Photon-initiated homolysis of peroxynitrous acid

Article abstract of DOI:10.1021/ic900614e

Laser flash photolysis of ONOOH at 355 nm and a pH of 4.0-5.5 causes homolysis of ONOOH nearly exclusively at the N-O bond rather than at the O-O bond (HO₂^?/HO^? > 25:1). All of the NO^? and HO₂^?

Full text of DOI:10.1021/ic900614e

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Inorg. Chem. 2009, 48, 7307–7312 7307
DOI: 10.1021/ic900614e
Photon-Initiated Homolysis of Peroxynitrous Acid
::
Manuel Sturzbecher-Hohne, Thomas Nauser, Reinhard Kissner, and Willem H. Koppenol*
Institute of Inorganic Chemistry, Department of Chemistry and Applied Biosciences, ETH Zurich, 8093 Zurich,
Switzerland
Received March 30, 2009
Laser flash photolysis of ONOOH at 355 nm and a pH of 4.0-5.5 causes homolysis of ONOOH nearly exclusively
•
•
at the N-O bond rather than at the O-O bond (HO2 /HO• > 25:1). All of the NO• and HO2 radicals formed
by photolysis subsequently recombine with a rate constant of (1.2 ( 0.2) ꢀ 1010 M-1 s-1 via second-order kinetics,
as demonstrated by the return of the UV/vis absorbance to initial levels. Excitation at 266 nm also yields HO2 and NO•,
•
but after recombination, the absorbance levels are lower than initial values, possibly because HO• produced by the
photolysis of water reacts with ONOOH. When NO3-, the product of the ONOOH isomerization, is photolyzed, the
ONOO- formed is rapidly protonated with a second-order rate constant of (1.7 ( 0.8) ꢀ 1010 M-1 s-1. The ONOOH
decays to the starting material, NO3-, with a first-order rate constant of 1.2 s-1. The quantum yield for the photon-
initiated homolysis is 15% for both ONOOH and ONOO-. We conclude that the ON-OOH and ON-OO- bond
dissociation energies are similar.
Introduction
Peroxynitrous acid, ONOOH,1 a strong oxidant as
with a rate constant of 1.2 s-1 (reaction 1).8 The pKa of
ONOOH is 6.5-6.8 (reaction 2), depending on the ionic
strength of the solution.8
well as a nitrating agent, may be formed in vivo2,3 by the
•-
diffusion-limited reaction of O2 and the cellular messen-
-
ONOOH ¼ NO3 þHþ
ð1Þ
ð2Þ
ð3Þ
ð4Þ
ð5Þ
ð6Þ
ger NO•, k = (1.6 ( 0.3) ꢀ 1010 M-1 s-1.4 ONOOH and
ONOO- react with a variety of biomolecules5,6 and are
implicated in a variety of diseases and disorders, includ-
ing atherosclerosis, inflammation, and neurodegenera-
tive disorders.5,7
ONOOH ¼ ONOO-þHþ
ONOO- is fairly stable in aqueous media at pH values
ONOO- ¼ NO•þO2
• -
-
above 10. In the protonated form, ONOOH decays to NO3
•
ONOOH ¼ NO2 þHO•
*To whom correspondence should be addressed. Phone: þ41-44-632-
2875. Fax þ41-44-632-1090. E-mail: koppenol@inorg.chem.ethz.ch.
(1) Formulae, systematic names, and trivial names (in italics): ONOO-,
oxidoperoxidonitrate(1-), peroxynitrite; ONOOH, (hydridodioxido)oxido-
ONOOH ¼ NO•þHO2
•
•
nitrogen, peroxynitrous acid; O2•-, dioxide(•1-), superoxide; HO2 , hydri-
dodioxygen(•), perhydroxyl radical or hydroperoxyl radical; O2, dioxygen;
O
•-, oxide(•1-); NO•, oxidonitrogen(•) or nitrogen monoxide, nitric oxide;
NO2•, dioxidonitrogen(•) or nitrogen dioxide, pernitric oxide; N2O4, tetra-
oxidodinitrogen, nitrogen tetroxide; NO3-, trioxidonitrate(1-), nitrate;
HO•, hydridooxygen(•), hydroxyl radical; Cl2•-, dichloride(•1-), dichlorine
radical anion. Connelly, N. G.; Damhus, T.; Hartshorn, R. M.; Hutton, A. T.
Nomenclature of Inorganic Chemistry. IUPAC Recommendations 2005; Royal
Society of Chemistry: Cambridge, U.K., 2005.
•
• -
HO2 ¼ O2 þHþ
Homolysis of ONOO- at the N-O bond (reaction 3) takes
place with a rate constant of 0.020 ( 0.001 s-1.9 From this rate
constant and the known rate constant for reaction 3, (1.6 ( 0.3)
ꢀ 1010 M-1 s-1 4
, we derived standard Gibbs energies of
(2) Beckman, J. S.; Beckman, T. W.; Chen, J.; Marshall, P. A.; Freeman,
B. A. Proc. Natl. Acad. Sci. U. S. A. 1990, 87, 1620–1624.
(3) Beckman, J. S.; Koppenol, W. H. Am. J. Physiol. Cell Physiol. 1996,
271, C1424–C1437.
formation of ONOO- (þ68 ( 1 kJ mol-1) and ONOOH
(4) Nauser, T.; Koppenol, W. H. J. Phys. Chem. A 2002, 106, 4084–4086.
(5) Alvarez, B.; Radi, R. Amino Acids 2003, 25, 295–311.
(6) Szabo, C.; Ischiropoulos, H.; Radi, R. Nat. Rev. Drug Discovery 2007,
6, 662–680.
(8) Kissner, R.; Nauser, T.; Bugnon, P.; Lye, P. G.; Koppenol, W. H.
Chem. Res. Toxicol. 1997, 10, 1285–1292.
(9) Sturzbecher, M.; Kissner, R.; Nauser, T.; Koppenol, W. H. Inorg.
Chem. 2007, 46, 10655–10658.
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r
2009 American Chemical Society
Published on Web 06/25/2009
pubs.acs.org/IC

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7308 Inorganic Chemistry, Vol. 48, No. 15, 2009
Sturzbecher-Hohne et al.
(þ31 ( 1 kJ mol-1), in excellent agreement with other pub-
lished values.10-13 We summarize the published thermody-
namic values in Figure 1. The Gibbs energies of formation of
ONOOH (þ31 ( 1 kJ mol-1),10,12 NO2• (63 ( 1 kJ mol-1),14
and HO• (26 ( 1 kJ mol-1 15,16
and the published rate constant
)
for reaction 4 (4.5 ꢀ 109 M-1 s-1 17,18
)
have been used to predict
a rate of homolysis of ONOOH at the O-O bond (reaction 4) of
0.38 ( 0.25 s-1 10,11,19
However, this approach is valid only
.
when no branching of the reaction coordinate occurs, which is
not the case here: the reaction of HO• with NO2 yields two
•
Figure 1. The reaction coordinate for the photolysis of NO3- showing
products, namely, ONOOH and NO3-/Hþ.18 Since we have
searched for, but have been unable to find, evidence of
significant levels of formation of HO• from ONOOH,20-24 we
attempted to stimulate reaction 4 by laser flash photolysis.
A pathway of homolysis of ONOOH, distinct from reaction
Gibbs energies of formation of ONOOH, ONOO-, and NO3 and
-
possible products of homolysis and experimentally determined activation
energy barriers. Activation energies and the barriers for homolysis of
ONOO- to NO• and O2•- and isomerization of ONOOH to NO3 are
-
shown in red. Gibbs energies of reaction are shown in black.
4, leads to the formation of NO• and HO2 (reaction 5). The
•
All reagent gases were obtained from PanGas (Dagmer-
sellen, CH). LiONOO was prepared as previously
described.33 All other chemicals were obtained at the
highest purity available and used as received. A Milli-
pore Milli-Q unit (Molsheim, F) was used to purify
deionized water.
energies of these products of reaction lie 20 kJ mol-1 above those
of the products of reaction 4 (Figure 1): ΔfG°(NO•) = 102 ( 0.2
kJ mol-1 is calculated from ΔfG°(NO•)g = þ86.57 kJ mol-1 25
and Henry’s constant, 1.92 ꢀ 10-3 M 0.100 MPa-1 26 at 25 °C;
•
ΔfG°(HO2 ) = 6.4 ( 1.0 kJ mol-1 is based on E°(O2/O2•-) =
350 ( 11 mV27 and pKa,6 = 4.8 ( 0.1 for HO2•.28 With k-5
=
Instrumentation. UV/vis spectra were recorded with a
double-beam Analytik Jena Specord 200 (Jena, D). Rapid
mixing was achieved by the stopped-flow mixing unit
from Applied Photophysics, SX 17MV (Leatherhead,
Surrey, Great Britain), that was operated in the sym-
metric mixing mode interfaced to an Applied Photo-
physics LKS 50 instrument (Leatherhead, Surrey, Great
Britain). The quartz cell was asymmetric with a 10 mm
optical path length, a 2 mm laser path length, and a total
volume of 0.08 cm3; measurements were performed at
25 °C, maintained with a thermostat. Optical changes
were collected and stored on a WaveRunner 64Xi
digital oscilloscope from Lecroy (Chestnut Ridge, NY).
The bandwidth chosen for these experiments was
20 MHz, and the sampling rate was 200 MHz. The third
(λ=355 nm) or fourth (λ=266 nm) harmonic of a Quantel
Brilliant B Nd:YAG laser (Les Ulis Cedex, F) was used
with a pulse duration of 6 ns and spot size of
9 mm. The laser energies were determined with a PRD-J
peak-reading joulemeter from Gentec Inc. (Sainte-Foy,
Quebec, Canada). The pH was measured with a Metrohm
glass electrode (Herisau, Switzerland) interfaced with a
901 microprocessor analyzer from Orion Research, Inc.
(Cambridge, MA).
3.2 ꢀ 109 M-1 s-129 and the Gibbs energies of formation of
ONOOH, NO•,andHO2•,wecalculatek5=(0.8(1.0) ꢀ 10-4s-1,
which is more than 4 orders of magnitude lower than k4.
•
•
We report here that NO• and HO2 , but not NO2 and HO•,
are formed by the photolysis of ONOOH. The observed
products recombine rapidly to form ONOOH.
Materials and Methods
Chemicals. (Me4N)ONOO was prepared from NO• and
(Me4N)O2 according to the method of Bohle et al.30-32
ꢀ
(10) Merenyi, G.; Lind, J. Chem. Res. Toxicol. 1998, 11, 243–246.
ꢀ
(11) Goldstein, S.; Czapski, G.; Lind, J.; Merenyi, G. Chem. Res. Toxicol.
2001, 14, 657–660.
ꢀ
(12) Merenyi, G.; Lind, J.; Czapski, G.; Goldstein, S. Inorg. Chem. 2003,
42, 3796–3800.
(13) Lymar, S. V.; Poskrebyshev, G. A. J. Phys. Chem. A 2003, 107, 7991–
7996.
(14) Stanbury, D. M. Adv. Inorg. Chem. 1989, 33, 69–138.
(15) Schwarz, H. A.; Dodson, R. W. J. Phys. Chem. 1984, 88, 3643–3647.
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(16) Klaning, U. K.; Sehested, K.; Holcman, J. J. Phys. Chem. 1985, 89,
760–763.
(17) Logager, T.; Sehested, K. J. Phys. Chem. 1993, 97, 6664–6669.
ꢀ
(18) Merenyi, G.; Lind, J.; Goldstein, S.; Czapski, G. J. Phys. Chem. A
1999, 103, 5685–5691.
(19) Lymar, S. V.; Khairutdinov, R. F.; Hurst, J. K. Inorg. Chem. 2003,
42, 5259–5266.
Methods. Photolysis of NO3-. Argon-saturated 1 M
(20) Kissner, R.; Koppenol, W. H. J. Am. Chem. Soc. 2002, 124, 234–239.
::
(21) Maurer, P.; Thomas, C. F.; Kissner, R.; Ruegger, H.; Greter, O.;
::
-
NO3 solutions were excited at 266 nm at pH 2.9-4.4,
Rothlisberger, U.; Koppenol, W. H. J. Phys. Chem. A 2003, 107, 1763–1769.
and subsequent reactions were followed at 260 and 330 nm.
Highly concentrated solutions of NO3- were used because
of the low absorbance of NO3- at 266 nm (ε266=1.5 M-1
cm-1). ONOO (M = (Me4N)þ, Liþ) solutions were freshly
prepared in 10 mM MOH (M = Kþ, Liþ), kept in the dark
on ice.
(22) Kissner, R.; Nauser, T.; Kurz, C.; Koppenol, W. H. IUBMB Life
2003, 55, 567–572.
(23) Kissner, R.; Thomas, C.; Hamsa, M. S. A.; van Eldik, R.; Koppenol,
W. H. J. Phys. Chem. A 2003, 107, 11261–11263.
(24) Kurz, C.; Zeng, X.; Hannemann, S.; Kissner, R.; Koppenol, W. H. J.
Phys. Chem. A 2005, 109, 965–969.
(25) Wagman, D. D.; Evans, W. H.; Parker, V. B.; Schumm, R. H.;
Halow, I.; Bailey, S. M.; Churney, K. L.; Nuttal, R. L. J. Phys. Chem. Ref.
Data 1982, 11(Suppl. 2), 37–38.
(26) Wilhelm, E.; Battino, R.; Wilcock, R. J. Chem. Rev. 1977, 77, 219–
262.
(27) Wardman, P. Free Radical Res. Commun. 1991, 14, 57–67.
(28) Bielski, B. H. J.; Cabelli, D. E.; Arudi, R. L. J. Phys. Chem. Ref. Data
1985, 14, 1041–1100.
Photolysis of ONOO-. Argon-saturated solutions of
ONOO- at ca. 2 mM in 10 mM MOH were mixed by
stopped-flow techniques at 25 °C with either 20 mM
H3PO4 or 100 mM pivalate buffer as a proton source to
final pH values of 2-5.5. After 5 ms of mixing, excitation
of ONOOH was achieved with a laser pulse of 266 or
355 nm with energies of 10-140 mJ/pulse (266 nm) or
(29) Goldstein, S.; Czapski, G. Free Radical Biol. Med. 1995, 19, 505–510.
(30) Bohle, D. S.; Hansert, B.; Paulson, S. C.; Smith, B. D. J. Am. Chem.
Soc. 1994, 116, 7423–7424.
(31) Bohle, D. S.; Glassbrenner, P. A.; Hansert, B. Methods Enzymol.
1996, 269, 302–311.
(32) Bohle, D. S.; Sagan, E. S. Inorg. Synth. 2004, 34, 36–42.
::
(33) Sturzbecher-Hohne, M.; Kissner, R.; Nauser, T.; Koppenol, W. H.
Chem. Res. Toxicol. 2008, 21, 2257–2259.

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Article
Inorganic Chemistry, Vol. 48, No. 15, 2009 7309
Table 1. Molar Absorptivities of Relevant Species, Used to Calculate Yields and
Spectra in Figures 3 and 5
Excitation λ (nm)
Observation λ (nm)
266, ε
355,
ε (M-1
260,
ε (M-1
cm-1
330,
ε (M-1
400,
ε (M-1
cm-1
(M-1
cm-1
)
cm-1
)
)
cm-1
)
)
ONOOH8
ONOO-8
HO•45
450
680
70
500
500
475
480
430
454
540
1940
>10
270
750
80
850
8
71
HO•46
HO•47
HO2
135
30
•-28,35
•-28,35
O2
NO•
NO2
>10
150
300
>10
200
•48
36,49
N2O4
10-160 mJ/pulse (355 nm), and reactions were followed
at 260, 330, and 400 nm. The pH was recorded after
mixing. Transient absorbance changes at pH 4.8 were
recorded by changing the wavelength settings of the
monochromator in 5 nm increments from 235-300 nm.
The concentration of HO•, generated in photolysis ex-
periments, was determined by mixing ONOO- in a 1:1
ratio with NaCl at different concentrations in 20 mM
HCl before excitation with detection of the transient
Cl2 at 340 nm (ε340 = 8800 M-1 cm-1 34); NaCl
•-
Figure 2. Photolysis (355 nm) of ONOOH, prepared by mixing
2.06 mM LiONOO in 10 mM LiOH, with a 0.1 M pivalate buffer, at
pH 4.1-5.4 and 25 °C. (a) Absorbance increase at 260 nm upon irradia-
tion andsubsequentdecay attributed tosecond-order recombination. The
signal distortion at 58 μs was also observed in the absence of ONOOH.
(b) Absorbance at 260 nm (red data points) measured directly after
irradiation (140 mJ/pulse), with the species distributions of HO2• (dashed
solutions photolyzed in 10 mM HCl served as controls.
Results
Photolysis of NO3-. We investigated the photolysis of
NO3-, the product of isomerization of ONOOH, in the
presence of the cations Liþ, Naþ, Kþ, Csþ, and (Me4N)þ.
We did not use RbNO3, because it is not sufficiently
•-
line) and O2 (solid line) calculated on the basis of pKa = 4.8. (c)
•
Concentration of HO2 /O2•-, based on absorbance at 260 nm, as a
function of laser energy.
-
soluble in aqueous solution. All alkali metal salts of NO3
behave similarly: during excitation at pH 2.9-4.4, an
initial increase in absorbance at 330 nm is observed
that is ascribed to the formation of ONOO-, followed
by a rapid decay that is linearly dependent on the Hþ
concentration and represents protonation of ONOO-.
The protonation takes place with a second-order rate
constant of (1.7 ( 0.8) ꢀ 1010 M-1 s-1 (data not shown).
From the absorbance of the initial product and reported
amount of 20 mM H3PO4 (final pH 2.1) and irradiated
with laser pulses at 355 nm. Only small changes in
absorbance were detected: the highest laser energy
used (142 mJ/pulse) resulted in an increase of 0.004
absorbance units at 260 nm relative to the preirradiation
level. Importantly, the subsequent decay in absorbance
to the preirradiation absorbance level obeys second-order
kinetics and is complete after 50 μs. The low signal-to-
noise ratio obviated meaningful evaluation of the traces.
We found larger absorbance changes at higher pH,
and Figure 2a shows data collected in a 100 mM pivalate
buffer at pH 4.1-5.4. The increase in absorbance at
260 nm as a function of pH approximates that expected
absorptivities for ONOO- at 260 nm (ε = 500 M-1 cm-1
)
and 330 nm (ε = 80 M-1 cm-1)8 (Table 1), we calculated
that excitation at 266 nm with a laser energy of 100 mJ/
pulse results in the formation of 10 μM ONOOH.
After the protonation step, we observe a first-order
decay to initial absorbance levels at 260 nm with a rate
constant of 1.2 s-1, which we assign to the isomeriza-
tion of ONOOH to NO3-/Hþ. The behavior of (Me4N)
NO3 was different from that of the alkali metal salts,
with a slower increase in absorbance followed by only a
single decay.
for HO2 , the pKa of which is 4.8 (Figure 2b).28 When the
•
irradiation energy is varied, a strictly linear correlation
•
between the concentration of HO2 /O2•- and laser energy
is found (Figure 2c). Kinetics data obtained in additional
experiments performed with (Me4N)ONOO were similar
to those obtained with LiONOO.
Photolysis of ONOOH. Laser flash photolysis ex-
periments were performed with Liþ as the counterion.
LiONOO (1.76 mM) was rapidly mixed with an equal
We plotted the absorbance measured at 5 nm intervals
between 235 and 300 nm after excitation (355 nm, 125 mJ/
pulse) at pH 4.83 (after mixing) in Figure 3a. At pH 4.8,
•
the concentrations of HO2 and O2•- are the same, and we
used molar absorptivities from the literature28,35 to eval-
uate the concentrations and rate constants to generate the
calculated mixed spectrum shown in Figure 3b. That
(34) Under acidic conditions, Cl- reacts quickly with HO• to form
ultimately the strongly absorbing Cl2•-. This is shown in the reactions
below. The yield of the transient Cl2•- is dependent on the Cl- concentration
and the laser energy. HO• þ Cl- = HOCl•-, HOCl•- þ Hþ = H2O þ Cl•, Cl•
þ Cl- = Cl2•-. Jayson, G. G.; Parsons, B. J.; Swallow, A. J. J. Chem. Soc.,
Faraday Trans. 1 1973, 69, 1597–1607.
(35) Bielski, B. H. J. Photochem. Photobiol. 1978, 28, 645–649.

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7310 Inorganic Chemistry, Vol. 48, No. 15, 2009
Sturzbecher-Hohne et al.
Figure 4. Trapping of HO•. Detection of Cl2 at 340 nm (ε = 8800
•-
M
-1cm-1) formed in the reaction of Cl- with HO• in acidic solution.
Solutions of 1 mM ONOOH (pH 2.1) were irradiated (355 nm, 135 mJ/
pulse) in the presence of 10 mM, 60 mM, and 110 mM chloride (blue
traces). The yields of Cl2•- increase with [Cl-] to a maximum of 0.9 μM.
Control experiments: photolysis of 1.0 mM ONOOH in the absence of
Cl-, orange trace, and photolysis of 10 mM Cl- in the absence of
ONOOH, green trace.
•
To probe for the formation of NO2 , we analyzed the
kinetics of the reaction at 330 and 400 nm, where the
•
absorptivity of NO2 and its dimer N2O4 is higher than
that of ONOOH.8,36 No change in absorbance was
detected at 400 nm, but we observed a transient decrease
in the absorbance at 330 nm that within 5 μs returned to
the initial level (data not shown).
We also attempted to find evidence for the for-
mation of HO• during the photolysis of ONOOH, via
its reaction with Cl- under acidic conditions.34 Solutions
of NaCl at various concentrations at pH 2 (HCl) were
mixed 1:1 with 2.03 mM ONOO- and irradiated. At the
highest Cl- concentration (110 mM) and a laser energy
Figure 3. Spectral changes upon photolysis (355 nm, 125 mJ/
pulse) of 1.1 mM (Me4N)ONOO at pH 4.8 and 25 °C, prepared
from equal volumes of 2.2 mM (Me4N)ONOO in 10 mM KOH
and a 0.1 M pivalate buffer. (a) Spectrum as a function of time after
the pulse, showing apparent second-order decay to initial absorbance
•-
of 135 mJ/pulse, the concentration of Cl2 generated
from HO• is only 0.9 μM. In the absence of Cl-, the
absorbance at 340 nm of a 1.01 mM ONOOH solu-
tion initially decreases, then returns to the initial
level (Figure 4). In the absence of ONOOH, but with
Cl- present, no absorbance changes after excitation at
355 nm were observed.
levelsvia recombinationofHO2 /O2•- withNO•, yield of HO2 /O2•- > 25
•
•
μM. (b) Raw spectral data (gray circles) recorded 0.07 μs after photolysis
•
of ONOOH at pH 4.8. Spectrum of HO2 /O2•- (green line) calculated for
pH = pKa,6 = 4.8;28,35 experimental spectrum of ONOOH (red line),
extrapolated to λ < 250 nm based on an assumed symmetric band;
•
difference spectrum, Δε (gray line) = ε (HO2 /O2•-, green line) - ε
We also performed photolysis of ONOOH at 266 nm
(100 mJ/pulse) at pH 4.83 (after mixing) and found
corroborating evidence for the formation of HO2
(ONOOH, red line). Spectral data corrected for the bleaching of
ONOOH (green points). Data points are derived from the data in a.
Note that the measured absorptivities and, hence, the shape of the
spectra are not dependent on the time elapsed after the flash over the
range of 0.07-5 μs. The data fit a model for exclusive formation of NO•
and HO2• (reaction 5). (c) The observed rate constant as a function of the
reciprocaleffective molarabsorptivity(thesumofthemolarabsorptivities
•
(Figure 5), followed by second-order decay. The concen-
•
•-
tration of HO2 /O2 after irradiation is 125 μM, much
higher than that found after excitation at 355 nm, as
expected since the molar absorptivity of ONOOH
at 266 nm (450 M-1 cm-1) is much higher than at
355 nm, (70 M-1 cm-1).8 While photolysis at 355 nm
is fully reversible, that at 266 nm is not: the absorbance
at 235-310 nm is significantly lower at the end of
the experiment, likely due to a fast reaction of ONOOH
with HO• or H• and eaq-, which originate from the
two-photon excitation of H2O37,38 and, to a minor
degree, by HO• formed in reaction 4. By the irradiation
of acidic Cl- solutions, we determined that HO• is formed
in concentrations up to 5 μM.
•
of HO2 , O2•-, and ONOOH at 235-300 nm) at pH 4.8: a second-order
rate constant of (1.2 ( 0.2) ꢀ 1010 M-1 s-1 for the reaction of HO2• and
NO• to form ONOOH is determined from the slope of the line. Rate
constants are derived from the data in a.
the absorbance maximum of the transient species lies near
•
•-
240-245 nm is evidence for the formation of HO2 /O2
which indicates that irradiation causes homolysis of the
,
N-O bond. The decay of the transient species is expected
to correspond to the reaction of HO2 with NO• to reform
•
ONOOH, which should follow second-order kinetics. The
•
•-
concentration of HO2 /O2 immediately following irra-
diation with a laser energy of 125 mJ/pulse is >25 μM.
The decay rates are plotted as a function of the reciprocal
of the extinction coefficient in Figure 3c, from which we
determine a second-order rate constant of (1.2 ( 0.2) ꢀ
When experiments were carried out over the
wavelength range 235-300 nm, we observed for both
excitation wavelengths that the rate constants increase
•
with decreasing molar absorptivities of HO2 and
1010 M-1 -1
for reaction 5.
s
(37) Reuther, A.; Lauberau, A.; Nikogosyan, D. N. J. Phys. Chem. 1996,
(36) Hug, G. L. Optical Spectra of Nonmetallic Inorganic Transient
Species in Aqueous Solution, NSRDS-NBS 69; National Bureau of Standards:
Washinton, DC, 1981.
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(38) Gorner, H.; Nikogosyan, D. N. J. Photochem. Photobiol. B: Biol.
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