Chapter XXIV: Note 9 ^{bis}). If there were more such (1)

well-investigated cases of solutions, our knowledge of double
salts, solutions, the influence of water, equilibria, isomorphous
mixtures, and such-like provinces of chemical relations might be
considerably advanced.

[24] The normal ferric salts are decomposed by heat and even by water,
forming basic salts, which may be prepared in various ways.
Generally ferric hydroxide is dissolved in solutions of ferric
nitrate; if it contains a double quantity of iron the basic salt
is formed which contains Fe_{2}O_{3} (in the form of hydroxide) +
2Fe_{2}(NO_{3})_{6} = 3Fe_{2}O(NO_{3})_{4}, a salt of the type
Fe_{2}OX_{4}. Probably water enters into its composition. With
considerable quantities of ferric oxide, insoluble basic salts are
obtained containing various amounts of ferric hydroxide. Thus when
a solution of the above-mentioned basic acid is boiled, a
precipitate is formed containing
4(Fe_{2}O_{3})_{8},2(N_{2}O_{5}),3H_{2}O, which probably contains
2Fe_{2}O_{2}(NO_{3})_{2} + 2Fe_{2}O_{3},3H_{2}O. If a solution of
basic nitrate be sealed in a tube and then immersed in boiling
water, the colour of the solution changes just in the same way as
if a solution of ferric acetate had been employed (Note 22). The
solution obtained smells strongly of nitric acid, and on adding a
drop of sulphuric or hydrochloric acid the insoluble variety of
hydrated ferric oxide is precipitated.

Normal ferric _orthophosphate_ is soluble in sulphuric,
hydrochloric, and nitric acids, but insoluble in others, such as,
for instance, acetic acid. The composition of this salt in the
anhydrous state is FePO_{4}, because in orthophosphoric acid there
are three atoms of hydrogen, and iron, in the ferric state,
replaces the three atoms of hydrogen. This salt is obtained from
ferric acetate, which, with disodium phosphate, forms a _white
precipitate_ of FePO_{4}, containing water. If a solution of
ferric chloride (yellowish-red colour) be mixed with a solution of
sodium acetate in excess, the liquid assumes an intense brown
colour which demonstrates the formation of a certain quantity of
ferric acetate; then the disodium phosphate directly forms a white
gelatinous precipitate of ferric phosphate. By this means the
whole of the iron may be precipitated, and the liquid which was
brown then becomes colourless. If this normal salt be dissolved in
orthophosphoric acid, the crystalline acid salt
FeH_{3}(PO_{4})_{2} is formed. If there be an excess of ferric
oxide in the solution, the precipitate will consist of the basic
salt. If ferric phosphate be dissolved in hydrochloric acid, and
ammonia be added, a salt is precipitated on heating which, after
continued washing in water and heating (to remove the water), has
the composition Fe_{4}P_{2}O_{11}--that is,
2Fe_{2}O_{3},P_{2}O_{5}. In an aqueous condition this salt may be
considered as ferric hydroxide, Fe_{2}(OH)_{6}, in which (OH)_{3}
is replaced by the equivalent group PO_{4}. Whenever ammonia is
added to a solution containing an excess of ferric salt and a
certain amount of phosphoric acid, a precipitate is formed
containing the whole of the phosphoric acid in the mass of the
ferric oxide.

Ferric oxide is characterised as a feeble base, and also by the
fact of its forming double salts--for instance, _potassium iron
alum_, which has a composition
Fe_{2}(SO_{4})_{3},K_{2}SO_{4},24H_{2}O or
FeK(SO_{4})_{2},12H_{2}O. It is obtained in the form of almost
colourless or light rose-coloured large octahedra of the regular
system by simply mixing solutions of potassium sulphate and the
ferric sulphate obtained by dissolving ferric oxide in sulphuric
acid.

[25] It would seem that all normal ferric salts are colourless, and
that the brown colour which is peculiar to the solutions is really
due to basic ferric salts. A remarkable example of the apparent
change of colour of salts is represented by the ferrous and ferric
oxalates. The former in a dry state has a yellow colour, although
as a rule the ferrous salts are green, and the latter is
colourless or pale green. When the normal ferric salt is dissolved
in water it is, like many salts, probably decomposed by the water
into acid and basic salts, and the latter communicates a brown
colour to the solution. Iron alum is almost colourless, is easily
decomposed by water, and is the best proof of our assertion. The
study of the phenomena peculiar to ferric nitrate might, in my
opinion, give a very useful addition to our knowledge of the
aqueous solutions of salts in general.

[25 bis] The reaction FeX_{3} + KI = FeX_{2} + KX + I proceeds
comparatively slowly in solutions, is not complete (depends upon
the mass), and is reversible. In this connection we may cite the
following data from Seubert and Rohrer's (1894) comprehensive
researches. The investigations were conducted with solutions
containing 1/10 gram--equivalent weights of Fe_{2}(SO_{4})_{3}
(_i.e._ containing 20 grams of salt per litre), and a
corresponding solution of KI; the amount of iodine liberated being
determined (after the addition of starch) by a solution (also 1/10
normal) of Na_{2}S_{2}O_{3} (_see_ Chapter XX., Note 42). The
progress of the reaction was expressed by the amount of liberated
iodine in percentages of the theoretical amount. For instance, the
following amount of iodide of potassium was decomposed when
Fe_{2}(SO_{4})_{3} + 2_n_KI was taken:

_n_ = 1 2 3 6 10 20
After 15´ 11·4 26·3 40·6 73·5 91·6 96·0
" 30´ 14·0 35·8 47·8 78·5 94·3 97·4
" 1 hour 19·0 42·7 56·0 84·0 95·7 97·6
" 10 " 32·6 56·0 75·7 93·2 96·5 97·6
" 48 " 39·4 67·7 82·6 93·4 96·6 97·6

Similar results were obtained for FeCl_{3}, but then the amount of
iodine liberated was somewhat greater. Similar results were also
obtained by increasing the mass of FeX_{3} per KI, and by
replacing it by HI (_see_ Chapter XXI., Note 26).

Iron forms one other oxide besides the ferric and ferrous oxides; this contains twice as much oxygen as the former, but is so very unstable that it can neither be obtained in the free state nor as a hydrate. Whenever such conditions of double decomposition occur as should allow of its separation in the free state, it decomposes into oxygen and ferric oxide. It is known in the state of salts, and is only stable in the presence of alkalis, and forms salts with them which have a decidedly alkaline reaction; it is therefore a feebly acid oxide. Thus when small pieces of iron are heated with nitre or potassium chlorate a potassium salt of the composition K_{2}FeO_{4} is formed, and therefore the hydrate corresponding with this salt should have the composition H_{2}FeO_{4}. It is called _ferric acid_. Its anhydride ought to contain FeO_{3} or Fe_{2}O_{6}--twice as much oxygen as ferric oxide. If a solution of potassium ferrate be mixed with acid, the free hydrate ought to be formed, but it immediately decomposes (2K_{2}FeO_{4} + 5H_{2}SO_{4} = 2K_{2}SO_{4} + Fe_{2}(SO_{4})_{3} + 5H_{2}O + O_{3}), oxygen being evolved. If a small quantity of acid be taken, or if a solution of potassium ferrate be heated with solutions of other metallic salts, ferric oxide is separated--for instance:

2CuSO_{4} + 2K_{2}FeO_{4} = 2K_{2}SO_{4} + O_{3} + Fe_{2}O_{3} + 2CuO.

Both these oxides are of course deposited in the form of hydrates. This shows that not only the hydrate H_{2}FeO_{4}, but also the salts of the heavy metals corresponding with this higher oxide of iron, are not formed by reactions of double decomposition. The solution of potassium ferrate naturally acts as a powerful oxidising agent; for instance, it transforms manganous oxide into the dioxide, sulphurous into sulphuric acid, oxalic acid into carbonic anhydride and water, &c.[26]

[26] If chlorine be passed through a strong solution of potassium
hydroxide in which hydrated ferric oxide is suspended, the turbid
liquid acquires a dark pomegranate-red colour and contains
potassium ferrate: 10KHO + Fe_{2}O_{3} + 3Cl_{2} = 2K_{2}FeO_{4} +
6KCl + 5H_{2}O. The chlorine must not be in excess, otherwise the
salt is again decomposed, although the mode of decomposition is
unknown; however, ferric chloride and potassium chlorate are
probably formed. Another way in which the above-described salt is
formed is also remarkable; a galvanic current (from 6 Grove
elements) is passed through cast-iron and platinum electrodes into
a strong solution of potassium hydroxide. The cast-iron electrode
is connected with the positive pole, and the platinum electrode is
surrounded by a porous earthenware cylinder. Oxygen would be
evolved at the cast-iron electrode, but it is used up in
oxidation, and a dark solution of potassium ferrate is therefore
formed about it. It is remarkable that the cast iron cannot be
replaced by wrought iron.

Iron thus combines with oxygen in three proportions: RO, R_{2}O_{3}, and RO_{3}. It might have been expected that there would be intermediate stages RO_{2} (corresponding to pyrites FeS_{2}) and R_{2}O_{5}, but for iron these are unknown.[26 bis] The lower oxide has a distinctly basic character, the higher is feebly acid. The only one which is stable in the free state is ferric oxide, Fe_{2}O_{3}; the suboxide, FeO, absorbs oxygen, and ferric anhydride, FeO_{3}, evolves it. It is also the same for other elements; the character of each is determined by the relative degree of stability of the known oxides. The salts FeX_{2} correspond with the suboxide, the salts FeX_{3} or Fe_{2}X_{6} with the sesquioxide, and FeX_{6} represents those of ferric acid, as its potassium salt is FeO_{2}(OK)_{2}, corresponding with K_{2}SO_{4}, K_{2}MnO_{4}, K_{2}CrO_{4}, &c. Iron therefore forms compounds of the types FeX_{2}, FeX_{3}, and FeX_{6}, but this latter, like the type NX_{5}, does not appear separately, but only when X represents heterogeneous elements or groups; for instance, for nitrogen in the form of NO_{2}(OH), NH_{4}Cl, &c., for iron in the form of FeO_{2}(OK)_{2}. But still the type FeX_{6} exists, and therefore FeX_{2} and FeX_{3} are compounds which, like ammonia, NH_{3}, are capable of further combinations up to FeX_{6}; this is also seen in the property of ferrous and ferric salts of forming compounds with water of crystallisation, besides double and basic salts, whose stability is determined by the quality of the elements included in the types FeX_{2} and FeX_{3}.[26 tri] It is therefore to be expected that there should be complex compounds derived from ferrous and ferric oxides. Amongst these the series of cyanogen compounds is particularly interesting; their formation and character is not only determined by the property which iron possesses of forming complex types, but also by the similar faculty of the cyanogen compounds, which, like nitriles (Chapter IX.), have clearly developed properties of polymerisation and in general of forming complex compounds.[27]

[26 bis] When Mond and his assistants obtained the remarkable volatile
compound Ni(CO)_{4} (described later, Chapter XXII.), it was shown
subsequently by Mond and Quincke (1891), and also by Berthelot,
that iron, under certain conditions, in a stream of carbonic
oxide, also volatilises and forms a compound like that given by
nickel. Roscoe and Scudder then showed that when water gas is
passed through and kept under pressure (8 atmospheres) in iron
vessels a portion of the iron volatilises from the sides of the
vessel, and that when the gas is burnt it deposits a certain
amount of oxides of iron (the same result is obtained with
ordinary coal gas which contains a small amount of CO). To obtain
the _volatile compound of iron with carbonic oxide_, Mond prepared
a finely divided iron by heating the oxalate in a stream of
hydrogen, and after cooling it to 80°-45° he passed CO over the
powder. The iron then formed (although very slowly) a volatile
compound containing Fe(CO)_{5} (as though it answered to a very
high type, FeX_{10}), which when cooled condenses into a liquid
(slightly coloured, probably owing to incipient decomposition),
sp. gr. 1·47, which solidifies at -21°, boils at about 103°, and
has a vapour density (about 6·5 with respect to air) corresponding
to the above formula; it decomposes at 180°. Water and dilute
acids do not act upon it, but it decomposes under the action of
light and forms a hard, non-volatile crystalline yellow compound
Fe_{2}(CO)_{7} which decomposes at 80° and again forms Fe(CO)_{5}.

[26 tri] When the molecular Fe_{2}Cl_{6} is produced instead of
FeCl_{3} this complication of the type also occurs.

[27] Some light may be thrown upon the faculty of Fe of forming various
compounds with CN, by the fact that Fe not only combines with
carbon but also with nitrogen. _Nitride of iron_ Fe_{2}N was
obtained by Fowler by heating finely powdered iron in a stream of
NH_{3} at the temperature of melting lead.

_In the cyanogen compounds of iron_, two degrees might be expected: Fe(CN)_{2}, corresponding with ferrous oxide, and Fe(CN)_{3}, corresponding with ferric oxide. There are actually, however, many other known compounds, intermediate and far more complex. They correspond with the double salts so easily formed by metallic cyanides. The two following double salts are particularly well known, very stable, often used, and easily prepared. _Potassium ferrocyanide_ or _yellow prussiate of potash_, a double salt of cyanide of potassium and ferrous cyanide, has the composition FeC_{2}N_{2},4KCN; its crystals contain 3 mol. of water: K_{4}FeC_{6}N_{6},3H_{2}O. The other is _potassium ferricyanide_ or _red prussiate of potash_. It is also known as _Gmelin's salt_, and contains cyanide of potassium with ferric cyanide; its composition is Fe(CN)_{3},3KCN or K_{3}FeC_{6}N_{6}. Its crystals do not contain water. It is obtained from the first by the action of chlorine, which removes one atom of the potassium. A whole series of other ferrocyanic compounds correspond with these ordinary salts.

Before treating of the preparation and properties of these two remarkable and very stable salts, it must be observed that with ordinary reagents neither of them gives the same double decompositions as the other ferrous and ferric salts, and they both present a series of remarkable properties. Thus these salts have a neutral reaction, are unchanged by air, dilute acids, or water, unlike potassium cyanide and even some of its double salts. When solutions of these salts are treated with caustic alkalis, they do not give a precipitate of ferrous or ferric hydroxides, neither are they precipitated by sodium carbonate. This led the earlier investigators to recognise special independent groupings in them. The yellow prussiate was considered to contain the complex radicle FeC_{6}N_{6} combined with potassium, namely with K_{4}, and K_{3} was attributed to the red prussiate. This was confirmed by the fact that whilst in both salts any other metal, even hydrogen, might be substituted for potassium, the iron remained unchangeable, just as nitrogen in cyanogen, ammonium, and nitrates does not enter into double decomposition, being in the state of the complex radicles CN, NH_{4}, NO_{2}. Such a representation is, however, completely superfluous for the explanation of the peculiarities in the reactions of such compounds as double salts. If a magnesium salt which can be precipitated by potassium hydroxide does not form a precipitate in the presence of ammonium chloride, it is very clear that it is owing to the formation of a soluble double salt which is not decomposed by alkalis. And there is no necessity to account for the peculiarity of reaction of a double salt by the formation of a new complex radicle. In the same way also, in the presence of an excess of tartaric acid, cupric salts do not form a precipitate with potassium hydroxide, because a double salt is formed. These peculiarities are more easily understood in the case of cyanogen compounds than in all others, because all cyanogen compounds, as unsaturated compounds, show a marked tendency to complexity. This tendency is satisfied in double salts. The appearance of a peculiar character in double cyanides is the more easily understood since in the case of potassium cyanide itself, and also in hydrocyanic acid, a great many peculiarities have been observed which are not encountered in those haloid compounds, potassium chloride and hydrochloric acid, with which it was usual to compare cyanogen compounds. These peculiarities become more comprehensible on comparing cyanogen compounds with ammonium compounds. Thus in the presence of ammonia the reactions of many compounds change considerably. If in addition to this it is remembered that the presence of many carbon (organic) compounds frequently completely disturbs the reaction of salts, the peculiarities of certain double cyanides will appear still less strange, because they contain carbon. The fact that the presence of carbon or another element in the compound produces a change in the reactions, may be compared to the action of oxygen, which, when entering into a combination, also very materially changes the nature of reactions. Chlorine is not detected by silver nitrate when it is in the form of potassium chlorate, KClO_{3}, as it is detected in potassium chloride, KCl. The iron in ferrous and ferric compounds varies in its reactions. In addition to the above-mentioned facts, consideration ought to be given to the circumstance that the easy mutability of nitric acid undergoes modification in its alkali salts, and in general the properties of a salt often differ much from those of the acid. Every double salt ought to be regarded as a peculiar kind of saline compound: potassium cyanide is, as it were, a basic, and ferrous cyanide an acid, element. They may be unstable in the separate state, but form a stable double compound when combined together; the act of combination disengages the energy of the elements, and they, so to speak, saturate each other. Of course, all this is not a definite explanation, but then the supposition of a special complex radicle can even less be regarded as such.

Potassium ferrocyanide, K_{4}FeC_{6}N_{6}, is very easily formed by mixing solutions of ferrous sulphate and potassium cyanide. First, a white precipitate of ferrous cyanide, FeC_{2}N_{2}, is formed, which becomes blue on exposure to air, but is soluble in an excess of potassium cyanide, forming the ferrocyanide. The same yellow prussiate is obtained on heating animal nitrogenous charcoal or animal matters--such as horn, leather cuttings, &c.--with potassium carbonate in iron vessels,[27 bis] the mass formed being afterwards boiled with water with exposure to air, potassium cyanide first appearing, which gives yellow prussiate. The animal charcoal may be exchanged for wood charcoal, permeated with potassium carbonate and heated in nitrogen or ammonia; the mass thus produced is then boiled in water with ferric oxide.[28] In this manner it is manufactured on the large scale, and is called 'yellow prussiate' ('prussiate de potasse,' Blutlaugensalz).

[27 bis] The sulphur of the animal refuse here forms the compound
FeKS_{2}, which by the action of potassium cyanide yields
potassium sulphide, thiocyanate, and ferrocyanide.

[28] Potassium ferrocyanide may also be obtained from Prussian blue by
boiling with a solution of potassium hydroxide, and from the
ferricyanide by the action of alkalis and reducing substances
(because the red prussiate is a product of oxidation produced by
the action of chlorine: a ferric salt is reduced to a ferrous
salt), &c. In many works (especially in Germany and France) yellow
prussiate is prepared from the mass, containing oxide of iron, and
employed for purifying coal gas (Vol. I., p. 361), which generally
contains cyanogen compounds. About 2 p.c. of the nitrogen
contained in coal is converted into cyanogen, which forms Prussian
blue and thiocyanates in the mass used for purifying the gas. On
evaporation the solution yields large yellow crystals containing 3
molecules of water, which is easily expelled by heating above
100°. 100 parts of water at the ordinary temperature are capable
of dissolving 25 parts of this salt; its sp. gr. is 1·83. When
ignited it forms potassium cyanide and iron carbide, FeC_{2}
(Chapter XIII., Note 12). Oxidising substances change it into
potassium ferricyanide. With strong sulphuric acid it gives
carbonic oxide, and with dilute sulphuric acid, when heated,
prussic acid is evolved according to the equation:
2K_{4}FeC_{6}N_{6} + 3H_{2}SO_{4} = K_{2}Fe_{2}C_{6}N_{6} +
3K_{2}SO_{4} + 6HCN; hence in the yellow prussiate K_{2} replaces
Fe.

It is easy to substitute other metals for the potassium in the yellow prussiate. The hydrogen salt or hydroferrocyanic acid, H_{4}FeC_{6}N_{6}, is obtained by mixing strong solutions of yellow prussiate and hydrochloric acid. If ether be added and the air excluded, the acid is obtained directly in the form of a white scarcely crystalline precipitate which becomes blue on exposure to air (as ferrous cyanide does from the formation of blue compounds of ferrous and ferric cyanides, and it is on this account used in cotton printing). It is soluble in water and alcohol, but not in ether, has marked acid properties, and decomposes carbonates, which renders it easily possible to prepare ferrocyanides of the metals of the alkalis and alkaline earths; these are readily soluble, have a neutral reaction, and resemble the yellow prussiate. Solutions of these salts form precipitates with the salts of other metals, because the ferrocyanides of the heavy metals are insoluble. Here either the whole of the potassium of the yellow prussiate, or only a part of it, is exchanged for an equivalent quantity of the heavy metal. Thus, when a cupric salt is added to a solution of yellow prussiate, a red precipitate is obtained which still contains half the potassium of the yellow prussiate:

K_{4}FeC_{6}N_{6} + CuSO_{4} = K_{2}CuFeC_{6}N_{6} + K_{2}SO_{4}.

But if the process be reversed (the salt of copper being then in excess) the whole of the potassium will be exchanged for copper, forming a reddish-brown precipitate, Cu_{2}FeC_{6}N_{6},9H_{2}O. This reaction and those similar to it are very sensitive and may be used for detecting metals in solution, more especially as the colour of the precipitate very often shows a marked difference when one metal is exchanged for another. Zinc, cadmium, lead, antimony, tin, silver, cuprous and aurous salts form _white_ precipitates; cupric, uranium, titanium and molybdenum salts _reddish-brown_; those of nickel, cobalt, and chromium, _green_ precipitates; _with ferrous salts_, ferrocyanide forms, as has been already mentioned, a _white_ precipitate--namely, Fe_{2}FeC_{6}N_{6}, or FeC_{2}N_{2}--which turns blue on exposure to air, and with ferric salts a _blue precipitate_ called _Prussian blue_. Here the potassium is replaced by iron, the reaction being expressed thus: 2Fe_{2}Cl_{6} + 3K_{4}FeC_{6}N_{6} = 12KCl + Fe_{4}Fe_{3}C_{18}N_{18}, the latter formula expressing the composition of Prussian blue. It is therefore the compound 4Fe(CN)_{3} + 3Fe(CN)_{2}. The yellow prussiate is prepared in chemical works on a large scale especially for the manufacture of this blue pigment, which is used for dyeing cloth and other fabrics and also as one of the ordinary blue paints. It is insoluble in water, and the stuffs are therefore dyed by first soaking them in a solution of a ferric salt and then in a solution of yellow prussiate. If however an excess of yellow prussiate be present complete substitution between potassium and iron does not occur, and _soluble Prussian blue_ is formed; KFe_{2}(CN)_{6} = KCN,Fe(CN)_{2},Fe(CN)_{3}. This blue salt is colloidal, is soluble in pure water, but insoluble and precipitated when other salts--for instance, potassium or sodium chloride--are present even in small quantities, and is therefore first obtained as a precipitate.[29]

[29] Skraup obtained this salt both from potassium ferrocyanide with
ferric chloride and from ferricyanide with ferrous chloride, which
evidently shows that it contains iron in both the ferric and
ferrous states. With ferrous chloride it forms Prussian blue, and
with ferric chloride Turnbull's blue.

Prussian blue was discovered in the beginning of the last century
by a Berlin manufacturer, Diesbach. It was then prepared, as it
sometimes is also at present, directly from potassium cyanide
obtained by heating animal charcoal with potassium carbonate. The
mass thus obtained is dissolved in water, alum is added to the
solution in order to saturate the free alkali, and then a solution
of green vitriol is added which has previously been sufficiently
exposed to the air to contain both ferric and ferrous salts. If
the solution of potassium cyanide be mixed with a solution
containing both salts, Prussian blue will be formed, because it is
a compound of ferrous cyanide, FeC_{2}N_{2}, and ferric cyanide,
Fe_{2}C_{6}N_{6}. A ferric salt with potassium ferrocyanide forms
a blue colour, because ferrous cyanide is obtained from the first
salt and ferric cyanide from the second. During the preparation of
this compound alkali must be avoided, as otherwise the precipitate
would contain oxides of iron. Prussian blue has not a crystalline
structure; it forms a blue mass with a copper-red metallic lustre.
Both acids and alkalis act on it. The action is at first confined
to the ferric salt it contains. Thus alkalis form ferric oxide and
ferrocyanide in solution: 2Fe_{2}C_{6}N_{6},3FeC_{2}N_{2} + 12KHO
= 2(Fe_{2}O_{3},3H_{2}O) + 3K_{4}FeC_{6}N_{6}. Various
ferrocyanides may thus be prepared. Prussian blue is soluble in an
aqueous solution of oxalic acid, forming blue ink. In air, when
exposed to the action of light, it fades; but in the dark again
absorbs oxygen and becomes blue, which fact is also sometimes
noticed in blue cloth. An excess of potassium ferrocyanide renders
Prussian blue soluble in water, although insoluble in various
saline solutions--that is, it converts it into the soluble
variety. Strong hydrochloric acid also dissolves Prussian blue.

Potassium ferricyanide, or _red prussiate_ of potash, K_{3}FeC_{6}N_{6}, is called 'Gmelin's salt,' because this savant obtained it by the action of chlorine on a solution of the yellow prussiate: K_{4}FeC_{6}N_{6} + Cl = K_{3}FeC_{6}N_{6} + KCl. The reaction is due to the ferrous salt being changed by the action of the chlorine into a ferric salt. It separates from solutions in anhydrous, well-formed prisms of a red colour, but the solution has an olive colour; 100 parts of water, at 10°, dissolve 37 parts of the salt, and at 100°, 78 parts.[30] The red prussiate gives a blue precipitate with ferrous salts, called _Turnbull's blue_, very much like Prussian blue (and the soluble variety), because it also contains ferrous cyanide and ferric cyanide, although in another proportion, being formed according to the equation: 3FeCl_{2} + 2K_{3}FeC_{6}N_{6} = 6KCl + Fe_{3}Fe_{2}C_{12}N_{12}, or 3FeC_{2}N_{2},Fe_{2}C_{6}N_{6}; in Prussian blue we have Fe_{7}Cy_{18}, and here Fe_{5}Cy_{12}. A ferric salt ought to form ferric cyanide Fe_{2}C_{6}N_{6}, with red prussiate, but ferric cyanide is soluble, and therefore no precipitate is obtained, and the liquid only becomes brown.[31]

[30] An excess of chlorine must not be employed in preparing this
compound, otherwise the reaction goes further. It is easy to find
out when the action of the chlorine on potassium ferrocyanide must
cease; it is only necessary to take a sample of the liquid and add
a solution of a ferric salt to it. If a precipitate of Prussian
blue is formed, more chlorine must be added, as there is still
some undecomposed ferrocyanide, for the ferricyanide does not give
a precipitate with ferric salts. Potassium ferricyanide, like the
ferrocyanide, easily exchanges its potassium for hydrogen and
various metals by double decomposition. With the salts of tin,
silver, and mercury it forms yellow precipitates, and with those
of uranium, nickel, cobalt, copper, and bismuth brown
precipitates. The lead salt under the action of sulphuretted
hydrogen forms lead sulphide and a hydrogen salt or acid,
H_{3}FeC_{6}N_{6}, corresponding with potassium ferricyanide,
which is soluble, crystallises in red needles, and resembles
hydroferrocyanic acid, H_{4}FeC_{6}N_{6}. Under the action of
reducing agents--for instance, sulphuretted hydrogen,
copper--potassium ferricyanide is changed into ferrocyanide,
especially in the presence of alkalis, and thus forms a rather
energetic _oxidising agent_--capable, for instance, of changing
manganous oxide into dioxide, bleaching tissues, &c.

[31] It is important to mention a series of readily crystallisable
salts formed by the action of nitric acid on potassium and other
ferrocyanides and ferricyanides. These salt contain the elements
of nitric oxide, and are therefore called _nitro-(nitroso)
ferricyanides_ (_nitroprussides_). Generally a crystalline sodium
salt is obtained, Na_{2}FeC_{5}N_{6}O,2H_{2}O. In its composition
this salt differs from the red sodium salt, Na_{3}FeC_{6}N_{6}, by
the fact that in it one molecule of sodium cyanide, NaCN, is
replaced by nitric oxide, NO. In order to prepare it, potassium
ferrocyanide in powder is mixed with five-sevenths of its weight
of nitric acid diluted with an equal volume of water. The mixture
is at first left at the ordinary temperature, and then heated on a
water-bath. Here ferricyanide is first of all formed (as shown by
the liquid giving a precipitate with ferrous chloride), which then
disappears (no precipitate with ferrous chloride), and forms a
green precipitate. The liquid, when cooled, deposits crystals of
nitre. The liquid is then strained off and mixed with sodium
carbonate, boiled, filtered, and evaporated; sodium nitrate and
the salt described are deposited in crystals. It separates in
prisms of a red colour. Alkalis and salts of the alkaline earths
do not give precipitates: they are soluble, but the salts of iron,
zinc, copper, and silver form precipitates where sodium is
exchanged with these metals. It is remarkable that the sulphides
of the alkali metals give with this salt an intense bright purple
coloration. This series of compounds was discovered by Gmelin and
studied by Playfair and others (1849).

This series to a certain extent resembles the nitro-sulphide
series described by Roussin. Here the primary compound consists of
black crystals, which are obtained as follows:--Solutions of
potassium hydrosulphide and nitrate are mixed, and the mixture is
agitated whilst ferric chloride is added, then boiled and
filtered; on cooling, _black crystals_ are deposited, having the
composition Fe_{6}S_{3}(NO)_{10},H_{2}O (Rosenberg), or, according
to Demel, FeNO_{2},NH_{2}S. They have a slightly metallic lustre,
and are soluble in water, alcohol, and ether. They absorb the
latter as easily as calcium chloride absorbs water. In the
presence of alkalis these crystals remain unchanged, but with
acids they evolve nitric oxides. There are several compounds which
are capable of interchanging, and correspond with Roussin's salt.
Here we enter into the series of the nitrogen compounds which have
been as yet but little investigated, and will most probably in
time form most instructive material for studying the nature of
that element. These series of compounds are as unlike the usual
saline compounds of inorganic chemistry as are organic
hydrocarbons. There is no necessity to describe these series in
detail, because their connection with other compounds is not yet
clear, and they have not yet any application.

If chlorine and sodium are representatives of independent groups of elements, the same may also be said of iron. Its nearest analogues show, besides a similarity in character, a likeness as regards physical properties and a proximity in atomic weight. Iron occupies a medium position amongst its nearest analogues, both with respect to properties and faculty of forming saline oxides, and also as regards atomic weight. On the one hand, cobalt, 58, and nickel, 59, approach iron, 56; they are metals of a more basic character, they do not form stable acids or higher degrees of oxidation, and are a transition to copper, 63, and zinc, 65. On the other hand, manganese, 55, and chromium, 52, are the nearest to iron; they form both basic and acid oxides, and are a transition to the metals possessing acid properties. In addition to having atomic weights approximately alike, chromium, manganese, iron, cobalt, nickel, and copper have also nearly the same specific gravity, so that the atomic volumes and the molecules of their analogous compounds are also near to one another (see table at the beginning of this volume). Besides this, the likeness between the above-mentioned elements is also seen from the following:

They form suboxides, RO, fairly energetic bases, isomorphous with magnesia--for instance, the salt RSO_{4},7H_{2}O, akin to MgSO_{4},7H_{2}O, and FeSO_{4},7H_{2}O, or to sulphates containing less water; with alkali sulphates all form double salts crystallising with 6H_{2}O; all are capable of forming ammonium salts, &c. The lower oxides, in the cases of nickel and cobalt, are tolerably stable, are not easily oxidised (the nickel compound with more difficulty than cobalt, a transition to copper); with manganese, and especially with chromium, they are more easily oxidised than with iron and pass into higher oxides. They also form oxides of the form R_{2}O_{3}, and with nickel, cobalt, and manganese this oxide is very unstable, and is more easily reduced than ferric oxide; but, in the case of chromium, it is very stable, and forms the ordinary kind of salts. It is isomorphous with ferric oxide, forms alums, is a feeble base, &c. Chromium, manganese, and iron are oxidised by alkali and oxidising agents, forming salts like Na_{2}SO_{4}; but cobalt and nickel are difficult to oxidise; their acids are not known with any certainty, and are, in all probability, still less stable than the ferrates. Cr, Mn and Fe form compounds R_{2}Cl_{6} which are like Fe_{2}Cl_{6} in many respects; in Co this faculty is weaker and in Ni it has almost disappeared. The cyanogen compounds, especially of manganese and cobalt, are very near akin to the corresponding ferrocyanides. The oxides of nickel and cobalt are more easily reduced to metal than those of iron, but those of manganese and chromium are not reduced so easily as iron, and the metals themselves are not easily obtained in a pure state; they are capable of forming varieties resembling cast iron. The metals Cr, Mn, Fe, Co, and Ni have a grey iron colour and are very difficult to melt, but nickel and cobalt can be melted in the reverberatory furnace and are more fusible than iron, whilst chromium is more difficult to melt than platinum (Deville). These metals decompose water, but with greater difficulty as the atomic weight rises, forming a transition to copper, which does not decompose water. All the compounds of these metals have various colours, which are sometimes very bright, especially in the higher stages of oxidation.

These metals of the iron group are often met with together in nature. Manganese nearly everywhere accompanies iron, and iron is always an ingredient in the ores of manganese. Chromium is found principally as chrome ironstone--that is, a peculiar kind of magnetic oxide in which Fe_{2}O_{3} is replaced by Cr_{2}O_{3}.

Nickel and cobalt are as inseparable companions as iron and manganese. The similarity between them even extends to such remote properties as magnetic qualities. In this series of metals we find those which are the most magnetic: iron, cobalt, and nickel. There is even a magnetic oxide among the chromium compounds, such being unknown in the other series. Nickel easily becomes passive in strong nitric acid. It absorbs hydrogen in just the same way as iron. In short, in the series Cr, Mn, Fe, Co, and Ni, there are many points in common although there are many differences, as will be seen still more clearly on becoming acquainted with cobalt and nickel.

In nature _cobalt_ is principally found in combination with arsenic and sulphur. _Cobalt arsenide_, or _cobalt speiss_, CoAs_{2}, is found in brilliant crystals of the regular system, principally in Saxony. _Cobalt glance_, CoAs_{2}CoS_{2}, resembles it very much, and also belongs to the regular system; it is found in Sweden, Norway, and the Caucasus. _Kupfernickel_ is a nickel ore in combination with arsenic, but of a different composition from cobalt arsenide, having the formula NiAs; it is found in Bohemia and Saxony. It has a copper-red colour and is rarely crystalline; it is so called because the miners of Saxony first mistook it for an ore of copper (_Kupfer_), but were unable to extract copper from it. _Nickel glance_, NiS_{2},NiAs_{2}, corresponding with cobalt glance, is also known. Nickel accompanies the ores of cobalt and cobalt those of nickel, so that both metals are found together. The ores of cobalt are worked in the Caucasus in the Government of Elizavetopolsk. Nickel ores containing aqueous hydrated nickel silicate are found in the Ural (Revdansk). Large quantities of a similar ore are exported into Europe from New Caledonia. Both ores contain about 12 per cent. Ni. _Garnierite_, (RO)_{5}(SiO_{2})_{4}1-1/2H_{2}O, where R = Ni and Mg, predominates in the New Caledonian ore. Large deposits of nickel have been discovered in Canada, where the ore (as nickelous pyrites) is free from arsenic. Cobalt is principally worked up into cobalt compounds, but nickel is generally reduced to the metallic state, in which it is now often used for alloys--for instance, for coinage in many European States, and for plating other metals, because it does not oxidise. Cobalt arsenide and cobalt glance are principally used for the preparation of cobalt compounds; they are first sorted by discarding the rocky matter, and then roasted. During this process most of the sulphur and arsenic disappears; the arsenious anhydride volatilises with the sulphurous anhydride and the metal also oxidises.[32] It is a simple matter to obtain nickel and cobalt from their oxides. In order to obtain the latter, solutions of their salts are treated with sodium carbonate and the precipitated carbonates are heated; the suboxides are thus obtained, and these latter are reduced in a stream of hydrogen, or even by heating with ammonium chloride. They easily oxidise when in the state of powder. When the chlorides of nickel and cobalt are heated in a stream of hydrogen, the metal is deposited in brilliant scales. _Nickel is always much more easily and quickly reduced than cobalt._ Nickel melts more easily than cobalt, and this even furnishes a means of testing the heating powers of a reverberatory furnace. Cobalt fuses at a temperature only a little lower than that at which iron does. In general, cobalt is nearer to iron than nickel, nickel being nearer to copper.[32 bis] Both nickel and cobalt have magnetic properties like iron, but Co is less magnetic than Fe, and Ni still less so. The specific gravity of nickel reduced by hydrogen is 9·1 and that of cobalt 8·9. Fused cobalt has a specific gravity of 8·5, the density of ordinary nickel being almost the same. Nickel has a greyish silvery-white colour; it is brilliant and very ductile, so that the finest wire may be easily drawn from it. This wire has a resistance to tension equal to iron wire. The beautiful colour of nickel, and the high polish which it is capable of receiving and retaining, as it does not oxidise, render it a useful metal for many purposes, and in many ways it resembles silver.[32 tri] It is now very common to cover other metals with a layer of nickel (nickel plating). This is done by a process of electro-plating, using a solution of a nickel salt. The colour of cobalt is dark and redder; it is also ductile, and has a greater tensile resistance than iron. Dilute acids act very slowly on nickel and cobalt; nitric acid may be considered as the best solvent for them. The solutions in every case contain salts corresponding with the ferrous salts--that is, the _salts_ CoX_{2}, NiX_{2}, _correspond with the suboxides_ of these metals. These salts in their types are similar to the magnesium salts. The salts of nickel when crystallising with water have a green colour, and form bright green solutions, but in the anhydrous state they most frequently have a yellow colour. The salts of cobalt are generally rose-coloured, and generally blue when in the anhydrous state. Their aqueous solutions are rose-coloured. Cobaltous chloride is easily soluble in alcohol, and forms a solution of an intense blue colour.[33]

[32] The residue from the roasting of cobalt ores is called _zafflor_,
and is often met with in commerce. From this the purer compounds
of cobalt may be prepared. The ores of nickel are also first
roasted, and the oxides dissolved in acid, nickelous salts being
then obtained.

The further treatment of cobalt and nickel ores is facilitated if
the arsenic can be almost entirely removed, which may be effected
by roasting the ore a second time with a small addition of nitre
and sodium carbonate; the nitre combines with the arsenic, forming
an arsenious salt, which may be extracted with water. The
remaining mass is dissolved in hydrochloric acid, mixed with a
small quantity of nitric acid. Copper, iron, manganese, nickel,
cobalt, &c., pass into solution. By passing hydrogen sulphide
through the solution, copper, bismuth, lead, and arsenic are
deposited as metallic sulphides; but iron, cobalt, nickel, and
manganese remain in solution. If an alkaline solution of bleaching
powder be then added to the remaining solution, the whole of the
manganese will first be deposited in the form of dioxide, then the
cobalt as hydrated cobaltic oxide, and finally the nickel also. It
is, however, impossible to rely on this method for effecting a
complete separation, the more so since the higher oxides of the
three above-mentioned metals have all a black colour; but, after a
few trials, it will be easy to find how much bleaching powder is
required to precipitate the manganese, and the amount which will
precipitate all the cobalt. The manganese may also be separated
from cobalt by precipitation from a mixture of the solutions of
both metals (in the form of the 'ous' salts) with ammonium
sulphide, and then treating the precipitate with acetic acid or
dilute hydrochloric acid, in which manganese sulphide is easily
soluble and cobalt sulphide almost insoluble. Further particulars
relating to the separation of cobalt from nickel may be found in
treatises on analytical chemistry. In practice it is usual to rely
on the rough method of separation founded on the fact that nickel
is more easily reduced and more difficult to oxidise than cobalt.
The New Caledonian ore is smelted with CaSO_{4} and CaCO_{3} on
coke, and a metallic regulus is obtained containing all the Ni,
Fe, and S. This is roasted with SiO_{2}, which converts all the
iron into slag, whilst the Ni remains combined with the S; this
residue on further roasting gives NiO, which is reduced by the
carbon to metallic Ni. The Canadian ore (a pyrites containing 11
p.c. Ni) is frequently treated in America (after a preliminary
dressing) by smelting it with Na_{2}SO_{4} and charcoal; the
resultant fusible Na_{2}S then dissolves the CuS and FeS_{2},
while the NiS is obtained in a bottom layer (Bartlett and
Thomson's process) from which Ni is obtained in the manner
described above.

For manufacturing purposes somewhat impure cobalt compounds are
frequently used, which are converted into _smalt_. This is glass
containing a certain amount of cobalt oxide; the glass acquires a
bright blue colour from this addition, so that when powdered it
may be used as a blue pigment; it is also unaltered at high
temperatures, so that it used to take the place now occupied by
Prussian blue, ultramarine, &c. At present smalt is almost
exclusively used for colouring glass and china. To prepare smalt,
ordinary impure cobalt ore (zaffre) is fused in a crucible with
quartz and potassium carbonate. A fused mass of cobalt glass is
thus formed, containing silica, cobalt oxide, and potassium oxide,
and a metallic mass remains at the bottom of the crucible,
containing almost all the other metals, arsenic, nickel, copper,
silver, &c. This metallic mass is called _speiss_, and is used as
nickel ore for the extraction of nickel. Smalt usually contains 70
p.c. of silica, 20 p.c. of potash and soda, and about 5 to 6 p.c.
of cobaltous oxide; the remainder consisting of other metallic
oxides.

[32 bis] All we know respecting the relations of Co and Ni to Fe and Cu
confirms the fact that Co is more closely related to Fe and Ni to
Cu; and as the atomic weight of Fe = 56 and of Cu = 63, then
according to the principles of the periodic system it would be
expected that the atomic weight of Co would be about 59-60, whilst
that of Ni should be greater than that of Co but less than that of
Cu, _i.e._ about 50·5-60·5. However, as yet the majority of the
determinations of the atomic weights of Co and Ni give a different
result and show that a lower atomic weight is obtained for Ni than
for Co. Thus K. Winkler (1894) obtained (employing metals
deposited electrolytically and determining the amount of iodine
which combined with them) Ni = 58·72 and Co = 59·37 (if H = 1 and
I = 126·53). In my opinion this should not be regarded as proving
that the principles of the periodic system cannot be applied in
this instance, nor as a reason for altering the position of these
elements in the system (_i.e._ by placing Ni after Fe, and Co next
to Cu), because in the first place the figures given by different
chemists (for instance, Zimmermann, Krüss, and others) are
somewhat divergent, and in the second place the majority of the
latest modes of determining the atomic weights of Co and Ni aim at
finding what weights of these metals react with known weights of
other elements without taking into account the faculty they have
of absorbing hydrogen; since this faculty is more developed in Ni
than in Co the hydrogen (occluded in Ni) should lower the atomic
weight of Ni more than that of Co. On the whole, the question of
the atomic weights of Co and Ni cannot yet be considered as
decided, notwithstanding the numerous researches which have been
made; still there can be no doubt that the atomic weights of these
two metals are very nearly equal, and greater than that of Fe, but
less than that of Cu. This question is of great interest, not only
for completing our knowledge of these metals, but also for
perfecting our knowledge of the periodic system of the elements.

[32 tri] For instance, the alkalis may be fused in nickel vessels as
well as in silver, because they have no action upon either metal.
Nickel, like silver, is not acted upon by dilute acids. Only
nitric acid dissolves both metals well. Nickel is harder, and
fuses at a higher temperature than silver. For castings, a small
quantity of magnesium (0·001 part by weight) is added to nickel to
render it more homogeneous (just as aluminium is added to steel).
Nickel forms many valuable alloys. Steel containing 3 p.c. Ni is
particularly valuable, its limit of elasticity is higher and its
hardness is greater; it is used for armour plate and other large
pieces. The alloys of nickel, especially with copper and zinc
(melchior, _see_ later), aluminium and silver, although used in
certain cases, are now replaced by nickel-plated or
nickel-deposited goods (deposited by electricity from a solution
of the ammonium salts).

[33] The change of colour is dependent in all probability on the
combination with water, or according to others on polymeric
transformation. It enables a solution of cobalt chloride to be
used as sympathetic ink. If something be written with cobalt
chloride on white paper, it will be invisible on account of the
feeble colour of the solution, and when dry nothing can be
distinguished. If, however, the paper be heated before the fire,
the rose-coloured salt will be changed into a less hydrous blue
salt, and the writing will become quite visible, but fade away
when cool.

The change of colour which takes place in solutions of CoCl_{2}
under the influence not only of solution in water or alcohol, but
also of a change of temperature, is a characteristic of all the
halogen salts of cobalt. Crystalline iodide of cobalt,
CoI_{2}6H_{2}O, gives a dark red solution between -22° and +20°;
above +20° the solution turns brown and passes from olive to
green, from +35° to 320° the solution remains green. According to
Étard the change of colour is due to the fact that at first the
solution contains the hydrate CoI_{2}H_{2}O, and that above 35° it
contains CoI_{2}4H_{2}O. These hydrates can be crystallised from
the solutions; the former at ordinary temperature and the latter
on heating the solution. The intermediate olive colour of the
solutions corresponds to the incipient decomposition of the
hexahydrated salt and its passage into CoI_{2}4H_{2}O. A solution
of the hexahydrated chloride of cobalt, CoCl_{2}6H_{2}O, is
rose-coloured between -22° and +25°; but the colour changes
starting from +25°, and passes through all the tints between red
and blue right up to 50°; a true blue solution is only obtained at
55° and remains up to 300°. This true blue solution contains
another hydrate, CoCl_{2}2H_{2}O.

The dependence between the solubility of the iodide and chloride
of cobalt and the temperature is expressed by two almost straight
lines corresponding to the hexa- and di-hydrates; the passage of
the one into the other hydrate being expressed by a curve. The
same character of phenomena is seen also in the variation of the
vapour tension of solutions of chloride of cobalt with the
temperature. We have repeatedly seen that aqueous solutions (for
instance, Chapter XXII., Note 23 for Fe_{2}Cl_{6}) deposit
different crystallo-hydrates at different temperatures, and that
the amount of water in the hydrate decreases as the temperature
_t_ rises, so that it is not surprising that CoCl_{2}2H_{2}O (or
according to Potilitzin CoCl_{2}H_{2}O) should separate out above
55° and CoCl_{2}6H_{2}O at 25° and below. Nor is it exceptional
that the colour of a salt varies according as it contains
different amounts of H_{2}O. But in this instance it is
characteristic that the change of colour takes place in solution
in the presence of an excess of water. This apparently shows that
the actual solution may contain either CoCl_{2}6H_{2}O or
CoCl_{2}2H_{2}O. And as we know that a solution may contain both
metaphosphoric PHO_{3} and orthophosphoric acid H_{3}PO_{4} =
HPO_{3} + H_{2}O, as well as certain other anhydrides, the
question of the state of substances in solutions becomes still
more complicated.

Nickel sulphate crystallises from neutral solutions at a
temperature of from 15° to 20° in _rhombic_ crystals containing
7H_{2}O. Its form approaches very closely to that of the salts of
zinc and magnesium. The planes of a vertical prism for magnesium
salts are inclined at an angle of 90° 30´, for zinc salts at an
angle of 91° 7´, and for nickel salts at an angle of 91° 10´. Such
is also the form of the zinc and magnesium selenates and
chromates. Cobalt sulphate containing 7 molecules of water is
deposited in crystals of the _monoclinic_ system, like the
corresponding salts of iron and manganese. The angle of a vertical
prism for the iron salt = 82° 20´, for cobalt = 82° 22´, and the
inclination of the horizontal pinacoid to the vertical prism for
the iron salt = 99° 2´, and for the cobalt salt 99° 36´. All the
isomorphous mixtures of the salts of magnesium, iron, cobalt,
nickel and manganese have the same form if they contain 7 mol.
H_{2}O and iron or cobalt predominate, whilst if there is a
preponderance of magnesium, zinc, or nickel, the crystals have a
rhombic form like magnesium sulphate. Hence these sulphates are
_dimorphous_, but for some the one form is more stable and for
others the other. Brooke, Moss, Mitscherlich, Rammelsberg, and
Marignac have explained these relations. Brooke and Mitscherlich
also supposed that NiSO_{4},7H_{2}O is not only capable of
assuming these forms, but also that of the _tetragonal_ system,
because it is deposited in this form from acid, and especially
from slightly-heated solutions (30° to 40°). But Marignac
demonstrated that the tetragonal crystals do not contain 7, but 6,
molecules of water, NiSO_{4},6H_{2}O. He also observed that a
solution evaporated at 50° to 70° deposits monoclinic crystals,
but of a different form from ferrous sulphate,
FeSO_{4},7H_{2}O--namely, the angle of the prism is 71° 52´, that
of the pinacoid 95° 6´. This salt appears to be the same with 6
molecules of water as the tetragonal. Marignac also obtained
magnesium and zinc salts with 6 molecules of water by evaporating
their solutions at a higher temperature, and these salts were
found to be isomorphous with the monoclinic nickel salt. In
addition to this it must be observed that the rhombic crystals of
nickel sulphate with 7H_{2}O become turbid under the influence of
heat and light, lose water, and change into the tetragonal salt.
The monoclinic crystals in time also become turbid, and change
their structure, so that the tetragonal form of this salt is the
most stable. Let us also add that nickel sulphate in all its
shapes forms very beautiful emerald green crystals, which, when
heated to 230°, assume a dirty greenish-yellow hue and then
contain one molecule of water.

Klobb (1891) and Langlot and Lenoir obtained anhydrous CoSO_{4}
and NiSO_{4} by igniting the hydrated salt with (NH_{4})_{2}SO_{4}
until the ammonium salt had completely volatilised and decomposed.

We may add that when equivalent aqueous solutions of NiX_{2}
(green) and CoX_{2} (red) are mixed together they give an almost
colourless (grey) solution, in which the green and red colour of
the component parts disappears owing to the combination of the
complementary colours.

A double salt NiKF_{3} is obtained by heating NiCl_{2} with KFHF
in a platinum crucible; KCoF_{3} is formed in a similar manner.
The nickel salt occurs in fine green plates, easily soluble in
water but scarcely soluble in ethyl and methyl alcohol. They
decompose into green oxide of nickel and potassium fluoride when
heated in a current of air. The analogous salt of cobalt
crystallises in crimson flakes.

If instead of potassium fluoride, CoCl_{2} or NiCl_{2} be fused
with ammonium fluoride, they also form double salts with the
latter. This gives the possibility of obtaining anhydrous
fluorides NiF_{2} and CoF_{2}. Crystalline fluoride of nickel,
obtained by heating the amorphous powder formed by decomposing the
double ammonium salt in a stream of hydrofluoric acid, occurs in
beautiful green prisms, sp. gr. 4·63, which are insoluble in
water, alcohol, and ether; sulphuric, hydrochloric, and nitric
acids also have no action upon them, even when heated; NiF_{2} is
decomposed by steam, with the formation of black oxide, which
retains the crystalline structure of the salt. Fluoride of cobalt,
obtained as a rose-coloured powder by decomposing the double
ammonium salt with the aid of heat in a stream of hydrofluoric
acid, fuses into a ruby-coloured mass which bears distinct signs
of a crystalline structure; sp. gr. 4·43. The molten salt only
volatilises at about 1400°, which forms a clear distinction
between CoF_{2} and the volatile NiF_{2}. Hydrochloric, sulphuric,
and nitric acids act upon CoF_{2} even in the cold, although
slowly, while when heated the reaction proceeds rapidly (Poulenc,
1892).