CASE STUDIES IN CONTROLLED POTENTIAL
6.3. CASE II. EVALUATING CATALYSTS – THERMODYNAMIC CONSIDERATIONS
k k ARed
CRed
BRed
Net reaction:
AOx
COx
BOx
AOx + BRed ARed + BOx
FIGURE 6.3 The reduction of AOxby reagent, BRed, as mediated by a catalyst, C. The thermodynamics of this cycling scheme require that the formal potential for C must lie between the formal potentials of A and B. In other words,E∘′
A>E∘′ C >E∘′
B.
potential of the complex. Because of this effect, inorganic chemists have studied the use of different ligand structures as a means of manipulating the redox potential of transition metal complexes for the purpose of making more efficient redox catalysts. A redox cata- lyst acts as a mediator in transferring electrons between two reactant molecules when a direct exchange is too slow. The general scheme is shown in Figure 6.3. Investigating lig- and effects on the formal potential of complexes with structures similar to the active site of a redox enzyme is also an indirect method for studying how that enzyme functions [5].
By making subtle changes in the ligands, one can tune the formal potential of the complex to reach an optimum value.
In order for a redox catalyst to work, thermodynamics require that its formal potential lie between the formal potentials of the two redox couples in the net reaction that it is designed to catalyze. This situation is diagramed in Figure 6.3. In order for the catalyst, C, to cycle, COx must be able to oxidize the reduced form of the reactant, BRed and the reduced form of the catalyst, CRed, must be easily oxidized by AOx. The formal potentials of these three redox pairs must follow the rankingEo′A>Eo′C>Eo′B. That is, AOxmust be a stronger oxidizer than COxwhich, in turn, must be a stronger oxidizer than BOx. Evaluating the formal potentials of individual complexes by cyclic voltammetry is an important step in testing complexes to see if they meet the thermodynamic criterion for catalyzing the reaction of interest.
The same criteria apply to redox enzymes. The influence of the ligands on the redox potential of the complex in the active site is a key to understanding how these enzymes work. Because the mass transport of proteins is relatively slow, direct voltammetry on enzymes is difficult. Instead, researchers have used complexes to model the environment of the active site of important enzymes. An example of this is a study of manganese and iron complexes that mimic the behavior of the active site of a group of enzymes called superoxide dismutases (SODs). Normal cell metabolism generates a small amount of superoxide, O2−, a very reactive free radical that can cause all sorts of damage, if left unchecked within the cell. The SOD enzymes protect the cell by hastening the dispro- portionation of the superoxide. The average oxidation state of the oxygen atoms in O2− is−0.5. It is an intermediate on the path between O2(oxidation state of 0) and hydrogen peroxide, H2O2, (at −1). Consequently, O2− can be oxidized back to O2 or reduced to H2O2[6].
O2− ⇌O2+e− (6.13)
O2−+e−+2H+⇌H2O2 (6.14)
k k In the enzyme system, O2− and O2 play the role of the redox couple B in the diagram
in Figure 6.3. Eq. (6.14) represents the reduction reaction for the redox couple A. The metal-containing active site of the SOD enzyme is represented by C in the diagram. For clarity, MSODrepresents the metal at the center of the enzyme’s active site in the two dif- ferent redox reactions in Eqs. (6.15, 6.16).
O2−+MIIISOD⇌O2+MIISOD (6.15) O2−+MIISOD+2H+ ⇌H2O2+MIIISOD (6.16) It has been estimated that the formal potential (at pH 7) for the half reaction in Eq. (6.13) is+0.89 V versus SHE and the formal potential for Eq. (6.14) is−0.16 V [6]. Consequently, the complex in the active site of the enzyme must have a formal potential between+0.89 and−0.16 V versus SHE at pH 7. Manganese(II) and iron(II) are commonly found in vari- ous forms of SOD. Small changes in the structure of the ligands that hold the metal in the active site are thought to be responsible for the apparent formal potential of various forms of the SOD enzymes. In order to explore this hypothesis, Sjửdin et al. created a series of iron and manganese complexes with ligands that mimic the environment of some known SOD enzymes [6]. The structures of these model complexes are represented in Figure 6.4.
It was hypothesized that the tendency of the substituent to donate electron density to the aromatic ring and, indirectly, to the ligand site would influence the formal potential of the complex. The electron- donating/-withdrawing effect of substituents was quantified by the work of Hammet beginning in the late 1930s. Hammet measured the pKaof para- substituted benzoic acids. He reasoned that an electron-withdrawing substituent would take electron density from the carboxylic acid group as well as the aromatic ring [7]. The effect would cause the acid to dissociate more easily. An electron-donating group would have the opposite effect. He quantified his comparisons by assigning each substituent a constant,𝜎, based on the difference between the pKaof the substituted molecule and the
N X
N
N
N X
N
N Mn
O X
O O
X O
Mn OH2
N NH
N
N NH
Mn(4′–X–terpy)22+
complexes
Mn(biap)(p–X–bz)2
complexes
FIGURE 6.4 General structures for a series of Mn(II) complexes that were studied as models of the active site for superoxide dismutase enzymes. (Similar Fe(II) complexes were also used.) The formal potential of the complex was measured for different ligand substituents, X. Source: Reproduced with permission from Sjửdin et al. [6]. Copyright 2008, American Chemical Society.
k k pKaof benzoic acid itself:
𝜎x=pKa,x−pKa,benzoic (6.17)
Positive values of𝜎are more electron-donating than a hydrogen atom, and negative val- ues are more electron-withdrawing. Sjửdin found that the formal potential of Fe and Mn complexes of the general form shown in Figure 6.4 could be tuned by changing the sub- stituents, X, in two places [6]. Furthermore, the shift in the formal potential was a linear function of the Hammet constant. A summary is shown graphically in Figure 6.5.
Many of the complexes used in this work were poorly soluble in water. Consequently, the experiments were performed in acetonitrile (with an electrolyte of 0.1 M tetrabutyl- ammonium/hexafluorophosphate). The workers used an aqueous Ag/AgCl reference electrode. Because the difference in media can lead to an undefined and slowly drifting junction potential, the authors used a secondary reference technique. After recording voltammograms of each complex, they spiked the solutions with ferrocene in order to determine theEo′of the ferrocenium/ferrocene redox couple under the same conditions.
The formal potential for the ferrocene provided a secondary reference that eliminated the need to know the absolute junction potential between the reference electrode and the working solution. All potentials were referenced to the Eo′ for the ferrocene. The data plotted in Figure 6.5 (and that for several other substituents) demonstrated that relatively small changes to parts of ligands that are distant from the coordinating site can influence the electron donating properties of the ligand and, in turn, shift the formal potential of these model complexes in a predictable manner. These results support the hypothesis that the reduction potential of the active site of a redox enzyme is also subject to subtle structural changes in the ligands that bind the metal.
0.2–2.0 0.4 0.6 0.8 1.0
–1.5
DMA
OH
H CI
CH3Ph
Pyrr
–1.0 –0.5 nσPara
E0 (V versus Fc+/Fc)
0.0 0.5
FIGURE 6.5 A plot of the formal potential for complexes versus Hammet constant,n𝜎, wheren is the number of X-substituents (n=2, in these cases). The experiments were carried out in acetoni- trile, and the formal potential was measured with respect to the average of the ferrocenium/ferrocene peak potentials in order to account for any drift of the potential of the aqueous Ag/AgCl (saturated KCl) reference electrode or of the junction potential. Open circles represent the [Fe(4′-X-terpy)2]3+/2+
complexes, and the closed circles are data for the [Mn(4′-X-terpy)2]3+/2+ complexes. Abbreviations for substituents: Pyrr=pyrrolidine, DMA=dimethylamine, OH=hydroxyl, CH3Ph=methyl phenyl, H=hydrogen, Cl=chlorine. Source: Adapted with permission from Sjửdin et al. [6]. Copyright 2008, American Chemical Society.