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Communications Gibbs Energy of Formation of Peroxynitrite T. Nauser, M. Merkofer, R. Kissner, and W. H. Koppenol* Laboratorium fu¨ r Anorganische Chemie, Eidgeno¨ ssische Technische Hochschule Zu¨ rich, Universita¨ tsstrasse 6, CH-8092 Zu¨ rich, Switzerland Received November 14, 2000
A standard Gibbs energy of formation of 16.6 kcal mol-1 has been reported for peroxynitrite [Mere´nyi, G., and Lind, J. (1998) Chem. Res. Toxicol. 11, 243-246]. This value is based on the rate constants for the forward and backward rate constants of the equilibrium O2•- + NO• S ONOO-. A rate constant of 0.017 s-1 for the backward rate constant was determined by observing the formation of C(NO2)3- when peroxynitrite was mixed with C(NO2)4. However, a similar rate constant is also observed in the presence of NO•, which indicates that formation of C(NO2)3- is due to a process other than the reduction of C(NO2)4 by O2•-. Additionally, copper(II) nitrilotriacetate enhances the decay of ONOO- at pH 9.3, without reduction of copper(II). The preferred thermodynamic values are therefore as follows: ∆fH°(ONOO-) ) -10 ( 2 kcal mol-1, ∆fG°(ONOO-) ) 14 ( 3 kcal mol-1, S°(ONOO-) ) 31 eu, and E°′(ONOOH/NO2•, H2O) ) 1.6 V at pH 7 [Koppenol, W. H., and Kissner, R. (1998) Chem. Res. Toxicol. 11, 87-90].
Introduction The standard Gibbs energy of formation of peroxynitrite is derived from the rate at which the nitroform anion is formed when peroxynitrite is mixed with tetranitromethane (1). The interpretation of the reduction of tetranitromethane is that the peroxynitrite undergoes homolysis to form superoxide and nitrogen monoxide (reaction 1)
ONOO- f NO• + O2•-
(1)
and that superoxide reduces tetranitromethane. The rate constant for this process, 0.017 s-1, when combined with the rate constant for the backward reaction, results in a Gibbs energy of formation of the peroxynitrite anion of 16.6 kcal mol-1.1 In a later publication (2), an uncertainty of 0.4 kcal mol-1 was added, although its origin is not clear and was not discussed. According to this Gibbs energy of formation, and the Gibbs energy of ionization, homolysis of peroxynitrous acid (reaction 2)
ONOOH f HO• + NO2•
(2)
is possible at the rate observed for isomerization to nitrate, 1.2 s-1 at 25 °C. We show here that formation of the nitroform anion is also observed in the presence of nitrogen monoxide, and that in the presence of copper(II) nitrilotriacetate per* Corresponding author. Telephone: 41-1-6342-2875. Fax: 41-1-6321090. E-mail:
[email protected]. 1The discussion about the thermodynamics of oxoperoxonitrate(1-) has been in terms of calories. To facilitate comparison, we use this unit, too.
oxynitrite decays at a rate faster than 0.017 s-1 without formation of copper(I).
Materials and Methods Chemicals. Peroxynitrite [oxoperoxonitrate(1-)] was prepared from solid potassium superoxide and nitrogen monoxide as previously described (3) and synthesized according to the method of Bohle et al. (4). The concentration of oxoperoxonitrate(1-) was calculated from the absorbance at 302 nm [ ) 1700 M-1 cm-1 (4)]. All other reagents were purchased and of the highest grade available. Water was purified with a Millipore Milli-Q unit fed with deionized water. The slightly acidic solution of tetranitromethane was stable for hours in the presence and absence of nitrogen monoxide. Stopped-Flow Spectrophotometry. Kinetic experiments were carried out at ambient pressure and 25 °C with an OLIS RSM 1000 and an Applied Photophysics SX 17MV stopped-flow spectrophotometer. Formation of the nitroform (trinitromethanide) anion was followed between 300 and 450 nm. Kinetic traces were analyzed at 360 nm. The pH was measured in separate experiments. In all experiments (with and without nitrogen monoxide) small, but systematic, deviations from firstorder kinetics were found directly after mixing. Nevertheless, traces were analyzed under the assumption that a first-order decay was valid to obtain results that could be compared to those of Mere´nyi and Lind (1). Six kinetic traces were averaged to extract a pseudo-first-order rate constant. The decay of 0.5 mM oxoperoxonitrate(1-) by 50 µM copper nitrilotriacetate was studied at 302 nm and between 570 and 720 nm in ammonia buffer [0.05 M (NH4)2SO4 and 0.1 M NH3] at pH 9.3. To detect copper(I), some experiments were carried out with 100 µM copper(II) nitrilotriacetate in ammonia buffer saturated with 2,2′-biquinolyl (biq) before mixing with 5 mM oxoperoxonitrate(1-). The visual detection limit for the copper(I)-bis(biq) complex color is ∼2 µM (545 ) 5500 M-1 cm-1). Ion Chromatograpy. The levels of nitrite and nitrate were determined by anion chromatography and conductivity detection (ICSeparation Center 733, Hamilton PRP X-100 column, IC
10.1021/tx000236g CCC: $20.00 © 2001 American Chemical Society Published on Web 03/16/2001
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Figure 2. Decomposition of 0.5 mM peroxynitrite at 302 nm. The buffer was 0.1 M NH4+, 0.1 M NH3, and 0.05 M SO42- (pH 9.3) with 0.05 mM nitrilotriacetate (- - -) and with 0.05 mM nitrilotriacetate and 0.05 mM Cu2+ (s).
Figure 1. Production of trinitromethanide (nitroform) at 360 nm and pH 12 (A) without NO• and with 40 µM ONOO- and 80 µM C(NO2)4 and (B) with 1.9 mM NO•, 4 µM ONOO-, and 10 µM C(NO2)4. If superoxide were responsible for trinitromethanide formation, then the rate of its formation in panel B should be at least 20 times smaller than that shown in panel A: (s) experimental trace, (b) simulation for a rate constant of 0.018 s-1, and (0) residuum × 10. The residuum shows a small but systematic deviation between the fit and experiment. detector 732, and IC pump 709; Metrohm AG). A phthalic acid solution [2.5 mM phthalic acid, 5% MeCN (pH 4.5), and TRIS] was used as the eluant.
Results Oxoperoxonitrate(1-) (1-30 µM) in 1-20 mM sodium hydroxide was mixed at room temperature with tetranitromethane (10 µM to 1.2 mM) in the presence and absence of nitrogen monoxide (0.95-1.9 mM). All concentrations are those after mixing. The solutions were stored on ice and stable for hours. Formation of trinitromethanide (Figure 1) is observed in the absence and presence of nitrogen monoxide. The experiments without nitrogen monoxide yield results that are identical to those published by Mere´nyi and Lind (1). With nitrogen monoxide, we find rates between 0.005 and 0.0185 s-1, depending on presence or absence of nitrite, the pH, and how the kinetic traces are analyzed. In all experiments (with and without nitrogen monoxide), we observed directly after mixing small but systematic deviations from first-order kinetics (Figure 1). If these deviations are ignored, the reaction is first-order in oxoperoxonitrate(1-) and zero-order in tetranitromethane and rate constants upward of 0.005 s-1 are found. Determination of the rate constant by the initial rate method yields values closer to 0.018 s-1. Given a rate constant for the hydrolysis of tetranitromethane at pH 12 of 10-3 s-1 (not shown) and at pH 11 of 10-4 s-1 (1), the formation of the trinitromethanide takes place at a rate between 0.005 and 0.017 s-1. As a control experiment, we mixed tetranitromethane (0.01 mM) with nitrite (0.1-10 mM, pH 5-10), a common contaminant of oxoperoxonitrate(1-) preparations, and found that trinitromethanide was formed at a rate of (1.0 ( 0.1) × 10-3 s-1, independent of pH or nitrite concentration. During these experiments, nitrogen monoxide reacts
with the nitrogen dioxide released after reduction of tetranitromethane and forms dinitrogen trioxide that rapidly hydrolyzes (5, 6). When copper nitrilotriacetate was mixed with oxoperoxonitrate(1-), an acceleration of the decay at 302 nm by a factor 4 was observed (Figure 2). The reaction is close to zero-order in copper nitrilotriacetate and first-order in oxoperoxonitrate(1-). When oxoperoxonitrate(1-) is present in excess, its decay cannot be fitted to a single exponential. The curve is biphasic; the first part (k ) 0.062 ( 0.008 s-1) ends when oxoperoxonitrate(1-) is no longer in excess, and the second (k ) 0.032 ( 0.010 s-1) is for the remainder of the decay. The rate constant derived from the second part is, within the error, identical to that obtained with copper(II) nitrilotriacetate in excess (k ) 0.040 ( 0.008 s-1). These values have been corrected for the decay of oxoperoxonitrate at pH 9.3 (0.010 ( 0.03 s-1). During the first part of the reaction, oxoperoxonitrate(1-) is decomposed to nitrite and dioxygen, and during the second part, the product is nitrate. Copper(I) was never detected during the course of the oxoperoxonitrate(1-) decomposition.
Discussion If oxoperoxonitrate(1-) undergoes homolysis, then nitrogen monoxide and tetranitromethane compete for superoxide. How much oxoperoxonitrate(1-) and trimethylmethanide are formed is determined by the products of the rate constants of superoxide with nitrogen monoxide and tetranitromethane, respectively. Using typical concentrations of 0.5 mM tetranitromethane and 1 mM nitrogen monoxide, and rate constants of 2 × 109 (7) and 1.9 × 1010 M-1 s-1 (3), we calculate that 95% of all superoxide would form oxoperoxonitrate(1-) back. This stands in contrast to the experimental observation in which the presence of nitrogen monoxide did not significantly influence the formation of trimethylmethanide. Said differently, given a Gibbs energy of reaction of ∼18 kcal mol-1 for homolysis, reaction 1, the presence of millimolar concentrations of nitrogen monoxide should have kept the equilibrium entirely on the oxoperoxonitrate(1-) side. Thus, the observation that the trinitromethanide anion is formed in the presence of up to 1.9 mM nitrogen monoxide argues strongly against the interpretation of Mere´ny and Lind (1) that the oxoperoxonitrate(1-) anion undergoes homolysis to form nitrogen monoxide and superoxide, and that the latter reduces tetranitromethane. A second argument against homolysis is the finding that the kinetics of the reduction of tetranitromethane
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do not strictly follow the first-order behavior that is expected for a rate-limiting homolysis, followed by a rapid second-order reduction of tetranitromethane by superoxide. It is intriguing that in the presence of nitrite, tetranitromethane forms trinitromethanide anions with a rate independent of the nitrite concentration when the latter is in excess. Again, in this case, there is no secondorder rate law, but rather a first-order one. The formation of trinitromethanide from a reaction with nitrite indicates that tetranitromethane is a fairly strong oxidant with a reduction potential near 1 V [E°(NO2•aq/NO2-) ) 1.04 V (8)], which implies that there are no thermodynamic barriers to oxidizing oxoperoxonitrate(1-) to nitrosodioxidanyl, ONOO•, which requires ∼0.2 V (see below). Oxoperoxonitrate(1-) is also readily oxidized by tetraoxomanganate(1-) under formation of tetraoxomanganate(2-) (9); this couple has a reduction potential of 0.56 V (10). The third argument follows from the copper(II) nitrilotriacetate experiments. The rate of decay of oxoperoxonitrate(1-) should have been 0.017 s-1 if homolysis occurred and superoxide was scavenged by copper(II) nitrilotriacetate. This complex catalyzes the dismutation of superoxide with a rate constant of ∼8 × 107 M-1 s-1 as determined by pulse radiolysis (T. Nauser, 2000, unpublished). Given a copper(II) nitrilotriacetate concentration of 50 µM, superoxide would have disappeared at a rate of 4 × 103 s-1, which makes the homolysis ratelimiting. Instead of the expected rate of 0.017 s-1, we found a rate of decay that was 2-3 times higher in the presence of the copper complex. One could argue that this faster reaction occurs parallel to the homolysis. To find out, we carried out experiments with 2,2′-biquinolyl which forms thermodynamically stable complexes with copper(I). Had superoxide been produced, it would have reduced copper(II) nitrilotriacetate. Due to the labile nature of copper, the thermodynamically stable copper(I)-2,2′-biquinolyl complex would have been detected. Since we did not observe this complex, we conclude that homolysis and decomposition do not run concurrently. Additional aspects of the accelerated decay of oxoperoxonitrate(1-) by copper(II) nitrilotriacetate will be discussed in a future publication. A fourth argument is based on the consequences of the reported value of 16.6 kcal mol-1 for the standard Gibbs energy of formation of oxoperoxonitrate(1-). This value implies a standard entropy of 23 eu for oxoperoxonitrate(1-), which is much lower than the corresponding 35 eu for nitrate (11), and implies a hydration entropy of -45 eu, much larger than expected for oxoperoxonitrate(1-), when compared with similar anions, e.g., -23 eu for nitrate. We recommended values of 31 eu for S°, -37 for
Communications
∆hydS° (12), -10 ( 2 kcal mol-1 for ∆fH°(ONOO-) (13, 14), 14 ( 3 kcal mol-1 for ∆fG°(ONOO-), and 1.6 ( 0.2 V at pH 7 for E°′(ONOOH/NO2•, H2O), from which a value of 0.2 V follows for E°(ONOO•/ONOO-); see the oxidation state diagram in ref 15. We are of the opinion that these values more accurately reflect the thermodynamic properties of oxoperoxonitrate(1-).
Acknowledgment. We thank Dr. P. L. Bounds for helpful discussions. Supported by ETHZ.
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) Mere´nyi, G., Lind, J., Goldstein, S., and Czapski, G. (1999) Mechanism and thermochemistry of peroxynitrite decomposition in water. J. Phys. Chem. A 103, 5685-5691. (3) Kissner, R., Nauser, T., Bugnon, P., Lye, P. G., and Koppenol, W. H. (1997) Formation and properties of peroxynitrite studied by laser flash photolysis, high-pressure stopped flow and pulse radiolysis. Chem. Res. Toxicol. 10, 1285-1292. (4) Bohle, D. S., Hansert, B., Paulson, S. C., and Smith, B. D. (1994) Biomimetic synthesis of the putative cytotoxin peroxynitrite, ONOO-, and its characterization as a tetramethylammonium salt. J. Am. Chem. Soc. 116, 7423-7424. (5) Gra¨tzel, M., Taniguchi, S., and Henglein, A. (1970) Pulsradiolytische Untersuchung der NO-Oxydation und des Gleichgewichts N2O3 S NO + NO2 in wa¨ssriger Lo¨sung. Ber. BunsenGes. 74, 488-492. (6) Park, J.-Y., and Lee, Y.-N. (1988) Solubility and decomposition kinetics of nitrous acid in aqueous solution. J. Phys. Chem. 92, 6294-6302. (7) Bielski, B. H. J., Cabelli, D. E., and Arudi, R. L. (1985) Reactivity of hydrodioxyl/superoxide radicals in aqueous solution. J. Phys. Chem. Ref. Data 14, 1041-1100. (8) Stanbury, D. M. (1989) Reduction potentials involving inorganic free radicals in aqueous solution. Adv. Inorg. Chem. 33, 69-138. (9) Gleu, K., and Roell, E. (1929) Die Einwirkung von Ozon auf Alkaliazid. Persalpetrige Sa¨ure I. Z. Anorg. Allg. Chem. 179, 233266. (10) Weast, R. C., and Astle, M. J., Eds. (1981) Handbook of Chemistry and Physics, p D-136, CRC Press, Boca Raton, FL. (11) Wagman, D. D., Evans, W. H., Parker, V. B., Schumm, R. H., Halow, I., Bailey, S. M., Churney, K. L., and Nuttal, R. L. (1982) Selected values for inorganic and C1 and C2 organic substances in SI units. J. Phys. Chem. Ref. Data 11 (Suppl. 2), 37-38. (12) Koppenol, W. H., and Kissner, R. (1998) Can OdNOOH undergo homolysis? Chem. Res. Toxicol. 11, 87-90. (13) 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. (14) Manuszak, M., and Koppenol, W. H. (1996) The enthalpy of isomerization of peroxynitrite to nitrate. Thermochim. Acta 273, 11-15. (15) 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.
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