Reactivity of Peroxynitric Acid (O2NOOH): A Pulse Radiolysis Study

Article abstract of DOI:10.1021/ic961186z

Peroxynitrate (O₂NOOH/O₂NOO^-) is formed within less than 2 ms after pulse irradiation of aerated solutions containing relatively low concentrations of formate and nitrate. The pK_a for peroxynitric acid was determined to be 5.9 ± 0.1 both from the pH-dependent absorbance of the anion at 310 nm and from the dependence of the decay kinetics on pH. An absorption spectrum was measured for the anion giving ε_max(290) = 1500 ± 100 M^-1 cm^-1. This method of generation of peroxynitrate is very useful for studying the mechanism of the oxidation of various substrates by peroxynitrate. The oxidation by peroxynitrate can take place either directly or indirectly. In the direct oxidation pathway, the reaction is first order in peroxynitrate and first order in the substrate, whereas in the indirect oxidation pathway, the reaction is zero order in the substrate. In both cases, the observed rate constants are highly pH-dependent. The results show that the direct oxidation pathway takes place through O₂NOOH. We suggest that the indirect oxidation takes place through reactive intermediates that are formed during the decomposition of peroxynitrate. In the presence of sufficient concentrations of the substrates, the oxidation yields approach 100% through the direct and indirect oxidation pathways.

Full text of DOI:10.1021/ic961186z

4156
Inorg. Chem. 1997, 36, 4156-4162
Reactivity of Peroxynitric Acid (O2NOOH): A Pulse Radiolysis Study
Sara Goldstein* and Gidon Czapski
Department of Physical Chemistry, The Hebrew University of Jerusalem, Jerusalem 91904, Israel
ReceiVed September 27, 1996X
Peroxynitrate (O2NOOH/O2NOO-) is formed within less than 2 ms after pulse irradiation of aerated solutions
containing relatively low concentrations of formate and nitrate. The pKa for peroxynitric acid was determined to
be 5.9 ( 0.1 both from the pH-dependent absorbance of the anion at 310 nm and from the dependence of the
decay kinetics on pH. An absorption spectrum was measured for the anion giving ꢀmax(290) ) 1500 ( 100 M-1
cm-1. This method of generation of peroxynitrate is very useful for studying the mechanism of the oxidation of
various substrates by peroxynitrate. The oxidation by peroxynitrate can take place either directly or indirectly.
In the direct oxidation pathway, the reaction is first order in peroxynitrate and first order in the substrate, whereas
in the indirect oxidation pathway, the reaction is zero order in the substrate. In both cases, the observed rate
constants are highly pH-dependent. The results show that the direct oxidation pathway takes place through
O2NOOH. We suggest that the indirect oxidation takes place through reactive intermediates that are formed
during the decomposition of peroxynitrate. In the presence of sufficient concentrations of the substrates, the
oxidation yields approach 100% through the direct and indirect oxidation pathways.
Introduction
study, we used the pulse radiolysis technique to study the
mechanism of the oxidation of various substrates by peroxyni-
trate. This method seems to be the best for this purpose because
it does not require extreme conditions.
Peroxynitric acid, O2NOOH, is formed in the gas phase by
•
•
the recombination of HO2 and NO2 radicals.1-3 In the gas
phase, at ordinary temperatures, the compound is in equilibrium
with its precursors, and it decays slowly due to the dismutation
Experimental Section
• 4,5
of HO2 .
In aqueous solutions, peroxynitrate (O2NOOH/
O2NOO-) decomposes mainly through a unimolecular dissocia-
tion of the anion into nitrite and oxygen.6-8 It has been
suggested that the decomposition of O2NOOH takes place
through its dissociation into HO2• and •NO2, but recently it was
argued that O2NOOH decomposes directly into HNO2 and O2.8
Peroxynitric acid is a strong oxidizing agent, reacting rapidly
with I-, Br-, Cl-, VO2+, and benzene.7,9 The oxidation
mechanism has not been investigated, as it is very difficult to
prepare O2NOOH in aqueous solutions. The reported methods
for its preparation are as follows: (i) 90% H2O2 and 70%
HNO3;6 (ii) 90% H2O2 and NO2BF4;6 (iii) HNO2 and excess
H2O2;9 (iv) pulse radiolysis of O2-saturated nitrite/nitrate solu-
tions.8
Nitrogen dioxide is one of the most important toxic compo-
nents of photochemical smog,10 and thus, understanding the
reactions that •NO2 undergoes in the lungs exposed to smoggy
air is of considerable importance. Peroxynitrate may be formed
in the lungs through the reaction of superoxide with •NO2,8 and
therefore the mechanism of its formation and decomposition
as well as its redox chemistry is of great importance. In this
Materials. All chemicals were of analytical grade and were used
as received. â-Nicotinamide adenine dinucleotide, reduced (NADH)
from Grade III yeast was obtained from Sigma. Solutions of NADH
were prepared immediately before use, and the concentration of NADH
was determined using ꢀ340 ) 6200 M-1 cm-1. Solutions were prepared
with deionized water that was distilled and purified using a Milli-Q
water purification system, and unless otherwise stated, they contained
100 µM EDTA. The pH was adjusted with the use of 1 mM acetate,
phosphate, or borate buffers. All experiments were carried out at 22
°C.
Methods. Pulse radiolysis experiments were carried out with the
Varian 7715 linear accelerator using 5 MeV electron pulses of 0.1-
1.5 µs and a 200 mA current. The dose per pulse was 3-29 Gy,
respectively, and was determined with a hexacyanoferrate(II) dosimeter
(5 mM K4Fe(CN)6 in N2O-saturated water) using Gꢀ(Fe(CN)63-) )
6.7 × 103 M-1 cm-1 at 420 nm.11 A 150 W Xe or a 200 W Xe-Hg
lamp produced the analyzing light. Appropriate filters were used to
minimize photochemistry. Irradiations were carried out in 1- or 4-cm-
long Spectrosil cells using one or three light passes.
Results
Formation of O2NOOH/O2NOO-. Logager and Sehested8
produced peroxynitrate by irradiating O2-saturated solutions
containing nitrite or nitrate. We modified this method somewhat
and irradiated air-saturated solutions containing nitrate and
formate. Under these conditions, the following reactions take
place:
* To whom all correspondence should be directed. Tel: 972-2-6586478.
Fax: 972-2-6586925. E-mail: SARAG@HUJI.VMS.AC.IL.
X Abstract published in AdVance ACS Abstracts, August 1, 1997.
(1) Niki, H.; Maker, P. D.; Savage, C. M.; Breitenbach, L. P. Chem. Phys.
Lett. 1977, 45, 564.
(2) Hanst, P. L.; Gay, B. W. EnViron. Sci. Technol. 1977, 11, 1105.
(3) Levine, S. Z.; Uselman, W. M.; Chan, W. H.; Calvert, J. G.; Shaw, J.
H. Chem. Phys. Lett. 1977, 48, 528.
(4) Uselman, W. M.; Levine, S. Z.; Chan, W. H.; Calvert, J. G.; Shaw, J.
H. Chem. Phys. Lett. 1978, 58, 437.
(5) Kuryolo, M. J.; Quellette, P. A. J. Phys. Chem. 1986, 90, 441-444.
(6) Keleny, R. A.; Trevor, P. L.; Lan, B. Y. J. Am. Chem. Soc. 1981,
103, 2203.
H2O
9
γ8 eaq- (2.6), •OH (2.7), H• (0.6), H2 (0.45),
H2O2 (0.7), H3O+ (2.6)
(1)
The numbers in parentheses are G values, which represent the
(7) Lammel, G.; Perner, D.; Warneck, P. J. Phys. Chem. 1990, 94, 6141.
(8) Logager, T.; Sehested, K. J. Phys. Chem. 1993, 97, 10047.
(9) Appelman, E. V.; Gosztola, D. J. Inorg. Chem. 1995, 34, 787.
(10) Pryor, W. H.; Stone, K. Ann. N.Y. Acad. Sci. 1993, 686, 12.
(11) Buxton, G. V.; Stuart, C. R. J. Chem. Soc., Faraday Trans. 1995, 91,
279.
S0020-1669(96)01186-X CCC: $14.00 © 1997 American Chemical Society

Reactivity of Peroxynitric Acid
Inorganic Chemistry, Vol. 36, No. 19, 1997 4157
number of molecules formed per 100 eV of energy absorbed
by pure water.
HO2•/O2•-, which subsequently yield O2NOOH/O2NOO- through
reactions 12 and 13.
•NO2 + O2•- f O2NOO-
k12 ) 4.5 × 109 M-1 s-1 8
(12)
(13)
(14)
•-
eaq- + O2 f O2
(2)
(3)
(4)
(5)
(6)
(7)
(8)
(9)
k2 ) 1.9 × 1010 M-1 s-1 12
•
•NO2 + HO2 f O2NOOH
eaq- + NO3- f NO3
2-
k13 ) 1.8 × 109 M-1 s-1 8
O2NOOH h H+ + O2NOO-
pKa ) 5.858
k3 ) 9.7 × 109 M-1 s-1 12
eaq- + H+ f H•
The change in the absorbance with time was monitored at
k4 ) 2.3 × 1010 M-1 s-1 12
250 nm, where ꢀ250(O2•-) ) 2250 M-1 cm-1 13
When O2
.
•-
•
was generated in excess over NO2 at pH >6.8, a fast first-
•-
order decay of O2 followed by a slower second-order decay
NO32- + H2O f •NO2 + 2OH-
was observed. During the fast decay observed at 250 nm, a
transient was formed with maximum absorbance around 290
nm (Figure 1). The yield of the transient as measured at 310
(where superoxide does not absorb) decreased with the decrease
in pH with an apparent pKa ) 5.8 ( 0.1 (Figure 2), assuming
that O2NOOH does not absorb at 310 nm.8
The total yield of peroxynitrate, G(O2NOOH)T, equals
G(•NO2) because under our experimental conditions there is
always an excess of superoxide over •NO2. The yield is given
by
k5 ) 5.5 × 104 s-1 12
NO32- + H+ f •NO2 + OH
k6 ) 2 × 1010 M-1 s-1 12
NO32- + O2 f NO3- + O2
•-
G(O2NOOH)T )
-
k3[NO3 ]
k5 + k6[H+]
Ge )
-
k3[NO3 ] + k2[O2] + k4[H+] k5 + k6[H+] + k7[O2]
k7/k5 ) 576 M-1 (see below)
k5 + k6[H+]
k5 + k6[H+] + k7[O2]
RGe (15)
•
H• + O2 f HO2
-
G(O2 )T ) GOH + GH + Ge - G(O2NOOH)T (16)
k8 ) 2 × 1010 M-1 s-1 12
As predicted by eq 15, the yield of peroxynitrate at pH 6.8
decreased with the increase in [O2] (Table 1). A plot of R/OD320
versus [O2] yields a straight line with slope/intercept ) k7/k5 )
576 M-1. Forni et al.14 determined k7/k5 ) 1.5 × 103 M-1 at
pH 6 (k5 ) 5.5 × 104 s-1, k7 ) 8.5 × 107 M-1 s-1) in aqueous
solutions containing 0.5 mM ABTS (2,2′-azinobis(3-ethylben-
zothiazoline-6-sulfonic acid), 1 M tert-butyl alcohol, and 0.5
M nitrate. In their study,14 the yield of ABTS+ decreased with
the increase in [O2], and the [ABTS•+]/[eaq-] values were 0.81
and 0.51 in air/N2 and O2/N2, respectively. From their results
we calculated that k7/k5 ) 772 M-1, which is closer to our value.
We were unable to determine how they calculated their reported
value.
•OH/H• + HCO2- f H2O/H2 + CO2
•-
kOH ) 3.5 × 109 M-1 s-1 12
kH ) 2.1 × 108 M-1 s-1 12
CO2•- + O2 f CO2 + O2
(10)
•
k10 ) 2.4 × 109 M-1 s-1 12
•
•-
HO2 h H+ + O2
(11)
Using our value for k7/k5 ) 576 M-1, eqs 15 and 16, and the
•-
experimental conditions of Figure 1, we calculated that [O2
]
0
) 10.3 µM and [•NO2]0 ) [O2NOO-]T ) 6.07 µM. The
pKa ) 4.813
•-
residual spectrum of peroxynitrate was corrected for excess O2
(Figure 1), resulting in ꢀ290max(O2NOO-) ) 1500 ( 150 M-1
cm-1, assuming that O2NOOH does not absorb at 290 nm and
its pKa equals 5.85.8
Under our experimental conditions ([KNO3] ) 0.375-40
mM, [NaHCO2] ) 0.01-0.4 M, pH > 3), all the primary free
radicals formed by the radiation are converted into NO2 and
•
(13) Bielski, B. H. J.; Cabelli, D. E.; Arudi, R. L.; Ross, A. B. J. Phys.
Chem. Ref. Data 1985, 14, 1041.
(14) Forni, L. G.; Mora-Arellano, V. O.; Packer, J. E.; Willson, R. L. J.
Chem. Soc. Perin Trans. 2 1986, 1.
(12) Ross, A. B.; Mallard, W. G.; Helman, W. P.; Buxton, J. V.; Huie, R.
E.; Neta, P. NIST Standard References Database 40, Version 2.0;
NIST: Washington, DC, 1994.

4158 Inorganic Chemistry, Vol. 36, No. 19, 1997
Goldstein and Czapski
Figure 1. Absorbances of a pulse-irradiated air-saturated solution
containing 0.01 M formate,11.25 mM nitrate, and 100 µM EDTA at
pH 6.8 (1 mM phosphate buffer): absorbance measured 40 µs after
the pulse (0); absorbance measured 1 ms after the pulse ((); residual
absorbance corrected for excess O2 (if all NO2 radicals react with
O2•-, [O2•-]0 ) 10.2 µM and [•NO2]0 ) [O2NOO-]T ) 6.07 µM) (b).
The optical path length was 12.1 cm, and the dose was 27.6 Gy.
Figure 3. Decay plots: observed rate constant for the decay of
peroxynitrate in the presence of 10 mM formate and 12-50 mM nitrate
as a function of pH (b). The solid curves were calculated using kd )
1.0 s-1 (alkaline) and 4.6 × 10-3 s-1 7 or 7 × 10-4 s-1 8 (acid) and pKa
) 5.9. Oxidation plots: observed rate constant for the oxidation of
•-
•
4-
0.1-1 mM Fe(CN)6 in the presence of 0.1 M formate and 30-50
mM nitrate (0) and of the oxidation of 0.1-0.275 mM NADH in the
presence of 30 mM formate and nitrate (+) by peroxynitrate as a
function of pH. The solid curve for the oxidation rates was calculated
using kf ) 1.0 s-1 (alkaline) and 0.05 s-1 (acid) and pKa ) 5.6.
Figure 2. Absorbance measured at 310 nm 1-2 ms after the end of
the pulse as a function of pH. All solutions were air-saturated and
contained 0.01 M formate, 100 µM EDTA, and 11.25 mM nitrate. The
optical path length was 12.1 cm, and the dose was 20.6 Gy. The solid
curve was calculated using pKa ) 5.8, OD310(O2NOO-) ) 0.066, and
OD310(O2NOOH) ) 0.
-
Table 1. Yield of O2NOO- as a Function of [O2]
Figure 4. Observed rate constant for the formation of I3 in the
oxidation of I- by peroxynitrate as a function of [I-]. All solutions
contained 30 mM nitrate and 0.1-0.4 M formate at pH 4.9 (b), 5.9
(1), 6.4 (]), and 6.8 (0). The dose was 7 Gy.
[O2], M
OD320
G(NO32-) ) RGe
R/OD320
2.4 × 10-4
4.5 × 10-4
7.5 × 10-4
1.2 × 10-3
0.069
0.061
0.051
0.046
0.985Ge
0.971Ge
0.953Ge
0.927Ge
14.27
15.92
18.68
21.22
Our results are is in agreement with the earlier reported values
of ꢀ285max(O2NOO-) ) 1650 ( 100 M-1 cm-1, pKa ) 5.85 (
0.1, and kd(alkali) ) 1.0 ( 0.2 s-1 8
.
a All solutions contained 10 mM formate and 30 mM nitrate at pH
6.8. The dose was 27.6 Gy, and the optical path length was 12.1 cm.
In conclusion, we have shown that peroxynitrate is formed
within less than 2 ms after the irradiation of aerated solutions
containing relatively low concentrations of formate and nitrate.
This method is very useful for studying the mechanism of the
oxidation of various substrates by peroxynitrate.
The decay of O2NOOH/O2NOO- was studied at pH 3.8-
10. Below pH 5 repetitive pulsing was used to produce
detectable amounts of O2NOOH at 264 nm. The decay rate
was first order and decreased with the decrease in pH (Figure
3), indicating that O2NOOH is relatively stable as compared to
O2NOO-. We determined the rate constant of the decay of
O2NOO- to be 1.0 ( 0.1 s-1. The best fit to the experimental
data given in Figure 3 was obtained for kd ) 1.0 s-1 (alkali), kd
) 7 × 10-4-4.6 × 10-3 s-1 (acid),7,8 and pKa ) 5.9 ( 0.1.
Oxidation of Iodide by Peroxynitrate. When aerated
solutions containing 30 mM nitrate, 0.1-0.4 M formate, and
-
0.5-20 mM iodide were irradiated, I3 was formed. The
stoichiometry and the kinetics of the reaction were studied by
following the formation of I3 at 352 nm, using ꢀ352(I3-) )
-
25 800 M-1 cm-1, and 710 M-1 for the stability constant of

Reactivity of Peroxynitric Acid
Inorganic Chemistry, Vol. 36, No. 19, 1997 4159
Figure 5. Observed rate constant for the oxidation of iodide by
peroxynitrate as a function of pH. All solutions contained 30 mM nitrate
and 0.1-0.4 M formate. The dose was 7 Gy. The solid curve was
calculated with the assumption that O2NOO- does not oxidize iodide
and by using k(O2NOOH + I-) ) 840 M-1 s-1 and pKa ) 6.
Figure 6. Oxidation yield of ferricyanide as a function of pH and
ferrocyanide concentrations at various pH’s. The dose was 21.5-27
Gy. The yield of peroxynitrate was calculated according to eq 15.
- 15
I3
.
The formation of I3- obeyed first-order kinetics and was
faster than the self-decomposition of peroxynitrate at pH 3.45-
7.3. The observed first-order rate constant was linearly de-
pendent on [I-]0 (Figure 4) and was highly pH dependent,
resulting in an apparent pKa ) 6.0 ( 0.1 (Figure 5). These
results show that the oxidation of iodide takes place through
O2NOOH, and that the rate constant of this reaction is 840 (
50 M-1 s-1
.
The stoichiometry of the reaction was determined in the
presence of 0.5-2 mM iodide at pH 3.45-4.9. The total yield
-
of I3 (G(I3-)T ) G(I2) + G(I3-)) was obtained by using the
stability constant of I3- or by adding 0.3 M iodide to the solution
after the irradiation to convert all I2 to I3-. The total yield of
-
I3 obtained by both methods was found to be identical, and
under various experimental conditions G(I3-)T ) G(O2NOOH)T.
Oxidation of Ferrocyanide by Peroxynitrate. The oxida-
tion of ferrocyanide by peroxynitrate was followed at 420 nm
(ꢀ420 ) 1000 M-1 cm-1). In this system, reactions 17-19 may
interfere with the determination of the stoichiometry and the
kinetics.
Figure 7. Oxidation yield of NADH as a function of [NADH]0 at pH
4.95 and 5.9. Aerated solutions contained 30 mM formate and 30 mM
formate. The dose was 27.6 Gy. The yield of peroxynitrate was
calculated according to eq 15.
Fe(CN)64- + •OH f Fe(CN)63- + OH-
(17)
(18)
(19)
k17 ) 1.1 × 1010 M-1 s-1 12
ferricyanide at acid pH’s increases slightly with the increase in
[Fe(CN)64-]0, reaching a plateau value at high yields. Under
these conditions, it was found to be highly pH-dependent (Figure
3). The yields of ferricyanide were dependent on pH and
[Fe(CN)64-]0 (Figure 6). The stoichiometry of the reaction was
determined under acidic solutions to be [Fe(CN)63-]/[O2NOOH]
) 2.0 ( 0.2.
-
Fe(CN)64- + •NO2 f Fe(CN)63- + NO2
k18 ) 2.1 × 106 M-1 s-1 16
•
-
Fe(CN)64- + HO2 f Fe(CN)63- + HO2
Oxidation of NADH by Peroxynitrate. The oxidation of
NADH by peroxynitrate was studied by following the decrease
in the absorbance at 340 nm using a cell with a 1 cm optical
path length. The concentration of NADH could not be raised
above 0.2 mM, due to total absorbance of the light by the
solution. In addition, the reaction could not be studied below
pH 4.5 because of decomposition of NADH in acidic solutions.
k19 ) 1.6 × 105 M-1 s-1 12
Under our experimental conditions where [HCO2-]/[Fe-
(CN)63-] > 30, reaction 17 can be neglected. However, because
of the competition of reactions 18 and 19 with reactions 12
and 13, the concentration of ferrocyanide must not exceed 1
mM at pH >4 and 0.4 mM at pH <4.
When aerated solutions containing 30 mM nitrate, 30 mM
formate, and 1.3 × 10-5-2 × 10-4 M NADH were irradiated
at pH 4.5-7.1; the rate of the decay of the absorbance at 340
nm was first order. The observed rate constants increased
slightly with the increase in [NADH]0, reaching a plateau value
The rate of formation of ferricyanide obeyed first-order
kinetics. The observed rate constant for the formation of
(15) Schwartz, H. A.; Bielski, B. H. J. J. Phys. Chem. 1986, 90, 1445.
(16) Goldstein, S.; Czapski, G. J. Am. Chem. Soc. 1995, 117, 12078.

4160 Inorganic Chemistry, Vol. 36, No. 19, 1997
Goldstein and Czapski
at high yields. The latter values were highly pH-dependent and
within experimental error identical to the values obtained in
the presence of ferrocyanide (Figure 3). The oxidation yields
increased with the increase in [NADH]0 and the decrease in
pH, as in the case of ferrocyanide, but did not reach plateau
values (Figure 7). In the presence of oxygen, a chain reaction
takes place and the stoichiometry cannot be determined. (NAD•,
which is the product of the one-electron oxidation of NADH,
Both mechanisms II and III are followed by reactions 28-30,
and mechanism III is also followed by reactions 9 and 10.
•-
I• + I- f I2
(28)
(29)
k28 ) 1.1 × 1010 M-1 s-1 12
•
reacts rapidly with oxygen to form NAD+ and HO2•,17 and HO2
oxidizes another molecule of NADH to NAD•.13)
I2•- + I2•- f I3- + I-
Discussion
Mechanism of Direct Oxidation by Peroxynitrate. Our
results show that the oxidation of iodide takes place directly
through O2NOOH and that the stoichiometry of this process is
given by eq 20. The mechanism of the oxidation of iodide by
k29 ) 2.3 × 109 M-1 s-1 12
-
I2 + I- h I3
(30)
O2NOOH + 3I- + H+ f I3- + NO3- + H2O (20)
K30 ) 710 M-1 15
peroxynitrate can take place through outer-sphere electron-
transfer mechanisms (mechanisms I-III) or through an inner-
sphere electron-transfer mechanism (mechanism IV).
The present results rule out mechanism III because, in the
presence of 0.1 M formate and 0.5 mM iodide, •NO3 would not
mechanism I
be scavenged by formate (k ) 2 × 105 M-1 s-1 12
) and 0.5 mM
-
O2NOOH + I- f I• + •HO2 + NO2
(21)
iodide would not compete efficiently with the hydrolysis of
•NO3 to •OH, which in the presence of excess of formate forms
•
superoxide. Because HO2 does not oxidize iodide,12
Since the oxidation of iodide by superoxide does not compete
efficiently with the dismutation of superoxide,12,13
G(O2NOOH)T/G(I3-)T will be 0.5, in contrast to the observed
value of unity. Therefore, reaction 21 cannot describe the first
stage of this process.
G(O2NOOH)T/G(I3-)T will not exceed 0.5, in contrast to the
observed value of unity.
mechanism IV
–
I
O2NOOH + I–
mechanism II
(31)
O2NOO
H
-
O2NOOH + I- f I• + •NO2 + HO2
k22 ) 840 M-1 s-1
(22)
(23)
(24)
k31 = 840 M–1 s–1
–
I
fast
–
O2NOO
HOI + NO3
H
HOI + I– + H+
(32)
I2 + H2O
-
•NO2 + I- f I• + NO2
K32 = 2 x 1012 M–2 15
k23 ) 1.1 × 105 M-1 s-1 12
In conclusion, the kinetics and the stoichiometry results
demonstrate that the direct oxidation of iodide by peroxynitrate
is consistent with either mechanism II or IV.
-
H2O2 + NO2- f H2O + NO3
Mechanism of Indirect Oxidation by Peroxynitrate. The
rate of the oxidation of ferrocyanide and NADH by peroxynitrate
in the presence of sufficient concentrations of these substrates
is zero order in the substrate concentration. The observed rate
constants at pH >6 are within experimental error identical to
those of the self-decay of peroxynitrate, whereas at acid solutions
they are considerably higher (Figure 3). The best fit obtained
for the observed oxidation rates is for kf ) 1.0 s-1 (alkaline),
0.05 s-1 (acid) and pKa ) 5.6 ( 0.1 (Figure 3), whereas the
best fit for the rate of the self-decay of peroxynitrate is for kd
) 1.0 s-1 (alkaline), 7 × 10-4-4.6 × 10-3 s-1 (acid)6-8 and
pKa ) 5.9 ( 0.1 (Figure 3).
k24 ) 4.6 × 103[H+] M-1 s-1 18
mechanism III
O2NOOH + I- f I• + •NO3 + OH-
k25 ) 840 M-1 s-1
(25)
(26)
(27)
-
•NO3 + I- f I• + NO3
k26 > 4.0 × 109 M-1 s-1 19
(17) Willson, R. L. Chem. Commun. 1970, 1005.
•NO3 + H2O f HNO3 + •OH
(18) Damschen, D. E.; Martin, L. R. Atmos. EnViron. 1983, 17, 2005.
(19) The rate constant of reaction 26 has not yet been determined, whereas
those for reactions of •NO3 with chloride and bromide were determined
to be 4 × 107 and 4 × 109 M-1 s-1, respectively.12 Therefore, k26 is
expected to be higher than the rate constant for the bromide reaction.
k27 ) 2.9 × 107 s-1 12

Reactivity of Peroxynitric Acid
Inorganic Chemistry, Vol. 36, No. 19, 1997 4161
It was previously suggested that the decomposition of
peroxynitrate takes place via the reactions8
oxidation yields on pH and [S]0 (S ) substrate) (Figures 6 and
7) indicates that the mechanism of the indirect oxidation by
peroxynitrate is not as simple as the radical mechanism. If the
O2NOO- f NO2- + O2
(33)
•
•
oxidation of a substrate takes place only through HO2 and -
NO2, the oxidation yields will decrease with the increase in pH.
Under the conditions where the rates of the oxidation of S by
k33 ) 1.0 s-1
•
•
HO2 and NO2 compete efficiently with the dismutation of
superoxide and with the hydrolysis of NO2, the yield of S+
•
O2NOOH f HNO2 + O2
(34)
will be given by eq 38. Thus, at pH > pKa, the oxidation yields
k34 ) 7 × 10-4 s-1
or via the radical mechanism6,7
G(S+)
)
G(O2NOOH)T
2k-13[H+]/(Ka + [H+])
k-13[H+]/(Ka + [H+]) + k33Ka/(Ka + [H+])
•
O2NOOH h •NO2 + HO2
O2NOO- f NO2- + O2
(-13)
(33)
)
2k-13[H+]
(38)
•
-
k-13[H+] + k33Ka
HO2 + O2•- f O2 + HO2
(35)
reduce to zero, independent of ferrocyanide concentrations,
which is in contrast to the experimental results given in Figure
6.
k35 pH-dependent13
2•NO2 h N2O4
We therefore suggest that peroxynitrate decomposes through
(36)
the formation of O2NOOH* and O2NOO-* as oxidizing
•
k36 ) 4.5 × 108 M-1 s-1
k-36 ) 6.9 × 103 s-1 12
intermediates, where the former is in equilibrium with HO2
•
and NO2 and the latter decomposes into nitrite and oxygen
(Scheme 1).
N2O4 + H2O f NO3- + NO2- + 2H+
(37)
Our suggested mechanism fits the results of Lammel et al.,7
who showed that the yield of nitrite decreased with the decrease
in pH. The yield of HNO2 in acidic solutions will depend on
the relative rates of reactions 40 and 41 because k-40 ) 1.8 ×
109 M-1 s-1,8 and reaction -40 competes efficiently with the
k37 ) 1 × 103 s-1 12
Both mechanisms predict that the rate of the decay of perox-
ynitrate will be first order and pH-dependent. However,
according to reactions 33 and 34, the products of the decom-
position process will be oxygen and nitrite at all pH’s, whereas
the radical mechanism predicts the decrease in the yield of nitrite
and oxygen and an increase in the yield of nitrate with a decrease
in pH.
•
hydrolysis of •NO2 and the dismutation of HO2 . The contribu-
tion of reaction 41 to the decomposition process requires a
detailed study on product yields (nitrite, nitrate, and oxygen)
as a function of pH. Our suggested mechanism also shows that
H2O2 is consumed by O2NOOH, though it predicts that less
than one molecule of H2O2 will be consumed by one molecule
of O2NOOH, which is in agreement with a recent study where
only about 0.6 mol of H2O2 was consumed by 1 mol of
O2NOOH.9
Lammel et al.7 found that, above pH 5, the decomposition of
peroxynitrate yields predominately nitrite, whereas in acidic
solutions the yield of nitrite decreased to about 8%. The yield
of nitrate increased with the decrease in pH, but the results are
inaccurate due to nitric acid impurity.7 This observation
suggests a change in the mechanism of decomposition when
the pH decreases, which supports the radical mechanism.
However, Logager and Sehested8 determined the oxidation
capacity of a mixture of excess of H2O2 over O2NOOH at pH
2 and found that one molecule of O2NOOH consumed one
molecule of H2O2. If the radical mechanism were correct, H2O2
would not be consumed (reactions -13, and 35-37 followed
by reaction 24), whereas in the nonradical mechanism (reaction
34 followed by reaction 24), one molecule of O2NOOH
consumes one molecule of H2O2. However, the nonradical
mechanism (reactions 33 and 34) cannot explain the decrease
in the yield of nitrite with the decrease in pH, and the indirect
oxidation of ferrocyanide and NADH by peroxynitrate as nitrite
and oxygen does not oxidize these substrates under our
experimental conditions.
In the presence of an efficient scavenger of O2NOOH*, kobs
•
) k39, and when the oxidation takes place through HO2 and
•NO2, kobs ) k39(k40 + k41)/(k-39 + k40 + k41).
Conclusions
Peroxynitrate is formed within less than 2 ms after the
irradiation of aerated solutions containing formate and nitrate.
This method is very useful for studying the mechanism of the
oxidation by peroxynitrate.
Our suggested mechanism for the decomposition of perox-
ynitrate (Scheme 1) explains the following features: (i) The
observed rate constant of the decay of peroxynitrate is first order
and highly pH-dependent. (ii) The decomposition of perox-
ynitrate at pH >5 yields mainly nitrite and oxygen. The yield
of nitrite decreases and that of nitrate increases in acidic
Scheme 1
Our results show that the observed rate constant of the
oxidation of ferrocyanide and NADH by O2NOOH equals 0.05
s-1, which is considerably higher than the rate of the self-
decomposition of peroxynitrate in acidic solutions (Figure 3).6-8
k41
HNO2 + O2
k40
k39
•NO2 + HO2
•
O2NOOH
pKa
O2NOOH*
pKa*
4-
Considering the radical mechanism, if a scavenger (Fe(CN)6
k–39
k–40
•
•
or NADH) reacts with NO2 and HO2 , the decay rate of
O2NOOH in the presence of sufficient concentrations of these
substrates may be determined by k-13 ) 0.05 s-1, and the
stoichiometry will be 2. However, the dependence of the
k42
k43
O2NOO–
O2NOO–*
NO2– + O2
k–42

4162 Inorganic Chemistry, Vol. 36, No. 19, 1997
Goldstein and Czapski
solutions. (iii) Peroxynitrate may oxidize the substrates directly
through O2NOOH in a reaction that is first order in peroxynitrate
and first order in the substrate, e.g., iodide. (iv) Peroxynitrate
may oxidize the substrates in a reaction that is first order in
peroxynitrate and zero order in the substrate, e.g., ferrocyanide
and NADH. The indirect oxidation by peroxynitrate may take
(v) The oxidation yields at sufficient concentrations of the
substrates approach 100% via the direct and indirect oxidation
pathways.
Acknowledgment. This research was supported by Grant
4129 from the council for Tobacco Research and by the Israel
Science Foundation.
•
•
place through O2NOO-*, O2NOOH*, NO2, or HO2 , all of
which are formed during the decomposition of peroxynitrate.
IC961186Z