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Cobalt(II) chloride

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"CoCl₂" redirects here. For the compound with molecular formula 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 Edit this at Wikidata
EC Number
  • 231-589-4
RTECS number
  • GF9800000
UNII
UN number 3288
  • InChI=1S/2ClH.Co/h2*1H;/q;;+2/p-2 checkY
    Key: GVPFVAHMJGGAJG-UHFFFAOYSA-L checkY
  • InChI=1/2ClH.Co/h2*1H;/q;;+2/p-2
    Key: GVPFVAHMJGGAJG-NUQVWONBAU
  • anhydrous: Cl[Co]Cl
  • hexahydrate: Cl[Co-4](Cl)([OH2+])([OH2+])([OH2+])[OH2+].O.O
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
GHS labelling:
GHS06: Toxic GHS08: Health hazard GHS09: Environmental hazard
NFPA 704 (fire diamond)
[3]
NFPA 704 four-colored diamondHealth 3: Short exposure could cause serious temporary or residual injury. E.g. chlorine gasFlammability 0: Will not burn. E.g. waterInstability 0: Normally stable, even under fire exposure conditions, and is not reactive with water. E.g. liquid nitrogenSpecial hazards (white): no code
3
0
0
Flash point Non-flammable
Lethal dose or concentration (LD, LC):
80 mg/kg (rat, oral)
Safety data sheet (SDS) ICSC 0783
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).
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Cobalt(II) chloride is an inorganic compound, a salt of cobalt and chlorine, with the formula CoCl
2. 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 anhydrous form is a blue crystalline solid; the dihydrate is purple and the hexahydrate is pink. Commercial samples are usually the hexahydrate, which is one of the most commonly used cobalt salts in the lab.[5]

Properties

[edit]

Anhydrous

[edit]

At room temperature, anhydrous cobalt chloride has the cadmium 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 mmHg at the melting point.[6]

Solutions

[edit]

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

[edit]
Subunit of CoCl2(H2O)2 lattice.

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.[citation needed] 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]

Preparation

[edit]

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 + H
2O
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.[4] On slow heating in an open container, so that the water vapor pressure over the solid is practically zero, 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 125°C, respectively. If the partial pressure of the water vapor is in equilibrium with the solid, as in a confined but not pressurized container, the decomposition occurs at about 115°C, 145°C, and 195°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

[edit]

In the laboratory, cobalt(II) chloride serves as a common precursor to other cobalt compounds. Generally, diluted 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.[citation needed]

Complexed chlorides

[edit]

The hexahydrate and the anhydrous salt are weak Lewis acids. The adducts are usually either octahedral or tetrahedral. It forms an octahedral complex with pyridine (C
5H
5N):[12]

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:[13]

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

The tetrachlorocobaltate ion [CoCl4]2− is the blue ion that forms upon addition of hydrochloric acid to aqueous solutions of hydrated cobalt chloride, which are pink.

Reduction

[edit]
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 cobaltocenium cation [Co(C
5H
5)
2]+
.

Oxidation to cobalt(III)

[edit]

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.[14]

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.[citation needed]

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.[citation needed]

Oxidation to cobalt(IV)

[edit]

Reaction of 1-norbornyllithium with the CoCl
2·THF in pentane produces the brown, thermally stable tetrakis(1-norbornyl)cobalt(IV)[15][16] — a rare example of a stable transition metal/saturated alkane compound,[5] different products are obtained in other solvents.[17]

Moisture indication

[edit]
The deep blue colour of this moisture indicating silica gel is due to cobalt chloride. When hydrated the colour changes to a light pink/purple.

Cobalt chloride is a common visual moisture indicator due to its distinct colour change when hydrated. The colour change is from some shade of blue when dry, to a pink when hydrated, although the shade of colour depends on the substrate and concentration. It is impregnated into paper to make test strips for detecting moisture in solutions, or more slowly, in air/gas. Desiccants such as silica gel can incorporate cobalt chloride to indicate when it is "spent" (i.e. hydrated).[18]

Health issues

[edit]

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.[19][20][21]

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.[22]

In 2005–06, cobalt chloride was the eighth-most-prevalent allergen in patch tests (8.4%).[23]

Other uses

[edit]
  • 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 reversibly changes to blue.[24]
  • Cobalt chloride is an established chemical inducer of hypoxia-like responses such as erythropoiesis.[citation needed] Cobalt supplementation is not banned and therefore would not be detected by current anti-doping testing.[25] Cobalt chloride is a banned substance under the Australian Thoroughbred Racing Board.[26]
  • Cobalt chloride is one method used to induce pulmonary arterial hypertension in animals for research and evaluation of treatment efficacy.

References

[edit]
  1. ^ "Cobalt muriate, CAS Number: 7646-79-9". www.chemindustry.com. Archived from the original on 28 May 2019. Retrieved 19 April 2018.
  2. ^ a b c Wojakowska, A.; Krzyżak, E.; Plińska, S. (2007). "Melting and high-temperature solid state transitions in cobalt(II) halides". Journal of Thermal Analysis and Calorimetry. 88 (2): 525–530. doi:10.1007/s10973-006-8000-9.
  3. ^ Santa Cruz Biotechnology: Cobalt(II) chloride
  4. ^ a b c d e 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. ^ a b 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 (2007). "Solvated and Unsolvated Anhydrous Metal Chlorides from Metal Chloride Hydrates". Inorganic Syntheses. Vol. 29. pp. 108–111. doi:10.1002/9780470132609.ch26. ISBN 9780470132609.
  12. ^ Long, Gary J.; Clarke, Peter J. (1978). "Crystal and Molecular Structures of trans-Tetrakis(pyridine)dichloroiron(II), -Nickel(II), and -Cobalt(II) and trans-Tetrakis(pyridine)dichloroiron(II) Monohydrate". Inorganic Chemistry. 17 (6): 1394–1401. doi:10.1021/ic50184a002.
  13. ^ Gill, N. S. & Taylor, F. B. (1967). "Tetrahalo Complexes of Dipositive Metals in the First Transition Series". Inorganic Syntheses. Vol. 9. pp. 136–142. doi:10.1002/9780470132401.ch37. ISBN 9780470132401.
  14. ^ Handbook of Chemistry and Physics, 71st edition, CRC Press, Ann Arbor, Michigan, 1990.
  15. ^ 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.
  16. ^ 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.
  17. ^ 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.
  18. ^ Solomon, M. E. (February 1945). "The use of cobalt salts as indicators of humidity and moisture". Annals of Applied Biology. 32 (1): 75–85. doi:10.1111/j.1744-7348.1945.tb06763.x. ISSN 0003-4746.
  19. ^ 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.
  20. ^ Barceloux, Donald G. & Barceloux, Donald (1999). "Cobalt". Clinical Toxicology. 37 (2): 201–216. doi:10.1081/CLT-100102420. PMID 10382556.
  21. ^ 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).
  22. ^ [PDF
  23. ^ 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.
  24. ^ "Making invisible ink". Royal Society of Chemistry. Retrieved 7 April 2023.
  25. ^ 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): 872–873. doi:10.1136/bjsm.2005.019232. PMC 1725077. PMID 16244201.
  26. ^ Bartley, Patrick (6 February 2015). "Cobalt crisis turns the eyes of the world onto Australian racing". The Sydney Morning Herald.
[edit]
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