Cobalt(II) chloride
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| Cobalt(II) chloride | |
|---|---|
anhydrous (left) and hexahydrate (right)  |
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| General | |
| Systematic name | Cobalt(II) chloride Cobalt dichloride |
| Other names | Cobaltous chloride |
| Molecular formula | CoCl2 |
| Molar mass | 129.84 g/mol (anhydrous) |
| Appearance | see text |
| CAS number | [7646-79-9] (anhydrous) |
| Properties | |
| Density and phase | 3.356 g/cm³, solid |
| Solubility in water | 45 g/100 ml (7 °C) 53 g/100ml (20 °C) |
| Melting point | 735°C |
| Boiling point | 1049°C (1322 K) |
| Structure | |
| Coordination geometry |
Octahedral |
| Crystal structure | CdCl2 structure |
| Hazards | |
| MSDS | External MSDS |
| EU classification | Toxic (T) Carc. Cat. 2 Dangerous for the environment (N) |
| NFPA 704 | |
| R-phrases | R49, R22, R42/43, R50/53 |
| S-phrases | S2, S22, S53, S45, S60, S61 |
| Flash point | non flammable |
| RTECS number | ? |
| Supplementary data page | |
| Structure and properties |
n, εr, etc. |
| Thermodynamic data |
Phase behaviour Solid, liquid, gas |
| Spectral data | UV, IR, NMR, MS |
| Related compounds | |
| Other anions | Cobalt(II) fluoride Cobalt(II) bromide Cobalt(II) iodide Cobalt(II) oxide |
| Other cations | Rhodium(III) chloride Iridium(III) chloride |
| Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa) Infobox disclaimer and references |
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Cobalt(II) chloride is the chemical compound with the formula CoCl2, although the term is used also to refer to the hexahydrate, which is a different chemical compound. CoCl2 is blue, and CoCl2·6H2O is deep magenta. Because of this dramatic color change and the ease of the hydration/dehydration reaction, "cobalt chloride" is used as an indicator for water. The magenta hexahydrate is probably the most familiar cobalt compound in the laboratory.
Aqueous solutions of both CoCl2 and the hydrate contain the species trans-[CoCl2(H2O)4]. This is maintained in the solid state by the hexahdrate, the remaining two water molecules in its formula unit being water of crystallization. This species dissolves readily in water and alcohol. It has the interesting property that a concentrated aqueous solution is red at room temperature, but becomes blue when heated. CoCl2·6H2O is deliquescent and the anhydrous salt CoCl2 is hygroscopic, readily converting to the hydrate.
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[edit] Chemical properties
CoCl2·6H2O and CoCl2 are weak Lewis acids that convert to many other complexes. These cobalt (II) complexes are usually either octahedral or tetrahedral. Examples include:
- CoCl2·6H2O + C5H5N → CoCl2(C5H5N)4 + 6 H2O
- CoCl2·6H2O + P(C6H5)3 → CoCl2{P(C6H5)3}2 + 6 H2O
- CoCl2 + 2 [(C2H5)4N]Cl → [(C2H5)4N)]2[CoCl4]
Otherwise, aqueous solutions of cobalt(II) chlorides behave like other cobalt(II) salts, such as precipitating CoS upon treatment with H2S.
[edit] Co(III) derivatives
In the presence of ammonia or amines, cobalt(II) is readily oxidised by atmospheric oxygen to give a variety of cobalt(III) complexes. For example:
The reaction is often performed in the presence of charcoal as a catalyst, or hydrogen peroxide is employed in place of air. 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.
[edit] CoCl3?
Regarding the existence of a simple cobalt(III) chloride, CoCl3, the literature is contradictory. The CRC 71 ed describes such a compound as known, but Greenwood & Earnshaw (which is more recent) states, "Apart from CoF3, the only known halides of cobalt are the dihalides."
[edit] Preparation
Cobalt(II) chloride may be prepared in its anhydrous form from cobalt metal and chlorine gas:
The hydrated form can be prepared from cobalt(II) hydroxide or cobalt(II) carbonate and hydrochloric acid.
[edit] Uses
A popular use for cobalt(II) chloride is for the detection of moisture, for example in drying agents such as silica gel. In the US calcium sulfate is sold as a drying agent under the trade name Drierite. When cobalt(II) chloride is added as an indicator, the drying agent is blue when still active, pink when exhausted, corresponding to the anhydrous and hydrated forms of CoCl2 respectively. Cobalt chloride paper is likewise used to detect the presence of water.
In the laboratory, cobalt(II) chloride serves as a starting point for the synthesis of a variety of cobalt compounds. For example, the reaction of 1-norbonyllithium with CoCl2 produces a brown, thermally stable cobalt(IV) tetralkyl — the only compound of its kind for which the detailed structure is fully known[3]:
Reaction of anhydrous CoCl2 with sodium cyclopentadienylide in THF gives the black sandwich compound cobaltocene. This 19-electron species is a good reducing agent, being readily oxidised to the yellow 18-electron cobaltacenium cation, which is isoelectronic with ferrocene.
[edit] References
- ↑ The Merck Index, 7th edition, Merck & Co, Rahway, New Jersey, USA, 1960.
- ↑ Handbook of Chemistry and Physics, 71st edition, CRC Press, Ann Arbor, Michigan, 1990.
- ↑ Greenwood, N. N.; A. Earnshaw (1997). Chemistry of the Elements, 2nd Edition, Oxford:Butterworth-Heinemann. ISBN 0-7506-3365-4.
- ↑ E. K. Byrne, D. S. Richeson, K. H. Theopold, Journal of the Chemical Society, Chemical Communications, 1491-2 (1986).
- ↑ A. F. Wells, Structural Inorganic Chemistry, 5th ed., Oxford University Press, Oxford, UK, 1984. ISBN 0-19-855370-6.
- ↑ Hill, Petrucci, McCreary, Perry, General Chemistry, 4th ed., Pearson/Prentice Hall, Upper Saddle River, New Jersey, USA. ISBN 0-13-140283-8.
