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Cobalt(II) chloride

Cobalt(II) chloride
IUPAC name Cobalt(II) chloride
Cobalt dichloride
Other names Cobaltous chloride
CAS number [7646-79-9] (anhydrous)
Molecular formula CoCl2
Molar mass Anhydrous 129.84 g/mol

Hexahydrate 237.93 g/mol

Appearance see text
Density 3.356 g/cm³, solid
Melting point


Boiling point

1049°C (1322 K)

Solubility in water 45 g/100 ml (7 °C)
53 g/100ml (20 °C)
Crystal structure CdCl2 structure
EU classification Toxic (T)
Carc. Cat. 2
Dangerous for
the environment (N)
R-phrases R49, R22, R42/43, R50/53
S-phrases (S2), S22, S53, S45,
S60, S61
Flash point non flammable
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

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.



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.


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.


  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.
This article is licensed under the GNU Free Documentation License. It uses material from the Wikipedia article "Cobalt(II)_chloride". A list of authors is available in Wikipedia.
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