Comment pubs.acs.org/JPCC
Reply to “Comment on ‘Buffer Effects in the Kinetics of Concerted Proton-Coupled Electron Transfer: The Electrochemical Oxidation of Glutathione Mediated by [IrCl6]2− at Variable Buffer pKa and Concentration’” Jonnathan Medina-Ramos,§ Olufemi Oyesanya,‡ and Julio C. Alvarez*,† †
Department of Chemistry, Virginia Commonwealth University, P.O. Box 842006, Richmond, Virginia 23284, United States Department of Chemistry, Norfolk State University, 700 Park Avenue, Norfolk, Virginia 23504, United States § Department of Chemistry and Biochemistry, University of Delaware, 102 Brown Laboratory, Newark, Delaware 19716, United States ‡
J. Phys. Chem. C 2013, 117 (2), 902−912. DOI: 10.1021/jp3111265 J. Phys. Chem. C 2013, 117. DOI: 10.1021/jp4062253
T
his is a Reply to the Comment by Stanbury et al.1 on our report about the buffer catalysis observed when glutathione (GSH) is oxidized by electrogenerated IrCl62−.2 We considered the reaction to occur in two stages, the first one being the rate-determining step (RDS) that produces the thiyl radical GS•, the protonated base BH+, and the reduced mediator IrCl63−, followed by radical dimerization to yield the disulfide GSSG.2 The RDS involves the transfer of an e− and a H+ that could occur in stepwise (via intermediates, GS•+ or GS−) or concerted manner. The base B can be the solvent H2O or a component of the buffer system B/BH+. Cyclic voltammetry and fitted simulations were used to study this reaction in buffers of pKaBH+ < pKaGSH = 8.7 and in the pH range 4.0−9.0.2 The proposed rate law (eq 1) was written regarding the RDS as a concerted step in which B also controls the reaction rate and k is the termolecular rate constant of the reaction.2 Rate = k[B][GSH][IrCl 6 2 −]
Figure 12). Therefore, GSH gets oxidized near the electrode by the electrogenerated Mn+1 in a chemical step subsequent to the electrochemical oxidation of the bulk Mn. This type of reaction sequence is known as EC′ (E = electrochemical, C = chemical).6 The shape of the CV for an EC′ scheme is very sensitive to the kinetics of the chemical step, and reaction zone diagrams have been established accordingly.7 Two extreme kinetic zones (“total” catalysis and pure kinetic conditions) with strikingly different shape of CV have been identified depending on the reactant (GSH or Mn+1) that produces pseudo f irst-order conditions.7 The transition between these two regimes at 25 °C is given by the kinetic dimensionless parameter λ = 0.02569 (kCi°/ν), in which k is the rate constant of the chemical step (kobs in our case, see below), ν the scan rate of the CV, and Ci° the bulk concentration of Mn or GSH.6,7 In our last report,3 we show that kobs increases as EMn°′ is raised (Table 33 and Figure 43) and the CV for GSH oxidation undergoes the expected transformation reflecting a higher value of kobs (Figure 13). Stanbury et al.8 also observed higher k’s for GSH oxidation when increasing the E°′ of the oxidant, which agrees with the Marcus Theory prediction that the rate of electron transfer (ET) increases when the reaction becomes more exergonic.6 When raising the pKaBH+ or the buffer concentration we observed the same CV transformation (Figures 23 and Figures S1 and S23) and a corresponding increment in kobs (Table 33 and Figure 33).3 Therefore, the claims made in our first report are further confirmed indicating that the rate between GSH and Mn+1 can be accelerated by raising either EMn°′ or pKaBH+.2,3 Nevertheless, the catalytic effect by buffer components can be easily overlooked when [B] is in excess or [B] changes very little during the course of the reaction.2,3 In those conditions, [B] ≈ [B]o in eq 1 and can be incorporated into the rate constant to yield kobs = k[B] and a rate that seems to be independent of [B] (eq 2).2,3 As a consequence, plateau regions are observed in the plots of IGSH vs [B] where the GSH oxidation current no longer increases upon further increments in [B] (Figures 2B2 and 53). These results are obtained without
(1)
Three key observations from experimental cyclic voltammograms (CVs) and without invoking any kinetic model support this claim:2 first, the negligible GSH oxidation current (IGSH) observed in unbuffered solution when H2O is the only proton acceptor present (Figures 1B2 and 2A2); second, the variation of IGSH as a function of pKaBH+ and buffer concentration at constant pH (Figures 2B2 and 53); and third, the dependence of the CV shape on [B] at differential scan rates (Figure 52). The current IGSH is the charge transferred per unit time and therefore is related to the reaction rate. These results were corroborated in our most recent report in which the GSH oxidation was studied as function of the redox formal potential (E°′) of the mediator Mn.3 The concerted nature of this reaction was postulated because the conditions studied would not allow the stepwise pathways to prevail. For instance, the intermediate GSH•+ would be too acidic to exist and produce a RDS with the observed isotopic effects. Likewise, given that pKaBH+ < pKaGSH, GS− is totally disfavored at pH 7.0 or lower, and reprotonation by BH+ is faster than oxidation by IrCl62−,2,4 as has been reported for thioglycolic acid by IrCl62−.5 On a glassy carbon (GC) electrode, GSH is essentially nonelectroactive because no CV peaks are detected (inset © 2013 American Chemical Society
Received: July 28, 2013 Revised: October 26, 2013 Published: December 6, 2013 743
dx.doi.org/10.1021/jp4075042 | J. Phys. Chem. C 2014, 118, 743−745
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Comment
implementing any kinetic model.2,3 We believe Stanbury et al.1,8 did not observe a change in the rate upon increasing the concentration of acetate (1.0−10 mM) and succinic (5.0−35 mM) buffers because in their conditions [IrCl62−] is so small (0.1−0.08 mM) while [GSH] = 3.0 mM that the extent of the reaction essentially maintains [B] ≈ [B]o.1,8
unlikely that Cu2+ impurities are dominating the effects in our experiments. Because the oxidation of GSH occurs directly on Pt and Au electrodes,11 the EC′ scheme is not applicable in those cases. Early on in our studies we conducted experiments at 0.1 M of ionic strength (μ) instead of 1.0 and observed the same trends. Given the pH dependency of the GSH charge and the multiple charges of the oxidants, we decided to increase μ to mitigate even more the double-layer effects and the electrostatic interaction between ions.6 In summary, we do not think our data are in conflict with the reports by Stanbury et al.1,8 The similarity in the pH profile of kobs for both studies confirms that.1 Changing the pH is another way of changing the concentration of each buffer component. In our opinion, the absence of buffer catalysis in their experiments is because in their conditions the GSH oxidation is limited by [IrCl62−] (0.08−0.1 mM) instead of [B]. Hence [B] ≈ [B]o throughout the reaction such that the rate and rate law appear pseudo-zeroth order in [B] (eq 2).1,8 To detect buffer catalysis, experiments must be conducted at constant pH over a concentration range in which the other reactants do not limit the reaction rate.9,10 Our simplified model only considers the monoanionic state of GSH that is dominant between pH 5.0 and 7.0, and we do not investigate final reaction products (GSO3−, etc.).2,3 Nevertheless, the model still works because the fate of GS• does not matter if the reaction that generates GS• in the first step is the RDS.2,3 Stanbury et al. also concluded that Mn + GSH = Mn+1 + GS• + H+ is the RDS (eq 98). Therefore, given that there are multiple reports about thiols undergoing buffer catalysis in nonredox reactions due to the concomitant release of H+,12 why would it not be possible for thiol oxidation when the oxidation step also produces H+ and happens to be the RDS in a complex mechanism?
Rate[B] ≈ [B]o = k[B]o [GSH][IrCl 6 2 −] = kobs[GSH][IrCl 6 2 −]
(2)
kobs = kobsBH+ + kobsB = kBH+[H 2PO4 −] + kB[HPO4 2 −] (3)
Equation 3 is the expression of kobs for the two components of phosphate buffer (PB). The bases H2PO4− and HPO42− are the only major phosphate species present in PB in the pH range 4.0 to 9.0 (Figure 32). Water and OH− are neglected in eq 3 because kH2O is very small, and although kOH- must be higher than kB, [OH−] is too low to contribute.2 Figure 1A shows the
Figure 1. (A) Plots of calculated kobs vs [PB]pH=7.0. (B) CV response for 1.0 mM IrCl62−, 3.0 mM GSH, at different [Cu2+], 10 mM malic buffer pH 5.1, 1.0 M NaCl, 0.1 V s−1.
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plots of calculated kobs values, employing the actual pKa’s of H2PO4− (2.1) and HPO42 (7.2) and assuming the values for kBH+ and kB. The plateaus were generated by making [B] and [BH+] constant at the threshold concentration indicated by our experiments (Figure 2B2). As expected, the slope for each phosphate is different because of the pKa difference,9 but the addition of kobsBH+ and kobsB to produce kobs (eq 3) generates a plot with two linear sections and slopes that correspond to the original kobsBH+ and kobsB. The experimental plot of kobs vs [B] (Figure 3B2) shows exactly this same trend, and the slope change coincides with the threshold concentration, 0.050 M PB, at which the rate becomes pseudo zeroth-order in HPO4− (Figure 2B2). Therefore, the nonlinearity in Figure 3B2 is a result of the contribution of more than one base (H2PO4− and HPO42−) and the onset of their corresponding pseudo zerothorder behavior. When plotting log kobs vs pH using the calculated values from Figure 1A, a plot exactly resembling the one in Figure 3A2 is obtained. Our kobs graphs agree well with the expected kinetic plots described in the classic presentations of acid−base catalysis by Jencks9 and Anslyn.10 To address the Cu2+ effect, we performed CVs in the presence of 0.0, 1.0, and 5.0 μM of CuSO4 as suggested by one of the reviewers at the time. This produced two trends: for malic buffer 10 mM pH 5.1 and unbuffered neutral solution, a new prewave oxidation peak appeared at ∼0.60 V, the size of which seemed to be correlated to the amount of Cu2+ added, but no change on IGSH was detected (Figure 1B). For phosphate buffer 10 mM pH 7.0, there was no appearance of a prewave peak, but IGSH decreased by ∼10%. Since none of these trends were observed when manipulating the buffers, we were led to conclude that although Cu2+ traces do have an effect on the GSH voltammetry it is
AUTHOR INFORMATION
Corresponding Author
*E-mail:
[email protected]. Notes
The authors declare no competing financial interest.
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ACKNOWLEDGMENTS The support of this work by the National Science Foundation through Grant CHE-0645494 is greatly appreciated. REFERENCES
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(6) Bard, A. J.; Faulkner, L. R. Electrochemical Methods: Fundamentals and Applications, 2nd ed.; John Wiley & Sons Inc.: New York, 2001; p 833. (7) Saveant, J. M. Elements of Molecular and Biomolecular Electrochemistry: An Electrochemical Approach to Electron Transfer Chemistry; John Wiley & Sons, & Inc.: Hoboken, NJ, 2006; p 485. (8) Bhattarai, N.; Stanbury, D. M. Oxidation of Glutathione by Hexachloroiridate (IV), Dicyanobis(bipyridine)iron(III), and Tetracyano(bipyridine)iron(III). Inorg. Chem. 2012, 51, 13303− 13311. (9) Jencks, W. P. Catalysis in Chemistry and Enzymology; McGrawHill: New York, 1969; p 644. (10) Anslyn, E. V.; Dougherty, D. A. Modern Physical Organic Chemistry; University Science Books: Sausalito, CA, 2006; p 1079. (11) Ossendorfova, N.; Pradac, J.; Koryta, J. Electrode Processes of the Sulfhydryl-Disulfide System, IV. Glutathione at Platinum and Gold Electrodes. J. Electroanal. Chem. 1970, 28, 311−316. (12) Jencks, W. P.; Carriuolo, J. Imidazole Catalysis, III. General Base Catalysis and the Reactions of Acetyl Imidazole with Thiols and Amines. J. Biol. Chem. 1959, 234, 1280−1285.
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