THE JOURNAL OF BIOLOGICAL CHEMISTRY
© 1997 by The American Society for Biochemistry and Molecular Biology, Inc.
Vol. 272, No. 6, Issue of February 7, pp. 3465–3470, 1997
Printed in U.S.A.
Metabolic Fate of Peroxynitrite in Aqueous Solution
REACTION WITH NITRIC OXIDE AND pH-DEPENDENT DECOMPOSITION TO NITRITE AND
OXYGEN IN A 2:1 STOICHIOMETRY*
(Received for publication, July 15, 1996, and in revised form, November 18, 1996)
Silvia Pfeiffer, Antonius C. F. Gorren, Kurt Schmidt, Ernst R. Werner‡, Bernhard Hansert§,
D. Scott Bohle§, and Bernd Mayer¶
From the Institut für Pharmakologie und Toxikologie, Karl-Franzens-Universität Graz, Universitätsplatz 2, A-8010
Graz, Austria, the ‡Institut für Medizinische Chemie und Biochemie, Universität Innsbruck, Fritz-Pregl-Strasse 3,
A-6020 Innsbruck, Austria, and the §Department of Chemistry, University of Wyoming, Laramie, Wyoming 82071-3838
Peroxynitrite, the reaction product of nitric oxide
(NO) and superoxide (O2. ) is assumed to decompose upon
protonation in a first order process via intramolecular
rearrangement to NO2
3 . The present study was carried
out to elucidate the origin of NO2
2 found in decomposed
peroxynitrite solutions. As revealed by stopped-flow
spectroscopy, the decay of peroxynitrite followed firstorder kinetics and exhibited a pKa of 6.8 6 0.1. The
reaction of peroxynitrite with NO was considered as one
possible source of NO2
2 , but the calculated second order
rate constant of 9.1 3 104 M21 s21 is probably too small to
explain NO2
2 formation under physiological conditions.
Moreover, pure peroxynitrite decomposed to NO2
2 without apparent release of NO. Determination of NO2
2 and
NO2
3 in solutions of decomposed peroxynitrite showed
that the relative amount of NO2
2 increased with increasing pH, with NO2
2 accounting for about 30% of decomposition products at pH 7.5 and NO2
3 being the sole metabolite at pH 3.0. Formation of NO2
2 was accompanied by
release of stoichiometric amounts of O2 (0.495 mol/mol of
2
2
NO2
2 ). The two reactions yielding NO2 and NO3 showed
distinct temperature dependences from which a difference in Eact of 26.2 6 0.9 kJ mol21 was calculated. The
present results demonstrate that peroxynitrite decomposes with significant rates to NO2
2 plus O2 at physiological pH. Through formation of biologically active intermediates, this novel pathway of peroxynitrite decomposition may contribute to the physiology and/or cytotoxicity of NO and superoxide.
The reaction between nitric oxide (NO) and superoxide anion
(O2. ) yields peroxynitrite with a second order rate constant near
the diffusion-controlled limit (k 5 4.3– 6.7 3 109 M21 s21) (1, 2).
The reaction constitutes an important sink for O2. because it is
about twice as fast as the maximum velocity of SOD.1 Consequently, peroxynitrite has been implicated in many patho-
* This work was supported by Fonds zur Förderung der Wissenschaftlichen Forschung in Österreich Grants P 10655, P 10859, P 11478
(to B. M.), P 11301 (to E. R. W.), and F 712 (to K. S.) and by National
Institutes of Health Grant GM53828 and Grant-in-aid 94 – 017-580 (to
D. S. B.). The costs of publication of this article were defrayed in part by
the payment of page charges. This article must therefore be hereby
marked “advertisement” in accordance with 18 U.S.C. Section 1734
solely to indicate this fact.
¶ To whom correspondence should be addressed: Institut für Pharmakologie und Toxikologie, Karl-Franzens-Universität Graz, Universitätsplatz 2, A-8010 Graz. Tel.: 43-316-380-5567; Fax: 43-316-380-9890;
and E-mail: mayer@kfunigraz.ac.at.
1
The abbreviations used are: SOD, superoxide dismutase; SIN-1,
3-(4-morpholinyl)-sydnoniminehydrochloride; HPLC; high performance
liquid chromatography; DTPA, diethylenetriaminepentaacetic acid.
This paper is available on line at http://www-jbc.stanford.edu/jbc/
logical conditions including stroke (3), heart disease (4), and
atherosclerosis (5, 6). The potential cellular targets for peroxynitrite cytotoxicity include the antioxidants ascorbate, a-tocopherol, and uric acid (7–10), protein and non-protein sulfhydryls (11), DNA (12), and membrane phospholipids (13).
Decomposition of peroxynitrite is complex (14, 15). The anion
is rather stable in alkaline solutions but decomposes rapidly
(t1/2 5 1 s at pH 7.4, 37 °C) upon protonation to peroxynitrous
acid (ONOOH) (pKa 5 6.8) (16). Two pathways of ONOOH
decomposition have been proposed. Some studies have argued
that ONOOH is cleaved homolytically to generate hydroxyl and
NO2 radicals. This hypothesis is based on the sensitivity to
hydroxyl radical scavengers of certain peroxynitrite-induced
reactions, including the formation of malondialdehyde from
deoxyribose and the hydroxylation on the benzene ring of sodium benzoate, phenylalanine, and tyrosine (16, 17). Studies on
decomposition of peroxynitrite by electron paramagnetic resonance spectroscopy with the spin traps 5,5-dimethyl-1-pyrroline N-oxide and 4-pyridyl-1-oxide-N-tert-butylnitrone also provided evidence for the formation of free hydroxyl radicals (18,
19). Against this, Koppenol et al. (15) concluded from molecular
dynamic calculations that homolytic cleavage of ONOOH is
highly improbable. This was reinforced by the independence of
the rate of ONOOH decomposition on solvent viscosity (20).
Based on these results, it was suggested that decomposition of
ONOOH to NO2
3 involves formation of an activated intermediate (ONOOH*), which might account for the hydroxyl radicallike properties of peroxynitrite (15, 21).
There are several methods for the detection of peroxynitrite
in biological systems. Since ONOOH decomposition yields an
intermediate that nitrates phenolic compounds (22, 23), presence of nitrotyrosine in proteins was proposed to be evidence of
peroxynitrite production in tissues (24). However, using both a
monoclonal antibody specifically recognizing peroxynitritemodified proteins (24) as well as a published HPLC method
(17), we failed to detect tyrosine nitration by authentic peroxynitrite at concentrations ,0.1 mM.2 Spectrophotometric determination of dihydrorhodamine 123 oxidation was described
as another sensitive assay for the specific detection of peroxynitrite at submicromolar concentrations (25), but in our hands,
interference of several redox-active compounds precluded application of this method in cell-free assay systems.3 Under
certain experimental conditions, indirect evidence for peroxynitrite production can be obtained by comparing NO release in
the absence and presence of SOD. The peroxynitrite donor
compound SIN-1, for example, does not release detectable
3465
2
3
S. Pfeiffer, and B. Mayer, unpublished observations.
P. Klatt, and B. Mayer, unpublished observations.
3466
Decomposition of Peroxynitrite
amounts of free NO unless SOD is present in amounts sufficient to outcompete the reaction with concomitantly produced
O2. (26). Based on similar results obtained with purified neuronal NO synthase, we suggested that the enzyme generates NO
and O2. simultaneously and hence functions as peroxynitrite
synthase if incubated in vitro (27). However, in contrast with
the widely held view that peroxynitrite decomposes exclusively
2
to NO2
3 , considerable amounts of NO2 were also found as a
major stable product of SIN-1 or NO synthase under physiological conditions.2 Similarly, excess NO2
2 formation was observed in peroxynitrite producing cells (28), suggesting that
additional as yet unidentified reactions contribute to peroxynitrite decomposition.
The present study was done to elucidate the fate of peroxynitrite in aqueous solution. Studies with the authentic compound, prepared in two different ways, identified a reaction
leading to release of NO2
2 and O2 in a 2:1 stoichiometry as a
route of peroxynitrite decomposition at pH $ 7.5.
EXPERIMENTAL PROCEDURES
Materials—NO solutions were prepared by dissolving NO gas (Linde
München, Germany, 99% pure) in deoxygenated water as described
previously (29). All solutions were prepared freshly each day with
Nano-pure water (Barnstead ultrafiltered type I, resistance .18
megaohms cm21). Sulfanilamide, sodium nitrite, cadmium, and the
Griess-Ilosvay reagent for postcolumn derivatization were from Merck,
Darmstadt, Germany. All other chemicals were from Sigma, Vienna,
Austria.
Synthesis of Peroxynitrite—Alkaline solutions of peroxynitrite (80 –
100 mM) were prepared from acidified NO2
2 and H2O2 according to the
Baeyer-Villinger reaction (30) and quantified spectrophotometrically
using an extinction coefficient of 1670 M21 cm21 (26, 30, 31). Stock
solutions were diluted with H2O to 10 mM immediately before the
experiments. The tetramethylammonium salt of peroxynitrite
([Me4N][ONOO]) was synthesized from [Me4N] [O2. ] and NO as described previously (32). Purity of the sample was ascertained spectrophotometrically in aqueous solution at pH 14, and magnetic susceptibility with a Faraday balance indicated that there were no detectable
paramagnetic (O2. ) impurities present. Purity of [Me4N][ONOO] was
also checked by 15N NMR spectroscopy, which indicated that no NO2
2
was present. The salt was dissolved in 1 M NaOH to give a 24 mM stock
solution, which was stored at 270 °C and diluted with H2O prior to
experiments. With the exception of stopped-flow kinetics, all experiments described here were initially performed with conventional preparations of peroxynitrite and then repeated with [Me4N][ONOO] to
exclude that the results were due to unidentified contaminants.
Kinetic Experiments—Peroxynitrite decomposition was studied by
stopped-flow absorbance spectroscopy at 302 nm (Bio-Sequential SX17MV stopped-flow ASVD spectrofluorimeter, Applied Photophysics,
Leatherhead, U. K.). For simple decomposition experiments, reservoir 1
contained peroxynitrite in 0.01 M NaOH, and reservoir 2 contained the
buffer solution (at pH 3.0 – 6.0, 1 M acetate buffer; at pH 5.0 –9.0, 1 M
phosphate buffer; at pH 8.0 –10.0, 1 M Tris/HCl; at pH $ 10, solutions of
NaOH). The NaOH concentration in reservoir 1 was, in some cases,
adapted to the requirements of the experiment: non-buffered experiments at pH 3.0, 10.0, and 11.0 were carried out with sufficiently low
concentrations of NaOH.
The reaction of peroxynitrite with NO was studied by sequential
stopped-flow, i.e. reservoirs 1 and 2 were premixed followed by mixing
with contents of reservoir 3 with short delay time (10 ms). Reservoir 4
was used to push the mixed contents of reservoirs 1 and 2 forward into
the main mixing chamber. Reservoir 1 contained buffer (pH 3.0 –11.0;
4 3 final concentration), reservoir 2 contained a solution of peroxynitrite in NaOH (4 3 final concentration; typical final [NaOH] 5 mM),
reservoir 3 contained a saturated solution of NO (;2 mM giving ;1 mM
final concentration), and reservoir 4 contained buffer (2 3 final concentration). To vary NO concentrations, experiments were also done with
2-fold diluted peroxynitrite in reservoir 3 and NO in reservoir 2. This
yields the same final concentration of peroxynitrite but a 2-fold lower
final concentration of NO (;0.5 mM). Samples of the NO solution were
taken with a plastic syringe under helium gas and transferred directly
into the stopped-flow reservoir. Experiments were carried out both with
air-containing buffers and with buffers that had been thoroughly degassed. Degassing made no difference.
Decomposition of Peroxynitrite and Determination of NO22 and
NO32—Unless indicated otherwise, peroxynitrite (1 mM or 0.5 mM) was
decomposed by incubation in 0.1 M phosphate buffer for 1 h at pH
3.0 –9.0. [Me4N][ONOO] (0.25 mM or 0.1 mM) was decomposed in 0.5 M
phosphate buffer under the same conditions. NO2
2 was determined by
the Griess assay. The samples (0.1 ml) were mixed with 10 ml of H2O
and 10 ml of an EDTA solution (0.5 M, pH 8.0), followed by addition of
0.12 ml of freshly prepared Griess reagent (20 mg N-(1-naphthyl)ethylenediamine and 0.2 g sulfanilamide dissolved in 20 ml of 5% (w/v)
phosphoric acid) and measurement of the absorbance at 546 nm. For
2
determination of NO2
2 1 NO3 , samples (0.2 ml) were adjusted to pH
;7.5 and mixed with 20 ml of an aqueous zinc suspension (100 mg/ml)
and 20 ml of an EDTA solution (0.5 M, pH 8.0). Samples were spun down
for 5 min, and 0.12 ml of the supernatant were mixed with 0.12 ml of the
Griess reagent, followed by determination of the absorbance at 546 nm.
2
Calibration curves were established with NO2
2 and NO3 (10 –50 mM
each). The calculated amount of NO2
2 present in stock solutions of
conventionally prepared peroxynitrite agreed well with NO2
2 measured
after decomposition at pH 3.0. This amount was subtracted from the
measured values.
2
The NO2
2 /NO3 data were confirmed by HPLC analysis according to
published protocols (33, 34). 50 ml samples were injected onto a 250 3
4 mm C18 reversed phase column (LiChrospher 100 RP-18, 5 mm
particle size, Merck, Vienna, Austria) and eluted with 5% (w/v) NH4Cl,
pH 7.0, at a flow rate of 0.7 ml/min. NO2
2 was detected by postcolumn
derivatization with the stable Griess-Ilosvay reagent (Merck) (0.7 ml/
min), heating to 60 °C, and measurement of the absorbance at 546 nm.
2
For determination of NO2
2 plus NO3 , samples were reduced with a
cadmium reactor (Cd, 0.3– 0.8 mm, 20 –50 mesh ASTM, Merck, washed
with 0.1 N HCl, and packed in a Pharmacia HR 5/5 glass column) prior
to postcolumn derivatization.
Electrochemical Detection of NO and Oxygen—NO and O2 were
measured with commercially available Clark-type electrodes (Iso-NO
and ISO2, World Precision Instruments, Mauer, Germany) (27). NO and
O2 meters were connected to an Apple Macintosh computer via an
analog to digital (A/D) converter (MacLab, World Precision Instruments). Release of O2 from peroxynitrite was determined in 1.8-ml
water-jacketed vials sealed with a rubber septum and maintained at
37 °C. Experiments were performed in phosphate buffer (0.1 M or 0.5 M,
pH 3.0 –9.0), which had been gassed with argon to reduce the O2
concentration to 20 – 40 mM. Aliquots of peroxynitrite stock solutions
were injected through the septum to give concentrations of 0.5 mM
(conventional peroxynitrite) or 0.25 mM ([Me4N][ONOO]), and the output current was recorded at 0.33 Hz under constant stirring. Two-point
calibration of the sensor was performed in air-saturated H2O at 37 °C
(6.9 ppm; 0.216 mM O2) and argon atmosphere (zero O2).
To study the reaction of peroxynitrite with NO, 4-ml aliquots of an ;2
mM NO solution were injected into 1.8-ml glass vials completely filled
with 0.1 M phosphate buffer, pH 7.4, and sealed with a septum. At the
indicated time points, 1.8 –3.6 ml of peroxynitrite solution (0.5 mM) were
applied to give concentrations of 0.25–1 mM. The output current was
recorded at 1.66 Hz under constant stirring. The sensor was calibrated
with NO2
2 standards according to manufacturer recommendations.
RESULTS
Decomposition of peroxynitrite was monitored as decrease in
absorbance at 302 nm at 20 °C. As expected, decomposition at
pH 3 was very fast and followed first order kinetics with a
calculated rate constant (kcalc) of 0.86 6 0.05 s21 but slowed
down at increasing pH. The kcalc values and corresponding Hill
coefficients summarized in Table I demonstrate that peroxynitrite decay was first order under most conditions although Hill
coefficients smaller than 1.0 were obtained at pH 8.0 (0.67 6
0.02) and pH 11.0 (0.5 6 0.1). Using the Hill equation for
overall kinetic analysis of decomposition at pH 3–11, we calculated a pKa of 6.8 6 0.1, which agrees well with published data
(35). The possible contribution of transition metals to peroxynitrite decomposition was studied with 0.6 mM peroxynitrite in
0.5 M phosphate buffer (pH 7.4) in the presence of Cu(NO3)2,
Fe(NH4)(SO4)2, Fe(NH4)2(SO4)2, and the metal chelator DTPA.
Rates of decomposition were affected neither by the metal salts
(0.1 mM each) nor by DTPA (0.1 and 1 mM). At a concentration
of 2.5 mM DTPA, the peroxynitrite decay rate was enhanced
10-fold.
Stopped-flow data showed that peroxynitrite decomposition
Decomposition of Peroxynitrite
TABLE I
Apparent first-order rate constants of peroxynitrite decomposition
as a function of pH
Decomposition of peroxynitrite (0.1– 0.7 mM) was measured spectrophotometrically at 20 °C as decrease in absorbance at 302 nm at the
av
indicated pH values. The kcalc
were obtained by averaging the apparent
first order rate constants that were calculated by dividing initial rates
by the peroxynitrite concentrations (mean 6 S.D.; n 5 8). Hill coefficients were calculated from the slope of plots of log v0 versus log
[peroxynitrite].
pH
3
4
5
6
7
8
9
10
11
av
kcalc
(s21)
0.86 6 0.05
0.82 6 0.01
0.71 6 0.02
0.61 6 0.01
0.39 6 0.02
0.08 6 0.009
0.0298 6 0.009
0.0033 6 0.0001
0.00008 6 0.00001
Hill coefficient
0.91 6 0.07
1.09 6 0.07
1.16 6 0.05
1.11 6 0.03
0.83 6 0.03
0.67 6 0.02
1.10 6 0.03
1.16 6 0.05
0.5 6 0.1
3467
TABLE II
2
Formation of NO2
2 and NO3 upon decomposition of peroxynitrite in
the presence of metal chelators
Peroxynitrite (1 mM) was decomposed in the presence of the metal
chelators in 0.1 M phosphate buffer, pH 7.4, at 37 °C for 1 h. NO2
2 and
NO2
3 were determined as described under “Experimental Procedures.”
Data are mean values 6 S.E. of three experiments performed in
duplicate.
Chelator
NO2
2
0.1 mM
%
Control
EDTA
Neocuproine
Cuprizone
Bathophenanthroline
DTPA
32.4 6 6.3
27.7 6 5.9
29.0 6 5.4
32.8 6 2.2
32.8 6 7.3
32.6 6 2.5
FIG. 1. Consumption of NO by peroxynitrite. At time point zero,
4 ml of a saturated NO solution (;2 mM) was added to 1.8 ml of 0.1 M
phosphate buffer, pH 7.4, at ambient temperature. The arrows indicate
the time points at which 2.7 ml of peroxynitrite stock solution (containing 1.35 nmol) were applied. Changes in NO concentration were monitored with an NO electrode over 2.5 min. NO autoxidation determined
in separate experiments is shown as a dotted line. The experiment
shown is representative for three.
was faster in the presence of ;1 mM NO and that the increase
in rate was dependent on the NO concentration. However,
calculation of rate constants was difficult because the exact NO
concentrations in these experiments were not known and the
effect of NO was observed only as a relatively small increase of
an already fast reaction. Therefore, we used an NO-sensitive
electrode to measure the consumption of NO by known
amounts of peroxynitrite. Fig. 1 shows a representative trace
obtained by addition of 4 ml of a saturated NO solution to 1.8 ml
of 0.1 M phosphate buffer, followed by two repetitive additions
of peroxynitrite yielding concentrations of 0.75 mM each. Peroxynitrite induced a rapid consumption of NO with initial rates
of 100 6 9 nM s21 and a stoichiometry close to 1:1 (0.75 mM
peroxynitrite consumed 0.66 6 0.06 mM NO). NO consumption
(initial NO concentration 1–2 mM) was linear in the range of
0.25–1 mM peroxynitrite with initial rates ranging from 20 to
167 nM s21 and a rate constant of 9.1 3 104 M21 s21.
We consistently observed that decomposition of peroxynitrite
or [Me4N][ONOO] resulted in formation of about 70% NO2
3 and
2
2
30% NO2
2 at pH 7.4 and 37 °C. As the NO2 /NO3 ratios were not
affected by known metal chelators (Table II), our results do not
support previous suggestions according to which formation of
NO2
2 is due to contamination of peroxynitrite solutions with
trace metals (36) but indicate that NO2
2 release results from an
as yet unrecognized pathway of peroxynitrite decomposition.
2
To address this issue, we measured NO2
2 and NO3 after peroxynitrite decomposition at pH 3–9 and found that the relative
amount of NO2
2 increased with increasing pH (Fig. 2A). Assum-
FIG. 2. Decomposition of peroxynitrite yields NO22 and oxygen. A, peroxynitrite (0.5 mM final initial concentration) was decomposed by incubation in 0.1 M phosphate buffer (pH 3.0 –9.0) at 37 °C for
2
1 h, followed by the determination of NO2
2 , NO3 , and O2 as described
2
under “Experimental Procedures”. NO2
and
NO
2
3 were determined
after measurement of O2 release in the same vials. Data are means 6
S.E. of six experiments. B, correlation between O2 and NO2
2 production
(slope 5 0.495, correlation coefficient 5 0.988).
ing that these results were not due to a reaction of peroxynitrite with contaminants in the stock solutions, our findings led
us to speculate that 2 mol of peroxynitrite decomposed to 2 mol
of NO2
2 and 1 mol of O2. Indeed, using a Clark-type O2 sensor,
we found that the pH-dependent formation of NO2
2 was accompanied by release of stoichiometric amounts of O2 (Fig. 2A). The
replot of the data (Fig. 2B) revealed a correlation coefficient of
0.988 and a slope of 0.495, suggesting that NO2
2 and O2 were
released in a 2:1 stoichiometry.
To corroborate these data and exclude possible artifacts, the
experiments were repeated with [Me4N][ONOO]. Fig. 3A shows
that formation of NO2
2 and O2 increased when [Me4N][ONOO]
(0.25 mM) was decomposed at increasing pH. The linear correlation of NO2
2 versus O2 shown in Fig. 3B yielded a slope of
0.657 and a correlation coefficient of 0.983. Although these data
nicely confirmed the results obtained with the conventional
preparation, two interesting differences were observed. First,
while release of NO2
2 and O2 was negligible when the BaeyerVillinger preparation of peroxynitrite was decomposed at pH
3468
Decomposition of Peroxynitrite
TABLE III
2
Effect of temperature on formation of NO2
2 and NO3
from [Me4N][ONOO]
[Me4N][ONOO] (0.1 mM) was decomposed in 0.1 M phosphate buffer,
2
pH 7.4, at the indicated temperatures for 1 h. NO2
2 and NO3 were
determined as described under “Experimental Procedures.” Data are
mean values 6 S.E. of three experiments performed in duplicate.
Temperature
NO2
2
NO2
3
°C
mM
mM
5
24
37
56
15.0 6 1.1
22.1 6 1.6
32.8 6 0.7
49.2 6 1.9
77.7 6 0.1
68.9 6 1.6
59.4 6 1.4
45.4 6 1.2
% NO2
2
16.2
24.3
35.6
52.1
FIG. 3. Decomposition of [Me4N][ONOO] yields NO2
2 and oxygen. A, [Me4N][ONOO] (0.25 mM final initial concentration) was decomposed by incubation in 0.5 M phosphate buffer (pH 3.0 –9.0) at 37 °C
2
for 1 h, followed by the determination of NO2
2 , NO3 , and O2 as described
2
under “Experimental Procedures.” NO2
2 and NO3 were determined
after measurement of O2 release in the same vials. Data are means 6
S.E. of six experiments. B, correlation between O2 and NO2
2 production
(slope 5 0.657, correlation coefficient 5 0.983).
#6.5 (cf. Fig. 2A), decomposition of [Me4N][ONOO] resulted in
formation of significant amounts of NO2
2 and O2, even at low
pH. Second, release of NO2
2 and O2 from [Me4N][ONOO] did
not level off at high pH values and appeared to account for
virtually 100% of the decomposition occurring at pH 9.0. In all
2
experiments, the measured sum of NO2
2 plus NO3 was close to
theoretical values.
The assumption that two different pathways of decomposi2
tion account for the formation of NO2
2 and NO3 was supported
2
by a pronounced temperature sensitivity of the NO2
2 /NO3 ratio. As shown in Table III, decomposition of 0.1 mM
[Me4N][ONOO] at pH 7.4 yielded 16.2 and 52.1% NO2
2 at 5 and
2
56 °C, respectively. A similar increase of the NO2
/NO
2
3 ratio
was observed with three concentrations (0.1, 0.5, and 1 mM) of
the Baeyer-Villinger preparation (not shown). We also determined the temperature dependence for the overall peroxynitrite decomposition rate between 5 and 50 °C by stopped-flow
spectroscopy. The Arrhenius plots showed a strictly linear relationship between ln kobs and T21 at pH 5.0 and 7.4 (Fig. 4A).
From the slope of the plots, values for Eact of 92.0 6 2 kJ mol21
and 90.0 6 0.8 kJ mol21 were calculated for decomposition at
2
pH 5.0 and 7.4, respectively. Assuming that the NO2
2 /NO3
ratios reflect the kinetic partitioning of the two pathways lead2
ing to NO2
2 and NO3 formation, the difference in Eact of the
reations (DEact) was estimated as 26.2 6 0.9 kJ mol21 (Fig. 4B).
2
The NO2
2 /NO3 ratio was independent of the initial peroxynitrite concentration. Decomposition of 0.01, 0.1, 0.5, and 1 mM
peroxynitrite at pH 7.4 and 37 °C resulted in formation of
29.6 6 1.3, 25.5 6 5.7, 27.3 6 1.2, and 28.9 6 1.5%, respectively, of NO2
2 (mean 6 S.E.; n 5 3 each).
DISCUSSION
The present study was carried out to identify the pathways of
formation of NO2
2 in the course of peroxynitrite decomposition.
Stopped-flow kinetic experiments confirmed that peroxynitrite
FIG. 4. Arrhenius plots of peroxynitrite decomposition yield2
ing NO2
2 and NO3 . A, peroxynitrite decomposition rates between 5 and
50 °C were calculated from first-order fits. Peroxynitrite (final initial
concentration 0.6 mM in 0.01 M NaOH) was mixed with 0.5 M acetate
buffer (pH 3.0) or 0.5 M phosphate buffer (pH 7.4). Data are mean values
2
of three experiments. B, the NO2
2 /NO3 data shown in Table I were
2
replotted as ln((k1)/(k2)) with ((k1)/(k2)) 5 ((%NO2
2 )/(%NO3 )) versus (1/
T). From the slope of the linear plot, the difference between the activa2
tion energies (DEact) of the two reactions yielding NO2
2 and NO3 was
calculated. Data are the mean values of three experiments.
decomposed rapidly upon protonation with a pKa of 6.8. The
first order rate constants calculated for peroxynitrite decomposition at different pH values agreed well with previously published data (11, 15, 37). Under physiological conditions (pH 7.4
and 37 °C), decomposition consistently yielded about 30% NO2
2,
whereas NO2
was
the
sole
product
at
pH
,5.0.
Some
studies
3
indicated that under certain experimental conditions, peroxynitrite does indeed decompose to NO2
2 , but this was attributed to minor side reactions catalyzed by contaminating trace
metals (36, 38). In our hands, NO2
2 release was not metalcatalyzed because it was affected neither by chelators nor by
metal ions (0.1 mM).
Previous studies showed that small amounts of NO are released upon peroxynitrite decomposition under certain conditions (39) and that peroxynitrite reacts with NO according to
Equation 1 (40, 41).
NO 1 ONOO2 3 NO2 1 NO2
2
(Eq. 1)
Lewis et al. (28) observed that activated macrophages release
more NO2
2 than expected and considered Equation 1 as a pos-
Decomposition of Peroxynitrite
FIG. 5. Hypothetical mechanism of peroxynitrite decomposition to nitrite and nitrate.
sible source for excess NO2
2 . To judge whether Equation 1 could
account for NO2
2 formation under physiological conditions, we
studied the reaction of peroxynitrite with NO using both
stopped-flow kinetic spectroscopy and an NO-sensitive electrode. These experiments yielded an apparent second order
rate constant of 9.1 3 104 M21 s21, which is in the same order
of magnitude as the rate constant determined for peroxynitrite
reacting with CO2 (42). However, the concentration of NO that
would be required to render the reaction as fast as peroxynitrite decomposition (0.69 s21 at pH 7.4 and 37 °C) is so high (7.6
mM) that the reaction with NO is probably insignificant under
most physiological conditions (43).
Since pure peroxynitrite decomposed to NO2
2 without detectable release of free NO, we considered additional pathways
that could account for NO2
2 formation. Assuming that there
were no other redox-active reaction partners of peroxynitrite
present, we speculated that two peroxynitrite molecules might
combine to release 2 molecules of NO2
2 and 1 molecule of O2
according to Equation 2.
ONOO2 1 ONOOH 3 HNO2 1 NO2
2 1 O2
(Eq. 2)
This hypothesis was corroborated by determination of NO2
2 and
O2 formed upon decomposition of two different peroxynitrite
preparations. With both products, we obtained linear correlations between NO2
2 and O2 release with slopes close to the
theoretical value of 0.50. At pH 3.0 –11.0, peroxynitrite decomposition was first order at most. Moreover, even though the
2
temperature dependence of the NO2
2 /NO3 ratio clearly indicated that the two pathways have different activation energies
(DEact 5 26.2 6 0.9 kJ mol21), the Arrhenius plot for overall
peroxynitrite decomposition at pH 7.4 (;30% NO2
2 ) was strictly
linear and yielded an Eact value that was virtually identical to
that observed at pH 5.0 (;10% NO2
2 ). The Eact of overall peroxynitrite decomposition was found to be rather high (92 6 2
and 90.0 6 0.8 kJ mol21 at pH 5.0 and 7.4, respectively). A
similar value (77.5 kJ mol21 at pH 5.0) was reported by
Koppenol et al. (15).
From these observations, we conclude that the rate-limiting
step in both reactions is the same, a conformational change of
ONOOH to an activated intermediate that either rearranges to
HNO3 (15, 35) or undergoes a reaction with peroxynitrite anion
to yield NO2
2 and O2 (this study). Any potential model must
account for the thoroughly characterized kinetics of peroxynitrite decomposition as well as the stoichiometries of the end
products. Further, a bimolecular rate law for either of the
product determining steps is excluded because the partitioning
of the two pathways does not depend on the concentration of
peroxynitrite. Fig. 5 shows a hypothetical pathway of peroxynitrite decomposition that appears to be most consistent with the
data presented both here and in the literature. According to
3469
this scheme, activated ONOOH can either isomerize to NO2
3 or
decompose to HO and NO2 radicals. At alkaline pH, the OH
radical may react with peroxynitrite anion yielding O2, NO,
and OH2, and NO could react with NO2 radicals to yield N2O3
and finally nitrite.
The novel pathway of peroxynitrite decomposition described
here could have important physiological consequences, as it
possibly involves generation of intermediates with biological
activities not attributed so far to peroxynitrite. In a recent
paper, it was reported that peroxynitrite decomposition could
lead to release of singlet O2 (44). If that observation were due
to the novel reaction proposed here, peroxynitrite-dependent
toxicity might be mediated by singlet O2 toxicity under certain
pathophysiological conditions. Alternatively, decomposition to
H NO2
2 and O2 may be responsible for the observed NO-like
biological activity of peroxynitrite. At pH 7.4, peroxynitrite
oxidizes hemoglobin to methemoglobin with an efficiency of
about 20% (26), and it is tempting to speculate that this reaction represents scavenging by hemoglobin of the NO that is
formed as intermediate during decomposition to NO2
2 and O2
(Fig. 5). Also, our working hypothesis involves intermediary
formation of N2O3, a potent nitrosating agent that could account for the observed peroxynitrite-induced nitrosation of
GSH, especially in light of our findings that the nitrosation
reaction has a pronounced pH dependence and does not occur
at significant rates below pH 7.5 (45). Accordingly, reactive
intermediates formed in the course decomposition to NO2
2 and
O2 could be responsible for stimulation of soluble guanylyl
cyclase by peroxynitrite (45), resulting in cyclic GMP-mediated
biological effects such as vascular smooth muscle relaxation
and inhibition of platelet aggregation (46, 47).
Acknowledgments—We thank an anonymous referee for suggesting
the scheme of peroxynitrite decomposition, Margit Rehn for excellent
technical assistance, and Dr. B. Hemmens for critical reading of this
manuscript.
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