APRIL 1998 VOLUME 11, NUMBER 4 © Copyright 1998 by the American Chemical Society
Communications Free Radical Formation in the Peroxynitrous Acid (ONOOH)/Peroxynitrite (ONOO-) System Ga´bor Mere´nyi* and Johan Lind Department of Chemistry, Nuclear Chemistry, The Royal Institute of Technology, S-10044 Stockholm 70, Sweden Received February 16, 1998
The rate constant of homolysis of peroxynitrite, ONOO-, into O2•- and NO• was determined to be 0.017 s-1 at 20 °C. In combination with other experimental data taken from the literature, this value yields the Gibbs free energy of formation of ONOO-, ∆fG°(ONOO-) ) 16.6 kcal/mol. On the basis of this result, we conclude that peroxynitrous acid homolyzes to yield nitrogen dioxide (NO2•) and hydroxyl (OH•) free radicals and derive ∆fG°(ONOOH) ) 7.7 kcal/mol. The rate constant of the reaction between NO• and ONOO- was found to be 5 × 10-2 M-1 s-1 at most. The implications of the two homolysis reactions are discussed.
Introduction Peroxynitrite, ONOO-, is considered to be a reactive oxygen species (1), since the recent recognition of its rapid formation from nitric oxide (NO•) and superoxide (O2•-) (1-3). In vivo generation of ONOO- is thought to be important in the context of the macrophage immune response (4-7), and under conditions of oxidative stress such as endotoxic shock and ischemia/reperfusion, ONOOand its conjugate acid, peroxynitrous acid (ONOOH), are known to react with a broad range of biological substrates. These reactions include the nitration of tyrosine residues in proteins (8) and oxidation of DNA (9), lipids (10), sulfhydryls (11), and methionine (12). Peroxynitrous acid isomerizes to nitrate but also forms a highly reactive intermediate in a parallel reaction. At 20 °C the overall rate constant (13) of ONOOH decay, kd, is 0.8 s-1 and the yield (14) of the reactive intermediate is ca. 40%; hence the rate constant of intermediate formation is approximately 0.4kd ) 0.3 s-1. The nature
of this species is hotly debated at present. As a recent review (15) demonstrates, there is abundant evidence both in favor of and against the assumption that this intermediate is a mixture of hydroxyl (OH•) and nitrogen dioxide (NO2•) free radicals. Recently, we estimated (16) the Gibbs free energies of formation of ONOOH and ONOO- and, by means of thermokinetic arguments, concluded that the reactive species is very likely a mixture of OH• and NO2•. However, due to the relatively large uncertainty in the estimation, we could not decide whether the OH• and NO2• free radicals are the only or even the major reactive species that form. In the present work we shall adopt a different approach. We shall assume that the reactive intermediate that forms initially is the NO2• and OH• free radicals. The simplest way to account for the simultaneous formation of NO3- and OH• and NO2• free radicals is to invoke a radical pair cage as a common intermediate preceding product formation. Such a model has been proposed in ref 15 and elaborated in more detail in ref 16. However, for the purposes of
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the present study, the detailed reaction mechanism is not important and only the assumption of free radical formation according to reaction 1 is of essence. •
•
ONOOH h OH + NO2
(1)
Then, k1 at 20 °C is 0.3 s-1. Since the rate constant of the reverse reaction (17) is k-1 ) 4.5 × 109 M-1 s-1, we calculate K1 ) 7 × 10-11 M at 20 °C. Utilizing the Gibbs free energies of formation (18) (in kcal/mol), ∆fG°(OH•) ) 6.2, ∆fG°(NO2•) ) 15.06, ∆fG°(O2•-) ) 7.6, and ∆fG°(NO•) ) 24.38, as well as pKa(ONOOH) ) 6.5-6.8 (15), we immediately calculate K2 ) (2-4) × 10-12 M. -
•
•-
ONOO h NO + O2
(2)
With k-2 ) 5.5 × 109 M-1 s-1, the average of two literature values (2,3), we obtain k2 ) 0.01-0.02 s-1. This value is significantly higher than kd, the rate constant of peroxynitrite decay at pH values well above pH 9 (19). kd has a limiting upper value of ca. 1 s-1 at low pH and decreases roughly by 1 decade for every pH increase by 1 unit above pKa(ONOOH). However, at intermediate pH values (9-11), these rates vary between different laboratories by a factor of ca. 2-3, and kd appears to settle around a limiting value of a few times 10-5 s-1 above pH 12 (19). We reason that a very convincing case would be made for the homolysis reaction 1 if we were able to measure the rate constant k2 above ca. pH 9 and find its magnitude in the predicted interval of (1-2) × 10-2 s-1. If reaction 2 occurs, it forms O2•- and NO•. A possible strategy to measure k2 would then be rapidly to intercept O2•- or NO• by means of an added scavenger molecule and thus to prevent reaction -2 from occurring. Wellknown O2•- scavengers are the enzyme superoxide dismutase (SOD) and tetranitromethane [C(NO2)4]. As SOD is unstable at high pH, we chose to work with C(NO2)4 in the present study.
Results and Discussion Formation and Interception of O2•-. Tetranitromethane, C(NO2)4, is an ideal candidate for O2•scavenging, as reaction 3 is fast and yields the strongly colored nitroform anion, C(NO2)3-.
O2•- + C(NO2)4 f O2 + NO2• + C(NO2)3-
(3)
k3 ) 1.9 × 109 M-1 s-1 (20) We prepared ONOO- by ozonolysis (21,22) at 0 °C of a 0.1 M NaN3 solution containing 10-2 M NaOH. The UV absorbance at 302 nm showed that the final solutions contained between 0.018 and 0.064 M ONOO-, the rest being N3-. These solutions were diluted with water containing NaOH when necessary and mixed in 1/1 ratio in a stopped-flow equipment with a solution of excess C(NO2)4, the latter being dissolved in water or ammonia buffer. We measured the buildup rate and size of the absorbance at 360 nm, due to the formation of C(NO2)3(360 nm is the shortest wavelength measurable with our stopped-flow equipment). As C(NO2)4 reacts slowly with OH-, we chose 10.7 as the highest pH for our measurements. At this pH the rate of formation of C(NO2)3- in the absence of ONOO- was found to be 2 × 10-4 s-1, i.e.,
Table 1. Dependence on pH and C(NO2)4 (TNM) and ONOO- (PN) Concentrations of the Rate Constants and Yields of C(NO2)3- (NF) Formation and PN Decaya pH
k (s-1)
4.8 0.8 4.8 0.8 6.8 1.0 6.8 0.9 9.3 0.020 9.3 0.003 10.7 0.018 10.7 0.024 10.7c ≈10-4
yield (%), TNM PN ∆ODb NF/PN (mM) (mM) (360 nm) 0
0.6 0 0.6 0 0.15 0 0.24 1.2 0
0 81 89 99
0.36 0.36 0.18 0.18 0.03 0.10 0.06 0.06 0.10
≈-0.03 ≈-0.03 ≈-0.05 ≈-0.05 0.31 ≈-0.03 0.68 0.76
buffer (20 mM) acetate acetate phosphate phosphate ammonium ammonium
a All concentrations are given in the final mixture. b ∆OD denotes change in optical density. c Measured at 302 nm in a spectrophotometer.
ca. 100 times slower than the corresponding rate in the presence of ONOO-. Control experiments in the absence of ONOO- also revealed that neither N3- nor the ammonia buffer had any effect on the buildup rate of C(NO2)3- at the concentrations used in this work. Our data are compiled in Table 1. Each entry gives the average of five measurements. It is seen that at pH 4.8 and 6.8 no C(NO2)3- forms. Instead, a slight bleaching is observed at 360 nm, which can be ascribed to the decay of ONOOH/ONOO- with roughly the expected rate constants, kd (19). However, at pH 9.3 and 10.7 we observe the buildup of the characteristic C(NO2)3- absorbance. With 12 800 M-1 cm-1, the extinction coefficient of C(NO2)3- at 360 nm, as determined by us in the present work, we find that, within the experimental accuracy of ca. 10%, one ONOO- molecule yields one C(NO2)3-. The buildup is described by a single exponential. The rates are first-order in [ONOO-] and essentially unaffected by variation of [C(NO2)4] or the rise in pH from 9.3 to 10.7. Equally important, at these pH values and in the absence of C(NO2)4, ONOO- disappears more slowly and with a strongly pH-dependent rate constant, kd (see Table 1). The latter values, which within a factor of ca. 3 agree with the kd values reported in the literature (19), were only measured to emphasize the difference in rate constants in the presence versus absence of C(NO2)4. For example, at pH 10.7 the two rate constants differ by a factor of ca. 200 [ca. 10-4 s-1 in the absence as compared to (1.8-2.4) × 10-2 s-1 in the presence of C(NO2)4]. Closer scrutiny reveals that, at this pH, the buildup rate of C(NO2)3- increases slightly but systematically with increasing C(NO2)4 concentration. Thus, upon varying the latter from 0.24 to 1.2 mM the buildup rate of C(NO2)3increases from 0.018 to 0.024 s-1. This increase can be ascribed to reaction 4.
ONOO- + C(NO2)4 f NO• + O2 + C(NO2)3- + NO2• (4) The observed rates can be presented as k ) k0 + k4[C(NO2)4], where k0 ) 0.017 s-1 is the rate constant extrapolated to zero [C(NO2)4] and k4 ) 5 M-1 s-1. At pH 9.3, kd in the absence of C(NO2)4 is 0.003 s-1. Then the rate constant k in the presence of 0.15 mM C(NO2)4 is k ) 0.020 ) k0 + kd + k4[C(NO2)4]. Inserting k4 ) 5 M-1 s-1 and kd ) 0.003 s-1, we obtain k0 ) 0.016 s-1, which, within experimental accuracy, is the same as k0 obtained at pH 10.7. In conclusion, the rate constant of C(NO2)3- buildup extrapolated to zero [C(NO2)4], 0.017
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Chem. Res. Toxicol., Vol. 11, No. 4, 1998 245
s-1, is independent of the pH. It is then reasonable to identify this rate constant with k2, the more so as its value is in the interval predicted on the assumption of reaction 1 occurring. Clearly, in the presence of 0.1 mM or higher C(NO2)4 concentrations, the rate of reaction 3 exceeds 2 × 105 s-1 and hence reaction 2 becomes the rate-limiting step in the formation of C(NO2)3-. Role of NO•. If O2•- and NO• form in reaction 2 at high pH with k2 ) 0.017 s-1, we are faced with a dilemma, since according to ref 23 NO• reacts with ONOO- in reaction 5, the reported k5 being as high as 9.1 × 104 M-1 s-1.
NO• + ONOO- f NO2• + NO2-
(5)
In such a case ONOO- would act as a NO• scavenger, and reaction 5 would compete with reaction -2. This would entail a dependence at high pH on [ONOO-] of the apparent kd of ONOO- decay, with the latter approaching a limiting value at high [ONOO-]. This is not what is observed. Thus, e.g., at pH 11 the measured kd is only ca. 10-4 s-1 and essentially independent of [ONOO-]. Such a high k5 value would also clearly preclude standard preparations of 0.1 M or so peroxynitrite stock solutions at high pH, in contradiction with routine laboratory practice. Therefore, if reaction 2 occurs reaction 5 must have a negligible rate. In view of this contradiction in terms we reexamined reaction 5. First, we observe that the authors of ref 23 studied reaction 5 by following the disappearance of NO• at pH 7.4, where kd of ONOO- is rather high and thus should strongly interfere with the measurements. In particular, the OH• and NO2• radicals which form in reaction 1 are known rapidly to react with and thus to consume NO•, the latter being present in excess. Therefore, we decided to study reaction 5 at pH 12, where kd is very slow. The experiments were done as follows. Two solutions were prepared: 0.25 mL of a 2.9 mM pH 12 stock solution of ONOO- was added to 2.25 mL of (a) a pH 12 aqueous solution and (b) a deaerated NO•-saturated pH 12 solution. In both solutions a and b the final concentration of ONOO- is thus 0.29 mM. Furthermore, solution b contains initially 1.7 mM NO• and 0.026 mM O2. Figure 1 shows a small difference between the two spectra recorded after 10 min. If we ascribe this difference completely to reaction 5, we calculate k5 ) 5 × 10-2 M-1 s-1. At any rate k5 cannot exceed this value. The smallness of k5 clearly shows that reaction 5 plays no role in the decay of ONOO- In conclusion, NO• has negligible reactivity toward ONOO-, which is exactly what is required if reaction 2 is to occur with the measured rate constant of 0.017 s-1. Consequences of the Homolysis Reactions 1 and 2. Given that reaction 5 is very slow, NO• and O2•-, formed in reaction 2, can only disappear in reactions 6 and 7.
2O2•- + H2O f O2 + HO2- + OH-
(6)
k6 ) 108[H+]/(10-4.8 + [H+]) M-1 s-1 (24) 4NO + O2 + 4OH- f 4NO2- + 2H2O
(7)
Figure 1. Spectra of 0.29 mM ONOO- in the presence and absence of NO at pH 12: s, spectra at 20 s and 10 min after mixing of ONOO- with water; b, spectrum observed at 10 min after mixing ONOO- with NO-saturated water.
4k7[O2] ≈ 3 × 103 M-1 s-1 in air-saturated solutions (25) However, as can be seen, the rate constants of the latter are much smaller than k-2. Applying a steadystate treatment for NO• and O2•-, we immediately obtain the relationship: keff ) K2(8k6k7[O2])1/2, where keff is the effective rate constant with which ONOO- would disappear via reaction 2. Let us calculate keff at, e.g., pH 11. Inserting the values k6 ) 63 M-1 s-1, k7[O2] ) 750 M-1 s-1, and K2 ) k2/k-2 ) 0.017/(5.5 × 109) ) 3.1 × 10-12 M, we obtain keff ) 2 × 10-9 s-1. As this is smaller by almost 5 orders of magnitude than the natural kd ≈ 10-4 s-1, it is clear that reaction 2 will not be seen kinetically. Since both the predicted magnitude of k2 as well as the prediction of negligible reactivity of NO• toward ONOOhave been confirmed in the present work, we conclude that reaction 1, i.e., the homolysis of ONOOH into OH• and NO2• free radicals, also takes place. As the experimental rate constant of ONOO- decay is described (15, 19) roughly by kd ) 0.8[H+]/(10-6.8 + [H+]), where 0.8 s-1 and 6.8 are the rate constant of decay and pKa, respectively, of ONOOH at 20 °C, it is readily seen that kd becomes equal to k2 around pH 8.4. We note that Beckman et al. (26) observed an effect on the nitration efficiency as well as the rate of nitration of 4-hydroxyphenylacetate by ONOO- upon addition of SOD to the solution at pH 8 and 37 °C. Given that SOD is an efficient O2•- scavenger, we believe that these observations might bear on the interception by SOD of O2•formed in reaction 2. However, as the system studied in ref 26 is complex and the measurements are indirect, quantitative treatment of these data is not possible. While reaction 2 does not influence the kinetics of ONOO- decomposition in the absence of scavengers, it is expected to affect the pH dependence of the product yield. This can be seen as follows. As k1 ) 0.4kd, k1 becomes equal to k2 at ca. pH 8. This means that well below pH 8, apart from the directly formed NO3-, only NO2• and OH• radicals are produced to any appreciable extent. Depending on the pH (i.e., the [ONOO-]/ [ONOOH] ratio) and the concentration of NO2- in the solution, OH• will be converted to NO2• or NO• in reactions
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8 and 9 (27).
OH• + NO2- f OH- + NO2•
(8)
OH• + ONOO- f OH- + O2 + NO•
(9)
Therefore, we shall always have an excess of NO2• over NO•, and thus the radicals will produce both NO3- and NO2- in reactions (28) 10 and 11.
2NO2• + H2O f NO3- + NO2- + 2H+
(10)
NO2• + NO• + H2O f 2NO2- + 2H+
(11)
However, well above pH 8 k2 is larger than k1. Here, a pseudoequilibrium between ONOO- and NO• + O2•will be established more rapidly than ONOO- has time to decay in reaction 1. Thus, whenever a NO2• radical forms in reaction 1, it will see an excess of NO• and O2•and will therefore be converted to NO2- either in reaction 11 or in reaction 12 (29). The crossover in the product yield should occur around pH 8, where all four radicals are produced at about the same rate and thus the system is at its most complex.
NO2• + O2•- f NO2- + O2
(12)
The two homolysis reactions 1 and 2 thus nicely explain the observation in ref 23, namely, that above ca. pH 8 the nitrite yield increases at the expense of the nitrate yield. Thermodynamics of the ONOOH/ONOO- System. Combination of k2 ) 0.017 s-1 with k-2 ) 5.5 × 109 M-1 s-1, the average of the two literature values (2,3), yields K2 ) 3.1 × 10-12 M. Then, by means of the free energies of formation (18) of NO• and O2•-, ∆fG°(ONOO-) ) 16.6 kcal/mol is calculated. As was pointed out before, we feel confident to conclude that k1 ) 0.4 × 0.8 ) 0.32 s-1, where 0.8 s-1 is the kd value at 20 °C of ONOOH and 0.4 is the fraction of free radical formation, i.e., the formation of OH• and NO2•. As k-1 is 4.5 × 109 M-1 s-1, we obtain K1 ) 7.1 × 10-11 M. Together with ∆fG°(OH•) ) 6.2 kcal/ mol and ∆fG°(NO2•) ) 15.06 kcal/mol, we obtain ∆fG°(ONOOH) ) 7.7 kcal/mol. From this we calculate pKa(ONOOH) ) 6.6. While this value is somewhat lower than 6.8, the value currently accepted, we have good evidence (30) that 6.6 is the correct thermodynamic pKa value. However, be it as it may, the consistency of the double-homolysis scheme is gratifying.
Acknowledgment. We thank Dr. Martin Ragnar for running the stopped-flow experiments and the Swedish Natural Science Research Council for financial support.
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