Chromium

Chromium (symbol Cr. atomic weight 52.1), one of the metallic chemical elements, the name being derived from the fine colour (Gr. χρῶμα) of its compounds. It is a member of the sixth group in the periodic classification of the elements, being included in the natural family of elements containing molybdenum, tungsten and uranium. The element is not found in the free state in nature, nor to any large extent in combination, occurring chiefly as chrome-ironstone, Cr2O3·FeO, and occasionally being found as crocoisite, PbCrO4, chrome-ochre, Cr2O3, and chrome-garnet, CaO·Cr2O3·3SiO2, while it is also the cause of the colour in serpentine, chrome-mica and the emerald. It was first investigated in 1789 by L.N. Vauquelin and Macquart, and in 1797 by Vauquelin, who found that the lead in crocoisite was in combination with an acid, which he recognized as the oxide of a new metal.

The metal can be obtained by various processes. Thus Sainte Claire Deville prepared it as a very hard substance of steel-grey colour, capable of taking a high polish, by strong ignition of chromic oxide and sugar charcoal in a lime crucible. F. Wöhler reduced the sesquioxide by zinc, and obtained a shining green powder of specific gravity 6.81, which tarnished in air and dissolved in hydrochloric acid and warm dilute sulphuric acid, but was unacted upon by concentrated nitric acid. H. Moissan (Comptes rendus, 1893, 116, p. 349; 1894, 119, p. 185) reduces the sesquioxide with carbon, in an electric furnace; the product so obtained (which contains carbon) is then strongly heated with lime, whereby most of the carbon is removed as calcium carbide, and the remainder by heating the purified product in a crucible lined with the double oxide of calcium and chromium. An easier process is that of H. Goldschmidt (Annalen, 1898, 301, p. 19) in which the oxide is reduced by metallic aluminium; and if care is taken to have excess of the sesquioxide of chromium present, the metal is obtained quite free from aluminium. The metal as obtained in this process is lustrous and takes a polish, does not melt in the oxyhydrogen flame, but liquefies in the electric arc, and is not affected by air at ordinary temperatures. Chromium as prepared by the Goldschmidt process is in a passive condition as regards dilute sulphuric acid and dilute hydrochloric acid at ordinary temperatures; but by heating the metal with the acid it passes into the active condition, the same effect being produced by heating the inactive form with a solution of an alkaline halide. W. Hittorf thinks that two allotropic forms of chromium exist (Zeit. für phys. Chem., 1898, 25, p. 729; 1899, 30, p. 481; 1900, 34, p. 385), namely active and inactive chromium; while W. Ostwald (ibid., 1900, 35, pp. 33, 204) has observed that on dissolving chromium in dilute acids, the rate of solution as measured by the evolution of gas is not continuous but periodic. It is largely made as ferro-chrome, an alloy containing about 60-70% of chromium, by reducing chromite in the electric furnace or by aluminium.

Chromium and its salts may be detected by the fact that they give a deep green bead when heated with borax, or that on fusion with sodium carbonate and nitre, a yellow mass of an alkaline chromate is obtained, which, on solution in water and acidification with acetic acid, gives a bright yellow precipitate on the addition of soluble lead salts. Sodium and potassium hydroxide solutions precipitate green chromium hydroxide from solutions of chromic salts; the precipitate is soluble in excess of the cold alkali, but is completely thrown down on boiling the solution. Chromic acid and its salts, the chromates and bichromates, can be detected by the violet coloration which they give on addition of hydrogen peroxide to their dilute acid solution, or by the fact that on distillation with concentrated sulphuric acid and an alkaline chloride, the red vapours of chromium oxychloride are produced. The yellow colour of normal chromates changes to red on the addition of an acid, but goes back again to yellow on making the solution alkaline. Normal chromates on the addition of silver nitrate give a red precipitate of silver chromate, easily soluble in ammonia, and with barium chloride a yellow precipitate of barium chromate, insoluble in acetic acid. Reducing agents, such as sulphurous acid and sulphuretted hydrogen, convert the chromates into chromic salts. Chromium in the form of its salts may be estimated quantitatively by precipitation from boiling solutions with a slight excess of ammonia, and boiling until the free ammonia is nearly all expelled. The precipitate obtained is filtered, well washed with hot water, dried and then ignited until the weight is constant. In the form of a chromate, it may be determined by precipitation, in acetic acid solution, with lead acetate; the lead chromate precipitate collected on a tared filter paper, well washed, dried at 100° C. and weighed; or the chromate may be reduced by means of sulphur dioxide to the condition of a chromic salt, the excess of sulphur dioxide expelled by boiling, and the estimation carried out as above.

The atomic weight of chromium has been determined by S.G. Rawson, by the conversion of pure ammonium bichromate into the trioxide (Journal of Chem. Soc., 1899, 55, p. 213), the mean value obtained being 52.06; and also by C. Meinecke, who estimated the amount of silver, chromium and oxygen in silver chromate, the amount of oxygen in potassium bichromate, and the amount of oxygen and chromium in ammonium bichromate (Ann., 1891, 261, p. 339), the mean value obtained being 51.99.

Chromium forms three series of compounds, namely the chromous salts corresponding to CrO, chromous oxide, chromic salts, corresponding to Cr2O3, chromium sesquioxide, and the chromates corresponding to CrO3, chromium trioxide or chromic anhydride. Chromium sesquioxide is a basic oxide, although like alumina it acts as an acid-forming oxide towards strong bases, forming salts called chromites. Various other oxides of chromium, intermediate in composition between the sesquioxide and trioxide, have been described, namely chromium dioxide, Cr2O3·CrO3, and the oxide CrO3·2Cr2O3.

Chromous oxide, CrO, is unknown in the free state, but in the hydrated condition as CrO·H2O or Cr(OH)2 it may be prepared by precipitating chromous chloride by a solution of potassium hydroxide in air-free water. The precipitate so obtained is a brown amorphous solid which readily oxidizes on exposure, and is decomposed by heat with liberation of hydrogen and formation of the sesquioxide. The sesquioxide, Cr2O3, occurs native, and can be artificially obtained in several different ways, e.g., by igniting the corresponding hydroxide, or chromium trioxide, or ammonium bichromate, or by passing the vapours of chromium oxychloride through a red-hot tube, or by ignition of mercurous chromate. In the amorphous state it is a dull green, almost infusible powder, but as obtained from chromium oxychloride it is deposited in the form of dark green hexagonal crystals of specific gravity 5.2. After ignition it becomes almost insoluble in acids, and on fusion with silicates it colours them green; consequently it is used as a pigment for colouring glass and china. By the fusion of potassium bichromate with boric acid, and extraction of the melt with water, a residue is left which possesses a fine green colour, and is used as a pigment under the name of Guignet’s green. In composition it approximates to Cr2O3·H2O, but it always contains more or less boron trioxide. Several forms of hydrated chromium sesquioxide are known; thus on precipitation of a chromic salt, free from alkali, by ammonia, a light blue precipitate is formed, which after drying over sulphuric acid, has the composition Cr2O3·7H2O, and this after being heated to 200° C. in a current of hydrogen leaves a residue of composition CrO·OH or Cr2O3·H2O which occurs naturally as chrome ochre. Other hydrated oxides such as Cr2O3·2H2O have also been described. Chromium trioxide, CrO3, is obtained by adding concentrated sulphuric acid to a cold saturated solution of potassium bichromate, when it separates in long red needles; the mother liquor is drained off and the crystals are washed with concentrated nitric acid, the excess of which is removed by means of a current of dry air. It is readily soluble in water, melts at 193° C., and is decomposed at a higher temperature into chromium sesquioxide and oxygen; it is a very powerful oxidizing agent, acting violently on alcohol, converting it into acetaldehyde, and in glacial acetic acid solution converting naphthalene and anthracene into the corresponding quinones. Heated with concentrated hydrochloric acid it liberates chlorine, and with sulphuric acid it liberates oxygen. Gaseous ammonia passed over the oxide reduces it to the sesquioxide with formation of nitrogen and water. Dissolved in hydrochloric acid at -20°, it yields with solutions of the alkaline chlorides compounds of the type MCl·CrOCl3, pointing to pentavalent chromium. For salts of this acid-forming oxide and for perchromic acid see Bichromates.

The chromites may be looked upon as salts of chromium sesquioxide with other basic oxides, the most important being chromite (q.v.).

Chromous chloride, CrCl2, is prepared by reducing chromic chloride in hydrogen; it forms white silky needles, which dissolve in water giving a deep blue solution, which rapidly absorbs oxygen, forming basic chromic salts, and acts as a very strong reducing agent. The bromide and iodide are formed in a similar manner by heating the metal in gaseous hydrobromic or hydriodic acids.

Chromous sulphate, CrSO4·7H2O, isomorphous with ferrous sulphate, results on dissolving the metal in dilute sulphuric acid or, better, by dissolving chromous acetate in dilute sulphuric acid, when it separates in blue crystals on cooling the solution. On pouring a solution of chromous chloride into a saturated solution of sodium acetate, a red crystalline precipitate of chromous acetate is produced; this is much more permanent in air than the other chromous salts and consequently can be used for their preparation. Chromic salts are of a blue or violet colour, and apparently the chloride and bromide exist in a green and violet form.

Chromic chloride, CrCl3, is obtained in the anhydrous form by igniting a mixture of the sesquioxide and carbon in a current of dry chlorine; it forms violet laminae almost insoluble in water, but dissolves rapidly in presence of a trace of chromous chloride; this action has been regarded as a catalytic action, it being assumed that the insoluble chromic chloride is first reduced by the chromous chloride to the chromous condition and the original chromous chloride converted into soluble chromic chloride, the newly formed chromous chloride then reacting with the insoluble chromic chloride. Solutions of chromic chloride in presence of excess of acid are green in colour. According to A. Werner, four hydrated chromium chlorides exist, namely the green and violet salts, CrCl3·6H2O, a hydrate, CrCl3·10H2O and one CrCl3·4H2O. The violet form gives a purple solution, and all its chlorine is precipitated by silver nitrate, the aqueous solution containing four ions, probably Cr(OH2)6 and three chlorine ions. The green salt appears to dissociate in aqueous solution into two ions, namely CrCl2(OH2)4 and one chlorine ion, since practically only one-third of the chlorine is precipitated by silver nitrate solution at 0° C. Two of the six water molecules are easily removed in a desiccator, and the salt formed, CrCl3·4H2O, resembles the original salt in properties, only one-third of the chlorine being precipitated by silver nitrate. In accordance with his theory of the constitution of salts Werner formulates the hexahydrate as CrCl2·(OH2)4·Cl·2H2O.

Chromic bromide, CrBr3, is prepared in the anhydrous form by the same method as the chloride, and resembles it in its properties. The iodide is unknown.

The fluoride, CrF3, results on passing hydrofluoric acid over the heated chloride, and sublimes in needles. The hydrated fluoride, CrF3·9H2O, obtained by adding ammonium fluoride to cold chromic sulphate solution, is sparingly soluble in water, and is decomposed by heat.

Oxyhalogen derivatives of chromium are known, the oxychloride, CrO2Cl2, resulting on heating potassium bichromate and common salt with concentrated sulphuric acid. It distils over as a dark red liquid of boiling point 117° C., and is to be regarded as the acid chloride corresponding to chromic acid, CrO2(OH)2. It dissolves iodine and absorbs chlorine, and is decomposed by water with formation of chromic and hydrochloric acids; it takes fire in contact with sulphur, ammonia, alcohol, &c., and explodes in contact with phosphorus; it also acts as a powerful oxidizing agent. Heated in a closed tube at 180° C. it loses chlorine and leaves a black residue of trichromyl chloride, Cr3O6Cl2, which deliquesces on exposure to air. Analogous bromine and iodine compounds are unknown, since bromides and iodides on heating with potassium bichromate and concentrated sulphuric acid give free bromine or free iodine.

The oxyfluoride, CrO2F2, is obtained in a similar manner to the oxychloride by using fluorspar in place of common salt. It may be condensed to a dark red liquid which is decomposed by moist air into chromic acid and chromic fluoride.

The semi-acid chloride, CrO2·Cl·OH, chlorochromic acid, is only known in the form of its salts, the chlorochromates.

Potassium chlorochromate, CrO2·Cl·OK, is produced when potassium bichromate is heated with concentrated hydrochloric acid and a little water, or from chromium oxychloride and saturated potassium chloride solution, when it separates as a red crystalline salt. By suspending it in ether and passing ammonia, potassium amidochromate, CrO2·NH2·OK, is obtained; on evaporating the ether solution, after it has stood for 24 hours, red prisms of the amidochromate separate; it is slowly decomposed by boiling water, and also by nitrous acid, with liberation of nitrogen.

Chromic sulphide, Cr2S3, results on heating chromium and sulphur or on strongly heating the trioxide in a current of sulphuretted hydrogen; it forms a dark green crystalline powder, and on ignition gives the sesquioxide.

Chromic sulphate, Cr2(SO4)3, is prepared by mixing the hydroxide with concentrated sulphuric acid and allowing the mixture to stand, a green solution is first formed which gradually changes to blue, and deposits violet-blue crystals, which are purified by dissolving in water and then precipitating with alcohol. It is soluble in cold water, giving a violet solution, which turns green on boiling. If the violet solution is allowed to evaporate slowly at ordinary temperatures the sulphate crystallizes out as Cr2(SO4)3·15H2O, but the green solution on evaporation leaves only an amorphous mass. Investigation has shown that the change is due to the splitting off of sulphuric acid during the process, and that green-coloured chrom-sulphuric acids are formed thus—

2Cr2(SO4)3 \+ H2O = H2SO4 \+ [Cr4O·(SO4)4]SO4

(violet) (green)  
  
---  
  
since, on adding barium chloride to the green solution, only one-third of the total sulphuric acid is precipitated as barium sulphate, whence it follows that only one-third of the original SO4 ions are present in the green solution. The green salt in aqueous solution, on standing, gradually passes back to the violet form. Several other complex chrom-sulphuric acids are known, e.g.

[Cr2(SO4)4]H2; [Cr2(SO4)5]H4; [Cr2(SO4)6]H6

(see A. Recoura, Annales de Chimie et de Physique, 1895 (7), 4, p. 505.)

Chromic sulphate combines with the sulphates of the alkali metals to form double sulphates, which correspond to the alums. Chrome alum, K2SO4·Cr2(SO4)3·24H2O, is best prepared by passing sulphur dioxide through a solution of potassium bichromate containing the calculated quantity of sulphuric acid,

K2Cr2O7 \+ 3SO2 \+ H2SO4 = H2O + K2SO4 \+ Cr2(SO4)3.

On evaporating the solution dark purple octahedra of the alum are obtained. It is easily soluble in warm water, the solution being of a dull blue tint, and is used in calico-printing, dyeing and tanning. Chromium ammonium sulphate, (NH4)2SO4·Cr2(SO4)3·24H2O, results on mixing equivalent quantities of chromic sulphate and ammonium sulphate in aqueous solution and allowing the mixture to crystallize. It forms red octahedra and is less soluble in water than the corresponding potassium compound. The salt CrClSO4·8H2O has been described. By passing ammonia over heated chromic chloride, the nitride, CrN, is formed as a brownish powder. By the action of concentrated sulphuric acid it is transformed into chromium ammonium sulphate.

The nitrate, Cr(NO3)3·9H2O, crystallizes in purple prisms and results on dissolving the hydroxide in nitric acid, its solution turns green on boiling. A phosphide, PCr, is known; it burns in oxygen forming the phosphate. By adding sodium phosphate to an excess of chrome alum the violet phosphate, CrPO4·6H2O, is precipitated; on heating to 100° C. it loses water and turns green. A green precipitate, perhaps CrPO4·3H2O, is obtained on adding an excess of sodium phosphate to chromic chloride solution.

Carbides of chromium are known; when the metal is heated in an electric furnace with excess of carbon, crystalline, C2Cr3, is formed; this scratches quartz and topaz, and the crystals are very resistant to the action of acids; CCr4 has also been described (H. Moissan, Comptes rendus, 1894, 119, p. 185).

Cyanogen compounds of chromium, analogous to those of iron, have been prepared; thus potassium chromocyanide, K4Cr(CN)6·2H2O, is formed from potassium cyanide and chromous acetate; on exposure to air it is converted into the chromicyanide, K3Cr(CN)6, which can also be prepared by adding chromic acetate solution to boiling potassium cyanide solution. Chromic thiocyanate, Cr(SCN)3, an amorphous deliquescent mass, is formed by dissolving the hydroxide in thiocyanic acid and drying over sulphuric acid. The double thiocyanate, Cr(SCN)3·3KCNS·4H2O, is also known.

Chromium salts readily combine with ammonia to form complex salts in which the ammonia molecule is in direct combination with the chromium atom. In many of these salts one finds that the elements of water are frequently found in combination with the metal, and further, that the ammonia molecule may be replaced by such other molecular groups as -NO2, &c. Of the types studied the following may be mentioned: the diammine chromium thiocyanates, M[Cr(NH3)2·(SCN)4], the chloraquotetrammine chromic salts, R¹2[Cr(NH3)4·H2O·Cl], the aquopentammine or roseo-chromium salts, R¹3[Cr(NH3)5·H2O], the chlorpentammine or purpureo-chromium salts, R¹2[Cr(NH3)5·Cl], the nitrito pentammine or xanthochromium salts, R¹2[NO2·(NH3)5·Cr], the luteo or hexammine chromium salts, R¹3[(NH3)6·Cr], and the rhodochromium salts: where R¹ = a monovalent acid radical and M = a monovalent basic radical. For the preparation and properties of these salts and a discussion on their constitution the papers of S.F. Jörgensen and of A. Werner in the Zeitschrift für anorganische Chemie from 1892 onwards should be consulted.

P. Pfeiffer (Berichte, 1904, 37, p. 4255) has shown that chromium salts of the type [Cr{C2H4(NH2)2}2X2]X exist in two stereo-isomeric forms, namely, the cis- and trans- forms, the dithiocyan-diethylene-diamine-chromium salts being the trans- salts. Their configuration was determined by their relationship to their oxalo-derivatives; the cis-dichloro chloride, [CrC2H4(NH2)2Cl2]Cl·H2O, compound with potassium oxalate gave a carmine red crystalline complex salt, [Cr{C2H4(NH2)2}C2O4][CrC2H4(NH2)2·(C2O4)2]1½H2O, while from the trans-chloride a red complex salt is obtained containing the unaltered trans-dichloro group [CrC2H4(NH2)2·Cl2].

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