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Chem. Res. Toxicol. 1998, 11, 712-713
Peroxynitrous Acid Homolyzes into •OH and •NO2 Radicals Ga´bor Mere´nyi,*,† Johan Lind,† Sara Goldstein,‡ and Gidon Czapski‡ Department of Chemistry, Nuclear Chemistry, The Royal Institute of Technology, Stockholm 10044, Sweden, and Department of Physical Chemistry, The Hebrew University of Jerusalem, Jerusalem 91904, Israel Received March 5, 1998
Introduction. In a recent publication (1) we studied equilibrium 1 and determined k1 ) 0.017 s-1 at 20 °C. Using the literature value of k-1 ) (4.3-6.7) × 109 M-1 s-1 (2, 3), we calculate K1 ) (3.1 ( 0.9) × 10-12 M and ∆G°1 ) 15.4 ( 0.2 kcal/mol.
Then, using the literature values of ∆fG°(•NO) ) 24.4 kcal/mol and ∆fG°(O2•-) ) 7.6 kcal/mol, we obtain ∆fG°(ONOO-) ) 16.6 ( 0.4 kcal/mol. With a pKa ) 6.6 ( 0.1, we calculate ∆fG°(ONOOH) ) 7.7 ( 0.4 kcal/mol. The pKa employed here is that obtained by absorption measurements (3) and is somewhat lower than the currently accepted value of 6.8, but we have good evidence to show that 6.6 is the correct thermodynamic value.1 Utilizing the Gibbs free energies of formation of •NO (15.1 kcal/mol) and •OH (6.2 kcal/mol), we calculate 2 ∆G°2 ) 13.6 ( 0.4 kcal/mol and hence K2 ) (7.6 ( 3.8) × 10-11 M. As k-2 ) 5 × 109 M-1 s-1 (4), we calculate k2 ) 0.38 ( 0.19 s-1 at 20 °C. The overall rate constant of ONOOH decay was determined to be ko ) 0.66 s-1 at 20 °C1 and the yield of the intermediate ca. 40% (5). Therefore, the rate constant of intermediate formation is 0.4ko ) 0.26 s-1 at 20 °C. The excellent agreement of the experimental and calculated rate constants shows that reaction 2 can take place, and as a consequence the •OH and •NO2 free radicals are implicated as the reactive intermediates. The present work will summarize further evidence for the occurrence of the homolysis of ONOOH into •OH and • NO2 free radicals. (i) Decay of ONOO- at High pH. Using K2 ) (7.6 ( 3.8) × 10-11 M, pKa(ONOOH) ) 6.6 ( 0.1, and pKa(•OH) ) 11.9, we calculate K3 ) (1.5-7.2) × 10-16 M.
We determined k-3 ) 3 × 109 M-1 s-1, using the pulse radiolysis technique,1 and from this we calculate k3 ) (0.5-2.2) × 10-6 s-1. It has been known for a long time that the measured rate constants of the self-decomposition of peroxynitrite, * To whom all correspondence should be directed. † The Royal Institute of Technology. ‡ The Hebrew University of Jerusalem. 1Mere ´ nyi, G., Lind, J., Goldstein, S., and Czapski, G. Thermodynamics and mechanism of peroxynitrite (ONOOH/ONOO-) decomposition. Submitted.
kd, at pH > 10 are substantially higher than those predicted assuming that only ONOOH decomposes into nitrate [kd ) ko[H+]/(Ka + [H+])] (6). In the present work we predict the existence and the magnitude of a limiting kd value at high pH. We determined kd ) 7.5 × 10-6 and 5.9 × 10-6 s-1 at pH 13 and 14, respectively,1 almost the same in this pH range. These rate constants were found to be independent of [ONOO-]o and hardly affected by the presence of 0.1 mM DTPA.1 The latter observation rules out catalysis of ONOO- decomposition by traces of metal impurities as suggested earlier (6). We suggest that the mechanism of the decomposition of ONOO- at pH 14 takes place via reactions 1 and 3, followed by reactions 4-8,
and results in the consumption of two molecules of ONOO- for every homolysis via reaction 3. Furthermore, at pH 14 the decay of ONOO- produces between 80% and 100% NO2-, indicating that 0-20% of the decomposition of ONOO- yields NO3-.1 Thus, the experimental kd should be divided by a correction factor of 2-2.5, and k3 is therefore (2.4-3.0) × 10-6 s-1. This value is in good agreement with the experimental value, k3 ) (0.5-2.2) × 10-6 s-1, which was derived from reaction 2 above by use of several equilibrium constants. (ii) Decay of ONOOH at pH 4. The temperature dependence of the decay of ONOOH was carefully measured in the temperature interval of 3.2-48.1 °C at pH 4 in the presence of 0.1 M acetate buffer.1 Using the Arrhenius equation [ko ) A exp(-Ea/RT)] or Eyring equation [ko ) (kT/h) exp(∆Sq/R) exp(-∆Hq/RT)], we obtained1 Ea ) 21.6 kcal/mol and A ) 9.3 × 1015 s-1 or ∆Hq ) 21.2 kcal/mol and ∆Sq ) 12.7 cal/mol‚K. These activation parameters are in good agreement with those found by Benton and Moore (7) and Pfeiffer et al. (8) but differ significantly from those determined by Koppenol et al. (9). While we do not understand the reasons for the discrepancy between our values and those in ref 9, we note that the frequency factor (or ∆Sq) reported in the present work is of the same order of magnitude as
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Forum: Reactive Species of Peroxynitrite
those found for homolysis reactions of peroxides in the gas phase and in nonpolar organic solvents, and therefore it provides strong support for the homolysis of ONOOH via reaction 2. These values also suggest that the formation of •NO2 and •OH free radicals (40%) and nitrate (60%) occurs from a common intermediate, namely, a radical pair cage. This can be seen as follows: If ONOOH decomposes directly in a nonradical reaction to yield 60% NO3- with a rate constant kN and in parallel 40% undergoes homolysis with k2, then kN ) 0.6ko and k2 ) 0.4ko. However, as kN is the rate constant of a nonradical reaction, its A value would be expected to be around 1013 s-1. In such a case the experimentally determined A value should be somewhere between 1013 and 1016 s-1. The finding1 that the observed A value is as high as 1016 s-1 is a clear indication of a common homolysis being the rate-determining step. (iii) Activation Volume. The decomposition of ONOOH is characterized by a significantly positive activation volume of 9.6 ( 1.0 cm3/mol (10), which is in favor of a bond breakage process. Bond breakage could, in principle, account for the observed ∆Vq value, but not when it is accompanied by charge creation, since this is expected to lead to a drastic increase in electrostriction and an overall volume collapse. Many reactions that involve bond breakage or homolysis have volumes of activation between 5 and 10 cm3/mol. Furthermore, the direct isomerization of ONOOH to nitrate cannot take place in one step as a significantly negative ∆Vq is expected due to charge creation. Thus, a significantly positive ∆Vq of 9.6 cm3/mol suggests that the formation of nitrate from ONOOH takes place through the formation of an intermediate, which is accompanied by a volume increase. This suggests that the formation of both •NO2 and •OH free radicals (40%) as well as that of nitrate (60%) occurs from a common intermediate, namely, a radical pair cage. (iv) Reaction of Peroxynitrite with •NO. Detailed study of the reaction of •NO with peroxynitrite reveals that N2O3 is the reactive intermediate.2 In alkaline solutions •NO has no effect on the decay rate of ONOOin the absence of O2. Around neutral pH the decay of peroxynitrite in the presence of excess of •NO becomes very fast and is unaffected by [•NO] and by the presence of O2, indicating that there is no direct reaction between • NO and ONOO-.2 These results can be rationalized only if ONOOH homolyzes according to reaction 2. In such a case, •NO scavenges •NO2, and a chain reaction takes place due to the reaction of N2O3 with ONOO- which most probably produces 2•NO2 and NO2-. Final Remarks. The observations reported in the present work support the homolysis of ONOOH into •OH and •NO2 free radicals. However, there exist two types of experiments in the literature, which appear to contradict free radical formation. These are yield measurements of end products in the presence of certain radical scavengers and a study of the decay rate of ONOOH as a function of viscosity (5). Competitive inhibition studies with various hydroxyl radical scavengers showed that for some cases these scavengers have no effect on the oxidation yield. In others, only partial inhibition was observed, far less than that predicted according to the 2Goldstein, S., Czapski, G., Mere ´ nyi, G., and Lind, J. The reaction of •NO with peroxynitrite: Evidence for the homolysis of ONOOH into •NO and •OH. Submitted. 2
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known rate constants of these scavengers with hydroxyl radical. However, there are cases where the extent of inhibition correlates well with the known rate constants of these scavengers with hydroxyl radical, suggesting that hydroxyl radical is the active intermediate (5). The results in the presence of •OH scavengers can only be interpreted unambiguously, if all possible reactions in the system with the assorted rate constants are accounted for. Furthermore, the very reactive ONOOH molecule is likely to participate in reactions with secondary species produced after the •OH radical reacting with the added scavenger. This may easily upset the expected product yield. As for the viscosity experiments, similar arguments can be brought forward. Specifically, the polyethers added in order to change the viscosity react with •OH radicals to produce reductive carbon-centered radicals. The latter are expected to reduce ONOOH, thus increasing the decay rate of ONOOH. Therefore, a possible decrease of the decay rate because of increased viscosity would be counteracted. If one wants to redo properly viscosity experiments, one should add substances that do not react fast with •OH and determine the effect of viscosity on the self-decay rate as well as on the oxidation yield (ca. 40%). The latter experiment is more conclusive as the amount of radicals escaping the cage decreases with increasing the viscosity. In conclusion, we feel that, unless direct experimental evidence, e.g., by means of spectroscopic characterization, can be obtained for an alternative intermediate, the hydroxyl and nitrogen dioxide free radicals produced by homolysis of ONOOH remain the simplest and most plausible candidates for the reactive intermediate.
Acknowledgment. G.M. and J.L. thank the Swedish Natural Science Research Council and S.G. and G.C. the Israel Academy of Sciences for their financial support.
References (1) Mere´nyi, G., and Lind, J. (1998) Free radical formation in the peroxynitrous acid (ONOOH)/peroxynitrite (ONOO-) system. Chem. Res. Toxicol. 11, 243-246. (2) Huie, R. E., and Padmaja, S. (1993) The reaction of NO with superoxide. Free Radical Res. Commun. 18, 195-199. (3) Goldstein, S., and Czapski, G. (1995) The reaction of NO with O2•- and HO2•. A pulse radiolysis study. Free Radical Biol. Med. 19, 505-510. (4) Mere´nyi, G., and Lind, J. (1997) Thermodynamics of peroxynitrite and its CO2 adduct. Chem. Res. Toxicol. 10, 1216-1220 and references therein. (5) Goldstein, S., Squadrito, G. L., Pryor, W. A., and Czapski, G. (1996) Direct and indirect oxidations of peroxynitrite, neither involving the hydroxyl radical. Free Radical Biol. Med. 21, 965974 and references therein. (6) Edwards, J. O., and Plumb, R. C. (1993) The chemistry of peroxynitrites. Prog. Inorg. Chem. 41, 599-635 and references therein. (7) Benton, D. J., and Moore, P. (1970) Kinetics and mechanism of the formation and decay of peroxynitrous acid in perchloric acid solutions. J. Chem. Soc. A 3179-3182. (8) Pfeiffer, S., Gorren, A. C. F., Schmidt, K., Werner, E. R., Hansert, B., Bohle, D. S., and Mayer, B. (1977) Metabolic fate of peroxynitrite in aqueous solution. J. Biol. Chem. 272, 3465-3470. (9) Koppenol, W. H., Moreno, J. J., Pryor, W. A., Ischiropoulos, H., and Beckman, J. S. (1992) Peroxynitrite, a cloaked oxidant formed by nitric oxide and superoxide. Chem. Res. Toxicol. 5, 834-842. (10) Goldstein, S., van Eldik, R., Meyerstein, D., and Czapski, G. (1997) Spontaneous reactions and reduction by iodide of peroxynitrite and peroxynitrate: Mechanistic insight from activation parameters. J. Phys. Chem. A 101, 7114-7118.
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