Cobalt(II) chloride

Cobalt(II) chloride is an inorganic compound of cobalt and chlorine, with the formula CoCl₂. It is usually supplied as the hexahydrate CoCl₂·6H₂O, which is one of the most commonly used cobalt compounds in the lab.^[2]

The hexahydrate is purple, whereas the anhydrous form is sky blue. 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

Aqueous solutions of both CoCl₂ and the hydrate contain the species [Co(H₂O)₆]²⁺. They also contain chloride ions. In the solid state CoCl₂·6H₂O consists of the molecule trans-[CoCl₂(H₂O)₄] and two molecules of water of crystallization.^[3] This species dissolves readily in water and alcohol. Concentrated aqueous solutions are red at room temperature but become blue at higher temperatures.^[4] CoCl₂·6H₂O is deliquescent, and the anhydrous salt CoCl₂ is hygroscopic, readily converting to the hydrate.

Preparation

Hydrated cobalt chloride is prepared from cobalt(II) hydroxide or cobalt(II) carbonate and hydrochloric acid:

CoCO₃ + 2 HCl + 5 H₂O → Co(H₂O)₆Cl₂ + CO₂

Upon heating, the hexahydrate dehydrates.^[5]

Reactions

Generally, aqueous solutions of cobalt(II) chlorides behave like other cobalt(II) salts since these solutions consist of the [Co(H₂O)₆]²⁺ ion regardless of the anion. Such solutions give a precipitate of CoS upon treatment with H₂S. CoCl₂·6H₂O and CoCl₂ are weak Lewis acids. The adducts are usually either octahedral or tetrahedral. With pyridine (C
5H
5N), one obtains the octahedral complex:

CoCl₂·6H₂O + 4 C₅H₅N → CoCl₂(C₅H₅N)₄ + 6 H₂O

With triphenylphosphine (P(C
6H
5)
3), a tetrahedral complex results:

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

Salts of the anionic complex CoCl₄²⁻ can be prepared using tetraethylammonium chloride:^[6]

CoCl₂ + 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.

The structure of a cobalt(IV) coordination complex with the norbornyl anion

In the laboratory, cobalt(II) chloride serves as a common precursor to other cobalt compounds. 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. Reaction of 1-norbonyllithium with the CoCl₂·THF in pentane produces the brown, thermally stable cobalt(IV) tetralkyl^[7]^[8] — a rare example of a stable transition metal/saturated alkane compound,^[2] different products are obtained in other solvents.^[9]

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 presence of ammonia triggers the oxidation of cobalt(II) chloride to hexamminecobalt(III) chloride:

4 CoCl₂·6H₂O + 4 NH₄Cl + 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

Anhydrous cobalt(II) chloride

The existence of cobalt(III) chloride, CoCl₃, is disputed, although this compound is listed in some compendia.^[10] According to Greenwood and Earnshaw, the only stable binary compounds of cobalt and the halogens, excluding CoF₃, are the dihalides.^[2] Stated differently, CoCl₂ is unreactive toward Cl₂. The stability of Co(III) in solution is considerably increased in the presence of ligands of greater Lewis basicity than chloride, such as amines.

Safety

Patch test

In 2005–06, cobalt chloride was the eighth-most-prevalent allergen in patch tests (8.4%).^[11]

In popular culture

The NZ National Film Unit PICTORIAL PARADE No. 23 features cobalt chloride as the missing component in the pumicelands near Rotorua. Without this compound, no cattle would be able to survive on this land—only wild horses.

Other uses

Invisible ink
Cobalt chloride is an established chemical inducer of hypoxia-like responses such as erythropoiesis.^[12] Cobalt supplementation is not banned and therefore would not be detected by current anti-doping testing.^[13] Cobalt chloride is a banned substance under the Australian Thoroughbred Racing

References

↑ http://www.chemindustry.com/chemicals/0519906.html.
1 2 3 Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 0-08-037941-9.
↑ Wells, A. F. (1984), Structural Inorganic Chemistry (5th ed.), Oxford: Clarendon Press, ISBN 0-19-855370-6
↑ The Merck Index, 7th edition, Merck & Co, Rahway, New Jersey, USA, 1960.
↑ John Dallas Donaldson, Detmar Beyersmann, "Cobalt and Cobalt Compounds" in Ullmann's Encyclopedia of Industrial Chemistry, Wiley-VCH, Weinheim, 2005. doi:10.1002/14356007.a07_281.pub2
↑ Gill, N. S. & Taylor, F. B. (1967). "Tetrahalo Complexes of Dipositive Metals in the First Transition Series". Inorg. Synth. 9: 136–142. doi:10.1002/9780470132401.ch37.
↑ Barton K. Bower & Howard G. Tennent (1972). "Transition metal bicyclo[2.2.1]hept-1-yls". J. Am. Chem. Soc. 94 (7): 2512–2514. doi:10.1021/ja00762a056.
↑ Erin K. Byrne; Darrin S. Richeson & Klaus H. Theopold (1986). "Tetrakis(1-norbornyl)cobalt, a low spin tetrahedral complex of a first row transition metal". J. Chem. Soc., Chem. Commun. (19): 1491–1492. doi:10.1039/C39860001491.
↑ Erin K. Byrne; Klaus H. Theopold (1989). "Synthesis, characterization, and electron-transfer reactivity of norbornyl complexes of cobalt in unusually high oxidation states". J. Am. Chem. Soc. 111 (11): 3887–3896. doi:10.1021/ja00193a021.
↑ Handbook of Chemistry and Physics, 71st edition, CRC Press, Ann Arbor, Michigan, 1990.
↑ Zug KA, Warshaw EM, Fowler JF Jr, Maibach HI, Belsito DL, Pratt MD, Sasseville D, Storrs FJ, Taylor JS, Mathias CG, Deleo VA, Rietschel RL, Marks J. Patch-test results of the North American Contact Dermatitis Group 2005–2006. Dermatitis. 2009 May–Jun;20(3):149-60.
↑ W. Jelkmann: The disparate roles of cobalt in erythropoiesis, and doping relevance. Open Journal of Hematology, 2012, 3–6. http://rossscience.org/ojhmt/2075-907X-3-6.php
↑ Lippi G, Franchini M, Guidi GC (November 2005). "Cobalt chloride administration in athletes: a new perspective in blood doping?". Br J Sports Med. 39 (11): 872–3. PMC 1725077 . PMID 16244201. doi:10.1136/bjsm.2005.019232.

External links

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Cobalt compounds
CoBr₂ Co(CN)₂ CoCO₃ CoC₂O₄ CoCl₂ CoCl₃ Co(ClO₃)₂ CoF₂ CoF₃ CoI₂ Co(NO₃)₂ Co(NO₃)₃ CoO Co(OH)₂ CoS CoSO₄ CoSe Co₂O₃ Co₂P Co₃O₄ Co₃(PO₄)₂

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