Potassium ferrocyanide, K₄Fe(CN)₆

Potassium ferrocyanide, K₄Fe(CN)₆.3H₂O, is the most important salt of ferrocyanic acid, and is known in commerce by the more familiar name of yellow prussiate of potash. It results when a solution of potassium cyanide is added in excess to one of a ferrous salt, so that the precipitate first formed completely redissolyes. It may also be prepared by allowing iron to dissolve in an air-free solution of potassium cyanide: -

6KCN + Fe + 2H₂O = K₄[Fe(CN)₆] + 2KOH + H₂, and by the action of potassium hydroxide solution on ferrous cyanide: - 3Fe(CN)₂ + 4KOH = K₄Fe(CN)₆ + 2Fe(OH)₂.

The old commercial method of preparing the salt lay in heating nitrogenous material, such as horn, wool, feathers, blood, etc., with potash and iron turnings. The mass was ultimately heated to fusion to complete the reaction, cooled, and extracted with boiling water. The solution contained potassium ferrocyanide, thiocyanate, carbonate, and sulphide. The first-named was crystallised out, but the yield was seldom more than 20 per cent, of the quantity theoretically obtainable from the nitrogen content of the organic material consumed.

The chemistry of the reactions involved has been made the subject of considerable study.

Another method consists in passing the vapour of trimethylamine into a retort at red heat. The resulting products are passed into sulphuric acid, whereby ammonium cyanide is converted into hydrogen cyanide, which is now absorbed in potash to yield the corresponding cyanide. Ferrous hydroxide, prepared by addition of milk of lime to a solution of ferrous chloride, is added to the cyanide solution, and the liquid, after filtering, deposits a relatively pure crop of potassium ferrocyanide.

Alkali thiocyanates may be made the starting-point for the preparation of ferrocyanides. The potassium salt is mixed with twice the weight of iron filings necessary to form ferrous sulphide, and with double the quantity of ferrous hydroxide, in a freshly precipitated Condition, to form ferrocyanide. The mixture is maintained for some twelve hours under agitation in a closed vessel at 110° to 120° C. Potassium ferrocyanide is formed, together with Prussian blue, and extracted from the residue with water.

From 1885 to 1895 potassium ferrocyanide was manufactured very largely from spent oxide of iron, used in purifying coal gas from hydrogen sulphide. Coal gas, as it leaves the retorts, contains hydrogen cyanide, formed by the action of ammonia on red-hot carbon. Thus: -

NH₃ + C = HCN + H₂.

The hydrogen cyanide yields ammonia again in contact with heated water vapour: -

HCN + H₂O = NH₃ + CO,

the ammonia uniting with the free acid to form ammonium cyanide.

The relative proportions of ammonia and cyanide in coal gas thus vary with the conditions.

The oxide of iron employed in gas works is the hydrated ferric compound. It absorbs the hydrogen sulphide in the coal gas, yielding ferrous and ferric sulphides; the ammonium cyanide, also present in the gas, reacts with the iron compounds, forming ferrocyanide and, ultimately, Prussian blue. The last-named was frequently converted into calcium ferrocyanide by intimate admixture with lime, and dissolved in cold water. Potassium chloride solution was next added, whereby the potassium calcium ferrocyanide, K₂CaFe(CN)₆, was precipitated, to be converted into potassium ferrocyanide by boiling with potassium carbonate.

Wet methods are now largely employed for the preparation of potassium ferrocyanide from coal gas, and, whilst several have been patented and found to work well, it will suffice, for present purposes, to mention one method only.

The crude coal gas is washed with ferrous sulphate solution, whereby the latter is converted into a suspension of ferrous sulphide in ammonium sulphate solution. This reacts with the ammonium cyanide, yielding ferrous ferrocyanide, Fe₂[Fe(CN)₆], or ammonium ferrous ferrocyanide, (NH₄)₂Fe[Fe(CN)₆], according to circumstances. Potassium ferrocyanide may be obtained from these by treatment with lime, as in the spent oxide process. By repeatedly dissolving in water and precipitating with alcohol, the salt can be obtained in a very pure state.

What appeared to be two forms of the salt, designated as α and β respectively, were prepared by Briggs. The former was obtained from the pure commercial salt by dissolving it in water with 1 per cent, of its weight of potassium cyanide. After twenty-four hours alcohol was . added, and a white crystalline precipitate of the α salt obtained. Upon recrystallisation, if the crystals were large, they were lemon-yellow in colour, but otherwise quite white; density at 20° C. 1.889. At 20° C. 100 grams of saturated solution contained 25.0 grams of the α trihydrated salt, K₄[Fe(CN)₆].3H₂O. The β salt was obtained by dissolving the pure commercial salt in water with 1 per cent, of its weight of dilute acetic acid (1 part acid, 10 parts water), and allowing to stand in the absence of air. After twenty-four hours the addition of alcohol yielded a cream-coloured precipitate. Upon recrystallisation orange-coloured crystals were obtained, of more intense colour than the a salt; of slightly less density, namely 1-882 at 20° C., and rather less soluble in water, 100 grams of saturated solution at 20° C. containing 24.6 grams of the salt.

Solutions of the α salt, upon prolonged standing, are converted into the β variety, a process that is hastened by the addition of 1 per cent, of acetic acid. Cyanides, alkalies, and ammonia induce the reverse transformation.

Bennett, however, having prepared these . two isomerides according to the directions given by Briggs, and measured their angles, concluded that the two forms are identical crystallographically. He drew attention to the fact that the more intense colour of the β salt might easily arise from slight decomposition induced by its acid method of preparation; in support of this, it is significant that the most striking difference of colour occurs in the ammonium salt, which is the most unstable of all, its aqueous solution being decomposed on simply warming. The differences in density and solubility are too slight to be very definite evidence either way. Again, Kolthoff found no difference in the physical properties of solutions of α and β ferrocyanide, except their colour.

Briggs therefore reinvestigated the matter, and showed that the supposed isomeride or β variety consists of mixed crystals of potassium ferrocyanide and aquopentacyanoferrite, K₃[Fe(CN)₅H₂O], the amount of the latter being too small to detect by qualitative analysis. Since the β salt is formed when the a variety is repeatedly recrystallised from water, it is concluded that the compounds are, in aqueous solution, in a state of equilibrium, thus: -

H₂O + K₄[Fe(CN)₆] = K₃[Fe(CN)₅.H₂O] + KCN.

It must not be overlooked, however, that isomerism has been definitely discovered in the case of tetramethyl ferrocyanide, (CH₃)₄Fe(CN)₆, and it is always possible that isomerides of inorganic salts of ferrocyanic acid may be capable of existence.

Potassium ferrocyanide crystallises with three molecules of water, which are completely expelled at 110° C., leaving the anhydrous salt in the form of a white powder. Crystals of potassium ferrocyanide, in common with the few other hydrated crystals which have been examined, are permeable to water vapour. This has been demonstrated by cementing crystals with wax into the necks of small flasks containing phosphorus pentoxide, and noting any alteration in weight after exposure to moist air. An increase was observed, which was slow but steady and continuous. It is not suggested that the crystals are porous in the ordinary mechanical sense in which, for example, unglazed earthenware is porous. Rather is it believed that the layers of the crystal inside the flask give up their moisture to the dry air in contact with the phosphorus pentoxide; these dehydrated layers take up water from the next layers, and so on, until the layers are reached in direct contact with the moist air. The net result is thus that the water passes through the crystal from the wet to the dry atmosphere.

The crystals of potassium ferrocyanide are tough and difficult to powder. They are non-poisonous, but act as an aperient. When the dry anhydrous salt is heated to incipient fusion in a vacuum no gas is evolved, potassium cyanide and ferrous potassium ferrocyanide being produced: -

2K₄Fe(CN)₆ = FeK₂Fe(CN)₆ + 6KCN.

At red heat the latter salt is further decomposed, evolving free cyanogen. Thus: -

FeK₂Fe(CN)₆ = 2Fe + 2KCN + 2C₂N₂.

This is a convenient method of preparing potassium cyanide, although wasteful in so far as the nitrogen is concerned, in that one-third of this element is lost as cyanogen.

The solubility of potassium ferrocyanide in water has been determined by a number of investigators, but the results do not harmonise. The only reliable figures appear to be those of Briggs, quoted above; of Harkins and Pearee, namely, that at 25° C. 100 grams of water dissolve 24.796 grams of K₄Fe(CN)₆; and of Grube, that one litre of saturated solution at 25° C. contains 319.4 grams of K₄Fe(CN)₆.3H₂O. The solubility is increased by the presence of sodium ferrocyanide.

The density of potassium ferrocyanide solution at 8.9° C., saturated and in contact with crystals of the salt, is 1.1191, and at 25° С. 1.09081. The contractions resulting when given volumes of potassium ferrocyanide solutions are mixed with equal volumes of water have been measured by Wade.

Solutions of potassium ferrocyanide are gradually decomposed by light, ferric hydroxide being precipitated. Addition of an alkali sulphide to a fresh solution of the ferrocyanide effects the gradual precipitation of ferrous sulphide in sunlight but not in the dark. The reason appears to be as follows: -

1. Potassium ferrocyanide dissociates normally in solution into potassium and ferrocyanogen ions: -
   K₄Fe(CN)₆ ⇔ 4K^• + Fe(CN)₆''''.
2. Under the influence of light the ferrocyanogen ion dissociates thus: -
   Fe(CN)₆'''' ⇔ Fe^•• + 6CN'.

This reaction is reversible, proceeding from right to left in the dark. It is these iron ions which are precipitated as ferrous sulphide, in £he presence of an alkali sulphide, or as ferric hydroxide in neutral or alkaline solution by the action of atmospheric oxygen.

Upon prolonged exposure to light, a solution of potassium ferrocyanide deposits Prussian blue, whilst on continued boiling ammonia is evolved. With ferrous salts it yields an immediate white precipitate of ferrous potassium ferrocyanide, K₂Fe[Fe(CN)₆], which readily absorbs oxygen, becoming blue. The presence of dilute hydrochloric or sulphuric acid, or the employment of excess of the ferrous salt, accelerates the formation of the blue colour, and the reaction is exceedingly delicate.

With ferric salts Prussian blue is obtained, but the reaction is a time-reaction, and is retarded by the presence both of acids and salts. In very dilute solution no colour may be produced, or only an indefinite green after several hours, although the blue colour may be produced by the addition of concentrated solutions of many salts to the dilute solutions of the reagents. Probably the explanation lies in the precipitation by the salts of Prussian blue, which is first formed in colourless colloidal form.

Carbon dioxide decomposes potassium ferrocyanide solution at 72° to 74° C., liberating hydrogen cyanide and precipitating ferrous potassium ferrocyanide. Continued passage of carbon dioxide through a boiling solution of potassium ferrocyanide results in the precipitation of ferric hydroxide and the formation of potassium carbonate and hydrogen cyanide, or its decomposition products, ammonia and formaldehyde.

When potassium ferrocyanide is heated with concentrated sulphuric acid, carbon monoxide is evolved. This reaction has been known for many years, but it was not until 1900 that the reaction was thoroughly investigated.

Concentrated sulphuric acid dissolves dry anhydrous potassium ferrocyanide, yielding potassium hydrogen sulphate and hydrogen ferrocyanide: -

K₄Fe(CN)₆ + 4H₂SO₄ = 4KHSO₄ + H₄Fe(CN)₆.

This solution is decomposed on warming, carbon monoxide being evolved, although even at 200° C. the rate of evolution is slow.

If a little water is present, carbon monoxide is readily formed on warming, the best result being attained with acid of concentration corresponding to H₂SO₄.2H₂O, when dry anhydrous potassium ferrocyanide is employed; the reaction is then complete at 180° C.: -

K₄[Fe(CN)₆] + 8(H₂SO₄.2H₂O) = 4KHSO₄ + FeSO₄ + 3(NH₄)₂SO₄ + 6CO + 10H₂O.

This is a very convenient method of preparing pure carbon monoxide.

When heated with dilute sulphuric acid hydrogen cyanide is evolved, and Everitt's salt remains behind: -

2K₄[Fe(CN)₆] + 8H₂SO₄ = 8K₂SO₄ + K₂Fe^••[Fe^••(CN)₆](Everitt's salt) + 6HCN

Its aqueous solution upon saturation with chlorine darkens in colour, and upon concentration yields potassium ferricyanide, K₃[Fe^•••(CN)₆]: -

2K₄[Fe^••(CN)₆] + Cl₂ ⇔ 2KCl + 2K₃[Fe^•••(CN)₆].

Potassium ferrocyanide is similarly oxidised by potassium bromate in acid solution; thus: -

6K₄Fe(CN)₆ + KBrO₃ + 6HCl = 6K₃Fe(CN)₆ + KBr + 6KCl + 3H₂O,

the reaction being quantitative under certain well-defined conditions. It probably takes place in three stages, namely: -

1. 2KBrO₃ + 12HCl = 2KCl + Br₂ + 5Cl₂ + 6H₂O.
2. 2K₄Fe(CN)₆ + Br₂ = 2KBr + 2K₃Fe(CN)₆.
3. 10K₄Fe(CN)₆ + 5Cl₂ = 10KCl + 10K₃Fe(CN)₆.

With potassium permanganate oxidation to ferricyanide proceeds quantitatively in acid solution, the reaction affording a useful volumetric method of estimating ferrocyanides if carried out under certain well-defined conditions.

Mixed solutions of potassium ferrocyanide and p-nitroso dimethyl aniline are at first yellow in colour, but become green very rapidly on exposure to light, due, perhaps, to the production of colloidal Prussian blue under the influence of a catalyst generated by the nitroso compound. Solutions of potassium ferrocyanide are catalytically decomposed when boiled with cuprous chloride in the presence of hydrochloric acid. The action appears to consist in the alternate formation of cuprous cyanide and regeneration of cuprous chloride, hydrogen cyanide being evolved. By collecting the evolved acid in alkali, and afterwards titrating excess of the latter, the amount of ferrocyanide originally present may be conveniently estimated.

If a few drops of potassium ferrocyanide solution are added to dilute hydrogen peroxide (1 per cent.), and kept in the dark, decomposition of the latter is exceedingly slow. On placing in direct sunlight for a few moments, however, brisk evolution of oxygen takes place and continues, even after removal from the light. The effect is not due to rise of temperature, but, presumably, to some catalyst generated under the influence of the light.

Potassium ferrocyanide finds application in commerce in the manufacture of Prussian blue, and also for case-hardening of steel.

When equal parts of potassium ferrocyanide and ammonium chloride are dissolved in water at 100° C. and allowed to cool, pale yellow hexagonal crystals are obtained of composition NH₄KH₂Fe(CN)₆.2NH₄Cl. The neutral salt, (NH₄)₃KFe(CN)₆.2NH₄Cl, has also been obtained.