compound, otherwise the reaction goes further. It is easy to find
out when the action of the chlorine on potassium ferrocyanide must
cease; it is only necessary to take a sample of the liquid and add
a solution of a ferric salt to it. If a precipitate of Prussian
blue is formed, more chlorine must be added, as there is still
some undecomposed ferrocyanide, for the ferricyanide does not give
a precipitate with ferric salts. Potassium ferricyanide, like the
ferrocyanide, easily exchanges its potassium for hydrogen and
various metals by double decomposition. With the salts of tin,
silver, and mercury it forms yellow precipitates, and with those
of uranium, nickel, cobalt, copper, and bismuth brown
precipitates. The lead salt under the action of sulphuretted
hydrogen forms lead sulphide and a hydrogen salt or acid,
H_{3}FeC_{6}N_{6}, corresponding with potassium ferricyanide,
which is soluble, crystallises in red needles, and resembles
hydroferrocyanic acid, H_{4}FeC_{6}N_{6}. Under the action of
reducing agents--for instance, sulphuretted hydrogen,
copper--potassium ferricyanide is changed into ferrocyanide,
especially in the presence of alkalis, and thus forms a rather
energetic _oxidising agent_--capable, for instance, of changing
manganous oxide into dioxide, bleaching tissues, &c.
[31] It is important to mention a series of readily crystallisable
salts formed by the action of nitric acid on potassium and other
ferrocyanides and ferricyanides. These salt contain the elements
of nitric oxide, and are therefore called _nitro-(nitroso)
ferricyanides_ (_nitroprussides_). Generally a crystalline sodium
salt is obtained, Na_{2}FeC_{5}N_{6}O,2H_{2}O. In its composition
this salt differs from the red sodium salt, Na_{3}FeC_{6}N_{6}, by
the fact that in it one molecule of sodium cyanide, NaCN, is
replaced by nitric oxide, NO. In order to prepare it, potassium
ferrocyanide in powder is mixed with five-sevenths of its weight
of nitric acid diluted with an equal volume of water. The mixture
is at first left at the ordinary temperature, and then heated on a
water-bath. Here ferricyanide is first of all formed (as shown by
the liquid giving a precipitate with ferrous chloride), which then
disappears (no precipitate with ferrous chloride), and forms a
green precipitate. The liquid, when cooled, deposits crystals of
nitre. The liquid is then strained off and mixed with sodium
carbonate, boiled, filtered, and evaporated; sodium nitrate and
the salt described are deposited in crystals. It separates in
prisms of a red colour. Alkalis and salts of the alkaline earths
do not give precipitates: they are soluble, but the salts of iron,
zinc, copper, and silver form precipitates where sodium is
exchanged with these metals. It is remarkable that the sulphides
of the alkali metals give with this salt an intense bright purple
coloration. This series of compounds was discovered by Gmelin and
studied by Playfair and others (1849).
This series to a certain extent resembles the nitro-sulphide
series described by Roussin. Here the primary compound consists of
black crystals, which are obtained as follows:--Solutions of
potassium hydrosulphide and nitrate are mixed, and the mixture is
agitated whilst ferric chloride is added, then boiled and
filtered; on cooling, _black crystals_ are deposited, having the
composition Fe_{6}S_{3}(NO)_{10},H_{2}O (Rosenberg), or, according
to Demel, FeNO_{2},NH_{2}S. They have a slightly metallic lustre,
and are soluble in water, alcohol, and ether. They absorb the
latter as easily as calcium chloride absorbs water. In the
presence of alkalis these crystals remain unchanged, but with
acids they evolve nitric oxides. There are several compounds which
are capable of interchanging, and correspond with Roussin's salt.
Here we enter into the series of the nitrogen compounds which have
been as yet but little investigated, and will most probably in
time form most instructive material for studying the nature of
that element. These series of compounds are as unlike the usual
saline compounds of inorganic chemistry as are organic
hydrocarbons. There is no necessity to describe these series in
detail, because their connection with other compounds is not yet
clear, and they have not yet any application.
If chlorine and sodium are representatives of independent groups of
elements, the same may also be said of iron. Its nearest analogues show,
besides a similarity in character, a likeness as regards physical
properties and a proximity in atomic weight. Iron occupies a medium
position amongst its nearest analogues, both with respect to properties
and faculty of forming saline oxides, and also as regards atomic weight.
On the one hand, cobalt, 58, and nickel, 59, approach iron, 56; they are
metals of a more basic character, they do not form stable acids or higher
degrees of oxidation, and are a transition to copper, 63, and zinc, 65.
On the other hand, manganese, 55, and chromium, 52, are the nearest to
iron; they form both basic and acid oxides, and are a transition to the
metals possessing acid properties. In addition to having atomic weights
approximately alike, chromium, manganese, iron, cobalt, nickel, and
copper have also nearly the same specific gravity, so that the atomic
volumes and the molecules of their analogous compounds are also near to
one another (see table at the beginning of this volume). Besides this,
the likeness between the above-mentioned elements is also seen from the
following:
They form suboxides, RO, fairly energetic bases, isomorphous with
magnesia--for instance, the salt RSO_{4},7H_{2}O, akin to
MgSO_{4},7H_{2}O, and FeSO_{4},7H_{2}O, or to sulphates containing less
water; with alkali sulphates all form double salts crystallising with
6H_{2}O; all are capable of forming ammonium salts, &c. The lower oxides,
in the cases of nickel and cobalt, are tolerably stable, are not easily
oxidised (the nickel compound with more difficulty than cobalt, a
transition to copper); with manganese, and especially with chromium, they
are more easily oxidised than with iron and pass into higher oxides. They
also form oxides of the form R_{2}O_{3}, and with nickel, cobalt, and
manganese this oxide is very unstable, and is more easily reduced than
ferric oxide; but, in the case of chromium, it is very stable, and forms
the ordinary kind of salts. It is isomorphous with ferric oxide, forms
alums, is a feeble base, &c. Chromium, manganese, and iron are oxidised
by alkali and oxidising agents, forming salts like Na_{2}SO_{4}; but
cobalt and nickel are difficult to oxidise; their acids are not known
with any certainty, and are, in all probability, still less stable than
the ferrates. Cr, Mn and Fe form compounds R_{2}Cl_{6} which are like
Fe_{2}Cl_{6} in many respects; in Co this faculty is weaker and in Ni it
has almost disappeared. The cyanogen compounds, especially of manganese
and cobalt, are very near akin to the corresponding ferrocyanides. The
oxides of nickel and cobalt are more easily reduced to metal than those
of iron, but those of manganese and chromium are not reduced so easily as
iron, and the metals themselves are not easily obtained in a pure state;
they are capable of forming varieties resembling cast iron. The metals
Cr, Mn, Fe, Co, and Ni have a grey iron colour and are very difficult to
melt, but nickel and cobalt can be melted in the reverberatory furnace
and are more fusible than iron, whilst chromium is more difficult to melt
than platinum (Deville). These metals decompose water, but with greater
difficulty as the atomic weight rises, forming a transition to copper,
which does not decompose water. All the compounds of these metals have
various colours, which are sometimes very bright, especially in the
higher stages of oxidation.
These metals of the iron group are often met with together in nature.
Manganese nearly everywhere accompanies iron, and iron is always an
ingredient in the ores of manganese. Chromium is found principally as
chrome ironstone--that is, a peculiar kind of magnetic oxide in which
Fe_{2}O_{3} is replaced by Cr_{2}O_{3}.
Nickel and cobalt are as inseparable companions as iron and manganese.
The similarity between them even extends to such remote properties as
magnetic qualities. In this series of metals we find those which are the
most magnetic: iron, cobalt, and nickel. There is even a magnetic oxide
among the chromium compounds, such being unknown in the other series.
Nickel easily becomes passive in strong nitric acid. It absorbs hydrogen
in just the same way as iron. In short, in the series Cr, Mn, Fe, Co, and
Ni, there are many points in common although there are many differences,
as will be seen still more clearly on becoming acquainted with cobalt and
nickel.
In nature _cobalt_ is principally found in combination with arsenic
and sulphur. _Cobalt arsenide_, or _cobalt speiss_, CoAs_{2}, is found in
brilliant crystals of the regular system, principally in Saxony. _Cobalt
glance_, CoAs_{2}CoS_{2}, resembles it very much, and also belongs to the
regular system; it is found in Sweden, Norway, and the Caucasus.
_Kupfernickel_ is a nickel ore in combination with arsenic, but of a
different composition from cobalt arsenide, having the formula NiAs; it
is found in Bohemia and Saxony. It has a copper-red colour and is rarely
crystalline; it is so called because the miners of Saxony first mistook
it for an ore of copper (_Kupfer_), but were unable to extract copper
from it. _Nickel glance_, NiS_{2},NiAs_{2}, corresponding with cobalt
glance, is also known. Nickel accompanies the ores of cobalt and cobalt
those of nickel, so that both metals are found together. The ores of
cobalt are worked in the Caucasus in the Government of Elizavetopolsk.
Nickel ores containing aqueous hydrated nickel silicate are found in the
Ural (Revdansk). Large quantities of a similar ore are exported into
Europe from New Caledonia. Both ores contain about 12 per cent. Ni.
_Garnierite_, (RO)_{5}(SiO_{2})_{4}1-1/2H_{2}O, where R = Ni and Mg,
predominates in the New Caledonian ore. Large deposits of nickel have
been discovered in Canada, where the ore (as nickelous pyrites) is free
from arsenic. Cobalt is principally worked up into cobalt compounds, but
nickel is generally reduced to the metallic state, in which it is now
often used for alloys--for instance, for coinage in many European States,
and for plating other metals, because it does not oxidise. Cobalt
arsenide and cobalt glance are principally used for the preparation of
cobalt compounds; they are first sorted by discarding the rocky matter,
and then roasted. During this process most of the sulphur and arsenic
disappears; the arsenious anhydride volatilises with the sulphurous
anhydride and the metal also oxidises.[32] It is a simple matter to
obtain nickel and cobalt from their oxides. In order to obtain the
latter, solutions of their salts are treated with sodium carbonate and
the precipitated carbonates are heated; the suboxides are thus obtained,
and these latter are reduced in a stream of hydrogen, or even by heating
with ammonium chloride. They easily oxidise when in the state of powder.
When the chlorides of nickel and cobalt are heated in a stream of
hydrogen, the metal is deposited in brilliant scales. _Nickel is always
much more easily and quickly reduced than cobalt._ Nickel melts more
easily than cobalt, and this even furnishes a means of testing the
heating powers of a reverberatory furnace. Cobalt fuses at a temperature
only a little lower than that at which iron does. In general, cobalt is
nearer to iron than nickel, nickel being nearer to copper.[32 bis] Both
nickel and cobalt have magnetic properties like iron, but Co is less
magnetic than Fe, and Ni still less so. The specific gravity of nickel
reduced by hydrogen is 9·1 and that of cobalt 8·9. Fused cobalt has a
specific gravity of 8·5, the density of ordinary nickel being almost the
same. Nickel has a greyish silvery-white colour; it is brilliant and very
ductile, so that the finest wire may be easily drawn from it. This wire
has a resistance to tension equal to iron wire. The beautiful colour of
nickel, and the high polish which it is capable of receiving and
retaining, as it does not oxidise, render it a useful metal for many
purposes, and in many ways it resembles silver.[32 tri] It is now very
common to cover other metals with a layer of nickel (nickel plating).
This is done by a process of electro-plating, using a solution of a
nickel salt. The colour of cobalt is dark and redder; it is also ductile,
and has a greater tensile resistance than iron. Dilute acids act very
slowly on nickel and cobalt; nitric acid may be considered as the best
solvent for them. The solutions in every case contain salts corresponding
with the ferrous salts--that is, the _salts_ CoX_{2}, NiX_{2},
_correspond with the suboxides_ of these metals. These salts in their
types are similar to the magnesium salts. The salts of nickel when
crystallising with water have a green colour, and form bright green
solutions, but in the anhydrous state they most frequently have a yellow
colour. The salts of cobalt are generally rose-coloured, and generally
blue when in the anhydrous state. Their aqueous solutions are
rose-coloured. Cobaltous chloride is easily soluble in alcohol, and forms
a solution of an intense blue colour.[33]
[32] The residue from the roasting of cobalt ores is called _zafflor_,
and is often met with in commerce. From this the purer compounds
of cobalt may be prepared. The ores of nickel are also first
roasted, and the oxides dissolved in acid, nickelous salts being
then obtained.
The further treatment of cobalt and nickel ores is facilitated if
the arsenic can be almost entirely removed, which may be effected
by roasting the ore a second time with a small addition of nitre
and sodium carbonate; the nitre combines with the arsenic, forming
an arsenious salt, which may be extracted with water. The
remaining mass is dissolved in hydrochloric acid, mixed with a
small quantity of nitric acid. Copper, iron, manganese, nickel,
cobalt, &c., pass into solution. By passing hydrogen sulphide
through the solution, copper, bismuth, lead, and arsenic are
deposited as metallic sulphides; but iron, cobalt, nickel, and
manganese remain in solution. If an alkaline solution of bleaching
powder be then added to the remaining solution, the whole of the
manganese will first be deposited in the form of dioxide, then the
cobalt as hydrated cobaltic oxide, and finally the nickel also. It
is, however, impossible to rely on this method for effecting a
complete separation, the more so since the higher oxides of the
three above-mentioned metals have all a black colour; but, after a
few trials, it will be easy to find how much bleaching powder is
required to precipitate the manganese, and the amount which will
precipitate all the cobalt. The manganese may also be separated
from cobalt by precipitation from a mixture of the solutions of
both metals (in the form of the 'ous' salts) with ammonium
sulphide, and then treating the precipitate with acetic acid or
dilute hydrochloric acid, in which manganese sulphide is easily
soluble and cobalt sulphide almost insoluble. Further particulars
relating to the separation of cobalt from nickel may be found in
treatises on analytical chemistry. In practice it is usual to rely
on the rough method of separation founded on the fact that nickel
is more easily reduced and more difficult to oxidise than cobalt.
The New Caledonian ore is smelted with CaSO_{4} and CaCO_{3} on
coke, and a metallic regulus is obtained containing all the Ni,
Fe, and S. This is roasted with SiO_{2}, which converts all the
iron into slag, whilst the Ni remains combined with the S; this
residue on further roasting gives NiO, which is reduced by the
carbon to metallic Ni. The Canadian ore (a pyrites containing 11
p.c. Ni) is frequently treated in America (after a preliminary
dressing) by smelting it with Na_{2}SO_{4} and charcoal; the
resultant fusible Na_{2}S then dissolves the CuS and FeS_{2},
while the NiS is obtained in a bottom layer (Bartlett and
Thomson's process) from which Ni is obtained in the manner
described above.
For manufacturing purposes somewhat impure cobalt compounds are
frequently used, which are converted into _smalt_. This is glass
containing a certain amount of cobalt oxide; the glass acquires a
bright blue colour from this addition, so that when powdered it
may be used as a blue pigment; it is also unaltered at high
temperatures, so that it used to take the place now occupied by
Prussian blue, ultramarine, &c. At present smalt is almost
exclusively used for colouring glass and china. To prepare smalt,
ordinary impure cobalt ore (zaffre) is fused in a crucible with
quartz and potassium carbonate. A fused mass of cobalt glass is
thus formed, containing silica, cobalt oxide, and potassium oxide,
and a metallic mass remains at the bottom of the crucible,
containing almost all the other metals, arsenic, nickel, copper,
silver, &c. This metallic mass is called _speiss_, and is used as
nickel ore for the extraction of nickel. Smalt usually contains 70
p.c. of silica, 20 p.c. of potash and soda, and about 5 to 6 p.c.
of cobaltous oxide; the remainder consisting of other metallic
oxides.
[32 bis] All we know respecting the relations of Co and Ni to Fe and Cu
confirms the fact that Co is more closely related to Fe and Ni to
Cu; and as the atomic weight of Fe = 56 and of Cu = 63, then
according to the principles of the periodic system it would be
expected that the atomic weight of Co would be about 59-60, whilst
that of Ni should be greater than that of Co but less than that of
Cu, _i.e._ about 50·5-60·5. However, as yet the majority of the
determinations of the atomic weights of Co and Ni give a different
result and show that a lower atomic weight is obtained for Ni than
for Co. Thus K. Winkler (1894) obtained (employing metals
deposited electrolytically and determining the amount of iodine
which combined with them) Ni = 58·72 and Co = 59·37 (if H = 1 and
I = 126·53). In my opinion this should not be regarded as proving
that the principles of the periodic system cannot be applied in
this instance, nor as a reason for altering the position of these
elements in the system (_i.e._ by placing Ni after Fe, and Co next
to Cu), because in the first place the figures given by different
chemists (for instance, Zimmermann, Krüss, and others) are
somewhat divergent, and in the second place the majority of the
latest modes of determining the atomic weights of Co and Ni aim at
finding what weights of these metals react with known weights of
other elements without taking into account the faculty they have
of absorbing hydrogen; since this faculty is more developed in Ni
than in Co the hydrogen (occluded in Ni) should lower the atomic
weight of Ni more than that of Co. On the whole, the question of
the atomic weights of Co and Ni cannot yet be considered as
decided, notwithstanding the numerous researches which have been
made; still there can be no doubt that the atomic weights of these
two metals are very nearly equal, and greater than that of Fe, but
less than that of Cu. This question is of great interest, not only
for completing our knowledge of these metals, but also for
perfecting our knowledge of the periodic system of the elements.
[32 tri] For instance, the alkalis may be fused in nickel vessels as
well as in silver, because they have no action upon either metal.
Nickel, like silver, is not acted upon by dilute acids. Only
nitric acid dissolves both metals well. Nickel is harder, and
fuses at a higher temperature than silver. For castings, a small
quantity of magnesium (0·001 part by weight) is added to nickel to
render it more homogeneous (just as aluminium is added to steel).
Nickel forms many valuable alloys. Steel containing 3 p.c. Ni is
particularly valuable, its limit of elasticity is higher and its
hardness is greater; it is used for armour plate and other large
pieces. The alloys of nickel, especially with copper and zinc
(melchior, _see_ later), aluminium and silver, although used in
certain cases, are now replaced by nickel-plated or
nickel-deposited goods (deposited by electricity from a solution
of the ammonium salts).
[33] The change of colour is dependent in all probability on the
combination with water, or according to others on polymeric
transformation. It enables a solution of cobalt chloride to be
used as sympathetic ink. If something be written with cobalt
chloride on white paper, it will be invisible on account of the
feeble colour of the solution, and when dry nothing can be
distinguished. If, however, the paper be heated before the fire,
the rose-coloured salt will be changed into a less hydrous blue
salt, and the writing will become quite visible, but fade away
when cool.
The change of colour which takes place in solutions of CoCl_{2}
under the influence not only of solution in water or alcohol, but
also of a change of temperature, is a characteristic of all the
halogen salts of cobalt. Crystalline iodide of cobalt,
CoI_{2}6H_{2}O, gives a dark red solution between -22° and +20°;
above +20° the solution turns brown and passes from olive to
green, from +35° to 320° the solution remains green. According to
Étard the change of colour is due to the fact that at first the
solution contains the hydrate CoI_{2}H_{2}O, and that above 35° it
contains CoI_{2}4H_{2}O. These hydrates can be crystallised from
the solutions; the former at ordinary temperature and the latter
on heating the solution. The intermediate olive colour of the
solutions corresponds to the incipient decomposition of the
hexahydrated salt and its passage into CoI_{2}4H_{2}O. A solution
of the hexahydrated chloride of cobalt, CoCl_{2}6H_{2}O, is
rose-coloured between -22° and +25°; but the colour changes
starting from +25°, and passes through all the tints between red
and blue right up to 50°; a true blue solution is only obtained at
55° and remains up to 300°. This true blue solution contains
another hydrate, CoCl_{2}2H_{2}O.
The dependence between the solubility of the iodide and chloride
of cobalt and the temperature is expressed by two almost straight
lines corresponding to the hexa- and di-hydrates; the passage of
the one into the other hydrate being expressed by a curve. The
same character of phenomena is seen also in the variation of the
vapour tension of solutions of chloride of cobalt with the
temperature. We have repeatedly seen that aqueous solutions (for
instance, Chapter XXII., Note 23 for Fe_{2}Cl_{6}) deposit
different crystallo-hydrates at different temperatures, and that
the amount of water in the hydrate decreases as the temperature
_t_ rises, so that it is not surprising that CoCl_{2}2H_{2}O (or
according to Potilitzin CoCl_{2}H_{2}O) should separate out above
55° and CoCl_{2}6H_{2}O at 25° and below. Nor is it exceptional
that the colour of a salt varies according as it contains
different amounts of H_{2}O. But in this instance it is
characteristic that the change of colour takes place in solution
in the presence of an excess of water. This apparently shows that
the actual solution may contain either CoCl_{2}6H_{2}O or
CoCl_{2}2H_{2}O. And as we know that a solution may contain both
metaphosphoric PHO_{3} and orthophosphoric acid H_{3}PO_{4} =
HPO_{3} + H_{2}O, as well as certain other anhydrides, the
question of the state of substances in solutions becomes still
more complicated.
Nickel sulphate crystallises from neutral solutions at a
temperature of from 15° to 20° in _rhombic_ crystals containing
7H_{2}O. Its form approaches very closely to that of the salts of
zinc and magnesium. The planes of a vertical prism for magnesium
salts are inclined at an angle of 90° 30´, for zinc salts at an
angle of 91° 7´, and for nickel salts at an angle of 91° 10´. Such
is also the form of the zinc and magnesium selenates and
chromates. Cobalt sulphate containing 7 molecules of water is
deposited in crystals of the _monoclinic_ system, like the
corresponding salts of iron and manganese. The angle of a vertical
prism for the iron salt = 82° 20´, for cobalt = 82° 22´, and the
inclination of the horizontal pinacoid to the vertical prism for
the iron salt = 99° 2´, and for the cobalt salt 99° 36´. All the
isomorphous mixtures of the salts of magnesium, iron, cobalt,
nickel and manganese have the same form if they contain 7 mol.
H_{2}O and iron or cobalt predominate, whilst if there is a
preponderance of magnesium, zinc, or nickel, the crystals have a
rhombic form like magnesium sulphate. Hence these sulphates are
_dimorphous_, but for some the one form is more stable and for
others the other. Brooke, Moss, Mitscherlich, Rammelsberg, and
Marignac have explained these relations. Brooke and Mitscherlich
also supposed that NiSO_{4},7H_{2}O is not only capable of
assuming these forms, but also that of the _tetragonal_ system,
because it is deposited in this form from acid, and especially
from slightly-heated solutions (30° to 40°). But Marignac
demonstrated that the tetragonal crystals do not contain 7, but 6,
molecules of water, NiSO_{4},6H_{2}O. He also observed that a
solution evaporated at 50° to 70° deposits monoclinic crystals,
but of a different form from ferrous sulphate,
FeSO_{4},7H_{2}O--namely, the angle of the prism is 71° 52´, that
of the pinacoid 95° 6´. This salt appears to be the same with 6
molecules of water as the tetragonal. Marignac also obtained
magnesium and zinc salts with 6 molecules of water by evaporating
their solutions at a higher temperature, and these salts were
found to be isomorphous with the monoclinic nickel salt. In
addition to this it must be observed that the rhombic crystals of
nickel sulphate with 7H_{2}O become turbid under the influence of
heat and light, lose water, and change into the tetragonal salt.
The monoclinic crystals in time also become turbid, and change
their structure, so that the tetragonal form of this salt is the
most stable. Let us also add that nickel sulphate in all its
shapes forms very beautiful emerald green crystals, which, when
heated to 230°, assume a dirty greenish-yellow hue and then
contain one molecule of water.
Klobb (1891) and Langlot and Lenoir obtained anhydrous CoSO_{4}
and NiSO_{4} by igniting the hydrated salt with (NH_{4})_{2}SO_{4}
until the ammonium salt had completely volatilised and decomposed.
We may add that when equivalent aqueous solutions of NiX_{2}
(green) and CoX_{2} (red) are mixed together they give an almost
colourless (grey) solution, in which the green and red colour of
the component parts disappears owing to the combination of the
complementary colours.
A double salt NiKF_{3} is obtained by heating NiCl_{2} with KFHF
in a platinum crucible; KCoF_{3} is formed in a similar manner.
The nickel salt occurs in fine green plates, easily soluble in
water but scarcely soluble in ethyl and methyl alcohol. They
decompose into green oxide of nickel and potassium fluoride when
heated in a current of air. The analogous salt of cobalt
crystallises in crimson flakes.
If instead of potassium fluoride, CoCl_{2} or NiCl_{2} be fused
with ammonium fluoride, they also form double salts with the
latter. This gives the possibility of obtaining anhydrous
fluorides NiF_{2} and CoF_{2}. Crystalline fluoride of nickel,
obtained by heating the amorphous powder formed by decomposing the
double ammonium salt in a stream of hydrofluoric acid, occurs in
beautiful green prisms, sp. gr. 4·63, which are insoluble in
water, alcohol, and ether; sulphuric, hydrochloric, and nitric
acids also have no action upon them, even when heated; NiF_{2} is
decomposed by steam, with the formation of black oxide, which
retains the crystalline structure of the salt. Fluoride of cobalt,
obtained as a rose-coloured powder by decomposing the double
ammonium salt with the aid of heat in a stream of hydrofluoric
acid, fuses into a ruby-coloured mass which bears distinct signs
of a crystalline structure; sp. gr. 4·43. The molten salt only
volatilises at about 1400°, which forms a clear distinction
between CoF_{2} and the volatile NiF_{2}. Hydrochloric, sulphuric,
and nitric acids act upon CoF_{2} even in the cold, although
slowly, while when heated the reaction proceeds rapidly (Poulenc,
1892).
If a solution of potassium hydroxide be added to a solution of a cobalt
salt, a blue precipitate of the basic salt will be formed. If a solution
of a cobalt salt be heated almost to the boiling-point, and the solution
be then mixed with a boiling solution of an alkali hydroxide, a _pink
precipitate of cobaltous hydroxide_, CoH_{2}O_{2}, will be formed. If air
be not completely excluded during the precipitation by boiling, the
precipitate will also contain brown cobaltic hydroxide formed by the
further oxidation of the cobaltous oxide.[34] Under similar circumstances
nickel salts form _a green precipitate of nickelous hydroxide_, the
formation of which is not hindered by the presence of ammonium salts, but
in that case only requires more alkali to completely separate the nickel.
The nickelous oxide obtained by heating the hydroxide, or from the
carbonate or nitrate, is a grey powder, easily soluble in acids and
easily reduced, but the same substance may be obtained in the crystalline
form as an ordinary product from the ores; it crystallises in regular
octahedra, with a metallic lustre, and is of a grey colour. In this state
the nickelous oxide almost resists the action of acids.[34 bis]
[34] Hydrated suboxide of cobalt (de Schulten, 1889) is obtained in the
following manner. A solution of 10 grams of CoCl_{2}6H_{2}O in 60
c.c. of water is heated in a flask with 250 grams of caustic
potash and a stream of coal gas is passed through the solution.
When heated the hydrate of the suboxide of cobalt which separates
out, dissolves in the caustic potash and forms a dark blue
solution. This solution is allowed to stand for 24 hours in an
atmosphere of coal gas (in order to prevent oxidation). The
crystalline mass which separates out has a composition Co(OH)_{2},
and to the naked eye appears as a violet powder, which is seen to
be crystalline under the microscope. The specific gravity of this
hydrate is 3·597 at 15°. It does not undergo change in the air;
warm acetic acid dissolves it, but it is insoluble in warm and
cold solutions of ammonia and sal-ammoniac.
[34 bis] The following reaction may be added to those of the cobaltous
and nickelous salts: potassium cyanide forms a precipitate with
cobalt salts which is soluble in an excess of the reagent and
forms a green solution. On heating this and adding a certain
quantity of acid, a double _cobalt cyanide_ is formed which
corresponds with potassium ferricyanide. Its formation is
accompanied with the evolution of hydrogen, and is founded upon
the property which cobalt has of oxidising in an alkaline
solution, the development of which has been observed in such a
considerable measure in the cobaltamine salts. The process which
goes on here may be expressed by the following equation;
CoC_{2}N_{2} + 4KCN first forms CoK_{4}C_{6}N_{6}, which salt with
water, H_{2}O, forms potassium hydroxide, KHO, hydrogen, H, and
the salt, K_{3}CoC_{6}N_{6}. Here naturally the presence of the
acid is indispensable in consequence of its being required to
combine with the alkali. From aqueous solutions this salt
crystallises in transparent, hexagonal prisms of a yellow colour,
easily soluble in water. The reactions of double decomposition,
and even the formation of the corresponding acid, are here
completely the same as in the case of the ferricyanide. If a
nickelous salt be treated in precisely the same manner as that
just described for a salt of cobalt, decomposition will occur.
It is interesting to note _the relation_ of the cobaltous and
nickelous hydroxides _to ammonia_; aqueous ammonia dissolves the
precipitate of cobaltous and nickelous hydroxide. The blue ammoniacal
solution of nickel resembles the same solution of cupric oxide, but has a
somewhat reddish tint. It is characterised by the fact that it dissolves
silk in the same way as the ammoniacal cupric oxide dissolves cellulose.
Ammonia likewise dissolves the precipitate of cobaltous hydroxide,
forming a brownish liquid, which becomes darker in air and finally
assumes a bright red hue, absorbing oxygen. The admixture of ammonium
chloride prevents the precipitation of cobalt salts by ammonia; when the
ammonia is added, a brown solution is obtained from which, as in the case
of the preceding solution, potassium hydroxide does not separate the
cobaltous oxide. Peculiar compounds are produced in this solution; they
are comparatively stable, containing ammonia and an excess of oxygen;
they bear the name cobaltoamine and cobaltiamine salts. They have been
principally investigated by Genth, Frémy, Jörgenson and others. Genth
found that when a cobalt salt, mixed with an excess of ammonium chloride,
is treated with ammonia and exposed to the air, after a certain lapse of
time, on adding hydrochloric acid and boiling, a red powder is
precipitated and the remaining solution contains an orange salt. The
study of these compounds led to the discovery of a whole series of
similar salts, some of which correspond with particular higher degrees of
oxidation of cobalt, which are described later.[35] Nickel does not
possess this property of absorbing the oxygen of the air when in an
ammoniacal solution. In order to understand this distinction, and in
general the relation of nickel, it is important to observe that cobalt
more easily forms a higher degree of oxidation--namely, _sesquioxide of
cobalt_, _cobaltic oxide_, Co_{2}O_{3}--than nickel, especially in the
presence of hypochlorous acid. If a solution of a cobalt salt be mixed
with barium carbonate and an excess of hypochlorous acid be added, or
chlorine gas be passed through it, then at the ordinary temperature on
shaking, the whole of the cobalt will be separated in the form of black
cobaltic oxide: 2CoSO_{4} + ClHO + 2BaCO_{3} = Co_{2}O_{3} + 2BaSO_{4} +
HCl + 2CO_{2}. Under these circumstances nickelous oxide does not
immediately form black sesquioxide, but after a considerable space of
time it also separates in the form of sesquioxide, Ni_{2}O_{3}, but
always later than cobalt. This is due to the relative difficulty of
further oxidation of the nickelous oxide. It is, however, possible to
oxidise it; if, for instance, the hydroxide NiH_{2}O_{2} be shaken in
water and chlorine gas be passed through it, then nickel chloride will be
formed, which is soluble in water, and insoluble nickelic oxide in the
form of a black precipitate: 3NiH_{2}O_{2} + Cl_{2} = NiCl_{2} +
Ni_{2}O_{3},3H_{2}O. Nickelic oxide may also be obtained by adding sodium
hypochlorite mixed with alkali to a solution of a nickel salt. Nickelic
and cobaltic hydrates are black. Nickelic oxide evolves oxygen with all
acids, and in consequence of this it is not separated as a precipitate in
the presence of acids; thus it evolves chlorine with hydrochloric acid,
exactly like manganese dioxide. When nickelic oxide is dissolved in
aqueous ammonia it liberates nitrogen, and an ammoniacal solution of
nickelous oxide is formed. When heated, nickelic oxide loses oxygen,
forming nickelous oxide. Cobaltic oxide, Co_{2}O_{3}, exhibits more
stability than nickelic oxide, and shows feeble basic properties; thus it
is dissolved in acetic acid without the evolution of oxygen.[35 bis] But
ordinary acids, especially on heating, evolve oxygen, forming a solution
of a cobaltous salt. The presence of a cobaltic salt in a solution of a
cobaltous salt may be detected by the brown colour of the solution and
the black precipitate formed by the addition of alkali, and also from the
fact that such solutions evolve chlorine when heated with hydrochloric
acid. Cobaltic oxide may not only be prepared by the above-mentioned
methods, but also by heating cobalt nitrate, after which a steel-coloured
mass remains which retains traces of nitric acid, but when heated further
to incandescence evolves oxygen, leaving a compound of cobaltic and
cobaltous oxides, similar to magnetic ironstone. Cobalt (but not nickel)
undoubtedly forms besides Co_{2}O_{3} a _dioxide_ CoO_{2}. This is
obtained[36] when the cobaltous oxide is oxidised by iodine or peroxide
of barium.[37]
[35] The cobalt salts may be divided into at least the following
classes, which repeat themselves for Cr, Ir, Rh (we shall not stop
to consider the latter, particularly as they closely resemble the
cobalt salts):--
(_a_) _Ammonium cobalt salts_, which are simply direct compounds
of the cobaltous salts CoX_{2} with ammonia, similar to various
other compounds of the salts of silver, copper, and even calcium
and magnesium, with ammonia. They are easily crystallised from an
ammoniacal solution, and have a pink colour. Thus, for instance,
when cobaltous chloride in solution is mixed with sufficient
ammonia to redissolve the precipitate first formed, octahedral
crystals are deposited which have a composition
CoCl_{2},H_{2}O,6NH_{3}. These salts are nothing else but
combinations with ammonia of crystallisation--if it may be so
termed--likening them in this way to combinations with water of
crystallisation. This similarity is evident both from their
composition and from their capability of giving off ammonia at
various temperatures. The most important point to observe is that
all these salts contain 6 molecules of ammonia to 1 atom of
cobalt, and this ammonia is held in fairly stable connection.
Water decomposes these salts. (Nickel behaves similarly without
forming other compounds corresponding to the true cobaltic.)
(_b_) The solutions of the above-mentioned salts are rendered
turbid by the action of the air; they absorb oxygen and become
covered with a crust of _oxycobaltamine salts_. The latter are
sparingly soluble in aqueous ammonia, have a brown colour, and are
characterised by the fact that with warm water _they evolve
oxygen_, forming salts of the following category: The nitrate may
be taken as an example of this kind of salt; its composition is
CoN_{2}O_{7},5NH_{3},H_{2}O. It differs from cobaltous nitrate,
Co(NO_{3})_{2}, in containing an extra atom of oxygen--that is, it
corresponds with cobalt dioxide, CoO_{2}, in the same way that the
first salts correspond with cobaltous oxide; they contain 5, and
not 6, molecules of ammonia, as if NH_{3} had been replaced by O,
but we shall afterwards meet compounds containing either 5NH_{3}
or 6NH_{3} to each atom of cobalt.
(_c_) _The luteocobaltic salts_ are thus called because they have
a yellow (luteus) colour. They are obtained from the salts of the
first kind by submitting them in dilute solution to the action of
the air; in this case salts of the second kind are not formed,
because they are decomposed by an excess of water, with the
evolution of oxygen and the formation of luteocobaltic salts. By
the action of ammonia the salts of the fifth kind (roseocobaltic)
are also converted into luteocobaltic salts. These last-named
salts generally crystallise readily, and have a yellow colour;
they are comparatively much more stable than the preceding ones,
and even for a certain time resist the action of boiling water.
Boiling aqueous potash liberates ammonia and precipitates hydrated
cobaltic oxide, Co_{2}O_{3},3H_{2}O, from them. This shows that
the luteocobaltic salts correspond with cobaltic oxide,
Co_{2}O_{3}, and those of the second kind with the dioxide. When a
solution of luteocobaltic sulphate,
Co_{2}(SO_{4})_{3},12NH_{3},4H_{2}O, is treated with baryta,
barium sulphate is precipitated, and the solution contains
luteocobaltic hydroxide, Co(OH)_{3},6NH_{3}, which is soluble in
water, is powerfully alkaline, absorbs the oxygen of the air, and
when heated is decomposed with the evolution of ammonia. This
compound therefore corresponds to a solution of cobaltic hydroxide
in ammonia. The luteocobaltic salts contain 2 atoms of cobalt and
12 molecules of ammonia--that is, 6NH_{3} to each atom of cobalt,
like the salts of the first kind. The CoX_{2} salts have a
metallic taste, whilst those of luteocobalt and others have a
purely saline taste, like the salts of the alkali metals. In the
luteo-salts all the X's react (are ionised, as some chemists say)
as in ordinary salts--for instance, all the Cl_{2} is precipitated
by a solution of AgNO_{3}; all the (SO_{4})_{3} gives a
precipitate with BaX_{2}, &c. The double salt formed with PtCl_{4}
is composed in the same manner as the potassium salt,
K_{2}PtCl_{4} = 2KCl + PtCl_{4}, that is, contains
(CoCl_{3},6NH_{3})_{2},3PtCl_{4}, or the amount of chlorine in the
PtCl_{4} is double that in the alkaline salt. In the rosepentamine
(_e_), and rosetetramine (_f_), salts, also all the X's react or
are ionised, but in the (_g_) and (_h_) salts only a portion of
the X's react, and they are equal to the (_e_) and (_f_) salts
minus water; this means that although the water dissolves them it
is not combined with them, as PHO_{3} differs from PH_{3}O_{3};
phenomena of this class correspond exactly to what has been
already (Chapter XXI., Note 7) mentioned respecting the green and
violet salts of oxide of chromium.
(_d_) _The fuscocobaltic salts._ An ammoniacal solution of cobalt
salts acquires a brown colour in the air, due to the formation of
these salts. They are also produced by the decomposition of salts
of the second kind; they crystallise badly, and are separated from
their solutions by addition of alcohol or an excess of ammonia.
When boiled they give up the ammonia and cobaltic oxide which they
contain. Hydrochloric and nitric acids give a yellow precipitate
with these salts, which turns red when boiled, forming salts of
the next category. The following is an example of the composition
of two of the fuscocobaltic salts,
Co_{2}O(SO_{4})_{2},8NH_{3},4H_{2}O and
Co_{2}OCl_{4},8NH_{3},3H_{2}O. It is evident that the
fuscocobaltic salts are ammoniacal compounds of basic cobaltic
salts. The normal cobaltic sulphate ought to have the composition
Co_{2}(SO_{4})_{3} = Co_{2}O_{3},3SO_{3}; the simplest basic salts
will be Co_{2}O(SO_{4})_{2} = Co_{2}O_{3})2SO_{3}, and
Co_{2}O_{2}(SO_{4}) = Co_{2}O_{3},SO_{3}. The fuscocobaltic salts
correspond with the first type of basic salts. They are changed
(in concentrated solutions) into oxycobaltamine salts by
absorption of one atom of oxygen, Co_{2}O_{2}(SO_{4})_{2}. The
whole process of oxidation will be as follows: first of all
Co_{2}X_{4}, a cobaltous salt, is in the solution (X a univalent
haloid, 2 molecules of the salt being taken), then Co_{2}OX_{4},
the basic cobaltic salt (4th series), then Co_{2}O_{2}X_{4}, the
salt of the dioxide (2nd series). The series of basic salts with
an acid, 2HX, forms water and a normal salt, Co_{2}X_{6} (in 3, 5,
6 series). These salts are combined with various amounts of water
and ammonia. Under many conditions the salts of fuscocobalt are
easily transformed into salts of the next series. The salts of the
series that has just been described contain 4 molecules of ammonia
to 1 atom of cobalt.
(_e_) _The roseocobaltic_ (or rosepentamine),
CoX_{2}H_{2}O,5NH_{3}, _salts_, like the luteocobaltic, correspond
with the normal cobaltic salts, but contain less ammonia, and an
extra molecule of water. Thus the sulphate is obtained from
cobaltous sulphate dissolved in ammonia and left exposed to the
air until transformed into a brown solution of the fuscocobaltic
salt; when this is treated with sulphuric acid a crystalline
powder of the roseocobaltic salt,
Co_{2}(SO_{4})_{3},10NH_{3},5H_{2}O, separates. The formation of
this salt is easily understood: cobaltous sulphate in the presence
of ammonia absorbs oxygen, and the solution of the fuscocobaltic
salt will therefore contain, like cobaltous sulphate, one part of
sulphuric acid to every part of cobalt, so that the whole process
of formation may be expressed by the equation: 10NH_{3} +
2CoSO_{4} + H_{2}SO_{4} + 4H_{2}O + O =
Co_{2}(SO_{4})_{3},10NH_{3},5H_{2}O. This salt forms tetragonal
crystals of a red colour, slightly soluble in cold, but readily
soluble in warm water. When the sulphate is treated with baryta,
roseocobaltic hydroxide is formed in the solution, which absorbs
the carbonic anhydride of the air. It is obtained from the next
series by the action of alkalis.
(_f_) The _rosetetramine cobaltic salts_ CoCl_{2},2H_{2}O,4NH_{3}
were obtained by Jörgenson, and belong to the type of the
luteo-salts, only with the substitution of 2NH_{3} for H_{2}O.
Like the luteo- and roseo-salts they give double salts with
PtCl_{4}, similar to the alkaline double salts, for instance
(Co_{2}H_{2}O,4NH_{3})2(SO_{4})_{2}Cl_{2}PtCl_{4}. They are darker
in colour than the preceding, but also crystallise well. They are
formed by dissolving CoCO_{3} in sulphuric acid (of a given
strength), and after NH_{3} and carbonate of ammonium have been
added, air is passed through the solution (for oxidation) until
the latter turns red. It is then evaporated with lumps of
carbonate of ammonium, filtered from the precipitate and
crystallised. A salt of the composition
Co_{2}(CO_{3})_{2}(SO_{4}),(2H_{2}O,4NH_{3})_{2} is thus obtained,
from which the other salts may be easily prepared.
(_g_) The _purpureocobaltic salts_, CoX_{3},5NH_{3}, are also
products of the direct oxidation of ammoniacal solutions of cobalt
salts. They are easily obtained by heating the roseocobaltic and
luteo-salts with strong acids. They are to all effects the same as
the roseocobaltic salts, only anhydrous. Thus, for instance, the
purpureocobaltic chloride, Co_{2}Cl_{6},10NH_{3}, or
CoCl_{3},5NH_{3}, is obtained by boiling the oxycobaltamine salts
with ammonia. There is the same distinction between these salts
and the preceding ones as between the various compounds of
cobaltous chloride with water. In the purpureocobaltic only X_{2}
out of the X_{3} react (are ionised). To the rosetetramine salts
(_f_) there correspond the _purpureotetramine_ salts,
CoX_{3}H_{2}O,4NH_{3}. The corresponding chromium
purpureopentamine salt, CrCl_{3},5NH_{3} is obtained with
particular ease (Christensen, 1893). Dry anhydrous chromium
chloride is treated with anhydrous liquid ammonia in a freezing
mixture composed of liquid CO_{2} and chlorine, and after some
time the mixture is taken out of the freezing mixture, so that the
excess of NH_{3} boils away; the violet crystals then immediately
acquire the red colour of the salt, CrCl_{3},5NH_{3}, which is
formed. The product is washed with water (to extract the
luteo-salt, CrCl_{3},6NH_{3}), which does not dissolve the salt,
and it is then recrystallised from a hot solution of hydrochloric
acid.
(_h_) The _praseocobaltic salts_, CoX_{3},4NH_{3}, are green, and
form, with respect to the rosetetramine salts (_f_), the products
of ultimate dehydration (for example, like metaphosphoric acid
with respect to orthophosphoric acid, but in dissolving in water
they give neither rosetetramine nor tetramine salts. (In my
opinion one should expect salts with a still smaller amount of
NH_{3}, of the blue colour proper to the low hydrated compounds of
cobalt; the green colour of the prazeo-salts already forms a step
towards the blue.) Jörgenson obtained salts for ethylene-diamine,
N_{2}H_{4}C_{2}H_{4} which replaces 2NH_{3}. After being kept a
long time in aqueous solution they give rosetetramine salts, just
as metaphosphoric acid gives orthophosphoric acid, while the
rosetetramine salts are converted into prazeo-salts by Ag_{2}O and
NaHO. Here only one X is ionised out of the X_{3}. There are also
basic salts of the same type; but the best known is the chromium
salt called the rhodozochromic salt,
Cr_{2}(OH)_{3}Cl_{3},6NH_{3},2H_{2}O, which is formed by the
prolonged action of water upon the corresponding roseo-salt.
The cobaltamine compounds differ essentially but little from the
ammoniacal compounds of other metals. The only difference is that
here the cobaltic oxide is obtained from the cobaltous oxide in
the presence of ammonia. In any case it is a simpler question than
that of the double cyanides. Those forces in virtue of which such
a considerable number of ammonia molecules are united with a
molecule of a cobalt compound, appertain naturally to the series
of those slightly investigated forces which exist even in the
highest degrees of combination of the majority of elements. They
are the same forces which lead to the formation of compounds
containing water of crystallisation, double salts, isomorphous
mixtures and complex acids (Chapter XXI., Note 8 bis). The
simplest conception, according to my opinion, of cobalt compounds
(much more so than by assuming special complex radicles, with
Schiff, Weltzien, Claus, and others), may be formed by comparing
them with other ammoniacal products. Ammonia, like water, combines
in various proportions with a multitude of molecules. Silver
chloride and calcium chloride, just like cobalt chloride, absorb
ammonia, forming compounds which are sometimes slightly stable,
and easily dissociated, sometimes more stable, in exactly the same
way as water combines with certain substances, forming fairly
stable compounds called hydroxides or hydrates, or less stable
compounds which are called compounds with water of
crystallisation. Naturally the difference in the properties in
both cases depends on the properties of those elements which enter
into the composition of the given substance, and on those kinds of
affinity towards which chemists have not as yet turned their
attention. If boron fluoride, silicon fluoride, &c., combine with
hydrofluoric acid, if platinic chloride, and even cadmium
chloride, combine with hydrochloric acid, these compounds may be
regarded as double salts, because acids are salts of hydrogen. But
evidently water and ammonia have the same saline faculty, more
especially as they, like haloid acids, contain hydrogen, and are
both capable of further combination--for instance, ammonia with
hydrochloric acid. Hence it is simpler to compare complex
ammoniacal with double salts, hydrates, and similar compounds, but
_the ammonio-metallic salts_ present a most complete qualitative
and quantitative resemblance to _the hydrated salts of metals_.
The composition of the latter is MX_{_n_}_m_H_{2}O, where M =
metal, X = the haloid, simple or complex, and _n_ and _m_ the
quantities of the haloid and so-called water of crystallisation
respectively. The composition of the ammoniacal salts of metals is
MX_{_n_}_m_NH_{3}. The water of crystallisation is held by the
salt with more or less stability, and some salts even do not
retain it at all; some part with water easily when exposed to the
air, others when heated, and then with difficulty. In the case of
some metals all the salts combine with water, whilst with others
only a few, and the water so combined may then be easily
disengaged. All this applies equally well to the ammoniacal salts,
and therefore the combination of ammonia may be termed _the
ammonia of crystallisation_. Just as the water which is combined
with a salt is held by it with different degrees of force, so it
is with ammonia. In combining with 2NH_{3},PtCl_{2} evolves 31,000
cals.; while CaCl_{2} only evolves 14,000 cals.; and the former
compound parts with its NH_{3} (together with HCl in this case)
with more difficulty, only above 200°, while the latter disengages
ammonia at 180°. ZnCl_{2},2NH_{3} in forming ZnCl_{2},4NH_{3}
evolves only 11,000 cals., and splits up again into its components
at 80°. The amount of combined ammonia is as variable as the
amount of water of crystallisation--for instance, SnI_{4}8NH_{3},
CrCl_{2}8NH_{3}, CrCl_{3}6NH_{3},
CrCl_{3}5NH_{3},PtCl_{2},4NH_{3}, &c. are known. Very often NH_{3}
is replaceable by OH_{2} and conversely. A colourless, anhydrous
cupric salt--for instance, cupric sulphate--when combined with
water forms blue and green salts, and violet when combined with
ammonia. If steam be passed through anhydrous copper sulphate the
salt absorbs water and becomes heated; if ammonia be substituted
for the water the heating becomes much more intense, and the salt
breaks up into a fine violet powder. With water CuSO_{4},5H_{2}O
is formed, and with ammonia CuSO_{4},5NH_{3}, the number of water
and ammonia molecules retained by the salt being the same in each
case, and as a proof of this, and that it is not an isolated
coincidence, the remarkable fact must be borne in mind that water
and ammonia consecutively, molecule for molecule, are capable of
supplanting each other, and forming the compounds
CuSO_{4},5H_{2}O, CuSO_{4},4H_{2}O,NH_{3};
CuSO_{4},3H_{2}O,2NH_{3}; CuSO_{4},2H_{2}O,3NH_{3};
CuSO_{4},H_{2}O,4NH_{3}, and CuSO_{4},5NH_{3}. The last of these
compounds was obtained by Henry Rose, and my experiments have
shown that more ammonia than this cannot be retained. By adding to
a strong solution of cupric sulphate sufficient ammonia to
dissolve the whole of the oxide precipitated, and then adding
alcohol, Berzelius obtained the compound CuSO_{4},H_{2}O,4NH_{3},
&c. The law of substitution also assists in rendering these
phenomena clearer, because a compound of ammonia with water forms
ammonium hydroxide, NH_{4}HO, and therefore these molecules
combining with one another may also interchange, as being of equal
value. In general, those salts form stable ammoniacal compounds
which are capable of forming stable compounds with water of
crystallisation; and as ammonia is capable of combining with
acids, and as some of the salts formed by slightly energetic bases
in their properties more closely resemble acids (that is, salts of
hydrogen) than those salts containing more energetic bases, we
might expect to find more stable and more easily-formed
ammonio-metallic salts with metals and their oxides having weaker
basic properties than with those which form energetic bases. This
explains why the salts of potassium, barium, &c., do not form
ammonio-metallic salts, whilst the salts of silver, copper, zinc,
&c., easily form them, and the salts RX_{3} still more easily and
with greater stability. This consideration also accounts for the
great stability of the ammoniacal compounds of cupric oxide
compared with those of silver oxide, since the former is displaced
by the latter. It also enables us to see clearly the distinction
which exists in the stability of the cobaltamine salts containing
salts corresponding with cobaltous oxide, and those corresponding
with higher oxides of cobalt, for the latter are weaker bases than
cobaltous oxides. _The nature of the forces and quality of the
phenomena occurring during the formation of the most stable
substances, and of such compounds as crystallisable compounds, are
one and the same, although perhaps exhibited in a different
degree._ This, in my opinion, may be best confirmed by examining
the compounds of carbon, because for this element the nature of
the forces acting during the formation of its compounds is well
known. Let us take as an example two unstable compounds of carbon.
Acetic acid, C_{2}H_{4}O_{2} (specific gravity 1·06), with water
forms the hydrate, C_{2}H_{4}O_{2},H_{2}O, denser (1·07) than
either of the components, but unstable and easily decomposed,
generally simply referred to as a solution. Such also is the
crystalline compound of oxalic acid, C_{2}H_{2}O_{4}, with water,
C_{2}H_{2}O_{4},2H_{2}O. Their formation might be predicted as
starting from the hydrocarbon C_{2}H_{6}, in which, as in any
other, the hydrogen may be exchanged for chlorine, the water
residue (hydroxyl), &c. The first substitution product with
hydroxyl, C_{2}H_{5}(HO), is stable; it can be distilled without
alteration, resists a temperature higher than 100°, and then does
not give off water. This is ordinary alcohol. The second,
C_{2}H_{4}(HO)_{2}, can also be distilled without change, but can
be decomposed into water and C_{2}H_{4}O (ethylene oxide or
aldehyde); it boils at about 197°, whilst the first hydrate boils
at 78°, a difference of about 100°. The compound
C_{2}H_{3}(HO)_{3} will be the third product of such substitution;
it ought to boil at about 300°, but does not resist this
temperature--it decomposes into H_{2}O and C_{2}H_{4}O_{2}, where
only one hydroxyl group remains, and the other atom of oxygen is
left in the same condition as in ethylene oxide, C_{2}H_{4}O.
There is a proof of this. Glycol, C_{2}H_{4}(HO)_{2}, boils at
197°, and forms water and ethylene oxide, which boils at 13°
(aldehyde, its isomeride, boils at 21°); therefore the product
disengaged by the splitting up of the hydrate boils at 184° lower
than the hydrate C_{2}H_{4}(HO)_{2}. Thus the hydrate
C_{2}H_{3}(HO)_{3}, which ought to boil at about 300°, splits up
in exactly the same way into water and the product
C_{2}H_{4}O_{2}, which boils at 117°--that is, nearly 183° lower
than the hydrate, C_{2}H_{3}(HO)_{3}. But this hydrate splits up
before distillation. The above-mentioned hydrate of acetic acid is
such a decomposable hydrate--that is to say, what is called a
solution. Still less stability may be expected from the following
hydrates. C_{2}H_{2}(HO)_{4} also splits up into water and a
hydrate (it contains two hydroxyl groups) called glycolic acid,
C_{2}H_{2}O(HO)_{2} = C_{2}H_{4}O_{3}. The next product of
substitution will be C_{2}H(HO)_{5}; it splits up into water,
H_{2}O, and glyoxylic acid, C_{2}H_{4}O_{4} (three hydroxyl
groups). The last hydrate which ought to be obtained from
C_{2}H_{6}, and ought to contain C_{2}(HO)_{6}, is the crystalline
compound of oxalic acid, C_{2}H_{2}O_{4} (two hydroxyl groups),
and water, 2H_{2}O, which has been already mentioned. The hydrate
C_{2}(HO)_{6} = C_{2}H_{2}O_{4},2H_{2}O, ought, according to the
foregoing reasoning, to boil at about 600° (because the hydrate,
C_{2}H_{4}(HO)_{2}, boils at about 200°, and the substitution of 4
hydroxyl groups for 4 atoms of hydrogen will raise the
boiling-point 400°). It does not resist this temperature, but at a
much lower point splits up into water, 2H_{2}O, and the hydrate
C_{2}O_{2}(HO)_{2}, which is also capable of yielding water.
Without going into further discussion of this subject, it may be
observed that the formation of the hydrates, or compounds with
water of crystallisation, of acetic and oxalic acids has thus
received an accurate explanation, illustrating the point we
desired to prove in affirming that compounds with water of
crystallisation are held together by the same forces as those
which act in the formation of other complex substances, and that
the easy displaceability of the water of crystallisation is only a
peculiarity of a local character, and not a radical point of
distinction. All the above-mentioned hydrates, C_{2}X_{6}, or
products of their destruction, are actually obtained by the
oxidation of the first hydrate, C_{2}H_{3}(HO), or common alcohol,
by nitric acid (Sokoloff and others). Hence the forces which
induce salts to combine with _n_H_{2}O or with NH_{3} are
undoubtedly of the same order as the forces which govern the
formation of ordinary 'atomic' and saline compounds. (A great
impediment in the study of the former was caused by the conviction
which reigned in the sixties and seventies, that 'atomic' were
essentially different from 'molecular' compounds like
crystallohydrates, in which it was assumed that there was a
combination of entire molecules, as though without the
participation of the atomic forces.) If the bond between chlorine
and different metals is not equally strong, so also the bond
uniting _n_H_{2}O and _n_NH_{3} is exceeding variable; there is
nothing very surprising in this. And in the fact that the
combination of different amounts of NH_{3} and H_{2}O alters the
capacity of the haloids X of the salts RX_{2} for reaction (for
instance, in the luteo-salts all the X_{3}, while in the purpureo,
only 2 out of the 3, and in the prazeo-salts only 1 of the 3 X's
reacts), we should see in the first place a phenomenon similar to
what we met with in Cr_{2}Cl_{6} (Chapter XXI., Note 7 bis), for
in both instances the essence of the difference lies in the
removal of water; a molecule RCl_{3},6H_{2}O or RCl_{3},6NH_{3}
contains the halogen in a perfectly mobile (ionised) state, while
in the molecule RCl_{3},5H_{2}O or RCl_{3},5NH_{3} a portion of
the halogen has almost lost its faculty for reacting with
AgNO_{3}, just as metalepsical chlorine has lost this faculty
which is fully developed in the chloranhydride. Until the reason
of this difference be clear, we cannot expect that ordinary points
of view and generalisation can give a clear answer. However, we
may assume that here the explanation lies in the nature and kind
of motion of the atoms in the molecules, although as yet it is not
clear how. Nevertheless, I think it well to call attention again
(Chapter I.) to the fact that the combination of water, and hence,
also, of any other element, leads to most diverse consequences;
the water in the gelatinous hydrate of alumina or in the
decahydrated Glauber salt is very mobile, and easily reacts like
water in a free state; but the same water combined with oxide of
calcium, or C_{2}H_{4} (for instance, in C_{2}H_{6}O and in
C_{4}H_{10}O), or with P_{2}O_{5}, has become quite different, and
no longer acts like water in a free state. We see the same
phenomenon in many other cases--for example, the chlorine in
chlorates no longer gives a precipitate of chloride of silver with
AgNO_{3}. Thus, although the instance which is found in the
difference between the roseo- and purpureo-salts deserves to be
fully studied on account of its simplicity, still it is far from
being exceptional, and we cannot expect it to be thoroughly
explained unless a mass of similar instances, which are
exceedingly common among chemical compounds, be conjointly
explained. (Among the researches which add to our knowledge
respecting the complex ammoniacal compounds, I think it
indispensable to call the reader's attention to Prof. Kournakoff's
dissertation 'On complex metallic bases,' 1893.)
Kournakoff (1894) showed that the solubility of the luteo-salt,
CoCl_{3},6NH_{3}, at 0° = 4·30 (per 100 of water), at 20° = 7·7,
that in passing into the roseo-salt, CoCl_{3}H_{2}O_{5}NH_{3}, the
solubility rises considerably, and at 0° = 16·4, and at 20° =
about 27, whilst the passage into the purpureo-salt,
CoCl_{3},5NH_{3}, is accompanied by a great fall in the
solubility, namely, at 0° = 0·23, and at 20° = about 0·5. And as
crystallohydrates with a smaller amount of water are usually more
soluble than the higher crystallohydrates (Le Chatelier), whilst
here we find that the solubility falls (in the purpureo-salt) with
a loss of water, that water which is contained in the roseo-salt
cannot be compared with the water of crystallisation. Kournakoff,
therefore, connects the fall in solubility (in the passage of the
roseo- into the purpureo-salts) with the accompanying loss in the
reactive capacity of the chlorine.
In conclusion, it may be observed that the elements of the eighth
group--that is, the analogues of iron and platinum--according to
my opinion, will yield most fruitful results when studied as to
combinations with whole molecules, as already shown by the
examples of complex ammoniacal, cyanogen, nitro-, and other
compounds, which are easily formed in this eighth group, and are
remarkable for their stability. This faculty of the elements of
the eighth group for forming the complex compounds alluded to, is
in all probability connected with the position which the eighth
group occupies with regard to the others. Following the seventh,
which forms the type RX_{7}, it might be expected to contain the
most complex type, RX_{8}. This is met with in OsO_{4}. The other
elements of the eighth group, however, only form the lower types
RX_{2}, RX_{3}, RX_{4} ... and these accordingly should be
expected to aggregate themselves into the higher types, which is
accomplished in the formation of the above-mentioned complex
compounds.
[35 bis] Marshall (1891) obtained cobaltic sulphate,
Co_{2}(SO_{4})_{3},18H_{2}O, by the action of an electric current
upon a strong solution of CoSO_{4}.
[36] The action of an alkaline hypochlorite or hypobromite upon a
boiling solution of cobaltous salts, according to Schroederer
(1889), produces oxides, whose composition varies between
Co_{3}O_{5} (Rose's compound) and Co_{2}O_{3}, and also between
Co_{5}O_{8} and Co_{12}O_{19}. If caustic potash and then bromine
be added to the liquid, only Co_{2}O_{3} is formed. The action of
alkaline hypochlorites or hypo-bromites, or of iodine, upon
cobaltic salts, gives a highly-coloured precipitate which has a
different colour to the hydrate of the oxide Co_{2}(OH)_{6}.
According to Carnot the precipitate produced by the hypochlorites
has a composition Co_{10}O_{16}, whilst that given by iodine in
the presence of an alkali contains a larger amount of oxygen.
Fortmann (1891) reinvestigated the composition of the higher
oxygen oxide obtained by iodine in the presence of alkali, and
found that the greenish precipitate (which disengages oxygen when
heated to 100°) corresponds to the formula CoO_{2}. The reaction
must be expressed by the equation: CoX_{2} + I_{2} + 4KHO =
CoO_{2} + 2KX + 2KI + 2H_{2}O.
[37] Prior to Fortmann, Rousseau (1889) endeavoured to solve the
question as to whether CoO_{2} was able to combine with bases. He
succeeded in obtaining a barium compound corresponding to this
oxide. Fifteen grams of BaCl_{2} or BaBr_{2} are triturated with
5-6 grams of oxide of barium, and the mixture heated to redness in
a closed platinum crucible; 1 gram of oxide of cobalt is then
gradually added to the fused mass. Each addition of oxide is
accompanied by a violent disengagement of oxygen. After a short
time, however, the mass fuses quietly, and a salt settles at the
bottom of the crucible, which, when freed from the residue,
appears as black hexagonal, very brilliant crystals. In dissolving
in water this substance evolves chlorine; its composition
corresponds to the formula 2(CoO_{2})BaO. If the original mass be
heated for a long time (40 hours), the amount of dioxide in the
resultant mass decreases. The author obtained a neutral salt
having the composition CoO_{2}BaO (this compound = BaO_{2}CoO) by
breaking up the mass as it agglomerates together, and bringing the
pieces into contact with the more heated surface of the crucible.
This salt is formed between the somewhat narrow limits of
temperature 1,000°-1,100°; above and below these limits compounds
richer or poorer in CoO_{2} are formed. The formation of CoO_{2}
by the action of BaO_{2}, and the easy decomposition of CoO_{2}
with the evolution of oxygen, give reason for thinking that it
belongs to the class of peroxides (like Cr_{2}O_{7}, CaO_{2},
&c.); it is not yet known whether they give peroxide of hydrogen
like the true peroxides. The fact that it is obtained by means of
iodine (probably through HIO), and its great resemblance to
MnO_{2}, leads rather to the supposition that CoO_{2} is a very
feeble saline oxide. The form CoO_{2} is repeated in the cobaltic
compounds (Note 35), and the existence of CoO_{2} should have long
ago been recognised upon this basis.
Nickel alloys possess qualities which render them valuable for technical
purposes, the alloy of nickel with iron being particularly remarkable.
This alloy is met with in nature as _meteoric iron_. The Pallasoffsky
mass of meteoric iron, preserved in the St. Petersburg Academy, fell in
Siberia in the last century; it weighs about 15 cwt. and contains 88 p.c.
of iron and about 10 p.c. of nickel, with a small admixture of other
metals. In the arts _German silver_ is most extensively used; it is an
alloy containing nickel, copper, and zinc in various proportions. It
generally consists of about 50 parts of copper, 25 parts of zinc, and 25
parts of nickel. This alloy is characterised by its white colour
resembling that of silver, and, like this latter metal, it does not rust,
and therefore furnishes an excellent substitute for silver in the
majority of cases where it is used. Alloys which contain silver in
addition to nickel show the properties of silver to a still greater
extent. Alloys of nickel are used for currency, and if rich deposits of
nickel are discovered a wide field of application lies before it, not
only in a pure state (because it is a beautiful metal and does not rust)
but also for use in alloys. Steel vessels (pressed or forged out of sheet
steel) covered with nickel have such practical merits that their
manufacture, which has not long commenced, will most probably be rapidly
developed, whilst nickel steel, which exceeds ordinary steel in its
tenacity, has already proved its excellent qualities for many purposes
(for instance, for armour plate).
Until 1890 no compound of cobalt or nickel was known of sufficient
volatility to determine the molecular weights of the compounds of these
metals; but in 1890 Mr. L. Mond, in conducting (together with Langer and
Quincke) his researches on the action of nickel upon carbonic oxide
(Chapter IX., Note 24 bis), observed that nickel gradually volatilises in
a stream of carbonic oxide; this only takes place at low temperatures,
and is seen by the coloration of the flame of the carbonic oxide. This
observation led to the discovery of a remarkable volatile _compound of
nickel and carbonic oxide_, having as molecular composition
Ni(CO)_{4},[38] as determined by the vapour density and depression of the
freezing point. Cobalt and many other metals do not form volatile
compounds under these conditions, but iron gives a similar product (Note
26 bis). Ni(CO)_{4} is prepared by taking finely divided Ni (obtained by
reducing NiO by heating it in a stream of hydrogen, or by igniting the
oxalate NiC_{2}O_{4})[39] and passing (at a temperature below 50°, for
even at 60° decomposition may take place and an explosion) a stream of CO
over it; the latter carries over the vapour of the compound, which
condenses (in a well-cooled receiver) into a perfectly colourless
extremely mobile liquid, boiling without decomposition at 43°, and
crystallising in needles at -25° (Mond and Nasini, 1891). Liquid
Ni(CO)_{4} has a sp. gr. 1·356 at 0°, is insoluble in water, dissolves in
alcohol and benzene, and burns with a very smoky flame due to the
liberation of Ni. The vapour when passed through a tube heated to 180°
and above deposits a brilliant coating of metal, and disengages CO. If
the tube be strongly heated the decomposition is accompanied by an
explosion. If Ni(CO)_{4} as vapour be passed through a solution of
CuCl_{2}, it reduces the latter to metal; it has the same action upon an
ammoniacal solution of AgCl, strong nitric acid oxidises Ni(CO)_{4},
dilute solutions of acids have no action; if the vapour be passed through
strong sulphuric acid, CO is liberated, chlorine gives NiCl and COCl_{2};
no simple reactions of double decomposition are yet known for Ni(CO)_{4},
however, so that its connection with other carbon compounds is not clear.
Probably the formation of this compound could be applied for extracting
nickel from its ores.[40]
[38] This compound is known as nickel tetra-carbonyl. It appears to me
yet premature to judge of the structure of such an extraordinary
compound as Ni(CO)_{4}. It has long been known that potassium
combines with CO forming K_{_n_}(CO)_{_n_} (Chapter IX., Note 31),
but this substance is apparently saline and non-volatile, and has
as little in common with Ni(CO)_{4} as Na_{2}H has with SbH_{3}.
However, Berthelot observed that when NiC_{4}O_{4} is kept in air,
it oxidises and gives a colourless compound,
Ni_{3}C_{2}O_{3},10H_{2}O, having apparently saline properties. We
may add that Schützenberger, on reducing NiCl_{2} by heating it in
a current of hydrogen, observed that a nickel compound partially
volatilises with the HCl and gives metallic nickel when heated
again. The platinum compound, PtCl_{2}(CO)_{3} (Chapter XXIII.,
Note 11), offers the greatest analogy to Ni(CO)_{4}. This compound
was obtained as a volatile substance by Schützenberger by
moderately heating (to 235°) metallic platinum in a mixture of
chlorine and carbonic oxide. If we designate CO by Y, and an atom
of chlorine by X, then taking into account that, according to the
periodic system, Ni is an analogue of Pt, a certain degree of
correspondence is seen in the composition NiY_{4} and
PtX_{2}Y_{2}. It would be interesting to compare the reactions of
the two compounds.
[39] According to its empirical formula oxalate of nickel also contains
nickel and carbonic oxide.
[40] The following are the thermo-chemical data (according to Thomsen,
and referred to gram weights expressed by the formula, in large
calories or thousand units of heat) for the formation of
corresponding compounds of Mn, Fe, Co, Ni, and Cu (+ Aq signifies
that the reaction proceeds in an excess of water):
R = Mn Fe Co Ni Cu
R + Cl_{2} + Aq 128 100 95 94 63
R + Br_{2} + Aq 106 78 73 72 41
R + I_{2} + Aq 76 48 43 41 32
R + O + H_{2}O 95 68 63 61 38
R + O_{2} + SO_{2} + _n_H_{2}O 193 169 163 163 130
RCl_{2} + Aq +16 18 18 19 11
These examples show that for analogous reactions the amount of
heat evolved in passing from Mn to Fe, Co, Ni, and Cu varies in
regular sequences as the atomic weight increases. A similar
difference is to be found in other groups and series, and proves
that thermo-chemical phenomena are subject to the periodic law.
CHAPTER XXIII
THE PLATINUM METALS
The six metals: ruthenium, Ru, rhodium, Rh, palladium, Pd, osmium, Os,
iridium, Ir, and platinum, Pt, are met with associated together in
nature. Platinum always predominates over the others, and hence they are
known as the _platinum metals_. By their chemical character their
position in the periodic system is in the eighth group, corresponding
with iron, cobalt, and nickel.
The natural transition from titanium and vanadium to copper and zinc by
means of the elements of the iron group is demonstrated by all the
properties of these elements, and in exactly the same manner a transition
from zirconium, niobium, and molybdenum to silver, cadmium, and indium,
through ruthenium, rhodium, and palladium, is in perfect accordance with
fact and with the magnitude of the atomic weights, as also is the
position of osmium, iridium, and platinum between tantalum and tungsten
on the one side, and gold and mercury on the other. In all these three
cases the elements of smaller atomic weight (chromium, molybdenum, and
tungsten) are able, in their higher grades of oxidation, to give acid
oxides having the properties of distinct but feebly energetic acids (in
the lower oxides they give bases), whilst the elements of greater atomic
weight (zinc, cadmium, mercury), even in their higher grades of
oxidation, only give bases, although with feebly developed basic
properties. The platinum metals present the same intermediate properties
such as we have already seen in iron and the elements of the eighth
group.
In the platinum metals the intermediate properties _of feebly acid and
feebly basic metals_ are developed with great clearness, so that there is
not one sharply-defined acid anhydride among their oxides, although there
is a great diversity in the grades of oxidation from the type RO_{4} to
R_{2}O. The feebleness of the chemical forces observed in the platinum
metals is connected with the ready decomposability of their compounds,
with the small atomic volume of the metals themselves, and with their
large atomic weight. The oxides of platinum, iridium, and osmium can
scarcely be termed either basic or acid; they are capable of combinations
of both kinds, each of which is feeble. They are all intermediate oxides.
The atomic weights of platinum, iridium, and osmium are nearly 191 to
196, and of palladium, rhodium, and ruthenium, 104 to 106. Thus, strictly
speaking, we have here two series of metals, which are, moreover,
perfectly parallel to each other; three members in the first series, and
three members in the second--namely, platinum presents an analogy to
palladium, iridium to rhodium, and osmium to ruthenium. As a matter of
fact, however, the whole _group_ of the platinum metals is characterised
by _a number of common properties_, both physical and chemical, and,
moreover, there are several points of resemblance between the members of
this group and those of the _iron_ group (Chapter XXII.) The atomic
volumes (Table III., column 18) of the elements of this group are _nearly
equal_ and _very small_. The iron metals have atomic volumes of nearly 7,
whilst that of the metals allied to palladium is nearly 9, and of those
adjacent to platinum (Pt, Ir, Os) nearly 9·4. This comparatively small
atomic volume corresponds with the great infusibility and tenacity proper
to all the iron and platinum metals, and to their small chemical energy,
which stands out very clearly in the heavy platinum metals. All the
platinum metals are very _easily reduced_ by ignition and by the action
of various reducing agents, in which process oxygen, or a haloid group,
is disengaged from their compounds and the metal left behind. This is a
property of the platinum metals which determines many of their reactions,
and the circumstance of their always being found in nature _in a native
state_. In Russia in the Urals (discovered in 1819) and in Brazil (1735)
platinum is obtained from alluvial deposits, but in 1892 Professor
Inostrantseff discovered a vein deposit of platinum in serpentine near
Tagil in the Urals.[1] The facility with which they are reduced is so
great that their chlorides are even decomposed by gaseous hydrogen,
especially when shaken up and heated under a certain pressure. Hence it
will be readily understood that such metals as zinc, iron, &c., separate
them from solutions with great ease, which fact is taken advantage of in
practice and in the chemical treatment of the platinum metals.[1 bis]
[1] Wells and Penfield (1888) have described a mineral sperryllite
found in the Canadian gold-bearing quartz and consisting of
platinum diarsenide, PtAs_{2}. It is a noticeable fact that this
mineral clearly confirms the position of platinum in the same group
as iron, because it corresponds in crystalline form (regular
octahedron) and chemical composition with iron pyrites, FeS_{2}.
[1 bis] Some light is thrown upon the facility with which the platinum
compounds decompose by Thomsen's data, showing that in an excess of
water (+ Aq) the formation from platinum, of such a double salt as
PtCl_{2},2KCl, is accompanied by a comparatively small evolution of
heat (_see_ Chapter XXI., Note 40), for instance, Pt + Cl_{2} +
2KCl + Aq only evolves about 33,000 calories (hence the reaction,
Pt + Cl_{2} + Aq, will evidently disengage still less, because
PtCl_{2} + 2KCl evolves a certain amount of heat), whilst on the
other hand, Fe + Cl_{2} + Aq gives 100,000 calories, and even the
reaction with copper (for the formation of the double salt) evolves
63,000 calories.
All the platinum metals, like those of the iron group, are grey, with a
comparatively feeble metallic lustre, and are very infusible. In this
respect they stand in the same order as the metals of the iron series;
nickel is more fusible and whiter than cobalt and iron, so also palladium
is whiter and more fusible than rhodium and ruthenium, and platinum is
comparatively more fusible and whiter than iridium or osmium. The saline
compounds of these metals are red or yellow, like those of the majority
of the metals of the iron series, and like the latter, the different
forms of oxidation present different colours. Moreover, certain complex
compounds of the platinum metals, like certain complex compounds of the
iron series, either have particular characteristic tints or else are
colourless.
The platinum metals are found _in nature associated together in alluvial
deposits_ in a few localities, from which they are washed, owing to their
very considerable density, which enables a stream of water to wash away
the sand and clay with which they are mixed. Platinum deposits are
chiefly known in the Urals, and also in Brazil and a few other
localities. The platinum ore washed from these alluvial deposits presents
the appearance of more or less coarse grains, and sometimes, as it were,
of semi-fused nuggets.[2]
[2] The largest amount of platinum is extracted in the Urals, about
five tons annually. A certain amount of gold is extracted from the
washed platinum by means of mercury, which does not dissolve the
platinum metals but dissolves the gold accompanying the platinum in
its ores. Moreover, the ores of platinum always contain metals of
the iron series associated with them. The washed and mechanically
sorted ore in the majority of cases contains about 70 to 80 p.c. of
platinum, about 5 to 8 p.c. of iridium, and a somewhat smaller
quantity of osmium. The other platinum metals--palladium, rhodium,
and ruthenium--occur in smaller proportions than the three above
named. Sometimes grains of almost pure osmium-iridium, containing
only a small quantity of other metals, are found in platinum ores.
This _osmium-iridium_ may be easily separated from the other
platinum metals, owing to its being nearly insoluble in aqua regia,
by which the latter are easily dissolved. There are grains of
platinum which are magnetic. The grains of osmium-iridium are very
hard and malleable, and are therefore used for certain purposes,
for instance, for the tips of gold pens.
All the platinum metals give compounds with the halogens, and the
highest haloid type of combination for all is RX_{4}. For the majority of
the platinum metals this type is exceedingly unstable; the lower
compounds corresponding to the type RX_{2}, which are formed by the
separation of X_{2}, are more stable. In the type RX_{2} the platinum
metals form more stable salts, which offer no little resemblance to the
kindred compounds of the iron series--for example, to nickelous chloride,
NiCl_{2}, cobaltous chloride, CoCl_{2}, &c. This even expresses itself in
a similarity of volume (platinous chloride, PtCl_{2}, volume, 46;
nickelous chloride, NiCl_{2} = 50), although in the type RX_{2} the true
iron metals give very stable compounds, whilst the platinum metals
frequently react after the manner of suboxides, decomposing into the
metal and higher types, 2RX_{2} = R + RX_{4}. This probably depends on
the facility with which RX_{2} decomposes into R and X_{2}, when X_{2}
combines with the remaining portion of RX_{2}.
As in the series iron, cobalt, nickel, nickel gives NiO and Ni_{2}O_{3},
whilst cobalt and iron give higher and varied forms of oxidation, so also
among the platinum metals, platinum and palladium only give the forms
RX_{2} and RX_{4}, whilst rhodium and iridium form another and
intermediate type, RX_{3}, also met with in cobalt, corresponding with
the oxide, having the composition R_{2}O_{3}, besides which they form an
acid oxide, like ferric acid, which is also known in the form of salts,
but is in every respect unstable. _Osmium_ and _ruthenium_, like
manganese, form still higher oxides, and in this respect exhibit the
greatest diversity. They not only give RX_{2}, RX_{3}, RX_{4}, and
RX_{6}, but also a still _higher form of oxidation_, RO_{4}, which is not
met with in any other series. This form is exceedingly characteristic,
owing to the fact that the oxides, OsO_{4} and RuO_{4}, are volatile and
have feebly acid properties. In this respect they most resemble
permanganic anhydride, which is also somewhat volatile.[3]
[3] In characterising the platinum metals according to their relation
to the iron metals, it is very important to add two more very
remarkable points. The platinum metals are capable of forming a
sort of unstable compound with _hydrogen_; they absorb it and only
part with it when somewhat strongly heated. This faculty is
especially developed in platinum and palladium, and it is very
characteristic that nickel, which exactly corresponds with platinum
and palladium in the periodic system, should exhibit the same
faculty for retaining a considerable quantity of hydrogen (Graham's
and Raoult's experiments). Another characteristic property of the
platinum metals consists in their easily giving (like cobalt which
forms the cobaltic salts) stable and characteristic saline
_compounds with ammonia_, and like Fe and Co, double salts with the
cyanides of the alkali metals, especially in their lower forms of
combination. All the above so clearly brings the elements of the
iron series in close relation to the platinum metals, that the
eighth group acquires as natural a character as can be required,
with a certain originality or individuality for each element.
When dissolved in aqua regia (PtCl_{4} is formed) and liberated from the
solution by sal-ammoniac ((NH_{4})_{2}PtCl_{6} is formed) and reduced by
ignition (which may be done by Zn and other reducing agents, direct from
a solution of PtCl_{4}) platinum[3 bis] forms a powdery mass, known as
spongy platinum or platinum black. If this powder of platinum be heated
and pressed, or hammered in a cylinder, the grains aggregate or forge
together, and form a continuous, though of course not entirely
homogeneous, mass. Platinum was formerly, and is even now, worked up in
this manner. The platinum money formerly used in Russia was made in this
way. Sainte-Claire Deville, in the fifties, for the first time melted
platinum in considerable quantities by employing a special furnace made
in the form of a small reverberatory furnace, and composed of two pieces
of lime, on which the heat of the oxyhydrogen flame has no action. Into
this furnace (shown in fig. 34, Vol. I. p. 175)--or, more strictly
speaking, into the cavity made in the pieces of lime--the platinum is
introduced, and two orifices are made in the lime; through one, the
upper, or side orifice, is introduced an oxyhydrogen gas burner, in which
either detonating gas or a mixture of oxygen and coal-gas is burnt,
whilst the other orifice serves for the escape of the products of
combustion and certain impurities which are more volatile than the
platinum, and especially the oxidised compounds of osmium, ruthenium, and
palladium, which are comparatively easily volatilised by heat. In this
manner the platinum is converted into a continuous metallic form by means
of fusion, and this method is now used for melting considerable masses of
platinum[4] and its alloys with iridium.
[3 bis] Platinum was first obtained in the last century from Brazil,
where it was called silver (platinus). Watson in 1750 characterised
platinum as a separate independent metal. In 1803 Wollaston
discovered palladium and rhodium in crude platinum, and at about
the same time Tennant distinguished iridium and osmium in it.
Professor Claus, of Kazan, in his researches on the platinum metals
(about 1840) discovered ruthenium in them, and to him are due many
important discoveries with regard to these elements, such as the
indication of the remarkable analogy between the series Pd--Rh--Ru
and Pt--Ir--Os.
_The treatment of platinum ore_ is chiefly carried on for the
extraction of the platinum itself and its alloys with iridium,
because these metals offer a greater resistance to the action of
chemical reagents and high temperatures than any of the other
malleable and ductile metals, and therefore the wire so often used
in the laboratory and for technical purposes is made from them, as
also are various vessels used for chemical purposes in the
laboratory and in works. Thus sulphuric acid is distilled in
platinum retorts, and many substances are fused, ignited, and
evaporated in the laboratory in platinum crucibles and on platinum
foil. Gold and many other substances are dissolved in dishes made
of iridium-platinum, because the alloys of platinum and iridium are
but slightly attacked when subjected to the action of aqua regia.
The comparatively high density (about 21·5), hardness, ductility,
and infusibility (it does not melt at a furnace heat, but only in
the oxyhydrogen flame or electric furnace), as well as the fact of
its resisting the action of water, air, and other reagents, renders
an alloy of 90 parts of platinum and 10 parts of iridium (Deville's
platinum-iridium alloy) a most valuable material for making
standard weights and measures, such as the metre, kilogram, and
pound, and therefore all the newest standards of most countries are
made of this alloy.
[4] This process has altered the technical treatment of platinum to a
considerable extent. It has in particular facilitated the
manufacture of alloys of platinum with iridium and rhodium from the
pure platinum ores, since it is sufficient to fuse the ore in order
for the greater amount of the osmium to burn off, and for the mass
to fuse into a homogeneous, malleable alloy, which can be directly
made use of. There is very little ruthenium in the ores of
platinum. If during fusion lead be added, it dissolves the platinum
(and other platinum metals) owing to its being able to form a very
characteristic alloy containing PtPb. If an alloy of the two metals
be left exposed to moist air, the excess of lead is converted into
carbonate (white lead) in the presence of the water and carbonic
acid of the air, whilst the above platinum alloy remains unchanged.
The white lead may be extracted by dilute acid, and the alloy PtPb
remains unaltered. The other platinum metals also give similar
alloys with lead. The fusibility of these alloys enables the
platinum metals to be separated from the gangue of the ore, and
they may afterwards be separated from the lead by subjecting the
alloy to oxidation in furnaces furnished with a bone ash bed,
because the lead is then oxidised and absorbed by the bone ash,
leaving the platinum metals untouched. This method of treatment was
proposed by H. Sainte-Claire Deville in the sixties, and is also
used in the analysis of these metals (_see_ further on).
To obtain pure platinum, the ore is treated with aqua regia in which
only the osmium and iridium are insoluble. The solution contains the
platinum metals in the form RCl_{4}, and in the lower forms of
chlorination, RCl_{3} and RCl_{2}, because some of these metals--for
instance, palladium and rhodium--form such unstable chlorides of the type
RX_{4} that they partially decompose even when diluted with water, and
pass into the stable lower type of combination; in addition to which the
chlorine is very easily disengaged if it comes in contact with substances
on which it can act. In this respect platinum resists the action of heat
and reducing agents better than any of its companions--that is, it passes
with greater difficulty from PtCl_{4} to the lower compound PtCl_{2}. On
this is based the method of preparation of more or less pure platinum.
Lime or sodium hydroxide is added to the solution in aqua regia until
neutralised, or only containing a very slight excess of alkali. It is
best to first evaporate and slightly ignite the solution, in order to
remove the excess of acid, and by heating it to partially convert the
higher chlorides of the palladium, &c., into the lower. The addition of
alkalis completes the reduction, because the chlorine held in the
compounds RX_{4} acts on the alkali like free chlorine, converting it
into a hypochlorite. Thus palladium chloride, PdCl_{4}, for example, is
converted into palladious chloride, PdCl_{2}, by this means, according to
the equation PdCl_{4} + 2NaHO = PdCl_{2} + NaCl + NaClO + H_{2}O. In a
similar manner iridic chloride, IrCl_{4}, is converted into the
trichloride, IrCl_{3}, by this method. When this conversion takes place
the platinum still remains in the form of platinic chloride, PtCl_{4}. It
is then possible to take advantage of a certain difference in the
properties of the higher and lower chlorides of the platinum metals. Thus
lime precipitates the lower chlorides of the members of the platinum
metals occurring in solution without acting on the platinic chloride,
PtCl_{4}, and hence the addition of a large proportion of lime
immediately precipitates the associated metals, leaving the platinum
itself in solution in the form of a soluble double salt,
PtCl_{4},CaCl_{2}. A far better and more perfect _separation_ is effected
_by means of ammonium chloride_, which gives, with platinic chloride, an
insoluble yellow precipitate, PtCl_{4},2NH_{4}Cl, whilst it forms soluble
double salts with the lower chlorides RCl_{2} and RCl_{3}, so that
ammonium chloride precipitates the platinum only from the solution
obtained by the preceding method. These methods are employed for
preparing the platinum which is used for the manufacture of platinum
articles, because, having platinum in solution as calcium
platinochloride, PtCaCl_{6}, or as the insoluble ammonium
platinochloride, Pt(NH_{4})_{2}Cl_{6}, the platinum compound in every
case, after drying or ignition, loses all the chlorine from the platinic
chloride and leaves finely-divided metallic platinum, which may be
converted into homogeneous metal by compression and forging, or by
fusion.[5]
[5] For the ultimate purification of platinum from palladium and
iridium the metals must be re-dissolved in aqua regia, and the
solution evaporated until the residue begins to evolve chlorine.
The residue is then re-precipitated with ammonium or potassium
chloride. The precipitate may still contain a certain amount of
iridium, which passes with greater difficulty from the
tetrachloride, IrCl_{4}, into the trichloride, IrCl_{3}, but it
will be quite free from palladium, because the latter easily loses
its chlorine and passes into palladious chloride, PdCl_{2}, which
gives an easily-soluble salt with potassium chloride. The
precipitate, containing a small quantity of iridium, is then heated
with sodium carbonate in a crucible, when the mass decomposes,
giving metallic platinum and iridium oxide. If potassium chloride
has been employed, the residue after ignition is washed with water
and treated with aqua regia. The iridium oxide remains undissolved,
and the platinum easily passes into solution. Only cold and dilute
aqua regia must be used. The solution will then contain pure
platinic chloride, which forms the starting-point for the
preparation of all platinum compounds. Pure platinum for accurate
researches (for instance, for the unit of light, according to
Violle's method) may be obtained (Mylius and Foerster, 1892) by
Finkener's method, by dissolving the impure metal in aqua regia (it
should be evaporated to drive off the nitrogen compounds), and
adding NaCl so as to form a double sodium salt, which is purified
by crystallising with a small amount of caustic soda, washing the
crystals with a strong solution of NaCl, and then dissolving them
in a hot 1 p.c. solution of soda, repeating the above and
ultimately igniting the double salt, previously dried at 120°, in a
stream of hydrogen; platinum black and NaCl are then formed. The
three following are very sensitive tests (to thousandths of a per
cent.) for the presence of Ir, Ru, Rh, Pd (osmium is not usually
present in platinum which has once been purified, since it easily
volatilises with Cl_{2} and CO_{2}, and in the first treatment of
the crude platinum either passes off as OsO_{4} or remains
undissolved), Fe, Cu, Ag, and Pb: (1) the assay is alloyed with 10
parts of pure lead, the alloy treated with dilute nitric acid (to
remove the greater part of the Pb), and dissolved in aqua regia;
the residue will consist of Ir and Ru; the Pb is precipitated from
the nitric acid solution by sulphuric acid, whilst the remaining
platinum metals are reduced from the evaporated solution by formic
acid, and the resultant precipitate fused with KHSO_{4}; the Pd and
Rh are thus converted into soluble salts, and the former is then
precipitated by HgC_{2}N_{2}. (2) Iron may be detected by the usual
reagents, if the crude platinum be dissolved in aqua regia, and the
platinum metals precipitated from the solution by formic acid. (3)
If crude platinum (as foil or sponge) be heated in a mixture of
chlorine and carbonic oxide it volatilises (with a certain amount
of Ir, Pd, Fe, &c.) as PtCl_{2},2CO (Note 11), whilst the whole of
the Rh, Ag, and Cu it may contain remains behind. Among other
characteristic reactions for the platinum metals, we may mention:
(1) that rhodium is precipitated from the solution obtained after
fusion with KHSO_{4} (in which Pt does not dissolve) by NH_{3},
acetic and formic acids; (2) that dilute aqua regia dissolves
precipitated Pt, but not Rh; (3) that if the insoluble residue of
the platinum metals (Ir, Ru, Os) obtained, after treating with aqua
regia, be fused with a mixture of 1 part of KNO_{3} and 3 parts of
K_{2}CO_{3} (in a gold crucible), and then treated with water, it
gives a solution containing the Ru (and a portion of the Ir), but
which throws it all down when saturated with chlorine and boiled;
(4) that if iridium be fused with a mixture of KHO and KNO_{3}, it
gives a soluble potassium salt, IrK_{2}O_{4} (the solution is
blue), which, when saturated with chlorine, gives IrCl_{4}, which
is precipitated by NH_{4}Cl (the precipitate is black), forming a
double salt, leaving metallic Ir after ignition; (5) that rhodium
mixed with NaCl and ignited in a current of chlorine gives a
soluble double salt (from which sal-ammoniac separates Pt and Ir),
which gives (according to Jörgensen) a difficultly soluble
purpureo-salt (Chapter XXII., Note 35), Rh_{2}Cl_{3},5NH_{3}, when
treated with NH_{3}; in this form the Rh may be easily purified and
obtained in a metallic form by igniting in hydrogen; and (6) that
palladium, dissolved in aqua regia and dried (NH_{4}Cl throws down
any Pt), gives soluble PdCl_{2}, which forms an easily
crystallisable yellow salt, PdCl_{2}NH_{3}, with ammonia; this salt
(Wilm) may be easily purified by crystallisation, and gives
metallic Pd when ignited. These reactions illustrate the method of
separating the platinum metals from each other.
Metallic _platinum_ in a fused state has a specific gravity of 21; it is
grey, softer than iron but harder than copper, exceedingly ductile, and
therefore easily drawn into wire and rolled into thin sheets, and may be
hammered into crucibles and drawn into thin tubes, &c. In the state in
which it is obtained by the ignition of its compounds, it forms a spongy
mass, known as spongy platinum, or else as powder (platinum black).[6] In
either case it is dull grey, and is characterised, as we already know, by
the faculty of absorbing hydrogen and other gases. Platinum is not acted
on by hydrochloric, hydriodic, nitric, and sulphuric acids, or a mixture
of hydrofluoric and nitric acids. Aqua regia, and any liquid containing
chlorine or able to evolve chlorine or bromine, dissolves platinum.
Alkalis are decomposed by platinum at a red heat, owing to the faculty of
the platinum oxide, PtO_{2}, formed to combine with alkaline bases,
inasmuch as it has a feebly-developed acid character (_see_ Note 8).
Sulphur, phosphorus (the phosphide, PtP_{2}, is formed), arsenic and
silicon all act more or less rapidly on platinum, under the influence of
heat. Many of the metals form alloys with it. Even charcoal combines with
platinum when it is ignited with it, and therefore carbonaceous matter
cannot be subjected to prolonged and powerful ignition in platinum
vessels. Hence a platinum crucible soon becomes dull on the surface in a
smoky flame. Platinum also forms alloys with zinc, lead, tin, copper,
gold, and silver.[7] Although mercury does not directly dissolve
platinum, still it forms a solution or amalgam with spongy platinum in
the presence of sodium amalgam; a similar amalgam is also formed by the
action of sodium amalgam on a solution of platinum chloride, and is used
for physical experiments.
[6] We have already become acquainted with the effect of finely-divided
platinum on many gaseous substances. It is best seen in the
so-called _platinum black_, which is a coal-black powder left by
the action of sulphuric acid on the alloy of zinc and platinum, or
which is precipitated by metallic zinc from a dilute solution of
platinum. In any case, finely-divided platinum absorbs gases more
powerfully and rapidly the more finely divided and porous it is.
Sulphurous anhydride, hydrogen, alcohol, and many organic
substances in the presence of such platinum are easily oxidised by
the oxygen of the air, although they do not combine with it
directly. The absorption of oxygen is as much as several hundred
volumes per one volume of platinum, and the oxidising power of such
absorbed oxygen is taken advantage of not only in the laboratory
but even in manufacturing processes. Asbestos or charcoal, soaked
in a solution of platinic chloride and ignited, is very useful for
this purpose, because by this means it becomes coated with platinum
black. If 50 grams of PtCl_{4} be dissolved in 60 c.c. of water,
and 70 c.c. of a strong (40 p.c.) solution of formic aldehyde
added, the mixture cooled, and then a solution of 50 grams of NaHO
in 50 grams of water added, the platinum is precipitated. After
washing with water the precipitate passes into solution and forms a
black liquid containing _soluble colloidal platinum_ (Loew, 1890).
If the precipitated platinum be allowed to absorb oxygen on the
filter, the temperature rises 40°, and a very porous _platinum
black_ is obtained which vigorously facilitates oxidation.
[7] It is necessary to remark that platinum when alloyed with silver,
or as amalgam, is soluble in nitric acid, and in this respect it
differs from gold, so that it is possible, by alloying gold with
silver, and acting on the alloy with nitric acid, to recognise the
presence of platinum in the gold, because nitric acid does not act
on gold alloyed with silver.
There are _two kinds_ of _platinum compounds_, PtX_{4} and PtX_{2}. The
former are produced by an excess of halogen in the cold, and the latter
by the aid of heat or by the splitting up of the former. The
starting-point for the platinum compounds is _platinum tetrachloride_,
_platinic chloride_, PtCl_{4}, obtained by dissolving platinum in aqua
regia.[7 bis] The solution crystallises in the cold, in a desiccator, in
the form of reddish-brown deliquescent crystals which contain
hydrochloric acid, PtCl_{4},2HCl,6H_{2}O, and behave like a true acid
whose salts correspond to the formula R_{2}PtCl_{6}--ammonium
platinochloride, for example.[7 tri] The hydrochloric acid is liberated
from these crystals by gently heating or evaporating the solution to
dryness; or, better still, after treatment with silver nitrate a
reddish-brown mass remains behind, which dissolves in water, and forms a
yellowish-red solution which on cooling deposits crystals of the
composition PtCl_{4},8H_{2}O. The _tendency_ of PtCl_{4} _to combine_
with hydrochloric acid and water--that is, _to form higher crystalline
compounds_--is evident in the platinum compounds, and must be taken into
account in explaining the properties of platinum and the formation of
many other of its complex compounds. Dilute solutions of platinic
chloride are yellow, and are completely reduced by hydrogen, sulphurous
anhydride, and many reducing agents, which first convert the platinic
chloride into the lower compound platinous chloride, PtCl_{2}. That
faculty which reveals itself in platinum tetrachloride of combining with
water of crystallisation and hydrochloric acid is distinctly marked in
its property, with which we are already acquainted, of giving
precipitates with the salts of potassium, ammonium, rubidium, &c. In
general it _readily forms double salts_, R_{2}PtCl_{6} = PtCl_{4} + 2RCl,
where R is a univalent metal such as potassium or NH_{4}. Hence the
addition of a solution of potassium or ammonium chloride to a solution of
platinic chloride is followed by the formation of a yellow precipitate,
which is sparingly soluble in water and almost entirely insoluble in
alcohol and ether (platinic chloride is soluble in alcohol, potassium
iridiochloride, IrK_{3}Cl_{6}, _i.e._ a compound of IrCl_{3}, is soluble
in water but not in alcohol). It is especially remarkable in this case,
that the potassium compounds here, as in a number of other instances,
separate in an anhydrous form, whilst the sodium compounds, which are
soluble in water and alcohol, form red crystals containing water. The
composition Na_{2}PtCl_{6},6H_{2}O exactly corresponds with the
above-mentioned hydrochloric compound. The compounds with barium,
BaPtCl_{6},4H_{2}O, strontium, SrPtCl_{6},8H_{2}O, calcium, magnesium,
iron, manganese, and many other metals are all soluble in water.[8]
[7 bis] PtCl_{4} is also formed by the action of a mixture of HCl
vapour and air, and by the action of gaseous chlorine upon
platinum.
[7 tri] Pigeon (1891) obtained fine yellow crystals of
PtH_{2}Cl_{6},4H_{2}O by adding strong sulphuric acid to a strong
solution of PtH_{2}Cl_{6},6H_{2}O. If crystals of
H_{2}PtCl_{6},6H_{2}O be melted in vacuo (60°) in the presence of
anhydrous potash, a red-brown solid hydrate is obtained containing
less water and HCl, which parts with the remainder at 200°, leaving
anhydrous PtCl_{4}. The latter does not disengage chlorine before
220°, and is perfectly soluble in water.
[8] Nilson (1877), who investigated the platinochlorides of various
metals subsequently to Bonsdorff, Topsöe, Clève, Marignac, and
others, found that univalent and bivalent metals--such as hydrogen,
potassium, ammonium ... beryllium, calcium, barium--give compounds
of such a composition that there is always twice as much chlorine
in the platinic chloride as in the combined metallic chloride; for
example, K_{2}Cl_{2},PtCl_{4}; BeCl_{2},PtCl_{4},8H_{2}O, &c. Such
trivalent metals as aluminium, iron (ferric), chromium, didymium,
cerium (cerous) form compounds of the type RCl_{3}PtCl_{4}, in
which the amounts of chlorine are in the ratio 3:4. Only indium and
yttrium give salts of a different composition--namely,
2InCl_{3},5PtCl_{4},36H_{2}O and 4YCl_{3},5PtCl_{4},51H_{2}O. Such
quadrivalent metals as thorium, tin, zirconium give compounds of
the type RCl_{4},PtCl_{4}, in which the ratio of the chlorine is
1:1. In this manner the valency of a metal may, to a certain
extent, be judged from the composition of the double salts formed
with platinic chloride.
Platinic bromide, PtBr_{4}, and iodide, PtI_{4}, are analogous to
the tetrachloride, but the iodide is decomposed still more easily
than the chloride. If sulphuric acid be added to platinic chloride,
and the solution evaporated, it forms a black porous mass like
charcoal, which deliquesces in the air, and has the composition
Pt(SO_{4})_{2}. But this, the only oxygen salt of the type PtX_{4},
is exceedingly unstable. This is due to the fact that _platinum
oxide_, the oxide of the type PtO_{2}, has a feeble acid character.
This is shown in a number of instances. Thus if a strong solution
of platinic chloride treated with sodium carbonate be exposed to
the action of light or evaporated to dryness and then washed with
water, a sodium platinate, Pt_{3}Na_{2}O_{7},6H_{2}O, remains. The
composition of this salt, if we regard it in the same sense as we
did the salts of silicic, titanic, molybdic and other acids, will
be PtO(ONa)_{2},2PtO_{2},6H_{2}O--that is, the same type is
repeated as we saw in the crystalline compounds of platinum
tetrachloride with sodium chloride, or with hydrochloric
acid--namely, the type PtX_{4}8Y, where Y is the molecule
H_{2}O,HCl, &c. Similar compounds are also obtained with other
alkalis. They will be platinates of the alkalis in which the
platinic oxide, PtO_{2}, plays the part of an acid oxide. Rousseau
(1889) obtained different grades of combination BaOPtO_{2},
3(BaO)2PtO_{2}, &c., by igniting a mixture of PtCl_{4} and caustic
baryta. If such an alkaline compound of platinum be treated with
acetic acid, the alkali combines with the latter, and a _platinic
hydroxide_, Pt(OH)_{4}, remains as a brown mass, which loses water
and oxygen when ignited, and in so doing decomposes with a slight
explosion. When slightly ignited this hydroxide first loses water
and gives the very unstable oxide PtO_{2}. Platinic sulphide,
PtS_{2}, belongs to the same type; it is precipitated by the action
of sulphuretted hydrogen on a solution of platinum tetrachloride.
The moist precipitate is capable of attracting oxygen, and is then
converted into the sulphate above mentioned, which is soluble in
water. This absorption of oxygen and conversion into sulphate is
another illustration of the basic nature of PtO_{2}, so that it
clearly exhibits both basic and acid properties. The latter appear,
for instance, in the fact that platinic sulphide, PtS_{2}, gives
crystalline compounds with the alkali sulphides.
_Platinous chloride_, PtCl_{2}, is formed when hydrogen platinochloride,
PtH_{2}Cl_{6}, is ignited at 300°, or when potassium is heated at 230° in
a stream of chlorine. The undecomposed tetrachloride is extracted from
the residue by washing it with water, and a greenish-grey or brown
insoluble mass of the dichloride (sp. gr. 5·9) is then obtained. It is
soluble in hydrochloric acid, giving an acid solution of the composition
PtCl_{2},2HCl, corresponding with the type of double salts PtR_{2}Cl_{4}.
Although platinous chloride decomposes below 500°, still it is formed to
a small extent at higher temperatures. Troost and Hautefeuille, and
Seelheim observed that when platinum was strongly ignited in a stream of
chlorine, the metal, as it were, slowly volatilised and was deposited in
crystals; a volatile chloride, probably platinous chloride, was evidently
formed in this case, and decomposed subsequently to its formation,
depositing crystals of platinum.
The properties of platinum above-described are repeated more or
less distinctly, or sometimes with certain modifications, in the
above-mentioned associates and analogues of this metal. Thus although
palladium forms PdCl_{4}, this form passes into PdCl_{2} with extreme
ease.[9] Whilst rhodium and iridium in dissolving in aqua regia also form
RhCl_{4} and IrCl_{4}, but they pass into RhCl_{3} and IrCl_{3}[9 bis]
very easily when heated or when acted upon by substances capable of
taking up chlorine (even alkalis, which form bleaching salts). Among the
platinum metals, ruthenium and osmium have the most acid character, and
although they give RuCl_{4} and OsCl_{4} they are easily oxidised to
RuO_{4}, and OsO_{4} by the action of chlorine in the presence of water;
the latter are volatile and may be distilled with the water and
hydrochloric acid, from a solution containing other platinum metals.[9
tri] Thus with respect to the types of combination, all the platinum
metals, under certain circumstances, give compounds of the type
RX_{4}--for instance, RO_{2}, RCl_{4}, &c. But this is the highest form
for only platinum and palladium. The remaining platinum metals further,
_like iron, give acids_ of the type RO_{3} or hydrates, H_{2}RO_{4} =
RO_{2}(HO)_{2} (the type of sulphuric acid); but they, like ferric and
manganic acids, are chiefly known in the form of salts of the composition
K_{2}RO_{4} or K_{2}R_{2}O_{7} (like the dichromate). These salts are
obtained, like the manganates and ferrates, by fusing the oxides, or even
the metals themselves, with nitric, or, better still, with potassium
peroxide. They are soluble in water, are easily deoxidised and do not
yield the acid anhydrides under the action of acids, but break up, either
(like the ferrate) forming oxygen and a basic oxide (iridium and rhodium
react in this manner, as they do not give higher forms of oxidation), or
passing into a lower and higher form of oxidation--that is, reacting like
a manganate (or partly like nitrite or phosphite). Osmium and ruthenium
react according to the latter form, as they are capable of giving _higher
forms of oxidation_, OsO_{4} and RuO_{4}, and therefore their reactions
of decomposition may be essentially represented by the equation: 2OsO_{3}
= OsO_{2} + OsO_{4}.[10]
[9] In comparing the characteristics of the platinum metals, it must be
observed that palladium in its form of combination PdX_{2} gives
saline compounds of considerable stability. Amongst them _palladous
chloride_ is formed by the direct action of chlorine or aqua regia
(not in excess or in dilute solutions) on palladium. It forms a
brown solution, which gives a black insoluble precipitate of
_palladous iodide_, PdI_{2}, with solutions of iodides (in this
respect, as in many others, palladium resembles mercury in the
mercuric compounds HgX_{2}). With a solution of mercuric cyanide it
gives a yellowish white precipitate, palladous cyanide,
PdC_{2}N_{2}, which is soluble in potassium cyanide, and gives
other double salts, M_{2}PdC_{4}N_{4}.
That portion of the platinum ore which dissolves in aqua regia and
is precipitated by ammonium or potassium chloride does not contain
palladium. It remains in solution, because the palladic chloride,
PdCl_{4}, is decomposed and the palladous chloride formed is not
precipitated by ammonium chloride; the same holds good for all the
other lower chlorides of the platinum metals. Zinc (and iron)
separates out all the unprecipitated platinum metals (and also
copper, &c.) from the solution. The palladium is found in these
platinum residues precipitated by zinc. If this mixture of metals
be treated with aqua regia, all the palladium will pass into
solution as palladous chloride with some platinic chloride. By this
treatment the main portion of the iridium, rhodium, &c. remains
almost undissolved, the platinum is separated from the mixture of
palladous and platinic chlorides by a solution of ammonium
chloride, and the solution of palladium is precipitated by
potassium iodide or mercuric cyanide. Wilm (1881) showed that
palladium may be separated from an impure solution by saturating it
with ammonia; all the iron present is thus precipitated, and, after
filtering, the addition of hydrochloric acid to the filtrate gives
a yellow precipitate of an ammonio-palladium compound,
PdCl_{2},2NH_{3}, whilst nearly all the other metals remain in
solution. _Metallic palladium_ is obtained by igniting the
ammonio-compound or the cyanide, PdC_{2}N_{2}. It occurs native,
although rarely, and is a metal of a whiter colour than platinum,
sp. gr. 11·4, melts at about 1,500°; it is much more volatile than
platinum, partially oxidises on the surface when heated (Wilm
obtained spongy palladium by igniting PdCl_{2},2NH_{3}, and
observed that it gives PdO when ignited in oxygen, and that on
further ignition this oxide forms a mixture of Pd_{2}O and Pd), and
loses its absorbed oxygen on a further rise of temperature. It does
not blacken or tarnish (does not absorb sulphur) in the air at the
ordinary temperature, and is therefore better suited than silver
for astronomical and other instruments in which fine divisions have
to be engraved on a white metal, in order that the fine lines
should be clearly visible. The most remarkable property of
palladium, discovered by Graham, consists in its capacity for
_absorbing_ a large amount of _hydrogen_. Ignited palladium absorbs
as much as 940 volumes of hydrogen, or about 0·7 p.c. of its own
weight, which closely approaches to the formation of the compound
Pd_{3}H_{2}, and probably indicates the formation of _palladium
hydride_, Pd_{2}H. This absorption also takes place at the ordinary
temperature--for example, when palladium serves as an electrode at
which hydrogen is evolved. In absorbing the hydrogen, the palladium
does not change in appearance, and retains all its metallic
properties, only its volume increases by about 10 p.c.--that is,
the hydrogen pushes out and separates the atoms of the palladium
from each other, and is itself compressed to 1/900 of its volume.
This compression indicates a great force of chemical attraction,
and is accompanied by the evolution of heat (Chapter II., Note 38).
The absorption of 1 grm. of hydrogen by metallic palladium (Favre)
is accompanied by the evolution of 4·2 thousand calories (for Pt
20, for Na 13, for K 10 thousand units of heat). Troost showed that
the dissociation pressure of palladium hydride is inconsiderable at
the ordinary temperature, but reaches the atmospheric pressure at
about 140°. This subject was subsequently investigated by A. A.
Cracow of St. Petersburg (1894), who showed that at first the
absorption of hydrogen by the palladium proceeds like solution,
according to the law of Dalton and Henry, but that towards the end
it proceeds like a dissociation phenomenon in definite compounds;
this forms another link between the phenomenon of solution and of
the formation of definite atomic compounds. Cracow's observations
for a temperature 18°, showed that the electro-conductivity and
tension vary until a compound Pd_{2}H is reached, and namely, that
the tension _p_ rises with the volume _v_ of hydrogen absorbed,
according to the law of Dalton and Henry--for instance, for
_p_ = 2·1 3·2 5·5 7·7 mm.
_v_ = 14 20 34 47
The maximum tension at 18° is 9 mm. At a temperature of about 140°
(in the vapour of xylene) the maximum tension is about 760 mm., and
when _v_ = 10-50 vols. the tension (according to Cracow's
experiments) stands at 90-450 mm.--that is, increases in proportion
to the volume of hydrogen absorbed. But from the point of view of
chemical mechanics it is especially important to remark that
Moutier clearly showed, through palladium hydride, the similarity
of the phenomena which proceed in evaporation and dissociation,
which fact Henri Sainte-Claire Deville placed as a fundamental
proposition in the theory of dissociation. It is possible upon the
basis of the second law of the theory of heat, according to the law
of the variation of the tension _p_ of evaporation with the
temperature T (counted from -273°), to calculate the latent heat of
evaporation L (_see_ works on physics) because 424L = T(1/_d_ -
1/D)_dp_/_dt_, where _d_ and D are the weights of cubic measures of
the gas (vapour) and liquid. (Thus, for instance, for water, when
_t_ = 100°, T = 373, _d_ = 0·605, D = 960, _dp_/_dt_ = 0·027 m.,
13,596 = 367, L = 536, whence 424L = 227,264, and the second
portion of the equation 226,144, which is sufficiently near, within
the limits of experimental error, _see_ Chapter I., Note 11.) The
same equation is applicable to the dissociation of Na_{2}H and
K_{2}H--(Chapter XII., Note 42)--but it has only been verified in
this respect for Pd_{2}H, since Moutier, by calculating the amount
of heat L evolved, for _t_ = 20, according to the variation of the
tension (_dp_/_dt_) obtained 4·1 thousand calories, which is very
near the figure obtained experimentally by Favre (_see_ Chapter
XII., Note 44). The absorbed hydrogen is easily disengaged by
ignition or decreased pressure. The resultant compound does not
decompose at the ordinary temperature, but when exposed to air the
metal sometimes glows spontaneously, owing to the hydrogen burning
at the expense of the atmospheric oxygen. The hydrogen absorbed by
palladium acts towards many solutions as a reducing agent; in a
word, everything here points to the formation of a definite
compound and at the same time of a physically-compressed gas, and
forms one of the best examples of the bond existing between
chemical and physical processes, to which we have many times drawn
attention. It must be again remembered that the other metals of the
eighth group, even copper, are, like palladium and platinum, able
to combine with hydrogen. The permeability of iron and platinum
tubes to hydrogen is naturally due to the formation of similar
compounds, but palladium is the most permeable.
[9 bis] _Rhodium_ is generally separated, together with iridium, from
the residues left after the treatment of native platinum, because
the palladium is entirely separated from them, and the ruthenium is
present in them in very small traces, whilst the osmium at any rate
is easily separated, as we shall soon see. The mixture of rhodium
and iridium which is left undissolved in dilute aqua regia is
dissolved in chlorine water, or by the action of chlorine on a
mixture of the metals with sodium chloride. In either case both
metals pass into solution. They may be separated by many methods.
In either case (if the action be aided by heat) the rhodium is
obtained in the form of the chloride RhCl_{3}, and the iridium as
iridious chloride, IrCl_{3}. They both form double salts with
sodium chloride which are soluble in water, but the iridium salt is
also partially soluble in alcohol, whilst the rhodium salt is not.
A mixture of the chlorides, when treated with dilute aqua regia,
gives iridic chloride, IrCl_{4}, whilst the rhodium chloride,
RhCl_{3}, remains unaltered; ammonium chloride then precipitates
the iridium as ammonium iridiochloride, Ir(NH_{4})2Cl_{6}, and on
evaporating the rose-coloured filtrate the rhodium gives a
crystalline salt, Rh(NH_{4})_{3}Cl_{6}. Rhodium and its various
oxides are dissolved when fused with potassium hydrogen sulphate,
and give a soluble double sulphate (whilst iridium remains unacted
on); this fact is very characteristic for this metal, which offers
in its properties many points of resemblance with the iron metals.
When fused with potassium hydroxide and chlorate it is oxidised
like iridium, but it is not afterwards soluble in water, in which
respect it differs from ruthenium. This is taken advantage of for
separating rhodium, ruthenium, and iridium. In any case, rhodium
under ordinary conditions always gives salts of the type RX_{3},
and not of any other type; and not only halogen salts, but also
oxygen salts, are known in this type, which is rare among the
platinum metals. Rhodium chloride, RhCl_{3}, is known in an
insoluble anhydrous and also in a soluble form (like CrX_{3} or
salts of chromic oxides), in which it easily gives double salts,
compounds with water of crystallisation, and forms rose-coloured
solutions. In this form rhodium easily gives double salts of the
two types RhM_{3}Cl_{6} and RhM_{2}Cl_{3}--for example,
K_{5}RhCl_{6},3H_{2}O and K_{2}RhCl_{5},H_{2}O. Solutions of the
salts (at least, the ammonium salt) of the first kind give salts of
the second kind when they are boiled. If a strong solution of
potash be added to a red solution of rhodium chloride and boiled, a
black precipitate of the hydroxide Rh(OH)_{3} is formed; but if the
solution of potash is added little by little, it gives a yellow
precipitate containing more water. This yellow hydrate of rhodium
oxide gives a yellow solution when it is dissolved in acids, which
only becomes rose-coloured after being boiled. It is obvious a
change here takes place, like the transmutations of the salts of
chromic oxide. It is also a remarkable fact that the black
hydroxide, like many other oxidised compounds of the platinoid
metals, does not dissolve in the ordinary oxygen acids, whilst the
yellow hydroxide is easily soluble and gives yellow solutions,
which deposit imperfectly crystallised salts. Metallic rhodium is
easily obtained by igniting its oxygen and other compounds in
hydrogen, or by precipitation with zinc. It resembles platinum, and
has a sp. gr. of 12·1. At the ordinary temperature it decomposes
formic acid into hydrogen and carbonic anhydride, with development
of heat (Deville). With the alkali sulphites, the salts of rhodium
and iridium of the type RX_{3} give sparingly-soluble precipitates
of double sulphites of the composition R(SO_{3}Na)_{3},H_{2}O, by
means of which these metals may be separated from solution, and
also may be separated from each other, for a mixture of these salts
when treated with strong sulphuric acid gives a soluble iridium
sulphate and leaves a red insoluble double salt of rhodium and
sodium. It may be remarked that the oxides Ir_{2}O_{3} and
Rh_{2}O_{3} are comparatively stable and are easily formed, and
that they also form different double salts (for instance,
IrCl_{3},3KCl_{3}H_{2}O, RhCl_{3},2NH_{4}Cl_{4}H_{2}O,
RhCl_{3},3NH_{4}Cl1-1/2H_{2}O) and compounds like the cobaltia
compounds (for instance, luteo-salts RhX_{3},6NH_{3}, roseo-salts,
RhX_{3}H_{2}O_{5}NH_{3}, and purpureo-salts IrX_{3},5NH_{3}, &c.)
_Iridious oxide_, Ir_{2}O_{3}, is obtained by fusing iridious
chloride and its compounds with sodium carbonate, and treating the
mass with water. The oxide is then left as a black powder, which,
when strongly heated, is decomposed into iridium and oxygen; it is
easily reduced, and is insoluble in acids, which indicates the
feeble basic character of this oxide, in many respects resembling
such oxides as cobaltic oxide, ceric or lead dioxide, &c. It does
not dissolve when fused with potassium hydrogen sulphate. Rhodium
oxide, Rh_{2}O_{3}, is a far more energetic base. It dissolves when
fused with potassium hydrogen sulphate.
From what has been said respecting the separation of platinum and
rhodium it will be understood how the compounds of _iridium_, which
is the main associate of platinum, are obtained. In describing the
treatment of osmiridium we shall again have an opportunity of
learning the method of extraction of the compounds of this metal,
which has in recent times found a technical application in the form
of its oxide, Ir_{2}O_{3}; this is obtained from many of the
compounds of iridium by ignition with water, is easily reduced by
hydrogen, and is insoluble in acids. It is used in painting on
china, for giving a black colour. Iridium itself is more
difficultly fusible than platinum, and when fused it does not
decompose acids or even aqua regia; it is extremely hard, and is
not malleable; its sp. gr. is 22·4. In the form of powder it
dissolves in aqua regia, and is even partially oxidised when heated
in air, sets fire to hydrogen, and, in a word, closely resembles
platinum. Heated in an excess of chlorine it gives iridic chloride,
IrCl_{4}, but this loses chlorine at 50°; it is, however, more
stable in the form of double salts, which have a characteristic
_black_ colour--for instance, Ir(NH_{4})_{2}Cl_{6}--but they give
iridious chloride, IrCl_{3}, when treated with sulphuric acid.
[9 tri] We have yet to become acquainted with the two remaining
associates of platinum--ruthenium and osmium--whose most important
property is that they are oxidised even when heated in air, and
that they are able to give _volatile_ oxides of the form RuO_{4}
and OsO_{4}; these have a powerful odour (like iodine and nitrous
anhydride). Both these higher oxides are solids; they volatilise
with great ease at 100°; the former is yellow and the latter white.
They are known as _ruthenic_ and _osmic anhydrides_, although their
aqueous solutions (they both slowly dissolve in water) do not show
an acid reaction, and although they do not even expel carbonic
anhydride from potassium carbonate, do not give crystalline salts
with bases, and their alkaline solutions partially deposit them
again when boiled (an excess of water decomposes the salts). The
formulæ OsO_{4} and RuO_{4} correspond with the vapour density of
these oxides. Thus Deville found the vapour density of osmic
anhydride to be 128 (by the formula 127·5) referred to hydrogen.
Tennant and Vauquelin discovered this compound, and Berzelius,
Wöhler, Fritzsche, Struvé, Deville, Claus, Joly, and others helped
in its investigation; nevertheless there are still many questions
concerning it which remain unsolved. It should be observed that
RO_{4} is the highest known form for an oxygen compound, and RH_{4}
is the highest known form for a compound of hydrogen; whilst the
highest forms of acid hydrates contain SiH_{4}O_{4}, PH_{3}O_{4},
SH_{2}O_{4}, ClHO_{4}--all with four atoms of oxygen, and therefore
in this number there is apparently the limit for the simple forms
of combination of hydrogen and oxygen. In combination with
_several_ atoms of an element, or several elements, there may be
more than O_{4} or H_{4}, but a molecule never contains more than
four atoms of either O or H to one atom of another element. Thus
the simplest forms of combination of hydrogen and oxygen are
exhausted by the list RH_{4}, RH_{3}, RH_{2}, RH, RO, RO_{2},
RO_{3}, RO_{4}. The extreme members are RH_{4} and RO_{4}, and are
only met with for such elements as carbon, silicon, osmium,
ruthenium, which also give RCl_{4} with chlorine. In these extreme
forms, RH_{4} and RO_{4}, the compounds are the least stable
(compare SiH_{4}, PH_{3}, SH_{2}, ClH, or RuO_{4}, MoO_{3},
ZrO_{2}, SrO), and easily give up part, or even all, their oxygen
or hydrogen.