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