## 2 Iron Chelators Design and Redox Activity of Iron  
### Redox Activity  
ROS is produced during aerobic respiration and involves free iron being utilised in the Fenton and Haber Weiss reactions, as shown in Figure 1 below. ROS levels are typically under tight regulation and are used beneficially in cell signalling using iron-dependent enzymes for cell survival and proliferation [15; 16]. However, if the reactions are uncontrolled, they can produce toxic levels of ROS, which can result in oxidative damage and may contribute to carcinogenesis [17; 18]. Thus, iron's pathological importance is due to this redox activity enabled by its ability to participate in Fenton chemistry, which results in ROS production. Redox cycling can occur if the oxidised complex is reduced, and the reduced complex can give an electron. The design of a chelator is important as chelators that can form redox-active metal complexes can produce cytotoxic ROS, killing cancerous cells. Another critical design consideration is thermodynamics, which predicts the direction of free radical reactions and the complexation of a metal centre, such as in the Fenton reaction and Haber Weiss reactions. The electromotive force is vital when calculating free radical reactions. In particular, the half-cell reduction potential shows if the reaction will occur in standard conditions, where the cell is referred to as a galvanic cell. The electromotive force for two half cells and Gibs free energy shows that the Fenton reaction is favourable in Figure 2. The term Gibbs free energy has been changed, favouring the term electrode potential defined by the IUPAC [19]. \[\Delta\mathrm{E}=\mathrm{E2}\ (\mathrm{acceptor\ of\ electrons})-\mathrm{E1}(\mathrm{donor\ of\ electrons})\] (A) \[\Delta\mathrm{G^{*}}=\Delta\mathrm{H^{*}}-\mathrm{T\Delta S^{*} }=\mathrm{RTInK}=-\mathrm{nF\Delta E}\] (B)  
Using the above equations, we see that the electrode potential of the hydrogen peroxide/hydroxyl radical (\(\mathrm{E}\)o = +0.32V, pH 7, 25 \({}^{\circ}\)C, NHE) is favourable with respect to Fe\({}^{3+}\)/Fe\({}^{2+}\) (\(\mathrm{E}\)o = +0.77 V, pH 7, 25 \({}^{\circ}\)C, NHE). Thus, it is likely to oxidise, producing ROS in standard conditions [20; 21]. However, the electrode potential of Fe\({}^{3+}\) (complex) + e can vary significantly depending on the type and strength of ligand from \(-1\) V to +1 V [22]. Consequently, the ligands are essential in the coordination sphere of Fe\({}^{2+}\), as the reactivity depends on the Fe complexes' speciation an important consideration in iron chelator design [23]. Iron chelators that can bind to both Fe (II) and Fe (III) can form redox-active complexes; this facilitates redox cycling resulting in the production of more toxic free radicals [24]. Iron chelators, such as traipine that form redox-active complexes, can produce toxic ROS levels, killing the cancer cells [25; 26]. Therefore, the design of a chelator is important in terms of its redox activity. The design determines if it can act in one of two ways: first, the increased production of ROS by using a redox-active chelator, or secondly, no production of ROS with the chelator's sole purpose being to remove the desired metal. Thus, the second design consideration is to act as a non-redox-active chelator. The second design consideration is of much more use in iron overload diseases where the production of ROS would be undesirable. Iron chelators may also affect physiochemical properties, not just directly interacting with redox reactions however resulting in a change to the electrode potential of the iron complex. The resulting iron complex may have free binding sites that  
Figure 2: Part (**A**) The electromotive force where E2 is the reduction potential for the half-cell reaction of the reduced species. E1 is the reduction potential for the half-cell reaction of the species oxidised if the cell’s potential is negative. The reaction can be reversed.