Siderophore complex

Many microorganisms, when in iron-depleted conditions, are able to mobilize environmental iron by secreting low-molecular weight, high-affinity ironchelating compounds called siderophores, as well as cell-surface receptors for the ferric-siderophore complex and other proteins necessary for their uptake. Over the last three decades we have learned a lot about the chemical nature of these... 

There are a variety of factors concerning siderophore architecture that can contribute to the stability of iron-siderophore complexes. The overall architecture for natural multidentate siderophores can be any one of a number of different types linear,... 

As mentioned previously, siderophores must selectively bind iron tightly in order to solubilize the metal ion and prevent hydrolysis, as well as effectively compete with other chelators in the system. The following discussion will address in more detail the effect of siderophore structure on the thermodynamics of iron binding, as well as different methods for measuring and comparing iron-siderophore complex stability. The redox potentials of the ferri-siderophore complexes will also be addressed, as ferri-siderophore reduction may be important in the iron uptake process in biological systems. 

Siderophore binding sites for iron(III) are for the most part negatively charged and therefore, in aqueous solution there is a competition between H+ and Fe3+ binding. Consequently, the equilibrium expression for the formation of the iron-siderophore complex must take into account proton participation in the reaction. 

Stability comparisons between siderophore complexes with different binding stoichiometries are complicated by the fact that the units for the concentration equilibrium constants are different. Also, since the Fe3+ binding moieties have different pKa values competition for binding with H+ differs, which will not be reflected in the pH-independent / mlh values. Therefore, it is important to have a scale for iron-siderophore complex... 

This variation in complex stability with change in pH is particularly important in the context of biological systems, as it can potentially play a role in the iron uptake mechanism of some organisms. In some cases, the iron-siderophore complex is taken... 

Electrochemical experiments allow the determination of complex stability constants for Fe2+ by measuring complex redox potentials over a range of pH values. The Fe34YFe2+ redox potential of the siderophore complex, as with the spectral characteristics of the complex, is dependent on the inner coordination environment of the iron. These considerations will be addressed later (Section III.D). 

The effect of the amino acid spacer on iron(III) affinity was investigated using a series of enterobactin-mimic TRENCAM-based siderophores (82). While TRENCAM (17) has structural similarities to enterobactin, in that it is a tripodal tris-catechol iron-binding molecule, the addition of amino acid spacers to the TRENCAM frame (Fig. 10) increases the stability of the iron(III) complexes of the analogs in the order ofbAla (19)complex stability is attributed to the intramolecular interactions of the additional amino acid side chains that stabilize the iron-siderophore complex slightly. 

Thermodynamic Stability Constants for Iron(III)-Siderophore Complexes,... 

As mentioned previously, in many siderophore complexes, a decrease in pH will result in protonation and dissociation of a donor group. This decrease in effective denticity of the chelator will lead to a corresponding decrease in complex stability and the opening of available coordination sites for the formation of ternary complexes, and/or exchange with other chelators. However, in the case of catecholamide donor group siderophores, such... 

One aspect of microbial iron metabolism that remains unclear in many cases is the mechanism for iron release from tight sequestration once the siderophore complex arrives at its... 

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). 

The kinetics and mechanism of iron-siderophore complex formation are influenced by the oxidation state and composition of the first coordination shell of iron. The iron sequestration... 

The implications of these mechanistic studies for our understanding of environmental iron sequestration by siderophores is as follows. The hydroxyl containing aqua ferric ions will tend to form ferri-siderophore complexes more rapidly than the hexaaqua ion and ferrous ion will be sequestered more rapidly than the ferric ion. However, once in a siderophore binding site the ferrous ion will be air oxidized to the ferric ion, due to the negative redox potentials (see Section III.D). This also means that Fe dissolution from rocks will be influenced by mineral composition (other donors in the first coordination shell) as well as surface reductases in contact with the rock, and of course surface area (4,13). 

The dissociation of iron from a tetradentate siderophore complex is more rapid than from the analogous hexadentate system (3). This may be a reason for some organisms to produce tetradentate siderophores instead of hexadentate siderophores despite the concentration effect noted in Section III. A. As was illustrated in Fig. 19, it is also thermodynamically easier to reduce iron(III) in tetradentate complexes than hexadentate complexes, making it easier to induce release of iron from the complex by a redox mechanism. 

It has also been observed that the ionic strength of the solution or the presence of SDS micelles will affect the rate of dissociation of iron-siderophore complexes (22,181). 

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