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 d
10
Cu
I
ion, the common geometries are two‐coordinate linear, three‐coordinate trigonal planar, and four‐coordinate tetrahedral. Cu
I
compounds are diamagnetic and colorless, except where color results from charge‐transfer bands or a counterion; these complexes are often fairly readily oxidized to Cu
II
compounds. The d
9
Cu
II
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 Cu
II
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. Cu
III
complexes are typically square planar and diamagnetic.
