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Chem. Res. Toxicol. 1998, 11, 714-715
Radical Nature of Peroxynitrite Reactivity Sergei V. Lymar† and James K. Hurst* Department of Chemistry, Washington State University, Pullman, Washington 99164-4630, and Brookhaven National Laboratory, Upton, New York 11973-5000 Received April 14, 1998
General Dynamical Properties. The peroxynitrite anion (ONOO-) is a chemically unstable strong oxidant (1). Its decomposition rate in alkaline solutions is slow but is dramatically increased upon protonation or binding of other Lewis acids (1-4). Both one-electron and twoelectron oxidations by these adducts have been observed, and pathways involving both rate-limiting bimolecular interaction of reactants and rate-limiting unimolecular activation of peroxynitrite have been identified (5-11). One distinctive characteristic of the unimolecular (or “indirect”) pathway is that product yields are always substantially less than stoichiometric limits. This behavior requires postulation of two discrete intermediates, both of which produce NO3- but only one of which oxidizes other compounds. As described below, there are two prevailing opinions concerning the nature of these intermediates. Inadequacies of Kinetic Models Based upon CisTrans Isomerization. Early papers describing the reactivity of ONOOH were interpreted in terms of intermediary formation of •OH or “hydroxyl radical-like” species (5). Formation of hydroxyl radical was subsequently discounted by thermodynamic calculations (12) and replaced by (i) models that ascribed reactivity differences to geometrical isomers formed by cis vs trans orientation about the ON-OOH bond (6) or (ii) models based upon cis-trans isomerization as the unimolecular activation step (12). In the latter case, the reactive form was thought to be an unstable “transoid” intermediate whose properties resembled the transition state for cistrans isomerization. It is instructive to apply these models to simple “outer-sphere” one-electron redox processes, whose reactivities can be understood within the context of the Marcus formalism (13). Assuming E°′(ONOOH,H+/•NO2) ) 1.6-1.7 V (pH 7) (14, 15) and E°(Fe(CN)64-/3-) ) 0.42 V, the electron-transfer reaction:
ONOOH + Fe(CN)64- f •NO2 + Fe(CN)63- + OH- (1) is exergonic at pH 7 by 8-10 kcal/mol. (Note that most of the overall driving force implied by the reduction potentials is obtained from protonation of OH-.) Ab initio calculations indicate that the cis and trans isomers of ONOOH are nearly isoenergetic (16). Thus, the driving force for reaction 1 is nearly the same for both isomers, and there is no thermodynamic basis for one being substantially more reactive than the other. Nonetheless, reaction 1 proceeds by the indirect pathway with maximally ∼35% of the ONOOH being capable of oxidizing Fe(CN)64- (6). One-electron outer-sphere oxidations of coordination compounds by the CO2 adduct (ONOOCO2-) behave * To whom correspondence should be addressed at Washington State University. † Brookhaven National Laboratory.
similarly (17). These reactions can be described by the general reaction scheme: ONOO– + CO2
"X"
"Y" + Mn+
•
NO2 + M(n +1)+ + CO32–
(2)
NO3– + CO2
For a series of coordination complexes, the initial reaction step, i.e., ONOOCO2- adduct formation, was found to be rate-limiting. If “X” and “Y” are stable geometrical isomers of ONOOCO2-, analogous to the cis and trans isomers of ONOOH, they should have comparable thermodynamic driving forces and electron-transfer reorganization energies, i.e., similar reactivities toward common reductants (13). However, although“Y” was able to oxidize strong oxidants such as Ru(bpy)32+ (E° ) 1.26 V), implying its reduction potential is g1.1 V, “X” was unable to oxidize even Fe(CN)64- (17). An alternative explanation for the reactivity difference, namely, that “Y” is a thermally activated high-energy intermediate with a correspondingly greater driving force, still leaves unexplained why “X” is unreactive toward relatively weak oxidants such as Fe(CN)64-, for which simple electron transfer is undoubtedly thermodynamically allowed. Most likely, “X” is simply too short-lived to engage in bimolecular reactions. Within this scenario, “X” and “Y” could still be stable geometrical isomers of ONOOCO2-, with “X” being properly oriented to undergo facile intramolecular bond rearrangement and “Y” being sufficiently long-lived to be scavenged by reductants. However, this reaction scheme fails completely when applied to ONOOH, i.e., when CO2 is replaced by H+ in eq 2. In this case, protonation of ONO2- to form “X” and “Y” is not rate-limiting, and the accumulated ONOOH undergoes relatively slow unimolecular decay (1). Any reaction scheme based upon an isomeric two-state model subject to the condition that “X” is short-lived would predict one or more of the following: (i) decay is biphasic, (ii) all of the peroxynitrite can react with oxidants, (iii) decay is very rapid. The reaction exhibits none of these features; consequently, it becomes difficult to write a chemically plausible mechanism based upon reactive and unreactive isomeric forms of the oxidant. Radical Mechanisms. The dynamical properties of these oxidants can be rationalized by assuming that the reactions are initiated by O-O bond homolysis. In this case, “X” are the geminate pairs {•NO2,•CO3-} for ONOOCO2- and {•NO2,•OH} for ONOOH and “Y” are the corresponding separated radicals. Rearrangement to NO3- can occur both within the cage or by recombination of the free radicals, but the cage is too short-lived to react directly with reductants. Maximal product yields are limited by the cage escape yields of the geminate radical pairs and should be independent of the identity of the reacting partner, as has been confirmed for reactions with
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Forum: Reactive Species of Peroxynitrite
Chem. Res. Toxicol., Vol. 11, No. 7, 1998 715 Scheme 1
Figure 1. Comparison of measured O2 yields (0) and results of simulations based upon rate constants for reactions of •OH and secondary radicals (b).
several one-electron reductants (15, 17). Implicit in this mechanism is the notion that bimolecular redox reactions of all isomeric forms of the precursor peroxides are slower than bond homolysis. Why should this be? As previously noted, most of the driving force for the overall reaction is obtained from protonation of the immediate product ions, OH- or CO32-. Furthermore, the self-exchange rate constant for •NO2 + NO2- is low (kex ≈ 0.1 M-1 s-1), indicating the presence of a large nuclear reorganizational barrier to electron-transfer (18). Both factors probably contribute to a high activation barrier for crossreactions between the peroxynitrites and one-electron reductants (13). Consequently, if alternative low-energy transition states exist, reaction will proceed by other pathways. The O-O bonds for both ONOOH [bond dissociation energy (BDE) ) 22 kcal/mol] and ONOOCO2- (BDE ) 9 kcal/mol) are exceptionally weak (16); for comparison, the BDE of HOOH is 51 kcal/mol. The only low-energy pathways for unimolecular isomerization to nitrate that were revealed by ab initio calculations (16) involved cleavage of the O-O bond; a configurational isomer that might represent a thermally accessible “transoid” intermediate could not be found. The thermodynamics of homolysis has also been reevaluated, with one group concluding that cleavage of the ONOOH O-O bond to form •NO2 and •OH could occur at a sufficient rate for these radicals to be the reactive intermediates (14) and another concluding that bond homolysis is (marginally) too slow to account for the observed reactivity of ONOOH (19). There appears to be no disagreement that formation of •NO2 and •CO3- from ONOOCO2- is a plausible unimolecular activation step (2, 6, 14-17). Is there physical evidence for •OH radical participation in reactions of ONOOH? In addition to NO3-, O2 and NO2- are formed in 1:2 proportion during decomposition of ONOOH in neutral to alkaline solutions (20). We have extended these observations1 to show that (i) O2 yields are diminished in the presence of various radical scavengers, (ii) excess NO2- reverses this inhibition at pH 9 but is itself inhibitory under more acidic conditions, and (iii) no O2 is formed upon decomposition of ONOOCO2-. The attenuation of O2 yields by the radical scavengers has provided the basis for competition kinetic analyses. Specifically, if free •OH is “Y”, a linear relationship should exist between the function, F/(1 - F) × [ONOO-]/[S], and the rate constant (ks) for reaction between •OH and the scavenger (S) (5, 21); F is defined as the fractional inhibition of O2 yield by the scavenger. The predicted relationship has been confirmed using eight common radical scavengers. Moreover, the slope of the line must
be the reciprocal of the rate constant for bimolecular reaction between •OH and ONOO-; from the data we obtained a slope of 4 × 109 M-1 s-1, which is nearly identical to the rate constant of 5 × 109 M-1 s-1 measured directly by pulse radiolysis (22). Thus, the relative reactivity of “Y” toward these compounds is the same as the relative reactivity exhibited by •OH. A more rigorous test of the mechanism is being made2 by simulating the oxygen yields assuming that free radicals are the reactant species. The individual reactions being considered are all of those known to be important under the reaction conditions (6, 23-25),3 included in Scheme 1 and eqs 3 and 4:
N2O3 + ONOO- f NO2- + 2•NO2
(3)
2•NO2 h N2O4 + H2O f NO2- + NO3- + 2H+ (4) A comparison to the experimental results is given in Figure 1.4 The close correspondence of experimental and calculated yields is remarkable because the simulation is based upon known rate constants for each of the elementary reaction steps and is therefore independent of the experimental data. Within this scheme, diminution of O2 yields by radical scavengers is attributed to their competitive reaction with •OH. This effect is reversed in alkaline solution by reaction of NO2- with •OH, yielding •NO (step 5), followed by •NO oxidation 2 2 of •O2- (step 4) which, in turn, is formed by ONO2homolysis (step 2). Although the pH dependence of the yield (Figure 1) has the appearance of a titration curve, this does not define the pKa of any particular species involved in the reaction but arises from changes in the relative rates of steps 1 and 2. Thus, all experimental data, including the effects of radical scavengers, are accurately duplicated by assuming that the intermediary oxidant has a chemical reactivity identical to that of radiolytically generated •OH. Debates concerning whether this species is actually “free” •OH or exists as a weakly H-bonded radical pair (7, 16) become moot at this point.
Acknowledgment. We are indebted to Drs. J. W. Coddington (Washington State University), G. Merenyi (The Royal Institute of Technology, Stockholm), and G. Czapski (The Hebrew University of Jerusalem) for communicating experimental results prior to their publication. This research is supported by grants from the National Institute of Allergy and Infectious Diseases (to J.K.H., No. AI15834) and the U.S. Department of Energy (to S.V.L., No. EMSP CH37SP23, and to Brookhaven National Laboratory, No. DE-AC02-98CH10886). Supporting Information Available: Cited references (3 pages). Ordering information can be found on any current masthead page. TX980076X 1J.
W. Coddington, unpublished results. V. Lymar, unpublished results. 3G. Merenyi and G. Czapski, personal communications. 4The data fit equally well to an alternative mechanism in which the radical product of step 3 in the scheme is •NO, rather than •O2-. 2S.
