Structure of anhydrous compound
Structure of hexahydrate
Muriate of cobalt
3D model (JSmol)
|Molar mass||129.839 g/mol (anhydrous) |
165.87 g/mol (dihydrate)
237.93 g/mol (hexahydrate)
|Appearance||blue crystals (anhydrous) |
rose red crystals (hexahydrate)
|Density||3.356 g/cm3 (anhydrous) |
2.477 g/cm3 (dihydrate)
1.924 g/cm3 (hexahydrate)
|Melting point|| 726 °C (1,339 °F; 999 K) ±2 (anhydrous) |
140 °C (monohydrate)
100 °C (dihydrate)
86 °C (hexahydrate)
|Boiling point||1,049 °C (1,920 °F; 1,322 K)|
|43.6 g/100 mL (0 °C) |
45 g/100 mL (7 °C)
52.9 g/100 mL (20 °C)
105 g/100 mL (96 °C)
|Solubility||38.5 g/100 mL (methanol) |
8.6 g/100 mL (acetone)
soluble in ethanol, pyridine, glycerol
|hexagonal (anhydrous) |
|Safety data sheet||ICSC 0783|
|Lethal dose or concentration (LD, LC):|
LD50 (median dose)
|80 mg/kg (rat, oral)|
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
|what is ?)(|
The compound forms several hydrates CoCl
2O, 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 CoCl
2O, which is one of the most commonly used cobalt compounds in the lab.
At room temperature, anhydrous cobalt chloride has the CdCl
2 structure (R3m) in which the cobalt(II) ions are octahedrally coordinated. At about 706 C (20 degrees below the melting point), the coordination is believed to change to tetrahedral. The vapor pressure has been reported as 7.6 mm Hg at the melting point.
Cobalt chloride is fairly soluble in water. Under atmospheric pressure, the mass concentration of a saturated solution of CoCl
2 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.
The crystal unit of the solid hexahydrate CoCl
2O contains the neutral molecule trans-CoCl
4 and two molecules of water of crystallization. This species dissolves readily in water and alcohol.
3 + 2 HCl(aq) → CoCl
2(aq) + CO
2 + 2 HCl(aq) → CoCl
2(aq) + 2H
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.
2O + 12 (CH
3SiCl → CoCl
2 + 6[(CH
2O + 12 HCl
The anhydrous compound can be purified by sublimation in vacuum.
In the laboratory, cobalt(II) chloride serves as a common precursor to other cobalt compounds. Generally, aqueous solutions of the salt behave like other cobalt(II) salts since these solutions consist of the [Co(H
ion regardless of the anion. For example, such solutions give a precipitate of Cobalt sulfide CoS upon treatment with hydrogen sulfide H
2O + 4 C
5N → CoCl
4 + 6 H
With triphenylphosphine (P(C
3), a tetrahedral complex results:
2O + 2 P(C
3 → CoCl
2 + 6 H
Salts of the anionic complex CoCl42− can be prepared using tetraethylammonium chloride:
2 + 2 [(C2H5)4N]Cl → [(C2H5)4N)]2[CoCl4]
The [CoCl4]2− ion is the blue ion that forms upon addition of hydrochloric acid to aqueous solutions of hydrated cobalt chloride, which are pink.
Reaction of the anhydrous compound with sodium cyclopentadienide gives cobaltocene Co(C
2. This 19-electron species is a good reducing agent, being readily oxidised to the yellow 18-electron cobaltacenium cation [Co(C
Oxidation to cobalt(III)Edit
Compounds of cobalt in the +3 oxidation state exist, such as cobalt(III) fluoride CoF
3, nitrate Co(NO
3, and sulfate Co
3; however, cobalt(III) chloride CoCl
3 is not stable in normal conditions, and would decompose immediately into CoCl
2 and chlorine.
On the other hand, cobalt(III) 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(II) chloride is readily oxidised by atmospheric oxygen to hexamminecobalt(III) chloride:
- 4 CoCl
2O + 4 NH
4Cl + 20 NH
3 + O
2 → 4 [Co(NH
3 + 26 H
Similar reactions occur with other amines. These reactions are often performed in the presence of charcoal as a catalyst, or with hydrogen peroxide H
2 substituted for atmospheric oxygen. Other highly basic ligands, including carbonate, acetylacetonate, and oxalate, induce the formation of Co(III) derivatives. Simple carboxylates and halides do not.
Unlike Co(II) complexes, Co(III) 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(III) compounds, work that led to an understanding of the structures of such coordination compounds.
Oxidation to cobalt(IV)Edit
Reaction of 1-norbonyllithium with the CoCl
2·THF in pentane produces the brown, thermally stable cobalt(IV) tetralkyl — a rare example of a stable transition metal/saturated alkane compound, different products are obtained in other solvents.
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