On the chemical and electrochemical one-electron reduction of peroxynitrous acid

Article abstract of DOI:10.1021/jp046123d

Peroxynitrous acid was reduced by cathodic linear sweep voltammetry at a gold electrode and by iodide at pH 3.2 and 5.6. The cathodic reduction wave was identified by measuring its decay in time, which was the same as observed by optical spectroscopy. The iodide oxidation was followed by optical measurement of the triiodide formation. Both reductions show one-electron stoichiometry, with the product n_αα = 0.23 ± 0.04 from the electrochemical experiments, in which α is the transfer coefficient and na the number of electrons transferred, and an diiodine yield of ca. 0.5 equiv per equivalent of peroxynitrous acid. The voltammetric reduction was irreversible up to scan rates of 80 V s^-1. Both reductions were pH independent in the range studied. The voltammetric reduction is most likely an irreversible elemental reaction followed by a chemical decay that cannot be observed directly. Because of the pH independence, we conclude that both reductions have a common short-lived intermediate, namely [HOONO]^-. We estimate the electrode potential of the likely ONOOH/ONOOH^- couple to be larger than 1 V. The commonly used electrode potential E°(ONOOH, H ⁺/NO₂, H₂O) does not describe the chemistry of peroxynitrous acid.

Full text of DOI:10.1021/jp046123d

J. Phys. Chem. A 2005, 109, 965-969
965
On the Chemical and Electrochemical One-Electron Reduction of Peroxynitrous Acid
†
‡
†
†
Christophe Kurz, Xiuqiong Zeng, Stefan Hannemann, Reinhard Kissner, and
Willem H. Koppenol*,†
Laboratorium f u¨ r Anorganische Chemie, Department of Chemistry and Applied Biosciences,
ETH Zurich, CH 8093 Zurich, Switzerland, and Department of Chemistry at Xixi Campus, Zhejiang UniVersity,
Hangzhou 310028, China
ReceiVed: August 27, 2004; In Final Form: NoVember 3, 2004
Peroxynitrous acid was reduced by cathodic linear sweep voltammetry at a gold electrode and by iodide at
pH 3.2 and 5.6. The cathodic reduction wave was identified by measuring its decay in time, which was the
same as observed by optical spectroscopy. The iodide oxidation was followed by optical measurement of the
triiodide formation. Both reductions show one-electron stoichiometry, with the product n
from the electrochemical experiments, in which R is the transfer coefficient and n the number of electrons
transferred, and an diiodine yield of ca. 0.5 equiv per equivalent of peroxynitrous acid. The voltammetric
R
R ) 0.23 ( 0.04
R
-
1
reduction was irreversible up to scan rates of 80 V s . Both reductions were pH independent in the range
studied. The voltammetric reduction is most likely an irreversible elemental reaction followed by a chemical
decay that cannot be observed directly. Because of the pH independence, we conclude that both reductions
•
-
have a common short-lived intermediate, namely [HOONO] . We estimate the electrode potential of the
•-
likely ONOOH/ONOOH couple to be larger than 1 V. The commonly used electrode potential E°(ONOOH,
+
•
2 2
H /NO , H O) does not describe the chemistry of peroxynitrous acid.
Introduction
Micromolar concentrations of both nitrogen monoxide and
1
superoxide are produced by activated macrophages during the
immune response. Given the diffusion-controlled rate of 1.6 ×
10
-1 -1 2
10
M
s , this is likely to lead to formation of peroxynitrite.
Because the pKa of peroxynitrous acid is ca. 6.8, depending on
3
-5
temperature, buffer composition, and concentration,
both
peroxynitrous acid and peroxynitrite can be involved in deleteri-
ous reactions. Thus, peroxynitrous acid and/or its deprotonated
5
-8
9-14
15,16
form oxidize thiols, nitrate tyrosines,
and tryptophan,
and initiate lipid peroxidation.17,18 Both one-electron oxidations
and oxygen-transfer reactions have been reported.1
9-21
The one-electron electrode potential of peroxynitrous acid is
a crucial parameter to predict whether peroxynitrous acid can
oxidize a compound specifically in a one-electron step, a process
that potentially could yield free radicals. Values of 2.14 and
2
2
2
.0 V were published by Mer e´ nyi and Lind and Koppenol
Figure 1. The equipment for stopped-flow electrochemistry consists
of two pumps, controlled by computer, an electrochemical cell, a
potentiostat, and a computer for the collection and evaluation of the
data.
3
and Kissner, respectively, and refer to standard conditions.
These values are based on determinations and estimates of the
standard Gibbs energy of formation of peroxynitrous acid. To
our knowledge, no direct electrochemical investigation of
peroxynitrous acid has been reported.
•
Louis, MO) were analytical grade or better. NO (Linde/
Unterschleissheim, Germany) was of 99.95% purity. Peroxyni-
The oxidation of the peroxynitrite anion has been studied
electrochemically and an E°′ value of 0.51 ( 0.02 V has been
24
trite was synthesized according to Koppenol et al. and Bohle
2
5
•
-
23
et al. The concentration of the peroxynitrite solution was
reported for the ONOO /ONOO couple.
determined by measuring the absorbance at 302 nm (ꢀ302 ) 1705
-
1
-1 25
M
cm ). Deionized water was purified further by a
Experimental Section
Millipore Milli-Q unit (Millipore/Bedford, MA). Buffers were
prepared from the reagents above.
Reagents. Na2HPO4, NaH2PO4, NaNO2 (Fluka/Buchs, Swit-
zerland), H3PO4, NaOH (Siegfried/Zofingen, Switzerland), NaI
Cyclic Voltammetry. Cyclic voltammograms were measured
by a stopped-flow technique. The apparatus setup is shown in
Figure 1. One pump (Dosimat 665 titration unit from Metrohm/
Herisau, Switzerland) was filled with a 0.2 M NaH2PO4 solution
and the second one with peroxynitrite solution in 0.01 M NaOH.
The pump units were activated simultaneously by a computer.
(Merck/Darmstadt, Germany), and KO2 (Sigma-Aldrich/St.
*
Corresponding author. E-mail: koppenol@inorg.chem.ethz.ch.
Phone: +41-1-6322875, Fax: +41-1-6321090.
†
‡
ETH Zurich.
Zhejiang University.
1
0.1021/jp046123d CCC: $30.25 © 2005 American Chemical Society
Published on Web 01/26/2005

966 J. Phys. Chem. A, Vol. 109, No. 6, 2005
Kurz et al.
Figure 3. Peak currents vs square root of the scan rate: 1 mM
peroxynitrite with 0.1 M phosphoric acid/phosphate buffers, gold
electrode; circles, pH 3.2; diamonds, pH 5.6.
Figure 2. Details of the measuring cell.
The titration devices were both set to a flow of 20 mL/min,
and the measuring cell (Figure 2) was flushed with 20 times its
volume. The measurements were started by a trigger signal sent
by the pump control unit to the measuring computer. The
measuring equipment consisted of a computer with a 60 kHz
12 bit analog to digital converter, an AMEL 2049 potentiostat,
and an AMEL 586 function generator (Amel/Milan, Italy). The
working electrode was a planar gold disk from Metrohm
2
26
(
Herisau, Switzerland) with an active surface of 1.2 mm . The
reference electrode was an Ag/AgCl electrode from Methrom
Herisau, Switzerland) filled with 0.1 M KCl and 0.1 M NaNO3
(
as bridge electrolyte, with a electrode potential of 0.270 V vs
normal hydrogen electrode (25 °C). The auxiliary electrode was
a glassy carbon electrode from Metrohm (Herisau, Switzerland).
Stopped-Flow Spectrophotometrie Studies. Kinetics experi-
ments were carried out at 25 °C and ambient pressure with an
Applied Photophysics SX 17MV (Leatherhead, Surrey, UK)
stopped-flow spectrophotometer operated in the symmetric
mixing mode. The reaction was initiated by mixing 10, 50, or
Figure 4. Reduction waves of peroxynitrous acid at pH 3.2 and 5.6
and several scan rates: 1 mM peroxynitrous acid, 0.1 M phosphoric
-
1
-1
acid/phosphate buffers; (a) pH 5.6, 1 V s ; (b) pH 3.2, 1 V s ; (c)
-
1
-1
-1
pH 5.6, 4 V s ; (d) pH 3.2, 4 V s ; (e) pH 5.6, 10 V s ; (f) pH 3.2,
-
1
1
0 V s .
1
00 µM peroxynitrite in 0.01 M sodium hydroxide with various
reduction waves at more negative potential. These waves are
most probably caused by the reduction of decomposition
products, possibly nitrogen dioxide and dinitrogen tetraoxide.
Nitrite was ruled out, because it yielded no cathodic wave under
sodium iodide concentrations (200 µM to 40 mM) in 0.2 M
phosphate buffer at pH 5.8 and at pH 3.2. The changes in
absorbance were followed at 350 nm and the pH of the mixture
was determined at the outlet of the stopped-flow apparatus with
a glass electrode. The equilibrium constant of the triiodide
formation I2 + I S I3 was determined to be 930 M in the
phosphate buffer system used, somewhat different from the 710
M
-
1
the same conditions. Cyclic scans up to 80 V s showed no
reoxidation wave. The negative peak potential shift for a 10-
-
-
-1
-
1
fold increase in scan rate from 1 to 10 V s was 190 mV at
pH 3.2 and 210 mV at pH 5.6. The peak currents were always
directly proportional to the initial peroxynitrous acid concentra-
tion and to the square root of the scan rate at both pH values
-1
27
found in the literature, which is a value obtained by
extrapolation to zero ionic strength. The extinction coefficient
of I3 at 350 nm was 28 100 M cm , in good agreement
with the literature.
-
-1
-1
(Figure 3). The peak current-scan rate relation indicates that
2
7
the reductions observed were caused by species under diffusion
control and not adsorbed to the electrode surface. The current
peaks at pH 3.2 were higher than at pH 5.6 (Figure 4) for the
same peroxynitrous acid concentrations and scan rates, but
hardly shifted by the different pH. The variation of maximum
current strength with pH was caused by a change in ohmic
resistance of the metal-solution interface, which depends on
the buffer composition. At pH 5.6, there is a considerable
concentration of monohydrogen phosphate present, while at pH
3.2, dihydrogen phosphate is the only anionic species.
Results
Voltammetry. Electrochemical reduction of peroxynitrous
acid at a polycrystalline gold electrode with a linear sweep
yielded a cathodic wave that could be assigned unambiguously
to peroxynitrous acid, since the peak current at a given scan
rate decreased with the time between the formation of the acid
by mixing peroxynitrite with dihydrogen phosphate/phosphoric
acid and the onset of the voltammetric scan. The rate constant
-
1
for the first-order decrease in peak current was (1.1 ( 0.1) s
Stopped-Flow. Peroxynitrous acid oxidizes iodide to diiodine
at pH 5.6 and (1.5 ( 0.1) s-1 at pH 3.2, quite typical for
like other peroxides.
28,29
However, unlike other peracids, it
4
peroxynitrous acid isomerization at room temperature. The
cannot be determined quantitatively by iodometric titration, since
the relative yield of the reaction depends on the concentrations
reduction wave of peroxynitrous acid was followed by further

One-Electron Reduction of ONOOH
J. Phys. Chem. A, Vol. 109, No. 6, 2005 967
+
+
HOONO + H S [H OONO]
(1)
2
which is supported by the slightly increased isomerization rate
3
3
of peroxynitrous acid at pH <3. An estimated pKa of about 2
for equilibrium 1 is compatible with this observation. The
hydronation equilibrium would be reached much faster than the
electrochemical reduction, and therefore the hydronation equi-
librium should shift the peak potential position by
∆
E ) (RT/Rn F) ln(K/K + 1)
(2)
R
with K ) Ì
+
/ÌHOONO
H OONO
2
from the value without a preequilibrium. K is the equilibrium
constant for a given buffer pH and pKa ) 2. With the assumption
of R ) 0.5, the peak potential at pH 3.2 should differ by about
200 mV from that at 5.6 with nR ) 1, or by 100 mV with nR )
2, at the same scan rates. Experimentally this was not found:
the peaks appeared at nearly identical potentials at both pH
values (Figure 4). The peak potential difference was only 15
Figure 5. Relative yields of diiodine from the oxidation of iodide by
peroxynitrous acid vs total iodide concentration: 0.1 M phosphoric
acid/phosphate buffers.; circles, [HOONO]initial ) 25 µM; squares,
-
1
mV at 10 V s and smaller at lower scan rates. The lack of a
significant pH-dependent peak shift rules out a hydronation
equilibrium. Another preceding equilibrium could be the ho-
molysis of peroxynitrous acid, which is claimed to take place
[HOONO]initial ) 5 µM.
3
4,35
involved. With iodide concentrations below 1 mM, but in at
least 5-fold excess over peroxynitrous acid, precisely 0.5 equiv
of diiodine was produced. Only with iodide concentrations
below 100 µM did the isomerization of the peroxynitrous acid
compete with the iodide oxidation, and the diiodine yield
dropped below 0.5 equiv. At iodide concentrations above 1 mM,
the diiodine yield rose above 0.5 equiv and approached, but
never reached, 1 equiv (Figure 5). The rate law of the reaction
was first order in both iodide and peroxynitrous acid, which
means that only these molecules are involved in the rate-
determining step. The second-order rate constant at pH 3.2 was
at a rate of one-third of the isomerization rate:
•
•
HOONO S NO + OH
(3)
2
The rate of this reaction would be almost pH independent below
the pKa of peroxynitrous acid. It is, however, rather unlikely
that the cathodic current measured was caused by the reduction
of hydroxyl radicals, when one considers that their steady-state
concentration would be very low due to fast recombination
reactions that result in nitric acid, peroxynitrous acid, dinitrogen
tetraoxide, and hydrogen peroxide. The shape of the cathodic
peak and its potential shift with scan rate imply a kinetically
inhibited reduction, which is not what one expects for hydroxyl
radicals. Since there were no other reactants present, a purely
irreversible mechanism of type a is consistent with the data
obtained. Linear sweep peaks of reaction type a can be analyzed
4
-1 -1
(
2.8 ( 0.3) × 10 M s , based on the rate of formation of
-
diiodine, which was obtained by converting the I3 absorbance
vs time trace with the law of mass action of the equilibrium I2
I S I3 into the diiodine concentration vs time. The reaction
-
-
+
rate increased somewhat from pH 5.6 to 3.2, and this small
change can be attributed to the fact that at pH 5.6 about 10%
3
6
for RnR by means of the relation |Ep - Ep/2| ) 1.857RT/RnRF.
-
The average of RnR, in which R is the transfer coefficient and
nR the number of electrons transferred, for 11 peaks recorded
at pH 3.2 yields RnR ) 0.26 ( 0.06, and for 9 peaks at pH 5.6,
RnR ) 0.27 ( 0.03. Plots of ln(ip) vs Ep have slopes with RnR
of the peroxynitrite is in the less reactive ONOO form.
Discussion
)
0.21 ( 0.01 for pH 5.6 and RnR ) 0.22 ( 0.02 for pH 3.2,
The complete lack of reoxidation wave and the linear
dependence of the peak current on the square root of the scan
in good agreement with the values derived from |Ep - Ep/2|.
An R of 0.5 implies a perfectly symmetric activation barrier
and is representative for most reversible electrochemical reac-
tions. An R below 0.4 is typical for irreversible reactions, and
an R below 0.2 is rarely observed. From the experimentally
observed product RnR ) 0.26 we conclude that nR equals 1.
The rate-determining step in the electrochemical reduction of
peroxynitrous acid is obviously the transfer of an electron to
the molecule. If the electron transfer would yield a moderately
stable product, the transfer would be at least quasireversible;
since this is not the case it must be followed by a very fast and
irreversible decay, eqs 4-6,
3
0
rate rule out a quasireversible electrochemical reaction. The
remaining conceivable mechanisms are (a) an irreversible
electrochemical reaction Ei, (b) an irreversible electrochemical
reaction Ei preceded by a chemical equilibrium, CrEi, or (c) a
reversible electrochemical reaction followed by an irreversible
chemical process, ErCi. Process c would show no reoxidation
wave only if the kinetic parameter λ ) (kc/V)(RT/nF), in which
kc is the chemical rate constant based on mole fractions and V
the scan rate, were larger than 0.1.31 The shift of the negative
peak potential with a 10-fold increase in V should be 29.6 mV/n
for a truly reversible electrode reaction. In the case of a
quasireversible reduction, the shift should not exceed 59.1 mV/
-
•-
HOONO + e S [HOONO]
(4)
(5)
3
2
n. The experimentally determined shift of about 200 mV is
far too high to account for any reduction of type c. The huge
peak potential shift suggests nR ) 1 for the rate-determining
electrochemical step and an R smaller than 0.5. Processes a and
b remain. As regards possibility b, a preceding equilibrium could
be the hydronation of peroxynitrous acid, eq 1,
[
HOONO]•- f OH + NO
-
•
2
[HOONO]•- + H f H O + NO
+
•
(6)
2
2
so that no quasireversible signal can be observed with ordinary

968 J. Phys. Chem. A, Vol. 109, No. 6, 2005
Kurz et al.
•
-
3
voltammetry equipment. The formation of [HOONO] , other-
Kissner, respectively. On the basis of the results presented here,
eq 11 does not reflect the chemistry of peroxynitrous acid.
Given the irreversibility of the electrode processes, it is not
possible to estimate the electrode potential of the couple
wise named (hydridodioxido)oxonitrate(•1-), as an intermediate
•
2-
is not that farfetched: it has been shown that NO2
can be
produced from nitrite and hydrated electrons.3
7-43
•-
The iodide oxidation by peroxynitrous acid is hardly pH
dependent in the pH range where peroxynitrite is extensively
hydronated. The rate law for the iodide oxidation is d[I2]/dt )
ONOOH/[ONOOH] strictly on the basis of the electrochemical
experiments. However, the observation that iodide and hexachlor-
oiridate(III) are oxidized by peroxynitrous acid suggests that
this electrode potential is larger than 1 V.
4
5
-
k[HOONO][I ]. This rate law, the rate constant, and the yields
are in agreement with earlier observations by Goldstein and
Czapski,28 but these relative yields are not compatible with the
Acknowledgment. Supported by the ETH and the Schweiz-
erische Nationalfonds.
29
mechanism they postulated later. If the iodide oxidation were
a two-electron process as suggested, a yield of 1 equiv of
diiodine per equivalent of peroxynitrous should be attainable
at sufficient iodide excess, which is not the case (Figure 5).
Since the same rate law was found under a variety of
experimental conditions, the combined observations imply that
a single electron transfer to peroxynitrous acid is rate limiting,
as is the case in the electrochemical reduction. Therefore, it is
rather probable that the same intermediate is formed as in the
electrochemical reduction, namely (hydridodioxido)oxonitrate-
References and Notes
•-
•
(
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