1216
Chem. Res. Toxicol. 1997, 10, 1216-1220
Communications Thermodynamics of Peroxynitrite and Its CO2 Adduct Ga´bor Mere´nyi* and Johan Lind Department of Chemistry, Nuclear Chemistry, The Royal Institute of Technology, S-10044, Stockholm 70, Sweden Received June 11, 1997X
The equilibrium constant, K3, of aqueous homolysis of peroxynitrous acid into hydroxyl and nitrogen dioxide free radicals was estimated to be 5 × 10-10 M. This value was derived from a thermodynamic cycle by use of the experimentally known ∆fH°(ONOO-,aq) ) -10.8 kcal/ mol and the enthalpy of ionic dissociation of ONOOH(aq), ∆H°1 ) 0 kcal/mol, as well as of the entropy of gaseous ONOOH, S°(ONOOH,g) ) 72 eu. Furthermore we assumed the entropy of hydration of ONOOH, ∆S°2, to be -25 eu, a value closely bracketed by the hydration entropies of analogous substances. The rate constant of radical recombination of OH• with NO2• to yield ONOOH, k-3, was resimulated from experimental data and found to be ca. 5 × 109 M-1 s-1. Together with the estimated K3, this yields the homolysis rate constant k3 ) 2.5 s-1. This value is close to 0.5 s-1, the rate constant of formation of a reactive intermediate during the isomerization of peroxynitrous acid to nitrate. Our thermodynamic estimate is therefore consistent with substantial amounts of OH• and NO2• free radicals being formed in this process. The thermodynamic implications for the carbon dioxide/peroxynitrite system are also discussed.
Introduction Having led a largely anonymous existence on the outer fringes of inorganic chemistry (1), peroxynitrite (ONOO-) and its conjugate acid, peroxynitrous acid (ONOOH), have entered the mainstream of biology during the last decade. This dramatic change in status is contingent on the realization that peroxynitrite is efficiently produced in living matter (2) through coupling of NO• with O2•and on the surmise that it may have deleterious biological effects (2, 3). During isomerization to nitrate, peroxynitrous acid forms a highly reactive intermediate (4, 5) capable of oxidizing a host of compounds. The yield of this intermediate was found (6, 7) to be ca. 40%, and in the absence of suitable substrates, it ends up forming predominantly nitrate. Initially, this species was believed (2, 5) to be the hydroxyl radical, coproduced with NO2• during homolysis of ONOOH in water. At present, whether or not hydroxyl radicals are formed is a subject for lively controversy. While results from several spintrap studies (8-10) strongly implicate OH• radicals, other spin-trap investigations (11-13) provide little evidence for OH•. Support for hydroxyl radical formation also comes from selectivity studies (14, 15) during aromatic hydroxylation by peroxynitrite. Several experiments with well-known OH• radical scavengers appear to preclude OH• radical formation on the grounds that the scavenger has no (16) or too little (7, 17-19) effect on the product yield. Again, there do exist OH• scavenger studies (2, 5, 20, 21) that fully bear out the production of OH•. On the other hand, the low sensitivity of the rate of peroxynitrite decomposition to variation of solvent viscosity (22) does not support radical formation being an important process. The present state of the art in peroxynitrite research has been thoroughly reviewed in X
Abstract published in Advance ACS Abstracts, November 1, 1997.
S0893-228x(97)00101-X CCC: $14.00
refs 23 and 24. While the scientific community is far from a consensus, the picture currently favored for the reactive intermediate is an unspecified high-energy isomer of ONOOH (23). In another type of approach, Koppenol et al. (25) treated the ONOOH problem thermodynamically. Not possessing all experimental data, they made certain assumptions, notably about the entropy of hydration of ONOO-, assuming the latter to be the same as for the nitrate anion. They did not indicate how accurate they deemed this estimate to be. We strongly question the validity of their assumption, given that the two anions are structurally very different, one having the negative charge symmetrically delocalized, while the other has it localized on the terminal oxygen atom. It should further be noted that estimating hydration entropies for ions is a difficult and by no means straightforward matter, as can be gleaned from ref 26. By use of the above estimation, Koppenol et al. (25) predicted homolysis of ONOOH in water to be an extremely unlikely process, contributing merely to the extent of 10-6-10-4% to the overall isomerization of ONOOH to NO3- + H+. Thus, if left unchallenged, this work would seem to deal the coup de grace to the OH• radical hypothesis. Our work will present an alternative thermodynamical approach, where we shall utilize a thermochemical cycle and estimate the entropy of hydration of neutral ONOOH. Arguably, the latter estimation is much safer than that for ions. We shall show that our predictions are fully consistent with the homolysis of ONOOH. Recently, the presence of bicarbonate in ONOO- solutions was reported (27) to alter the oxidative properties of the latter, and a tentative mechanism was suggested for this effect. In a subsequent publication (28) it was shown that CO2, rather than bicarbonate, is the active species. The presence of CO2 dramatically enhances the rate (28, 29) of isomerization of ONOO-, itself fairly © 1997 American Chemical Society
Communications
Chem. Res. Toxicol., Vol. 10, No. 11, 1997 1217 Scheme 1
Table 1. Entropies of Hydration, ∆S°hyd, of Small and Medium Size Neutral Speciesa
Scheme 2
species
∆S°hyd
ref
CH4 C2H6 CH3OH CH3NH2 HC(O)OH CH3C(O)OH NO2• HNO2
-24 -26 -25 -28 -21 -24 -25 -24
47 47 47 47 47 47 b b
a The entropies are given in eu (cal/mol‚K) units. b The entropies were calculated from data in ref 48.
Table 2. Entropies of Hydration, ∆S°hyd, of Some Hydroperoxides and a Peracida
sluggish. The mechanism of this reaction has been thoroughly investigated in ref 28, where possible intermediates were suggested, and a thermochemical appraisal was also made. The addition product of ONOOto CO2, presumably ONOOC(O)O-, produces a reactive intermediate, with a yield of ca. 40% (30). Interestingly, this yield is about the same as the one found with ONOOH. The CO2/ONOO- system brings about more efficient nitration of tyrosine (31) than does peroxynitrite, but in contrast to the latter, it appears unable to convert methionine to methionine oxide (32). On the basis of its oxidizing capability [E0 above 1.3 V (30)], the intermediate is believed to consist of CO3•- and NO2• free radicals, (30) formed during homolysis of ONOOC(O)O-. As the thermochemistry of the latter is intimately connected with that of ONOO-, we shall also estimate the properties of this adduct, particularly in relation to homolysis.
Results and Discussion The ONOOH/ONOO- System. The thermochemical steps taken in this work are present in Schemes 1 and 2. All values in the sequel will refer to room temperature, i.e., 298.15 K. Just as in ref 25, the starting point for the chemical cycle is the experimental heat of isomerization of ONOO-(aq) into NO3-(aq), yielding ∆fH°(ONOO-,aq) ) -10.8 kcal/mol (33). This value has been confirmed in a subsequent publication (34). Another important experimental parameter is ∆H°1, i.e., the enthalpy of ionization of ONOOH(aq), which is close to 0 (25, 35).
ONOOH(aq) H ONOO-(aq) + H+
(1)
We thus obtain ∆fH°(ONOOH,aq) ) ∆fH°(ONOO-,aq) - ∆H°1 ) -10.8 - 0 ) -10.8 kcal/mol. According to Benson (36) the entropy of the ONOO• radical in the gas phase is 70 eu. The same author (37) finds the entropy of CH3OOH higher than that of CH3OO• by 2 eu. Therefore, we set S°(ONOOH,g) ) 72 eu. Next we address the entropy of hydration of ONOOH, i.e., ∆S°2.
ONOOH(g) H ONOOH(aq)
(2)
First, we note that the entropies of hydration of small and medium size neutral species are remarkably similar, having values between -20 and -30 eu. A representative selection is presented in Table 1. In particular, the entropies of hydration of hydroperoxides are close to -25 eu, as transpires from Table 2. We thus set ∆S°2 ) -25 eu, the average of the values in Table 2. Combining the above figures, we obtain S°(ONOOH,aq) ) S°(ONOOH,g) + ∆S°2 ) 72 - 25 ) 47 eu. This
species
∆S°hyd
ref
H2O2 CH3OOH CH3CH2OOH CH3C(O)OOH
-27 -24 -28 -22
b b b b
a Entropy units are given in eu (cal/mol‚K). b The values were calculated from data in ref 49.
yields ∆fS°(ONOOH,aq) ) -65 eu, from whence ∆fG°(ONOOH,aq) ) 8.6 kcal/mol is calculated. At this point it is worthwhile discussing the error limits inherent in this estimation. First, the experimental accuracy of ∆fH°(ONOO-,aq) is given as (1.0 kcal/mol (34). Judging by the data in ref 35, the error margin in ∆H°1 should also be (1.0 kcal/mol. Finally, the uncertainty in the entropy estimations in the present work is thought to be (5 eu. Taken together this makes for a combined uncertainty of (3.5 kcal/mol in the Gibbs free energy of formation of ONOOH(aq), i.e., ∆fG°(ONOOH,aq) ) 8.6 ( 3.5 kcal/mol. In combination with ∆fG°(NO2•,aq) ) 15.06 kcal/mol (38) and ∆fG°(OH•,aq) ) 6.2 kcal/mol (38), both probably being accurate within (0.1 kcal/mol, we obtain ∆G°3, the Gibbs free energy change for the homolysis reaction 3, to be 12.7 ( 3.7 kcal/mol. This yields the equilibrium constant of homolysis, K3, to be 5 × 10-10 M, or log K3 ) -9.3 ( 2.7.
ONOOH(aq) H NO2•(aq) + OH•(aq)
(3)
By means of this estimated equilibrium constant, we can now probe the issue of the probability of OH• radical formation during isomerization of ONOOH(aq). In order to clarify the connection between the thermodynamic and kinetic parameters, including a radical pair cage, we present Scheme 3, which is similar to the radical cage model suggested in ref 23. From their experimental data in ref 39 and assuming that the reaction of OH• with NO2• only produces ONOOH, the authors simulated k-3 to be (4.5 ( 1.0) × 109 M-1 s-1. We resimulated these data using the model of Scheme 3. It turns out that k-3 is very insensitive to the assumed nitrate yield in the NO2• + OH• reaction. We obtained a somewhat better fit to the data by assuming kN ) k-cag, i.e., the ratio [NO3-]/ [ONOOH] ) 1, an assumption we feel to be very reasonable. The overall second-order rate constant between OH• and NO2•, ksec ) k-diff(k-cag + kN)/(k-cag + kN + kdiff), comes out as 1 × 1010 M-1 s-1 and k-3 ) kseck-cag/(k-cag + kN) ) ksec/2 ) 5 × 109 M-1 s-1 is obtained. Clearly, while the simulation rather accurately predicts k-3, it is not sufficiently sensitive to reveal how much NO3- forms (or whether it forms at all) in the NO2• + OH• reaction. We note, however, that a recent theoretical calculation finds
1218 Chem. Res. Toxicol., Vol. 10, No. 11, 1997
Communications
Scheme 3
Table 3. Calculated Two- and One-Electron Reduction Potentials of Peroxynitrous Acid, ONOOHa reaction 2H+
2e-
ONOOH + + h HNO2 + H2O ONOOH + H+ + e- h NO2• + H2O
E°
E°7
1.68 2.14
1.37 1.70
a All potentials are expressed in V and are given vs NHE. E° denotes standard potentials, while E°7 signifies the corresponding potentials at pH 7, allowance being made for pKa(ONOOH) ) 6.8 and pKa(HNO2) ) 3.2, respectively.
formation of HNO3 and ONOOH equally facile, i.e., no activation energy, in the gaseous recombination of NO2• with OH• (40). Finally, we calculate k3 ) K3 × k-3 ) 2.5 s-1. Now, the rate constant of formation of the reactive intermediate during the isomerization of ONOOH is 0.4 × kdis ) 0.4 × 1.3 ) 0.5 s-1, where kdis is the overall firstorder experimental rate constant of ONOOH disappearance and 0.4 is the fraction yielding the reactive intermediate. Given the good agreement between the experimental rate constant of intermediate formation (0.5 s-1) and the predicted rate constant of ONOOH homolysis (2.5 s-1), we conclude that, rather than disproving it, thermodynamics supports the formation of substantial amounts of OH• and NO2• free radicals during isomerization of ONOOH. From the error margin in our estimation of K3, the upper limit of k3 is as high as 103 s-1, while the lower limit is 5 × 10-3 s-1. Of course, k3 cannot exceed the experimental value of 0.5 s-1. However, not even the most unfavorable case, i.e., when k3 is taken to be the lower limit, 0.005 s-1, prohibits OH• radical formation. Indeed, this value is sufficiently close to 0.5 s-1 and reveals that during the isomerization of ONOOH we can expect no less than 0.4% OH• radicals to be generated. This is in sharp contrast to the estimations of Koppenol et al. (25), according to which no more than 10-4% OH• radicals can be expected to form. In view of the above considerations we shall henceforth use values based on experiment. We shall thus identify kdis ) 1.3 s-1 with kexp in Scheme 3, which implies kdiff/(kdiff + kN) ) 0.4 and k3 ) 0.5 s-1. This will yield an equilibrium constant of homolysis: K3 ) k3/k-3 ) 10-10 M and ∆G°3 ) 13.6 kcal/mol. Then, we calculate the ∆fG°(ONOOH,aq) ) 7.7 kcal/mol. Utilizing 6.8, the pKa of ONOOH(aq), we obtain ∆fG°(ONOO-,aq) ) 17.0 kcal/ mol. Finally, from the above values we derive the reduction potentials presented in Table 3. We note that the two-electron reduction potential of ONOOH is close to those of alkyl hydroperoxides (41) and peracids (41). On the other hand, the corresponding one-electron reduc-
tion potential of ONOOH exceeds those of alkyl hydroperoxides by ca. 0.95 V (42). It is also higher by ca. 0.8 V (42) than the one-electron reduction potentials of peracids. Some comments are in order. Our values imply that only one type of reactive intermediate, i.e., OH• and NO2•, forms. Clearly, while the calculations support efficient OH• radical formation, they do not preclude generation of other reactive intermediates as well. Indeed, a lot of experimental findings are consistent with the assumption that both OH• + NO2• and something else may form. However, we are also aware of how very difficult such experiments are, and we feel that several factors may bedevil the interpretation of the primary chemical events. A case in point is the possibility of ONOOH- and/or substrate-mediated chain reactions, initiated by secondary radicals, with the latter being produced from OH• and the substrate. Until all such points have been resolved experimentally, we feel that homolysis according to Scheme 3 remains the simplest and most reasonable mechanism of formation of the reactive intermediate. The ONOO-/CO2 System. The Gibbs free energies of formation of CO2(aq), CO3•-(aq) (43), and NO2•(aq) (38) are known. In combination with ∆fG°(ONOO-,aq) ) 17.0 kcal/mol, they yield K4 ) 0.3 and ∆G°4 ) 0.7 kcal/mol for the overall O•- transfer reaction 4.
ONOO-(aq) + CO2(aq) H NO2• + CO3•-
(4)
While no adduct ONOOC(O)O- has ever been observed, its intermediacy is nevertheless likely. We therefore believe reaction 4 to be composed of two other reactions, 5 and 6.
CO2(aq) + ONOO-(aq) H ONOOC(O)O-
(5)
ONOOC(O)O- H NO2• + CO3•-
(6)
We shall now estimate K5. From ref 44 it transpires that, with carbonyl compounds such as aldehydes, hydroperoxides form adducts ca. 10 times more strongly than does pure water. It is then reasonable to assume that this ratio also holds true in the case of CO2(aq). As K7 ) 1.5 × 10-3 (where the standard state of water is taken at unit mole fraction), we shall assume K8 ) 1.5 × 10-2 M-1.
CO2(aq) + H2O H HOC(O)OH
(7)
CO2(aq) + ONOOH H ONOOC(O)OH
(8)
Now, the acidities of formic acid and carbonic acid are very close, pKa(HOC(O)OH) ) 3.6 being only slightly lower than pKa(HC(O)OH ) 3.75. This comparison suggests that not even the strongly electron-withdrawing ONOO group is likely to depress the pKa by more than ca. 0.5 unit. We thus assume pKa(ONOOC(O)OH) to be ca. 3.0. Then we obtain K5 ) K8 × 10(pKa(ONOOH)-pKa(ONOOC(O)OH) ) 1.5 × 10-2 × 103.8 ) 102 M-1. Together with K4 this yields K6 ) 3 × 10-3 M. Let us consider what these values have to say about the radical model and in particular about the lifetime of ONOOC(O)O-. Since reaction 6 is a homolysis, just like reaction 3, we can employ a similar scheme, namely, Scheme 4, to describe it. Here, kN refers to the irreversible production of NO3- + CO2 in the radical pair cage [NO2•‚‚‚CO3•-]. It is known that CO3•- reacts with NO2•
Communications
Chem. Res. Toxicol., Vol. 10, No. 11, 1997 1219 Scheme 4
oxidizing intermediate, which forms in ca. 40% yield, is an equimolar mixture of CO3•- and NO2• free radicals. According to the present estimations looking for the adduct should be futile, its short lifetime preventing it from accumulating under most experimental conditions.
Acknowledgment. The authors are grateful for financial support given by the Swedish Natural Science Research Council. with an overall second-order rate constant, ksec, of ca. 109 M-1 s-1 and that the product of recombination is NO3and CO2 (45). We recall that data in ref 39 are consistent with similar efficiencies of O-O and O-N coupling (i.e., 50% each, where O-N coupling results in NO3- formation) during the reaction of OH• + NO2•. Expecting a similar O-O coupling efficiency in the [NO2•‚‚‚CO3•-] cage (where N-O coupling implies Otransfer to yield NO3- + CO2), we suggest ONOOC(O)Oto be an important intermediate in the reaction of CO3•with NO2, i.e., k-6 ) 0.5 × 109 ) 5 × 108 M-1 s-1. k6 ) k-6 × K6 ≈ 5 × 108 × 3 × 10-3 ) 1.5 × 106 s-1, while the overall rate constant of adduct destruction is ca. k6/0.4 ) 4 × 106 s-1, where 0.4 is kdiff/(kdiff + kN), the fraction of the free radical formation. The homolysis model thus predicts a lifetime for ONOOC(O)O- on the order of submicroseconds. It should therefore be impossible to accumulate the adduct in reaction 5, where the experimental second-order rate constant (28) k5 is as low as ca. 3 × 104 M-1 s-1 and the calculated k-5 ) 3 × 102 s-1 is far too low to compete with the rate of product formation. Furthermore, this short lifetime also explains why pulse radiolysis experiments (45), where CO3•- and NO2• radical concentrations on the order of 10-5 M were produced, never disclosed a distinct intermediate. Finally, comparison of K3 with K6 reveals that the O-O bond in ONOOC(O)O- is weaker than in ONOOH by ca. 10 kcal/mol. Gratifyingly, a recent theoretical calculation (40) predicts a difference of 13 kcal/mol between the O-O bonds in cis-ONOOH and ONOOOC(O)O-. Although the calculation refers to the gas phase, we feel that the good agreement between the values is significant. Also, a decrease of the O-O bond strength by ca. 10 kcal/mol is well in line with the known fact that O-O bonds get weaker upon replacing H in hydroperoxides by alkyl or acyl groups (46).
Conclusions The present study has shown that with the assumption of a reasonable and justifiable value for the hydration entropy of ONOOH a strong thermodynamic case could be made for the hypothesis that the reactive intermediate that forms with a yield of ca. 40% during the isomerization of peroxynitrous acid is, in fact, an equimolar mixture of hydroxyl and nitrogen dioxide free radicals. While this model has yet to be verified or disproved experimentally, the fact that it makes good thermodynamic sense and by its simplicity well complies with the criterion of Occam’s razor would seem to place the major burden of proof on the shoulders of the model’s opponents. The conclusions about the putative adduct between CO2 and ONOO- in water follow more or less directly from the predicted properties of peroxynitrite and are less prone to controversy. The adduct appears to be extremely labile against homolysis into free CO3•- and NO2• radicals. The predictions are in agreement with experimentally based suggestions to the effect that the strongly
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