![]() Transactions of the American Electrochemical Society, volume 45, page 161. McDermott, and Adelle Bergman (1983): "Infrared spectra of the matrix-isolated chlorides of iron, cobalt, and nickel." Journal of Molecular Spectroscopy, volume 98, issue 1, pages 111-124. Corrosion Science, volume 15, issues 6–12, pages 603-625. Halstead (1975): "A review of saturated vapour pressures and allied data for the principal corrosion products of iron, chromium, nickel and cobalt in flue gases". Zeitschrift für anorganische und allgemeine Chemie, volume 268, issue 1‐2, pages 25-34. ^ a b Harald Schäfer and Kurt Krehl (1952): "Das gasförmige Kobalt(III)‐chlorid und seine thermochemischen Eigenschaften".Inorganic and Nuclear Chemistry Letters, volume 5, issue 4, pages 277-283. (1969): "The interaction of cobalt(III) with chloride ion in acetic acid". Another classical example is tris(ethylenediamine)cobalt(III) chloride Co(H In particular, hexamminecobalt(III) chloride Co(NHģ is the archetypal Werner complex and has uses in biological research. Trichlorides of cobalt(III) complexed with various ligands, such as organic amines, can be quite stable. In solutions of cobalt(III) salts with chloride ions, the anionic complexes (H The hexachlorocobaltate(III) anion CoCl 3−Ħ has been identified in preparations of cobalt(III) salts and hydrochloric acid HCl in glacial acetic acid. In a 1932 report, the compound was claimed to arise in the electrolysis of cobalt(II) chloride in anhydrous ethanol. Ī report from 1969 claims that treatment of solid cobalt(III) hydroxide CoOOHĢO with anhydrous ether saturated with HCl at −20 ☌ produces a green solution (stable at −78 ☌) with the characteristic spectrum of CoCl Ĭobalt trichloride, in amounts sufficient to study spectroscopically, was obtained by Green and others in 1983, by sputtering cobalt electrodes with chlorine atoms and trapping the resulting molecules in frozen argon at 14 K. ![]() At 1073 K, the partial pressures were 7.3 and 31.3 mm Hg, respectively. However, equilibrium shifts to the left at higher temperatures. The trichloride is formed through the equilibriumĪt 918 K (below the melting point of CoClĢ, 999 K), the trichloride was the predominant cobalt species in the vapor, with partial pressure of 0.72 mm Hg versus 0.62 for the dichloride. Preparation Ĭobalts trichlorides was detected in 1952 by Schäfer and Krehl in the gas phase when cobalt(II) chloride CoClĢ is heated in an atmosphere of chlorine ClĢ. Īerodynamic properties for the gas phase have been determined by the Glushko Thermocenter of the Russian Academy of Sciences. Ī Scientific study of the stability of this and other metal trihalides at 50 ☌ was published by Nelsoon and Sharpe in 1956. The infrared spectrum of the compound in frozen argon indicates that the isolated CoClģ molecule is planar with D 3h symmetry. Those earlier reports claim that it gives green solutions in anhydrous solvents such as ethanol and diethyl ether, and that it is stable only a very low temperatures (below −60 ☌). Some articles from the 1920s and 1930s claim the synthesis of bulk amounts of this compound in pure form however, those results do not seem to have been reproduced, or have been attributed to other substances like the hexachlorocobaltate(III) anion CoCl 3−Ħ. It has also been found to be stable at very low temperatures, dispersed in a frozen argon matrix. The compound has been reported to exist in the gas phase at high temperatures, in equilibrium with cobalt(II) chloride and chlorine gas. ![]() In this compound, the cobalt atoms have a formal charge of +3. Cobalt(III) chloride or cobaltic chloride is an unstable and elusive compound of cobalt and chlorine with formula CoClģ.
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