TY - JOUR
T1 - Gibbs energy of formation of peroxynitrite - Order restored
AU - Goldstein, S.
AU - Czapski, G.
AU - Lind, J.
AU - Merényi, G.
PY - 2001
Y1 - 2001
N2 - In a recent publication [Nauser et al. (2001) Chem. Res. Toxicol. 14, 248-350], the authors estimated a value of 14 ± 3 kcal/mol for the standard Gibbs energy of formation of ONOO- and argued that the experimental value of 16.6 kcal/mol [Merényi, G., and Lind, J. (1998) Chem. Res. Toxicol. 11, 243-246] is in error. The lower value would suggest that the yield of free radicals during decomposition of ONOOH into nitrate is negligibly low, i.e., less than 0.5%, though within the large error limit given, the radical yield might vary between 0.003% and ca. 80%. The experimental value of 16.6 ± 0.4 kcal/mol was based on the determination of the rate constant of the forward reaction in the equilibrium ONOO- ⇋ ·NO and O2·- by use of C(NO2)4, an efficient scavenger of O2·- which yields C(NO2)3-. Nauser et al. reported that addition of ·NO has no significant effect on the rate of formation of C(NO2)3-, and therefore the formation of C(NO2)3- is due to a process other then reduction of C(NO2)4 by O2·-. In addition, they argued that Cu(II) nitrilotriacetate enhances the rate of peroxynitrite decomposition at pH 9.3 without reduction of Cu(II). In the present paper, we show that the formation of C(NO2)3- due to the presence peroxynitrite is completely blocked upon addition of ·NO. Furthermore, the acceleration of the rate of peroxynitrite decomposition at pH 9 in the presence of catalytic concentrations of SOD ([ONOO-]/[SOD] > 30) results in the same rate constant as that obtained in the presence of C(NO2)4. These results can only be rationalized by assuming that ONOO- homolyses into ·NO and O2·- with k = 0.02 s-1 at 25 °C. Thus, the critical experiments suggested by Nauser et al. fully support the currently accepted thermodynamics as well as the mode of decomposition of the ONOOH/ONOO- system.
AB - In a recent publication [Nauser et al. (2001) Chem. Res. Toxicol. 14, 248-350], the authors estimated a value of 14 ± 3 kcal/mol for the standard Gibbs energy of formation of ONOO- and argued that the experimental value of 16.6 kcal/mol [Merényi, G., and Lind, J. (1998) Chem. Res. Toxicol. 11, 243-246] is in error. The lower value would suggest that the yield of free radicals during decomposition of ONOOH into nitrate is negligibly low, i.e., less than 0.5%, though within the large error limit given, the radical yield might vary between 0.003% and ca. 80%. The experimental value of 16.6 ± 0.4 kcal/mol was based on the determination of the rate constant of the forward reaction in the equilibrium ONOO- ⇋ ·NO and O2·- by use of C(NO2)4, an efficient scavenger of O2·- which yields C(NO2)3-. Nauser et al. reported that addition of ·NO has no significant effect on the rate of formation of C(NO2)3-, and therefore the formation of C(NO2)3- is due to a process other then reduction of C(NO2)4 by O2·-. In addition, they argued that Cu(II) nitrilotriacetate enhances the rate of peroxynitrite decomposition at pH 9.3 without reduction of Cu(II). In the present paper, we show that the formation of C(NO2)3- due to the presence peroxynitrite is completely blocked upon addition of ·NO. Furthermore, the acceleration of the rate of peroxynitrite decomposition at pH 9 in the presence of catalytic concentrations of SOD ([ONOO-]/[SOD] > 30) results in the same rate constant as that obtained in the presence of C(NO2)4. These results can only be rationalized by assuming that ONOO- homolyses into ·NO and O2·- with k = 0.02 s-1 at 25 °C. Thus, the critical experiments suggested by Nauser et al. fully support the currently accepted thermodynamics as well as the mode of decomposition of the ONOOH/ONOO- system.
UR - http://www.scopus.com/inward/record.url?scp=0034972048&partnerID=8YFLogxK
U2 - 10.1021/tx010066n
DO - 10.1021/tx010066n
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AN - SCOPUS:0034972048
SN - 0893-228X
VL - 14
SP - 657
EP - 660
JO - Chemical Research in Toxicology
JF - Chemical Research in Toxicology
IS - 6
ER -