Copper Chemistry

Copper has an atomic number 29 and atomic mass of 63.55. It belongs to Group I-B transition metals. The melting point of copper is 1084.6�C. Copper occurs naturally in the cuprous (I, Cu⁺) and cupric (II, Cu²⁺) valence states. There is a single electron in the outer 4s orbital. The 3d10 orbital does not effectively shield this outer electron from the positive nuclear charge, and therefore the 4s1 electron is difficult to remove from the Cu atom (1). The first ionization potential is 7.72 eV and the second is 20.29 eV. Because the second ionization potential is much higher than the first, a variety of stable Cu⁺ species exist (2). The ionization state of copper depends on the physical environment, the solvent, and the concentration of ligands present. In solution, copper is present as Cu²⁺ or complexes of this ion. The cuprous ion Cu¹⁺ is unstable in aqueous solutions at concentrations greater than 10^-7 M (3). However, in wet soils, Cu¹⁺ is moderately stable at typically expected conditions (10^-6 to 10^-7 M). Under such conditions, hydrated Cu¹⁺ would be the dominant copper species (1). Copper can exist as two natural isotopes, ⁶³Cu and ⁶⁵Cu, with relative abundances of 69.09 and 30.91%, respectively (4). In the Earth's crust, copper is present as stable sulfides in minerals rather than silicates or oxides (3). The Cu¹⁺ ion is present more commonly in minerals formed at considerable depth, whereas Cu²⁺ is present close to the Earth's surface (3).

The transition metals are noted for the variety of complexes they form with bases. In these complexes, Cu¹⁺ and Cu²⁺ act as electron acceptors. Chelating bases are so named because they have two or more electron donor sites (often on O, S, or N atoms) that form a ‘claw’ around the copper ion (1). Such complexes are important in soil chemistry and in plant nutrition. The Cu¹⁺ ion forms strong complexes with bases containing S, but Cu²⁺ does not. In the presence of these bases, Cu²⁺ acts as a strong oxidant (2).