Chemical elements
  Iron
    History of Iron
    Mineralogy
    Isotopes
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    Application
    Physical Properties
    Chemical Properties
      Iron Hydride
      Ferrous fluoride
      Aluminium pentafluoferrite
      Ferric fluoride
      Ammonium ferrifluoride
      Barium ferrifluoride
      Potassium ferrifluoride
      Sodium ferrifluoride
      Thallous ferrifluoride
      Ferrous diferrifluoride
      Ferrous monoferrifluoride
      Ferrous chloride
      Ammonium tetrachlorferrite
      Ferric chloride
      Tetrachlorferrates
      Pentachlorferrates
      Ferroso-ferric chloride
      Ferrous perchlorate
      Ferric perchlorate
      Ferrous chlorate
      Ferric chlorate
      Ferrous Oxychlorides
      Ferrous bromide
      Ferric bromide
      Ferric chloro-bromide
      Ferrous bromate
      Ferrous iodide
      Ferric iodide
      Ferric iodate
      Ferrous oxide
      Ferrous hydroxide
      Triferric tetroxide
      Ferric oxide
      Ferrous acid
      Calcium ferrite
      Cobalt ferrite
      Cupric ferrite
      Cuprous ferrite
      Magnesium ferrite
      Nickel ferrite
      Potassium ferrite
      Sodium ferrite
      Zinc ferrite
      Barium ferrate
      Strontium ferrate
      Barium perferrate
      Calcium perferrate
      Potassium perferrate
      Sodium perferrate
      Strontium perferrate
      Iron Subsulphides
      Ferrous sulphide
      Ferric sulphide
      Potassium ferric sulphide
      Sodium ferric sulphide
      Cuprous ferric sulphide
      Iron disulphide
      Ferrous sulphite
      Ferric sulphite
      Potassium ferri-tetrasulphite
      Potassium ferri-disulphite
      Potassium ferri-sulphite
      Ammonium ferri-sulphite
      Sodium ferri-disulphite
      Sodium hydrogen ferri-tetrasulphite
      Ferrous sulphate
      Ferrous copper sulphate Fe
      Ferrous ammonium sulphate
      Ferrous potassium sulphate
      Ferrous aluminium sulphate
      Basic ferrous sulphate
      Ferric sulphate
      Ammonium ferri-disulphate
      Trisodium ferri-trisulphate
      Ferric Alums
      Ferric ammonium alum
      Ferric potassium alum
      Ferric rubidium alum
      Ferroso-ferric sulphate
      Ferrous amido-sulphonate
      Ferric amido-sulphonate
      Ferrous thiosulphate
      Ferrous pyrosulphate
      Ferrous tetrathionate
      Ferric selenide
      Iron diselenide
      Iron Selenites
      Ferrous selenate
      Ferric rubidium selenium alum
      Ferric caesium selenium alum
      Ferric tellurite
      Ferrous chromite
      Ferrous chromate
      Iron nitride
      Nitro-Iron
      Ferrous nitrate
      Ferric nitrate
      Ferrous Nitroso Salts
      Potassium ferro-heptanitroso sulphide
      Sodium ferro-heptanitroso sulphide
      Ammonium ferro-heptanitroso sulphide
      Tetramethyl ammonium ferro-heptanitroso sulphide
      Ferro-dinitroso Sulphides
      Potassium ferro-dinitroso thiosulphate
      Triferro phosphide
      Diferro phosphide
      Iron monophosphide
      Iron sesqui-phosphide
      Ferrous hypophosphite
      Ferric hypophosphite
      Ferrous phosphite
      Ferric phosphite
      Ferrous orthophosphate
      Ferrous hydrogen orthophosphate
      Ferrous dihydrogen orthophosphate
      Ferric orthophosphate
      Sodium ferri-diorthophosphate
      Ammonium ferri-diorthophosphate
      Sodium ferri-triorthophosphate
      Ferric dihydrogen orthophosphate
      Acid ferric orthophosphate
      Ferrous metaphosphate
      Ferric metaphosphate
      Ferrous pyrophosphate
      Ferric pyrophosphate
      Hydrogen ferri-pyrophosphate
      Sodium ferro-pyrophosphate
      Ferrous thio-orthophosphite
      Ferrous thio-orthophosphate
      Ferrous thio-pyrophosphite
      Ferrous thio-pyrophosphate
      Iron sub-arsenide
      Iron mon-arsenide
      Iron sesqui-arsenide
      Iron di-arsenide
      Iron thio-arsenide
      Ferrous met-arsenite
      Ferric arsenite
      Ferrous ortho-arsenate
      Ferric ortho-arsenate
      Ferro mono-antimonide
      The di-antimonide
      Ferrous thio-antimonite
      Ferric ortho-antimonate
      Triferro carbide
      Diferro carbide
      Iron dicarbide
      Iron pentacarbonyl
      Diferro nonacarbonyl
      Iron tetracarbonyl
      Ferrous carbonate
      Ferrous bicarbonate
      Ferrous potassium carbonate
      Complex Iron Carbonates
      Ferrous thiocarbonate
      Ferrous thiocarbonate hexammoniate
      Ferrous cyanide
      Ferro-cyanic acid
      Aluminium ferrocyanide
      Aluminium ammonium ferrocyanide
      Ammonium ferrocyanide
      Barium ferrocyanide
      Calcium ferrocyanide
      Calcium ammonium ferrocyanide
      Cobalt ferrocyanide
      Copper ferrocyanide
      Ammonium cuproferrocyanide
      Barium cuproferrocyanide
      Lithium cuproferrocyanide
      Magnesium cuproferrocyanide
      Potassium cuproferrocyanide
      Sodium cuproferrocyanide
      Ammonium cupriferrocyanide
      Potassium cupriferrocyanide
      Potassium ferrous cupriferrocyanide
      Sodium cupriferrocyanide
      Strontium cupriferrocyanide
      Lithium ferrocyanide
      Magnesium ferrocyanide
      Magnesium ammonium ferrocyanide
      Manganese ferrocyanide
      Nickel ferrocyanide
      Potassium ferrocyanide
      Potassium aluminium ferrocyanide
      Potassium barium ferrocyanide
      Potassium calcium ferrocyanide
      Potassium cerium ferrocyanide
      Potassium magnesium ferrocyanide
      Potassium mercuric ferrocyanide
      Silver ferrocyanide
      Sodium ferrocyanide
      Sodium cerium ferrocyanide
      Strontium ferrocyanide
      Thallium ferrocyanide
      Zinc potassium ferrocyanide
      Ferricyanic acid
      Ammonium ferricyanide
      Barium ferricyanide
      Barium potassium ferricyanide
      Calcium ferricyanide
      Calcium potassium ferricyanide
      Cobalt ferricyanide
      Copper ferricyanide
      Lead ferricyanide
      Magnesium ferricyanide
      Mercuric ferricyanide
      Mercurous ferricyanide
      Potassium ferricyanide
      Sodium ferricyanide
      Strontium ferricyanide
      Zinc ferricyanide
      Ferrous hydrogen ferrocyanide
      Ferrous potassium ferrocyanide
      Prussian Blues
      Ferrous ferrocyanide
      Ferric ammonium ferrocyanide
      Nitroprussic acid
      Sodium nitroprusside
      Ammonium nitroprusside
      Barium nitroprusside
      Cobalt nitroprusside
      Nickel nitroprusside
      Potassium nitroprusside
      Carbonyl Penta-Ferrocyanides
      Carbonyl ferrocyanic acid
      Barium carbonyl ferrocyanide
      Copper carbonyl ferrocyanide
      Ferric carbonyl ferrocyanide
      Potassium carbonyl ferrocyanide
      Silver carbonyl ferrocyanide
      Sodium carbonyl ferrocyanide
      Strontium carbonyl ferrocyanide
      Uranyl carbonyl ferrocyanide
      Sodium ammonio ferrocyanide
      Potassium aquo ferrocyanide
      Potassium aquo ferricyanide
      Sodium aquo penta-ferricyanide
      Potassium sulphito ferrocyanide
      Ferrous thiocyanate
      Ferric thiocyanate
      Sodium ferrothiocyanate
      Sodium ferrithiocyanate
      Potassium ferrithiocyanate
      Iron subsilicide
      Iron monosilicide
      Iron disilicide
      Triferro disilicide
      Ferrous orthosilicate
      Ferrous magnesium orthosilicate
      Ferrous metasilicate
      Ferric silicate
      Diferro boride
      Iron monoboride
      Iron diboride
      Ferrous chlorborate
      Ferrous bromborate
    Corrosion
    Iron Salts
    PDB 101m-1aeb
    PDB 1aed-1awd
    PDB 1awp-1beq
    PDB 1bes-1c53
    PDB 1c6o-1ci6
    PDB 1cie-1cry
    PDB 1csu-1dfx
    PDB 1dgb-1dry
    PDB 1ds1-1e08
    PDB 1e0z-1ehj
    PDB 1ehk-1f5o
    PDB 1f5p-1fnp
    PDB 1fnq-1fzi
    PDB 1g08-1gnl
    PDB 1gnt-1h43
    PDB 1h44-1hdb
    PDB 1hds-1i5u
    PDB 1i6d-1iwh
    PDB 1iwi-1jgx
    PDB 1jgy-1k2o
    PDB 1k2r-1kw6
    PDB 1kw8-1lj0
    PDB 1lj1-1m2m
    PDB 1m34-1mko
    PDB 1mkq-1mun
    PDB 1muy-1n9x
    PDB 1naz-1nx4
    PDB 1nx7-1ofe
    PDB 1off-1p3t
    PDB 1p3u-1pmb
    PDB 1po3-1qmq
    PDB 1qn0-1ra0
    PDB 1ra5-1rxg
    PDB 1ry5-1smi
    PDB 1smj-1t71
    PDB 1t85-1u8v
    PDB 1u9m-1uyu
    PDB 1uzr-1vxf
    PDB 1vxg-1wri
    PDB 1wtf-1xlq
    PDB 1xm8-1y4r
    PDB 1y4t-1ygd
    PDB 1yge-1z01
    PDB 1z02-2a9e
    PDB 2aa1-2azq
    PDB 2b0z-2boz
    PDB 2bpb-2ca3
    PDB 2ca4-2cz7
    PDB 2czs-2dyr
    PDB 2dys-2ewk
    PDB 2ewu-2fwl
    PDB 2fwt-2gl3
    PDB 2gln-2hhb
    PDB 2hhd-2ibn
    PDB 2ibz-2jb8
    PDB 2jbl-2mgh
    PDB 2mgi-2o01
    PDB 2o08-2ozy
    PDB 2p0b-2q0i
    PDB 2q0j-2r1h
    PDB 2r1k-2spm
    PDB 2spn-2vbd
    PDB 2vbp-2vzb
    PDB 2vzm-2wiv
    PDB 2wiy-2xj5
    PDB 2xj6-2ylj
    PDB 2yrs-2zon
    PDB 2zoo-3a17
    PDB 3a18-3aes
    PDB 3aet-3bnd
    PDB 3bne-3cir
    PDB 3ciu-3dax
    PDB 3dbg-3e1p
    PDB 3e1q-3eh4
    PDB 3eh5-3fll
    PDB 3fm1-3gas
    PDB 3gb4-3h57
    PDB 3h58-3hrw
    PDB 3hsn-3ir6
    PDB 3ir7-3k9y
    PDB 3k9z-3l4p
    PDB 3l61-3lxi
    PDB 3lyq-3mm8
    PDB 3mm9-3n62
    PDB 3n63-3nlo
    PDB 3nlp-3o0f
    PDB 3o0r-3p6o
    PDB 3p6p-3prq
    PDB 3prr-3sel
    PDB 3sik-3una
    PDB 3unc-4blc
    PDB 4cat-4erg
    PDB 4erm-4nse
    PDB 4pah-8cat
    PDB 8cpp-9nse

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.


Double Salts of Potassium ferrocyanide

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.
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