(Received for publication, July 15, 1996, and in revised form, November 18, 1996)
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 NO3
. The present study was
carried out to elucidate the origin of NO2
found in decomposed
peroxynitrite solutions. As revealed by stopped-flow spectroscopy, the
decay of peroxynitrite followed first-order kinetics and exhibited a
pKa of 6.8 ± 0.1. The reaction of
peroxynitrite with NO was considered as one possible source of
NO2
, but the calculated second
order rate constant of 9.1 × 104
M
1 s
1 is probably too small to
explain NO2
formation under
physiological conditions. Moreover, pure peroxynitrite decomposed to
NO2
without apparent release of
NO. Determination of NO2
and
NO3
in solutions of decomposed
peroxynitrite showed that the relative amount of
NO2
increased with increasing
pH, with NO2
accounting for
about 30% of decomposition products at pH 7.5 and
NO3
being the sole metabolite
at pH 3.0. Formation of NO2
was
accompanied by release of stoichiometric amounts of O2
(0.495 mol/mol of NO2
). The two
reactions yielding NO2
and
NO3
showed distinct temperature
dependences from which a difference in Eact of
26.2 ± 0.9 kJ mol
1 was calculated. The present
results demonstrate that peroxynitrite decomposes with significant
rates to NO2
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 = 4.3-6.7 × 109 M
1 s
1) (1,
2). The reaction constitutes an important sink for O
2 because
it is about twice as fast as the maximum velocity of
SOD.1 Consequently, peroxynitrite has
been implicated in many pathological conditions including stroke
(3), heart disease (4), and atherosclerosis (5, 6). The potential
cellular targets for peroxynitrite cytotoxicity include the
antioxidants ascorbate,
-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 = 1 s at pH 7.4, 37 °C) upon
protonation to peroxynitrous acid (ONOOH) (pKa = 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
NO3 involves formation of an activated
intermediate (ONOOH*), which might account for the hydroxyl
radical-like 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 peroxynitrite-modified 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 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
O
2 simultaneously and hence functions as peroxynitrite
synthase if incubated in vitro (27). However, in contrast
with the widely held view that peroxynitrite decomposes exclusively to
NO3
, 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
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 and O2 in a 2:1
stoichiometry as a route of peroxynitrite decomposition at pH
7.5.
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 cm1). 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.
Alkaline solutions of
peroxynitrite (80-100 mM) were prepared from acidified
NO2 and H2O2
according to the Baeyer-Villinger reaction (30) and quantified
spectrophotometrically using an extinction coefficient of 1670 M
1 cm
1 (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] [O
2] 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
(O
2) impurities present. Purity of
[Me4N][ONOO] was also checked by 15N NMR
spectroscopy, which indicated that no
NO2
was present. The salt was
dissolved in 1 M NaOH to give a 24 mM stock
solution, which was stored at
70 °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.
Peroxynitrite decomposition was studied
by stopped-flow absorbance spectroscopy at 302 nm (Bio-Sequential
SX-17MV 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 × final concentration), reservoir 2 contained a solution of peroxynitrite in NaOH (4 × 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 × 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 NO2Unless
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 was
determined by the Griess assay. The samples (0.1 ml) were mixed with 10 µl of H2O and 10 µl 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 determination of
NO2
+ NO3
, samples (0.2 ml) were adjusted to
pH ~7.5 and mixed with 20 µl of an aqueous zinc suspension (100 mg/ml) and 20 µl 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. Calibration curves were established with
NO2
and
NO3
(10-50 µM each).
The calculated amount of NO2
present
in stock solutions of conventionally prepared peroxynitrite agreed well
with NO2
measured after decomposition
at pH 3.0. This amount was subtracted from the measured values.
The
NO2/NO3
data were confirmed by HPLC analysis according to published protocols
(33, 34). 50 µl samples were injected onto a 250 × 4 mm C18
reversed phase column (LiChrospher 100 RP-18, 5 µm 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
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. For determination of
NO2
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.
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 µM. 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-µl 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 µl of
peroxynitrite solution (0.5 mM) were applied to give
concentrations of 0.25-1 µM. The output current was
recorded at 1.66 Hz under constant stirring. The sensor was calibrated
with NO2 standards according to
manufacturer recommendations.
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 ± 0.05 s1 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 ± 0.02) and pH 11.0 (0.5 ± 0.1). Using the Hill equation for overall kinetic analysis of decomposition at pH 3-11, we calculated a
pKa of 6.8 ± 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 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 µl 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 µM each. Peroxynitrite induced a rapid consumption
of NO with initial rates of 100 ± 9 nM
s1 and a stoichiometry close to 1:1 (0.75 µM peroxynitrite consumed 0.66 ± 0.06 µM NO). NO consumption (initial NO concentration 1-2 µM) was linear in the range of 0.25-1 µM
peroxynitrite with initial rates ranging from 20 to 167 nM
s
1 and a rate constant of 9.1 × 104
M
1 s
1.
We consistently observed that decomposition of peroxynitrite or
[Me4N][ONOO] resulted in formation of about 70%
NO3 and 30%
NO2
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
is
due to contamination of peroxynitrite solutions with trace metals (36)
but indicate that NO2
release results
from an as yet unrecognized pathway of peroxynitrite decomposition. To
address this issue, we measured NO2
and NO3
after peroxynitrite
decomposition at pH 3-9 and found that the relative amount of
NO2
increased with increasing pH (Fig.
2A). Assuming 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
and 1 mol
of O2. Indeed, using a Clark-type O2 sensor, we
found that the pH-dependent formation of
NO2
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
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 and O2 increased when
[Me4N][ONOO] (0.25 mM) was decomposed at increasing pH. The linear correlation of
NO2
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
and O2 was negligible
when the Baeyer-Villinger preparation of peroxynitrite was decomposed
at pH
6.5 (cf. Fig. 2A), decomposition of
[Me4N][ONOO] resulted in formation of significant
amounts of NO2
and O2,
even at low pH. Second, release of NO2
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 experiments, the measured sum
of NO2
plus
NO3
was close to theoretical
values.
The assumption that two different pathways of decomposition account for
the formation of NO2 and
NO3
was supported by a pronounced
temperature sensitivity of the NO2
/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
at 5 and 56 °C,
respectively. A similar increase of the
NO2
/NO3
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 T
1
at pH 5.0 and 7.4 (Fig. 4A). From the slope
of the plots, values for Eact of 92.0 ± 2 kJ mol
1 and 90.0 ± 0.8 kJ mol
1 were
calculated for decomposition at pH 5.0 and 7.4, respectively. Assuming
that the
NO2
/NO3
ratios reflect the kinetic partitioning of the two pathways leading to
NO2
and
NO3
formation, the difference in
Eact of the reations
(
Eact) was estimated as 26.2 ± 0.9 kJ
mol
1 (Fig. 4B).
|
The
NO2/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 ± 1.3, 25.5 ± 5.7, 27.3 ± 1.2, and 28.9 ± 1.5%,
respectively, of NO2
(mean ± S.E.; n = 3 each).
The present study was carried out to identify the pathways of
formation of NO2 in the course of
peroxynitrite decomposition. Stopped-flow kinetic experiments confirmed
that peroxynitrite 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
,
whereas NO3
was the sole product at pH
<5.0. Some studies indicated that under certain experimental
conditions, peroxynitrite does indeed decompose to
NO2
, but this was attributed to minor
side reactions catalyzed by contaminating trace metals (36, 38). In our
hands, NO2
release was not
metal-catalyzed 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).
![]() |
(Eq. 1) |
Since pure peroxynitrite decomposed to
NO2 without detectable release of free
NO, we considered additional pathways that could account for
NO2
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
and 1 molecule of
O2 according to Equation 2.
![]() |
(Eq. 2) |
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 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 this scheme,
activated ONOOH can either isomerize to
NO3
or decompose to HO and
NO2 radicals. At alkaline pH, the OH radical may react with
peroxynitrite anion yielding O2, NO, and OH
,
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 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
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
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).
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.