Chem. Res. Toxicol. 1998, 11, 87-90
87
Communications Can OdNOOH Undergo Homolysis? Willem H. Koppenol* and Reinhard Kissner Laboratorium fu¨ r Anorganische Chemie, Eidgeno¨ ssische Technische Hochschule-Zu¨ rich, Universita¨ tstrasse 6, CH-8092 Zu¨ rich, Switzerland Received October 31, 1997
Recent thermodynamic calculations of Mere´nyi and Lind [(1997) Chem. Res. Toxicol. 10, 1216-1220] suggest that OdNOOH can undergo homolysis to form the hydroxyl radical and nitrogen dioxide. This result is based in part on our statement that the enthalpy of ionization of OdNOOH is close to zero [Koppenol et al. (1992) Chem. Res. Toxicol. 5, 834-842]. As the ionization of OdNOOH is sensitive to the milieu and the rate of isomerization (to nitrate) to the total concentration of OdNOOH and OdNOO- [Kissner et al. (1997) Chem. Res. Toxicol. 10, 1285-1292], we reinvestigated the temperature dependence of the ionization constant and determined a ∆H° of 4 ( 2 kcal mol-1. This results in a standard Gibbs energy of homolysis of 16 kcal mol-1 and a rate of homolysis of 1 × 10-2 s-1. Given the uncertainty in the Gibbs energy of homolysis, upper and lower rates are 1 × 10-4 and 0.6 s-1, slower than the rate of isomerization, 1.2 s-1 at 25 °C. The recombination of the homolysis products NO2• and HO• is known to lead to mainly peroxynitrous acid. If one assumes that a few percent of the recombinations lead to nitrate instead, then the rate of homolysis must be much higher than the rate of isomerization. We conclude therefore that homolysis is unlikely.
Introduction To determine the Gibbs energy of formation of peroxynitrite [oxoperoxonitrate(1-)] one needs an estimate for the entropy; the enthalpy is known, -10 ( 1 kcal mol-1 (1, 2). In 1992 we published an estimate of 45 eu for S°(OdNOO-aq) (3). That value was based on the published entropy for nitrate, 35 eu (4), to which the difference in entropy between the gaseous species NO3• and OdNOO•, 10 eu (5),was added. A consequence of that estimate was a standard Gibbs energy of formation of 10 kcal mol-1 and the conclusion that homolysis, reaction 1:
OdNOOH f HO• + NO2•
(1)
has a ∆rxnG° of 21 kcal/mol in water. Via the relation ∆rxnG° ) -RT ln k1/k-1, with k-1 ) 4.5 × 109 M-1 s-1 (6), it can be estimated that homolysis takes place at a rate between 10-6 and 10-8 s-1, much lower than the isomerization of OdNOOH to NO3- and H+ where k ) 1.2 s-1. For this reason we excluded formation of the hydroxyl radical as an intermediate in the isomerization reaction (3). Our estimate of the absolute entropy was recently criticized by Mere´nyi and Lind (7) who argued that our approach implied that the entropies of solution for nitrate and oxoperoxonitrate(1-) are the same. Because nitrate and oxoperoxonitrate(1-) have different shapes, the * To whom correspondence should be addressed. Tel: 41-1-632-2875 or -2852 (Ms. R. Pfister, secretary). Fax: 41-1-632-1090. E-mail:
[email protected].
entropies of solution are indeed not necessarily identical. Mere´nyi and Lind (7) used a cycle as depicted in Scheme 1 to derive an entropy (S°) for the oxoperoxonitrate anion of only 15 eu. This value implies an entropy of hydration of -53 eu, which seems much too large in comparison to entropies of solution of other anions. In Table 1 we show that entropies of solutions of halides decrease with increasing atomic weight and that these entropies can be approximated by the difference in entropy of the neutral halogen atom in the gas phase and that of the corresponding ion in solution plus 3 eu. Thus, ∆solS°(F-) ) S°aq(F-) - S°g(F-) ) -37 eu, and the approximation yields the same result: ∆solS°(F-) ≈ S°aq(F-) - S°g(F•) + 3 eu ) -37 eu. Applied to the ions nitrite and nitrate, correction by 1 eu instead of 3 eu gives a better result: these approximations yield entropies of solution of -27 and -24 eu, identical to the values found in the literature (8). We therefore settle for a correction by 2 eu. In addition, as argued by Mere´nyi and Lind (7), it is reasonable to assume that the entropy of HX is 2 eu higher than that of X•. Entropies of solution for the halides listed in Table 1 appear to decrease with increasing atomic weight. Entropies of solutions of multiatomic ions vary between -47 and -21 eu (see Table 2, again with a trend of decreasing entropies of solution versus atomic weight. Exceptions appear to be the ions HO2and IO3-, which have exceptionally large entropies of solution. However, neither of these entropies of solution is as large as that proposed by Mere´nyi and Lind (7) for OdNOO-. Thus, to summarize the dispute at this point: Mere´nyi and Lind (7) have a valid criticism of our earlier estimate of 45 eu for S°(OdNOO-); their thermo-
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88 Chem. Res. Toxicol., Vol. 11, No. 2, 1998
Communications
Scheme 1. Estimates of S°(ONOO-aq)a
a Results obtained by Mere ´ nyi and Lind (7) are given in italics; those obtained in the present paper are in bold.
Table 1. Standard Entropies of Solvation (in eu)a XF-
ClBrIa
S°g(X-) S°g(X•) S°aq(X-) ∆hydS° S°aq (X-) - S°g(X•) + 3 eu 35 37 38 40
38 39 42 43
-2 13 20 25
-37 -24 -18 -15
-37 -23 -19 -15
Data taken from the literature (4, 15).
Table 2. Standard Entropies of Solvation (in eu) of Di-, Tri-, and Tetraatomic Speciesa species g
aq
∆hydS°
MW
CNO2•HO2• N3HNCO NO2SCNNO3ClO2• ClO3BrO3IO3-
CNO2•HO2N3NCONO2SCNNO3ClO2ClO3BrO3IO3-
-25 (8) -33 (16) -47 -25 (8) -27 -27 (8) -21 (8) -24 (8) -35 -24 (8) -28 (8) -41 (8)
26 32 33 42 43 46 58 62 67.5 83.5 128 175
a Data were taken from the literature, as indicated, and estimated from the NBS Tables (4) on the assumption that S°g(XYZ-) ≈ S°g(XYZ•) - 2 eu, and S°g(HXYZ) ≈ S°g(XYZ•) + 2 eu.
dynamic approach to estimate a new value is very reasonable, but the result, an entropy of hydration of -53 eu, seems too large.
Methods Mere´nyi and Lind (7) used an enthalpy of ionization of 0 kcal mol-1, based on our earlier published observation that the pKa does not appear to be temperature-dependent (3). However, we discovered recently that the isomerization of oxoperoxonitrate(1-) to nitrate is subject to general acid catalysis and that the rate of isomerization is influenced by the formation of an adduct between oxoperoxonitrate(1-) and its protonated form (9). We therefore redetermined the dependence of the pKa on the temperature at low total oxoperoxonitrate(1-) concentrations in phosphate buffer at a constant ionic strength of 0.1 M across the pH range used. The pH was measured at the outlet of the Applied Photophysics SX 17MV stopped-flow spectrophotometer operated in the symmetric mixing mode. As the pH could not be measured in the mixing cell, it was calculated from the pH measured at room temperature with the aid of tabulated ionization constants for phosphoric acid and dihydrogen phosphate between 0 and 50 °C (10). The values at 55 °C were estimated by extrapolation. The temperature was measured with an external thermocouple mounted near the mixing cell; the temperatures reported by the integrated thermocouple of the stopped-flow instrument were deemed inaccurate: too low
Figure 1. Rate constants for the decay of hydrogen oxoperoxonitrate as a function of temperature and pH. All points represent averages of nine determinations. Rate constants determined at the same temperature have been connected with straight lines for clarity: a, 55 °C; b, 45 °C; c, 35 °C; d, 25 °C; e, 15 °C; f, 5 °C. The ionic strength from the phosphate buffer was kept constant at 0.1 M across the pH range. at high temperatures and too high at low temperatures.
Results The rate constants measured as a function of pH and temperature are shown in Figure 1. The points in this figure represent the average of nine determinations each; those rate constants determined at a particular temperature have been connected for clarity. From such a set of rate constants a pKa was calculated as previously described (9). The pKa’s and their errors (2σ) are shown in Figure 2. The relatively large error in the Ka determined at 55 °C is caused by the fact that the peroxynitrite was already partially decayed before mixing. At 5 °C the reaction proceeds very slowly, such that rates of isomerization could not be determined with the desired accuracy. The enthalpy of ionization, 4.1 kcal mol-1, and its error of 1.6 kcal mol-1 are determined by the steepest and shallowest slopes of the lines that can be drawnthrough the points with their 2σ error bars in a plot of ln Ka vs T-1, as shown in Figure 2. The pKa of hydrogen oxoperoxonitrate at 25 °C and 0.1 M phosphate buffer is 6.8; at lower phosphate concentrations the pKa approaches 6.5 (9). Based on the value of 6.5 the standard Gibbs energy of ionization is 9.1 kcal mol-1. With a standard enthalpy of ionization of +4.1 kcal mol-1, the standard entropy of ionization is -16 eu, which results in an entropy of solution of -37 eu and a S°(OdNOO-aq) of 31 eu, approximately the same as that of nitrate, 35 eu. Compared to other entropies of solution (Table 2), this value is high but within the realm of reasonable values. From a S°(OdNOO-aq) of 31 eu, we calculate a standard Gibbs energy of formation of 14.2 kcal mol-1. The uncertainty in the standard Gibbs energy originates from the error in the enthalpy of formation, 1 kcal mol-1 (2), and that in the T∆fS° term, 2.2 kcal mol-1 (7). As these errors are unrelated,
Communications
Chem. Res. Toxicol., Vol. 11, No. 2, 1998 89
Discussion
Figure 2. Dependence of Ka on temperature. The ln Ka values and their 2σ errors are shown. The steepest and shallowest slopes of the lines that can be drawn through these points and their errors determine the enthalpy of ionization, 4.1 kcal mol-1, and its error of 1.6 kcal mol-1. Table 3. Thermodynamic Data Related to OdNOO- and OdNOOH Formation OdNOO-aq OdNOOHaq OdNOOHg
∆fH° (kcal mol-1)
∆fG° (kcal mol-1)
S° (eu)
-10 ( 1 -14 ( 2 1(3
14 ( 3 5(3 13 ( 3a
31 ( 7 47 ( 5 72 ( 2
Ionization ∆ionH°
∆ionG° (kcal mol-1)
∆ionS° (eu)
4(2
9.1 ( 0.2
-16 ( 5
Isomerization ∆isoH° (kcal mol-1)
∆isoG° (kcal mol-1)
∆isoS° (eu)
-39 ( 1
-40 ( 3
3(8
Homolysis ∆homG° (kcal mol-1) ONOOHaq f NO2•aq + HO•aq ONOOHg f NO2•g + HO•g ONOO-aq f NO•aq + O2•-aq
16 7.2 18
OdNOO- + OdNOOH f adduct
Reduction Potentials E°(ONOOH, H+/NO2•, H2O) E°′(ONOOH, H+/NO2•, H2O) E°′(ONOOH, H+/NO2-, H2O)
From the new thermodynamic quantities, we calculate a new standard Gibbs energy of homolysis (reaction 1) of 16.0 ( 2.4 kcal mol-1 and a new rate constant of homolysis of 1 × 10-2 s-1, with, due to the 2.4 kcal mol-1 error, limits of 0.6 and 1 × 10-4 s-1. Given a rate constant of the isomerization of 1.2 s-1 at 25 °C (9), it would appear that homolysis is unlikely, but it cannot be completely excluded. However, pulse radiolysis studies (6, 11) do not mention that the hydroxyl radical reacts with nitrogen dioxide to form nitrate; hydrogen oxoperoxonitrate is the only or the main product. If, for the sake of argument, we assume that between 0% and 10% of the hydroxyl and nitrogen dioxide radical react to form nitrate, then the rate of homolysis has to be at least a factor of 10 larger than the rate of isomerization. Such a rate is not compatible with our estimation of the rate of homolysis. In addition, the isomerization reaction has both a small entropy of activation of 3 ( 2 eu (3) and a small volume of activation of 1.7 ( 1.0 cm3 mol-1 (9); these values are not compatible with homolysis. Other published values for the entropy (12) and volume of activation (13) are each based on only two data points, rendering them far less reliable than the values we reported. Mere´nyi and Lind (7) advocated the use of Occam’s razor in deciding whether the hydroxyl radical could be formed. Using this principle, we conclude that formation of the hydroxyl radical from hydrogen oxoperoxonitrate is very unlikely; it requires the use of Occam’s broom (14) to come to a different interpretation. If one assumes that the standard Gibbs energy of hydration of OdNOOH is similar to that of HOOH, then one can also calculate the standard Gibbs energy of homolysis in the gas phase (see Table 3). The value is 7 kcal mol-1, much lower than in water. Consequently, formation of hydroxyl radicals as an intermediate is quite feasible in the gas phase. Hydroxyl radical formation in water is thermodynamically possible via reactions 2 and 3, and we have experimental evidence to support the formation of an adduct at pH values near the pKa and higher and at combined concentrations of oxoperoxonitrate(1-) and hydrogen oxoperoxonitrate that exceed 0.1 mM (9). Adduct formation should be taken into account when bolus additions of oxoperoxonitrate(1-) are made near neutral pH.
V
pH
2.0 ( 0.1 1.6 ( 0.1 1.3 ( 0.1
0 7 7
a Based on the assumption that the standard Gibbs energy of solution of hydrogen oxoperoxonitrate is similar to that of hydrogen peroxide (16).
quadratic addition results in an error in the standard Gibbs energy of formation of 2.4 kcal mol-1. We therefore report a standard Gibbs energy of formation of oxoperoxonitrate(1-) of 14 ( 2 kcal mol-1. Given the pKa of 6.5 (9), the standard Gibbs energy of formation of OdNOOH is 5 ( 2 kcal mol-1. Derived thermodynamic properties, such as reduction potentials, are given in Table 3.
(2)
∆rxnG°′ ) -5 kcal mol-1 adduct f NO3- + HO• + NO2•
(3)
∆rxnG° ′ ) -24 kcal mol-1 Whether reaction 3 takes placesparallel to the decay to nitrite and dioxygensneeds to be investigated.
Acknowledgment. We thank Dr. G. Mere´nyi for frank discussions and Dr. P. L. Bounds for linguistic advice. These studies were supported by the Swiss Nationalfonds and ETH.
References (1) 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.
90 Chem. Res. Toxicol., Vol. 11, No. 2, 1998 (2) Manuszak, M., and Koppenol, W. H. (1996) The enthalpy of isomerization of peroxynitrite to nitrate. Thermochim. Acta 273, 11-15. (3) 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. (4) 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. (5) Benson, S. W. (1976) Thermochemical Kinetics, p 293, John Wiley & Sons, New York. (6) Logager, T., and Sehested, K. (1993) Formation and decay of peroxynitrous acid: A pulse radiolysis study. J. Phys. Chem. 97, 6664-6669. (7) Mere´nyi, G., and Lind, J. (1997) Thermodynamics of peroxynitrite and its CO2 adduct. Chem. Res. Toxicol. 10, 1216-1220. (8) Marcus, Y. (1986) The hydration entropies of ions and their effects on the structure of water. J. Chem. Soc., Faraday Trans. 1 82, 233-242. (9) 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.
Communications (10) Handbook of Chemistry and Physics, 62nd ed. (1981) p D-145, CRC Press, Boca Raton, FL. (11) Gra¨tzel, M., Henglein, A., and Taniguchi, S. (1970) Pulsradiolytische Beobachtungen u¨ber die Reduktion des NO3--Ions und u¨ber die Bildung und Zerfall der persalpetrigen Sa¨ure in wa¨ssriger Lo¨sung. (Pulse radiolysis observations regarding the reduction of the NO3- ion and the formation and decay of pernitric acid in aqueous solution.) Ber. Bunsen-Ges. physik. Chem. 94, 292298. (12) Benton, D. J., and Moore, P. (1970) Kinetics and mechanism of the formation and decay of peroxynitrous acid in perchloric acid solutions. J. Chem. Soc. A 3179-3182. (13) Goldstein, S., Meyerstein, D., van Eldik, R., and Czapski, G. (1997) Spontaneous reactions and reduction by iodide of peroxynitrite and peroxynitrate: Mechanistic insight from activation parameters. J. Phys. Chem. A 101, 7114-7118. (14) Brenner, S. (1997) Loose Ends. Curr. Biol. (15) Moore, W. J. (1972) Physical Chemistry, p 538, Prentice Hall, Inc., Englewood Cliffs, NJ. (16) Koppenol, W. H. (1989) Generation and thermodynamic properties of oxyradicals. In Focus on Membrane Lipid Oxidation (VigoPelfrey, C., Ed.) Vol. I, pp 1-13, CRC Press, Boca Raton, FL.
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