1.1 Iron Chemistry
In the Earthâs crust, iron is the fourth most abundant element and the second most abundant metal (the most abundant is aluminium). Situated in the Periodic Table in the middle of the first transition series (characterised by having incompletely filled d orbitals), iron has access to a number of oxidation states (from âII to +VI), the principal being II (d6) and III (d5). A number of iron-dependent monooxygenases are able to generate high-valent Fe(IV) or Fe(V) reactive intermediates during their catalytic cycle. Whereas, Fe2+ is extremely water-soluble, Fe3+ is quite insoluble in water (Ksp = 10â39 M and at pH 7.0, [Fe3+] = 10â18 M) and significant concentrations of water-soluble Fe3+ species can be attained only by strong complex formation with appropriate ligands.
The interaction between Fe2+ and Fe3+ and ligand donor atoms will depend on the strength of the chemical bond formed between them. An idea of the strength of such bonds can be found in the concept of âhardâ and âsoftâ acids and bases (HSAB) (Pearson, 1963). âSoftâ bases have donor atoms of high polarisability with empty, low-energy orbitals; they usually have low electronegativity and are easily oxidised. In contrast, âhardâ bases have donor atoms of low polarisability, and only have vacant orbitals of high energy; they have high electronegativity and are difficult to oxidise. Metal ions are âsoftâ acids if they are of low charge density, have a large ionic radius. and have easily excited outer electrons. âHardâ acid metal ions have high charge density, a small ionic radius, and no easily excited outer electrons. In general, âhardâ acids prefer âhardâ bases and âsoftâ acids form more stable complexes with âsoftâ bases (Pearson, 1963). Fe(III) with an ionic radius of 0.067 nm and a charge of 3+ is a âhardâ acid and will prefer âhardâ oxygen ligands such as phenolate and carboxylate, compared to imidazole or thiolate. In contrast, Fe(II) with an ionic radius of 0.083 nm and a charge of only 2+ is on the borderline between âhardâ and âsoft,â favouring nitrogen (imidazole and pyrrole) and sulphur ligands (thiolate and methionine) over oxygen ligands.
The coordination number of 6 is the most frequently found for both Fe(II) and Fe(III) giving octahedral stereochemistry, although four-coordinate (tetrahedral) and particularly five-coordinate complexes (trigonal bipyramidal or square pyrimidal) are also found. For octahedral complexes, two different spin states1 can be observed. Strong-field ligands (e.g. Fe3+ OHâ), where the crystal field splitting is high and hence electrons are paired, give low-spin complexes, while weak-field ligands (e.g. CO, CNâ), where crystal field splitting is low, favour a maximum number of unpaired electrons and give high-spin complexes Changes of spin state affect the ion size of both Fe(II) and Fe(III), the high-spin ion being significantly larger than the low-spin ion. As we will see in Chapter 2, this is put to good use as a trigger for the cooperative binding of dioxygen to haemoglobin. High-spin complexes are kinetically labile, while low-spin complexes are exchange-inert. For both oxidation states only high-spin tetrahedral complexes are formed, and both oxidation states are Lewis acids, particularly the ferric state.
The unique biological role of iron comes from the extreme variability of the Fe2+/Fe3+ redox potential, which can be fine-tuned by well-chosen ligands, so that iron sites can encompass almost the entire biologically significant range of redox potentials, from about â0.5 V to about +0.6 V. However, as we will see in Chapter 13, copper allows access to an even higher range of redox potentials (0 V to +0.8 V), which turned out to be of crucial importance in the Earthâs rapidly evolving aerobic environment, following the arrival of water-splitting, oxygen-generating photosynthetic organisms.