Cobalt(II) chloride is an inorganic compound, a salt of cobalt and chlorine, with the formula CoCl 2. The compound forms several hydratesCoCl 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]
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]
Cobalt chloride is fairly soluble in water. Under atmospheric pressure, the mass concentration of a saturated solutionofCoCl 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]
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 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 contained, the decomposition occurs at about 115°C, 145°C, and 195°C, respectively.[4]
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 sulfideCoS upon treatment with hydrogen sulfideH 2S.[citation needed]
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.
Reaction of the anhydrous compound with sodium cyclopentadienide gives cobaltoceneCo(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]+ .
Compounds of cobalt in the +3 oxidation state exist, such as cobalt(III) fluorideCoF 3, nitrateCo(NO 3) 3, and sulfateCo 2(SO 4) 3; however, cobalt(III) chlorideCoCl 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 oxygentohexamminecobalt(III) chloride:
Similar reactions occur with other amines. These reactions are often performed in the presence of charcoal as a catalyst, or with hydrogen peroxideH 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]
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]
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]
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]
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.
^ abcWojakowska, 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.
^ abcdeM. T. Saugier, M. Noailly, R. Cohen-Adad, F. Paulik, and J. Paulik (1977): "Equilibres solide ⇄ liquide ⇆ vapeur du systeme binaire CoCl 2-H 2O" 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.
^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
^The Merck Index, 7th edition, Merck & Co, Rahway, New Jersey, USA, 1960.
^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.
^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
^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. ISBN9780470132609.
^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.
^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.
^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.
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