Cobalt(II) chloride

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 rose. Because of this dramatic color change and the ease of the hydration/dehydration reaction, "cobalt chloride" is used as an indicator for water. The rose hexahydrate is one the most common cobalt compounds in the laboratory.

Aqueous solutions of both CoCl₂ and the hydrate contain the species [Co(H₂O)₆]²⁺. In the solid state CoCl₂·6H₂O contains trans-[CoCl₂(H₂O)₄]·2H₂O, 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.^[1] CoCl₂·6H₂O is deliquescent and the anhydrous salt CoCl₂ is hygroscopic, readily converting to the hydrate.

Additional recommended knowledge

Daily Sensitivity Test

Recognize and detect the effects of electrostatic charges on your balance

Essential Laboratory Skills Guide

Preparation

Cobalt(II) chloride can be prepared in its anhydrous form from cobalt metal and chlorine gas:

Co(s) + Cl₂(g) → CoCl₂(s)

The hydrated form can be prepared from cobalt(II) hydroxide or cobalt(II) carbonate and hydrochloric acid.

Uses

A common use for cobalt(II) chloride was for the detection of moisture, for example in drying agents such as silica gel this use was disscontinued due to the carcenogenic nature of cobalt salts.^{[citation needed]} 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 anhydrous and hydrated CoCl₂, respectively. Similarly, paper impregnated with cobalt chloride, known as "cobalt chloride paper" is used to detect the presence of water.

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 + 4 C₅H₅N → CoCl₂(C₅H₅N)₄ + 6 H₂O

CoCl₂·6H₂O + 2 P(C₆H₅)₃ → CoCl₂{P(C₆H₅)₃}₂ + 6 H₂O

CoCl₂ + 2 [(C₂H₅)₄N]Cl → [(C₂H₅)₄N)]₂[CoCl₄]^[2]

Otherwise, aqueous solutions of cobalt(II) chlorides behave like other cobalt(II) salts, such as precipitating CoS upon treatment with H₂S.

In the laboratory, cobalt(II) chloride serves as a standard precursor for the synthesis of other cobalt compounds. For example, the reaction of 1-norbonyllithium with CoCl₂ produces the brown, thermally stable cobalt(IV) tetralkyl^[3][4] — the only compound of its kind for which the detailed structure is fully known^[5]:

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.

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:

4 CoCl₂·6H₂O + 4 [NH₄]Cl + 20 NH₃ + O₂ → 4 [Co(NH₃)₆]Cl₃ + 26 H₂O

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.

Instability of CoCl₃

The existence of cobalt(III) chloride, CoCl₃, is disputed, although it is listed in some compendia.^[6] An authoritative monograph^[5] states, "Apart from CoF₃, the only known halides of cobalt are the dihalides." The reduction potential for Co³⁺ + e^- → Co²⁺ is more favorable (+1.92 V) than the reduction Cl₂ to Cl^- (+1.36 V). This analysis suggests also that the naked cation Co³⁺ would oxidize chloride to chlorine, precluding the formation of CoCl₃. Stated differently, CoCl₂ is unreactive toward Cl₂. This analysis changes considerably in the presence of ligands of greater Lewis basicity than chloride, such as amines.

References

1. ^ The Merck Index, 7th edition, Merck & Co, Rahway, New Jersey, USA, 1960.
2. ^ Gill, N. S. and Taylor, F. B., "Tetrahalo Complexes of Dipositive Metals in the First Transition Series", Inorganic Syntheses, 1967, volume 9, pages 136-142.
3. ^ Barton K. Bower and Howard G. Tennent "Transition metal bicyclo[2.2.1]hept-1-yls" Journal of American Chemical Society 1972, Volume 94, pp 2512 - 2514; DOI: 10.1021/ja00762a056
4. ^ Erin K. Byrne, Darrin S. Richeson and Klaus H. Theopold (1986). "Tetrakis(1-norbornyl)cobalt, a low spin tetrahedral complex of a first row transition metal". J. Chem. Soc., Chem. Commun.: 1491 - 1492. doi:10.1039/C39860001491.
5. ^ ^{a b} Greenwood, N. N.; Earnshaw, A. (1997). Chemistry of the Elements, 2nd Edition, Oxford:Butterworth-Heinemann. ISBN 0-7506-3365-4. 
6. ^ Handbook of Chemistry and Physics, 71st edition, CRC Press, Ann Arbor, Michigan, 1990.