Copper: Inorganic & Coordination Chemistry
Published Online: 15 DEC 2011
Copyright © 2011 John Wiley & Sons, Ltd. All rights reserved.
Encyclopedia of Inorganic and Bioinorganic Chemistry
How to Cite
Conry, R. R. 2011. Copper: Inorganic & Coordination Chemistry . Encyclopedia of Inorganic and Bioinorganic Chemistry. .
- Published Online: 15 DEC 2011
Copper was one of the first metals used widely because the metal is fairly plentiful (among the 25 most abundant elements in the earth's crust) and can be found in its metallic state. In addition, the metal and its alloys have a number of beneficial qualities including ductility, malleability, strength, corrosion resistance, and high thermal and electrical conductivity, combined with an attractive appearance. Copper is also an essential trace nutrient for organisms ranging from bacteria to mammals.
Copper exhibits a rich coordination chemistry with complexes known in oxidation states ranging from 0 to +4, although the +2 (cupric) and the +1 (cuprous) oxidation states are by far the most common, with +2 predominating. Compounds of copper have found extensive practical usage, including as catalysts in both homogeneous and heterogeneous reactions, as fungicides, pesticides, and wood preservatives, as pigments for paints and glasses, and in the so-called high-temperature superconductors.
The coordination numbers and geometries of copper complexes vary with oxidation state. For the spherically symmetric d10 CuI ion, the common geometries are two-coordinate linear, three-coordinate trigonal planar, and four-coordinate tetrahedral. CuI compounds are diamagnetic and colorless, except where color results from charge-transfer bands or a counterion; these complexes are often fairly readily oxidized to CuII compounds. The d9 CuII ion is usually found in a tetragonal coordination environment, with four short equatorial bonds and another one or two longer axial bonds although complexes with other structures are known, including tetrahedral, square planar, and trigonal bipyramidal geometries. Most of the CuII compounds are blue or green because of d–d absorptions in the 600 to 900-nm region; exceptions generally also have charge-transfer bands tailing into the visible, causing a red or brown appearance. CuIII complexes are typically square planar and diamagnetic.