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Chromic Chloride »
Chromic Chloride, CrCl3
Chromic Chloride, CrCl3 may be obtained by a variety of methods, for example,
Anhydrous chromic chloride forms large, lustrous, unctuous plates of the colour of peach blossom, of density 2.757 at 15° C. It has a specific heat of 0.143. Its vapour density just above 1065° C., at which temperature it volatilises, is 6.135 (air = 1), whilst at 1200° C. it approximates to the theoretical value, 5.478, required by the formula CrCl3. At much higher temperatures partial decomposition takes place, with consequent diminution of the vapour density.
Chromic chloride is reduced at a red heat by hydrogen to chromous chloride, and by zinc or magnesium to metallic chromium. Heated in air, green chromium sesquioxide is formed; in oxygen or moist chlorine, chromyl chloride is produced. On strongly heating in dry ammonia, the nitride, CrN, is obtained, whilst in presence of ammonium chloride, the lower chloride is formed; in hydrogen sulphide, the black crystalline sesquisulphide Cr2S3, and with phosphorous pentachloride, the double chloride 2CrCl3.PCl5, are produced. Mineral acids, including even aqua regia, have no action on the anhydrous salt; fused alkali hydroxide or carbonate, in presence of nitrate, yields alkali chromate and chloride. Anhydrous chromic chloride is almost insoluble in cold water, but dissolves readily in presence of very slight traces (0.00001 per cent.) of chromous chloride, either previously added or formed in solution by the addition of a reducing agent such as tin, zinc, ferrous or cuprous chlorides. It has not been proved that the catalytic influence of these reagents is wholly due to the formation of chromous chloride. Chromic chloride is sparingly soluble in boiling water.
A solution of chromic chloride is readily obtained by the action of dilute hydrochloric acid on chromic hydroxide, or on chromium sesquioxide which has not been ignited; this solution on evaporation yields an amorphous, deliquescent, green mass, soluble in water and in alcohol. The product on heating in hydrogen chloride or chlorine at temperatures above 250° C. becomes anhydrous, at the same time assuming the usual colour of the anhydrous salt. Basic chlorides are formed by heating the hydrated chloride in air. A number of hydrates, one at least appearing to exist in three isomeric forms, have been prepared. The hydrates in solution all yield the same hydroxide, Cr(OH)3, on treatment with an alkali hydroxide.
Dilute solutions of the chloride are violet in colour, while more concentrated or acidulated solutions are green, the colour depending on the equilibrium established between the green and violet modifications of the salt which is in solution as the hexahydrate, CrCl3.6H2O. The amount of the green salt present at equilibrium increases with concentration. The violet modification is precipitated before the green when hydrogen chloride is passed into a boiled solution of the mixture. From dehydration experiments Werner concluded that the molecule of the green hydrate contained 4H2O as water of constitution and 2H2O as water of crystallisation, while the violet modification contained 6H2O as water of constitution. Olie confirmed these results at ordinary temperatures, but found that at 100° C. both varieties lost approximately 4H2O. Further, Werner found that from solutions of the violet hydrate, the whole, and from solutions of the green hydrate only one- third, of the chlorine could be precipitated by a soluble silver salt. Although the amount of chlorine precipitated appears to depend to a certain extent on the conditions of the experiment, yet it is evident that in the two salts the chlorine is not all similarly combined. The isomerism was explained by Werner according to the co-ordination theory, and the violet and green chlorides were considered to possess the formulae [Cr(H2O)6]Cl3 and | CrCl2(H2O)4]Cl.(H2O)2 respectively, the former being termed the hexaquo salt and the latter dichlortetraquo- chromic chloride, or simply dichlorchromium chloride.
It has been suggested that the greyish-blue or violet chromic chloride is bimolecular, 2[Cr(H2O)6Cl3], while the green variety is uni- molecular. This is not accepted by Bjerrum, who considered that, in order to interpret the equilibrium changes of a concentrated solution of chromic chloride, it was necessary to assume the presence of a third isomeric hexahydrate in the solution. This he succeeded in isolating, as very deliquescent pale green crystals, by adding ether saturated with hydrogen chloride to the solution remaining after precipitating the violet salt with hydrogen chloride. Two-thirds of the chlorine present in this compound may be precipitated by means of silver nitrate, so that it may be regarded as monochlorpentaquochromic chloride, [CrCl (H2O)5]Cl2.H2O.
In dilute solution the dark green hexahydrate changes rapidly to the light green salt, and then more slowly to the violet salt, thus:
[CrCl2(H2O)4]Cl.(H2O)2, dark green, → [CrCl(H2O)5]Cl2.H2O, light green, → [Cr(H2O)6]Cl3, violet.
The influence of light and of the presence of neutral chlorides on the transformation has been studied.
By dehydration of the green hydrate, Bjerrum obtained two red substances, 2CrCl3.3H2O and 2CrCl3.H2O. Other hydrates are the deca- hydrate, CrCl3.10H2O, and tetrahydrate, CrCl3.4H2O. The former yields brilliant green triclinic crystals, strongly dichroic, and can be obtained by triturating the hexahydrate with the calculated amount of water. Werner and Gubser assign the formula
to the decahydrate, which in dry air loses water, being converted first into the hexahydrate and finally into the tetrahydrate, a pale green, slightly hygroscopic powder. A hydrate, 2CrCl3.9H2O, has also been said to exist. An alcoholate, CrCl3.3C2H5OH, in the form of red needles, stable in dry air, has been obtained by the action of dry hydrogen chloride in absolute alcohol upon metallic chromium. Double salts with alkali chlorides, with antimony pentachloride, with organic bases, ammonia (for example, 2CrCl3.12NH3.2H2O and 2CrCl3.10NH3), and substituted ammonias are known. Complex halogen-halides and chlorsulphates have also been obtained.
On progressive hydrolysis, chromic chloride gives rise to two soluble basic chlorides, Cr(OH)Cl2 and Cr(OH)2Cl, and an insoluble grey-green hydroxide. The compound, Cr(OH)2Cl, is never present in any large proportion, and the hydroxide is only formed upon addition of alkali. The formulae of the three compounds are probably [Cr(H2O)5(OH)]Cl2, [Cr(H2O)4(OH)2]Cl, and Cr(H2O)3(OH)3 respectively.
- by heating metallic chromium at 600° C. in a stream of chlorine;
- by passing a stream of chlorine over chromium sesquioxide at 440° C.;
- by heating a mixture of the sesquioxide and carbon in a stream of chlorine;
- by the action of a mixture of carbon monoxide and chlorine on chromium sesquioxide at a red heat:
Cr2O3 + 3CO + 3Cl2 = 2CrCl3 + 3CO2; - by the interaction of phosphorus trichloride and chromyl chloride, CrO2Cl2; or of chromyl chloride, carbon monoxide, and chlorine.
- Undoubtedly the most convenient method is that of Bourion, who finds that sulphur chloride alone is preferable to a mixture of that substance with chlorine for the chlorination of the oxide. Precipitated calcined chromic oxide is attacked by sulphur chloride at a temperature rather above 400° C., but below red heat:
6S2Cl2 + 2Cr2O3 = 4CrCl3 + 3SO2 + 9S.
Anhydrous chromic chloride forms large, lustrous, unctuous plates of the colour of peach blossom, of density 2.757 at 15° C. It has a specific heat of 0.143. Its vapour density just above 1065° C., at which temperature it volatilises, is 6.135 (air = 1), whilst at 1200° C. it approximates to the theoretical value, 5.478, required by the formula CrCl3. At much higher temperatures partial decomposition takes place, with consequent diminution of the vapour density.
Chromic chloride is reduced at a red heat by hydrogen to chromous chloride, and by zinc or magnesium to metallic chromium. Heated in air, green chromium sesquioxide is formed; in oxygen or moist chlorine, chromyl chloride is produced. On strongly heating in dry ammonia, the nitride, CrN, is obtained, whilst in presence of ammonium chloride, the lower chloride is formed; in hydrogen sulphide, the black crystalline sesquisulphide Cr2S3, and with phosphorous pentachloride, the double chloride 2CrCl3.PCl5, are produced. Mineral acids, including even aqua regia, have no action on the anhydrous salt; fused alkali hydroxide or carbonate, in presence of nitrate, yields alkali chromate and chloride. Anhydrous chromic chloride is almost insoluble in cold water, but dissolves readily in presence of very slight traces (0.00001 per cent.) of chromous chloride, either previously added or formed in solution by the addition of a reducing agent such as tin, zinc, ferrous or cuprous chlorides. It has not been proved that the catalytic influence of these reagents is wholly due to the formation of chromous chloride. Chromic chloride is sparingly soluble in boiling water.
A solution of chromic chloride is readily obtained by the action of dilute hydrochloric acid on chromic hydroxide, or on chromium sesquioxide which has not been ignited; this solution on evaporation yields an amorphous, deliquescent, green mass, soluble in water and in alcohol. The product on heating in hydrogen chloride or chlorine at temperatures above 250° C. becomes anhydrous, at the same time assuming the usual colour of the anhydrous salt. Basic chlorides are formed by heating the hydrated chloride in air. A number of hydrates, one at least appearing to exist in three isomeric forms, have been prepared. The hydrates in solution all yield the same hydroxide, Cr(OH)3, on treatment with an alkali hydroxide.
Dilute solutions of the chloride are violet in colour, while more concentrated or acidulated solutions are green, the colour depending on the equilibrium established between the green and violet modifications of the salt which is in solution as the hexahydrate, CrCl3.6H2O. The amount of the green salt present at equilibrium increases with concentration. The violet modification is precipitated before the green when hydrogen chloride is passed into a boiled solution of the mixture. From dehydration experiments Werner concluded that the molecule of the green hydrate contained 4H2O as water of constitution and 2H2O as water of crystallisation, while the violet modification contained 6H2O as water of constitution. Olie confirmed these results at ordinary temperatures, but found that at 100° C. both varieties lost approximately 4H2O. Further, Werner found that from solutions of the violet hydrate, the whole, and from solutions of the green hydrate only one- third, of the chlorine could be precipitated by a soluble silver salt. Although the amount of chlorine precipitated appears to depend to a certain extent on the conditions of the experiment, yet it is evident that in the two salts the chlorine is not all similarly combined. The isomerism was explained by Werner according to the co-ordination theory, and the violet and green chlorides were considered to possess the formulae [Cr(H2O)6]Cl3 and | CrCl2(H2O)4]Cl.(H2O)2 respectively, the former being termed the hexaquo salt and the latter dichlortetraquo- chromic chloride, or simply dichlorchromium chloride.
It has been suggested that the greyish-blue or violet chromic chloride is bimolecular, 2[Cr(H2O)6Cl3], while the green variety is uni- molecular. This is not accepted by Bjerrum, who considered that, in order to interpret the equilibrium changes of a concentrated solution of chromic chloride, it was necessary to assume the presence of a third isomeric hexahydrate in the solution. This he succeeded in isolating, as very deliquescent pale green crystals, by adding ether saturated with hydrogen chloride to the solution remaining after precipitating the violet salt with hydrogen chloride. Two-thirds of the chlorine present in this compound may be precipitated by means of silver nitrate, so that it may be regarded as monochlorpentaquochromic chloride, [CrCl (H2O)5]Cl2.H2O.
In dilute solution the dark green hexahydrate changes rapidly to the light green salt, and then more slowly to the violet salt, thus:
[CrCl2(H2O)4]Cl.(H2O)2, dark green, → [CrCl(H2O)5]Cl2.H2O, light green, → [Cr(H2O)6]Cl3, violet.
The influence of light and of the presence of neutral chlorides on the transformation has been studied.
By dehydration of the green hydrate, Bjerrum obtained two red substances, 2CrCl3.3H2O and 2CrCl3.H2O. Other hydrates are the deca- hydrate, CrCl3.10H2O, and tetrahydrate, CrCl3.4H2O. The former yields brilliant green triclinic crystals, strongly dichroic, and can be obtained by triturating the hexahydrate with the calculated amount of water. Werner and Gubser assign the formula
to the decahydrate, which in dry air loses water, being converted first into the hexahydrate and finally into the tetrahydrate, a pale green, slightly hygroscopic powder. A hydrate, 2CrCl3.9H2O, has also been said to exist. An alcoholate, CrCl3.3C2H5OH, in the form of red needles, stable in dry air, has been obtained by the action of dry hydrogen chloride in absolute alcohol upon metallic chromium. Double salts with alkali chlorides, with antimony pentachloride, with organic bases, ammonia (for example, 2CrCl3.12NH3.2H2O and 2CrCl3.10NH3), and substituted ammonias are known. Complex halogen-halides and chlorsulphates have also been obtained.
On progressive hydrolysis, chromic chloride gives rise to two soluble basic chlorides, Cr(OH)Cl2 and Cr(OH)2Cl, and an insoluble grey-green hydroxide. The compound, Cr(OH)2Cl, is never present in any large proportion, and the hydroxide is only formed upon addition of alkali. The formulae of the three compounds are probably [Cr(H2O)5(OH)]Cl2, [Cr(H2O)4(OH)2]Cl, and Cr(H2O)3(OH)3 respectively.
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