Transition metal perchlorate complexes are coordination complexes with one or more perchlorate ligands. Perchlorate can bind to metals through one, two, three, or all four oxygen atoms. Usually however, perchlorate is a counterion, not a ligand.
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Homoleptic complexes
editHomoleptic complexes, i.e. complexes where all the ligands are the same (in this case perchlorate), are of fundamental interest because of their simple stoichiometries.
Several anhydrous metal diperchlorate complexes are known but most are not molecular (and hence, not complexes). For example, many compounds with the formula M(ClO4)2 are coordination polymers (M = Mn, Fe, Co, Ni, Cu). An exception to this pattern is palladium(II) perchlorate Pd(ClO4)2, which is a square planar complex consisting of a pair of bidentate perchlorate ligands. Furthermore, anhydrous Cu(ClO4)2 is sublimable, which implies the existence of molecular Cu(ClO4)2.[1]
Titanium(IV) perchlorate and zirconium(IV) perchlorate are molecular, featuring four bidentate perchlorate ligands. They are volatile.
Mixed ligand complexes
editMore common than homoleptic complexes are those with two or more types of ligands. A classic case is the dicationic complex pentamminecobalt(III) perchlorate, which had resisted formation by conventional substitution reactions.[2] It was prepared by oxidation of the azide complex:[3]
- [Co(NH3)5N3]2+ + HClO4 + NO+ → [Co(NH3)5OClO3]2+ + N2 + N2O
Another mixed ligand complex is the perchlorate complex of the ferric derivative of octaethylporphyrin.[4]
Perchlorate as a counterion
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Being the conjugate base of the strongly acidic perchloric acid, perchlorate is very weakly basic. It is more commonly encountered as a counterion in coordination chemistry. Illustrative of its low basicity is the ability of water to outcompete perchlorate as a ligand for metal ions is indicated by the multitude of aquo complexes with noncoordinated perchlorate. Ferrous perchlorate, cobalt(II) perchlorate, chromium(III) perchlorate, manganese(II) perchlorate, nickel(II) perchlorate, and copper(II) perchlorate are commonly encountered as their hexaaquo complexes.[5]
Synthesis
editThe preparation of perchlorate complexes can be challenging because perchlorate is a weakly ligand.
Chlorine trioxide is an important precursor to anhydrous perchlorate complexes. It serves as a source of ClO+2 and ClO−4. It reacts with vanadium pentoxide (V2O5) to give VO2(ClO4) and VO(ClO4)3. Hydrated mercury and cadmium perchlorates can be dehydrated with Cl2O6, affording anhydrous compounds.[6]
- MCl2 + 2Cl2O6 → ClO2M(ClO4)3 + 2 ClO2 + Cl2
- ClO2M(ClO4)3 → M(ClO4)2 + ClO2
In some cases, chlorine trioxide serves both as an oxidant and a dehydrating agent:
- M(H2O)6Cl2 + 2Cl2O6 → [M(H2O)6](ClO4)2 + 2 ClO2
- [M(H2O)6](ClO4)2 + 6 Cl2O6 → M(ClO4)2 + 6 HClO4 + 6 HClO3
Silver perchlorate, which has some solubility in noncoordinating solvents, reacts with some metal chlorides to give the corresponding perchlorate complex.[4]
Reactions
editAnhydrous perchlorate complexes are susceptible to hydrolysis:
- Cu(ClO4)2 + 6 H2O → [Cu(H2O)6](ClO4)2
Upon heating, perchlorate complexes yield oxides, evolving chlorine oxides in the process. For example, thermolysis of titanium perchlorate gives TiO2, ClO2, and O2 The titanyl species TiO(ClO4)2 is an intermediate in this decomposition.[7]
- Ti(ClO4)4 → TiO2 + 4ClO2 + 3O2 ΔH = +6 kcal/mol (25 kJ/mol)
Safety
editPerchlorate complexes and the reagents used to prepare them are often dangerously explosive intrinsically and especially in contact with organic compounds.[6]
References
edit- ^ Favier, Frederic; Barguès, Stephane; Pascal, Jean Louis; Belin, Claude; Tillard-Charbonnel, Monique (1994). "Crystal and molecular structure of anhydrous copper(II) perchlorate". J. Chem. Soc., Dalton Trans. (21): 3119–3121. doi:10.1039/DT9940003119.
- ^ Jones, W. E.; Swaddle, T. W. (1967). "Concerning the existence of perchloratopentamminecobalt(III) perchlorate". Canadian Journal of Chemistry. 45 (22): 2647–2650. doi:10.1139/v67-433.
- ^ Harrowfield, J. Macb.; Sargeson, A. M.; Singh, B.; Sullivan, J. C. (1975). "Trapping of Labile Cobalt(III) Complexes. Characterization of the Perchloratopentaamminecobalt(III) Ion". Inorganic Chemistry. 14 (11): 2864–2865. doi:10.1021/ic50153a059.
- ^ a b Masuda, Hideki; Taga, Tooru; Osaki, Kenji; Sugimoto, Hiroshi; Yoshida, Zenichi; Ogoshi, Hisanobu (1980). "Crystal and molecular structure of (Octaethylporphinato)iron(III) perchlorate. Anomalous magnetic properties and structural aspects". Inorganic Chemistry. 19 (4): 950–955. doi:10.1021/ic50206a031.
- ^ Gallucci, J. C.; Gerkin, R. E. (1989). "Structure of copper(II) perchlorate hexahydrate". Acta Crystallographica. 45 (9): 1279–1284. Bibcode:1989AcCrC..45.1279G. doi:10.1107/S0108270189000818. PMID 2557867.
- ^ a b Pascal, Jean-Louis; Favier, Frédéric (1998). "Inorganic Perchlorato Complexes". Coordination Chemistry Reviews. 178–180: 865–902. doi:10.1016/S0010-8545(98)00102-7.
- ^ Babaeva, V. P.; Rosolovskii, V. (1974). "Volatile titanium perchlorate". Bulletin of the Academy of Sciences of the USSR Division of Chemical Science. 23 (11): 2330–2334. doi:10.1007/BF00922105. ISSN 0568-5230.