, and we shall therefore not consider them more
fully.
_Tellurium_ is met with still more rarely than selenium (it is known in
Saxony) in combination with gold, silver, lead, and antimony in the
so-called foliated tellurium ore. Bismuth telluride and silver telluride
have been found in Hungary and in the Altai. Tellurium is extracted from
bismuth telluride by mixing the finely-powdered ore with potassium and
charcoal in as intimate a mixture as possible, and then heating in a
covered crucible. Potassium telluride, K_{2}Te, is then formed, because
the charcoal reduces potassium tellurite. As potassium telluride is
soluble in water, forming a red-brown solution which is decomposed by the
oxygen of the atmosphere (K_{2}Te + O + H_{2}O = 2KHO + Te), the mass
formed in the crucible is treated with boiling water and filtered as
rapidly as possible, and the resultant solution exposed to the air, by
which means the tellurium is precipitated.[80] In a free state tellurium
has a perfectly _metallic appearance_; it is of a silver-white colour,
crystallises very easily in long brilliant needles; is very brittle, so
that it can be easily reduced to powder; but it is a bad conductor of
heat and electricity, and in this respect, as in many others, it forms a
transition from the metals to the non-metals. Its specific gravity is
6·18, it melts at an incipient red heat, and takes fire when heated in
air, like selenium and sulphur, burning with a blue flame, evolving white
fumes of tellurous anhydride, TeO_{2}, and emitting an acrid smell if no
selenium be present; but if it be, the odour of the latter preponderates.
Alkalis dissolve tellurium when boiled with it, potassium telluride,
K_{2}Te, and potassium tellurite, K_{2}TeO_{3}, being formed. The
solution is of a red colour, owing to the presence of the telluride,
K_{2}Te; but the colour disappears when the solution is cooled or
diluted, the tellurium being all precipitated: 2K_{2}Te + K_{2}TeO_{3} +
3H_{2}O = 6KHO + 3Te.[81]
[80] The tellurium thus prepared is impure, and contains a large amount
of selenium. The latter may be removed by converting the mixture
into the salts of potassium, and treating this with nitric acid
and barium nitrate, when barium selenate only is precipitated,
whilst the barium tellurate remains in solution. This method does
not, however, give a pure product, and it appears to be best to
separate the selenium from the tellurium in a metallic form; this
is done by boiling the impure potassium tellurate with
hydrochloric acid, which converts it into potassium tellurite,
from which the tellurium is reduced by sulphurous anhydride. The
metal thus obtained is then fused and distilled in a stream of
hydrogen; the selenium volatilises first, and then the tellurium,
owing to its being much less volatile than the former.
Nevertheless, tellurium is also volatile, and may be separated in
this manner from less volatile metals, such as antimony. Brauner
determined the atomic weight of pure tellurium, and found it to be
125, but showed (1889) that tellurium purified by the usual
method, even after distillation, contains a large amount of
impurities.
[81] The decomposition proceeds in the above order in the cold, but in
a hot solution with an excess of potassium hydroxide it proceeds
inversely. A similar phenomenon takes place when tellurium is
fused with alkalis, and it is therefore necessary in order to
obtain potassium telluride to add charcoal.
Selenium and tellurium form higher compounds with chlorine with
comparative ease. For selenium, SeCl_{2} and SeCl_{4} are known,
and for tellurium TeCl_{2} and TeCl_{4}. The tetrachlorides of
selenium and tellurium are formed by passing chlorine over these
elements. Selenium tetrachloride, SeCl_{4}, is a crystalline,
volatile mass which gives selenious anhydride and hydrochloric
acid with water. Tellurium tetrachloride is much less volatile,
fuses easily, and is also decomposed by water. Both elements form
similar compounds with bromine. Tellurium tetrabromide is red,
fuses to a brown liquid, volatilises, and gives a crystalline
salt, K_{2}TeBr_{6},3H_{2}O, with an aqueous solution of potassium
bromide.
CHAPTER XXI
CHROMIUM, MOLYBDENUM, TUNGSTEN, URANIUM, AND MANGANESE
Sulphur, selenium, and tellurium belong to the uneven series of the sixth
group. In the even series of this group there are known _chromium,
molybdenum, tungsten, and uranium_; these give acid oxides of the type
RO_{3}, like SO_{3}. Their acid properties are less sharply defined than
those of sulphur, selenium, and tellurium, as is the case with all
elements of the even series as compared with those of the uneven series
in the same group. But still the oxides CrO_{3}, MoO_{3}, WO_{3}, and
even UO_{3}, have clearly defined acid properties, and form salts of the
composition MO,_n_RO_{3} with bases MO. In the case of the heavy
elements, and especially of uranium, the type of oxide, UO_{3}, is less
acid and more basic, because in the even series of oxides the element
with the highest atomic weight always acquires a more and more pronounced
basic character. Hence UO_{3} shows the properties of a base, and gives
salts UO_{2}X_{2}. The basic properties of chromium, molybdenum,
tungsten, and uranium are most clearly expressed in the lower oxides,
which they all form. Thus chromic oxide, Cr_{2}O_{3}, is as distinct a
base as alumina, Al_{2}O_{3}.
Of all these elements _chromium_ is the most widely distributed and the
most frequently used. It gives chromic anhydride, CrO_{3}, and chromic
oxide, Cr_{2}O_{3}--two compounds whose relative amounts of oxygen stand
in the ratio 2 : 1. Chromium is, although somewhat rarely, met with in
nature as a compound of one or the other type. The red chromium ore of
the Urals, lead chromate or crocoisite PbCrO_{4}, was the source in which
chromium was discovered by Vauquelin, who gave it this name (from the
Greek word signifying colour) owing to the brilliant colours of its
compounds; the chromates (salts of chromic anhydride) are red and yellow,
and the chromic salts (from Cr_{2}O_{3}) green and violet. The red lead
chromate is, however, a rare chromium ore found only in the Urals and in
a few other localities. Chromic oxide, Cr_{2}O_{3}, is more frequently
met with. In small quantities it forms the colouring matter of many
minerals and rocks--for example, of some serpentines. The commonest ore,
and the chief source of the chromium compounds, is the _chrome iron ore_
or chromite, which occurs in the Urals[1] and Asia Minor, California,
Australia, and other localities. This is magnetic iron ore,
FeO,Fe_{2}O_{3}, in which the ferric oxide is replaced by chromic oxide,
its composition being FeO,Cr_{2}O_{3}. Chrome iron ore crystallises in
octahedra of sp. gr. 4·4; it has a feeble metallic lustre, is of a
greyish-black colour, and gives a brown powder. It is very feebly acted
on by acids, but when fused with potassium acid sulphate it gives a
soluble mass, which contains a chromic salt, besides potassium sulphate
and ferrous sulphate. In practice the treatment of chrome iron ore is
mainly carried on for the preparation of chromates, and not of chromic
salts, and therefore we will trace the history of the element by
beginning with chromic acid, and especially with the working up of the
chrome iron ore into _potassium dichromate_, K_{2}Cr_{2}O_{7}, as the
most common salt of this acid. It must be remarked that chromic
anhydride, CrO_{3}, is only obtained in an anhydrous state, and is
distinguished for its capacity for easily giving anhydro-salts with the
alkalis, containing one, two, and even three equivalents of the anhydride
to one equivalent of base. Thus among the potassium salts there is known
the normal or yellow chromate, K_{2}CrO_{4}, which corresponds to, and is
perfectly isomorphous with, potassium sulphate, easily forms isomorphous
mixtures with it, and is not therefore suitable for a process in which it
is necessary to separate the salt from a mixture containing sulphates. As
in the presence of a certain excess of acid, the dichromate,
K_{2}Cr_{2}O_{7} = 2K_{2}CrO_{4} + 2HX - 2KX - H_{2}O, is easily formed
from K_{2}CrO_{4}, the object of the manufacturer is to produce such a
dichromate, the more so as it contains a larger proportion of the
elements of chromic acid than the normal salt. Finely-ground chrome iron
ore, when heated with an alkali, absorbs oxygen almost as easily (Chapter
III., Note 7) as a mixture of the oxides of manganese with an alkali.
This absorption is due to the presence of chromic oxide, which is
oxidised into the anhydride, and then combines with the alkali
Cr_{2}O_{3} + O_{3} = 2CrO_{3}. As the oxidation and formation of the
chromate proceeds, the mass turns _yellow_. The iron is also oxidised,
but does not give ferric acid, because the capacity of the chromium for
oxidation is incomparably greater than that of the iron.
[1] The working of the Ural chrome iron ore into chromium compounds has
been firmly established in Russia, thanks to the endeavours of P.
K. Ushakoff, who constructed large works for this purpose on the
river Kama, near Elabougi, where as much as 2,000 tons of ore are
treated yearly, owing to which the importation of chromium
preparations into Russia has ceased.
A mixture of lime (sometimes with potash) and chrome iron ore is heated
in a reverberatory furnace, with free access of air and at a red heat for
several hours, until the mass becomes yellow; it then contains normal
calcium chromate, CaCr_O_{4}, which is insoluble in water in the presence
of an excess of lime.[1 bis] The resultant mass is ground up, and treated
with water and sulphuric acid. The excess of lime forms gypsum, and the
soluble calcium dichromate, CaCr_{2}O_{7}, together with a certain amount
of iron, pass into solution. The solution is poured off, and chalk added
to it; this precipitates the ferric oxide (the ferrous oxide is converted
into ferric oxide in the furnace) and forms a fresh quantity of gypsum,
while the chromic acid remains in solution--that is, it does not form the
sparingly-soluble normal salt (1 part soluble in 240 parts of water). The
solution then contains a fairly pure calcium dichromate, which by double
decomposition gives other chromates; for example, with a solution of
potassium sulphate it gives a precipitate of calcium sulphate and a
solution of potassium dichromate, which crystallises when evaporated.[2]
[1 bis] But the calcium chromate is soluble in water in the presence of
an excess of chromic acid, as may be seen from the fact that a
solution of chromic acid dissolves lime.
[2] There are many variations in the details of the manufacturing
processes, and these must be looked for in works on technical
chemistry. But we may add that the chromate may also be obtained by
slightly roasting briquettes of a mixture of chrome iron and lime,
and then leaving the resultant mass to the action of moist air
(oxygen is absorbed, and the mass turns yellow).
_Potassium dichromate_, K_{2}Cr_{2}O_{7}, easily crystallises from acid
solutions in red, well-formed prismatic crystals, which fuse at a red
heat and evolve oxygen at a very high temperature, leaving chromic oxide
and the normal salt, which undergoes no further change: 2K_{2}Cr_{2}O_{7}
= 2K_{2}CrO_{4} + Cr_{2}O_{3} + O_{3}. At the ordinary temperature 100
parts of water dissolve 10 parts of this salt, and the solubility
increases as the temperature rises. It is most important to note that the
dichromate does not contain water, it is K_{2}CrO_{4} + CrO_{3}; the acid
salt corresponding to potassium acid sulphate, KHSO_{4}, does not exist.
It does not even evolve heat when dissolving in water, but on the
contrary produces cold, _i.e._ it does not form a very stable compound
with water. The solution and the salt itself are poisonous, and act as
powerful oxidising agents, which is the character of chromic acid in
general. When heated with sulphur or organic substances, with sulphurous
anhydride, hydrogen sulphide, &c., this salt is deoxidised, yielding
chromic compounds.[2 bis] Potassium dichromate[3] is used in the arts and
in chemistry as a source for the preparation of all other chromium
compounds. It is converted into yellow pigments by means of double
decomposition with salts of lead, barium, and zinc. When solutions of the
salts of these metals are mixed with potassium dichromate (in dyeing
generally mixed with soda, in order to obtain normal salts), they are
precipitated as insoluble normal salts; for example, 2BaCl_{2} +
K_{2}Cr_{2}O_{7} + H_{2}O = 2BaCrO_{4} + 2KCl + 2HCl. It follows from
this that these salts are insoluble in dilute acids, but the
precipitation is not complete (as it would be with the normal salt). The
barium and zinc salts are of a lemon yellow colour; the lead salt has a
still more intense colour passing into orange. Yellow cotton prints are
dyed with this pigment. The silver salt, Ag_{2}CrO_{4}, is of a bright
red colour.
[2 bis] The oxidising action of potassium dichromate on organic
substances at the ordinary temperature is especially marked under
the action of light. Thus it acts on gelatin, as Poutven
discovered; this is applied to photography in the processes of
photogravure, photo-lithography, pigment printing, &c. Under the
action of light this gelatin is oxidised, and the chromic anhydride
deoxidised into chromic oxide, which unites with the gelatin and
forms a compound insoluble in warm water, whilst where the light
has not acted, the gelatin remains soluble, its properties being
unaffected by the presence of chromic acid or potassium dichromate.
[3] Ammonium and sodium dichromates are now also prepared on a large
scale. The sodium salts may be prepared in exactly the same manner
as those of potassium. The normal salt combines with ten
equivalents of water, like Glauber's salt, with which it is
isomorphous. Its solution above 30° deposits the anhydrous salt.
Sodium dichromate crystals contain Na_{2}Cr_{2}O_{7},2H_{2}O. The
_ammonium salts of chromic acid_ are obtained by saturating the
anhydride itself with ammonia. The dichromate is obtained by
saturating one part of the anhydride with ammonia, and then adding
a second part of anhydride and evaporating under the receiver of an
air-pump. On ignition, the normal and acid salts leave chromic
oxide. Potassium ammonium chromate, NH_{4}KCrO_{4}, is obtained in
yellow needles from a solution of potassium dichromate in aqueous
ammonia; it not only loses ammonia and becomes converted into
potassium dichromate when ignited, but also by degrees at the
ordinary temperature. This shows the feeble energy of chromic acid,
and its tendency to form stable dichromates. Magnesium chromate is
soluble in water, as also is the strontium salt. The calcium salt
is also somewhat soluble, but the barium salt is almost insoluble.
The isomorphism with sulphuric acid is shown in the chromates by
the fact that the magnesium and ammonium salts form double salts
containing six equivalents of water, which are perfectly
isomorphous with the corresponding sulphates. The magnesium salt
crystallises in large crystals containing seven equivalents of
water. The beryllium, cerium, and cobalt salts are insoluble in
water. Chromic acid dissolves manganous carbonate, but on
evaporation the solution deposits manganese dioxide, formed at the
expense of the oxygen of the chromic acid. Chromic acid also
oxidises ferrous oxide, and ferric oxide is soluble in chromic
acid.
One of the chromates most used by the dyer is the insoluble yellow
lead chromate, PbCrO_{4} (Chapter XVIII., Note 46), which is
precipitated on mixing solutions of PbX_{2} with soluble chromates.
It easily forms a basic salt, having the composition PbO,PbCrO_{4},
as a crystalline powder, obtained by fusing the normal salt with
nitre and then rapidly washing in water. The same substance is
obtained, although impure and in small quantity, by treating lead
chromate with neutral potassium chromate, especially on boiling the
mixture; and this gives the possibility of attaining, by means of
these materials, various tints of lead chromate, from yellow to
red, passing through different orange shades. The decomposition
which takes place (incompletely) in this case is as follows:
2PbCrO_{4} + K_{2}CrO_{4} = PbCrO_{4},PbO + K_{2}Cr_{2}O_{7}--that
is, potassium dichromate is formed in solution.
When potassium dichromate is mixed with potassium hydroxide or carbonate
(carbonic anhydride being disengaged in the latter case) it forms the
_normal_ salt, K_{2}CrO_{4}, known as _yellow chromate of potassium_. Its
specific gravity is 2·7, being almost the same as that of the dichromate.
It absorbs heat in dissolving; one part of the salt dissolves in 1·75
part of water at the ordinary temperature, forming a yellow solution.
When mixed even with such feeble acids as acetic, and more especially
with the ordinary acids, it gives the dichromate, and Graham obtained a
trichromate, K_{2}Cr_{3}O_{10} = K_{2}CrO_{4},2CrO_{3}, by mixing a
solution of the latter salt with an excess of nitric acid.
_Chromic anhydride_ is obtained by preparing a saturated solution of
potassium dichromate at the ordinary temperature, and pouring it in a
thin stream into an equal volume of pure sulphuric acid.[4] On mixing,
the temperature naturally rises; when slowly cooled, the solution
deposits chromic anhydride in needle-shaped crystals of a red colour
sometimes several centimetres long. The crystals are freed from the
mother liquor by placing them on a porous tile.[4 bis] It is very
important at this point to call attention to the fact that a hydrate of
chromic anhydride is never obtained in the decomposition of chromic
compounds, but always the _anhydride_, CrO_{3}. The corresponding
hydrate, CrO_{4}H_{2}, or any other hydrate, is not even known.
Nevertheless, it must be admitted that chromic acid is bibasic, because
it forms salts isomorphous or perfectly analogous with the salts formed
by sulphuric acid, which is the best example of a bibasic acid. A clear
proof of the bibasicity of CrO_{3} is seen in the fact that the anhydride
and salts give (when heated with sodium chloride and sulphuric acid) a
volatile chloranhydride, CrO_{2}Cl_{2}, containing two atoms of chlorine
as a bibasic acid should.[5] Chromic anhydride is a red crystalline
substance, which is converted into a black mass by heat; it fuses at
190°, and disengages oxygen above 250°, leaving a residue of chromium
dioxide, CrO_{2},[6] and, on still further heating, chromic oxide,
Cr_{2}O_{3}. Chromic anhydride is exceedingly soluble in water, and even
attracts moisture from the air, but, as was mentioned above, it does not
form any definite compound with water. The specific gravity of its
crystals is 2·7, and when fused it has a specific gravity 2·6. The
solution presents perfectly defined acid properties. It liberates
carbonic anhydride from carbonates; gives insoluble precipitates of the
chromates with salts of barium, lead, silver, and mercury.
[4] The sulphuric acid should not contain any lower oxides of nitrogen,
because they reduce chromic anhydride into chromic oxide. If a
solution of a chromate be heated with an excess of acid--for
instance, sulphuric or hydrochloric acid--oxygen or chlorine is
evolved, and a solution of a chromic salt is formed. Hence, under
these circumstances, chromic acid cannot be obtained from its
salts. One of the first methods employed consisted in converting
its salts into volatile _chromium hexafluoride_, CrF_{6}. This
compound, obtained by Unverdorben, may be prepared by mixing lead
chromate with fluor spar in a dry state, and treating the mixture
with fuming sulphuric acid in a platinum vessel: PbCrO_{4} +
3CaF_{2} + 4H_{2}SO_{4} = PbSO_{4} + 3CaSO_{4} + 4H_{2}O + CrF_{6}.
Fuming sulphuric acid is taken, and in considerable excess, because
the chromium fluoride which is formed is very easily decomposed by
water. It is volatile, and forms a very caustic, poisonous vapour,
which condenses when cooled in a dry platinum vessel into a red,
exceedingly volatile liquid, which fumes powerfully in air. The
vapours of this substance when introduced into water are decomposed
into hydrofluoric acid and chromic anhydride: CrF_{6} + 3H_{2}O =
CrO_{3} + 6HF. If very little water be taken the hydrofluoric acid
volatilises, and chromic anhydride separates directly in crystals.
The chloranhydride of chromic acid, CrO_{2}Cl_{2} (Note 5), is also
decomposed in the same manner. A solution of chromic acid and a
precipitate of barium sulphate are formed by treating the insoluble
barium chromate with an equivalent quantity of sulphuric acid. If
carefully evaporated, the solution yields crystals of chromic
anhydride. Fritzsche gave a very convenient method of preparing
chromic anhydride, based on the relation of chromic to sulphuric
acid. At the ordinary temperature the strong acid dissolves both
chromic anhydride and potassium chromate, but if a certain amount
of water is added to the solution the chromic anhydride separates,
and if the amount of water be increased the precipitated chromic
anhydride is again dissolved. The chromic anhydride is almost all
separated from the solution when it contains two equivalents of
water to one equivalent of sulphuric acid. Many methods for the
preparation of chromic anhydride are based on this fact.
[4 bis] They cannot be filtered through paper or washed, because the
chromic anhydride is reduced by the filter-paper, and is dissolved
during the process of washing.
[5] Berzelius observed, and Rose carefully investigated, this
remarkable reaction, which occurs between chromic acid and sodium
chloride in the presence of sulphuric acid. If 10 parts of common
salt be mixed with 12 parts of potassium dichromate, fused, cooled,
and broken up into lumps, and placed in a retort with 20 parts of
fuming sulphuric acid, it gives rise to a violent reaction,
accompanied by the formation of brown fumes of _chromic
chloranhydride_, or _chromyl chloride_, CrO_{2}Cl_{2}, according to
the reaction: CrO_{3} + 2NaCl + H_{2}SO_{4} = Na_{2}SO_{4} + H_{2}O
+ CrO_{2}Cl_{2}. The addition of an excess of sulphuric acid is
necessary in order to retain the water. The same substance is
always formed when a metallic chloride is heated with chromic acid,
or any of its salts, in the presence of sulphuric acid. The
formation of this volatile substance is easily observed from the
brown colour which is proper to its vapour. On condensing the
vapour in a dry receiver a liquid is obtained having a sp. gr. of
1·9, boiling at 118°, and giving a vapour whose density, compared
with hydrogen, is 78, which corresponds with the above formula.
Chromyl chloride is decomposed by heat into chromic oxide, oxygen,
and chlorine: 2CrO_{2}Cl_{2} = Cr_{2}O_{3} + 2Cl_{2} + O; so that
it is able to act simultaneously as a powerful oxidising and
chlorinating agent, which is taken advantage of in the
investigation of many, and especially of organic, substances. When
treated with water, this substance first falls to the bottom, and
is then decomposed into hydrochloric and chromic acids, like all
chloranhydrides: CrO_{2}Cl_{2} + H_{2}O = CrO_{3} + 2HCl. When
brought into contact with inflammable substances it sets fire to
them; it acts thus, for instance, on phosphorus, sulphur, oil of
turpentine, ammonia, hydrogen, and other substances. It attracts
moisture from the atmosphere with great energy, and must therefore
be kept in closed vessels. It dissolves iodine and chlorine, and
even forms a solid compound with the latter, which depends upon the
faculty of chromium to form its higher oxide, Cr_{2}O_{7}. The
close analogy in the physical properties of the chloranhydrides,
CrO_{2}Cl_{2} and SO_{2}Cl_{2}, is very remarkable, although
sulphurous anhydride is a gas, and the corresponding oxide,
CrO_{2}, is a non-volatile solid. It may be imagined, therefore,
that chromium dioxide (which will be mentioned in the following
note) presents a polymerised modification of the substance having
the composition CrO_{2}; in fact, this is obvious from the method
of its formation.
If three parts of potassium dichromate be mixed with four parts of
strong hydrochloric acid and a small quantity of water, and gently
warmed, it all passes into solution, and no chlorine is evolved; on
cooling, the liquid deposits red prismatic crystals, known as
_Peligot's salt_, very stable in air. This has the composition
KCl,CrO_{3}, and is formed according to the equation
K_{2}Cr_{2}O_{7} + 2HCl = 2KCl,CrO_{3} + H_{2}O. It is evident that
this is the first chloranhydride of chromic acid, HCrO_{3}Cl, in
which the hydrogen is replaced by potassium. It is decomposed by
water, and on evaporation the solution yields potassium dichromate
and hydrochloric acid. This is a fresh instance of the reversible
reactions so frequently encountered. With sulphuric acid Peligot's
salt forms chromyl chloride. The latter circumstance, and the fact
that Geuther produced Peligot's salt from potassium chromate and
chromyl chloride, give reason for thinking that it is a compound of
these two substances: 2KCl,CrO_{3} = K_{2}CrO_{4} + CrO_{2}Cl_{2}.
It is also sometimes regarded as potassium dichromate in which one
atom of oxygen is replaced by chlorine--that is,
K_{2}Cr_{2}O_{6}Cl_{2}, corresponding with K_{2}Cr_{2}O_{7}. When
heated it parts with all its chlorine, and on further heating gives
chromic oxide.
[6] This intermediate degree of oxidation, CrO_{2}, may also be
obtained by mixing solutions of chromic salts with solutions of
chromates. The brown precipitate formed contains a compound,
Cr_{2}O_{3},CrO_{3}, consisting of equivalent amounts of chromic
oxide and anhydride. The brown precipitate of chromium dioxide
contains water. The same substance is formed by the imperfect
deoxidation of chromic anhydride by various reducing agents.
Chromic oxide, when heated, absorbs oxygen, and appears to give the
same substance. Chromic nitrate, when ignited, also gives this
substance. When this substance is heated it first disengages water
and then oxygen, chromic oxide being left. It corresponds with
manganese dioxide, Cr_{2}O_{3},CrO_{3} = 3CrO_{2}. Krüger treated
chromium dioxide with a mixture of sodium chloride and sulphuric
acid, and found that chlorine gas was evolved, but that chromyl
chloride was not formed. Under the action of light, a solution of
chromic acid also deposits the brown dioxide. At the ordinary
temperature chromic anhydride leaves a brown stain upon the skin
and tissues, which probably proceeds from a decomposition of the
same kind. Chromic anhydride is soluble in alcohol containing
water, and this solution is decomposed in a similar manner by
light. Chromium dioxide forms K_{2}CrO_{4} when treated with
H_{2}O_{2} in the presence of KHO.
The action of hydrogen peroxide on a solution of chromic acid or of
potassium dichromate gives a blue solution, which very quickly becomes
colourless with the disengagement of oxygen. Barreswill showed that this
is due to the formation of a _perchromic anhydride_, Cr_{2}O_{7},
corresponding with sulphur peroxide. This peroxide is remarkable from the
fact that it very easily dissolves in ether and is much more stable in
this solution, so that, by shaking up hydrogen peroxide mixed with a
small quantity of chromic acid, with ether, it is possible to transfer
all the blue substance formed to the ether.[6 bis]
[6 bis] Now that persulphuric acid H_{2}S_{2}O_{8} is well known it
might be supposed that perchromic anhydride, Cr_{2}O_{7}, would
correspond to perchromic acid, H_{2}Cr_{2}O_{8}, but as yet it is
not certain whether corresponding salts are formed. Péchard (1891)
on adding an excess of H_{2}O_{2} and baryta water to a dilute
solution of CrO_{2} (8 grm. per litre), observed the formation of a
yellow precipitate, but oxygen was disengaged at the same time and
the precipitate (which easily exploded when dried) was found to
contain, besides an admixture of BaO_{2}, a compound BaCrO_{5}, and
this = BaO_{2} + CrO_{3}, and does not correspond to perchromic
acid. The fact of its decomposing with an explosion, and the mode
of its preparation, proves, however, that this is a similar
derivative of peroxide of hydrogen to persulphuric acid (Chapter
XX.)
With oxygen acids, chromic acid evolves oxygen; for example, with
sulphuric acid the following reaction takes place: 2CrO_{3} +
3H_{2}SO_{4} = Cr_{2}(SO_{4})_{3} + O_{3} + 3H_{2}O. It will be readily
understood from this that _a mixture of chromic acid_ or _of its salts
with sulphuric acid_ forms an excellent _oxidising agent_, which is
frequently employed in chemical laboratories and even for technical
purposes as a means of oxidation. Thus hydrogen sulphide and sulphurous
anhydride are converted into sulphuric acid by this means. Chromic acid
is able to act as a powerful oxidising agent because it passes into
chromic oxide, and in so doing disengages half of the oxygen contained in
it: 2CrO_{3} = Cr_{2}O_{3} + O_{3}. Thus chromic anhydride itself is a
powerful oxidising agent, and is therefore employed instead of nitric
acid in galvanic batteries (as a depolariser), the hydrogen evolved at
the carbon being then oxidised, and the chromic acid converted into a
non-volatile product of deoxidation, instead of yielding, as nitric acid
does, volatile lower oxides of offensive odour. Organic substances are
more or less perfectly oxidised by means of chromic anhydride, although
this generally requires the aid of heat, and does not proceed in the
presence of alkalis, but generally _in the presence of acids_. In acting
on a solution of potassium iodide, chromic acid, like many oxidising
agents, liberates iodine; the reaction proceeds in proportion to the
amount of CrO_{3} present, and may serve for determining the amount of
CrO_{3}, since the amount of iodine liberated can be accurately
determined by the iodometric method (Chapter XX., Note 42). If chromic
anhydride be ignited in a stream of ammonia, it gives chromic oxide,
water, and nitrogen. In all cases when chromic acid acts as an oxidising
agent in the presence of acids and under the action of heat, the product
of its deoxidation is a chromic salt, CrX_{3}, which is characterised by
the green colour of its solution, so that the _red_ or yellow _solution_
of a salt of chromic acid is then transformed into a _green solution_ of
a chromic salt, derived from chromic oxide, Cr_{2}O_{3}, which is closely
analogous to Al_{2}O_{3}, Fe_{2}O_{3}, and other bases of the composition
R_{2}O_{3}. This analogy is seen in the insolubility of the anhydrous
oxide, in the gelatinous form of the colloidal hydrate, in the formation
of alums,[7] of a volatile chloride of chromium, &c.[7 bis]
[7] As a mixture of potassium dichromate and sulphuric acid is usually
employed for oxidation, the resultant solution generally contains a
double sulphate of potassium and chromium--that is, _chrome alum_,
isomorphous with ordinary alum--K_{2}Cr_{2}O_{7} + 4H_{2}SO_{4} +
20H_{2}O = O_{3} + K_{2}Cr_{2}(SO_{4})_{4},24H_{2}O or
2(KCr(SO_{4})_{2},12H_{2}O). It is prepared by dissolving potassium
dichromate in dilute sulphuric acid; alcohol is then added and the
solution slightly heated, or sulphurous anhydride is passed through
it. On the addition of alcohol to a cold mixture of potassium
dichromate and sulphuric acid, the gradual disengagement of
pleasant-smelling volatile products of the oxidation of alcohol,
and especially of aldehyde, C_{2}H_{4}O, is remarked. If the
temperature of decomposition does not exceed 35°, a _violet_
solution of chrome alum is obtained, but if the temperature be
higher, a solution of the same alum is obtained of a _green_
colour. As chrome alum requires for solution 7 parts of water at
the ordinary temperature, it follows that if a somewhat strong
solution of potassium dichromate be taken (4 parts of water and
1-1/2 of sulphuric acid to 1 part of dichromate), it will give so
concentrated a solution of chrome alum that on cooling, the salt
will separate without further evaporation. _If the liquid_,
prepared as above or in any instance of the deoxidation of chromic
acid, _be heated_ (the oxidation naturally proceeds more rapidly)
somewhat strongly, for instance, to the boiling-point of water, or
if the violet solution already formed be raised to the same
temperature, it acquires a bright _green colour_, and on
evaporation the same mixture, which at lower temperatures so easily
gives cubical crystals of chrome alum, _does not give any crystals
whatever_. _If the green solution be kept_, however, _for several
weeks_ at the ordinary temperature, it deposits _violet crystals_
of chrome alum. The green solution, when evaporated, gives a
non-crystalline mass, and the violet crystals lose water at 100°
and turn green. It must be remarked that the transition of the
green modification into the violet is accompanied by a decrease in
volume (Lecoq de Boisbaudran, Favre). If the green mass formed at
the higher temperature be evaporated to dryness and heated at 30°
in a current of air, it does not retain more then 6 equivalents of
water. Hence Löwel, and also Schrötter, concluded that the green
and violet modifications of the alum depend on different degrees of
combination with water, which may be likened to the different
compounds of sodium sulphate with water and to the different
hydrates of ferric oxide.
However, the question in this case is not so simple, as we shall
afterwards see. Not chrome alum alone, but _all the chromic salts_,
give two, if not three, _varieties_. At least, there is no doubt
about the existence of two--a _green_ and a _violet modification_.
The green chromic salts are obtained by heating solutions of the
violet salts, the violet solutions are produced on keeping
solutions of the green salts for a long time. The conversion of the
violet salts into green by the action of heat is itself an
indication of the possibility of explaining the different
modifications by their containing different proportions (or states)
of water, and, moreover, by the green salts having a less amount of
water than the violet. However, there are other explanations.
Chromic oxide is a base like alumina, and is therefore able to give
both acid and basic salts. It is supposed that the difference
between the green and violet salts is due to this fact. This
opinion of Krüger is based on the fact that alcohol separates out a
salt from the green solution which contains less sulphuric acid
than the normal violet salt. On the other hand, Löwel showed that
all the acid cannot be separated from the green chromic salts by
suitable reagents, as easily as it can be from the same solution of
the violet salts; thus barium salts do not precipitate all the
sulphuric acid from solutions of the green salts. According to
other researches the cause of the varieties of the chromic salts
lies in a difference in the bases they contain--that is, it is
connected with a modification of the properties of the oxide of
chromium itself. This only refers to the hydroxides, but as
hydroxides themselves are only special forms of salts, the
differences observed as yet in this direction between the
hydroxides only confirm the generality of the difference observed
in the chromic compounds (_see_ Note 7 bis).
The salts of chromic oxide, like those of alumina, are easily
decomposed, give basic and double salts, and have an acid reaction,
as chromic oxide is a feeble base. Potassium and sodium hydroxides
give a _precipitate_ of the hydroxide with chromic salts, CrX_{3}.
The violet and green salts give a _hydroxide soluble in an excess
of the reagent_; but the hydroxide is held in solution by very
feeble affinities, so that it is partially separated by heat and
dilution with water, and completely so on boiling. In an alkaline
solution, chromic hydroxide is easily converted into chromic acid
by the action of lead dioxide, chlorine, and other oxidising
agents. If the chromic oxide occurs together with such oxides as
magnesia, or zinc oxide, then on precipitation it separates out
from its solution in combination with these oxides, forming, for
example, ZnO,Cr_{2}O_{3}. Viard obtained compounds of Cr_{2}O_{3}
with the oxides of Mg, Zn, Cd, &c.) On precipitating the violet
solution of chrome alum with ammonia, a precipitate containing
Cr_{2}O_{3},6H_{2}O is obtained, whilst the precipitate from the
boiling solution with caustic potash was a hydrate containing four
equivalents of water. When fused with borax chromic salts give a
green glass. The same coloration is communicated to ordinary glass
by the presence of traces of chromic oxide. A chrome glass
containing a large amount of chromic oxide may be ground up and
used as a green pigment. Among the hydrates of oxide of chromium
_Guignet's green_ forms one of the widely-used green pigments which
have been substituted for the poisonous arsenical copper pigments,
such as Schweinfurt green, which formerly was much used. Guignet's
green has an extremely bright green colour, and is distinguished
for its great stability, not only under the action of light but
also towards reagents; thus it is not altered by alkaline
solutions, and even nitric acid does not act on it. This pigment
remains unchanged up to a temperature of 250°; it contains
Cr_{2}O_{3},2H_{2}O, and generally a small amount of alkali. It is
prepared by fusing 3 parts of boric acid with 1 part of potassium
dichromate; oxygen is disengaged, and a green glass, containing a
mixture of the borates of chromium and potassium, is obtained. When
cool this glass is ground up and treated with water, which extracts
the boric acid and alkali and leaves the above-named chromic
hydroxide behind. This hydroxide only parts with its water at a red
heat, leaving the anhydrous oxide.
The chromic hydroxides lose their water by ignition, and in so
doing become spontaneously incandescent, like the ordinary ferric
hydroxide (Chapter XXII.). It is not known, however, whether all
the modifications of chromic oxide show this phenomenon. The
anhydrous _chromic oxide_, Cr_{2}O_{3}, is exceedingly difficultly
soluble in acids, if it has passed through the above recalescence.
But if it has parted with its water, or the greater part of it, and
not yet undergone this self-induced incandescence (has not lost a
portion of its energy), then it is soluble in acids. It is not
reduced by hydrogen. It is easily obtained in various crystalline
forms by many methods. The chromates of mercury and ammonium give a
very convenient method for its preparation, because when ignited
they leave chromic oxide behind. In the first instance oxygen and
mercury are disengaged, and in the second case nitrogen and water:
2Hg_{2}CrO_{4} = Cr_{2}O_{3} + O_{5} + 4Hg or
(NH_{4})_{2}Cr_{2}O_{7} = Cr_{2}O_{3} + 4H_{2}O + N_{2}. The second
reaction is very energetic, and the mass of salt burns
spontaneously if the temperature be sufficiently high. A mixture of
potassium sulphate and chromic oxide is formed by heating potassium
dichromate with an equal weight of sulphur: K_{2}Cr_{2}O_{7} + S =
K_{2}SO_{4} + Cr_{2}O_{3}. The sulphate is easily extracted by
water, and there remains a bright green residue of the oxide, whose
colour is more brilliant the lower the temperature of the
decomposition. The oxide thus obtained is used as a green pigment
for china and enamel. The anhydrous chromic oxide obtained from
chromyl chloride, CrO_{2}Cl_{2}, has a specific gravity of 5·21,
and forms almost black crystals, which give a green powder. They
are hard enough to scratch glass, and have a metallic lustre. The
crystalline form of chromic oxide is identical with that of the
oxide of iron and alumina, with which it is isomorphous.
[7 bis] The most important of the compounds corresponding with chromic
oxide is _chromic chloride_, Cr_{2}Cl_{6}, which is known in an
anhydrous and in a hydrated form. It resembles ferric and aluminic
chlorides in many respects. There is a great difference between the
anhydrous and the hydrated chlorides; the former is insoluble in
water, the latter easily dissolves, and on evaporation its solution
forms a hygroscopic mass which is very unstable and easily evolves
hydrochloric acid when heated with water. The anhydrous form is of
a violet colour, and Wöhler gives the following method for its
preparation: an intimate mixture is prepared of the anhydrous
chromic oxide with carbon and organic matter, and charged into a
wide infusible glass or porcelain tube which is heated in a
combustion furnace; one extremity of the tube communicates with an
apparatus generating chlorine which is passed through several
bottles containing sulphuric acid in order to dry it perfectly
before it reaches the tube. On heating the portion of the tube in
which the mixture is placed and passing chlorine through, a
slightly volatile sublimate of chromic chloride, CrCl_{3} or
Cr_{2}Cl_{6}, is formed. This substance forms _violet tabular
crystals_, which may be distilled in dry chlorine without change,
but which, however, require a red heat for their volatilisation.
These crystals are greasy to the touch and insoluble in water, but
if they be powdered and boiled in water for a long time they pass
into _a green solution_. Strong sulphuric acid does not act on the
anhydrous salt, or only acts with exceeding slowness, like water.
Even aqua regia and other acids do not act on the crystals, and
alkalis only show a very feeble action. The specific gravity of the
crystals is 2·99. When fused with sodium carbonate and nitre they
give sodium chloride and potassium chromate, and when ignited in
air they form green chromic oxide and evolve chlorine. On ignition
in a stream of ammonia, chromic chloride forms sal-ammoniac and
chromium nitride, CrN (analogous to the nitrides BN, AlN). Mosberg
and Peligot showed that when chromic chloride is ignited in
hydrogen, it parts with one-third of its chlorine, forming chromous
chloride, CrCl_{2}--that is, there is formed from a compound
corresponding with chromic oxide, Cr_{2}O_{3}, a compound answering
to the _suboxide_, chromous oxide, CrO--just as hydrogen converts
ferric chloride into ferrous chloride with the aid of heat.
_Chromous chloride_, CrCl_{2}, forms colourless crystals easily
soluble in water, which in dissolving evolve a considerable amount
of heat, and form a blue liquid, capable of absorbing oxygen from
the air with great facility, being converted thereby into a chromic
compound.
The blue solution of chromous chloride may also be obtained by the
action of metallic zinc on the green solution of the hydrated
chromic chloride; the zinc in this case takes up chlorine just as
the hydrogen did. It must be employed in a large excess. Chromic
oxide is also formed in the action of zinc on chromic chloride, and
if the solution remain for a long time in contact with the zinc the
whole of the chromium is converted into chromic oxychloride. Other
chromic salts are also reduced by zinc into _chromous salts_,
CrX_{2}, just as the ferric salts FeX_{3} are converted into
ferrous salts FeX_{2} by it. The chromous salts are exceedingly
unstable and easily oxidise and pass into chromic salts; hence the
reducing power of these salts is very great. From cupric salts they
separate cuprous salts, from stannous salts they precipitate
metallic tin, they reduce mercuric salts into mercurous and ferric
into ferrous salts. Moreover, they absorb oxygen from the air
directly. With potassium chromate they give a brown precipitate of
chromium dioxide or of chromic oxide, according to the relative
amounts of the substances taken: CrO_{3} + CrO = 2CrO_{2} or
CrO_{3} + 3CrO = 2Cr_{2}O_{3}. Aqueous ammonia gives a blue
precipitate, and in the presence of ammoniacal salts a blue liquid
is obtained which turns red in the air from oxidation. This is
accompanied by the formation of compounds analogous to those given
by cobalt (Chapter XXII.). A solution of chromous chloride with a
hot saturated solution of sodium acetate, C_{2}H_{3}NaO_{2}, gives,
on cooling, transparent red crystals of chromous acetate,
C_{4}H_{6}CrO_{4},H_{2}O. This salt is also a powerful reducing
agent, but may be kept for a long time in a vessel full of carbonic
anhydride.
The insoluble anhydrous _chromic chloride_ CrCl_{3} very easily
_passes into solution_ in the presence of a trace (0·004) of
_chromous chloride_ CrCl_{2}. This remarkable phenomenon was
observed by Peligot and explained by Löwel in the following manner:
chromous chloride, as a lower stage of oxidation, is capable of
absorbing both oxygen and chlorine, combining with various
substances. It is able to decompose many chlorides by taking up
chlorine from them; thus it precipitates mercurous chloride from a
solution of mercuric chloride, and in so doing passes into chromic
chloride: 2CrCl_{2} + 2HgCl_{2} = Cr_{2}Cl_{6} + 2HgCl. Let us
suppose that the same phenomenon takes place when the anhydrous
chromic chloride is mixed with a solution of chromous chloride. The
latter will then take up a portion of the chlorine of the former,
and pass into a soluble hydrate of chromic chloride (hydrochloride
of oxide of chromium), and the original anhydrous chromic chloride
will pass into chromous chloride. The chromous chloride re-formed
in this manner will then act on a fresh quantity of the chromic
chloride, and in this manner transfer it entirely into solution as
hydrate. This view is confirmed by the fact that other chlorides,
capable of absorbing chlorine like chromous chloride, also induce
the solution of the insoluble chromic chloride--for example,
ferrous chloride, FeCl_{2}, and cuprous chloride. The presence of
zinc also aids the solution of chromic chloride, owing to its
converting a portion of it into chromous chloride. The solution of
chromic chloride in water obtained by these methods is perfectly
identical with that which is formed by dissolving chromic hydroxide
in hydrochloric acid. On evaporating the _green solution_ obtained
in this manner, it gives a green mass, containing water. On further
heating it leaves a soluble chromic oxychloride, and when ignited
it first forms an insoluble oxychloride and then chromic oxide; but
no anhydrous chromic chloride, Cr_{2}Cl_{6}, is formed by heating
the aqueous solution of chromic chloride, which forms an important
fact in support of the view that the green solution of chromic
chloride is nothing else but hydrochloride of oxide of chromium. At
100° the composition of the green hydrate is Cr_{2}Cl_{6},9H_{2}O,
and on evaporation at the ordinary temperature over H_{2}SO_{4}
crystals are obtained with 12 equivalents of water; the red mass
obtained at 120° contains Cr_{2}O_{3},4Cr_{2}Cl_{6},24H_{2}O. The
greater portion of it is soluble in water, like the mass which is
formed at 150°. The latter contains
Cr_{2}O_{3},2Cr_{2}Cl_{6},9H_{2}O = 3(Cr_{2}OCl_{4},3H_{2}O)--that
is, it presents the same composition as chromic chloride in which
one atom of oxygen replaces two of chlorine. And if the hydrate of
chromic chloride be regarded as Cr_{2}O_{3},6HCl, the substance
which is obtained should be regarded as Cr_{2}O_{3},4HCl combined
with water, H_{2}O. The addition of alkalis--for example,
baryta--to a solution of chromic chloride immediately produces a
precipitate, which, however, re-dissolves on shaking, owing to the
formation of one of the oxychlorides just mentioned, which may be
regarded as _basic salts_. Thus we may represent the product of the
change produced on chromic chloride under the influence of water
and heat by the following formulæ: first Cr_{2}O_{3},6HCl or
Cr_{2}Cl_{6},3H_{2}O is formed, then Cr_{2}O_{3},4HCl,H_{2}O or
Cr_{2}OCl_{4},3H_{2}O, and lastly Cr_{2}O_{3},2HCl,2H_{2}O or
Cr_{2}O_{2}Cl_{2},3H_{2}O. In all three cases there are 2
equivalents of chromium to at least 3 equivalents of water. These
compounds may be regarded as being intermediate between chromic
hydroxide and chloride; chromic chloride is Cr_{2}Cl_{6}, the first
oxychloride Cr_{2}(OH)_{2}Cl_{4}, the second Cr_{2}(OH)_{4}Cl_{2},
and the hydrate Cr_{2}(OH)_{6}--that is, the chlorine is replaced
by hydroxyl.
It is very important to remark two circumstances in respect to
this: (1) That the whole of the chlorine in the above compounds is
not precipitated from their solutions by silver nitrate; thus the
normal salt of the composition Cr_{2}Cl_{6},9H_{2}O only gives up
two-thirds of its chlorine; therefore Peligot supposes that the
normal salt contains the oxychloride combined with hydrochloric
acid: Cr_{2}Cl_{6} + 2H_{2}O = Cr_{2}O_{2}Cl_{2},4HCl, and that the
chlorine held as hydrochloric acid reacts with the silver, whilst
that held in the oxychloride does not enter into reaction, just as
we observe a very feebly-developed faculty for reaction in the
anhydrous chromic chloride; and (2) if the green aqueous solution
of CrCl_{3} be left to stand for some time, it ultimately turns
violet; in this form the whole of the chlorine is precipitated by
AgNO_{3}, whilst boiling re-converts it into the green variety.
Löwel obtained the violet solution of hydrochloride of chromic
oxide by decomposing the violet chromic sulphate with barium
chloride. Silver nitrate precipitates all the chlorine from this
violet modification; but if the violet solution be boiled and so
converted into the green modification, silver nitrate then only
precipitates a portion of the chlorine.
Recoura (1890-1893) obtained a crystallohydrate of violet chromium
sulphate, Cr_{2}(SO_{4})_{3}, with 18 or 15 H_{2}O. By boiling a
solution of this crystallohydrate, he converted it into the green
salt, which, when treated with alkalis, gave a precipitate of
Cr_{2}O_{3},2H_{2}O, soluble in 2H_{2}SO_{4} (and not 3), and only
forming the basic salt, Cr_{2}(OH)_{2}(SO_{4})_{2}. He therefore
concludes that the green salts are basic salts. The cryoscopic
determinations made by A. Speransky (1892) and Marchetti (1892)
give a greater 'depression' for the violet than the green salts,
that is, indicate a greater molecular weight for the green salts.
But as Étard, by heating the violet sulphate to 100°, converted it
into a green salt of the same composition, but with a smaller
amount of H_{2}O, it follows that the formation of a basic salt
alone is insufficient to explain the difference between the green
and violet varieties, and this is also shown by the fact that
BaCl_{2} precipitates the whole of the sulphuric acid of the violet
salt, and only a portion of that of the green salt. A. Speransky
also showed that the molecular electro-conductivity of the green
solutions is less than that of the violet. It is also known that
the passage of the former into the latter is accompanied by an
increase of volume, and, according to Recoura, by an evolution of
heat also.
Piccini's researches (1894) throw an important light upon the
peculiarities of the green chromium trichloride (or chromic
chloride); he showed (1) that AgF (in contradistinction to the
other salts of silver) precipitates all the chlorine from an
aqueous solution of the green variety; (2) that solutions of green
CrCl_{3},6H_{2}O in ethyl alcohol and acetone precipitate all their
chlorine when mixed with a similar solution of AgNO_{3}; (3) that
the rise of the boiling-point of the ethyl alcohol and acetone
green solutions of CrCl_{3},6H_{2}O (Chapter VII., Note 27 bis)
shows that i in this case (as in the aqueous solutions of MgSO_{4}
and HgCl_{2}) is nearly equal to 1, that is, that they are like
solutions of non-conductors; (4) that a solution of green CrCl_{3}
in methyl alcohol at first precipitates about 7/8 of its chlorine
(an aqueous solution about 2/3) when treated with AgNO_{3}, but
after a time the whole of the chlorine is precipitated; and (5)
that an aqueous solution of the green variety gradually passes into
the violet, while a methyl alcoholic solution preserves its green
colour, both of itself and also after the whole of the chlorine has
been precipitated by AgNO_{3}. If, however, in an aqueous or methyl
alcoholic solution only a portion of the chlorine be precipitated,
the solution gradually turns violet. In my opinion the general
meaning of all these observations requires further elucidation and
explanation, which should be in harmony with the theory of
solutions. Recoura, moreover, obtained compounds of the green salt,
Cr_{2}(SO_{4})_{3}, with 1, 2, and 3 molecules of H_{2}SO_{4},
K_{2}SO_{4}, and even a compound Cr_{2}(SO_{4})_{3}H_{2}CrO_{4}. By
neutralising the sulphuric acid of the compounds of
Cr_{2}(SO_{4})_{3} and H_{2}SO_{4} with caustic soda, Recoura
obtained an evolution of 33 thousand calories per each 2NaHO, while
free H_{2}SO_{4} only gives 30·8 thousand calories. Recoura is of
opinion that special _chromo sulphuric acids_, for instance
(CrSO_{4})H_{2}SO_{4} = 1/2Cr_{2}(SO_{4})_{3}H_{2}SO_{4}, are
formed. With a still larger excess of sulphuric acid, Recoura
obtained salts containing a still greater number of sulphuric acid
radicles, but even this method does not explain the difference
between the green and violet salts.
These facts must naturally be taken into consideration in order to
arrive at any complete decision as to the cause of the different
modifications of the chromic salts. We may observe that the green
modification of chromic chloride does not give double salts with
the metallic chlorides, whilst the violet variety forms compounds
Cr_{2}Cl_{6},2RCl (where R = an alkali metal), which are obtained
by heating the chromates with an excess of hydrochloric acid and
evaporating the solution until it acquires a violet colour. As the
result of all the existing researches on the green and violet
chromic salts, it appears to me most probable that their difference
is determined by the feeble basic character of chromic oxide, by
its faculty of giving basic salts, and by the colloidal properties
of its hydroxide (these three properties are mutually connected),
and moreover, it seems to me that the relation between the green
and violet salts of chromic oxide best answers to the relation of
the purpureo to the luteo cobaltic salts (Chapter XXII., Note 35).
This subject cannot yet be considered as exhausted (_see_ Note 7).
We may here observe that with tin the chromic salts, CrX_{3}, give
at low temperatures CrX_{2} and SnX_{2}, whilst at high
temperatures, on the contrary, CrX_{2} reduces the metal from its
salts SnX_{2}. The reaction, therefore, belongs to the class of
reversible reactions (Beketoff).
Poulenc obtained anhydrous CrF_{3} (sp. gr. 3·78) and CrF_{2} (sp.
gr. 4·11) by the action of gaseous HF upon CrCl_{2}. A solution of
fluoride of chromium is employed as a mordant in dyeing. Recoura
(1890) obtained green and violet varieties of Cr_{2}Br_{6},6H_{2}O.
The green variety can only be kept in the presence of an excess of
HBr in the solution; if alone its solution easily passes into the
violet variety with evolution of heat.
_Chromic oxide_, Cr_{2}O_{3}, rarely found, and in small quantities,
in chrome ochre, is formed by the oxidation of chromium and its lower
oxides, by the reduction of chromates (for example, of ammonium or
mercuric chromate) and by the decomposition (splitting up) of the saline
compounds of the oxide itself, CrX_{3} or Cr_{2}X_{6}, like alumina,
which it resembles in forming a feeble base easily giving double and
basic salts, which are either green or violet.
The reduction of chromic oxide--for instance, in a solution by zinc
and sulphuric acid--leads to the formation of chromous oxide, CrO, and
its salts, CrX_{2}, of a blue colour (_see_ Notes 7 and 7 ^{bis}). The
further reduction[8] of oxide of chromium and its corresponding compounds
gives _metallic chromium_. Deville obtained it (probably containing
carbon) by reducing chromic oxide with carbon, at a temperature near the
melting point of platinum, about 1750°, but the metal itself does not
fuse at this temperature. Chromium has a steel-grey colour and is very
hard (sp. gr. 5·9), takes a good polish, and dissolves in hydrochloric
acid, but cold dilute sulphuric and nitric acids have no action upon it.
Bunsen obtained metallic chromium by decomposing a solution of chromic
chloride, Cr_{2}Cl_{6}, by a galvanic current, as scales of a grey colour
(sp. gr. 7·3). Wöhler obtained crystalline chromium by igniting a mixture
of the anhydrous chromic chloride Cr_{2}Cl_{6} (_see_ Note 7 bis) with
finely-divided zinc, and sodium and potassium chlorides, at the
boiling-point of zinc. When the resultant mass has cooled the zinc may be
dissolved in dilute nitric acid, and grey crystalline chromium (sp. gr.
6·81) is left behind. Frémy also prepared crystalline chromium by the
action of the vapour of sodium on anhydrous chromic chloride in a stream
of hydrogen, using the apparatus shown in the accompanying drawing, and
placing the sodium and the chromic chloride in separate porcelain boats.
The tube containing these boats is only heated when it is quite full of
dry hydrogen. The crystals of metallic chromium obtained in the tube are
grey cubes having a considerable hardness and withstanding the action of
powerful acids, and even of aqua regia. The chromium obtained by Wöhler
by the action of a galvanic current is, on the contrary, acted on under
these conditions. The reason of this difference must be looked for in the
presence of impurities, and in the crystalline structure. But in any
case, among the properties of metallic chromium, the following may be
considered established: it is white in colour, with a specific gravity of
about 6·7, is extremely hard in a crystalline form, is not oxidised by
air at the ordinary temperature, and with carbon it forms alloys like
cast iron and steel.
[8] The reduction of metallic chromium proceeds with comparative ease
in aqueous solutions. Thus the action of sodium amalgams upon a
strong solution of Cr_{2}Cl_{6} gives (first CrCl_{2}) an amalgam
of chromium from which the mercury may be easily driven off by
heating (in hydrogen to avoid oxidation), and there remains a
spongy mass of easily oxidizable chromium. Plaset (1891), by
passing an electric current through a solution of chrome alum mixed
with a small amount of H_{2}SO_{4} and K_{2}SO_{4}, obtained hard
scales of chromium of a bluish-white colour possessing great
hardness and stability (under the action of water, air, and acids).
Glatzel (1890) reduced a mixture of 2KCl + Cr_{2}Cl_{6} by heating
it to redness with shavings of magnesium. The metallic chromium
thus obtained has the appearance of a fine light-grey powder which
is seen to be crystalline under the microscope; its sp. gr. at 16°
is 6·7284. It fuses (with anhydrous borax) only at the highest
temperatures, and after fusion presents a silver-white fracture.
The strongest magnet has no action upon it.
Moissan (1893) obtained chromium by reducing the oxide Cr_{2}O_{3}
with carbon in the electrical furnace (Chapter VIII., Note 17) in
9-10 minutes with a current of 350 ampères and 50 volts. The
mixture of oxide and carbon gives a bright ingot weighing 100-110
grams. A current of 100 ampères and 50 volts completes the
experiment upon a smaller quantity of material in 15 minutes; a
current of 30 ampères and 50 volts gave an ingot of 10 grams in
30-40 minutes. The resultant carbon alloy is more or less rich in
chromium (from 87·37-91·7 p.c.). To obtain the metal free from
carbon, the alloy is broken into large lumps, mixed with oxide of
chromium, put into a crucible and covered with a layer of oxide.
This mixture is then heated in the electric furnace and the pure
metal is obtained. This reduction can also be carried on with
chrome iron ore FeOCr_{2}O_{3} which occurs in nature. In this case
a homogeneous alloy of iron and chromium is obtained. If this alloy
be thrown in lumps into molten nitre, it forms insoluble
sesquioxide of iron and a soluble alkaline chromate. This alloy of
iron and chromium dissolved in molten steel (chrome steel) renders
it hard and tough, so that such steel has many valuable
applications. The alloy, containing about 3 p.c. Cr and about 1·3
p.c. carbon, is even harder than the ordinary kinds of tempered
steel and has a fine granular fracture. The usual mode of preparing
the ferrochromes for adding to steel is by fusing powdered chrome
iron ore under fluxes in a graphite crucible.
[Illustration: FIG. 92.--Apparatus for the preparation of metallic
chromium by igniting chromic chloride and sodium in a stream of
hydrogen.]
The two analogues of chromium, _molybdenum_ and _tungsten_ (or wolfram),
are of still rarer occurrence in nature, and form acid oxides, RO_{3},
which are still less energetic than CrO_{3}. Tungsten occurs in the
somewhat rare minerals, _scheelite_, CaWO_{4}, and _wolfram_; the latter
being an isomorphous mixture of the normal tungstates of iron and
manganese, (MnFe)WO_{4}. Molybdenum is most frequently met with as
_molybdenite_, MoS_{2}, which presents a certain resemblance to graphite
in its physical properties and softness. It also occurs, but much more
rarely, as a yellow lead ore, PbMoO_{4}. In both these forms molybdenum
occurs in the primary rocks, in granites, gneiss, &c., and in iron and
copper ores in Saxony, Sweden, and Finland. Tungsten ores are sometimes
met with in considerable masses in the primary rocks of Bohemia and
Saxony, and also in England, America, and the Urals. The preliminary
treatment of the ore is very simple; for example, the sulphide, MoS_{2},
is roasted, and thus converted into sulphurous anhydride and molybdic
anhydride, MoO_{3}, which is then dissolved in alkalis, generally in
ammonia. The ammonium molybdate is then treated with acids, when the
sparingly soluble molybdic acid is precipitated. Wolfram is treated in a
different manner. Most frequently the finely-ground ore is repeatedly
boiled with hydrochloric and nitric acids, and the resultant solutions
(of salts of manganese and iron) poured off, until the dark brown mass of
ore disappears, whilst the tungstic acid remains, mixed with silica, as
an insoluble residue; it is treated also with ammonia, and is thus
converted into soluble ammonium tungstate, which passes into solution and
yields tungstic acid when treated with acids. This hydrate is then
ignited, and leaves tungstic anhydride. The general character of molybdic
and tungstic anhydrides is analogous to that of chromic anhydride; they
are anhydrides of a feebly acid character, which easily give polyacid
salts and colloid solutions.[8 bis]
[8 bis] The atomic composition of the tungsten and molybdenum compounds
is taken as being identical with that of the compounds of sulphur
and chromium, because (1) both these metals give two oxides in
which the amounts of oxygen per given amount of metal stand in the
ratio 2 : 3; (2) the higher oxide is of the latter kind, and, like
chromic and sulphuric anhydrides, it has an acid character; (3)
certain of the molybdates are isomorphous with the sulphates; (4)
the specific heat of tungsten is 0·0334, consequently the product
of the atomic weight and specific heat is 6·15, like that of the
other elements--it is the same with molybdenum, 96·0 × 0·0722 =
6·9; (5) tungsten forms with chlorine not only compounds WCl_{6},
WCl_{5}, and WOCl_{4}, but also WO_{2}Cl_{2}, a volatile substance
the analogue of chromyl chloride, CrO_{2}Cl_{2}, and sulphuryl
chloride, SO_{2}Cl_{2}. Molybdenum gives the chlorine compounds,
MOCl_{2}, MOCl_{3}(?), MOCl_{4} (fuses at 194°, boils at 268°;
according to Debray it contains MOCl_{5}), MoOCl_{4},
MoO_{2}Cl_{2}, and MoO_{2}(OH)Cl. The existence of tungsten
hexachloride, WCl_{6}, is an excellent proof of the fact that the
type SX_{6} appears in the analogues of sulphur as in SO_{3}; (6)
the vapour density accurately determined for the chlorine compounds
MoCl_{4}, WCl_{6}, WCl_{5}, WOCl_{4} (Roscoe) leaves no doubt as to
the molecular composition of the compounds of tungsten and
molybdenum, for the observed and calculated results entirely agree.
Tungsten is sometimes called scheele in honour of Scheele, who
discovered it in 1781 and molybdenum in 1778. Tungsten is also
known as wolfram; the former name was the name given to it by
Scheele, because he extracted it from the mineral then known as
tungsten and now called scheelite, CaWO_{4}. The researches of
Roscoe, Blomstrand and others have subsequently thrown considerable
light on the whole history of the compounds of molybdenum and
tungsten.
The ammonium salts of tungsten and molybdic acids when ignited
leave the anhydrides, which resemble each other in many respects.
_Tungsten anhydride_, WO_{3}, is a yellowish substance, which only
fuses at a strong heat, and has a sp. gr. of 6·8. It is insoluble
both in water and acid, but solutions of the alkalis, and even of
the alkali carbonates, dissolve it, especially when heated, forming
alkaline salts. _Molybdic anhydride_, MoO_{3}, is obtained by
igniting the acid (hydrate) or the ammonium salt, and forms a white
mass which fuses at a red heat, and solidifies to a yellow
crystalline mass of sp. gr. 4·4; whilst on further heating in open
vessels or in a stream of air this anhydride _sublimes_ in pearly
scales--this enables it to be obtained in a tolerably pure state.
Water dissolves it in small quantities--namely, 1 part requires 600
parts of water for its solution. The hydrates of molybdic anhydride
are _soluble also in acids_ (a hydrate, H_{2}MoO_{4}, is obtained
from the nitric acid solution of the ammonium salt), which forms
one of their distinctions from the tungstic acids. But after
ignition molybdic anhydride is insoluble in acids, like tungstic
anhydride; alkalis dissolve this anhydride, easily forming
molybdates. Potassium bitartrate dissolves the anhydride with the
aid of heat. None of the acids yet considered by us form so many
different salts with one and the same base (alkali) as molybdic and
tungstic acids. The composition of these salts, and their
properties also, vary considerably. The most important discovery in
this respect was made by Marguerite and Laurent, who showed that
the salts which contain a large proportion of tungstic acid are
easily soluble in water, and ascribed this property to the fact
that tungstic acid may be obtained _in several states_. The common
tungstates, obtained with an excess of alkali, have an alkaline
reaction, and on the addition of sulphuric or hydrochloric acid
first deposit an acid salt and then a hydrate of tungstic acid,
which is insoluble both in water and acids; but if instead of
sulphuric or hydrochloric acids, we add acetic or phosphoric acid,
or if the tungstate be saturated with a fresh quantity of tungstic
acid, which may be done by boiling the solution of the alkali salt
with the precipitated tungstic acid, a solution is obtained which,
on the addition of sulphuric or a similar acid, does not give a
precipitate of tungstic acid at the ordinary or at higher
temperatures. The solution then contains peculiar salts of tungstic
acid, and if there be an excess of acid it also contains tungstic
acid itself; Laurent, Riche, and others called it _metatungstic
acid_, and it is still known by this name. Those salts which with
acids immediately give the insoluble tungstic acid have the
composition R_{2}WO_{4},RHWO_{4}, whilst those which give the
soluble metatungstic acid contain a far greater proportion of the
acid elements. Scheibler obtained the (soluble) metatungstic acid
itself by treating the soluble barium (meta) tetratungstate,
BaO,4WO_{3}, with sulphuric acid. Subsequent research showed the
existence of a similar phenomenon for molybdic acid. There is no
doubt that this is a case of colloidal modifications.
Many chemists have worked on the various salts formed by molybdic
and tungstic acids. The tungstates have been investigated by
Marguerite, Laurent, Marignac, Riche, Scheibler, Anthon, and
others. The molybdates were partially studied by the same chemists,
but chiefly by Struvé and Svanberg, Delafontaine, and others. It
appears that for a given amount of base the salts contain one to
eight equivalents of molybdic or tungstic anhydride; _i.e._ if the
base have the composition RO, then the highest proportion of base
will be contained by the salts of the composition ROWO_{3} or
ROMoO_{3}--that is, by those salts which correspond with the normal
acids H_{2}WO_{4} and H_{2}MoO_{4}, of the same nature as sulphuric
acid; but there also exist salts of the composition RO,2WO_{3},
RO,3WO_{3} ... RO,8WO_{3}. The water contained in the composition
of many of the acid salts is often not taken into account in the
above. The properties of the salts holding different proportions of
acids vary considerably, but one salt may be converted into another
by the addition of acid or base with great facility, and the
greater the proportion of the elements of the acid in a salt, the
more stable, within a certain limit, is its solution and the salt
itself.
The most common ammonium molybdate has the composition
(NH_{4}HO)_{6},H_{2}O,7MoO_{3} (or, according to Marignac and
others, NH_{4}HMoO_{4}), and is prepared by evaporating an
ammoniacal solution of molybdic acid. It is used in the laboratory
for precipitating phosphoric acid, and is purified for this purpose
by mixing its solution with a small quantity of magnesium nitrate,
in order to precipitate any phosphoric acid present, filtering, and
then adding nitric acid and evaporating to dryness. A pure ammonium
molybdate free from phosphoric acid may then be extracted from the
residue.
Phosphoric acid forms insoluble compounds with the oxides of
uranium and iron, tin, bismuth, &c., having feeble basic and even
acid properties. This perhaps depends on the fact that the atoms of
hydrogen in phosphoric acid are of a very different character, as
we saw above. Those atoms of hydrogen which are replaced with
difficulty by ammonium, sodium, &c., are probably easily replaced
by feebly energetic acid groups--that is, the formation of
particular complex substances may be expected to take place at the
expense of these atoms of the hydrogen of phosphoric acid and of
certain feeble metallic acids; and these substances will still be
acids, because the hydrogen of the phosphoric acids and metallic
acids, which is easily replaced by metals, is not removed by their
mutual combination, but remains in the resultant compound. Such a
conclusion is verified in the _phosphomolybdic acids_ obtained
(1888) by Debray. If a solution of ammonium molybdate be acidified,
and a small amount of a solution (it may be acid) containing
orthophosphoric acid or its salts be added to it (so that there are
at least 40 parts of molybdic acid present to 1 part of phosphoric
acid), then after a period of twenty-four hours the whole of the
phosphoric acid is separated as a yellow precipitate, containing,
however, not more than 3 to 4 p.c. of phosphoric anhydride, about 3
p.c. of ammonia, about 90 p.c. of molybdic anhydride, and about 4
p.c. of water. The formation of this precipitate is so distinct and
so complete that this method is employed for the discovery and
separation of the smallest quantities of phosphoric acid.
Phosphoric acid was found to be present in the majority of rocks by
this means. The precipitate is soluble in ammonia and its salts, in
alkalis and phosphates, but is perfectly insoluble in nitric,
sulphuric, and hydrochloric acids in the presence of ammonium
molybdate. The composition of the precipitate appears to vary under
the conditions of its precipitation, but its nature became clear
when the acid corresponding with it was obtained. If the
above-described yellow precipitate be boiled in aqua regia, the
ammonia is destroyed, and an acid is obtained in solution, which,
when evaporated in the air, crystallises out in yellow oblique
prisms of approximately the composition
P_{2}O_{5},20MoO_{3},26H_{2}O. Such an unusual proportion of
component parts is explained by the above-mentioned considerations.
We saw above that molybdic acid easily gives salts
R_{2}O_n_MoO_{3}_m_H_{2}O, which we may imagine to correspond to a
hydrate MoO_{2}(HO)_{2}_n_Mo_{3}_m_H_{2}O. And suppose that such a
hydrate reacts on orthophosphoric acid, forming water and compounds
of the composition MoO_{2}(HPO_{4})_n_MoO_{3}_m_H_{2}O or
MoO_{3}(H_{2}PO_{4})_{2}_n_MoO_{3}_m_H_{2}O; this is actually the
composition of phosphomolybdic acid. Probably it contains a portion
of the hydrogen replaceable by metals of both the acids H_{3}PO_{4}
and of H_{2}MoO_{4}. The crystalline acid above is probably
H_{3}MoPO_{7},9MoO_{3},12H_{2}O. This acid is really tribasic,
because its aqueous solution precipitates salts of potassium,
ammonium, rubidium (but not lithium and sodium) _from acid
solutions_, and gives a _yellow_ precipitate of the composition
R_{3}MoPO_{7},9MoO_{3},8H_{2}O, where R = NH_{4}. Besides these,
salts of another composition may be obtained, as would be expected
from the preceding. These salts are only stable in acid solutions
(which is naturally due to their containing an excess of acid
oxides), whilst under the action of alkalis they give _colourless_
phosphomolybdates of the composition R_{3}MoPO_{3},MoO_{2},3H_{2}O.
The corresponding salts of potassium, silver, ammonium, are easily
soluble in water and crystalline.
Phosphomolybdic acid is an example of the _complex inorganic acids_
first obtained by Marignac and afterwards generalised and studied
in detail by Gibbs. We shall afterwards meet with several examples
of such acids, and we will now turn attention to the fact that they
are usually formed by weak polybasic acids (boric, silicic,
molybdic, &c.), and in certain respects resemble the cobaltic and
such similar complex compounds, with which we shall become
acquainted in the following chapter. As an example we will here
mention certain complex compounds containing molybdic and tungstic
acids, as they will illustrate the possibility of a considerable
complexity in the composition of salts. The action of ammonium
molybdate upon a dilute solution of purpureocobaltic salts (_see_
Chapter XXII.) acidulated with acetic acid gives a salt which after
drying at 100° has the composition
Co_{2}O_{3}10NH_{3}7MoO_{3}3H_{2}O. After ignition this salt leaves
a residue having the composition 2CoO_{7}MoO_{3}. An analogous
compound is also obtained for tungstic acid, having the composition
Co_{2}O_{3}10NH_{3}10WO_{3}9H_{2}O. In this case after ignition
there remains a salt of the composition CoO_{5}WO_{3} (Carnot,
1889). Professor Kournakoff, by treating a solution of potassium
and sodium molybdates, containing a certain amount of suboxide of
cobalt, with bromine obtained salts having the composition:
3K_{2}OCo_{2}O_{3}12MoO_{3}20H_{2}O (light green) and
3K_{2}OCo_{2}O_{3}10Mo_{3}10H_{2}O (dark green). Péchard (1893)
obtained salts of the four complex phosphotungstic acids by
evaporating equivalent mixtures of solutions of phosphoric acid and
metatungstic acid (_see_ further on): phosphotrimetatungstic acid
P_{2}O_{5}12WO_{3}48H_{2}O, phosphotetrametatungstic acid
P_{2}O_{5}16WO_{3}69H_{2}O, phosphopentametatungstic acid
P_{2}O_{5}20WO_{3}H_{2}O, and phosphohexametatungstic acid
P_{2}O_{5}24WO_{3}59H_{2}O. Kehrmann and Frankel described still
more complex salts, such as:
3Ag_{2}O_{4}BaOP_{2}O_{5}22WO_{3}H_{2}O,5BaO_{2}
K_{2}OP_{2}O_{3}22WO_{3}48H_{2}O.
Analogous double salts with 22WO_{3} were also obtained with KSr,
KHg, BaHg, and NH_{4}Pb. Kehrmann (1892) considers the possibility
of obtaining an unlimited number of such salts to be a general
characteristic of such compounds. Mahom and Friedheim (1892)
obtained compounds of similar complexity for molybdic and arsenic
acids.
For tungstic acid there are known: (1) Normal salts--for example,
K_{2}WO_{4}; (2) the so-called acid salts have a composition like
3K_{2}O,7WO_{3},6H_{2}O or K_{6}H_{8}(WO_{4})_{7},2H_{2}O; (3) the
tritungstates like Na_{3}O,3WO_{3},3H_{2}O =
Na_{2}H_{4}(WO_{4})_{3},H_{2}O. All these three classes of salts
are soluble in water, but are precipitated by barium chloride, and
with acids in solution give an insoluble hydrate of tungstic acid;
whilst those salts which are enumerated below do not give a
precipitate either with acids or with the salts of the heavy
metals, because they form soluble salts even with barium and lead.
They are generally called metatungstates. They all contain water
and a larger proportion of acid elements than the preceding salts;
(4) the tetratungstates, like Na_{2}O,4WO_{3},10H_{2}O and
BaO,4WO_{3},9H_{2}O for example; (5) the octatungstates--for
example, Na_{2}O,8WO_{3},24H_{2}O. Since the metatungstates lose so
much water at 100° that they leave salts whose composition
corresponds with an acid, 3H_{2}O,4WO_{3}--that is,
H_{6}W_{4}O_{15}--whilst in the meta salts only 2 hydrogens are
replaced by metals, it is assumed, although without much ground,
that these salts contain a particular soluble metatungstic acid of
the composition H_{6}W_{4}O_{15}.
As an example we will give a short description of the sodium salts.
The normal salt, Na_{2}WO_{4}, is obtained by heating a strong
solution of sodium carbonate with tungstic acid to a temperature of
80°; if the solution be filtered hot, it crystallises in rhombic
tabular crystals, having the composition Na_{2}WO_{4},2H_{2}O,
which remain unchanged in the air and are easily soluble in water.
When this salt is fused with a fresh quantity of tungstic acid, it
gives a ditungstate, which is soluble in water and separates from
its solution in crystals containing water. The same salt is
obtained by carefully adding hydrochloric acid to the solution of
the normal salt so long as a precipitate does not appear, and the
liquid still has an alkaline reaction. This salt was first supposed
to have the composition Na_{2}W_{2}O_{7},4H_{2}O, but it has since
been found to contain (at 100°) Na_{6}W_{7}O_{24},16H_{2}O--that
is, it corresponds with the similar salt of molybdic acid.
(If this salt be heated to a red heat in a stream of hydrogen, it
loses a portion of its oxygen, acquires a metallic lustre, and
turns a golden yellow colour, and, after being treated with water,
alkali, and acid, leaves golden yellow leaflets and cubes which are
very like gold. This very remarkable substance, discovered by
Wöhler, has, according to Malaguti's analysis, the composition
Na_{2}W_{3}O_{9}; that is, it, as it were, contains a double
tungstate of tungsten oxide, WO_{2}, and of sodium,
Na_{2}WO_{4},WO_{2}WO_{3}. The decomposition of the fused sodium
salt is best effected by finely-divided tin. This substance has a
sp. gr. 6·6; it conducts electricity like metals, and like them has
a metallic lustre. When brought into contact with zinc and
sulphuric acid it disengages hydrogen, and it becomes covered with
a coating of copper in a solution of copper sulphate in the
presence of zinc--that is, notwithstanding its complex composition
it presents to a certain extent the appearance and reactions of the
metals. It is not acted on by aqua regia or alkaline solutions, but
it is oxidised when ignited in air.)
The ditungstate mentioned above, deprived of water (having
undergone a modification similar to that of metaphosphoric acid),
after being treated with water, leaves an anhydrous, sparingly
soluble tetratungstate, Na_{2}WO_{4},3WO_{3}, which, when heated at
120° in a closed tube with water, passes into an easily soluble
metatungstate. It may therefore be said that the metatungstates are
hydrated compounds. On boiling a solution of the above-mentioned
salts of sodium with the yellow hydrate of tungstic acid they give
a solution of metatungstate, which is the hydrated tetratungstate.
Its crystals contain Na_{2}W_{4}O_{13},10H_{2}O. After the hydrate
of tungstic acid (obtained from the ordinary tungstates by
precipitation with an acid) has stood a long time in contact with a
solution (hot or cold) of sodium tungstate, it gives a solution
which is not precipitated by hydrochloric acid; this must be
filtered and evaporated over sulphuric acid in a desiccator (it is
decomposed by boiling). It first forms a very dense solution
(aluminium floats in it) of sp. gr. 3·0, and octahedral crystals of
_sodium metatungstate_, Na_{2}W_{4}O_{13},10H_{2}O, sp. gr. 3·85,
then separate. It effloresces and loses water, and at 100° only two
out of the ten equivalents of water remain, but the properties of
the salt remain unaltered. If the salt be deprived of water by
further heating, it becomes insoluble. At the ordinary temperature
one part of water dissolves ten parts of the metatungstate. The
other metatungstates are easily obtained from this salt. Thus a
strong and hot solution, mixed with a like solution of barium
chloride, gives on cooling crystals of barium metatungstate,
BaW_{4}O_{13},9H_{2}O. These crystals are dissolved without change
in water containing hydrochloric acid, and also in hot water, but
they are partially decomposed by cold water, with the formation of
a solution of metatungstic acid and of the normal barium salt
BaWO_{4}.
In order to explain the difference in the properties of the salts
of tungstic acid, we may add that a mixture of a solution of
tungstic acid with a solution of silicic acid does not coagulate
when heated, although the silicic acid alone would do so; this is
due to the formation of a silicotungstic acid, discovered by
Marignac, which presents a fresh example of a complex acid. A
solution of a tungstate dissolves gelatinous silica, just as it
does gelatinous tungstic acid, and when evaporated deposits a
crystalline salt of silicotungstic acid. This solution is not
precipitated either by acids (a clear analogy to the
metatungstates) or by sulphuretted hydrogen, and corresponds with a
series of salts. These salts contain one equivalent of silica and 8
equivalents of hydrogen or metals, in the same form as in salts, to
12 or 10 equivalents of tungstic anhydride; for example the
crystalline potassium salt has the composition
K_{8}W_{12}SiO_{42},14H_{2}O = 4K_{2}O,12WO_{3},SiO_{2},14H_{2}O.
Acid salts are also known in which half of the metal is replaced by
hydrogen. The complexity of the composition of such complex acids
(for example, of the phosphomolybdic acid) involuntarily leads to
the idea of polymerisation, which we were obliged to recognise for
silica, lead oxide, and other compounds. This polymerisation, it
seems to me, may be understood thus: a hydrate A (for example,
tungstic acid) is capable of combining with a hydrate B (for
example, silica or phosphoric acid, with or without the
disengagement of water), and by reason of this faculty it is
capable of polymerisation--that is, A combines with A--combines
with itself--just as aldehyde, C_{2}H_{4}O, or the cyanogen
compounds are able to combine with hydrogen, oxygen, &c., and are
liable to polymerisation. On this view the molecule of tungstic
acid is probably much more complex than we represent it; this
agrees with the easy volatility of such compounds as the
chloranhydrides, CrO_{2}Cl_{2}, MoO_{2}Cl_{2}, the analogues of the
volatile sulphuryl chloride, SO_{2}Cl_{2}, and with the
non-volatility, or difficult volatility, of chromic and molybdic
anhydrides, the analogues of the volatile sulphuric anhydride. Such
a view also finds a certain confirmation in the researches made by
Graham on the _colloidal_ state of tungstic acid, because colloidal
properties only appertain to compounds of a very complex
composition. The observations made by Graham on the colloidal state
of tungstic and molybdic acids introduced much new matter into the
history of these substances. When sodium tungstate, mixed in a
dilute solution with an equivalent quantity of dilute hydrochloric
acid, is placed in a dialyser, hydrochloric acid and sodium
chloride pass through the membrane, and a solution of tungstic acid
remains in the dialyser. Out of 100 parts of tungstic acid about 80
parts remain in the dialyser. The solution has a bitter, astringent
taste, and does not yield gelatinous tungstic acid (hydrogel)
either when heated or on the addition of acids or salts. It may
also be evaporated to dryness; it then forms a vitreous mass of the
_hydrosol_ of _tungstic acid_, which adheres strongly to the walls
of the vessel in which it has been evaporated, and is perfectly
soluble in water. It does not even lose its solubility after having
been heated to 200°, and only becomes insoluble when heated to a
red heat, when it loses about 2-1/2 p.c. of water. The dry acid,
dissolved in a small quantity of water, forms a gluey mass, just
like gum arabic, which is one of the representatives of the
hydrosols of colloidal substances. The solution, containing 5 p.c
of the anhydride, has a sp. gr. of 1·047; with 20 p.c., of 1·217;
with 50 p.c., of 1·80; and with 80 p.c., of 3·24. The presence of a
polymerised trioxide in the form of hydrate, H_{2}OW_{3}O_{9} or
H_{2}O_{4}WO_{3}, must then be recognised in the solution: this is
confirmed by Sabaneeff's cryoscopic determinations (1889). A
similar stable solution of molybdic acid is obtained by the
dialysis of a mixture of a strong solution of sodium molybdate with
hydrochloric acid (the precipitate which is formed is
re-dissolved). If MoCl_{4} be precipitated by ammonia and washed
with water, a point is reached at which perfect solution takes
place, and the molybdic acid forms a colloid solution which is
precipitated by the addition of ammonia (Muthmann). The addition of
alkali to the solutions of the hydrosols of tungstic and molybdic
acids immediately results in the re-formation of the ordinary
tungstates and molybdates. There appears to be no doubt but that
the same transformation is accomplished in the passage of the
ordinary tungstates into the metatungstates as takes place in the
passage of tungstic acid itself from an insoluble into a soluble
state; but this may be even actually proved to be the case, because
Scheibler obtained a solution of tungstic acid, before Graham, by
decomposing barium metatungstate (BaO_{4}WO_{3},9H_{2}O) with
sulphuric acid. By treating this salt with sulphuric acid in the
amount required for the precipitation of the baryta, Scheibler
obtained a solution of metatungstic acid which, when containing
43·75 p.c. of acid, had a sp. gr. of 1·634, and with 27·61 p.c. a
sp. gr. of 1·327--that is, specific gravities corresponding with
those found by Graham.
Péchard found that as much heat is evolved by neutralising
metatungstic acid as with sulphuric acid.
Questions connected with the metamorphoses or modifications of
tungstic and molybdic acids, and the polymerisation and colloidal
state of substances, as well as the formation of complex acids,
belong to that class of problems the solution of which will do much
towards attaining a true comprehension of the mechanism of a number
of chemical reactions. I think, moreover, that questions of this
kind stand in intimate connection with the theory of the formation
of solutions and alloys and other so-called indefinite compounds.
Hydrogen (which does not directly form compounds with Cr, Mo, and W)
reduces molybdic and tungstic anhydride at a red heat; and this forms the
means of obtaining metallic molybdenum and tungsten. _Both metals_ are
infusible, and both under the action of heat form compounds with carbon
and iron (the addition of tungsten to steel renders the latter ductile
and hard).[9] Molybdenum forms a grey powder, which scarcely aggregates
under a most powerful heat, and has a specific gravity of 8·5. It is not
acted on by the air at the ordinary temperature, but when ignited it is
first converted into a brown, and then into a blue oxide, and lastly into
molybdic anhydride. Acids do not act on it--that is, it does not liberate
hydrogen from them, not even from hydrochloric acid--but strong sulphuric
acid disengages sulphurous anhydride, forming a brown mass, containing a
lower oxide of molybdenum. Alkalis in solution do not act on molybdenum,
but when fused with it hydrogen is given off, which shows, as does its
whole character, the acid properties of the metal. The properties of
tungsten are almost identical; it is infusible, has an iron-grey colour,
is exceedingly hard, so that it even scratches glass. Its specific
gravity is 19·1 (according to Roscoe), so that, like uranium, platinum,
&c., it is one of the heaviest metals.[9 bis] Just as sulphur and
chromium have their corresponding persulphuric and perchromic acids,
H_{2}S_{2}O_{8} and H_{2}CrO_{8}, having the properties of peroxides, and
corresponding to peroxide of hydrogen, so also molybdenum and tungsten
are known to give _permolybdic_ and _pertungstic_ acids, H_{2}Mo_{2}O_{8}
and H_{2}W_{2}O_{8}, which have the properties of true peroxides, _i.e._
easily disengage iodine from KI and chlorine from HCl, easily part with
their oxygen, and are formed by the action of peroxide of hydrogen, into
which they are readily reconverted (hence they may be regarded as
compounds of H_{2}O_{2} with 2MoO_{3} and 2WO_{3}), &c. Their formation
(Boerwald 1884, Kemmerer 1891) is at once seen in the coloration (not
destroyed by boiling), which is obtained on mixing a solution of the
salts with peroxide of hydrogen, and on treating, for instance, molybdic
acid with a solution of peroxide of hydrogen (Péchard 1892). The acid
then forms an orange-coloured solution, which after evaporation in vacuo
leaves Mo_{2}H_{2}O_{8}4H_{2}O as a crystalline powder, and loses 4H_{2}O
at 100°, beyond which it decomposes with the evolution of oxygen.[9 tri]
[9] Moissan (1893) studied the compounds of Mo and W formed with carbon
in the electrical furnace (they are extremely hard) from a mixture
of the anhydrides and carbon. Poleck and Grützner obtained definite
compounds FeW_{2} and FeW_{2}C_{3} for tungsten. Metallic W and Mo
displace Ag from its solutions but not Pb. There is reason for
believing that the sp. gr. of pure molybdenum is higher than that
(8·5) generally ascribed to it.
[9 bis] We may conclude our description of tungsten and molybdenum by
stating that their sulphur compounds have an acid character, like
carbon bisulphide or stannic sulphide. If sulphuretted hydrogen be
passed through a solution of a molybdate it does not give a
precipitate unless sulphuric acid be present, when a dark brown
precipitate of _molybdenum trisulphide_, MoS_{3}, is formed. When
this sulphide is ignited without access of air it gives the
bisulphide MoS_{2}; the latter is not able to combine with
potassium sulphide like the trisulphide MoS_{3}, which forms a
salt, K_{2}MoS_{4}, corresponding with K_{2}MoO_{4}. This is
soluble in water, and separates out from its solution in red
crystals, which have a metallic lustre and reflect a green light.
It is easily obtained by heating the native bisulphide, MoS_{2},
with potash, sulphur, and a small amount of charcoal, which serves
for deoxidising the oxygen compounds. Tungsten gives similar
compounds, R_{2}WS_{4}, where R = NH_{4}, K, Na. They are
decomposed by acids, with the separation of tungsten trisulphide,
WS_{3}, and molybdenum trisulphide, MoS_{3}. Rideal (1892) obtained
W_{2}N_{3} by heating WO_{3} in NH_{3}. This compound exhibited the
general properties of metallic nitrides.
[9 tri] When peroxide of hydrogen acts upon a solution of potassium
molybdate well-formed yellow crystals belonging to the triclinic
system separate out in the cold. When these crystals are heated in
vacuo they first lose water and then decompose, leaving a residue
composed of the salt originally taken. They are soluble in water
but insoluble in alcohol. Their composition is represented by the
formula K_{2}Mo_{2}O_{8}2H_{2}O. An ammonium salt is obtained by
evaporating peroxide of hydrogen with ammonium molybdate. The
following salts have also been obtained by the action of peroxide
of hydrogen upon the corresponding molybdates:
Na_{2}Mo_{2}O_{6}6H_{2}O--in yellow prismatic crystals;
MgMo_{2}O_{8}10H_{2}O--stellar needles; BaMoO_{8}2H_{2}O--in
microscopic yellow octahedra. A corresponding sodium pertungstate
has been obtained by Péchard by boiling sodium tungstate with a
solution of peroxide of hydrogen for several minutes. The solution
rapidly turns yellow, and no longer gives a precipitate of tungstic
anhydride when treated with nitric acid. When evaporated in vacuo
the solution leaves a thick syrupy liquid from which ray-like
crystals separate out; these crystals are more soluble in water
than the salt originally taken. When heated they also lose water
and oxygen. Their composition answers to the formula
M_{2}W_{2}O_{8}2H_{2}O, where M = Na, NH_{4}, &c. The permolybdates
and pertungstates have similar properties. When treated with oxygen
acids they give peroxide of hydrogen, and disengage chlorine and
iodine from hydrochloric acid and potassium iodide.
Piccini (1891) showed that peroxide of hydrogen not only combines
with the oxygen compounds of Mo and W, but also with their
fluo-compounds, among which ammonium fluo-molybdate
MoO_{2}F_{2}2NH_{4} and others have long been known. (A few new
salts of similar composition have been obtained by F. Moureu in
1893.) The action of peroxide of hydrogen upon these compounds
gives salts containing a larger amount of oxygen; for instance, a
solution of MoO_{2}F_{2}2KFH_{2}O with peroxide of hydrogen gives a
yellow solution which after cooling separates out yellow
crystalline flakes of MoO_{3}F_{2}2KFH_{2}O, resembling the salt
originally taken in their external appearance. By employing a
similar method Piccini also obtained:
MoO_{3}F_{2}2RbFH_{2}O--yellow monoclinic crystals;
MoO_{3}F_{2},2CsFH_{2}O,--yellow flakes, and the corresponding
tungstic compounds. All these salts react like peroxide of
hydrogen.
In speaking of these compounds I for my part think it may be well
to call attention to the fact that, in the first place, the
composition of Piccini's oxy-fluo compounds does not correspond to
that of permolybdic and pertungstic acid. If the latter be
expressed by formulæ with one equivalent of an element, they will
be HMoO_{4} and HWO_{4}, and the oxy-fluo form corresponding to
them should have the composition MoO_{3}F and WO_{3}F while it
contains MO_{3}F_{2} and WO_{3}F_{2}, _i.e._ answers as it were to
a higher degree of oxidation, MoH_{2}O_{3} and W_{3}HO_{3}. But if
permolybdic acid be regarded as 2MoO_{3} + H_{2}O_{2}, _i.e._ as
containing the elements of peroxide of hydrogen, then Piccini's
compound will also be found to contain the original salts + H_{2}O;
for example, from MoO_{2}F_{2}2KFH_{2}O there is obtained a
compound MoO_{2}F_{2}2KFH_{2}O_{2}, _i.e._ instead of H_{2}O they
contain H_{2}O_{2}. In the second place the capacity of the salts
of molybdenum and tungsten to retain a further amount of oxygen or
H_{2}O_{2} probably bears some relation to their property of giving
complex acids and of polymerising which has been considered in Note
8 bis. There is, however, a great chemical interest in the
accumulation of data respecting these high peroxide compounds
corresponding to molybdic and tungstic acids. With regard to the
peroxide form of uranium, _see_ Chapter XX., Note 66.
_Uranium_, U = 240, has the highest atomic weight of all the analogues
of chromium, and indeed of all the elements yet known. Its highest
salt-forming oxide, UO_{3}, shows very feeble acid properties. Although
it gives sparingly-soluble yellow compounds with alkalis, which fully
correspond with the dichromates--for example, Na_{2}U_{2}O_{7} =
Na_{2}O,2UO_{3},[10]--yet it more frequently and easily reacts with
acids, HX, forming fluorescent yellowish-green salts of the composition
UO_{2}X_{2}, and in this respect uranic trioxide, UO_{3}, differs from
chromic anhydride, CrO_{3}, although the latter is able to give the
oxychloride, CrO_{2}Cl_{2}. In molybdenum and tungsten, however, we see a
clear transition from chromium to uranium. Thus, for example, chromyl
chloride, CrO_{2}Cl_{2}, is a brown liquid which volatilises without
change, and is completely decomposed by water; molybdenum oxychloride,
MoO_{2}Cl_{2}, is a crystalline substance of a yellow colour, which is
volatile and soluble in water (Blomstrand), like many salts. Tungsten
oxychloride, WO_{2}Cl_{2}, stands still nearer to uranyl chloride in its
properties; it forms yellow scales on which water and alkalis act, as
they do on many salts (zinc chloride, ferric chloride, aluminium
chloride, stannic chloride, &c.), and perfectly corresponds with the
difficultly volatile salt, UO_{2}Cl_{2} (obtained by Peligot by the
action of chlorine on ignited uranium dioxide, UO_{2}), which is also
yellow and gives a yellow solution with water, like all the salts
UO_{2}X_{2}. The property of uranic oxide, UO_{3}, of forming salts
UO_{2}X_{2} is shown in the fact that the hydrated oxide of uranium,
UO_{2}(HO)_{2}, which is obtained from the nitrate, carbonate, and other
salts by the loss of the elements of the acid, is easily soluble in
acids, as well as in the fact that the lower grades of oxidation of
uranium are able, when treated with nitric acid, to form an easily
crystallisable uranyl nitrate, UO_{2}(NO_{3})_{2},6H_{2}O; this is the
most commonly occurring uranium salt.[11]
[10] Uranium trioxide, or uranic oxide, shows its feeble basic and acid
properties in a great number of its reactions. (1) Solutions of
uranic salts give yellow precipitates with alkalis, but these
precipitates do not contain the hydrate of the oxide, but
compounds of it with bases; for example, 2UO_{2}(NO_{3})_{2} +
6KHO = 4KNO_{3} + 3H_{2}O + K_{2}U_{2}O_{7}. There are other
_urano-alkali compounds_ of the same constitution; for example,
(NH_{4})_{2}U_{2}O_{7} (known commercially as uranic oxide),
MgU_{2}O_{7}, BaU_{2}O_{7}. They are the analogues of the
dichromates. Sodium uranate is the most generally used under the
name of uranium yellow, Na_{2}U_{2}O_{7}. It is used for imparting
the characteristic yellow-green tint to glass and porcelain.
Neither heat nor water nor acids are able to extract the alkali
from sodium uranate, Na_{2}U_{2}O_{7}, and therefore it is a true
insoluble salt, of a yellow colour, and clearly indicates the acid
character (although feeble) of uranic oxide. (2) The carbonates of
the alkaline earths (for instance, barium carbonate) precipitate
uranic oxide from its salts, as they do all the salts of feeble
bases; for example, R_{2}O_{3}. (3) The _alkaline carbonates_,
when added to solutions of uranic salts, give a _precipitate,
which is soluble in_ _an excess of the reagent_, and particularly
so if the acid carbonates be taken. This is due to the fact that
(4) the uranyl salts _easily form double salts_ with the salts of
the alkali metals, including the salts of ammonium. Uranium, in
the form of these double salts, often gives salts of well-defined
crystalline form, although the simple salts are little prone to
appear in crystals. Such, for example, are the salts obtained by
dissolving potassium uranate, K_{2}U_{2}O_{7}, in acids, with the
addition of potassium salts of the same acids. Thus, with
hydrochloric acid and potassium chloride a well-formed crystalline
salt, K_{2}(UO_{2})Cl_{4},2H_{2}O, belonging to the monoclinic
system, is produced. This salt decomposes in dissolving in pure
water. Among these double salts we may mention the double
carbonate with the alkalis, R_{4}(UO_{2})(CO_{3})_{3} (equal to
2R_{2}CO_{3} + UO_{2}CO_{3}); the acetates,
R(UO_{2})(C_{2}H_{3}O_{2})_{3}--for instance, the sodium salt,
Na(UO_{2})(C_{2}H_{3}O_{2})_{3}, and the potassium salt,
K(UO_{2})(C_{2}H_{3}O_{2})_{3},H_{2}O; the sulphates,
R_{2}(UO_{2})(SO_{4})_{3},2H_{2}O, &c. In the preceding formula R
= K, Na, NH_{4}, or R_{2} = Mg, Ba, &c. _This property of giving
comparatively stable double salts indicates feebly developed basic
properties_, because double salts are mainly formed by salts of
distinctly basic metals (these form, as it were, the basic element
of a double salt) and salts of feebly energetic bases (these form
the acid element of a double salt), just as the former also give
acid salts; the acid of the acid salts is replaced in the double
salts by the salt of the feebly energetic base, which, like water,
belongs to the class of intermediate bases. For this reason barium
does not give double salts with alkalis as magnesium does, and
this is why double salts are more easily formed by potassium than
by lithium in the series of the alkali metals. (5) The most
remarkable property, proving the feeble energy of uranic oxide as
a base, is seen in the fact that when their composition is
compared with that of other salts those of uranic oxide _always
appear as basic salts_. It is well known that a normal salt,
R_{2}X_{6}, corresponds with the oxide R_{2}O_{3}, where X = Cl,
NO_{3}, &c., or X_{2} = SO_{4}, CO_{3}, &c.; but there also exist
basic salts of the same type where X = HO or X_{2} = O. We saw
salts of all kinds among the salts of aluminium, chromium, and
others. With uranic oxide no salts are known of the types UX_{6}
(UCl_{6}, U(SO_{4})_{3}, alums, &c., are not known), nor even
salts, U(HO)_{2}X_{4} or UOX_{4}, but it always forms salts of the
type U(HO)_{4}X_{2}, or UO_{2}X_{2}. Judging from the fact that
nearly all the salts of uranic oxide retain water in crystallising
from their solutions, and that this water is difficult to separate
from them, it may be thought to be water of hydration. This is
seen in part from the fact that the composition of many of the
salts of uranic oxide may then be expressed without the presence
of water of crystallisation; for instance, U(HO)_{4}K_{2}Cl_{4}
(and the salt of NH_{4}, U(HO)_{4}K_{2}(SO_{4})_{2},
U(HO)_{4}(C_{2}H_{3}O_{2})_{2}. Sodium uranyl acetate however does
not contain water.
[11] _Uranyl nitrate_, or uranium nitrate, UO_{2}(NO_{3})_{2},6H_{2}O,
crystallises from its solutions in transparent yellowish-green
prisms (from an acid solution), or in tabular crystals (from a
neutral solution), which effloresce in the air and are easily
soluble in water, alcohol, and ether, have a sp. gr. of 2·8, and
fuse when heated, losing nitric acid and water in the process. If
the salt itself (Berzelius) or its alcoholic solution (Malaguti)
be heated up to the temperature at which oxides of nitrogen are
evolved, there then remains a mass which, after being evaporated
with water, leaves uranyl hydroxide, UO_{2}(HO)_{2} (sp. gr.
5·93), whilst if the salt be ignited there remains the dioxide,
UO_{2}, as a brick-red powder, which on further heating loses
oxygen and forms the dark olive uranoso-uranic oxide, U_{3}O_{8}.
The solution of the nitrate obtained from the ore is purified in
the following manner: sulphurous anhydride is first passed through
it in order to reduce the arsenic acid present into arsenious
acid; the solution is then heated to 60°, and sulphuretted
hydrogen passed through it; this precipitates the lead, arsenic,
and tin, and certain other metals, as sulphides, insoluble in
water and dilute nitric acid. This liquid is then filtered and
evaporated with nitric acid to crystallisation, and the crystals
are dissolved in ether. Or else the solution is first treated with
chlorine in order to convert the ferrous chloride (produced by the
action of the hydrogen sulphide) into ferric chloride, the oxides
are then precipitated by ammonia, and the resultant precipitate,
containing the oxides Fe_{2}O_{3}, UO_{3}, and compounds of the
latter with potash, lime, ammonia, and other bases present in the
solution (the latter being due to the property of uranic oxide of
combining with bases), is washed and dissolved in a strong,
slightly-heated solution of ammonium carbonate, which dissolves
the uranic oxide but not the ferric oxide. The solution is
filtered, and on cooling deposits a well-crystallising _uranyl
ammonium carbonate_, UO_{2}(NH_{4})_{4}(CO_{3})_{3}, in brilliant
monoclinic crystals which on exposure to air slowly give off
water, carbonic anhydride, and ammonia; the same decomposition is
readily effected at 300°, the residue then consisting of uranic
oxide. This salt is not very soluble in water, but is readily so
in ammonium carbonate; it is obvious that it may readily be
converted into all the other salts of oxides of uranium. Uranium
salts are also purified in the form of _acetate_, which is very
sparingly soluble, and is therefore directly precipitated from a
strong solution of the nitrate by mixing it with acetic acid.
We may also mention the _uranyl phosphate_, HUPO_{6}, which must
be regarded as an orthophosphate in which two hydrogens are
replaced by the radicle uranyl, UO_{2}, _i.e._ as H(UO_{2})PO_{4}.
This salt is formed as a hydrated gelatinous yellow precipitate,
on mixing a solution of uranyl nitrate with disodium phosphate.
The precipitation occurs in the presence of acetic acid, but not
in the presence of hydrochloric acid. If moreover an excess of an
ammonium salt be present, the ammonia enters into the composition
of the bright yellow gelatinous precipitate formed, in the
proportion UO_{2}NH_{4}PO_{4}. This precipitate is not soluble in
water and acetic acid, and its solution in inorganic acids when
boiled entirely expels all the phosphoric acid. This fact is taken
advantage of for removing phosphoric acids from solutions--for
instance, from those containing salts of calcium and magnesium.
_Uranium_, which gives an oxide, UO_{3}, and the corresponding salt
UO_{2}X_{2} and dioxide UO_{2}, to which the salts UX_{4} correspond, is
rarely met with in nature. Uranite or the double orthophosphate of uranic
oxide, R(UO_{2})H_{2}P_{2}O_{8},7H_{2}O, where R = Cu or Ca,
uranium-vitriol U(SO_{4})_{2},H_{2}O, samarakite, and æschynite, are very
rarely found, and then only in small quantities. Of more frequent and
abundant occurrence is the non-crystalline, earthy brown uranium ore
known as _pitchblende_ (sp. gr. 7·2), which is mainly composed of the
intermediate oxide, U_{3}O_{8} = UO_{2},2UO_{3}. This ore is found at
Joachimsthal in Bohemia and in Cornwall. It usually contains a number of
different impurities, chiefly sulphides and arsenides of lead and iron,
as well as lime and silica compounds. In order to expel the arsenic and
sulphur it is roasted, ground, washed with dilute hydrochloric acid,
which does not dissolve the uranoso-uranic oxide, U_{3}O_{8}, and the
residue is dissolved in nitric acid, which transforms the uranium oxide
into the nitrate, UO_{2}(NO_{3})_{2}.
It must be observed that the oxide of uranium, first distinguished by
Klaproth (1789), was for a long time regarded as able to give metallic
uranium under the action of charcoal and other reducing agents (with the
aid of heat). But the substance thus obtained was only the _uranium
dioxide_, UO_{2}. The compound nature of this dioxide,[12] or the
presence of oxygen in it, was demonstrated by Peligot (1841), by igniting
it with charcoal in a stream of chlorine. He thus obtained a volatile
_uranium tetrachloride_, UCl_{4},[13] which, when heated with sodium,
gave _metallic uranium_ as a grey metal, having a specific gravity of
18·7, and liberating hydrogen from acids, with the formation of green
uranous salts, UX_{4}, which act as powerful reducing agents.[14]
[12] Uranium dioxide, or _uranyl_, UO_{2}, which is contained in the
salts UO_{2}X_{2}, has the appearance and many of the properties
of a metal. Uranic oxide may be regarded as uranyl oxide,
(UO_{2})O, its salts as salts of this uranyl; its hydroxide,
(UO_{2})H_{2}O_{2}, is constituted like CaH_{2}O_{2}. The green
oxide of uranium, uranoso-uranic oxide (easily formed from uranic
salts by the loss of oxygen), U_{3}O_{8} = UO_{2},2UO_{3}, when
ignited with charcoal or hydrogen (dry) gives a brilliant
crystalline substance of sp. gr. about 11·0 (Urlaub), whose
appearance resembles that of metals, and decomposes steam at a red
heat with the evolution of hydrogen; it does not, however,
decompose hydrochloric or sulphuric acid, but is oxidised by
nitric acid. The same substance (i.e. uranium dioxide UO_{2}) is
also obtained by igniting the compound (UO_{2})K_{2}Cl_{4} in a
stream of hydrogen, according to the equation UO_{2}K_{4}Cl_{4} +
H_{2} = UO_{2} + 2HCl + 2KCl. It was at first regarded as the
metal. In 1841 Peligot found that it contained oxygen, because
carbonic oxide and anhydride were evolved when it was ignited with
charcoal in a stream of chlorine, and from 272 parts of the
substance which was considered to be metal he obtained 382 parts
of a volatile product containing 142 parts of chlorine. From this
it was concluded that the substance taken contained an equivalent
amount of oxygen. As 142 parts of chlorine correspond with 32
parts of oxygen, it followed that 272 - 32 = 240 parts of metal
were combined in the substance with 32 parts of oxygen, and also
in the chlorine compound obtained with 142 parts of chlorine.
These calculations have been made for the now accepted atomic
weight of uranium (U = 240, _see_ Note 14). Peligot took another
atomic weight, but this does not alter the principle of the
argument.
[13] _Uranium tetrachloride_, uranous chloride, UCl_{4}, corresponds
with uranous oxide as a base. It was obtained by Peligot by
igniting uranic oxide mixed with charcoal in a stream of _dry_
chlorine: UO_{3} + 3C + 2Cl_{2} = UCl_{4} + 3CO. This green
volatile compound (Note 12) crystallises in regular octahedra, is
very hygroscopic, easily soluble in water, with the development of
a considerable amount of heat, and no longer separates out from
its solution in an anhydrous state, but disengages hydrochloric
acid when evaporated. The solution of uranous chloride in water is
green. It is also formed by the action of zinc and copper (forming
cuprous chloride) on a solution of uranyl chloride, UO_{2}Cl_{2},
especially in the presence of hydrochloric acid and sal-ammoniac.
Solutions of uranyl salts are converted into uranous salts by the
action of various reducing agents, and among others by organic
substances or by the action of light, whilst the salts UX_{4} are
converted into uranyl salts, UO_{2}X_{2}, by exposure to air or by
oxidising agents. Solutions of the green uranyl salts act as
powerful reducing agents, and give a brown precipitate of the
uranous hydroxide, UH_{4}O_{4}, with potash and other alkalis.
This hydroxide is easily soluble in acids but not in alkalis. On
ignition it does not form the oxide UO_{2}, because it decomposes
water, but when the higher oxides of uranium are ignited in a
stream of hydrogen or with charcoal they yield uranous oxide. Both
it and the chloride UCl_{4}, dissolve in strong sulphuric acid,
forming a green salt, U(SO_{4})_{2},2H_{2}O. The same salt,
together with uranyl sulphate, UO_{2}(SO_{4}), is formed when the
green oxide, U_{3}O_{8}, is dissolved in hot sulphuric acid. The
salts obtained in the latter instance may be separated by adding
alcohol to the solution, which is left exposed to the light; the
alcohol reduces the uranyl salt to uranous salt, an excess of acid
being required. An excess of water decomposes this salt, forming a
basic salt, which is also easily produced under other
circumstances, and contains UO(SO_{4}),2H_{2}O (which corresponds
to the uranic salt).
[14] The atomic weight of uranium was formerly taken as half the
present one, U = 120, and the oxides U_{2}O_{3}, suboxide UO, and
green oxide U_{3}O_{4}, were of the same types as the oxides of
iron. With a certain resemblance to the elements of the iron
group, uranium presents many points of distinction which do not
permit its being grouped with them. Thus uranium forms a very
stable oxide, U_{2}O_{3}(U = 120), but does not give the
corresponding chloride U_{2}Cl_{6} (Roscoe, however, in 1874
obtained UCl_{5}, like MoCl_{5} and WCl_{5}), and under those
circumstances (the ignition of oxide of uranium mixed with
charcoal, in a stream of chlorine), when the formation of this
compound might be expected, it gives (U = 120) the chloride
UCl_{2}, which is characterised by its volatility; this is not a
property, to such an extent, of any of the bichlorides, RCl_{2},
of the iron group.
The alteration or doubling of the atomic weight of uranium--_i.e._
the recognition of U = 240--was made for the first time in the
first (Russian) edition of this work (1871), and in my memoir of
the same year in Liebig's _Annalen_, on the ground that with an
atomic weight 120, uranium could not be placed in the periodic
system. I think it will not be superfluous to add the following
remarks on this subject: (1) In the other groups (K--Rb--Cs,
Ca--Sr--Ba, Cl--Br--I) the acid character of the oxides decreases
and their basic character increases with the rise of atomic
weight, and therefore we should expect to find the same in the
group Cr--Mo--W--U, and if CrO_{3}, MoO_{3}, WO_{3} be the
anhydrides of acids then we indeed find a decrease in their acid
character, and therefore uranium trioxide, UO_{3}, should be a
very feeble anhydride, but its basic properties should also be
very feeble. Uranic oxide does indeed show these properties, as
was pointed out above (Note 10). (2) Chromium and its analogues,
besides the oxides RO_{3}, also form lower grades of oxidation
RO_{2}, R_{2}O_{3}, and the same is seen in uranium; it forms
UO_{3}, UO_{2}, U_{2}O_{3} and their compounds. (3) Molybdenum and
tungsten, in being reduced from RO_{3}, easily and frequently give
an intermediate oxide of a blue colour, and uranium shows the same
property; giving the so-called green oxide which, according to
present views, must be regarded as U_{3}O_{8} = UO_{2}2UO_{3},
analogous to Mo_{3}O_{8}. (4) The higher chlorides, RCl_{6},
possible for the elements of this group, are either unstable
(WCl_{6}) or do not exist at all (Cr); but there is one single
lower volatile compound, which is decomposed by water, and liable
to further reduction into a non-volatile chlorine product and the
metal. The same is observed in uranium, which forms an easily
volatile chloride, UCl_{4}, decomposed by water. (5) The high sp.
gr. of uranium (18·6) is explained by its analogy to tungsten (sp.
gr. 19·1). (6) For uranium, as for chromium and tungsten, yellow
tints predominate in the form RO_{3}, whilst the lower forms are
green and blue. (7) Zimmermann (1881) determined the vapour
densities of uranous bromide, UBr_{4}, and chloride, UCl_{4} (19·4
and 13·2), and they were found to correspond to the formulæ given
above--that is, they confirmed the higher atomic weight U = 240.
Roscoe, a great authority on the metals of this group, was the
first to accept the proposed atomic weight of uranium, U = 240,
which since Zimmermann's work has been generally recognised.
As the salts of uranic oxide are reduced in the absence of organic matter
by the action of light, and as they impart a characteristic coloration to
glass,[15] they find a certain application in photography and glass work.
[15] Uranium glass, obtained by the addition of the yellow salt
K_{2}U_{2}O_{7} to glass, has a green yellow fluorescence, and is
sometimes employed for ornaments; it absorbs the violet rays, like
the other salts of uranic oxide--that is, it possesses an
absorption spectrum in which the violet rays are absent. The index
of refraction of the absorbed rays is altered, and they are given
out again as greenish-yellow rays; hence, compounds of uranic
acid, when placed in the violet portion of the spectrum, emit a
greenish-yellow light, and this forms one of the best examples
(another is found in a solution of quinine sulphate) of the
phenomenon of fluorescence. The rays of light which pass through
uranic compounds do not contain the rays which excite the
phenomena of fluorescence and of chemical transformation, as the
researches of Stokes prove.
If we compare together the highly acid elements, sulphur, selenium, and
tellurium, of the uneven series, with chromium, molybdenum, tungsten, and
uranium of the even series, we find that the resemblance of the
properties of the higher form RO_{3} does not extend to the lower forms,
and even entirely disappears in the elements, for there is not the
smallest resemblance between sulphur and chromium and their analogues in
a free state. In other words, this means that the small periods, like Na,
Mg, Al, Si, P, S, Cl, containing seven elements, do not contain any near
analogues of chromium, molybdenum, &c., and therefore their true position
among the other elements must be looked for only in those large periods
which contain two small periods, and whose type is seen in the period
containing: K, Ca, Sc, Ti, V, Cr, Mn, Fe, Co, Ni, Cu, Zn, Ga, Ge, As, Se,
Br. These large periods contain Ca and Zn, giving RO, Sc, and Ga of the
third group, Ti and Ge giving RO_{2}, V and As forming R_{2}O_{5}, Cr and
Se of the sixth group, Mn and Br of the seventh group, and the remaining
elements, Fe, Co, Ni, form connective members of the intermediate eighth
group, to the description of the representatives of which we shall turn
in the following chapters. We will now proceed to describe _manganese_,
Mn = 55, as an element of the seventh group of the even series, directly
following after Cr = 52, which corresponds with Br = 80 to the same
degree that Cr does with Se = 79. For chromium, selenium, and bromine
very close analogues are known, but for manganese as yet none have been
obtained--that is, it is the only representative of the even series in
the seventh group. In placing manganese with the halogens in one group,
the periodic system of the elements only requires that it should bear an
analogy to the halogens in the higher type of oxidation--_i.e._ in the
salts and acids--whilst it requires that as great a difference should be
expected in the lower types and elements as there exists between chromium
or molybdenum and sulphur or selenium. And this is actually the case. The
elements of the seventh group form a higher salt-forming oxide,
R_{2}O_{7}, and its corresponding hydrate, HRO_{4}, and salts--for
example, KClO_{4}. Manganese in the form of potassium permanganate,
KMnO_{4}, actually presents a great analogy in many respects to potassium
perchlorate, KClO_{4}. The analogy of the crystalline form of both salts
was shown by Mitscherlich. The salts of permanganic acid are also nearly
all soluble in water, like those of perchloric acid, and if the silver
salt of the latter, AgClO_{4}, be sparingly soluble in water, so also is
silver permanganate, AgMnO_{4}. The specific volume of potassium
perchlorate is equal to 55, because its specific gravity = 2·54; the
specific volume of potassium permanganate is equal to 58, because its
specific gravity = 2·71. So that the volumes of equivalent quantities are
in this instance approximately the same whilst the atomic volumes of
chlorine (35·5/1·3 = 27) and manganese (55/7·5) are in the ratio 4 : 1.
In a free state the higher acids HClO_{4} and HMnO_{4} are both soluble
in water and volatile, both are powerful oxidisers--in a word, their
analogy is still closer than that of chromic and sulphuric acids, and
those points of distinction which they present also appear among the
nearest analogues--for example, in sulphuric and telluric acids, in
hydrochloric and hydriodic acids, &c. Besides Mn_{2}O_{7} manganese gives
a lower grade of oxidation, MnO_{3}, analogous to sulphuric and chromic
trioxides, and with it corresponds potassium manganate, K_{2}MnO_{4},
isomorphous with potassium sulphate.[16] In the still lower grades of
oxidation, Mn_{2}O_{3} and MnO, there is hardly any similarity to
chlorine, whilst every point of resemblance disappears when we come to
the elements themselves--_i.e._ to manganese and chlorine--for manganese
is a metal, like iron, which combines directly with chlorine to form a
saline compound, MnCl_{2}, analogous to magnesium chloride.[17]
[16] The comparison of potassium permanganate with potassium
perchlorate, or of potassium manganate with potassium sulphate,
shows directly that many of the physical and chemical properties
of substances do not depend on the nature of the elements, but on
the atomic types in which they appear, on the kind of movements,
or on the positions in which the atoms forming the molecule occur.
[17] If, however, we compare the spectra (Vol. I. p. 565) of chlorine,
bromine, and iodine with that of manganese, a certain resemblance
or analogy is to be found connecting manganese both to iron and to
chlorine, bromine, and iodine.
Manganese belongs to the number of metals widely distributed in nature,
especially in those localities where iron occurs, whose ores frequently
contain compounds of manganous oxide, MnO, which presents a resemblance
to ferrous oxide, FeO, and to magnesia. In many minerals magnesia and the
oxides allied to it are replaced by manganous oxide; calcspars and
magnesites--_i.e._ R´´CO_{3} in general--are frequently met with
containing manganous carbonate, which also occurs in a separate state,
although but rarely. The soil also and the ash of plants generally
contain a small quantity of manganese. In the analysis of minerals it is
generally found that manganese occurs together with magnesia, because,
like it, manganous oxide remains in solution in the presence of
ammoniacal salts, not being precipitated by reagents. The property of
this manganous oxide, MnO, of passing into the higher grades of oxidation
under the influence of heat, alkalis, and air, gives an easy means not
only of discovering the presence of manganese in admixture with magnesia,
but also of separating these two analogous bases. Magnesia is not able to
give higher grades of oxidation, whilst manganese gives them with great
facility. Thus, for instance, an _alkaline_ solution of sodium
hypochlorite produces a precipitate of manganese dioxide in a solution of
a manganous salt: MnCl_{2} + NaClO + 2NaHO = MnO_{2} + H_{2}O + 3NaCl;
whilst magnesia is not changed under these circumstances, and remains in
the form of MgCl_{2}. If the magnesia be precipitated owing to the
presence of alkali, it may be dissolved in acetic acid, in which
manganese dioxide is insoluble. The presence of small quantities of
manganese may also be recognised by the green coloration which alkalis
acquire when heated with manganese compounds in the air. This green
coloration depends on the property of manganese of giving a green
alkaline manganate: MnCl_{2} + 4KHO + O_{2} = K_{2}MnO_{4} + 2KCl +
2H_{2}O. Thus _the faculty of oxidising in the presence of alkalis_ forms
an essential character of manganese. The higher grades of oxidation
containing Mn_{2}O_{7} and MnO_{3} are quite unknown in nature, and even
MnO_{2} is not so widely spread in nature as the ores composed of
manganous compounds which are met with nearly everywhere. The most
important ore of manganese is its dioxide, or so-called _peroxide_,
MnO_{2}, which is known in mineralogy as _pyrolusite_. Manganese also
occurs as an oxide corresponding with magnetic iron ore, MnO,Mn_{2}O_{3}
= Mn_{3}O_{4}, forming the mineral known as _hausmannite_. The oxide
Mn_{2}O_{3} also occurs in nature as the anhydrous mineral _braunite_,
and in a hydrated form, Mn_{2}O_{3},H_{2}O, called _manganite_. Both of
these often occur as an admixture in pyrolusite. Besides which, manganese
is met with in nature as a rose-coloured mineral, _rhodonite_, or
silicate, MnSiO_{3}. Very fine and rich deposits of manganese ores have
been found in the Caucasus, the Urals, and along the Dnieper. Those at
the Sharapansky district of the Government of Kutais and at Nicopol on
the Dnieper are particularly rich. A large quantity of the ore (as much
as 100,000 tons yearly) is exported from these localities.
Thus manganese gives oxides of the following forms: MnO, manganous oxide,
and manganous salts, MnX_{2}, corresponding with the base, which
resembles magnesia and ferrous oxide in many respects; Mn_{2}O_{3}, a
very feeble base, giving salts, MnX_{3}, analogous to the aluminium and
ferric salts, easily reduced to MnX_{2}; MnO_{2}, dioxide, generally
called peroxide, an almost indifferent oxide, or feebly acid;[18]
MnO_{3}, manganic anhydride, which forms salts resembling potassium
sulphate;[18 bis] Mn_{2}O_{7}, permanganic anhydride, giving salts
analogous to the perchlorates.
[18] The name 'peroxide' should only be retained for those _highest_
oxides (and MnO_{2} stands between MnO and MnO_{3}) which either
by a direct method of double decomposition are able to give
hydrogen peroxide or contain a larger proportion of oxygen than
the base or the acid, just as hydrogen peroxide contains more
oxygen than water. Their type will be H_{2}O_{2}, and they are
exemplified by barium peroxide, BaO_{2}, and sulphur peroxide,
S_{2}O_{7}, &c. Such a dioxide as MnO_{2} is, in all probability,
a salt--that is, a manganous manganate, MnO_{3}MnO, and also, as a
basic salt of a feeble base, capable of combining with alkalis and
acids. Hence the name of manganese peroxide should be abandoned,
and replaced by manganese dioxide. PbO_{2} is better termed lead
dioxide than peroxide. Bisulphide of manganese, MnS_{2},
corresponding to iron pyrites, FeS_{2}, sometimes occurs in nature
in fine octahedra (and cube combinations), for instance, in
Sicily; it is called Hauerite.
[18 bis] On comparing the manganates with the permanganates--for
example, K_{2}MnO_{4} with KMnO_{4}--we find that they differ in
composition by the abstraction of one equivalent of the metal.
Such a relation in composition produced by oxidation is of
frequent occurrence--for instance, K_{4}Fe(CN)_{6} in oxidising
gives K_{3}Fe(CN)_{6}; H_{2}SO_{4} in oxidising gives persulphuric
acid, HSO_{4}, or H_{2}S_{7}O_{8}; H_{2}O forms HO or
H_{2}O_{2}, &c.
_All the oxides of manganese when heated with acids give salts_, MnX_{2},
corresponding with the lower grade of oxidation, _manganous oxide_, MnO.
Manganic oxide, Mn_{2}O_{3}, is a feebly energetic base; it is true that
it dissolves in hydrochloric acid and gives a dark solution containing
the salt MnCl_{3}, but the latter when heated evolves chlorine and gives
a salt corresponding with manganous oxide MnCl_{2}--_i.e._ at first:
Mn_{2}O_{3} + 6HCl = 3H_{2}O + Mn_{2}Cl_{6}, and then the Mn_{2}Cl_{6}
decomposes into 2MnCl_{2} + Cl_{2}. None of the remaining higher grades
of oxidation have a basic character, but _act as oxidising agents in the
presence of acids_, disengaging oxygen and passing into salts of the
lower grade of oxidation of manganese, MnO. Owing to this circumstance,
_the manganous salts_ are often obtained; they are, for instance, left in
the residue when the dioxide is used for the preparation of oxygen and
chlorine.[19]
[19] In the preparation of oxygen from the dioxide by means of
H_{2}SO_{4}, MnSO_{4} is formed; in the preparation of chlorine
from HCl and MnO_{2}, MnCl_{2} is obtained. These two manganous
salts may be taken as examples of compounds MnX_{2}. Manganous
sulphate generally contains various impurities, and also a large
amount of iron salt (from the native MnO_{2}), from which it
cannot be freed by crystallisation. Their removal may, however, be
effected by mixing a portion of the liquid with a solution of
sodium carbonate; a precipitate of manganous carbonate is then
formed. This precipitate is collected and washed, and then added
to the remaining mass of the impure solution of manganous
sulphate; on heating the solution with this precipitate, the whole
of the iron is precipitated as oxide. This is due to the fact that
in the solution of the manganese dioxide in sulphuric acid the
whole of the iron is converted into the ferric state (because the
dioxide acts as an oxidising agent), which, as an exceedingly
feeble base precipitated by calcium carbonate and other kindred
salts, is also precipitated by manganous carbonate. After being
treated in this manner, the solution of manganous sulphate is
further purified by crystallisation. If it be a bright red colour,
it is due to the presence of higher grades of oxidation of
manganese; they may be destroyed by boiling the solution, when the
oxygen from the oxides of manganese is evolved and a very faintly
coloured solution of manganous sulphate is obtained. This salt is
remarkable for the facility with which it gives various
combinations with water. By evaporating the almost colourless
solution of _manganous sulphate_ at very low temperatures, and by
cooling the saturated solution at about 0°, crystals are obtained
containing 7 atoms of water of crystallisation, MnSO_{4},7H_{2}O,
which are isomorphous with cobaltous and ferrous sulphates. These
crystals, even at 10°, lose 5 p.c. of water, and completely
effloresce at 15°, losing about 20 p.c. of water. By evaporating a
solution of the salt at the ordinary temperature, but not above
20°, crystals are obtained containing 5 mol. H_{2}O, which are
isomorphous with copper sulphate; whilst if the crystallisation be
carried on between 20° and 30°, large transparent prismatic
crystals are formed containing 4 mol. H_{2}O (see Nickel). A
boiling solution also deposits these crystals together with
crystals containing 3 mol. H_{2}O, whilst the first salt, when
fused and boiled with alcohol, gives crystals containing 2 mol.
H_{2}O. Graham obtained a monohydrated salt by drying the salt at
about 200°. The last atom of water is eliminated with difficulty,
as is the case with all salts like MnSO_{4}nH_{2}O. The crystals
containing a considerable amount of water are rose-coloured, and
the anhydrous crystals are colourless. The solubility of
MnSO_{4},4H_{2}O (Chapter I., Note 24) per 100 parts of water is:
at 10°, 127 parts; at 37°·5, 149 parts; at 75°, 145 parts; and at
101°, 92 parts. Whence it is seen that at the boiling-point this
salt is less soluble than at lower temperatures, and therefore a
solution saturated at the ordinary temperature becomes turbid when
boiled. Manganous sulphate, being analogous to magnesium sulphate,
is decomposed, like the latter, when ignited, but it does not then
leave manganous oxide, but the intermediate oxide, Mn_{3}O_{4}. It
gives double salts with the alkali sulphates. With aluminium
sulphate it forms fine radiated crystals, whose composition
resembles that of the alums--namely,
MnAl_{2}(SO_{4})_{4},24H_{2}O. This salt is easily soluble in
water, and occurs in nature.
_Manganous chloride_, MCl_{2}, crystallises with 4 mol. H_{2}O,
like the ferrous salt, and not with 6 mol. H_{2}O like many
kindred salts--for example, those of cobalt, calcium, and
magnesium; 100 parts of water dissolve 38 parts of the anhydrous
salt at 10° and 55 parts at 62°. Alcohol also dissolves manganous
chloride, and the alcoholic solution burns with a red flame. This
salt, like magnesium chloride, readily forms double salts. A
solution of borax gives a dirty rose-coloured precipitate having
the composition MnH_{4}(BO_{3})_{2}H_{2}O, which is used as a
drier in paint-making. Potassium cyanide produces a yellowish-grey
precipitate, MnC_{2}N_{2}, with manganous salts, soluble in an
excess of the reagent, a double salt, K_{4}MnC_{6}N_{6},
corresponding with potassium ferrocyanide, being formed. On
evaporation of this solution, a portion of the manganese is
oxidised and precipitated, whilst a salt corresponding to Gmelin's
red salt, K_{3},MnC_{6}N_{6} (_see_ Chapter XXII.), remains in
solution. Sulphuretted hydrogen does not precipitate salts of
manganese, not even the acetate, but ammonium sulphide gives a
flesh-coloured precipitate, MnS; at 320° this sulphide of
manganese passes into a green variety (Antony). Oxalic acid in
strong solutions of manganous salts gives a white precipitate of
the oxalate, MnC_{2}O_{4}. This precipitate is insoluble in water,
and is used for the preparation of manganous oxide itself because
it decomposes like oxalic acid when ignited (in a tube without
access of air), with the formation of carbonic anhydride, carbonic
oxide, and manganous oxide. _Manganous oxide_ thus obtained is a
green powder, which however oxidises with such facility that it
burns in air when brought into contact with an incandescent
substance, and passes into the red intermediate oxide Mn_{3}O_{4}.
In solutions of manganous salts, alkalis produce a precipitate of
the hydroxide MnH_{2}O_{2}, which rapidly absorbs oxygen in the
presence of air and gives the brown intermediate oxide, or, more
correctly speaking, its hydrate.
Manganous oxide, besides being obtained by the above-described
method from manganous oxalate, may also be obtained by igniting
the higher oxides in a stream of hydrogen, and also from manganese
carbonate. The manganous oxide ignited in the presence of hydrogen
acquires a great density, and is no longer so easily oxidised. It
may also be obtained in a crystalline form, if during the ignition
of the carbonate or higher oxide a trace of dry hydrochloric acid
gas be passed into the current of hydrogen. It is thus obtained in
the form of transparent emerald green crystals of the regular
system, and in this state is easily soluble in acids.
Manganous oxide in oxidising gives the _red oxide of manganese_,
Mn_{5}O_{4}. This is the most stable of all the oxides of
manganese; it is not only stable at the ordinary but also at a
high temperature--that is, it does not absorb or disengage oxygen
spontaneously. When ignited, all the higher oxides of manganese
pass into it by losing oxygen, and manganous oxide by absorbing
oxygen. This oxide does not give any distinct salts, but it
dissolves in sulphuric acid, forming a dark red solution, which
contains both manganous and manganic (of the _oxide_, Mn_{2}O_{3})
sulphates. The latter with potassium sulphate gives a manganese
alum, in which the alumina is replaced by its isomorphous oxide of
manganese. But this alum, like the solution of the intermediate
oxide in sulphuric acid, evolves oxygen and leaves a manganous
salt when slightly heated.
_Manganese dioxide_ is still less basic than the oxide, and
disengages oxygen or a halogen in the presence of acids, forming
manganous salts, like the oxide. However, if it be suspended in
ether, and hydrochloric acid gas passed into the mixture, which is
kept cool, the ether acquires a green colour, owing to the
formation of tetrachloride of manganese, MnCl_{4}, corresponding
with the dioxide which passes into solution. It is however very
unstable, being exceedingly easily decomposed with the evolution
of chlorine. The corresponding fluoride, MnF_{4}, obtained by
Nicklés is much more stable. At all events, manganese dioxide does
not exhibit any well-defined basic character, but has rather an
acid character, which is particularly shown in the compounds
MnF_{4} and MnCl_{4} just mentioned, and in the property of
manganese dioxide of combining with alkalis. If the higher grades
of oxidation of manganese be deoxidised in the presence of
alkalis, they frequently give the dioxide combined with the
alkali--for example, in the presence of potash a compound is
formed which contains K_{2}O,5MnO_{2}, which shows the weak acid
character of this oxide. When ignited in the presence of sodium
compounds manganese dioxide frequently forms Na_{2}O,8MnO_{2} and
Na_{2}O,12MnO_{2}, and lime when heated with MnO_{2} gives from
CaO,3MnO_{2} to (CaO)_{2},MnO_{2} (Rousseau) according to the
temperature. Besides which, perhaps, MnO_{2} is a saline compound,
containing MnOMnO_{3} or (MnO)_{3}Mn_{2}O_{7}, and there are
reactions which support such a view (Spring, Richards, Traube, and
others); for instance it is known that manganous chloride and
potassium permanganate give the dioxide in the presence of
alkalis.
Manganese dioxide may be obtained from manganous salts by the
action of oxidising agents. If manganous hydroxide or carbonate be
shaken up in water through which chlorine is passed, the
hypochlorite of the metal is not formed, as is the case with
certain other oxides, but manganese dioxide is precipitated:
2MnO_{2}H_{2} + Cl_{2} = MnCl_{2} + MnO_{2},H_{2}O + H_{2}O. Owing
to this fact, hypochlorites in the presence of alkalis and acetic
acid when added to a solution of manganous salts give hydrated
manganese dioxide, as was mentioned above. Manganous nitrate also
leaves manganese dioxide when heated to 200°. It is also obtained
from manganous and manganic salts of the alkalis, when they are
decomposed in the presence of a small amount of acid; the
practical method of converting the salts MnX_{2} into the higher
grades of oxidation is given in Chapter II., Note 6.
As the salts of manganous oxide MnX_{2} closely resemble (and are
isomorphous with) the salts of magnesia MgX_{2} in many respects (with
the exception of the fact that MnX_{2} are rose coloured and are easily
oxidised in the presence of alkalis), we will not dwell upon them, but
limit ourselves to illustrating the chemical character of manganese by
describing the metal and its corresponding acids. The fact alone that the
oxides of manganese are not reduced to the metal when ignited in hydrogen
(whilst the oxides of iron give metallic iron under these circumstances),
but only to manganous oxide, MnO, shows that manganese has a considerable
affinity for oxygen--that is, it is difficult to reduce. This may be
effected, however, by means of charcoal or sodium at a very high
temperature. A mixture of one of the oxides of manganese with charcoal or
organic matter gives fused _metallic manganese_ under the powerful heat
developed by coke with an artificial draught. The metal was obtained for
the first time in this manner by Gahn, after Pott, and more especially
Scheele, had in the last century shown the difference between the
compounds of iron and manganese (they were previously regarded as being
the same). Manganese is prepared by mixing one of its oxides in a
finely-divided state with oil and soot; the resultant mass is then first
ignited in order to decompose the organic matter, and afterwards strongly
heated in a charcoal crucible. The manganese thus obtained, however,
contains, as a rule, a considerable amount of silicon and other
impurities. Its specific gravity varies between 7·2 and 8·0. It has a
light grey colour, a feebly metallic lustre, and although it is very hard
it can be scratched by a file. It rapidly oxidises in air, being
converted into a black oxide; water acts on it with the evolution of
hydrogen--this decomposition proceeds very rapidly with boiling water,
and if the metal contain carbon.[20]
[20] Other chemists have obtained manganese by different methods, and
attributed different properties to it. This difference probably
depends on the presence of carbon in different proportions.
Deville obtained manganese by subjecting the pure dioxide, mixed
with pure charcoal (from burnt sugar), to a strong heat in a lime
crucible until the resultant metal fused. The metal obtained had a
rose tint, like bismuth, and like it was very brittle, although
exceedingly hard. It decomposed water at the ordinary temperature.
Brunner obtained manganese having a specific gravity of about 7·2,
which decomposed water very feebly at the ordinary temperature,
did not oxidise in air, and was capable of taking a bright polish,
like steel; it had the grey colour of cast iron, was very brittle,
and hard enough to scratch steel and glass, like a diamond.
Brunner's method was as follows: He decomposed the manganese
fluoride (obtained as a soluble compound by the action of
hydrofluoric acid on manganese carbonate) with sodium, by mixing
these substances together in a crucible and covering the mixture
with a layer of salt and fluor spar; after which the crucible was
first gradually heated until the reaction began, and then strongly
heated in order to fuse the metal separated. Glatzel (1889)
obtained 25 grms. of manganese, having a grey colour and sp. gr.
7·39, by heating a mixture of 100 grms. of MnCl_{2} with 200 grms.
KCl and 15 grms. Mg to a bright white heat. Moissan and others, by
heating the oxides of manganese with carbon in the electric
furnace, obtained carbides of manganese--for example, Mn_{3}C--and
remarked that the metal volatilised in the heat of the voltaic
arc. Metallic manganese is, however, not prepared on a large
scale, but only its alloys with carbon (they readily and rapidly
oxidise) and _ferro-manganese_ or a coarsely crystalline alloy of
iron, manganese and carbon, which is smelted in blast-furnaces
like pig-iron (_see_ Chapter XXII.) This ferro-manganese is
employed in the manufacture of steel by Bessemer's and other
processes (see Chapter XXII.) and for the manufacture of manganese
bronze. However, in America, Green and Wahl (1895) obtained almost
pure metallic manganese on a large scale. They first treat the ore
of MnO_{2} with 30 p.c. sulphuric acid (which extracts all the
oxides of iron present in the ore), and then heat it in a reducing
flame to convert it into MnO, which they mix with a powder of Al,
lime and CaF_{2} (as a flux), and heat the mixture in a crucible
lined with magnesia; a reaction immediately takes place at a
certain temperature, and a metal of specific gravity 7·3 is
obtained, which only contains a small trace of iron.
Manganese gives two compounds with _nitrogen_, Mn_{5}N_{2} and
Mn_{3}N_{2}. They were obtained by Prelinger (1894) from the
amalgam of manganese Mn_{2}Hg_{5} (obtained on a mercury anode by
the action of an electric current upon a solution of MnCl_{2});
the mercury may be removed from this amalgam by heating it in an
atmosphere of hydrogen, and then metallic manganese is obtained as
a grey porous mass of specific gravity 7·42. If this amalgam be
heated in dry nitrogen it gives Mn_{5}N_{2} (grey powder, sp. gr.
6·58), but if heated in an atmosphere of NH_{3} it gives (as also
does Mn_{5}N_{2}) Mn_{3}N_{2}, (a dark mass with a metallic
lustre, sp. gr. 6·21), which, when heated in nitrogen is converted
into Mn_{5}N_{2}, and if heated in hydrogen evolves NH_{3} and
disengages hydrogen from a solution of NH_{4}Cl. At all events,
manganese is a metal which decomposes water more easily than iron,
nickel, and cobalt.
It has been shown above that if manganese dioxide, or any lower oxide of
manganese, be heated with an alkali in the presence of air, the mixture
absorbs oxygen,[21] and forms an alkaline manganate of a green colour:
2KHO + MnO_{2} + O = K_{2}MnO_{4} + H_{2}O. Steam is disengaged during
the ignition of the mixture, and if this does not take place there is no
absorption of oxygen. The oxidation proceeds much more rapidly if, before
igniting in air, potassium chlorate or nitre be added to the mixture, and
this is the method of preparing _potassium manganate_, K_{2}MnO_{4}. The
resultant mass dissolved in a small quantity of water gives a dark green
solution, which, when evaporated under the receiver of an air-pump over
sulphuric acid, deposits green crystals of exactly the same form as
potassium sulphate--namely, six-sided prisms and pyramids. The
composition of the product is not changed by being redissolved, if
perfectly pure water free from air and carbonic acid be taken. But in the
presence of even very feeble acids the solution of this salt changes its
colour and becomes red, and deposits manganese dioxide. The same
decomposition takes place when the salt is heated with water, but when
diluted with a large quantity of unboiled water manganese dioxide does
not separate, although the solution turns red. This change of colour
depends on the fact that potassium manganate, K_{2}MnO_{4}, whose
solution is green, is transformed into potassium permanganate, KMnO_{4},
whose solution is of a red colour. The reaction proceeding under the
influence of acids and a large quantity of water is expressed in the
following manner: 3K_{2}MnO_{4} + 2H_{2}O = 2KMnO_{4} + MnO_{2} + 4KHO.
If there is a large proportion of acid and the decomposition is aided by
heat, the manganese dioxide and potassium permanganate are also
decomposed, with formation of manganous salt. Exactly the same
decomposition as takes place under the action of acids is also
accomplished by magnesium sulphate, which reacts in many cases like an
acid. When water holding atmospheric oxygen in solution acts on a
solution of potassium manganate, the oxygen combines directly with the
manganate and forms potassium permanganate, without precipitating
manganese dioxide, 2K_{2}MnO_{4} + O + H_{2}O = 2KMnO_{4} + 2KHO. Thus a
solution of potassium manganate undergoes a very characteristic change in
colour and passes from green to red; hence this salt received the name of
_chameleon mineral_.[22]
[21] Volume I. p. 157, Note 7.
[22] It was known to the alchemists by this name, but the true
explanation of the change in colour is due to the researches of
Chevillot, Edwards, Mitscherlich, and Forchhammer. The change in
colour of potassium manganate is due to its instability and to its
splitting up into two other manganese compounds, a higher and a
lower: 3MnO_{3} = Mn_{2}O_{7} + MnO_{2}. Manganese trioxide is
really decomposed in this manner by the action of water (see
later): 3MnO_{3} + H_{2}O = 2MnHO_{4} + MnO_{2} (Franke, Thorpe,
and Humbly). The instability of the salt is proved by the fact of
its being deoxidised by organic matter, with the formation of
manganese dioxide and alkali, so that, for instance, a solution of
this salt cannot be filtered through paper. The presence of an
excess of alkali increases the stability of the salt; when heated
it breaks up in the presence of water, with the evolution of
oxygen.
The method of preparing _potassium permanganate_ will be
understood from the above. There are many recipes for preparing
this substance, as it is now used in considerable quantities both
for technical and laboratory purposes. But in all cases the
essence of the methods is one and the same: a mixture of alkali
with any oxide of manganese (even manganous hydroxide, which may
be obtained from manganous chloride) is first heated in the
presence of air or of an oxidising substance (for the sake of
rapidity, with potassium chlorate); the resultant mass is then
treated with water and heated, when manganese dioxide is
precipitated and potassium permanganate remains in solution. This
solution may be boiled, as the liquid will contain free alkali;
but the solution cannot be evaporated to dryness, because a strong
solution, as well as the solid salt, is decomposed by heat.
By adding a dilute solution of manganous sulphate to a boiling
mixture of lead dioxide and dilute nitric acid, the whole of the
manganese may be converted into permanganic acid (Crum).
_Potassium permanganate_, KMnO_{4}, crystallises in well-formed, long
red prisms with a bright green metallic lustre. In the arts the potash is
frequently replaced by soda, and by other alkaline bases, but no salt of
permanganic acid crystallises so well as the potassium salt, and
therefore this salt is exclusively used in chemical laboratories. One
part of the crystalline salt dissolves in 15 parts of water at the
ordinary temperature. The solution is of a very deep _red colour_, which
is so intense that it is still clearly observable after being highly
diluted with water. In a solid state it is decomposed by heat, with
evolution of oxygen, a residue consisting of the lower oxides of
manganese and potassium oxide being left.[22 bis] A mixture of
permanganate of potassium, phosphorous and sulphur takes fire when struck
or rubbed, a mixture of the permanganate with carbon only takes fire when
heated, not when struck. The instability of the salt is also seen in the
fact that its solution is decomposed by peroxide of hydrogen, which at
the same time it decomposes. A number of substances reduce potassium
permanganate to manganese dioxide (in which case the red solution becomes
colourless).[23] Many organic substances (although far from all, even
when boiled in a solution of permanganate) act in this manner, being
oxidised at the expense of a portion of its oxygen. Thus, a solution of
sugar decomposes a cold solution of potassium permanganate. In the
presence of an excess of alkali, with a small quantity of sugar, the
reduction leads to the formation of potassium manganate, because
2KMnO_{4} + 2KHO = O + 2K_{2}MnO_{4} + H_{2}O. With a considerable amount
of sugar and a more prolonged action, the solution turns brown and
precipitates manganese dioxide or even oxide. In the oxidation of many
organic bodies by an alkaline solution of KMnO_{4} generally
three-eighths of the oxygen in the salt are utilised for oxidation:
2KMnO_{4} = K_{2}O + 2MnO_{2} + O_{3}. A portion of the alkali liberated
is retained by the manganese dioxide, and the other portion generally
combines with the substance oxidised, because the latter most frequently
gives an acid with an excess of alkali. A solution of potassium iodide
acts in a similar manner, being converted into potassium iodate at the
expense of the three atoms of oxygen disengaged by two molecules of
potassium permanganate.
[22 bis] The solution of this salt with an excess of impure commercial
alkali generally acquires a green tint.
[23] A solution of potassium permanganate gives a beautiful absorption
spectrum (Chapter XIII.) If the light in passing through this
solution loses a portion of its rays in it (if one may so account
for it), this is partially explained by the increased oxidising
power which the solution then acquires. We may here also remark
that a dilute solution of permanganate of potassium forms a
colourless solution with nickel salts, because the green colour of
the solution of nickel salts is complementary to the red. Such a
decolorised solution, containing a large proportion of nickel and
a small proportion of manganese, decomposes after a time, throws
down a precipitate, and re-acquires the green colour proper to the
nickel salts. The addition of a solution of a cobalt salt
(rose-red) to the nickel salt also destroys the colour of both
salts.
_In the presence of acids, potassium permanganate acts as an oxidising
agent_ with still greater energy than in the presence of alkalis. At any
rate, a greater proportion of oxygen is then available for oxidation,
namely, not 3/8, as in the presence of alkalis, but 5/8, because in the
first instance manganese dioxide is formed, and in the second case
manganous oxide, or rather the salt, MnX_{2}, corresponding with it.
Thus, for instance, in the presence of an excess of sulphuric acid, the
decomposition is accomplished in the following manner: 2KMnO_{4} +
3H_{2}SO_{4} = K_{2}SO_{4} + 2MnSO_{4} + 3H_{2}O + 5O. This
decomposition, however, does not proceed directly on mixing a solution of
the salt with sulphuric acid, and crystals of the salt even dissolve in
oil of vitriol without the evolution of oxygen, and this solution only
decomposes by degrees after a certain time. This is due to the fact that
sulphuric acid liberates free permanganic acid from the permanganate,[24]
which acid is stable in solution. But if, in the presence of acids and a
permanganate, there is a substance capable of absorbing oxygen--for
instance, capable of passing into a higher grade of oxidation--then the
reduction of the permanganic acid into manganous oxides sometimes
proceeds directly at the ordinary temperature. This reduction is very
clearly seen, because the solutions of potassium permanganate are red
whilst the manganous salts are almost colourless. Thus, for instance,
nitrous acid and its salts are converted into nitric acid and decolorise
the acid solution of the permanganate. Sulphurous anhydride and its salts
immediately decolorise potassium permanganate, forming sulphuric acid.
Ferrous salts, and in general salts of lower grades of oxidation capable
of being oxidised in solution, act in exactly the same manner.
Sulphuretted hydrogen is also oxidised to sulphuric acid; even mercury is
oxidised at the expense of permanganic acid, and decolorises its
solution, being converted into mercuric oxide. Moreover, the end point of
these reactions may easily be seen, and therefore, having first
determined the amount of active oxygen in one volume of a solution of
potassium permanganate, and knowing how many volumes are required to
effect a given oxidation, it is easy to determine the amount of an
oxidisable substance in a solution from the amount of permanganate
expended (Marguerite's method).
[24] If sulphuric acid is allowed to act on potassium permanganate
without any special precautions, a large amount of oxygen is
evolved (it may even explode and inflame), and a violet spray of
the decomposing permanganic acid is given off. But if the pure
salt (_i.e._ free from chlorine) be dissolved in pure well-cooled
sulphuric acid, without any rise in temperature, a green-coloured
liquid settles at the bottom of the vessel. This liquid does not
contain any sulphuric acid, and consists of permanganic anhydride,
Mn_{2}O_{7} (Aschoff, Terreil). It is impossible to prepare any
considerable quantity of the anhydride by this method, as it
decomposes with an explosion as it collects, evolving oxygen and
leaving red oxide of manganese. _Permanganic anhydride_,
Mn_{2}O_{7}, in dissolving in sulphuric acid, gives a green
solution, which (according to Franke, 1887) contains a compound
Mn_{2}SO_{10} = (MnO_{3})_{2}SO_{4}--that is, sulphuric acid in
which both hydrogens are replaced by the group MnO_{3}, which is
combined with OK in permanganate of potassium. This mixture with a
small quantity of water gives Mn_{2}O_{7}, according to the
equation: (MnO_{3})_{2}SO_{4} + H_{2}O = H_{2}SO_{4} +
Mn_{2}O_{7}, and when heated to 30° it gives _manganese trioxide_,
(MnO_{3})_{2}SO_{4} + H_{2}O = 2MnO_{2} + H_{2}SO_{4} + O. Pure
manganese trioxide is obtained if the solution of
(MnO_{3})_{2}SO_{4} be poured in drops on to sodium carbonate.
Then, together with carbonic anhydride, a spray of manganese
trioxide passes over, which may be collected in a well-cooled
receiver, and this shows that the reaction proceeds according to
the equation: (MnO_{3})_{2}SO_{4} + Na_{2}CO_{3} = Na_{2}SO_{4} +
2MnO_{3} + CO_{2} + O (Thorpe). The trioxide is decomposed by
water, forming manganese dioxide and a solution of _permanganic
acid_: 3MnO_{3} + H_{2}O = MnO_{2} + 2HMnO_{4}. The same acid is
obtained by dissolving permanganic anhydride in water.
Barium permanganate when treated with sulphuric acid gives the
same acid. This barium salt may be prepared by the action of
barium chloride on the difficultly soluble silver permanganate,
AgMnO_{4}, which is precipitated on mixing a strong solution of
the potassium salt with silver nitrate. The solution of
permanganic acid forms a bright red liquid which reflects a dark
violet tint. A dilute solution has exactly the same colour as that
of the potassium salt. It deposits manganese dioxide when exposed
to the action of light, and also when heated above 60°, and this
proceeds the more rapidly the more dilute the solution. It shows
its oxidising properties in many cases, as already mentioned. Even
hydrogen gas is absorbed by a solution of permanganic acid; and
charcoal and sulphur are also oxidised by it, as they are by
potassium permanganate. This may be taken advantage of in
analysing gunpowder, because when it is treated with a solution of
potassium permanganate, all the sulphur is converted into
sulphuric acid and all the charcoal into carbonic anhydride.
Finely-divided platinum immediately decomposes permanganic acid.
With potassium iodide it liberates iodine (which may afterwards be
oxidised into iodic acid) (Mitscherlich, Fromherz, Aschoff, and
others). Ammonia does not form a corresponding salt with free
permanganic acid, because it is oxidised with evolution of
nitrogen. The oxidising action of permanganic acid in a strong
solution may be accompanied by flame and the formation of violet
fumes of permanganic acid; thus a strong solution of it takes fire
when brought into contact with paper, alcohol, alkaline sulphides,
fats, &c.
We may add that, according to Franke, 1 part of potassium
permanganate with 13 parts of sulphuric acid at 100° gives brown
crystals of the salt Mn_{2}(SO_{4})_{3},H_{2}SO_{4},4H_{2}O, which
gives a precipitate of hydrated manganese dioxide, H_{2}MnO_{3} =
MnO_{2}H_{2}O, when treated with water.
Spring, by precipitating potassium permanganate with sodium
sulphite and washing the precipitate by decantation, obtained a
soluble colloidal manganese oxide, whose composition was the mean
between Mn_{2}O_{3} and MnO_{2}--namely,
Mn_{2}O_{3},4(MnO_{2}H_{2}O).
The oxidising action of KMnO_{4}, like all other chemical reactions, is
not accomplished instantaneously, but only gradually. And, as the course
of the reaction is here easily followed by determining the amount of salt
unchanged in a sample taken at a given moment,[25] the oxidising reaction
of potassium permanganate, in an acid liquid, was employed by Harcourt
and Esson (1865) as one of the first cases for the investigation of the
laws of the _rate of chemical change_[26] as a subject of great
importance in chemical mechanics. In their experiments they took oxalic
acid, C_{2}H_{2}O_{4}, which in oxidising gives carbonic anhydride,
whilst, with an excess of sulphuric acid, the potassium permanganate is
converted into manganous sulphate, MnSO_{4}, so that the ultimate
oxidation will be expressed by the equation: 5C_{2}H_{2}O_{4} + 2MnKO_{4}
+ 3H_{2}SO_{4} = 10CO_{2} + K_{2}SO_{4} + 2MnSO_{4} + 8H_{2}O. The
influence of the relative amount of sulphuric acid is seen from the
annexed table, which gives the measure of reaction _p_ per 100 parts of
potassium permanganate, taken four minutes after mixing, using n
molecules of sulphuric acid, H_{2}SO_{4}, per 2KMnO_{4} +
5C_{2}H_{2}O_{4}:
_n_ = 2 4 6 8 12 16 22
_p_ = 22 36 51 63 77 86 92
showing that in a given time (4 minutes) the oxidation is the more
perfect the greater the amount of sulphuric acid taken for given amounts
of KMnO_{4} and C_{2}H_{2}O_{4}. It is obvious also that the temperature
and relative amount of every one of the acting and resulting substances
should show its influence on the relative velocity of reaction; thus, for
instance, direct experiment showed the influence of the admixture of
manganous sulphate. When a large proportion of oxalic acid (108
molecules) was taken to a large mass of water and to 2 molecules of
permanganate 14 molecules of manganous sulphate were added, the quantity
x of the potassium permanganate acted on (in percentages of the potassium
permanganate taken) in t minutes (at 16°) was as follows:
_t_ = 2 5 8 11 14 44 47 53 61 68
_x_ = 5·2 12·1 18·7 25·1 31·3 68·4 71·7 75·8 79·8 83·0
These figures show that the rate of reaction--that is, the quantity of
permanganate changed in one minute--decreases proportionally to the
decrease in the amount of unchanged potassium permanganate. At the
commencement, about 2·6 per cent. of the salt taken was decomposed in the
course of one minute, whilst after an hour the rate was about 0·5 per
cent. The same phenomena are observed in every case which has been
investigated, and this branch of theoretical or physical chemistry, now
studied by many,[27] promises to explain the course of chemical
transformations from a fresh point of view, which is closely allied to
the doctrine of affinity, because the rate of reaction, without doubt, is
connected with the magnitude of the affinities acting between the
reacting substances.
[25] For rapid and accurate determinations of this kind, advantage is
taken of those methods of chemical analysis which are known as
'titrations' (volumetric analysis), and consist in measuring the
volume of solutions of known strength required for the complete
conversion of a given substance. Details respecting the theory and
practice of titration, in which potassium permanganate is very
frequently employed, must be looked for in works on analytical
chemistry.
[26] The measurements of velocity and acceleration serve for
determining the measure of forces in mechanics, but in that case
the velocities are magnitudes of length or paths passed over in a
unit of time. The velocity of chemical change embodies a
conception of quite another kind. In the first place, the
velocities of reactions are magnitudes of the masses which have
entered into chemical transformations; in the second place, these
velocities can only be relative quantities. Hence the conception
of 'velocity' has quite a different meaning in chemistry from what
it has in mechanics. Their only common factor is time. If _dt_ be
the increment of time and _dx_ the quantity of a substance changed
in this space of time, then the fraction (or quotient) _dx/dt_
will express the rate of the reaction. The natural conclusion,
come to both by Harcourt and Esson, and previously to them (1850)
by Wilhelmj (who investigated the rate of conversion, or
inversion, of sugar in its passage into glucose), consists in
establishing that this velocity is proportional to the quantity of
substances still unchanged--_i.e._ that _dx/dt_ = C(A - _x_),
where C is a constant coefficient of proportionality, and where A
is the quantity of a substance taken for reaction at the moment
when _t_ = 0 and _x_ = 0--that is, at the beginning of the
experiment, from which the time _t_ and quantity _x_ of substance
changed is counted. On integrating the preceding equation we
obtain log(A/(A - _x_)) = _kt_, where _k_ is a new constant, if we
take ordinary (and not natural) logarithms. Hence, knowing A, _x_,
and _t_, for each reaction, we find _k_, and it proves to be a
constant quantity. Thus from the figures cited in the text for the
reaction 2KMnO_{4} + 108C_{2}H_{2}O_{4} + 14MnSO_{4}, it may be
calculated that _k_ = 0·0114; for example, _t_ = 44, _x_ = 68·4 (A
= 100), whence _kt_ = 0·5004 and _k_ = 0·0114, (_see also_ Chapter
XIV., Note 3, and Chapter XVII., Note 25 bis).
[27] The researches made by Hood, Van't Hoff, Ostwald, Warder,
Menschutkin, Konovaloff, and others have a particular significance
in this direction. Owing to the comparative novelty of this
subject, and the absence of applicable as well as indubitable
deductions, I consider it impossible to enter into this province
of theoretical chemistry, although I am quite confident that its
development should lead to very important results, especially in
respect to chemical equilibria, for Van't Hoff has already shown
that the limit of reaction in reversible reactions is determined
by the attainment of equal velocities for the opposite reactions.
CHAPTER XXII
IRON, COBALT, AND NICKEL
Judging from the atomic weights, and the forms of the higher oxides of
the elements already considered, it is easy to form an idea of the seven
groups of the periodic system. Such are, for instance, the typical series
Li, Be, B, C, N, O, F, or the third series, Na, Mg, Al, Si, P, S, Cl. The
seven usual types of oxides from R_{2}O to R_{2}O_{7} correspond with
them (Chapter XV.) The position of the eighth group is quite separate,
and is determined by the fact that, as we have already seen, in each
group of metals having a greater atomic weight than potassium a
distinction ought to be made between the elements of the even and uneven
series. The series of even elements, commencing with a strikingly
alkaline element (potassium, rubidium, cæsium), together with the uneven
series following it, and concluding with a haloid (chlorine, bromine,
iodine), forms a large period, the properties of whose members repeat
themselves in other similar periods. The elements of the eighth group are
situated between the elements of the even series and the elements of the
uneven series following them. And for this reason elements of the eighth
group are found in the middle of each large period. The properties of the
elements belonging to it, in many respects independent and striking, are
shown with typical clearness in the case of iron, the well-known
representative of this group.
_Iron_ is one of those elements which are not only widely diffused in the
crust of the earth, but also throughout the entire universe. Its oxides
and their various compounds are found in the most diverse portions of the
earth's crust; but here iron is always found combined with some other
element. Iron is not found on the earth's surface in a free state,
because it easily oxidises under the action of air. It is occasionally
found in the native state in meteorites, or aerolites, which fall upon
the earth.
_Meteoric iron_ is formed outside the earth.[1] Meteorites are
fragments which are carried round the sun in orbits, and fall upon the
earth when coming into proximity with it during their motion in space.
The meteoric dust, on passing through the upper parts of the atmosphere,
and becoming incandescent from friction with the gases, produces that
phenomenon which is familiar under the name of falling stars.[2] Such is
the doctrine concerning meteorites, and therefore the fact of their
containing rocky (siliceous) matter and metallic iron shows that outside
the earth the elements and their aggregation are in some degree the same
as upon the earth itself.
[1] The composition of meteoric iron is variable. It generally contains
nickel, phosphorus, carbon, &c. The schreibersite of meteoric
stones contains Fe_{4}Ni_{2}P.
[2] Comets and the rings of Saturn ought now to be considered as
consisting of an accumulation of such meteoric cosmic particles.
Perhaps the part played by these minute bodies scattered throughout
space is much more important in the formation of the largest
celestial bodies than has hitherto been imagined. The investigation
of this branch of astronomy, due to Schiaparelli, has a bearing on
the whole of natural science.
The question arises as to why the iron in meteorites is in a free
state, whilst on earth it is in a state of combination. Does not
this tend to show that the condition of our globe is very different
from that of the rest? My answer to this question has been already
given in Volume I. p. 377, Note 57. It is my opinion that inside
the earth there is a mass similar in composition to
meteorites--that is, containing rocky matter and metallic iron,
partly carburetted. In conclusion, I consider it will not be out of
place to add the following explanations. According to the theory of
the distribution of pressures (see my treatise, _On Barometrical
Levelling_, 1876, pages 48 _et seq._) in an atmosphere of mixed
gases, it follows that two gases, whose densities are _d_ and
_d__{1}, and whose relative quantities or partial pressures at a
certain distance from the centre of gravity are _h_ and _h__{1},
will, when at a greater distance from the centre of attraction,
present a different ratio of their masses _x_ : _x__{1}--that is,
of their partial pressures--which may be found by the equation
_d__{1}(log(_h_) - log(_x_)) = _d_(log(_h__{1}) - log(_x__{1})).
If, for instance, _d_ : _d__{1} = 2 : 1, and _h_ = _h__{1} (that is
to say, the masses are equal at the lower height) = 1000, then when
_x_ = 10 the magnitude of _x__{1} will not be 10 (_i.e._ the mass
of a gas at a higher level whose density = 1 will not be equal to
the mass of a gas whose density = 2, as was the case at a lower
level), but much greater--namely, _x__{1} = 100--that is, the
lighter gas will predominate over a heavier one at a higher level.
Therefore, when the whole mass of the earth was in a state of
vapour, the substances having a greater vapour density accumulated
about the centre and those with a lesser vapour density at the
surface. And as the vapour densities depend on the atomic and
molecular weights, those substances which have small atomic and
molecular weights ought to have accumulated at the surface, and
those with high atomic and molecular weights, which are the least
volatile and the easiest to condense, at the centre. Thus it
becomes apparent why such light elements as hydrogen, carbon,
nitrogen, oxygen, sodium, magnesium, aluminium, silicon,
phosphorus, sulphur, chlorine, potassium, calcium, and their
compounds predominate at the surface and largely form the earth's
crust. There is also now much iron in the sun, as spectrum analysis
shows, and therefore it must have entered into the composition of
the earth and other planets, but would have accumulated at the
centre, because the density of its vapour is certainly large and it
easily condenses. There was also oxygen near the centre of the
earth, but not sufficient to combine with the iron. The former, as
a much lighter element, principally accumulated at the surface,
where we at the present time find all oxidised compounds and even a
remnant of free oxygen. This gives the possibility not only of
explaining in accordance with cosmogonic theories the predominance
of oxygen compounds on the surface of the earth, with the
occurrence of unoxidised iron in the interior of the earth and in
meteorites, but also of understanding why the density of the whole
earth (over 5) is far greater than that of the rocks (1 to 3)
composing its crust. And if all the preceding arguments and
theories (for instance the supposition that the sun, earth, and all
the planets were formed of an elementary homogeneous mass, formerly
composed of vapours and gases) be true, it must be admitted that
the interior of the earth and other planets contains metallic
(unoxidised) iron, which, however, is only found on the surface as
aerolites. And then assuming that aerolites are the fragments of
planets which have crumbled to pieces so to say during cooling
(this has been held to be the case by astronomers, judging from the
paths of aerolites), it is readily understood why they should be
composed of metallic iron, and this would explain its occurrence in
the depths of the earth, which we assumed as the basis of our
theory of the formation of naphtha (Chapter VIII., Notes 57-60).
The most widely diffused terrestrial compound of iron is iron bisulphide,
FeS_{2}, or _iron pyrites_. It occurs in formations of both aqueous and
igneous origin, and sometimes in enormous masses. It is a substance
having a greyish-yellow colour, with a metallic lustre, and a specific
gravity of 5·0; it crystallises in the regular system.[2 bis]
[2 bis] Immense deposits of iron pyrites are known in various parts of
Russia. On the river Msta, near Borovitsi, thousands of tons are
yearly collected from the detritus of the neighbouring rocks. In
the Governments of Toula, Riazan, and in the Donets district
continuous layers of pyrites occur among the coal seams. Very thick
beds of pyrites are also known in many parts of the Caucasus. But
the deposits of the Urals are particularly vast, and have been
worked for a long time. Amongst these I will only indicate the
deposits on the Soymensky estate near the Kishteimsky works; the
Kaletinsky deposits near the Virhny-Isetsky works (containing 1-2
p.c. Cu); on the banks of the river Koushaivi near Koushvi (3-5
p.c. Cu), and the deposits near the Bogoslovsky works (3-5 p.c.
Cu). Iron pyrites (especially that containing copper which is
extracted after roasting) is now chiefly employed for roasting, as
a source of SO_{2}, for the manufacture of chamber sulphuric acid
(Vol. I. p. 291), but the remaining oxide of iron is perfectly
suitable for smelting into pig iron, although it gives a sulphurous
pig iron (the sulphur may be easily removed by subsequent
treatment, especially with the aid of ferro-manganese in Bessemer's
process). The great technical importance of iron pyrites leads to
its sometimes being imported from great distances; for instance,
into England from Spain. Besides which, when heated in closed
retorts FeS_{2} gives sulphur, and if allowed to oxidise in damp
air, green vitriol, FeSO_{4}.
The oxides are the principal ores used for producing metallic iron. The
majority of the ores contain ferric oxide, Fe_{2}O_{3}, either in a free
state or combined with water, or else in combination with ferrous oxide,
FeO. The species and varieties of iron ores are numerous and diverse.
Ferric oxide in a separate form appears sometimes as crystals of the
rhombohedric system, having a metallic lustre and greyish steel colour;
they are brittle, and form a red powder, specific gravity about 5·25.
Ferric oxide in type of oxidation and properties resembles alumina; it
is, however, although with difficulty, soluble in acids even when
anhydrous. The crystalline oxide bears the name of _specular iron ore_,
but ferric oxide most often occurs in a non-crystalline form, in masses
having a red fracture, and is then known as _red hæmatite_. In this form,
however, it is rather a rare ore, and is principally found in veins. The
hydrates of ferric oxide, ferric hydroxides,[3] are most often found in
aqueous or stratified formations, and are known as _brown hæmatites_;
they generally have a brown colour, form a yellowish-brown powder, and
have no metallic lustre but an earthy appearance. They easily dissolve in
acids and diffuse through other formations, especially clays (for
instance, ochre); they sometimes occur in reniform and similar masses,
evidently of aqueous origin. Such are, for instance, the so-called bog or
lake and peat ores found at the bottom of marshes and lakes, and also
under and in peat beds. This ore is formed from water containing ferrous
carbonate in solution, which, after absorbing oxygen, deposits ferric
hydroxide. In rivers and springs, iron is found in solution as ferrous
carbonate through the agency of carbonic acid: hence the existence of
chalybeate springs containing FeCO_{3}. This ferrous carbonate, or
_siderite_, is either found as a non-crystalline product of evidently
aqueous origin, or as a crystalline spar called _spathic iron ore_. The
reniform deposits of the former are most remarkable; they are called
spherosiderites, and sometimes form whole strata in the jurassic and
carboniferous formations. _Magnetic iron ore_, Fe_{3}O_{4} =
FeO,Fe_{2}O_{3}, in virtue of its purity and practical uses, is a very
important ore; it is a compound of the ferrous and ferric oxides, is
naturally magnetic, has a specific gravity of 5·1, crystallises in
well-formed crystals of the regular system, is with difficulty soluble in
acids, and sometimes forms enormous masses, as, for instance, Mount
Blagodat in the Ural. However, in most cases--for instance, at
Korsak-Mogila (to the north of Berdiansk and Nogaiska, near the Sea of
Azov), or at Krivoi Rog (to the west of Ekaterinoslav)--the magnetic iron
ore is mixed with other iron ores. In the Urals, the Caucasus (without
mentioning Siberia), and in the districts adjoining the basin of the Don,
Russia possesses the richest iron ores in the world. To the south of
Moscow, in the Governments of Toula and Nijninovgorod, in the Olonetz
district, and in the Government of Orloffsky (near Zinovieff in the
district of Kromsky), and in many other places, there are likewise
abundant supplies of iron ores amongst the deposited aqueous formations;
the siderite of Orloffsky, for instance, is distinguished by its great
purity.[4]
[3] The hydrated ferric oxide is found in nature in a dual form. It is
somewhat rarely met with in the form of a crystalline mineral
called _göthite_, whose specific gravity is 4·4 and composition
Fe_{2}H_{3}O_{4}, or FeHO_{2}--that is, one of oxide of iron to one
of water, Fe_{2}O_{3},H_{2}O; frequently found as brown ironstone,
forming a dense mass of fibrous, reniform deposits containing
2Fe_{2}O_{3},3H_{2}O--that is, having a composition
Fe_{4}H_{6}O_{9}. In bog ore and other similar ores we most often
find a mixture of this hydrated ferric oxide with clay and other
impurities. The specific gravity of such formations is rarely as
high as 4·0.
[4] The ores of iron, similarly to all substances extracted from veins
and deposits, are worked according to mining practice by means of
vertical, horizontal, or inclined shafts which reach and penetrate
the veins and strata containing the ore deposits. The mass of ore
excavated is raised to the surface, then sorted either by hand or
else in special sorting apparatus (generally acting with water to
wash the ore), and is subjected to roasting and other treatment. In
every case the ore contains foreign matter. In the extraction of
iron, which is one of the cheapest metals, the dressing of an ore
is in most cases unprofitable, and only ores rich in metal are
worked--namely, those containing at least 20 p.c. It is often
profitable to transport very rich and pure ores (with as much as 70
p.c. of iron) from long distances. The details concerning the
working and extraction of metals will be found in special treatises
on metallurgy and mining.
Iron is also found in the form of various other compounds--for instance,
in certain silicates, and also in some phosphates; but these forms are
comparatively rare in nature in a pure state, and have not the industrial
importance of those natural compounds of iron previously mentioned. In
small quantities iron enters into the composition of every kind of _soil_
and all rocky formations. As ferrous oxide, FeO, is isomorphous with
magnesia, and ferric oxide, Fe_{2}O_{3}, with alumina, isomorphous
substitution is possible here, and hence minerals are not unfrequently
found in which the quantity of iron varies considerably; such, for
instance, are pyroxene, amphibole, certain varieties of mica, &c.
Although much iron oxide is deleterious to the growth of vegetation,
still plants do not flourish without iron; it enters as an indispensable
component into the composition of all higher _organisms_; in the ash of
plants we always find more or less of its compounds. It also occurs in
blood, and forms one of the colouring matters in it; 100 parts of the
blood of the highest organisms contain about 0·05 of iron.
The _reduction_ of the ores of iron into metallic iron is in principle
very simple, because when the oxides of iron are strongly heated with
charcoal, hydrogen, carbonic oxide, and other reducing agents,[5] they
easily give metallic iron. But the matter is rendered more difficult by
the fact that the iron does not melt at the heat developed by the
combustion of the charcoal, and therefore it does not separate from those
mechanically mixed impurities which are found in the iron ore. This is
obviated by the following very remarkable property of iron: at a high
temperature it is capable of combining with a small quantity (from 2 to 5
p.c.) of carbon, and then forms _cast iron_, which easily _melts_ in the
heat developed by the combustion of charcoal in air. For this reason
metallic iron is not obtained directly from the ore, but is only formed
after the further treatment of the cast iron; the first product extracted
from the ore being cast iron. The fused mass disposes itself in the
furnace below the slag--that is, the impurities of the ore fused by the
heat of the furnace. If these impurities did not fuse they would block up
the furnace in which the ore was being smelted, and the continuous
smelting of the cast iron would not be possible;[6] it would be necessary
periodically to cool the furnace and heat it up again, which means a
wasteful expenditure of fuel, and hence in the production of cast iron,
the object in view is to obtain all the earthy impurities of the ore in
the shape of a fused mass or slag. Only in rare cases does the ore itself
form a mass which fuses at the temperature employed, and these cases are
objectionable if much iron oxide is carried away in the slag. The
impurities of the ores most often consist of certain mixtures--for
instance, a mixture of clay and sand, or a mixture of limestone and clay,
or quartz, &c. These impurities do not separate of themselves, or do not
fuse. The difficulty of the industry lies in forming an easily-fusible
slag, into which the whole of the foreign matter of the ore would pass
and flow down to the bottom of the furnace above the heavier cast iron.
This is effected by mixing certain _fluxes_ with the ore and charcoal. A
flux is a substance which, when mixed with the foreign matter of the ore,
forms a fusible vitreous mass or slag. The flux used for silica is
limestone with clay; for limestone a definite quantity of silica is used,
the best procedure having been arrived at by experiment and by long
practice in iron smelting and other metallurgical processes.[7]
[5] The reduction of iron oxides by hydrogen belongs to the order of
reversible reactions (Chapter II.), and is therefore determined by
a limit which is here expressed by the attainment of the same
pressure as in the case where hydrogen acts on iron oxides, and as
in the case where (at the same temperature) water is decomposed by
metallic iron. The calculations referring to this matter were made
by Henri Sainte-Claire Deville (1870). Spongy iron was placed in a
tube having a temperature _t_, one end of which was connected with
a vessel containing water at 0° (vapour tension = 4·6 mm.) and the
other end with a mercury pump and pressure gauge which determined
the limiting tension attained by the dry hydrogen _p_ (subtracting
the tension of the water vapour from the tension observed). A tube
was then taken containing an excess of iron oxide. It was filled
with hydrogen, and the tension _p__{1} observed of the residual
hydrogen when the water was condensed at 0°.
_t_ = 200° 440° 860° 1040°
_p_ = 95·9 25·8 12·8 9·2 mm.
_p_{1} = -- -- 12·8 9·4 mm.
The equality of the pressure (tension) of the hydrogen in the two
cases is evident. The hydrogen here behaves like the vapour of iron
or of its oxide.
By taking ferric oxide, Fe_{2}O_{3}, Moissan observed that at 350°
it passed into magnetic oxide, Fe_{3}O_{4}, at 500° into ferrous
oxide, FeO, and at 600° into metallic iron. Wright and Luff (1878),
whilst investigating the reduction of oxides, found that (_a_) the
temperature of reaction depends on the condition of the oxide
taken--for instance, precipitated ferric oxide is reduced by
hydrogen at 85°, that obtained by oxidising the metal or from its
nitrate at 175°; (_b_) when other conditions are the same, the
reduction by carbonic oxide commences earlier than that by
hydrogen, and the reduction by hydrogen still earlier than that by
charcoal; (_c_) the reduction is effected with greater facility
when a greater quantity of heat is evolved during the reaction.
Ferric oxide obtained by heating ferrous sulphate to a red heat
begins to be reduced by carbonic oxide at 202°, by hydrogen at
260°, by charcoal at 430°, whilst for magnetic oxide, Fe_{3}O_{4},
the temperatures are 200°, 290°, and 450° respectively.
[6] The primitive methods of iron manufacture were conducted by
intermittent processes in hearths resembling smiths' fires. As
evidenced by the uninterrupted action of the steam boiler, or the
process of lime burning, and the continuous preparation and
condensation of sulphuric acid or the uninterrupted smelting of
iron, every industrial process becomes increasingly profitable and
complete under the condition of the continuous action, as far as
possible, of all agencies concerned in the production. This
continuous method of production is the first condition for the
profitable production on the large scale of nearly all industrial
products. This method lessens the cost of labour, simplifies the
supervision of the work, renders the product uniform, and
frequently introduces a very great economy in the expenditure of
fuel and at the same time presents the simplicity and perfection of
an equilibrated system. Hence every manufacturing operation should
be a continuous one, and the manufacture of pig iron and sulphuric
acid, which have long since become so, may be taken as examples in
many respects. A study of these two manufactures should form the
commencement of an acquaintance with all the contemporary methods
of manufacturing both from a technical and economical point of
view.
[7] The composition of slag suitable for iron smelting most often
approaches the following: 50 to 60 p.c. SiO_{2}, 5 to 20
Al_{2}O_{3}, the rest of the mass consisting of MgO, CaO, MnO, FeO.
Thus the most fusible slag (according to the observations of
Bodeman) contains the alloy Al_{2}O_{3},4CaO,7SiO_{2}. On altering
the quantity of magnesia and lime, and especially of the alkalis
(which increases the fusibility) and of silica (which decreases
it), the temperature of fusion changes with the relation between
the total quantity of oxygen and that in the silica. Slags of the
composition RO,SiO_{2} are easily fusible, have a vitreous
appearance, and are very common. Basic slags approach the
composition 2RO,SiO_{2}. Hence, knowing the composition and
quantity of the foreign matter in the ore, it is at once easy to
find the quantity and quality of the flux which must be added to
form a suitable slag. The smelting of iron is rendered more complex
by the fact that the silica, SiO_{2}, which enters into the slag
and fluxes is capable of forming a slag with the iron oxides. In
order that the least quantity of iron may pass into the slag, it is
necessary for it to be reduced before the temperature is attained
at which the slags are formed (about 1000°), which is effected by
reducing the iron, not with charcoal itself, but with carbonic
oxide. From this it will be understood how the progress of the
whole treatment may be judged by the properties of the slags.
Details of this complicated and well-studied subject will be found
in works on metallurgy.
Thus the following materials have to be introduced into the furnace
where the smelting of the iron ore is carried on: (1) the iron ore,
composed of oxide of iron and foreign matter; (2) the flux required to
form a fusible slag with the foreign matter; (3) the carbon which is
necessary (_a_) for reducing, (_b_) for combining with the reduced iron
to form cast iron, (_c_) principally for the purpose of combustion and
the heat generated thereby, necessary not only for reducing the iron and
transforming it into cast iron, but also for melting the slag, as well as
the cast iron--and (4) the air necessary for the combustion of the
charcoal. The air is introduced after a preparatory heating in order to
economise fuel and to obtain the highest temperature. The air is forced
in under pressure by means of a special blast arrangement. This permits
of an exact regulation of the heat and rate of smelting. All these
component parts necessary for the smelting of iron must be contained in a
vertical, that is, _shaft furnace_, which at the base must have a
receptacle for the accumulation of the slag and cast iron formed, in
order that the operation may proceed without interruption. The walls of
such a furnace ought to be built of fireproof materials if it be designed
to serve for the continuous production of cast iron by charging the ore,
fuel, and flux into the mouth of the furnace, forcing a blast of air into
the lower part, and running out the molten iron and slag from below. The
whole operation is conducted in furnaces known as _blast furnaces_. The
annexed illustration, fig. 93 (which is taken by kind permission from
Thorpe's Dictionary of Applied Chemistry), represents the vertical
section of such a furnace. These furnaces are generally of large
dimensions--varying from 50 to 90 feet in height. They are sometimes
built against rising ground in order to afford easy access to the top
where the ore, flux, and charcoal or coke are charged.[8]
[8] The section of a blast furnace is represented by two truncated
cones joined at their bases, the upper cone being longer than the
lower one; the lower cone is terminated by the hearth, or almost
cylindrical cavity in which the cast iron and slag collect, one
side being provided with apertures for drawing off the iron and
slag. The air is blown into the blast furnace through special
pipes, situated over the hearth, as shown in the section. The air
previously passes through a series of cast-iron pipes, heated by
the combustion of the carbonic oxide obtained from the upper parts
of the furnace, where it is formed as in a 'gas-producer.' The
blast furnace acts continuously until it is worn out; the iron is
tapped off twice a day, and the furnace is allowed to cool a little
from time to time so as not to be spoilt by the increasing heat,
and to enable it to withstand long usage.
Blast furnaces worked with charcoal fuel are not so high, and in
general give a smaller yield than those using coke, because the
latter are worked with heavier charges than those in which charcoal
is employed. Coke furnaces yield 20,000 tons and over of pig iron a
year. In the United States there are blast furnaces 30 metres high,
and upwards of 600 cubic metres capacity, yielding as much as
130,000 tons of pig iron, requiring a blast of about 750 cubic
metres of air per minute, heated to 600°, and consuming about 0·85
part of coke per 1 part of pig iron produced. At the present time
the world produces as much as 30 million tons of pig iron a year,
about 9/10 of which is converted into wrought iron and steel. The
chief producers are the United States (about 10 million tons a
year) and England (about 9 million tons a year); Russia yields
about 1-1/5 million tons a year. The world's production has doubled
during the last 20 years, and in this respect the United States
have outrun all other countries. The reason of this increase of
production must be looked for in the increased demand for iron and
steel for railway purposes, for structures (especially
ship-building), and in the fact that: (_a_) the cost of pig iron
has fallen, thanks to the erection of large furnaces and a fuller
study of the processes taking place in them, and (_b_) that every
kind of iron ore (even sulphurous and phosphoritic) can now be
converted into a homogeneous steel.
[Illustration: FIG. 93.--Vertical section of a modern Cleveland
blast furnace capable of producing 300 to 1,000 tons of pig iron
weekly. The outer casing is of riveted iron plates, the furnace
being lined with refractory fire-brick. It is closed at the top by
a 'cap and cone' arrangement, by means of which the charge can be
fed into the furnace at suitable intervals by lowering the moveable
cone.]
In order to more thoroughly grasp the chemical process which takes
place in blast furnaces, it is necessary to follow the course of
the material charged in at the top and of the air passing through
the furnace. From 50 to 200 parts of carbon are expended on 100
parts of iron. The ore, flux, and coke are charged into the top of
the furnace, in layers, as the cast iron is formed in the lower
parts and flowing down to the bottom causes the whole contents of
the furnace to subside, thus forming an empty space at the top,
which is again filled up with the afore-mentioned mixture. During
its downward course this mixture is subjected to increasing heat.
This rise of temperature first drives off the moisture of the ore
mixture, and then leads to the formation of the products of the dry
distillation of coal or charcoal. Little by little the subsiding
mass attains a temperature at which the heated carbon reacts with
the carbonic anhydride passing upwards through the furnace and
transforms it into carbonic oxide. This is the reason why carbonic
anhydride is not evolved from the furnace, but only carbonic oxide.
As regards the ore itself, on being heated to about 600° to 800° it
is reduced at the expense of the carbonic oxide ascending the
furnace, and formed by the contact of the carbonic anhydride with
the incandescent charcoal, so that the reduction in the blast
furnace is without doubt brought about _by_ the formation and
decomposition of _carbonic oxide_ and not by carbon itself--thus,
Fe_{2}O_{3} + 3CO = Fe_{2} + 3CO_{2}. The reduced iron, on further
subsidence and contact with carbon, forms cast iron, which flows to
the bottom of the furnace. In these lower layers, where the
temperature is highest (about 1,300°), the foreign matter of the
ore finally forms slag, which also is fusible, with the aid of
fluxes. The air blown in from below, through the so-called
_tuyeres_, encounters carbon in the lower layers of the furnace,
and burns it, converting it into carbonic anhydride. It is evident
that this develops the highest temperature in these lower layers of
the furnace, because here the combustion of the carbon is effected
by heated and compressed air. This is very essential, for it is by
virtue of this high temperature that the process of forming the
slag and of forming and fusing the cast iron are effected
simultaneously in these lower portions of the furnace. The carbonic
acid formed in these parts rises higher, encounters incandescent
carbon, and forms with it carbonic oxide. This heated carbonic
oxide acts as a reducing agent on the iron ore, and is reconverted
by it into carbonic anhydride; this gas meets with more carbon, and
again forms carbonic oxide, which again acts as a reducing agent.
The final transformation of the carbonic anhydride into carbonic
oxide is effected in those parts of the furnace where the reduction
of the oxides of iron does not take place, but where the
temperature is still high enough to reduce the carbonic anhydride.
The ascending mixture of carbonic oxide and nitrogen, CO_{2}, &c.,
is then withdrawn through special lateral apertures formed in the
upper cold parts of the furnace walls, and is conducted through
pipes to those stoves which are used for heating the air, and also
sometimes into other furnaces used for the further processes of
iron manufacture. The fuel of blast furnaces consists of wood
charcoal (this is the most expensive material, but the pig iron
produced is the purest, because charcoal does not contain any
sulphur, while coke does), anthracite (for instance, in
Pennsylvania, and in Russia at Pastouhoff's works in the Don
district), coke, coal, and even wood and peat. It must be borne in
mind that the utilisation of naphtha and naphtha refuse would
probably give very profitable results in metallurgical processes.
The process just described is accompanied by a series of other
processes. Thus, for instance, in the blast furnace a considerable
quantity of cyanogen compounds are formed. This takes place because
the nitrogen of the air blast comes into contact with incandescent
carbon and various alkaline matters contained in the foreign matter
of the ores. A considerable quantity of potassium cyanide is formed
when wood charcoal is employed for iron smelting, as its ash is
rich in potash.
The _cast iron_ formed in blast furnaces is not always of the same
quality. When slowly cooled it is soft, has a grey colour, and is not
completely soluble in acids. When treated with acids a residue of
graphite remains; it is known as _grey_ or soft cast iron. This is the
general form of the ordinary cast iron used for casting various objects,
because in this state it is not so brittle as in the shape of _white cast
iron_, which does not leave particles of graphite when dissolved, but
yields its carbon in the form of hydrocarbons. This white cast iron is
characterised by its whitish-grey colour, dull lustre, the crystalline
structure of its fracture (more homogeneous than that of grey iron), and
such hardness that a file will hardly cut it. When white cast iron is
produced (from manganese ore) at high temperatures (and with an excess of
lime), and containing little sulphur and silica but a considerable amount
of carbon (as much as 5 p.c.), it acquires a coarse crystalline structure
which increases in proportion to the amount of manganese, and it is then
known under the name of 'spiegeleisen' (and 'ferro-manganese').[9]
[9] The specific gravity of white cast iron is about 7·5. Grey cast
iron has a much lower specific gravity, namely, 7·0. Grey cast iron
generally contains less manganese and more silica than white; but
both contain from 2 to 3 p.c. of carbon. The difference between the
varieties of cast iron depends on the condition of the carbon which
enters into the composition of the iron. In white cast iron the
carbon is in combination with the iron--in all probability, as the
compound CFe_{4} (Abel and Osmond and others extracted this
compound, which is sometimes called 'carbide,' from tempered steel,
which stands to unannealed steel as white cast iron does to grey),
but perhaps in the state of an indefinite chemical compound
resembling a solution. In any case the compound of the iron and
carbon in white cast iron is chemically very unstable, because when
slowly cooled it decomposes, with separation of graphite, just as a
solution when slowly cooled yields a portion of the substance
dissolved. The separation of carbon in the form of graphite on the
conversion of white cast iron into grey is never complete, however
slowly the separation be carried on; part of the carbon remains in
combination with the iron in the same state in which it exists in
white cast iron. Hence when grey cast iron is treated with acids,
the whole of the carbon does not remain in the form of graphite,
but a part of it is separated as hydrocarbons, which proves the
existence of chemically-combined carbon in grey cast iron. It is
sufficient to re-melt grey cast iron and to cool it quickly to
transform it into white cast iron. It is not carbon alone that
influences the properties of cast iron; when it contains a
considerable amount of sulphur, cast iron remains white even after
having been slowly cooled. The same is observed in cast iron very
rich in manganese (5 to 7 p.c.), and in this latter case the
fracture is very distinctly crystalline and brilliant. When cast
iron contains a large amount of manganese, the quantity of carbon
may also be increased. Crystalline varieties of cast iron rich in
manganese are in practice called ferro-manganese (p. 310), and are
prepared for the Bessemer process. Grey cast iron not having an
uniform structure is much more liable to various changes than dense
and thoroughly uniform white cast iron, and the latter oxidises
much more slowly in air than the former. White cast iron is not
only used for conversion into wrought iron and steel, but also in
those cases where great hardness is required, although it be
accompanied by a certain brittleness; for instance, for making
rollers, plough-shares, &c.
Cast iron is a material which is either suitable for direct
application for casting in moulds or else for working up into _wrought
iron_ and _steel_. The latter principally differ from cast iron in their
containing less carbon--thus, steel contains from 1 p.c. to 0·5 p.c. of
carbon and far less silicon and manganese than cast iron; wrought iron
does not generally contain more than 0·25 p.c. of carbon and not more
than 0·25 p.c. of the other impurities. Thus the essence of the working
up of cast iron into steel and wrought iron consists in the removal of
the greater part of the carbon and other elements, S, P, Mn, Si, &c. This
is effected by means of oxidation, because the oxygen of the atmosphere,
oxidising the iron at a high temperature, forms solid oxides with it; and
the latter, coming into contact with the carbon contained in the cast
iron, are deoxidised, forming wrought iron and carbonic oxide, which is
evolved from the mass in a gaseous form. It is evident that the oxidation
must be carried on with a molten mass in a state of agitation, so that
the oxygen of the air may be brought into contact with the whole mass of
carbon contained in the cast iron, or else the operation is effected by
means of the addition of oxygen compounds of iron (oxides, ores, as in
Martin's process). Cast iron melts much more easily than wrought iron and
steel, and, therefore, as the carbon separates, the mass in the furnace
(in puddling) or hearth (in the bloomery process) becomes more and more
solid; moreover the degree of hardness forms, to a certain extent, a
measure of the amount of carbon separated, and the operation may
terminate either in the formation of steel or wrought iron.[10] In any
case, the iron used for industrial purposes contains impurities.
_Chemically pure iron_ may be obtained by precipitating iron from a
solution (a mixture of ferrous sulphate with magnesium sulphate or
ammonium chloride) by the prolonged action of a feeble galvanic current;
the iron may be then obtained as a dense mass. This method, proposed by
Böttcher and applied by Klein, gives, as R. Lenz showed, iron containing
occluded hydrogen, which is disengaged on heating. This galvanic
deposition of iron is used for making galvanoplastic _clichés_, which are
distinguished for their great hardness. Electro-deposited iron is
brittle, but if heated (after the separation of the hydrogen) it becomes
soft. If pure ferric hydroxide, which is easily prepared by the
precipitation of solutions of ferric salts by means of ammonia, be heated
in a stream of hydrogen, it forms, first of all, a dull black powder
which ignites spontaneously in air (pyrophoric iron), and then a grey
powder of pure iron. The powdery substance first obtained is an iron
suboxide; when thrown into the air it ignites, forming the oxide
Fe_{3}O_{4}. If the heating in hydrogen be continued, more water and pure
iron, which does not ignite spontaneously, will be obtained. If a small
quantity of iron be fused in the oxyhydrogen flame (with an excess of
oxygen) in a piece of lime and mixed with powdered glass, pure molten
iron will be formed, because in the oxyhydrogen flame iron melts and
burns, but the substances mixed with the iron oxidise first. The oxidised
impurities here either disappear (carbonic anhydride) in a gaseous form,
or turn into slag (silica, manganese, oxide, and others)--that is, fuse
with the glass. Pure iron has a silvery white colour and a specific
gravity of 7·84; it melts at a temperature higher than the melting-points
of silver, gold, nickel, and steel, _i.e._ about 1400°-1500° and below
the melting point of platinum (1750°).[11] But pure iron becomes soft at
a temperature considerably below that at which it melts, and may then be
easily forged, welded, and rolled or drawn into sheets and wire.[11 bis]
Pure iron may be rolled into an exceedingly thin sheet, weighing less
than a sheet of ordinary paper of the same size. This ductility is the
most important property of iron in all its forms, and is most marked with
sheet iron, and least so with cast iron, whose ductility, compared with
wrought iron, is small, but it is still very considerable when compared
with other substances--such, for instance, as rocks.[12]
[10] This direct process of separating the carbon from cast iron is
termed _puddling_. It is conducted in reverberatory furnaces. The
cast iron is placed on the bed of the furnace and melted; through
a special aperture, the puddler stirs up the oxidising mass of
cast iron, pressing the oxides into the molten iron. This
resembles kneading dough, and the process introduced in England
became known as puddling. It is evident that the puddled mass, or
bloom, is a heterogeneous substance obtained by mixing, and hence
one part of the mass will still be rich in carbon, another will be
poor, some parts will contain oxide not reduced, &c. The further
treatment of the puddled mass consists in hammering and drawing it
out into flat pieces, which on being hammered become more
homogeneous, and when several pieces are welded together and again
hammered out a still more homogeneous mass is obtained. The
quality of the steel and iron thus formed depends principally on
their uniformity. The want of uniformity depends on the oxides
remaining inside the mass, and on the variable distribution of the
carbon throughout the mass. In order to obtain a more homogeneous
metal for manufacturing articles out of steel, it is drawn into
thin rods, which are tied together in bundles and then again
hammered out. As an example of what may be attained in this
direction, imitation Damascus steel may be cited; it consists of
twisted and plaited wire, which is then hammered into a dense
mass. (Real damascened wootz steel may be made by melting a
mixture of the best iron with graphite (1/12) and iron rust; the
article is then corroded with acid, and the carbon remains in the
form of a pattern.)
Steel and wrought iron are manufactured from cast iron by
puddling. They are, however, obtained not only by this method but
also by the _bloomery process_, which is carried out in a fire
similar to a blacksmith's forge, fed with charcoal and provided
with a blast; a pig of cast iron is gradually pushed into the
fire, and portions of it melt and fall to the bottom of the
hearth, coming into contact with an air blast, and are thus
oxidised. The bloom thus formed is then squeezed and hammered. It
is evident that this process is only available when the charcoal
used in the fire does not contain any foreign matter which might
injure the quality of the iron or steel--for instance, sulphur or
phosphorus--and therefore only wood charcoal may be used with
impunity, from which it follows that this process can only be
carried on where the manufacture of iron can be conducted with
this fuel. Coal and coke contain the above-mentioned impurities,
and would therefore produce iron of a brittle nature, and thus it
would be necessary to have recourse to puddling, where the fuel is
burnt on a special hearth, separate from the cast iron, whereby
the impurities of the fuel do not come into contact with it. The
manufacture of steel from cast iron may also be conducted in
fires; but, in addition to this, it is also now prepared by many
other methods. One of the long-known processes is called
_cementation_, by which steel is prepared from wrought iron but
not from cast iron. For this process strips of iron are heated
red-hot for a considerable time whilst immersed in powdered
charcoal; during this operation the iron at the surface combines
with the charcoal, which however does not penetrate; after this
the iron strips are re-forged, drawn out again, and cemented anew,
repeating this process until a steel of the desired quality is
formed--that is, containing the requisite proportion of carbon.
The _Bessemer_ process occupies the front rank among the newer
methods (since 1856); it is so called from the name of its
inventor. This process consists in running melted cast iron into
converters (holding about 6 tons of cast iron)--that is,
egg-shaped receivers, fig. 94, capable of revolving on trunnions
(in order to charge in the cast iron and discharge the steel), and
forcing a stream of air through small apertures at a considerable
pressure. Combustion of the iron and carbon at an elevated
temperature then takes place, resulting from the bubbles of oxygen
thus penetrating the mass of the cast iron. The carbon, however,
burns to a greater extent than the iron, and therefore a mass is
obtained which is much poorer in carbon than cast iron. As the
combustion proceeds very rapidly in the mass of metal, the
temperature rises to such an extent that even the wrought iron
which may be formed remains in a molten condition, whilst the
steel, being more fusible than the wrought iron, remains very
liquid. In half an hour the mass is ready. The purest possible
cast iron is used in the Bessemer process, because sulphur and
phosphorus do not burn out like carbon, silicon, and manganese.
[Illustration: FIG. 94.--Bessemer converter, constructed of iron
plate and lined with ganister. The air is carried by the tubes, L,
O, D to the bottom, M, from which it passes by a number of holes
into the converter. The converter is rotated on the trunnion _d_
by means of the rack and pinion H, when it is required either to
receive molten cast iron from the melting furnaces or to pour out
the steel.]
The presence of manganese enables the sulphur to be removed with
the slag, and the presence of lime or magnesia, which are
introduced into the lining of the converter, facilitates the
removal of the phosphorus. This basic Bessemer process, or _Thomas
Gilchrist process_, introduced about 1880, enables ores containing
a considerable amount of phosphorus, which had hitherto only been
used for cast iron, to be used for making wrought iron and steel.
Naturally the greatest uniformity will be obtained by re-melting
the metal. Steel is re-melted in small wind furnaces, in masses
not exceeding 30 kilos; a liquid metal is formed, which may be
cast in moulds. A mixture of wrought and cast iron is often used
for making cast steel (the addition of a small amount of metallic
Al improves the homogeneity of the castings, by facilitating the
passage of the impurities into slag). Large steel castings are
made by simultaneous fusion in several furnaces and crucibles; in
this way, castings up to 80 tons or more, such as large ordnance,
may be made. This molten, and therefore homogeneous, steel is
called _cast steel_. Of late years the _Martin's process_ for the
manufacture of steel has come largely into use; it was invented in
France about 1860, and with the use of regenerative furnaces it
enables large quantities of cast steel to be made at a time. It is
based on the melting of cast iron with iron oxides and iron
itself--for instance, pure ores, scrap, &c. There the carbon of
the cast iron and the oxygen of the oxide form carbonic oxide, and
the carbon therefore burns out, and thus cast steel is obtained
from cast iron, providing, naturally, that there is a requisite
proportion and corresponding degree of heat. The advantage of this
process is that not only do the carbon, silicon, and manganese,
but also a great part of the sulphur and phosphorus of the cast
iron burn out at the expense of the oxygen of the iron oxides.
During the last decade the manufacture of steel and its
application for rails, armour plate, guns, boilers, &c., has
developed to an enormous extent, thanks to the invention of cheap
processes for the manufacture of large masses of homogeneous cast
steel. Wrought iron may also be melted, but the heat of a blast
furnace is insufficient for this. It easily melts in the
oxyhydrogen flame. It may be obtained in a molten state directly
from cast iron, if the latter be melted with nitre and
sufficiently stirred up. Considerable oxidation then takes place
inside the mass of cast iron, and the temperature rises to such an
extent that the wrought iron formed remains liquid. A method is
also known for obtaining wrought iron directly from rich iron ores
by the action of carbonic oxide: the wrought iron is then formed
as a spongy mass (which forms an excellent filter for purifying
water), and may be worked up into wrought iron or steel either by
forging or by dissolving in molten cast iron.
Everybody is more or less familiar with the _difference in the
properties of steel and wrought iron_. Iron is remarkable for its
softness, pliability, and small elasticity, whilst steel may be
characterised by its capability of attaining elasticity and
hardness if it be cooled suddenly after having been heated to a
definite temperature, or, as it is termed, _tempered_. But if
tempered steel be re-heated and slowly cooled, it becomes as soft
as wrought iron, and can then be cut with the file and forged, and
in general can be made to assume any shape, like wrought iron. In
this soft condition it is called _annealed steel_. The transition
from tempered to annealed steel thus takes place in a similar way
to the transition from white to grey cast iron. Steel, when
homogeneous, has considerable lustre, and such a fine granular
structure that it takes a very high polish. Its fracture clearly
shows the granular nature of its structure. The possibility of
tempering steel enables it to be used for making all kinds of
cutting instruments, because annealed steel can be forged, turned,
drawn (under rollers, for instance, for making rails, bars, &c.),
filed, &c., and it may then be tempered, ground and polished. The
method and temperature of tempering and annealing steel determine
its hardness and other qualities. Steel is generally tempered to
the required degree of hardness in the following manner: It is
first strongly heated (for instance, up to 600°), and then plunged
into water--that is, hardened by rapid cooling (it then becomes as
brittle as glass). It is then heated until the surface assumes a
definite colour, and finally cooled either quickly or slowly. When
steel is heated up to 220°, its surface acquires a yellow colour
(surgical instruments); it first of all becomes straw-coloured
(razors, &c.), and then gold-coloured; then at a temperature of
250° it becomes brown (scissors), then red, then light blue at
285° (springs), then indigo at 300° (files), and finally sea-green
at about 340°. These colours are only the tints of thin films,
like the hues of soap bubbles, and appear on the steel because a
thin layer of oxides is formed over its surface. Steel rusts more
slowly than wrought iron, and is more soluble in acids than cast
iron, but less so than wrought iron. Its specific gravity is about
7·6 to 7·9.
As regards the formation of steel, it was a long time before the
process of cementation was thoroughly understood, because in this
case infusible charcoal permeates unfused wrought iron. Caron
showed that this permeation depends on the fact that the charcoal
used in the process contains alkalis, which, in the presence of
the nitrogen of the air, form metallic cyanides; these being
volatile and fusible, permeate the iron, and, giving up their
carbon to it, serve as the material for the formation of steel.
This explanation is confirmed by the fact that charcoal without
alkalis or without nitrogen will not cement iron. The charcoal
used for cementation acts badly when used over again, as it has
lost alkali. The very volatile ammonium cyanide easily conduces to
the formation of steel. Although steel is also formed by the
action of cyanogen compounds, nevertheless it does not contain
more nitrogen than cast or wrought iron (0·01 p.c.), and these
latter contain it because their ores contain titanium, which
combines directly with nitrogen. Hence the part played by nitrogen
in steel is but an insignificant one. It may be useful here to add
some information taken from Caron's treatise concerning the
influence of foreign matter on the quality of steel. The principal
properties of steel are those of tempering and annealing. The
compounds of iron with silicon and boron have not these
properties. They are more stable than the carbon compound, and
this latter is capable of changing its properties; because the
carbon in it either enters into combination or else is disengaged,
which determines the condition of hardness or softness of steel,
as in white and grey cast iron. When slowly cooled, steel splits
up into a mixture of soft and carburetted iron; but, nevertheless,
the carbon does not separate from the iron. If such steel be again
heated, it forms a uniform compound, and hardens when rapidly
cooled. If the same steel as before be taken and heated a long
time, then, after being slowly cooled, it becomes much more
soluble in acid, and leaves a residue of pure carbon. This shows
that the combination between the carbon and iron in steel becomes
destroyed when subjected to heat, and the steel becomes iron mixed
with carbon. Such _burnt_ steel cannot be tempered, but may be
corrected by continued forging in a heated condition, which has
the effect of redistributing the carbon equally throughout the
whole mass. After the forging, if the iron is pure and the carbon
has not been burnt out, steel is again formed, which may be
tempered. If steel be repeatedly or strongly heated, it becomes
burnt through and cannot be tempered or annealed; the carbon
separates from the iron, and this is effected more easily if the
steel contains other impurities which are capable of forming
stable combinations with iron, such as silicon, sulphur, or
phosphorus. If there be much silicon, it occupies the place of the
carbon, and then continued forging will not induce the carbon once
separated to re-enter into combination. Such steel is easily burnt
through and cannot be corrected; when burnt through, it is hard
and cannot be annealed--this is tough steel, an inferior kind.
Iron which contains sulphur and phosphorus cements badly, combines
but little with carbon, and steel of this kind is brittle, both
hot and cold. Iron in combination with the above-mentioned
substances cannot be annealed by slow cooling, showing that these
compounds are more stable than those of carbon and iron, and
therefore they prevent the formation of the latter. Such metals as
tin and zinc combine with iron, but not with carbon, and form a
brittle mass which cannot be annealed and is deleterious to steel.
Manganese and tungsten, on the contrary, are capable of combining
with charcoal; they do not hinder the formation of steel, but even
remove the injurious effects of other admixtures (by transforming
these admixed substances into new compounds and slags), and are
therefore ranked with the substances which act beneficially on
steel; but, nevertheless, the best steel, which is capable of
renewing most often its primitive qualities after burning or hot
forging, is the purest. The addition of Ni, Cr, W, and certain
other metals to steel renders it very suitable for certain special
purposes, and is therefore frequently made use of.
It is worthy of attention that steel, besides temper, possesses
many variable properties, a review of which may be made in the
classification of the _sorts of steel_ (1878, Cockerell). (1)
_Very mild steel_ contains from 0·05 to 0·20 p.c. of carbon,
breaks with a weight of 40 to 50 kilos per square millimetre, and
has an extension of 20 to 30 p.c.; it may be welded, like wrought
iron, but cannot be tempered; is used in sheets for boilers,
armour plate and bridges, nails, rivets, &c., as a substitute for
wrought iron; (2) _mild steel_, from 0·20 to 0·35 p.c. of carbon,
resistance to tension 50 to 60 kilos, extension 15 to 20 p.c., not
easily welded, and tempers badly, used for axles, rails, and
railway tyres, for cannons and guns, and for parts of machines
destined to resist bending and torsion; (3) _hard steel_, carbon
0·35 to 0·50 p.c., breaking weight 60 to 70 kilos per square
millimetre, extension 10 to 15 p.c., cannot be welded, takes a
temper; used for rails, all kinds of springs, swords, parts of
machinery in motion subjected to friction, spindles of looms,
hammers, spades, hoes, &c.; (4) _very hard steel_, carbon 0·5 to
0·65 p.c., tensile breaking weight 70 to 80 kilos, extension 5 to
10 p.c., does not weld, but tempers easily; used for small
springs, saws, files, knives and similar instruments.
The properties of ordinary _wrought iron_ are well known. The best
iron is the most tenacious--that is to say, that which does not
break up when struck with the hammer or bent, and yet at the same
time is sufficiently hard. There is, however, a distinction
between hard and soft iron. Generally the softest iron is the most
tenacious, and can best be welded, drawn into wire, sheets, &c.
Hard, especially tough, iron is often characterised by its
breaking when bent, and is therefore very difficult to work, and
objects made from it are less serviceable in many respects. Soft
iron is most adapted for making wire and sheet iron and such small
objects as nails. Soft iron is characterised by its attaining a
fibrous fracture after forging, whilst tough iron preserves its
granular structure after this operation. Certain sorts of iron,
although fairly soft at the ordinary temperature, become brittle
when heated and are difficult to weld. These sorts are less
suitable for being worked up into small objects. The variety of
the properties of iron depends on the impurities which it
contains. In general, the iron used in the arts still contains
carbon and always a certain quantity of silicon, manganese,
sulphur, phosphorus, &c. A variety in the proportion of these
component parts changes the quality of the iron. In addition to
this the change which soft wrought iron, having a fibrous
structure, undergoes when subjected to repeated blows and
vibrations is considerable; it then becomes granular and brittle.
This to a certain degree explains the want of stability of some
iron objects--such as truck axles, which must be renewed after a
certain term of service, otherwise they become brittle. It is
evident that there are innumerable intermediate transitions from
wrought iron to steel and cast iron.
At the present day the greater part of the cast iron manufactured
is converted into steel, generally cast steel (Bessemer's and
Martin's). I may add the Urals, Donetz district, and other parts
of Russia offer the greatest advantages for the development of an
iron industry, because these localities not only contain vast
supplies of excellent iron ore, but also coal, which is necessary
for smelting it.
[11] According to information supplied by A. T. Skinder's experiments
at the Oboukoff Steel Works, 140 volumes of liquid molten steel
give 128 volumes of solid metal. By means of a galvanic current of
great intensity and dense charcoal as one electrode, and iron as
the other, Bernadoss welded iron and fused holes through sheet
iron. Soft wrought iron, like steel and soft malleable cast iron,
may be melted in Siemens' regenerative furnaces, and in furnaces
heated with naphtha.
[11 bis] Gore (1869), Tait, Barret, Tchernoff, Osmond, and others
observed that at a temperature approaching 600°--that is, between
dark and bright red heat--all kinds of wrought iron undergo a
peculiar change called _recalescence_, _i.e._ a spontaneous rise
of temperature. If iron be considerably heated and allowed to
cool, it may be observed that at this temperature the cooling
stops--that is, latent heat is disengaged, corresponding with a
change in condition. The specific heat, electrical conductivity,
magnetic, and other properties then also change. In tempering, the
temperature of recalescence must not be reached, and so also in
annealing, &c. It is evident that a change of the internal
condition is here encountered, exactly similar to the transition
from a solid to a liquid, although there is no evident physical
change. It is probable that attentive study would lead to the
discovery of a similar change in other substances.
[12] The particles of steel are linked together or connected more
closely than those of the other metals; this is shown by the fact
that it only breaks with a tensile strain of 50-80 kilos per sq.
mm., whilst wrought iron only withstands about 30 kilos, cast iron
10, copper 35, silver 23, platinum 30, wood 8. The elasticity of
iron, steel, and other metals is expressed by the so-called
_coefficient of elasticity_. Let a rod be taken whose length is L;
if a weight, P, be hung from the extremity of it, it will lengthen
to _l_. The less it lengthens under other equal conditions, the
more elastic the material, if it resumes its original length when
the weight is removed. It has been shown by experiment that the
increase in length _l_, due to elasticity, is directly
proportional to the length L and the weight P, and inversely
proportional to the section, but changes with the material. The
coefficient of elasticity expresses that weight (in kilos per sq.
mm.) under which a rod having a square section taken as 1 (we take
1 sq. mm.) acquires double the length by tension. Naturally in
practice materials do not withstand such a lengthening, under a
certain weight they attain a limit of elasticity, _i.e._ they
stretch permanently (undergo deformation). Neglecting fractions
(as the elasticity of metals varies not only with the temperature,
but also with forging, purity, &c.), the coefficient of elasticity
of steel and iron is 20,000, copper and brass 10,000, silver
7,000, glass 6,000, lead 2,000, and wood 1,200.
_The chemical properties of iron_ have been already repeatedly
mentioned in preceding chapters. Iron _rusts_ in air at the ordinary
temperature--that is to say, it becomes covered with a layer of iron
oxides. Here, without doubt, the moisture of the air plays a part,
because in dry air iron does not oxidise at all, and also because, more
particularly, ammonia is always found in iron rust; the ammonia must
arise from the action of the hydrogen of the water, at the moment of its
separation, on the nitrogen of the air. Highly-polished steel does not
rust nearly so readily, but if moistened with water, it easily becomes
coated with rust. As rust depends on the access of moisture, iron may be
preserved from rust by coating it with substances which prevent the
moisture having access to it. Thus arises the practice of covering iron
objects with paraffin,[13] varnish, oil, paints, or enamelling it with a
glassy-looking flux possessing the same coefficient of expansion as iron,
or with a dense scoria (formed by the heat of superheated steam), or with
a compact coating of various metals. Wrought iron (both as sheet iron and
in other forms), cast iron, and steel are often coated with tin, copper,
lead, nickel, and similar metals, which prevent contact with the air.
These metals preserve iron very effectually from rust if they form a
completely compact surface, but in those places where the iron becomes
exposed, either accidentally or from wear, rust appears much more quickly
than on a uniform iron surface, because, towards these metals (and also
towards the rust), the iron will then behave as an electro-positive pole
in a galvanic couple, and hence will attract oxygen. A coating of zinc
does not produce this inconvenience, because iron is electro-negative
with reference to zinc, in consequence of which galvanised iron does not
easily rust, and even an iron boiler containing some lumps of zinc rusts
less than one without zinc.[14] Iron oxidises at a high temperature,
forming _iron scale_, Fe_{3}O_{4}, composed of ferrous and ferric oxides,
and, as has been seen, decomposes water and acids with the evolution of
hydrogen. It is also capable of decomposing salts and oxides of other
metals, which property is applied in the arts for the extraction of
copper, silver, lead, tin, &c. For this reason iron is soluble in the
solutions of many salts--for instance, in cupric sulphate, with
precipitation of copper and formation of ferrous sulphate.[15] When iron
_acts on acids_ it always _forms compounds_ FeX_{2}--that is,
corresponding to the suboxide FeO--and answering to magnesium
compounds--and hence two atoms of hydrogen are replaced by one atom of
iron. Strongly oxidising acids like nitric acid may transform the ferrous
salt which is forming into the higher degree of oxidation or ferric salt
(corresponding with the sesquioxide, Fe_{2}O_{3}), but this is a
secondary reaction. Iron, although easily soluble in dilute nitric acid,
loses this property when plunged into strong fuming nitric acid; after
this operation it even loses the property of solubility in other acids
until the external coating formed by the action of the strong nitric acid
is mechanically removed. This condition of iron is termed the passive
state. _The passive condition_ of iron depends on the formation, on its
surface, of a coating of oxide due to the iron being acted on by the
lower oxides of nitrogen contained in the fuming nitric acid.[16] Strong
nitric acid which does not contain these lower oxides, does not render
iron passive, but it is only necessary to add some alcohol or other
reducing agent which forms these lower oxides in the nitric acid, and the
iron will assume the passive state.
[13] Paraffin is one of the best preservatives for iron against
oxidation in the air. I found this by experiments about 1860, and
immediately published the fact. This method is now very generally
applied.
[14] See Chapter XVIII., Note 34 bis. Based on the rapid oxidation of
iron and its increase in volume in the presence of water and salts
of ammonium, a packing is used for water mains and steam pipes
which is tightly hammered into the socket joints. This packing
consists of a mixture of iron filings and a small quantity of
sal-ammoniac (and sulphur) moistened with water; after a certain
lapse of time, especially after the pipes have been used, this
mass swells to such an extent that it hermetically seals the
joints of the pipes.
[15] Here, however, a ferric salt may also be formed (when all the iron
has dissolved and the cupric salt is still in excess), because the
cupric salts are reduced by ferrous salts. Cast iron is also
dissolved.
[16] Powdery reduced iron is passive with regard to nitric acid of a
specific gravity of 1·37, but when heated the acid acts on it.
This passiveness disappears in the magnetic field. Saint-Edme
attributes the passiveness of iron (and nickel) to the formation
of nitride of iron on the surface of the metal, because he
observed that when heated in dry hydrogen ammonia is evolved by
passive iron.
Remsen observed that if a strip of iron be immersed in acid and
placed in the magnetic field, it is principally dissolved at its
middle part--that is, the acid acts more feebly at the poles.
According to Étard (1891) strong nitric acid dissolves iron in
making it passive, although the action is a very slow one.
Iron readily combines with non-metals--for instance, with chlorine,
iodine, bromine, sulphur, and even with phosphorus and carbon; but on the
other hand the property of combining with metals is but little developed
in it--that is to say, it does not easily form alloys. Mercury, which
acts on most metals, does not act directly on iron, and the _iron
amalgam_, or solution of iron in mercury, which is used for electrical
machines, is only obtained in a particular way--namely, with the
co-operation of a sodium amalgam, in which the iron dissolves and by
means of which it is reduced from solutions of its salts.
When iron acts on acids it forms ferrous salts of the type FeX_{2}, and
in the presence of air and oxidising agents they change by degrees into
ferric salts of the type FeX_{3}. This faculty of passing from the
ferrous to the ferric state is still further developed in ferrous
hydroxide. If sodium hydroxide be added to a solution of ferrous sulphate
or green vitriol, FeSO_{4},[17] a white precipitate of ferrous hydroxide,
FeH_{2}O_{2}, is obtained; but on exposure to the air, even under water,
it turns green, becomes grey, and finally turns brown, which is due to
the oxidation that it undergoes. Ferrous hydroxide is very sparingly
soluble in water; the solution has, however, a distinct alkaline
reaction, which is due to its being a fairly energetic basic oxide. In
any case, ferrous oxide is far more energetic than ferric oxide, so that
if ammonia be added to a solution containing a mixture of a ferrous and
ferric salt, at first ferric hydroxide only will be precipitated. If
barium carbonate, BaCO_{3}, be shaken up in the cold with ferrous salts,
it does not precipitate them--that is, does not change them into ferrous
carbonate; but it completely separates all the iron from the ferric salts
in the cold, according to the equation Fe_{2}Cl_{6} + 3BaCO_{3} + 3H_{2}O
= Fe_{2}O_{3},3H_{2}O + 3BaCl_{2} + 3CO_{2}. If ferrous hydroxide be
boiled with a solution of potash, the water is decomposed, hydrogen is
evolved, and the ferrous hydroxide is oxidised. The ferrous salts are in
all respects similar to the salts of magnesium and zinc; they are
isomorphous with them, but differ from them in that the ferrous hydroxide
is not soluble either in aqueous potash or ammonia. In the presence of an
excess of ammonium salts, however, a certain proportion of the iron is
not precipitated by alkalis and alkali carbonates, which fact points to
the formation of double ammonium salts.[18] The ferrous salts have a dull
_greenish_ colour, and form solutions also of a pale green colour, whilst
the ferric salts have a _brown_ or reddish-brown colour. The ferrous
salts, being capable of oxidation, form very active reducing agents--for
instance, under their action gold chloride, AuCl_{3}, deposits metallic
gold, nitric acid is transformed into lower oxides, and the highest
oxides of manganese also pass into the lower forms of oxidation. All
these reactions take place with especial ease in the presence of an
excess of acid. This depends on the fact that the ferrous oxide, FeO (or
salt), acting as a reducing agent, turns into ferric oxide, Fe_{2}O_{3}
(or salt), and in the ferric state it requires more acid for the
formation of a normal salt than in the ferrous condition. Thus in the
normal ferrous sulphate, FeSO_{4}, there is one equivalent of iron to one
equivalent of sulphur (in the sulphuric radicle), but in the neutral
ferric salt, Fe_{2}(SO_{4})_{3}, there is one equivalent of iron to one
and a half of sulphur in the form of the elements of sulphuric acid.[19]
[17] _Iron vitriol_ or _green vitriol_, sulphate of iron or ferrous
sulphate, generally crystallises from solutions, like magnesium
sulphate, with seven molecules of water, FeSO_{4},7H_{2}O. This
salt is not only formed by the action of iron on sulphuric acid,
but also by the action of moisture and air on iron pyrites,
especially when previously roasted (FeS_{2} + O_{2} = FeS +
SO_{2}), and in this condition it easily absorbs the oxygen of
damp air (FeS + O_{4} = FeSO_{4}). Green vitriol is obtained in
many processes as a by-product. Ferrous sulphate, like all the
ferrous salts, has a pale greenish colour hardly perceptible in
solution. If it be desired to preserve it without change--that is,
so as not to contain ferric compounds--it is necessary to keep it
hermetically sealed. This is best done by expelling the air by
means of sulphurous anhydride or ether; sulphurous anhydride,
SO_{2}, removes oxygen from ferric compounds, which might be
formed, and is itself changed into sulphuric acid, and hence the
oxidation of the ferrous compound does not take place in its
presence. Unless these precautions are taken, green vitriol turns
brown, partly changing into the ferric salt. When turned brown, it
is not completely soluble in water, because during its oxidation a
certain amount of free insoluble ferric oxide is formed: 6FeSO_{4}
+ O_{3} = 2Fe_{2}(SO_{4})_{3} + Fe_{2}O_{3}. In order to cleanse
such mixed green vitriol from the oxide, it is necessary to add
some sulphuric acid and iron and boil the mixture; the ferric salt
is then transformed into the ferrous state: Fe_{2}(SO_{4})_{3} +
Fe = 3FeSO_{4}.
Green vitriol is used for the manufacture of Nordhausen sulphuric
acid (Chapter XX.), for preparing ferric oxide, in many dye works
(for preparing the indigo vats and reducing blue indigo to white),
and in many other processes; it is also a very good disinfectant,
and is the cheapest salt from which other compounds of iron may be
obtained.
The other ferrous salts (excepting the yellow prussiate, which
will be mentioned later) are but little used, and it is therefore
unnecessary to dwell upon them. We will only mention _ferrous
chloride_, which, in the crystalline state, has the composition
FeCl_{2},4H_{2}O. It is easily prepared; for instance, by the
action of hydrochloric acid on iron, and in the anhydrous state by
the action of hydrochloric acid gas on metallic iron at a red
heat. The anhydrous ferrous chloride then volatilises in the form
of colourless cubic crystals. Ferrous oxalate (or the double
potassium salt) acts as a powerful reducing agent, and is
frequently employed in photography (as a developer).
[18] Ferrous sulphate, like magnesium sulphate, easily forms double
salts--for instance, (NH_{4})_{2}SO_{4},FeSO_{4},6H_{2}O. This
salt does not oxidise in air so readily as green vitriol, and is
therefore used for standardising KMnO_{4}.
[19] The transformation of ferrous oxide into ferric oxide is not
completely effected in air, as then only a part of the suboxide is
converted into ferric oxide. Under these circumstances the
so-called magnetic oxide of iron is generally produced, which
contains atomic quantities of the suboxide and oxide--namely,
FeO,Fe_{2}O_{3} = Fe_{3}O_{4}. This substance, as already
mentioned, is found in nature and in iron scale. It is also formed
when most ferrous and ferric salts are heated in air; thus, for
instance, when ferrous carbonate, FeCO_{3} (native or the
precipitate given by soda in a solution of FeX_{2}), is heated it
loses the elements of carbonic anhydride, and magnetic oxide
remains. This oxide of iron is attracted by the magnet, and is on
this account called magnetic oxide, although it does not always
show magnetic properties. If magnetic oxide be dissolved in any
acid--for instance, hydrochloric--which does not act as an
oxidising agent, a ferrous salt is first formed and ferric oxide
remains, which is also capable of passing into solution. The best
way of preparing the hydrate of the magnetic oxide is by
decomposing a mixture of ferrous and ferric salts with ammonia; it
is, however, indispensable to pour this mixture into the ammonia,
and not _vice versâ_, as in that case the ferrous oxide would at
first be precipitated alone, and then the ferric oxide. The
compound thus formed has a bright green colour, and when dried
forms a black powder. Other combinations of ferrous with ferric
oxide are known, as are also compounds of ferric oxide with other
bases. Thus, for instance, compounds are known containing 4
molecules of ferrous oxide to 1 of ferric oxide, and also 6 of
ferrous to 1 of ferric oxide. These are also magnetic, and are
formed by heating iron in air. The magnesium compound
MgO,Fe_{2}O_{3} is prepared by passing gaseous hydrochloric acid
over a heated mixture of magnesia and ferric oxide. Crystalline
magnesium oxide is then formed, and black, shiny, octahedral
crystals of the above-mentioned composition. This compound is
analogous to the aluminates--for instance, to spinel. Bernheim
(1888) and Rousseau (1891) obtained many similar compounds of
ferric oxide, and their composition apparently corresponds to the
hydrates (Note 22) known for the oxide.
The most simple oxidising agent for transforming ferrous into ferric
salts is chlorine in the presence of water--for instance, 2FeCl_{2} +
Cl_{2} = Fe_{2}Cl_{6}, or, generally speaking, 2FeO + Cl_{2} + H_{2}O =
Fe_{2}O_{3} + 2HCl. When such a transformation is required it is best to
add potassium chlorate and hydrochloric acid to the ferrous solution;
chlorine is formed by their mutual reaction and acts as an oxidising
agent. Nitric acid produces a similar effect, although more slowly.
Ferrous salts may be completely and rapidly oxidised into ferric salts by
means of chromic acid or permanganic acid, HMnO_{4}, in the presence of
acids--for example, 10FeSO_{4} + 2KMnO_{4} + 8H_{2}SO_{4} =
5Fe_{2}(SO_{4})_{3} + 2MnSO_{4} + K_{2}SO_{4} + 8H_{2}O. This reaction is
easily observed by the change of colour, and its termination is easily
seen, because potassium permanganate forms solutions of a bright red
colour, and when added to a solution of a ferrous salt the above reaction
immediately takes place _in the presence of acid_, and the solution then
becomes colourless, because all the substances formed are only faintly
coloured in solution. Directly all the ferrous compound has passed into
the ferric state, any excess of permanganate which is added communicates
a red colour to the liquid (see Chapter XXI.)
Thus when ferrous salts are acted on by oxidising agents, they pass into
the ferric form, and under the action of reducing agents the reverse
reaction occurs. Sulphuretted hydrogen may, for instance, be used for
this complete transformation, for under its influence ferric salts are
reduced with separation of sulphur--for example, Fe_{2}Cl_{6} + H_{2}S =
2FeCl_{2} + 2HCl + S. Sodium thiosulphate acts in a similar way:
Fe_{2}Cl_{6} + Na_{2}S_{2}O_{3} + H_{2}O = 2FeCl_{2} + Na_{2}SO_{4} +
2HCl + S. Metallic iron or zinc,[20] in the presence, of acids, or sodium
amalgam, &c., acts like hydrogen, and has also a similar reducing action,
and this furnishes the best method for reducing ferric salts to ferrous
salts--for instance, Fe_{2}Cl_{6} + Zn = 2FeCl_{2} + ZnCl_{2}. Thus _the
transition from ferrous salts to ferric salts and vice versâ is always
possible_.[21]
[20] Copper and cuprous salts also reduce ferric oxide to ferrous
oxide, and are themselves turned into cupric salts. The essence of
the reactions is expressed by the following equations: Fe_{2}O_{3}
+ Cu_{2}O = 2FeO + 2CuO; Fe_{2}O_{3} + Cu = 2FeO + CuO. This fact
is made use of in analysing copper compounds, the quantity of
copper being ascertained by the amount of ferrous salt obtained.
An excess of ferric salt is required to complete the reaction.
Here we have an example of reverse reaction; the ferrous oxide or
its salt in the presence of alkali transforms the cupric oxide
into cuprous oxide and metallic copper, as observed by Lovel,
Knopp, and others.
[21] We will here mention the reactions by means of which it may be
ascertained whether the ferrous compound has been entirely
converted into a ferric compound or _vice versâ_. There are two
substances which are best employed for this purpose: potassium
ferricyanide, FeK_{3}C_{6}N_{6}, and potassium thiocyanate, KCNS.
The first salt gives with ferrous salts a blue precipitate of an
insoluble salt, having a composition Fe_{5}C_{12}N_{12}; but with
ferric salts it does not form any precipitate, and only gives a
brown colour, and therefore when transforming a ferrous salt into
a ferric salt, the completion of the transformation may be
detected by taking a drop of the liquid on paper or on a porcelain
plate and adding a drop of the ferricyanide solution. If a blue
precipitate be formed, then part of the ferrous salt still
remains; if there is none, the transformation is complete. The
thiocyanate does not give any marked coloration with ferrous
salts; but with ferric salts in the most diluted state it forms a
bright red soluble compound, and therefore when transforming a
ferric salt into a ferrous salt we must proceed as before, testing
a drop of the solution with thiocyanate, when the absence of a red
colour will prove the total transformation of the ferric salt into
the ferrous state, and if a red colour is apparent it shows that
the transformation is not yet complete.
_Ferric oxide_, or _sesquioxide of iron_, Fe_{2}O_{3}, is found in
nature, and is artificially prepared in the form of a red powder by many
methods. Thus after heating green vitriol a red oxide of iron remains,
called colcothar, which is used as an oil paint, principally for painting
wood. The same substance in the form of a very fine powder (rouge) is
used for polishing glass, steel, and other objects. If a mixture of
ferrous sulphate with an excess of common salt be strongly heated,
crystalline ferric oxide will be formed, having a dark violet colour, and
resembling some natural varieties of this substance. When iron pyrites is
heated for preparing sulphurous anhydride, ferric oxide also remains
behind; it is used as a pigment. On the addition of alkalis to a solution
of ferric salts, a brown precipitate of ferric hydroxide is formed, which
when heated (even when boiled in water, that is, at about 100°, according
to Tomassi) easily parts with the water, and leaves red anhydrous ferric
oxide. Pure ferric oxide does not show any magnetic properties, but when
heated to a white heat it loses oxygen and is converted into the magnetic
oxide. Anhydrous ferric oxide which has been heated to a high temperature
is with difficulty soluble in acids (but it is soluble when heated in
strong acids, and also when fused with potassium hydrogen sulphate),
whilst ferric hydroxide, at all events that which is precipitated from
salts by means of alkalis, is very readily soluble in acids. The
precipitated _ferric hydroxide_ has the composition 2Fe_{2}O_{3}3H_{2}O,
or Fe_{4}H_{6}O_{9}. If this ordinary hydroxide be rendered anhydrous (at
100°), at a certain moment it becomes incandescent--that is, loses a
certain quantity of heat. This self-incandescence depends on internal
displacement produced by the transition of the easily-soluble (in acids)
variety into the difficultly-soluble variety, and does not depend on the
loss of water, since the anhydrous oxide undergoes the same change. In
addition to this there exists a ferric hydroxide, or hydrated oxide of
iron, which, like the strongly-heated anhydrous iron oxide, is
difficultly soluble in acids. This hydroxide on losing water, or after
the loss of water, does not undergo such self-incandescence, because no
such state of internal displacement occurs (loss of energy or heat) with
it as that which is peculiar to the ordinary oxide of iron. The ferric
hydroxide which is difficultly soluble in acids has the composition
Fe_{2}O_{3},H_{2}O. This hydroxide is obtained by a prolonged ebullition
of water in which ferric hydroxide prepared by the oxidation of ferrous
oxide is suspended, and also sometimes by similar treatment of the
ordinary hydroxide after it has been for a long time in contact with
water. The transition of one hydroxide to another is apparent by a change
of colour; the easily-soluble hydroxide is redder, and the
sparingly-soluble hydroxide more yellow in colour.[22]
[22] The two ferric hydroxides are not only characterised by the
above-mentioned properties, but also by the fact that the first
hydroxide forms immediately with potassium ferrocyanide,
K_{4}FeC_{6}N_{6}, a blue colour depending on the formation of
Prussian blue, whilst the second hydroxide does not give any
reaction whatever with this salt. The first hydroxide is entirely
soluble in nitric, hydrochloric, and all other acids; whilst the
second sometimes (not always) forms a brick-coloured liquid, which
appears turbid and does not give the reactions peculiar to the
ferric salts (Péan de Saint-Gilles, Scheurer-Kestner). In addition
to this, when the smallest quantity of an alkaline salt is added
to this liquid, ferric oxide is precipitated. Thus a colloidal
solution is formed (hydrosol), which is exactly similar to silica
hydrosol (Chapter XVII.), according to which example the hydrosol
of ferric oxide may be obtained.
If ordinary ferric hydroxide be dissolved in acetic acid, a
solution of the colour of red wine is obtained, which has all the
reactions characteristic of ferric salts. But if this solution
(formed in the cold) be heated to the boiling-point, its colour is
very rapidly intensified, a smell of acetic acid becomes apparent,
and the solution then contains a new variety of ferric oxide. If
the boiling of the solution be continued, acetic acid is evolved,
and the modified ferric oxide is precipitated. If the evaporation
of the acetic acid be prevented (in a closed or sealed vessel),
and the liquid be heated for some time, the whole of the ferric
hydroxide then passes into the insoluble form, and if some
alkaline salt be added (to the hydrosol formed), the whole of the
ferric oxide is then precipitated in its insoluble form. This
method may be applied for separating ferric oxide from solutions
of its salts.
All phenomena observed respecting ferric oxide (colloidal
properties, various forms, formation of double basic salts)
demonstrate that this substance, like silica, alumina, lead
hydroxide, &c., is polymerised, that the composition is
represented by (Fe_{2}O_{3})_{_n_}.
The normal salts of the composition Fe_{2}X_{6} or FeX_{3} correspond
with ferric oxide--for example, the exceedingly volatile _ferric
chloride_, Fe_{2}Cl_{6}, which is easily prepared in the anhydrous state
by the action of chlorine on heated iron.[23] Such also is the _normal
ferric nitrate_, Fe_{2}(NO_{3})_{6}; it is obtained by dissolving iron in
an excess of nitric acid, taking care as far as possible to prevent any
rise of temperature.[24] The normal salt separates from the brown
solution when it is concentrated under a bell jar over sulphuric acid.
This salt, Fe_{2}(NO_{3})_{6},9H_{2}O, then crystallises in well-formed
and perfectly colourless crystals,[25] which deliquesce in air, melt at
35°, and are soluble in and decomposed by water. The decomposition may be
seen from the fact that the solution is brown and does not yield the
whole of the salt again, but gives partly basic salt. The normal salt
(only stable in the presence of an excess of HNO_{3}) is completely
decomposed with great facility by heating with water, even at 130°, and
this is made use of for removing iron (and also certain other oxides of
the form R_{2}O_{3}) from many other bases (of the form RO) whose
nitrates are far more stable. The ferric salts, FeX_{3}, in passing into
ferrous salts, act as oxidising agents, as is seen from the fact that
they not only liberate S from SH_{2}, but also iodine from KI like many
oxidising agents.[25 bis]
[23] The ferric compound which is most used in practice (for instance,
in medicine, for cauterising, stopping bleeding, &c.--Oleum
Martis) is _ferric chloride_, Fe_{2}Cl_{6}, easily obtained by
dissolving the ordinary hydrated oxide of iron in hydrochloric
acid. It is obtained in the anhydrous state by the action of
chlorine on heated iron. The experiment is carried on in a
porcelain tube, and a solid _volatile substance_ is then formed in
the shape of brilliant violet scales which very readily absorb
moisture from the air, and when heated with water decompose into
crystalline ferric oxide and hydrochloric acid: Fe_{2}Cl_{6} +
3H_{2}O = 6HCl + Fe_{2}O_{3}. Ferric chloride is so volatile that
the density of its vapour may be determined. At 440° it is equal
to 164·0 referred to hydrogen; the formula Fe_{2}Cl_{6}
corresponds with a density of 162·5. An aqueous solution of this
salt has a brown colour. On evaporating and cooling this solution,
crystals separate containing 6 or 12 molecules of H_{2}O. Ferric
chloride is not only soluble in water, but also in alcohol
(similarly to magnesium chloride, &c.) and in ether. If the latter
solutions are exposed to the rays of the sun they become
colourless, and deposit ferrous chloride, FeCl_{2}, chlorine being
disengaged. After a certain lapse of time, the aqueous solutions
of ferric chloride decompose with precipitation of a basic salt,
thus demonstrating the instability of ferric chloride, like the
other salts of ferric oxide (Note 22). This salt is much more
stable in the form of double salts, like all the ferric salts and
also the salts of many other feeble bases. Potassium or ammonium
chloride forms with it very beautiful red crystals of a double
salt, having the composition Fe_{2}Cl_{6},4KCl,2H_{2}O. When a
solution of this salt is evaporated it decomposes, with separation
of potassium chloride.
B. Roozeboom (1892) studied in detail (as for CaCl_{2}, Chapter
XIV., Note 50) the separation of different hydrates from saturated
solutions of Fe_{2}Cl_{6} at various concentrations and
temperatures; he found that there are 4 crystallohydrates with 12,
7, 5, and 4 molecules of water. An orange yellow only slightly
hygroscopic hydrate, Fe_{2}Cl_{6},12H_{2}O, is most easily and
usually obtained, which melts at 37°; its solubility at different
temperatures is represented by the curve BCD in the accompanying
figure, where the point B corresponds to the formation, at -55°,
of a cryohydrate containing about Fe_{2}Cl_{6} + 36H_{2}O, the
point C corresponds to the melting-point (+37°) of the hydrate
Fe_{2}Cl_{6},12H_{2}O, and the curve CD to the fall in the
temperature of crystallisation with an increase in the amount of
salt, or decrease in the amount of water (in the figure the
temperatures are taken along the axis of abscissæ, and the amount
of _n_ in the formula _n_Fe_{2}Cl_{6} + 100H_{2}O along the axis
of ordinates). When anhydrous Fe_{2}Cl_{6} is added to the above
hydrate (12H_{2}O), or some of the water is evaporated from the
latter, very hygroscopic crystals of Fe_{2}Cl_{6},5H_{2}O
(Fritsche) are formed; they melt at 56°, their solubility is
expressed by the curve HJ, which also presents a small branch at
the end J. This again gives the fall in the temperature of
crystallisation with an increase in the amount of Fe_{2}Cl_{6}.
Besides these curves and the solubility of the anhydrous salt
expressed by the line KL (up to 100°, beyond which chlorine is
liberated), Roozeboom also gives the two curves, EFG and JK,
corresponding to the crystallohydrates, Fe_{2}Cl_{6},7H_{2}O
(melts at +32°·5, that is lower than any of the others) and
Fe_{2}Cl_{6},4H_{2}O (melts at 73°·5), which he discovered by a
systematic research on the solutions of ferric chloride. The curve
AB represents the separation of ice from dilute solutions of the
salt.
[Illustration: FIG. 95.--Diagram of the solubility of
Fe_{2}Cl_{6}.]
The researches of the same Dutch chemist upon the conditions of
the formation of crystals from the double salt
(NH_{4}Cl)_{4}Fe_{2}Cl_{6},2H_{2}O are even more perfect. This
salt was obtained in 1839 by Fritsche, and is easily formed from a
strong solution of Fe_{2}Cl_{6} by adding sal-ammoniac, when it
separates in crimson rhombic crystals, which, after dissolving in
water, only deposit again on evaporation, together with the
sal-ammoniac.
Roozeboom (1892) found that when the solution contains _b_
molecules of Fe_{2}Cl_{6}, and _a_ molecules of NH_{4}Cl, per 100
molecules H_{2}O, then at 15° one of the following separations
takes place: (1) crystals, Fe_{2}Cl_{6},12H_{2}O, when _a_ varies
between 0 and 11, and _b_ between 4·65 and 4·8, or (2) a mixture
of these crystals and the double salt, when _a_ = 1·36, and _b_ =
4·47, or (3) the double salt, Fe_{2}Cl_{6},4NH_{4}Cl,2H_{2}O, when
_a_ varies between 2 and 11·8, and _b_ between 3·1 and 4·56, or
(4) a mixture of sal-ammoniac with the iron salt (it crystallises
in separate cubes, Retgers, Lehmann), when _a_ varies between 7·7
and 10·9, and _b_ is less than 3·38, or (5) sal-ammoniac, when _a_
= 11·88. And as in the double salt, _a_ : _b_ :: 4 : 1 it is
evident that the double salt only separates out when the ratio _a_
: _b_ is less than 4 : 1 (_i.e._ when Fe_{2}Cl_{6} predominates).
The above is seen more clearly in the accompanying figure, where
_a_, or the number of molecules of NH_{4}Cl per 100H_{2}O, is
taken along the axis of abscissæ, and _b_, or the number of
molecules of Fe_{2}Cl_{6}, along the ordinates. The curves ABCD
correspond to saturation and present an iso-therm of 15°. The
portion AB corresponds to the separation of chloride of iron (the
ascending nature of this curve shows that the solubility of
Fe_{2}Cl_{6} is increased by the presence of NH_{4}Cl, while that
of NH_{4}Cl decreases in the presence of Fe_{2}Cl_{6}), the
portion BC to the double salt, and the portion CD to a mixture of
sal-ammoniac and ferric chloride, while the straight line OF
corresponds to the ratio Fe_{2}Cl_{6},4NH_{4}Cl, or _a_ : _b_ :: 4
: 1. The portion CE shows that more double salt may be introduced
into the solution without decomposition, but then the solution
deposits a mixture of sal-ammoniac and ferric chloride (_see_
Chapter XXIV. Note 9 ^{bis}). If there were more such
well-investigated cases of solutions, our knowledge of double
salts, solutions, the influence of water, equilibria, isomorphous
mixtures, and such-like provinces of chemical relations might be
considerably advanced.
[Illustration: FIG. 96.--Diagram of the formation, at 15°, of the
double salt Fe_{2}Cl_{6}4NH_{4}Cl_{2}H_{2}O or
Fe(NH_{4})_{2}Cl_{5}H_{2}O. (After Roozeboom.)]
[24] The normal ferric salts are decomposed by heat and even by water,
forming basic salts, which may be prepared in various ways.
Generally ferric hydroxide is dissolved in solutions of ferric
nitrate; if it contains a double quantity of iron the basic salt
is formed which contains Fe_{2}O_{3} (in the form of hydroxide) +
2Fe_{2}(NO_{3})_{6} = 3Fe_{2}O(NO_{3})_{4}, a salt of the type
Fe_{2}OX_{4}. Probably water enters into its composition. With
considerable quantities of ferric oxide, insoluble basic salts are
obtained containing various amounts of ferric hydroxide. Thus when
a solution of the above-mentioned basic acid is boiled, a
precipitate is formed containing
4(Fe_{2}O_{3})_{8},2(N_{2}O_{5}),3H_{2}O, which probably contains
2Fe_{2}O_{2}(NO_{3})_{2} + 2Fe_{2}O_{3},3H_{2}O. If a solution of
basic nitrate be sealed in a tube and then immersed in boiling
water, the colour of the solution changes just in the same way as
if a solution of ferric acetate had been employed (Note 22). The
solution obtained smells strongly of nitric acid, and on adding a
drop of sulphuric or hydrochloric acid the insoluble variety of
hydrated ferric oxide is precipitated.
Normal ferric _orthophosphate_ is soluble in sulphuric,
hydrochloric, and nitric acids, but insoluble in others, such as,
for instance, acetic acid. The composition of this salt in the
anhydrous state is FePO_{4}, because in orthophosphoric acid there
are three atoms of hydrogen, and iron, in the ferric state,
replaces the three atoms of hydrogen. This salt is obtained from
ferric acetate, which, with disodium phosphate, forms a _white
precipitate_ of FePO_{4}, containing water. If a solution of
ferric chloride (yellowish-red colour) be mixed with a solution of
sodium acetate in excess, the liquid assumes an intense brown
colour which demonstrates the formation of a certain quantity of
ferric acetate; then the disodium phosphate directly forms a white
gelatinous precipitate of ferric phosphate. By this means the
whole of the iron may be precipitated, and the liquid which was
brown then becomes colourless. If this normal salt be dissolved in
orthophosphoric acid, the crystalline acid salt
FeH_{3}(PO_{4})_{2} is formed. If there be an excess of ferric
oxide in the solution, the precipitate will consist of the basic
salt. If ferric phosphate be dissolved in hydrochloric acid, and
ammonia be added, a salt is precipitated on heating which, after
continued washing in water and heating (to remove the water), has
the composition Fe_{4}P_{2}O_{11}--that is,
2Fe_{2}O_{3},P_{2}O_{5}. In an aqueous condition this salt may be
considered as ferric hydroxide, Fe_{2}(OH)_{6}, in which (OH)_{3}
is replaced by the equivalent group PO_{4}. Whenever ammonia is
added to a solution containing an excess of ferric salt and a
certain amount of phosphoric acid, a precipitate is formed
containing the whole of the phosphoric acid in the mass of the
ferric oxide.
Ferric oxide is characterised as a feeble base, and also by the
fact of its forming double salts--for instance, _potassium iron
alum_, which has a composition
Fe_{2}(SO_{4})_{3},K_{2}SO_{4},24H_{2}O or
FeK(SO_{4})_{2},12H_{2}O. It is obtained in the form of almost
colourless or light rose-coloured large octahedra of the regular
system by simply mixing solutions of potassium sulphate and the
ferric sulphate obtained by dissolving ferric oxide in sulphuric
acid.
[25] It would seem that all normal ferric salts are colourless, and
that the brown colour which is peculiar to the solutions is really
due to basic ferric salts. A remarkable example of the apparent
change of colour of salts is represented by the ferrous and ferric
oxalates. The former in a dry state has a yellow colour, although
as a rule the ferrous salts are green, and the latter is
colourless or pale green. When the normal ferric salt is dissolved
in water it is, like many salts, probably decomposed by the water
into acid and basic salts, and the latter communicates a brown
colour to the solution. Iron alum is almost colourless, is easily
decomposed by water, and is the best proof of our assertion. The
study of the phenomena peculiar to ferric nitrate might, in my
opinion, give a very useful addition to our knowledge of the
aqueous solutions of salts in general.
[25 bis] The reaction FeX_{3} + KI = FeX_{2} + KX + I proceeds
comparatively slowly in solutions, is not complete (depends upon
the mass), and is reversible. In this connection we may cite the
following data from Seubert and Rohrer's (1894) comprehensive
researches. The investigations were conducted with solutions
containing 1/10 gram--equivalent weights of Fe_{2}(SO_{4})_{3}
(_i.e._ containing 20 grams of salt per litre), and a
corresponding solution of KI; the amount of iodine liberated being
determined (after the addition of starch) by a solution (also 1/10
normal) of Na_{2}S_{2}O_{3} (_see_ Chapter XX., Note 42). The
progress of the reaction was expressed by the amount of liberated
iodine in percentages of the theoretical amount. For instance, the
following amount of iodide of potassium was decomposed when
Fe_{2}(SO_{4})_{3} + 2_n_KI was taken:
_n_ = 1 2 3 6 10 20
After 15´ 11·4 26·3 40·6 73·5 91·6 96·0
" 30´ 14·0 35·8 47·8 78·5 94·3 97·4
" 1 hour 19·0 42·7 56·0 84·0 95·7 97·6
" 10 " 32·6 56·0 75·7 93·2 96·5 97·6
" 48 " 39·4 67·7 82·6 93·4 96·6 97·6
Similar results were obtained for FeCl_{3}, but then the amount of
iodine liberated was somewhat greater. Similar results were also
obtained by increasing the mass of FeX_{3} per KI, and by
replacing it by HI (_see_ Chapter XXI., Note 26).
Iron forms one other oxide besides the ferric and ferrous oxides; this
contains twice as much oxygen as the former, but is so very unstable that
it can neither be obtained in the free state nor as a hydrate. Whenever
such conditions of double decomposition occur as should allow of its
separation in the free state, it decomposes into oxygen and ferric oxide.
It is known in the state of salts, and is only stable in the presence of
alkalis, and forms salts with them which have a decidedly alkaline
reaction; it is therefore a feebly acid oxide. Thus when small pieces of
iron are heated with nitre or potassium chlorate a potassium salt of the
composition K_{2}FeO_{4} is formed, and therefore the hydrate
corresponding with this salt should have the composition H_{2}FeO_{4}. It
is called _ferric acid_. Its anhydride ought to contain FeO_{3} or
Fe_{2}O_{6}--twice as much oxygen as ferric oxide. If a solution of
potassium ferrate be mixed with acid, the free hydrate ought to be
formed, but it immediately decomposes (2K_{2}FeO_{4} + 5H_{2}SO_{4} =
2K_{2}SO_{4} + Fe_{2}(SO_{4})_{3} + 5H_{2}O + O_{3}), oxygen being
evolved. If a small quantity of acid be taken, or if a solution of
potassium ferrate be heated with solutions of other metallic salts,
ferric oxide is separated--for instance:
2CuSO_{4} + 2K_{2}FeO_{4} = 2K_{2}SO_{4} + O_{3} + Fe_{2}O_{3} + 2CuO.
Both these oxides are of course deposited in the form of hydrates. This
shows that not only the hydrate H_{2}FeO_{4}, but also the salts of the
heavy metals corresponding with this higher oxide of iron, are not formed
by reactions of double decomposition. The solution of potassium ferrate
naturally acts as a powerful oxidising agent; for instance, it transforms
manganous oxide into the dioxide, sulphurous into sulphuric acid, oxalic
acid into carbonic anhydride and water, &c.[26]
[26] If chlorine be passed through a strong solution of potassium
hydroxide in which hydrated ferric oxide is suspended, the turbid
liquid acquires a dark pomegranate-red colour and contains
potassium ferrate: 10KHO + Fe_{2}O_{3} + 3Cl_{2} = 2K_{2}FeO_{4} +
6KCl + 5H_{2}O. The chlorine must not be in excess, otherwise the
salt is again decomposed, although the mode of decomposition is
unknown; however, ferric chloride and potassium chlorate are
probably formed. Another way in which the above-described salt is
formed is also remarkable; a galvanic current (from 6 Grove
elements) is passed through cast-iron and platinum electrodes into
a strong solution of potassium hydroxide. The cast-iron electrode
is connected with the positive pole, and the platinum electrode is
surrounded by a porous earthenware cylinder. Oxygen would be
evolved at the cast-iron electrode, but it is used up in
oxidation, and a dark solution of potassium ferrate is therefore
formed about it. It is remarkable that the cast iron cannot be
replaced by wrought iron.
Iron thus combines with oxygen in three proportions: RO, R_{2}O_{3},
and RO_{3}. It might have been expected that there would be intermediate
stages RO_{2} (corresponding to pyrites FeS_{2}) and R_{2}O_{5}, but for
iron these are unknown.[26 bis] The lower oxide has a distinctly basic
character, the higher is feebly acid. The only one which is stable in the
free state is ferric oxide, Fe_{2}O_{3}; the suboxide, FeO, absorbs
oxygen, and ferric anhydride, FeO_{3}, evolves it. It is also the same
for other elements; the character of each is determined by the relative
degree of stability of the known oxides. The salts FeX_{2} correspond
with the suboxide, the salts FeX_{3} or Fe_{2}X_{6} with the sesquioxide,
and FeX_{6} represents those of ferric acid, as its potassium salt is
FeO_{2}(OK)_{2}, corresponding with K_{2}SO_{4}, K_{2}MnO_{4},
K_{2}CrO_{4}, &c. Iron therefore forms compounds of the types FeX_{2},
FeX_{3}, and FeX_{6}, but this latter, like the type NX_{5}, does not
appear separately, but only when X represents heterogeneous elements or
groups; for instance, for nitrogen in the form of NO_{2}(OH), NH_{4}Cl,
&c., for iron in the form of FeO_{2}(OK)_{2}. But still the type FeX_{6}
exists, and therefore FeX_{2} and FeX_{3} are compounds which, like
ammonia, NH_{3}, are capable of further combinations up to FeX_{6}; this
is also seen in the property of ferrous and ferric salts of forming
compounds with water of crystallisation, besides double and basic salts,
whose stability is determined by the quality of the elements included in
the types FeX_{2} and FeX_{3}.[26 tri] It is therefore to be expected
that there should be complex compounds derived from ferrous and ferric
oxides. Amongst these the series of cyanogen compounds is particularly
interesting; their formation and character is not only determined by the
property which iron possesses of forming complex types, but also by the
similar faculty of the cyanogen compounds, which, like nitriles (Chapter
IX.), have clearly developed properties of polymerisation and in general
of forming complex compounds.[27]
[26 bis] When Mond and his assistants obtained the remarkable volatile
compound Ni(CO)_{4} (described later, Chapter XXII.), it was shown
subsequently by Mond and Quincke (1891), and also by Berthelot,
that iron, under certain conditions, in a stream of carbonic
oxide, also volatilises and forms a compound like that given by
nickel. Roscoe and Scudder then showed that when water gas is
passed through and kept under pressure (8 atmospheres) in iron
vessels a portion of the iron volatilises from the sides of the
vessel, and that when the gas is burnt it deposits a certain
amount of oxides of iron (the same result is obtained with
ordinary coal gas which contains a small amount of CO). To obtain
the _volatile compound of iron with carbonic oxide_, Mond prepared
a finely divided iron by heating the oxalate in a stream of
hydrogen, and after cooling it to 80°-45° he passed CO over the
powder. The iron then formed (although very slowly) a volatile
compound containing Fe(CO)_{5} (as though it answered to a very
high type, FeX_{10}), which when cooled condenses into a liquid
(slightly coloured, probably owing to incipient decomposition),
sp. gr. 1·47, which solidifies at -21°, boils at about 103°, and
has a vapour density (about 6·5 with respect to air) corresponding
to the above formula; it decomposes at 180°. Water and dilute
acids do not act upon it, but it decomposes under the action of
light and forms a hard, non-volatile crystalline yellow compound
Fe_{2}(CO)_{7} which decomposes at 80° and again forms Fe(CO)_{5}.
[26 tri] When the molecular Fe_{2}Cl_{6} is produced instead of
FeCl_{3} this complication of the type also occurs.
[27] Some light may be thrown upon the faculty of Fe of forming various
compounds with CN, by the fact that Fe not only combines with
carbon but also with nitrogen. _Nitride of iron_ Fe_{2}N was
obtained by Fowler by heating finely powdered iron in a stream of
NH_{3} at the temperature of melting lead.
_In the cyanogen compounds of iron_, two degrees might be expected:
Fe(CN)_{2}, corresponding with ferrous oxide, and Fe(CN)_{3},
corresponding with ferric oxide. There are actually, however, many other
known compounds, intermediate and far more complex. They correspond with
the double salts so easily formed by metallic cyanides. The two following
double salts are particularly well known, very stable, often used, and
easily prepared. _Potassium ferrocyanide_ or _yellow prussiate of
potash_, a double salt of cyanide of potassium and ferrous cyanide, has
the composition FeC_{2}N_{2},4KCN; its crystals contain 3 mol. of water:
K_{4}FeC_{6}N_{6},3H_{2}O. The other is _potassium ferricyanide_ or _red
prussiate of potash_. It is also known as _Gmelin's salt_, and contains
cyanide of potassium with ferric cyanide; its composition is
Fe(CN)_{3},3KCN or K_{3}FeC_{6}N_{6}. Its crystals do not contain water.
It is obtained from the first by the action of chlorine, which removes
one atom of the potassium. A whole series of other ferrocyanic compounds
correspond with these ordinary salts.
Before treating of the preparation and properties of these two
remarkable and very stable salts, it must be observed that with ordinary
reagents neither of them gives the same double decompositions as the
other ferrous and ferric salts, and they both present a series of
remarkable properties. Thus these salts have a neutral reaction, are
unchanged by air, dilute acids, or water, unlike potassium cyanide and
even some of its double salts. When solutions of these salts are treated
with caustic alkalis, they do not give a precipitate of ferrous or ferric
hydroxides, neither are they precipitated by sodium carbonate. This led
the earlier investigators to recognise special independent groupings in
them. The yellow prussiate was considered to contain the complex radicle
FeC_{6}N_{6} combined with potassium, namely with K_{4}, and K_{3} was
attributed to the red prussiate. This was confirmed by the fact that
whilst in both salts any other metal, even hydrogen, might be substituted
for potassium, the iron remained unchangeable, just as nitrogen in
cyanogen, ammonium, and nitrates does not enter into double
decomposition, being in the state of the complex radicles CN, NH_{4},
NO_{2}. Such a representation is, however, completely superfluous for the
explanation of the peculiarities in the reactions of such compounds as
double salts. If a magnesium salt which can be precipitated by potassium
hydroxide does not form a precipitate in the presence of ammonium
chloride, it is very clear that it is owing to the formation of a soluble
double salt which is not decomposed by alkalis. And there is no necessity
to account for the peculiarity of reaction of a double salt by the
formation of a new complex radicle. In the same way also, in the presence
of an excess of tartaric acid, cupric salts do not form a precipitate
with potassium hydroxide, because a double salt is formed. These
peculiarities are more easily understood in the case of cyanogen
compounds than in all others, because all cyanogen compounds, as
unsaturated compounds, show a marked tendency to complexity. This
tendency is satisfied in double salts. The appearance of a peculiar
character in double cyanides is the more easily understood since in the
case of potassium cyanide itself, and also in hydrocyanic acid, a great
many peculiarities have been observed which are not encountered in those
haloid compounds, potassium chloride and hydrochloric acid, with which it
was usual to compare cyanogen compounds. These peculiarities become more
comprehensible on comparing cyanogen compounds with ammonium compounds.
Thus in the presence of ammonia the reactions of many compounds change
considerably. If in addition to this it is remembered that the presence
of many carbon (organic) compounds frequently completely disturbs the
reaction of salts, the peculiarities of certain double cyanides will
appear still less strange, because they contain carbon. The fact that the
presence of carbon or another element in the compound produces a change
in the reactions, may be compared to the action of oxygen, which, when
entering into a combination, also very materially changes the nature of
reactions. Chlorine is not detected by silver nitrate when it is in the
form of potassium chlorate, KClO_{3}, as it is detected in potassium
chloride, KCl. The iron in ferrous and ferric compounds varies in its
reactions. In addition to the above-mentioned facts, consideration ought
to be given to the circumstance that the easy mutability of nitric acid
undergoes modification in its alkali salts, and in general the properties
of a salt often differ much from those of the acid. Every double salt
ought to be regarded as a peculiar kind of saline compound: potassium
cyanide is, as it were, a basic, and ferrous cyanide an acid, element.
They may be unstable in the separate state, but form a stable double
compound when combined together; the act of combination disengages the
energy of the elements, and they, so to speak, saturate each other. Of
course, all this is not a definite explanation, but then the supposition
of a special complex radicle can even less be regarded as such.
Potassium ferrocyanide, K_{4}FeC_{6}N_{6}, is very easily formed by
mixing solutions of ferrous sulphate and potassium cyanide. First, a
white precipitate of ferrous cyanide, FeC_{2}N_{2}, is formed, which
becomes blue on exposure to air, but is soluble in an excess of potassium
cyanide, forming the ferrocyanide. The same yellow prussiate is obtained
on heating animal nitrogenous charcoal or animal matters--such as horn,
leather cuttings, &c.--with potassium carbonate in iron vessels,[27 bis]
the mass formed being afterwards boiled with water with exposure to air,
potassium cyanide first appearing, which gives yellow prussiate. The
animal charcoal may be exchanged for wood charcoal, permeated with
potassium carbonate and heated in nitrogen or ammonia; the mass thus
produced is then boiled in water with ferric oxide.[28] In this manner it
is manufactured on the large scale, and is called 'yellow prussiate'
('prussiate de potasse,' Blutlaugensalz).
[27 bis] The sulphur of the animal refuse here forms the compound
FeKS_{2}, which by the action of potassium cyanide yields
potassium sulphide, thiocyanate, and ferrocyanide.
[28] Potassium ferrocyanide may also be obtained from Prussian blue by
boiling with a solution of potassium hydroxide, and from the
ferricyanide by the action of alkalis and reducing substances
(because the red prussiate is a product of oxidation produced by
the action of chlorine: a ferric salt is reduced to a ferrous
salt), &c. In many works (especially in Germany and France) yellow
prussiate is prepared from the mass, containing oxide of iron, and
employed for purifying coal gas (Vol. I., p. 361), which generally
contains cyanogen compounds. About 2 p.c. of the nitrogen
contained in coal is converted into cyanogen, which forms Prussian
blue and thiocyanates in the mass used for purifying the gas. On
evaporation the solution yields large yellow crystals containing 3
molecules of water, which is easily expelled by heating above
100°. 100 parts of water at the ordinary temperature are capable
of dissolving 25 parts of this salt; its sp. gr. is 1·83. When
ignited it forms potassium cyanide and iron carbide, FeC_{2}
(Chapter XIII., Note 12). Oxidising substances change it into
potassium ferricyanide. With strong sulphuric acid it gives
carbonic oxide, and with dilute sulphuric acid, when heated,
prussic acid is evolved according to the equation:
2K_{4}FeC_{6}N_{6} + 3H_{2}SO_{4} = K_{2}Fe_{2}C_{6}N_{6} +
3K_{2}SO_{4} + 6HCN; hence in the yellow prussiate K_{2} replaces
Fe.
It is easy to substitute other metals for the potassium in the yellow
prussiate. The hydrogen salt or hydroferrocyanic acid, H_{4}FeC_{6}N_{6},
is obtained by mixing strong solutions of yellow prussiate and
hydrochloric acid. If ether be added and the air excluded, the acid is
obtained directly in the form of a white scarcely crystalline precipitate
which becomes blue on exposure to air (as ferrous cyanide does from the
formation of blue compounds of ferrous and ferric cyanides, and it is on
this account used in cotton printing). It is soluble in water and
alcohol, but not in ether, has marked acid properties, and decomposes
carbonates, which renders it easily possible to prepare ferrocyanides of
the metals of the alkalis and alkaline earths; these are readily soluble,
have a neutral reaction, and resemble the yellow prussiate. Solutions of
these salts form precipitates with the salts of other metals, because the
ferrocyanides of the heavy metals are insoluble. Here either the whole of
the potassium of the yellow prussiate, or only a part of it, is exchanged
for an equivalent quantity of the heavy metal. Thus, when a cupric salt
is added to a solution of yellow prussiate, a red precipitate is obtained
which still contains half the potassium of the yellow prussiate:
K_{4}FeC_{6}N_{6} + CuSO_{4} = K_{2}CuFeC_{6}N_{6} + K_{2}SO_{4}.
But if the process be reversed (the salt of copper being then in excess)
the whole of the potassium will be exchanged for copper, forming a
reddish-brown precipitate, Cu_{2}FeC_{6}N_{6},9H_{2}O. This reaction and
those similar to it are very sensitive and may be used for detecting
metals in solution, more especially as the colour of the precipitate very
often shows a marked difference when one metal is exchanged for another.
Zinc, cadmium, lead, antimony, tin, silver, cuprous and aurous salts form
_white_ precipitates; cupric, uranium, titanium and molybdenum salts
_reddish-brown_; those of nickel, cobalt, and chromium, _green_
precipitates; _with ferrous salts_, ferrocyanide forms, as has been
already mentioned, a _white_ precipitate--namely, Fe_{2}FeC_{6}N_{6}, or
FeC_{2}N_{2}--which turns blue on exposure to air, and with ferric salts
a _blue precipitate_ called _Prussian blue_. Here the potassium is
replaced by iron, the reaction being expressed thus: 2Fe_{2}Cl_{6} +
3K_{4}FeC_{6}N_{6} = 12KCl + Fe_{4}Fe_{3}C_{18}N_{18}, the latter formula
expressing the composition of Prussian blue. It is therefore the compound
4Fe(CN)_{3} + 3Fe(CN)_{2}. The yellow prussiate is prepared in chemical
works on a large scale especially for the manufacture of this blue
pigment, which is used for dyeing cloth and other fabrics and also as one
of the ordinary blue paints. It is insoluble in water, and the stuffs are
therefore dyed by first soaking them in a solution of a ferric salt and
then in a solution of yellow prussiate. If however an excess of yellow
prussiate be present complete substitution between potassium and iron
does not occur, and _soluble Prussian blue_ is formed; KFe_{2}(CN)_{6} =
KCN,Fe(CN)_{2},Fe(CN)_{3}. This blue salt is colloidal, is soluble in
pure water, but insoluble and precipitated when other salts--for
instance, potassium or sodium chloride--are present even in small
quantities, and is therefore first obtained as a precipitate.[29]
[29] Skraup obtained this salt both from potassium ferrocyanide with
ferric chloride and from ferricyanide with ferrous chloride, which
evidently shows that it contains iron in both the ferric and
ferrous states. With ferrous chloride it forms Prussian blue, and
with ferric chloride Turnbull's blue.
Prussian blue was discovered in the beginning of the last century
by a Berlin manufacturer, Diesbach. It was then prepared, as it
sometimes is also at present, directly from potassium cyanide
obtained by heating animal charcoal with potassium carbonate. The
mass thus obtained is dissolved in water, alum is added to the
solution in order to saturate the free alkali, and then a solution
of green vitriol is added which has previously been sufficiently
exposed to the air to contain both ferric and ferrous salts. If
the solution of potassium cyanide be mixed with a solution
containing both salts, Prussian blue will be formed, because it is
a compound of ferrous cyanide, FeC_{2}N_{2}, and ferric cyanide,
Fe_{2}C_{6}N_{6}. A ferric salt with potassium ferrocyanide forms
a blue colour, because ferrous cyanide is obtained from the first
salt and ferric cyanide from the second. During the preparation of
this compound alkali must be avoided, as otherwise the precipitate
would contain oxides of iron. Prussian blue has not a crystalline
structure; it forms a blue mass with a copper-red metallic lustre.
Both acids and alkalis act on it. The action is at first confined
to the ferric salt it contains. Thus alkalis form ferric oxide and
ferrocyanide in solution: 2Fe_{2}C_{6}N_{6},3FeC_{2}N_{2} + 12KHO
= 2(Fe_{2}O_{3},3H_{2}O) + 3K_{4}FeC_{6}N_{6}. Various
ferrocyanides may thus be prepared. Prussian blue is soluble in an
aqueous solution of oxalic acid, forming blue ink. In air, when
exposed to the action of light, it fades; but in the dark again
absorbs oxygen and becomes blue, which fact is also sometimes
noticed in blue cloth. An excess of potassium ferrocyanide renders
Prussian blue soluble in water, although insoluble in various
saline solutions--that is, it converts it into the soluble
variety. Strong hydrochloric acid also dissolves Prussian blue.
Potassium ferricyanide, or _red prussiate_ of potash, K_{3}FeC_{6}N_{6},
is called 'Gmelin's salt,' because this savant obtained it by the action
of chlorine on a solution of the yellow prussiate: K_{4}FeC_{6}N_{6} + Cl
= K_{3}FeC_{6}N_{6} + KCl. The reaction is due to the ferrous salt being
changed by the action of the chlorine into a ferric salt. It separates
from solutions in anhydrous, well-formed prisms of a red colour, but the
solution has an olive colour; 100 parts of water, at 10°, dissolve 37
parts of the salt, and at 100°, 78 parts.[30] The red prussiate gives a
blue precipitate with ferrous salts, called _Turnbull's blue_, very much
like Prussian blue (and the soluble variety), because it also contains
ferrous cyanide and ferric cyanide, although in another proportion, being
formed according to the equation: 3FeCl_{2} + 2K_{3}FeC_{6}N_{6} = 6KCl +
Fe_{3}Fe_{2}C_{12}N_{12}, or 3FeC_{2}N_{2},Fe_{2}C_{6}N_{6}; in Prussian
blue we have Fe_{7}Cy_{18}, and here Fe_{5}Cy_{12}. A ferric salt ought
to form ferric cyanide Fe_{2}C_{6}N_{6}, with red prussiate, but ferric
cyanide is soluble, and therefore no precipitate is obtained, and the
liquid only becomes brown.[31]
[30] An excess of chlorine must not be employed in preparing this
compound, otherwise the reaction goes further. It is easy to find
out when the action of the chlorine on potassium ferrocyanide must
cease; it is only necessary to take a sample of the liquid and add
a solution of a ferric salt to it. If a precipitate of Prussian
blue is formed, more chlorine must be added, as there is still
some undecomposed ferrocyanide, for the ferricyanide does not give
a precipitate with ferric salts. Potassium ferricyanide, like the
ferrocyanide, easily exchanges its potassium for hydrogen and
various metals by double decomposition. With the salts of tin,
silver, and mercury it forms yellow precipitates, and with those
of uranium, nickel, cobalt, copper, and bismuth brown
precipitates. The lead salt under the action of sulphuretted
hydrogen forms lead sulphide and a hydrogen salt or acid,
H_{3}FeC_{6}N_{6}, corresponding with potassium ferricyanide,
which is soluble, crystallises in red needles, and resembles
hydroferrocyanic acid, H_{4}FeC_{6}N_{6}. Under the action of
reducing agents--for instance, sulphuretted hydrogen,
copper--potassium ferricyanide is changed into ferrocyanide,
especially in the presence of alkalis, and thus forms a rather
energetic _oxidising agent_--capable, for instance, of changing
manganous oxide into dioxide, bleaching tissues, &c.
[31] It is important to mention a series of readily crystallisable
salts formed by the action of nitric acid on potassium and other
ferrocyanides and ferricyanides. These salt contain the elements
of nitric oxide, and are therefore called _nitro-(nitroso)
ferricyanides_ (_nitroprussides_). Generally a crystalline sodium
salt is obtained, Na_{2}FeC_{5}N_{6}O,2H_{2}O. In its composition
this salt differs from the red sodium salt, Na_{3}FeC_{6}N_{6}, by
the fact that in it one molecule of sodium cyanide, NaCN, is
replaced by nitric oxide, NO. In order to prepare it, potassium
ferrocyanide in powder is mixed with five-sevenths of its weight
of nitric acid diluted with an equal volume of water. The mixture
is at first left at the ordinary temperature, and then heated on a
water-bath. Here ferricyanide is first of all formed (as shown by
the liquid giving a precipitate with ferrous chloride), which then
disappears (no precipitate with ferrous chloride), and forms a
green precipitate. The liquid, when cooled, deposits crystals of
nitre. The liquid is then strained off and mixed with sodium
carbonate, boiled, filtered, and evaporated; sodium nitrate and
the salt described are deposited in crystals. It separates in
prisms of a red colour. Alkalis and salts of the alkaline earths
do not give precipitates: they are soluble, but the salts of iron,
zinc, copper, and silver form precipitates where sodium is
exchanged with these metals. It is remarkable that the sulphides
of the alkali metals give with this salt an intense bright purple
coloration. This series of compounds was discovered by Gmelin and
studied by Playfair and others (1849).
This series to a certain extent resembles the nitro-sulphide
series described by Roussin. Here the primary compound consists of
black crystals, which are obtained as follows:--Solutions of
potassium hydrosulphide and nitrate are mixed, and the mixture is
agitated whilst ferric chloride is added, then boiled and
filtered; on cooling, _black crystals_ are deposited, having the
composition Fe_{6}S_{3}(NO)_{10},H_{2}O (Rosenberg), or, according
to Demel, FeNO_{2},NH_{2}S. They have a slightly metallic lustre,
and are soluble in water, alcohol, and ether. They absorb the
latter as easily as calcium chloride absorbs water. In the
presence of alkalis these crystals remain unchanged, but with
acids they evolve nitric oxides. There are several compounds which
are capable of interchanging, and correspond with Roussin's salt.
Here we enter into the series of the nitrogen compounds which have
been as yet but little investigated, and will most probably in
time form most instructive material for studying the nature of
that element. These series of compounds are as unlike the usual
saline compounds of inorganic chemistry as are organic
hydrocarbons. There is no necessity to describe these series in
detail, because their connection with other compounds is not yet
clear, and they have not yet any application.
If chlorine and sodium are representatives of independent groups of
elements, the same may also be said of iron. Its nearest analogues show,
besides a similarity in character, a likeness as regards physical
properties and a proximity in atomic weight. Iron occupies a medium
position amongst its nearest analogues, both with respect to properties
and faculty of forming saline oxides, and also as regards atomic weight.
On the one hand, cobalt, 58, and nickel, 59, approach iron, 56; they are
metals of a more basic character, they do not form stable acids or higher
degrees of oxidation, and are a transition to copper, 63, and zinc, 65.
On the other hand, manganese, 55, and chromium, 52, are the nearest to
iron; they form both basic and acid oxides, and are a transition to the
metals possessing acid properties. In addition to having atomic weights
approximately alike, chromium, manganese, iron, cobalt, nickel, and
copper have also nearly the same specific gravity, so that the atomic
volumes and the molecules of their analogous compounds are also near to
one another (see table at the beginning of this volume). Besides this,
the likeness between the above-mentioned elements is also seen from the
following:
They form suboxides, RO, fairly energetic bases, isomorphous with
magnesia--for instance, the salt RSO_{4},7H_{2}O, akin to
MgSO_{4},7H_{2}O, and FeSO_{4},7H_{2}O, or to sulphates containing less
water; with alkali sulphates all form double salts crystallising with
6H_{2}O; all are capable of forming ammonium salts, &c. The lower oxides,
in the cases of nickel and cobalt, are tolerably stable, are not easily
oxidised (the nickel compound with more difficulty than cobalt, a
transition to copper); with manganese, and especially with chromium, they
are more easily oxidised than with iron and pass into higher oxides. They
also form oxides of the form R_{2}O_{3}, and with nickel, cobalt, and
manganese this oxide is very unstable, and is more easily reduced than
ferric oxide; but, in the case of chromium, it is very stable, and forms
the ordinary kind of salts. It is isomorphous with ferric oxide, forms
alums, is a feeble base, &c. Chromium, manganese, and iron are oxidised
by alkali and oxidising agents, forming salts like Na_{2}SO_{4}; but
cobalt and nickel are difficult to oxidise; their acids are not known
with any certainty, and are, in all probability, still less stable than
the ferrates. Cr, Mn and Fe form compounds R_{2}Cl_{6} which are like
Fe_{2}Cl_{6} in many respects; in Co this faculty is weaker and in Ni it
has almost disappeared. The cyanogen compounds, especially of manganese
and cobalt, are very near akin to the corresponding ferrocyanides. The
oxides of nickel and cobalt are more easily reduced to metal than those
of iron, but those of manganese and chromium are not reduced so easily as
iron, and the metals themselves are not easily obtained in a pure state;
they are capable of forming varieties resembling cast iron. The metals
Cr, Mn, Fe, Co, and Ni have a grey iron colour and are very difficult to
melt, but nickel and cobalt can be melted in the reverberatory furnace
and are more fusible than iron, whilst chromium is more difficult to melt
than platinum (Deville). These metals decompose water, but with greater
difficulty as the atomic weight rises, forming a transition to copper,
which does not decompose water. All the compounds of these metals have
various colours, which are sometimes very bright, especially in the
higher stages of oxidation.
These metals of the iron group are often met with together in nature.
Manganese nearly everywhere accompanies iron, and iron is always an
ingredient in the ores of manganese. Chromium is found principally as
chrome ironstone--that is, a peculiar kind of magnetic oxide in which
Fe_{2}O_{3} is replaced by Cr_{2}O_{3}.
Nickel and cobalt are as inseparable companions as iron and manganese.
The similarity between them even extends to such remote properties as
magnetic qualities. In this series of metals we find those which are the
most magnetic: iron, cobalt, and nickel. There is even a magnetic oxide
among the chromium compounds, such being unknown in the other series.
Nickel easily becomes passive in strong nitric acid. It absorbs hydrogen
in just the same way as iron. In short, in the series Cr, Mn, Fe, Co, and
Ni, there are many points in common although there are many differences,
as will be seen still more clearly on becoming acquainted with cobalt and
nickel.
In nature _cobalt_ is principally found in combination with arsenic
and sulphur. _Cobalt arsenide_, or _cobalt speiss_, CoAs_{2}, is found in
brilliant crystals of the regular system, principally in Saxony. _Cobalt
glance_, CoAs_{2}CoS_{2}, resembles it very much, and also belongs to the
regular system; it is found in Sweden, Norway, and the Caucasus.
_Kupfernickel_ is a nickel ore in combination with arsenic, but of a
different composition from cobalt arsenide, having the formula NiAs; it
is found in Bohemia and Saxony. It has a copper-red colour and is rarely
crystalline; it is so called because the miners of Saxony first mistook
it for an ore of copper (_Kupfer_), but were unable to extract copper
from it. _Nickel glance_, NiS_{2},NiAs_{2}, corresponding with cobalt
glance, is also known. Nickel accompanies the ores of cobalt and cobalt
those of nickel, so that both metals are found together. The ores of
cobalt are worked in the Caucasus in the Government of Elizavetopolsk.
Nickel ores containing aqueous hydrated nickel silicate are found in the
Ural (Revdansk). Large quantities of a similar ore are exported into
Europe from New Caledonia. Both ores contain about 12 per cent. Ni.
_Garnierite_, (RO)_{5}(SiO_{2})_{4}1-1/2H_{2}O, where R = Ni and Mg,
predominates in the New Caledonian ore. Large deposits of nickel have
been discovered in Canada, where the ore (as nickelous pyrites) is free
from arsenic. Cobalt is principally worked up into cobalt compounds, but
nickel is generally reduced to the metallic state, in which it is now
often used for alloys--for instance, for coinage in many European States,
and for plating other metals, because it does not oxidise. Cobalt
arsenide and cobalt glance are principally used for the preparation of
cobalt compounds; they are first sorted by discarding the rocky matter,
and then roasted. During this process most of the sulphur and arsenic
disappears; the arsenious anhydride volatilises with the sulphurous
anhydride and the metal also oxidises.[32] It is a simple matter to
obtain nickel and cobalt from their oxides. In order to obtain the
latter, solutions of their salts are treated with sodium carbonate and
the precipitated carbonates are heated; the suboxides are thus obtained,
and these latter are reduced in a stream of hydrogen, or even by heating
with ammonium chloride. They easily oxidise when in the state of powder.
When the chlorides of nickel and cobalt are heated in a stream of
hydrogen, the metal is deposited in brilliant scales. _Nickel is always
much more easily and quickly reduced than cobalt._ Nickel melts more
easily than cobalt, and this even furnishes a means of testing the
heating powers of a reverberatory furnace. Cobalt fuses at a temperature
only a little lower than that at which iron does. In general, cobalt is
nearer to iron than nickel, nickel being nearer to copper.[32 bis] Both
nickel and cobalt have magnetic properties like iron, but Co is less
magnetic than Fe, and Ni still less so. The specific gravity of nickel
reduced by hydrogen is 9·1 and that of cobalt 8·9. Fused cobalt has a
specific gravity of 8·5, the density of ordinary nickel being almost the
same. Nickel has a greyish silvery-white colour; it is brilliant and very
ductile, so that the finest wire may be easily drawn from it. This wire
has a resistance to tension equal to iron wire. The beautiful colour of
nickel, and the high polish which it is capable of receiving and
retaining, as it does not oxidise, render it a useful metal for many
purposes, and in many ways it resembles silver.[32 tri] It is now very
common to cover other metals with a layer of nickel (nickel plating).
This is done by a process of electro-plating, using a solution of a
nickel salt. The colour of cobalt is dark and redder; it is also ductile,
and has a greater tensile resistance than iron. Dilute acids act very
slowly on nickel and cobalt; nitric acid may be considered as the best
solvent for them. The solutions in every case contain salts corresponding
with the ferrous salts--that is, the _salts_ CoX_{2}, NiX_{2},
_correspond with the suboxides_ of these metals. These salts in their
types are similar to the magnesium salts. The salts of nickel when
crystallising with water have a green colour, and form bright green
solutions, but in the anhydrous state they most frequently have a yellow
colour. The salts of cobalt are generally rose-coloured, and generally
blue when in the anhydrous state. Their aqueous solutions are
rose-coloured. Cobaltous chloride is easily soluble in alcohol, and forms
a solution of an intense blue colour.[33]
[32] The residue from the roasting of cobalt ores is called _zafflor_,
and is often met with in commerce. From this the purer compounds
of cobalt may be prepared. The ores of nickel are also first
roasted, and the oxides dissolved in acid, nickelous salts being
then obtained.
The further treatment of cobalt and nickel ores is facilitated if
the arsenic can be almost entirely removed, which may be effected
by roasting the ore a second time with a small addition of nitre
and sodium carbonate; the nitre combines with the arsenic, forming
an arsenious salt, which may be extracted with water. The
remaining mass is dissolved in hydrochloric acid, mixed with a
small quantity of nitric acid. Copper, iron, manganese, nickel,
cobalt, &c., pass into solution. By passing hydrogen sulphide
through the solution, copper, bismuth, lead, and arsenic are
deposited as metallic sulphides; but iron, cobalt, nickel, and
manganese remain in solution. If an alkaline solution of bleaching
powder be then added to the remaining solution, the whole of the
manganese will first be deposited in the form of dioxide, then the
cobalt as hydrated cobaltic oxide, and finally the nickel also. It
is, however, impossible to rely on this method for effecting a
complete separation, the more so since the higher oxides of the
three above-mentioned metals have all a black colour; but, after a
few trials, it will be easy to find how much bleaching powder is
required to precipitate the manganese, and the amount which will
precipitate all the cobalt. The manganese may also be separated
from cobalt by precipitation from a mixture of the solutions of
both metals (in the form of the 'ous' salts) with ammonium
sulphide, and then treating the precipitate with acetic acid or
dilute hydrochloric acid, in which manganese sulphide is easily
soluble and cobalt sulphide almost insoluble. Further particulars
relating to the separation of cobalt from nickel may be found in
treatises on analytical chemistry. In practice it is usual to rely
on the rough method of separation founded on the fact that nickel
is more easily reduced and more difficult to oxidise than cobalt.
The New Caledonian ore is smelted with CaSO_{4} and CaCO_{3} on
coke, and a metallic regulus is obtained containing all the Ni,
Fe, and S. This is roasted with SiO_{2}, which converts all the
iron into slag, whilst the Ni remains combined with the S; this
residue on further roasting gives NiO, which is reduced by the
carbon to metallic Ni. The Canadian ore (a pyrites containing 11
p.c. Ni) is frequently treated in America (after a preliminary
dressing) by smelting it with Na_{2}SO_{4} and charcoal; the
resultant fusible Na_{2}S then dissolves the CuS and FeS_{2},
while the NiS is obtained in a bottom layer (Bartlett and
Thomson's process) from which Ni is obtained in the manner
described above.
For manufacturing purposes somewhat impure cobalt compounds are
frequently used, which are converted into _smalt_. This is glass
containing a certain amount of cobalt oxide; the glass acquires a
bright blue colour from this addition, so that when powdered it
may be used as a blue pigment; it is also unaltered at high
temperatures, so that it used to take the place now occupied by
Prussian blue, ultramarine, &c. At present smalt is almost
exclusively used for colouring glass and china. To prepare smalt,
ordinary impure cobalt ore (zaffre) is fused in a crucible with
quartz and potassium carbonate. A fused mass of cobalt glass is
thus formed, containing silica, cobalt oxide, and potassium oxide,
and a metallic mass remains at the bottom of the crucible,
containing almost all the other metals, arsenic, nickel, copper,
silver, &c. This metallic mass is called _speiss_, and is used as
nickel ore for the extraction of nickel. Smalt usually contains 70
p.c. of silica, 20 p.c. of potash and soda, and about 5 to 6 p.c.
of cobaltous oxide; the remainder consisting of other metallic
oxides.
[32 bis] All we know respecting the relations of Co and Ni to Fe and Cu
confirms the fact that Co is more closely related to Fe and Ni to
Cu; and as the atomic weight of Fe = 56 and of Cu = 63, then
according to the principles of the periodic system it would be
expected that the atomic weight of Co would be about 59-60, whilst
that of Ni should be greater than that of Co but less than that of
Cu, _i.e._ about 50·5-60·5. However, as yet the majority of the
determinations of the atomic weights of Co and Ni give a different
result and show that a lower atomic weight is obtained for Ni than
for Co. Thus K. Winkler (1894) obtained (employing metals
deposited electrolytically and determining the amount of iodine
which combined with them) Ni = 58·72 and Co = 59·37 (if H = 1 and
I = 126·53). In my opinion this should not be regarded as proving
that the principles of the periodic system cannot be applied in
this instance, nor as a reason for altering the position of these
elements in the system (_i.e._ by placing Ni after Fe, and Co next
to Cu), because in the first place the figures given by different
chemists (for instance, Zimmermann, Krüss, and others) are
somewhat divergent, and in the second place the majority of the
latest modes of determining the atomic weights of Co and Ni aim at
finding what weights of these metals react with known weights of
other elements without taking into account the faculty they have
of absorbing hydrogen; since this faculty is more developed in Ni
than in Co the hydrogen (occluded in Ni) should lower the atomic
weight of Ni more than that of Co. On the whole, the question of
the atomic weights of Co and Ni cannot yet be considered as
decided, notwithstanding the numerous researches which have been
made; still there can be no doubt that the atomic weights of these
two metals are very nearly equal, and greater than that of Fe, but
less than that of Cu. This question is of great interest, not only
for completing our knowledge of these metals, but also for
perfecting our knowledge of the periodic system of the elements.
[32 tri] For instance, the alkalis may be fused in nickel vessels as
well as in silver, because they have no action upon either metal.
Nickel, like silver, is not acted upon by dilute acids. Only
nitric acid dissolves both metals well. Nickel is harder, and
fuses at a higher temperature than silver. For castings, a small
quantity of magnesium (0·001 part by weight) is added to nickel to
render it more homogeneous (just as aluminium is added to steel).
Nickel forms many valuable alloys. Steel containing 3 p.c. Ni is
particularly valuable, its limit of elasticity is higher and its
hardness is greater; it is used for armour plate and other large
pieces. The alloys of nickel, especially with copper and zinc
(melchior, _see_ later), aluminium and silver, although used in
certain cases, are now replaced by nickel-plated or
nickel-deposited goods (deposited by electricity from a solution
of the ammonium salts).
[33] The change of colour is dependent in all probability on the
combination with water, or according to others on polymeric
transformation. It enables a solution of cobalt chloride to be
used as sympathetic ink. If something be written with cobalt
chloride on white paper, it will be invisible on account of the
feeble colour of the solution, and when dry nothing can be
distinguished. If, however, the paper be heated before the fire,
the rose-coloured salt will be changed into a less hydrous blue
salt, and the writing will become quite visible, but fade away
when cool.
The change of colour which takes place in solutions of CoCl_{2}
under the influence not only of solution in water or alcohol, but
also of a change of temperature, is a characteristic of all the
halogen salts of cobalt. Crystalline iodide of cobalt,
CoI_{2}6H_{2}O, gives a dark red solution between -22° and +20°;
above +20° the solution turns brown and passes from olive to
green, from +35° to 320° the solution remains green. According to
Étard the change of colour is due to the fact that at first the
solution contains the hydrate CoI_{2}H_{2}O, and that above 35° it
contains CoI_{2}4H_{2}O. These hydrates can be crystallised from
the solutions; the former at ordinary temperature and the latter
on heating the solution. The intermediate olive colour of the
solutions corresponds to the incipient decomposition of the
hexahydrated salt and its passage into CoI_{2}4H_{2}O. A solution
of the hexahydrated chloride of cobalt, CoCl_{2}6H_{2}O, is
rose-coloured between -22° and +25°; but the colour changes
starting from +25°, and passes through all the tints between red
and blue right up to 50°; a true blue solution is only obtained at
55° and remains up to 300°. This true blue solution contains
another hydrate, CoCl_{2}2H_{2}O.
The dependence between the solubility of the iodide and chloride
of cobalt and the temperature is expressed by two almost straight
lines corresponding to the hexa- and di-hydrates; the passage of
the one into the other hydrate being expressed by a curve. The
same character of phenomena is seen also in the variation of the
vapour tension of solutions of chloride of cobalt with the
temperature. We have repeatedly seen that aqueous solutions (for
instance, Chapter XXII., Note 23 for Fe_{2}Cl_{6}) deposit
different crystallo-hydrates at different temperatures, and that
the amount of water in the hydrate decreases as the temperature
_t_ rises, so that it is not surprising that CoCl_{2}2H_{2}O (or
according to Potilitzin CoCl_{2}H_{2}O) should separate out above
55° and CoCl_{2}6H_{2}O at 25° and below. Nor is it exceptional
that the colour of a salt varies according as it contains
different amounts of H_{2}O. But in this instance it is
characteristic that the change of colour takes place in solution
in the presence of an excess of water. This apparently shows that
the actual solution may contain either CoCl_{2}6H_{2}O or
CoCl_{2}2H_{2}O. And as we know that a solution may contain both
metaphosphoric PHO_{3} and orthophosphoric acid H_{3}PO_{4} =
HPO_{3} + H_{2}O, as well as certain other anhydrides, the
question of the state of substances in solutions becomes still
more complicated.
Nickel sulphate crystallises from neutral solutions at a
temperature of from 15° to 20° in _rhombic_ crystals containing
7H_{2}O. Its form approaches very closely to that of the salts of
zinc and magnesium. The planes of a vertical prism for magnesium
salts are inclined at an angle of 90° 30´, for zinc salts at an
angle of 91° 7´, and for nickel salts at an angle of 91° 10´. Such
is also the form of the zinc and magnesium selenates and
chromates. Cobalt sulphate containing 7 molecules of water is
deposited in crystals of the _monoclinic_ system, like the
corresponding salts of iron and manganese. The angle of a vertical
prism for the iron salt = 82° 20´, for cobalt = 82° 22´, and the
inclination of the horizontal pinacoid to the vertical prism for
the iron salt = 99° 2´, and for the cobalt salt 99° 36´. All the
isomorphous mixtures of the salts of magnesium, iron, cobalt,
nickel and manganese have the same form if they contain 7 mol.
H_{2}O and iron or cobalt predominate, whilst if there is a
preponderance of magnesium, zinc, or nickel, the crystals have a
rhombic form like magnesium sulphate. Hence these sulphates are
_dimorphous_, but for some the one form is more stable and for
others the other. Brooke, Moss, Mitscherlich, Rammelsberg, and
Marignac have explained these relations. Brooke and Mitscherlich
also supposed that NiSO_{4},7H_{2}O is not only capable of
assuming these forms, but also that of the _tetragonal_ system,
because it is deposited in this form from acid, and especially
from slightly-heated solutions (30° to 40°). But Marignac
demonstrated that the tetragonal crystals do not contain 7, but 6,
molecules of water, NiSO_{4},6H_{2}O. He also observed that a
solution evaporated at 50° to 70° deposits monoclinic crystals,
but of a different form from ferrous sulphate,
FeSO_{4},7H_{2}O--namely, the angle of the prism is 71° 52´, that
of the pinacoid 95° 6´. This salt appears to be the same with 6
molecules of water as the tetragonal. Marignac also obtained
magnesium and zinc salts with 6 molecules of water by evaporating
their solutions at a higher temperature, and these salts were
found to be isomorphous with the monoclinic nickel salt. In
addition to this it must be observed that the rhombic crystals of
nickel sulphate with 7H_{2}O become turbid under the influence of
heat and light, lose water, and change into the tetragonal salt.
The monoclinic crystals in time also become turbid, and change
their structure, so that the tetragonal form of this salt is the
most stable. Let us also add that nickel sulphate in all its
shapes forms very beautiful emerald green crystals, which, when
heated to 230°, assume a dirty greenish-yellow hue and then
contain one molecule of water.
Klobb (1891) and Langlot and Lenoir obtained anhydrous CoSO_{4}
and NiSO_{4} by igniting the hydrated salt with (NH_{4})_{2}SO_{4}
until the ammonium salt had completely volatilised and decomposed.
We may add that when equivalent aqueous solutions of NiX_{2}
(green) and CoX_{2} (red) are mixed together they give an almost
colourless (grey) solution, in which the green and red colour of
the component parts disappears owing to the combination of the
complementary colours.
A double salt NiKF_{3} is obtained by heating NiCl_{2} with KFHF
in a platinum crucible; KCoF_{3} is formed in a similar manner.
The nickel salt occurs in fine green plates, easily soluble in
water but scarcely soluble in ethyl and methyl alcohol. They
decompose into green oxide of nickel and potassium fluoride when
heated in a current of air. The analogous salt of cobalt
crystallises in crimson flakes.
If instead of potassium fluoride, CoCl_{2} or NiCl_{2} be fused
with ammonium fluoride, they also form double salts with the
latter. This gives the possibility of obtaining anhydrous
fluorides NiF_{2} and CoF_{2}. Crystalline fluoride of nickel,
obtained by heating the amorphous powder formed by decomposing the
double ammonium salt in a stream of hydrofluoric acid, occurs in
beautiful green prisms, sp. gr. 4·63, which are insoluble in
water, alcohol, and ether; sulphuric, hydrochloric, and nitric
acids also have no action upon them, even when heated; NiF_{2} is
decomposed by steam, with the formation of black oxide, which
retains the crystalline structure of the salt. Fluoride of cobalt,
obtained as a rose-coloured powder by decomposing the double
ammonium salt with the aid of heat in a stream of hydrofluoric
acid, fuses into a ruby-coloured mass which bears distinct signs
of a crystalline structure; sp. gr. 4·43. The molten salt only
volatilises at about 1400°, which forms a clear distinction
between CoF_{2} and the volatile NiF_{2}. Hydrochloric, sulphuric,
and nitric acids act upon CoF_{2} even in the cold, although
slowly, while when heated the reaction proceeds rapidly (Poulenc,
1892).
If a solution of potassium hydroxide be added to a solution of a cobalt
salt, a blue precipitate of the basic salt will be formed. If a solution
of a cobalt salt be heated almost to the boiling-point, and the solution
be then mixed with a boiling solution of an alkali hydroxide, a _pink
precipitate of cobaltous hydroxide_, CoH_{2}O_{2}, will be formed. If air
be not completely excluded during the precipitation by boiling, the
precipitate will also contain brown cobaltic hydroxide formed by the
further oxidation of the cobaltous oxide.[34] Under similar circumstances
nickel salts form _a green precipitate of nickelous hydroxide_, the
formation of which is not hindered by the presence of ammonium salts, but
in that case only requires more alkali to completely separate the nickel.
The nickelous oxide obtained by heating the hydroxide, or from the
carbonate or nitrate, is a grey powder, easily soluble in acids and
easily reduced, but the same substance may be obtained in the crystalline
form as an ordinary product from the ores; it crystallises in regular
octahedra, with a metallic lustre, and is of a grey colour. In this state
the nickelous oxide almost resists the action of acids.[34 bis]
[34] Hydrated suboxide of cobalt (de Schulten, 1889) is obtained in the
following manner. A solution of 10 grams of CoCl_{2}6H_{2}O in 60
c.c. of water is heated in a flask with 250 grams of caustic
potash and a stream of coal gas is passed through the solution.
When heated the hydrate of the suboxide of cobalt which separates
out, dissolves in the caustic potash and forms a dark blue
solution. This solution is allowed to stand for 24 hours in an
atmosphere of coal gas (in order to prevent oxidation). The
crystalline mass which separates out has a composition Co(OH)_{2},
and to the naked eye appears as a violet powder, which is seen to
be crystalline under the microscope. The specific gravity of this
hydrate is 3·597 at 15°. It does not undergo change in the air;
warm acetic acid dissolves it, but it is insoluble in warm and
cold solutions of ammonia and sal-ammoniac.
[34 bis] The following reaction may be added to those of the cobaltous
and nickelous salts: potassium cyanide forms a precipitate with
cobalt salts which is soluble in an excess of the reagent and
forms a green solution. On heating this and adding a certain
quantity of acid, a double _cobalt cyanide_ is formed which
corresponds with potassium ferricyanide. Its formation is
accompanied with the evolution of hydrogen, and is founded upon
the property which cobalt has of oxidising in an alkaline
solution, the development of which has been observed in such a
considerable measure in the cobaltamine salts. The process which
goes on here may be expressed by the following equation;
CoC_{2}N_{2} + 4KCN first forms CoK_{4}C_{6}N_{6}, which salt with
water, H_{2}O, forms potassium hydroxide, KHO, hydrogen, H, and
the salt, K_{3}CoC_{6}N_{6}. Here naturally the presence of the
acid is indispensable in consequence of its being required to
combine with the alkali. From aqueous solutions this salt
crystallises in transparent, hexagonal prisms of a yellow colour,
easily soluble in water. The reactions of double decomposition,
and even the formation of the corresponding acid, are here
completely the same as in the case of the ferricyanide. If a
nickelous salt be treated in precisely the same manner as that
just described for a salt of cobalt, decomposition will occur.
It is interesting to note _the relation_ of the cobaltous and
nickelous hydroxides _to ammonia_; aqueous ammonia dissolves the
precipitate of cobaltous and nickelous hydroxide. The blue ammoniacal
solution of nickel resembles the same solution of cupric oxide, but has a
somewhat reddish tint. It is characterised by the fact that it dissolves
silk in the same way as the ammoniacal cupric oxide dissolves cellulose.
Ammonia likewise dissolves the precipitate of cobaltous hydroxide,
forming a brownish liquid, which becomes darker in air and finally
assumes a bright red hue, absorbing oxygen. The admixture of ammonium
chloride prevents the precipitation of cobalt salts by ammonia; when the
ammonia is added, a brown solution is obtained from which, as in the case
of the preceding solution, potassium hydroxide does not separate the
cobaltous oxide. Peculiar compounds are produced in this solution; they
are comparatively stable, containing ammonia and an excess of oxygen;
they bear the name cobaltoamine and cobaltiamine salts. They have been
principally investigated by Genth, Frémy, Jörgenson and others. Genth
found that when a cobalt salt, mixed with an excess of ammonium chloride,
is treated with ammonia and exposed to the air, after a certain lapse of
time, on adding hydrochloric acid and boiling, a red powder is
precipitated and the remaining solution contains an orange salt. The
study of these compounds led to the discovery of a whole series of
similar salts, some of which correspond with particular higher degrees of
oxidation of cobalt, which are described later.[35] Nickel does not
possess this property of absorbing the oxygen of the air when in an
ammoniacal solution. In order to understand this distinction, and in
general the relation of nickel, it is important to observe that cobalt
more easily forms a higher degree of oxidation--namely, _sesquioxide of
cobalt_, _cobaltic oxide_, Co_{2}O_{3}--than nickel, especially in the
presence of hypochlorous acid. If a solution of a cobalt salt be mixed
with barium carbonate and an excess of hypochlorous acid be added, or
chlorine gas be passed through it, then at the ordinary temperature on
shaking, the whole of the cobalt will be separated in the form of black
cobaltic oxide: 2CoSO_{4} + ClHO + 2BaCO_{3} = Co_{2}O_{3} + 2BaSO_{4} +
HCl + 2CO_{2}. Under these circumstances nickelous oxide does not
immediately form black sesquioxide, but after a considerable space of
time it also separates in the form of sesquioxide, Ni_{2}O_{3}, but
always later than cobalt. This is due to the relative difficulty of
further oxidation of the nickelous oxide. It is, however, possible to
oxidise it; if, for instance, the hydroxide NiH_{2}O_{2} be shaken in
water and chlorine gas be passed through it, then nickel chloride will be
formed, which is soluble in water, and insoluble nickelic oxide in the
form of a black precipitate: 3NiH_{2}O_{2} + Cl_{2} = NiCl_{2} +
Ni_{2}O_{3},3H_{2}O. Nickelic oxide may also be obtained by adding sodium
hypochlorite mixed with alkali to a solution of a nickel salt. Nickelic
and cobaltic hydrates are black. Nickelic oxide evolves oxygen with all
acids, and in consequence of this it is not separated as a precipitate in
the presence of acids; thus it evolves chlorine with hydrochloric acid,
exactly like manganese dioxide. When nickelic oxide is dissolved in
aqueous ammonia it liberates nitrogen, and an ammoniacal solution of
nickelous oxide is formed. When heated, nickelic oxide loses oxygen,
forming nickelous oxide. Cobaltic oxide, Co_{2}O_{3}, exhibits more
stability than nickelic oxide, and shows feeble basic properties; thus it
is dissolved in acetic acid without the evolution of oxygen.[35 bis] But
ordinary acids, especially on heating, evolve oxygen, forming a solution
of a cobaltous salt. The presence of a cobaltic salt in a solution of a
cobaltous salt may be detected by the brown colour of the solution and
the black precipitate formed by the addition of alkali, and also from the
fact that such solutions evolve chlorine when heated with hydrochloric
acid. Cobaltic oxide may not only be prepared by the above-mentioned
methods, but also by heating cobalt nitrate, after which a steel-coloured
mass remains which retains traces of nitric acid, but when heated further
to incandescence evolves oxygen, leaving a compound of cobaltic and
cobaltous oxides, similar to magnetic ironstone. Cobalt (but not nickel)
undoubtedly forms besides Co_{2}O_{3} a _dioxide_ CoO_{2}. This is
obtained[36] when the cobaltous oxide is oxidised by iodine or peroxide
of barium.[37]
[35] The cobalt salts may be divided into at least the following
classes, which repeat themselves for Cr, Ir, Rh (we shall not stop
to consider the latter, particularly as they closely resemble the
cobalt salts):--
(_a_) _Ammonium cobalt salts_, which are simply direct compounds
of the cobaltous salts CoX_{2} with ammonia, similar to various
other compounds of the salts of silver, copper, and even calcium
and magnesium, with ammonia. They are easily crystallised from an
ammoniacal solution, and have a pink colour. Thus, for instance,
when cobaltous chloride in solution is mixed with sufficient
ammonia to redissolve the precipitate first formed, octahedral
crystals are deposited which have a composition
CoCl_{2},H_{2}O,6NH_{3}. These salts are nothing else but
combinations with ammonia of crystallisation--if it may be so
termed--likening them in this way to combinations with water of
crystallisation. This similarity is evident both from their
composition and from their capability of giving off ammonia at
various temperatures. The most important point to observe is that
all these salts contain 6 molecules of ammonia to 1 atom of
cobalt, and this ammonia is held in fairly stable connection.
Water decomposes these salts. (Nickel behaves similarly without
forming other compounds corresponding to the true cobaltic.)
(_b_) The solutions of the above-mentioned salts are rendered
turbid by the action of the air; they absorb oxygen and become
covered with a crust of _oxycobaltamine salts_. The latter are
sparingly soluble in aqueous ammonia, have a brown colour, and are
characterised by the fact that with warm water _they evolve
oxygen_, forming salts of the following category: The nitrate may
be taken as an example of this kind of salt; its composition is
CoN_{2}O_{7},5NH_{3},H_{2}O. It differs from cobaltous nitrate,
Co(NO_{3})_{2}, in containing an extra atom of oxygen--that is, it
corresponds with cobalt dioxide, CoO_{2}, in the same way that the
first salts correspond with cobaltous oxide; they contain 5, and
not 6, molecules of ammonia, as if NH_{3} had been replaced by O,
but we shall afterwards meet compounds containing either 5NH_{3}
or 6NH_{3} to each atom of cobalt.
(_c_) _The luteocobaltic salts_ are thus called because they have
a yellow (luteus) colour. They are obtained from the salts of the
first kind by submitting them in dilute solution to the action of
the air; in this case salts of the second kind are not formed,
because they are decomposed by an excess of water, with the
evolution of oxygen and the formation of luteocobaltic salts. By
the action of ammonia the salts of the fifth kind (roseocobaltic)
are also converted into luteocobaltic salts. These last-named
salts generally crystallise readily, and have a yellow colour;
they are comparatively much more stable than the preceding ones,
and even for a certain time resist the action of boiling water.
Boiling aqueous potash liberates ammonia and precipitates hydrated
cobaltic oxide, Co_{2}O_{3},3H_{2}O, from them. This shows that
the luteocobaltic salts correspond with cobaltic oxide,
Co_{2}O_{3}, and those of the second kind with the dioxide. When a
solution of luteocobaltic sulphate,
Co_{2}(SO_{4})_{3},12NH_{3},4H_{2}O, is treated with baryta,
barium sulphate is precipitated, and the solution contains
luteocobaltic hydroxide, Co(OH)_{3},6NH_{3}, which is soluble in
water, is powerfully alkaline, absorbs the oxygen of the air, and
when heated is decomposed with the evolution of ammonia. This
compound therefore corresponds to a solution of cobaltic hydroxide
in ammonia. The luteocobaltic salts contain 2 atoms of cobalt and
12 molecules of ammonia--that is, 6NH_{3} to each atom of cobalt,
like the salts of the first kind. The CoX_{2} salts have a
metallic taste, whilst those of luteocobalt and others have a
purely saline taste, like the salts of the alkali metals. In the
luteo-salts all the X's react (are ionised, as some chemists say)
as in ordinary salts--for instance, all the Cl_{2} is precipitated
by a solution of AgNO_{3}; all the (SO_{4})_{3} gives a
precipitate with BaX_{2}, &c. The double salt formed with PtCl_{4}
is composed in the same manner as the potassium salt,
K_{2}PtCl_{4} = 2KCl + PtCl_{4}, that is, contains
(CoCl_{3},6NH_{3})_{2},3PtCl_{4}, or the amount of chlorine in the
PtCl_{4} is double that in the alkaline salt. In the rosepentamine
(_e_), and rosetetramine (_f_), salts, also all the X's react or
are ionised, but in the (_g_) and (_h_) salts only a portion of
the X's react, and they are equal to the (_e_) and (_f_) salts
minus water; this means that although the water dissolves them it
is not combined with them, as PHO_{3} differs from PH_{3}O_{3};
phenomena of this class correspond exactly to what has been
already (Chapter XXI., Note 7) mentioned respecting the green and
violet salts of oxide of chromium.
(_d_) _The fuscocobaltic salts._ An ammoniacal solution of cobalt
salts acquires a brown colour in the air, due to the formation of
these salts. They are also produced by the decomposition of salts
of the second kind; they crystallise badly, and are separated from
their solutions by addition of alcohol or an excess of ammonia.
When boiled they give up the ammonia and cobaltic oxide which they
contain. Hydrochloric and nitric acids give a yellow precipitate
with these salts, which turns red when boiled, forming salts of
the next category. The following is an example of the composition
of two of the fuscocobaltic salts,
Co_{2}O(SO_{4})_{2},8NH_{3},4H_{2}O and
Co_{2}OCl_{4},8NH_{3},3H_{2}O. It is evident that the
fuscocobaltic salts are ammoniacal compounds of basic cobaltic
salts. The normal cobaltic sulphate ought to have the composition
Co_{2}(SO_{4})_{3} = Co_{2}O_{3},3SO_{3}; the simplest basic salts
will be Co_{2}O(SO_{4})_{2} = Co_{2}O_{3})2SO_{3}, and
Co_{2}O_{2}(SO_{4}) = Co_{2}O_{3},SO_{3}. The fuscocobaltic salts
correspond with the first type of basic salts. They are changed
(in concentrated solutions) into oxycobaltamine salts by
absorption of one atom of oxygen, Co_{2}O_{2}(SO_{4})_{2}. The
whole process of oxidation will be as follows: first of all
Co_{2}X_{4}, a cobaltous salt, is in the solution (X a univalent
haloid, 2 molecules of the salt being taken), then Co_{2}OX_{4},
the basic cobaltic salt (4th series), then Co_{2}O_{2}X_{4}, the
salt of the dioxide (2nd series). The series of basic salts with
an acid, 2HX, forms water and a normal salt, Co_{2}X_{6} (in 3, 5,
6 series). These salts are combined with various amounts of water
and ammonia. Under many conditions the salts of fuscocobalt are
easily transformed into salts of the next series. The salts of the
series that has just been described contain 4 molecules of ammonia
to 1 atom of cobalt.
(_e_) _The roseocobaltic_ (or rosepentamine),
CoX_{2}H_{2}O,5NH_{3}, _salts_, like the luteocobaltic, correspond
with the normal cobaltic salts, but contain less ammonia, and an
extra molecule of water. Thus the sulphate is obtained from
cobaltous sulphate dissolved in ammonia and left exposed to the
air until transformed into a brown solution of the fuscocobaltic
salt; when this is treated with sulphuric acid a crystalline
powder of the roseocobaltic salt,
Co_{2}(SO_{4})_{3},10NH_{3},5H_{2}O, separates. The formation of
this salt is easily understood: cobaltous sulphate in the presence
of ammonia absorbs oxygen, and the solution of the fuscocobaltic
salt will therefore contain, like cobaltous sulphate, one part of
sulphuric acid to every part of cobalt, so that the whole process
of formation may be expressed by the equation: 10NH_{3} +
2CoSO_{4} + H_{2}SO_{4} + 4H_{2}O + O =
Co_{2}(SO_{4})_{3},10NH_{3},5H_{2}O. This salt forms tetragonal
crystals of a red colour, slightly soluble in cold, but readily
soluble in warm water. When the sulphate is treated with baryta,
roseocobaltic hydroxide is formed in the solution, which absorbs
the carbonic anhydride of the air. It is obtained from the next
series by the action of alkalis.
(_f_) The _rosetetramine cobaltic salts_ CoCl_{2},2H_{2}O,4NH_{3}
were obtained by Jörgenson, and belong to the type of the
luteo-salts, only with the substitution of 2NH_{3} for H_{2}O.
Like the luteo- and roseo-salts they give double salts with
PtCl_{4}, similar to the alkaline double salts, for instance
(Co_{2}H_{2}O,4NH_{3})2(SO_{4})_{2}Cl_{2}PtCl_{4}. They are darker
in colour than the preceding, but also crystallise well. They are
formed by dissolving CoCO_{3} in sulphuric acid (of a given
strength), and after NH_{3} and carbonate of ammonium have been
added, air is passed through the solution (for oxidation) until
the latter turns red. It is then evaporated with lumps of
carbonate of ammonium, filtered from the precipitate and
crystallised. A salt of the composition
Co_{2}(CO_{3})_{2}(SO_{4}),(2H_{2}O,4NH_{3})_{2} is thus obtained,
from which the other salts may be easily prepared.
(_g_) The _purpureocobaltic salts_, CoX_{3},5NH_{3}, are also
products of the direct oxidation of ammoniacal solutions of cobalt
salts. They are easily obtained by heating the roseocobaltic and
luteo-salts with strong acids. They are to all effects the same as
the roseocobaltic salts, only anhydrous. Thus, for instance, the
purpureocobaltic chloride, Co_{2}Cl_{6},10NH_{3}, or
CoCl_{3},5NH_{3}, is obtained by boiling the oxycobaltamine salts
with ammonia. There is the same distinction between these salts
and the preceding ones as between the various compounds of
cobaltous chloride with water. In the purpureocobaltic only X_{2}
out of the X_{3} react (are ionised). To the rosetetramine salts
(_f_) there correspond the _purpureotetramine_ salts,
CoX_{3}H_{2}O,4NH_{3}. The corresponding chromium
purpureopentamine salt, CrCl_{3},5NH_{3} is obtained with
particular ease (Christensen, 1893). Dry anhydrous chromium
chloride is treated with anhydrous liquid ammonia in a freezing
mixture composed of liquid CO_{2} and chlorine, and after some
time the mixture is taken out of the freezing mixture, so that the
excess of NH_{3} boils away; the violet crystals then immediately
acquire the red colour of the salt, CrCl_{3},5NH_{3}, which is
formed. The product is washed with water (to extract the
luteo-salt, CrCl_{3},6NH_{3}), which does not dissolve the salt,
and it is then recrystallised from a hot solution of hydrochloric
acid.
(_h_) The _praseocobaltic salts_, CoX_{3},4NH_{3}, are green, and
form, with respect to the rosetetramine salts (_f_), the products
of ultimate dehydration (for example, like metaphosphoric acid
with respect to orthophosphoric acid, but in dissolving in water
they give neither rosetetramine nor tetramine salts. (In my
opinion one should expect salts with a still smaller amount of
NH_{3}, of the blue colour proper to the low hydrated compounds of
cobalt; the green colour of the prazeo-salts already forms a step
towards the blue.) Jörgenson obtained salts for ethylene-diamine,
N_{2}H_{4}C_{2}H_{4} which replaces 2NH_{3}. After being kept a
long time in aqueous solution they give rosetetramine salts, just
as metaphosphoric acid gives orthophosphoric acid, while the
rosetetramine salts are converted into prazeo-salts by Ag_{2}O and
NaHO. Here only one X is ionised out of the X_{3}. There are also
basic salts of the same type; but the best known is the chromium
salt called the rhodozochromic salt,
Cr_{2}(OH)_{3}Cl_{3},6NH_{3},2H_{2}O, which is formed by the
prolonged action of water upon the corresponding roseo-salt.
The cobaltamine compounds differ essentially but little from the
ammoniacal compounds of other metals. The only difference is that
here the cobaltic oxide is obtained from the cobaltous oxide in
the presence of ammonia. In any case it is a simpler question than
that of the double cyanides. Those forces in virtue of which such
a considerable number of ammonia molecules are united with a
molecule of a cobalt compound, appertain naturally to the series
of those slightly investigated forces which exist even in the
highest degrees of combination of the majority of elements. They
are the same forces which lead to the formation of compounds
containing water of crystallisation, double salts, isomorphous
mixtures and complex acids (Chapter XXI., Note 8 bis). The
simplest conception, according to my opinion, of cobalt compounds
(much more so than by assuming special complex radicles, with
Schiff, Weltzien, Claus, and others), may be formed by comparing
them with other ammoniacal products. Ammonia, like water, combines
in various proportions with a multitude of molecules. Silver
chloride and calcium chloride, just like cobalt chloride, absorb
ammonia, forming compounds which are sometimes slightly stable,
and easily dissociated, sometimes more stable, in exactly the same
way as water combines with certain substances, forming fairly
stable compounds called hydroxides or hydrates, or less stable
compounds which are called compounds with water of
crystallisation. Naturally the difference in the properties in
both cases depends on the properties of those elements which enter
into the composition of the given substance, and on those kinds of
affinity towards which chemists have not as yet turned their
attention. If boron fluoride, silicon fluoride, &c., combine with
hydrofluoric acid, if platinic chloride, and even cadmium
chloride, combine with hydrochloric acid, these compounds may be
regarded as double salts, because acids are salts of hydrogen. But
evidently water and ammonia have the same saline faculty, more
especially as they, like haloid acids, contain hydrogen, and are
both capable of further combination--for instance, ammonia with
hydrochloric acid. Hence it is simpler to compare complex
ammoniacal with double salts, hydrates, and similar compounds, but
_the ammonio-metallic salts_ present a most complete qualitative
and quantitative resemblance to _the hydrated salts of metals_.
The composition of the latter is MX_{_n_}_m_H_{2}O, where M =
metal, X = the haloid, simple or complex, and _n_ and _m_ the
quantities of the haloid and so-called water of crystallisation
respectively. The composition of the ammoniacal salts of metals is
MX_{_n_}_m_NH_{3}. The water of crystallisation is held by the
salt with more or less stability, and some salts even do not
retain it at all; some part with water easily when exposed to the
air, others when heated, and then with difficulty. In the case of
some metals all the salts combine with water, whilst with others
only a few, and the water so combined may then be easily
disengaged. All this applies equally well to the ammoniacal salts,
and therefore the combination of ammonia may be termed _the
ammonia of crystallisation_. Just as the water which is combined
with a salt is held by it with different degrees of force, so it
is with ammonia. In combining with 2NH_{3},PtCl_{2} evolves 31,000
cals.; while CaCl_{2} only evolves 14,000 cals.; and the former
compound parts with its NH_{3} (together with HCl in this case)
with more difficulty, only above 200°, while the latter disengages
ammonia at 180°. ZnCl_{2},2NH_{3} in forming ZnCl_{2},4NH_{3}
evolves only 11,000 cals., and splits up again into its components
at 80°. The amount of combined ammonia is as variable as the
amount of water of crystallisation--for instance, SnI_{4}8NH_{3},
CrCl_{2}8NH_{3}, CrCl_{3}6NH_{3},
CrCl_{3}5NH_{3},PtCl_{2},4NH_{3}, &c. are known. Very often NH_{3}
is replaceable by OH_{2} and conversely. A colourless, anhydrous
cupric salt--for instance, cupric sulphate--when combined with
water forms blue and green salts, and violet when combined with
ammonia. If steam be passed through anhydrous copper sulphate the
salt absorbs water and becomes heated; if ammonia be substituted
for the water the heating becomes much more intense, and the salt
breaks up into a fine violet powder. With water CuSO_{4},5H_{2}O
is formed, and with ammonia CuSO_{4},5NH_{3}, the number of water
and ammonia molecules retained by the salt being the same in each
case, and as a proof of this, and that it is not an isolated
coincidence, the remarkable fact must be borne in mind that water
and ammonia consecutively, molecule for molecule, are capable of
supplanting each other, and forming the compounds
CuSO_{4},5H_{2}O, CuSO_{4},4H_{2}O,NH_{3};
CuSO_{4},3H_{2}O,2NH_{3}; CuSO_{4},2H_{2}O,3NH_{3};
CuSO_{4},H_{2}O,4NH_{3}, and CuSO_{4},5NH_{3}. The last of these
compounds was obtained by Henry Rose, and my experiments have
shown that more ammonia than this cannot be retained. By adding to
a strong solution of cupric sulphate sufficient ammonia to
dissolve the whole of the oxide precipitated, and then adding
alcohol, Berzelius obtained the compound CuSO_{4},H_{2}O,4NH_{3},
&c. The law of substitution also assists in rendering these
phenomena clearer, because a compound of ammonia with water forms
ammonium hydroxide, NH_{4}HO, and therefore these molecules
combining with one another may also interchange, as being of equal
value. In general, those salts form stable ammoniacal compounds
which are capable of forming stable compounds with water of
crystallisation; and as ammonia is capable of combining with
acids, and as some of the salts formed by slightly energetic bases
in their properties more closely resemble acids (that is, salts of
hydrogen) than those salts containing more energetic bases, we
might expect to find more stable and more easily-formed
ammonio-metallic salts with metals and their oxides having weaker
basic properties than with those which form energetic bases. This
explains why the salts of potassium, barium, &c., do not form
ammonio-metallic salts, whilst the salts of silver, copper, zinc,
&c., easily form them, and the salts RX_{3} still more easily and
with greater stability. This consideration also accounts for the
great stability of the ammoniacal compounds of cupric oxide
compared with those of silver oxide, since the former is displaced
by the latter. It also enables us to see clearly the distinction
which exists in the stability of the cobaltamine salts containing
salts corresponding with cobaltous oxide, and those corresponding
with higher oxides of cobalt, for the latter are weaker bases than
cobaltous oxides. _The nature of the forces and quality of the
phenomena occurring during the formation of the most stable
substances, and of such compounds as crystallisable compounds, are
one and the same, although perhaps exhibited in a different
degree._ This, in my opinion, may be best confirmed by examining
the compounds of carbon, because for this element the nature of
the forces acting during the formation of its compounds is well
known. Let us take as an example two unstable compounds of carbon.
Acetic acid, C_{2}H_{4}O_{2} (specific gravity 1·06), with water
forms the hydrate, C_{2}H_{4}O_{2},H_{2}O, denser (1·07) than
either of the components, but unstable and easily decomposed,
generally simply referred to as a solution. Such also is the
crystalline compound of oxalic acid, C_{2}H_{2}O_{4}, with water,
C_{2}H_{2}O_{4},2H_{2}O. Their formation might be predicted as
starting from the hydrocarbon C_{2}H_{6}, in which, as in any
other, the hydrogen may be exchanged for chlorine, the water
residue (hydroxyl), &c. The first substitution product with
hydroxyl, C_{2}H_{5}(HO), is stable; it can be distilled without
alteration, resists a temperature higher than 100°, and then does
not give off water. This is ordinary alcohol. The second,
C_{2}H_{4}(HO)_{2}, can also be distilled without change, but can
be decomposed into water and C_{2}H_{4}O (ethylene oxide or
aldehyde); it boils at about 197°, whilst the first hydrate boils
at 78°, a difference of about 100°. The compound
C_{2}H_{3}(HO)_{3} will be the third product of such substitution;
it ought to boil at about 300°, but does not resist this
temperature--it decomposes into H_{2}O and C_{2}H_{4}O_{2}, where
only one hydroxyl group remains, and the other atom of oxygen is
left in the same condition as in ethylene oxide, C_{2}H_{4}O.
There is a proof of this. Glycol, C_{2}H_{4}(HO)_{2}, boils at
197°, and forms water and ethylene oxide, which boils at 13°
(aldehyde, its isomeride, boils at 21°); therefore the product
disengaged by the splitting up of the hydrate boils at 184° lower
than the hydrate C_{2}H_{4}(HO)_{2}. Thus the hydrate
C_{2}H_{3}(HO)_{3}, which ought to boil at about 300°, splits up
in exactly the same way into water and the product
C_{2}H_{4}O_{2}, which boils at 117°--that is, nearly 183° lower
than the hydrate, C_{2}H_{3}(HO)_{3}. But this hydrate splits up
before distillation. The above-mentioned hydrate of acetic acid is
such a decomposable hydrate--that is to say, what is called a
solution. Still less stability may be expected from the following
hydrates. C_{2}H_{2}(HO)_{4} also splits up into water and a
hydrate (it contains two hydroxyl groups) called glycolic acid,
C_{2}H_{2}O(HO)_{2} = C_{2}H_{4}O_{3}. The next product of
substitution will be C_{2}H(HO)_{5}; it splits up into water,
H_{2}O, and glyoxylic acid, C_{2}H_{4}O_{4} (three hydroxyl
groups). The last hydrate which ought to be obtained from
C_{2}H_{6}, and ought to contain C_{2}(HO)_{6}, is the crystalline
compound of oxalic acid, C_{2}H_{2}O_{4} (two hydroxyl groups),
and water, 2H_{2}O, which has been already mentioned. The hydrate
C_{2}(HO)_{6} = C_{2}H_{2}O_{4},2H_{2}O, ought, according to the
foregoing reasoning, to boil at about 600° (because the hydrate,
C_{2}H_{4}(HO)_{2}, boils at about 200°, and the substitution of 4
hydroxyl groups for 4 atoms of hydrogen will raise the
boiling-point 400°). It does not resist this temperature, but at a
much lower point splits up into water, 2H_{2}O, and the hydrate
C_{2}O_{2}(HO)_{2}, which is also capable of yielding water.
Without going into further discussion of this subject, it may be
observed that the formation of the hydrates, or compounds with
water of crystallisation, of acetic and oxalic acids has thus
received an accurate explanation, illustrating the point we
desired to prove in affirming that compounds with water of
crystallisation are held together by the same forces as those
which act in the formation of other complex substances, and that
the easy displaceability of the water of crystallisation is only a
peculiarity of a local character, and not a radical point of
distinction. All the above-mentioned hydrates, C_{2}X_{6}, or
products of their destruction, are actually obtained by the
oxidation of the first hydrate, C_{2}H_{3}(HO), or common alcohol,
by nitric acid (Sokoloff and others). Hence the forces which
induce salts to combine with _n_H_{2}O or with NH_{3} are
undoubtedly of the same order as the forces which govern the
formation of ordinary 'atomic' and saline compounds. (A great
impediment in the study of the former was caused by the conviction
which reigned in the sixties and seventies, that 'atomic' were
essentially different from 'molecular' compounds like
crystallohydrates, in which it was assumed that there was a
combination of entire molecules, as though without the
participation of the atomic forces.) If the bond between chlorine
and different metals is not equally strong, so also the bond
uniting _n_H_{2}O and _n_NH_{3} is exceeding variable; there is
nothing very surprising in this. And in the fact that the
combination of different amounts of NH_{3} and H_{2}O alters the
capacity of the haloids X of the salts RX_{2} for reaction (for
instance, in the luteo-salts all the X_{3}, while in the purpureo,
only 2 out of the 3, and in the prazeo-salts only 1 of the 3 X's
reacts), we should see in the first place a phenomenon similar to
what we met with in Cr_{2}Cl_{6} (Chapter XXI., Note 7 bis), for
in both instances the essence of the difference lies in the
removal of water; a molecule RCl_{3},6H_{2}O or RCl_{3},6NH_{3}
contains the halogen in a perfectly mobile (ionised) state, while
in the molecule RCl_{3},5H_{2}O or RCl_{3},5NH_{3} a portion of
the halogen has almost lost its faculty for reacting with
AgNO_{3}, just as metalepsical chlorine has lost this faculty
which is fully developed in the chloranhydride. Until the reason
of this difference be clear, we cannot expect that ordinary points
of view and generalisation can give a clear answer. However, we
may assume that here the explanation lies in the nature and kind
of motion of the atoms in the molecules, although as yet it is not
clear how. Nevertheless, I think it well to call attention again
(Chapter I.) to the fact that the combination of water, and hence,
also, of any other element, leads to most diverse consequences;
the water in the gelatinous hydrate of alumina or in the
decahydrated Glauber salt is very mobile, and easily reacts like
water in a free state; but the same water combined with oxide of
calcium, or C_{2}H_{4} (for instance, in C_{2}H_{6}O and in
C_{4}H_{10}O), or with P_{2}O_{5}, has become quite different, and
no longer acts like water in a free state. We see the same
phenomenon in many other cases--for example, the chlorine in
chlorates no longer gives a precipitate of chloride of silver with
AgNO_{3}. Thus, although the instance which is found in the
difference between the roseo- and purpureo-salts deserves to be
fully studied on account of its simplicity, still it is far from
being exceptional, and we cannot expect it to be thoroughly
explained unless a mass of similar instances, which are
exceedingly common among chemical compounds, be conjointly
explained. (Among the researches which add to our knowledge
respecting the complex ammoniacal compounds, I think it
indispensable to call the reader's attention to Prof. Kournakoff's
dissertation 'On complex metallic bases,' 1893.)
Kournakoff (1894) showed that the solubility of the luteo-salt,
CoCl_{3},6NH_{3}, at 0° = 4·30 (per 100 of water), at 20° = 7·7,
that in passing into the roseo-salt, CoCl_{3}H_{2}O_{5}NH_{3}, the
solubility rises considerably, and at 0° = 16·4, and at 20° =
about 27, whilst the passage into the purpureo-salt,
CoCl_{3},5NH_{3}, is accompanied by a great fall in the
solubility, namely, at 0° = 0·23, and at 20° = about 0·5. And as
crystallohydrates with a smaller amount of water are usually more
soluble than the higher crystallohydrates (Le Chatelier), whilst
here we find that the solubility falls (in the purpureo-salt) with
a loss of water, that water which is contained in the roseo-salt
cannot be compared with the water of crystallisation. Kournakoff,
therefore, connects the fall in solubility (in the passage of the
roseo- into the purpureo-salts) with the accompanying loss in the
reactive capacity of the chlorine.
In conclusion, it may be observed that the elements of the eighth
group--that is, the analogues of iron and platinum--according to
my opinion, will yield most fruitful results when studied as to
combinations with whole molecules, as already shown by the
examples of complex ammoniacal, cyanogen, nitro-, and other
compounds, which are easily formed in this eighth group, and are
remarkable for their stability. This faculty of the elements of
the eighth group for forming the complex compounds alluded to, is
in all probability connected with the position which the eighth
group occupies with regard to the others. Following the seventh,
which forms the type RX_{7}, it might be expected to contain the
most complex type, RX_{8}. This is met with in OsO_{4}. The other
elements of the eighth group, however, only form the lower types
RX_{2}, RX_{3}, RX_{4} ... and these accordingly should be
expected to aggregate themselves into the higher types, which is
accomplished in the formation of the above-mentioned complex
compounds.
[35 bis] Marshall (1891) obtained cobaltic sulphate,
Co_{2}(SO_{4})_{3},18H_{2}O, by the action of an electric current
upon a strong solution of CoSO_{4}.
[36] The action of an alkaline hypochlorite or hypobromite upon a
boiling solution of cobaltous salts, according to Schroederer
(1889), produces oxides, whose composition varies between
Co_{3}O_{5} (Rose's compound) and Co_{2}O_{3}, and also between
Co_{5}O_{8} and Co_{12}O_{19}. If caustic potash and then bromine
be added to the liquid, only Co_{2}O_{3} is formed. The action of
alkaline hypochlorites or hypo-bromites, or of iodine, upon
cobaltic salts, gives a highly-coloured precipitate which has a
different colour to the hydrate of the oxide Co_{2}(OH)_{6}.
According to Carnot the precipitate produced by the hypochlorites
has a composition Co_{10}O_{16}, whilst that given by iodine in
the presence of an alkali contains a larger amount of oxygen.
Fortmann (1891) reinvestigated the composition of the higher
oxygen oxide obtained by iodine in the presence of alkali, and
found that the greenish precipitate (which disengages oxygen when
heated to 100°) corresponds to the formula CoO_{2}. The reaction
must be expressed by the equation: CoX_{2} + I_{2} + 4KHO =
CoO_{2} + 2KX + 2KI + 2H_{2}O.
[37] Prior to Fortmann, Rousseau (1889) endeavoured to solve the
question as to whether CoO_{2} was able to combine with bases. He
succeeded in obtaining a barium compound corresponding to this
oxide. Fifteen grams of BaCl_{2} or BaBr_{2} are triturated with
5-6 grams of oxide of barium, and the mixture heated to redness in
a closed platinum crucible; 1 gram of oxide of cobalt is then
gradually added to the fused mass. Each addition of oxide is
accompanied by a violent disengagement of oxygen. After a short
time, however, the mass fuses quietly, and a salt settles at the
bottom of the crucible, which, when freed from the residue,
appears as black hexagonal, very brilliant crystals. In dissolving
in water this substance evolves chlorine; its composition
corresponds to the formula 2(CoO_{2})BaO. If the original mass be
heated for a long time (40 hours), the amount of dioxide in the
resultant mass decreases. The author obtained a neutral salt
having the composition CoO_{2}BaO (this compound = BaO_{2}CoO) by
breaking up the mass as it agglomerates together, and bringing the
pieces into contact with the more heated surface of the crucible.
This salt is formed between the somewhat narrow limits of
temperature 1,000°-1,100°; above and below these limits compounds
richer or poorer in CoO_{2} are formed. The formation of CoO_{2}
by the action of BaO_{2}, and the easy decomposition of CoO_{2}
with the evolution of oxygen, give reason for thinking that it
belongs to the class of peroxides (like Cr_{2}O_{7}, CaO_{2},
&c.); it is not yet known whether they give peroxide of hydrogen
like the true peroxides. The fact that it is obtained by means of
iodine (probably through HIO), and its great resemblance to
MnO_{2}, leads rather to the supposition that CoO_{2} is a very
feeble saline oxide. The form CoO_{2} is repeated in the cobaltic
compounds (Note 35), and the existence of CoO_{2} should have long
ago been recognised upon this basis.
Nickel alloys possess qualities which render them valuable for technical
purposes, the alloy of nickel with iron being particularly remarkable.
This alloy is met with in nature as _meteoric iron_. The Pallasoffsky
mass of meteoric iron, preserved in the St. Petersburg Academy, fell in
Siberia in the last century; it weighs about 15 cwt. and contains 88 p.c.
of iron and about 10 p.c. of nickel, with a small admixture of other
metals. In the arts _German silver_ is most extensively used; it is an
alloy containing nickel, copper, and zinc in various proportions. It
generally consists of about 50 parts of copper, 25 parts of zinc, and 25
parts of nickel. This alloy is characterised by its white colour
resembling that of silver, and, like this latter metal, it does not rust,
and therefore furnishes an excellent substitute for silver in the
majority of cases where it is used. Alloys which contain silver in
addition to nickel show the properties of silver to a still greater
extent. Alloys of nickel are used for currency, and if rich deposits of
nickel are discovered a wide field of application lies before it, not
only in a pure state (because it is a beautiful metal and does not rust)
but also for use in alloys. Steel vessels (pressed or forged out of sheet
steel) covered with nickel have such practical merits that their
manufacture, which has not long commenced, will most probably be rapidly
developed, whilst nickel steel, which exceeds ordinary steel in its
tenacity, has already proved its excellent qualities for many purposes
(for instance, for armour plate).
Until 1890 no compound of cobalt or nickel was known of sufficient
volatility to determine the molecular weights of the compounds of these
metals; but in 1890 Mr. L. Mond, in conducting (together with Langer and
Quincke) his researches on the action of nickel upon carbonic oxide
(Chapter IX., Note 24 bis), observed that nickel gradually volatilises in
a stream of carbonic oxide; this only takes place at low temperatures,
and is seen by the coloration of the flame of the carbonic oxide. This
observation led to the discovery of a remarkable volatile _compound of
nickel and carbonic oxide_, having as molecular composition
Ni(CO)_{4},[38] as determined by the vapour density and depression of the
freezing point. Cobalt and many other metals do not form volatile
compounds under these conditions, but iron gives a similar product (Note
26 bis). Ni(CO)_{4} is prepared by taking finely divided Ni (obtained by
reducing NiO by heating it in a stream of hydrogen, or by igniting the
oxalate NiC_{2}O_{4})[39] and passing (at a temperature below 50°, for
even at 60° decomposition may take place and an explosion) a stream of CO
over it; the latter carries over the vapour of the compound, which
condenses (in a well-cooled receiver) into a perfectly colourless
extremely mobile liquid, boiling without decomposition at 43°, and
crystallising in needles at -25° (Mond and Nasini, 1891). Liquid
Ni(CO)_{4} has a sp. gr. 1·356 at 0°, is insoluble in water, dissolves in
alcohol and benzene, and burns with a very smoky flame due to the
liberation of Ni. The vapour when passed through a tube heated to 180°
and above deposits a brilliant coating of metal, and disengages CO. If
the tube be strongly heated the decomposition is accompanied by an
explosion. If Ni(CO)_{4} as vapour be passed through a solution of
CuCl_{2}, it reduces the latter to metal; it has the same action upon an
ammoniacal solution of AgCl, strong nitric acid oxidises Ni(CO)_{4},
dilute solutions of acids have no action; if the vapour be passed through
strong sulphuric acid, CO is liberated, chlorine gives NiCl and COCl_{2};
no simple reactions of double decomposition are yet known for Ni(CO)_{4},
however, so that its connection with other carbon compounds is not clear.
Probably the formation of this compound could be applied for extracting
nickel from its ores.[40]
[38] This compound is known as nickel tetra-carbonyl. It appears to me
yet premature to judge of the structure of such an extraordinary
compound as Ni(CO)_{4}. It has long been known that potassium
combines with CO forming K_{_n_}(CO)_{_n_} (Chapter IX., Note 31),
but this substance is apparently saline and non-volatile, and has
as little in common with Ni(CO)_{4} as Na_{2}H has with SbH_{3}.
However, Berthelot observed that when NiC_{4}O_{4} is kept in air,
it oxidises and gives a colourless compound,
Ni_{3}C_{2}O_{3},10H_{2}O, having apparently saline properties. We
may add that Schützenberger, on reducing NiCl_{2} by heating it in
a current of hydrogen, observed that a nickel compound partially
volatilises with the HCl and gives metallic nickel when heated
again. The platinum compound, PtCl_{2}(CO)_{3} (Chapter XXIII.,
Note 11), offers the greatest analogy to Ni(CO)_{4}. This compound
was obtained as a volatile substance by Schützenberger by
moderately heating (to 235°) metallic platinum in a mixture of
chlorine and carbonic oxide. If we designate CO by Y, and an atom
of chlorine by X, then taking into account that, according to the
periodic system, Ni is an analogue of Pt, a certain degree of
correspondence is seen in the composition NiY_{4} and
PtX_{2}Y_{2}. It would be interesting to compare the reactions of
the two compounds.
[39] According to its empirical formula oxalate of nickel also contains
nickel and carbonic oxide.
[40] The following are the thermo-chemical data (according to Thomsen,
and referred to gram weights expressed by the formula, in large
calories or thousand units of heat) for the formation of
corresponding compounds of Mn, Fe, Co, Ni, and Cu (+ Aq signifies
that the reaction proceeds in an excess of water):
R = Mn Fe Co Ni Cu
R + Cl_{2} + Aq 128 100 95 94 63
R + Br_{2} + Aq 106 78 73 72 41
R + I_{2} + Aq 76 48 43 41 32
R + O + H_{2}O 95 68 63 61 38
R + O_{2} + SO_{2} + _n_H_{2}O 193 169 163 163 130
RCl_{2} + Aq +16 18 18 19 11
These examples show that for analogous reactions the amount of
heat evolved in passing from Mn to Fe, Co, Ni, and Cu varies in
regular sequences as the atomic weight increases. A similar
difference is to be found in other groups and series, and proves
that thermo-chemical phenomena are subject to the periodic law.
CHAPTER XXIII
THE PLATINUM METALS
The six metals: ruthenium, Ru, rhodium, Rh, palladium, Pd, osmium, Os,
iridium, Ir, and platinum, Pt, are met with associated together in
nature. Platinum always predominates over the others, and hence they are
known as the _platinum metals_. By their chemical character their
position in the periodic system is in the eighth group, corresponding
with iron, cobalt, and nickel.
The natural transition from titanium and vanadium to copper and zinc by
means of the elements of the iron group is demonstrated by all the
properties of these elements, and in exactly the same manner a transition
from zirconium, niobium, and molybdenum to silver, cadmium, and indium,
through ruthenium, rhodium, and palladium, is in perfect accordance with
fact and with the magnitude of the atomic weights, as also is the
position of osmium, iridium, and platinum between tantalum and tungsten
on the one side, and gold and mercury on the other. In all these three
cases the elements of smaller atomic weight (chromium, molybdenum, and
tungsten) are able, in their higher grades of oxidation, to give acid
oxides having the properties of distinct but feebly energetic acids (in
the lower oxides they give bases), whilst the elements of greater atomic
weight (zinc, cadmium, mercury), even in their higher grades of
oxidation, only give bases, although with feebly developed basic
properties. The platinum metals present the same intermediate properties
such as we have already seen in iron and the elements of the eighth
group.
In the platinum metals the intermediate properties _of feebly acid and
feebly basic metals_ are developed with great clearness, so that there is
not one sharply-defined acid anhydride among their oxides, although there
is a great diversity in the grades of oxidation from the type RO_{4} to
R_{2}O. The feebleness of the chemical forces observed in the platinum
metals is connected with the ready decomposability of their compounds,
with the small atomic volume of the metals themselves, and with their
large atomic weight. The oxides of platinum, iridium, and osmium can
scarcely be termed either basic or acid; they are capable of combinations
of both kinds, each of which is feeble. They are all intermediate oxides.
The atomic weights of platinum, iridium, and osmium are nearly 191 to
196, and of palladium, rhodium, and ruthenium, 104 to 106. Thus, strictly
speaking, we have here two series of metals, which are, moreover,
perfectly parallel to each other; three members in the first series, and
three members in the second--namely, platinum presents an analogy to
palladium, iridium to rhodium, and osmium to ruthenium. As a matter of
fact, however, the whole _group_ of the platinum metals is characterised
by _a number of common properties_, both physical and chemical, and,
moreover, there are several points of resemblance between the members of
this group and those of the _iron_ group (Chapter XXII.) The atomic
volumes (Table III., column 18) of the elements of this group are _nearly
equal_ and _very small_. The iron metals have atomic volumes of nearly 7,
whilst that of the metals allied to palladium is nearly 9, and of those
adjacent to platinum (Pt, Ir, Os) nearly 9·4. This comparatively small
atomic volume corresponds with the great infusibility and tenacity proper
to all the iron and platinum metals, and to their small chemical energy,
which stands out very clearly in the heavy platinum metals. All the
platinum metals are very _easily reduced_ by ignition and by the action
of various reducing agents, in which process oxygen, or a haloid group,
is disengaged from their compounds and the metal left behind. This is a
property of the platinum metals which determines many of their reactions,
and the circumstance of their always being found in nature _in a native
state_. In Russia in the Urals (discovered in 1819) and in Brazil (1735)
platinum is obtained from alluvial deposits, but in 1892 Professor
Inostrantseff discovered a vein deposit of platinum in serpentine near
Tagil in the Urals.[1] The facility with which they are reduced is so
great that their chlorides are even decomposed by gaseous hydrogen,
especially when shaken up and heated under a certain pressure. Hence it
will be readily understood that such metals as zinc, iron, &c., separate
them from solutions with great ease, which fact is taken advantage of in
practice and in the chemical treatment of the platinum metals.[1 bis]
[1] Wells and Penfield (1888) have described a mineral sperryllite
found in the Canadian gold-bearing quartz and consisting of
platinum diarsenide, PtAs_{2}. It is a noticeable fact that this
mineral clearly confirms the position of platinum in the same group
as iron, because it corresponds in crystalline form (regular
octahedron) and chemical composition with iron pyrites, FeS_{2}.
[1 bis] Some light is thrown upon the facility with which the platinum
compounds decompose by Thomsen's data, showing that in an excess of
water (+ Aq) the formation from platinum, of such a double salt as
PtCl_{2},2KCl, is accompanied by a comparatively small evolution of
heat (_see_ Chapter XXI., Note 40), for instance, Pt + Cl_{2} +
2KCl + Aq only evolves about 33,000 calories (hence the reaction,
Pt + Cl_{2} + Aq, will evidently disengage still less, because
PtCl_{2} + 2KCl evolves a certain amount of heat), whilst on the
other hand, Fe + Cl_{2} + Aq gives 100,000 calories, and even the
reaction with copper (for the formation of the double salt) evolves
63,000 calories.
All the platinum metals, like those of the iron group, are grey, with a
comparatively feeble metallic lustre, and are very infusible. In this
respect they stand in the same order as the metals of the iron series;
nickel is more fusible and whiter than cobalt and iron, so also palladium
is whiter and more fusible than rhodium and ruthenium, and platinum is
comparatively more fusible and whiter than iridium or osmium. The saline
compounds of these metals are red or yellow, like those of the majority
of the metals of the iron series, and like the latter, the different
forms of oxidation present different colours. Moreover, certain complex
compounds of the platinum metals, like certain complex compounds of the
iron series, either have particular characteristic tints or else are
colourless.
The platinum metals are found _in nature associated together in alluvial
deposits_ in a few localities, from which they are washed, owing to their
very considerable density, which enables a stream of water to wash away
the sand and clay with which they are mixed. Platinum deposits are
chiefly known in the Urals, and also in Brazil and a few other
localities. The platinum ore washed from these alluvial deposits presents
the appearance of more or less coarse grains, and sometimes, as it were,
of semi-fused nuggets.[2]
[2] The largest amount of platinum is extracted in the Urals, about
five tons annually. A certain amount of gold is extracted from the
washed platinum by means of mercury, which does not dissolve the
platinum metals but dissolves the gold accompanying the platinum in
its ores. Moreover, the ores of platinum always contain metals of
the iron series associated with them. The washed and mechanically
sorted ore in the majority of cases contains about 70 to 80 p.c. of
platinum, about 5 to 8 p.c. of iridium, and a somewhat smaller
quantity of osmium. The other platinum metals--palladium, rhodium,
and ruthenium--occur in smaller proportions than the three above
named. Sometimes grains of almost pure osmium-iridium, containing
only a small quantity of other metals, are found in platinum ores.
This _osmium-iridium_ may be easily separated from the other
platinum metals, owing to its being nearly insoluble in aqua regia,
by which the latter are easily dissolved. There are grains of
platinum which are magnetic. The grains of osmium-iridium are very
hard and malleable, and are therefore used for certain purposes,
for instance, for the tips of gold pens.
All the platinum metals give compounds with the halogens, and the
highest haloid type of combination for all is RX_{4}. For the majority of
the platinum metals this type is exceedingly unstable; the lower
compounds corresponding to the type RX_{2}, which are formed by the
separation of X_{2}, are more stable. In the type RX_{2} the platinum
metals form more stable salts, which offer no little resemblance to the
kindred compounds of the iron series--for example, to nickelous chloride,
NiCl_{2}, cobaltous chloride, CoCl_{2}, &c. This even expresses itself in
a similarity of volume (platinous chloride, PtCl_{2}, volume, 46;
nickelous chloride, NiCl_{2} = 50), although in the type RX_{2} the true
iron metals give very stable compounds, whilst the platinum metals
frequently react after the manner of suboxides, decomposing into the
metal and higher types, 2RX_{2} = R + RX_{4}. This probably depends on
the facility with which RX_{2} decomposes into R and X_{2}, when X_{2}
combines with the remaining portion of RX_{2}.
As in the series iron, cobalt, nickel, nickel gives NiO and Ni_{2}O_{3},
whilst cobalt and iron give higher and varied forms of oxidation, so also
among the platinum metals, platinum and palladium only give the forms
RX_{2} and RX_{4}, whilst rhodium and iridium form another and
intermediate type, RX_{3}, also met with in cobalt, corresponding with
the oxide, having the composition R_{2}O_{3}, besides which they form an
acid oxide, like ferric acid, which is also known in the form of salts,
but is in every respect unstable. _Osmium_ and _ruthenium_, like
manganese, form still higher oxides, and in this respect exhibit the
greatest diversity. They not only give RX_{2}, RX_{3}, RX_{4}, and
RX_{6}, but also a still _higher form of oxidation_, RO_{4}, which is not
met with in any other series. This form is exceedingly characteristic,
owing to the fact that the oxides, OsO_{4} and RuO_{4}, are volatile and
have feebly acid properties. In this respect they most resemble
permanganic anhydride, which is also somewhat volatile.[3]
[3] In characterising the platinum metals according to their relation
to the iron metals, it is very important to add two more very
remarkable points. The platinum metals are capable of forming a
sort of unstable compound with _hydrogen_; they absorb it and only
part with it when somewhat strongly heated. This faculty is
especially developed in platinum and palladium, and it is very
characteristic that nickel, which exactly corresponds with platinum
and palladium in the periodic system, should exhibit the same
faculty for retaining a considerable quantity of hydrogen (Graham's
and Raoult's experiments). Another characteristic property of the
platinum metals consists in their easily giving (like cobalt which
forms the cobaltic salts) stable and characteristic saline
_compounds with ammonia_, and like Fe and Co, double salts with the
cyanides of the alkali metals, especially in their lower forms of
combination. All the above so clearly brings the elements of the
iron series in close relation to the platinum metals, that the
eighth group acquires as natural a character as can be required,
with a certain originality or individuality for each element.
When dissolved in aqua regia (PtCl_{4} is formed) and liberated from the
solution by sal-ammoniac ((NH_{4})_{2}PtCl_{6} is formed) and reduced by
ignition (which may be done by Zn and other reducing agents, direct from
a solution of PtCl_{4}) platinum[3 bis] forms a powdery mass, known as
spongy platinum or platinum black. If this powder of platinum be heated
and pressed, or hammered in a cylinder, the grains aggregate or forge
together, and form a continuous, though of course not entirely
homogeneous, mass. Platinum was formerly, and is even now, worked up in
this manner. The platinum money formerly used in Russia was made in this
way. Sainte-Claire Deville, in the fifties, for the first time melted
platinum in considerable quantities by employing a special furnace made
in the form of a small reverberatory furnace, and composed of two pieces
of lime, on which the heat of the oxyhydrogen flame has no action. Into
this furnace (shown in fig. 34, Vol. I. p. 175)--or, more strictly
speaking, into the cavity made in the pieces of lime--the platinum is
introduced, and two orifices are made in the lime; through one, the
upper, or side orifice, is introduced an oxyhydrogen gas burner, in which
either detonating gas or a mixture of oxygen and coal-gas is burnt,
whilst the other orifice serves for the escape of the products of
combustion and certain impurities which are more volatile than the
platinum, and especially the oxidised compounds of osmium, ruthenium, and
palladium, which are comparatively easily volatilised by heat. In this
manner the platinum is converted into a continuous metallic form by means
of fusion, and this method is now used for melting considerable masses of
platinum[4] and its alloys with iridium.
[3 bis] Platinum was first obtained in the last century from Brazil,
where it was called silver (platinus). Watson in 1750 characterised
platinum as a separate independent metal. In 1803 Wollaston
discovered palladium and rhodium in crude platinum, and at about
the same time Tennant distinguished iridium and osmium in it.
Professor Claus, of Kazan, in his researches on the platinum metals
(about 1840) discovered ruthenium in them, and to him are due many
important discoveries with regard to these elements, such as the
indication of the remarkable analogy between the series Pd--Rh--Ru
and Pt--Ir--Os.
_The treatment of platinum ore_ is chiefly carried on for the
extraction of the platinum itself and its alloys with iridium,
because these metals offer a greater resistance to the action of
chemical reagents and high temperatures than any of the other
malleable and ductile metals, and therefore the wire so often used
in the laboratory and for technical purposes is made from them, as
also are various vessels used for chemical purposes in the
laboratory and in works. Thus sulphuric acid is distilled in
platinum retorts, and many substances are fused, ignited, and
evaporated in the laboratory in platinum crucibles and on platinum
foil. Gold and many other substances are dissolved in dishes made
of iridium-platinum, because the alloys of platinum and iridium are
but slightly attacked when subjected to the action of aqua regia.
The comparatively high density (about 21·5), hardness, ductility,
and infusibility (it does not melt at a furnace heat, but only in
the oxyhydrogen flame or electric furnace), as well as the fact of
its resisting the action of water, air, and other reagents, renders
an alloy of 90 parts of platinum and 10 parts of iridium (Deville's
platinum-iridium alloy) a most valuable material for making
standard weights and measures, such as the metre, kilogram, and
pound, and therefore all the newest standards of most countries are
made of this alloy.
[4] This process has altered the technical treatment of platinum to a
considerable extent. It has in particular facilitated the
manufacture of alloys of platinum with iridium and rhodium from the
pure platinum ores, since it is sufficient to fuse the ore in order
for the greater amount of the osmium to burn off, and for the mass
to fuse into a homogeneous, malleable alloy, which can be directly
made use of. There is very little ruthenium in the ores of
platinum. If during fusion lead be added, it dissolves the platinum
(and other platinum metals) owing to its being able to form a very
characteristic alloy containing PtPb. If an alloy of the two metals
be left exposed to moist air, the excess of lead is converted into
carbonate (white lead) in the presence of the water and carbonic
acid of the air, whilst the above platinum alloy remains unchanged.
The white lead may be extracted by dilute acid, and the alloy PtPb
remains unaltered. The other platinum metals also give similar
alloys with lead. The fusibility of these alloys enables the
platinum metals to be separated from the gangue of the ore, and
they may afterwards be separated from the lead by subjecting the
alloy to oxidation in furnaces furnished with a bone ash bed,
because the lead is then oxidised and absorbed by the bone ash,
leaving the platinum metals untouched. This method of treatment was
proposed by H. Sainte-Claire Deville in the sixties, and is also
used in the analysis of these metals (_see_ further on).
To obtain pure platinum, the ore is treated with aqua regia in which
only the osmium and iridium are insoluble. The solution contains the
platinum metals in the form RCl_{4}, and in the lower forms of
chlorination, RCl_{3} and RCl_{2}, because some of these metals--for
instance, palladium and rhodium--form such unstable chlorides of the type
RX_{4} that they partially decompose even when diluted with water, and
pass into the stable lower type of combination; in addition to which the
chlorine is very easily disengaged if it comes in contact with substances
on which it can act. In this respect platinum resists the action of heat
and reducing agents better than any of its companions--that is, it passes
with greater difficulty from PtCl_{4} to the lower compound PtCl_{2}. On
this is based the method of preparation of more or less pure platinum.
Lime or sodium hydroxide is added to the solution in aqua regia until
neutralised, or only containing a very slight excess of alkali. It is
best to first evaporate and slightly ignite the solution, in order to
remove the excess of acid, and by heating it to partially convert the
higher chlorides of the palladium, &c., into the lower. The addition of
alkalis completes the reduction, because the chlorine held in the
compounds RX_{4} acts on the alkali like free chlorine, converting it
into a hypochlorite. Thus palladium chloride, PdCl_{4}, for example, is
converted into palladious chloride, PdCl_{2}, by this means, according to
the equation PdCl_{4} + 2NaHO = PdCl_{2} + NaCl + NaClO + H_{2}O. In a
similar manner iridic chloride, IrCl_{4}, is converted into the
trichloride, IrCl_{3}, by this method. When this conversion takes place
the platinum still remains in the form of platinic chloride, PtCl_{4}. It
is then possible to take advantage of a certain difference in the
properties of the higher and lower chlorides of the platinum metals. Thus
lime precipitates the lower chlorides of the members of the platinum
metals occurring in solution without acting on the platinic chloride,
PtCl_{4}, and hence the addition of a large proportion of lime
immediately precipitates the associated metals, leaving the platinum
itself in solution in the form of a soluble double salt,
PtCl_{4},CaCl_{2}. A far better and more perfect _separation_ is effected
_by means of ammonium chloride_, which gives, with platinic chloride, an
insoluble yellow precipitate, PtCl_{4},2NH_{4}Cl, whilst it forms soluble
double salts with the lower chlorides RCl_{2} and RCl_{3}, so that
ammonium chloride precipitates the platinum only from the solution
obtained by the preceding method. These methods are employed for
preparing the platinum which is used for the manufacture of platinum
articles, because, having platinum in solution as calcium
platinochloride, PtCaCl_{6}, or as the insoluble ammonium
platinochloride, Pt(NH_{4})_{2}Cl_{6}, the platinum compound in every
case, after drying or ignition, loses all the chlorine from the platinic
chloride and leaves finely-divided metallic platinum, which may be
converted into homogeneous metal by compression and forging, or by
fusion.[5]
[5] For the ultimate purification of platinum from palladium and
iridium the metals must be re-dissolved in aqua regia, and the
solution evaporated until the residue begins to evolve chlorine.
The residue is then re-precipitated with ammonium or potassium
chloride. The precipitate may still contain a certain amount of
iridium, which passes with greater difficulty from the
tetrachloride, IrCl_{4}, into the trichloride, IrCl_{3}, but it
will be quite free from palladium, because the latter easily loses
its chlorine and passes into palladious chloride, PdCl_{2}, which
gives an easily-soluble salt with potassium chloride. The
precipitate, containing a small quantity of iridium, is then heated
with sodium carbonate in a crucible, when the mass decomposes,
giving metallic platinum and iridium oxide. If potassium chloride
has been employed, the residue after ignition is washed with water
and treated with aqua regia. The iridium oxide remains undissolved,
and the platinum easily passes into solution. Only cold and dilute
aqua regia must be used. The solution will then contain pure
platinic chloride, which forms the starting-point for the
preparation of all platinum compounds. Pure platinum for accurate
researches (for instance, for the unit of light, according to
Violle's method) may be obtained (Mylius and Foerster, 1892) by
Finkener's method, by dissolving the impure metal in aqua regia (it
should be evaporated to drive off the nitrogen compounds), and
adding NaCl so as to form a double sodium salt, which is purified
by crystallising with a small amount of caustic soda, washing the
crystals with a strong solution of NaCl, and then dissolving them
in a hot 1 p.c. solution of soda, repeating the above and
ultimately igniting the double salt, previously dried at 120°, in a
stream of hydrogen; platinum black and NaCl are then formed. The
three following are very sensitive tests (to thousandths of a per
cent.) for the presence of Ir, Ru, Rh, Pd (osmium is not usually
present in platinum which has once been purified, since it easily
volatilises with Cl_{2} and CO_{2}, and in the first treatment of
the crude platinum either passes off as OsO_{4} or remains
undissolved), Fe, Cu, Ag, and Pb: (1) the assay is alloyed with 10
parts of pure lead, the alloy treated with dilute nitric acid (to
remove the greater part of the Pb), and dissolved in aqua regia;
the residue will consist of Ir and Ru; the Pb is precipitated from
the nitric acid solution by sulphuric acid, whilst the remaining
platinum metals are reduced from the evaporated solution by formic
acid, and the resultant precipitate fused with KHSO_{4}; the Pd and
Rh are thus converted into soluble salts, and the former is then
precipitated by HgC_{2}N_{2}. (2) Iron may be detected by the usual
reagents, if the crude platinum be dissolved in aqua regia, and the
platinum metals precipitated from the solution by formic acid. (3)
If crude platinum (as foil or sponge) be heated in a mixture of
chlorine and carbonic oxide it volatilises (with a certain amount
of Ir, Pd, Fe, &c.) as PtCl_{2},2CO (Note 11), whilst the whole of
the Rh, Ag, and Cu it may contain remains behind. Among other
characteristic reactions for the platinum metals, we may mention:
(1) that rhodium is precipitated from the solution obtained after
fusion with KHSO_{4} (in which Pt does not dissolve) by NH_{3},
acetic and formic acids; (2) that dilute aqua regia dissolves
precipitated Pt, but not Rh; (3) that if the insoluble residue of
the platinum metals (Ir, Ru, Os) obtained, after treating with aqua
regia, be fused with a mixture of 1 part of KNO_{3} and 3 parts of
K_{2}CO_{3} (in a gold crucible), and then treated with water, it
gives a solution containing the Ru (and a portion of the Ir), but
which throws it all down when saturated with chlorine and boiled;
(4) that if iridium be fused with a mixture of KHO and KNO_{3}, it
gives a soluble potassium salt, IrK_{2}O_{4} (the solution is
blue), which, when saturated with chlorine, gives IrCl_{4}, which
is precipitated by NH_{4}Cl (the precipitate is black), forming a
double salt, leaving metallic Ir after ignition; (5) that rhodium
mixed with NaCl and ignited in a current of chlorine gives a
soluble double salt (from which sal-ammoniac separates Pt and Ir),
which gives (according to Jörgensen) a difficultly soluble
purpureo-salt (Chapter XXII., Note 35), Rh_{2}Cl_{3},5NH_{3}, when
treated with NH_{3}; in this form the Rh may be easily purified and
obtained in a metallic form by igniting in hydrogen; and (6) that
palladium, dissolved in aqua regia and dried (NH_{4}Cl throws down
any Pt), gives soluble PdCl_{2}, which forms an easily
crystallisable yellow salt, PdCl_{2}NH_{3}, with ammonia; this salt
(Wilm) may be easily purified by crystallisation, and gives
metallic Pd when ignited. These reactions illustrate the method of
separating the platinum metals from each other.
Metallic _platinum_ in a fused state has a specific gravity of 21; it is
grey, softer than iron but harder than copper, exceedingly ductile, and
therefore easily drawn into wire and rolled into thin sheets, and may be
hammered into crucibles and drawn into thin tubes, &c. In the state in
which it is obtained by the ignition of its compounds, it forms a spongy
mass, known as spongy platinum, or else as powder (platinum black).[6] In
either case it is dull grey, and is characterised, as we already know, by
the faculty of absorbing hydrogen and other gases. Platinum is not acted
on by hydrochloric, hydriodic, nitric, and sulphuric acids, or a mixture
of hydrofluoric and nitric acids. Aqua regia, and any liquid containing
chlorine or able to evolve chlorine or bromine, dissolves platinum.
Alkalis are decomposed by platinum at a red heat, owing to the faculty of
the platinum oxide, PtO_{2}, formed to combine with alkaline bases,
inasmuch as it has a feebly-developed acid character (_see_ Note 8).
Sulphur, phosphorus (the phosphide, PtP_{2}, is formed), arsenic and
silicon all act more or less rapidly on platinum, under the influence of
heat. Many of the metals form alloys with it. Even charcoal combines with
platinum when it is ignited with it, and therefore carbonaceous matter
cannot be subjected to prolonged and powerful ignition in platinum
vessels. Hence a platinum crucible soon becomes dull on the surface in a
smoky flame. Platinum also forms alloys with zinc, lead, tin, copper,
gold, and silver.[7] Although mercury does not directly dissolve
platinum, still it forms a solution or amalgam with spongy platinum in
the presence of sodium amalgam; a similar amalgam is also formed by the
action of sodium amalgam on a solution of platinum chloride, and is used
for physical experiments.
[6] We have already become acquainted with the effect of finely-divided
platinum on many gaseous substances. It is best seen in the
so-called _platinum black_, which is a coal-black powder left by
the action of sulphuric acid on the alloy of zinc and platinum, or
which is precipitated by metallic zinc from a dilute solution of
platinum. In any case, finely-divided platinum absorbs gases more
powerfully and rapidly the more finely divided and porous it is.
Sulphurous anhydride, hydrogen, alcohol, and many organic
substances in the presence of such platinum are easily oxidised by
the oxygen of the air, although they do not combine with it
directly. The absorption of oxygen is as much as several hundred
volumes per one volume of platinum, and the oxidising power of such
absorbed oxygen is taken advantage of not only in the laboratory
but even in manufacturing processes. Asbestos or charcoal, soaked
in a solution of platinic chloride and ignited, is very useful for
this purpose, because by this means it becomes coated with platinum
black. If 50 grams of PtCl_{4} be dissolved in 60 c.c. of water,
and 70 c.c. of a strong (40 p.c.) solution of formic aldehyde
added, the mixture cooled, and then a solution of 50 grams of NaHO
in 50 grams of water added, the platinum is precipitated. After
washing with water the precipitate passes into solution and forms a
black liquid containing _soluble colloidal platinum_ (Loew, 1890).
If the precipitated platinum be allowed to absorb oxygen on the
filter, the temperature rises 40°, and a very porous _platinum
black_ is obtained which vigorously facilitates oxidation.
[7] It is necessary to remark that platinum when alloyed with silver,
or as amalgam, is soluble in nitric acid, and in this respect it
differs from gold, so that it is possible, by alloying gold with
silver, and acting on the alloy with nitric acid, to recognise the
presence of platinum in the gold, because nitric acid does not act
on gold alloyed with silver.
There are _two kinds_ of _platinum compounds_, PtX_{4} and PtX_{2}. The
former are produced by an excess of halogen in the cold, and the latter
by the aid of heat or by the splitting up of the former. The
starting-point for the platinum compounds is _platinum tetrachloride_,
_platinic chloride_, PtCl_{4}, obtained by dissolving platinum in aqua
regia.[7 bis] The solution crystallises in the cold, in a desiccator, in
the form of reddish-brown deliquescent crystals which contain
hydrochloric acid, PtCl_{4},2HCl,6H_{2}O, and behave like a true acid
whose salts correspond to the formula R_{2}PtCl_{6}--ammonium
platinochloride, for example.[7 tri] The hydrochloric acid is liberated
from these crystals by gently heating or evaporating the solution to
dryness; or, better still, after treatment with silver nitrate a
reddish-brown mass remains behind, which dissolves in water, and forms a
yellowish-red solution which on cooling deposits crystals of the
composition PtCl_{4},8H_{2}O. The _tendency_ of PtCl_{4} _to combine_
with hydrochloric acid and water--that is, _to form higher crystalline
compounds_--is evident in the platinum compounds, and must be taken into
account in explaining the properties of platinum and the formation of
many other of its complex compounds. Dilute solutions of platinic
chloride are yellow, and are completely reduced by hydrogen, sulphurous
anhydride, and many reducing agents, which first convert the platinic
chloride into the lower compound platinous chloride, PtCl_{2}. That
faculty which reveals itself in platinum tetrachloride of combining with
water of crystallisation and hydrochloric acid is distinctly marked in
its property, with which we are already acquainted, of giving
precipitates with the salts of potassium, ammonium, rubidium, &c. In
general it _readily forms double salts_, R_{2}PtCl_{6} = PtCl_{4} + 2RCl,
where R is a univalent metal such as potassium or NH_{4}. Hence the
addition of a solution of potassium or ammonium chloride to a solution of
platinic chloride is followed by the formation of a yellow precipitate,
which is sparingly soluble in water and almost entirely insoluble in
alcohol and ether (platinic chloride is soluble in alcohol, potassium
iridiochloride, IrK_{3}Cl_{6}, _i.e._ a compound of IrCl_{3}, is soluble
in water but not in alcohol). It is especially remarkable in this case,
that the potassium compounds here, as in a number of other instances,
separate in an anhydrous form, whilst the sodium compounds, which are
soluble in water and alcohol, form red crystals containing water. The
composition Na_{2}PtCl_{6},6H_{2}O exactly corresponds with the
above-mentioned hydrochloric compound. The compounds with barium,
BaPtCl_{6},4H_{2}O, strontium, SrPtCl_{6},8H_{2}O, calcium, magnesium,
iron, manganese, and many other metals are all soluble in water.[8]
[7 bis] PtCl_{4} is also formed by the action of a mixture of HCl
vapour and air, and by the action of gaseous chlorine upon
platinum.
[7 tri] Pigeon (1891) obtained fine yellow crystals of
PtH_{2}Cl_{6},4H_{2}O by adding strong sulphuric acid to a strong
solution of PtH_{2}Cl_{6},6H_{2}O. If crystals of
H_{2}PtCl_{6},6H_{2}O be melted in vacuo (60°) in the presence of
anhydrous potash, a red-brown solid hydrate is obtained containing
less water and HCl, which parts with the remainder at 200°, leaving
anhydrous PtCl_{4}. The latter does not disengage chlorine before
220°, and is perfectly soluble in water.
[8] Nilson (1877), who investigated the platinochlorides of various
metals subsequently to Bonsdorff, Topsöe, Clève, Marignac, and
others, found that univalent and bivalent metals--such as hydrogen,
potassium, ammonium ... beryllium, calcium, barium--give compounds
of such a composition that there is always twice as much chlorine
in the platinic chloride as in the combined metallic chloride; for
example, K_{2}Cl_{2},PtCl_{4}; BeCl_{2},PtCl_{4},8H_{2}O, &c. Such
trivalent metals as aluminium, iron (ferric), chromium, didymium,
cerium (cerous) form compounds of the type RCl_{3}PtCl_{4}, in
which the amounts of chlorine are in the ratio 3:4. Only indium and
yttrium give salts of a different composition--namely,
2InCl_{3},5PtCl_{4},36H_{2}O and 4YCl_{3},5PtCl_{4},51H_{2}O. Such
quadrivalent metals as thorium, tin, zirconium give compounds of
the type RCl_{4},PtCl_{4}, in which the ratio of the chlorine is
1:1. In this manner the valency of a metal may, to a certain
extent, be judged from the composition of the double salts formed
with platinic chloride.
Platinic bromide, PtBr_{4}, and iodide, PtI_{4}, are analogous to
the tetrachloride, but the iodide is decomposed still more easily
than the chloride. If sulphuric acid be added to platinic chloride,
and the solution evaporated, it forms a black porous mass like
charcoal, which deliquesces in the air, and has the composition
Pt(SO_{4})_{2}. But this, the only oxygen salt of the type PtX_{4},
is exceedingly unstable. This is due to the fact that _platinum
oxide_, the oxide of the type PtO_{2}, has a feeble acid character.
This is shown in a number of instances. Thus if a strong solution
of platinic chloride treated with sodium carbonate be exposed to
the action of light or evaporated to dryness and then washed with
water, a sodium platinate, Pt_{3}Na_{2}O_{7},6H_{2}O, remains. The
composition of this salt, if we regard it in the same sense as we
did the salts of silicic, titanic, molybdic and other acids, will
be PtO(ONa)_{2},2PtO_{2},6H_{2}O--that is, the same type is
repeated as we saw in the crystalline compounds of platinum
tetrachloride with sodium chloride, or with hydrochloric
acid--namely, the type PtX_{4}8Y, where Y is the molecule
H_{2}O,HCl, &c. Similar compounds are also obtained with other
alkalis. They will be platinates of the alkalis in which the
platinic oxide, PtO_{2}, plays the part of an acid oxide. Rousseau
(1889) obtained different grades of combination BaOPtO_{2},
3(BaO)2PtO_{2}, &c., by igniting a mixture of PtCl_{4} and caustic
baryta. If such an alkaline compound of platinum be treated with
acetic acid, the alkali combines with the latter, and a _platinic
hydroxide_, Pt(OH)_{4}, remains as a brown mass, which loses water
and oxygen when ignited, and in so doing decomposes with a slight
explosion. When slightly ignited this hydroxide first loses water
and gives the very unstable oxide PtO_{2}. Platinic sulphide,
PtS_{2}, belongs to the same type; it is precipitated by the action
of sulphuretted hydrogen on a solution of platinum tetrachloride.
The moist precipitate is capable of attracting oxygen, and is then
converted into the sulphate above mentioned, which is soluble in
water. This absorption of oxygen and conversion into sulphate is
another illustration of the basic nature of PtO_{2}, so that it
clearly exhibits both basic and acid properties. The latter appear,
for instance, in the fact that platinic sulphide, PtS_{2}, gives
crystalline compounds with the alkali sulphides.
_Platinous chloride_, PtCl_{2}, is formed when hydrogen platinochloride,
PtH_{2}Cl_{6}, is ignited at 300°, or when potassium is heated at 230° in
a stream of chlorine. The undecomposed tetrachloride is extracted from
the residue by washing it with water, and a greenish-grey or brown
insoluble mass of the dichloride (sp. gr. 5·9) is then obtained. It is
soluble in hydrochloric acid, giving an acid solution of the composition
PtCl_{2},2HCl, corresponding with the type of double salts PtR_{2}Cl_{4}.
Although platinous chloride decomposes below 500°, still it is formed to
a small extent at higher temperatures. Troost and Hautefeuille, and
Seelheim observed that when platinum was strongly ignited in a stream of
chlorine, the metal, as it were, slowly volatilised and was deposited in
crystals; a volatile chloride, probably platinous chloride, was evidently
formed in this case, and decomposed subsequently to its formation,
depositing crystals of platinum.
The properties of platinum above-described are repeated more or
less distinctly, or sometimes with certain modifications, in the
above-mentioned associates and analogues of this metal. Thus although
palladium forms PdCl_{4}, this form passes into PdCl_{2} with extreme
ease.[9] Whilst rhodium and iridium in dissolving in aqua regia also form
RhCl_{4} and IrCl_{4}, but they pass into RhCl_{3} and IrCl_{3}[9 bis]
very easily when heated or when acted upon by substances capable of
taking up chlorine (even alkalis, which form bleaching salts). Among the
platinum metals, ruthenium and osmium have the most acid character, and
although they give RuCl_{4} and OsCl_{4} they are easily oxidised to
RuO_{4}, and OsO_{4} by the action of chlorine in the presence of water;
the latter are volatile and may be distilled with the water and
hydrochloric acid, from a solution containing other platinum metals.[9
tri] Thus with respect to the types of combination, all the platinum
metals, under certain circumstances, give compounds of the type
RX_{4}--for instance, RO_{2}, RCl_{4}, &c. But this is the highest form
for only platinum and palladium. The remaining platinum metals further,
_like iron, give acids_ of the type RO_{3} or hydrates, H_{2}RO_{4} =
RO_{2}(HO)_{2} (the type of sulphuric acid); but they, like ferric and
manganic acids, are chiefly known in the form of salts of the composition
K_{2}RO_{4} or K_{2}R_{2}O_{7} (like the dichromate). These salts are
obtained, like the manganates and ferrates, by fusing the oxides, or even
the metals themselves, with nitric, or, better still, with potassium
peroxide. They are soluble in water, are easily deoxidised and do not
yield the acid anhydrides under the action of acids, but break up, either
(like the ferrate) forming oxygen and a basic oxide (iridium and rhodium
react in this manner, as they do not give higher forms of oxidation), or
passing into a lower and higher form of oxidation--that is, reacting like
a manganate (or partly like nitrite or phosphite). Osmium and ruthenium
react according to the latter form, as they are capable of giving _higher
forms of oxidation_, OsO_{4} and RuO_{4}, and therefore their reactions
of decomposition may be essentially represented by the equation: 2OsO_{3}
= OsO_{2} + OsO_{4}.[10]
[9] In comparing the characteristics of the platinum metals, it must be
observed that palladium in its form of combination PdX_{2} gives
saline compounds of considerable stability. Amongst them _palladous
chloride_ is formed by the direct action of chlorine or aqua regia
(not in excess or in dilute solutions) on palladium. It forms a
brown solution, which gives a black insoluble precipitate of
_palladous iodide_, PdI_{2}, with solutions of iodides (in this
respect, as in many others, palladium resembles mercury in the
mercuric compounds HgX_{2}). With a solution of mercuric cyanide it
gives a yellowish white precipitate, palladous cyanide,
PdC_{2}N_{2}, which is soluble in potassium cyanide, and gives
other double salts, M_{2}PdC_{4}N_{4}.
That portion of the platinum ore which dissolves in aqua regia and
is precipitated by ammonium or potassium chloride does not contain
palladium. It remains in solution, because the palladic chloride,
PdCl_{4}, is decomposed and the palladous chloride formed is not
precipitated by ammonium chloride; the same holds good for all the
other lower chlorides of the platinum metals. Zinc (and iron)
separates out all the unprecipitated platinum metals (and also
copper, &c.) from the solution. The palladium is found in these
platinum residues precipitated by zinc. If this mixture of metals
be treated with aqua regia, all the palladium will pass into
solution as palladous chloride with some platinic chloride. By this
treatment the main portion of the iridium, rhodium, &c. remains
almost undissolved, the platinum is separated from the mixture of
palladous and platinic chlorides by a solution of ammonium
chloride, and the solution of palladium is precipitated by
potassium iodide or mercuric cyanide. Wilm (1881) showed that
palladium may be separated from an impure solution by saturating it
with ammonia; all the iron present is thus precipitated, and, after
filtering, the addition of hydrochloric acid to the filtrate gives
a yellow precipitate of an ammonio-palladium compound,
PdCl_{2},2NH_{3}, whilst nearly all the other metals remain in
solution. _Metallic palladium_ is obtained by igniting the
ammonio-compound or the cyanide, PdC_{2}N_{2}. It occurs native,
although rarely, and is a metal of a whiter colour than platinum,
sp. gr. 11·4, melts at about 1,500°; it is much more volatile than
platinum, partially oxidises on the surface when heated (Wilm
obtained spongy palladium by igniting PdCl_{2},2NH_{3}, and
observed that it gives PdO when ignited in oxygen, and that on
further ignition this oxide forms a mixture of Pd_{2}O and Pd), and
loses its absorbed oxygen on a further rise of temperature. It does
not blacken or tarnish (does not absorb sulphur) in the air at the
ordinary temperature, and is therefore better suited than silver
for astronomical and other instruments in which fine divisions have
to be engraved on a white metal, in order that the fine lines
should be clearly visible. The most remarkable property of
palladium, discovered by Graham, consists in its capacity for
_absorbing_ a large amount of _hydrogen_. Ignited palladium absorbs
as much as 940 volumes of hydrogen, or about 0·7 p.c. of its own
weight, which closely approaches to the formation of the compound
Pd_{3}H_{2}, and probably indicates the formation of _palladium
hydride_, Pd_{2}H. This absorption also takes place at the ordinary
temperature--for example, when palladium serves as an electrode at
which hydrogen is evolved. In absorbing the hydrogen, the palladium
does not change in appearance, and retains all its metallic
properties, only its volume increases by about 10 p.c.--that is,
the hydrogen pushes out and separates the atoms of the palladium
from each other, and is itself compressed to 1/900 of its volume.
This compression indicates a great force of chemical attraction,
and is accompanied by the evolution of heat (Chapter II., Note 38).
The absorption of 1 grm. of hydrogen by metallic palladium (Favre)
is accompanied by the evolution of 4·2 thousand calories (for Pt
20, for Na 13, for K 10 thousand units of heat). Troost showed that
the dissociation pressure of palladium hydride is inconsiderable at
the ordinary temperature, but reaches the atmospheric pressure at
about 140°. This subject was subsequently investigated by A. A.
Cracow of St. Petersburg (1894), who showed that at first the
absorption of hydrogen by the palladium proceeds like solution,
according to the law of Dalton and Henry, but that towards the end
it proceeds like a dissociation phenomenon in definite compounds;
this forms another link between the phenomenon of solution and of
the formation of definite atomic compounds. Cracow's observations
for a temperature 18°, showed that the electro-conductivity and
tension vary until a compound Pd_{2}H is reached, and namely, that
the tension _p_ rises with the volume _v_ of hydrogen absorbed,
according to the law of Dalton and Henry--for instance, for
_p_ = 2·1 3·2 5·5 7·7 mm.
_v_ = 14 20 34 47
The maximum tension at 18° is 9 mm. At a temperature of about 140°
(in the vapour of xylene) the maximum tension is about 760 mm., and
when _v_ = 10-50 vols. the tension (according to Cracow's
experiments) stands at 90-450 mm.--that is, increases in proportion
to the volume of hydrogen absorbed. But from the point of view of
chemical mechanics it is especially important to remark that
Moutier clearly showed, through palladium hydride, the similarity
of the phenomena which proceed in evaporation and dissociation,
which fact Henri Sainte-Claire Deville placed as a fundamental
proposition in the theory of dissociation. It is possible upon the
basis of the second law of the theory of heat, according to the law
of the variation of the tension _p_ of evaporation with the
temperature T (counted from -273°), to calculate the latent heat of
evaporation L (_see_ works on physics) because 424L = T(1/_d_ -
1/D)_dp_/_dt_, where _d_ and D are the weights of cubic measures of
the gas (vapour) and liquid. (Thus, for instance, for water, when
_t_ = 100°, T = 373, _d_ = 0·605, D = 960, _dp_/_dt_ = 0·027 m.,
13,596 = 367, L = 536, whence 424L = 227,264, and the second
portion of the equation 226,144, which is sufficiently near, within
the limits of experimental error, _see_ Chapter I., Note 11.) The
same equation is applicable to the dissociation of Na_{2}H and
K_{2}H--(Chapter XII., Note 42)--but it has only been verified in
this respect for Pd_{2}H, since Moutier, by calculating the amount
of heat L evolved, for _t_ = 20, according to the variation of the
tension (_dp_/_dt_) obtained 4·1 thousand calories, which is very
near the figure obtained experimentally by Favre (_see_ Chapter
XII., Note 44). The absorbed hydrogen is easily disengaged by
ignition or decreased pressure. The resultant compound does not
decompose at the ordinary temperature, but when exposed to air the
metal sometimes glows spontaneously, owing to the hydrogen burning
at the expense of the atmospheric oxygen. The hydrogen absorbed by
palladium acts towards many solutions as a reducing agent; in a
word, everything here points to the formation of a definite
compound and at the same time of a physically-compressed gas, and
forms one of the best examples of the bond existing between
chemical and physical processes, to which we have many times drawn
attention. It must be again remembered that the other metals of the
eighth group, even copper, are, like palladium and platinum, able
to combine with hydrogen. The permeability of iron and platinum
tubes to hydrogen is naturally due to the formation of similar
compounds, but palladium is the most permeable.
[9 bis] _Rhodium_ is generally separated, together with iridium, from
the residues left after the treatment of native platinum, because
the palladium is entirely separated from them, and the ruthenium is
present in them in very small traces, whilst the osmium at any rate
is easily separated, as we shall soon see. The mixture of rhodium
and iridium which is left undissolved in dilute aqua regia is
dissolved in chlorine water, or by the action of chlorine on a
mixture of the metals with sodium chloride. In either case both
metals pass into solution. They may be separated by many methods.
In either case (if the action be aided by heat) the rhodium is
obtained in the form of the chloride RhCl_{3}, and the iridium as
iridious chloride, IrCl_{3}. They both form double salts with
sodium chloride which are soluble in water, but the iridium salt is
also partially soluble in alcohol, whilst the rhodium salt is not.
A mixture of the chlorides, when treated with dilute aqua regia,
gives iridic chloride, IrCl_{4}, whilst the rhodium chloride,
RhCl_{3}, remains unaltered; ammonium chloride then precipitates
the iridium as ammonium iridiochloride, Ir(NH_{4})2Cl_{6}, and on
evaporating the rose-coloured filtrate the rhodium gives a
crystalline salt, Rh(NH_{4})_{3}Cl_{6}. Rhodium and its various
oxides are dissolved when fused with potassium hydrogen sulphate,
and give a soluble double sulphate (whilst iridium remains unacted
on); this fact is very characteristic for this metal, which offers
in its properties many points of resemblance with the iron metals.
When fused with potassium hydroxide and chlorate it is oxidised
like iridium, but it is not afterwards soluble in water, in which
respect it differs from ruthenium. This is taken advantage of for
separating rhodium, ruthenium, and iridium. In any case, rhodium
under ordinary conditions always gives salts of the type RX_{3},
and not of any other type; and not only halogen salts, but also
oxygen salts, are known in this type, which is rare among the
platinum metals. Rhodium chloride, RhCl_{3}, is known in an
insoluble anhydrous and also in a soluble form (like CrX_{3} or
salts of chromic oxides), in which it easily gives double salts,
compounds with water of crystallisation, and forms rose-coloured
solutions. In this form rhodium easily gives double salts of the
two types RhM_{3}Cl_{6} and RhM_{2}Cl_{3}--for example,
K_{5}RhCl_{6},3H_{2}O and K_{2}RhCl_{5},H_{2}O. Solutions of the
salts (at least, the ammonium salt) of the first kind give salts of
the second kind when they are boiled. If a strong solution of
potash be added to a red solution of rhodium chloride and boiled, a
black precipitate of the hydroxide Rh(OH)_{3} is formed; but if the
solution of potash is added little by little, it gives a yellow
precipitate containing more water. This yellow hydrate of rhodium
oxide gives a yellow solution when it is dissolved in acids, which
only becomes rose-coloured after being boiled. It is obvious a
change here takes place, like the transmutations of the salts of
chromic oxide. It is also a remarkable fact that the black
hydroxide, like many other oxidised compounds of the platinoid
metals, does not dissolve in the ordinary oxygen acids, whilst the
yellow hydroxide is easily soluble and gives yellow solutions,
which deposit imperfectly crystallised salts. Metallic rhodium is
easily obtained by igniting its oxygen and other compounds in
hydrogen, or by precipitation with zinc. It resembles platinum, and
has a sp. gr. of 12·1. At the ordinary temperature it decomposes
formic acid into hydrogen and carbonic anhydride, with development
of heat (Deville). With the alkali sulphites, the salts of rhodium
and iridium of the type RX_{3} give sparingly-soluble precipitates
of double sulphites of the composition R(SO_{3}Na)_{3},H_{2}O, by
means of which these metals may be separated from solution, and
also may be separated from each other, for a mixture of these salts
when treated with strong sulphuric acid gives a soluble iridium
sulphate and leaves a red insoluble double salt of rhodium and
sodium. It may be remarked that the oxides Ir_{2}O_{3} and
Rh_{2}O_{3} are comparatively stable and are easily formed, and
that they also form different double salts (for instance,
IrCl_{3},3KCl_{3}H_{2}O, RhCl_{3},2NH_{4}Cl_{4}H_{2}O,
RhCl_{3},3NH_{4}Cl1-1/2H_{2}O) and compounds like the cobaltia
compounds (for instance, luteo-salts RhX_{3},6NH_{3}, roseo-salts,
RhX_{3}H_{2}O_{5}NH_{3}, and purpureo-salts IrX_{3},5NH_{3}, &c.)
_Iridious oxide_, Ir_{2}O_{3}, is obtained by fusing iridious
chloride and its compounds with sodium carbonate, and treating the
mass with water. The oxide is then left as a black powder, which,
when strongly heated, is decomposed into iridium and oxygen; it is
easily reduced, and is insoluble in acids, which indicates the
feeble basic character of this oxide, in many respects resembling
such oxides as cobaltic oxide, ceric or lead dioxide, &c. It does
not dissolve when fused with potassium hydrogen sulphate. Rhodium
oxide, Rh_{2}O_{3}, is a far more energetic base. It dissolves when
fused with potassium hydrogen sulphate.
From what has been said respecting the separation of platinum and
rhodium it will be understood how the compounds of _iridium_, which
is the main associate of platinum, are obtained. In describing the
treatment of osmiridium we shall again have an opportunity of
learning the method of extraction of the compounds of this metal,
which has in recent times found a technical application in the form
of its oxide, Ir_{2}O_{3}; this is obtained from many of the
compounds of iridium by ignition with water, is easily reduced by
hydrogen, and is insoluble in acids. It is used in painting on
china, for giving a black colour. Iridium itself is more
difficultly fusible than platinum, and when fused it does not
decompose acids or even aqua regia; it is extremely hard, and is
not malleable; its sp. gr. is 22·4. In the form of powder it
dissolves in aqua regia, and is even partially oxidised when heated
in air, sets fire to hydrogen, and, in a word, closely resembles
platinum. Heated in an excess of chlorine it gives iridic chloride,
IrCl_{4}, but this loses chlorine at 50°; it is, however, more
stable in the form of double salts, which have a characteristic
_black_ colour--for instance, Ir(NH_{4})_{2}Cl_{6}--but they give
iridious chloride, IrCl_{3}, when treated with sulphuric acid.
[9 tri] We have yet to become acquainted with the two remaining
associates of platinum--ruthenium and osmium--whose most important
property is that they are oxidised even when heated in air, and
that they are able to give _volatile_ oxides of the form RuO_{4}
and OsO_{4}; these have a powerful odour (like iodine and nitrous
anhydride). Both these higher oxides are solids; they volatilise
with great ease at 100°; the former is yellow and the latter white.
They are known as _ruthenic_ and _osmic anhydrides_, although their
aqueous solutions (they both slowly dissolve in water) do not show
an acid reaction, and although they do not even expel carbonic
anhydride from potassium carbonate, do not give crystalline salts
with bases, and their alkaline solutions partially deposit them
again when boiled (an excess of water decomposes the salts). The
formulæ OsO_{4} and RuO_{4} correspond with the vapour density of
these oxides. Thus Deville found the vapour density of osmic
anhydride to be 128 (by the formula 127·5) referred to hydrogen.
Tennant and Vauquelin discovered this compound, and Berzelius,
Wöhler, Fritzsche, Struvé, Deville, Claus, Joly, and others helped
in its investigation; nevertheless there are still many questions
concerning it which remain unsolved. It should be observed that
RO_{4} is the highest known form for an oxygen compound, and RH_{4}
is the highest known form for a compound of hydrogen; whilst the
highest forms of acid hydrates contain SiH_{4}O_{4}, PH_{3}O_{4},
SH_{2}O_{4}, ClHO_{4}--all with four atoms of oxygen, and therefore
in this number there is apparently the limit for the simple forms
of combination of hydrogen and oxygen. In combination with
_several_ atoms of an element, or several elements, there may be
more than O_{4} or H_{4}, but a molecule never contains more than
four atoms of either O or H to one atom of another element. Thus
the simplest forms of combination of hydrogen and oxygen are
exhausted by the list RH_{4}, RH_{3}, RH_{2}, RH, RO, RO_{2},
RO_{3}, RO_{4}. The extreme members are RH_{4} and RO_{4}, and are
only met with for such elements as carbon, silicon, osmium,
ruthenium, which also give RCl_{4} with chlorine. In these extreme
forms, RH_{4} and RO_{4}, the compounds are the least stable
(compare SiH_{4}, PH_{3}, SH_{2}, ClH, or RuO_{4}, MoO_{3},
ZrO_{2}, SrO), and easily give up part, or even all, their oxygen
or hydrogen.
The primary source from which the compounds of ruthenium and osmium
are obtained is either _osmiridium_ (the osmium predominates, from
IrOs to IrOs_{4}, sp. gr. from 16 to 21), which occurs in platinum
ores (it is distinguished from the grains of platinum by its
crystalline structure, hardness, and insolubility in aqua regia),
or else those insoluble residues which are obtained, as we saw
above, after treating platinum with aqua regia. Osmium predominates
in these materials, which sometimes contain from 30 p.c. to 40 p.c.
of it, and rarely more than 4 p.c. to 5 p.c. of ruthenium. The
process for their treatment is as follows: they are first fused
with 6 parts of zinc, and the zinc is then extracted with dilute
hydrochloric acid. The osmiridium thus treated is, according to
Fritzsche and Struvé's method, then added to a fused mixture of
potassium hydroxide and chlorate in an iron crucible; the mass as
it begins to evolve oxygen acts on the metal, and the reaction
afterwards proceeds spontaneously. The dark product is treated with
water, and gives a solution of osmium and ruthenium in the form of
soluble salts, R_{2}OsO_{4} and R_{2}RuO_{4}, whilst the insoluble
residue contains a mixture of oxides of iridium (and some osmium,
rhodium, and ruthenium), and grains of metallic iridium still
unacted on. According to Frémy's method the lumps of osmiridium are
straightway heated to whiteness in a porcelain tube in a stream of
air or oxygen, when the very volatile osmic anhydride is obtained
directly, and is collected in a well-cooled receiver, whilst the
ruthenium gives a crystalline sublimate of the dioxide, RuO_{2},
which is, however, very difficultly volatile (it volatilises
together with osmic anhydride), and therefore remains in the cooler
portions of the tube; this method does not give volatile ruthenic
anhydride, and the iridium and other metals are not oxidised or
give non-volatile products. This method is simple, and at once
gives dry, pure osmic anhydride in the receiver, and ruthenium
dioxide in the sublimate. The air which passes through the tube
should be previously passed through sulphuric acid, not only in
order to dry it, but also to remove the organic and reducing dust.
The vapour of osmic anhydride must be powerfully cooled, and
ultimately passed over caustic potash. A third mode of treatment,
which is most frequently employed, was proposed by Wöhler, and
consists in slightly heating (in order that the sodium chloride
should not melt) an intimate mixture of osmiridium and common salt
in a stream of moist chlorine. The metals then form compounds with
chlorine and sodium chloride, whilst the osmium forms the chloride,
OsCl_{4}, which reacts with the moisture, and gives osmic
anhydride, which is condensed. The ruthenium in this, as in the
other processes, does not directly give ruthenic anhydride, but is
always extracted as the soluble ruthenium salt, K_{2}RuO_{4},
obtained by fusion with potassium hydroxide and chlorate or
nitrate. When the orange-coloured ruthenate, K_{2}RuO_{4}, is mixed
with acids, the liberated ruthenic acid immediately decomposes into
the volatile ruthenic anhydride and the insoluble ruthenic oxide:
2K_{2}RuO_{4} + 4HNO_{3} = RuO_{4} + RuO_{2},2H_{2}O + 4KNO_{3}.
When once one of the above compounds of ruthenium or osmium is
procured it is easy to obtain all the remaining compounds, and by
reduction (by metals, hydrogen, formic acid, &c.) the metals
themselves.
Osmic anhydride, OsO_{4}, is very easily deoxidised by many
methods. It blackens organic substances, owing to reduction, and is
therefore used in investigating vegetable and animal, and
especially nerve, preparations under the microscope. Although osmic
anhydride may be distilled in hydrogen, still complete reduction is
accomplished when a mixture of hydrogen and osmic anhydride is
slightly ignited (just before it inflames). If osmium be placed in
the flame it is oxidised, and gives vapours of osmic anhydride,
which become reduced, and the flame gives a brilliant light. Osmic
anhydride deflagrates like nitre on red-hot charcoal; zinc, and
even mercury and silver, reduce osmic anhydride from its aqueous
solutions into the lower oxides or metal; such reducing agents as
hydrogen sulphide, ferrous sulphate, or sulphurous anhydride,
alcohol, &c., act in the same manner with great ease.
The lower oxides of osmium, ruthenium, and of the other elements of
the platinum series are not volatile, and it is noteworthy that the
other elements behave differently. On comparing SO_{2}, SO_{3};
As_{2}O_{3}, As_{2}O_{5}; P_{2}O_{3}, P_{2}O_{5}; CO, CO_{2}, &c.,
we observe a converse phenomenon; the higher oxides are less
volatile than the lower. In the case of osmium all the oxides, with
the exception of the highest, are non-volatile, and it may
therefore be thought that this higher form is more simply
constituted than the lower. It is possible that osmic oxide,
OsO_{2}, stands in the same relation to the anhydride as C_{2}H_{4}
to CH_{4}--_i.e._ the lower oxide is perhaps Os_{2}O_{4}, or is
still more polymerised, which would explain why the lower oxides,
having a greater molecular weight, are less volatile than the
higher oxides, just as we saw in the case of the nitrogen oxides,
N_{2}O and NO.
_Ruthenium and osmium_, obtained by the ignition or reduction of
their compounds in the form of powder, have a density considerably
less than in the fused form, and differ in this condition in their
capacity for reaction; they are much more difficultly fused than
platinum and iridium, although ruthenium is more fusible than
osmium. Ruthenium in powder has a specific gravity of 8·5, the
fused metal of 12·2; osmium in powder has a specific gravity of
20·0, and when semi-fused--or, more strictly speaking,
agglomerated--in the oxyhydrogen flame, of 21·4, and fused 22·5.
The powder of slightly-heated osmium oxidises very easily in the
air, and when ignited burns like tinder, directly forming the
odoriferous osmic anhydride (hence its name, from the Greek word
signifying odour); ruthenium also oxidises when heated in air, but
with more difficulty, forming the oxide RuO_{2}. The oxides of the
types RO, R_{2}O_{3}, and RO_{2} (and their hydrates) obtained by
reduction from the higher oxides, and also from the chlorides, are
analogous to those given by the other platinum metals, in which
respect osmium and ruthenium closely resemble them. We may also
remark that ruthenium has been found in the platinum deposits of
Borneo in the form of _laurite_, Ru_{2}S_{3}, in grey octahedra of
sp. gr. 7·0.
For osmium, Moraht and Wischin (1893) obtained free osmic acid,
H_{2}OsO_{4}, by decomposing K_{2}OsO_{4} with water, and
precipitating with alcohol in a current of hydrogen (because in air
volatile OsO_{4} is formed); with H_{2}S, osmic acid gives
OsO_{3}(HS)_{2} at the ordinary temperature.
Debray and Joly showed that ruthenic anhydride, RuO_{4}, fuses at
25°, boils at 100°, and evolves oxygen when dissolved in potash,
forming the salt KRuO_{4} (not isomorphous with potassium
permanganate).
Joly (1891), who studied the ruthenium compounds in greater detail,
showed that the easily-formed KRuO_{4} gives RuKO_{4}RuO_{3} when
ignited, but it resembles KMnO_{4} in many respects. In general, Ru
has much in common with Mn. Joly (1889) also showed that if KNO_{3}
be added to a solution of RuCl_{3} containing HCl, the solution
becomes hot, and a salt, RuCl_{3}NO_{2}KCl, is formed, which enters
into double decomposition and is very stable. Moreover, if RuCl_{3}
be treated with an excess of nitric acid, it forms a salt,
RuCl_{3}NOH_{2}O, after being heated (to boiling) and the addition
of HCl. The vapour density of RuO_{4}, determined by Debray and
Joly, corresponds to that formula.
[10] Although palladium gives the same types of combination (with
chlorine) as platinum, its reduction to RX_{2} is incomparably
easier than that of platinic chloride, and in the case of iridium
it is also very easy. Iridic chloride, IrCl_{4}, acts as an
oxidising agent, readily parts with a fourth of its chlorine to a
number of substances, readily evolves chlorine when heated, and it
is only at low temperatures that chlorine and aqua regia convert
iridium into iridic chloride. In disengaging chlorine iridium more
often and easily gives the very stable iridious chloride, IrCl_{3}
(perhaps this substance is Ir_{2}Cl_{6} = IrCl_{2},IrCl_{4},
insoluble in water, but soluble in potassium chloride, because it
forms the double salt K_{3}IrCl_{6}), than the dichloride, IrCl_2.
This compound, corresponding to IrX_{2}, is very stable, and
corresponds with the _basic oxide_, Ir_{2}O_{3}, resembling the
oxides Fe_{2}O_{3}, Co_{2}O_{3}. To this form there correspond
ammoniacal compounds similar to those given by cobaltic oxide.
Although iridium also gives an acid in the form of the salt
K_{2}Ir_{2}O_{7}, it does not, like iron (and chromium), form the
corresponding chloride, IrCl_{6}. In general, in this as in the
other elements, it is impossible to predict the chlorine compounds
from those of oxygen. Just as there is no chloride SCl_{6}, but
only SCl_{2}, so also, although IrO_{3} exists, IrCl_{6} is
wanting, the only chloride being IrCl_{4}, and this is unstable,
like SCl_{2}, and easily parts with its chlorine. In this respect
rhodium is very much like iridium (as platinum is like palladium).
For RhCl_{4} decomposes with extreme ease, whilst rhodium
chloride, RhCl_{3}, is very stable, like many of the salts of the
type RhX_{3}, although like the platinum elements these salts are
easily reduced to metal by the action of heat and powerful
reagents. There is as close a resemblance between osmium and
ruthenium. Osmium when submitted to the action of dry chlorine
gives osmic chloride, OsCl_{4}, but the latter is converted by
water (as is osmium by moist chlorine) into osmic anhydride,
although the greater portion is then decomposed into Os(HO)_{4}
and 4HCl, like a chloranhydride of an acid. In general this acid
character is more developed in osmium than in platinum and
iridium. Having parted with chlorine, osmic chloride, OsCl_{4},
gives the unstable trichloride, OsCl_{3}, and the stable soluble
dichloride, OsCl_{2}, which corresponds with platinous chloride in
its properties and reactions. The relation of ruthenium to the
halogens is of the same nature.
Platinum and its analogues, like iron and its analogues, are able to form
complex and comparatively stable cyanogen and ammonia compounds,
corresponding with the ferrocyanides and the ammoniacal compounds of
cobalt, which we have already considered in the preceding chapter.
If platinous chloride, PtCl_{2} (insoluble in water), be added by degrees
to a solution of potassium cyanide, it is completely dissolved (like
silver chloride), and on evaporating the solution deposits rhombic prisms
of _potassium platinocyanide_, PtK_{2}(CN)_{4},3H_{2}O. This salt, like
all those corresponding with it, has a remarkable play of colours, due to
the phenomena of dichromism, and even polychromism, natural to all the
platinocyanides. Thus it is yellow and reflects a bright blue light. It
is easily soluble in water, effloresces in air, then turns red, and at
100° orange, when it loses all its water. The loss of water does not
destroy its stability--that is, it still remains unchanged, and its
stability is further shown by the fact that it is formed when potassium
ferrocyanide, K_{4}Fe(CN)_{6}, is heated with platinum black. This salt,
first obtained by Gmelin, shows a neutral reaction with litmus; it is
exceedingly stable under the action of air, like potassium ferrocyanide,
which it resembles in many respects. Thus the platinum in it cannot be
detected by reagents such as sulphuretted hydrogen; the potassium may be
replaced by other metals by the action of their salts, so that it
corresponds with a whole series of compounds, R_{2}Pt(CN)_{4}, and it is
stable, although the potassium cyanide and platinous salts, of which it
is composed, individually easily undergo change. When treated with
oxidising agents it passes, like the ferrocyanide, into a higher form of
combination of platinum. If salts of silver be added to its solution, it
gives a heavy white precipitate of silver platinocyanide,
PtAg_{2}(CN)_{4}, which when suspended in water and treated with
sulphuretted hydrogen, enters into double decomposition with the latter
and forms insoluble silver sulphide, Ag_{2}S, and soluble
_hydroplatinocyanic acid_, H_{2}Pt(CN)_{4}. If potassium platinocyanide
be mixed with an equivalent quantity of sulphuric acid, the
hydroplatinocyanic acid liberated may be extracted by a mixture of
alcohol and ether. The ethereal solution, when evaporated in a
desiccator, deposits bright red crystals of the composition
PtH_{2}(CN)_{4},5H_{2}O. This acid colours litmus paper, liberates
carbonic anhydride from sodium carbonate, and saturates alkalis, so that
it presents an analogy to hydroferrocyanic acid.[11]
[11] This acid character is explained by the influence of the platinum
on the hydrogen, and by the attachment of the cyanogen groups.
Thus cyanuric acid, H_{3}(CN)_{3}O_{3}, is an energetic acid
compared with cyanic acid, HCNO. And the formation of a compound
with five molecules of water of crystallisation,
(PtH_{2}(CN)_{4},5H_{2}O), confirms the opinion that platinum is
able to form compounds of still higher types than that expressed
in its saline compounds, and, moreover, the combination of
hydroplatinocyanic acid with water does not reach the limit of the
compounds which appears in PtCl_{4},2HCl,6H_{2}O.
A whole series of _platinocyanides_ of the common type
PtR_{2}(CN)_{4}_n_H_{2}O is obtained by means of double
decomposition with the potassium or hydrogen or silver salts. For
example, the salts of sodium and lithium contain, like the
potassium salt, three molecules of water. The sodium salt is
soluble in water and alcohol. The ammonium salt has the
composition Pt(NH_{4})_{2}(CN)_{4},2H_{2}O and gives crystals
which reflect blue and rose-coloured light. This ammonium salt
decomposes at 300°, with evolution of water and ammonium cyanide,
leaving a greenish _platinum dicyanide_, Pt(CN)_{2}, which is
insoluble in water and acid but dissolves in potassium cyanide,
hydrocyanic acid, and other cyanides. The same platinous cyanide
is obtained by the action of sulphuric acid on the potassium salts
in the form of a reddish-brown amorphous precipitate. The most
characteristic of the platinocyanides are those of the alkaline
earths. The magnesium salt PtMg(CN)_{4},7H_{2}O crystallises in
regular prisms, whose side faces are of a metallic green colour
and terminal planes dark blue. It shows a carmine-red colour along
the main axis, and dark red along the lateral axes; it easily
loses water, (2H_{2}O), at 40°, and then turns blue (it then
contains 5H_{2}O, which is frequently the case with the
platinocyanides). Its aqueous solution is colourless, and an
alcoholic solution deposits yellow crystals. The remainder of the
water is given off at 230°. It is obtained by saturating
platinocyanic acid with magnesia, or else by double decomposition
between the barium salt and magnesium sulphate. The strontium
salt, SrPt(CN)_{4},4H_{2}O crystallises in milk-white plates
having a violet and green iridescence. When it effloresces in a
desiccator, its surfaces have a violet and metallic green
iridescence. A colourless solution of the barium salt
PtBa(CN)_{4},4H_{2}O is obtained by saturating a solution of
hydroplatinocyanic acid with baryta, or by boiling the insoluble
copper platinocyanide in baryta water. It crystallises in
monoclinic prisms of a yellow colour, with blue and green
iridescence; it loses half its water at 100°, and the whole at
150°. The ethyl salt, Pt(C_{2}H_{5})_{2}(CN)_{4},2H_{2}0, is also
very characteristic; its crystals are isomorphous with those of
the potassium salt, and are obtained by passing hydrochloric acid
into an alcoholic solution of hydroplatinocyanic acid. The
facility with which they crystallise, the regularity of their
forms, and their remarkable play of colours, renders the
preparation of the platinocyanides one of the most attractive
lessons of the laboratory.
By the action of chlorine or dilute nitric acid, the
platinocyanides are converted into salts of the composition
PtM_{2}(CN)_{5}, which corresponds with Pt(CN)_{3},2KCN--that is,
they express the type of a non-existent form of oxidation of
platinum, PtX_{3} (_i.e._ oxide Pt_{2}O_{3}), just as potassium
ferricyanide (FeCy_{3},3KCy) corresponds with ferric oxide, and
the ferrocyanide corresponds with the ferrous oxide. The potassium
salt of this series contains PtK_{2}(CN)_{5},3H_{2}O, and forms
brown regular prisms with a metallic lustre, and is soluble in
water but insoluble in alcohol. Alkalis re-convert this compound
into the ordinary platinocyanide K_{2}Pt(CN)_{4}, taking up the
excess of cyanogen. It is remarkable that the salts of the type
PtM_{2}Cy_{5} contain the same amount of water of crystallisation
as those of the type PtM_{2}Cy_{4}. Thus the salts of potassium
and lithium contain three, and the salt of magnesium seven,
molecules of water, like the corresponding salts of the type of
platinous oxide. Moreover, neither platinum nor any of its
associates gives any cyanogen compound corresponding with the
oxide, _i.e._ having the composition PtK_{2}Cy_{6}, just as there
are no compounds higher than those which correspond to
RCy_{3}_n_MCy_{3} for cobalt or iron. This would appear to
indicate the absence of any such cyanides, and indeed, for no
element are there yet known any poly-cyanides containing more than
three equivalents of cyanogen for one equivalent of the element.
The phenomenon is perhaps connected with the faculty of cyanogen
of giving tricyanogen polymerides, such as cyanuric acid, solid
cyanogen chloride, &c. Under the action of an excess of chlorine,
a solution of PtK_{2}(CN)_{4} gives (besides PtK_{2}Cy_{5}) a
product PtK_{2}Cy_{4}Cl_{2}, which evidently contains the form
PtX_{4}, but at first the action of the chlorine (or the
electrolysis of, or addition of dilute peroxide of hydrogen to, a
solution of PtK_{2}Cy_{4}, acidulated with hydrochloric acid)
produces an easily soluble intermediate salt which crystallises in
thin copper-red needles (Wilm, Hadow, 1889). It only contains a
small amount of chlorine, and apparently corresponds to a compound
5PtK_{2}Cy_{4} + PtK_{2}Cy_{4}Cl_{2} + 24H_{2}O. Under the action
of an excess of ammonia both these chlorine products are converted
either completely or in part (according to Wilm ammonia does not
act upon PtK_{2}Cy_{4}) into PtCy_{2},2NH_{3}, _i.e._ a
platino-ammonia compound (_see_ further on). It is also necessary
to pay attention to the fact that ruthenium and osmium--which, as
we know, give higher forms of oxidation than platinum--are also
able to combine with a larger proportion of potassium cyanide (but
not of cyanogen) than platinum. Thus ruthenium forms a crystalline
_hydroruthenocyanic acid_, RuH_{4}(CN)_{6}, which is soluble in
water and alcohol, and corresponds with the salts M_{4}Ru(CN)_{6}.
There are exactly similar osmic compounds--for example,
K_{4}Os(CN)_{6},3H_{2}O. The latter is obtained in the form of
colourless, sparingly-soluble regular tablets on evaporating the
solution obtained from a fused mixture of potassium osmiochloride,
K_{2}OsCl_{6}, and potassium cyanide. These osmic and ruthenic
compounds fully correspond with potassium ferrocyanide,
K_{4}Fe(CN)_{6},3H_{2}O, not only in their composition but also in
their crystalline form and reactions, which again demonstrates the
close analogy between iron, ruthenium, and osmium, which we have
shown by giving these three elements a similar position (in the
eighth group) in the periodic system. For rhodium and iridium only
salts of the same type as the ferricyanides, M_{3}RCy_{6}, are
known, and for palladium only of the type M_{2}PdCy_{4}, which are
analogous to the platinum salts. In all these examples a
_constancy of the types_ of the double cyanides is apparent. In
the eighth group we have iron, cobalt, nickel, copper, and their
analogues ruthenium, rhodium, palladium, silver, and also osmium,
iridium, platinum, gold. The double cyanides of iron, ruthenium,
osmium have the type K_{4}R(CN)_{6}; of cobalt, rhodium, iridium,
the type K_{3}R(CN)_{6}; of nickel, palladium, platinum the type
K_{2}R(CN)_{4} and K_{2}R(CN)_{5}; and for copper, silver, gold
there are known KR(CN)_{2}, so that the presence of 4, 3, 2, and 1
atoms of potassium corresponds with the order of the elements in
the periodic system. Those types which we have seen in the
ferrocyanides and ferricyanides of iron repeat themselves in all
the platinoid metals, and this naturally leads to the conclusion
that the formation of similar so-called double salts is of exactly
the same nature as that of the ordinary salts. If, in expressing
the union of the elements in the oxygen salts, the existence of an
_aqueous residue_ (hydroxyl group) be admitted, in which the
hydrogen is replaced by a metal, we have then only to apply this
mode of expression to the double salts and the analogy will be
obvious, if only we remember that Cl_{2}, (CN)_{2}, SO_{4}, &c.,
are equivalent to O, as we see in RO, RCl_{2}, RSO_{4}, &c. They
all = X_{2}, and, therefore, in point of fact, wherever X (= Cl or
OH, &c.) can be placed, there (Cl_{2}H), (SO_{4}H), &c., can also
stand. And as Cl_{2}H = Cl + HCl and SO_{4}H = OH + SO_{3}, &c.,
it follows that molecules HCl or SO_{3}, or, in general, whole
molecules--for instance, NH_{3}, H_{2}O, salts, &c., can annex
themselves to a compound containing X. (This is an indirect
consequence of the law of substitution which explains the origin
of double salts, ammonia compounds, compounds with water of
crystallisation, &c., by one general method.) Thus the double salt
MgSO_{4},K_{2}SO_{4}, according to this reasoning, _may be_
considered as a substance of the same type as MgCl_{2}, namely, as
= Mg(SO_{4}K)_{2}, and the alums as derived from Al(OH)(SO_{4}),
namely, as Al(SO_{4}K)(SO_{4}). Without stopping to pursue this
digression further, we will apply these considerations to the type
of the ferrocyanides and ferricyanides and their platinum
analogues. Such a salt as K_{2}PtCy_{4} may accordingly be
regarded as Pt(Cy_{2}K)_{2}, like Pt(OH)_{2}; and such a salt as
PtK_{2}Cy_{5} as PtCy(Cy_{2}K)_{2}, the analogue of PtX(OH)_{2},
or AlX(OH)_{2}, and other compounds of the type RX_{3}. Potassium
ferricyanide and the analogous compounds of cobalt, iridium, and
rhodium, belong to the same type, with the same difference as
there is between RX(OH)_{2} and R(OH)_{3}, since FeK_{3}Cy_{6} =
Fe(Cy_{2}K)_{3}. Limiting myself to these considerations, which
may partially elucidate the nature of double salts, I will now
pass again to the complex saline compounds known for platinum.
(_A_) On mixing a solution of potassium thiocyanate with a
solution of potassium platinosochloride, K_{2}PtCl_{4}, they form
a double thiocyanate, PtK_{2}(CNS)_{4}, which is easily soluble in
water and alcohol, crystallises in red prisms, and gives an
orange-coloured solution, which precipitates salts of the heavy
metals. The action of sulphuric acid on the lead salt of the same
type gives the acid itself, PtH_{2}(SCN)_{4}, which corresponds
with these salts. The type of these compounds is evidently the
same as that of the cyanides.
(_B_) _Platinous chloride_, PtCl_{2}, which is insoluble in water,
forms _double salts with the metallic chlorides_. These double
chlorides are soluble in water, and capable of crystallising.
Hence when a hydrochloric acid solution of platinous chloride is
mixed with solutions of metallic salts and evaporated it forms
crystalline salts of a red or yellow colour. Thus, for example,
the potassium salt, PtK_{2}Cl_{4}, is red, and easily soluble in
water; the sodium salt is also soluble in alcohol; the barium
salt, PtBaCl_{4},3H_{2}O, is soluble in water, but the silver
salt, PtAg_{2}Cl_{4}, is insoluble in water, and may be used for
obtaining the remaining salts by means of double decomposition
with their chlorides.
(_C_) A remarkable example of the complex compounds of platinum
was observed by Schützenberger (1868). He showed that
finely-divided platinum in the presence of chlorine and carbonic
oxide at 250°-300° gives phosgene and a volatile compound
containing platinum. The same substance is formed by the action of
carbonic oxide on platinous chloride. It decomposes with an
explosion in contact with water. Carbon tetrachloride dissolves a
portion of this substance, and on evaporation gives crystals of
2PtCl_{2},3CO, whilst the compound PtCl_{2},2CO remains
undissolved. When fused and sublimed it gives yellow needles of
PtCl_{2},CO, and in the presence of an excess of carbonic oxide
PtCl_{2},2CO is formed. These compounds are fusible (the first at
250°, the second at 142°, and the third at 195°). In this case (as
in the double cyanides) combination takes place, because both
carbonic oxide and platinous chloride are unsaturated compounds
capable of further combination. The carbon tetrachloride solution
absorbs NH_{3} and gives PtCl_{2},CO,2NH_{3}, and
PtCl_{2},2CO,2NH_{3}, and these substances are analogous
(Foerster, Zeisel, Jörgensen) to similar compounds containing
complex amines (for instance, pyridine, C_{5}H_{5}N), instead of
NH_{3}, and ethylene, &c., instead of CO, so that here we have a
whole series of complex platino-compounds. The compound PtCl_{2}CO
dissolves in hydrochloric acid without change, and the solution
disengages all the carbonic oxide when KCN is added to it, which
shows that those forces which bind 2 molecules of KCN to PtCl_{2}
can also bind the molecule CO, or 2 molecules of CO. When the
hydrochloric acid solution of PtCl_{2}CO is mixed with a solution
of sodium acetate or acetic acid, it gives a precipitate of PtOCO,
_i.e._ the Cl_{2} is replaced by oxygen (probably because the
acetate is decomposed by water). This oxide, PtOCO, splits up into
Pt + CO_{2} at 350°. PtSCO is obtained by the action of
sulphuretted hydrogen upon PtCl_{2}CO. All this leads to the
conclusion that the group PtCO is able to assimilate X_{2} =
Cl_{2}, S, O, &c. (Mylius, Foerster, 1891). Pullinger (1891), by
igniting spongy platinum at 250°, first in a stream of chlorine,
and then in a stream of carbonic oxide, obtained (besides volatile
products) a non-volatile yellow substance which remained unchanged
in air and disengaged chlorine and phosgene gas when ignited; its
composition was PtCl_{6}(CO)_{2}, which apparently proves it to be
a compound of PtCl_{2} and 2COCl_{2}, as PtCl_{2} is able to
combine with oxychlorides, and forms somewhat stable compounds.
(_D_) The faculty of platinous chloride for forming stable
compounds with divers substances shows itself in the formation of
the compound PtCl_{2},PCl_{3} by the action of phosphorus
pentachloride at 250° on platinum powder (Pd reacts in a similar
manner, according to Fink, 1892). The product contains both
phosphorus pentachloride and platinum, whilst the presence of
PtCl_{2} is shown in the fact that the action of water produces
_chlorplatino-phosphorous acid_, PtCl_{2}P(OH)_{3}.
(_E_) After the cyanides, the _double salts_ of platinum _formed
by sulphurous acid_ are most distinguished for their stability and
characteristic properties. This is all the more instructive, as
sulphurous acid is only feebly energetic, and, moreover, in these,
as in all its compounds, it exhibits a dual reaction. The salts of
sulphurous acid, R_{2}SO_{3}, either react as salts of a feeble
bibasic acid, where the group SO_{3} presents itself as bivalent,
and consequently equal to X_{2}, or else they react after the
manner of salts of a monobasic acid containing the same residue,
RSO_{3}, as occurs in the salts of sulphuric acid. In sulphurous
acid this residue is combined with hydrogen, H(SO_{3}H), whilst in
sulphuric acid it is united with the aqueous residue (hydroxyl),
OH(SO_{3}H). These two forms of action of the sulphites appear in
their reactions with the platinum salts--that is to say, salts of
both kinds are formed, and they both correspond with the type
PtH_{2}X_{4}. The one series of salts contain PtH_{2}(SO_{3})_{2},
and their reactions are due to the bivalent residue of sulphurous
acid, which replaces X_{2}. The others, which have the composition
PtR_{2}(SO_{3}H)_{4}, contain sulphoxyl. The latter salts will
evidently react like acids; they are formed simultaneously with
the salts of the first kind, and pass into them. These salts are
obtained either by directly dissolving platinous oxide in water
containing sulphurous acid, or by passing sulphurous anhydride
into a solution of platinous chloride in hydrochloric acid. If a
solution of platinous chloride or platinous oxide in sulphurous
acid be saturated with sodium carbonate, it forms a white,
sparingly soluble precipitate containing
PtNa_{2}(SO_{3}Na)_{4},7H_{2}O. If this precipitate be dissolved
in a small quantity of hydrochloric acid and left to evaporate at
the ordinary temperature, it deposits a salt of the other type,
PtNa_{2}(SO_{3})_{2},H_{2}O, in the form of a yellow powder, which
is sparingly soluble in water. The potassium salt analogous to the
first salt, PtK_{2}(SO_{3}K)_{4},2H_{2}O, is precipitated by
passing sulphurous anhydride into a solution of potassium sulphite
in which platinous oxide is suspended. A similar salt is known for
ammonium, and with hydrochloric acid it gives a salt of the second
kind, Pt(NH_{4})_{2}(SO_{3})_{2},H_{2}O. If ammonio-chloride of
platinum be added to an aqueous solution of sulphurous anhydride,
it is first deoxidised, and chlorine is evolved, forming a salt of
the type PtX_{2}; a double decomposition then takes place with the
ammonium sulphite, and a salt of the composition
Pt(NH_{4})_{2}Cl_{3}(SO_{3}H) is formed (in a desiccator). The
acid character of this substance is explained by the fact that it
contains the elements SO_{3}H--sulphoxyl, with the hydrogen not
yet displaced by a metal. On saturating a solution of this acid
with potassium carbonate it gives orange-coloured crystals of a
potassium salt of the composition Pt(NH_{4})_{2}Cl_{3}(SO_{3}K).
Here it is evident that an equivalent of chlorine in
Pt(NH_{4})_{2})Cl_{4} is replaced by the univalent residue of
sulphurous acid. Among these salts, that of the composition
Pt(NH_{4})_{2})Cl_{2}(SO_{3}H)_{2},H_{2}O is very readily formed,
and crystallises in well-formed colourless crystals; it is
obtained by dissolving ammonium platinosochloride,
Pt(NH_{4})_{2}Cl_{4}, in an aqueous solution of sulphurous acid.
The difficulty with which sulphurous anhydride and platinum are
separated from these salts indicates the same basic character in
these compounds as is seen in the double cyanides of platinum. In
their passage into a complex salt, the metal platinum and the
group SO_{2} modify their relations (compared with those of
PtX_{2} or SO_{2}X_{2}), just as the chlorine in the salts KClO,
KClO_{3}, and KClO_{4} is modified in its relations as compared
with hydrochloric acid or potassium chloride.
(_F_) No less characteristic are the _platinonitrites_ formed by
platinous oxide. They correspond with nitrous acid, whose salts,
RNO_{2}, contain the univalent radicle, NO_{2}, which is capable
of replacing chlorine, and therefore the salts of this kind should
form a common type PtR_{2}(NO_{2})_{4}, and such a salt of
potassium has actually been obtained by mixing a solution of
potassium platinosochloride with a solution of potassium nitrite,
when the liquid becomes colourless, especially if it be heated,
which indicates the change in the chemical distribution of the
elements. As the liquid decolorises it gradually deposits
sparingly soluble, colourless prisms of the potassium salt
K_{2}Pt(NO_{2})_{4}, which does not contain any water. With silver
nitrate a solution of this salt gives a precipitate of silver
platinonitrite, PtAg_{2}(NO_{2})_{4}. The silver of this salt may
be replaced by other metals by means of double decomposition with
metallic chlorides. The sparingly soluble barium salt, when
treated with an equivalent quantity of sulphuric acid, gives a
soluble acid, which separates, under the receiver of an air-pump,
in red crystals; this acid has the composition
PtH_{2}(NO_{2})_{4}. To the potassium salt, K_{2}Pt(NO_{2})_{4},
there correspond (Vèzes, 1892) K_{2}Pt(NO_{2})_{4}Br_{2} and
K_{2}Pt(NO_{2})_{4}Cl_{2} and other compounds of the same type
K_{2}PtX_{6}, where X is partly replaced by Cl or Br and partly by
(NO_{2}), showing a transition towards the type of the double
salts like the platino-ammoniacal salts. (The corresponding double
sodium nitrite salt of cobalt is soluble in water, while the
K,NH_{4} and many other salts are insoluble in water, as I was
informed by Prof. K. Winkler in 1894).
In all the preceding complex compounds of Pt we see a common type
PtX_{2},2MX (_i.e._ of double salts corresponding to PtO) or
PtM_{2}X_{4} = Pt(MX_{2})_{2}, corresponding to Pt(HO)_{2} with
the replacement of O by its equivalent X_{2}. Two other facts must
also be noted. In the first place these X's generally correspond
to elements (like chlorine) or groups (like CN, NO_{2}, SO_{3},
&c.), which are capable of further combination. In the second
place all the compounds of the type PtM_{2}X_{4} are capable of
combining with chlorine or similar elements, and thus passing into
compounds of the types PtX_{3} or PtX_{4}.
Ammonia, like potassium cyanide, has the faculty of easily reacting
with platinum dichloride, forming compounds similar to the platinocyanide
and cobaltia compounds, which are comparatively stable. But as ammonia
does not contain any hydrogen easily replaceable by metals, and as
ammonia itself is able to combine with acids, the PtX_{2} plays, as it
were, the part of an acid with reference to the ammonia. Owing to the
influence of the ammonia, the X_{2} in the resultant compound will
represent the same character as it has in ammoniacal salts; consequently,
the ammoniacal compounds produced from PtX_{2} will be salts in which X
will be replaceable by various other haloids, just as the metal is
replaced in the cyanogen salts; such is the nature of the
_platino-ammonium compounds_. PtX_{2} forms compounds with 2NH_{3} and
with 4NH_{3}, and so also PtX_{4} gives (not directly from PtX_{4} and
ammonia, but from the compounds of PtX_{2} by the action of chlorine,
&c.) similar compounds with 2NH_{3} and with 4NH_{3}.[12]
[12] The platinum salt and ammonia, when once combined together, are no
longer subject to their ordinary reactions but form compounds
which are comparatively very stable. The question at once suggests
itself to all who are acquainted with these phenomena, as to what
is the relation of the elements contained in these compounds. The
first explanation is that these compounds are salts of ammonium in
which the hydrogen is partially replaced by platinum. This is the
view, with certain shades of difference, held by many respecting
the platino-ammonium compounds. They were regarded in this light
by Gerhardt, Schiff, Kolbe, Weltzien, and many others. If we
suppose the hydrogen in 2NH_{4}X to be replaced by bivalent
platinum (as in the salts PtX_{2}), we shall obtain
NH_{3} X
Pt
NH_{3} X
--that is, the compound PtX_{2},2NH_{3}. The compound with 4NH_{3}
will then be represented by a further substitution of the hydrogen
in ammonia by ammonium itself--_i.e._ as NH_{2}(NH_{4}X)_{2}Pt or
PtX_{2},4NH_{3}. A modification of this view is found in that
representation of compounds of this kind which is based on
atomicity. As platinum in PtX_{2} is bivalent, has two affinities,
and ammonia, NH_{3}, is also bivalent, because nitrogen is
quinquivalent and is here only combined with H_{3}, it is evident
what bonds should be represented in PtX_{2},2NH_{3} and in
PtX_{2},4NH_{3}. In the former, Pt(NH_{3}Cl)_{2}, the nitrogen of
each atom of ammonia is united by three affinities with H_{3}, by
one with platinum, and by the fifth with chlorine. The other
compound is Pt(NH_{3}.NH_{3}Cl)_{2}--that is, the N is united by
one affinity with the other N, whilst the remaining bonds are the
same as in the first salt. It is evident that this union or chain
of ammonias has no obvious limit, and the most essential fault of
such a mode of representation is that it does not indicate at all
what number of ammonias are capable of being retained by platinum.
Moreover, it is hardly possible to admit the bond between nitrogen
and platinum in such stable compounds, for these kinds of
affinities are, at all events, feeble, and cannot lead to
stability, but would rather indicate explosive and
easily-decomposed compounds. Moreover, it is not clear why this
platinum, which is capable of giving PtX_{4}, does not act with
its remaining affinities when the addition of ammonia to PtX_{2}
takes place. These, and certain other considerations which
indicate the imperfection of this representation of the structure
of the platino-ammonium salts, cause many chemists to incline more
to the representations of Berzelius, Claus, Gibbs, and others, who
suppose that NH_{3} is able to combine with substances, to adjoin
itself or pair itself with them (this kind of combination is
called 'Paarung') without altering the fundamental capacity of a
substance for further combinations. Thus, in PtX_{2},2NH_{3}, the
ammonia is the associate of PtX_{2}, as is expressed by the
formula N_{2}H_{6}PtX_{2}. Without enlarging on the exposition of
the details of this doctrine, we will only mention that it, like
the first, does not render it possible to foresee a limit to the
compounds with ammonia; it isolates compounds of this kind into a
special and artificial class; does not show the connection between
compounds of this and of other kinds, and therefore it essentially
only expresses the fact of the combination with ammonia and the
modification in its ordinary reactions. For these reasons we do
not hold to either of these proposed representations of the
ammonio-platinum compounds, but regard them from the point of view
cited above with reference to double salts and water of
crystallisation--that is, we embrace all these compounds under the
representation of compounds of complex types. The type of the
compound PtX_{2},2NH_{3} is far more probably the same as that of
PtX_{2},2Z--_i.e._ as PtX_{4}, or, still more accurately and
truly, it is a compound of the same type as PtX_{2},2KX or
PtX_{2},2H_{2}O, &c. Although the platinum first entered into
PtK_{2}X_{4} as the type PtX_{2}, yet its character has changed in
the same manner as the character of sulphur changes when from
SO_{2} the compound SO_{2}(OH)_{2} is obtained, or when KClO_{4},
the higher form, is obtained from KCl. For us as yet there is no
question as to _what_ affinities hold X_{2} and what hold 2NH_{3},
because this is a question which arises from the supposition of
the existence of different affinities in the atoms, which there is
no reason for taking as a common phenomenon. It seems to me that
it is most important _as a commencement_ to render clear the
analogy in the formation of various complex compounds, and it is
this analogy of the ammonia compounds with those of water of
crystallisation and double salts that forms the main object of the
primary generalisation. We recognise in platinum, at all events,
not only the four affinities expressed in the compound PtCl_{4},
but a much larger number of them, if only the _summation of
affinities_ is actually possible. Thus, in sulphur we recognise
not two but a much greater number of affinities; it is clear that
at least six affinities can act. So also among the analogues of
platinum: osmic anhydride, OsO_{4}, Ni(CO)_{4}, PtH_{2}Cl_{6}, &c.
indicate the existence of at least eight affinities; whilst, in
chlorine, judging from the compound KClO_{4} = ClO_{3}(OK) =
ClX_{7}, we must recognise at least seven affinities, instead of
the one which is accepted. The latter mode of calculating
affinities is a tribute to that period of the development of
science when only the simplest hydrogen compounds were considered,
and when all complex compounds were entirely neglected (they were
placed under the class of molecular compounds). This is
insufficient for the present state of knowledge, because we find
that, in complex compounds as in the most simple, the same
constant types or modes of equilibrium are repeated, and the
character of certain elements is greatly modified in the passage
from the most simple into very complex compounds.
Judging from the most complex platino-ammonium compounds
PtCl_{4},4NH_{3}, we should admit the possibility of the formation
of compounds of the type PtX_{4}Y_{4}, where Y_{4} = 4X_{2} =
4NH_{3}, and this shows that those forces which form such a
characteristic series of double platinocyanides
PtK_{2}(CN)_{4},3H_{2}O, probably also determine the formation of
the higher ammonia derivatives, as is seen on comparing--
PtCl_{2} NH_{3} Cl_{2} 3NH_{3}
Pt(CN)_{2} KCN KCN 3H_{2}O.
Moreover, it is obviously much more natural to ascribe the faculty
for combination with _n_Y to the whole of the acting
elements--that is, to PtX_{2} or PtX_{4}, and not to platinum
alone. Naturally such compounds are not produced with any Y. With
certain X's there only combine certain Y's. The best known and
most frequently-formed compounds of this kind are those with
water--that is, compounds with water of crystallisation. Compounds
with salts are double salts; also we know that similar compounds
are also frequently formed by means of ammonia. Salts of zinc,
ZnX_{2}, copper, CuX_{2}, silver, AgX, and many others give
similar compounds, but these and many other _ammonio-metallic_
saline compounds are unstable, and readily part with their
combined ammonia, and it is only in the elements of the platinum
group and in the group of the analogues of iron, that we observe
the faculty to form stable ammonio-metallic compounds. It must be
remembered that the metals of the platinum and iron groups are
able to form several high grades of oxidation which have an acid
character, and consequently in the lower degrees of combination
there yet remain affinities capable of retaining other elements,
and they probably retain ammonia, and hold it the more stably,
because all the properties of the platinum compounds are rather
acid than basic--that is, PtX_{n} recalls rather HX or SnX_{n} or
CX_{n} than KX, CaX_{2}, BaX_{2}, &c., and ammonia naturally will
rather combine with an acid than with a basic substance. Further,
a dependence, or certain connection of the forms of oxidation with
the ammonia compounds, is seen on comparing the following
compounds:
PdCl_{2},2NH_{3},H_{2}O PdCl_{2},4NH_{3},H_{2}O
PtCl_{2},2NH_{3} PtCl_{4},4NH_{3}
RhCl_{3},5NH_{3} RuCl_{2},4NH_{3},3H_{2}O
IrCl_{3},5NH_{3} OsCl_{2},4NH_{3},2H_{2}O
We know that platinum and palladium give compounds of lower types
than iridium and rhodium, whilst ruthenium and osmium give the
highest forms of oxidation; this shows itself in this case also.
We have purposely cited the same compounds with 4NH_{3} for osmium
and ruthenium as we have for platinum and palladium, and it is
then seen that Ru and Os are capable of retaining 2H_{2}O and
3H_{2}O, besides Cl_{2} and NH_{3}, which the compounds of
platinum and palladium are unable to do. The same ideas which were
developed in Note 35, Chapter XXII. respecting the cobaltia
compounds are perfectly applicable to the present case, _i.e._ to
the _platinia_ compounds or ammonia compounds of the platinum
metals, among which Rh and Ir give compounds which are perfectly
analogous to the cobaltia compounds.
Iridium and rhodium, which easily give compounds of the type
RX_{3}, give compounds (Claus) of the type IrX_{3},5NH_{3}, of a
rose colour, and RhX_{3},5NH_{3}, of a yellow colour. Jörgensen,
in his researches on these compounds, showed their entire analogy
with the cobalt compounds, as was to be expected from the periodic
system.
If ammonia acts on a boiling solution of platinous chloride in
hydrochloric acid, it produces the green _salt of Magnus_ (1829),
PtCl_{2},2NH_{3}, insoluble in water and hydrochloric acid. But, judging
by its reactions, this salt has twice this formula. Thus, Gros (1837), on
boiling Magnus's salt with nitric acid, observed that half the chlorine
was replaced by the residue of nitric acid and half the
platinum was disengaged: 2PtCl_{2}(NH_{3})_{2} + 2HNO_{3} =
PtCl_{2}(NO_{3})_{2}(NH_{3})_{4} + 2PtCl_{2}. The Gros's salt thus
obtained, PtCl_{2}(NO_{3})_{2}4NH_{3} (if Magnus's salt belongs to the
type PtX_{2}, then Gros's salt belongs to the type PtX_{4}), is soluble
in water, and the elements of nitric acid, but not the chlorine,
contained in it are capable of easily submitting themselves to double
saline decomposition. Thus silver nitrate does not enter into double
decomposition with the chlorine of Gros's salt. Most instructive was the
circumstance that Gros, by acting on his salt with hydrochloric acid,
succeeded in substituting the residue of nitric acid in it by chlorine,
and the chlorine thus introduced, easily reacted with silver nitrate.
Thus it appeared that Gros's salt contained two varieties of
chlorine--one which reacts readily, and the other which reacts with
difficulty. The composition of Gros's first salt is
PtCl_{2}(NH_{3})_{4}(NO_{3})_{2}; it may be converted into
PtCl_{2}(NH_{3})_{4}(SO_{4}), and in general into
PtCl_{2}(NH_{3})_{4}X_{2}.[13]
[13] Subsequently, a whole series of such compounds was obtained with
various elements in the place of the (non-reacting) chlorine, and
nevertheless they, like the chlorine, reacted with difficulty,
whilst the second portion of the X's introduced into such salts
easily underwent reaction. This formed the most important reason
for the interest which the study of the composition and structure
of the platino-ammonium salts subsequently presented to many
chemists, such as Reiset, Blomstrand, Peyrone, Raeffski, Gerhardt,
Buckton, Clève, Thomsen, Jörgensen, Kournakoff, Verner, and
others. The salts PtX_{4},2NH_{3}, discovered by Gerhardt, also
exhibited several different properties in the two pairs of X's. In
the remaining platino-ammonium salts all the X's appear to react
alike.
The quality of the X's, retainable in the platino-ammonium salts,
may be considerably modified, and they may frequently be wholly or
partially replaced by hydroxyl. For example, the action of ammonia
on the nitrate of Gerhardt's base, Pt(NO_{3})_{4},2NH_{3}, in a
boiling solution, gradually produces a yellow crystalline
precipitate which is nothing else than a _basic hydrate_ or
_alkali_, Pt(OH)_{4},2NH_{3}. It is sparingly soluble in water,
but gives directly soluble salts PtX_{4},2NH_{3} with acids. The
stability of this hydroxide is such that potash does not expel
ammonia from it, even on boiling, and it does not change below
130°. Similar properties are shown by the hydroxide
Pt(OH)_{2},2NH_{3} and the oxide PtO,2NH_{3} of Reiset's second
base. But the hydroxides of the compounds containing 4NH_{3} are
particularly remarkable. The presence of ammonia renders them
soluble and energetic. The brevity of this work does not permit
us, however, to mention many interesting particulars in connection
with this subject.
The salt of Magnus when boiled with a solution of ammonia gives the
salt (of Reiset's first base) PtCl_{2}(NH_{3})_{4}, and this, when
treated with bromine, forms the salt PtCl_{2}Br_{2}(NH_{3})_{4}, which
has the same composition and reactions as Gros's salt. To Reiset's salts
there corresponds a soluble, colourless, crystalline _hydroxide_,
Pt(OH)_{2}(NH_{3})_{4}, having the properties of a powerful and very
energetic _alkali_; it attracts carbonic anhydride from the atmosphere,
precipitates metallic salts like potash, saturates active acids, even
sulphuric, forming colourless (with nitric, carbonic, and hydrochloric
acids), or yellow (with sulphuric acid), salts of the type
PtX_{2}(NH_{3})_{4}.[14] The comparative stability (for instance, as
compared with AgCl and NH_{3}) of such compounds, and the existence of
many other compounds analogous to them, endows them with a particular
chemical interest. Thus Kournakoff (1889) obtained a series of
corresponding compounds containing thiocarbamide, CSN_{2}H_{4}, in the
place of ammonia, PtCl_{2},4CSN_{2}H_{4}, and others corresponding with
Reiset's salts. Hydroxylamine, and other substances corresponding with
ammonia, also give similar compounds. The common properties and
composition of such compounds show their entire analogy to the cobaltia
compounds (especially for ruthenium and iridium) and correspond to the
fact that both the platinum metals and cobalt occur in the same, eighth,
group.
[14] Hydroxides are known corresponding with Gros's salts, which
contain one hydroxyl group in the place of that chlorine or haloid
which in Gros's salts reacts with difficulty, and these hydroxides
do not at once show the properties of alkalis, just as the
chlorine which stands in the same place does not react distinctly;
but still, after the prolonged action of acids, this hydroxyl
group is also replaced by acids. Thus, for example, the action of
nitric acid on Pt(NO_{3})_{2}Cl_{2},4NH_{3} causes the non-active
chlorine to react, but in the product all the chlorine is not
replaced by NO_{3}, but only half, and the other half is replaced
by the hydroxyl group: Pt(NO_{3})_{2}Cl_{2},4NH_{3} + HNO_{3} +
H_{2}O = Pt(NO_{3})_{3}(OH),4NH_{3} + 2HCl; and this is
particularly characteristic, because here the hydroxyl group has
not reacted with the acid--an evident sign of the non-alkaline
character of this residue. I think it may be well to call
attention to the fact that the composition of the
ammonio-metallosalts very often exhibits a correspondence between
the amount of X's and the amount of NH_{3}, of such a nature that
we find they contain either XNH_{3} or the grouping X_{2}NH_{3};
for example, Pt(XNH_{3})_{2} and Pt(X_{2}NH_{3})_{2},
Co(X_{2}NH_{3})_{3}, Pt(XNH_{3})_{4}, &c. Judging from this, the
view of the constitution of the double cyanides of platinum given
in Note 11 finds some confirmation here, but, in my opinion, all
questions respecting the composition (and structure) of the
ammoniacal, double, complex, and crystallisation compounds stand
connected with the solution of questions respecting the formation
of compounds of various degrees of stability, among which a theory
of solutions must be included, and therefore I think that the time
has not yet come for a complete generalisation of the data which
exist for these compounds; and here I again refer the reader to
Prof. Kournakoff's work cited in Chapter XXII., Note 35. However,
we may add a few individual remarks concerning the platinia
compounds.
To the common properties of the platino-ammonium salts, we must
add not only their _stability_ (feeble acids and alkalis do not
decompose them, the ammonia is not evolved by heating, &c.), but
also the fact that the ordinary reactions of platinum are
concealed in them to as great an extent as those of iron in the
ferricyanides. Thus neither alkalis nor hydrogen sulphide will
separate the platinum from them. For example, sulphuretted
hydrogen in acting on Gros's salts gives sulphur, removes half the
chlorine by means of its hydrogen, and forms salts of Reiset's
first base. This may be understood or explained by considering the
platinum in the molecule as covered, walled up by the ammonia, or
situated in the centre of the molecule, and therefore inaccessible
to reagents. On this assumption, however, we should expect to find
clearly-expressed ammoniacal properties, and this is not the case.
Thus ammonia is easily decomposed by chlorine, whilst in acting on
the platino-ammonium salts containing PtX_{2} and 2NH_{3} or
4NH_{3}, chlorine combines and does not destroy the ammonia; it
converts Reiset's salts into those of Gros and Gerhardt. Thus from
PtX_{2},2NH_{3} there is formed PtX_{2}Cl_{2},2NH_{3}, and from
PtX_{2},4NH_{3} the salt of Gros's base PtX_{2}Cl_{2},4NH_{3}.
This shows that the amount of chlorine which combines is not
dependent on the amount of ammonia present, but is due to the
basic properties of platinum. Owing to this some chemists suppose
the ammonia to be inactive or passive in certain compounds. It
appears to me that these relations, these modifications, in the
usual properties of ammonia and platinum are explained directly by
their mutual combination. Sulphur, in sulphurous anhydride,
SO_{2}, and hydrogen sulphide, SH_{2}, is naturally one and the
same, but if we only knew of it in the form of hydrogen sulphide,
then, having obtained it in the form of sulphurous anhydride, we
should consider its properties as hidden. The oxygen in magnesia,
MgO, and in nitric peroxide, NO_{2}, is so different that there is
no resemblance. Arsenic no longer reacts in its compounds with
hydrogen as it reacts in its compounds with chlorine, and in their
compounds with nitrogen all metals modify both their reactions and
their physical properties. We are accustomed to judge the metals
by their saline compounds with haloid groups, and ammonia by its
compounds with acid substances, and here, in the
platino-compounds, if we assume the platinum to be bound to the
entire mass of the ammonia--to its hydrogen and nitrogen--we shall
understand that both the platinum and ammonia modify their
characters. Far more complicated is the question why a portion of
the chlorine (and other haloid simple and complex groups) in
Gros's salts acts in a different manner from the other portion,
and why only half of it acts in the usual way. But this also is
not an exclusive case. The chlorine in potassium chlorate or in
carbon tetrachloride does not react with the same ease with metals
as the chlorine in the salts corresponding with hydrochloric acid.
In this case it is united to oxygen and carbon, whilst in the
platino-ammonium compounds it is united partly to platinum and
partly to the platino-ammonium group. Many chemists, moreover,
suppose that a part of the chlorine is united directly to the
platinum and the other part to the nitrogen of the ammonia, and
thus explain the difference of the reactions; but chlorine united
to platinum reacts as well with a silver salt as the chlorine of
ammonium chloride, NH_{4}Cl, or nitrosyl chloride, NOCl, although
there is no doubt that in this case there is a union between the
chlorine and nitrogen. Hence it is necessary to explain the
absence of a facile reactive capacity in a portion of the chlorine
by the conjoint influence of the platinum and ammonia on it,
whilst the other portion may be admitted as being under the
influence of the platinum only, and therefore as reacting as in
other salts. By admitting a certain kind of stable union in the
platino-ammonium grouping, it is possible to imagine that the
chlorine does not react with its customary facility, because
access to a portion of the atoms of chlorine in this complex
grouping is difficult, and the chlorine union is not the same as
we usually meet in the saline compounds of chlorine. These are the
grounds on which we, in refuting the now accepted explanations of
the reactions and formation of the platino-compounds, pronounce
the following opinion as to their structure.
In characterising the platino-ammonium compounds, it is necessary
to bear in mind that compounds which already contain PtX_{4} do
not combine directly with NH_{3}, and that such compounds as
PtX_{4},4NH_{3} only proceed from PtX_{2}, and therefore it is
natural to conclude that those affinities and forces which cause
PtX_{2} to combine with X_{2} also cause it to combine with
2NH_{3}. And having the compound PtX_{2},2NH_{3}, and supposing
that in subsequently combining with Cl_{2} it reacts with those
affinities which produce the compounds of platinic chloride,
PtCl_{4}, with water, potassium chloride, potassium cyanide,
hydrochloric acid, and the like, we explain not only the fact of
combination, but also many of the reactions occurring in the
transition of one kind of platino-ammonium salts into another.
Thus by this means we explain the fact that (1) PtX_{2},2NH_{3}
combines with 2NH_{3}, forming salts of Reiset's first base; (2)
and the fact that this compound (represented as follows for
distinctness), PtX_{2},2NH_{3},2NH_{3}, when heated, or even when
boiled in solution, again passes into PtX_{2},2NH_{3} (which
resembles the easy disengagement of water of crystallisation,
&c.); (3) the fact that PtX_{2},2NH_{3} is capable of absorbing,
under the action of the same forces, a molecule of chlorine,
PtX_{2},2NH_{3},Cl_{2}, which it then retains with energy, because
it is attracted, not only by the platinum, but also by the
hydrogen of the ammonia; (4) the fact that this chlorine held in
this compound (of Gerhardt) will have a position unusual in salts,
which will explain a certain (although very feebly-marked)
difficulty of reaction; (5) the fact that this does not exhaust
the faculty of platinum for further combination (we need only
recall the compound PtCl_{4},2HCl,16H_{2}O), and that therefore
both PtX_{2},2NH_{3},Cl_{2} and PtX_{2},2NH_{3},2NH_{3} are still
capable of combination, whence the latter, with chlorine, gives
PtX_{2},2NH_{3},2NH_{3},Cl_{2}, after the type of PtX_{4}Y_{4}
(and perhaps higher); (6) the fact that Gros's compounds thus
formed are readily reconverted into the salts of Reiset's first
base when acted on by reducing agents; (7) the fact that in Gros's
salts, PtX_{2},2NH_{3}(NH_{3}X)_{2}, the newly-attached chlorine
or haloid will react with difficulty with salts of silver, &c.,
because it is attached both to the platinum and to the ammonia,
for both of which it has an attraction; (8) the fact that the
faculty for further combination is not even yet exhausted in the
type of Gros's salts, and that we actually have a compound of
Gros's chlorine salt with platinous chloride and with platinic
chloride; the salt PtSO_{4},2NH_{3},2NH_{3},SO_{4} combines
further also with H_{2}O; (9) the fact that such a faculty of
combination with new molecules is naturally more developed in the
lower forms of combination than in the higher. Hence the salts of
Reiset's first base--for example, PtCl_{2},2NH_{3},2NH_{3}--both
combine with water and give precipitates (soluble in water but not
in hydrochloric acid) of double salts with many salts of the heavy
metals--for example, with lead chloride, cupric chloride, and also
with platinic and platinous chlorides (Buckton's salts). The
latter compounds will have the composition
PtCl_{2},2NH_{3},2NH_{3},PtCl_{2}--that is, the same composition
as the salts of Reiset's second base, but it cannot be identical
with it. Such an interesting case does actually exist. The first
salt, PtCl_{2},4NH_{3},PtCl_{2}, is green, insoluble in water and
in hydrochloric acid, and is known as _Magnus's salt_, and the
second, PtCl_{2},2NH_{3}, is Reiset's yellow, sparingly soluble
(in water). They are polymeric, namely, the first contains twice
the number of elements held in the second, and at the same time
they easily pass into each other. If ammonia be added to a hot
hydrochloric acid solution of platinous chloride, it forms the
salt PtCl_{2},4NH_{3}, but in the presence of an excess of
platinous chloride it gives Magnus's salt. On boiling the latter
in ammonia it gives a colourless soluble salt of Reiset's first
base, PtCl_{2},4NH_{3}, and if this be boiled with water, ammonia
is disengaged, and a salt of Reiset's second base,
PtCl_{2},2NH_{3}, is obtained.
A class of platino-ammonium isomerides (obtained by Millon and
Thomsen) are also known. Buckton's salts--for example, the copper
salt--were obtained by them from the salts of Reiset's first base,
PtCl_{2},4NH_{3}, by treatment with a solution of cupric chloride,
&c., and therefore, according to our method of expression,
Buckton's copper salt will be PtCl_{2},4NH_{3},CuCl_{2}. This salt
is soluble in water, but not in hydrochloric acid. In it the
ammonia must be considered as united to the platinum. But if
cupric chloride be dissolved in ammonia, and a solution of
platinous chloride in ammonium chloride is added to it, a violet
precipitate is obtained of the same composition as Buckton's salt,
which, however, is insoluble in water, but soluble in hydrochloric
acid. In this a portion, if not all, of the ammonia must be
regarded as united to the copper, and it must therefore be
represented as CuCl_{2},4NH_{3},PtCl_{2}. This form is identical
in composition but different in properties (is isomeric) with the
preceding salt (Buckton's). The salt of Magnus is intermediate
between them, PtCl_{2},4NH_{3},PtCl_{2}; it is insoluble in water
and hydrochloric acid. These and certain other instances of
isomeric compounds in the series of the platino-ammonium salts
throw a light on the nature of the compounds in question, just as
the study of the isomerides of the carbon compounds has served and
still serves as the chief cause of the rapid progress of organic
chemistry. In conclusion, we may add that (according to the law of
substitution) we must necessarily expect all kinds of intermediate
compounds between the platino and analogous ammonia derivatives on
the one hand, and the complex compounds of nitrous acid on the
other. Perhaps the instance of the reaction of ammonia upon osmic
anhydride, OsO_{4}, observed by Fritsche, Frémy, and others, and
more fully studied by Joly (1891), belongs to this class. The
latter showed that when ammonia acts upon an alkaline solution of
OsO_{4} the reaction proceeds according to the equation: OsO_{4} +
KHO + NH_{3} = OsNKO_{3} + 2H_{2}O. It might be imagined that in
this case the ammonia is oxidised, probably forming the residue of
nitrous acid (NO), while the type OsO_{4} is deoxidised into
OsO_{2}, and a salt, OsO(NO)(KO), of the type OsX_{4} is formed.
This salt crystallises well in light yellow octahedra. It
corresponds to _osmiamic acid_, OsO(ON)(HO), whose anhydride
[OsO(NO)]_{2}, has the composition Os_{2}N_{2}O_{5}, which equals
2Os + N_{2}O_{5} to the same extent as the above-mentioned
compound PtCO_{2} equals Pt + CO_{2} (_see_ Note 11).
CHAPTER XXIV
COPPER, SILVER, AND GOLD
That degree of analogy and difference which exists between iron, cobalt,
and nickel repeats itself in the corresponding triad ruthenium, rhodium,
and palladium, and also in the heavy platinum metals, osmium, iridium,
and platinum. These nine metals form Group VIII. of the elements in the
periodic system, being the intermediate group between the even elements
of the large periods and the uneven, among which we know zinc, cadmium,
and mercury in Group II. Copper, silver, and gold complete[1] this
transition, because their properties place them in proximity to nickel,
palladium, and platinum on the one hand, and to zinc, cadmium, and
mercury on the other. Just as Zn, Cd, and Hg; Fe, Ru, and Os; Co, Rh, and
Ir; Ni, Pd, and Pt, resemble each other in many respects, so also do Cu,
Ag, and Au. Thus, for example, the atomic weight of copper Cu = 63, and
in all its properties it stands between Ni = 59 and Zn = 65. But as the
transition from Group VIII. to Group II., where zinc is situated, cannot
be otherwise than through Group I., so in copper there are certain
properties of the elements of Group I. Thus it gives a suboxide, Cu_{2}O,
and salts, CuX, like the elements of Group I., although at the same time
it forms an oxide, CuO, and salts CuX_{2}, like nickel and zinc. In the
state of the oxide, CuO, and the salts, CuX_{2}, copper is analogous to
zinc, judging from the insolubility of the carbonates, phosphates, and
similar salts, and by the isomorphism, and other characters.[2] In the
cuprous salts there is undoubtedly a great resemblance to the silver
salts--thus, for example, silver chloride, AgCl, is characterised by its
insolubility and capacity of combining with ammonia, and in this respect
cuprous chloride closely resembles it, for it is also insoluble in water,
and combines with ammonia and dissolves in it, &c. Its composition is
also RCl, the same as AgCl, NaCl, KCl, &c., and silver in many compounds
resembles, and is even isomorphous with, sodium, so that this again
justifies their being brought together. Silver chloride, cuprous
chloride, and sodium chloride crystallise in the regular system. Besides
which, the specific heats of copper and silver require that they should
have the atomic weights ascribed to them. To the oxides Cu_{2}O and
Ag_{2}O there are corresponding sulphides Ag_{2}S and Cu_{2}S. They both
occur in nature in crystals of the rhombic system, and, what is most
important, copper glance contains an isomorphous mixture of them both,
and retains the form of copper glance with various proportions of copper
and silver, and therefore has the composition R_{2}S where R = Cu, Ag.
[1] The perfectly unique position held by copper, silver, and gold in
the periodic system of the elements, and the degree of affinity
which is found between them, is all the more remarkable, as nature
and practice have long isolated these metals from all others by
having employed them--for example, for coinage--and determined
their relative importance and value in conformity with the order
(silver between copper and gold) of their atomic weights, &c.
[2] Cupric sulphate contains 5 molecules of water, CuSO_{4},5H_{2}O,
and the isomorphous mixtures with ZnSO_{4},7H_{2}O contain either 5
or 7 equivalents, according to whether copper or zinc predominates
(Vol. II. p. 6). If there be a large proportion of copper, and if
the mixture contain 5H_{2}O, the form of the isomorphous mixture
(triclinic) will be isomorphous with cupric sulphate,
CuSO_{4},5H_{2}O, but if a large amount of zinc (or magnesium,
iron, nickel, or cobalt) be present the form (rhombic or
monoclinic) will be nearly the same as that of zinc sulphate,
ZnSO_{4},7H_{2}O. Supersaturated solutions of each of these salts
crystallise in that form and with that amount of water which is
contained in a crystal of one or other of the salts brought in
contact with the solution (Chapter XIV., Note 27).
Notwithstanding the resemblance in the atomic composition of the cuprous
compounds, CuX, and silver compounds, AgX, with the compounds of the
alkali metals KX, NaX, there is a considerable degree of difference
between these two series of elements. This difference is clearly seen in
the fact that the alkali metals belong to those elements which combine
with extreme facility with oxygen, decompose water, and form the most
alkaline bases; whilst silver and copper are oxidised with difficulty,
form less energetic oxides, and do not decompose water, even at a rather
high temperature. Moreover, they only displace hydrogen from very few
acids. The difference between them is also seen in the dissimilarity of
the properties of many of the corresponding compounds. Thus cuprous
oxide, Cu_{2}O, and silver oxide, Ag_{2}O, are insoluble in water: the
cuprous and silver carbonates, chlorides, and sulphates are also
sparingly soluble in water. The oxides of silver and copper are also
easily reduced to metal. This difference in properties is in intimate
relation with that difference in the density of the metals which exists
in this case. The alkali metals belong to the lightest, and copper and
silver to the heaviest, and therefore the distance between the molecules
in these metals is very dissimilar--it is greater for the former than the
latter (tables in Chapter XV.). From the point of view of the periodic
law, this difference between copper and silver and such elements of Group
I. as potassium and rubidium, is clearly seen from the fact that copper
and silver stand in the middle of those large periods (for example, K,
Ca, Sc, Ti, V, Cr, Mn, Fe, Co, Ni, Cu, Zn, Ga, Ge, As, Se, Br) which
start with the true metals of the alkalis--that is to say, the analogy
and difference between potassium and copper are of the same nature as
that between chromium and selenium, or vanadium and arsenic.
_Copper_ is one of the few metals which have long been known in a
metallic form. The Greeks and Romans imported copper chiefly from the
island of Cyprus--whence its Latin name, _cuprum_. It was known to the
ancients before iron, and was used, especially when alloyed with other
metals, for arms and domestic utensils. That copper was known to the
ancients will be understood from the fact that it occurs, although
rarely, in a _native state_, and is easily extracted from its other
natural compounds. Among the latter are the oxygen compounds of copper.
When ignited with charcoal, they easily give up their oxygen to it, and
yield metallic copper; hydrogen also easily takes up the oxygen from
copper oxide when heated. Copper occurs in a native state, sometimes in
association with other ores, in many parts of the Urals and in Sweden,
and in considerable masses in America, especially in the neighbourhood of
the great American lakes; and also in Chili, Japan, and China. The oxygen
compounds of copper are also of somewhat common occurrence in certain
localities; in this respect certain deposits of the Urals are especially
famous. The geological period of the Urals (Permian) is characterised by
a considerable distribution of copper ores. Copper is met with in the
form of _cuprous oxide_, or _suboxide of copper_, Cu_{2}O, and is then
known as _red copper ore_, because it forms red masses which not
unfrequently are crystallised in the regular system. It is found much
more rarely in the state of _cupric oxide_, CuO, and is then called
_black copper ore_. The most common of the oxygenised compounds of copper
are the _basic carbonates_ corresponding with the oxides. That these
compounds are undoubtedly of aqueous origin is apparent, not only from
the fact that specimens are frequently found of a gradual transition from
the metallic, sulphuretted, and oxidised copper into its various
carbonates, but also from the presence of water in their composition, and
from the laminar, reniform structure which many of them present. In this
respect _malachite_ is particularly well known; it is used as a green
paint and also for ornaments, owing to the diversity of the shades of
colour presented by the different layers of deposited malachite. The
composition of malachite corresponds with the basic carbonate containing
one molecule of cupric carbonate to one of hydroxide:
CuCO_{3},CuH_{2}O_{2}. In this form the copper frequently occurs in
admixture with various sedimentary rocks, forming large strata, which
confirms the aqueous origin of these compounds. There are many such
localities in the Perm and other Governments bounding the Urals. _Blue
carbonate of copper_, or _azurite_, is also often met with in the same
localities; it contains the same ingredients as malachite, but in a
different proportion, its composition being CuH_{2}O_{2},2CuCO_{3}. Both
these substances may be obtained artificially by the action of the alkali
carbonates on solutions of cupric salts at various temperatures. These
native carbonates are often used for the extraction of copper, all the
more as they very readily give metallic copper, evolving water and
carbonic anhydride when ignited, and leaving the easily-reducible cupric
oxide. Copper is, however, still more often met with in the form of the
sulphides. The sulphides of copper generally occur in chemical
combination with the sulphides of iron.[3] These copper-sulphur compounds
(copper pyrites CuFeS_{2}, variegated copper ore Cu_{3}FeS_{3}, &c.)
generally occur in veins in a rock gangue.
[3] Iron pyrites, FeS_{2}, very often contain a small quantity of
copper sulphide (_see_ Chapter XXII., Note 2 bis), and on burning
the iron pyrites for sulphurous anhydride the copper oxide remains
in the residue, from which the copper is often extracted with
profit. For this purpose the whole of the sulphur is not burnt off
from the iron pyrites, but a portion is left behind in the ore,
which is then slowly ignited (roasted) with access of air. Cupric
sulphate is then formed, and is extracted by water; or what is
better and more frequently done, the residue from the roasting of
the pyrites is roasted with common salt, and the solution of cupric
chloride obtained by lixiviating is precipitated with iron. A far
greater amount of copper is obtained from other sulphuretted ores.
Among these _copper glance_, Cu_{2}S, is more rarely met with. It
has a metallic lustre, is grey, generally crystalline, and is
obtained in admixture with organic matter; so that there is no
doubt that its origin is due to the reducing action of the latter
on solutions of cupric sulphate. _Variegated copper ore_, which
crystallises in octahedra, not infrequently forms an admixture in
copper glance; it has a metallic lustre, and is reddish-brown; it
has a superficial play of colours, due to oxidation proceeding on
its surface. Its composition is Cu_{3}FeS_{3}. But the most common
and widely-distributed copper ore is _copper pyrites_, which
crystallises in regular octahedra; it has a metallic lustre, a sp.
gr. of 4·0, and yellow colour. Its composition is CuFeS_{2}. It
must be remarked that the sulphurous ores of copper are oxidised in
the presence of water containing oxygen in solution, and form
cupric sulphate, blue vitriol, which is easily soluble in water. If
this water contains calcium carbonate, gypsum and cupric carbonate
are formed by double decomposition: CuSO_{4} + CaCO_{3} = CuCO_{3}
+ CaSO_{4}. Hence copper sulphide in the form of different ores
must be considered as the primary product, and the many other
copper ores as secondary products, formed by water. This is
confirmed by the fact that at the present time the water extracted
from many copper mines contains cupric sulphate in solution. From
this liquid it is easy to extract cupric oxide by the action of
organic matter and various impurities of water. Hence metallic
copper is sometimes found in natural products of the modification
of copper sulphide and is probably deposited by the action of
organic matter present in the water.
_The extraction of copper from its oxide ores_ does not present any
difficulty, because the copper, when ignited with charcoal and melted, is
reduced from the impurities which accompany it. This mode of smelting
copper ores is carried on in cupola or cylindrical furnaces, fluxes
forming a slag being added to the mixture of ore and charcoal. The
smelted copper still contains sulphur, iron, and other metallic
impurities, from which it is freed by fusion in reverberatory furnaces,
with access of air to the surface of the molten metal, as the iron and
sulphur are more easily oxidised than the copper. The iron then separates
as oxides, which collect in the slag.[4]
[4] Copper ores rich in oxygen are very rare; the sulphur ores are of
more common occurrence, but the extraction of the copper from them
is much more difficult. The problem here not only consists in the
removal of the sulphur, but also in the removal of the iron
combined with the sulphur and copper. This is attained by a whole
series of operations, after which there still sometimes remains the
extraction of the metallic silver which generally accompanies the
copper, although in but small quantity. These processes commence
with the roasting--_i.e._ calcination--of the ore with access of
air, by which means the sulphur is converted into sulphurous
anhydride. It should here be remarked that iron sulphide is more
easily oxidised than copper sulphide, and therefore the greater
part of the iron in the residue from roasting is no longer in the
form of sulphide but of oxide of iron. The roasted ore is mixed
with charcoal, and siliceous fluxes, and smelted in a cupola
furnace. The iron then passes into the slag, because its oxide
gives an easily-fusible mass with the silica, whilst the copper, in
the form of sulphide, fuses and collects under the slag. The
greater part of the iron is removed from the mass by this smelting.
The resultant _coarse metal_ is again roasted in order to remove
the greater part of the sulphur from the copper sulphide, and to
convert the metal into oxide, after which the mass is again
smelted. These processes are repeated several times, according to
the richness of the ore. During these smeltings a portion of the
copper is already obtained in a metallic form, because copper
sulphide gives metallic copper with the oxide (CuS + 2CuO = 3Cu +
SO_{2}). We will not here describe the furnaces used or the details
of this process, but the above remarks include the explanation of
those chemical processes which are accomplished in the various
technical operations which are made use of in the process (for
details _see_ works on metallurgy).
Besides the smelting of copper there also exist methods for its
extraction from solutions in the wet way, as it is called. Recourse
is generally had to these methods for poor copper ores. The copper
is brought into solution, from which it is separated by means of
metallic iron or by other methods (by the action of an electric
current). The sulphides are roasted in such a manner that the
greater part of the copper is oxidised into cupric sulphate, whilst
at the same time the corresponding iron salts are as far as
possible decomposed. This process is based on the fact that the
copper sulphides absorb oxygen when they are calcined in the
presence of air, forming cupric sulphate. The roasted ore is
treated with water, to which acid is sometimes added, and after
lixiviation the resultant solution containing copper is treated
either with metallic iron or with milk of lime, which precipitates
cupric hydroxide from the solution. Copper oxide ores poor in metal
may be treated with dilute acids in order to obtain the copper
oxides in solution, from which the copper is then easily
precipitated either by iron or as hydroxide by lime. According to
Hunt and Douglas's method, the copper in the ore is converted by
calcination into the cupric oxide, which is brought into solution
by the action of a mixture of solutions of ferrous sulphate and
sodium chloride; the oxide converts the ferrous chloride into
ferric oxide, forming copper chlorides, according to the equation
3CuO + 2FeCl_{2} = CuCl_{2} + 2CuCl + Fe_{2}O_{3}. The cupric
chloride is soluble in water, whilst the cuprous chloride is
dissolved in the solution of sodium chloride, and therefore all the
copper passes into solution, from which it is precipitated by iron.
The same American metallurgists give the following wet method for
extracting the Ag and Au occurring in many copper ores, especially
in sulphurous ores: (1) The Cu_{2}S is first converted into oxide
by roasting in a calciner; (2) the CuO is extracted by the dilute
sulphuric acid obtained in the fourth process, the Cu then passes
into solution, while the Ag, Au and oxides of iron remain behind in
the residue (from which the noble metals may be extracted); (3) a
portion of the copper in solution is converted into CuCl_{2} (and
CaSO_{4} precipitated) by means of the CaCl_{2} obtained in the
fifth process; (4) the mixture of solutions of CuSO_{4} and
CuCl_{2} is converted into the insoluble CuCl (salt of the
suboxide) by the action of the SO_{2} obtained by roasting the ore
(in the first operation), sulphuric acid is then formed in the
solution, according to the equation: CuSO_{4} + CuCl_{2} + SO_{2} +
2H_{2}O = 2H_{2}SO_{4} + 2CuCl; (5) the precipitated CuCl is
treated with lime and water, and gives CuCl_{2} in solution and CuO
in the residue; and lastly (6) the Cu_{2}O is reduced to metallic
Cu by carbon in a furnace. According to Crooke's method the impure
copper regulus obtained by roasting and smelting the ore is broken
up and immersed repeatedly in molten lead, which extracts the Ag
and Au occurring in the regulus. The regulus is then heated in a
reverberatory furnace to run off the lead, and is then smelted for
Cu.
The copper brought into the market often contains small quantities
of various impurities. Among these there are generally present
iron, lead, silver, arsenic, and sometimes small quantities of
oxides of copper. As copper, when mixed with a small amount of
foreign substances, loses its tenacity to a certain degree, the
manufacture of very thin sheet copper requires the use of Chili
copper, which is distinguished for its great softness, and
therefore when it is desired to have pure copper, it is best to
take thin sheet copper, like that which is used in the manufacture
of cartridges. But the purest copper is electrolytic copper--that
is, that which is deposited from a solution by the action of an
electric current.
If the copper contains silver, as is often the case, it is used in
gold refineries for the precipitation of silver from its solutions
in sulphuric acid. Iron and zinc reduce copper salts, but copper
reduces mercury and silver salts. The precipitate contains not only
the silver which was previously in solution, but also all that
which was in the copper. The silver solutions in sulphuric acid are
obtained in the separation of silver from gold by treating their
alloys with sulphuric acid, which only dissolves the silver.
Copper is characterised by its red colour, which distinguishes it from
all other metals. Pure copper is soft, and may be beaten out by a hammer
at the ordinary temperature, and when hot may be rolled into very thin
sheets. Extremely thin leaves of copper transmit a green light. The
tenacity of copper is also considerable, and next to iron it is one of
the most durable metals in this respect. Copper wire of 1 sq. millimetre
in section only breaks under a weight of 45 kilograms. The specific
gravity of copper is 8·8, unless it contains cavities due to the fact
that molten copper absorbs oxygen from the air, which is disengaged on
cooling, and therefore gives a porous mass whose density is much less.
Rolled copper, and also that which is deposited by the electric current,
has a comparatively high density. Copper melts at a bright red heat,
about 1050°, although below the temperature at which many kinds of cast
iron melt. At a high temperature it is converted into vapour, which
communicates a green colour to the flame. Both native copper and that
cooled from a molten state crystallise in regular octahedra. Copper is
not oxidised in dry air at the ordinary temperature, but when calcined it
becomes coated with a layer of oxide, and it does not burn even at the
highest temperature. Copper, when calcined in air, forms either the red
cuprous oxide or the black cupric oxide, according to the temperature and
quantity of air supplied. In air at the ordinary temperature, copper--as
everyone knows--becomes coated with a brown layer of oxides or a green
coating of basic salts, due to the action of the damp air containing
carbonic acid. If this action continue for a prolonged time, the copper
is covered with a thick coating of basic carbonate, or the so-called
verdigris (the _ærugo nobilis_ of ancient statues). This is due to the
fact that copper, although scarcely capable of oxidising by itself,[5]
_in the presence of water and acids_--even very feeble acids, like
carbonic acid--_absorbs oxygen from the air and forms salts_, which is a
very characteristic property of it (and of lead).[6] _Copper does not
decompose water_, and therefore does not disengage hydrogen from it
either at the ordinary or at high temperatures. Nor does copper liberate
hydrogen from the oxygen acids; these act on it in two ways: they either
give up a portion of their oxygen, forming lower grades of oxidation, or
else only react in the presence of air. Thus, when nitric acid acts on
copper it evolves nitric oxide, the copper being oxidised at the expense
of the nitric acid. In the same way copper converts sulphuric acid into
the lower grade of oxidation--into sulphurous anhydride, SO_{2}. In these
cases the copper is oxidised to copper oxide, which combines with the
excess of acid taken, and therefore forms a cupric salt, CuX_{2}. Dilute
nitric acid does not act on copper at the ordinary temperature, but when
heated it reacts with great ease; dilute sulphuric acid does not act on
copper except in presence of air.
[5] Schützenberger showed that when the basic carbonate of copper is
decomposed by an electric current it gives, besides the ordinary
copper, an allotropic form which grows on the negative platinum
electrode, if its surface be smaller than that of the positive
copper electrode, in the form of brittle crystalline growths of sp.
gr. 8·1. It differs from ordinary copper by giving not nitric oxide
but nitrous oxide when treated with nitric acid, and in being very
easily oxidised in air, and coated with red shades of colour. It is
possible that this is copper hydride, or copper which has occluded
hydrogen. Spring (1892) observed that copper reduced from the oxide
by hydrogen at the lowest possible temperature was pulverulent,
while that reduced from CuCl_{2} at a somewhat high temperature
appeared in bright crystals. The same difference occurs with many
other metals, and is probably partly due to the volatility of the
metallic chlorides.
[6] This is taken advantage of in practice; for instance, by pouring
dilute acids over copper turnings on revolving tables in the
preparation of copper salts, such as verdigris, or the basic
acetate 2C_{4}H_{6}CuO_{4},CuH_{2}O_{2},5H_{2}O, which is so much
used as an oil paint (_i.e._ with boiled oil). The capacity of
copper for absorbing oxygen in the presence of acids is so great
that it is possible by this means (by taking, for example, thin
copper shavings moistened with sulphuric acid) to take up all the
oxygen from a given volume of air, and this is even employed for
the analysis of air.
The combination of copper with oxygen is not only aided by acids
but also by alkalis, although cupric oxide does not appear to have
an acid character. Alkalis do not act on copper except in the
presence of air, when they produce cupric oxide, which does not
appear to combine with such alkalis as caustic potash or soda. But
the _action of ammonia_ is particularly distinct (Chapter V., Note
2). In the action of a solution of ammonia not only is oxygen
absorbed by the copper, but it also acts on the ammonia, and a
definite quantity of ammonia is always acted on simultaneously with
the passage of the copper into solution. The ammonia is then
converted into nitrous acid, according to the reaction: NH_{3} +
O_{3} = NHO_{2} + H_{2}O, and the nitrous acid thus formed passes
into the state of ammonium nitrite, NH_{4}NO_{2}. In this manner
three equivalents of oxygen are expended on the oxidation of the
ammonia, and six equivalents of oxygen pass over to the copper,
forming six atoms of cupric oxide. The latter does not remain in
the state of oxide, but combines with the ammonia.
A strong solution of common salt does not act on copper, but a
dilute solution of the salt corrodes copper, converting it into
oxychloride--that is, in the presence of air. This action of salt
water is evident in those cases where the bottoms of ships are
coated with sheet copper. From what has been said above it will be
evident that copper vessels should not be employed in the
preparation of food, because this contains salts and acids which
act on copper in the presence of air, and give copper salts, which
are poisonous, and therefore the food prepared in untinned copper
vessels may be poisonous. Hence tinned vessels are employed for
this purpose--that is, copper vessels coated with a thin layer of
tin, on which acid and saline solutions do not act.
Both the oxides of copper, Cu_{2}O and CuO, are unacted on by air, and,
as already mentioned, they both occur in nature.[6 bis] However, in the
majority of cases copper is obtained in the form of cupric oxide and its
salts--and the copper compounds used industrially generally belong to
this type. This is due to the fact that the _cuprous compounds absorb
oxygen_ from the air and pass into cupric compounds. The cupric compounds
may serve as the source for the preparation of cuprous oxide, because
many reducing agents are capable of deoxidising the oxide into the
suboxide. Organic substances are most generally employed for this
purpose, and especially saccharine substances, which are able, in the
presence of alkalis, to undergo oxidation at the expense of the oxygen of
the cupric oxide, and to give acids which combine with the alkali: 2CuO -
O = Cu_{2}O. In this case the deoxidation of the copper may be carried
further and metallic copper obtained, if only the reaction be aided by
heat. Thus, for example, a fine powder of metallic copper may be obtained
by heating an ammoniacal solution of cupric oxide with caustic potash and
grape sugar. But if the reducing action of the saccharine substance
proceed in the presence of a sufficient quantity of alkali in solution,
and at not too high a temperature, cuprous oxide is obtained. To see this
reaction clearly, it is not sufficient to take any cupric salt, because
the alkali necessary for the reaction might precipitate cupric oxide--it
is necessary to add previously some substance which will prevent this
precipitation. Among such substances, tartaric acid, C_{4}H_{6}O_{6}, is
one of the best. In the presence of a sufficient quantity of tartaric
acid, any amount of alkali may be added to a solution of cupric salt
without producing a precipitate, because a soluble double salt of cupric
oxide and alkali is then formed. If glucose (for instance, honey or
molasses) be added to such an alkaline tartaric solution, and the
temperature be slightly raised, it first gives a yellow precipitate (this
is cuprous hydroxide, CuHO), and then, on boiling, a red precipitate of
(anhydrous) cuprous oxide. If such a mixture be left for a long time at
the ordinary temperature, it deposits well-formed crystals of anhydrous
cuprous oxide belonging to the regular system.[7]
[6 bis] Copper, besides the cuprous oxide, Cu_{2}O, and cupric oxide,
CuO, gives two known higher forms of oxidation, but they have
scarcely been investigated, and even their composition is not well
known. _Copper dioxide_ (CuO_{2}, or CuO_{2},H_{2}O, perhaps
CuOH_{2}O_{2}) is obtained by the action of hydrogen peroxide on
cupric hydroxide, when the green colour of the latter is changed to
yellow. It is very unstable, and is decomposed even by boiling
water, with the evolution of oxygen, whilst the action of acids
gives cupric salts, oxygen being also disengaged. A still higher
_copper peroxide_ is formed by heating a mixture of caustic potash,
nitre, and metallic copper to a red heat, and by dissolving cupric
hydroxide in solutions of the hypochlorites of the alkali metals. A
slight heating of the soluble salt formed is enough for it to be
decomposed into oxygen and copper dioxide, which is precipitated.
Judging from Frémy's researches, the composition of the
copper-potassic compound should be K_{2}CuO_{4}. Perhaps this is a
compound of the peroxides of potassium, K_{2}O_{2}, and of copper,
CuO_{2}.
[7] Colourless solutions of cuprous salts may also be obtained by the
action of sulphurous or phosphorous acid and similar lower grades
of oxidation on the blue solutions of the cupric salts. This is
very clearly and easily effected by means of sodium thiosulphate,
Na_{2}S_{2}O_{3}, which is oxidised in the process. Cuprous oxide
can not only be obtained by the deoxidation of cupric oxide, but
also directly from metallic copper itself, because the latter, in
oxidising at a red heat in air, first gives cuprous oxide. It is
prepared in this manner on a large scale by heating sheet copper
rolled into spirals in reverberatory furnaces. Care must be taken
that the air is not in great excess, and that the coating of red
cuprous oxide formed does not begin to pass into the black cupric
oxide. If the oxidised spiral sheet is then unbent, the brittle
cuprous oxide falls away from the soft metal. The suboxide obtained
in this manner fuses with ease. It is necessary to prevent the
access of air during the fusion, and if the mass contains cupric
oxide it must be mixed with charcoal, which reduces the latter.
Cuprous chloride, CuCl, corresponding with cuprous oxide (as sodium
chloride corresponds with sodium oxide), when calcined with sodium
carbonate, gives sodium chloride and cuprous oxide, carbonic
anhydride being evolved, because it does not combine with the
cuprous oxide under these conditions. The reaction can be expressed
by the following equation: 2CuCl + Na_{2}CO_{3} = Cu_{2}O + 2NaCl +
CO_{2}. The cupric oxide itself, when calcined with finely-divided
copper, this copper powder may be obtained by many methods--for
instance, by immersing zinc in a solution of a copper salt, or by
igniting cupric oxide in hydrogen), gives the fusible cuprous
oxide: Cu + CuO = Cu_{2}O. Both the native and artificial cuprous
oxide have a sp. gr. of 5·6. It is insoluble in water, and is not
acted on by (dry) air. When heated with acids the suboxide forms a
solution of a cupric salt and metallic copper--for example, Cu_{2}O
+ H_{2}SO_{4} = Cu + CuSO_{4} + H_{2}O. However, strong
hydrochloric acid does not separate metallic copper on dissolving
cuprous oxide, which is due to the fact that the cuprous chloride
formed is soluble in strong hydrochloric acid. Cuprous oxide also
dissolves in a solution of ammonia, and in the absence of air gives
a colourless solution, which turns blue in the air, absorbing
oxygen, owing to the conversion of the cuprous oxide into cupric
oxide. The blue solution thus formed may be again reconverted into
a colourless cuprous solution by immersing a copper strip in it,
because the metallic copper then deoxidises the cupric oxide in the
solution into cuprous oxide. Cuprous oxide is characterised by the
fact that it gives red glasses when fused with glass or with salts
forming vitreous alloys. Glass tinted with cuprous oxide is used
for ornaments. The access of air must be avoided during its
preparation, because the colour then becomes green, owing to the
formation of cupric oxide, which colours glass blue. This may even
be taken advantage of in testing for copper under the blow-pipe by
heating the copper compound with borax in the flame of a blow-pipe;
a red glass is obtained in the reducing flame, and a blue glass in
the oxidising flame, owing to the conversion of the cuprous into
cupric oxide.
Étard (1882), by passing sulphurous anhydride into a solution of
cupric acetate, obtained a white precipitate of cuprous sulphite,
Cu_{2}SO_{3},H_{2}O, whilst he obtained the same salt, of a red
colour, from the double salt of sodium and copper; but there are
not any convincing proofs of isomerism in this case.
Cupric chloride, CuCl_{2}, when ignited, gives _cuprous chloride_,
CuCl--_i.e._ the salt corresponding with suboxide of copper--and
therefore cuprous chloride is always formed when copper enters into
reaction with chlorine at a high temperature. Thus, for example, when
copper is calcined with mercuric chloride, it forms cuprous chloride and
vapours of mercury. The same substance is obtained on heating metallic
copper in hydrochloric acid, hydrogen being disengaged; but this reaction
only proceeds with finely-divided copper, as hydrochloric acid acts very
feebly on compact masses of copper, and, in the presence of air, gives
cupric chloride. The green solution of cupric chloride is decolorised by
metallic copper, cuprous chloride being formed; but this reaction is only
accomplished with ease when the solution is very concentrated and in the
presence of an excess of hydrochloric acid to dissolve the cuprous
chloride. The addition of water to the solution precipitates the cuprous
chloride, because it is less soluble in dilute than in strong
hydrochloric acid. Many reducing agents which are able to take up half
the oxygen from cupric oxide are able, in the presence of hydrochloric
acid, to form cuprous chloride. Stannous salts, sulphurous anhydride,
alkali sulphites, phosphorous and hypophosphorous acids, and many similar
reducing agents, act in this manner. The usual method of preparing
cuprous chloride consists in passing sulphurous anhydride into a very
strong solution of cupric chloride: 2CuCl_{2} + SO_{2} + 2H_{2}O = 2CuCl
+ 2HCl + H_{2}SO_{4}. Cuprous chloride forms colourless cubic crystals
which are insoluble in water. It is easily fusible, and even volatile.
Under the action of oxidising agents, it passes into the cupric salt, and
it absorbs oxygen from moist air, forming cupric oxychloride,
Cu_{2}Cl_{2}O. _Aqueous ammonia_ easily _dissolves_ cuprous chloride as
well as cuprous oxide; the solution also turns blue on exposure to the
air. Thus an ammoniacal solution of cuprous chloride serves as an
excellent absorbent for oxygen; but this solution absorbs not only
oxygen, but also certain other gases--for example, carbonic oxide and
acetylene.[8]
[8] The solubility of cuprous chloride in ammonia is due to the
formation of compounds between the ammonia and the chloride. In a
warm solution the compound NH_{3},2CuCl is formed, and at the
ordinary temperature CuCl,NH_{3}. This salt is soluble in
hydrochloric acid, and then forms a corresponding double salt of
cuprous chloride and ammonium chloride. By the action of a certain
excess of ammonia on a hydrochloric acid solution of cuprous
chloride, very well formed crystals, having the composition
CuCl,NH_{3},H_{2}O, are obtained. Cuprous chloride is not only
soluble in ammonia and hydrochloric acid, but it also dissolves in
solutions of certain other salts--for example, in sodium chloride,
potassium chloride, sodium thiosulphate, and certain others. All
the solutions of cuprous chloride act in many cases as very
powerful deoxidising substances; for example, it is easy, by means
of these solutions, to precipitate gold from its solutions in a
metallic form, according to the equation AuCl_{3} + 3CuCl = Au +
3CuCl_{2}.
Among the other compounds corresponding with cuprous oxide,
_cuprous iodide_, CuI, is worthy of remark. It is a colourless
substance which is insoluble in water and sparingly soluble in
ammonia (like silver iodide), but capable of absorbing it, and in
this respect it resembles cuprous chloride. It is remarkable from
the fact that it is exceedingly easily formed from the
corresponding cupric compound CuI_{2}. A solution of cupric iodide
easily decomposes into iodine and cuprous iodide, even at the
ordinary temperature, whilst cupric chloride only suffers a similar
change on ignition. If a solution of a cupric salt be mixed with a
solution of potassium iodide the cupric iodide formed immediately
decomposes into free iodine and cuprous iodide, which separates out
as a precipitate. In this case the cupric salt acts in an oxidising
manner, like, for example, nitrous acid, ozone, and other
substances which liberate iodine from iodides, but with this
difference, that it only liberates half, whilst they set free the
whole of the iodine from potassium iodide: 2KI + CuCl_{2} = 2KCl +
CuI + I.
It must also be remarked that cuprous oxide, when treated with
hydrofluoric acid, gives an insoluble cuprous fluoride, CuF.
Cuprous cyanide is also insoluble in water, and is obtained by the
addition of hydrocyanic acid to a solution of cupric chloride
saturated with sulphurous anhydride. This cuprous cyanide, like
silver cyanide, gives a double soluble salt with potassium cyanide.
The double cyanide of copper and potassium is tolerably stable in
the air, and enters into double decompositions with various other
salts, like those double cyanides of iron with which we are already
acquainted.
_Copper hydride_, CuH, also belongs to the number of the cuprous
compounds. It was obtained by Würtz by mixing a hot (70°) solution
of cupric sulphate with a solution of hypophosphorous acid,
H_{3}PO_{2}. The addition of the reducing hypophosphorous acid must
be stopped when a brown precipitate makes its appearance, and when
gas begins to be evolved. The brown precipitate is the hydrated
cuprous hydride. When gently heated it disengages hydrogen; it
gives cuprous oxide when exposed to the air, burns in a stream of
chlorine, and liberates hydrogen with hydrochloric acid: CuH + HCl
= CuCl + H_{2}. Zinc, silver, mercury, lead, and many other heavy
metals do not form such a compound with hydrogen, neither under
these circumstances nor under the action of hydrogen at the moment
of the decomposition of salts by a galvanic current. The greatest
resemblance is seen between cuprous hydride and the hydrogen
compounds of potassium, sodium, Pd, Ca, and Ba.
When copper is oxidised with a considerable quantity of oxygen at a
high temperature, or at the ordinary temperature in the presence of
acids, and also when it decomposes acids, converting them into lower
grades of oxidation (for example, when submitted to the action of nitric
and sulphuric acids), it forms _cupric oxide_, CuO, or, in the presence
of acids, cupric salts. Copper rust, or that black mass which forms on
the surface of copper when it is calcined, consists of cupric oxide. The
coating of the oxidised copper is very easily separated from the metallic
copper, because it is brittle and very easily peels off, when it is
struck or immersed in water. Many copper salts (for instance, the nitrite
and carbonate) leave oxide of copper[8 bis] in the form of friable black
powder, after being ignited. If the ignition be carried further, Cu_{2}O
may be formed from the CuO.[8 tri] Anhydrous cupric oxide is very easily
dissolved in acids, forming cupric salts, CuX_{2}. They are analogous to
the salts MgX_{2}, ZnX_{2}, NiX_{2}, FeX_{2}, in many respects. On adding
potassium or ammonium hydroxide to a solution of a cupric salt, it forms
a gelatinous blue precipitate of the hydrated oxide of copper,
CuH_{2}O_{2}, insoluble in water. The resultant precipitate _is
redissolved by an excess of ammonia_, and gives a very beautiful azure
blue solution, of so intense a colour that the presence of small traces
of cupric salts may be discovered by this means.[9] An excess of
potassium or sodium hydroxide does not dissolve cupric hydroxide. A hot
solution gives a black precipitate of the anhydrous oxide instead of the
blue precipitate, and the precipitate of the hydroxide of copper becomes
granular, and turns black when the solution is heated. This is due to the
fact that the blue hydroxide is exceedingly unstable, and when slightly
heated it loses the elements of water and gives the black anhydrous
cupric oxide: CuH_{2}O_{2} = CuO + H_{2}O.
[8 bis] The oxide of copper obtained by igniting the nitrate is
frequently used for organic analyses. It is hygroscopic and retains
nitrogen (1·5 c.c. per gram) when the nitrate is heated in vacuo
(Richards and Rogers, 1893).
[8 tri] Oxide of copper is also capable of dissociating when heated.
Debray and Joannis showed that it then disengages oxygen, whose
maximum tension is constant for a given temperature, providing that
fusion does not take place (the CuO then dissolves in the molten
Cu_{2}O); that this loss of oxygen is followed by the formation of
suboxide, and that on cooling, the oxygen is again absorbed,
forming CuO.
[9] Cupric oxide and many of its salts are able to give definite,
although unstable, _compounds with ammonia_. This faculty already
shows itself in the fact that cupric oxide, as well as the salts of
copper, dissolves in aqueous ammonia, and also in the fact that
salts of copper absorb ammonia gas. If ammonia be added to a
solution of any cupric salt, it first forms a precipitate of cupric
hydroxide, which then dissolves in an excess of ammonia. The
solution thus formed, when evaporated or on the addition of
alcohol, frequently deposits crystals of salts containing both the
elements of the salt of copper taken and of ammonia. Several such
compounds are generally formed. Thus cupric chloride, CuCl_{2},
according to Deherain, forms four compounds with ammonia--namely,
with one, two, four, and six molecules of ammonia. Thus, for
example, if ammonia gas be passed into a boiling saturated solution
of cupric chloride, on cooling, small octahedral crystals of a blue
colour separate out, containing CuCl_{2},2NH_{3},H_{2}O. At 150°
this substance loses half the ammonia and all the water contained
in it, leaving the compound CuCl_{2},NH_{3}. Nitrate of copper
forms the compound Cu(NO_{3})_{2},2NH_{3}· This compound remains
unchanged on evaporation. Dry cupric sulphate absorbs ammonia gas,
and gives a compound containing five molecules of ammonia to one of
sulphate (Vol. I., p. 257, and Chapter XXII., Note 35). If this
compound is dissolved in aqueous ammonia, on evaporation it
deposits a crystalline substance containing
CuSO_{4},4NH_{3},H_{2}O. At 150° this substance loses the molecule
of water and one-fourth of its ammonia. On ignition all these
compounds part with the remaining ammonia in the form of an
ammoniacal salt, so that the residue consists of cupric oxide. Both
the hydrated and anhydrous cupric oxide are soluble in aqueous
ammonia.
The solution obtained by the action of aqueous ammonia and air on
copper turnings (Note 6) is remarkable for its faculty of
_dissolving cellulose_, which is insoluble in water, dilute acids,
and alkalis. Paper soaked in such a solution acquires the property
of not rotting, of being difficultly combustible, and waterproof,
&c. It has therefore been applied, especially in England, to many
practical purposes--for example, to the construction of temporary
buildings, for covering roofs, &c. The composition of the substance
held in solution is Cu(HO)_{2},4NH_{3}.
If dry ammonia gas be passed over cupric oxide heated to 265°, a
portion of the oxide of copper remains unaltered, whilst the other
portion gives _copper nitride_, the oxygen of the copper oxide
combining with the hydrogen and forming water. The oxide of copper
which remains unchanged is easily removed by washing the resultant
product with aqueous ammonia. Copper nitride is very stable, and is
insoluble; it has the composition Cu_{3}N (_i.e._ the copper is
monatomic here as in Cu_{2}O), and is an amorphous green powder,
which is decomposed when strongly ignited, and gives cuprous
chloride and ammonium chloride when treated with hydrochloric acid.
Like the other nitrides, copper nitride, Cu_{3}N, has scarcely been
investigated. Granger (1892), by heating copper in the vapour of
phosphorus, obtained hexagonal prisms of Cu_{5}P, which passed into
Cu_{6}P (previously obtained by Abel) when heated in nitrogen.
Arsenic is easily absorbed by copper, and its presence (like P),
even in small quantities, has a great influence upon the properties
of copper--for instance, pure copper wire 1 sq. mm. in section
breaks under a load of 35 kilos, while the presence of O·22 p.c. of
arsenic raises the breaking load to 42 kilos.
Cupric oxide fuses at a strong heat, and on cooling forms a heavy
crystalline mass, which is black, opaque, and somewhat tenacious. It is a
feebly energetic base, so that not only do the oxides of the metals of
the alkalis and alkaline earths displace it from its compounds, but even
such oxides as those of lead and silver precipitate it from solutions,
which is partially due to these oxides being soluble, although but
slightly so, in water. However, cupric oxide, and especially the
hydroxide, easily combines with even the least energetic acids, and does
not give any compounds with bases; but, on the other hand, _it easily
forms basic salts_,[9 bis] and in this respect outstrips magnesium and
recalls the oxides of lead or mercury. Hence the hydroxide of copper
dissolves in solutions of neutral cupric salts. The cupric salts are
generally blue or green, because cupric hydroxide itself is coloured. But
some of the salts in the anhydrous state are colourless.[10]
[9 bis] As a comparatively feeble base, oxide of copper easily forms
both basic and double salts. As an instance we may mention the
double salts composed of the dichloride CuCl_{2},2H_{2}O and
potassium chloride. The double salt CuK_{2}Cl_{4},2H_{2}O
crystallises from solutions in _blue_ plates, but when heated alone
or with substances taking up water easily gives _brown_ needles
CuKCl_{3} and at the same time KCl, and this reaction is reversible
at 92° as Meyerhoffer (1889) showed (_i.e._ above 92° the simpler
double salt is formed and below 92° the more complex salt). With an
excess of the copper salt, KCl gives another double salt,
Cu_{2}KCl_{5},4H_{2}O, the transition temperature of which is 55°.
The instances of equilibria which are encountered in such complex
relations (_see_ Chapter XIV., Note 25, astrakhanite, and Chapter
XXII., Note 23) are embraced by the _law of phases_ given by Gibbs
(Transactions of the Connecticut Academy of Sciences, 1875-1878, in
J. Willard Gibbs' memoir 'On the equilibrium of heterogeneous
substances:' and in a clearer and more accessible form in H. W.
Bakhuis Roozeboom's papers, Rec. trav. chim., Vol. VI., and in W.
Meyerhoffer's memoir _Die Phasenregel und ihre Anwendungen_, 1893,
to which sources we refer those desiring fuller information
respecting this law). Gibbs calls '_bodies_' substances (simple or
compound) capable of forming homogeneous complexes (for instance,
solutions or intercombinations) of a varied composition; a
_phase_--a mechanically separable portion of such bodies or of
their homogeneous complexes (for instance, a vapour, liquid or
precipitated solid), _perfect equilibrium_--such a state of bodies
and of their complexes as is characterised by a constant pressure
at a constant temperature even under a change in the amount of one
of the component parts (for instance, of a salt in a saturated
solution), while an _imperfect equilibrium_ is such a one for which
such a change corresponds with a change of pressure (for instance,
an unsaturated solution). The law of phases consists in the fact
that: _n bodies only give a perfect equilibrium when n + 1 phases
participate in that equilibrium_--for example, in the equilibrium
of a salt in its saturated solution in water there are two bodies
(the salt and water) and three phases, namely, the salt, solution,
and vapour, which can be mechanically separated from each other,
and to this equilibrium there corresponds a definite tension. At
the same time, _n bodies may occur in n + 2 phases, but only at one
definite temperature and one pressure_; a change of one of these
may bring about another state (perfect or not--equilibrium stable
or unstable). Thus water when liquid at the ordinary temperature
offers two phases (liquid and vapour) and is in perfect equilibrium
(as also is ice below 0°), but water, ice, and vapour (three phases
and only one body) can only be in equilibrium at 0°, and at the
ordinary pressure; with a change of _t_ there will remain either
only ice and vapour or only liquid water and vapour; whilst with a
rise of pressure not only will the vapour pass into the liquid
(there again only remain two phases) but also the temperature of
the formation of ice will fall (by about 7° per 1000 atmospheres).
The same laws of phases are applicable to the consideration of the
formation of simple or double salts from saturated solutions and to
a number of other purely chemical relations. Thus, for example, in
the above-mentioned instance, when the bodies are KCl, CuCl_{2},
and H_{2}O, perfect equilibrium (which here has reference to the
solubility) consisting of four phases, corresponds to the following
seven cases, considering only the phases (above 0°) A =
CuCl_{2},2KCl,2H_{2}O; B = CuCl_{2}KCl; C = CuCl_{2},2H_{2}O,KCl,
solution and vapour: (1) A + B + solution + vapour; (2) A + C +
solution + vapour; (3) A + KCl + solution + vapour; (4) A + B + C +
vapour (it follows that B + KCl + solution gives A); (5) A + C +
KCl + vapour; (6) B + C + solution + vapour; and (7) B + KCl +
solution + vapour. Thus above 92° A gives B + KCl. The law of
phases by bringing complex instances of chemical reaction under
simple physical schemes, facilitates their study in detail and
gives the means of seeking the simplest chemical relations dealing
with solutions, dissociation, double decompositions and similar
cases, and therefore deserves consideration, but a detailed
exposition of this subject must be looked for in works on physical
chemistry.
[10] The normal _cupric nitrate_, CuN_{2}O_{6},3H_{2}O, is obtained as
a deliquescent salt of a blue colour (soluble in water and in
alcohol) by dissolving copper or cupric oxide in nitric acid. It
is so easily decomposed by the action of heat that it is
impossible to drive off the water of crystallisation from it
before it begins to decompose. During the ignition of the normal
salt the cupric oxide formed enters into combination with the
remaining undecomposed normal salt, and gives a basic salt,
CuN_{2}O_{6},2CuH_{2}O_{2}. The same basic salt is obtained if a
certain quantity of alkali or cupric hydroxide or carbonate be
added to the solution of the normal salt, which is even decomposed
when boiled with metallic copper, and forms the basic salt as a
green powder, which easily decomposes under the action of heat and
leaves a residue of cupric oxide. The basic salt, having the
composition CuN_{2}O_{6},3CuH_{2}O_{2}, is nearly insoluble in
water.
The normal _carbonate of copper_, CuCO_{3}, occurs in nature,
although extremely rarely. If solutions of cupric salts be mixed
with solutions of alkali carbonates, then, as in the case of
magnesium, carbonic anhydride is evolved and basic salts are
formed, which vary in composition according to the temperature and
conditions of the reaction. By mixing cold solutions, a voluminous
blue precipitate is formed, containing an equivalent proportion of
cupric hydroxide and carbonate (after standing or heating, its
composition is the same as malachite, sp. gr. 3·51: 2CuSO_{4} +
2Na_{2}CO_{3} + H_{2}O = CuCO_{3},CuH_{2}O_{2} + 2Na_{2}SO_{4} +
CO_{2}. If the resultant blue precipitate be heated in the liquid,
it loses water and is transformed into a granular green mass of
the composition Cu_{2}CO_{4}--_i.e._ into a compound of the normal
salt with anhydrous cupric oxide. This salt of the oxide
corresponds with orthocarbonic acid, C(OH)_{4} = CH_{4}O_{4} where
4H is replaced by 2Cu. On further boiling this salt loses a
portion of the carbonic acid, forming black cupric oxide, so
unstable is the compound of copper with carbonic anhydride.
Another basic salt which occurs in nature, 2CuCO_{3},CuH_{2}O_{2},
is known as azurite, or blue carbonate of copper; it also loses
carbonic acid when boiled with water. On mixing a solution of
cupric sulphate with sodium sesquicarbonate no precipitate is at
first obtained, but after boiling a precipitate is formed having
the composition of malachite. Debray obtained artificial azurite
by heating cupric nitrate with chalk.
The commonest normal salt is _blue vitriol_--_i.e._ the normal cupric
sulphate. It generally contains five molecules of water of
crystallisation, CuSO_{4},5H_{2}O. It forms the product of the action of
strong sulphuric acid on copper, sulphurous anhydride being evolved. The
same salt is obtained in practice by carefully roasting sulphuretted ores
of copper, and also by the action of water holding oxygen in solution on
them: CuS + O_{4} = CuSO_{4}. This salt forms a by-product, obtained in
gold refineries, when the silver is precipitated from the sulphuric acid
solution by means of copper. It is also obtained by pouring dilute
sulphuric acid over sheet copper in the presence of air, or by heating
cupric oxide or carbonate in sulphuric acid. The crystals of this salt
belong to the triclinic system, have a specific gravity of 2·19, are of a
beautiful blue colour, and give a solution of the same colour. 100 parts
of water at 0° dissolve 15, at 25° 23, and at 100° about 45 parts of
cupric sulphate, CuSO_{4}.[10 bis] At 100° this salt loses a portion of
its water of crystallisation, which it only parts with entirely at a high
temperature (220°) and then gives a white powder of the anhydrous
sulphate; and the latter, on further calcination, loses the elements of
sulphuric anhydride, leaving cupric oxide, like all the cupric salts. The
anhydrous (colourless) cupric sulphate is sometimes used for absorbing
water; it turns blue in the process. It offers the advantage that it
retains both hydrochloric acid and water, but not carbonic anhydride.[11]
Cupric sulphate is used for steeping seed corn; this is said to prevent
the growth of certain parasites on the plants. In the arts a considerable
quantity of cupric sulphate is also used in the preparation of other
copper salts--for instance, of certain pigments[11 bis]--and a
particularly large quantity is used _in the galvanoplastic process_,
which consists in the deposition of copper from a solution of cupric
sulphate by the action of a galvanic current, when the metallic copper is
deposited on the negative pole and takes the shape of the latter. The
description of the processes of galvanoplastic art introduced by Jacobi
in St. Petersburg forms a part of applied physics, and will not be
touched on here, and we will only mention that, although first introduced
for small articles, it is now used for such articles as type moulds
(_clichés_), for maps, prints, &c., and also for large statues, and for
the deposition of iron, zinc, nickel, gold, silver, &c. on other metals
and materials. The beginning of the application of the galvanic current
to the practical extraction of metals from solutions has also been
established, especially since the dynamo-electric machines of Gramme,
Siemens, and others have rendered it possible to cheaply convert the
mechanical motion of the steam engine into an electric current. It is to
be expected that the application of the electric current, which has long
since given such important results in chemistry, will, in the near
future, play an important part in technical processes, the example being
shown by electric lighting.
[10 bis] Although sulphate of copper usually crystallises with 5H_{2}O,
that is, differently to the sulphates of Mg, Fe, and Mn, it is
nevertheless perfectly isomorphous with them, as is seen not only
in the fact that it gives isomorphous mixtures with them,
containing a similar amount of water of crystallisation, but also
in the ease with which it forms, like all bases analogous to MgO,
double salts, R_{2}Cu(SO_{4})_{2},6H_{2}O, where R = K, Rb, Cs, of
the monoclinic system.
Salts of this kind, like CuCl_{2},2KCl,2H_{2}O,PtK_{2}Cy_{4}, &c.,
present a composition CuX_{2} if the representation of double
salts given in Chapter XXIII., Note 11, be admitted, because they,
like Cu(HO)_{2}, contain Cu(X_{2}K)_{2}, where X_{2} = SO_{4},
_i.e._ the residue of sulphuric acid, which combines with H_{2},
and is therefore able to replace the H_{2} by X_{2} or O. A
detailed study of the crystalline forms of these salts, made by
Tutton (1893) (_see_ Chapter XIII., Note 1), showed: (1) that 22
investigated salts of the composition R_{2}M(SO_{4}),6H_{2}O,
where R = K, Rb, Cs, and M = Mg, Zn, Cd, Mn, Fe, Co, Ni, Cu,
present a complete crystallographic resemblance; (2) that in all
respects the Rb salts present a transition between the K and Cs
salts; (3) that the Cs salts form crystals most easily, and the K
salts the most difficultly, and that for the K salts of Cd and Mn
it was even impossible to obtain well-formed crystals; (4) that
notwithstanding the closeness of their angles, the general
appearance (habit) of the potassium compound differs very clearly
from the Cs salts, while the Rb salts present a distinct
transition in this respect; (5) that the angle of the inclination
of one of the axes to the plane of the two other axes showed that
in the K salts (angle from 75° to 75° 38´) the inclination is
least, in the Cs salts (from 72° 52´ to 73° 50´) greatest, and in
the Rb salts (from 73° 57´ to 74° 42´) intermediate between the
two; the replacement of Mg ... Cu produces but a very small change
in this angle; (6) that the other angles and the ratio of the axes
of the crystals exhibit a similar variation; and (7) that thus the
variation of the form is chiefly determined by the atomic weight
of the alkaline metal. As an example we cite the magnitude of the
inclination of the axes of R_{2}M(SO_{4})_{2},6H_{2}O.
R = K Rb Cs
M = Mg 75° 12´ 74° 1´ 72° 54´
Zn 75° 12´ 74° 7´ 72° 59´
Cd -- 74° 7´ 72° 49´
Mn -- 73° 3´ 72° 53´
Fe 75° 28´ 74° 16´ 73° 8´
Co 75° 5´ 73° 59´ 72° 52´
Ni 75° 0´ 73° 57´ 72° 58´
Cu 75° 32´ 74° 42´ 73° 50´
This shows clearly (within the limits of possible error, which may
be as much as 30´) the almost perfect identity of the independent
crystalline forms notwithstanding the difference of the atomic
weights of the diatomic elements, M = Mg, Cu.
[11] In addition to what has been said (Chapter I., Note 65, and
Chapter XXII., Note 35) respecting the combination of CuSO_{4}
with water and ammonia, we may add that Lachinoff (1893) showed
that CuSO_{4},5H_{2}O loses 4-3/4H_{2}O at 180°, that
CuSO_{4},5NH_{3} also loses 4-3/4NH_{3} at 320°, and that only
1/4H_{2}O and 1/4NH_{3} remain in combination with the CuSO_{4}.
The last 1/4H_{2}O can only be driven off by heating to 200°, and
the last 1/4NH_{3} by heating to 360°. Ammonia displaces water
from CuSO_{4},5H_{2}O, but water cannot displace the ammonia from
CuSO_{4},5NH_{3}. If hydrochloric acid gas be passed over
CuSO_{4},5H_{2}O at the ordinary temperature, it first forms
CuSO_{4},5H_{2}O,3HCl, and then CuSO_{4},2H_{2}O,2HCl. When air is
passed over the latter compound it passes into CuSO_{4}H_{2}O with
a small amount of HCl (about 1/8HCl). At 100° CuSO_{4},5H_{2}O in
a stream of hydrochloric acid gas gives CuSO_{4},1/4H_{2}O,2HCl,
and then CuSO_{4},1/4H_{2}O,HCl, whilst after prolonged heating
CuSO_{4} remains, which rapidly passes into CuSO_{4},5H_{2}O when
placed under a bell jar over water. Over sulphuric acid, however,
CuSO_{4},5H_{2}O only parts with 3H_{2}O, and if CuSO_{4},2H_{2}O
be placed over water it again forms CuSO_{4},5H_{2}O, and so on.
[11 bis] Commercial blue vitriol generally contains ferrous sulphate.
The salt is purified by converting the ferrous salt into a ferric
salt by heating the solution with chlorine or nitric acid. The
solution is then evaporated to dryness, and the unchanged cupric
sulphate extracted from the residue, which will contain the larger
portion of the ferric oxide. The remainder will be separated if
cupric hydroxide is added to the solution and boiled; the cupric
oxide, CuO, then precipitates the ferric oxide, Fe_{2}O_{3}, just
as it is itself precipitated by silver oxide. But the solution
will contain a small proportion of a basic salt of copper, and
therefore sulphuric acid must be added to the filtered solution,
and the salt allowed to crystallise. Acid salts are not formed,
and cupric sulphate itself has an acid reaction on litmus paper.
_The alloys of copper_ with certain metals, and especially with zinc and
tin, are easily formed by directly melting the metals together. They are
easily cast into moulds, forged, and worked like copper, whilst they are
much more durable in the air, and are therefore frequently used in the
arts. Even the ancients used exclusively alloys of copper, and not pure
copper, but its alloys with tin or different kinds of bronze (Chapter
XVIII., Note 35). The alloys of copper with zinc are called _brass_ or
'yellow metal.' Brass contains about 32 p.c. of zinc; generally, however,
it does not contain more than 65 p.c. of copper. The remainder is
composed of lead and tin, which usually occur, although in small
quantities, in brass. Yellow metal contains about 40 p.c. of zinc.[12]
The addition of zinc to copper changes the colour of the latter to a
considerable degree; with a certain amount of zinc the colour of the
copper becomes yellow, and with a still larger proportion of zinc an
alloy is formed which has a greenish tint. In those alloys of zinc and
copper which contain a larger amount of zinc than of copper, the yellow
colour disappears and is replaced by a greyish colour. But when the
amount of zinc is diminished to about 20 p.c., the alloy is red and hard,
and is called 'tombac.' A contraction takes place in alloying copper with
zinc, so that the volume of the alloy is less than that of either metal
individually. The zinc volatilises on prolonged heating at a high
temperature and the excess of metallic copper remains behind. When heated
in the air, the zinc oxidises before the copper, so that all the zinc
alloyed with copper may be removed from the copper by this means. An
important property of brass containing about 30 p.c. of zinc is that it
is soft and malleable in the cold, but becomes somewhat brittle when
heated. We may also mention that ordinary copper coins contain, in order
to render them hard, tin, zinc, and iron (Cu = 95 p.c.); that it is now
customary to add a small amount of phosphorus to copper and bronze, for
the same purpose; and also that copper is added to silver and gold in
coining, &c. to render it hard; moreover, in Germany, Switzerland, and
Belgium, and other countries, a silver-white alloy (melchior, German
silver, &c.), for base coinage and other purposes, is prepared from brass
and nickel (from 10 to 20 p.c. of nickel; 20 to 30 p.c. zinc: 50 to 70
p.c. copper), or directly from copper and nickel, or, more rarely, from
an alloy containing silver, nickel, and copper.[12 bis]
[12] Among the alloys of copper resembling brass, _delta metal_,
invented by A. Dick (London) is largely used (since 1883). It
contains 55 p.c. Cu, and 41 p.c. Zn, the remaining 4 p.c. being
composed of iron (as much as 3-1/2 p.c., which is first alloyed
with zinc), or of cobalt, and manganese, and certain other metals.
The sp. gr. of delta metal is 8·4. It melts at 950°, and then
becomes so fluid that it fills up all the cavities in a mould and
forms excellent castings. It has a tensile strength of 70 kilos
per sq. mm. (gun metal about 20, phosphor bronze about 30). It is
very soft, especially when heated to 600°, but after forging and
rolling it becomes very hard; it is more difficultly acted upon by
air and water than other kinds of brass, and preserves its golden
yellow colour for any length of time, especially if well polished.
It is used for making bearings, screw propellers, valves, and many
other articles. In general the alloys of Cu and Zn containing
about 2/3 p.c. by weight of copper were for a long time almost
exclusively made in Sweden and England (Bristol, Birmingham).
These alloys for the most part are cheaper, harder, and more
fusible than copper alone, and form good castings. The alloys
containing 45-80 p.c. Cu crystallise in cubes if slowly cooled (Bi
also gives crystals). By washing the surface of brass with dilute
sulphuric acid, Zn is removed and the article acquires the colour
of copper. The alloys approaching Zn_{2}Cu_{3} in their
composition exhibit the greatest resistance (under other equal
conditions; of purity, forging, rolling, &c.) The addition of 3
p.c. Al, or 5 p.c. Sn, improves the quality of brass. Respecting
aluminium bronze _see_ Chapter XVII. p. 88.
[12 bis] Ball (also Kamensky), 1888, by investigating the electrical
conductivity of the alloys of antimony and copper with lead, came
to the conclusion that only two definite compounds of antimony and
copper exist, whilst the other alloys are either alloys of these
two together or with antimony or with copper. These compounds are
Cu_{2}Sb and Cu_{4}Sb--one corresponds with the maximum, and the
other with the minimum, electrical resistance. In general, the
resistance offered to an electrical current forms one of the
methods by which the composition of definite alloys (for example,
Pb_{2}Zn_{7}) is often established, whilst the electromotive force
of alloys affords (Laurie, 1888) a still more accurate method--for
instance, several definite compounds were discovered by this
method among the alloys of copper with zinc and tin; but we will
not enter into any details of this subject, because we avoid all
references to electricity, although the reader is recommended to
make himself acquainted with this branch of science, which has
many points in common with chemistry. The study of alloys regarded
as solid solutions should, in my opinion, throw much light upon
the question of solutions, which is still obscure in many aspects
and in many branches of chemistry.
Copper, in its cuprous compounds, is so analogous to _silver_, that were
there no cupric compounds, or if silver gave stable compounds of the
higher oxide, AgO, the resemblance would be as close as that between
chlorine and bromine or zinc and cadmium; but silver compounds
corresponding to AgO are quite unknown. Although silver peroxide--which
was regarded as AgO, but which Berthelot (1880) recognised as the
sesquioxide Ag_{2}O_{3}--is known, still it does not form any true salts,
and consequently cannot be placed along with cupric oxide. In distinction
to copper, silver as a metal does not oxidise under the influence of
heat; and its oxides, Ag_{2}O and Ag_{2}O_{3}, easily lose oxygen (_see_
Note 8 tri). Silver does _not oxidise_ in air at the ordinary pressure,
and is therefore classed among the so-called _noble metals_. It has a
white colour, which is much purer than that of any other known metal,
especially when the metal is chemically pure. In the arts silver is
always used alloyed, because chemically-pure silver is so soft that it
wears exceedingly easily, whilst when fused with a small amount of
copper, it becomes very hard, without losing its colour.[13]
[13] There are not many soft metals; lead, tin, copper, silver, iron,
and gold are somewhat soft, and potassium and sodium very soft.
The metals of the alkaline earths are sonorous and hard, and many
other metals are even brittle, especially bismuth and antimony.
But the very slight significance which these properties have in
determining the fundamental chemical properties of substances
(although, however, of immense importance in the practical
applications of metals) is seen from the example shown by zinc,
which is hard at the ordinary temperature, soft at 100°, and
brittle at 200°.
[Illustration: FIG. 95.--Cupel for silver assaying.]
[Illustration: FIG. 96.--Clay muffle.]
[Illustration: FIG. 97.--Portable muffle furnace.]
As the value of silver depends exclusively on its purity, and as
there is no possibility of telling the amount of impurities
alloyed with it from its external appearance, it is customary in
most countries to mark an article with the amount of pure silver
it contains after an accurately-made analysis known as the assay
of the silver. In France the assay of silver shows the amount of
pure silver in 1,000 parts by weight; in Russia the amount of pure
silver in 96 parts--that is, the assay shows the number of
zolotniks (4·26 grams) of pure silver in one pound (410 grams) of
alloyed silver. Russian silver is generally 84 assay--that is,
contains 84 parts by weight of pure silver and 12 parts of copper
and other metals. French money contains 90 p.c. (in the Russian
system this will be 86·4 assay) by weight of silver [English coins
and jewellery contain 92·5 p.c. of silver]; the silver rouble is
of 83-1/3 assay--that is, it contains 86·8 p.c. of silver--and the
smaller Russian silver coinage is of 48 assay, and therefore
contains 50 p.c. of silver. Silver ornaments and articles are
usually made in Russia of 84 and 72 assay. As the alloys of silver
and copper, especially after being subjected to the action of
heat, are not so white as pure silver, they generally undergo a
process known as 'blanching' (or 'pickling') after being worked
up. This consists in removing the copper from the surface of the
article by subjecting it to a dark-red heat and then immersing it
in dilute acid. During the calcination the copper on the surface
is oxidised, whilst the silver remains unchanged; the dilute acid
then dissolves the copper oxides formed, and pure silver is left
on the surface. The surface is dull after this treatment, owing to
the removal of a portion of the metal by the acid. After being
polished the article acquires the desired lustre and colour, so as
to be indistinguishable from a pure silver object. In order to
test a silver article, a portion of its mass must be taken, not
from the surface, but to a certain depth. The methods of assay
used in practice are very varied. The commonest and most often
used is that known as _cupellation_. It is based on the difference
in the oxidisability of copper, lead, and silver. The cupel is a
porous cup with thick sides, made by compressing bone ash. The
porous mass of bone ash absorbs the fused oxides, especially the
lead oxide, which is easily fusible, but it does not absorb the
unoxidised metal. The latter collects into a globule under the
action of a strong heat in the cupel, and on cooling solidifies
into a button, which may then be weighed. Several cupels are
placed in a muffle. A muffle is a semi-cylindrical clay vessel,
shown in the accompanying drawing. The sides of the muffle are
pierced with several orifices, which allow the access of air into
it. The muffle is placed in a furnace, where it is strongly
heated. Under the action of the air entering the muffle the copper
of the silver alloy is oxidised, but as the oxide of copper is
infusible, or, more strictly speaking, difficultly fusible, a
certain quantity of lead is added to the alloy; the lead is also
oxidised by the air at the high temperature of the muffle, and
gives the very fusible lead oxide. The copper oxide then fuses
with the lead oxide, and is absorbed by the cupel, whilst the
silver remains as a bright white globule. If the weight of the
alloy taken and of the silver left on the cupel be determined, it
is possible to calculate the composition of the alloy. Thus the
essence of cupellation consists in the separation of the
oxidisable metals from silver, which does not oxidise under the
action of heat. A more accurate method, based on the precipitation
of silver from its solutions in the form of silver chloride, is
described in detail in works on analytical chemistry.
Silver occurs in _nature_, both in a native state and in certain
compounds. Native silver, however, is of rather rare occurrence. A far
greater quantity of silver occurs in combination with sulphur, and
especially in the form of _silver sulphide_, Ag_{2}S, with lead sulphide
or copper sulphide, or the ores of various other metals. The largest
amount of silver is extracted from the lead in which it occurs. If this
lead be calcined in the presence of air, it oxidises, and the resultant
lead oxide, PbO ('litharge' or 'silberglätte,' as it is called), melts
into a mobile liquid, which is easily removed. The silver remains in an
unoxidised metallic state.[14] This process is called _cupellation_.
[14] In America, whence the largest amount of silver is now obtained,
ores are worked containing not more than 1/5 p.c. of silver,
whilst at 1/2 p.c. its extraction is very profitable. Moreover,
the extraction of silver from ores containing not more than 0·01
p.c. of this metal is sometimes profitable. The majority of the
lead smelted from galena contains silver, which is extracted from
it. Thus near Arras, in France, an ore is worked which contains
about 65 parts of lead and 0·088 part of silver in 100 parts of
ore, which corresponds with 136 parts of silver in 100,000 parts
of lead. At Freiberg, in Saxony, the ore used (enriched by
mechanical dressing) contains about 0·9 of silver, 160 of lead,
and 2 of copper in 10,000 parts. In every case the lead is first
extracted in the manner described in Chapter XVIII., and this lead
will contain all the silver. Not unfrequently other ores of silver
are mixed with lead ores, in order to obtain an argentiferous lead
as the product. The extraction of small quantities of silver from
lead is facilitated by the fact (Pattinson's process) that molten
argentiferous lead in cooling first deposits crystals of pure
lead, which fall to the bottom of the cooling vessel, whilst the
proportion of silver in the unsolidified mass increases owing to
the removal of the crystals of lead. The lead is enriched in this
manner until it contains 1/400 part of silver, and is then
subjected to cupellation on a larger scale. According to Park's
process, zinc is added to the molten argentiferous lead, and the
alloy of Pb and Zn, which first separates out on cooling, is
collected. This alloy is found to contain all the silver
previously contained in the lead. The addition of 0·5 p.c. of
aluminium to the zinc (Rossler and Edelman) facilitates the
extraction of the Ag from the resultant alloy besides preventing
oxidation; for, after re-melting, nearly all the lead easily runs
off (remains fluid), and leaves an alloy containing about 30 p.c.
Ag and about 70 p.c. Zn. This alloy may be used as an anode in a
solution of ZnCl_{2}, when the Zn is deposited on the cathode,
leaving the silver with a small amount of Pb, &c. behind. The
silver can be easily obtained pure by treating it with dilute
acids and cupelling.
The ores of silver which contain a larger amount of it are: silver
glance, Ag_{2}S (sp. gr. 7·2); argentiferous-copper glance, CuAgS;
horn silver or chloride of silver, AgCl; argentiferous grey copper
ore; polybasite, M_{9}RS_{6} (where M = Ag, Cu, and R = Sb, As),
and argentiferous gold. The latter is the usual form in which gold
is found in alluvial deposits and ores. The crystals of gold from
the Berezoffsky mines in the Urals contain 90 to 95 of gold and 5
to 9 of silver, and the Altai gold contains 50 to 65 of gold and
36 to 38 of silver. The proportion of silver in native gold varies
between these limits in other localities. Silver ores, which
generally occur in veins, usually contain native silver and
various sulphur compounds. The most famous mines in Europe are in
Saxony (Freiberg), which has a yearly output of as much as 26 tons
of silver, Hungary, and Bohemia (41 tons). In Russia, silver is
extracted in the Altai and at Nerchinsk (17 tons). The richest
silver mines known are in America, especially in Chili (as much as
70 tons), Mexico (200 tons), and more particularly in the Western
States of North America. The richness of these mines may be judged
from the fact that one mine in the State of Nevada (Comstock, near
Washoe and the cities of Gold Hill and Virginia), which was
discovered in 1859, gave an output of 400 tons in 1866. In place
of cupellation, chlorination may also be employed for extracting
silver from its ores. The method of chlorination consists in
converting the silver in an ore into silver chloride. This is
either done by a wet or by a dry method, roasting the ore with
NaCl. When the silver chloride is formed, the extraction of the
metal is also done by two methods. The first consists in the
silver chloride being reduced to metal by means of iron in
rotating barrels, with the subsequent addition of mercury which
dissolves the silver, but does not act on the other metals. The
mercury holding the silver in solution is distilled, when the
silver remains behind. This method is called _amalgamation_. The
other method is less frequently used, and consists in dissolving
the silver chloride in sodium chloride or in sodium thiosulphate,
and then precipitating the silver from the solution. The
amalgamation is then carried on in rotating barrels containing the
roasted ore mixed with water, iron, and mercury. The iron reduces
the silver chloride by taking up the chlorine from it. The
technical details of these processes are described in works on
metallurgy. The extraction of AgCl by the wet method is carried on
(Patera's process) by means of a solution of hyposulphite of
sodium which dissolves AgCl (_see_ Note 23), or by lixiviating
with a 2 p.c. solution of a double hyposulphite of Na and Cu
(obtained by adding CuSO_{4} to Na_{2}S_{2}O_{3}). The resultant
solution of AgCl is first treated with soda to precipitate
PbCO_{3}, and then with Na_{2}S, which precipitates the Ag and Au.
The process should be carried on rapidly to prevent the
precipitation of Cu_{2}S from the solution of CuSO_{4} and
Na_{2}S_{2}O_{3}.
Commercial silver generally contains copper, and, more rarely, other
metallic impurities also. Chemically _pure silver_ is obtained either by
cupellation or by subjecting ordinary silver to the following treatment.
The silver is first dissolved in nitric acid, which converts it and the
copper into nitrates, Cu(NO_{3})_{2} and AgNO_{3}; hydrochloric acid is
then added to the resultant solution (green, owing to the presence of the
cupric salt), which is considerably diluted with water in order to retain
the lead chloride in solution if the silver contained lead. The copper
and many other metals remain in solution, whilst the silver is
precipitated as silver chloride. The precipitate is allowed to settle,
and the liquid is decanted off; the precipitate is then washed and fused
with sodium carbonate. A double decomposition then takes place, sodium
chloride and silver carbonate being formed; but the latter decomposes
into metallic silver, because the silver oxide is decomposed by heat:
Ag_{2}CO_{3} = Ag_{2} + O + CO_{2}. The silver chloride may also be mixed
with metallic zinc, sulphuric acid, and water, and left for some time,
when the zinc removes the chlorine from the silver chloride and
precipitates the silver as a powder. This finely-divided silver is called
'molecular silver.'[15]
[15] There is another practical method which is also suitable for
separating the silver from the solutions obtained in photography,
and consists in precipitating the silver by oxalic acid. In this
case the amount of silver in the solution must be known, and 23
grams of oxalic acid dissolved in 400 grams of water must be added
for every 60 grams of silver in solution in a litre of water. A
precipitate of silver oxalate, Ag_{2}C_{2}O_{4}, is then obtained,
which is insoluble in water but soluble in acids. Hence, if the
liquid contain any free acid it must be previously freed from it
by the addition of sodium carbonate. The resultant precipitate of
silver oxalate is dried, mixed with an equal weight of dry sodium
carbonate, and thrown into a gently-heated crucible. The
separation of the silver then proceeds without an explosion,
whilst the silver oxalate if heated alone decomposes with
explosion.
According to Stas, the best method for obtaining silver from its
solutions is by the reduction of silver chloride dissolved in
ammonia by means of an ammoniacal solution of cuprous
thiosulphate; the silver is then precipitated in a crystalline
form. A solution of ammonium sulphite may be used instead of the
cuprous salt.
Chemically-pure silver has an exceeding pure white colour, and a specific
gravity of 10·5. Solid silver is lighter than the molten metal, and
therefore a piece of silver floats on the latter. The fusing-point of
silver is about 950° C., and at the high temperature attained by the
combustion of detonating gas it volatilises.[16] By employing silver
reduced from silver chloride by milk sugar and caustic potash, and
distilling it, Stas obtained silver purer than that obtained by any other
means; in fact, this was perfectly pure silver. The vapour of silver has
a very beautiful green colour, which is seen when a silver wire is placed
in an oxyhydrogen flame.[17]
[16] Silver is very malleable and ductile; it may be beaten into leaves
0·002 mm. in thickness. Silver wire may be made so fine that 1
gram is drawn into a wire 2-1/2 kilometres long. In this respect
silver is second only to gold. A wire of 2 mm. diameter breaks
under a strain of 20 kilograms.
[17] In melting, silver absorbs a considerable amount of oxygen, which
is disengaged on solidifying. One volume of molten silver absorbs
as much as 22 volumes of oxygen. In solidifying, the silver forms
cavities like the craters of a volcano, and throws off metal,
owing to the evolution of the gas; all these phenomena recall a
volcano on a miniature scale (Dumas). Silver which contains a
small quantity of copper or gold, &c., does not show this property
of dissolving oxygen.
The absorption of oxygen by molten silver is, however, an
oxidation, but it is at the same time a phenomenon of solution.
One cubic centimetre of molten silver can dissolve twenty-two
cubic centimetres of oxygen, which, even at 0°, only weighs 0·03
gram, whilst 1 cubic centimetre of silver weighs at least 10
grams, and therefore it is impossible to suppose that the
absorption of the oxygen is attended by the formation of any
definite compound (rich in oxygen) of silver and oxygen (about 45
atoms of silver to 1 of oxygen) in any other but a dissociated
form, and this is the state in which substances in solution must
be regarded (Chapter I.)
Le Chatelier showed that at 300° and 15 atmospheres pressure
silver absorbs so much oxygen that it may be regarded as having
formed the compound Ag_{4}O, or a mixture of Ag_{2} and Ag_{2}O.
Moreover, silver oxide, Ag_{2}O, only decomposes at 300° under low
pressures, whilst at pressures above 10 atmospheres there is no
decomposition at 300° but only at 400°.
Stas showed that silver is oxidised by air in the presence of
acids. V. d. Pfordten confirmed this, and showed that an acidified
solution of potassium permanganate rapidly dissolves silver in the
presence of air.
It has long been known (Wöhler) that when nitrate of silver, AgNO_{3},
reacts as an oxidising agent upon citrates and tartrates, it is able
under certain conditions to give either a salt of suboxide of silver (see
Note 19) or a red solution, or to give a precipitate of metallic silver
reduced at the expense of the organic substances. In 1889 Carey Lea, in
his researches on this class of reactions, showed that _soluble silver_
is here formed, which he called _allotropic silver_. It may be obtained
by taking 200 c.c. of a 10 per cent. solution of AgNO_{3} and quickly
adding a mixture (neutralised with NaHO) of 200 c.c. of a 30 per cent.
solution of FeSO_{4} and 200 c.c. of a 40 per cent. solution of sodium
citrate. A lilac precipitate is obtained, which is collected on a filter
(the precipitate becomes blue) and washed with a solution of
NH_{4}NO_{3}. It then becomes soluble in pure water, forming a red
perfectly transparent solution from which the dissolved silver is
precipitated on the addition of many soluble foreign bodies. Some of the
latter--for instance, NH_{4}NO_{3}, alkaline sulphates, nitrates, and
citrates--give a precipitate which redissolves in pure water, whilst
others--for instance, MgSO_{4}, FeSO_{4}, K_{2}Cr_{2}O_{7}, AgNO_{3},
Ba(NO_{3})_{2} and many others--convert the precipitated silver into a
new variety, which, although no longer soluble in water, regains its
solubility in a solution of borax and is soluble in ammonia. Both the
soluble and insoluble silver are rapidly converted into the ordinary
grey-metallic variety by sulphuric acid, although nothing is given off in
the reaction; the same change takes place on ignition, but in this case
CO_{2} is disengaged; the latter is formed from the organic substances
which remain (to the amount of 3 per cent.) in the modified silver (they
are not removed by soaking in alcohol or water). If the precipitated
silver be slightly washed and laid in a smooth thin layer on paper or
glass, it is seen that the soluble variety is red when moist and a fine
blue colour when dry, whilst the insoluble variety has a blue reflex.
Besides these, under special conditions[18] a golden yellow variety may
be obtained, which gives a brilliant golden yellow coating on glass; but
it is easily converted into the ordinary grey-metallic state by friction
or trituration. There is no doubt[18 bis] that there is the same relation
between ordinary silver which is perfectly insoluble in water and the
varieties of silver obtained by Carey Lea[18 tri] as there is between
quartz and soluble silica or between CuS and As_{2}S_{2} in their
ordinary insoluble forms and in the state of the colloid solution of
their hydrosols (_see_ Chapter I., Note 57, and Chapter XVII., Note 25
bis). Here, however, an important step in advance has been made in this
respect, that we are dealing with the solution of a simple body, and
moreover of a metal--_i.e._ of a particularly characteristic state of
matter. And as boron, gold, and certain other simple bodies have already
been obtained in a soluble (colloid) form, and as numerous organic
compounds (albuminous substances, gum, cellulose, starch, &c.) and
inorganic substances are also known in this form, it might be said that
the colloid state (of hydrogels and hydrosols) can be acquired, if not by
every substance, at all events by substances of most varied chemical
character under particular conditions of formation from solutions. And
this being the case, we may hope that a further study of soluble colloid
compounds, which apparently present various transitions towards
emulsions, may throw a new light upon the complex question of solutions,
which forms one of the problems of the present epoch of chemical science.
Moreover, we may remark that Spring (1890) clearly proved the colloid
state of soluble silver by means of dialysis as it did not pass through
the membrane.
[18] When solutions of AgNO_{3}, FeSO_{4}, sodium citrate, and NaHO are
mixed together in the manner described above, they throw down a
precipitate of a beautiful lilac colour; when transferred to a
filter paper the precipitate soon changes colour, and becomes dark
blue. To obtain the substance as pure as possible it is washed
with a 5-10 p.c. solution of ammonium nitrate; the liquid is
decanted, and 150 c.c. of water poured over the precipitate. It
then dissolves entirely in the water. A small quantity of a
saturated solution of ammonium nitrate is added to the solution,
and the silver in solution again separates out as a precipitate.
These alternate solutions and precipitations are repeated seven or
eight times, after which the precipitate is transferred to a
filter and washed with 95 p.c. alcohol until the filtrate gives no
residue on evaporation. An analysis of the substance so obtained
showed that it contained from 97·18 p.c. to 97·31 p.c. of metallic
silver. It remained to discover what the remaining 2-3 p.c. were
composed of. Are they merely impurities, or is the substance some
compound of silver with oxygen or hydrogen, or does it contain
citric acid in combination which might account for its solubility?
The first supposition is set aside by the fact that no gases are
disengaged by the precipitate of silver, either under the action
of gases or when heated. The second supposition is shown to be
impossible by the fact that there is no definite relation between
the silver and citric acid. A determination of the amount of
silver in solution showed that the amount of citric acid varies
greatly for one and the same amount of silver, and there is no
simple ratio between them. Among other methods of preparing
soluble silver given by Carey Lea, we may mention the method
published by him in 1891. AgNO_{3} is added to a solution of
dextrine in caustic soda or potash; at first a precipitate of
brown oxide of silver is thrown down, but the brown colour then
changes into a reddish chocolate, owing to the reduction of the
silver by the dextrine, and the solution turns a deep red. A few
drops of this solution turn water bright red, and give a perfectly
transparent liquid. The dextrine solution is prepared by
dissolving 40 grams of caustic soda and the same amount of
ordinary brown dextrine in two litres of water. To this solution
is gradually added 28 grams of AgNO_{3} dissolved in a small
quantity of water.
The insoluble allotropic silver is obtained, as was mentioned
above, from a solution of silver prepared in the manner described,
by the addition of sulphate of copper, iron, barium, magnesium,
&c. In one experiment Lea succeeded in obtaining the insoluble
allotropic Ag in a crystalline form. The red solution, described
above, after standing several weeks, deposits crystals
spontaneously in the form of short black needles and thin prisms,
the liquid becoming colourless. This insoluble variety, when
rubbed upon paper, has the appearance of bright shining green
flakes, which polarise light.
The gold variety is obtained in a different manner to the two
other varieties. A solution is prepared containing 200 c.c. of a
10 p.c. solution of nitrate of silver, 200 c.c. of a 20 p.c.
solution of Rochelle salt, and 800 c.c. of water. Just as in the
previous case the reaction consisted in the reduction of the
citrate of silver, so in this case it consists in the reduction of
the tartrate, which here first forms a red, and then a black
precipitate of allotropic Ag, which, when transferred to the
filter, appears of a beautiful bronze colour. After washing and
drying, this precipitate acquires the lustre and colour peculiar
to polished gold, and this is especially remarked where the
precipitate comes into contact with glass or china. An analysis of
the golden variety gave a percentage composition of 98·750 to
98·749 Ag. Both the insoluble varieties (the blue and gold) have a
different specific gravity from ordinary silver. Whilst that of
fused silver is 10·50, and of finely-divided silver 10·62, the
specific gravity of the blue insoluble variety is 9·58, and of the
gold variety 8·51. The gold variety passes into ordinary Ag with
great ease. This transition may even be remarked on the filter in
those places which have accidentally not been moistened with
water. A simple shock, and therefore friction of one particle upon
another, is enough to convert the gold variety into normal white
silver. Carey Lea sent samples of the gold variety for a long
distance by rail packed in three tubes, in which the silver
occupied about the quarter of their volume; in one tube only he
filled up this space with cotton-wool. It was afterwards found
that the shaking of the particles of Ag had completely converted
it into ordinary white silver, and that only the tube containing
the cotton-wool had preserved the golden variety intact.
The soluble variety of Ag also passes into the ordinary state with
great ease, the heat of conversion being, as Prange showed in
1890, about +60 calories.
[18 bis] The opinion of the nature of soluble silver given below was
first enunciated in the _Journal of the Russian Chemical Society_,
February 1, 1890, Vol. XXII., Note 73. This view is, at the
present time, generally accepted, and this silver is frequently
known as the 'colloid' variety. I may add that Carey Lea observed
the solution of ordinary molecular silver in ammonia without the
access of air.
[18 tri] It is, however, noteworthy that ordinary metallic lead has
long been considered soluble in water, that boron has been
repeatedly obtained in a brown solution, and that observations
upon the development of certain bacteria have shown that the
latter die in water which has been for some time in contact with
metals. This seems to indicate the passage of small quantities of
metals into water (however, the formation of peroxide of hydrogen
may be supposed to have some influence in these cases).
As regards the capacity of silver for chemical reactions, it is
remarkable for its small capacity for combination with oxygen and for its
considerable energy of combination with sulphur, iodine, and certain
kindred non-metals. _Silver does not oxidise_ at any temperature, and its
oxide, Ag_{2}O, is decomposed by heat. It is also a very important fact
that silver is not oxidised by oxygen either in the presence of alkalis,
even at exceedingly high temperatures, or in the presence of acids--at
least, of dilute acids--which properties render it a very important metal
in chemical industry for the fusion of alkalis, and also for many
purposes in everyday life; for instance, for making spoons, salt-cellars,
&c. Ozone, however, oxidises it. Of all acids nitric acid has the
greatest action on silver. The reaction is accompanied by the formation
of oxides of nitrogen and silver nitrate, AgNO_{3}, which dissolves in
water and does not, therefore, hinder the further action of the acid on
the metal. The halogen acids, especially hydriodic acid, act on silver,
hydrogen being evolved; but this action soon stops, owing to the halogen
compounds of silver being insoluble in water and only very slightly
soluble in acids; they therefore preserve the remaining mass of metal
from the further action of the acid; in consequence of this the action of
the halogen acids is only distinctly seen with finely-divided silver.
Sulphuric acid acts on silver in the same manner that it does on copper,
only it must be concentrated and at a higher temperature. Sulphurous
anhydride, and not hydrogen, is then evolved, but there is no action at
the ordinary temperature, even in the presence of air. Among the various
salts, sodium chloride (in the presence of moisture, air, and carbonic
acid) and potassium cyanide (in the presence of air) act on silver more
decidedly than any others, converting it respectively into silver
chloride and a double cyanide.
Although silver does not directly combine with oxygen, still three
different grades of combination with oxygen may be obtained indirectly
from the salts of silver. They are all, however, unstable, and decompose
into oxygen and metallic silver when ignited. These three oxides of
silver have the following composition: _silver suboxide_, Ag_{4}O,[19]
corresponding with the (little investigated) suboxides of the alkali
metals; _silver oxide_, Ag_{2}O, corresponding with the oxides of the
alkali metals and the ordinary salts of silver, AgX; and _silver
peroxide_, AgO,[19 bis] or, judging from Berthelot's researches,
Ag_{2}O_{3}. _Silver oxide_ is obtained as a brown precipitate (which
when dried does not contain water) by adding potassium hydroxide to a
solution of a silver salt--for example, of silver nitrate. The
precipitate formed seems, however, to be an hydroxide, AgHO, _i.e._
AgNO_{3} + KHO = KNO_{3} + AgHO, and the formation of the anhydrous
oxide, 2AgHO = Ag_{2}O + H_{2}O, may be compared with the formation of
the anhydrous cupric oxide by the action of potassium hydroxide on hot
cupric solutions. Silver hydroxide decomposes into water and silver
oxide, even at low temperatures; at least, the hydroxide no longer exists
at 60°, but forms the anhydrous oxide, Ag_{2}O.[19 tri] Silver oxide is
almost insoluble in water; but, nevertheless, it is undoubtedly a rather
powerful basic oxide, because it displaces the oxides of many metals from
their soluble salts, and saturates such acids as nitric acid, forming
with them neutral salts, which do not act on litmus paper.[20]
Undoubtedly water dissolves a small quantity of silver oxide, which
explains the possibility of its action on solutions of salts--for
example, on solutions of cupric salts. Water in which silver oxide is
shaken up has a distinctly alkaline reaction. The oxide is distinguished
by its great instability when heated, so that it loses all its oxygen
when slightly heated. Hydrogen reduces it at about 80°.[20 bis] The
feebleness of the affinity of silver for oxygen is shown by the fact that
silver oxide decomposes under the action of light, so that it must be
kept in opaque vessels. The silver _salts_ are colourless and decompose
when heated, leaving metallic silver if the elements of the acid are
volatile.[20 tri] They have a peculiar metallic taste, and are
exceedingly poisonous; the majority of them are acted on by light,
especially in the presence of organic substances, which are then
oxidised. The alkaline carbonates give a white precipitate of silver
carbonate, Ag_{2}CO_{3}, which is insoluble in water, but soluble in
ammonia and ammonium carbonate. Aqueous ammonia, added to a solution of a
normal silver salt, first acts like potassium hydroxide, but the
precipitate dissolves in an excess of the reagent, like the precipitate
of cupric hydroxide.[21] Silver oxalate and the halogen compounds of
silver are insoluble in water; hydrochloric acid and soluble chlorides
give, as already repeatedly observed, a white precipitate of silver
chloride in solutions of silver salts. Potassium iodide gives a yellowish
precipitate of silver iodide. Zinc separates all the silver in a metallic
form from solutions of silver salts. Many other metals and reducing
agents--for example, organic substances--also reduce silver from the
solutions of its salts.
[19] Silver suboxide (Ag_{4}O) or argentous oxide is obtained from
argentic citrate by heating it to 100° in a stream of hydrogen.
Water and argentous citrate are then formed, and the latter,
although but slightly soluble in water, gives a reddish-brown
solution of colloid silver (Note 18), and when boiled this
solution becomes colourless and deposits metallic silver, the
argentic salt being again formed. Wöhler, who discovered this
oxide, obtained it as a black precipitate by adding potassium
hydroxide to the above solution of argentous citrate. With
hydrochloric acid the suboxide gives a brown compound, Ag_{2}Cl.
Since the discovery of soluble silver the above data cannot be
regarded as perfectly trustworthy; it is probable that a mixture
of Ag_{2} and Ag_{2}O was being dealt with, so that the actual
existence of Ag_{4}O is now doubtful, but there can be no doubt as
to the formation of a subchloride, Ag_{2}Cl (_see_ Note 25),
corresponding to the suboxide. The same compound is obtained by
the action of light on the higher chloride. Other acids do not
combine with silver suboxide, but convert it into an argentic salt
and metallic silver. In this respect cuprous oxide presents a
certain resemblance to these suboxides. But copper forms a
suboxide of the composition Cu_{4}O, which is obtained by the
action of an alkaline solution of stannous oxide on cupric
hydroxide, and is decomposed by acids into cupric salts and
metallic copper. The problems offered by the suboxides, as well as
by the peroxides, cannot be considered as fully solved.
[19 bis] _Silver peroxide_, AgO or Ag_{2}O_{3}, is obtained by the
decomposition of a dilute (10 p.c.) solution of silver nitrate by
the action of a galvanic current (Ritter). On the positive pole,
where oxygen is usually evolved in the decomposition of salts,
brittle grey needles with a metallic lustre, which occasionally
attain a somewhat considerable size, are then formed. They are
insoluble in water, and decompose with the evolution of oxygen
when they are dried and heated, especially up to 150°, and, like
lead dioxide, barium peroxide, &c., their action is strongly
oxidising. When treated with acids, oxygen is evolved and a salt
of the oxide formed. Silver peroxide absorbs sulphurous anhydride
and forms silver sulphate. Hydrochloric acid evolves chlorine;
ammonia reduces the silver, and is itself oxidised, forming water
and gaseous nitrogen. Analyses of the above-mentioned crystals
show that they contain silver nitrate, peroxide, and water.
According to Fisher, they have the composition
4AgO,AgNO_{3},H_{2}O, and, according to Berthelot,
4Ag_{2}O_{5},2AgNO_{3},H_{2}O.
[19 tri] According to Carey Lea, however, oxide of silver still retains
water even at 100°, and only parts with it together with the
oxygen. Oxide of silver is used for colouring glass yellow.
[20] The reaction of Pb(OH)_{2} upon AgHO in the presence of NaHO leads
to the formation of a compound of both oxides, PbO_n_Ag_{2}O, from
which the oxide of lead cannot be removed by alkalies (Wöhler,
Leton). Wöhler, Welch, and others obtained crystalline double
salts, R_{2}AgX_{3}, by the action of strong solutions of RX of
the halogen salts of the alkaline metals upon AgX, where R = Cs,
Rb, K.
[20 bis] According to Müller, ferric oxide is reduced by hydrogen
(_see_ Chapter XXII., Note 5) at 295° (into what ?), cupric oxide
at 140°, Ni_{2}O_{3} at 150°; nickelous oxide, NiO, is reduced to
the suboxide, Ni_{2}O, at 195°, and to nickel at 270°; zinc oxide
requires so high a temperature for its reduction that the glass
tube in which Müller conducted the experiment did not stand the
heat; antimony oxide requires a temperature of 215° for its
reduction; yellow mercuric oxide is reduced at 130° and the red
oxide at 230°; silver oxide at 85°, and platinum oxide even at the
ordinary temperature.
[20 tri] A silica compound, Ag_{2}OSiO_{2} is obtained by fusing
AgNO_{3} with silica; this salt is able to decompose with the
evolution of oxygen, leaving Ag + SiO_{2}.
[21] If a solution of a silver salt be precipitated by sodium
hydroxide, and aqueous ammonia is added drop by drop until the
precipitate is completely dissolved, the liquid when evaporated
deposits a violet mass of crystalline silver oxide. If moist
silver oxide be left in a strong solution of ammonia it gives a
black mass, which easily decomposes with a loud explosion,
especially when struck. This black substance is called fulminating
silver. Probably this is a compound like the other compounds of
oxides with ammonia, and in exploding the oxygen of the silver
oxide forms water with the hydrogen of the ammonia, which is
naturally accompanied by the evolution of heat and formation of
gaseous nitrogen, or, as Raschig states, fulminating silver
contains NAg_{3} or one of the amides (for instance, NHAg_{2} =
NH_{3} + Ag_{2}O - H_{2}O). Fulminating silver is also formed when
potassium hydroxide is added to a solution of silver nitrate in
ammonia. The dangerous explosions which are produced by this
compound render it needful that great care be taken when salts of
silver come into contact with ammonia and alkalis (_see_ Chapter
XVI., Note 26).
_Silver nitrate_, AgNO_{3}, is known by the name of lunar caustic (or
_lapis infernalis_); it is obtained by dissolving metallic silver in
nitric acid. If the silver be impure, the resultant solution will contain
a mixture of the nitrates of copper and silver. If this mixture be
evaporated to dryness and the residue carefully fused at an incipient red
heat, all the cupric nitrate is decomposed, whilst the greater part of
the silver nitrate remains unchanged. On treating the fused mass with
water the latter is dissolved, whilst the cupric oxide remains insoluble.
If a certain amount of silver oxide be added to the solution containing
the nitrates of silver and copper, it displaces all the cupric oxide. In
this case it is of course not necessary to take pure silver oxide, but
only to pour off some of the solution and to add potassium hydroxide to
one portion, and to mix the resultant precipitate of the hydroxides,
Cu(OH)_{2} and AgOH, with the remaining portion.[22] By these methods all
the copper can be easily removed and pure silver nitrate obtained (its
solution is colourless, while the presence of Cu renders it blue), which
may be ultimately purified by crystallisation. It crystallises in
colourless transparent prismatic plates, which are not acted on by air.
They are anhydrous. Its sp. gr. is 4·34; it dissolves in half its weight
of water at the ordinary temperature.[22 bis] The pure salt is not acted
on by light, but it easily acts in an oxidising manner on the majority of
organic substances, which it generally blackens. This is due to the fact
that the organic substance is oxidised by the silver nitrate, which is
reduced to metallic silver; the latter is thus obtained in a
finely-divided state, which causes the black stain. This peculiarity is
taken advantage of for marking linen. Silver nitrate is for the same
reason used for _cauterising wounds_ and various growths on the body.
Here again it acts by virtue of its oxidising capacity in destroying the
organic matter, which it oxidises, as is seen from the separation of a
coating of black metallic powdery silver from the part cauterised.[22
tri] From the description of the preparation of silver nitrate it will
have been seen that this salt, which fuses at 218°, does not decompose at
an incipient red heat; when cast into sticks it is usually employed for
cauterising. On further heating, the fused salt undergoes decomposition,
first forming silver nitrite and then metallic silver. With ammonia,
silver nitrate forms, on evaporation of the solution, colourless crystals
containing AgNO_{3},2HN_{3} (Marignac). In general the salts of silver,
like cuprous, cupric, zinc, &c. salts, are able to give several compounds
with ammonia; for example, silver nitrate in a dry state absorbs three
molecules (Rose). The ammonia is generally easily expelled from these
compounds by the action of heat.
[22] So that we here encounter the following phenomena: copper
displaces silver from the solutions of its salts, and silver oxide
displaces copper oxide from cupric salts. Guided by the
conceptions enunciated in Chapter XV., we can account for this in
the following manner: The atomic volume of silver = 10·3, and of
copper = 7·2, of silver oxide = 32, and of copper oxide = 13. A
greater contraction has taken place in the formation of cupric
oxide, CuO, than in the formation of silver oxide, Ag_{2}O, since
in the former (13 - 7 = 6) the volume after combination with the
oxygen has increased by very little, whilst the volume of silver
oxide is considerably greater than that of the metal it contains
[32 - (2 × 10·3) = 11·4]. Hence silver oxide is less compact than
cupric oxide, and is therefore less stable; but, on the other
hand, there are greater intervals between the atoms in silver
oxide than in cupric oxide, and therefore silver oxide is able to
give more stable compounds than those of copper oxide. This is
verified by the figures and data of their reactions. It is
impossible to calculate for cupric nitrate, because this salt has
not yet been obtained in an anhydrous state; but the sulphates of
both oxides are known. The specific gravity of copper sulphate in
an anhydrous state is 3·53, and of silver sulphate 5·36; the
molecular volume of the former is 45, and of the latter 58. The
group SO_{3} in the copper occupies, as it were, a volume 45 - 13
= 32, and in the silver salt a volume 58 - 32 = 26; hence a
smaller contraction has taken place in the formation of the copper
salt from the oxide than in the formation of the silver salt, and
consequently the latter should be more stable than the former.
Hence silver oxide is able to decompose the salt of copper oxide,
whilst with respect to the metals both salts have been formed with
an almost identical contraction, since 58 volumes of the silver
salt contain 21 volumes of metal (difference = 37), and 45 volumes
of the copper salt contain 7 volumes of copper (difference = 38).
Besides which, it must be observed that copper oxide displaces
iron oxide, just as silver oxide displaces copper oxide. Silver,
copper, and iron, in the form of oxides, displace each other in
the above order, but in the form of metals in a reverse order
(iron, copper, silver). The cause of this order of the
displacement of the oxides lies, amongst other things, in their
composition. They have the composition Ag_{2}O, Cu_{2}O_{2},
Fe_{2}O_{3}; the oxide containing a less proportion of oxygen
displaces that containing a larger proportion, because the basic
character diminishes with the increase of contained oxygen.
Copper also displaces mercury from its salts. It may here be
remarked that Spring (1888), on leaving a mixture of dry mercurous
chloride and copper for two hours, observed a distinct reduction,
which belongs to the category of those phenomena which demonstrate
the existence of a mobility of parts (_i.e._ atoms and molecules)
in solid substances.
[22 bis] The reaction of 1 part by weight of AgNO_{3} requires
(according to Kremers) the following amounts of water: at 0°, 0·82
part, at 19°·5, 0·41 part, at 54°, 0·20 part, at 110°, 0·09 part,
and, according to Tilden, at 125°, 0·0617 part, and at 133°,
0·0515 part.
[22 tri] It may be remarked that the black stain produced by the
reduction of metallic silver disappears under the action of a
solution of mercuric chloride or of potassium cyanide, because
these salts act on finely-divided silver.
Nitrate of silver easily forms double salts like AgNO_{3}2NaNO_{3} and
AgNO_{3}KNO_{3}. Silver nitrate under the action of water and a halogen
gives nitric acid (_see_ Vol. I. p. 280, formation of N_{2}O_{5}), a
halogen salt of silver, and a silver salt of an oxygen acid of the
halogen. Thus, for example, a solution of chlorine in water, when mixed
with a solution of silver nitrate, gives silver chloride and chlorate. It
is here evident that the reaction of the silver nitrate is identical with
the reaction of the caustic alkalis, as the nitric acid is all set free
and the silver oxide only reacts in exactly the same way in which aqueous
potash acts on free chlorine. Hence the reaction may be expressed in the
following manner: 6AgNO_{3} + 3Cl_{2} + 3H_{2}O = 5AgCl + AgClO_{3} +
6NHO_{3}.
Silver nitrate, like the nitrates of the alkalis, does not contain any
water of crystallisation. Moreover the other salts of silver almost
always separate out without any water of crystallisation. The silver
salts are further characterised by the fact that they _give neither basic
nor acid salts_, owing to which the formation of silver salts generally
forms the means of determining the true composition of acids--thus, to
any acid H_{n}X there corresponds a salt Ag_{n}X--for instance,
Ag_{3}PO_{4} (Chapter XIX., Note 15).
_Silver_ gives insoluble and exceedingly stable _compounds with the
halogens_. They are obtained by double decomposition with great facility
whenever a silver salt comes in contact with halogen salts. Solutions of
nitrate, sulphate, and all other kindred salts of silver give a
precipitate of silver chloride or iodide in solutions of chlorides and
iodides and of the halogen acids, because the halogen salts of silver are
insoluble both in water[23] and in other acids. _Silver chloride_, AgCl,
is then obtained as a white flocculent precipitate, silver bromide forms
a yellowish precipitate, and silver iodide has a very distinct yellow
colour. These halogen compounds sometimes occur in nature; they are
formed by a dry method--by the action of halogen compounds on silver
compounds, especially under the influence of heat. Silver chloride easily
fuses at 451° on cooling from a molten state; it forms a somewhat soft
horn-like mass which can be cut with a knife and is known as _horn
silver_. It volatilises at a higher temperature. Its ammoniacal solution,
on the evaporation of the ammonia, deposits crystalline chloride of
silver, in octahedra. Bromide and iodide of silver also appear in forms
of the regular system, so that in this respect the halogen salts of
silver resemble the halogen salts of the alkali metals.[24]
[23] Silver chloride is almost perfectly insoluble in water, but is
somewhat soluble in water containing sodium chloride or
hydrochloric acid, or other chlorides, and many salts, in
solution. Thus at 100°, 100 parts of water saturated with sodium
chloride dissolve 0·4 part of silver chloride. Bromide and iodide
of silver are less soluble in this respect, as also in regard to
other solvents. It should be remarked that _silver chloride
dissolves in solutions of ammonia, potassium cyanide, and of
sodium thiosulphate_, Na_{2}S_{2}O_{3}. Silver bromide is almost
perfectly analogous to the chloride, but silver iodide is nearly
insoluble in a solution of ammonia. Silver chloride even absorbs
dry ammonia gas, forming very unstable ammoniacal compounds. When
heated, these compounds (Vol. I. p. 250, Note 8) evolve the
ammonia, as they also do under the action of all acids. Silver
chloride enters into double decomposition with potassium cyanide,
forming a soluble double cyanide, which we shall presently
describe; it also forms a soluble double salt, NaAgS_{2}O_{3},
with sodium thiosulphate.
Silver chloride offers different modifications in the structure of
its molecule, as is seen in the variations in the consistency of
the precipitate, and in the differences in the action of light
which partially decomposes AgCl (_see_ Note 25). Stas and Carey
Lea investigated this subject, which has a particular importance
in photography, because silver bromide also gives _photo-salts_.
There is still much to be discovered in this respect, since Abney
showed that perfectly dry AgCl placed in a vacuum in the dark is
not in the least acted upon when subsequently exposed to light.
[24] _Silver bromide_ and _iodide_ (which occur as the minerals bromite
and iodite) resemble the chloride in many respects, but the degree
of affinity of silver for iodine is greater than that for chlorine
and bromine, although less heat is evolved (_see_ Note 28 bis).
Deville deduced this fact from a number of experiments. Thus
silver chloride, when treated with hydriodic acid, evolves
hydrochloric acid, and forms silver iodide. Finely-divided silver
easily liberates hydrogen when treated with hydriodic acid; it
produces the same decomposition with hydrochloric acid, but in a
considerably less degree and only on the surface. The difference
between silver chloride and iodide is especially remarkable, since
the formation of the former is attended with a greater contraction
than that of the latter. The volume of AgCl = 26; of chlorine 27,
of silver 10, the sum = 37, hence a contraction has ensued; and in
the formation of silver iodide an expansion takes place, for the
volume of Ag is 10, of I 26, and of AgI 39 instead of 36 (density,
AgCl, 5·59; AgI, 5·67). The atoms of chlorine have united with the
atoms of silver without moving asunder, whilst the atoms of iodine
must have moved apart in combining with the silver. It is
otherwise with respect to the metal; the distance between its
atoms in the metal = 2·2, in silver chloride = 3·0, and in silver
iodide = 3·5; hence its atoms have moved asunder considerably in
both cases. It is also very remarkable, as Fizeau observed, that
the density of silver iodide increases with a rise of
temperature--that is, a contraction takes place when it is heated
and an expansion when it is cooled.
In order to explain the fact that in silver compounds the iodide
is more stable than the chloride and oxide, Professor N. N.
Beketoff, in his 'Researches on the Phenomena of Substitutions'
(Kharkoff, 1865), proposed the following original hypothesis,
which we will give in almost the words of the author:--In the case
of aluminium, the oxide, Al_{2}O_{3}, is more stable than the
chloride, Al_{2}Cl_{6}, and the iodide, Al_{2}I_{6}. In the oxide
the amount of the metal is to the amount of the element combined
with it as 54·8 (Al = 27·3) is to 48, or in the ratio 112 : 100;
for the chloride the ratio is = 25 : 100; for the iodide it = 7 :
100. In the case of silver the oxide (ratio = 1350 : 100) is less
stable than the chloride (ratio = 304 : 100), and the iodide
(ratio of the weight of metal to the weight of the halogen = 85 :
100) is the most stable. From these and similar examples it
follows that the most stable compounds are those in which the
weights of the combined substances are equal. This may be partly
explained by the attraction of similar molecules even after their
having passed into combination with others. This attraction is
proportional to the product of the acting masses. In silver oxide
the attraction of Ag_{2} for Ag_{2} = 216 × 216 = 46,656, and the
attraction of Ag_{2} for O = 216 × 16 = 3,456. The attraction of
like molecules thus counteracts the attraction of the unlike
molecules. The former naturally does not overcome the latter,
otherwise there would be a disruption, but it nevertheless
diminishes the stability. In the case of an equality or proximity
of the magnitude of the combining masses, the attraction of the
like parts will counteract the stability of the compound to the
least extent--in other words, with an inequality of the combined
masses, the molecules have an inclination to return to an
elementary state, to decompose, which does not exist to such an
extent where the combined masses are equal. There is, therefore, a
tendency for large masses to combine with large, and for small
masses to combine with small. Hence Ag_{2}O + 2KI gives K_{2}O +
2AgI. The influence of an equality of masses on the stability is
seen particularly clearly in the effect of a rise of temperature.
Argentic, mercuric, auric and other oxides composed of unequal
masses, are somewhat readily decomposed by heat, whilst the oxides
of the lighter metals (like water) are not so easily decomposed by
heat. Silver chloride and iodide approach the condition of
equality, and are not decomposed by heat. The most stable oxides
under the action of heat are those of magnesium, calcium, silicon,
and aluminium, since they also approach the condition of equality.
For the same reason hydriodic acid decomposes with greater
facility than hydrochloric acid. Chlorine does not act on magnesia
or alumina, but it acts on lime and silver oxide, &c. This is
partially explained by the fact that by considering heat as a mode
of motion, and knowing that the atomic heats of the free elements
are equal, it must be supposed that the amount of the motion of
atoms (their _vis viva_) is equal, and as it is equal to the
product of the mass (atomic weight) into the square of the
velocity, it follows that the greater the combining weight the
smaller will be the square of the velocity, and if the combining
weights be nearly equal, then the velocities also will be nearly
equal. Hence the greater the difference between the weights of the
combined atoms the greater will be the difference between their
velocities. The difference between the velocities will increase
with the temperature, and therefore the temperature of
decomposition will be the sooner attained the greater be the
original difference--that is, the greater the difference of the
weights of the combined substances. The nearer these weights are
to each other, the more analogous the motion of the unlike atoms,
and consequently, the more stable the resultant compound.
The instability of cupric chloride and nitric oxide, the absence
of compounds of fluorine with oxygen, whilst there are compounds
of oxygen with chlorine, the greater stability of the oxygen
compounds of iodine than those of chlorine, the stability of boron
nitride, and the instability of cyanogen, and a number of similar
instances, where, judging from the above argument, one would
expect (owing to the closeness of the atomic weights) a stability,
show that Beketoff's addition to the mechanical theory of chemical
phenomena is still far from sufficient for explaining the true
relations of affinities. Nevertheless, in his mode of explaining
the relative stabilities of compounds, we find an exceedingly
interesting treatment of questions of primary importance. Without
such efforts it would be impossible to generalise the complex data
of experimental knowledge.
_Fluoride of silver_, AgF, is obtained by dissolving Ag_{2}O or
Ag_{2}CO_{3} in hydrofluoric acid. It differs from the other
halogen salts of silver in being soluble in water (1 part of salt
in 0·55 of water). It crystallises from its solution in prisms,
AgFH_{2}O (Marignac), or AgF_{2}H_{2}O (Pfaundler), which lose
their water in vacuo. Güntz (1891), by electrolising a saturated
solution of Ag_{2}F, obtained _polyfluoride of silver_, Ag_{2}F,
which is decomposed by water into AgF + Ag. It is also formed by
the action of a strong solution of AgF upon finely-divided
(precipitated) silver.
Silver chloride may be decomposed, with the separation of silver oxide,
by heating it with a solution of an alkali, and if an organic substance
be added to the alkali the chloride can easily be reduced o metallic
silver, the silver oxide being reduced in the oxidation of the organic
substance. Iron, zinc, and many other metals reduce silver chloride in
the presence of water. Cuprous and mercurous chlorides and many organic
substances are also able to reduce the silver from chloride of silver.
This shows the rather easy decomposability of the halogen compounds of
silver. Silver iodide is much more stable in this respect than the
chloride. The same is also observed with respect to the _action of light_
upon moist AgCl. White silver chloride soon acquires a violet colour when
exposed to the action of light, and especially under the direct action of
the sun's rays. After being acted upon by light it is no longer entirely
soluble in ammonia, but leaves metallic silver undissolved, from which it
might be assumed that the action of light consisted in the decomposition
of the silver chloride into chlorine and metallic silver and in fact the
silver chloride becomes in time darker and darker. Silver bromide and
iodide are much more slowly acted on by light, and, according to certain
observations, when pure they are even quite unacted on; at least they do
not change in weight,[24 bis] so that if they are acted on by light, the
change they undergo must be one of a change in the structure of their
parts and not of decomposition, as it is in silver chloride. The silver
chloride under the action of light changes in weight, which indicates the
formation of a volatile product, and the deposition of metallic silver on
dissolving in ammonia shows the loss of chlorine. The change does
actually occur under the action of light, but the decomposition does not
go as far as into chlorine and silver, but only to the formation of a
subchloride of silver, Ag_{2}Cl, which is of a brown colour and is easily
decomposed into metallic silver and silver chloride, Ag_{2}Cl = AgCl +
Ag. This change of the chemical composition and structure of the halogen
salts of silver under the action of light forms the basis of
_photography_, because the halogen compounds of silver, after having been
exposed to light, give a precipitate of finely-divided silver, of a black
colour, when treated with reducing agents.[25]
[24 bis] The changes brought about by the action of light necessitate
distinguishing the photo-salts of silver.
[25] In photography these are called 'developers.' The most common
developers are: solutions of ferrous sulphate, pyrogallol, ferrous
oxalate, hydroxylamine, potassium sulphite, hydroquinone (the last
acts particularly well and is very convenient to use), &c. The
chemical processes of photography are of great practical and
theoretical interest; but it would be impossible in this work to
enter into this special branch of chemistry, which has as yet been
very little worked out from a theoretical point of view.
Nevertheless, we will pause to consider certain aspects of this
subject which are of a purely chemical interest, and especially
the facts concerning _subchloride of silver_, Ag_{2}Cl (_see_ Note
19), and the photo-salts (Note 23). There is no doubt that under
the action of light, AgCl becomes darker in colour, decreases in
weight, and probably forms a mixture of AgCl, Ag_{2}Cl, and Ag.
But the isolation of the subchloride has only been recently
accomplished by Güntz by means of the Ag_{2}F, discovered by him
(_see_ Note 24). Many chemists (and among them Hodgkinson) assumed
that an oxychloride of silver was formed by the decomposition of
AgCl under the action of light. Carey Lea's (1889) and A.
Richardson's (1891) experiments showed that the product formed
does not, however, contain any oxygen at all, and the change in
colour produced by the action of light upon AgCl is most probably
due to the formation of Ag_{2}Cl. This substance was isolated by
Güntz (1891) by passing HCl over crystals of Ag_{2}F. He also
obtained Ag_{2}I in a similar manner by passing HI, and Ag_{2}S by
passing H_{2}S over Ag_{2}F. Ag_{2}Cl is best prepared by the
action of phosphorus trichloride upon Ag_{2}F. At the temperature
of its formation Ag_{2}Cl has an easily changeable tint, with
shades of violet red to violet black. Under the action of light a
similar (isomeric) substance is obtained, which splits up into
AgCl + Ag when heated. With potassium cyanide Ag_{2}Cl gives Ag +
AgCN + KCl, whence it is possible to calculate the heat of
formation of Ag_{2}Cl; it = 29·7, whilst the heat of formation of
AgCl = 29·2--_i.e._ the reaction 2AgCl = Ag_{2}Cl + Cl corresponds
to an absorption of 28·7 major calories. If we admit the formation
of such a compound by the action of light, it is evident that the
energy of the light is consumed in the above reaction. Carey Lea
(1892) subjected AgCl, AgBr, and AgI to a pressure (of course in
the dark) of 3,000 atmospheres, and to trituration with water in a
mortar, and observed a change of colour indicating incipient
decomposition, which is facilitated under the action of light by
the molecular currents set up (Lermontoff, Egoroff). The change of
colour of the halogen salts of silver under the action of light,
and their faculty of subsequently giving a visible photographic
image under the action of 'developers,' must now be regarded as
connected with the decomposition of AgX, leading to the formation
of Ag_{2}X, and the different tinted photo-salts must be
considered as systems containing such Ag_{2}X's. Carey Lea
obtained photo-salts of this kind not only by the action of light
but also in many other ways, which we will enumerate to prove that
they contain the products of an incomplete combination of Ag with
the halogens, (for the salts Ag_{2}X must be regarded as such).
The photo-salts have been obtained (1) by the imperfect
chlorination of silver; (2) by the incomplete decomposition of
Ag_{2}O or Ag_{2}CO_{3} by alternately heating and treating with a
halogen acid; (3) by the action of nitric acid or Na_{2}S_{2}O_{3}
upon Ag_{2}Cl; (4) by mixing a solution of AgNO_{3} with the
hydrates of FeO, MnO and CrO, and precipitating by HCl; (5) by the
action of HCl upon the product obtained by the reduction of
citrate of silver in hydrogen (Note 19), and (6) by the action of
milk sugar upon AgNO_{3} together with soda and afterwards
acidulating with HCl. All these reactions should lead to the
formation of products of imperfect combination with the halogens
and give photo-salts of a similar diversity of colour to those
produced by the action of developers upon the halogen salts of
silver after exposure to light.
The insolubility of the halogen compounds of silver forms the basis of
many methods used in practical chemistry. Thus by means of this reaction
it is possible to obtain salts of other acids from a halogen salt of a
given metal, for instance, RCl_{2} + 2AgNO_{3} = R(NO_{3})_{2} + 2AgCl.
The formation of the halogen compounds of silver is very frequently used
in the investigation of organic substances; for example, if any product
of metalepsis containing iodine or chlorine be heated with a silver salt
or silver oxide, the silver combines with the halogen and gives a halogen
salt, whilst the elements previously combined with the silver replace the
halogen. For instance, ethylene dibromide, C_{2}H_{4}Br_{2}, is
transformed into ethylene diacetate, C_{2}H_{4}(C_{2}H_{3}O_{2})_{2}, and
silver bromide by heating it with silver acetate, 2C_{2}H_{3}O_{2}Ag. The
insolubility of the halogen compounds of silver is still more frequently
taken advantage of in determining the amount of silver and halogen in a
given solution. If it is required, for instance, to determine the
quantity of chlorine present in the form of a metallic chloride in a
given solution, a solution of silver nitrate is added to it so long as it
gives a precipitate. On _shaking or stirring_ the liquid, the silver
chloride easily settles in the form of heavy flakes. It is possible in
this way to precipitate the whole of the chlorine from a solution,
without adding an excess of silver nitrate, since it can be easily seen
whether the addition of a fresh quantity of silver nitrate produces a
precipitate in the clear liquid. In this manner it is possible to add to
a solution containing chlorine, as much silver as is required for its
entire precipitation, and to calculate the amount of chlorine previously
in solution from the amount of the solution of silver nitrate consumed,
if the quantity of silver nitrate in this solution has been previously
determined.[25 bis] The atomic proportions and preliminary experiments
with a pure salt--for example, with sodium chloride--will give the amount
of chlorine from the quantity of silver nitrate. Details of these methods
will be found in works on analytical chemistry.[25 tri]
[25 bis] In order to determine when the reaction is at an end, a few
drops of a solution of K_{2}CrO_{4} are added to the solution of
the chloride. Before all the chlorine is precipitated as AgCl, the
precipitate (after shaking) is white (since Ag_{2}CrO_{4} with
2RCl gives 2AgCl); but when all the chlorine is thrown down
Ag_{2}CrO_{4} is formed, which colours the precipitate
reddish-brown. In order to obtain accurate results the liquid
should be neutral to litmus.
[25 tri] _Silver cyanide_, AgCN, is closely analogous to the haloid
salts of silver. It is obtained, in similar manner to silver
chloride, by the addition of potassium cyanide to silver nitrate.
A white precipitate is then formed, which is almost insoluble in
boiling water. It is also, like silver chloride, insoluble in
dilute acids. However, it is dissolved when heated with nitric
acid, and both hydriodic and hydrochloric acids act on it,
converting it into silver chloride and iodide. Alkalis, however,
do not act on silver cyanide, although they act on the other
haloid salts of silver. Ammonia and solutions of the cyanides of
the alkali metals dissolve silver cyanide, as they do the
chloride. In the latter case double cyanides are formed--for
example, KAgC_{2}N_{2}. This salt is obtained in a crystalline
state on evaporating a solution of silver cyanide in potassium
cyanide. It is much more stable than silver cyanide itself. It has
a neutral reaction, does not change in the air, and does not smell
of hydrocyanic acid. Many acids, in acting on a solution of this
double salt, precipitate the insoluble silver cyanide. Metallic
silver dissolves in a solution of potassium cyanide in the
presence of air, with formation of the same double salt and
potassium hydroxide, and when silver chloride dissolves in
potassium cyanide it forms potassium chloride, besides the salt
KAgC_{2}N_{2}. This double salt of silver is used in silver
plating. For this purpose potassium cyanide is added to its
solution, as otherwise silver cyanide, and not metallic silver, is
deposited by the electric current. If two electrodes--one positive
(silver) and the other negative (copper)--be immersed in such a
solution, silver will be deposited upon the latter, and the silver
of the positive electrode will be dissolved by the liquid, which
will thus preserve the same amount of metal in solution as it
originally contained. If instead of the negative electrode a
copper object be taken, well cleaned from all dirt, the silver
will be deposited in an even coating; this, indeed, forms the mode
of _silver plating by the wet method_, which is most often used in
practice. A solution of one part of silver nitrate in 30 to 50
parts of water, and mixed with a sufficient quantity of a solution
of potassium cyanide to redissolve the precipitate of silver
cyanide formed, gives a dull coating of silver, but if twice as
much water be used the same mixture gives a bright coating.
Silver plating in the wet way has now replaced to a considerable
extent the old process of _dry silvering_, because this process,
which consists in dissolving silver in mercury and applying the
amalgam to the surface of the objects, and then vaporising the
mercury, offers the great disadvantage of the poisonous mercury
fumes. Besides these, there is another method of silver plating,
based on the direct displacement of silver from its salts by other
metals--for example, by copper. The copper reduces the silver from
its compounds, and the silver separated is deposited upon the
copper. Thus a solution of silver chloride in sodium thiosulphate
deposits a coating of silver upon a strip of copper immersed in
it. It is best for this purpose to take pure _silver sulphite_.
This is prepared by mixing a solution of silver nitrate with an
excess of ammonia, and adding a saturated solution of sodium
sulphite and then alcohol, which precipitates silver sulphite from
the solution. The latter and its solutions are very easily
decomposed by copper. Metallic iron produces the same
decomposition, and iron and steel articles may be very readily
silver-plated by means of the thiosulphate solution of silver
chloride. Indeed, copper and similar metals may even be
silver-plated by means of silver chloride; if the chloride of
silver, with a small amount of acid, be rubbed upon the surface of
the copper, the latter becomes covered with a coating of silver,
which it has reduced.
Silver plating is not only applicable to metallic objects, but
also to glass, china, &c. Glass is silvered for various
purposes--for example, glass globes silvered internally are used
for ornamentation, and have a mirrored surface. Common
looking-glass silvered upon one side forms a mirror which is
better than the ordinary mercury mirrors, owing to the truer
colours of the image due to the whiteness of the silver. For
optical instruments--for example, telescopes--concave mirrors are
now made of silvered glass, which has first been ground and
polished into the required form. The _silvering of glass_ is based
on the fact that silver which is reduced from certain solutions
deposits itself uniformly in a perfectly homogeneous and
continuous but very thin layer, forming a bright reflecting
surface. Certain organic substances have the property of reducing
silver in this form. The best known among these are certain
aldehydes--for instance, ordinary acetaldehyde, C_{2}H_{4}O, which
easily oxidises in the air and forms acetic acid, C_{2}H_{4}O_{2}.
This oxidation also easily takes place at the expense of silver
oxide, when a certain amount of ammonia is added to the mixture.
The oxide of silver gives up its oxygen to the aldehyde, and the
silver reduced from it is deposited in a metallic state in a
uniform bright coating. The same action is produced by certain
saccharine substances and certain organic acids, such as tartaric
acid, &c.
Accurate experiments, and more especially the _researches of Stas_ at
Brussels, show the proportion in which silver reacts with metallic
chlorides. These researches have led to the determination of the
_combining weights_ of silver, sodium, potassium, chlorine, bromine,
iodine, and other elements, and are distinguished for their model
exactitude, and we will therefore describe them in some detail. As sodium
chloride is the chloride most generally used for the precipitation of
silver, since it can most easily be obtained in a pure state, we will
here cite the quantitative observations made by Stas for showing the
co-relation between the quantities of chloride of sodium and silver which
react together. In order to obtain perfectly pure sodium chloride, he
took pure rock salt, containing only a small quantity of magnesium and
calcium compounds and a small amount of potassium salts. This salt was
dissolved in water, and the saturated solution evaporated by boiling. The
sodium chloride separated out during the boiling, and the mother liquor
containing the impurities was poured off. Alcohol of 65 p.c. strength and
platinic chloride were added to the resultant salt, in order to
precipitate all the potassium and a certain part of the sodium salts. The
resultant alcoholic solution, containing the sodium and platinum
chlorides, was then mixed with a solution of pure ammonium chloride in
order to remove the platinic chloride. After this precipitation, the
solution was evaporated in a platinum retort, and then separate portions
of this purified sodium chloride were collected as they crystallised. The
same salt was prepared from sodium sulphate, tartrate, nitrate, and from
the platinochloride, in order to have sodium chloride prepared by
different methods and from different sources, and in this manner ten
samples of sodium chloride thus prepared were purified and investigated
in their relation to silver. After being dried, weighed quantities of all
ten samples of sodium chloride were dissolved in water and mixed with a
solution in nitric acid of a weighed quantity of perfectly pure silver. A
slightly greater quantity of silver was taken than would be required for
the decomposition of the sodium chloride, and when, after pouring in all
the silver solution, the silver chloride had settled, the amount of
silver remaining in excess was determined by means of a solution of
sodium chloride of known strength. This solution of sodium chloride was
added so long as it formed a precipitate. In this manner Stas determined
how many parts of sodium chloride correspond to 100 parts by weight of
silver. The result of ten determinations was that for the entire
precipitation of 100 parts of silver, from 54·2060 to 54·2093 parts of
sodium chloride were required. The difference is so inconsiderable that
it has no perceptible influence on the subsequent calculations. The mean
of ten experiments was that 100 parts of silver react with 54·2078 parts
of sodium chloride. In order to learn from this the relation between the
chlorine and silver, it was necessary to determine the quantity of
chlorine contained in 54·2078 parts of sodium chloride, or, what is the
same thing, the quantity of chlorine which combines with 100 parts of
silver. For this purpose Stas made a series of observations on the
quantity of silver chloride obtained from 100 parts of silver. Four
syntheses were made by him for this purpose. The first synthesis
consisted in the formation of silver chloride by the action of chlorine
on silver at a red heat. This experiment showed that 100 parts of silver
give 132·841, 132·843 and 132·843 of silver chloride. The second method
consisted in dissolving a given quantity of silver in nitric acid and
precipitating it by means of gaseous hydrochloric acid passed over the
surface of the liquid; the resultant mass was evaporated in the dark to
drive off the nitric acid and excess of hydrochloric acid, and the
remaining silver chloride was fused first in an atmosphere of
hydrochloric acid gas and then in air. In this process the silver
chloride was not washed, and therefore there could be no loss from
solution. Two experiments made by this method showed that 100 parts of
silver give 132·849 and 132·846 parts of silver chloride. A third series
of determinations was also made by precipitating a solution of silver
nitrate with a certain excess of gaseous hydrochloric acid. The amount of
silver chloride obtained was altogether 132·848. Lastly, a fourth
determination was made by precipitating dissolved silver with a solution
of ammonium chloride, when it was found that a considerable amount of
silver (0·3175) had passed into solution in the washing; for 100 parts of
silver there was obtained altogether 132·8417 of silver chloride. Thus
from the mean of seven determinations it appears that 100 parts of silver
give 132·8445 parts of silver chloride--that is, that 32·8445 parts of
chlorine are able to combine with 100 parts of silver and with that
quantity of sodium which is contained in 54·2078 parts of sodium
chloride. These observations show that 32·8445 parts of chlorine combine
with 100 parts of silver and with 21·3633 parts of sodium. From these
figures expressing the relation between the combining weights of
chlorine, silver, and sodium, it would be possible to determine their
atomic weights--that is, the combining quantity of these elements with
respect to one part by weight of hydrogen or 16 parts of oxygen, if there
existed a series of similarly accurate determinations for the reactions
between hydrogen or oxygen and one of these elements--chlorine, sodium,
or silver. If we determine the quantity of silver chloride which is
obtained from silver chlorate, AgClO_{3}, we shall know the relation
between the combining weights of silver chloride and oxygen, so that,
taking the quantity of oxygen as a constant magnitude, we can learn from
this reaction the combining weight of silver chloride, and from the
preceding numbers the combining weights of chlorine and silver. For this
purpose it was first necessary to obtain pure silver chlorate. This Stas
did by acting on silver oxide or carbonate, suspended in water, with
gaseous chlorine.[26]
[26] The phenomenon which then takes place is described by Stas as
follows, in a manner which is perfect in its clearness and
accuracy: if silver oxide or carbonate be suspended in water, and
an excess of water saturated with chlorine be added, all the
silver is converted into chloride, just as is the case with oxide
or carbonate of mercury, and the water then contains, besides the
excess of chlorine, only pure hypochlorous acid without the least
trace of chloric or chlorous acid. If a stream of chlorine be
passed into water containing _an excess of silver oxide_ or silver
carbonate while the liquid is continually agitated, the reaction
is the same as the preceding; silver chloride and hypochlorous
acid are formed. But this acid does not long remain in a free
state: it gradually acts on the silver oxide and gives silver
hypochlorite, _i.e._ AgClO. If, after some time, the current of
chlorine be stopped but the shaking continued, the liquid loses
its characteristic odour of hypochlorous acid, while preserving
its energetic decolorising property, because the silver
hypochlorite which is formed is easily soluble in water. In the
presence of an excess of silver oxide this salt can be kept for
several days without decomposition, but it is exceedingly unstable
when no excess of silver oxide or carbonate is present. So long as
the solution of silver hypochlorite is shaken up with the silver
oxide, it preserves its transparency and bleaching property, but
directly it is allowed to stand, and the silver oxide settles, it
becomes rapidly cloudy and deposits large flakes of silver
chloride, so that the black silver oxide which had settled becomes
covered with the white precipitate. The liquid then loses its
bleaching properties and contains silver chlorate, _i.e._
AgClO_{3}, in solution, which has a slightly alkaline reaction,
owing to the presence of a small amount of dissolved oxide. In
this manner the reactions which are consecutively accomplished may
be expressed by the equations:
6Cl_{2} + 3Ag_{2}O + 3H_{2}O = 6AgCl + 6HClO;
6HClO + 3Ag_{2}O = 3H_{2}O + 6AgClO;
6AgClO = 4AgCl + 2AgClO_{3}.
Hence, Stas gives the following method for the preparation of
silver chlorate: A slow current of chlorine is caused to act on
oxide of silver, suspended in water which is kept in a state of
continual agitation. The shaking is continued after the supply of
chlorine has been stopped, in order that the free hypochlorous
acid should pass into silver hypochlorite, and the resultant
solution of the hypochlorite is drawn off from the sediment of the
excess of silver oxide. This solution decomposes spontaneously
into silver chloride and chlorate. The pure silver chlorate,
AgClO_{3}, does not change under the action of light. The salt is
prepared for further use by drying it in dry air at 150°. It is
necessary during drying to prevent the access of any organic
matter; this is done by filtering the air through cotton wool, and
passing it over a layer of red-hot copper oxide.
The decomposition of the silver chlorate thus obtained was accomplished
by the action of a solution of sulphurous anhydride on it. The salt was
first fused by carefully heating it at 243°. The solution of sulphurous
anhydride used was one saturated at 0°. Sulphurous anhydride in dilute
solutions is oxidised at the expense of silver chlorate, even at low
temperatures, with great ease if the liquid be continually shaken,
sulphuric acid and silver chloride being formed: AgClO_{3} + 3SO_{2} +
3H_{2}O = AgCl + 3H_{2}SO_{4}. After decomposition, the resultant liquid
was evaporated, and the residue of silver chloride weighed. Thus the
process consisted in taking a known weight of silver chlorate, converting
it into silver chloride, and determining the weight of the latter. The
analysis conducted in this manner gave the following results, which, like
the preceding, designate the weight in a vacuum calculated from the
weights obtained in air: In the first experiment it appeared that
138·7890 grams of silver chlorate gave 103·9795 parts of silver chloride,
and in the second experiment that 259·5287 grains of chlorate gave
194·44515 grams of silver chloride, and after fusion 194·4435 grams. The
mean result of both experiments, converted into percentages, shows that
100 parts of silver chlorate contain 74·9205 of silver chloride and
25·0795 parts of oxygen. From this it is possible to calculate the
combining weight of silver chloride, because in the decomposition of
silver chlorate there are obtained three atoms of oxygen and one molecule
of silver chloride: AgClO_{3} = AgCl + 3O. Taking the weight of an atom
of oxygen to be 16, we find from the mean result that the equivalent
weight of silver chloride is equal to 143·395. Thus if O = 16, AgCl =
143·395, and as the preceding experiments show that silver chloride
contains 32·8445 parts of chlorine per 100 parts of silver, the weight of
the atom of silver[26 bis] must be 107·94 and that of chlorine 35·45. The
weight of the atom of sodium is determined from the fact that 21·3633
parts of sodium chloride combine with 32·8445 parts of chlorine;
consequently Na = 23·05. This conclusion, arrived at by the analysis of
silver chlorate, was verified by means of the analysis of potassium
chlorate by decomposing it by heat and determining the weight of the
potassium chloride formed, and also by effecting the same decomposition
by igniting the chlorate in a stream of hydrochloric acid. The combining
weight of potassium chloride was thus determined, and another series of
determinations confirmed the relation between chlorine, potassium, and
silver, in the same manner as the relation between sodium, chlorine, and
silver was determined above. Consequently, the combining weights of
sodium, chlorine, and potassium could be deduced by combining these data
with the analysis of silver chlorate and the synthesis of silver
chloride. The agreement between the results showed that the
determinations made by the last method were perfectly correct, and did
not depend in any considerable degree on the methods which were employed
in the preceding determinations, as the combining weights of chlorine and
silver obtained were the same as before. There was naturally a
difference, but so small a one that it undoubtedly depended on the errors
incidental to every process of weighing and experiment. The atomic weight
of silver was also determined by Stas by means of the synthesis of silver
sulphide and the analysis of silver sulphate. The combining weight
obtained by this method was 107·920. The synthesis of silver iodide and
the analysis of silver iodate gave the figure 107·928. The synthesis of
silver bromide with the analysis of silver bromate gave the figure
107·921. The synthesis of silver chloride and the analysis of silver
chlorate gave a mean result of 107·937. Hence there is no doubt that the
combining weight of silver is at least as much as 107·9--greater than
107·90 and less than 107·95, and probably equal to the mean = 107·92.
Stas determined the combining weights of many other elements in this
manner, such as lithium, potassium, sodium, bromine, chlorine, iodine,
and also nitrogen, for the determination of the amount of silver nitrate
obtained from a given amount of silver gives directly the combining
weight of nitrogen. Taking that of oxygen as 16, he obtained the
following combining weights for these elements: nitrogen 14·04, silver
107·93, chlorine 35·46, bromine 79·95, iodine 126·85, lithium 7·02,
sodium 23·04, potassium 39·15. These figures differ slightly from those
which are usually employed in chemical investigations. They must be
regarded as the result of the best observations, whilst the figures
usually used in practical chemistry are only approximate--are, so to
speak, round numbers for the atomic weights which differ so little from
the exact figures (for instance, for Ag 108 instead of 107·92, for Na 23
instead of 23·04) that in ordinary determinations and calculations the
difference falls within the limits of experimental error inseparable from
such determinations.
[26 bis] The results given by Stas' determinations have recently
been recalculated and certain corrections have been introduced. We
give in the context the average results of van der Plaats and
Thomsen's calculations, as well as in Table III. neglecting the
doubtful thousandths.
The exhaustive investigations conducted by Stas on the atomic weights
of the above-named elements have great significance in the solution of
the problem as to whether the atomic weights of the elements can be
expressed in whole numbers if the unit taken be the atomic weight of
hydrogen. Prout, at the beginning of this century, stated that this was
the case, and held that the atomic weights of the elements are multiples
of the atomic weight of hydrogen. The subsequent determinations of
Berzelius, Penny, Marchand, Marignac, Dumas, and more especially of Stas,
proved this conclusion to be untenable; since a whole series of elements
proved to have fractional atomic weights--for example, chlorine, about
35·5. On account of this, Marignac and Dumas stated that the atomic
weights of the elements are expressed in relation to hydrogen, either by
whole numbers or by numbers with simple fractions of the magnitudes 1/2
and 1/4. But Stas's researches refute this supposition also. Even between
the combining weight of hydrogen and oxygen, there is not, so far as is
yet known, that simple relation which is required by _Prout's
hypothesis_,[27] _i.e._, taking O = 16, the atomic weight of hydrogen is
equal not to 1 but to a greater number somewhere between 1·002 and 1·008
or mean 1·005. Such a conclusion arrived at by direct experiment cannot
but be regarded as having greater weight than Prout's supposition
(hypothesis) that the atomic weights of the elements are in multiple
proportion to each other, which would give reason for surmising (but not
asserting) a complexity of nature in the elements, and their common
origin from a single primary material, and for expecting their mutual
conversion into each other. All such ideas and hopes must now, thanks
more especially to Stas, be placed in a region void of any experimental
support whatever, and therefore not subject to the discipline of the
positive data of science.
[27] This hypothesis, for the establishment or refutation of which so
many researches have been made, is exceedingly important, and
fully deserves the attention which has been given to it. Indeed,
if it appeared that the atomic weights of all the elements could
be expressed in whole numbers with reference to hydrogen, or if
they at least proved to be commensurable with one another, then it
could be affirmed with confidence that the elements, with all
their diversity, were formed of one material condensed or grouped
in various manners into the stable, and, under known conditions,
undecomposable groups which we call the atoms of the elements. At
first it was supposed that all the elements were nothing else but
condensed hydrogen, but when it appeared that the atomic weights
of the elements could not be expressed in whole numbers in
relation to hydrogen, it was still possible to imagine the
existence of a certain material from which both hydrogen and all
the other elements were formed. If it should transpire that four
atoms of this material form an atom of hydrogen, then the atom of
chlorine would present itself as consisting of 142 atoms of this
substance, the weight of whose atom would be equal to 0·25. But in
this case the atoms of all the elements should be expressed in
whole numbers with respect to the weight of the atom of this
original material. Let us suppose that the atomic weight of this
material is equal to unity, then all the atomic weights should be
expressible in whole numbers relatively to this unit. Thus the
atom of one element, let us suppose, would weigh _m_, and of
another _n_, but, as both _m_ and _n_ must be whole numbers, it
follows that the atomic weights of all the elements would be
commensurable. But it is sufficient to glance over the results
obtained by Stas, and to be assured of their accuracy, especially
for silver, in order to entirely destroy, or at least strongly
undermine, this attractive hypothesis. We must therefore refuse
our assent to the doctrine of the building up from a single
substance of the elements known to us. This hypothesis is not
supported either by any known transformation (for one element has
never been converted into another element), or by the
commensurability of the atomic weights of the elements. Although
the hypothesis of the formation of all the elements from a single
substance (for which Crookes has suggested the name protyle) is
most attractive in its comprehensiveness, it can neither be denied
nor accepted for want of sufficient data. Marignac endeavoured,
however, to overcome Stas's conclusions as to the
incommensurability of the atomic weights by supposing that in his,
as in the determinations of all other observers, there were
unperceived errors which were quite independent of the mode of
observation--for example, silver nitrate might be supposed to be
an unstable substance which changes, under the heatings,
evaporations, and other processes to which it is subjected in the
reactions for the determination of the combining weight of silver.
It might be supposed, for instance, that silver nitrate contains
some impurity which cannot be removed by any means; it might also
be supposed that a portion of the elements of the nitric acid are
disengaged in the evaporation of the solution of silver nitrate
(owing to the decomposing action of water), and in its fusion, and
that we have not to deal with normal silver nitrate, but with a
slightly basic salt, or perhaps an excess of nitric acid which
cannot be removed from the salt. In this case the observed
combining weight will not refer to an actually definite chemical
compound, but to some mixture for which there does not exist any
perfectly exact combining relations. Marignac upholds this
proposition by the fact that the conclusions of Stas and other
observers respecting the combining weights determined with the
greatest exactitude very nearly agree with the proposition of the
commensurability of the atomic weights--for example, the combining
weight of silver was shown to be equal to 107·93, so that it only
differs by 0·08 from the whole number 108, which is generally
accepted for silver. The combining weight of iodine proved to be
equal to 126·85--that is, it differs from 127 by 0·15. The
combining weights of sodium, nitrogen, bromine, chlorine, and
lithium are still nearer to the whole or round numbers which are
generally accepted. But Marignac's proposition will hardly bear
criticism. Indeed if we express the combining weights of the
elements determined by Stas in relation to hydrogen, the
approximation of these weights to whole numbers disappears,
because one part of hydrogen in reality does not combine with 16
parts of oxygen, but with 15·92 parts, and therefore we shall
obtain, taking H = 1, not the above-cited figures, but for silver
107·38, for bromine 79·55, magnitudes which are still further
removed from whole numbers. Besides which, if Marignac's
proposition were true the combining weight of silver determined by
one method--_e.g._ by the analysis of silver chlorate combined
with the synthesis of silver chloride--would not agree well with
the combining weight determined by another method--_e.g._ by means
of the analysis of silver iodate and the synthesis of silver
iodide. If in one case a basic salt could be obtained, in the
other case an acid salt might be obtained. Then the analysis of
the acid salt would give different results from that of the basic
salt. Thus Marignac's arguments cannot serve as a support for the
vindication of Prout's hypothesis.
In conclusion, I think it will not be out of place to cite the
following passage from a paper I read before the Chemical Society
of London in 1889 (Appendix II.), referring to the hypothesis of
the complexity of the elements recognised in chemistry, owing to
the fact that many have endeavoured to apply the periodic law to
the justification of this idea 'dating from a remote antiquity,
when it was found convenient to admit the existence of many gods
but only one matter.'
'When we try to explain the origin of the idea of a unique primary
matter, we easily trace that, in the absence of deductions from
experiment, it derives its origin from the scientifically
philosophical attempt at discovering some kind of unity in the
immense diversity of individualities which we see around. In
classical times such a tendency could only be satisfied by
conceptions about the immaterial world. As to the material world,
our ancestors were compelled to resort to some hypothesis, and
they adopted the idea of unity in the formative material, because
they were not able to evolve the conception of any other possible
unity in order to connect the multifarious relations of matter.
Responding to the same legitimate scientific tendency, natural
science has discovered throughout the universe a unity of plan, a
unity of forces, and a unity of matter; and the convincing
conclusions of modern science compel every one to admit these
kinds of unity. But while we admit unity in many things, we none
the less must also explain the individuality and the apparent
diversity which we cannot fail to trace everywhere. It was said of
old "Give us a fulcrum and it will become easy to displace the
earth." So also we must say, "Give us something that is
individualised, and the apparent diversity will be easily
understood." Otherwise, how could unity result in a multitude.
'After a long and painstaking research, natural science has
discovered the individualities of the chemical elements, and
therefore it is now capable, not only of analysing, but also of
synthesising; it can understand and grasp generality and unity, as
well as the individualised and multifarious. The general and
universal, like time and space, like force and motion, vary
uniformly. The uniform admit of interpolations, revealing every
intermediate phase; but the multitudinous, the
individualised--such as ourselves, or the chemical elements, or
the members of a peculiar periodic function of the elements, or
Dalton's multiple proportions--is characterised in another way. We
see in it--side by side with a general connecting
principle--leaps, breaks of continuity, points which escape from
the analysis of the infinitely small--an absence of complete
intermediate links. Chemistry has found an answer to the question
as to the causes of multitudes, and while retaining the conception
of many elements, all submitted to the discipline of a general
law, it offers an escape from the Indian Nirvana--the absorption
in the universal--replacing it by the individualised. However, the
place for individuality is so limited by the all-grasping,
all-powerful universal, that it is merely a point of support for
the understanding of multitude in unity.'
Among the platinum metals ruthenium, rhodium, and palladium, by their
atomic weights and properties, approach silver, just as iron and its
analogues (cobalt and nickel) approach copper in all respects. _Gold_
stands in exactly the same position in relation to the heavy platinum
metals, osmium, iridium, and platinum, as copper and silver do to the two
preceding series. The atomic weight of gold is nearly equal to their
atomic weights;[28] it is dense like these metals. It also gives various
grades of oxidation, which are feeble, both in a basic and an acid sense.
Whilst near to osmium, iridium, and platinum, gold at the same time is
able, like copper and silver, to form compounds which answer to the type
RX--that is, oxides of the composition R_{2}O. Cuprous chloride, CuCl,
silver chloride, AgCl, and aurous chloride, AuCl, are substances which
are very much alike in their physical and chemical properties.[28 bis]
They are insoluble in water, but dissolve in hydrochloric acid and
ammonia, in potassium cyanide, sodium thiosulphate, &c. Just as copper
forms a link between the iron metals and zinc, and as silver unites the
light platinum metals with cadmium, so also gold presents a transition
from the heavy platinum metals to mercury. Copper gives saline compounds
of the types CuX and CuX_{2}, silver of the type AgX, whilst gold,
besides compounds of the type AuX, very easily and most frequently forms
those of the type AuCl_{3}. The compounds of this type frequently pass
into those of the lower type, just as PtX_{4} passes into PtX_{2}, and
the same is observable in the elements which, in their atomic weights,
follow gold. Mercury gives HgX_{2} and HgX, thallium gives TlX_{3} and
TlX, lead gives PbX_{4} and PbX_{2}. On the other hand, gold in a
qualitative respect differs from silver and copper in the _extreme ease_
with which all its compounds are _reduced to metal_ by many means. This
is not only accomplished by many reducing agents, but also by the action
of heat. Thus its chlorides and oxides lose their chlorine and oxygen
when heated, and, if the temperature be sufficiently high, these elements
are entirely expelled and metallic gold alone remains. Its compounds,
therefore, act as oxidising agents.[29]
[28] It might be expected from the periodic law and analogies with the
series iron, cobalt, nickel, copper, zinc, that the atomic weights
of the elements of the series osmium, iridium, platinum, gold,
mercury, would rise in this order, and at the time of the
establishment of the periodic law (1869), the determinations of
Berzelius, Rose, and others gave the following values for the
atomic weights: Os = 200, Ir = 197, Pt = 198, Au = 196, Hg = 200.
The fulfilment of the expectations of the periodic law was given
in the first place by the fresh determinations (Seubert, Dittmar,
and Arthur) of the atomic weight of platinum, which proved to be
nearly 196, if O = 16 (as Marignac, Brauner, and others propose);
in the second place, by the fact that Seubert proved that the
atomic weight of osmium is really less than that of platinum, and
approximately Os = 191; and, in the third place, by the fact that
after the researches of Krüss, Thorpe, and Laurie there was no
doubt that the atomic weight of gold is greater than that of
platinum--namely, nearly 197.
[28 bis] In Chapter XXII., Note 40, we gave the thermal data for
certain of the compounds of copper of the type CuX_{2}; we will
now cite certain data for the cuprous compounds of the type CuX,
which present an analogy to the corresponding compounds AgX and
AuX, some of which were investigated by Thomsen in his classical
work, 'Thermochemische Untersuchungen' (Vol. iii., 1883). The data
are given in the same manner as in the above-mentioned note:
R = Cu Ag Au
R + Cl +33 +29 +6
R + Br +25 +23 0
R + I +16 +14 -6
R + O +41 + 6 -?
Thus we see in the first place that gold, which possesses a much
smaller affinity than Ag, evolves far less heat than an equivalent
amount of copper, giving the same compound, and in the second
place that the combination of copper with one atom of oxygen
disengages more heat than its combination with one atom of a
halogen, whilst with silver the reverse is the case. This is
connected with the fact that Cu_{2}O is more stable under the
action of heat than Ag_{2}O.
[29] Heavy atoms and molecules, although they may present many points
of analogy, are more easily isolated; thus C_{16}H_{32}, although,
like C_{2}H_{4}, it combines with Br_{2}, and has a similar
composition, yet reacts with much greater difficulty than
C_{2}H_{4}, and in this it resembles gold; the heavy atoms and
molecules are, so to say, inert, and already saturated by
themselves. Gold in its higher grade of oxidation, Au_{2}O_{3},
presents feeble basic properties and weakly-developed acid
properties, so that this oxide of gold, Au_{2}O_{3}, may be
referred to the class of feeble acid oxides, like platinic oxide.
This is not the case in the highest known oxides of copper and
silver. But in the lower grade of oxidation, aurous oxide,
Au_{2}O, gold, like silver and copper, presents basic properties,
although they are not very pronounced. In this respect it stands
very close in its properties, although not in its types of
combination (AuX and AuX_{3}), to platinum (PtX_{2} and PtX_{4})
and its analogues.
As yet the general chemical characteristics of gold and its
compounds have not been fully investigated. This is partly due to
the fact that very few researches have been undertaken on the
compounds of this metal, owing to its inaccessibility for working
in large quantities. As the atomic weight of gold is high (Au =
197), the preparation of its compounds requires that it should be
taken in large quantities, which forms an obstacle to its being
fully studied. Hence the facts concerning the history of this
metal are rarely distinguished by that exactitude with which many
facts have been established concerning other elements more
accessible, and long known in use.
_In nature_ gold occurs in the primary and chiefly in quartzose rocks,
and especially in quartz veins, as in the Urals (at Berezoffsk), in
Australia, and in California. The native gold is extracted from these
rocks by subjecting them to a mechanical treatment consisting of crushing
and washing.[29 bis] Nature has already accomplished a similar
disintegration of the hard rocky matter containing gold.[30] These
disintegrated rocks, washed by rain and other water, have formed
gold-bearing deposits, which are known as _alluvial gold deposits_.
Gold-bearing soil is sometimes met with on the surface and sometimes
under the upper soil, but more frequently along the banks of dried-up
water-courses and running streams. The sand of many rivers contains,
however, a very small amount of gold, which it is not profitable to work;
for example, that of the Alpine rivers contains 5 parts of gold in
10,000,000 parts of sand. The richest gold deposits are those of Siberia,
especially in the southern parts of the Government of Yeniseisk, the
South Urals, Mexico, California, South Africa, and Australia, and then
the comparatively poorer alluvial deposits of many countries (Hungary,
the Alps, and Spain in Europe). The extraction of the gold from alluvial
deposits is based on the principle of levigation; the earth is washed,
while constantly agitated, by a stream of water, which carries away the
lighter portion of the earth, and leaves the coarser particles of the
rock and heavier particles of the gold, together with certain substances
which accompany it, in the washing apparatus. The extraction of this
_washed_ gold only necessitates mechanical appliances,[31] and it is not
therefore surprising that gold was known to savages and in the most
remote period of history. It sometimes occurs in crystals belonging to
the regular system, but in the majority of cases in nuggets or grains of
greater or less magnitude. It always contains silver (from very small
quantities up to 30 p.c., when it is called 'electrum') and certain other
metals, among which lead and rhodium are sometimes found.
[29 bis] Sonstadt (1872) showed that sea water, besides silver, always
contains gold. Munster (1892) showed that the water of the
Norwegian fiords contains about 5 milligrams of gold per ton (or 5
milliardths)--_i.e._ a quantity deserving practical attention, and
I think it may be already said that, considering the immeasurable
amount of sea water, in time means will be discovered for
profitably extracting gold from sea water by bringing it into
contact with substances capable of depositing gold upon their
surface. The first efforts might be made upon the extraction of
salt from sea water, and as the total amount of sea water may be
taken as about 2,000,000,000,000,000,000 tons, it follows that it
contains about 10,000 million tons of gold. The yearly production
of gold is about 200 tons for the whole world, of which about one
quarter is extracted in Russia. It is supposed that gold is
dissolved in sea water owing to the presence of iodides, which,
under the action of animal organisms, yield free iodine. It is
thought (as Professor Konovaloff mentions in his work upon 'The
Industries of the United States,' 1894) that iodine facilitates
the solution of the gold, and the organic matter its
precipitation. These facts and considerations to a certain extent
explain the distribution of gold in veins or rock fissures,
chiefly filled with quartz, because there is sufficient reason for
supposing that these rocks once formed the ocean bottom. R.
Dentrie, and subsequently Wilkinson, showed that organic
matter--for instance, cork--and pyrites are able to precipitate
gold from its solutions in that metallic form and state in which
it occurs in quartz veins, where (especially in the deeper parts
of vein deposits) gold is frequently found on the surface of
pyrites, chiefly arsenical pyrites. Kazantseff (in Ekaterinburg,
1891) even supposes, from the distribution of the gold in these
pyrites, that it occurred in solution as a compound of sulphide of
gold and sulphide of arsenic when it penetrated into the veins. It
is from such considerations that the origin of vein and pyritic
gold is, at the present time, attributed to the reaction of
solutions of this metal, the remains of which are seen in the gold
still present in sea water.
[30] However, in recent times, especially since about 1870, when
chlorine (either as a solution of the gas or as bleaching powder)
and bromine began to be applied to the extraction of
finely-divided gold from poor ores (previously roasted in order to
drive off arsenic and sulphur, and oxidise the iron), the
extraction of gold from quartz and pyrites, by the wet method,
increases from year to year, and begins to equal the amount
extracted from alluvial deposits. Since the nineties the _cyanide
process_ (Chapter XIII., Note 13 bis) has taken an important place
among the wet methods for extracting gold from its ores. It
consists in pouring a dilute solution of cyanide of potassium
(about 500 parts of water and 1 to 4 parts of cyanide of potassium
per 1,000 parts of ore, the amount of cyanide depending upon the
richness of the ore) and a mixture of it with NaCN, (_see_ Chapter
XIII., Note 12) over the crushed ore (which need not be roasted,
whilst roasting is indispensable in the chlorination process, as
otherwise the chlorine is used up in oxidising the sulphur,
arsenic, &c.) The gold is dissolved very rapidly even from
pyrites, where it generally occurs on the surface in such fine and
adherent particles that it either cannot be mechanically washed
away, or, more frequently is carried away by the stream of water,
and cannot be caught by mechanical means or by the mercury used
for catching the gold in the sluices. Chlorination had already
given the possibility of extracting the finest particles of gold;
but the cyanide process enables such pyrites to be treated as
could be scarcely worked by other means. The treatment of the
crushed ore by the KCN is carried on in simple wooden vats (coated
with paraffin or tar) with the greatest possible rapidity (in
order that the KCN solution should not have time to change) by a
method of systematic lixiviation, and is completed in 10 to 12
hours. The resultant solution of gold, containing AuK(CN)_{2}, is
decomposed either with freshly-made zinc filings (but when the
gold settles on the Zn, the cyanide solution reacts upon the Zn
with the evolution of H_{2} and formation of ZnH_{2}O_{2}) or by
sodium amalgam prepared at the moment of reaction by the action of
an electric current upon a solution of NaHO poured into a vessel
partially immersed in mercury (the NaCN is renewed continually by
this means). The silver in the ore passes into solution, together
with the gold, as in amalgamation.
[31] But the particles of gold are sometimes so small that a large
amount is lost during the washing. It is then profitable to have
recourse to the extraction by chlorine and KCN (Note 30).
In speaking of the extraction of gold the following remarks may
not be out of place:
In California advantage is taken of water supplied from high
altitudes in order to have a powerful head of water, with which
the rocks are directly washed away, thus avoiding the greater
portion of the mechanical labour required for the exploitation of
these deposits.
The last residues of gold are sometimes extracted from sand by
washing them with mercury, which dissolves the gold. The sand
mixed with water is caused to come into contact with mercury
during the washing. The mercury is then distilled.
Many sulphurous ores, even pyrites, contain a small amount of
gold. Compounds of gold with bismuth, BiAu_{2}, tellurium,
AuTe_{2} (calverite), &c., have been found, although rarely.
Among the minerals which accompany gold, and from which the
presence of gold may be expected, we may mention white quartz,
titanic and magnetic iron ores, and also the following, which are
of rarer occurrence: zircon, topaz, garnet, and such like. The
concentrated gold washings first undergo a mechanical treatment,
and the impure gold obtained is treated for pure gold by various
methods. If the gold contain a considerable amount of foreign
metals, especially lead and copper, it is sometimes cupelled, like
silver, so that the oxidisable metals may be absorbed by the cupel
in the form of oxides, but in every case the gold is obtained
together with silver, because the latter metal also is not
oxidised. Sometimes the gold is extracted by means of mercury,
that is, by amalgamation (and the mercury subsequently driven off
by distillation), or by smelting it with lead (which is afterwards
removed by oxidation) and processes like those employed for the
extraction of silver, because gold, like silver, does not oxidise,
is dissolved by lead and mercury, and is non-volatile. If copper
or any other metal contain gold and it be employed as an anode,
pure copper will be deposited upon the cathode, while all the gold
will remain at the anode as a slime. This method often amply
repays the whole cost of the process, since it gives, besides the
gold, a pure electrolytic copper.
_The separation of the silver_ from gold is generally carried on with
great precision, as the presence of the silver in the gold does not
increase its value for exchange, and it can be substituted by other less
valuable metals, so that the extraction of the silver, as a precious
metal, from its alloy with gold, is a profitable operation. This
separation is conducted by different methods. Sometimes the argentiferous
gold is melted in crucibles, together with a mixture of common salt and
powdered bricks. The greater portion of the silver is thus converted into
the chloride, which fuses and is absorbed by the slags, from which it may
be extracted by the usual methods. The silver is also extracted from gold
by treating it with boiling sulphuric acid, which does not act on the
gold but dissolves the silver. But if the alloy does not contain a large
proportion of silver it cannot be extracted by this method or at all
events the separation will be imperfect, and therefore a fresh amount of
silver is added (by fusion) to the gold, in such quantity that the alloy
contains twice as much silver as gold. The silver which is added is
preferably such as contains gold, which is very frequently the case. The
alloy thus formed is poured in a thin stream into water, by which means
it is obtained in a granulated form; it is then boiled with strong
sulphuric acid, three parts of acid being used to one part of alloy. The
sulphuric acid extracts all the silver without acting on the gold. It is
best, however, to pour off the first portion of the acid, which has
dissolved the silver, and then treat the residue of still imperfectly
pure gold with a fresh quantity of sulphuric acid. The gold is thus
obtained in the form of powder, which is washed with water until it is
quite free from silver. The silver is precipitated from the solution by
means of copper, so that cupric sulphate and metallic silver are
obtained. This process is carried out in many countries, as in Russia, at
the Government mints.
Gold is generally used alloyed with copper; since pure gold, like pure
silver, is very soft, and therefore soon worn away. In assaying or
determining the amount of pure gold in such an alloy it is usual to add
silver to the gold in order to make up an alloy containing three parts of
silver to one of gold (this is known as quartation because the alloy
contains 1/4 of gold), and the resultant alloy is treated with nitric
acid. If the silver be not in excess over the gold, it is not all
dissolved by the nitric acid, and this is the reason for the quartation.
The amount of pure gold (assay) is determined by weighing the gold which
remains after this treatment. English gold (= 22 carats) coinage is
composed of an alloy containing 91·66 p.c. of gold, but for many articles
gold is frequently used containing a larger amount of foreign metals.
_Pure gold_ may be obtained from gold alloys by dissolving in aqua
regia, and then adding ferrous sulphate to the solution or heating it
with a solution of oxalic acid. These deoxidising agents reduce the gold,
but not the other metals. The chlorine combined with the gold then acts
like free chlorine. The gold, thus reduced, is precipitated as an
exceedingly fine brown powder.[31 bis] It is then washed with water, and
fused with nitre or borax. Pure gold reflects a yellow light, and in the
form of very thin sheets (gold leaf), into which it can be hammered and
rolled,[31 tri] it transmits a bluish-green light. The specific gravity
of gold is about 19·5, the sp. gr. of gold coin is about 17·1. It fuses
at 1090°--at a higher temperature than silver--and can be drawn into
exceedingly fine wires or hammered into thin sheets. With its softness
and ductility, gold is distinguished for its tenacity, and a gold wire
two millimetres thick breaks only under a load of 68 kilograms. Gold
vaporises even at a furnace heat, and imparts a greenish colour to a
flame passing over it in a furnace. Gold alloys with copper almost
without changing its volume.[32] In its chemical aspect, gold presents,
as is already seen from its general characteristics given above, an
example of the so-called noble metals--_i.e._ it is incapable of being
oxidised at any temperature, and its oxide is decomposed when calcined.
Only chlorine and bromine combine directly with it at the ordinary
temperature, but many other metals and non-metals combine with it at a
red heat--for example, sulphur, phosphorus, and arsenic. Mercury
dissolves it with great ease. It dissolves in potassium cyanide in the
presence of air; a mixture of sulphuric acid with nitric acid dissolves
it with the aid of heat, although in small quantity. It is also soluble
in aqua regia and in selenic acid. Sulphuric, hydrochloric, nitric, and
hydrofluoric acids and the caustic alkalis do not act on gold, but a
mixture of hydrochloric acid with such oxidising agents as evolve
chlorine naturally dissolves it like aqua regia.[32 bis]
[31 bis] Schottländer (1893) obtained gold in a soluble colloid form
(the solution is violet) by the action of a mixture of solutions
of cerium acetate and NaHO upon a solution of AuCl_{3}. The gold
separates out from such a solution in exactly the same manner as
Ag does from the solution of colloid silver mentioned above. There
always remains a certain amount of a higher oxide of cerium,
CeO_{2}, in the solution--_i.e._ the gold is reduced by converting
the cerium into a higher grade of oxidation. Besides which Krüss
and Hofmann showed that sulphide of gold precipitated by the
action of H_{2}S upon a solution of AuKCy_{2} mixed with HCl
easily passes into a colloid solution after being properly washed
(like As_{2}S_{3}, CuS, &c., Chapter I., Note 57).
[31 tri] Gold-leaf is used for gilding wood (leather, cardboard, and
suchlike, upon which it is glued by means of varnish, &c.), and is
about 0·003 millimetre thick. It is obtained from thin sheets
(weighing at first about 1/4 grm. to a square inch), rolled
between gold rollers, by gradually hammering them (in packets of a
number at once) between sheets of moist (but not wet) parchment,
and then, after cutting them into four pieces, between a specially
prepared membrane, which, when at the right degree of moisture,
does not tear or stick together under the blows of the hammer.
[32] The formation of the alloys Cu + Zn, Cu + Sn, Cu + Bi, Cu + Sb,
Pb + Sb, Ag + Pb, Ag + Sn, Au + Zn, Au + Sn, &c., is accompanied
by a contraction (and evolution of heat). The formation of the
alloys Fe + Sb, Fe + Pb, Cu + Pb, Pb + Sn, Pb + Sb, Zn + Sb, Ag +
Cu, Au + Cu, Au + Pb, takes place with a certain increase in
volume. With regard to the alloys of gold, it may be mentioned
that gold is only slightly dissolved by mercury (about 0·06 p.c.,
Dudley, 1890); the remaining portion forms a granular alloy, whose
composition has not been definitely determined. Aluminium (and
silicon) also have the capacity of forming alloys with gold. The
presence of a small amount of aluminium lowers the melting point
of gold considerably (Roberts-Austen, 1892); thus the addition of
4 p.c. of aluminium lowers it by 14°·28, the addition of 10 p.c.
Al by 41°·7. The latter alloy is white. The alloy AuAl_{2} has a
characteristic purple colour, and its melting point is 32°·5 above
that of gold, which shows it to be a definite compound of the two
metals. The melting points of alloys richer in Al gradually fall
to 660°--that is, below that of aluminium (665°).
Heycock and Neville (1892), in studying the triple alloys of Au,
Cd, and Sn, observed a tendency in the gold to give compounds with
Cd, and by sealing a mixture of Au and Cd in a tube, from which
the air had been exhausted, and heating it, they obtained a grey
crystalline brittle definite alloy AuCd.
[32 bis] Calderon (1892), at the request of some jewellers,
investigated the cause of a peculiar alteration sometimes found on
the surface of dead-gold articles, there appearing brownish and
blackish spots, which widen and alter their form in course of
time. He came to the conclusion that these spots are due to the
appearance and development of peculiar micro-organisms
(Aspergillus niger and Micrococcus cimbareus) on the gold, spores
of which were found in abundance on the cotton-wool in which the
gold articles had been kept.
As regards the compounds of gold, they belong, as was said above, to the
types AuX_{3} and AuX. _Auric chloride_ or _gold trichloride_, AuCl_{3},
which is formed when gold is dissolved in aqua regia, belongs to the
former and higher of these types. The solution of this substance in water
has a yellow colour, and it may be obtained pure by evaporating the
solution in aqua regia to dryness, but not to the point of decomposition.
If the evaporation proceed to the point of crystallisation, a compound of
gold chloride and hydrochloric acid, AuHCl_{4}, is obtained, like the
allied compounds of platinum; but it easily parts with the acid and
leaves auric chloride, which fuses into a red-brown liquid, and then
solidifies to a crystalline mass. If dry chlorine be passed over gold in
powder it forms a mixture of aurous and auric chlorides, but the aurous
chloride is also decomposed by water into gold and auric chloride. Auric
chloride crystallises from its solutions as AuCl_{3},2H_{2}O, which
easily loses water, and the dry chloride loses two-thirds of its chlorine
at 185°, forming aurous chloride, whilst above 300° the latter chloride
also loses its chlorine and leaves metallic gold. Auric chloride is the
usual form in which gold occurs in solutions, and in which its salts are
used in the arts and for chemical purposes. It is soluble in water,
alcohol, and ether. Light has a reducing action on these solutions, and
after a time metallic gold is deposited upon the sides of vessels
containing the solution. Hydrogen when nascent, and even in a gaseous
form, reduces gold from this solution to a metallic state. The reduction
is more conveniently and usually effected by ferrous sulphate, and in
general by the action of ferrous salts.[33]
[33] Stannous chloride as a reducing agent also acts on auric chloride,
and gives a red precipitate known as _purple of Cassius_. This
substance, which probably contains a mixture or compound of aurous
oxide and tin oxide, is used as a red pigment for china and glass.
Oxalic acid, on heating, reduces metallic gold from its salts, and
this property may be taken advantage of for separating it from its
solutions. The oxidation which then takes place in the presence of
water may be expressed by the following equation: 2AuCl_{3} +
3C_{2}H_{2}O_{4} = 2Au + 6HCl + 6CO_{2}. Nearly all organic
substances have a reducing action on gold, and solutions of gold
leave a violet stain on the skin.
Auric chloride, like platinic chloride, is distinguished for its
clearly-developed property of forming double salts. These double
salts, as a rule, belong to the type AuMCl_{4}. The compound of
auric chloride with hydrochloric acid mentioned above evidently
belongs to the same type. The compounds 2KAuCl_{4},5H_{2}O,
NaAuCl_{4},2H_{2}O, AuNH_{4}Cl_{4},H_{2}O,
Mg(AuCl_{4})_{2},2H_{2}O, and the like are easily crystallised in
well-formed crystals. Wells, Wheeler, and Penfield (1892) obtained
RbAuCl_{4} (reddish yellow) and CsAuCl_{4} (golden yellow), and
corresponding bromides (dark coloured). AuBr_{3} is extremely like
the chloride. Auric cyanide is obtained easily in the form of a
double salt of potassium, KAu(CN)_{4} by mixing saturated and hot
solutions of potassium cyanide with auric chloride and then
cooling.
If a solution of potassium hydroxide be added to a solution of auric
chloride, a precipitate is first formed, which re-dissolves in an excess
of the alkali. On being evaporated under the receiver of an air-pump,
this solution yields yellow crystals, which present the same composition
as the double salts AuMCl_{4}, with the substitution of the chlorine by
oxygen--that is to say, _potassium aurate_, AuKO_{2}, is formed in
crystals containing 3H_{2}O. The solution has a distinctly alkaline
reaction. _Auric oxide_, Au_{2}O_{3}, separates when this alkaline
solution is boiled with an excess of sulphuric acid. But it then still
retains some alkali; however, it may be obtained in a pure state as a
brown powder by dissolving in nitric acid and diluting with water. The
brown powder decomposes below 250° into gold and oxygen. It is insoluble
in water and in many acids, but it dissolves in alkalis, which shows the
acid character of this oxide. An hydroxide, Au(OH)_{3} may be obtained as
a brown powder by adding magnesium oxide to a solution of auric chloride
and treating the resultant precipitate of magnesium aurate with nitric
acid. This hydroxide loses water at 100°, and gives auric oxide.[34]
[34] If ammonia be added to a solution of auric chloride, it forms a
yellow precipitate of the so-called fulminating gold, which
contains gold, chlorine, hydrogen, nitrogen, and oxygen, but its
formula is not known with certainty. It is probably a sort of
ammonio-metallic compound, Au_{2}O_{3},4NH_{3}, or amide (like the
mercury compound). This precipitate explodes at 140°, but when
left in the presence of solutions containing ammonia it loses all
its chlorine and becomes non-explosive. In this form the
composition Au_{2}O_{3},2NH_{3},H_{2}O is ascribed to it, but this
is uncertain. Auric sulphide, Au_{2}S_{3}, is obtained by the
action of hydrogen sulphide on a solution of auric chloride, and
also directly by fusing sulphur with gold. It has an acid
character, and therefore dissolves in sodium and ammonium
sulphides.
The starting-point of the compounds of the type AuX[35] is _gold
monochloride_ or _aurous chloride_, AuCl, which is formed, as mentioned
above, by heating auric chloride at 185°. Aurous chloride forms a
yellowish-white powder; this, when heated with water, is decomposed into
metallic gold and auric chloride, which passes into solution: 3AuCl =
AuCl_{3} + 2Au. This decomposition is accelerated by the action of light.
Hence it is obvious that the compounds corresponding with aurous oxide
are comparatively unstable. But this only refers to the simple compounds
AuX; some of the complex compounds, on the contrary, form the most stable
compounds of gold. Such, for example, is the cyanide of gold and
potassium, AuK(CN)_{2}. It is formed, for instance, when finely-divided
gold dissolves in the presence of air in a solution of potassium cyanide:
4KCN + 2Au + H_{2}O + O = 2KAu(CN)_{2} + 2KHO (this reaction also
proceeds with solid pieces of gold, although very slowly). The same
compound is formed in solution when many compounds of gold are mixed with
potassium cyanide, because if a higher compound of gold be taken, it is
reduced by the potassium cyanide into aurous oxide, which dissolves in
potassium cyanide and forms KAu(CN)_{2}. This substance is soluble in
water, and gives a colourless solution, which can be kept for a long
time, and is employed in electro-gilding--that is, for coating other
metallic objects with a layer of gold, which is deposited if the object
be connected with the negative pole of a battery and the positive pole
consist of a gold plate. When an electric current is passed between them,
the gold from the latter will dissolve, whilst a coating of gold from the
solution will be deposited on the object.
[35] Many double salts of suboxide of gold belong to the type AuX--for
instance, the cyanide corresponding to the type AuKX_{2}, like
PtK_{2}X_{4}, with which we became acquainted in the last chapter.
We will enumerate several of the representatives of this class of
compounds. If auric chloride, AuCl_{3}, be mixed with a solution
of sodium thiosulphate, the gold passes into a colourless
solution, which deposits colourless crystals, containing a double
thiosulphate of gold and sodium, which are easily soluble in water
but are precipitated by alcohol. The composition of this salt is
Na_{3}Au(S_{2}O_{3})_{2},2H_{2}O. If the sodium thiosulphate be
represented as NaS_{2}O_{3}Na, the double salt in question will be
AuNa(S_{2}O_{3}Na)_{2},2H_{2}O, according to the type AuNaX_{2}.
The solution of this colourless and easily crystallisable salt has
a sweet taste, and the gold is not separated from it either by
ferrous sulphate or oxalic acid. This salt, which is known as
_Fordos and Gelis's salt_, is used in medicine and photography. In
general, aurous oxide exhibits a distinct inclination to the
formation of similar double salts, as we saw also with
PtX_{2}--for example, it forms similar salts with sulphurous acid.
Thus if a solution of sodium sulphite be gradually added to a
solution of oxide of gold in sodium hydroxide, the precipitate at
first formed re-dissolves to a colourless solution, which contains
the double salt Na_{3}Au(SO_{3})_{2} = AuNa(SO_{3}Na)_{2}. The
solution of this salt, when mixed with barium chloride, first
forms a precipitate of barium sulphite, and then a red barium
double salt which corresponds with the above sodium salt.
The oxygen compound of the type AuX, _aurous oxide_, Au_{2}O, is
obtained as a greenish violet powder on mixing aurous chloride
with potassium chloride in the cold. With hydrochloric acid this
oxide gives gold and auric chloride, and when heated it easily
splits up into oxygen and metallic gold.
APPENDIX I
AN ATTEMPT TO APPLY TO CHEMISTRY ONE OF THE PRINCIPLES
OF NEWTON'S NATURAL PHILOSOPHY
BY PROFESSOR MENDELÉEFF
A LECTURE DELIVERED AT THE ROYAL INSTITUTION OF GREAT BRITAIN
ON FRIDAY, MAY 31, 1889
Nature, inert to the eyes of the ancients, has been revealed to us as
full of life and activity. The conviction that motion pervaded all
things, which was first realised with respect to the stellar universe,
has now extended to the unseen world of atoms. No sooner had the human
understanding denied to the earth a fixed position and launched it along
its path in space, than it was sought to fix immovably the sun and the
stars. But astronomy has demonstrated that the sun moves with unswerving
regularity through the star-set universe at the rate of about 50
kilometres per second. Among the so-called fixed stars are now discerned
manifold changes and various orders of movement. Light, heat,
electricity--like sound--have been proved to be modes of motion; to the
realisation of this fact modern science is indebted for powers which have
been used with such brilliant success, and which have been expounded so
clearly at this lecture table by Faraday and by his successors. As, in
the imagination of Dante, the invisible air became peopled with spiritual
beings, so before the eyes of earnest investigators, and especially
before those of Clerk Maxwell, the invisible mass of gases became peopled
with particles: their rapid movements, their collisions, and impacts
became so manifest that it seemed almost possible to count the impacts
and determine many of the peculiarities or laws of their collisions. The
fact of the existence of these invisible motions may at once be made
apparent by demonstrating the difference in the rate of diffusion through
porous bodies of the light and rapidly moving atoms of hydrogen and the
heavier and more sluggish particles of air. Within the masses of liquid
and of solid bodies we have been forced to acknowledge the existence of
persistent though limited motion of their ultimate particles, for
otherwise it would be impossible to explain, for example, the celebrated
experiments of Graham on diffusion through liquid and colloidal
substances. If there were, in our times, no belief in the molecular
motion in solid bodies, could the famous Spring have hoped to attain any
result by mixing carefully-dried powders of potash, saltpetre and sodium
acetate, in order to produce, by pressure, a chemical reaction between
these substances through the interchange of their metals, and have
derived, for the conviction of the incredulous, a mixture of two
hygroscopic though solid salts--sodium nitrate and potassium acetate?
In these invisible and apparently chaotic movements, reaching from the
stars to the minutest atoms, there reigns, however, a harmonious order
which is commonly mistaken for complete rest, but which is really a
consequence of the conservation of that dynamic equilibrium which was
first discerned by the genius of Newton, and which has been traced by his
successors in the detailed analysis of the particular consequences of the
great generalisation, namely, relative immovability in the midst of
universal and active movement.
But the unseen world of chemical changes is closely analogous to the
visible world of the heavenly bodies, since our atoms form distinct
portions of an invisible world, as planets, satellites, and comets form
distinct portions of the astronomer's universe; our atoms may therefore
be compared to the solar systems, or to the systems of double or of
single stars: for example, ammonia (NH_{3}) may be represented in the
simplest manner by supposing the sun, nitrogen, surrounded by its planets
of hydrogen; and common salt (NaCl) may be looked on as a double star
formed of sodium and chlorine. Besides, now that the indestructibility of
the elements has been acknowledged, chemical changes cannot otherwise be
explained than as changes of motion, and the production by chemical
reactions of galvanic currents, of light, of heat, of pressure, or of
steam power, demonstrates visibly that the processes of chemical reaction
are inevitably connected with enormous though unseen displacements,
originating in the movements of atoms in molecules. Astronomers and
natural philosophers, in studying the visible motions of the heavenly
bodies and of matter on the earth, have understood and have estimated the
value of this store of energy. But the chemist has had to pursue a
contrary course. Observing in the physical and mechanical phenomena which
accompany chemical reactions the quantity of energy manifested by the
atoms and molecules, he is constrained to acknowledge that within the
molecules there exist atoms in motion, endowed with an energy which, like
matter itself, is neither being created nor capable of being destroyed.
Therefore, in chemistry, we must seek dynamic equilibrium not only
between the molecules, but also in their midst among their component
atoms. Many conditions of such equilibrium have been determined, but much
remains to be done, and it is not uncommon, even in these days, to find
that some chemists forget that there is the possibility of motion in the
interior of molecules, and therefore represent them as being in a
condition of death-like inactivity.
Chemical combinations take place with so much ease and rapidity, possess
so many special characteristics, and are so numerous, that their
simplicity and order were for a long time hidden from investigators.
Sympathy, relationship, all the caprices or all the fancifulness of human
intercourse, seemed to have found complete analogies in chemical
combinations, but with this difference, that the characteristics of the
material substances--such as silver, for example, or of any other
body--remain unchanged in every subdivision from the largest masses to
the smallest particles, and consequently these characteristics must be
properties of the particles. But the world of heavenly luminaries
appeared equally fanciful at man's first acquaintance with it, so much
so, that the astrologers imagined a connection between the
individualities of men and the conjunctions of planets. Thanks to the
genius of Lavoisier and of Dalton, man has been able, in the unseen world
of chemical combinations, to recognise laws of the same simple order as
those which Copernicus and Kepler proved to exist in the planetary
universe. Man discovered, and continues every hour to discover, _what_
remains unchanged in chemical evolution, and _how_ changes take place in
combinations of the unchangeable. He has learned to predict, not only
what possible combinations may take place, but also the very existence of
atoms of unknown elementary substances, and has besides succeeded in
making innumerable practical applications of his knowledge to the great
advantage of his race, and has accomplished this notwithstanding that
notions of sympathy and affinity still preserve a strong vitality in
science. At present we cannot apply Newton's principles to chemistry,
because the soil is only being now prepared. The invisible world of
chemical atoms is still waiting for the creator of chemical mechanics.
For him our age is collecting a mass of materials, the inductions of
well-digested facts, and many-sided inferences similar to those which
existed for Astronomy and Mechanics in the days of Newton. It is well
also to remember that Newton devoted much time to chemical experiments,
and while considering questions of celestial mechanics, persistently kept
in view the mutual action of those infinitely small worlds which are
concerned in chemical evolutions. For this reason, and also to maintain
the unity of laws, it seems to me that we must, in the first instance,
seek to harmonise the various phases of contemporary chemical theories
with the immortal principles of the Newtonian natural philosophy, and so
hasten the advent of true chemical mechanics. Let the above
considerations serve as my justification for the attempt which I propose
to make to act as a champion of the universality of the Newtonian
principles, which I believe are competent to embrace every phenomenon in
the universe, from the rotation of the fixed stars to the interchanges of
chemical atoms.
In the first place I consider it indispensable to bear in mind that, up
to quite recent times, only a one-sided affinity has been recognised in
chemical reactions. Thus, for example, from the circumstance that red-hot
iron decomposes water with the evolution of hydrogen, it was concluded
that oxygen had a greater affinity for iron than for hydrogen. But
hydrogen, in presence of red-hot iron scale, appropriates its oxygen and
forms water, whence an exactly opposite conclusion may be formed.
During the last ten years a gradual, scarcely perceptible, but most
important change has taken place in the views, and consequently in the
researches, of chemists. They have sought everywhere, and have always
found, systems of conservation or dynamic equilibrium substantially
similar to those which natural philosophers have long since discovered in
the visible world, and in virtue of which the position of the heavenly
bodies in the universe is determined. There where one-sided affinities
only were at first detected, not only secondary or lateral ones have been
found, but even those which are diametrically opposite; yet among these,
dynamical equilibrium establishes itself not by excluding one or other of
the forces, but regulating them all. So the chemist finds in the flame of
the blast furnace, in the formation of every salt, and, with especial
clearness, in double salts and in the crystallisation of solutions, not a
fight ending in the victory of one side, as used to be supposed, but the
conjunction of forces; the peace of dynamic equilibrium resulting from
the action of many forces and affinities. Carbonaceous matters, for
example, burn at the expense of the oxygen of the air, yielding a
quantity of heat, and forming products of combustion, in which it was
thought that the affinities of the oxygen with the combustible elements
were satisfied. But it appeared that the heat of combustion was competent
to decompose these products, to dissociate the oxygen from the
combustible elements, and therefore to explain combustion fully it is
necessary to take into account the equilibrium between opposite
reactions, between those which evolve and those which absorb heat.
In the same way, in the case of the solution of common salt in water, it
is necessary to take into account, on the one hand, the formation of
compound particles generated by the combination of salt with water, and,
on the other, the disintegration or scattering of the new particles
formed, as well as of these originally contained. At present we find two
currents of thought, apparently antagonistic to each other, dominating
the study of solutions: according to the one, solution seems a mere act
of building up or association; according to the other, it is only
dissociation or disintegration. The truth lies, evidently, between these
views; it lies, as I have endeavoured to prove by my investigations into
aqueous solutions, in the dynamic equilibrium of particles tending to
combine and also to fall asunder. The large majority of chemical
reactions which appeared to act victoriously along one line have been
proved capable of acting as victoriously even along an exactly opposite
line. Elements which utterly decline to combine directly may often be
formed into comparatively stable compounds by indirect means, as, for
example, in the case of chlorine and carbon; and consequently the
sympathies and antipathies which it was thought to transfer from human
relations to those of atoms should be laid aside until the mechanism of
chemical relations is explained. Let us remember, however, that chlorine,
which does not form with carbon the chloride of carbon, is strongly
absorbed, or, as it were, dissolved, by carbon, which leads us to suspect
incipient chemical action even in an external and purely surface contact,
and involuntarily gives rise to conceptions of that unity of the forces
of nature which has been so energetically insisted on by Sir William
Grove and formulated in his famous paradox. Grove noticed that platinum,
when fused in the oxyhydrogen flame, during which operation water is
formed, when allowed to drop into water decomposes the latter and
produces the explosive oxyhydrogen mixture. The explanation of this
paradox, as of many others which arose during the period of chemical
renaissance, has led, in our time, to the promulgation by Henri
Sainte-Claire Deville of the conception of dissociation and of
equilibrium, and has recalled the teaching of Berthollet, which,
notwithstanding its brilliant confirmation by Heinrich Rose and Dr.
Gladstone, had not, up to that period, been included in received chemical
views.
Chemical equilibrium in general, and dissociation in particular, are now
being so fully worked out in detail, and supplied in such various ways,
that I do not allude to them to develop, but only use them as examples by
which to indicate the correctness of a tendency to regard chemical
combinations from points of view differing from those expressed by the
term hitherto appropriated to define chemical forces, namely, 'affinity.'
Chemical equilibria, dissociation, the speed of chemical reactions,
thermochemistry, spectroscopy, and, more than all, the determination of
the influence of masses and the search for a connection between the
properties and weights of atoms and molecules--in one word, the vast mass
of the most important chemical researches of the present day--clearly
indicate the near approach of the time when chemical doctrines will
submit fully and completely to the doctrine which was first announced in
the _Principia_ of Newton.
In order that the application of these principles may bear fruit it is
evidently insufficient to assume that statical equilibrium reigns alone
in chemical systems or chemical molecules: it is necessary to grasp the
conditions of possible states of dynamical equilibria, and to apply to
them kinetic principles. Numerous considerations compel us to renounce
the idea of statical equilibrium in molecules, and the recent yet
strongly-supported appeals to dynamic principles constitute, in my
opinion, the foundation of the modern teaching relating to atomicity, or
the valency of the elements, which usually forms the basis of
investigations into organic or carbon compounds.
This teaching has led to brilliant explanations of very many chemical
relations and to cases of isomerism, or the difference in the properties
of substances having the same composition. It has been so fruitful in its
many applications and in the foreshadowing of remote consequences,
especially respecting carbon compounds, that it is impossible to deny its
claims to be ranked as a great achievement of chemical science. Its
practical application to the synthesis of many substances of the most
complicated composition entering into the structure of organised bodies,
and to the creation of an unlimited number of carbon compounds, among
which the colours derived from coal tar stand prominently forward,
surpass the synthetical powers of Nature itself. Yet this teaching, as
applied to the structure of carbon compounds, is not on the face of it
directly applicable to the investigation of other elements, because in
examining the first it is possible to assume that the atoms of carbon
have always a definite and equal number of affinities, whilst in the
combinations of other elements this is evidently inadmissible. Thus, for
example, an atom of carbon yields only one compound with four atoms of
hydrogen and one with four atoms of chlorine in the molecule, whilst the
atoms of chlorine and hydrogen unite only in the proportions of one to
one. Simplicity is here evident, and forms a point of departure from
which it is easy to move forward with firm and secure tread. Other
elements are of a different nature. Phosphorus unites with three and with
five atoms of chlorine, and consequently the simplicity and sharpness of
the application of structural conceptions are lost. Sulphur unites only
with two atoms of hydrogen, but with oxygen it enters into higher orders
of combination. The periodic relationship which exists among all the
properties of the elements--such, for example, as their ability to enter
into various combinations--and their atomic weights, indicate that this
variation in atomicity is subject to one perfectly exact and general law,
and it is only carbon and its near analogues which constitute cases of
permanently preserved atomicity. It is impossible to recognise as
constant and fundamental properties of atoms, powers which, in substance,
have proved to be variable. But by abandoning the idea of permanence, and
of the constant saturation of affinities--that is to say, by
acknowledging the possibility of free affinities--many retain a
comprehension of the atomicity of the elements 'under given conditions;'
and on this frail foundation they build up structures composed of
chemical molecules, evidently only because the conception of manifold
affinities gives, at once, a simple statical method of estimating the
composition of the most complicated molecules.
I shall enter neither into details, nor into the various consequences
following from these views, nor into the disputes which have sprung up
respecting them (and relating especially to the number of isomerides
possible on the assumption of free affinities), because the foundation or
origin of theories of this nature suffers from the radical defect of
being in opposition to dynamics. The molecule, as even Laurent expressed
himself, is represented as an architectural structure, the style of which
is determined by the fundamental arrangement of a few atoms, whilst the
decorative details, which are capable of being varied by the same forces,
are formed by the elements entering into the combination. It is on this
account that the term 'structural' is so appropriate to the contemporary
views of the above order, and that the 'structuralists' seek to justify
the tetrahedric, plane, or prismatic disposition of the atoms of carbon
in benzene. It is evident that the consideration relates to the statical
position of atoms and molecules and not to their kinetic relations. The
atoms of the structural type are like the lifeless pieces on a chess
board: they are endowed but with the voices of living beings, and are not
those living beings themselves; acting, indeed, according to laws, yet
each possessed of a store of energy which, in the present state of our
knowledge, must be taken into account.
In the days of Haüy, crystals were considered in the same statical and
structural light, but modern crystallographers, having become more
thoroughly acquainted with their physical properties and their actual
formation, have abandoned the earlier views, and have made their
doctrines dependent on dynamics.
The immediate object of this lecture is to show that, starting with
Newton's third law of motion, it is possible to preserve to chemistry all
the advantages arising from structural teaching, without being obliged to
build up molecules in solid and motionless figures, or to ascribe to
atoms definite limited valencies, directions of cohesion, or affinities.
The wide extent of the subject obliges me to treat only a small portion
of it, namely of _substitutions_, without specially considering
combinations and decompositions, and even then limiting myself to the
simplest examples, which, however, will throw open prospects embracing
all the natural complexity of chemical relations. For this reason, if it
should prove possible to form groups similar, for example, to H_{4} or
CH_{6} as the remnants of molecules CH_{4} or C_{2}H_{7} we shall not
pause to consider them, because, as far as we know, they fall asunder
into two parts, H_{2} + H_{2} or CH_{4} + H_{2}, as soon as they are even
temporarily formed, and are incapable of separate existence, and
therefore can take no part in the elementary act of substitution. With
respect to the simplest molecules which we shall select--that is to say,
those of which the parts have no separate existence, and therefore cannot
appear in substitutions--we shall consider them according to the periodic
law, arranging them in direct dependence on the atomic weight of the
elements.
Thus, for example, the molecules of the simplest hydrogen compounds--
HF H_{2}O H_{3}N H_{4}C
hydrofluoric acid water ammonia methane
correspond with elements the atomic weights of which decrease
consecutively
F = 19, O = 16, N = 14, C = 12.
Neither the arithmetical order (1, 2, 3, 4 atoms of hydrogen) nor the
total information we possess respecting the elements will permit us to
interpolate into this typical series one more additional element; and
therefore we have here, for hydrogen compounds, a natural base on which
are built up those simple chemical combinations which we take as typical.
But even they are competent to unite with each other, as we see, for
instance, in the property which hydrofluoric acid has of forming a
hydrate--that is, of combining with water; and a similar attribute of
ammonia, resulting in the formation of a caustic alkali, NH_{3},H_{2}O,
or NH_{4}OH.
Having made these indispensable preliminary observations, I may now
attack the problem itself and attempt to explain the so-called structure
or rather construction, of molecules--that is to say, their constitution
and transformations--without having recourse to the teaching of
'structuralists,' but on Newton's dynamical principles.
Of Newton's three laws of motion, only the third can be applied directly
to chemical molecules when regarded as systems of atoms among which it
must be supposed that there exist common influences or forces, and
resulting compounded relative motions. Chemical reactions of every kind
are undoubtedly accomplished by changes in these internal movements,
respecting the nature of which nothing is known at present, but the
existence of which the mass of evidence collected in modern times forces
us to acknowledge as forming part of the common motion of the universe,
and as a fact further established by the circumstance that chemical
reactions are always characterised by changes of volume or the relations
between the atoms or the molecules. Newton's third law, which is
applicable to every system, declares that, 'action is also associated
with reaction, and is equal to it.' The brevity of conciseness of this
axiom was, however, qualified by Newton in a more expanded statement,
'the action of bodies one upon another are always equal, and in opposite
directions.' This simple fact constitutes the point of departure for
explaining dynamic equilibrium--that is to say, systems of conservancy.
It is capable of satisfying even the dualists, and of explaining, without
additional assumptions, the preservation of those chemical types which
Dumas, Laurent, and Gerhardt created unit types, and those views of
atomic combinations which the structuralists express by atomicity or the
valency of the elements, and, in connection with them, the various
numbers of affinities. In reality, if a system of atoms or a molecule be
given, then in it, according to the third law of Newton, each portion of
atoms acts on the remaining portion in the same manner, and with the same
force as the second set of atoms acts on the first. We infer directly
from this consideration that both sets of atoms, forming a molecule, are
not only equivalent with regard to themselves, as they must be according
to Dalton's law, but also that they may, if united, replace each other.
Let there be a molecule containing atoms A B C, it is clear that,
according to Newton's law, the action of A on B C must be equal to the
action of B C on A, and if the first action is directed on B C, then the
second must be directed on A, and consequently then, where A can exist in
dynamic equilibrium, B C may take its place and act in a like manner. In
the same way the action of C is equal to the action of A B. In one word
every two sets of atoms forming a molecule are equivalent to each other,
and may take each other's place in other molecules, or, having the power
of balancing each other, the atoms or their complements are endowed with
the power of replacing each other. Let us call this consequence of an
evident axiom 'the principle of substitution,' and let us apply it to
those typical forms of hydrogen compounds which we have already
discussed, and which, on account of their simplicity, and regularity,
have served as starting-points of chemical argument long before the
appearance of the doctrine of structure.
In the type of hydrofluoric acid, HF, or in systems of double stars, are
included a multitude of the simplest molecules. It will be sufficient for
our purpose to recall a few: for example, the molecules of chlorine,
Cl_{2}, and of hydrogen, H_{2}, and hydrochloric acid, HCl, which is
familiar to all in aqueous solution as spirits of salt, and which has
many points of resemblance with HF, HBr, HI. In these cases division into
two parts can only be made in one way, and therefore the principle of
substitution renders it probable that exchanges between the chlorine and
the hydrogen can take place, if they are competent to unite with each
other. There was a time when no chemist would even admit the idea of any
such action; it was then thought that the power of combination indicated
a polar difference of the molecules in combination, and this thought set
aside all idea of the substitution of one component element by another.
Thanks to the observations and experiments of Dumas and Laurent fifty
years ago, such fallacies were dispelled, and in this manner the
principle of substitution was exhibited. Chlorine and bromine acting on
many hydrogen compounds, occupy immediately the place of their hydrogen,
and the displaced hydrogen, with another atom of chlorine or bromine,
forms hydrochloric acid or bromide of hydrogen. This takes place in all
typical hydrogen compounds. Thus chlorine acts on this principle on
gaseous hydrogen--reaction, under the influence of light, resulting in
the formation of hydrochloric acid. Chlorine acting on the alkalis,
constituted similarly to water, and even on water itself--only, however,
under the influence of light and only partially because of the
instability of HClO--forms by this principle bleaching salts, which are
the same as the alkalis, but with their hydrogen replaced by chlorine. In
ammonia and in methane, chlorine can also replace the hydrogen. From
ammonia is formed in this manner the so-called chloride of nitrogen,
NCl_{3}, which decomposes very readily with violent explosion on account
of the evolved gases, and falls asunder as chlorine and nitrogen. Out of
marsh gas, or methane, CH_{4}, may be obtained consecutively, by this
method, every possible substitution, of which chloroform, CHCl_{3}, is
the best known, and carbon tetrachloride, CCl_{4}, the most instructive.
But by virtue of the fact that chlorine and bromine act, in the manner
shown, on the simplest typical hydrogen compounds, their action on the
more complicated ones may be assumed to be the same. This can be easily
demonstrated. The hydrogen of benzene, C_{6}H_{6}, reacts feebly under
the influence of light on liquid bromine, but Gustavson has shown that
the addition of the smallest quantity of metallic aluminium causes
energetic action and the evolution of large volumes of hydrogen bromide.
If we pass on to the second typical hydrogen compound--that is to say,
water--its molecule, HOH, may be split up in two ways: either into an
atom of hydrogen and a semi-molecule of hydrogen peroxide, HO, or into
oxygen, O, and two atoms of hydrogen, H; and therefore, according to the
principle of substitution, it is evident that one atom of hydrogen can
exchange with hydrogen oxide, HO, and two atoms of hydrogen, H, with one
atom of oxygen, O.
Both these forms of substitution will constitute methods of
oxidation--that is to say, of the entrance of oxygen into the compound--a
reaction which is so common in nature as well as in the arts, taking
place at the expense of the oxygen of the air or by the aid of various
oxidising substances or bodies which part easily with their oxygen. There
is no occasion to reckon up the unlimited number of cases of such
oxidising reactions. It is sufficient to state that in the first of these
oxygen is directly transferred, and the position, the chemical function,
which hydrogen originally occupied, is, after the substitution, occupied
by the hydroxyl. Thus ammonia, NH_{3}, yields hydroxylamine, NH_{2}(OH),
a substance which retains many of the properties of ammonia.
Methane and a number of other hydrocarbons yield, by substitution of the
hydrogen by its oxide, methyl alcohol, CH_{3}(OH), and other alcohols.
The substitution of one atom of oxygen for two atoms of hydrogen is
equally common with hydrogen compounds. By this means alcoholic liquids
containing ethyl alcohol, or spirits of wine, C_{2}H_{5}(OH), are
oxidised until they become vinegar, or acetic acid, C_{2}H_{3}O(OH). In
the same way caustic ammonia, or the combination of ammonia with water,
NH_{3},H_{2}O, or NH_{4}(OH), which contains a great deal of hydrogen, by
oxidation exchanges four atoms of hydrogen for two atoms of oxygen, and
becomes converted into nitric acid, NO_{2}(OH). This process of
conversion of ammonium salts into saltpetre goes on in the fields every
summer, and with especial rapidity in tropical countries. The method by
which this is accomplished, though complex, though involving the agency
of all-permeating micro-organisms, is, in substance, the same as that by
which alcohol is converted into acetic acid, or glycol,
C_{2}H_{4}(OH)_{2}, into oxalic acid, if we view the process of oxidation
in the light of the Newtonian principles.
But while speaking of the application of the principle of substitution
to water, we need not multiply instances, but must turn our attention to
two special circumstances which are closely connected with the very
mechanism of substitutions.
In the first place, the replacement of two atoms of hydrogen by one atom
of oxygen may take place in two ways, because the hydrogen molecule is
composed of two atoms, and therefore, under the influence of oxygen, the
molecule forming water may separate before the oxygen has time to take
its place. It is for this reason that we find, during the conversion of
alcohol into acetic acid, that there is an interval during which is
formed aldehyde, C_{2}H_{4}O, which, as its very name implies, is
'alcohol dehydrogenatum,' or alcohol deprived of hydrogen. Hence aldehyde
combined with hydrogen yields alcohol; and united to oxygen, acetic acid.
For the same reason there should be, and there actually are, intermediate
products between ammonia and nitric acid, NO_{2}(HO), containing either
less hydrogen than ammonia, less oxygen than nitric acid, or less water
than caustic ammonia. Accordingly we find, among the products of the
deoxidation of nitric acid and the oxidation of ammonia, not only
hydroxylamine, but also nitrous oxide, nitrous and nitric anhydrides.
Thus, the production of nitrous acid results from the removal of two
atoms of hydrogen from caustic ammonia and the substitution of the oxygen
for the hydrogen, NO(OH); or by the substitution, in ammonia, of three
atoms of hydrogen by hydroxyl, N(OH)_{3}, and by the removal of water:
N(OH)_{3} - H_{2}O = NO(OH). The peculiarities and properties of nitrous
acid--as, for instance, its action on ammonia and its conversion, by
oxidation, into nitric acid--are thus clearly revealed.
On the other hand, in speaking of the principle of substitution as
applied to water, it is necessary to observe that hydrogen and hydroxyl,
H and OH, are not only competent to unite, but also to form combinations
with themselves, and thus become H_{2} and H_{2}O_{2}; and such are
hydrogen and the peroxide thereof. In general, if a molecule A B exists,
then molecules A A and B B can exist also. A direct reaction of this kind
does not, however, take place in water, therefore undoubtedly, at the
moment of formation, hydrogen reacts on hydrogen peroxide, as we can show
at once by experiment; and further because hydrogen peroxide, H_{2}O_{2},
exhibits a structure containing a molecule of hydrogen, H_{2}, and one of
oxygen, O_{2}, either of which is capable of separate existence. The
fact, however, may now be taken as thoroughly established, that, at the
moment of combustion of hydrogen or of the hydrogen compounds, hydrogen
peroxide is always formed, and not only so, but in all probability its
formation invariably precedes the formation of water. This was to be
expected as a consequence of the law of Avogadro and Gerhardt, which
leads us to expect this sequence in the case of equal interactions of
volumes of vapours and gases; and in hydrogen peroxide we actually have
such equal volumes of the elementary gases.
The instability of hydrogen peroxide--that is to say, the ease with
which it decomposes into water and oxygen, even at the mere contact of
porous substances--accounts for the circumstance that it does not form a
permanent product of combustion, and is not produced during the
decomposition of water. I may mention this additional consideration that,
with respect to hydrogen peroxide, we may look for its effecting still
further substitutions of hydrogen by means of which we may expect to
obtain still more highly oxidised water compounds, such as H_{2}O_{3} and
H_{2}O_{4}. These Schönbein and Bunsen have long been seeking, and
Berthelot is investigating them at present. It is probable, however, that
the reaction will stop at the last compound, because we find that, in a
number of cases, the addition of four atoms of oxygen seems to form a
limit. Thus, OsO_{4}, KClO_{4}, KMnO_{4}, K_{2}SO_{4}, Na_{3}PO_{4}, and
such like, represent the highest grades of oxidation.[1]
[1] Because more than four atoms of hydrogen never unite with one atom
of the elements, and because the hydrogen compounds (_e.g._ HCl,
H_{2}S, H_{3}P, H_{4}Si) always form their highest oxides with four
atoms of oxygen, and as the highest forms of oxides (OsO_{4},
RuO_{4}) also contain four of oxygen, and eight groups of the
periodic system, corresponding to the highest basic oxides R_{2}O,
RO, R_{2}O_{3}, RO_{2}, R_{2}O_{5}, RO_{3}, R_{2}O_{7}, and RO_{4},
imply the above relationship, and because of the nearest analogues
among the elements--such as Mg, Zn, Cd, and Hg; or Cr, Mo, W, and
U; or Si, Ge, Sn, and Pt; or F, Cl, Br, and I, and so forth--not
more than four are known, it seems to me that in these
relationships there lies a deep interest and meaning with regard to
chemical mechanics. But because, to my imagination, the idea of
unity of design in Nature, either acting in complex celestial
systems or among chemical molecules, is very attractive, especially
because the atomic teaching at once acquires its true meaning, I
will recall the following facts relating to the solar system. There
are eight major planets, of which the four inner ones are not only
separated from the four outer by asteroids, but differ from them in
many respects, as, for example, in the smallness of their diameters
and their greater density. Saturn with his ring has eight
satellites, Jupiter and Uranus have each four. It is evident that
in the solar systems also we meet with these higher numbers four
and eight which appear in the combination of chemical molecules.
As for the last forty years, from the times of Berzelius, Dumas, Liebig,
Gerhardt, Williamson, Frankland, Kolbe, Kekulé, and Butleroff, most
theoretical generalisations have centred round organic or carbon
compounds, we will, for the sake of brevity, leave out the discussion of
ammonia derivatives, notwithstanding their simplicity with respect to the
doctrine of substitutions; we will dwell more especially on its
application to carbon compounds, starting from methane, CH_{4}, as the
simplest of the hydrocarbons, containing in its molecule one atom of
carbon. According to the principles enumerated we may derive from CH_{4}
every combination of the form CH_{3}X, CH_{2}X_{2}, CHX_{3}, and CX_{4},
in which X is an element, or radicle, equivalent to hydrogen--that is to
say, competent to take its place or to combine with it. Such are the
chlorine substitutes already mentioned, such is wood-spirit, CH_{3}(OH),
in which X is represented by the residue of water, and such are numerous
other carbon derivatives. If we continue, with the aid of hydroxyl,
further substitutions of the hydrogen of methane we shall obtain
successively CH_{2}(OH)_{2}, CH(OH)_{3}, and C(OH)_{4}. But if, in
proceeding thus, we bear in mind that CH_{2}(OH)_{2} contains two
hydroxyls in the same form as hydrogen peroxide, H_{2}O_{2} or (OH)_{2},
contains them--and moreover not only in one molecule, but together,
attached to one and the same atom of carbon--so here we must look for the
same decomposition as that which we find in hydrogen peroxide, and
accompanied also by the formation of water as an independently existing
molecule; therefore CH_{2}(OH)_{2} should yield, as it actually does,
immediately water and the oxide of methylene, CH_{2}O, which is methane
with oxygen substituted for two atoms of hydrogen. Exactly in the same
manner out of CH(OH)_{3} are formed water and formic acid, CHO(OH), and
out of C(OH)_{4} is produced water and carbonic acid, or directly
carbonic anhydride, CO_{2}, which will therefore be nothing else than
methane with the double replacement of pairs of hydrogen by oxygen. As
nothing leads to the supposition that the four atoms of hydrogen in
methane differ one from the other, so it does not matter by what means we
obtain any one of the combinations indicated--they will be identical;
that is to say, there will be no case of actual isomerism, although there
may easily be such cases of isomerism as have been distinguished by the
term metamerism.
Formic acid, for example, has two atoms of hydrogen, one attached to the
carbon left from the methane, and the other attached to the oxygen which
has entered in the form of hydroxyl, and if one of them be replaced by
some substance X it is evident that we shall obtain substances of the
same composition, but of different construction, or of different orders
of movement among the molecules, and therefore endowed with other
properties and reactions. If X be methyl, CH_{4}--that is to say, a group
capable of replacing hydrogen because it is actually contained with
hydrogen in methane itself--then by substituting this group for the
original hydrogen we obtain acetic acid, CCH_{3}O(OH), out of formic, and
by substitution of the hydrogen in its oxide or hydroxyl we obtain methyl
formate, CHO(OCH_{3}). These substances differ so much from each other
physically and chemically that at first sight it is hardly possible to
admit that they contain the same atoms in identically the same
proportions. Acetic acid, for example, boils at a higher temperature than
water, and has a higher specific gravity than it, whilst its metameride,
methyl formate, is lighter than water, and boils at 30°--that is to say,
it evaporates very easily.
Let us now turn to carbon compounds containing two atoms of carbon to the
molecule, as in acetic acid, and proceed to evolve them from methane by
the principle of substitution. This principle declares at once that
methane can only be split up in the four following ways:--
1. Into a group CH_{3} equivalent with H. Let us call changes of this
nature methylation.
2. Into a group CH_{2} and H_{2}. We will call this order of
substitutions methylenation.
3. Into CH and H_{3}, which commutations we will call acetylenation.
4. Into C and H_{4}, which may be called carbonation.
It is evident that hydrocarbon compounds containing two atoms of carbon
can only proceed from methane, CH_{4}, which contains four atoms of
hydrogen by the first three methods of substitution; carbonation would
yield free carbon if it could take place directly, and if the molecule of
free carbon--which is in reality very complex, that is to say strongly
polyatomic, as I have long since been proving by various means--could
contain only C_{2} like the molecules O_{2}, H_{2}, N_{2}, and so on.
By methylation we should evidently obtain from marsh gas, ethane,
CH_{3}CH_{3} = C_{2}H_{6}.
By methylenation--that is, by substituting group CH_{2} for
H_{2}--methane forms ethylene, CH_{2}CH_{2} = C_{2}H_{4}.
By acetylenation--that is, by substituting three atoms of hydrogen,
H_{3}, in methane--by the remnant CH, we get acetylene, CHCH =
C_{2}H_{2}.
If we have applied the principles of Newton correctly, there should not
be any other hydrocarbons containing two atoms of carbon in the molecule.
All these combinations have long been known, and in each of them we can
not only produce those substitutions of which an example has been given
in the case of methane, but also all the phases of other substitutions,
as we shall find from a few more instances, by the aid of which I trust
that I shall be able to show the great complexity of those derivatives
which, on the principle of substitution, can be obtained from each
hydrocarbon. Let us content ourselves with the case of ethane,
CH_{3}CH_{3}, and the substitution of the hydrogen by hydroxyl. The
following are the possible changes:--
1. CH_{3}CH_{2}(OH): this is nothing more than spirit of wine, or ethyl
alcohol, C_{2}H_{5}(OH) or C_{2}H_{6}O.
2. CH_{2}(OH)CH_{2}(OH): this is the glycol of Würtz, which has shed so
much light on the history of alcohol. Its isomeride may be
CH_{3}CH(OH)_{2}, but as we have seen in the case of CH(OH)_{2}, it
decomposes, giving off water, and forming aldehyde, CH_{3}CHO, a
substance capable of yielding alcohol by uniting with hydrogen, and of
yielding acetic acid by uniting with oxygen.
If glycol, CH_{2}(OH)CH_{2}(OH), loses its water, it may be seen at once
that it will not now yield aldehyde, CH_{3}CHO, but its isomeride,
CH_{2}CH_{2}/O, the oxide of ethylene. I have here indicated in a special
manner the oxygen which has taken the place of two atoms of the hydrogen
of ethane taken from different atoms of the carbon.
3. CH_{3}C(OH)_{3} decomposed as CH(OH)_{3}, forming water and acetic
acid, CH_{3}CO(OH). It is evident that this acid is nothing else than
formic acid, CHO(OH), with its hydrogen replaced by methyl. Without
examining further the vast number of possible derivatives, I will direct
your attention to the circumstance that in dissolving acetic acid in
water we obtain the maximum contraction and the greatest viscosity when
to the molecule CH_{3}CO(OH) is added a molecule of water, which is the
proportion which would form the hydrate CH_{3}C(OH)_{3}. It is probable
that the doubling of the molecule of acetic acid at temperatures
approaching its boiling-point has some connection with this power of
uniting with one molecule of water.
4. CH_{2}(OH)C(OH)_{3} is evidently an alcoholic acid, and indeed this
compound, after losing water, answers to glycolic acid, CH_{2}(OH)CO(OH).
Without investigating all the possible isomerides, we will note only that
the hydrate CH(OH)_{2}CH(OH)_{2} has the same composition as
CH_{2}(OH)C(OH)_{3}, and although corresponding to glycol, and being a
symmetrical substance, it becomes, on parting with its water, the
aldehyde of oxalic acid, or the glyoxal of Debus, CHOCHO.
5. CH(OH)_{2}C(OH_{3}), from the tendency of all the preceding,
corresponds with glyoxylic acid, an aldehyde acid, CHOCO(OH), because the
group CO(OH), or carboxyl, enters into the compositions of organic acids,
and the group CHO defines the aldehyde function.
6. C(OH)_{3}C(OH)_{3} through the loss of 2H_{2}O yields the bibasic
oxalic acid CO(OH)CO(OH), which generally crystallises with 2H_{2}O,
following thus the normal type of hydration characteristic of ethane.[2]
[2] One more isomeride, CH_{2}CH(OH), is possible--that is, secondary
vinyl alcohol, which is related to ethylene, CH_{2}CH_{2}, but
derived by the principle of substitution from CH_{4}. Other
isomerides, of the composition C_{2}H_{4}O, such, for example, as
CCH_{3}(OH), are impossible, because it would correspond with the
hydrocarbon CHCH_{3} = C_{2}H_{4}, which is isomeric with ethylene,
and it cannot be derived from methane. If such an isomeride existed
it would be derived from CH_{2}, but such products are, up to the
present, unknown. In such cases the insufficiency of the points of
departure of the statical structural teaching is shown. It first
admits constant atomicity and then rejects it, the facts serving to
establish either one or the other view; and therefore it seems to
me that we must come to the conclusion that the structural method
of reasoning, having done a service to science, has outlived the
age, and must be regenerated, as in their time was the teaching of
the electro-chemists, the radicalists, and the adherents of the
doctrine of types. As we cannot now lean on the views above stated,
it is time to abandon the structural theory. They will all be
united in chemical mechanics, and the principle of substitution
must be looked on only as a preparation for the coming epoch in
chemistry, where such cases as the isomerism of fumaric and maleic
acids, when explained dynamically, as proposed by Le Bel and Van't
Hoff, may yield points of departure.
Thus, by applying the principle of substitution, we can, in the simplest
manner, derive not only every kind of hydrocarbon compound, such as the
alcohols, the aldehyde-alcohols, aldehydes, alcohol-acids, and the acids,
but also combinations analogous to hydrated crystals which usually are
disregarded.
But even those unsaturated substances, of which ethylene, CH_{2}CH_{2},
and acetylene, CHCH, are types, may be evolved with equal simplicity.
With respect to the phenomena of isomerism, there are many possibilities
among the hydrocarbon compounds containing two atoms of carbon, and
without going into details it will be sufficient to indicate that the
following formulæ, though not identical, will be isomeric substantially
among themselves:--CH_{3}CHX_{2} and CH_{2}XCH_{2}X, although both
contain C_{2}H_{4}X_{2}; or CH_{2}CX_{2} and CHXCHX, although both
contain C_{2}H_{2}X_{2}, if by X we indicate chlorine or generally an
element capable of replacing one atom of hydrogen, or capable of uniting
with it. To isomerism of this kind belongs the case of aldehyde and the
oxide of ethylene, to which we have already referred, because both have
the composition C_{2}H_{4}O.
What I have said appears to me sufficient to show that the principle of
substitution adequately explains the composition, the isomerism, and all
the diversity of combination of the hydrocarbons, and I shall limit the
further development of these views to preparing a complete list of every
possible hydrocarbon compound containing three atoms of carbon in the
molecule. There are eight in all, of which only five are known at
present.[3]
[3] Conceding variable atomicity, the structuralists must expect an
incomparably larger number of isomerides, and they cannot now
decline to acknowledge the change of atomicity, were it only for
the examples HgCl and HgCl_{2}, CO and CO_{2}, PCl_{3} and PCl_{5}.
Among those possible for C_{3}H_{6} there should be two isomerides,
propylene and trimethylene, and they are both already known. For
C_{3}H_{4} there should be three isomerides: allylene and allene are
known, but the third has not yet been discovered; and for C_{3}H_{2}
there should be two isomerides, though neither of them is known as yet.
Their composition and structure are easily deduced from ethane, ethylene,
and acetylene, by methylation, by methylenation, by acetylenation and by
carbonation.
1. C_{3}H_{8} = CH_{3}CH_{2}CH_{3} out of CH_{3}CH_{3} by methylation.
This hydrocarbon is named propane.
2. C_{3}H_{6} = CH_{3}CHCH_{2} out of CH_{3}CH_{3} by methylenation. This
substance is propylene.
3. C_{3}H_{6} = CH_{2}CH_{2}CH_{2} out of CH_{3}CH_{3} by methylenation.
This substance is trimethylene.
4. C_{3}H_{4} = CH_{3}CCH out of CH_{3}CH_{3} by acetylenation or from
CHCH by methylation. This hydrocarbon is named allylene.
5. C_{3}H_{4} = CHCH/CH_{2} out of CH_{3}CH_{3} by acetylenation, or from
CH_{2}CH_{2} by methylenation, because CH_{2}CH/CH = CHCH/CH_{2}. This
body is as yet unknown.
6. C_{3}H_{4} = CH_{2}CCH_{2} out of CH_{2}CH_{2} by methylenation. This
hydrocarbon is named allene, or iso-allylene.
7. C_{3}H_{2} = CHCH/C out of CH_{3}CH_{3} by symmetrical carbonation, or
out of CH_{2}CH_{2} by acetylenation. This compound is unknown.
8. C_{3}H_{2} = CC/CH_{2} out of CH_{3}CH_{3} by carbonation, or out of
CHCH by methylenation. This compound is unknown.
If we bear in mind that for each hydrocarbon serving as a type in the
above tables there are a number of corresponding derivatives, and that
every compound obtained may, by further methylation, methylenation,
acetylenation, and carbonation, produce new hydrocarbons, and these may
be followed by a numerous suite of derivatives and an immense number of
isomeric substances, it is possible to understand the limitless number of
carbon compounds, although they all have the one substance, methane, for
their origin. The number of substances is so enormous that it is no
longer a question of enlarging the possibilities of discovery, but rather
of finding some means of testing them analogous to the well-known two
which for a long time have served as gauges for all carbon compounds.
I refer to the law of even numbers and to that of limits, the first
enunciated by Gerhardt some forty years ago, with respect to
hydrocarbons, namely, that their molecules always contain an even number
of atoms of hydrogen. But by the method which I have used of deriving all
the hydrocarbons from methane, CH_{4}, this law may be deduced as a
direct consequence of the principle of substitutions. Accordingly, in
methylation, CH_{3} takes the place of H, and therefore CH_{2} is added.
In methylenation the number of atoms of hydrogen remains unchanged, and
at each acetylenation it is reduced by two, and in carbonation by four,
atoms--that is to say, an even number of atoms of hydrogen is always
added or removed. And because the fundamental hydrocarbon, methane,
CH_{4}, contains an even number of atoms of hydrogen, all its derivative
hydrocarbons will also contain even numbers of hydrogen, and this
constitutes the law of even numbers.
The principle of substitutions explains with equal simplicity the
conception of the limiting compositions of hydrocarbons C_{_n_}H_{2_n_ +
2}, which I derived, in 1861,[4] in an empirical manner from accumulated
materials available at that time, and on the basis of the limits to
combinations worked out by Dr. Frankland for other elements.
[4] 'Essai d'une théorie sur les limites des combinaisons organiques,'
par D. Mendeléeff, 2/11 août 1861, _Bulletin de l'Académie i. d.
Sc. de St. Pétersbourg_, t. v
Of all the various substitutions the highest proportion of hydrogen is
yielded by methylation, because in that operation alone does the quantity
of hydrogen increase; hence, taking methane as a point of departure, if
we imagine methylation effected (_n_ - 1) times we obtain hydrocarbon
compounds containing the highest quantities of hydrogen. It is evident
that they will contain CH_{4} + (_n_ - 1)CH_{2}, or C_{_n_}H_{2_n_} +
{2}, because methylation leads to the addition of CH_{2} to the compound.
It will thus be seen that by the principle of substitution--that is to
say, by the third law of Newton--we are able to deduce, in the simplest
manner, not only the individual composition, the isomerism, and relations
of substances, but also the general laws which govern their most complex
combinations without having recourse either to statical constructions, to
the definition of atomicities, to the exclusion of free affinities, or to
the recognition of those single, double or treble bonds which are so
indispensable to structuralists in the explanation of the composition and
construction of hydrocarbon compounds. And yet, by the application of the
dynamical principles of Newton, we can attain to that chief and
fundamental object, the comprehension of isomerism in hydrocarbon
compounds, and the forecasting of the existence of combinations as yet
unknown, by which the edifice raised by structural teaching is
strengthened and supported. Besides--and I count this for a circumstance
of special importance--the process which I advocate will make no
difference in those special cases which have been already so well worked
out, such as, for example, the isomerism of the hydrocarbons and
alcohols, even to the extent of not interfering with the nomenclature
which has been adopted, and the structural system will retain all the
glory of having worked up, in a thoroughly scientific manner, the store
of information which Gerhardt had accumulated about the middle of the
fifties, and the still higher glory of establishing the rational
synthesis of organic substances. Nothing will be lost to the structural
doctrine except its statical origin; and as soon as it will embrace the
dynamic principles of Newton, and suffer itself to be guided by them, I
believe that we shall attain for chemistry that unity of principle which
is now wanting. Many an adept will be attracted to that brilliant and
fascinating enterprise, the penetration into the unseen world of the
kinetic relations of atoms, to the study of which the last twenty-five
years have contributed so much labour and such high inventive faculties.
D'Alembert found in mechanics that if inertia be taken to represent
force, dynamic equations may be applied to statical questions, which are
thereby rendered more simple and more easily understood.
The structural doctrine in chemistry has unconsciously followed the same
course, and therefore its terms are easily adopted; they may retain their
present forms provided that a truly dynamical--that is to say,
Newtonian--meaning be ascribed to them.
Before finishing my task and demonstrating the possibility of adapting
structural doctrines to the dynamics of Newton, I consider it
indispensable to touch on one question which naturally arises, and which
I have heard discussed more than once. If bromine, the atom of which is
eighty times heavier than that of hydrogen, takes the place of hydrogen,
it would seem that the whole system of dynamic equilibrium must be
destroyed.
Without entering into the minute analysis of this question, I think it
will be sufficient to examine it by the light of two well-known
phenomena, one of which will be found in the department of chemistry and
the other in that of celestial mechanics, and both will serve to
demonstrate the existence of that unity in the plan of creation which is
a consequence of the Newtonian doctrines. Experiments demonstrate that
when a heavy element is substituted for a light one in a chemical
compound--for example, for magnesium, in the oxide of that metal, an atom
of mercury, which is 8-1/3 times heavier--the chief chemical
characteristics or properties are generally, though not always,
preserved.
The substitution of silver for hydrogen, than which it is 108 times
heavier, does not affect all the properties of the substance, though it
does some. Therefore chemical substitutions of this kind--the
substitution of light for heavy atoms--need not necessarily entail
changes in the original equilibrium; and this point is still further
elucidated by the consideration that the periodic law indicates the
degree of influence of an increment of weight in the atom as affecting
the possible equilibria, and also what degree of increase in the weight
of the atoms reproduces some, though not all, of the properties of the
substance.
This tendency to repetition--these periods--may be likened to those
annual or diurnal periods with which we are so familiar on the earth.
Days and years follow each other, but, as they do so, many things change;
and in like manner chemical evolutions, changes in the masses of the
elements, permit of much remaining undisturbed, though many properties
undergo alteration. The system is maintained according to the laws of
conservation in nature, but the motions are altered in consequence of the
change of parts.
Next, let us take an astronomical case--such, for example, as the earth
and the moon--and let us imagine that the mass of the latter is
constantly increasing. The question is, what will then occur? The path of
the moon in space is a wave-line similar to that which geometricians have
named epicycloidal, or the locus of a point in a circle rolling round
another circle. But in consequence of the influence of the moon it is
evident that the path of the earth itself cannot be a geometric ellipse,
even supposing the sun to be immovably fixed; it must be an epicycloidal
curve, though not very far removed from the true ellipse--that is to say,
it will be impressed with but faint undulations. It is only the common
centre of gravity of the earth and the moon which describes a true
ellipse round the sun. If the moon were to increase, the relative
undulations of the earth's path would increase in amplitude, those of the
moon would also change, and when the mass of the moon had increased to an
equality with that of the earth, the path would consist of epicycloidal
curves crossing each other, and having opposite phases. But a similar
relation exists between the sun and the earth, because the former is also
moving in space. We may apply these views to the world of atoms, and
suppose that in their movements, when heavy ones take the place of those
that are lighter, similar changes take place, provided that the system or
the molecule is preserved throughout the change.
It seems probable that in the heavenly systems, during incalculable
astronomical periods, changes have taken place and are still going on
similar to those which pass rapidly before our eyes during the chemical
reaction of molecules, and the progress of molecular mechanics may--we
hope will--in course of time permit us to explain those changes in the
stellar world which have more than once been noticed by astronomers, and
which are now so carefully studied. A coming Newton will discover the
laws of these changes. Those laws, when applied to chemistry, may exhibit
peculiarities, but these will certainly be mere variations on the grand
harmonious theme which reigns in nature. The discovery of the laws which
produce this harmony in chemical evolution will only be possible, it
seems to me, under the banner of Newtonian dynamics, which has so long
waved over the domains of mechanics, astronomy, and physics. In calling
chemists to take their stand under its peaceful and catholic shadow I
imagine that I am aiding in establishing that scientific union which the
managers of the Royal Institution wish to effect, who have shown their
desire to do so by the flattering invitation which has given me--a
Russian--the opportunity of laying before the countrymen of Newton an
attempt to apply to chemistry one of his immortal principles.
APPENDIX II
THE PERIODIC LAW OF THE CHEMICAL ELEMENTS
BY PROFESSOR MENDELÉEFF
FARADAY LECTURE DELIVERED BEFORE THE FELLOWS OF THE
CHEMICAL SOCIETY IN THE THEATRE OF THE ROYAL INSTITUTION,
ON TUESDAY, JUNE 4, 1889
The high honour bestowed by the Chemical Society in inviting me to pay a
tribute to the world-famed name of Faraday by delivering this lecture has
induced me to take for its subject the Periodic Law of the Elements--this
being a generalisation in chemistry which has of late attracted much
attention.
While science is pursuing a steady onward movement, it is convenient
from time to time to cast a glance back on the route already traversed,
and especially to consider the new conceptions which aim at discovering
the general meaning of the stock of facts accumulated from day to day in
our laboratories. Owing to the possession of laboratories, modern science
now bears a new character, quite unknown, not only to antiquity, but even
to the preceding century. Bacon's and Descartes' idea of submitting the
mechanism of science simultaneously to experiment and reasoning has been
fully realised in the case of chemistry, it having become not only
possible but always customary to experiment. Under the all-penetrating
control of experiment, a new theory, even if crude, is quickly
strengthened, provided it be founded on a sufficient basis; the
asperities are removed, it is amended by degrees, and soon loses the
phantom light of a shadowy form or of one founded on mere prejudice; it
is able to lead to logical conclusions, and to submit to experimental
proof. Willingly or not, in science we all must submit not to what seems
to us attractive from one point of view or from another, but to what
represents an agreement between theory and experiment; in other words, to
demonstrated generalisation and to the approved experiment. Is it long
since many refused to accept the generalisations involved in the law of
Avogadro and Ampère, so widely extended by Gerhardt? We still may hear
the voices of its opponents; they enjoy perfect freedom, but vainly will
their voices rise so long as they do not use the language of demonstrated
facts The striking observations with the spectroscope which have
permitted us to analyse the chemical constitution of distant worlds,
seemed, at first, applicable to the task of determining the nature of the
atoms themselves; but the working out of the idea in the laboratory soon
demonstrated that the characters of spectra are determined, not directly
by the atoms, but by the molecules into which the atoms are packed; and
so it became evident that more verified facts must be collected before it
will be possible to formulate new generalisations capable of taking their
place beside those ordinary ones based upon the conception of simple
substances and atoms. But as the shade of the leaves and roots of living
plants, together with the relics of a decayed vegetation, favour the
growth of the seedling and serve to promote its luxurious development, in
like manner sound generalisations--together with the relics of those
which have proved to be untenable--promote scientific productivity, and
ensure the luxurious growth of science under the influence of rays
emanating from the centres of scientific energy. Such centres are
scientific associations and societies. Before one of the oldest and most
powerful of these I am about to take the liberty of passing in review the
twenty years' life of a generalisation which is known under the name of
the Periodic Law. It was in March 1869 that I ventured to lay before the
then youthful Russian Chemical Society the ideas upon the same subject
which I had expressed in my just written 'Principles of Chemistry.'
Without entering into details, I will give the conclusions I then arrived
at in the very words I used:--
'1. The elements, if arranged according to their atomic weights, exhibit
an evident _periodicity_ of properties.
'2. Elements which are similar as regards their chemical properties have
atomic weights which are either of nearly the same value (_e.g._
platinum, iridium, osmium) or which increase regularly (_e.g._ potassium,
rubidium, cæsium).
'3. The arrangement of the elements, or of groups of elements, in the
order of their atomic weights, corresponds to their so-called _valencies_
as well as, to some extent, to their distinctive chemical properties--as
is apparent, among other series, in that of lithium, beryllium, barium,
carbon, nitrogen, oxygen, and iron.
'4. The elements which are the most widely diffused have _small_ atomic
weights.
'5. The _magnitude_ of the atomic weight determines the character of the
element, just as the magnitude of the molecule determines the character
of a compound.
'6. We must expect the discovery of many yet _unknown_ elements--for
example, elements analogous to aluminium and silicon, whose atomic weight
would be between 65 and 75.
'7. The atomic weight of an element may sometimes be amended by a
knowledge of those of the contiguous elements. Thus, the atomic weight of
tellurium must lie between 123 and 126, and cannot be 128.
'8. Certain characteristic properties of the elements can be foretold
from their atomic weights.
'The aim of this communication will be fully attained if I succeed in
drawing the attention of investigators to those relations which exist
between the atomic weights of dissimilar elements, which, so far as I
know, have hitherto been almost completely neglected. I believe that the
solution of some of the most important problems of our science lies in
researches of this kind.'
To-day, twenty years after the above conclusions were formulated, they
may still be considered as expressing the essence of the now well-known
periodic law.
Reverting to the epoch terminating with the sixties, it is proper to
indicate three series of data without the knowledge of which the periodic
law could not have been discovered, and which rendered its appearance
natural and intelligible.
In the first place, it was at that time that the numerical value of
atomic weights became definitely known. Ten years earlier such knowledge
did not exist, as may be gathered from the fact that in 1860 chemists
from all parts of the world met at Karlsruhe in order to come to some
agreement, if not with respect to views relating to atoms, at any rate as
regards their definite representation. Many of those present probably
remember how vain were the hopes of coming to an understanding, and how
much ground was gained at that Congress by the followers of the unitary
theory so brilliantly represented by Cannizzaro. I vividly remember the
impression produced by his speeches, which admitted of no compromise, and
seemed to advocate truth itself, based on the conceptions of Avogadro,
Gerhardt, and Regnault, which at that time were far from being generally
recognised. And though no understanding could be arrived at, yet the
objects of the meeting were attained, for the ideas of Cannizzaro proved,
after a few years, to be the only ones which could stand criticism, and
which represented an atom as--'the smallest portion of an element which
enters into a molecule of its compound.' Only such real atomic
weights--not conventional ones--could afford a basis for generalisation.
It is sufficient, by way of example, to indicate the following cases in
which the relation is seen at once and is perfectly clear:--
K = 39 Rb = 85 Cs = 133
Ca = 40 Sr = 87 Ba = 137
whereas with the equivalents then in use--
K = 39 Rb = 85 Cs = 133
Ca = 20 Sr = 43·5 Ba = 68·5
the consecutiveness of change in atomic weight, which with the true
values is so evident, completely disappears.
Secondly, it had become evident during the period 1860-70, and even
during the preceding decade, that the relations between the atomic
weights of analogous elements were governed by some general and simple
laws. Cooke, Cremers, Gladstone, Gmelin, Lenssen, Pettenkofer, and
especially Dumas, had already established many facts bearing on that
view. Thus Dumas compared the following groups of analogous elements with
organic radicles:--
Diff. Diff. Diff. Diff.
Mg = 12} P = 31} O = 8}
}8 }44 }8
Li = 7 } Ca = 20} As= 75} S = 16}
}16 }3 × 8 }44 }3 × 8
Na = 23} Sr = 44} Sb = 119} Se = 40}
}16 }3 × 8 }2 × 44 }3 × 8
K = 39 } Ba = 68} Bi = 207} Te = 64}
and pointed out some really striking relationships, such as the
following:--
F = 19.
Cl = 35·5 = 19 + 16·5.
Br = 80 = 19 + 2 × 16·5 + 28.
I = 127 = 2 x 19 + 2 × 16·5 + 2 × 28.
A. Strecker, in his work 'Theorien und Experimente zur Bestimmung der
Atomgewichte der Elemente' (Braunschweig, 1859), after summarising the
data relating to the subject, and pointing out the remarkable series of
equivalents--
Cr = 26·2 Mn = 27·6 Fe = 28 Ni = 29 Co = 30 Cu = 31·7 Zn = 32·5
remarks that: 'It is hardly probable that all the above-mentioned
relations between the atomic weights (or equivalents) of chemically
analogous elements are merely accidental. We must, however, leave to the
future the discovery of the _law_ of the relations which appears in these
figures.'[1]
[1] 'Es ist wohl kaum anzunehmen, dass alle im Vorhergehenden
hervorgehobenen Beziehungen zwischen den Atomgewichten (oder
Aequivalenten) in chemischen Verhältnissen einander ähnliche
Elemente bloss zufällig sind. Die Auffindung der in diesen Zahlen
_gesetzlichen_ Beziehungen müssen wir jedoch der Zukunft
überlassen.'
In such attempts at arrangement and in such views are to be recognised
the real forerunners of the periodic law; the ground was prepared for it
between 1860 and 1870, and that it was not expressed in a determinate
form before the end of the decade may, I suppose, be ascribed to the fact
that only analogous elements had been compared. The idea of seeking for a
relation between the atomic weights of all the elements was foreign to
the ideas then current, so that neither the _vis tellurique_ of De
Chancourtois, nor the _law of octaves_ of Newlands, could secure
anybody's attention. And yet both De Chancourtois and Newlands like Dumas
and Strecker, more than Lenssen and Pettenkofer, had made an approach to
the periodic law and had discovered its germs. The solution of the
problem advanced but slowly, because the facts, but not the law, stood
foremost in all attempts; and the law could not awaken a general interest
so long as elements, having no apparent connection with each other, were
included in the same octave, as for example:--
1st octave of | | | | | | | |
Newlands | H | F | Cl | Co & Ni | Br | Pd | I | Pt & Ir
7th Ditto | O | S | Fe | Se | Rh & Ru | Te | Au | Os or Th
Analogies of the above order seemed quite accidental, and the more so as
the octave contained occasionally ten elements instead of eight, and when
two such elements as Ba and V, Co and Ni, or Rh and Ru, occupied one
place in the octave.[2] Nevertheless, the fruit was ripening, and I now
see clearly that Strecker, De Chancourtois, and Newlands stood foremost
in the way towards the discovery of the periodic law, and that they
merely wanted the boldness necessary to place the whole question at such
a height that its reflection on the facts could be clearly seen.
[2] To judge from J. A. R. Newlands's work, _On the Discovery of the
Periodic Law_, London, 1884, p. 149; 'On the Law of Octaves' (from
the _Chemical News_, 12, 83, August 18, 1865).
A third circumstance which revealed the periodicity of chemical elements
was the accumulation, by the end of the sixties, of new information
respecting the rare elements, disclosing their many-sided relations to
the other elements and to each other. The researches of Marignac on
niobium, and those of Roscoe on vanadium, were of special moment. The
striking analogies between vanadium and phosphorus on the one hand, and
between vanadium and chromium on the other, which became so apparent in
the investigations connected with that element, naturally induced the
comparison of V = 51 with Cr = 52, Nb = 94 with Mo = 96, and Ta = 192
with W = 194; while, on the other hand, P = 31 could be compared with S =
32, As = 75 with Se = 79, and Sb = 120 with Te = 125. From such
approximations there remained but one step to the discovery of the law of
periodicity.
The law of periodicity was thus a direct outcome of the stock of
generalisations and established facts which had accumulated by the end of
the decade 1860-1870; it is an embodiment of those data in a more or less
systematic expression. Where, then, lies the secret of the special
importance which has since been attached to the periodic law, and has
raised it to the position of a generalisation which has already given to
chemistry unexpected aid, and which promises to be far more fruitful in
the future and to impress upon several branches of chemical research a
peculiar and original stamp? The remaining part of my communication will
be an attempt to answer this question.
In the first place we have the circumstance that, as soon as the law made
its appearance, it demanded a revision of many facts which were
considered by chemists as fully established by existing experience. I
shall return, later on, briefly to this subject, but I wish now to remind
you that the periodic law, by insisting on the necessity for a revision
of supposed facts, exposed itself at once to destruction in its very
origin. Its first requirements, however, have been almost entirely
satisfied during the last 20 years; the supposed facts have yielded to
the law, thus proving that the law itself was a legitimate induction from
the verified facts. But our inductions from data have often to do with
such details of a science so rich in facts, that only generalisations
which cover a wide range of important phenomena can attract general
attention. What were the regions touched on by the periodic law? This is
what we shall now consider.
The most important point to notice is, that periodic functions, used for
the purpose of expressing changes which are dependent on variations of
time and space, have been long known. They are familiar to the mind when
we have to deal with motion in closed cycles, or with any kind of
deviation from a stable position, such as occurs in
pendulum-oscillations. A like periodic function became evident in the
case of the elements, depending on the mass of the atom. The primary
conception of the masses of bodies, or of the masses of atoms, belongs to
a category which the present state of science forbids us to discuss,
because as yet we have no means of dissecting or analysing the
conception. All that was known of functions dependent on masses derived
its origin from Galileo and Newton, and indicated that such functions
either decrease or increase with the increase of mass, like the
attraction of celestial bodies. The numerical expression of the phenomena
was always found to be proportional to the mass, and in no case was an
increase of mass followed by a recurrence of properties such as is
disclosed by the periodic law of the elements. This constituted such a
novelty in the study of the phenomena of nature that, although it did not
lift the veil which conceals the true conception of mass, it nevertheless
indicated that the explanation of that conception must be searched for in
the masses of the atoms; the more so, as all masses are nothing but
aggregations, or additions, of chemical atoms which would be best
described as chemical individuals. Let me remark, by the way, that though
the Latin word 'individual' is merely a translation of the Greek word
'atom,' nevertheless history and custom have drawn a sharp distinction
between the two words, and the present chemical conception of atoms is
nearer to that defined by the Latin word than by the Greek, although this
latter also has acquired a special meaning which was unknown to the
classics. The periodic law has shown that our chemical individuals
display a harmonic periodicity of properties dependent on their masses.
Now natural science has long been accustomed to deal with periodicities
observed in nature, to seize them with the vice of mathematical analysis,
to submit them to the rasp of experiment. And these instruments of
scientific thought would surely, long since, have mastered the problem
connected with the chemical elements, were it not for a new feature which
was brought to light by the periodic law, and which gave a peculiar and
original character to the periodic function.
If we mark on an axis of abscissæ a series of lengths proportional to
angles, and trace ordinates which are proportional to sines or other
trigonometrical functions, we get periodic curves of a harmonic
character. So it might seem, at first sight, that with the increase of
atomic weights the function of the properties of the elements should also
vary in the same harmonious way. But in this case there is no such
continuous change as in the curves just referred to, because the periods
do not contain the infinite number of points constituting a curve, but a
_finite_ number only of such points. An example will better illustrate
this view. The atomic weights--
Ag = 108 Cd = 112 In = 113 Sn = 118 Sb = 120
Te = 125 I = 127
steadily increase, and their increase is accompanied by a modification of
many properties which constitutes the essence of the periodic law. Thus,
for example, the densities of the above elements decrease steadily, being
respectively--
10·5 8·6 7·4 7·2 6·7 6·4 4·9
while their oxides contain an increasing quantity of oxygen--
Ag_{2}O Cd_{2}O_{2} In_{2}O_{3} Sn_{2}O_{4} Sb_{2}O_{5}
Te_{2}O_{6} I_{2}O_{7}
But to connect by a curve the summits of the ordinates expressing any of
these properties would involve the rejection of Dalton's law of multiple
proportions. Not only are there no intermediate elements between silver,
which gives AgCl, and cadmium, which gives CdCl_{2}, but, according to
the very essence of the periodic law, there can be none; in fact a
uniform curve would be inapplicable in such a case, as it would lead us
to expect elements possessed of special properties at any point of the
curve. The periods of the elements have thus a character very different
from those which are so simply represented by geometers. They correspond
to points, to numbers, to sudden changes of the masses, and not to a
continuous evolution. In these sudden changes destitute of intermediate
steps or positions, in the absence of elements intermediate between, say,
silver and cadmium, or aluminium and silicon, we must recognise a problem
to which no direct application of the analysis of the infinitely small
can be made. Therefore, neither the trigonometrical functions proposed by
Ridberg and Flavitzky, nor the pendulum-oscillations suggested by
Crookes, nor the cubical curves of the Rev. Mr. Haughton, which have been
proposed for expressing the periodic law, from the nature of the case,
can represent the periods of the chemical elements. If geometrical
analysis is to be applied to this subject, it will require to be modified
in a special manner. It must find the means of representing in a special
way, not only such long periods as that comprising
K Ca Sc Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br,
but short periods like the following:--
Na Mg Al Si P S Cl.
In the theory of numbers only do we find problems analogous to ours, and
two attempts at expressing the atomic weights of the elements by
algebraic formulæ seem to be deserving of attention, although neither of
them can be considered as a complete theory, nor as promising finally to
solve the problem of the periodic law. The attempt of E. J. Mills (1886)
does not even aspire to attain this end. He considers that all atomic
weights can be expressed by a logarithmic function,
15(_n_ - 0·9375^_t_),
in which the variables _n_ and _t_ are _whole numbers_. Thus, for oxygen,
_n_ = 2, and _t_ = 1, whence its atomic weight is = 15·94; in the case of
chlorine, bromine, and iodine, _n_ has respective values of 3, 6, and 9,
whilst _t_ = 7, 6, and 9; in the case of potassium, rubidium, and cæsium,
_n_ = 4, 6, and 9, and _t_ = 14, 18, and 20.
Another attempt was made in 1888 by B. N. Tchitchérin. Its author places
the problem of the periodic law in the first rank, but as yet he has
investigated the alkali metals only. Tchitchérin first noticed the simple
relations existing between the atomic volumes of all alkali metals; they
can be expressed, according to his views, by the formula
A(2 - 0·00535A_n_),
where A is the atomic weight, and _n_ is equal to 8 for lithium and
sodium, to 4 for potassium, to 3 for rubidium, and to 2 for cæsium. If
_n_ remained equal to 8 during the increase of A, the volume would become
zero at A = 46-2/3, and it would reach its maximum at A = 23-1/3. The
close approximation of the number 46-2/3 to the differences between the
atomic weights of analogous elements (such as Cs - Rb, I - Br, and so
on); the close correspondence of the number 23-1/3 to the atomic weight
of sodium; the fact of _n_ being necessarily a whole number, and several
other aspects of the question, induce Tchitchérin to believe that they
afford a clue to the understanding of the nature of the elements; we
must, however, await the full development of his theory before
pronouncing judgment on it. What we can at present only be certain of is
this: that attempts like the two above named must be repeated and
multiplied, because the periodic law has clearly shown that the masses of
the atoms increase abruptly, by steps, which are clearly connected in
some way with Dalton's law of multiple proportions; and because the
periodicity of the elements finds expression in the transition from RX to
RX_{2}, RX_{3}, RX_{4}, and so on till RX_{8}, at which point, the energy
of the combining forces being exhausted, the series begins anew from RX
to RX_{2}, and so on.
While connecting by new bonds the theory of the chemical elements with
Dalton's theory of multiple proportions, or atomic structure of bodies,
the periodic law opened for natural philosophy a new and wide field for
speculation. Kant said that there are in the world 'two things which
never cease to call for the admiration and reverence of man: the moral
law within ourselves, and the stellar sky above us.' But when we turn our
thoughts towards the nature of the elements and the periodic law, we must
add a third subject, namely, 'the nature of the elementary individuals
which we discover everywhere around us.' Without them the stellar sky
itself is inconceivable; and in the atoms we see at once their peculiar
individualities, the infinite multiplicity of the individuals, and the
submission of their seeming freedom to the general harmony of Nature.
Having thus indicated a new mystery of Nature, which does not yet yield
to rational conception, the periodic law, together with the revelations
of spectrum analysis, have contributed to again revive an old but
remarkably long-lived hope--that of discovering, if not by experiment, at
least by a mental effort, the _primary matter_--which had its genesis in
the minds of the Grecian philosophers, and has been transmitted, together
with many other ideas of the classic period, to the heirs of their
civilisation. Having grown, during the times of the alchemists up to the
period when experimental proof was required, the idea has rendered good
service; it induced those careful observations and experiments which
later on called into being the works of Scheele, Lavoisier, Priestley,
and Cavendish. It then slumbered awhile, but was soon awakened by the
attempts either to confirm or to refute the ideas of Prout as to the
multiple proportion relationship of the atomic weights of all the
elements. And once again the inductive or experimental method of studying
Nature gained a direct advantage from the old Pythagorean idea: because
atomic weights were determined with an accuracy formerly unknown. But
again the idea could not stand the ordeal of experimental test, yet the
prejudice remains and has not been uprooted, even by Stas; nay, it has
gained a new vigour, for we see that all which is imperfectly worked out,
new and unexplained, from the still scarcely studied rare metals to the
hardly perceptible nebulæ, have been used to justify it. As soon as
spectrum analysis appears as a new and powerful weapon of chemistry, the
idea of a primary matter is immediately attached to it. From all sides we
see attempts to constitute the imaginary substance _helium_[3] the so
much longed for primary matter. No attention is paid to the circumstance
that the helium line is only seen in the spectrum of the solar
protuberances, so that its universality in Nature remains as problematic
as the primary matter itself; nor to the fact that the helium line is
wanting amongst the Fraunhofer lines of the solar spectrum, and thus does
not answer to the brilliant fundamental conception which gives its real
force to spectrum analysis.
[3] That is, a substance having a wave-length equal to 0·0005875
millimetre.
And finally, no notice is even taken of the indubitable fact that the
brilliancies of the spectral lines of the simple substances vary under
different temperatures and pressures; so that all probabilities are in
favour of the helium line simply belonging to some long since known
element placed under such conditions of temperature, pressure, and
gravity as have not yet been realised in our experiments. Again, the idea
that the excellent investigations of Lockyer of the spectrum of iron can
be interpreted in favour of the compound nature of that element,
evidently must have arisen from some misunderstanding. The spectrum of a
compound certainly does not appear as a sum of the spectra of its
components; and therefore the observations of Lockyer can be considered
precisely as a proof that iron undergoes no other changes at the
temperature of the sun than those which it experiences in the voltaic
arc--provided the spectrum of iron is preserved. As to the shifting of
some of the lines of the spectrum of iron while the other lines maintain
their positions, it can be explained, as shown by M. Kleiber ('Journal of
the Russian Chemical and Physical Society,' 1885, 147), by the relative
motion of the various strata of the sun's atmosphere, and by Zöllner's
laws of the relative brilliancies of different lines of the spectrum.
Moreover, it ought not to be forgotten that if iron were really proved to
consist of two or more unknown elements, we should simply have an
increase in the number of our elements--not a reduction, and still less a
reduction of all of them to one single primary matter.
Feeling that spectrum analysis will not yield a support to the
Pythagorean conception, its modern promoters are so bent upon its being
confirmed by the periodic law, that the illustrious Berthelot, in his
work 'Les origines de l'Alchimie,' 1885, 313, has simply mixed up the
fundamental idea of the law of periodicity with the ideas of Prout, the
alchemists, and Democritus about primary matter.[4] But the periodic law,
based as it is on the solid and wholesome ground of experimental
research, has been evolved independently of any conception as to the
nature of the elements; it does not in the least originate in the idea of
a unique matter; and it has no historical connection with that relic of
the torments of classical thought, and therefore it affords no more
indication of the unity of matter or of the compound character of our
elements, than the law of Avogadro, or the law of specific heats, or even
the conclusions of spectrum analysis. None of the advocates of a unique
matter have ever tried to explain the law from the standpoint of ideas
taken from a remote antiquity when it was found convenient to admit the
existence of many gods--and of a unique matter.
[4] He maintains (on p. 309) that the periodic law requires two new
analogous elements, having atomic weights of 48 and 64, occupying
positions between sulphur and selenium, although nothing of the
kind results from any of the different readings of the law.
When we try to explain the origin of the idea of a unique primary matter,
we easily trace that in the absence of inductions from experiment it
derives its origin from the scientifically philosophical attempt at
discovering some kind of unity in the immense diversity of
individualities which we see around. In classical times such a tendency
could only be satisfied by conceptions about the immaterial world. As to
the material world, our ancestors were compelled to resort to some
hypothesis, and they adopted the idea of unity in the formative material,
because they were not able to evolve the conception of any other possible
unity in order to connect the multifarious relations of matter.
Responding to the same legitimate scientific tendency, natural science
has discovered throughout the universe a unity of plan, a unity of
forces, and a unity of matter, and the convincing conclusions of modern
science compel every one to admit these kinds of unity. But while we
admit unity in many things, we none the less must also explain the
individuality and the apparent diversity which we cannot fail to trace
everywhere. It has been said of old, 'Give us a fulcrum, and it will
become easy to displace the earth.' So also we must say, 'Give us
something that is individualised, and the apparent diversity will be
easily understood.' Otherwise, how could unity result in a multitude?
After a long and painstaking research, natural science has discovered
the individualities of the chemical elements, and therefore it is now
capable not only of analysing, but also of synthesising; it can
understand and grasp generality and unity, as well as the individualised
and the multifarious. The general and universal, like time and space,
like force and motion, vary uniformly; the uniform admit of
interpolations, revealing every intermediate phase. But the
multitudinous, the individualised--such as ourselves, or the chemical
elements, or the members of a peculiar periodic function of the elements,
or Dalton's multiple proportions--is characterised in another way: we see
in it, side by side with a connecting general principle, leaps, breaks of
continuity, points which escape from the analysis of the infinitely
small--an absence of complete intermediate links. Chemistry has found an
answer to the question as to the causes of multitudes; and while
retaining the conception of many elements, all submitted to the
discipline of a general law, it offers an escape from the Indian
Nirvana--the absorption in the universal, replacing it by the
individualised. However, the place for individuality is so limited by the
all-grasping, all-powerful universal, that it is merely a point of
support for the understanding of multitude in unity.
Having touched upon the metaphysical bases of the conception of a unique
matter which is supposed to enter into the composition of all bodies. I
think it necessary to dwell upon another theory, akin to the above
conception--the theory of the compound character of the elements now
admitted by some--and especially upon one particular circumstance which,
being related to the periodic law, is considered to be an argument in
favour of that hypothesis.
Dr. Pelopidas, in 1883, made a communication to the Russian Chemical and
Physical Society on the periodicity of the hydrocarbon radicles, pointing
out the remarkable parallelism which was to be noticed in the change of
properties of hydrocarbon radicles and elements when classed in groups.
Professor Carnelley, in 1886, developed a similar parallelism. The idea
of M. Pelopidas will be easily understood if we consider the series of
hydrocarbon radicles which contain, say, 6 atoms of carbon:--
I. II. III. IV. V.
C_{6}H_{13} C_{6}H_{12} C_{6}H_{11} C_{6}H_{10} C_{6}H_{9}
VI. VII. VIII.
C_{6}H_{8} C_{6}H_{7} C_{6}H_{6}
The first of these radicles, like the elements of the 1st group, combines
with Cl, OH, and so on, and gives the derivatives of hexyl alcohol,
C_{6}H_{13}(OH); but, in proportion as the number of hydrogen atoms
decreases, the capacity of the radicles of combining with, say, the
halogens increases. C_{6}H_{12} already combines with 2 atoms of
chlorine; C_{6}H_{11}, with 3 atoms, and so on. The last members of the
series comprise the radicles of acids: thus C_{6}H_{8}, which belongs to
the 6th group, gives, like sulphur, a bibasic acid,
C_{6}H_{8}O_{2}(OH)_{2}, which is homologous with oxalic acid. The
parallelism can be traced still further, because C_{6}H_{5} appears as a
monovalent radicle of benzene, and with it begins a new series of
aromatic derivatives, so analogous to the derivatives of the aliphatic
series. Let me also mention another example from among those which have
been given by M. Pelopidas. Starting from the alkaline radicle of
monomethylammonium, N(CH_{3})H_{3}, or NCH_{6}, which presents many
analogies with the alkaline metals of the 1st group, he arrives, by
successively diminishing the number of the atoms of hydrogen, at a 7th
group which contains cyanogen, CN, which has long since been compared to
the halogens of the 7th group.
The most important consequence which, in my opinion, can be drawn from
the above comparison is that the periodic law, so apparent in the
elements, has a wider application than might appear at first sight; it
opens up a new vista of chemical evolutions. But, while admitting the
fullest parallelism between the periodicity of the elements and that of
the compound radicles, we must not forget that in the periods of the
hydrocarbon radicles we have a _decrease_ of mass as we pass from the
representatives of the first group to the next, while in the periods of
the elements the mass _increases_ during the progression. It thus becomes
evident that we cannot speak of an identity of periodicity in both cases,
unless we put aside the ideas of mass and attraction, which are the real
corner-stones of the whole of natural science, and even enter into those
very conceptions of simple substances which came to light a full hundred
years later than the immortal principles of Newton.[5]
[5] It is noteworthy that the year in which Lavoisier was born
(1743)--the author of the idea of elements and of the
indestructibility of matter--is later by exactly one century than
the year in which the author of the theory of gravitation and mass
was born (1643 N.S.). The affiliation of the ideas of Lavoisier and
those of Newton is beyond doubt.
From the foregoing, as well as from the failures of so many attempts at
finding in experiment and speculation a proof of the compound character
of the elements and of the existence of primordial matter, it is evident,
in my opinion, that this theory must be classed among mere utopias. But
utopias can only be combated by freedom of opinion, by experiment, and by
new utopias. In the republic of scientific theories freedom of opinions
is guaranteed. It is precisely that freedom which permits me to criticise
openly the widely-diffused idea as to the unity of matter in the
elements. Experiments and attempts at confirming that idea have been so
numerous that it really would be instructive to have them all collected
together, if only to serve as a warning against the repetition of old
failures. And now as to new utopias which may be helpful in the struggle
against the old ones, I do not think it quite useless to mention a
_phantasy_ of one of my students who imagined that the weight of bodies
does not depend upon their mass, but upon the character of the motion of
their atoms. The atoms, according to this new utopian, may all be
homogeneous or heterogeneous, we know not which; we know them in motion
only, and that motion they maintain with the same persistence as the
stellar bodies maintain theirs. The weights of atoms differ only in
consequence of their various modes and quantity of motion; the heaviest
atoms may be much simpler than the lighter ones: thus an atom of mercury
may be simpler than an atom of hydrogen--the manner in which it moves
causes it to be heavier. My interlocutor even suggested that the view
which attributes the greater complexity to the lighter elements finds
confirmation in the fact that the hydrocarbon radicles mentioned by
Pelopidas, while becoming lighter as they lose hydrogen, change their
properties periodically in the same manner as the elements change theirs,
according as the atoms grow heavier.
The French proverb, _La critique est facile, mais l'art est difficile_,
however, may well be reversed in the case of all such ideal views, as it
is much easier to formulate than to criticise them. Arising from the
virgin soil of newly-established facts, the knowledge relating to the
elements, to their masses, and to the periodic changes of their
properties has given a motive for the formation of utopian hypotheses,
probably because they could not be foreseen by the aid of any of the
various metaphysical systems, and exist, like the idea of gravitation, as
an independent outcome of natural science, requiring the acknowledgment
of general laws, when these have been established with the same degree of
persistency as is indispensable for the acceptance of a thoroughly
established fact. Two centuries have elapsed since the theory of
gravitation was enunciated, and although we do not understand its cause,
we still must regard gravitation as a fundamental conception of natural
philosophy, a conception which has enabled us to perceive much more than
the metaphysicians did or could with their seeming omniscience. A hundred
years later the conception of the elements arose; it made chemistry what
it now is; and yet we have advanced as little in our comprehension of
simple substances since the times of Lavoisier and Dalton as we have in
our understanding of gravitation. The periodic law of the elements is
only twenty years old; it is not surprising, therefore, that, knowing
nothing about the causes of gravitation and mass, or about the nature of
the elements, we do not comprehend the _rationale_ of the periodic law.
It is only by collecting established laws--that is, by working at the
acquirement of truth--that we can hope gradually to lift the veil which
conceals from us the causes of the mysteries of Nature and to discover
their mutual dependency. Like the telescope and the microscope, laws
founded on the basis of experiment are the instruments and means of
enlarging our mental horizon.
In the remaining part of my communication I shall endeavour to show, and
as briefly as possible, in how far the periodic law contributes to
enlarge our range of vision. Before the promulgation of this law the
chemical elements were mere fragmentary, incidental facts in Nature;
there was no special reason to expect the discovery of new elements, and
the new ones which were discovered from time to time appeared to be
possessed of quite novel properties. The law of periodicity first enabled
us to perceive undiscovered elements at a distance which formerly was
inaccessible to chemical vision; and long ere they were discovered new
elements appeared before our eyes possessed of a number of well-defined
properties. We now know three cases of elements whose existence and
properties were foreseen by the instrumentality of the periodic law. I
need but mention the brilliant discovery of _gallium_, which proved to
correspond to eka-aluminium of the periodic law, by Lecoq de Boisbaudran;
of _scandium_, corresponding to ekaboron, by Nilson; and of _germanium_,
which proved to correspond in all respects to ekasilicon, by Winkler.
When, in 1871, I described to the Russian Chemical Society the
properties, clearly defined by the periodic law, which such elements
ought to possess, I never hoped that I should live to mention their
discovery to the Chemical Society of Great Britain as a confirmation of
the exactitude and the generality of the periodic law. Now that I have
had the happiness of doing so, I unhesitatingly say that, although
greatly enlarging our vision, even now the periodic law needs further
improvements in order that it may become a trustworthy instrument in
further discoveries.[6]
[6] I foresee some more new elements, but not with the same certitude
as before. I shall give one example, and yet I do not see it quite
distinctly. In the series which contains Hg = 204, Pb = 206, and Bi
= 208, we can imagine the existence (at the place VI-11) of an
element analogous to tellurium, which we can describe as
dvi-tellurium, Dt, having an atomic weight of 212, and the property
of forming the oxide DtO_{3}. If this element really exists, it
ought in the free state to be an easily fusible, crystalline,
non-volatile metal of a grey colour, having a density of about 9·3,
capable of giving a dioxide, DtO_{2}, equally endowed with feeble
acid and basic properties. This dioxide must give on active
oxidation an unstable higher oxide, DtO_{3}, which should resemble
in its properties PbO_{2} and Bi_{2}O_{5}. Dvi-tellurium hydride,
if it be found to exist, will be a less stable compound than even
H_{2}Te. The compounds of dvi-tellurium will be easily reduced, and
it will form characteristic definite alloys with other metals.
I will venture to allude to some other matters which chemistry has
discerned by means of its new instrument, and which it could not have
made out without a knowledge of the law of periodicity, and I will
confine myself to simple substances and to oxides.
Before the periodic law was formulated the atomic weights of the elements
were purely empirical numbers, so that the magnitude of the equivalent,
and the atomicity, or the value in substitution possessed by an atom,
could only be tested by critically examining the methods of
determination, but never directly by considering the numerical values
themselves; in short, we were compelled to move in the dark, to submit to
the facts, instead of being masters of them. I need not recount the
methods which permitted the periodic law at last to master the facts
relating to atomic weights, and I would merely call to mind that it
compelled us to modify the valencies of _indium_ and _cerium_, and to
assign to their compounds a different molecular composition.
Determinations of the specific heats of these two metals fully confirmed
the change. The trivalency of _yttrium_, which makes us now represent its
oxide as Y_{2}O_{3} instead of as YO, was also foreseen (in 1870) by the
periodic law, and it has now become so probable that Clève, and all other
subsequent investigators of the rare metals, have not only adopted it,
but have also applied it without any new demonstration to substances so
imperfectly known as those of the cerite and gadolinite group, especially
since Hillebrand determined the specific heats of lanthanum and didymium
and confirmed the expectations suggested by the periodic law. But here,
especially in the case of didymium, we meet with a series of difficulties
long since foreseen through the periodic law, but only now becoming
evident, and chiefly arising from the relative rarity and insufficient
knowledge of the elements which usually accompany didymium.
Passing to the results obtained in the case of the rare elements
_beryllium_, _scandium_, and _thorium_, it is found that these have many
points of contact with the periodic law. Although Avdéeff long since
proposed the magnesia formula to represent beryllium oxide, yet there was
so much to be said in favour of the alumina formula, on account of the
specific heat of the metals and the isomorphism of the two oxides, that
it became generally adopted and seemed to be well established. The
periodic law, however, as Brauner repeatedly insisted ('Berichte,' 1878,
872; 1881, 53), was against the formula Be_{2}O_{3}; it required the
magnesia formula BeO--that is, an atomic weight of 9--because there was
no place in the system for an element like beryllium having an atomic
weight of 13·5. This divergence of opinion lasted for years, and I often
heard that the question as to the atomic weight of beryllium threatened
to disturb the generality of the periodic law, or, at any rate, to
require some important modifications of it. Many forces were operating in
the controversy regarding beryllium, evidently because a much more
important question was at issue than merely that involved in the
discussion of the atomic weight of a relatively rare element: and during
the controversy the periodic law became better understood, and the mutual
relations of the elements became more apparent than ever before. It is
most remarkable that the victory of the periodic law was won by the
researches of the very observers who previously had discovered a number
of facts in support of the trivalency of beryllium. Applying the higher
law of Avogadro, Nilson and Petterson have finally shown that the density
of the vapour of the beryllium chloride, BeCl_{2}, obliges us to regard
beryllium as bivalent in conformity with the periodic law.[7] I consider
the confirmation of Avdéeff's and Brauner's view as important in the
history of the periodic law as the discovery of scandium, which, in
Nilson's hands, confirmed the existence of ekaboron.
[7] Let me mention another proof of the bivalency of beryllium which
may have passed unnoticed, as it was only published in the Russian
chemical literature. Having remarked (in 1884) that the density of
such solutions of chlorides of metals, MCl_{_n_}, as contain 200
mols. of water (or a large and constant amount of water) regularly
increases as the molecular weight of the dissolved salt increases,
I proposed to one of our young chemists, M. Burdakoff, that he
should investigate beryllium chloride. If its molecule be BeCl_{2}
its weight must be = 80; and in such a case it must be heavier than
the molecule of KCl = 74·5, and lighter than that of MgCl_{2}, =
93. On the contrary, if beryllium chloride is a trichloride,
BeCl_{3} = 120, its molecule must be heavier than that of CaCl_{2}
= 111, and lighter than that of MnCl_{2} = 126. Experiment has
shown the correctness of the former formula, the solution BeCl_{2}
+ 200H_{2}O having (at 15°/4°) a density of 1·0138, this being a
higher density than that of the solution KCl + 200H_{2}O (=
1·0121), and lower than that of MgCl_{2} + 200H_{2}O (= 1·0203).
The bivalency of beryllium was thus confirmed in the case both of
the dissolved and the vaporised chloride.
The circumstance that _thorium_ proved to be quadrivalent, and Th = 232,
in accordance with the views of Chydenius and the requirements of the
periodic law, passed almost unnoticed, and was accepted without
opposition, and yet both thorium and uranium are of great importance in
the periodic system, as they are its last members, and have the highest
atomic weights of all the elements.
The alteration of the atomic weight of _uranium_ from U = 120 into U =
240 attracted more attention, the change having been made on account of
the periodic law, and for no other reason. Now that Roscoe, Rammelsberg,
Zimmermann, and several others have admitted the various claims of the
periodic law in the case of uranium, its high atomic weight is received
without objection, and it endows that element with a special interest.
While thus demonstrating the necessity for modifying the atomic weights
of several insufficiently known elements, the periodic law enabled us
also to detect errors in the determination of the atomic weights of
several elements whose valencies and true position among other elements
were already well known. Three such cases are especially noteworthy:
those of tellurium, titanium and platinum. Berzelius had determined the
atomic weight of _tellurium_ to be 128, while the periodic law claimed
for it an atomic weight below that of iodine, which had been fixed by
Stas at 126·5, and which was certainly not higher than 127. Brauner then
undertook the investigation, and he has shown that the true atomic weight
of tellurium is lower than that of iodine, being near to 125. For
_titanium_ the extensive researches of Thorpe have confirmed the atomic
weight of Ti = 48, indicated by the law, and already foreseen by Rose,
but contradicted by the analyses of Pierre and several other chemists. An
equally brilliant confirmation of the expectations based on the periodic
law has been given in the case of the series osmium, iridium, platinum,
and gold. At the time of the promulgation of the periodic law, the
determinations of Berzelius, Rose, and many others gave the following
figures:--
Os = 200; Ir = 197; Pt = 198; Au = 196.
The expectations of the periodic law[8] have been confirmed, first, by
new determinations of the atomic weight of _platinum_ (by Seubert,
Dittmar, and M'Arthur, which proved to be near to 196 (taking O = 16, as
proposed by Marignac, Brauner, and others); secondly, by Seubert having
proved that the atomic weight of _osmium_ is really lower than that of
platinum, being near to 191; and thirdly, by the investigations of Krüss,
Thorpe and Laurie, proving that the atomic weight of _gold_ exceeds that
of platinum, and approximates to 197. The atomic weights which were thus
found to require correction were precisely those which the periodic law
had indicated as affected with errors; and it has been proved, therefore,
that the periodic law affords a means of testing experimental results. If
we succeed in discovering the exact character of the periodic
relationships between the increments in atomic weights of allied elements
discussed by Ridberg in 1885, and again by Bazaroff in 1887, we may
expect that our instrument will give us the means of still more closely
controlling the experimental data relating to atomic weights.
[8] I pointed them out in the _Liebig's Annalen_, Supplement Band.,
viii. 1871, p. 211.
Let me next call to mind that, while disclosing the variation of chemical
properties,[9] the periodic law, has also enabled us to systematically
discuss many of the physical properties of elementary bodies, and to show
that these properties are also subject to the law of periodicity. At the
Moscow Congress of Russian Naturalists in August, 1869, I dwelt upon the
relations which existed between density and the atomic weight of the
elements. The following year Professor Lothar Meyer, in his well-known
paper,[10] studied the same subject in more detail, and thus contributed
to spread information about the periodic law. Later on, Carnelley,
Laurie, L. Meyer, Roberts-Austen, and several others applied the periodic
system to represent the order in the changes of the magnetic properties
of the elements, their melting points, the heats of formation of their
haloid compounds, and even of such mechanical properties as the
coefficient of elasticity, the breaking stress, &c., &c. These
deductions, which have received further support in the discovery of new
elements endowed not only with chemical but even with physical
properties, which were foreseen by the law of periodicity, are well
known; so I need not dwell upon the subject, and may pass to the
consideration of oxides.[11]
[9] Thus, in the typical small period of
Li, Be, B, C, N, O, F,
we see at once the progression from the alkali metals to the acid
non-metals, such as are the halogens.
[10] _Liebig's Annalen_, Supplement Band., vii. 1870.
[11] A distinct periodicity can also be discovered in the spectra of
the elements. Thus the researches of Hartley, Ciamician, and
others have disclosed, first, the homology of the spectra of
analogous elements: secondly, that the alkali metals have simpler
spectra than the metals of the following groups; and thirdly, that
there is a certain likeness between the complicated spectra of
manganese and iron on the one hand, and the no less complicated
spectra of chlorine and bromine on the other hand, and their
likeness corresponds to the degree of analogy between those
elements which is indicated by the periodic law.
In indicating that the gradual increase of the power of elements of
combining with oxygen is accompanied by a corresponding decrease in their
power of combining with hydrogen, the periodic law has shown that there
is a limit of oxidation, just as there is a well-known limit to the
capacity of elements for combining with hydrogen. A single atom of an
element combines with at most four atoms of either hydrogen or oxygen;
and while CH_{4} and SiH_{4} represent the highest hydrides, so RuO_{4}
and OsO_{4} are the highest oxides. We are thus led to recognise types of
oxides, just as we have had to recognise types of hydrides.[12]
[12] Formerly it was supposed that, being a bivalent element, oxygen
can enter into any grouping of the atoms, and there was no limit
foreseen as to the extent to which it could further enter into
combination. We could not explain why bivalent sulphur, which
forms compounds such as
O O
/ \ / \
S | and S O,
\ / \ /
O O
could not also form oxides such as--
O--O O--O
/ \ / \
S | or S O,
\ / \ /
O--O O--O
while other elements, as, for instance, chlorine, form compounds
such as--
Cl--O--O--O--O--K
The periodic law has demonstrated that the maximum extent to which
different non-metals enter into combination with oxygen is determined by
the extent to which they combine with hydrogen, and that the sum of the
number of equivalents of both must be equal to 8. Thus chlorine, which
combines with 1 atom or 1 equivalent of hydrogen, cannot fix more than 7
equivalents of oxygen, giving Cl_{2}O_{7}; while sulphur, which fixes 2
equivalents of hydrogen, cannot combine with more than 6 equivalents or 3
atoms of oxygen. It thus becomes evident that we cannot recognise as a
fundamental property of the elements the atomic valencies deduced from
their hydrides; and that we must modify, to a certain extent, the theory
of atomicity if we desire to raise it to the dignity of a general
principle capable of affording an insight into the constitution of all
compound molecules. In other words, it is only to carbon, which is
quadrivalent with regard both to oxygen and hydrogen, that we can apply
the theory of constant valency and of bond, by means of which so many
still endeavour to explain the structure of compound molecules. But I
should go too far if I ventured to explain in detail the conclusions
which can be drawn from the above considerations. Still, I think it
necessary to dwell upon one particular fact which must be explained from
the point of view of the periodic law in order to clear the way to its
extension in that particular direction.
The higher oxides yielding salts the formation of which was foreseen by
the periodic system--for instance, in the short series beginning with
sodium--
Na_{2}O, MgO, Al_{2}O_{3}, SiO_{2}, P_{2}O_{5}, SO_{3}, Cl_{2}O_{7},
must be clearly distinguished from the higher degrees of oxidation which
correspond to hydrogen peroxide and bear the true character of peroxides.
Peroxides such as Na_{2}O_{2}, BaO_{2}, and the like have long been
known. Similar peroxides have also recently become known in the case of
chromium, sulphur, titanium, and many other elements, and I have
sometimes heard it said that discoveries of this kind weaken the
conclusions of the periodic law in so far as it concerns the oxides. I do
not think so in the least, and I may remark, in the first place, that all
these peroxides are endowed with certain properties obviously common to
all of them, which distinguish them from the actual, higher, salt-forming
oxides, especially their easy decomposition by means of simple contact
agencies; their incapability of forming salts of the common type; and
their capability of combining with other peroxides (like the faculty
which hydrogen peroxide possesses of combining with barium peroxide,
discovered by Schoene). Again, we remark that some groups are especially
characterised by their capacity of generating peroxides. Such is, for
instance, the case in the sixth group, where we find the well-known
peroxides of sulphur, chromium, and uranium; so that further
investigation of peroxides will probably establish a new periodic
function, foreshadowing that molybdenum and tungsten will assume peroxide
forms with comparative readiness. To appreciate the constitution of such
peroxides, it is enough to notice that the peroxide form of sulphur
(so-called persulphuric acid) stands in the same relation to sulphuric
acid as hydrogen peroxide stands to water:--
H(OH), or H_{2}O, responds to (OH)(OH), or H_{2}O_{2},
and so also--
H(HSO_{4}), or H_{2}SO_{4}, responds to
(HSO_{4})(HSO_{4}), or H_{2}S_{2}O_{8}.
Similar relations are seen everywhere, and they correspond to the
principle of substitutions which I long since endeavoured to represent as
one of the chemical generalisations called into life by the periodic law.
So also sulphuric acid, if considered with reference to hydroxyl, and
represented as follows:--
HO(SO_{2}OH),
has its corresponding compound in dithionic acid--
(SO_{2}OH)(SO_{2}OH), or H_{2}S_{2}O_{6}.
Therefore, also, phosphoric acid, HO(POH_{2}O_{2}), has, in the same
sense, its corresponding compound in the subphosphoric acid of Saltzer:--
(POH_{2}O_{2})(POH_{2}O_{2}), or H_{4}P_{2}O_{6};
and we must suppose that the peroxide compound corresponding to
phosphoric acid, if it be discovered, will have the following
structure:--
(H_{2}PO_{4})_{2} or H_{4}P_{2}O_{8} = 2H_{2}O + 2PO_{3}.[13]
So far as is known at present, the highest form of peroxides is met with
in the peroxide of uranium, UO_{4}, prepared by Fairley;[14] while
OsO_{4} is the highest oxide giving salts. The line of argument which is
inspired by the periodic law, so far from being weakened by the discovery
of peroxides, is thus actually strengthened, and we must hope that a
further exploration of the region under consideration will confirm the
applicability to chemistry generally of the principles deduced from the
periodic law.
[13] In this sense, oxalic acid, (COOH)_{2}, also corresponds to
carbonic acid, OH(COOH), in the same way that dithionic acid
corresponds to sulphuric acid, and subphosphoric acid to
phosphoric; hence, if a peroxide corresponding to carbonic acid be
obtained, it will have the structure of (HCO_{3})_{2}, or
H_{2}C_{2}O_{6} = H_{2}O + C_{2}O_{5}. So also lead must have a
real peroxide, Pb_{2}O_{5}.
[14] The compounds of uranium prepared by Fairley seem to me especially
instructive in understanding the peroxides. By the action of
hydrogen peroxide on uranium oxide, UO_{3}, a peroxide of uranium,
UO_{4},4H_{2}O, is obtained (U = 240) if the solution be acid; but
if hydrogen peroxide act on uranium oxide in the presence of
caustic soda, a crystalline deposit is obtained which has the
composition Na_{4}UO_{8},4H_{2}O, and evidently is a combination
of sodium peroxide, Na_{2}O_{2}, with uranium peroxide, UO_{4}. It
is possible that the former peroxide, UO_{4},4H_{2}O, contains the
elements of hydrogen peroxide and uranium peroxide, U_{2}O_{7}, or
even U(OH)_{6},H_{2}O_{2}, like the peroxide of tin recently
discovered by Spring, which has the constitution
Sn_{2}O_{5},H_{2}O_{2}.
Permit me now to conclude my rapid sketch of the oxygen compounds by the
observation that the periodic law is especially brought into evidence in
the case of the oxides which constitute the immense majority of bodies at
our disposal on the surface of the earth.
The oxides are evidently subject to the law, both as regards their
chemical and their physical properties, especially if we take into
account the cases of polymerism which are so obvious when comparing
CO_{2}, with Si_{_n_}O_{2_n_}. In order to prove this I give the
densities s and the specific volumes v of the higher oxides of two short
periods. To render comparison easier, the oxides are all represented as
of the form R_{2}O_{_n_}. In the column headed [Delta] the differences
are given between the volume of the oxygen compound and that of the
parent element, divided by _n_--that is, by the number of atoms of oxygen
in the compound:--[15]
_s._ _v._ [Delta]
Na_{2}O 2·6 24 -22
Mg_{2}O_{2} 3·6 22 -3
Al_{2}O_{3} 4·0 26 +1·3
Si_{2}O_{4} 2·65 45 5·2
P_{2}O_{5} 2·39 59 6·2
S_{2}O_{6} 1·96 82 8·7
K_{2}O 2·7 35 -55
Sc_{2}O_{3} 3·86 35 0
Li_{2}O_{4} 4·2 38 +5
V_{2}O_{5} 3·49 52 6·7
Cr_{2}O_{6} 2·74 73 9·5
[15] [Delta] thus represents the average increase of volume for each
atom of oxygen contained in the higher salt-forming oxide. The
acid oxides give, as a rule, a higher value of [Delta], while in
the case of the strongly alkaline oxides its value is usually
negative.
I have nothing to add to these figures, except that like relations appear
in other periods as well. The above relations were precisely those which
made it possible for me to be certain that the relative density of
ekasilicon oxide would be about 4·7; germanium oxide, actually obtained
by Winkler, proved, in fact, to have the relative density 4·703.
The foregoing account is far from being an exhaustive one of all that
has already been discovered by means of the periodic law telescope in the
boundless realms of chemical evolution. Still less is it an exhaustive
account of all that may yet be seen, but I trust that the little which I
have said will account for the philosophical interest attached in
chemistry to this law. Although but a recent scientific generalisation,
it has already stood the test of laboratory verification, and appears as
an instrument of thought which has not yet been compelled to undergo
modification; but it needs not only new applications, but also
improvements, further development, and plenty of fresh energy. All this
will surely come, seeing that such an assembly of men of science as the
Chemical Society of Great Britain has expressed the desire to have the
history of the periodic law described in a lecture dedicated to the
glorious name of Faraday.
APPENDIX III
ARGON, A NEW CONSTITUENT OF THE ATMOSPHERE
WRITTEN BY PROFESSOR MENDELÉEFF IN FEBRUARY 1895
The remarks made in Chapter V., Note 16 bis respecting the newly
discovered constituent of the atmosphere are here supplemented by data
(taken from the publications of the Royal Society of London) given by the
discoverers Lord Rayleigh and Professor Ramsay in January 1895, together
with observations made by Crookes and Olszewsky upon the same subject.
This gas, which was discovered by Rayleigh and Ramsay in atmospheric
nitrogen, was named _argon_[1] by them, and upon the supposition of its
being an element, they gave it the symbol A. But its true chemical nature
is not yet fully known, for not only has no compound of it been yet
obtained, but it has not even been brought into any reaction. From all
that is known about it at the present time, we may conclude with the
discoverers that argon belongs to those gases which are permanent
constituents of the atmosphere, and that it is a new element. The latter
statement, however, requires confirmation. We shall presently see,
however, that the negative chemical character of argon (its incapacity to
react with any substance), and the small amount of it present in the
atmosphere (about 1-1/4 per cent. by volume in the nitrogen of air, and
consequently about 1 per cent. by volume in air), as well as the recent
date of its discovery (1894) and the difficulty of its preparation, are
quite sufficient reasons for the incompleteness of the existing knowledge
respecting this element. But since, so far as is yet known, we are
dealing with a normal constituent of the atmosphere[1 bis], the existing
data, notwithstanding their insufficiently definite nature, should find a
place even in such an elementary work as the present, all the more as the
names of Rayleigh, Ramsay, Crookes and Olszewsky, who have worked upon
argon, are among the highest in our science, and their researches among
the most difficult.[2] These researches, moreover, were directed straight
to the goal, which was only partly reached owing to the unusual
properties of argon itself.
[1] From the Greek [Greek: Argon]--inert.
[1 bis] In Note 16 bis, Chapter V., I mentioned that, judging from the
specific gravity of argon, it might possibly be polymerised
nitrogen, N_{3}, bearing the same relationship to nitrogen, N_{2},
that ozone, O_{3}, bears to ordinary oxygen. If this idea were
confirmed, still one would not imagine that argon was formed from
the atmospheric nitrogen by those reactions by which it was
obtained by Rayleigh and Ramsay, but rather that it arises from the
nitrogen of the atmosphere under natural conditions. Although this
proposition is not quite destroyed by the more recent results,
still it is contradicted by the fact that the ratio of the specific
heats of argon was found to be 1·66, which, as far as is now known,
could not be the case for a gas containing 3 atoms in its molecule,
since such gases (_see_ Chapter XIV., Note 7) give the ratio
approximately 1·3 (for example, CO_{2}). In abstaining from further
conclusions, for they must inevitably be purely conjectural, I
consider it advisable to suggest that in conducting further
researches upon argon it might be well to subject it to as high a
temperature as possible. And the possibility of nitrogen
polymerising is all the more admissible from the fact that the
aggregation of its atoms in the molecule is not at all unlikely,
and that polymerised nitrogen, judging from many examples, might be
inert if the polymerisation were accompanied by the evolution of
heat. In the following footnotes I frequently return to this
hypothesis, not only because I have not yet met any facts
definitely contradictory to it, but also because the chief
properties of argon agree with it to a certain extent.
[2] The chief difficulty in investigating argon lies in the fact that
its preparation requires the employment of a large quantity of air,
which has to be treated with a number of different reagents, whose
perfect purity (especially that of magnesium) will always be
doubtful, and argon has not yet been transferred to a substance in
which it could be easily purified. Perhaps the considerable
solubility of argon in water (or in other suitable liquids, which
have not apparently yet been tried) may give the means of doing so,
and it may be possible, by collecting the air expelled from boiling
water, to obtain a richer source of argon than ordinary air.
When it became known (Chapter V., Note 4 bis) that the nitrogen
obtained from air (by removing the oxygen, moisture and CO_{2}, by
various reagents) has a greater density than that obtained from the
various (oxygen, hydrogen and metallic) compounds of nitrogen, it was a
plausible explanation that the latter contained an admixture of hydrogen,
or of some other light gas lowering the density of the mixture. But such
an assumption is refuted not only by the fact that the nitrogen obtained
from its various compounds (after purification) has always the same
density (although the supposed impurities mixed with it should vary), but
also by Rayleigh and Ramsay's experiment of artificially adding hydrogen
to nitrogen, and then passing the mixture over red-hot oxide of copper,
when it was found that the nitrogen regained its original density, _i.e._
that the whole of the hydrogen was removed by this treatment. Therefore
the difference in the density of the two varieties of nitrogen had to be
explained by the presence of a heavier gas in admixture with the nitrogen
obtained from the atmosphere. This hypothesis was confirmed by the fact
that Rayleigh and Ramsay having obtained purified nitrogen (by removing
the O_{2}, CO_{2}, and H_{2}O), both from ordinary air and from air which
had been previously subjected to atmolysis, that is which had been passed
through porous tubes (of burnt clay, _e.g._ pipe-stem), surrounded by a
rarefied space, and so deprived of its lighter constituents (chiefly
nitrogen), found that the nitrogen from the air which had been subjected
to atmolysis was heavier than that obtained from air which had not been
so treated. This experiment showed that the nitrogen of air contains an
admixture of a gas which, being heavier than nitrogen itself,[3] diffuses
more slowly than nitrogen through the porous material. It remained,
therefore, to separate this impurity from the nitrogen. To do this
Rayleigh and Ramsay adopted two methods, converting the nitrogen into
solid and liquid substances, either by absorbing the nitrogen by heated
magnesium (Chapter V., Note 6, and Chapter XIV., Note 14), with the
formation of nitride of magnesium, or else by converting it into nitric
acid by the action of electric sparks or the presence of an excess of air
and alkali, as in Cavendish's method.[3 bis] In both cases the nitrogen
entered into reaction, while the heavier gas mixed with it remained
inert, and was thus able to be isolated. That is, the argon could be
separated by these means from the excess of atmospheric nitrogen
accompanying it.[4] As an illustration we will describe how argon was
obtained from the atmospheric nitrogen by means of magnesium.[5] To begin
with, it was discovered that when atmospheric nitrogen was passed through
a tube containing metallic magnesium heated to redness, its specific
gravity rose to 14·88. As this showed that part of the gas was absorbed
by the magnesium, a mercury gasometer filled with atmospheric nitrogen
was taken, and the gas drawn over soda-lime, P_{2}O_{5}, heated
magnesium[6] and then through tubes containing red-hot copper oxide,
soda-lime and phosphoric anhydride to a second mercury gasometer. Every
time the gas was repassed through the tubes, it decreased in volume and
increased in density. After repeating this for ten days 1,500 c.c. of gas
were reduced to 200 cc., and the density increased to 16·1 (if that of
H_{2} = 1 and N_{2} = 14). Further treatment of the remainder brought the
density up to 19·09. After adding a small quantity of oxygen and
repassing the gas through the apparatus, the density rose to 20·0. To
obtain argon by this process Ramsay and Rayleigh (employing a mercury air
pump and mercury gasometers) once treated about 150 litres of atmospheric
nitrogen. On another occasion they treated 7,925 c.c. of air by the
oxidation method and obtained 65 c.c. of argon, which corresponds to 0·82
per cent. The density of the argon obtained by this means was nearly
19·7, while that obtained by the magnesium method varied between 19·09
and 20·38.
[3] It might also be supposed that this heavy gas is separated by the
copper when the latter absorbs the oxygen of the air; but such a
supposition is not only improbable in itself, but does not agree
with the fact that nitrogen may be obtained from air by absorbing
the oxygen by various other substances in solution (for instance,
by the lower oxides of the metals, like FeO) besides red-hot
copper, and that the nitrogen obtained is always just as heavy.
Besides which, nitrogen is also set free from its oxides by copper,
and the nitrogen thus obtained is lighter. Therefore it is not the
copper which produces the heavy gas--_i.e._ argon.
[3 bis] It is worthy of note that Cavendish obtained a small residue of
gas in converting nitrogen into nitric acid; but he paid no
attention to it, although probably he had in his hands the very
argon recently discovered.
[4] When in these experiments, instead of atmospheric nitrogen the gas
obtained from its compound was taken, an inert residue of a heavy
gas, having the properties of argon, was also remarked, but its
amount was very small. Rayleigh and Ramsay ascribe the formation of
this residue to the fact that the gas in these experiments was
collected over water, and a portion of the dissolved argon in it
might have passed into the nitrogen. As the authors of this
supposition did not prove it by any special experiments, it forms a
weak point in their classical research. If it be admitted that
argon is N_{3}, the fact of its being obtained from the nitrogen of
compounds might be explained by the polymerisation of a portion of
the nitrogen in the act of reaction, although it is impossible to
refute Rayleigh and Ramsay's hypothesis of its being evolved from
the water employed in the manipulation of the gases. Three thousand
volumes of nitrogen extracted from its compounds gave about three
volumes of argon, while thirty volumes were yielded by the same
amount of atmospheric nitrogen.
[5] The preparation of argon by the conversion of nitrogen into nitric
acid is complicated by the necessity of adding a large proportion
of oxygen and alkali, of passing an electric discharge through the
mixture for a long period, and then removing the remaining oxygen.
All this was repeatedly done by the authors, but this method is far
more complex, both in practice and theory, than the preparation of
argon by means of magnesium. From 100 volumes of air subjected to
conversion into HNO_{3}, 0·76 volume of argon were obtained after
absorbing the excess of oxygen.
[6] In these and the following experiments the magnesium was placed in
an ordinary hard glass tube, and heated in a gas furnace to a
temperature almost sufficient to soften the glass. The current of
gas must be very slow (a tube containing a small quantity of
sulphuric acid served as a meter), as otherwise the heat evolved in
the formation of the Mg_{3}N_{2} (Chapter XIV., Note 14) will melt
the tube.
Thus the first positive and very important fact respecting argon is that
its specific gravity is nearly 20--that is, that it is 20 times heavier
than hydrogen, while nitrogen is only 14 times and oxygen 16 times
heavier than hydrogen. This explains the difference observed by Rayleigh
between the densities of nitrogen obtained from its compounds and from
the atmosphere (Chapter V., Note 4 bis). At 0° and 760 mm. a litre of the
former gas weighs 1·2505 grm., while a litre of the latter weighs 1·2572,
or taking H = 1, the density of the first = 13·916, and of the latter =
13·991. If the density of argon be taken as 20, it is contained in
atmospheric nitrogen to the extent of about 1·23 per cent. by volume,
whilst air contains about 0·97 per cent. by volume.
When argon had been isolated the question naturally arose, was it a new
homogeneous substance having definite properties or was it a mixture of
gases? The former may now be positively asserted, namely, that argon is a
peculiar gas previously unknown to chemistry. Such a conviction is in the
first place established by the fact that argon has a greater number of
negative properties, a smaller capacity for reaction, than any other
simple or compound body known. The most inert gas known is nitrogen, but
argon far exceeds it in this respect. Thus nitrogen is absorbed at a red
heat by many metals, with the formation of nitrides, while argon, as is
seen in the mode of its preparation and by direct experiment, does not
possess this property. Nitrogen, under the action of electric sparks,
combines with hydrogen in the presence of acids and with oxygen in the
presence of alkalis, while argon is unable to do so, as is seen from the
method of separation from nitrogen. Rayleigh and Ramsay also proved that
argon is unable to react with chlorine (dry or moist) either directly or
under the action of an electric discharge, or with phosphorus or sulphur,
at a red heat. Sodium, potassium, and tellurium may be distilled in an
atmosphere of argon without change. Fused caustic soda, incandescent
soda-lime, molten nitre, red-hot peroxide of sodium, and the
polysulphides of calcium and sodium also do not react with argon.
Platinum black does not absorb it, and spongy platinum is unable to
excite its reaction with oxygen or chlorine. Aqua regia, bromine water,
and a mixture of hydrochloric acid and KMnO_{4} were also without action
upon argon. Besides which it is evident from the method of its
preparation that it is not acted upon by red-hot oxide of copper. All
these facts exclude any possibility of argon containing any already known
body, and prove it to be the most inert of all the gases known. But
besides these negative points, the independency of argon is confirmed by
four observed positive properties possessed by it, which are:--
1. The spectrum of argon observed by Crookes under a low pressure (in
Geissler-Plücker tubes) distinguishes it from other gases.[7] It was
proved by this means that the argon obtained by means of magnesium is
identical with that which remains after the conversion of the atmospheric
nitrogen into nitric acid. Like nitrogen, argon presents two spectra
produced at different potentials of the induced current, one being
orange-red, the other steel-blue; the latter is obtained under a higher
degree of rarefaction and with a battery of Leyden jars. Both the spectra
of argon (in contradistinction to those of nitrogen) are distinguished by
clearly defined lines.[8] The red (ordinary) spectrum of argon has two
particularly brilliant and characteristic red lines (not far from the
bright red line of lithium, on the opposite side to the orange band)
having wave-lengths 705·64 and 696·56 (_see_ Vol. I., p. 565). Between
these bright lines there are in addition lines with wave lengths 603·8,
565·1, 561·0, 555·7, 518·58, 516·5, 450·95, 420·10, 415·95 and 394·85.
Altogether 80 lines have been observed in this spectrum and 119 in the
blue spectrum, of which 26 are common to both spectra.[9]
[7] The greatest brilliancy of the spectrum of argon is obtained at a
tension of 3 mm., while for nitrogen it is about 75 mm. (Crookes).
In Chapter V., Note 16 bis, it is said that the same blue line
observed in the spectrum of argon is also observed in the spectrum
of nitrogen. This is a mistake, since there is no coincidence
between the blue lines of the argon and nitrogen spectra. However,
we may add that for nitrogen the following moderately bright lines
are known of wave-lengths 585, 574, 544, 516, 457, 442, 436, and
426, which are repeated in the spectra (red and blue) of argon,
judging by Crookes' researches (1895); but it is naturally
impossible to assert that there is perfect identity until some
special comparative work has been done in this subject, which is
very desirable, and more especially for the bluish-violet portion
of the spectrum, more particularly between the lines 442-436, as
these lines are distinguished by their brilliancy in both the argon
and nitrogen spectra. The above-mentioned supposition of argon
being polymerised nitrogen (N_{3}), formed from nitrogen (N_{2}),
with the evolution of heat, might find some support should it be
found after careful comparison that even a limited number of
spectral lines coincided.
[8] At first the spectrum of argon exhibits the nitrogen lines, but
after a certain time these lines disappear (under the influence of
the platinum, and also of Al and Mg, but with the latter the
spectrum of hydrogen appears) and leave a pure argon spectrum. It
does not appear clear to me whether a polymerisation here takes
place or a simple absorption. Perhaps the elucidation of this
question would prove important in the history of argon. It would be
desirable to know, for instance, whether the volume of argon
changes when it is first subjected to the action of the electric
discharge.
[9] Crookes supposes that argon contains a mixture of two gases, but as
he gives no reasons for this, beyond certain peculiarities of a
spectroscopic character, we will not consider this hypothesis
further.
2. According to Rayleigh and Ramsay the solubility of argon in water is
approximately 4 volumes in 100 volumes of water at 13°. Thus argon is
nearly 2-1/2 times more soluble than nitrogen, and its solubility
approaches that of oxygen. Direct experiment proves that nitrogen
obtained from air from boiled water is heavier than that obtained
straight from the atmosphere. This again is an indirect proof of the
presence of argon in air.
3. The ratio _k_ of the two specific heats (at a constant pressure and
at a constant volume) of argon was determined by Rayleigh and Ramsay by
the method of the velocity of sound (_see_ Chapter XIV., Note 7 and
Chapter VII., Note 26) and was found to be nearly 1·66, that is greater
than for those gases whose molecules contain two atoms (for instance, CO,
H_{2}, N_{2}, air, &c., for which _k_ is nearly 1·4) or those whose
molecules contain three atoms (for instance, CO_{2}, N_{2}O, &c., for
which _k_ is about 1·3), but closely approximate to the ratio of the
specific heats of mercury vapour (Kundt and Warburg, _k_ = 1·67). And as
the molecule of mercury vapour contains one atom, so it may be said that
argon is a simple gaseous body whose molecule contains one atom.[10] A
compound body should give a smaller ratio. The experiments upon the
liquefaction of argon, which we shall presently describe, speak against
the supposition that argon is a mixture of two gases. The importance of
the results in question makes one wish that the determinations of the
ratio of the specific heats (and other physical properties) might be
confirmed with all possible accuracy.[11] If we admit, as we are obliged
to do for the present, that argon is a new element, its density shows
that its atomic weight must be nearly 40, that is, near to that of K = 39
and Ca = 40, which does not correspond to the existing data respecting
the periodicity of the properties of the elements in dependence upon
their atomic weights, for there is no reason on the basis of existing
data for admitting any intermediate elements between Cl = 35·5 and K =
39, and all the positions above potassium in the periodic system are
occupied. This renders it very desirable that the velocity of sound in
argon should be re-determined.[12]
[10] This portion of Rayleigh and Ramsay's researches deserves
particular attention as, so far, no gaseous substance is known
whose molecule contains but one atom. Were it not for the above
determinations, it might be thought that argon, having a density
20, has a complex molecule, and may be a compound or polymerised
body, for instance, N_{3} or NX_{_n_}, or in general X_{_n_}; but
as the matter stands, it can only be said that either (1) argon is
a new, peculiar, and quite unusual elementary substance, since
there is no reason for assuming it to contain two simple gases, or
(2) the magnitude, _k_ (the ratio of the specific heats) does not
only depend upon the number of atoms contained in the molecules,
but also upon the store of internal energy (internal motion of the
atoms in the molecule). Should the latter be admitted, it would
follow that the molecules of very active gaseous elements would
correspond to a smaller _k_ than those of other gases having an
equal number of atoms in their molecule. Such a gas is chlorine,
for which _k_ = 1·33 (Chapter XIV., Note 7). For gases having a
small chemical energy, on the contrary, a larger magnitude would
be expected for _k_. I think these questions might be partially
settled by determining _k_ for ozone (O_{3}) and sulphur (S_{6})
(at about 500°). In other words, I would suggest, though only
provisionally, that the magnitude, _k_ = 1·6, obtained for argon
might prove to agree with the hypothesis that argon is N_{3},
formed from N_{2} with the evolution of heat or loss of energy.
Here argon gives rise to questions of primary importance, and it
is to be hoped that further research will throw some light upon
them. In making these remarks, I only wish to clear the road for
further progress in the study of argon, and of the questions
depending on it. I may also remark that if argon is N_{3} formed
with the evolution of heat, its conversion into nitrogen, N_{2},
and into nitride compounds (for instance, boron nitride or nitride
of titanium) might only take place at a very high temperature.
[11] Without having the slightest reason for doubting the accuracy of
Rayleigh and Ramsay's determinations, I think it necessary to say
that as yet (February 1895) I am only acquainted with the short
memoir of the above chemists in the 'Proceedings of the Royal
Society,' which does not give any description of the methods
employed and results obtained, while at the end (in the general
conclusions) the authors themselves express some doubt as to the
simple nature of argon. Moreover, it seems to me that (Note 10)
there must be a dependence of _k_ upon the chemical energy.
Besides which, it is not clear what density of the gas Rayleigh
and Ramsay took in determining _k_. (If argon be N_{3}, its
density would be near to 21.) Hence I permit myself to express
some doubt as to whether the molecule of argon contains but one
atom.
[12] If it should be found that _k_ for argon is less than 1·4, or that
_k_ is dependent upon the chemical energy, it would be possible to
admit that the molecule of argon contains not one, but several
atoms--for instance, either N_{3} (then the density would be 21,
which is near to the observed density) or X_{6}, if X stand for an
element with an atomic weight near to 6·7. No elements are known
between H = 1 and Li = 7, but perhaps they may exist. The
hypothesis A = 40 does not admit argon into the periodic system.
If the molecule of argon be taken as A_{2}--_i.e._ the atomic
weight as A = 20--argon apparently finds a place in Group VIII.,
between F = 19 and Na = 23; but such a position could only be
justified by the consideration that elements of small atomic
weight belong to the category of typical elements which offer many
peculiarities in their properties, as is seen on comparing N with
the other elements of Group V., or O with those of Group VI. Apart
from this there appears to me to be little probability, in the
light of the periodic law, in the position of an inert substance
like argon in Group VIII., between such active elements as
fluorine and sodium, as the representatives of this group by their
atomic weights and also by their properties show distinct
transitions from the elements of the last groups of the uneven
series to the elements of the first groups of the even series--for
instance,
Group VI. VII. VIII. I. II.
Cr Mn Fe, Co, Ni Cu Zn
While if we place argon in a similar manner,
VI. VII. VIII. I. II.
O = 16 F = 19 A = 20 Na = 23 Mg = 24
although from a numerical point of view there is a similar
sequence to the above, still from a chemical and physical point of
view the result is quite different, as there is no such
resemblance between the properties of O, F and Na, Mg, as between
Cr, Mn, and Cu, Zn. I repeat that only the typical character of
the elements with small atomic weights can justify the atomic
weight A = 20, and the placing of argon in Group VIII. amongst the
typical elements; then N, O, F, A are a series of gases.
It appears to me simpler to assume that argon contains N_{3},
especially as argon is present in nitrogen and accompanies it,
and, as a matter of fact, none of the observed properties of argon
are contradictory to this hypothesis.
These observations were written by me in the beginning of February
1895, and on the 29th of that month I received a letter, dated
February 25, from Professor Ramsay informing me that 'the periodic
classification entirely corresponds to its (argon's) atomic
weight, and that it even gives a fresh proof of the periodic law,'
judging from the researches of my English friends. But in what
these researches consisted, and how the above agreement between
the atomic weight of argon and the periodic system was arrived at,
is not referred to in the letter, and we remain in expectation of
a first publication of the work of Lord Rayleigh and Professor
Ramsey. [For more complete information see papers read before the
Royal Society, January 31, 1895, February 13, March 10, and May
21, 1896, and a paper published in the Chemical Society's
Transactions, 1895, p. 684. For abstracts of these and other
papers on argon and helium, and correspondence, see 'Nature,' 1895
and 1896.
4. Argon was liquefied by Professor Olszewsky, who is well known for
his classical researches upon liquefied gases. These researches have an
especial interest since they show that argon exhibits a perfect constancy
in its properties in the liquid and critical states, which almost[13]
disposes of the supposition that it contains a mixture of two or more
unknown gases. As the first experiments showed, argon remains a gas under
a pressure of 100 atmospheres and at a temperature of -90°; this
indicated that its critical temperature was probably below this
temperature, as was indeed found to be the case when the temperature was
lowered to -128°·6[14] by means of liquid ethylene. At this temperature
argon easily liquefies to a colourless liquid under 38 atmospheres. The
meniscus begins to disappear at between -119°·8 and -121°·6, mean -121°
at a pressure of 50·6 atmospheres. The vapour tension of liquid argon at
-128°·6, is 38·0 atmospheres, at -187° it is one atmosphere, and at
-189°·6 it solidifies to a colourless substance like ice. The specific
gravity of liquid argon at about -187° is nearly 1·5, which is far above
that of other liquefied gases of very low absolute boiling point.
The discovery of argon is one of the most remarkable chemical
acquisitions of recent times, and we trust that Lord Rayleigh and
Professor Ramsay, who made this wonderful discovery, will further
elucidate the true nature of argon, as this should widen the fundamental
principles of chemistry, to which the chemists of Great Britain have from
early times made such valuable contributions. It would be premature now
to give any definite opinions upon so new a subject. Only one thing can
be said; argon is so inert that its rôle in nature cannot be
considerable, notwithstanding its presence in the atmosphere. But as the
atmosphere itself plays such a vast part in the life of the surface of
the earth, every addition to our knowledge of its composition must
directly or indirectly react upon the sum total of our knowledge of
nature.
[13] There only remains the very remote possibility that argon consists
of a mixture of two gases having very nearly the same properties.
[14] The following data, given by Olszewsky, supplement the data given
in Chapter II., Note 29, upon liquefied gases.
(_tc_) (_pc_) _t_ _t__{1} _s_
N_{2} -146° 35 -194°·4 -214° 0·885
CO -139°·5 35·5 -190° -207 ?
A -121° 50·6 -187° -189°·6 1·5
O_{2} -118°·8 50·8 -182°·7 ? 1·124
NO -93°·5 71·2 -153°·6 -167° ?
CH_{4} -81°·8 54·9 -164° -158°·8 0·415
where _tc_ is the absolute (critical) boiling point, _pc_ the
pressure (critical) in atmospheres corresponding to it, _t_ the
boiling point (under a pressure of 760 mm.), _t_{1}_ the melting
point, and _s_ the specific gravity in a liquid state at _t_.
The above shows that argon in its properties in a liquid state
stands near to oxygen (as it also does in its solubility), but
that all the temperatures relating to it (_tc_, _t_, and _t_{1}_)
are higher than for nitrogen. This fully answers, not only to the
higher density of argon, but also to the hypothesis that it
contains N_{3}. And as the boiling point of argon differs from
that of nitrogen and oxygen by less than 10°, and its amount is
small, it is easy to understand how Dewar (1894), who tried to
separate it from liquid air and nitrogen by fractional
distillation, was unable to do so. The first and last portions
were identical, and nitrogen from air showed no difference in its
liquefaction from that obtained from its compounds, or from that
which had been passed through a tube containing incandescent
magnesium. Still, it is not quite clear why both kinds of
nitrogen, after being passed over the magnesium in Dewar's
experiments, exhibited an almost similar alteration in their
properties, independent of the appearance of a small quantity of
hydrogen in them.
_Concluding Remarks_ (March 31, 1895).--The 'Comptes rendus' of
the Paris Academy of Sciences of March 18, 1895, contains a memoir
by Berthelot upon the reaction of argon with the vapour of benzene
under the action of a silent discharge. In his experiments,
Berthelot succeeded in treating 83 per cent. of the argon taken
for the purpose, and supplied to him by Ramsay (37 c.c. in all).
The composition of the product could not be determined owing to
the small amount obtained, but in its outward appearance it quite
resembled the product formed under similar conditions by nitrogen.
This observation of the famous French chemist to some extent
supports the supposition that argon is a polymerised variety of
nitrogen whose molecule contains N_{3}, while ordinary nitrogen
contains N_{2}. Should this supposition be eventually verified,
the interest in argon will not only not lessen, but become
greater. For this, however, we must wait for further observations
and detailed experimental data from Rayleigh and Ramsay.
The latest information obtained by me from London is that
Professor Ramsay, by treating cleveite (containing PbO, UO_{3},
Y_{2}O_{3}, &c.) with sulphuric acid, obtained argon, and, judging
by the spectrum, helium also. The accumulation of similar data
may, after detailed and diversified research, considerably
increase the stock of chemical knowledge which, constantly
widening, cannot be exhaustively treated in these 'Principles of
Chemistry,' although very probably furnishing fresh proof of the
'periodicity of the elements.'
* * * * *
INDEX OF AUTHORITIES
Abasheff, i. 75
Abel, ii. 56, 326, 410
Acheson, ii. 107
Adie, ii. 186
Alexéeff, i. 75, 94
Alluard, i. 458
Amagat, i. 132, 135, 140
Amat, ii. 171
Ammermüller, i. 504
Ampère, i. 309
Andréeff, i. 251
Andrews, i. 136, 203
Angeli, i. 266
Ansdell, i. 451
Arfvedson, i. 575
Arrhenius, i. 89, 92, 389
Aschoff, ii. 313
Askenasy, i. 508
Aubel, ii. 45
Aubin, i. 238
Avdéeff, i. 618; ii. 484
Avogadro, i. 309
Babo, v., i. 93, 200, 203
Bach, i. 394
Bachmetieff, ii. 31
Baeyer, v., i. 507
Bagouski, i. 384
Bailey, i. 449; ii. 29
Baker, i. 318, 403
Balard, i. 480, 494, 495, 505
Ball, ii. 414
Bannoff, i. 506
Barfoed, ii. 53
Baroni, i. 331
Barreswill, ii. 282
Baudrimont, ii. 35
Baumé, i. 193
Baumgauer, ii. 20
Baumhauer, i. 495
Bayer, ii. 76, 159
Bazaroff, i. 409; ii. 24, 68, 486
Becher, i. 17
Becker, i. 16
Beckmann, i. 91, 496; ii. 156
Becquerel, i. 228; ii. 97, 220
Beilby, i. 71
Beilstein, i. 373; ii. 188
Beketoff, i. 120, 122, 124, 146, 403, 459, 466, 534, 541, 574, 577;
ii. 87, 102, 289, 429
Bender, i. 476
Benedict, ii. 65
Berglund, ii. 229
Bergman, i. 27, 435; ii. 100
Berlin, i. 95
Bernouilli, i. 81
Bernthsen, ii. 228
Bert, i. 86, 153
Bertheim, ii. 337
Berthelot, i. 171, 173, 189, 199, 229, 230, 258, 264, 266, 267, 272,
83, 289, 351, 372, 393, 394, 405, 415, 424, 438, 457, 463, 502,
506, 507, 518, 529, 537, 582; ii. 23, 57, 207, 209, 251, 253, 259,
345, 367
Berthier, ii. 8
Berthollet, i. 27, 31, 105, 433, 434, 459, 470, 502, 609
Berzelius, i. 131, 148, 194, 255, 379; ii. 8, 100, 102, 147, 148,
219, 281, 300, 485
Besson, i. 288; ii. 67, 70, 105, 179
Beudant, ii. 7, 8
Bineau, i. 100, 271, 452, 504; ii. 239
Binget, i. 75
Blaese, ii. 188
Blagden, i. 91, 428
Blake, ii. 30
Blitz, ii. 184
Blomstrand, ii. 299
Boerwald, ii. 279
Böttger, i. 595
Bogorodsky, i. 574
Boilleau, i. 415
Boisbaudran, L. de, i. 97, 102, 572, 600; ii. 6, 26, 90, 82, 284, 483
Bornemann, i. 509
Botkin, ii. 30
Bouchardat, ii. 45
Boullay, ii. 55
Bourdiakoff, i. 584, 617
Boussingault, i. 131, 157, 233, 235, 525, 615
Boyle, i. 124
Brand, ii. 150
Brandau, i. 481
Brandes, i. 72
Bravais, i. 233
Brauner, i. 490, 491; ii. 26, 59, 94, 96, 97, 134, 144, 194, 271, 483
Brewster, i. 569
Brigham, ii. 193
Brodie, i. 212, 351, 405; ii. 252
Brooke, ii. 357
Brown, i. 81, 88
Brugellmann, i. 616
Brunn, ii. 182, 189
Bruyn, i. 262
Brühl, i. 263, 337
Brunner, i. 124, 146, 263; ii. 230, 309, 534
Buchner, i. 615
Buckton, ii. 143
Buff, ii. 103
Bunge, i. 288
Bunsen, i. 43, 69, 78, 117, 180, 465, 568, 575, 576, 577; ii. 27, 289
Bussy, i. 75, 594, 619
Butleroff, i. 143
Bystrom, i. 585
Cagniard de Latour, i. 135, 345
Cahours, ii. 143, 173
Cailletet, i. 132, 138; ii. 45
Calderon, i. 596
Callender, i. 134
Calvert, i. 484; ii. 45
Cannizzaro, i. 584, 587
Carey-Lea, ii. 420, 424, 425, 432
Carius, i. 69, 481
Carnelley, i. 483, 515, 555; ii. 22, 29, 30, 31, 64, 143, 486
Carnot, ii. 294, 361
Caron, i. 595, 604, 610; ii. 336
Carrara, i. 213
Cass, ii. 85
Castner, i. 431, 535, 541
Cavazzi, ii. 160, 172, 182
Cavendish, i. 113, 125, 228; ii. 493
Chabrié, i. 229
Chappuis, i. 50, 199, 205, 264
Chapuy, i. 59
Cheltzoff, i. 393, 457, 582; ii. 41, 247
Cherikoff, ii. 102
Chertel, ii. 245
Chevillot, ii. 311
Chevreul, i. 530
Chigeffsky, ii. 62
Christomanos, i. 511
Chroustchoff, i. 353, 444; ii. 122
Chydenius, ii. 148, 485
Ciamician, i. 565, 573; ii. 486
Clark, i. 26
Classen, ii. 146
Clausius, i. 81, 93, 140, 142, 212, 309, 491
Clement, i. 494
Clève, ii. 26, 94, 97, 484
Cloez, i. 207, 246, 377
Clowes, i. 242
Collendar, i. 134
Comaille, i. 596
Comb, ii. 81
Connell, i. 508
Coppet, i. 91, 428, 601
Corenwinder, i. 501
Cornu, i. 565
Courtois, i. 494
Cracow, ii. 380
Crafts, i. 380 ; ii. 80, 83
Cremers, ii. 100
Croissier, i. 251
Crompton, i. 247
Crookes, i. 229, 617; ii. 20, 91, 96, 440, 491
Crum, ii. 79, 311
Cundall, i. 611
Curtius, i. 258, 265
Dahl, ii. 59
Dalton, i. 29, 78, 81, 109, 206, 271, 322
Dana, ii. 8
Davies, i. 484
Davy, i. 37, 114, 195, 255, 364, 460, 463, 484, 489, 494, 533, 541,
594, 604, 617
Deacon, i. 599
Debray, i. 609; ii. 45, 122, 291, 293, 384, 385
De Chancourtois, ii. 20, 26
De Forcrand, ii. 106, 211
De Haën, ii. 189
De Heen, i. 140
Delafontaine, ii. 97, 148, 198
De la Rive, i. 198; ii. 226
Del-Rio, ii. 197
De Saussure, i. 235, 240
De Schulten, ii. 48
Deville, St.-Claire, i. 4, 36, 118, 143, 179, 180, 227, 239, 280, 281,
301, 320, 392, 393, 399, 459, 467, 476, 477, 500, 534, 595, 608, 609;
ii. 48, 80, 83, 85, 102, 147, 156, 198, 289, 309, 321, 352, 373, 374,
429
De Vries, i. 62, 64, 429
Dewar, i. 3, 5, 135, 139, 163, 297, 563, 565, 569, 585; ii. 176, 220
Dick, ii. 414
Dingwall, i. 486
Ditte, i. 72, 403, 430, 457, 509, 539, 618; ii. 64, 65, 85, 189, 249
Dittmar, i. 100, 452; ii. 240
Divers, i. 274, 294; ii. 54
Dixon, i. 171
Döbereiner, i. 145
Dokouchaeff, i. 344
Donny, i. 534
Dossios, i. 502
Draper, i. 465
Drawe, ii. 161
Drebbel, i. 294
Dulong, i. 131, 148, 437
Dumas, i. 28, 131, 148, 150, 233, 302, 320, 379, 471, 476, 584, 586,
604; ii. 22, 37, 62, 101, 156, 420
Dumont, ii. 197
Ebelmann, ii. 65
Eder, i. 566
Edron, ii. 95
Edwards, ii. 311
Egoreff, i. 569
Eissler, i. 553
Elbers, ii. 221
Emich, i. 286, 287
Emilianoff, ii. 126
Engel, i. 457; ii. 130, 132, 189, 206
Engelhardt, i. 530
Eötvös, i. 333
Erdmann, i. 150
Ernst, i. 399
Eroféeff, i. 352
Esson, ii. 314
Étard, i. 72, 516, 615; ii. 288, 335, 356
Ettinger, i. 53, 312
Famintzin, i. 611
Faraday, i. 134, 177, 296, 385, 463, 464
Favorsky, i. 373
Favre, i. 120, 172, 267, 582; ii. 83, 259, 284, 380
Fick, i. 62
Fisher, ii. 424
Fizeau, ii. 31, 429
Flavitzky, i. 21
Fleitmann, ii. 170
Foerster, ii. 375, 389
Forchhammer, ii. 311
Fordos, ii. 257
Fortmann, ii. 230, 366
Fourcroy, i. 114
Fowler, i. 449
Frank, ii. 88
Franke, ii. 311, 313
Frankel, ii. 294
Frankenheim, ii. 7
Frankland, i. 178, 357, 486; ii. 16, 143
Fraunhofer, i. 563
Frémy, i. 228, 489, 492; ii. 74, 131, 133, 142, 229, 290, 359
Freyer, i. 171, 488
Friedheim, ii. 197, 294
Friedel, i. 353, 472; ii. 80, 83, 103, 122
Friedrich, i. 49; ii. 144
Fritzsche, i. 94, 285, 600, 612; ii. 125, 218, 280, 341
Fromherz, ii. 313
Fürst, i. 484
Galileo, i. 7
Garni, i. 582
Garzarolli-Thurnlackh, i. 481
Gattermann, i. 596; ii. 102, 104
Gautier, i. 585
Gavaloffsky, i. 160
Gay-Lussac, i. 40, 61, 71, 93, 170, 302, 307, 406, 412, 460, 463, 464,
467, 500, 506, 508, 511, 515, 534, 539; ii. 8, 56, 256
Geber, i. 17
Gélis, ii. 257
Genth, ii. 359
Georgi, ii. 197
Georgiewics, ii. 64
Gerberts, i. 528
Gerhardt, i. 196, 309, 357, 388
Gerlach, i. 525
Gernez, i. 97; ii. 205
Geuther, i. 281, 283, 285; ii. 176
Gibbs, i. 140, 464; ii. 293, 410
Girault, i. 498
Gladstone, i. 337, 438, 573; ii. 213
Glatzel, ii. 213, 289, 309
Glauber, i. 17, 26, 193, 432
Glinka, i. 607
Goldberg, i. 93
Gooch, i. 484
Gore, i. 489, 492, 493
Graham, i. 62, 63, 98, 143, 155, 388, 429, 518, 601; ii. 77, 114, 131,
163, 170, 296, 307
Granger, ii. 157, 410
Grassi, i. 88
Green, ii. 310
Greshoff, i. 403
Griffiths, i. 135
Grimaldi, i. 537
Groth, ii. 10
Grouven, i. 615
Grove, i. 118, 119
Grünwald, i. 573
Grützner, ii. 296
Guckelberger, ii. 84
Guibourt, ii. 53
Guldberg, i. 439, 464
Güntz, i. 575; ii. 430
Gustavson, i. 443, 444, 472, 505, 547; ii. 29, 175
Guthrie, i. 99, 428, 601
Guy, i. 136
Habermann, ii. 210
Hagebach, i. 573
Hagen, i. 337
Haitinger, i. 593
Hammerl, i. 613
Hanisch, ii. 233
Hannay, i. 352; ii. 135
Harcourt, ii. 314
Hargreaves, i. 515
Harris, ii. 52
Hartley, i. 573; ii. 486
Hartog, ii. 268
Hasselberg, i. 566
Haüy, ii. 7
Haughton, ii. 20
Häussermann, i. 483
Hautefeuille, i. 199, 205, 264, 409, 414, 476, 477, 501, 538; ii. 102,
122, 379
Hayter, ii. 175
Hemilian, i. 132
Hempel, i. 59, 524
Henkoff, i. 530
Henneberg, ii. 170
Henning, ii. 3
Henry, i. 78, 81
Hérard, ii. 191
Hermann, ii. 8, 47, 197
Hermes, i. 529
Hertz, ii. 156
Hess, i. 178, 588
Heycock, i. 537; ii. 128, 448
Hillebrand, ii. 26, 93, 94, 484
Hintze, ii. 10
Hirtzel, ii. 55
Hittorf, ii. 155
Hodgkinson, ii. 432
Höglund, ii. 94
Hofmann, i. 302; ii. 146, 218, 447
Holtzmann, i. 505
Hoppé-Seyler, i. 611
Horstmann, i. 408
Houzeau, i. 202
Hughes, ii. 212
Hugo, ii. 21
Humboldt, i. 170
Humbly, i. 493; ii. 311
Hutchinson, i. 491
Huth, ii. 20
Huyghens, i. 569
Ikeda, ii. 152
Ilosva, i. 202
Inostrantzeff, i. 345; ii. 4
Isambert, i. 250, 257, 408; ii. 41
Ittner, i. 412
Janssen, i. 569
Jawein, ii. 170
Jay, i. 258
Jeannel, i. 104
Joannis, i. 251, 255, 405, 537, 559
Jörgensen, i. 498; ii. 359, 361, 376
Johnson, ii. 45
Jolly, i. 233
Joly, ii. 384, 385
Kamensky, ii. 414
Kammerer, i. 286, 462, 509; ii. 297
Kane, ii. 57
Kapoustin, i. 403
Karsten, i. 427, 428, 541, 599
Kassner, i. 158
Kayander, i. 133, 384; ii. 46
Keiser, i. 150
Kekulé, i. 358, 369, 507; ii. 294
Keyser, ii. 33
Khichinsky, i. 440
Kimmins, i. 510
Kirchhoff, i. 567
Kirmann, ii. 268
Kirpicheff, i. 132
Kjeldahl, i. 249; ii. 249
Klaproth, ii. 7, 145, 147, 301
Kleiber, i. 570
Klimenko, i. 465
Klobb, ii. 357
Klodt, i. 426
Knopp, ii. 338
Knox, i. 489
Kobb, ii. 125
Kobell, ii. 197
Koch, i. 44
Kohlrausch, i. 245, 525
Kolbe, i. 506
Kolotoff, i. 263
Konovaloff, i. 39, 65, 90, 93, 100, 140, 142, 172, 322; ii. 235, 268
Kopp, i. 586, 587, 612; ii. 3, 37
Koucheroff, i. 373
Kouriloff, i. 209, 247, 274; ii. 41
Kournakoff, i. 393; ii. 294, 365, 396
Kraevitch, i. 133, 134
Kraft, i. 65, 88, 537
Krafts, i. 393
Kreisler, i. 233
Kremers, i. 87, 443; ii. 244, 427
Kreider, i. 484
Krönig, i. 81
Krüger, ii. 282, 284
Krüss, ii. 355, 442, 447, 486
Kubierschky, ii. 213
Kühlmann, i. 608
Kuhnheim, i. 612
Kundt, i. 328, 589; ii. 496
Kvasnik, ii. 57
Kynaston, i. 522
Lachinoff, i. 116, 457; ii. 410
Ladenburg, ii, 103
Lamy, ii. 91
Landolt, i. 7, 337
Lang, i. 399
Langer, i. 226, 459, 462
Langlois, i. 570 ; ii. 257
Latchinoff (_see_ Lachinoff), i. 103, 352
Laurent, i. 28, 196, 388, 471, 526; ii. 7, 9, 10, 117, 292
Laurie, i. 106; ii. 32, 442, 486
Lavenig, i. 140
Lavoisier, i. 7, 29, 49, 114, 131, 155, 379, 459
Leblanc, ii. 8
Le Chatelier, i. 158, 172, 350, 393, 399, 585, 588, 611; ii. 51, 65,
420
Le Duc, i. 131, 170
Lémery, i. 125
Lemoine, i. 501; ii. 155
Lerch, i. 405
Leroy, i. 285
Lesc[oe]ur, i. 103
Leton, ii. 425
Levy, ii. 102
Lewes, i. 371
Lewy, i. 232
Lidoff, ii. 209
Liebig, i. 195, 388, 495, 527; ii. 56
Linder, ii. 223
Liés-Bodart, i. 604, 612
Lisenko, i. 373
Liveing, i. 563, 569
Lockyer, i. 565, 569
Loew, ii. 376
Löwel, i. 525, 600; ii. 45, 284, 286
Loewig, i. 528; ii. 77
Loewitz, i. 96
Lossen, i. 262
Louget, i. 489
Louguinine, i. 360
Louise, ii. 81
Lovel, i. 515; ii. 338
Lubavin, i. 593
Lubbert, ii. 85, 170
Ludwig, i. 463
Luedeking, ii. 194
Luff, ii. 321
Lunge, ii. 244, 246
Lüpke, ii. 157
Lvoff, i. 358
Maack, i. 596
Mac Cobb, i. 612
Mac Laurin, i. 553
McLeod, ii. 180
Magnus, i. 93, 510
Mailfert, i. 199
Malaguti, i. 437; ii. 300
Mallard, i. 172, 393, 588; ii. 4
Mallet, i. 493
Maquenne, i. 349, 620, 621
Marchand, i. 150
Marchetti, ii. 288
Maresca, i. 534
Marguerite, ii. 292
Marignac, i. 198, 233, 428, 430, 453, 454, 518, 525, 600, 601; ii. 6,
9, 95, 101, 194, 197, 198, 199, 234, 239, 241, 244, 292, 293, 295,
357, 440, 486
Markleffsky, i. 273
Markovnikoff, i. 373
Maroffsky, ii. 138
Marshall, ii. 253, 365
Matigon, i. 258, 266
Maumené, i. 258
Maxwell, i. 81
Mayow, i. 17
Mendeléeff, i. 99, 132, 133, 136, 141, 275, 321, 357, 373, 377, 406,
426, 427, 428, 506, 587, 596; ii. 27, 33, 93, 94
Menschutkin, i. 171
Mente, ii. 270
Mermé, i. 462
Merz, i. 505
Meselan, i. 463
Metzner, ii. 189
Metchikoff, i. 44
Meusnier, i. 114
Meyer (Lothar), i. 226, 321, 403; ii. 21, 24, 26, 29, 33, 486
Meyer (Victor), i. 135, 171, 294, 303, 320, 427, 459, 462, 467, 488,
506, 508, 558; ii. 43, 48, 52, 80, 129, 184
Meyerhoffer, ii. 410
Miasnikoff, i. 372
Michaelis, ii. 175
Michel, i. 65, 88
Millon, i. 481, 484, 508
Mills, ii. 20
Mitchell, i. 156
Mitscherlich, i. 428, 527; ii. 1, 5, 8, 156, 184, 311, 313
Moissan, i. 202, 349, 353, 490, 564, 585, 621; ii. 66, 67, 70, 88, 100,
107, 147, 174, 196, 289, 295, 309, 311, 313, 321
Mond, i. 129, 400, 405; ii. 345, 367
Monge, i. 114
Monnier, i. 611
Montemartini, i. 279
Moraht, ii. 384
Moreau, ii. 298
Morel, i. 549
Mosander, ii. 97
Mühlhäuser, ii. 66, 107
Muir, ii. 193
Mulder, i. 515
Müller-Erzbach, i. 103
Müller, i. 427; ii. 425
Munster, ii. 443
Müntz, i. 238, 241, 420, 553
Muthmann, ii. 273
Mylius, ii. 375, 389
Naschold, i. 483
Nasini, i. 496; ii. 156
Natanson, i. 282, 409
Natterer, i. 132, 135, 141, 385
Naumann, i. 399, 408
Nernst, i. 62, 148; ii. 3, 50
Nensky, i. 245
Neville, i. 537; ii. 128, 448
Newlands, ii. 21, 26
Newth, i. 505
Newton, i. 7, 29
Nicklès, ii. 10
Nikolukin, i. 491; ii. 144
Nilson, i. 618; ii. 26, 37, 80, 83, 91, 94, 95, 271, 378, 483
Nordenskiöld, i. 241
Norton, i. 76; ii. 94
Nuricsán, ii. 264
Odling, ii. 52
Offer, i. 99
Ogier, i. 321, 509; ii. 159, 182
Olszewski, i. 139, 569; ii. 491, 497
Oppenheim, i. 506
Ordway, ii. 80
Osmond, ii. 326
Ossovetsky, ii. 137
Ostwald, i. 89, 92, 389, 441, 443
Oumoff, i. 62
Pallard, i. 491; ii. 83
Panfeloff, i. 603
Paracelsus, i. 17, 125, 129, 379
Parkinson, i. 596,
Pashkoffsky, i. 595
Pasteur, i. 44, 241, 242
Paterno, i. 496; ii. 156
Pattison Muir, i. 436
Pebal, i. 315, 484
Péchard, ii. 282, 294, 296, 297
Pekatoros, i. 465
Peligot, ii. 299, 301
Pelopidas, ii. 22, 481
Pelouze, i. 463, 464, 480, 610; ii. 229
Penfield, i. 545; ii. 370
Perkin, i. 558; ii. 244
Perman, i. 537
Personne, i. 75, 506, 537
Petit, i. 584, 586
Petrieff, i. 440
Pettenkofer, ii. 22
Pettersson, i. 618, 619; ii. 37, 80, 83, 91, 197, 484
Pfaundler, i. 445; ii. 241, 430
Pfeiffer, i. 64
Pfordten, V. der, ii. 420
Phipson, i. 596; ii. 59
Piccini, ii. 23, 146, 197, 288, 298
Pici, ii. 57
Pickering, i. 88, 91, 99, 104, 106, 272, 333, 452, 517, 525, 529, 613;
ii. 241, 245, 246, 247
Pictet, i. 81, 129, 137; ii. 31, 241
Picton, ii. 223
Pierre, i. 452, 495; ii. 226, 485
Pierson, i. 93
Pigeon, ii. 377
Pionchon, i. 585
Pistor, i. 399
Plantamour, ii. 5
Plaset, ii. 289
Plessy, ii. 257
Plücker, i. 572
Poggiale, i. 427
Poiseuille, i. 355
Poleck, ii. 296
Poluta, ii. 30
Popp, ii. 232
Potilitzin, i. 96, 97, 98, 445, 486, 499, 502, 509, 612; ii. 29, 357
Pott, ii. 100
Poulenc, ii. 174, 289
Prange, ii. 422
Prelinger, ii. 310
Priestley, i. 17, 154, 159, 297, 379, 402
Pringsheim, i. 465
Prost, i. 98, 486
Prout, i. 31; ii. 439
Puchot, i. 452
Pullinger, ii. 389
Quincke, i. 427, 495
Rammelsberg, i. 430, 510, 525; ii. 26, 161, 485
Ramsay, i. 133, 140, 141, 232, 247, 333, 495, 496, 581; ii. 128, 491
Rantsheff, ii. 20
Raoult, i. 91, 274, 330, 331, 332, 429
Rascher, ii. 85
Raschig, i. 263; ii. 229
Rathke, i. 399
Ray, i. 17
Rayleigh, i. 226, 232, 491
Rebs, ii. 213, 217
Recoura, i. 332; ii. 289
Regnault, i. 40, 53, 54, 90, 93, 131, 133, 297, 443, 495, 584, 587,
588; ii. 50, 208, 238
Reich, ii. 91
Reiset, i. 238
Remsen, ii. 335
Retgers, ii. 157, 158, 180
Reychler, ii. 65
Reynolds, i. 581
Richards, i. 526, 585; ii. 32, 432
Riche, i. 509; ii. 127, 292
Richter, i. 193, 194; ii. 91
Ridberg, ii. 21, 24, 486
Riddle, i. 135
Rideal, ii. 297
Roberts-Austen, ii. 486
Robinson, i. 515
Rodger, ii. 213, 263
Rodwell, i. 17
Roebuck, i. 294
Röggs, ii. 119
Rohrer, ii. 343
Roozeboom, i. 106, 452, 453, 464, 496, 506, 511, 599, 613; ii. 3, 226,
341, 410
Roscoe, i. 80, 100, 101, 379, 452, 463, 485, 486, 568, 572; ii. 26,
194, 196, 197, 297, 303, 485
Rose, i. 436, 437, 518, 525, 608, 612; ii. 8, 230, 235, 248, 281, 363,
428, 485
Rosenberg, ii. 351
Rossetti, i. 428
Rouart, Le, ii. 86
Rousseau, i. 354; ii. 337, 366, 378
Roux, ii. 81
Rudberg, ii. 136
Rücker, i. 142
Rüdorff, i. 91, 428, 598, 601
Rybalkin, i. 455
Sabanéeff, i. 371
Sabatier, i. 284; ii. 66
Saint Edmé, ii. 335
Saint Gilles, i. 431
Sakurai, i. 331
Salzer, ii. 161
Sarasin, ii. 122
Sarrau, i. 140, 142
Saunders, ii. 189
Scharples, i. 576
Scheele, i. 155, 161, 412, 459, 462, 608; ii. 100, 150, 291
Scheffer, i. 453
Scheibler, ii. 292, 296
Scherer, ii. 8
Schiaparelli, ii. 318
Schidloffsky, i. 238
Schiloff, i. 212
Schlamp, i. 332
Schiff, i. 430, 588; ii. 106, 267
Schloesing, i. 238, 239, 240, 553, 610
Schmidt, i. 539
Schneider, i. 89
Schöne, i. 208, 209, 211, 394, 617; ii. 15, 72, 219, 251, 488
Schönebein, i. 198, 202, 208, 212, 509; ii. 228, 463
Schottländer, ii. 447
Schröder, i. 75
Schroederer, ii. 366
Schrötter, ii. 153, 284
Schützenberger, i. 511, 579; ii. 102, 107, 228, 367, 389
Schuliachenko, i. 608
Schuller, ii. 180
Schultz, i. 518; ii. 273
Schulze, i. 98; ii. 215
Schuster, i. 572
Schwicker, ii. 227, 230
Scott, i. 405, 537, 558
Sechenoff, i. 80, 86
Seelheim, ii. 379
Sefström, ii. 197
Selivanoff, i. 476, 507, 508
Senderens, i. 284
Serullas, i. 485
Setterberg, i. 576
Seubert, ii. 27, 83, 343, 442
Sewitsch, i. 372
Shaffgotsch, i. 555
Shapleigh, ii. 95
Shenstone, i. 611
Shields, i. 333
Shishkoff, i. 276; ii. 56
Silberman, i. 120, 172; ii. 259
Sims, ii. 268
Skraup, ii. 346
Smith, i. 271
Smithson, ii. 100
Snyders, ii. 100
Sokoloff, ii. 85, 122
Solet, i. 509
Sonstadt, ii. 443
Sorby, i. 88
Soret, i. 66, 202, 203, 427
Spring, i. 38, 98, 434, 486; ii. 45, 50, 133, 223, 258, 288, 314, 423,
427
Stadion, i. 485
Stahl, i. 16
Stas, i. 7, 233, 379, 428, 498, 581; ii. 420, 434, 485
Staudenmaier, ii. 168
Stcherbakoff, i. 97, 428, 458, 601
Stohmann, i. 359, 360, 396
Stokes, i. 355
Stortenbeker, i. 511
Stromeyer, ii. 47
Struvé, i. 208, 612
Tait, i. 203
Tammann, i. 91, 148; ii. 170, 247
Tanatar, i. 511
Tchitchérin, ii. 21
Terreil, ii. 313
Than, i. 317
Thénard, i. 207, 229, 460, 464, 534, 539; ii. 251
Thillot, ii. 170
Thilorier, i. 385
Thomsen, i. 111, 120, 124, 131, 173, 189, 267, 359, 389, 396, 441, 453,
466, 472, 494, 502, 515, 529, 555, 582; ii. 9, 32, 50, 55, 105, 165,
208, 224, 264, 368, 370, 438, 442
Thorpe, i. 142, 285, 445, 493; ii. 27, 160, 173, 213, 259, 263, 268,
301, 313, 442, 486
Thoune, i. 294, 295
Tiemman, i. 213
Tilden, i. 516
Timeraséeff, i. 170
Timoféeff, i. 78
Tessié du Motay, i. 158
Tissandier, i. 78
Titherley, i. 539
Tivoli, ii. 183
Tomassi, ii. 339
Topsöe, i. 506
Tourbaba, i. 88; ii. 247
Trapp, i. 511
Traubé, i. 312, 611; ii. 270
Troost, i. 64, 274, 281, 320, 409, 414, 500, 538; ii. 80, 83, 102, 147,
156, 254, 379
Tscherbacheff, i. 577
Tutton, i. 543; ii. 160, 174, 412
Umoff, i. 429
Unverdorben, ii. 280
Urlaub, ii. 301
Valentine, i. 17
Van der Heyd, i. 599
Van der Plaats, i. 496; ii. 438
Van der Waals, i. 82, 140
Van Deventer, i. 599
Van Helmont, i. 379
Van Marum, i. 198
Van't Hoff, i. 64, 65, 331, 599; ii. 3
Vare, ii. 55
Vauquelin, i. 114, 619; ii. 7
Veeren, i. 612; ii. 45
Veley, i. 279
Verneuille, ii. 225
Vernon, ii. 151
Vèzes, ii. 391
Viard, ii. 285
Vignon, ii. 126, 131
Villard, i. 106, 296, 297
Villiers, ii. 259
Violette, i. 342, 345
Violle, i. 301
Vogt, i. 611
Volkovitch, ii. 201
Voskresensky, i. 345
Waage, i. 439
Wachter, i. 508
Wagner, i. 357
Wahl, ii. 310
Walden, ii. 57
Walker, ii. 143
Walmer, i. 573
Walter, ii. 256
Walters, ii. 234
Wanklyn, i. 100, 539
Warburg, i. 589; ii. 496
Warder, i. 450
Warren, ii. 102
Watson, i. 527; ii. 169
Watts, i. 526
Weber, i. 280, 583; ii. 83, 129, 131, 186, 230, 233, 234, 249
Weith, i. 502
Weitz, ii. 57
Welch, ii. 425
Weller, ii. 146
Wells, i. 477, 545; ii. 57, 370
Welsbach, ii. 96, 97
Weltzien, i. 204, 595
Wenzel, i. 193
Weruboff (_see_ Wyruboff), ii. 4
Weselski, i. 507
Weyl, i. 255
Wheeler, i. 545
Wichelhaus, ii. 179
Wiedemann, i. 439, 588
Wilhelmj, ii. 315
Willgerodt, i. 508; ii. 29
Williamson, ii. 268
Wilm, ii. 376, 388
Winkler, i. 78, 79, 577, 594, 621; ii. 25, 30, 66, 97, 102, 124, 125,
147, 234, 246, 355, 483
Wischin, ii. 384
Wislicenus, i. 267, 294
Witt, ii. 3
Wöhler, i. 410, 619; ii. 85, 103, 107, 146, 285, 289, 420, 425
Wollaston, i. 8
Wreden, i. 507
Wright, ii. 321
Wroblewski, i. 79, 80, 106, 139, 387; ii. 226
Wülfing, ii. 119
Wülner, i. 91, 572
Würtz, i. 301, 476; ii. 171, 173, 213, 267
Wyruboff, ii. 4, 9
Young, i. 134, 136, 140, 141, 247, 495, 496
Zaboudsky, i. 354
Zaencheffsky, i. 140
Zimmermann, ii. 26, 303, 355, 485
Zinin, i. 276
Zörensen, i. 284
Zorn, i. 295
* * * * *
SUBJECT INDEX
Acid, acetic sp. gr. of solutions of, i. 59
-- arsenic, ii. 181
-- bismuthic, ii. 190
-- boric, ii. 64
-- carbamic, i. 408
-- chamber, i. 294
-- chloric, i. 482
-- chloroplatino-phosphorous, ii. 390
-- chlorosulphonic, ii. 268
-- chlorous, i. 481
-- chromic, i. 208; ii. 282
-- chromo-sulphuric, ii. 288
-- cyanic, i. 409
-- cyanuric, i. 409
-- dithionic, ii. 256
-- ferric, ii. 344
-- fluoboric, ii. 69
-- graphitic, i. 351
-- hydriodic, i. 501, 503, 505, 506
-- hydroborofluoric, ii. 69
-- hydrobromic, i. 80, 503, 505, 506
-- hydrochloric, i. 448, 451, 453
-- hydrocyanic, i. 406, 411
-- hydroferrocyanic, ii. 348
-- hydrofluoric, i. 49
-- hydrofluosilic, ii. 106
-- hydroplatinocyanic, ii. 386
-- hydrosulphurous, ii. 228
-- hydroruthenocyanic, ii. 388
-- hypochlorous, i. 479, 481
-- hyponitrous, i. 265, 294
-- hypophosphoric, ii. 161
-- hypophosphorous, ii. 172
-- iodic, i. 100, 508
-- isethionic, ii. 250
-- metantimonic, ii. 188
-- metaphosphoric, ii. 162, 169
-- metastannic, ii. 131
-- molybdic, ii. 292
-- nitric, i. 268, 272
-- Nordhausen, ii. 233
-- orthophosphoric, ii. 162
-- osmic, ii. 384
-- pentathionic, ii. 257
-- percarbonic, i. 394
-- perchloric, i. 484
-- periodic, i. 510
-- permanganic, ii. 313
-- permolybdic, ii. 297
-- pernitric, i. 264
-- persulphuric, ii. 251
-- pertungstic, ii. 297
-- phosphamic, ii. 179
-- phosphamolybdic, ii. 293
-- phosphorous, ii. 171
-- polysilicic, ii. 117
-- pyrophosphoric, ii. 169
-- pyrosulphuric, ii. 234
-- silenic, ii. 272
-- silicotungstic, ii. 295
-- stannic, ii. 130
-- sulphonic, ii. 249
-- sulphuric, i. 76, 77, 89, 111, 290, 294; ii. 235, 238, 241
-- telluric, ii. 272
-- tetrathionic, ii. 257
-- thiocarbonic, ii. 263
-- thiocyanic, ii. 263
-- thionic, ii. 255
-- thiosulphuric, ii. 230
-- trithionic, ii. 257
-- tungstic, ii. 292, 294
-- vanadic, ii. 196
Acids, i. 185
-- avidity of, i. 389, 442
-- basicity of, i. 387
-- complex, i. 197; ii. 293
-- fuming, i. 102
-- organic, i. 394, 396, 405
Acetylene, i. 372
Actinium, ii. 59
Affinity, chemical, i. 26, 389
Air, i. 131, 231, 233
Alchemy, i. 14
Alcohol, i. 53, 88
Alkali, metals, i. 558, 577
-- waste, ii. 204
Alkalis, i. 186
Allotropism, i. 207
Alloys, i. 537; ii. 128
Alumina, ii. 75
Aluminium, ii. 70, 85
-- bromide, ii. 84
-- bronze, ii. 88
-- carbide, ii. 88
-- chloride, ii. 80, 83
-- double chlorides, ii. 84
-- fluoride, ii. 83
-- hydroxide, ii. 75
-- iodide, ii. 85
-- nitrate, ii. 80
-- sulphate, ii. 82
Alums, ii. 5, 82, 343
Alunite, ii. 80
Amalgams, ii. 58
Amides, i. 258, 406
Amidogen, i. 258
-- hydrate, i. 258
Amines, i. 416
Ammonia, i. 229, 246
-- of crystallisation, i. 257
-- heat of solution of, i. 74
-- in air, i. 240
-- liquefaction of, i. 250
-- salts, i. 254
-- soda process, i. 524
-- solutions of, i. 80, 252
Ammonium, i. 254
-- amalgam, i. 255
-- bicarbonate, i. 527
-- carbamate, i. 407, 408
-- carbonate, i. 407
-- cobalt salts, ii. 359
-- dichromate, ii. 279
-- molybdate, ii. 292
-- nitrate, i. 273, 274
-- nitrite, i. 284
-- phosphates, ii. 167
-- sulphate, ii. 269
-- sulphide, ii. 218
Analogy of elements, i. 573, 578
Anthracite, i. 345
Antimoniuretted hydrogen, ii. 189
Antimony, ii. 186
-- chlorides, ii. 189
-- oxides, ii. 187, 188
-- sulphides, ii. 221
Aqua Regia, i. 467
Aqueous radicle, i. 213
Argon, i. 226, 232; App. III.
Arsenic, ii. 179
-- anhydride, ii. 181
-- sulphides, ii. 221
-- tribromide, ii. 181
-- trichloride, ii. 180
-- trifluoride, ii. 181
Arsenious anhydride, ii. 184
-- oxychloride, ii. 180
Arsenites, ii. 185
Arseniuretted hydrogen, ii. 182
Astrakhanite, i. 59
Atmolysis, i. 156
Atomic theory, i. 216
-- volumes, ii. 33
-- weights, i. 21
Atoms and molecules, i. 322
Barium, i. 614, 617
-- chlorate, i. 483
-- chloride, i. 615
-- hydroxide, i. 616
-- metatungstate, ii. 295
-- nitrate, i. 615
-- oxide, i. 616
-- peroxide, i. 157, 160, 209, 617
-- sulphate, i. 614, 615
Bauxite, ii. 76
Benzalazine, i. 258
Berthollet's doctrine, i. 433
Beryllium, i. 618
-- atomic weight of, i. 325, 618
-- chloride, i. 584
-- oxide, i. 619
Binary theory, i. 195
Bismuth, ii. 189
-- nitrates, ii. 192
-- oxides, ii. 190, 191
Blast furnace, ii. 324
Bleaching, i. 469
-- powder, i. 162, 477
Boiling point, absolute, i. 130
Borates, ii. 63
Borax, ii. 61
Boric anhydride, ii. 64
Boron, ii. 60, 66
-- chloride, ii. 69
-- fluoride, ii. 67, 68
-- iodide, ii. 70
-- nitride, i. 227; ii. 67
-- oxide, ii. 60
-- specific heat of, i. 585
-- sulphide, ii. 62
Bromides, ii. 32
Bromine, i. 494
Bronze, ii. 127
Butyl alcohol, solubility of, i. 75
Cadmium, ii. 47
-- iodide, ii. 48
-- oxide, ii. 48
-- sulphide, ii. 47
Cæsium, i. 576
Calcium, i. 590, 604
-- carbonate, i. 592, 608, 609, 610
-- chloride, i. 237, 612
-- -- crystallohydrates of, i. 613
-- fluoride, i. 491
-- hypochlorite, i. 162
-- iodide, i. 604
-- peroxide, i. 607
-- phosphate, ii. 167
-- sulphate, i. 611
-- sulphide, ii. 220
Calomel, ii. 54
Carbamide, i. 409
Carbides, i. 349, 353
Carbon, i. 338
-- bisulphide, ii. 258
-- molecule of, i. 354
-- oxysulphide, ii. 264
-- tetrachloride, i. 473
Carbonic anhydride, i. 379
-- -- assimilation of by plants, i. 393
-- -- dissociation of, i. 392, 393, 399
-- -- in air, i. 238, 242
-- -- liquid, i. 385
-- -- solutions of, i. 80, 86
-- -- specific heat of, i. 393
Carbonic oxide, i. 396
-- -- and nickel, i. 405
Carborundum, ii. 107
Carboxyl, i. 395
Carnallite, i. 421, 544, 560
Catalytic phenomena, i. 211
Caustic potash, i. 550
-- soda, i. 529
Cements, ii. 122
Cerite metals, ii. 93
Cerium, ii. 93
Chamber crystals, i. 290; ii. 230
Charcoal, i. 343
Chemical change, rate of, ii. 314
-- transformations, i. 3
Chloranhydrides, i. 468; ii. 174, 175, 177
Chlorates, i. 482
Chlorides, i. 455, 466; ii. 31
Chlorine, i. 463
-- compounds, heat of formation of, i. 44
-- crystallohydrates of, i. 464
-- oxides, i. 479
-- preparation of, i. 460
-- solubility of, i. 463
Chloroform, i. 473
Chlorophosphamide, ii. 179
Chloryl compounds, i. 476
Chrome alum, ii. 283
Chromic acid, i. 208
-- anhydride, ii. 280
-- oxide, ii. 284, 285
Chromium, ii. 276, 289
-- chlorides, ii. 285
-- fluorides, ii. 280, 289
Chromyl chloride, ii. 281
Chryseone, ii. 108
Clay, ii. 70
Coal, i. 345
Cobalt, ii. 353
-- dioxide, ii. 366
-- fluoride, ii. 358
Cobaltamine salts, ii. 359
Cobaltic oxide, ii. 362
Cobalto-amine, ii. 359
Cobaltous hydroxide, ii. 358
Cohesion of liquids, i. 52
Coke, i. 345
Collodion cotton, i. 275
Colloids, i. 63; ii. 77, 423
Combination, chemical, i. 3
Combining weights, i. 21; ii. 439
Combustion, imperfect, i. 341
-- heat of, i. 172, 176, 399, 400
Compounds, definite and indefinite, i. 31
-- types of, ii. 10
Compressibility of solutions, i. 88
Conductivity, electro-molecular, i. 389
Contact reactions, i. 163, 290
Copper, ii. 400
-- carbonate, ii. 411
-- complex salts of, ii. 412
-- nitrate, ii. 411
-- nitride, ii. 409
-- sulphate, ii. 413
Corundum, ii. 75
Critical points, i. 141
Cryohydrates, i. 99
Cryoscopic investigations of solutions, i. 90, 332
Crystals, i. 51
Crystalline form, ii. 7
Crystallo-hydrates, i. 102
Crystalloids, i. 63
Cupellation, ii. 417
Cyanides, i. 406
Cyanogen, i. 406, 414
-- chloride, ii. 176
Decomposition, chemical, i. 4
Deliquescence, i. 104
Delta metal, ii. 414
Desiccator, i. 58
Detonating gas, i. 115, 170, 173
Depression of freezing point of solutions, i. 90, 92, 330
Dialysis, i. 63; ii. 114
Diamond, i. 350, 353
Didymium, ii. 93
Diffusion, rate of, i. 63
Dimorphism, i. 610, ii. 178
Disinfectants, i. 245
Disodium orthophosphate, ii. 166
Dissociation, i. 36, 282, 608
Distillation, dry, i. 4, 247, 342
Dust, atmospheric, i. 241
Efflorescence, i. 103
Ekacadmium, ii. 59
Ekasilicon, ii. 25
Electro-chemical theory, i. 195
Electric energy and thermal units, i. 582
Electrolysis, i. 116
Elements, i. 20
-- grouping of, ii. 1
-- typical, ii. 19
Emulsions, i. 98
Energy, chemical, i. 29
Equations, chemical, i. 278
Equivalents, law of, i. 194
Equivalent weights, i. 581
Ethane, i. 366
Ether, critical points of, i. 141
Ethylene, i. 370
Ethyl silicates, i. 104
Euchlorine, i. 484
Eudiometer, i. 169
Expansion, linear, of elements, ii. 31
Explosion, rate of transmission of, i. 171
Explosives, i. 275, 276
Felspar, ii. 122
Fermentation, i. 242
Ferric chloride, i. 558; ii. 340
-- hydrates, ii. 339
-- nitrate, ii. 340
-- orthophosphate, ii. 342
-- oxide, ii. 339
Ferrous chloride, ii. 335
-- sulphate, ii. 335
-- -- solubility of, i. 72
-- sulphide, ii. 210
Flame, i. 177, 179
Fluoborates, ii. 69
Fluorides, i. 491, 493
Fluorine, i. 203, 489
Fluorspar, i. 491
Formula, chemical, i. 151, 326
Freezing mixtures, i. 76
Fuel, calorific capacity of, i. 360
Furnace, electrical, i. 352
Fusco-cobaltic salts, ii. 360
Gadolinite metals, ii. 93
Gallium, ii. 88, 90
Gas, illuminating, i. 361
-- producers, i. 397
Gases, absorption of, i. 348
-- diffusion of, i. 83
-- expansion of, i. 133
-- liquefaction of, i. 134, 135, 137
-- measurement of, i. 78, 300
-- solution of, i. 68, 78, 86
-- theory of, i. 81, 83, 140
Germanium, ii. 26, 124
-- chloride, ii. 125
-- oxide, ii. 125
Glass, i. 123
-- soluble, ii. 110
Glauber's salt, i. 517
Glycols, ii. 117
Gold, ii. 442
-- alloys, i. 446, 447
-- chlorides, ii. 448, 450
-- colloid, ii. 447
-- cyanide, ii. 450
-- extraction of, ii. 444, 445
-- fulminating, ii. 450
-- oxides, ii. 448
-- refining, ii. 446
Graduators, i. 424
Graphite, i. 350, 351
Gros' salt, ii. 393
Guignet's green, ii. 285
Gunpowder, i. 557
Gypsum, i. 593, 611
Halogens, i. 445, 487, 499
Halogen compounds, heat of formation of, i. 494, 502; ii. 32
-- -- boiling-points of, i. 502
Hausmannite, ii. 10
Helium, i. 570; ii. 498
Hemimorphism, ii. 9
Homeomorphism, ii. 8
Homologous compounds, i. 358
Humus, i. 344
Hydrates, i. 109, 185
Hydrazine, i. 258
Hydrides, i. 621; ii. 23
Hydrocarbons, i. 355, 359
Hydrogen, i. 123, 129, 130, 142, 143, 146
-- pentasulphide, ii. 217
-- peroxide, i. 207, 312
Hydrosols, i. 98
Hydroxyl, i. 192, 213
Hydroxylamine, i. 262
Hypochlorites, i. 481
Hyponitrites, i. 294
Imides, i. 258
Indium, ii. 27, 37, 88, 97
Iodates, i. 509
Iodides, ii. 32
-- of nitrogen, i. 507
Iodine, i. 320, 321, 496, 497, 498
-- chlorides of, i. 511
Iodosobenzol, i. 508
Iridious oxide, ii. 382
Iridium, ii. 382
Iron, i. 585; ii. 317, 322
-- and carbonic oxide, ii. 345
-- cast, ii. 325
-- nitride, ii. 346
-- ores, ii. 319
-- sulphate, ii. 335
Isethionic acid, ii. 250
Isomorphism, i. 203, 368; ii. 1, 4, 8
Kaolin, ii. 70
Lakes, ii. 77
Lanthanum, ii. 93
Laughing gas, ii. 297
Law of Avogadro-Gerhardt, i. 309
-- -- Berthollet, i. 445
-- -- Boyle and Mariotte, i. 132
-- -- combining weights, i. 221
-- -- Dulong and Petit, i. 584
-- -- equivalents, i. 194
-- -- even numbers, i. 357
-- -- Gay Lussac, i. 133, 304, 307
-- -- Guldberg and Waage, i. 441
-- -- Henry and Dalton, i. 78
-- -- indestructibility of matter, i. 6
-- -- Kirchoff, i. 568
-- -- limits, i. 357
-- -- maximum work, i. 120
-- -- multiple proportions, i. 109, 214
-- -- partial pressures, i. 82
-- -- periodic, ii. 17
-- -- phases, ii. 410
-- -- reversed spectra, i. 568
-- -- specific heats, i. 584
-- -- substitution, i. 260, 365
-- -- volumes, i. 304
Lead, ii. 134
-- acetate, ii. 137
-- carbonate, ii. 140
-- chloride, ii. 139
-- chromate, ii. 136, 279
-- dioxide, ii. 142
-- nitrate, ii. 139
-- oxide, ii. 137
-- red, ii. 142
-- salts of, i. 491
-- tetrachloride, ii. 144
-- tetrafluoride, ii. 144
-- white, ii. 140
Leucone, ii. 107
Levigation, ii. 72
Light, chemical action of, i. 465
Lime, i. 605
Liquids, boiling points of, i. 135
Lithium, i. 574
-- carbonate, i. 575
Litharge, ii. 137
Litmus, i. 185
Lixiviation, methodical, i. 521
Luteocobaltic salts, ii. 359
Magnus' salt, ii. 392
Magnesia, i. 597
Magnesium, i. 590, 594
-- carbonate, i. 592, 602
-- chloride, i. 602
-- crystallohydrates of, i. 601
-- double salts of, i. 597
-- nitride, i. 595
-- silicide, ii. 102
-- sulphate, i. 600
Manganese, ii. 303
-- nitrides, ii. 310
-- oxides, ii. 306, 307, 308, 313
-- peroxide, i. 159; ii. 305
-- sulphate, ii. 307
Mass, influence of, i. 32, 436
Matches, ii. 154
Matter, primary, ii. 440
-- transmutability of, i. 14
Mercury, ii. 48
-- ammonia compounds, ii. 57
-- basic salts of, ii. 54
-- chlorides, ii. 52, 53, 54
-- compounds, heat of formation, ii. 50
-- cyanide, ii. 55
-- fulminating, ii. 56
-- iodide, ii. 55
-- nitrates, ii. 51
-- nitrides, ii. 56
-- oxides, ii. 53
-- sulphate, ii. 57
-- sulphides, ii. 221
Metalepsis, i. 28, 471
Metalloids, i. 23
Metals, i. 23
-- of alkaline earths, i. 64, 590, 591
-- of alkalis, i. 543
-- displacement of, ii. 427
Methane, i. 360
Moisture, determination of, in gases, i. 40
-- influence upon reaction, i. 403
Molecular volumes, ii. 37
-- weight and boiling point, i. 331
-- -- -- coefficient of refraction, i. 336
-- -- -- latent heat, i. 329
-- -- -- specific gravity of solutions, i. 335
-- -- -- surface tension, i. 334
Molecules, i. 319, 322
Molybdates, ii. 292
Molybdenum, ii. 290
-- anhydride, ii. 291
-- fluo-compounds, ii. 298
-- sulphides, ii. 297
Monophosphamide, ii. 178
Monosodium orthophosphate, ii. 167
Morphotropy, ii. 10
Naphtha, i. 373, 377
Nascent state, i. 33, 145, 146
Neodymium, ii. 97
Nickel, ii. 353
-- alloys, ii. 367
-- and carbonic oxide, ii. 367
-- fluoride, ii. 358
-- hydroxide, ii. 358
-- oxide, ii. 365
-- sulphate, i. 97; ii. 359
-- tetra-carboxyl, ii. 367
Niobium, i. 199; ii. 194, 198
Nitrates, i. 273
Nitres, i. 268, 555
Nitric anhydride, i. 280
-- oxide, i. 286
Nitrides, i. 227, 258, 620
Nitriles, i. 406
Nitrites, i. 284
Nitro-cellulose, i. 275
Nitro-compounds, i. 274
Nitrogen, i. 223, 225, 475
-- chloride, i. 476
-- iodide, i. 507
-- oxides of, i. 267, 280, 284, 294, 295
-- sulphide, ii. 270
Nitro-prussides, ii. 351
Nitroso-compounds, i. 288
Nitrosulphates, ii. 229
Nitrosyl chloride, ii. 176
Norwegium, ii. 59
Occlusion, i. 143
Olefiant gas, i. 370
Organo-metallic compounds, i. 356
Osmium, ii. 372, 382, 384
Osmotic pressure, i. 64
Osmuridium, ii. 383
Oxamide, i. 406
Oxidation, i. 16
Oxides, i. 183; ii. 36
Oxycobaltamine salts, ii. 359
Oxygen, i. 152, 157, 158, 163
-- compounds, heat of formation of, i. 120, 466
Ozone, i. 198, 229
Palladium, ii. 369
-- hydride, i. 143; ii. 380
Palladous chloride, ii. 379
-- iodide, ii. 379
Paracyanogen, i. 414
Paramorphism, ii. 9
Parasulphatammon, ii. 269
Peat, i. 344
Peligot's salt, ii. 281
Percentage composition, i. 326
Perchloric anhydride, ii. 282
Periodates, i. 510
Permanganic anhydride, ii. 313
Permolybdates, ii. 297
Peroxide, chloric, i. 484
Peroxides, i. 159; ii. 15, 23
Perstannic oxide, ii. 133
Persulphates, ii. 253
Petroleum, i. 373
Phenol, solubility of, i. 75
Phlogiston, i. 17
Phosgene gas, ii. 175
Phospham, ii. 178
Phosphides, ii. 157
Phosphine, ii. 158, 160
Phosphonium iodide, ii. 159
Phosphoric anhydride, ii. 161
Phosphorous anhydride, ii. 160
Phosphorus, ii. 149
-- ammonium compounds, ii. 178
-- chlorides, ii. 174
-- fluorides, ii. 173
-- iodides, i. 505, 506; ii. 172
-- oxychlorides, ii. 175
-- sulphides, ii. 213
-- sulphochloride, ii. 213
-- thermo-chemical data for, ii. 153
Phosphuretted hydrogen, ii. 158, 160
Photography, ii. 431
Photo-salts, ii. 432
Plants, chemical reactions in, i. 547
-- and nitrogen, i. 230
Platinic chloride, ii. 377
-- hydroxide, ii. 379
Platino-ammonium compounds, ii. 391
-- chlorides, i. 467; ii. 378
-- cyanides, ii. 386
-- nitrites, ii. 390
-- sulphites, ii. 390
Platinous chloride, ii. 379
Platinum, ii. 376
-- alloys, ii. 373
-- black, ii. 376
-- metals, ii. 369, 375
-- oxide, ii. 378
Poly-haloid salts, i. 545
Polymerism, i. 207, 367
Polysulphides, ii. 217
Potassium, i. 544, 558
-- aurate, ii. 449
-- bromide, i. 550
-- carbonate, i. 549
-- chlorate, i. 161, 482
-- chloride, i. 72, 543
-- chromate, ii. 280
-- cyanide, i. 412, 551
-- dichromate, ii. 278
-- ferricyanide, ii. 346
-- ferrocyanide, i. 346, 412
-- hydrosulphide, ii. 219
-- hydroxide, i. 548
-- iodide, i. 550
-- manganate, ii. 310
-- nitrate, i. 553
-- oxides, i. 559
-- permanganate, ii. 311
-- stannate, ii. 133
-- sulphate, i. 72, 549
-- sulphide, ii. 219
-- telluride, ii. 274
Praseocobaltic salts, ii. 361
Praseodidymium, ii. 97
Proteid substances, i. 224
Prout's hypothesis, ii. 439
Prussian blue, i. 419; ii. 349
Purpureocobaltic salts, ii. 361
Purpureotetramine salts, ii. 361
Pyrocollodion, i. 275
Pyronaphtha, i. 375
Pyrosulphuryl chloride, i. 321; ii. 235
Reactions, chemical, i. 3
-- -- conditions for, i. 34
-- -- contact, i. 39
-- -- endothermal, i. 30
-- -- exothermal, i. 30
-- -- limit of, i. 437
-- -- rate of, ii. 152
Recalescence, ii. 333
Reduction, i. 16
Refraction equivalent, i. 336
Regenerative furnaces, i. 398
Reiset's salts, ii. 394
Respiration, i. 152, 154, 387
Rhodium, ii. 381
Rock salt, i. 421
Roseocobaltic salts, ii. 360
Rosetetramine salts, ii. 361
Rubidium, i. 576
Ruthenium, ii. 372, 382, 384
Sal-ammoniac, i. 248, 318, 457
-- solubility of, i. 458
-- vapour density of, i. 317
Salts, i. 187, 419
-- acid, i. 193, 533
-- basic, i. 193, 533; ii. 54
-- double, i. 598
-- electrolysis of, i. 191
-- heat of formation, i. 189
-- melting points of, i. 135
-- pyro, i. 193
-- theory of, i. 193
Saponification, i. 530
Scandium, ii. 94
Selenium, ii. 273
-- chlorides, ii. 275
Selenious anhydride, ii. 271
Silica, i. 100; ii. 108
-- soluble, ii. 113
Silicates, i. 544; ii. 116
Silicon, ii. 99
-- chloride, ii. 103, 104
-- chloroform, ii. 103
-- bromide, ii. 104
-- fluoride, ii. 105
-- hydride, ii. 102, 103
-- iodide, ii. 105
-- iodoform, ii. 105
Silver, ii. 418
-- allotropic varieties of, ii. 421
-- bromide, ii. 429
-- chlorate, ii. 437
-- chloride, ii. 429
-- cyanide, ii. 433
-- fluoride, ii. 430
-- fulminating, ii. 426
-- hyponitrite, i. 294
-- iodide, ii. 429
-- nitrate, ii. 426
-- nitrite, i. 284
-- orthophosphate, ii. 164
-- oxides, ii. 424
-- peroxide, ii. 422
-- plating, ii. 434
-- soluble, ii. 420
-- subchloride, ii. 432
Slags, ii. 323
Smalt, ii. 354
Soaps, i. 531
Soda ash, i. 519
-- caustic, i. 527
-- manufacture of, i. 459
-- waste, i. 522
Soda lime, i. 237
Sodamide, i. 539
Sodium, i. 513, 533
-- alloys, i. 559
-- amalgams, i. 537
-- bicarbonate, i. 526
-- carbonate, i. 519, 525
-- -- crystallohydrates of, i. 108
-- -- manufacture of, i. 523
-- -- solutions of, i. 525
-- chloride, i. 419
-- -- double salts of, i. 430
-- -- solutions of, i. 88, 99, 429
-- hydride, i. 537
-- hydroxide, i. 528, 529
-- -- solutions of, i. 529
-- nitrate, i. 269
-- -- solutions of, i. 72
-- organo compounds of, i. 540
-- oxides, i. 540, 541
-- phosphates, ii. 166
-- platinate, ii. 378
-- pyrosulphate, i. 518
-- sesquicarbonate, i. 526
-- stannate, ii. 133
-- subchloride, i. 540
-- sulphate, i. 513
-- -- acid salt, i. 518
-- -- crystallohydrates of, i. 515
-- -- solutions of, i. 73, 515, 516
-- sulphite, ii. 226
-- thiosulphate, ii. 230
-- -- solutions of, i. 74
-- tungstate, ii. 294
Soils, i. 344; ii. 73
Solubility, coefficient of, i. 67, 71
Solutions, i. 330
-- aqueous, i. 59
-- boiling points of, i. 94, 100
-- crystallisation of, i. 427
-- colour of, i. 95
-- diffusion of, i. 61, 429
-- of double salts, i. 599
-- formation of ice from, i. 91, 428
-- heat of formation of, i. 74, 75, 76
-- of gases, i. 68
-- isotonic, i. 64
-- saturated, i. 65
-- specific gravity of, i. 429, 584
-- supersaturated, i. 96
-- theory of, i. 64, 89, 92, 97, 106, 215, 323, 608; ii. 3, 164
-- vapour tension of, i. 90, 92
-- volumes of, i. 87
-- Specific heat, i. 585, 586, 588
Spectra absorption, i. 566
Spectrum analysis, i. 560, 561
Stannic chloride, ii. 132
-- fluoride, ii. 132
-- oxide, ii. 130
-- sulphide, ii. 132
Stannous chloride, ii. 130
-- oxide, ii. 129
-- salts, ii. 129
Steam, vapour tension of, i. 54
Steel, ii. 327, 328, 330
Strontium, i. 615
-- chloride, i. 615
-- hydroxide, i. 615
-- nitrate, i. 615
-- oxide, i. 617
Substitution chemical, i. 5
Sulphamide, ii. 270
Sulphatammon, ii. 269
Sulphates, ii. 248
Sulphides, i. 98; ii. 213
Sulphonitrites, ii. 229
Sulphoxyl, ii. 250
Sulphur, ii. 200
-- chlorides of, ii. 264
Sulphuretted hydrogen, ii. 208
Sulphuric anhydride, ii. 232
-- peroxide, ii. 251
Sulphurous anhydride, ii. 224
Sulphuryl chloride, ii. 268
Superphosphates, ii. 168
Tantalum, ii. 194, 198
Tellurium, ii. 274
-- bromide, ii. 275
-- chlorides, ii. 275
Tellurous anhydride, ii. 271
Temperature, critical, i. 131
Test papers, i. 185
Thallium, ii. 88, 91
Thallic oxide, ii. 93
Thallous hydroxide, ii. 92
-- oxide, ii. 92
Thiocarbonates, ii. 262
Thionyl chloride, ii. 267
Thiophosgene, ii. 262
Thiophosphoryl fluoride, ii. 263
Theory, atomic, i. 216
-- unitary, i. 195
-- vortex, i. 217
Thermochemistry, i. 173
Thorium, ii. 148
Tin, ii. 125
-- alloys, ii. 127
Titanium, ii. 144
-- chloride, ii. 145
-- nitride, ii. 146
-- nitrocyanide, ii. 146
-- oxides, ii. 145
Tripoli, ii. 110
Trisodium orthophosphate, ii. 166
Tungstates, ii. 292
Tungsten, ii. 290
-- anhydride, ii. 291
-- nitride, ii. 297
-- sulphide, ii. 297
Turnbull's blue, ii. 350
Types of combination, ii. 10
Ultramarine, ii. 84
Uranium, ii. 30, 297
-- atomic weight of, ii. 26
-- dioxide, ii. 301
-- oxides, ii. 298
-- tetrachloride, ii. 301
Urano-alkali compounds, ii. 298
Uranyl, ii. 301
-- ammonium carbonate, ii. 300
-- nitrate, ii. 300
-- phosphate, ii. 300
Urea, i. 409
Valency of elements, i. 404, 418, 581
Van der Waal's formula, i. 82, 140
Vanadic anhydride, ii. 196
Vanadium, ii. 194
-- oxychloride, ii. 195
Vapour density, determination of, i. 301
Ventilation, i. 244
Viscosity, i. 355
Volumes, molecular, ii. 4
-- gases, i. 300
Water, i. 40
-- composition of, i. 114, 118, 148, 169, 305, 333
-- compressibility of, i. 53
-- of constitution, i. 109
-- of crystallisation, i. 95, 510
-- dissociation of, i. 118
-- expansion of, i. 53
-- gas, i. 129, 400, 401
-- hard, i. 47
-- hygroscopic, i. 56
-- mineral, i. 45
-- rain, i. 43
-- river, i. 43
-- sea, i. 46
-- specific heat of, i. 52
-- -- gravity of, i. 50
-- spring, i. 44
Wave lengths, i. 564
Wood, i. 339
Ytterbium, ii. 93
Yttrium, ii. 93
Zinc, ii. 39
-- ammonia-chlorides, ii. 41
-- chloride, ii. 40, 41
-- compounds, heat of formation of, ii. 51
-- oxide, ii. 39, 40
-- sulphate, ii. 39
Zirconium, ii. 146
-- chloride, ii. 147
-- hydroxide, ii. 147
-- oxide, ii. 147
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Transcriber's Notes
The half-title has been deleted.
Spelling and punctuation of the original text were taken over
and obvious errors have been corrected silently.
Inconsistant hyphenation was kept as it is.
Table III has been split in three parts.
Text in italics has been marked with underscores (_text_).
The ligature oe has been marked as [oe].
The asterism sign has been marked as [asterism].
The 3root sign has been marked as [3root].
The sign ^ has been used as a superscript.
End of the Project Gutenberg EBook of The Principles of Chemistry. Volume II
(of 2), by D. Mendeléeff
*** END OF THE PROJECT GUTENBERG EBOOK 54210 ***