Cobalt(II) sulfate: Difference between revisions
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The hexahydrate is a [[metal aquo complex]] consisting of [[octahedral molecular geometry|octahedral]] [Co(H<sub>2</sub>O)<sub>6</sub>]<sup>2+</sup> ions associated with [[sulfate]] [[anions]].<ref>Elerman, Y. "Refinement of the crystal structure of CoSO<sub>4</sub>*6H<sub>2</sub>O" Acta Crystallographica Section C 1988, volume 44, p599-p601. {{doi|10.1107/S0108270187012447}}</ref> |
The hexahydrate is a [[metal aquo complex]] consisting of [[octahedral molecular geometry|octahedral]] [Co(H<sub>2</sub>O)<sub>6</sub>]<sup>2+</sup> ions associated with [[sulfate]] [[anions]].<ref>Elerman, Y. "Refinement of the crystal structure of CoSO<sub>4</sub>*6H<sub>2</sub>O" Acta Crystallographica Section C 1988, volume 44, p599-p601. {{doi|10.1107/S0108270187012447}}</ref> |
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==Uses== |
==Uses and reactions== |
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Cobalt sulfates are important as intermediates in the extraction of cobtal from its ores. Thus, crushed, partially refined ores are treated with sulfuric acid to give red-colored solutions containing cobalt sulfate.<ref name=Ullmann>{{Ullmann|doi=10.1002/14356007.a07_281.pub2|title=Cobalt and Cobalt Compounds|year=2005|last1=Donaldson|first1=John Dallas|last2=Beyersmann|first2=Detmar|isbn=3527306730}}</ref> |
Cobalt sulfates are important as intermediates in the extraction of cobtal from its ores. Thus, crushed, partially refined ores are treated with sulfuric acid to give red-colored solutions containing cobalt sulfate.<ref name=Ullmann>{{Ullmann|doi=10.1002/14356007.a07_281.pub2|title=Cobalt and Cobalt Compounds|year=2005|last1=Donaldson|first1=John Dallas|last2=Beyersmann|first2=Detmar|isbn=3527306730}}</ref> |
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Hydrated cobalt(II) sulfate is used in the preparation of [[pigment]]s, as well as in the manufacture of other cobalt salts. Cobalt pigment is used in porcelains and glass. Cobalt(II) sulfate is used in storage [[Battery (electricity)|batteries]] and [[electroplating]] baths, [[sympathetic ink]]s, and as an additive to soils and animal feeds. For these purposes, the cobalt sulfate is produced by treating cobalt oxide with sulfuric acid.<ref name=Ullmann/> |
Hydrated cobalt(II) sulfate is used in the preparation of [[pigment]]s, as well as in the manufacture of other cobalt salts. Cobalt pigment is used in porcelains and glass. Cobalt(II) sulfate is used in storage [[Battery (electricity)|batteries]] and [[electroplating]] baths, [[sympathetic ink]]s, and as an additive to soils and animal feeds. For these purposes, the cobalt sulfate is produced by treating cobalt oxide with sulfuric acid.<ref name=Ullmann/> |
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Being commonly available commercially, the heptahydrate is a routine source of cobalt in [[coordination chemistry]].<ref>{{cite journal |doi=10.1002/9780470132371.ch58|title=Resolution of the Tris(ethylenediamine)cobalt(III) Ion|journal=Inorganic Syntheses|year=1950|volume=6|last1=Broomhead|first1=J. A.|last2=Dwyer|first2=F. P.|last3=Hogarth|first3=J. W.|pages=183–186}}</ref> |
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==Natural occurrence== |
==Natural occurrence== |
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Revision as of 13:15, 13 March 2021
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| Names | |
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| IUPAC name
Cobalt(II) sulfate
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| Other names
Cobaltous sulfate
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| Identifiers | |
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3D model (JSmol)
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| ChEBI | |
| ChemSpider | |
| ECHA InfoCard | 100.030.291 |
| EC Number |
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| KEGG | |
PubChem CID
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| RTECS number |
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| UNII |
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CompTox Dashboard (EPA)
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| Properties | |
| CoSO4·(H2O)7 | |
| Molar mass | 154.996 g/mol (anhydrous) 173.01 g/mol (monohydrate) 263.08 g/mol (hexahydrate) 281.103 g/mol (heptahydrate) |
| Appearance | reddish crystalline (anhydrous, monohydrate) pink salt (hexahydrate) |
| Odor | odorless (heptahydrate) |
| Density | 3.71 g/cm3 (anhydrous) 3.075 g/cm3 (monohydrate) 2.019 g/cm3 (hexahydrate) 1.948 g/cm3 (heptahydrate) |
| Melting point | 735 °C (1,355 °F; 1,008 K) |
| anhydrous: 36.2 g/100 mL (20 °C) 38.3 g/100 mL (25 °C) 84 g/100 mL (100 °C) heptahydrate: 60.4 g/100 mL (3 °C) 67 g/100 mL (70 °C) | |
| Solubility | anhydrous: 1.04 g/100 mL (methanol, 18 °C) insoluble in ammonia heptahydrate: 54.5 g/100 mL (methanol, 18 °C) |
| +10,000·10−6 cm3/mol | |
Refractive index (nD)
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1.639 (monohydrate) 1.540 (hexahydrate) 1.483 (heptahydrate) |
| Structure | |
| orthorhombic (anhydrous) monoclinic (monohydrate, heptahydrate) | |
| Hazards | |
| GHS labelling: | |
  
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| Danger | |
| H302, H317, H334, H341, H350, H360, H410 | |
| P201, P202, P261, P264, P270, P272, P273, P280, P281, P285, P301+P312, P302+P352, P304+P341, P308+P313, P321, P330, P333+P313, P342+P311, P363, P391, P405, P501 | |
| NFPA 704 (fire diamond) | |
| Flash point | Non-flammable |
| Lethal dose or concentration (LD, LC): | |
LD50 (median dose)
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424 mg/kg (oral, rat) |
| Safety data sheet (SDS) | JT Baker MSDS |
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) sulfate is any of the inorganic compounds with the formula CoSO4(H2O)x. Usually cobalt sulfate refers to the hydrate CoSO4.7H2O, which is one of the most commonly available salts of cobalt.
The heptahydrate is a red solid that is soluble in water and methanol. Since cobalt(II) has an odd number of electrons, it salts are paramagnetic.
Preparation, and structure
It forms by the reaction of metallic cobalt, its oxide, hydroxide, or carbonate with aqueous sulfuric acid:[1]
- Co + H2SO4 + 7 H2O → CoSO4(H2O)7 + H2
- CoO + H2SO4 + 6 H2O → CoSO4(H2O)7
The hexahydrate is a metal aquo complex consisting of octahedral [Co(H2O)6]2+ ions associated with sulfate anions.[2]
Uses and reactions
Cobalt sulfates are important as intermediates in the extraction of cobtal from its ores. Thus, crushed, partially refined ores are treated with sulfuric acid to give red-colored solutions containing cobalt sulfate.[1]
Hydrated cobalt(II) sulfate is used in the preparation of pigments, as well as in the manufacture of other cobalt salts. Cobalt pigment is used in porcelains and glass. Cobalt(II) sulfate is used in storage batteries and electroplating baths, sympathetic inks, and as an additive to soils and animal feeds. For these purposes, the cobalt sulfate is produced by treating cobalt oxide with sulfuric acid.[1]
Being commonly available commercially, the heptahydrate is a routine source of cobalt in coordination chemistry.[3]
Natural occurrence
Rarely, cobalt(II) sulfate is found in form of few crystallohydrate minerals, occurring among oxidation zones containing primary Co minerals (like skutterudite or cobaltite). These minerals are: biebierite (heptahydrate), moorhouseite (Co,Ni,Mn)SO4.6H2O, aplowite (Co,Mn,Ni)SO4.4H2O and cobaltkieserite (monohydrate).[4][5][6][5] the hexahydrate - as moorhouseite,[7][5]
Health issues
Cobalt is an essential mineral for mammals, but more than a few micrograms per 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 a cobalt compound 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.[8][9][10]
Furthermore, cobalt(II) sulfate is suspected of causing cancer (i.e., possibly carcinogenic, IARC Group 2B) as per the International Agency for Research on Cancer (IARC) Monographs.[11]
References
- ^ a b c Donaldson, John Dallas; Beyersmann, Detmar (2005). "Cobalt and Cobalt Compounds". Ullmann's Encyclopedia of Industrial Chemistry. Weinheim: Wiley-VCH. doi:10.1002/14356007.a07_281.pub2. ISBN 3527306730.
- ^ Elerman, Y. "Refinement of the crystal structure of CoSO4*6H2O" Acta Crystallographica Section C 1988, volume 44, p599-p601. doi:10.1107/S0108270187012447
- ^ Broomhead, J. A.; Dwyer, F. P.; Hogarth, J. W. (1950). "Resolution of the Tris(ethylenediamine)cobalt(III) Ion". Inorganic Syntheses. 6: 183–186. doi:10.1002/9780470132371.ch58.
- ^ https://www.mindat.org/min-26414.html
- ^ a b c https://www.ima-mineralogy.org/Minlist.htm
- ^ https://www.mindat.org/min-664.html
- ^ https://www.mindat.org/min-2776.html
- ^ 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.
- ^ Barceloux, Donald G. & Barceloux, Donald (1999). "Cobalt". Clinical Toxicology. 37 (2): 201–216. doi:10.1081/CLT-100102420. PMID 10382556.
- ^ 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).
- ^ "Cobalt in Hard Metals and Cobalt Sulfate, Gallium Arsenide, Indium Phosphide and Vanadium Pentoxide" (PDF). IARC Monographs on the Evaluation of Carcinogenic Risks to Humans.