Ligand siderophores

Figure 2 Structures of representative catecholate and mixed-ligand siderophores (a) enterobactin (b) parabactin (c) mycobacins (R = various alkyl chains, = R = R = MeorH.R = alkyl chains or H ) (d) fluorescent chromophore of pseudobactins and pyoverdins...

Fig. 2. Structures of catechol-containing and mixed ligand siderophores...

Z. Yehuda, M. Shenker, V. Romheld, H. Mar.schner, Y. Hadar, and Y. Chen. The role of ligand exchange in uptake of iron from microbial siderophores by graminaceous plants. Plant Physiol. 112 [213 (1996). 

Another factor that relates complex stability and siderophore architecture is the chelate effect. The chelate effect is represented by an increase in complex stability for a multidentate ligand when compared to complexes with homologous donor atoms of lower denticity. The effect can be observed when comparing the stability of complexes of mono-hydroxamate ligands to their tris-hydroxamate analogs, such as ferrichrome (6) or desferrioxamine B (4). However, the increase in stability alone is not sufficient to explain the preponderance of hexadentate siderophores over tetradentate or bidentate siderophores in nature, and the chelate effect is not observed to a great extent in some siderophore structures (10,22,50,51). 

Complexes formed by tetradentate siderophores involve stepwise complex formation and therefore, have somewhat different equilibria from their hexadentate analogs. Initial chelation will occur with a tetracoordinate FeL complex forming. A subsequent equilibrium then occurs, where the FeL complexes will react in a 2 1 stoichiometry with free ligands in solution to form a single Fe2L3 complex (coordinated water and charges not shown for clarity). 

Using linear regression, it is possible to estimate the protonation constants of the Fe(II) complexes of siderophore complexes where the redox potentials have been measured over a range of pH values (59). This also explains the variation in reversibility of reduction as the pH changes, as the stability of the ferro-siderophore complex is much lower than the ferric complex, and the increased lability of ligand exchange and increased binding site competition from H+ may result in dissociation of the complex before the iron center can be reoxidized. 

Another factor that can possibly affect the redox potential in biological systems is the presence of secondary chelating agents that can participate in coupled equilibria (3). When other chelators are present, coupled equilibria involving iron-siderophore redox occur and a secondary ligand will cause the siderophore complex effective redox potential to shift. The decrease in stability of the iron-siderophore complex upon reduction results in a more facile release of the iron. Upon release, the iron(II) is available for complexation by the secondary ligand, which results in a corresponding shift in the redox equilibrium toward production of iron(II). In cases where iron(II) is stabilized by the secondary chelators, there is a shift in the redox potential to more positive values, as shown in Eqs. (42)—(45). 

In Eq. (45), KFe(II)L is the stability constant for iron(II) complexation by the competing ligand, KFe(II)sid the stability constant for the complex formed between iron(II) and the siderophore, n the number of electrons transferred, Erxn the observed redox potential for the iron(III)-siderophore system coupled with iron(II) chelation, and EFJ m sld the redox potential of the iron(III)-siderophore complex. 

Another possible route for reduction of the iron center is photoreduction. This has been studied in a variety of marine siderophore systems, such as aquachelin, marinobactin, and aerobactin (2), where it was demonstrated that photolytic reduction was due to a ligand-to-metal charge transfer band of the Fe(III)-siderophore complex, eventually resulting in reduction ofiron(III) and cleavage of the siderophore (31,154,155). This suggests a possible role for iron reduction in iron release (71,155). 

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