132
Chem. Res. Toxicol. 1999, 12, 132-136
Effect of •NO on the Decomposition of Peroxynitrite: Reaction of N2O3 with ONOOSara Goldstein,*,† Gidon Czapski,† Johan Lind,‡ and Ga¨bor 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 November 17, 1998
Nitric oxide reacts slowly with ONOO- (k < 1.3 × 10-3 M-1 s-1), and therefore does not affect the stability of peroxynitrite at pH >12. A chain consumption of peroxynitrite by •NO takes place at pH pKa, provided that its reaction with ONOO- competes efficiently with the hydrolysis of N2O3. In this study, we show that •NO affects the decomposition of peroxynitrite through the very fast reaction of ONOO- with N2O3, where the latter is formed during autoxidation of •NO and/or during the spontaneous decomposition of peroxynitrite.
Experimental Procedures Chemicals. Peroxynitrite was synthesized through the reaction of nitrite with acidified hydrogen peroxide in a quench flow as described recently (11). The peroxynitrite concentration was determined using an 302 1670 M-1 cm-1 (12). •NO solutions in 1-100 mM phosphate buffer were prepared as described elsewhere (10), and were diluted with deaerated buffered solution to the desired concentrations by the syringe technique. Apparatus. Kinetic measurements at 25 °C were taken using the Bio SX-17MV Sequential Stopped-Flow apparatus from Applied Photophysics, UK, and a HP 8452A diode array spectrophotometer coupled with a thermostat (HP 89075C Programmable Multicell Transport). Modeling of the experimental results was carried out using a noncommercial program developed at Brookhaven National Laboratories by H. A. Schwarz.
10.1021/tx9802522 CCC: $18.00 © 1999 American Chemical Society Published on Web 01/22/1999
Communications
Chem. Res. Toxicol., Vol. 12, No. 2, 1999 133
Figure 1. Reaction of 166 µM ONOO- with 1.36 mM •NO at pH 11.5 in the presence of 69 µM O2 (lower trace) and in the absence of O2 (upper trace).
Results The decay of peroxynitrite was followed at 302 nm under the conditions where deaerated or aerated solutions of peroxynitrite at pH 12-13 were mixed with •NO solutions in 1:5 or 2:5 ratios to yield final concentrations of 0.7-1.58 mM •NO, 20-335 µM ONOO-, 0.71-83.3 mM phosphate, and 0, 40, or 69 µM O2. At pH 12.45, the decay of 154 µM ONOO- during the first 500 s was unaffected, within our experimental accuracy of ca. 10%, by 1.36 mM •NO in the presence of 69 µM O , as well as in the absence 2 of O2. As the decay rate constant of ONOO- at this pH is 1.8 × 10-5 s-1, an upper limit of 1.3 × 10-3 M-1 s-1 can be set for the reaction of •NO with ONOO-. At pH 8. Typical kinetic traces of the first fast reaction in the absence and presence of O2 at pH 11.5 are given in Figure 1. The observed first-order rate constant of the first fast decay was almost unaffected by the pH and depended linearly on [•NO]o2 (Table 1), resulting in a third-order rate constant of (1.9 ( 0.2) × 106 M-2 s-1 at pH >11 (Table 1). The rate constant of the second process was highly pH-dependent, and was similar to that of the spontaneous decomposition of peroxynitrite at the same pH. The amount of peroxynitrite consumed, ∆[ONOO-]consumed, via the first fast process increased upon decreasing the pH and increasing [ONOO-]o (Table 1). Around neutral pH, the kinetics of the decomposition of peroxynitrite in the presence of excess of •NO was highly dependent on peroxynitrite and phosphate concentrations. The effect of •NO on the decomposition rate of peroxynitrite was similar to that observed in alkaline solutions, but only for relatively low peroxynitrite concentrations. Figure 2 shows that 1.36 mM •NO affects the decomposition of 58 µM peroxynitrite at pH 7.4 (71.4 mM phosphate) only in the presence of oxygen. When the peroxynitrite concentration is increased or phosphate concentration is decreased, some peroxynitrite was consumed rapidly, and the rest decayed via the two sequential first-order processes as described above in alkaline
solutions. For example, under the experimental conditions of Figure 2, but with 154 µM peroxynitrite, ca. 50% of peroxynitrite was consumed rapidly, and the rest decayed via two sequential first-order reactions with a kd(1) of 4.8 ( 0.3 s-1 and a kd(2) of 0.29 ( 0.01 s-1. In the presence of 311 µM peroxynitrite, as much as 85% decayed within the mixing time. At pH 6.9 (5.7 mM phosphate), all 150 µM peroxynitrite was consumed within the mixing time in the presence of 0.34-1.36 mM •NO with or without O . When the peroxynitrite concen2 tration is greater than the •NO concentration, the consumption of •NO occurred within the mixing time, and the remainder of peroxynitrite decayed spontaneously (Figure 3). The decomposition of ONOOH at pH 4.45 (29 mM acetate buffer) was followed at 280 nm, and the kd was determined to be 1.18 ( 0.02 s-1. When •NO was added, kd increased slightly, displaying a mild dependence on [ONOOH]o. For example, in the presence of 1.36 mM •NO and 69 µM O , k increased from 1.34 ( 0.01 to 1.52 2 d ( 0.02 s-1 upon increasing [ONOOH]o from 55 to 388 µM.•
Discussion •NO
reacts very slowly with ONOO- (2), and in this study, an upper limit of 1.3 × 10-3 M-1 s-1 was set for the rate constant of this reaction. It is also shown that •NO accelerates the decomposition of peroxynitrite in alkaline solutions only in the presence of oxygen. When the •NO concentration is much greater than the O2 concentration, the rate of this process was found to be second-order with respect to •NO concentration, firstorder with respect to ONOO- concentration, and hardly affected by the pH, resulting in a third-order rate constant of (1.9 ( 0.2) × 106 M-2 s-1 at pH >10. This rate constant is in excellent agreement with the literature value for the autoxidation of •NO in alkaline solutions (10, 13). It is therefore concluded that •NO affects the stability of ONOO- in alkaline solutions through intermediates that are formed during autoxidation of •NO. The mechanism of the autoxidation of •NO is given by the sequence of reactions 3-7 (10, 14, 15): •
NO + O2 h ONOO•
(3)
ONOO• + •NO f ONOONO
(4)
fast
ONOONO 98 2•NO2 •
NO2 + •NO h N2O3
(5) (6)
k6 ) 1.1 × 109 M-1 s-1 k-6 ) 8.4 × 104 s-1 (7) N2O3 + H2O f 2NO2- + 2H+
(7)
k7 ) ko + kOH[OH-] + kp[phosphate] According to this mechanism, -d[O2]/dt ) kauto[•NO]2[O2], and kauto ) k4K3 ) 2-2.9 × 106 M-1 s-1 (10, 13-16). It has been shown in several studies that oxidation and nitrosation of various substrates take place in oxygenated •NO solutions either through •NO or through N O (10, 2 2 3 14-16). Since the rate constant of the reaction of •NO2 with ONOO- does not exceed 2 × 104 M-1 s-1 (17), we suggest that N2O3 is the intermediate reacting with
134 Chem. Res. Toxicol., Vol. 12, No. 2, 1999
Communications
Table 1. Rate Constant of Peroxynitrite Consumption under Various Conditions [•NO]o (mM)
[O2]o (µM)
[phosphate] (mM)
pH
[ONOO-]o (µM)
kdecay (s-1)
exptl ∆[ONOO-]consumed (µM)
calcda ∆[ONOO-]consumed (µM)
1.36 1.36 1.36 1.58 1.58 1.36 1.05 0.7 0.7 1.36 1.58 0.79 1.58 1.58 1.58
69 69 69 40 40 69 40 40 40 69 40 40 40 40 40
0.71 0.71 0.71 8.33 8.33 0.71 8.33 8.33 8.33 0.71 83.3 83.3 83.3 83.3 83.3
11.61 11.5 11.5 11.4 11.4 11.4 11.45 11.45 11.4 11.0 10.45 9.65 9.48 8.50 8.08
267 166 60 160 84 112 335 335 180 142 96 150 170 107 42
3.0 ( 0.2 2.9 ( 0.1 3.1 ( 0.4 5.2 ( 0.4 4.8 ( 0.2 3.0 ( 0.2 1.9 ( 0.2 0.84 ( 0.04 0.77 ( 0.02 3.6 ( 0.1 4.9 ( 0.2 1.7 ( 0.5 5.5 ( 0.2 5.4 ( 0.6 6.2 ( 0.6
28 ( 2 20 ( 2 8(1 18 ( 2 9(1 15 ( 2 31 ( 2 33 ( 2 21 ( 2 47 ( 1 36 ( 1 66 ( 4 72 ( 4 66 ( 6 19 ( 1
32 ( 4 24 ( 3 8(1 18 ( 2 9(1 20 ( 2 42 ( 6 42 ( 6 21 ( 2 65 ( 6 28 ( 7 71 ( 15 88 ( 20 46 ( 13 15 ( 4
a The rate of consumption of peroxynitrite was calculated using simulations of reactions 3-8, for which k ) (2 × 103) + 108[OH-] + 7 [(8 ( 2) × 105[phosphate]] s-1 and k8 ) (3.1 ( 0.3) × 108 M-1 s-1.
Reaction 8 is similar to reaction 9, where ONOO- was identified as the final product of the nitrosation of H2O2 by •NO/O2 (10).
N2O3 + HO2- f ONOO- + NO2- + H+
Figure 2. Effect of oxygen on the decomposition of 58 µM peroxynitrite in the presence of 1.36 mM •NO at pH 7.4 (71.4 mM phosphate). For the upper curve, [O2] ) 0 and kd ) 0.27 s-1 (the same rate constant as for the spontaneous decomposition of peroxynitrite at this pH). For the lower curve, [O2] ) 69 µM, kd(1) ) 5.8 s-1, and kd(2) ) 0.28 s-1.
Figure 3. Decomposition of 154 µM peroxynitrite at pH 6.9 (5.7 mM phosphate buffer) in the presence of 0.113 mM •NO and 69 µM O2 (ca. 50% of peroxynitrite is consumed within the mixing time).
ONOO-. This suggestion is in keeping with the wellknown electrophilicity of N2O3 (18). The nitrosation of ONOO- by N2O3 can be described by reaction 8, followed by reaction 5, thus yielding 2•NO2.
N2O3 + ONOO- f ONOONO + NO2-
(8)
(9)
Further evidence for the occurrence of reaction 8 is the effect of the pH on the rate of the initial consumption of ONOO-, which decreases upon increasing the pH (Table 1), in agreement with a base-catalyzed hydrolysis of N2O3 (9). Under our experimental conditions, the rate of the autoxidation of •NO is considerably faster than that of the self-decay of peroxynitrite, and the two processes are well separated above pH 8. Simulations of the first fast decay of ONOO- in the presence of excess of •NO over O2 (reactions 3-8) fit the experimental kinetic traces for different ratios of k8/k7. The rate constant of the reaction of N2O3 with phosphate (kp) has been determined to be (6.4-9.4) × 105 M-1 s-1 at pH 7.4 (14, 16). Therefore, under our experimental conditions, k8/k7 ) k8/kOH[OH-] only at pH >10.5, where we determined k8/k7 ) (2.2 ( 0.1) × 103, (1.2 ( 0.15) × 103, (8.3 ( 0.7) × 102, and (6.9 ( 0.7) × 102 M-1 at pH 11.0, 11.4, 11.5, and 11.61, respectively. A plot of k7/k8 as a function of OH- concentration yields a straight line, the slope kOH/k8 being 0.32 ( 0.03. As kOH ) 1 × 108 M-1 s-1 (9), one calculates k8 to be (3.1 ( 0.3) × 108 M-1 s-1. Using a ko of 2 × 103 s-1 (9), a kOH of 1 × 108 M-1 s-1 (9), and a k8 of (3.1 ( 0.3) × 108 M-1 s-1 (this study), we determine that kp ) (8 ( 2) × 105 M-1 s-1, in excellent agreement with the literature value (14, 16). By using the above rate constants, we calculated ∆[ONOO-]consumed (Table 1, last entry) and found it in good agreement with experiment. The finding that k8 is somewhat smaller than k9 is consistent with ONOO- being a weaker base than HO2-. The rate of the autoxidation of •NO is hardly affected by the pH (10, 13), whereas the rate of spontaneous decomposition of peroxynitrite increases upon decreasing the pH (1, 2), and becomes significant as compared to the autoxidation of •NO. Figure 2 shows that at pH 7.4, the two processes are not well separated, and the firstorder rate constants for the two processes were determined using a double-exponential fit. The rapid consumption of peroxynitrite around neutral pH occurs even in the absence of O2. As •NO reacts very slowly with ONOO-, we have to assume that the rapid consumption of peroxynitrite around neutral pH takes place via a
Communications
Chem. Res. Toxicol., Vol. 12, No. 2, 1999 135
different mechanism; i.e., N2O3 must be formed during the spontaneous decomposition of peroxynitrite. This can occur when ONOOH homolyzes into •OH and •NO2 (reaction 1), and •NO2 is scavenged by •NO to yield N2O3. Reactions 1 and 3-8 constitute a chain reaction with reactions 1, 3, and 4 being the initiation, 5, 6, and 8 the propagation, and 7 the termination. Since two •NO2 molecules are formed in reactions 8 and 5, a branched chain reaction takes place. The explosive nature of this reaction sequence is clearly reflected in the extremely fast (essentially immeasurable) rate of destruction of ONOOaround neutral pH in the presence of high concentrations of peroxynitrite or relatively low concentrations of phosphate. A rate equation (eq 10) can be obtained for the decomposition of peroxynitrite (PN) by assuming the steady-state approximation for all intermediates.
-
{
d[PN] ) kd + (k1[PN] + dt
}
k8 [PN] (10) 2kauto[•NO]2[O2]) k7 - k8[PN]
Here, k1 ≈ kd/3 as 60-70% of ONOOH forms NO3directly (5, 6, and references therein). In alkaline solutions, kd and hence k1 can be ignored, and eq 11 is obtained for the consumption yield of peroxynitrite, which agrees well with the experimental yields given in Table 1.
k8[ONOO-]o ∆[ONOO-]consumed ≈ 2[O2]o k7 - k8[ONOO-]o
(11)
Equations 10 and 11 show qualitatively how explosion can occur when k7 - k8[ONOO-]o approaches zero. Clearly, however, the equations are invalid when k7 < k8[ONOO-], as in this regime the steady-state approximation is not applicable. Reaction 8 is slowed upon decreasing the pH as ONOOH does not react with N2O3. This is in keeping with the much poorer nucleophilicity of a hydroperoxide compared to that of its anion. A similar difference in reactivity was also observed for H2O2 versus HO2- (10). In view of these considerations, •NO should have but a slight effect on the decomposition rate of peroxynitrite in acidic solutions, as no matter how much peroxynitrite is present (given that ONOOH is inert toward N2O3), k7 remains always much larger than k8[ONOO-] under realistic experimental conditions (see eq 10). Pfeiffer et al. (1) interpreted their results in terms of a rate constant of 9.1 × 104 M-1 s-1 for the reaction of •NO with ONOO-, whereas we determined in this study an upper limit of 1.3 × 10-3 M-1 s-1 for this rate constant. They used a •NO electrode to measure •NO consumption when 0.25-1 µM ONOO- was added to 1-2 µM •NO at pH 7.4 (0.1 M phosphate). Under these conditions, the autoxidation of •NO (under aerated conditions) can be ignored, but •NO is an efficient scavenger of •OH and •NO , which are formed in reaction 1 with a yield of ca. 2 30-40% (5, 6, and references therein). Pfeiffer et al. (1) reported that 0.75 µM peroxynitrite reacting with 1-2 µM •NO induced a rapid consumption of 0.66 ( 0.06 µM •NO with initial rates of 100 ( 9 nM s-1. In this case, the hydrolysis of N2O3 competes efficiently with its reaction with ONOO-, and therefore, the consumption
of •NO should account for ca. 60-80% of added peroxynitrite, in agreement with their experimental observations. In addition, according to our model the rate of •NO consumption is first-order, and therefore from the initial rates reported, one calculates a kconsumption of 100 ( 9 nM s-1/0.75 µM (0.13 ( 0.01 s-1), which is within experimental accuracy of the rate constant of the decomposition of peroxynitrite at pH 7.4 and ambient temperature.
Conclusions N2O3 can be formed during the spontaneous decomposition of peroxynitrite at pH >8 as reactions 1 and 2 produce •NO2 and •NO, respectively. This should affect the kinetics of the decomposition of peroxynitrite, and a deviation from first-order kinetics is expected at between pH 8 and 11, provided that the ONOO- concentration is sufficiently high. Under physiological conditions ([O2] < 0.24 mM and [HCO3-] ) 12-30 mM), the rate of peroxynitrite decomposition in the absence (1, 3, 4) as well as in the presence of CO2 (19) exceeds by orders of magnitude that of the nitrosation of ONOO- by •NO/O2 (13-16). Therefore, the decomposition of peroxynitrite is not affected by the autoxidation of •NO, and the latter can only act as an efficient scavenger of •NO2, which can be formed in reaction 1 or as an intermediate during the reaction of peroxynitrite with CO2 (20). In this case, N2O3 will be formed, but its reaction with ONOO- will take place only if it can compete with the hydrolysis of N2O3, i.e., if k8[ONOO-] is not much smaller than k7. The rate constants for the reaction of N2O3 with bicarbonate, phosphate, and chloride are 1.9 × 106, 6 × 105, and 1.4 × 105 M-1 s-1, respectively (16, 21). Therefore, under physiological conditions where bicarbonate, phosphate, and chloride concentrations are 12-30, 5-80, and 5-110 mM, respectively (21), reaction 8 will compete with reaction 7 if the peroxynitrite concentration is >100 µM. As the peroxynitrite concentration is certainly much smaller than that in vivo, we conclude that under physiological conditions the reaction of •NO with peroxynitrite cannot give rise to chain consumption of peroxynitrite.
Acknowledgment. This research was supported by The Israel Science Foundation. G.M. and J.L. thank the Swedish Natural Science Research Council for financial support.
References (1) Pfeiffer, S., Gorren, A. C. F., Schmidt, K., Werner, E. R., Hansert, B., Bohle, D. S., and Mayer, B. (1997) Metabolic fate of peroxynitrite in aqueous solution. J. Biol. Chem. 272, 3465-3470. (2) Mere´nyi, G., and Lind, J. (1998) Free radical formation in the peroxynitrous acid (ONOOH)/peroxynitrite (ONOO-) system. Chem. Res. Toxicol. 11, 243-246. (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) 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. (5) Lymar, S. V., and Hurst, J. K. (1998) Radical nature of peroxynitrite reactivity. Chem. Res. Toxicol. 11, 714-715. (6) 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.
136 Chem. Res. Toxicol., Vol. 12, No. 2, 1999 (7) Graetzel, M., Taniguchi, S., and Henglein, A. (1970) Pulsradiolytische untersuchung der NO-oxydation und des gleichgewichts N2O3 > NO + NO2 in waessriger loesung. Ber. Bunsen-Ges. Phys. Chem. 74, 488-492. (8) Logager, T., and Sehested, K. (1993) Formation and decay of peroxynitric acid. J. Phys. Chem. 97, 10047-10052. (9) Treinin, A., and Hayon, E. (1970) Absorption spectra and reaction kinetics of NO2, N2O3, and N2O4 in aqueous solution. J. Am. Chem. Soc. 92, 5821-5828. (10) Goldstein, S., and Czapski, G. (1996) The formation of peroxynitrite from the nitrosation of hydrogen peroxide by oxygenated nitric oxide solution. Inorg. Chem. 35, 5935-5940. (11) Saha, A., Goldstein, S., Cabelli, D., and Czapski, G. (1988) Determination of the optimal conditions for synthesis of peroxynitrite by mixing acidified hydrogen peroxide with nitrite. Free Radical Biol. Med. 23, 653-659. (12) 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. (13) Awad, H. H., and Stanbury, D. M. (1993) Autoxidation of NO in aqueous solution. Int. J. Chem. Kinet. 25, 375-381. (14) Goldstein, S., and Czapski, G. (1996) Mechanism of nitrosation of thiols and amines by NO/O2 in aqueous solution: The nature of the nitrosating intermediate. J. Am. Chem. Soc. 118, 34193425.
Communications (15) Goldstein, S., and Czapski, G. (1995) Kinetics of autoxidation of NO in the absence and presence of various reductants: The nature of the oxidizing intermediates. J. Am. Chem. Soc. 117, 12078-12084. (16) Lewis, R. S., Tannenbaum, S. R., and Deen, W. M. (1995) Kinetics of N-nitrosation in oxygenated nitric oxide solutions at physiological pH. Role of nitrous anhydride, phosphate and chloride. J. Am. Chem. Soc. 117, 3933-3939. (17) Goldstein, S., Saha, A., Lymar, S., and Czapski, G. (1998) Oxidation of peroxynitrite by inorganic radicals: A pulse radiolysis study. J. Am. Chem. Soc. 120, 5549-5554. (18) Williams, D. L. H. (1988) Nitrosation, Cambridge University Press, Cambridge, U.K. (19) Lymar, S. V., and Hurst, J. K. (1995) Rapid reaction between peroxynitrite ion and carbon dioxide: Implications for biological activity. J. Am. Chem. Soc. 117, 8867-8868. (20) Goldstein, S., and Czapski, G. (1998) Formation of peroxynitrate from the reaction of peroxynitrite with CO2: Evidence for carbonate radical production. J. Am. Chem. Soc. 120, 3458-3463. (21) Caulfield, J. L., Singh, S. P., Wishnok, J. S., Deen, W. M., and Tannenbaum, S. R. (1996) Bicarbonate inhibits N-nitrosation in oxygenated nitric oxide solutions. J. Biol. Chem. 271, 25859-25863.
TX9802522
