Oxidation of peroxynitrite by inorganic radicals: A pulse radiolysis study

Article abstract of DOI:10.1021/ja9735362

Reactivity of the peroxynitrite ion toward a number of inorganic radicals was determined by using the pulse radiolysis technique. The rate constants for the oxidation of the ONOO^- ion by CO₃·^-, ·N₃, and ClO₂· radicals were determined from their decay kinetics to be (7.7 ± 1.2) x 10⁶ (I = 0.6 M), (7.2 ± 0.9) x 10⁸, and (3.2 ± 0.3) x 10⁴ M^-1 s^-(l), respectively. For the ·OH radical, the rate constant of (4.8 ± 0.8) x 10⁹ M^-1 s^-1 was obtained by using competition kinetic analysis. The oxidation potential of the ONOO^- ion was estimated as 0.8 V from the kinetic data. Although thermodynamically favorable, oxidation of ONOO^- by the °NO₂ radical was not observed; an upper limit of 2.5 x 10⁴ M^-1 s^-1 could be set for this reaction. Contribution from some of these reactions to the decomposition of peroxynitrite in the presence and absence of CO₂ is discussed.

Full text of DOI:10.1021/ja9735362

J. Am. Chem. Soc. 1998, 120, 5549-5554
5549
Oxidation of Peroxynitrite by Inorganic Radicals: A Pulse Radiolysis
Study
Sara Goldstein,*,‡ Abhijit Saha,‡ Sergei V. Lymar,§ and Gidon Czapski‡
Contribution from the Department of Physical Chemistry, The Hebrew UniVersity of Jerusalem,
Jerusalem 91904, Israel, and Chemistry Department, BrookhaVen National Laboratory,
Upton, New York 11973
ReceiVed October 8, 1997
Abstract: Reactivity of the peroxynitrite ion toward a number of inorganic radicals was determined by using
the pulse radiolysis technique. The rate constants for the oxidation of the ONOO- ion by CO3•-, N3, and
•
ClO2• radicals were determined from their decay kinetics to be (7.7 ( 1.2) × 106 (I ) 0.6 M), (7.2 ( 0.9) ×
108, and (3.2 ( 0.3) × 104 M-1 s-1, respectively. For the •OH radical, the rate constant of (4.8 ( 0.8) × 109
M-1 s-1 was obtained by using competition kinetic analysis. The oxidation potential of the ONOO- ion was
estimated as 0.8 V from the kinetic data. Although thermodynamically favorable, oxidation of ONOO- by
•
the NO2 radical was not observed; an upper limit of 2.5 × 104 M-1 s-1 could be set for this reaction.
Contribution from some of these reactions to the decomposition of peroxynitrite in the presence and absence
of CO2 is discussed.
•-
Introduction
These intermediates have been suggested to be •NO2 and CO3
radicals, which are strong oxidants formed by the homolytic
cleavage of the peroxo O-O bond of the ONOOCO2- adduct.8-10
At low peroxynitrite concentration and in the absence of
oxidizable compounds, the major pathway for the decay of these
Peroxynitrite (ONOOH/ONOO-) is a powerful oxidant that
can be produced in biological systems from superoxide and nitric
oxide. This reaction is fast (k ) (4.3-6.7) × 109 M-1 s-1 1,2
)
•-
and is expected to be efficient at physiological levels of O2
radicals is their recombination via O- transfer producing NO3
-
•
and NO. The reactivity of peroxynitrite toward biological
molecules3 and its very high toxicity toward cells4 are presently
under intense investigation as potential causes of a number of
debilitating diseases.
and CO2.10,11 However, thermodynamic estimate12 suggests that
the oxidation potential of the ONOO- ion lies below 1 V, NHE.
It is therefore possible that both •NO2 and CO3•- can also oxidize
ONOO-; when the latter is present at high concentration these
reactions may effectively compete with radical recombinations.13
In summary, substantial evidence has been accumulated that
oxidation by both ONOOH and ONOOCO2-, as well as their
decompositions, proceed through formation of strongly oxidizing
intermediates. Whether or not these intermediates can, in turn,
oxidize their precursor, the ONOO- ion, thus complicating
overall kinetics and mechanisms, remains unexplored. In this
study we determine for the first time the rate constants for the
reactions of various inorganic radicals, including CO3•-, •NO2,
Peroxynitrite ion is fairly stable, but its conjugate peroxyni-
trous acid (ONOOH, pKa ) 6.8)3 decomposes rapidly (τ1/2
)
0.53 s at 25 °C); isomerization to nitrate is the major decay
route in acidic media. On its way to NO3-, a significant portion
(∼40%) of ONOOH produces a highly oxidizing intermediate
with the reactivity similar to that of the hydroxyl radical.3,5 It
has been suggested that this intermediate can oxidize the
ONOO- ion and that this reaction may be responsible for
generation of nitrite and oxygen during peroxynitrite decom-
position above pH 5.6 Lymar and Hurst7 have shown that the
peroxynitrite ion reacts very fast with carbon dioxide, apparently
•
and OH, with ONOO- and discuss the possible role of these
reactions in peroxynitrite decomposition.
-
forming the ONOOCO2 adduct. We have recently reported
Experimental Section
that decomposition of this adduct generates reactive intermedi-
ates capable of oxidizing organic and inorganic compounds.8-10
Chemicals. All chemicals were of analytical grade and were used
as received. Sodium chlorite (Fluka) contained about 80% NaClO2,
* To whom all correspondence should be directed. Tel. 972-2-6586478.
Fax: 972-2-6586925. E-mail: SARAG@HUJI.VMS.AC.IL.
‡ The Hebrew University of Jerusalem.
(11) Pryor, W. A.; Lemercier, J.-N.; Zhang, H.; Uppu, R. M. Free Radical
Biol. Med. 1997, 23, 331.
§ Brookhaven National Laboratory.
(12) The gas phase heat of formation of the ONOO• radical was recently
estimated by McKee12a as ∆fH° ) 33 kcal/mol. From this value and an
absolute gas-phase entropy of 68 cal/(mol K)12b for this radical we calculate
(1) Huie, R. E.; Padmaja, S. Free Radical Res. Commun. 1993, 18, 195.
(2) Goldstein, S.; Czapski, G. Free Radical Biol. Med. 1995, 19, 505.
(3) Pryor, W. A.; Squadrito, G. L. Am. J. Physiol. (Lung Cell. Mol.
Physiol.) 1995, 268, L699.
(4) Hurst, J. K.; Lymar, S. V. Chem. Res. Toxicol. 1997, 10, 802.
(5) Goldstein, S.; Squadrito, G. L.; Pryor, W. A.; Czapski, G. Free
Radical Biol. Med. 1996, 21, 965 and references therein.
(6) Pfeiffer, S.; Gorren, A. C. F.; Schmidt, K.; Werner, E. R.; Hansert,
B.; Bohle, D. S.; Mayer, B. J. Biol. Chem. 1997, 272, 3465.
(7) Lymar, S. V.; Hurst, J. K. J. Am. Chem. Soc. 1995, 117, 8867.
(8) Lymar, S. V.; Jiang, Q.; Hurst, J. K. Biochemistry 1996, 35, 7855.
(9) Goldstein, S.; Czapski, G. Inorg. Chem. 1997, 36, 5113.
(10) Lymar, S. V.; Hurst, J. K. Inorg. Chem. 1998, 37, 294.
•
∆fG°(ONOOg ) ) 42 kcal/mol. Assuming that terminal oxygen atoms of
the ONOO• radical can form 4 hydrogen bonds we estimate its solution
free energy as -4 kcal/mol,12c which results in ∆fG°(ONOOaq•) ) 38 kcal/
mol. From this value and the recent estimate by Merenyi and Lind12d for
∆fG°(ONOOaq-) ) 17 kcal/mol, we obtained E°(ONOO•/ONOO-) ) 0.9
V. (a) McKee, M. L. J. Am. Chem. Soc. 1995, 117, 1629. (b) Guillory, W.
A.; Johnston, H. S. J. Chem. Phys. 1965, 42, 2457. (c) Schwarz, H. A.;
Dodson, R. W. J. Phys. Chem. 1984, 88, 3643. (d) Merenyi, G.; Lind, J.
Chem. Res. Toxicol. 1997, 10, 1246.
(13) Neta, P.; Huie, R. E.; Ross, A. B. J. Phys. Chem. Ref. Data 1988,
17, 1027.
S0002-7863(97)03536-1 CCC: $15.00 © 1998 American Chemical Society
Published on Web 05/21/1998

5550 J. Am. Chem. Soc., Vol. 120, No. 22, 1998
Goldstein et al.
H• + OH- f eaq
k2 ) 2.2 × 107 M-1 s-1 (ref 19) (2)
-
the rest being NaCl and NaClO3. Water for solution preparations was
distilled and purified with a Milli-Q purification system.
Fresh solutions of peroxynitrite were prepared daily by reacting nitrite
with acidified hydrogen peroxide at room temperature in a quenched-
flow system. This system was optimized, as we recently described,14
to produce the high yield of ONOO- with only a minor known
e-aq + N2O f N2 + OH- + OH•
k3 ) 9.1 × 109 M-1 s-1 (ref 19) (3)
•OH + CO32- f CO3•- + OH-
-
contamination by residual NO2 and H2O2. Minimization of these
contaminants is crucial for studying oxidation of ONOO- by free
radicals because both NO2- and H2O2 can compete for these radicals.13
A Syringe pump (WPI, model SP 230IW) was used to inject either
(preparation A) 0.63 M NaNO2 and 0.60 M H2O2 in 0.7 M HClO4 or
(preparation B) 0.60 M NaNO2 and 0.60 M H2O2 in 0.7 M HClO4 into
the first mixing chamber through tangential inlets. The combined
solutions were allowed to react in a delay line connected to a second
mixing chamber where 3.6 M NaOH was injected with the same flow
rate (45 mL/min) to quench the reaction.14 The yield of ONOO-
depends on the time of quenching14 and was determined from its
k4 ) 3.9 × 109 M-1 s-1 (ref 19) (4)
O
•- + CO32- f CO3•- + OH-
(5)
Although the rate constant for reaction 5 has not been
determined,19 it should be smaller than that for reaction 4 due
to electrostatic repulsion. The effective rate constant of the
reaction of hydroxyl radicals with CO32- is pH dependent and
given by
absorption at 302 nm with ꢀ ) 1670 M-1 cm-1 15
The residual H2O2
.
in the stock peroxynitrite solutions was determined iodometrically16
after diluting these solutions in 0.2 M acetate buffer at pH 4.6 and, in
parallel, at pH 2. Because residual NO2- destroys H2O2 at acidic pH,
k4[H+]
KOH + [H+] KOH + [H+]
k5KOH
-
the yield of I3 at pH 2 was used as a blank, which was subtracted
k6 )
+
(6)
-
from the yield of I3 obtained at pH 4.6. It was determined that
preparation A contained 0.18 M ONOO- and 0.32-0.49 mM H2O2.
Preparation B contained 0.12 M ONOO- and 0.94-1.1 mM H2O2.
These yields of peroxynitrite and residual H2O2 are in good agreement
with those calculated from kinetic simulation.14 Specifically, simulation
for preparation A gave 0.188 M ONOO-, 52 mM NO3-, 12.3 mM
NO2-, and 0.29 mM H2O2, while preparation B gave 0.122 M ONOO-,
0.117 M NO3-, 1.19 mM NO2-, and 1.19 mM H2O2.
•
where KOH is the acid dissociation constant of the OH radical
(pKa ) 11.9).19
In the absence of peroxynitrite, the decay of CO3•- monitored
at 600 nm was second order:
Methods. Pulse radiolysis experiments were carried out mainly in
Jerusalem with a 5 MeV Varian 7715 linear accelerator (0.2-1.5 µs
electron pulses, 200 mA current). Some experiments were done at
Brookhaven with a 2 MeV Van de Graaff accelerator (0.05-0.2 µs
pulses). All measurements were made at room temperature in a 4-cm
spectrosil cell using three light passes (optical path length 12.1 cm).
The yields and decay kinetics of CO3•-, •NO2, •N3, and ClO2• radicals
CO3•- + CO3•- f products
(7)
with 2k7 ) (3.2 ( 0.4) × 107 M-1 s-1 at I ) 0.6 M and (1.8
( 0.2) × 107 M-1 s-1 at I ) 0.04 M, in agreement with the
literature values.13 We found that addition of up to 0.2 mM
NO3-, which is the product of the ONOOH isomerization,3,4
were measured from their absorption with ꢀ600 ) 1860 M-1 cm-1 13
ꢀ400 ) 200 M-1 cm-1 13 ꢀ275 ) 1690 M-1 cm-1 17 and ꢀ358 ) 1250 M-1
cm-1 13
respectively.
,
•-
had no effect on the yield and decay rates of CO3 at pH 12
,
,
•-
(data not shown). The yield of CO3 was also unaffected by
,
the addition of up to 1.3 mM of ONOO-, indicating that, under
these conditions, carbonate ions scavenged all hydroxyl radicals
produced. However, the decay rate of CO3•- was enhanced in
the presence of ONOO-, and changed from second-order to
first-order kinetics (Figure 1). When [ONOO-]o . [CO3•-]o,
Results and Discussion
All reactions between ONOO- and the oxidizing radicals
were studied at pH g12, where the ONOO- ion is relatively
stable.3,4 Less than 5% decomposition of peroxynitrite occurred
within 15 min after solution preparation, which was the average
duration of each experiment.
•-
the observed first-order rate constant of CO3 decay was
independent of [CO3•-]o and depended linearly on the concen-
tration of added ONOO- (Figure 2, upper line). Assuming that
Carbonate Radical. The rate constants for the reactions of
the observed effect is due to oxidation of ONOO- by CO3
,
•-
CO3 with NO2-, H2O2, and HO2 are 6.6 × 105, 8 × 105,
and (1-5.6) × 107 M-1 s-1, respectively.13 Concentration of
residual H2O2 (pKa ) 11.7)18 in the peroxynitrite preparation
must therefore be kept as low as possible to minimize interfer-
ence from HO2-. For this reason we used preparation A for
these experiments as in this preparation contamination by H2O2
is less than 0.3% of ONOO- (see Experimental Section).
Carbonate radical was generated by irradiating N2O-saturated
(∼25 mM) aqueous solutions containing 0.2 M sodium carbon-
ate via the following reactions (given in parentheses are the
species radiation yields):
•-
-
CO3•- + ONOO- f ONOO• + CO3
(8)
2-
we determined k8 ) (9.6 ( 1.4) × 106 M-1 s-1 from the slope
of the upper line in Figure 2.
Since the value of k8 is small, we wanted to verify that the
•-
observed acceleration of CO3 decay was, indeed, due to
•-
reaction 8 and was not a result of CO3 scavenging by
contaminants, such as HO2-, introduced with ONOO-. The
lower line in Figure 2 shows the observed first-order rate
constant of CO3•- decay measured repeatedly over a period of
time, during which the initially added 1 mM ONOO- partially
decayed; the concentration of ONOO- remaining in solution
was determined spectrophotometrically prior to each kinetic
measurement. Again, kobs was linearly dependent on [ONOO-]
with the corresponding value of k8 ) (7.7 ( 1.2) × 106 M-1
H2O f e-aq(2.6), •OH (2.7), H• (0.6), H3O+ (2.6) (1)
(14) Saha, A.; Goldstein, S.; Cabelli, D.; Czapski, G. Free Radical Biol.
Med. 1998, 24, 653.
(15) Hughes, M. N.; Nicklin, H. G. J. Chem. Soc. A 1968, 450.
(16) Klassen, N. V.; Marchington, D.; McGowan, H. C. E. Anal. Chem.
1994, 66, 2921.
(17) Goldstein, S.; Czapski, G. Inorg. Chem. 1996, 35, 7735.
(18) Smith, R. M.; Martell, A. E. Critical Stability Constants; Plenum
Press: New York, 1976; Vol. 4.
(19) Buxton, J. V.; Greenstock, C. L.; Helman, W. P.; Ross, A. B. J.
Phys. Chem. Ref. Data 1988, 17, 513.

Reactions of Inorganic Radicals with Peroxynitrite
J. Am. Chem. Soc., Vol. 120, No. 22, 1998 5551
Figure 1. Kinetic traces of CO3•- radical decay monitored at 600
nm in the absence of peroxynitrite (a) and in the presence of 290 µM
(b) and 1.3 mM peroxynitrite (c). All solutions were N2O saturated
and contained 0.2 M Na2CO3 at pH 12. The concentration of CO3•- at
the end of the radiation pulse varied between 3.5 and 5.4 µM.
Figure 3. Reciprocal of the maximum absorption at 600 nm as a
function of [ONOO-] in the presence of: (b) 3 mM CO32- at pH 12,
2-
[CO3•-]o ) 15 µM; (9) 10 mM CO3 at pH 12, [CO3•-]o ) 15 µM,
2-
and (0) 20 mM CO3 at pH 12.6, [CO3•-]o ) 18 µM. All solutions
were N2O saturated.
k(O-+HO2) ) 4.0 × 108 M-1 s-1),19 but relatively slowly with
-
H2O2 (k(OH+H2O2) ) 2.7 × 107 M-1 s-1 19
)
and NO3
(k(OH+NO3-) ) 1.4 × 108 M-1 s-1).19 To minimize the effects
of these impurities we used peroxynitrite preparation B where
-
contamination by both residual H2O2 and NO2 was less than
1% of the ONOO- concentration (see Experimental Section).
The rate constant for the reaction between OH/O•- and
•
ONOO- was determined at pH 12 and 12.6 from competition
between ONOO- and CO32- (reactions 4 and 5) for the hydroxyl
•
radical. If both OH and O•- oxidize ONOO-,
•OH + ONOO- f ONOO• + OH-
•- + ONOO- f ONOO• + OH-
(9)
O
(10)
the effective rate constant is:
Figure 2. Observed first-order rate constant for the decay of CO3•- as
a function of the peroxynitrite concentration. Variation of [ONOO-]
by (b) diluting the stock peroxynitrite solution, [CO3•-]o ∼ 9 µM, and
(0) letting 0.97 mM ONOO- decay (details in the text), [CO3•-]o ∼ 3
µM. All solutions were N2O saturated and contained 0.2 M Na2CO3 at
pH 12.
k9[H+]
k10KOH
k11
)
+
(11)
KOH + [H+] KOH + [H+]
•-
The yield of CO3 observed following pulse irradiation
decreased with increasing the concentration of added ONOO-
as shown in Figure 3. Under our experimental conditions, the
s-1, somewhat lower than the value determined from the upper
line in Figure 2.
•-
half-life of the CO3 formation was shorter than that of the
A small intercept is expected for both lines in Figure 2 due
•-
2-
reaction between CO3
and ONOO- (k6[CO3
]
.
o
•-
k8[ONOO-]o), therefore the latter reaction had no effect on the
to the contribution from the bimolecular self-decay of CO3
and some scavenging by contaminating NO2-. However, if
reactive impurities are present, a larger intercept is expected
for the lower line. The lack of measurable intercepts indicates
that there is no significant contribution from contaminants to
the decay of CO3•-. The slope of the lower line in Figure 2
should not be affected by the impurities (provided they do not
decay simultaneously with ONOO-), whereas reactive impurities
should result in a larger slope of the upper line. Therefore, the
value obtained from the lower curve in Figure 2 seems to be
more reliable. The rate constant for reaction 8 at zero ionic
strength k8 ) (3.7 ( 0.6) × 106 M-1 s-1 was calculated from
the Debye-Huckel-Bronsted-Davies expression for the pri-
mary kinetic salt effect: log(k/ko) ) 1.02ZAZBI1/2/(1 + I1/2) -
0.2I.
•-
yield of CO3
. Considering competition for the hydroxyl
radicals between CO3 (reactions 4 and 5) and ONOO-
2-
(reactions 9 and 10) we obtain the following for the yield of
•-
CO3
:
[CO3•-] ) [CO3 ]ok6[CO32-]/(k6[CO32-]o + k11[ONOO-]o)
-
(12)
or, after rearrangement:
k11[ONOO-]o
k6[CO3•-]o[CO3
1
1
)
+
(13)
•-
•-
2-
[CO3
]
[CO3
]
]
o
o
Hydroxyl Radical. This radical reacts rapidly with NO2
where [CO3•-]o is the yield of CO3•- in the absence of ONOO-.
Equation 13 predicts a linear plot of 1/OD600 vs [ONOO-]o,
as observed (Figure 3), with the slope-to-intercept ratio S/I )
-
(k(OH+NO2-) ) 1.1 × 1010 M-1 s-1; k(O-+NO2-) ) 3.1 ×
108 M-1 s-1 19
)
and HO2- (k(OH+HO2) ) 7.5 × 109 M-1 s-1
;

5552 J. Am. Chem. Soc., Vol. 120, No. 22, 1998
Goldstein et al.
N2O4 + H2O f NO3- + NO2- + 2H+
k18 ) 1 × 103 s-1 (ref 20) (18)
This half-life was not shortened by the addition of up to 2.1
mM peroxynitrite. Taking into account that less than 20%
acceleration of •NO2 decay would be readily detectable, we can
estimate an upper limit for the rate constant of the reaction of
•NO2 with ONOO- as 2.5 × 104 M-1 s-1. This result is in
quantitative agreement with the observation by Lymar and
Hurst10 that the decomposition of ONOO- is not accelerated
•
when the NO2 radical is produced in the ONOO- - CO2 -
NO2- system. Based on kinetic simulation of their data, a very
similar upper limit of 2 × 104 M-1 s-1 can be set for the rate
•
constant of the reaction between NO2 and ONOO-.
Azide Radical. We used peroxynitrite preparation B (see
Experimental Section) to study the oxidation of ONOO- by this
-
radical, because minimization of both residual H2O2 and NO2
may be equally important in this case. The azide radical reacts
rapidly with HO2- (k ) 3.2 × 109 M-1 s-1).13 The rate constant
for oxidation of NO2- by •N3 has not been reported,13 but it is
not expected to be larger than that for oxidation of ClO2-, which
Figure 4. Dependence of the observed first-order rate constant for
the formation of CO3 on the carbonate concentration at pH 12 (b)
•-
and 12.6 (9). All solutions were N2O saturated. The maximum
•-
concentrations of CO3 varied between 3 and 11 µM.
is ca. 2 × 109 M-1 s-1 9
.
k11/k6[CO32-]o. From data in Figure 3 for pH 12 we obtained
k11/k6 ) 15.8 and 17.3 in the presence of 3 and 10 mM
carbonate, respectively. The latter value appears to be more
reliable, because the loss of hydroxyl radicals due to their
recombination can be neglected in more concentrated carbonate.
In 20 mM carbonate at pH 12.6 the ratio k11/k6 ) 22 was
calculated from the data in Figure 3. Values of k6 ) (1.2 (
0.1) × 108 and (3.4 ( 0.3) × 107 M-1 s-1 at pH 12 and 12.6,
respectively, were determined in a separate experiment by
When an N2O-saturated solution containing 20 mM NaN3 at
•
pH 12 is irradiated, N3 is formed via reactions 1-3 and 19:
•OH/O•- + N3- f •N3 + OH-
k19 ) 5.4 × 109 M-1 s-1 (pH 12) (ref 19) (19)
About 3 µM •N3 was produced by an electron pulse. The decay
of •N3 followed at 275 nm was second order with 2k ) (9.8 (
1.0) × 109 M-1 s-1, in agreement with the literature values.13
The second-order kinetics turned into first order upon addition
of 17-80 µM ONOO-. The observed first-order rate constant
was linearly dependent on the concentration of ONOO- (Table
1), yielding the rate constant k20 ) (7.2 ( 0.9) × 108 M-1 s-1
for the reaction
•-
following the rate of the formation of CO3 as a function of
[CO32-] at low doses (Figure 4). Using these values we
obtained k11 ) (2.1 ( 0.2) × 109 and (7.5 ( 0.7) × 108 M-1
s-1 at pH 12 and 12.6, respectively. From these rate constants
and eq 11 we calculated k9 ) (4.8 ( 0.8) × 109 M-1 s-1 and
k10 < 1.7 × 108 M-1 s-1
.
Nitrogen Dioxide. Approximately 20-23 µM of •NO2
radicals were generated by pulse irradiation of an Ar-saturated
solution containing 0.46 M NO3- and 0.05 M NO2- at pH 12,
where all the primary radicals formed in reaction 1 were
•N3 + ONOO- f ONOO• + N3
(20)
-
This value of k20 was, within experimental error, identical to
the value obtained from the kinetics of ONOO- bleaching
observed at 302 nm.
•
converted into NO2 via reactions 2 and 14-16:
e-aq + NO3- f NO3
2-
Chlorine Dioxide. Preparation B of peroxynitrite (see
Experimental) was used to study the oxidation of ONOO- by
this radical. Peroxynitrite was relatively stable in the presence
of NaClO2 at pH 12; a loss of about 8% and 14% of ONOO-
was detected within 30 min upon addition of 8 and 40 mM
chlorite, respectively. The average time of each experiment was,
therefore, kept under 10 min.
k14 ) 9.7 × 109 M-1 s-1 (ref 19) (14)
NO32- + H2O f •NO2 + 2OH-
k15 ) 5.5 × 104 M-1 s-1 (ref 20) (15)
•
The ClO2 radical was generated by the pulse irradiation of
•OH/O•- + NO2- f •NO2 + OH-
N2O-saturated solutions containing 24 mM NaClO4 at pH 12
through reactions 1-3 and 21:
k16 ) 6.4 × 109 M-1 s-1 (pH 12) (ref 19) (16)
•
•OH/O•- + ClO2- f ClO2 + OH-
In the absence of added ONOO-, the decay kinetics of •NO2
monitored at 400 nm corresponded to the well-established
pathway (reactions 17 and 18) with a measured half-life of 2.7
ms.
k21 ) 3.0 × 109 M-1 s-1 (pH 12) (ref 19) (21)
•
The formation and decay of ClO2 was followed at 420 nm,
where ONOO- does not absorb. In the presence of ONOO-,
•
the ClO2 radical, which was otherwise stable, decayed expo-
•NO2 + •NO2 h N2O4
nentially, apparently in the reaction
k17 ) 4.5 × 108 M-1 s-1; k-17 ) 6.9 × 103 s-1 (ref 20)
ClO2 + ONOO- f ONOO• + ClO2
(22)
•
-
(17)

Reactions of Inorganic Radicals with Peroxynitrite
J. Am. Chem. Soc., Vol. 120, No. 22, 1998 5553
equilibrium constant, and ln f12 ) (ln K12)2/4 ln(k11k22/1022)),26
the following ratios were derived:
Table 1. The Observed First-Order Rate Constants for the
•
•
Reactions of N3 and ClO2 with ONOO-
kobs, s-1
•
-
•
[ONOO-], M
azide radicala
chlorine dioxideb
k122(HO2-/ClO2 )
K12(HO2 /ClO2 )
) C
(23)
(24)
1.7 × 10-5
3.7 × 10-5
6.4 × 10-5
8.0 × 10-5
1.2 × 10-4
3.6 × 10-4
7.1 × 10-4
1.2 × 10-3
1.9 × 10-3
(3.0 ( 0.3) × 104
(4.7 ( 0.3) × 104
(5.8 ( 0.5) × 104
(7.9 ( 0.4) × 104
k122(ONOO-/ClO2 )
K12(ONOO-/ClO2 )
•
•
•
-
-
•
k11(HO2 /HO2 )
k11(ONOO•/ONOO-) f12(ONOO-/ClO2 )
f12(HO2 /ClO2 )
C )
3.0 ( 0.3
11.8 ( 1.8
22.5 ( 0.4
33.9 ( 1.2
64.5 ( 1.5
•
When the reaction driving force is not very large, factor f12 is
-
close to unity, e.g., f12 ) 0.71 for the oxidation of HO2 by
•
a Measured from the decays of N3 and ONOO- followed at 275
and 302 nm, respectively, in pulse-irradiated (dose 5 Gy) N2O-saturated
solutions containing 20 mM azide at pH 12. b Measured from the decay
•
ClO2 . If the magnitudes of k11 and K12 are not vastly different
-
for HO2 and ONOO-, factor C should also be close to unity.
With this assumption, we calculated E°(ONOO•/ONOO-) ) 0.8
V. Note that this result is not very sensitive to the value of C
chosen; variation of C between 0.1 and 10 changes E°(ONOO•/
ONOO-) from 0.86 to 0.74 V. From the corresponding data
in Table 2 for the •N3 radical, we obtained E°(ONOO•/ONOO-)
) 0.83 V, close to the former estimate. Both estimates are also
in good agreement with the thermodynamic estimate of 0.9 V.12
•
of ClO2 . followed at 420 nm in pulse-irradiated (dose 26 Gy) N2O-
saturated solutions containing 24 mM chlorite at pH 12.
Table 2. Rate Constants and Redox Potentials of Various
Radicals with ONOO- and HO2
-
E°(R•/R-) k(R•/R-)a k(R•+ONOO-)
k(R• + HO2
)
-
R•/R-
(V)
(M-1 s-1
30023
)
(M-1 s-1
)
(M-1 s-1
)
•OH/OH-
1.921
4.8 × 109
7.0 × 109 13
Implications for the Peroxynitrite Decomposition. In the
presence of CO2, the decomposition of peroxynitrite yields 30-
CO3•-/CO3
1.5922 0.423
3.7 × 106 (I ) 0) (1.0-5.6) × 107 13
2-
-
•N3/N3
1.3321 3.7 × 106 24 7.2 × 108
3.2 × 109 13
35% of CO3 and NO2.8-10 Because the rate constants for
oxidation of ONOO- by both CO3•- and •NO2 are rather small
(Table 2), the contribution from these reactions to the overall
decomposition is predicted to be minor even under the most
favorable catalytic conditions, i.e., when [ONOO-] > [CO2].
This conclusion is supported by numerical simulations of the
CO2-catalyzed peroxynitrite decomposition kinetics, which we
performed using previously reported data7-10 and the rate
constants from Table 2.
•-
•
-
•NO2•/NO2
1.0421 9.624
<2 × 104
n.d.
-
ClO2 /ClO2
0.93421 3.3 × 104 24 3.2 × 104
8.0 × 104 13
0.7521 1725
•
-
HO2 /HO2
a The self-exchange rate constants have been calculated in the cited
references by applying the Marcus cross relationship to the reactions
between the corresponding species and a series of coordination
compounds.
From the dependence of kobs on the concentration of ONOO-
given in Table 1, the rate constant k22 ) (3.2 ( 0.3) × 104
M-1 s-1 was determined.
A strong case was recently made by Merenyi and Lind12d
that the decomposition of ONOOH proceeds via the homolytic
cleavage of its peroxy bond producing a geminate pair of the
•OH and •NO2 radicals. It was concluded that the observed ca.
40% yield of the indirect oxidation of various substrates by
peroxynitrite5 represents the cage escape probability for the
Oxidation Potential of ONOO-. Rate constants for the
oxidation of ONOO- by the radicals determined in this study,
the corresponding literature values for the oxidation of HO2
-
,
•
•
geminate pair of the OH and NO2 radicals. The large rate
constant for reaction 9 determined in this study suggests that,
and the redox potentials of the radicals are summarized in Table
2. It is apparent that these rate constants do not reflect the trend
in the reduction potentials of the radicals. For instance, the
rate constants for oxidation of both ONOO- and HO2- by •N3
are much larger than those for oxidation by CO3•-, which is a
stronger oxidant. A similar effect was observed for the
oxidation of many other substrates by these radicals,13 which,
most probably, reflects a much higher self-exchange rate for
the •N3/N3- couple (Table 2). Although the reduction potential
•
in the absence of other reactants, OH will be efficiently
scavenged by ONOO-, yielding ONOO• even at pH as low as
5. The fate of the ONOO• is unexplored. By analogy with the
HO2 radical, ONOO• may disproportionate yielding oxygen
•
(reaction 25) with subsequent hydrolysis of •NO2 (reactions 17
and 18).
ONOO• + ONOO• f 2 •NO2 + O2
(25)
of ClO2 is somewhat lower than that of •NO2, the rate constants
•
•
for the one-electron oxidation by ClO2 are generally larger than
Alternatively, the ONOO• radical may undergo decomposition
(reaction 26) followed by reactions 27 and 28.
those by •NO2, also due to the higher self-exchange rate for the
•
-
ClO2 /ClO2 couple (Table 2). On the basis of this trend, one
expects that the reactivity of ClO2• toward ONOO- will exceed
ONOO• h •NO + O2
(26)
•
that of NO2, in agreement with our results.
(20) Graetzel, M.; Henglein, A.; Lilie, J.; Beck, G. Ber. Bunsen-Ges.
Phys. Chem. 1969, 73, 646.
(21) Stanbury, D. M. AdV. Inorg. Chem. 1989, 33, 69.
(22) Huie, R. E.; Clifton, C. L.; Neta, P. Radiat. Phys. Chem. 1991, 38,
477.
(23) Schindler, S.; Castner, E. W., Jr.; Creutz, C.; Sutin, N. Inorg. Chem.
1993, 32, 4200.
(24) Avad, H. H.; Stanbury, D. M. J. Am. Chem. Soc. 1993, 115, 3636.
(25) Lind, J.; Shen, X.; Merenyi, G.; Jonsson, B. O. J. Am. Chem. Soc.
1989, 111, 7654.
Assuming that reaction 22 represents a simple one-electron-
transfer we can estimate the oxidation potential of ONOO- using
Marcus formalism by comparing the rate constant for ONOO-
oxidation obtained here with the corresponding value for the
similar reaction of HO2- (Table 2). From the Marcus equation
k12 ) (k11k22K12f12)1/2 (here k12 is the electron-transfer rate
constant for the cross reaction, k11 and k22 are the self-exchange
rate constants for the reactants, K12 is the cross reaction
(26) Marcus, R. A. J. Phys. Chem. 1963, 67, 853.

5554 J. Am. Chem. Soc., Vol. 120, No. 22, 1998
•NO2 + •NO h N2O3
Goldstein et al.
possible that reactions 25 and 26 are, at least in part, responsible
for the observed O2 generation.27,28
k27 ) 1.1 × 109 M-1 s-1; k-27 ) 8.4 × 104 s-1 (ref 27)
(27)
Acknowledgment. This research was supported by Grant
No. 4129 from The Council For Tobacco Research and by The
Israel Science Foundation, and partially by the U.S. Department
of Energy under contract DE-AC02-76CH00016. S.L. is grateful
for a financial support by LDRD Grant No. 97-29 from BNL.
The authors wish to thank Drs. G. Merenyi and H. Schwarz for
helpful discussions.
N2O3 + H2O f 2NO2- + 2H+
k28 ) (2 × 103) + (1 × 108)[OH-] s-1 (ref 28) (28)
Indeed, ab initio calculations by McKee12a suggest that the
ONOO• radical is unstable by about 11.5 kcal/mol with respect
to reaction 26. In addition, it was recently reported that molecular
oxygen is found among the products of peroxynitrite decom-
position at pH 6-9.6,10 Concentrations of ONOO- used in these
studies were high enough to make reaction 9 an important
pathway for the •OH radical decay. It is therefore
JA9735362
(27) Von Graetzel, M.; Taniguchi, S.; Henglein, A. Ber. Bunsen-Ges.
Phys. Chem. 1970, 74, 488.
(28) Treinin, A.; Hayon, E. J. Am. Chem. Soc. 1970, 92, 5821.