Cobalt(II) perchlorate
 Hexahydrate
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| Names | |
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| Systematic IUPAC name
Cobalt(II) diperchlorate | |
Other names
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| Identifiers | |
3D model (JSmol)
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| ChemSpider | |
| ECHA InfoCard | 100.033.307 |
| EC Number |
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PubChem CID
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CompTox Dashboard (EPA)
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| Properties | |
| Co(ClO4)2 | |
| Molar mass | 257.83 g/mol |
| Appearance | Pink solid (anhydrous) Dark-red crystals (hexahydrate) |
| Density | 3.33 g/cm3 |
| Melting point | 170 °C (338 °F; 443 K)[1] (decomposition, hexahydrate) |
| 113 g/100 mL (25 °C) | |
| Solubility | Insoluble in ethanol and acetone |
| Hazards | |
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| Danger | |
| NFPA 704 (fire diamond) | |
| Safety data sheet (SDS) | Fisher SDS |
| Related compounds | |
Other cations
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Iron(II) perchlorate Nickel(II) perchlorate |
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) perchlorate is an inorganic chemical compound with the formula Co(ClO4)2·nH2O (n = 0,6). The pink anhydrous and red hexahydrate forms are both hygroscopic solids.
Preparation and reactions
Cobalt(II) perchlorate hexahydrate is produced by reacting cobalt metal or cobalt(II) carbonate with perchloric acid, followed by the evaporation of the solution:[1]
- CoCO3 + 2 HClO4 → Co(ClO4)2 + H2O + CO2
The anhydrous form cannot be produced from the hexahydrate by heating, as it instead decomposes to cobalt(II,III) oxide at 170 °C. Instead, anhydrous cobalt(II) perchlorate is produced from the reaction of dichlorine hexoxide and cobalt(II) chloride, followed by heating in a vacuum at 75 °C.[1][2]
Structure
The anhydrous form consists of octahedral Co(ClO4)6 centers, with tridentate perchlorate ligands.[3] On the other hand, the orthorhombic hexahydrate consists of isolated [Co(H2O)6]2+ octahedrons and perchlorate anions with lattice constants a = 7.76 Å, b = 13.44 Å and c = 5.20 Å. The hexahydrate undergoes phase transitions at low temperatures.[4][5]
References
- ^ a b c F. Solymosi; J. Raskó (1977). "Study of the thermal decompositions of some transition metal perchlorates". Journal of Thermal Analysis and Calorimetry. 11 (2): 289–304. doi:10.1007/bf01909967.
- ^ Jean-Louis Pascal; Jacqueline Potier; Cheng Shan Zhang (1985). "Chlorine trioxide, Cl2O6, a most efficient perchlorating reagent in new syntheses of anhydrous metal perchlorates, chloryl and nitryl perchloratometalates of cobalt(II), nickel(II), and copper(II). Reactivity of chlorine trioxide with anhydrous or hydrated chlorides and nitrates". Journal of the Chemical Society, Dalton Transactions. 2: 297–305. doi:10.1039/DT9850000297.
- ^ J. L. Pascal; J. Potier; D. J. Jones; J. Roziere; A. Michalowicz (1985). "Structural approach to the behavior of perchlorate as a ligand in transition-metal complexes using EXAFS, IR, and Raman spectroscopy. 2. Crystal structure of M(ClO4)2 (M = Co, Ni). A novel mode of perchlorate coordination". Inorganic Chemistry. 24 (2): 238–241. doi:10.1021/ic00196a026.
- ^ M.B. Patel; Sushama Patel; D.P. Khandelwal; H.D. Bist (1983). "Vibrational studies and phase transitions in Co(ClO4)2·6H2O and Mn(ClO4)2·6H2O". Chemical Physics Letters. 101 (1): 93–99. doi:10.1016/0009-2614(83)80311-X.
- ^ A. K. Jain; G. C. Upreti (1975). "On the anomalous paramagnetism of Co(II) perchlorate hexahydrate at low temperatures". Journal of Physics C: Solid State Physics. 8 (12): 1884–1888. doi:10.1088/0022-3719/8/12/013.