Cobalt(II) chloride


Cobalt chloride is an inorganic compound of cobalt and chlorine, with the formula. It is a sky blue crystalline solid.
The compound forms several hydrates •n, for n = 1, 2, 6, and 9. Claims of the formation of tri- and tetrahydrates have not been confirmed. The dihydrate is purple and hexahydrate is pink. It is usually supplied as the hexahydrate ·6, which is one of the most commonly used cobalt compounds in the lab.
Because of the ease of the hydration/dehydration reaction, and the resulting color change, cobalt chloride is used as an indicator for water in desiccants.
Niche uses of cobalt chloride include its role in organic synthesis and electroplating objects with cobalt metal.
Cobalt chloride has been classified as a substance of very high concern by the European Chemicals Agency as it is a suspected carcinogen.

Properties

Anhydrous

At room temperature, anhydrous cobalt chloride has the cadmium chloride structure in which the cobalt ions are octahedrally coordinated. At about 706 °C, the coordination is believed to change to tetrahedral. The vapor pressure has been reported as 7.6 mm Hg at the melting point.

Solutions

Cobalt chloride is fairly soluble in water. Under atmospheric pressure, the mass concentration of a saturated solution of in water is about 54% at the boiling point, 120.2 °C; 48% at 51.25 °C; 35% at 25 °C; 33% at 0 °C; and 29% at −27.8 °C.
Diluted aqueous solutions of contain the species, besides chloride ions. Concentrated solutions are red at room temperature but become blue at higher temperatures.

Hydrates

The crystal unit of the solid hexahydrate •6 contains the neutral molecule trans- and two molecules of water of crystallization. This species dissolves readily in water and alcohol.
The anhydrous salt is hygroscopic and the hexahydrate is deliquescent.

Preparation

Cobalt chloride can be prepared in aqueous solution from cobalt hydroxide or cobalt carbonate and hydrochloric acid:
The solid dihydrate and hexahydrate can be obtained by evaporation. Cooling saturated aqueous solutions yields the dihydrate between 120.2 °C and 51.25 °C, and the hexahydrate below 51.25 °C. Water ice, rather than cobalt chloride, will crystallize from solutions with concentration below 29%. The monohydrate and the anhydrous forms can be obtained by cooling solutions only under high pressure, above 206 °C and 335 °C, respectively.
The anhydrous compound can be prepared by heating the hydrates. On rapid heating or in a closed container, each of the 6-, 2-, and 1- hydrates partially melts into a mixture of the next lower hydrate and a saturated solution -- at 51.25 °C, 206 °C, and 335 °C, respectively. On slow heating in an open container, water evaporates out of each of the solid 6-, 2-, and 1- hydrates, leaving the next lower hydrate -- at about 40 °C, 89 °C, and 126 °C, respectively.
Dehydration can also be effected with trimethylsilyl chloride:
The anhydrous compound can be purified by sublimation in vacuum.

Reactions

In the laboratory, cobalt chloride serves as a common precursor to other cobalt compounds. Generally, aqueous solutions of the salt behave like other cobalt salts since these solutions consist of the ion regardless of the anion. For example, such solutions give a precipitate of cobalt sulfide upon treatment with hydrogen sulfide.

Complexed chlorides

The hexahydrate and the anhydrous salt are weak Lewis acids. The adducts are usually either octahedral or tetrahedral. With pyridine, one obtains an octahedral complex:
With triphenylphosphine, a tetrahedral complex results:
Salts of the anionic complex CoCl42− can be prepared using tetraethylammonium chloride:
The 2− ion is the blue ion that forms upon addition of hydrochloric acid to aqueous solutions of hydrated cobalt chloride, which are pink.

Reduction

Reaction of the anhydrous compound with sodium cyclopentadienide gives cobaltocene. This 19-electron species is a good reducing agent, being readily oxidised to the yellow 18-electron cobaltacenium cation.

Oxidation to cobalt(III)

Compounds of cobalt in the +3 oxidation state exist, such as cobalt fluoride, nitrate, and sulfate ; however, cobalt chloride is not stable in normal conditions, and would decompose immediately into and chlorine.
On the other hand, cobalt chlorides can be obtained if the cobalt is bound also to other ligands of greater Lewis basicity than chloride, such as amines. For example, in the presence of ammonia, cobalt chloride is readily oxidised by atmospheric oxygen to hexamminecobalt chloride:
Similar reactions occur with other amines. These reactions are often performed in the presence of charcoal as a catalyst, or with hydrogen peroxide substituted for atmospheric oxygen. Other highly basic ligands, including carbonate, acetylacetonate, and oxalate, induce the formation of Co derivatives. Simple carboxylates and halides do not.
Unlike Co complexes, Co complexes are very slow to exchange ligands, so they are said to be kinetically inert. The German chemist Alfred Werner was awarded the Nobel prize in 1913 for his studies on a series of these cobalt compounds, work that led to an understanding of the structures of such coordination compounds.

Oxidation to cobalt(IV)

Reaction of 1-norbonyllithium with the ·THF in pentane produces the brown, thermally stable cobalt tetralkyl — a rare example of a stable transition metal/saturated alkane compound, different products are obtained in other solvents.

Health issues

Cobalt is essential for most higher forms of life, but more than a few milligrams each day is harmful. Although poisonings have rarely resulted from cobalt compounds, their chronic ingestion has caused serious health problems at doses far less than the lethal dose. In 1966, the addition of cobalt compounds to stabilize beer foam in Canada led to a peculiar form of toxin-induced cardiomyopathy, which came to be known as beer drinker's cardiomyopathy.
Furthermore, cobalt chloride is suspected of causing cancer as per the International Agency for Research on Cancer Monographs.
In 2005–06, cobalt chloride was the eighth-most-prevalent allergen in patch tests.

Other uses