COBALT (symbol Co, atomic weight 59), one of the metallic chemical elements. The term "cobalt" is met with in the writings of Paracelsus, Agricola and Basil Valentine, being used to denote substances which, although resembling metallic ores, gave no metal on smelting. At a later date it was the name given to the mineral used for the production of a blue colour in glass. In 1735 G. Brandt prepared an impure cobalt metal, which was magnetic and very infusible. Cobalt is usually found associated with nickel, and frequently with arsenic, the chief ores being speiss-cobalt, (Co,Ni,Fe)As2, cobaltite (q.v.), wad, cobalt bloom, linnaeite, Co3S4, and skutterudite, CoAs3. Its presence has also been detected in the Sun and in meteoric iron. For the technical preparation of cobalt, and its separation from nickel, see Nickel. The metal is chiefly used, as the oxide, for colouring glass and porcelain.
Metallic cobalt may be obtained by reduction of the oxide or chloride in a current of hydrogen at a red heat, or by heating the oxalate, under a layer of powdered glass. As prepared by the reduction of the oxide it is a grey powder. In the massive state it has a colour resembling polished iron, and is malleable and very tough. It has a specific gravity of 8.8, and it melts at 1530° C. (H. Copaux). Its mean specific heat between 9° and 97° C. is 0.10674 (H. Kopp). It is permanent in dry air, but in the finely divided state it rapidly combines with oxygen, the compact metal requiring a strong heating to bring about this combination. It decomposes steam at a red heat, and slowly dissolves in dilute hydrochloric and sulphuric acids, but more readily in nitric acid. Cobalt burns in nitric oxide at 150° C. giving the monoxide. It may be obtained in the pure state, according to C. Winkler (Zeit. für anorg. Chem., 1895, 8, p. 1), by electrolysing the pure sulphate in the presence of ammonium sulphate and ammonia, using platinum electrodes, any occluded oxygen in the deposited metal being removed by heating in a current of hydrogen.
Three characteristic oxides of cobalt are known, the monoxide, CoO, the sesquioxide, Co2O3, and tricobalt tetroxide, Co3O4; besides these there are probably oxides of composition CoO2, Co8O9, Co6O7 and Co4O5. Cobalt monoxide, CoO, is prepared by heating the hydroxide or carbonate in a current of air, or by heating the oxide Co3O4 in a current of carbon dioxide. It is a brown coloured powder which is stable in air, but gives a higher oxide when heated. On heating in hydrogen, ammonia or carbon monoxide, or with carbon or sodium, it is reduced to the metallic state. It is readily soluble in warm dilute mineral acids forming cobaltous salts. Cobaltous hydroxide, Co(OH)2, is formed when a cobaltous salt is precipitated by caustic potash in the absence of air. A blue basic salt is precipitated first, which, on boiling, rapidly changes to the rose-coloured hydroxide. It dissolves in acids forming cobaltous salts, and on exposure to air it rapidly absorbs oxygen, turning brown in colour. A. de Schulten (Comptes Rendus, 1889, 109, p. 266) has obtained it in a crystalline form; the crystals have a specific gravity of 3.597, and are easily soluble in warm ammonium chloride solution. Cobalt sesquioxide, Co2O3, remains as a dark-brown powder when cobalt nitrate is gently heated. Heated at 190-300° in a current of hydrogen it gives the oxide Co3O4, while at higher temperatures the monoxide is formed, and ultimately cobalt is obtained. Cobaltic hydroxide, Co(OH)3, is formed when a cobalt salt is precipitated by an alkaline hypochlorite, or on passing chlorine through water containing suspended cobaltous hydroxide or carbonate. It is a brown-black powder soluble in hydrochloric acid, chlorine being simultaneously liberated. This hydroxide is soluble in well cooled acids, forming solutions which contain cobaltic salts, one of the most stable of which is the acetate. Cobalt dioxide, CoO2, has not yet been isolated in the pure state; it is probably formed when iodine and caustic soda are added to a solution of a cobaltous salt. By suspending cobaltous hydroxide in water and adding hydrogen peroxide, a strongly acid liquid is obtained (after filtering) which probably contains cobaltous acid, H2CoO3. The barium and magnesium salts of this acid are formed when baryta and magnesia are fused with cobalt sesquioxide. Tricobalt tetroxide, Co3O4, is produced when the other oxides, or the nitrate, are heated in air. By heating a mixture of cobalt oxalate and sal-ammoniac in air, it is obtained in the form of minute hard octahedra, which are not magnetic, and are only soluble in concentrated sulphuric acid.
The cobaltous salts are formed when the metal, cobaltous oxide, hydroxide or carbonate, are dissolved in acids, or, in the case of the insoluble salts, by precipitation. The insoluble salts are rose-red or violet in colour. The soluble salts are, when in the hydrated condition, also red, but in the anhydrous condition are blue. They are precipitated from their alkaline solutions as cobalt sulphide by sulphuretted hydrogen, but this precipitation is prevented by the presence of citric and tartaric acids; similarly the presence of ammonium salts hinders their precipitation by caustic alkalis. Alkaline carbonates give precipitates of basic carbonates, the formation of which is also retarded by the presence of ammonium salts. For the action of ammonia on the cobaltous salts in the presence of air see Cobaltammines (below). On the addition of potassium cyanide they give a brown precipitate of cobalt cyanide, Co(CN)2, which dissolves in excess of potassium cyanide to a green solution.
Cobalt chloride, CoCl2, in the anhydrous state, is formed by burning the metal in chlorine or by heating the sulphide in a current of the same gas. It is blue in colour and sublimes readily. It dissolves easily in water, forming the hydrated chloride, CoCl2·6H2O, which may also be prepared by dissolving the hydroxide or carbonate in hydrochloric acid. The hydrated salt forms rose-red prisms, readily soluble in water to a red solution, and in alcohol to a blue solution. Other hydrated forms of the chloride, of composition CoCl2·2H2O and CoCl2·4H2O have been described (P. Sabatier, Bull. Soc. Chim. 51, p. 88; Bersch, Jahresb. d. Chemie, 1867, p. 291). Double chlorides of composition CoCl2·NH4Cl·6H2O; CoCl2·SnCl4·6H2O and CoCl2·2CdCl2·12H2O are also known. By the addition of excess of ammonia to a cobalt chloride solution in absence of air, a greenish-blue precipitate is obtained which, on heating, dissolves in the solution, giving a rose-red liquid. This solution, on standing, deposits octahedra of the composition CoCl2·6NH3. These crystals when heated to 120° C. lose ammonia and are converted into the compound CoCl2·2NH3 (E. Frémy). The bromide, CoBr2, resembles the chloride, and may be prepared by similar methods. The hydrated salt readily loses water on heating, forming at 100° C. the hydrate CoBr2·2H2O, and at 130° C. passing into the anhydrous form. The iodide, CoI2, is produced by heating cobalt and iodine together, and forms a greyish-green mass which dissolves readily in water forming a red solution. On evaporating this solution the hydrated salt CoI2·6H2O is obtained in hexagonal prisms. It behaves in an analogous manner to CoBr2·6H2O on heating.
Cobalt fluoride, CoF2·2H2O, is formed when cobalt carbonate is evaporated with an excess of aqueous hydrofluoric acid, separating in rose-red crystalline crusts. Electrolysis of a solution in hydrofluoric acid gives cobaltic fluoride, CoF3.
Sulphides of cobalt of composition Co4S3, CoS, Co3S4, Co2S3 and CoS2 are known. The most common of these sulphides is cobaltous sulphide, CoS, which occurs naturally as syepoorite, and can be artificially prepared by heating cobaltous oxide with sulphur, or by fusing anhydrous cobalt sulphate with barium sulphide and common salt. By either of these methods, it is obtained in the form of bronze-coloured crystals. It may be prepared in the amorphous form by heating cobalt with sulphur dioxide, in a sealed tube, at 200° C. In the hydrated condition it is formed by the action of alkaline sulphides on cobaltous salts, or by precipitating cobalt acetate with sulphuretted hydrogen (in the absence of free acetic acid). It is a black amorphous powder soluble in concentrated sulphuric and hydrochloric acids, and when in the moist state readily oxidizes on exposure.
Cobaltous sulphate, CoSO4·7H2O, is found naturally as the mineral bieberite, and is formed when cobalt, cobaltous oxide or carbonate are dissolved in dilute sulphuric acid. It forms dark red crystals isomorphous with ferrous sulphate, and readily soluble in water. By dissolving it in concentrated sulphuric acid and warming the solution, the anhydrous salt is obtained. Hydrated sulphates of composition CoSO4·6H2O, CoSO4·4H2O and CoSO4·H2O are also known. The heptahydrated salt combines with the alkaline sulphates to form double sulphates of composition CoSO4·M2SO4·6H2O (M = K, NH4, etc.).
The cobaltic salts corresponding to the oxide Co2O3 are generally unstable compounds which exist only in solution. H. Marshall (Proc. Roy. Soc. Edin. 59, p. 760) has prepared cobaltic sulphate Co2(SO4)3·18H2O, in the form of small needles, by the electrolysis of cobalt sulphate. In a similar way potassium and ammonium cobalt alums have been obtained. A cobaltisulphurous acid, probably H6[(SO3)6·Co2] has been obtained by E. Berglund (Berichte, 1874, 7, p. 469), in aqueous solution, by dissolving ammonium cobalto-cobaltisulphite (NH4)2Co2[(SO3)6·Co2]·14H2O in dilute hydrochloric or nitric acids, or by decomposition of its silver salt with hydrochloric acid. The ammonium cobalto-cobaltisulphite is prepared by saturating an air-oxidized ammoniacal solution of cobaltous chloride with sulphur dioxide. The double salts containing the metal in the cobaltic form are more stable than the corresponding single salts, and of these potassium cobaltinitrite, Co2(NO2)6·6KNO2·3H2O, is best known. It may be prepared by the addition of potassium nitrite to an acetic acid solution of cobalt chloride. The yellow precipitate obtained is washed with a solution of potassium acetate and finally with dilute alcohol. The reaction proceeds according to the following equation: 2CoCl2 + 10KNO2 + 4HNO2 = Co2(NO2)6·6KNO2 + 4KCl + 2NO + 2H2O (A. Stromeyer, Annalen, 1855, 96, p. 220). This salt may be used for the separation of cobalt and nickel, since the latter metal does not form a similar double nitrite, but it is necessary that the alkaline earth metals should be absent, for in their presence nickel forms complex nitrites containing the alkaline earth metal and the alkali metal. A sodium cobaltinitrite is also known.
Cobalt nitrate, Co(NO3)2·6H2O, is obtained in dark-red monoclinic tables by the slow evaporation of a solution of the metal, its hydroxide or carbonate, in nitric acid. It deliquesces in the air and melts readily on heating. By the addition of excess of ammonia to its aqueous solution, in the complete absence of air, a blue precipitate of a basic nitrate of the composition 6CoO·N2O5·5H2O is obtained.
By boiling a solution of cobalt carbonate in phosphoric acid, the acid phosphate CoHPO4·3H2O is obtained, which when heated with water to 250° C. is converted into the neutral phosphate Co3(PO4)2·2H2O (H. Debray, Ann. de chimie, 1861,  61, p. 438). Cobalt ammonium phosphate, CoNH4PO4·12H2O, is formed when a soluble cobalt salt is digested for some time with excess of a warm solution of ammonium phosphate. It separates in the form of small rose-red crystals, which decompose on boiling with water.
Cobaltous cyanide, Co(CN)2·3H2O, is obtained when the carbonate is dissolved in hydrocyanic acid or when the acetate is precipitated by potassium cyanide. It is insoluble in dilute acids, but is readily soluble in excess of potassium cyanide. The double cyanides of cobalt are analogous to those of iron. Hydrocobaltocyanic acid is not known, but its potassium salt, K4Co(CN)6, is formed when freshly precipitated cobalt cyanide is dissolved in an ice-cold solution of potassium cyanide. The liquid is precipitated by alcohol, and the washed and dried precipitate is then dissolved in water and allowed to stand, when the salt separates in dark-coloured crystals. In alkaline solution it readily takes up oxygen and is converted into potassium cobalticyanide, K3Co(CN)6, which may also be obtained by evaporating a solution of cobalt cyanide, in excess of potassium cyanide, in the presence of air, 8KCN + 2Co(CN)2 + H2O + O = 2K3Co(CN)6 + 2KHO. It forms monoclinic crystals which are very soluble in water. From its aqueous solution, concentrated hydrochloric acid precipitates hydrocobalticyanic acid, H3Co(CN)6, as a colourless solid which is very deliquescent, and is not attacked by concentrated hydrochloric and nitric acids. For a description of the various salts of this acid, see P. Wesselsky, Berichte, 1869, 2, p. 588.
Cobaltammines. A large number of cobalt compounds are known, of which the empirical composition represents them as salts of cobalt to which one or more molecules of ammonia have been added. These salts have been divided into the following series: -
Triammine Series, [Co(NH3)3]X3. Here X = Cl, NO3, NO2, SO4, etc.
Tetrammine Series. This group may be divided into the
Praseo-salts [R2Co(NH3)4]X, where X = Cl.
Croceo-salts [(NO2)2Co(NH3)4]X, which may be considered as a subdivision of the praseo-salts.
Tetrammine purpureo-salts [RCo(NH3)4·H2O]X2.
Tetrammine roseo-salts [Co(NH3)4·(H2O)2]X3.
Pentammine purpureo-salts [R·Co(NH3)5]X2 where X = Cl, Br, NO3, N02, SO4, etc.
Pentammine roseo-salts [Co(NH3)5·H2O]X2.
Hexammine or Luteo Series [Co(NH3)6]X3.
The hexammine salts are formed by the oxidizing action of air on dilute ammoniacal solutions of cobaltous salts, especially in presence of a large excess of ammonium chloride. They form yellow or bronze-coloured crystals, which decompose on boiling their aqueous solution. On boiling their solution in caustic alkalis, ammonia is liberated. The pentammine purpureo-salts are formed from the luteo-salts by loss of ammonia, or from an air slowly oxidized ammoniacal cobalt salt solution, the precipitated luteo-salt being filtered off and the filtrate boiled with concentrated acids. They are violet-red in colour, and on boiling or long standing with dilute acids they pass into the corresponding roseo-salts.
The pentammine nitrito salts are known as the xanthocobalt salts and have the general formula [NO2·Co·(NH3)5]X2. They are formed by the action of nitrous fumes on ammoniacal solutions of cobaltous salts, or purpureo-salts, or by the mutual reaction of chlorpurpureo-salts and alkaline nitrites. They are soluble in water and give characteristic precipitates with platinic and auric chlorides, and with potassium ferrocyanide. The pentammine roseo-salts can be obtained from the action of concentrated acids, in the cold, on air-oxidized solutions of cobaltous salts. They are of a reddish colour and usually crystallize well; on heating with concentrated acids are usually transformed into the purpureo-salts. Their alkaline solutions liberate ammonia on boiling. They give a characteristic pale red precipitate with sodium pyrophosphate, soluble in an excess of the precipitant; they also form precipitates on the addition of platinic chloride and potassium ferrocyanide. For methods of preparation of the tetrammine and triammine salts, see O. Dammer's Handbuch der anorganischen Chemie, vol. 3 (containing a complete account of the preparation of the cobaltammine salts). The diammine salts are prepared by the action of alkaline nitrites on cobaltous salts in the presence of much ammonium chloride or nitrate; they are yellow or brown crystalline solids, not very soluble in cold water.
The above series of salts show striking differences in their behaviour towards reagents; thus, aqueous solutions of the luteo chlorides are strongly ionized, as is shown by their high electric conductivity; and all their chlorine is precipitated on the addition of silver nitrate solution. The aqueous solution, however, does not show the ordinary reactions of cobalt or of ammonia, and so it is to be presumed that the salt ionizes into [Co(NH3)6] and 3Cl'. The purpureo chloride has only two-thirds of its chlorine precipitated on the addition of silver nitrate, and the electric conductivity is much less than that of the luteo chloride; again in the praseo-salts only one-third of the chlorine is precipitated by silver nitrate, the conductivity again falling; while in the triammine salts all ionization has disappeared. For the constitution of these salts and of the "metal ammonia" compounds generally, see A. Werner, Zeit. für anorg. Chemie, 1893 et seq., and Berichte, 1895, et seq.; and S. Jörgensen, Zeit. für anorg. Chemie, 1892 et seq.
The oxycobaltammines are a series of compounds of the general type [Co2O3·H2(NH3)10]X4 first observed by L. Gmelin, and subsequently examined by E. Frémy, W. Gibbs and G. Vortmann (Monatshefte für Chemie, 1885, 6, p. 404). They result from the cobaltammines by the direct taking up of oxygen and water. On heating, they decompose, forming basic tetrammine salts.
The atomic weight of cobalt has been frequently determined, the earlier results not being very concordant (see R. Schneider, Pog. Ann., 1857, 101, p. 387; C. Marignac, Arch. Phys. Nat. , 1, p. 373; W. Gibbs, Amer. Jour. Sci. , 25, p. 483; J. B. Dumas, Ann. Chim. Phys., 1859 , 55, p. 129; W. J. Russell, Jour. Chem. Soc., 1863, 16, p. 51). C. Winkler, by the analysis of the chloride, and by the action of iodine on the metal, obtained the values 59.37 and 59.07, whilst W. Hempel and H. Thiele (Zeit. f. anorg. Chem., 1896, II, p. 73), by reducing cobalto-cobaltic oxide, and by the analysis of the chloride, have obtained the values 58.56 and 58.48. G. P. Baxter and others deduced the value 58.995 (O = 16).
Cobalt salts may be readily detected by the formation of the black sulphide, in alkaline solution, and by the blue colour they produce when fused with borax. For the quantitative determination of cobalt, it is either weighed as the oxide, Co3O4, obtained by ignition of the precipitated monoxide, or it is reduced in a current of hydrogen and weighed as metal. For the quantitative separation of cobalt and nickel, see E. Hintz (Zeit. f. anal. Chem., 1891, 30, p. 227), and also Nickel.
Note - this article incorporates content from Encyclopaedia Britannica, Eleventh Edition, (1910-1911)