Chem. Res. Toxicol. 2001, 14, 657-660
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Gibbs Energy of Formation of PeroxynitritesOrder Restored Sara Goldstein,*,† Gidon Czapski,† Johan Lind,‡ and Gabor Mere´nyi‡ Department of Physical Chemistry, University of Jerusalem, Jerusalem 91904, Israel, and Department of Chemistry, Nuclear Chemistry, The Royal Institute of Technology, S-10044 Stockholm 70, Sweden Received March 26, 2001
In a recent publication [Nauser et al. (2001) Chem. Res. Toxicol. 14, 248-350], the authors estimated a value of 14 ( 3 kcal/mol for the standard Gibbs energy of formation of ONOOand argued that the experimental value of 16.6 kcal/mol [Mere´nyi, G., and Lind, J. (1998) Chem. Res. Toxicol. 11, 243-246] is in error. The lower value would suggest that the yield of free radicals during decomposition of ONOOH into nitrate is negligibly low, i.e., less than 0.5%, though within the large error limit given, the radical yield might vary between 0.003% and ca. 80%. The experimental value of 16.6 ( 0.4 kcal/mol was based on the determination of the rate constant of the forward reaction in the equilibrium ONOO- h •NO and O2•- by use of C(NO2)4, an efficient scavenger of O2•- which yields C(NO2)3-. Nauser et al. reported that addition of •NO has no significant effect on the rate of formation of C(NO2)3-, and therefore the formation of C(NO2)3- is due to a process other then reduction of C(NO2)4 by O2•-. In addition, they argued that Cu(II) nitrilotriacetate enhances the rate of peroxynitrite decomposition at pH 9.3 without reduction of Cu(II). In the present paper, we show that the formation of C(NO2)3- due to the presence peroxynitrite is completely blocked upon addition of •NO. Furthermore, the acceleration of the rate of peroxynitrite decomposition at pH 9 in the presence of catalytic concentrations of SOD ([ONOO-]/[SOD] > 30) results in the same rate constant as that obtained in the presence of C(NO2)4. These results can only be rationalized by assuming that ONOO- homolyses into •NO and O2•- with k ) 0.02 s-1 at 25 °C. Thus, the critical experiments suggested by Nauser et al. fully support the currently accepted thermodynamics as well as the mode of decomposition of the ONOOH/ONOO- system.
Introduction Peroxynitrous acid [pKa ) 6.6 (1, 2)] decays almost entirely into nitrate with kd ) 1.2-1.3 s-1 at 25 °C (24), whereas ONOO- is essentially stable (4). Starting with the seminal work of Mahoney (5), the homolysis of ONOOH to yield substantial amounts of free •NO2 and •OH was confirmed in several studies and consensus has recently settled around a free radical yield of 28 ( 4% (6, 7).
ONOOH h •NO2 + •OH
(1)
The above studies (5-7) had the common feature of probing the selectivity and yield of the radicals generated during the decomposition of ONOOH. In a completely different type of study, Mere´nyi and Lind (8) demonstrated the occurrence of equilibrium 2 at pH > pKa(ONOOH).
ONOO- h •NO + O2•-
(2)
Briefly, the characteristic absorption of C(NO2)3- was found when ONOO- was mixed with excess of C(NO2)4 in alkaline solutions and 20 °C. The rate constant of * To whom all correspondence should be addressed. Phone: 972-26586478. Fax: 972-2-6586925. E-mail:
[email protected]. † University of Jerusalem. ‡ The Royal Institute of Technology.
formation of C(NO2)3- was determined to be 0.017 s-1, independent of [C(NO2)4], and the yield of C(NO2)3- was within the experimental accuracy ca. 100%, i.e., [ONOO-]/ [C(NO2)3-] ≈ 1 (8). As C(NO2)4 is known to be an efficient scavenger of O2•- (reaction 3), it was concluded that k2 ) 0.017 s-1 at 20 °C (8).
C(NO2)4 + O2•- f C(NO2)3- + •NO2 + O2 k3 ) 2 × 109 M-1 s-1 (9) (3) The value of k2, in combination with other experimental data taken from the literature, i.e., k-2 ) (4.3-6.7) × 109 M-1 s-1 (10, 11), ∆fG° (•NO) ) 24.4 kcal/mol (12) and ∆fG° (O2•-) ) 7.6 kcal/mol (12), yields the standard Gibbs energy of formation of ONOO-, ∆fG° (ONOO-) ) 16.6 ( 0.4 kcal/mol (8, 13). According to this Gibbs energy of formation, coupled with the Gibbs energy of ionization of ONOOH, reaction 1 should occur at a rate that is in agreement with the yield of free radicals observed during decomposition of ONOOH into nitrate, i.e., between 10% and ca. 40% (4, 8, 13). Equilibrium 1 and 2 presuppose each other, i.e., thermodynamics demands that if one equilibrium occurs, so must the other, and therefore if one turns out to be untenable, so will the other. In a recent rebuttal paper, Nauser et al. (14) presented data that allegedly negate the occurrence of equilibrium 2. They argued that the rate of formation of C(NO2)3- at pH 12 is not affected by the presence of •NO and that
10.1021/tx010066n CCC: $20.00 © 2001 American Chemical Society Published on Web 05/15/2001
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Figure 1. Reaction of peroxynitrite with various concentrations of C(NO2)4 at pH 12. The formation of C(NO2)3- was followed at 350 nm when C(NO2)4 (35-108 µM) in 0.1 mM HClO4 was mixed with 44 µM peroxynitrite in 20 mM NaOH at a 1:1 ratio.
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Figure 3. Reaction of 28 µM C(NO2)4 with 22 µM peroxynitrite with and without 1.8 mM •NO. The formation of C(NO2)3- was followed at 350 nm when 56 µM C(NO2)4 in 0.1 mM HClO4 was mixed with 20 mM NaOH (trace a) or of 44 µM peroxynitrite in 20 mM NaOH (trace b) at a 1:1 ratio. Trace c is the same as trace b, but with both solutions saturated with •NO.
trite is completely blocked upon addition of •NO. Furthermore, the acceleration of the rate of peroxynitrite decomposition in the presence of catalytic concentrations of SOD results in the same rate constant as that obtained in the presence of C(NO2)4.
Materials and Methods
Figure 2. Reaction of ONOO- with C(NO2)4 at 350 nm and pH 9.9. C(NO2)4 (70 µM) in 100 mM ammonium buffer containing 0, 10, or 100 µM NaN3 was mixed with peroxynitrite (40 µM) in 10 mM NaOH. The experimental trace in the presence of 100 µM NaN3 is given in dot points, and the line is the perfect fit of the results to first-order kinetics resulting in k ) 0.022 s-1.
the acceleration of the decomposition of peroxynitrite in the presence of Cu(II) nitrilotriacetate does not involve the reduction of Cu(II). On the basis of the experimental value ∆fH° (ONOO-) ) -10 ( 2 kcal/mol (15, 16) and their estimated value S° (ONOO-) ) 31 eu (17), they obtained ∆fG° (ONOO-) ) 14 ( 3 kcal/mol, which would suggest that the yield of free radicals during decomposition of ONOOH into nitrate is negligibly low, i.e., less than 0.5%. However, within the large error limit given ((3 kcal/mol), the radical yield might vary between 0.003% and ca. 80%, which leaves the mechanism of ONOOH decomposition in limbo. The present paper will show that the experiments as described in the rebuttal paper are irreproducible, i.e., the formation of C(NO2)3- due to the presence peroxyni-
Materials. All chemicals were of analytical grade and were used as received. Solutions were prepared with deionized water that was distilled and purified using a Milli-Q water purification system. C(NO2)4 (Aldrich) was dissolved in ethanol and diluted in water. Bovine Cu,Zn SOD was purchased from Sigma, and its concentration was determined from its absorption at 260 nm using ) 1.13 × 104 M-1cm-1 (18). Peroxynitrite was synthesized through the reaction of nitrite with acidified H2O2 using a quenched-flow with a computerized syringe pump (“World Precision Instruments” model SP 230IW) as described elsewhere (19). Briefly, 0.63 M nitrite was mixed with 0.60 M H2O2 in 0.70 M HClO4, and the mixture was quenched with 3 M NaOH at room temperature. The stock solution contained 0.11 M peroxynitrite, 0.77 M NaOH, 10 mM nitrite, and practically no residual H2O2. The yield of peroxynitrite was determined from its absorption at 302 nm using ) 1670 M-1cm-1 (20). •NOsaturated solutions in gastight syringes (1.8 mM at 22 °C and 690 mm Hg) were prepared by passing it through a series of scrubbing bottles containing 50% NaOH and purified water after they were purged with Helium for an hour. C(NO2)4 solutions saturated with •NO were prepared by adding small volumes of concentrated and aerated solutions of C(NO2)4 (0.22-0.64 cm3) to •NO-saturated solutions (40 cm3) in order to avoid evaporation of C(NO2)4. The addition of 0.22-0.64 cm3 of aerated solutions to 40 cm3 •NO-saturated solutions has only a minor effect on [•NO], which can be ignored. However, these solutions contain some nitrite, which react slowly with C(NO2)4 to form C(NO2)3(14). Thus, the latter solutions contain some C(NO2)3-, i.e., trace c in Figure 3 does not go through the origin. In some cases solutions containing peroxynitrite in 20 mM NaOH were purged with Helium for 30 min and then purged with •NO for another 10 min. As expected, no loss of peroxynitrite due to the presence of •NO was recorded (21). Methods. Stopped-flow kinetic measurements were carried out using the Bio SX-17MV Sequential Stopped-Flow from Applied Photophysics with a 1-cm optical path. Solutions of
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peroxynitrite in 20 mM NaOH (either aerated, deaeraed or saturated with •NO) were mixed with C(NO2)4 in 0.1 mM HClO4 (either aerated, deaerated, or saturated with •NO) in a 1:1 ratio. The formation of C(NO2)3- was followed at 350 nm ( ) 14 400 M-1 s-1). Another series of experiments was carried out at lower pHs by mixing peroxynitrite in 10 mM NaOH with 100 mM ammonium buffer at a 1:1 ratio. The final pH in each experiment was measured at the outlet of the stopped-flow. All experiments were carried out at 25 °C.
Results The reaction of 22 µM peroxynitrite with 17.5-56 µM C(NO2)4 was studied at pH 12. The kinetic traces are given in Figure 1. The rate constant of the hydrolysis of C(NO2)4 under these conditions was determined to be (2.6 ( 0.1) × 10-3 s-1. Nauser et al. (14) reported a value of 1 × 10-3 s-1 at pH 12, whereas Mere´nyi and Lind (8) reported a value of 2 × 10-4 s-1 at pH 10.7. Thus, the contribution of the hydrolysis of C(NO2)4 at pH 12 cannot be ignored, and is the reason why the reaction could not be fitted to first-order kinetics. Nevertheless, at relatively low [C(NO2)4], i.e., the lowest trace in Figure 1, one can obtain the upper limit for this reaction, i.e., k ) 0.0270.028 s-1. The discontinuities in the two lowest kinetic traces in Figure 1 are due to peroxynitrite being in excess over C(NO2)4. Similar experiments were carried out at pH 9.9, where the contribution of the hydrolysis of C(NO2)4 can be ignored. As seen in Figure 2, the formation of C(NO2)3- obeys perfect first-order kinetics when 20 µM peroxynitrite reacts with 30 µM C(NO2)4 at pH 9.9 (50 mM ammonium buffer) in the absence and presence of up to 100 µM azide, resulting in k ) 0.023 ( 0.001 s-1. We note that in order to get a perfect firstorder kinetics and 100% yield of C(NO2)3-, it is essential to scavenge N2O3. The latter is formed via the reaction of •NO (formed in reaction 2) with •NO2 (formed in reaction 3), and reacts rapidly with ONOO- to form 2•NO2 and NO2- (21). Azide ion is known as an efficient scavenger of N2O3 [k ) 2.1 × 109 M-1 s-1 (22)]. Mere´nyi and Lind (8) prepared ONOO- by ozonolysis of azide, and therefore their solutions already contained azide ions. The effect of 1.8 mM •NO on the rate of formation of C(NO2)3- was studied when 22 µM peroxynitrite reacted with 28 µM C(NO2)4 at pH 12 (Figure 3). The effect of 0.9 mM •NO was also studied when 4.2 µM peroxynitrite reacted with 9.6 µM C(NO2)4 at pH 12 (Figure 4). The results shown in Figures 3 and 4 demonstrate that the formation of C(NO2)3- in the presence of •NO is solely due to the hydrolysis of C(NO2)4. The effect of catalytic concentrations of SOD (0.363.25 µM) on the decomposition rate of peroxynitrite (90110 µM) was studied at pH 9 in the presence of 50 mM ammonium buffer. The rate constant in the absence of SOD was determined to be 0.023 ( 0.001 s-1. The rate of peroxynitrite decomposition is accelerated in the presence of SOD (Figure 5) and is zero-order in SOD and first-order in peroxynitrite. The rate constant in the presence of 0.36, 1.1, and 3.25 µM SOD was determined to be 0.042 ( 0.002 s-1.
Discussion If ONOO- undergoes homolysis (reaction 2), then both NO and C(NO2)4 compete for superoxide. Since the reaction of superoxide with •NO is faster than with •
Figure 4. Reaction of 9.6 µM C(NO2)4 with 4.2 µM peroxynitrite with and without 0.9 mM •NO. The formation of C(NO2)3was followed at 350 nm when 19.2 µM C(NO2)4 in 0.1 mM HClO4 was mixed with 20 mM NaOH (trace a) or of 8.4 µM peroxynitrite in 20 mM NaOH (trace b) at a 1:1 ratio. Trace c is the same as trace b, but with deaerated solution of peroxynitrite and •NO-saturated solution of C(NO2)4.
Figure 5. Effect of SOD on the decomposition rate of peroxynitrite at 302 nm and pH 9 (50 mM ammonium nitrate buffer). Upper trace: 110 µM peroxynitrite; Lower trace: 92 µM peroxynitrite and 0.36 µM SOD.
C(NO2)4 (k-2 > k3), one expects that excess of •NO over C(NO2)4 will slow considerably the rate of formation of C(NO2)3-, i.e., the formation of C(NO2)3- should be solely due to the hydrolysis of C(NO2)4 as shown in Figures 3 and 4. Nauser et al. (14) showed that 1.9 mM •NO had no effect on the rate of formation of C(NO2)3- when 4 µM peroxynitrite reacted with 10 µM C(NO2)4 at pH 12, though they stated that in other cases they found with •NO “rates between 0.005 and 0.0185 s-1 depending on presence and absence of nitrite, the pH, and how the kinetics was analyzed”. They chose to present only the results with the higher rate constants as evidence against the occurrence of equilibrium 2. It is not clear to us, why they preferred to omit the data with the lower rate constants, given that the latter can be rationalized. They
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argue that the formation of C(NO2)3- is due to a process other then reduction of C(NO2)4 by O2•-, but refrain from providing any alternative explanation. They also reported that in all experiments they observed directly after mixing small but systematic deviations from first-order kinetics. Such deviations were not obtained in our system (Figures 1-4). Finally, in view of the fact that the experiments reported in the rebutted paper were performed at pH e10.7 (8), it is surprising that Nauser et al. (14) chose to present their results at pH 12, where the hydrolysis of C(NO2)4 clearly interferes. Another argument of Nauser et al. (14) against homolysis was based on their experiments with Cu(II) nitrilotriacetate at pH 9.3. They did observe the acceleration of peroxynitrite decomposition, which was close to zero-order in Cu(II), but did not detect Cu(I) using 2,2′biquinolyl, which forms thermodynamically stable complexes with Cu(I). However, 2,2′-biquinolyl is all but insoluble in water. Furthermore, catalysis of O2•- dismutation by Cu(II) complexes can take place via oscillation between Cu(II) and Cu(III) rather than Cu(II) and Cu(I) (23, 24). SOD is by far a more preferable choice than Cu(II) nitrilitriacetate, because it catalyzes O2•much faster [compare (2-3) × 109 M-1 s-1 (9) with ∼8 × 107 M-1 s-1 (14)], and its mechanism of catalysis is well established (18, 23, 24). We found that SOD at catalyic concentrations (0.36-3.25 µM) accelerates the decomposition rate of peroxynitrite (90-100 µM), and that the reaction is zero-order in SOD and first-order in peroxynitrite. The difference between the rate constants in the presence and absence of SOD yields k ) 0.019 ( 0.003 s-1, which is in excellent agreement with the rate constant of the formation of C(NO2)3- at pH 9.9, i.e., k ) 0.023 ( 0.001 s-1. Let us now summarize the thermochemistry of peroxynitrite. The overall first-order rate constant for the selfdecomposition of ONOOH has been determined to be kd ) 1.25 ( 0.05 s-1 at 25 °C (2-4), and the yield of the radicals 28 ( 4% (6, 7). Hence, k1 ) 1.25 × 0.28 ) 0.35 ( 0.03 s-1. The rate constant of the reverse reaction has been determined to be k-1 ) (4.5 ( 1.0) × 109 M-1 s-1 using the pulse radiolysis technique (4). Thus, K1 ) (7.8 ( 2.7) × 10-11 M and ∆G°1 ) 13.9 ( 0.2 kcal/mol at 25 °C. Combination of the latter value with ∆fG° (•OH) ) 6.2 ( 0.2 kcal/mol (12) and ∆fG° (‚NO2) ) 15.1 ( 0.1 kcal/ mol (12) yields ∆fG° (ONOOH) ) 7.4 ( 0.3 kcal/mol. Finally, taking pKa(ONOOH) ) 6.6 ( 0.1 (1, 2), we arrive at ∆fG° (ONOO-) ) 16.4 ( 0.3 kcal/mol. As can be seen, this value is in excellent agreement with our more directly determined value of 16.6 kcal/mol (8), and leaves little doubt about the thermodynamics of peroxynitrite chemistry. Taking the value of ∆fG° (ONOO-) ) 16.4 ( 0.3 kcal/mol and ∆fH° ) -10 ( 2 kcal/mol (15, 16), we obtain S° (ONOO-) ) 23 ( 7 eu. As expected and discussed elsewhere (25), this value is significantly lower than S° ) 35 eu for NO3-. In conclusion, the critical experiments suggested by Nauser et al. (14) fully support the currently accepted thermodynamics as well as the mode of decomposition of the ONOOH/ONOO- system.
Acknowledgment. S.G. and G.C. thank The Israel Science Foundation, and G.M. and J.L. thank The Swed-
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ish Natural Science Research Council and for their financial support.
References (1) Logager, T., and Sehested, K. (1993) Formation and decay of peroxynitrous acid: A pulse radiolysis study. J. Phys. Chem. 97, 6664-6669. (2) Kissner, R., Nauser, T., Bugnon, P., Lye, P. G., and Koppenol, W. H. (1997) Formation and properties of peroxynitrite studies by flash photolysis, high-pressure stopped flow and pulse radiolysis. Chem. Res. Toxicol. 10, 1285-1292. (3) Pryor, W. A., and Squadrito, G. L. (1995) The chemistry of peroxynitrite: a product from the reaction of nitric oxide with superoxide. Am. J. Physiol. 268, L699-L722. (4) Merenyi, G., Lind, J., Goldstein, S., and Czapski, G. (1999) Mechanism and thermochemistry of peroxynitrite decomposition in water. J. Phys. Chem. 103, 5685-5691. (5) Mahoney, L. R. (1970) Evidence for the formation of hydroxyl radicals in the isomerization of pernitrous acid to nitric acid in aqueous solution. J. Am. Chem. Soc. 92, 5262-5263. (6) Gerasimov, O. V., and Lymar, S. V. (1999) The yield of hydroxyl radical from the decomposition of peroxynitrous acid. Inorg. Chem. 38, 4317-4321. (7) Hodges, G. R., and Ingold, K. U. J. (1999) Cage-escape of geminate radical pairs can produce peroxynitrate from peroxynitrite under a wide variety of experimental conditions J. Am. Chem. Soc. 121, 10695-10701. (8) Mere´nyi, G., and Lind, J. (1998) Free radical formation in the peroxynitrous acid (ONOOH)/peroxynitrite (ONOO-) system. Chem. Res. Toxicol. 11, 243-246. (9) Mallard, W. G., Ross, A. B., and Helman, W. P. (1998) NIST Standard References Database 40, Version 3.0. (10) Huie, R. E., and Padmaja, S. (1993) The reaction of NO with superoxide. Free Radical Res. Commun. 18, 195-199. (11) 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. (12) Stanbury, D. M. (1989) Reaction potentials involving inorganic free radicals in aqueous solution. Adv. Inorg. Chem. 33, 69-138. (13) Merenyi, G., Lind, J., Goldstein, S., and Czapski, G. (1998) Peroxynitrous acid homolyses into ‚OH and ‚NO2 radicals. Chem. Res. Toxicol. 11, 712-713. (14) Nauser, T., Merkofer, M., Kissner, R., and Koppenol, W. H. (2001) Gibbs energey of formation of peroxynitrite. Chem. Res. Toxicol. 14, 348-350. (15) Ray, J. D. (1962) Heat of isomerization of peroxynitrite to nitrate and kinetics of isomerization of peroxynitrous acid to nitric acid. J. Inorg. Nucl. Chem. 24, 1159-1162. (16) Manuszak, M., and Koppenol, W. H, (1996) The enthalpy of isomerization of peroxynitrite to nitrate. Thermochim. Acta 273, 11-15. (17) Koppenol, W. H., and Kissner, R. (1998) Can OdNOOH undergo homolysis? Chem. Res. Toxicol. 11, 87-90. (18) Rabani, J., Klug, D., and Fridovich, I. (1972) Decay of the HO2 and O2•- radicals catalyzed by superoxide dismutase. A pulse radiolysis study. Isr. J. Chem. 10, 1095-1106. (19) Saha, A., Goldstein, S., Cabelli, D., and Czapski, G. (1998) Determination of the optimal conditions for synthesis of peroxynitrite by mixing acidified hydrogen peroxide with nitrite. Free Radical Biol. Med. 23, 653-659. (20) Hughes, M. N., and Nicklin, H. G. (1968) The chemistry of pernitrites. Part I. Kinetics and decomposition of pernitrous acid. J. Chem. Soc. A 450-452. (21) Goldstein, S., Czapski, G., Mere´nyi, G., and Lind, J. (1998) Effect of •NO on the decomposition of peroxynitrite: Reaction of N2O3 with ONOO-. Chem. Res. Toxicol. 12, 132-136. (22) Williams, D. L. H. (1988) Nitrosation, Cambridge University Press, and references therein. (23) Goldstein, S., and Czapski, G. (1987) The uniqueness of SOD Why most copper compounds cannot substitute SOD in vivo? Free Radical Res. Commun. 4, 225-229. (24) Goldstein, S., Czapski, G., and Meyerstein, D. (1990) A mechanistic study of the reaction of Cu(II)-peptides catalyzed superoxide dismutation. A pulse radiolysis study. J. Am. Chem. Soc. 112, 6489-6492. (25) Mere´nyi, G., and Lind, J. (1997) Thermodynamics of peroxynitrite and its CO2 adduct. Chem. Res. Toxicol. 10, 1216-1220.
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