Iron: Inorganic & Coordination Chemistry
Published Online: 15 MAR 2006
Copyright © 2006 John Wiley & Sons, Ltd
Encyclopedia of Inorganic Chemistry
How to Cite
Burgess, J. and Twigg, M. V. 2006. Iron: Inorganic & Coordination Chemistry . Encyclopedia of Inorganic Chemistry. .
- Published Online: 15 MAR 2006
Iron, the most abundant transition element in the biosphere, is essential in higher forms of life and its compounds have numerous industrial applications. The element has catalytic properties. Several simple iron compounds are industrially important, for example, large amounts of iron oxides are produced for pigment and magnetic recording media applications. Oxidation states ranging from −2 to +6 are known, the most familiar are +2 (ferrous iron) and +3 (ferric iron). Most common complexes are octahedral, although examples of higher (e.g. seven coordinate) and lower (e.g. tetrahedral and square planar) geometries are well known, as are a variety of mixed oxidation state compounds. Depending on ligands involved high- and low-spin complexes can be formed, and temperature-dependent spin-state crossover complexes are well documented. The thermodynamics and kinetics of reactions of low-spin iron(II) complexes with ligands such as 1,4-diimines and CN− have been extensively studied. While low-spin iron(II) is present in biological systems such as oxygenated heme centered oxygen carriers and thiolate clusters, iron(III) is more common. Thiolate-containing iron clusters form readily, and here mixed oxidation states are a prominent feature that is exploited in some important electron transfer proteins. Microorganism iron uptake is facilitated by tris-catecholate and tris-hydroxamate ligands like enterobactin that have an extraordinarily high affinity for iron(III) that results in a dramatic enhancement of the 3+ relative to the 2+ state in terms of redox potential.
- Mössbauer spectroscopy;
- Prussian blue;
- spin crossover;