Cobalt(II) chloride

Cobalt(II) chloride is an inorganic compound of cobalt and chlorine, with the formula CoCl
2. It is a sky blue crystalline solid.

"CoCl2" redirects here. For the compound COCl2, see Phosgene.
Cobalt(II) chloride
Anhydrous
Hexahydrate

Structure of anhydrous compound

Structure of hexahydrate
Names
IUPAC name
Cobalt(II) chloride
Other names
Cobaltous chloride
Cobalt dichloride
Muriate of cobalt[1]
Identifiers
3D model (JSmol)
ChEBI
ChemSpider
ECHA InfoCard 100.028.718
EC Number
  • 231-589-4
RTECS number
  • GF9800000
UNII
UN number 3288
Properties
CoCl2
Molar mass 129.839 g/mol (anhydrous)
165.87 g/mol (dihydrate)
237.93 g/mol (hexahydrate)
Appearance blue crystals (anhydrous)
violet-blue (dihydrate)
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)[2]
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
+12,660·10−6 cm3/mol
Structure
CdCl2 structure
hexagonal (anhydrous)
monoclinic (dihydrate)
Octahedral (hexahydrate)
Hazards
Safety data sheet ICSC 0783
GHS pictograms
NFPA 704 (fire diamond)
Flash point Non-flammable
Lethal dose or concentration (LD, LC):
80 mg/kg (rat, oral)
Related compounds
Other anions
 Cobalt(II) fluoride
Cobalt(II) bromide
Cobalt(II) iodide
Other cations
 Rhodium(III) chloride
Iridium(III) chloride
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
N verify (what is YN ?)
Infobox references

The compound forms several hydrates CoCl
2•nH
2O, for n = 1, 2, 6, and 9. Claims of the formation of tri- and tetrahydrates have not been confirmed.[4] The dihydrate is purple and hexahydrate is pink. It is usually supplied as the hexahydrate CoCl
2·6H
2O, which is one of the most commonly used cobalt compounds in the lab.[5]

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

Anhydrous

At room temperature, anhydrous cobalt chloride has the cadmium(II) chloride structure (CdCl
2) (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.[2] The vapor pressure has been reported as 7.6 mm Hg at the melting point.[6]

Solutions

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.[4]

Diluted aqueous solutions of CoCl
2 contain the species [Co(H
2O)
6]2+
, besides chloride ions. Concentrated solutions are red at room temperature but become blue at higher temperatures.[7]

Hydrates

The crystal unit of the solid hexahydrate CoCl
2•6H
2O contains the neutral molecule trans-CoCl
2(H
2O)
4 and two molecules of water of crystallization.[8] This species dissolves readily in water and alcohol.

The anhydrous salt is hygroscopic and the hexahydrate is deliquescent. The dihydrate, CoCl2(H2O)2, is a coordination polymer. Each Co center is coordinated to four doubly bridging chloride ligands. The octahedron is completed by a pair of mutually trans aquo ligands.[9]

Subunit of CoCl2(H2O)2 lattice.

Preparation

Cobalt chloride can be prepared in aqueous solution from cobalt(II) hydroxide or cobalt(II) carbonate and hydrochloric acid:

CoCO
3 + 2 HCl(aq) → CoCl
2(aq) + CO
2
Co(OH)
2 + 2 HCl(aq) → CoCl
2(aq) + 2H
2O

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.[4]

The anhydrous compound can be prepared by heating the hydrates.[10] 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.[4]

Dehydration can also be effected with trimethylsilyl chloride:[11]

CoCl
2•6H
2O + 12 (CH
3)
3SiCl → CoCl
2 + 6[(CH
3)
3SiCl]
2O + 12 HCl

The anhydrous compound can be purified by sublimation in vacuum.[2]

Reactions

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
2O)
6]2+
 ion regardless of the anion. For example, such solutions give a precipitate of cobalt sulfide CoS upon treatment with hydrogen sulfide H
2S.

Complexed chlorides

The hexahydrate and the anhydrous salt are weak Lewis acids. The adducts are usually either octahedral or tetrahedral. With pyridine (C
5H
5N), one obtains an octahedral complex:

CoCl
2·6H
2O + 4 C
5H
5N → CoCl
2(C
5H
5N)
4 + 6 H
2O

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

CoCl
2·6H
2O + 2 P(C
6H
5)
3 → CoCl
2[P(C
6H
5)
3]
2 + 6 H
2O

Salts of the anionic complex CoCl42− can be prepared using tetraethylammonium chloride:[12]

CoCl
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.

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

Reaction of the anhydrous compound with sodium cyclopentadienide gives cobaltocene Co(C
5H
5)
2. This 19-electron species is a good reducing agent, being readily oxidised to the yellow 18-electron cobaltacenium cation [Co(C
5H
5)
2]+
.

Oxidation to cobalt(III)

Compounds of cobalt in the +3 oxidation state exist, such as cobalt(III) fluoride CoF
3, nitrate Co(NO
3)
3, and sulfate Co
2(SO
4)
3; however, cobalt(III) chloride CoCl
3 is not stable in normal conditions, and would decompose immediately into CoCl
2 and chlorine.[13]

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
2·6H
2O + 4 NH
4Cl + 20 NH
3 + O
2 → 4 [Co(NH
3)
6]Cl
3 + 26 H
2O

Similar reactions occur with other amines. These reactions are often performed in the presence of charcoal as a catalyst, or with hydrogen peroxide H
2O
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)

Reaction of 1-norbonyllithium with the CoCl
2·THF in pentane produces the brown, thermally stable cobalt(IV) tetralkyl[14][15] — a rare example of a stable transition metal/saturated alkane compound,[5] different products are obtained in other solvents.[16]

Health issues

Cobalt is essential for most higher forms of life, but more than a few milligrams each day is harmful. Although poisonings have rarely resulted from cobalt compounds, their chronic ingestion has caused serious health problems at doses far less than the lethal dose. In 1966, the addition of cobalt compounds to stabilize beer foam in Canada led to a peculiar form of toxin-induced cardiomyopathy, which came to be known as beer drinker's cardiomyopathy.[17][18][19]

Furthermore, cobalt(II) chloride is suspected of causing cancer (i.e., possibly carcinogenic, IARC Group 2B) as per the International Agency for Research on Cancer (IARC) Monographs.[20]

In 2005–06, cobalt chloride was the eighth-most-prevalent allergen in patch tests (8.4%).[21]

Other uses
  • Invisible ink: when suspended in solution, cobalt(II) chloride can be made to appear invisible on a surface; when that same surface is subsequently exposed to significant heat (such as from a handheld heat gun or lighter) the ink permanently/ irreversibly changes to blue.
  • Cobalt chloride is an established chemical inducer of hypoxia-like responses such as erythropoiesis. Cobalt supplementation is not banned and therefore would not be detected by current anti-doping testing.[22] Cobalt chloride is a banned substance under the Australian Thoroughbred Racing Board.[23]
  • Cobalt chloride is one method used to induce pulmonary arterial hypertension in animals for research and evaluation of treatment efficacy.

References
  1. "Cobalt muriate, CAS Number: 7646-79-9". www.chemindustry.com. Retrieved 19 April 2018.
  2. Wojakowska, A., Krzyżak, E. and Plińska, S. (2007): "Melting and high-temperature solid state transitions in cobalt(II) halides". Journal of Thermal Analysis and Calorimetry, volume 88, issue 2, pages 525-530. doi:10.1007/s10973-006-8000-9
  3. Santa Cruz Biotechnology: Cobalt(II) chloride
  4. M. T. Saugier, M. Noailly, R. Cohen-Adad, F. Paulik, and J. Paulik (1977): "Equilibres solide ⇄ liquide ⇆ vapeur du systeme binaire CoCl
    2    -H
    2    O    " Journal of Thermal Analysis, volume 11, issue 1, pages 87–100. doi:10.1007/BF02104087 Note: the lowest point of fig.6 is inconsistent with fig.7; probably should be at -27.8 C instead of 0 C.
  5. Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 978-0-08-037941-8.
  6. Yuzo Saeki, Ryoko Matsuzaki, Naomi Aoyama (1977): "The vapor pressure of cobalt dichloride". Journal of the Less Common Metals, volume 55, issue 2, pages 289-291. doi:10.1016/0022-5088(77)90204-1
  7. The Merck Index, 7th edition, Merck & Co, Rahway, New Jersey, USA, 1960.
  8. Wells, A. F. (1984), Structural Inorganic Chemistry (5th ed.), Oxford: Clarendon Press, ISBN 0-19-855370-6
  9. Morosin, B.; Graeber, E. J. (1965). "Crystal structures of manganese(II) and iron(II) chloride dihydrate". Journal of Chemical Physics. 42 (3): 898–901. Bibcode:1965JChPh..42..898M. doi:10.1063/1.1696078.
  10. 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
  11. Philip Boudjouk, Jeung-Ho So (1992). "Solvated and Unsolvated Anhydrous Metal Chlorides from Metal Chloride Hydrates". Inorganic Syntheses. Inorg. Synth. Inorganic Syntheses. 29. pp. 108–111. doi:10.1002/9780470132609.ch26. ISBN 9780470132609.CS1 maint: uses authors parameter (link)
  12. Gill, N. S. & Taylor, F. B. (1967). Tetrahalo Complexes of Dipositive Metals in the First Transition Series. Inorg. Synth. Inorganic Syntheses. 9. pp. 136–142. doi:10.1002/9780470132401.ch37. ISBN 9780470132401.
  13. Handbook of Chemistry and Physics, 71st edition, CRC Press, Ann Arbor, Michigan, 1990.
  14. 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.
  15. 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.
  16. 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.
  17. Morin Y; Tětu A; Mercier G (1969). "Quebec beer-drinkers' cardiomyopathy: Clinical and hemodynamic aspects". Annals of the New York Academy of Sciences. 156 (1): 566–576. Bibcode:1969NYASA.156..566M. doi:10.1111/j.1749-6632.1969.tb16751.x. PMID 5291148. S2CID 7422045.
  18. Barceloux, Donald G. & Barceloux, Donald (1999). "Cobalt". Clinical Toxicology. 37 (2): 201–216. doi:10.1081/CLT-100102420. PMID 10382556.
  19. 11.1.5 The unusual type of myocardiopathy recognized in 1965 and 1966 in Quebec (Canada), Minneapolis (Minnesota), Leuven (Belgium), and Omaha (Nebraska) was associated with episodes of acute heart failure (e/g/, 50 deaths among 112 beer drinkers).
  20. [PDF
  21. 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.
  22. 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. doi:10.1136/bjsm.2005.019232. PMC 1725077. PMID 16244201.
  23. Bartley, Patrick (6 February 2015). "Cobalt crisis turns the eyes of the world onto Australian racing". The Sydney Morning Herald.
Salts and covalent derivatives of the chloride ion
HCl He
LiCl BeCl2 BCl3
B2Cl4		 CCl4 NCl3
ClN3		 Cl2O
ClO2
Cl2O7		 ClF
ClF3
ClF5		 Ne
NaCl MgCl2 AlCl
AlCl3		 SiCl4 P2Cl4
PCl3
PCl5		 S2Cl2
SCl2
SCl4		 Cl2 Ar
KCl CaCl
CaCl2		 ScCl3 TiCl2
TiCl3
TiCl4		 VCl2
VCl3
VCl4
VCl5		 CrCl2
CrCl3
CrCl4		 MnCl2 FeCl2
FeCl3		 CoCl2
CoCl3		 NiCl2 CuCl
CuCl2		 ZnCl2 GaCl2
GaCl3		 GeCl2
GeCl4		 AsCl3
AsCl5		 Se2Cl2
SeCl4		 BrCl KrCl
RbCl SrCl2 YCl3 ZrCl3
ZrCl4		 NbCl3
NbCl4
NbCl5		 MoCl2
MoCl3
MoCl4
MoCl5
MoCl6		 TcCl3
TcCl4		 RuCl3 RhCl3 PdCl2 AgCl CdCl2 InCl
InCl2
InCl3		 SnCl2
SnCl4		 SbCl3
SbCl5		 Te3Cl2
TeCl4		 ICl
ICl3		 XeCl
XeCl2
XeCl4
CsCl BaCl2   HfCl4 TaCl5 WCl2
WCl3
WCl4
WCl5
WCl6		 Re3Cl9
ReCl4
ReCl5
ReCl6		 OsCl4 IrCl2
IrCl3
IrCl4		 PtCl2
PtCl4		 AuCl
AuCl3		 Hg2Cl2,
HgCl2		 TlCl PbCl2,
PbCl4		 BiCl3 PoCl2,
PoCl4		 AtCl RnCl2
FrCl RaCl2   RfCl4 Db Sg Bh Hs Mt Ds Rg Cn Nh Fl Mc Lv Ts Og
↓
LaCl3 CeCl3 PrCl3 NdCl2,
NdCl3		 PmCl3 SmCl2,
SmCl3		 EuCl2,
EuCl3		 GdCl3 TbCl3 DyCl2,
DyCl3		 HoCl3 ErCl3 TmCl2
TmCl3		 YbCl2
YbCl3		 LuCl3
AcCl3 ThCl4 PaCl5 UCl3
UCl4
UCl5
UCl6		 NpCl4 PuCl3 AmCl2
AmCl3		 CmCl3 BkCl3 CfCl3 EsCl3 Fm Md No LrCl3
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