Reaction of cyclic nitroxides with nitrogen dioxide: The intermediacy of the oxoammonium cations

Article abstract of DOI:10.1021/ja035286x

Piperidine and pyrrolidine nitroxides, such as 2,2,6,6-tetramethylpiperidinoxyl (TPO) and 3-carbamoylproxyl (3-CP), respectively, are cell-permeable stable radicals, which effectively protect cells, tissues, isolated organs, and laboratory animals from radical-induced damage. The kinetics and mechanism of their reactions with ^.OH, superoxide, and carbon-centered radicals have been extensively studied, but not with ^.NO₂, although the latter is a key intermediate in cellular nitrosative stress. In this research, ^.NO₂ was generated by pulse radiolysis, and its reactions with TPO, 4-OH-TPO, 4-oxo-TPO, and 3-CP were studied by fast kinetic spectroscopy, either directly or by using ferrocyanide or 2,2′-azinobis(3-ethylbenzothiazoline-6-sulfonate), which effectively scavenge the product of this reaction, the oxoammonium cation. The rate constants for the reactions of ^.NO₂ with these nitroxides were determined to be (7-8) × 10⁸ M^-1 s^-1, independent of the pH over the range 3.9-10.2. These are among the highest rate constants measured for ^.NO₂ and are close to that of the reaction of ^.NO₂ with ^.NO, that is, 1.1 × 10⁹ M^-1 s^-1. The hydroxylamines TPO-H and 4-OH-TPO-H are less reactive toward ^.NO₂, and an upper limit for the rate constant for these reactions was estimated to be 1 × 105 M^-1 s^-1. The kinetics results demonstrate that the reaction of nitroxides with ^.NO₂ proceeds via an inner-sphere electron-transfer mechanism to form the respective oxoammonium cation, which is reduced back to the nitroxide through the oxidation of nitrite to ^.NO₂. Hence, the nitroxide slows down the decomposition of ^.NO₂ into nitrite and nitrate and could serve as a reservoir of ^.NO₂ unless the respective oxoammonium is rapidly scavenged by other reductant. This mechanism can contribute toward the protective effect of nitroxides against reactive nitrogen-derived species, although the oxoammonium cations themselves might oxidize essential cellular targets if they are not scavenged by common biological reductants, such as thiols.

Full text of DOI:10.1021/ja035286x

Published on Web 06/17/2003
Reaction of Cyclic Nitroxides with Nitrogen Dioxide: The
Intermediacy of the Oxoammonium Cations
Sara Goldstein,*,† Amram Samuni,‡ and Angelo Russo§
Contribution from the Department of Physical Chemistry, The Hebrew UniVersity of Jerusalem,
Jerusalem 91904, Israel, Department of Molecular Biology, Hebrew UniVersity-Hadassah
Medical School, P.O. Box 12000, Jerusalem 91120, Israel, and Radiation Biology Branch,
National Cancer Institute, NIH, Bethesda, Maryland 20892
Received March 23, 2003; E-mail: sarag@vms.huji.ac.il
Abstract: Piperidine and pyrrolidine nitroxides, such as 2,2,6,6-tetramethylpiperidinoxyl (TPO) and
3-carbamoylproxyl (3-CP), respectively, are cell-permeable stable radicals, which effectively protect cells,
tissues, isolated organs, and laboratory animals from radical-induced damage. The kinetics and mechanism
‚
of their reactions with OH, superoxide, and carbon-centered radicals have been extensively studied, but
‚
‚
not with NO2, although the latter is a key intermediate in cellular nitrosative stress. In this research, NO2
was generated by pulse radiolysis, and its reactions with TPO, 4-OH-TPO, 4-oxo-TPO, and 3-CP were
studied by fast kinetic spectroscopy, either directly or by using ferrocyanide or 2,2′-azinobis(3-ethylben-
zothiazoline-6-sulfonate), which effectively scavenge the product of this reaction, the oxoammonium cation.
The rate constants for the reactions of ‚NO2 with these nitroxides were determined to be (7-8) × 108 M-1
s-1, independent of the pH over the range 3.9-10.2. These are among the highest rate constants measured
for ‚NO2 and are close to that of the reaction of ‚NO2 with ‚NO, that is, 1.1 × 109 M-1 s-1. The hydroxylamines
‚
TPO-H and 4-OH-TPO-H are less reactive toward NO2, and an upper limit for the rate constant for these
reactions was estimated to be 1 × 105 M-1 s-1. The kinetics results demonstrate that the reaction of nitroxides
with ‚NO2 proceeds via an inner-sphere electron-transfer mechanism to form the respective oxoammonium
‚
cation, which is reduced back to the nitroxide through the oxidation of nitrite to NO2. Hence, the nitroxide
slows down the decomposition of ‚NO2 into nitrite and nitrate and could serve as a reservoir of ‚NO2 unless
the respective oxoammonium is rapidly scavenged by other reductant. This mechanism can contribute
toward the protective effect of nitroxides against reactive nitrogen-derived species, although the oxoam-
monium cations themselves might oxidize essential cellular targets if they are not scavenged by common
biological reductants, such as thiols.
attributed to their ability to catalyze superoxide dismutation,1,8
as well as detoxification of carbon-centered radicals9-11 and
Introduction
The increasing knowledge of the roles played by free radical
in disease processes has expanded the search for more efficient
antioxidants that will diminish radical-induced damage. Cyclic
nitroxides, such as 2,2,6,6-tetramethylpiperidinoxyl (TPO), are
cell-permeable stable radicals, which effectively protect cells,
tissues, isolated organs, and laboratory animals from radical-
induced damage.1-7 Their protective effects have, in part, been
alkoxyl and peroxyl radicals.12,13 The kinetics and mechanisms
of the reactions of nitroxides with nitrogen-derived reactive
species are far from being elucidated. It has been shown that
‚NO reacts with nitronyl nitroxides14-16 and aromatic nitrox-
ides,17,18 but there is no agreement on the mechanism of this
reaction. Likewise, there is no agreement regarding the reactivity
(8) Samuni, A.; Krishna, C. M.; Riesz, P.; Finkelstein, E.; Russo, A. J. Biol.
Chem. 1988, 263, 17921-17924.
† The Hebrew University of Jerusalem.
‡ Hebrew University-Hadassah Medical School.
§ National Cancer Institute.
(9) Beckwith, A. L. J.; Bowry, V. W.; Ingold, K. U. J. Am. Chem. Soc. 1992,
114, 4983-4992.
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Samuni, U.; Russo, A. Biochemistry 1990, 29, 2802-2807.
(2) Samuni, A.; Godinger, D.; Aronovitch, J.; Russo, A.; Mitchell, J. B.
Biochemistry 1991, 30, 555-561.
(10) Bowry, V. W.; Ingold, K. U. J. Am. Chem. Soc. 1992, 114, 4992-4996.
(11) Brede, O.; Beckert, D.; Windolph, C.; Gottinger, H. A. J. Phys. Chem. A
1998, 102, 1457-1464.
(12) Nilsson, U. A.; Carlin, G.; Bylundfellenius, A. C. Chem.-Biol. Interact.
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V. N. Anti-Cancer Drugs 1991, 2, 591-595.
(13) Offer, T.; Samuni, A. Free Radical Biol. Med. 2002, 32, 872-881.
(14) Akaike, T.; Yoshida, M.; Miyamoto, Y.; Sato, K.; Kohno, M.; Sasamoto,
K.; Miyazaki, K.; Ueda, S.; Maeda, H. Biochemistry 1993, 32, 827-832.
(15) Singh, R. J.; Hogg, N.; McHaourab, H. S.; Kalyanaraman, B. Biochim.
Biophys. Acta 1994, 1201, 437-441.
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C.; Leverenz, V.; Misra, I. C.; Peters, J. L. Exp. Eye Res. 1993, 56, 543-
554.
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148-156.
(6) Howard, B. J.; Yatin, S.; Hensley, K.; Allen, K. L.; Kelly, J. P.; Carney,
J.; Butterfield, D. A. J. Neurochem. 1996, 67, 2045-2050.
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K. N.; Mota-Filipe, H.; Thiemermann, C. Kidney Int. 2000, 58, 658-
673.
(16) Hogg, N.; Singh, R. J.; Joseph, J.; Neese, F.; Kalyanaraman, B. Free Radical
Res. 1995, 22, 47-56.
(17) Weber, H.; Grzesiok, A.; Sustmann, R.; Korth, H. G. Z. Naturforsch., B:
Chem. Sci. 1994, 49, 1041-1050.
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1139-1144.
9
8364
J. AM. CHEM. SOC. 2003, 125, 8364-8370
10.1021/ja035286x CCC: $25.00 © 2003 American Chemical Society

Reaction of Cyclic Nitroxides with Nitrogen Dioxide
A R T I C L E S
reversing the applied potential, suggesting that 4-oxo-TPO+ is rapidly
converted into an unreactive species (see Discussion). Electrochemical
oxidation of 4-OH-TPO yielded a mixture of 4-OH-TPO+ and the same
end product as found during the electrooxidation of 4-oxo-TPO. The
yield of 4-OH-TPO+ did not exceed 70%, and it was stable for only
several minutes.
‚
of nitroxides toward NO2, which is a key biological oxidant
derived, for example, from decomposition of peroxynitrite in
the absence and presence of CO2,19,20 from peroxidase-catalyzed
oxidation of nitrite,21-23 or from autoxidation of ‚NO.24-26 The
reaction of ‚NO2 with 4-OH-TPO and TPO to form the
corresponding oxoammonium cations has been suggested to
occur during the reduction of tetranitromethane by these
nitroxides, and the rate constants were determined indirectly to
be 1.2 × 105 and 2 × 105 M-1 s-1, respectively.27 Carroll et
al.28 proposed that during the decomposition of peroxynitrite
the nitrosation of phenols takes place through the reaction of
‚NO2 with 4-OH-TPO, whereas Bonini et al.29 assumed that this
reaction does not take place because of the relatively low
reduction potential of the ‚NO2/NO2- couple (1.04 V30). It was
also assumed that ‚NO2 does not react with nitronyl nitroxides14-16
because a large excess of ‚NO2 had no effect on the EPR spectra
and signal intensities of these nitroxides.14
Electrooxidation of 1 mM 4-oxo-TPO in D2O containing 4 mM K2-
1
SO4 was carried out for the H NMR study. The oxidized nitroxides
along with the respective hydroxylamines TPO-H and 4-oxo-TPO-H
were scanned using a Mercury 300 NMR spectrometer of Varian
Instruments. The proton parameters were obtained while suppressing
the peak of water protons, and each of the samples was repeatedly
pulsed at least 32 times. Product analysis was also carried out by HPLC-
MS (LCQ ion trap equipped for electrospray, Finnigan Instrumentation).
Pulse radiolysis experiments were carried out using a 5-MeV Varian
7715 linear accelerator (0.05-1.5 µs electron pulses, 200 mA current).
The dose per pulse was 2-29 Gy as determined with N2O-saturated
solutions containing 5 mM KSCN or ferrocyanide. A 200 W Xe lamp
produced the analyzing light. Appropriate cutoff filters were used to
minimize photochemistry. All measurements were made at room
temperature by using a 4-cm spectrosil cell and applying three light
passes (optical path length 12.1 cm).
In the present study, ‚NO2 was generated by pulse radiolysis,
and its reaction with piperidine and pyrrolidine nitroxides was
investigated. The nitroxides were found to react rapidly with
‚NO2 to form the respective oxoammonium cations, which can
Results
oxidize nitrite, ferrocyanide, or ABTS2-
.
Nitrogen dioxide was generated by irradiating N2O-saturated
(∼25 mM) aqueous solutions containing 1-6 mM sodium nitrite
and 4-8 mM phosphate (pH 6-8) or 1 mM acetate (pH 4.8)
buffers through reactions 1-3 (in parentheses are the radiation-
chemical yields of the species, which are defined as the number
of species produced by 100 eV of absorbed energy. Yields are
somewhat higher in the presence of high solute concentrations):
Experimental Section
Materials and Methods. All chemicals were of analytical grade
and were used as received. Water for preparation of the solutions was
distilled and purified using a Milli-Q purification system. The nitroxides
TPO, 4-OH-TPO, and 3-carbamoylproxyl (3-CP) were purchased from
Aldrich and 4-oxo-TPO, and its corresponding hydroxylamine, 4-oxo-
TPO-H, were from Alexis Biochemicals. Fresh solutions of potassium
ferrocyanide and 2,2′-azinobis(3-ethylbenzothiazoline-6-sulfonate)
(ABTS2-) purchased from Sigma were prepared daily. Reduced
â-nicotinamide adenine dinucleotide (NADH) grade III from yeast was
obtained from Sigma. The concentration of NADH was determined
spectrophotometrically using ꢀ340 ) 6200 M-1 cm-1. The hydroxy-
lamines 4-OH-TPO-H and TPO-H were prepared by catalytic reduction
using H2 bubbled over Pt powder or by bubbling HCl gas through
ethanolic solution of the nitroxide followed by drying. Fresh solutions
of TPO-H and 4-OH-TPO-H were prepared immediately before each
experiment to minimize oxidation to the nitroxide. The oxoammonium
cations were prepared in aerated solutions containing 0.2 mM nitroxide
in 4 mM phosphate buffer (PB) at pH 6.8 using an electrochemical
reactor as previously described.31 The yield of TPO+ and 3-CP+
exceeded 94% as determined using ferrocyanide, which was im-
mediately oxidized to ferricyanide (ꢀ420 ) 1000 M-1 cm-1). These
oxoammonium cations were stable for at least 2 h. The final electrolyzed
product in the case of 4-oxo-TPO did not oxidize ferrocyanide.
Furthermore, the spectrum of the nitroxide was not restored upon
H2O
9
γ8 eaq- (2.6), ‚OH (2.7), H‚ (0.6),
H3O+ (2.6), H2O2 (0.72) (1)
eaq- + N2O f N2 + OH- + ‚OH
k2 ) 9.1 × 109 M-1 s-1 32 (2)
‚OH + NO2- f ‚NO2 + OH-
k3 ) 5.3 × 109 M-1 s-1 19 (3)
The yield of H‚ is ca. 10% of the total radical yield at pH > 3,
-
and it adds rapidly to NO2 (k ) 1.6 × 109 M-1 s-1) to form
‚NO in a process catalyzed by acids and bases.33
The decay of ‚NO2 was followed at 400 nm (ꢀmax ) 200 M-1
cm-1) and corresponded to the well-established pathway (reac-
tions 4 and 5), where N2O4 hardly absorbs at this wavelength.34
‚NO2 + ‚NO2 h N2O4
k4 ) 4.5 × 108 M-1 s-1
(19) Merenyi, G.; Lind, J.; Goldstein, S.; Czapski, G. J. Phys. Chem. A 1999,
103, 5685-5691.
(20) Goldstein, S.; Czapski, G. J. Am. Chem. Soc. 1998, 120, 3458-3463.
(21) Byun, J.; Mueller, D. M.; Fabjan, J. S.; Heinecke, J. W. Febs Lett. 1999,
455, 243-246.
k-4 ) 6.9 × 103 s-1 (4)
(22) Reszka, K. J.; Matuszak, Z.; Chignell, C. F.; Dillon, J. Free Radical Biol.
Med. 1999, 26, 669-678.
N2O4 + H2O f NO3- + NO2- + 2H+
(23) Burner, U.; Furtmuller, P. G.; Kettle, A. J.; Koppenol, W. H.; Obinger, C.
J. Biol. Chem. 2000, 275, 20597-20601.
k5 ) 1 × 103 s-1 (5)
(24) Ford, P. C.; Wink, D. A.; Stanbury, D. M. Febs Lett. 1993, 326, 1-3.
(25) Lewis, R. S.; Tannenbaum, S. R.; Deen, W. M. J. Am. Chem. Soc. 1995,
117, 3933-3939.
The half-life of ‚NO2 was not shortened by the addition of 500
µM TPO-H or 4-OH-TPO-H at pH 6.8. Because less than 20%
(26) Goldstein, S.; Czapski, G. Inorg. Chem. 1996, 35, 5935-5940.
(27) Petrov, A. N.; Kozlov, Y. N. Russ. J. Phys. Chem. 1986, 60, 195-198.
(28) Carroll, R. T.; Galatsis, P.; Borosky, S.; Kopec, K. K.; Kumar, V.; Althaus,
J. S.; Hall, E. D. Chem. Res. Toxicol. 2000, 13, 294-300.
(29) Bonini, M. G.; Mason, R. P.; Augusto, O. Chem. Res. Toxicol. 2002, 15,
506-511.
(32) Mallard, W. G.; Ross, A. B.; Helman, W. P. NIST Standard References
Database 40, Version 3.0, 1998.
(33) Lymar, S. V.; Schwarz, H. A.; Czapski, G. J. Phys. Chem. A 2002, 106,
7245-7250.
(30) Stanbury, D. M. AdV. Inorg. Chem. 1989, 33, 69-138.
(31) Goldstein, S.; Merenyi, G.; Russo, A.; Samuni, A. J. Am. Chem. Soc. 2003,
125, 789-795.
(34) Gratzel, M.; Henglein, A.; Lilie, J.; Beck, G. Ber. Bunsen-Ges. Phys. Chem.
1969, 73, 646.
9
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A R T I C L E S
Goldstein et al.
were accompanied by the appearance of a transient observed
in the UV region. Typical kinetic traces at 290 and 400 nm are
shown in Figure 3 for 3-CP, where ∆ꢀ290 ) 670 ( 50 M-1
cm-1
.
In the case of TPO, the decay of the transient absorption at
290 nm was too slow to study by our setup. In the case of the
other nitroxides, the decay of the absorption exhibited second-
-2
order kinetics, and kobs was linearly dependent on [nitroxide]o
-
at a constant [NO2-]o and on [NO2
]
o
2 at a constant [nitroxide]o
(Figures 4 and 5).
‚
The decay of NO2 was unaffected by the number of pulses
delivered into the solution, which suggests that there is no loss
of the nitroxides during this process. After repetitive pulsing
(340 Gy) of N2O-saturated solution containing 200 µM nitroxide
and 2 mM nitrite at pH 6.8, the solution was transferred to a
diode array spectrophotometer, and the spectrum was recorded
within less than 1 min after the irradiation. No appreciable
spectral changes were observed for 4-OH-TPO, 4-oxo-TPO, and
3-CP. In the case of TPO, the product of the irradiated sample
was identified as TPO+, which decayed slowly with a first half-
life of about 30 min in the presence of 2 mM nitrite.
Figure 1. The decay of ‚NO2 with and without 4-OH-TPO. The decay of
‚NO2 was followed at 400 nm upon pulse-irradiation (29 Gy/pulse) of N2O-
saturated solutions containing 2 mM nitrite and 4 mM PB at pH 6.8 in the
absence (a) and in the presence of 150 µM 4-OH-TPO (b). The optical
path length was 12.1 cm.
Reaction of ‚NO2 with Nitroxide in the Presence of
ABTS2- or Ferrocyanide. The rate constants for the reaction
‚
of NO2 with ABTS2- and ferrocyanide have been previously
determined to be 2.2 × 107 and 3 × 106 M-1 s-1, respectively.32
‚
The NO2-induced formation of ABTS‚- (ꢀ660 ) 12 000 M-1
cm-1) or ferricyanide (ꢀ420 ) 1000 M-1 cm-1) was followed in
the absence and presence of various concentrations of the
nitroxide upon pulse-irradiation of N2O-saturated solutions
containing 4 mM nitrite at pH 6.8 ([‚NO2]o ) 1.4-4 µM). The
yields of ABTS‚- or ferricyanide were unaffected upon the
addition of the various nitroxides, whereas the formation rate
increased linearly with the increase in the concentration of the
nitroxide. The observed first-order rate constant of the formation
of ABTS‚- as a function of [nitroxide]o is given in Figure 6
‚
Figure 2. The observed rate constant of the decay of NO2 as a function
-
of [nitroxide]. All solutions were N2O-saturated and contained 2 mM NO2
and 4 mM PB at pH 6.8. The dose was 12-29 Gy/pulse.
4-
‚
(only a few data points obtained in the presence of Fe(CN)6
were included to avoid overloading the figure).
acceleration of NO2 decay would be readily detectable, the
‚
upper limit for the rate constant of the reaction of NO2 with
these hydroxylamines can be estimated to be 1 × 105 M-1 s-1
.
The formation rate was almost unaffected by [ABTS2-]o or
[Fe(CN)64-]o, for example, less than 10% increase in kobs
when [ABTS2-]o was increased from 100 to 500 µM in the
The establishment of equilibrium 4 is a relatively fast second-
order process, which upon the addition of 80-300 µM nitroxide
turned into a fast first-order reaction as a result of reaction 6.
‚
presence of 100 µM nitroxide. The results indicate that NO2
reacts with the nitroxide to form a reactive intermediate, which
oxidizes ABTS2- or Fe(CN)64- with a diffusion-controlled rate
constant. Hence, the rate-determining step is the reaction of
‚NO2 with the nitroxide (reaction 6), and the rates of forma-
tion of ABTS‚- and ferricyanide in the presence of all nitrox-
ides tested are the same (Figure 6), both being characterized
by k6. From the slope of the lines in Figure 6, we calculated k6
(Table 1).
RNO‚ + ‚NO2 f products
(6)
‚
Typical kinetic traces for the decay of ca. 16 µM NO2 alone
and in the presence of 150 µM 4-OH-TPO are shown in Fig-
ure 1.
‚
The decay of NO2 in the presence of TPO and 4-OH-TPO
was first-order, and kobs was linearly dependent on [nitroxide]o
(Figure 2) and independent of the pH over the range 3.9-10.2.
In the presence of 4-oxo-TPO and 3-CP, the fast first-order
The Case of 4-oxo-TPO. In contrast to TPO, 4-OH-TPO,
and 3-CP, the electrochemical oxidation of 4-oxo-TPO forms a
relatively unreactive product, which does not oxidize ABTS2-
or ferrocyanide. We have recently shown that TPO+ and 3-CP+
react with NADH via a two-electron-transfer mechanism to form
the corresponding hydroxylamine and NAD+.31 In contrast, the
product accumulated during electrooxidation of 4-oxo-TPO was
stable for at least 3 h in the presence of NADH at pH 6.8.
However, about 4 µM NADH was consumed when 4 µM ‚NO2
reacted with 50 or 100 µM 4-oxo-TPO in the presence of 100
µM NADH and 4 mM nitrite at pH 6.8. The decay of NADH
‚
decay of NO2 was followed by a slower second-order decay.
The observed first-order rate constant was linearly dependent
on [nitroxide]o (Figure 2), whereas the rate of the slow second-
order decay increased upon decreasing [nitroxide]o and increas-
ing [NO2-]o. The rate constant of reaction 6 was calculated from
the slopes of the lines in Figure 2, and the results are
summarized in Table 1.
In the presence of each of the four nitroxides tested, the for-
‚
mation and decay of NO2, which were followed at 400 nm,
9
8366 J. AM. CHEM. SOC. VOL. 125, NO. 27, 2003

Reaction of Cyclic Nitroxides with Nitrogen Dioxide
A R T I C L E S
Table 1. Summary of All Rate Constants and Reduction Potentials Determined in the Present Study
TPO
4-OH-TPO
4-oxo-TPO
3-CP
k6, M-1 s-1
(Figure 2, direct)
k6, M-1 s-1
(Figure 6, ABTS2-
k6, M-1 s-1
(Figure 6, Fe(CN)6
(5.6 ( 0.3) × 108
(7.1 ( 0.2) × 108
(6.7 ( 0.2) × 108
(7.5 ( 0.4) × 108
(8.7 ( 0.2) × 108
(7.8 ( 0.2) × 108
(6.3 ( 0.2) × 108
(7.1 ( 0.2) × 108
(6.7 ( 0.3) × 108
(6.5 ( 0.3) × 108
(7.1 ( 0.2) × 108
(6.7 ( 0.3) × 108
)
4-
a
)
K6
(8 ( 4) × 104 b
(9 ( 4) × 103
0.7531
(2.9 ( 0.2) × 103
(2.7 ( 0.4) × 105
0.84
63 ( 7
450 ( 20
(1.5 ( 0.2) × 106
0.88
k-6, M-1 s-1
(1.1 ( 0.2) × 107
E°(RN+dO/RNO‚), V
k6, M-1 s-1
0.93
5 × 107 c
4 × 106 d
NDe
1 × 107 c
2 × 106 c
6 × 106 c
4 × 105 d
ND
(Marcus equation)
9 × 105 d
2 × 105 d
k9, M-1 s-1
ND
(7 ( 1) × 108
a The rate constant was calculated from the slopes of the lines as seen in Figure 6, although not all of the data points were included. The results in the
4-
presence of Fe(CN)6 are within experimental error identical to those determined in the presence of ABTS2-. We included in Figure 6 only a few data
points for each nitroxide to avoid overloading the figure. b Calculated using k7 and E°(TPO+/TPO) ) 0.75 ( 0.01 V.31 c Assuming that the self-exchange
rate constant for RNO‚/RN+ is 1 × 1010 M-1 s-1
.
d Assuming that the self-exchange rate constant for RNO‚/RN+ is 5 × 107 M-1 s-1 42
.
e ND - not
determined.
Figure 4. The effect of nitroxide concentrations on the decay rate of the
transient absorption at 290 nm. The effect of [3-CP]o and [4-oxo-TPO]o in
the presence of 2 mM nitrite and that of [4-OH-TPO]o in the presence of
4 mM nitrite were studied. All solutions were N2O-saturated and contained
4 mM PB at pH 6.8. The dose was 29 Gy/pulse.
‚
Figure 3. The reaction of 3-CP with NO2. The formation and decay
‚
of NO2 were followed at 400 nm (A), whereas those of the transient
formed during this reaction were monitored at 290 nm (B). N2O-satu-
rated solutions containing 140 µM 3-CP and 2 mM nitrite at pH 6.8 (4
mM PB) were pulse-irradiated (29 Gy/pulse). The optical path length was
12.1 cm.
at 385 nm (ꢀ ) 830 M-1 cm-1) was first-order with kobs ) (1.4
( 0.1) × 104 or (2.3 ( 0.1) × 104 s-1, respectively. These
results show that the reaction of ‚NO2 with 4-oxo-TPO forms a
reactive intermediate, which oxidizes NADH, and that the
former reaction is not rate-limiting as in the case of ferrocyanide
Figure 5. The effect of nitrite concentrations on the decay rate of the
-
transient absorption at 290 nm. The effect of [NO2
] in the presence of
o
140 µM 3-CP or 150 µM 4-oxo-TPO was studied. All solutions were
N2O-saturated and contained 4 mM PB at pH 6.8. The dose was 29 Gy/
pulse.
or ABTS2-
.
cation, which decays during the electrolysis process into a
relatively unreactive species. We used HPLC-MS to compare
the spectra of 4-oxo-TPO with that of its electrooxidation
product. We failed to detect the anticipated peak of 170+1 due
to 4-oxo-TPO itself, presumably because of a poor pickup of a
proton by the parent nitroxide. However, for the electrooxidized
product, we observed a peak of 170 that could not be attributed
We have previously shown that TPO and 3-CP were almost
fully recovered (ca. 94%) when their electrooxidized solution
was subjected to electroreduction.31 Conversely, the spectrum
of 4-oxo-TPO was not recovered upon reduction, and the
spectrum of the reduced product was hardly affected upon
subsequent electrooxidation. We therefore conclude that the
electrooxidation of 4-oxo-TPO forms an unstable oxoammonium
9
J. AM. CHEM. SOC. VOL. 125, NO. 27, 2003 8367

A R T I C L E S
Goldstein et al.
Figure 6. The observed rate constant for the formation of ABTS.- (open
symbols) or Fe(CN)63- (solid symbols) as a function of nitroxide concentra-
tion. All solutions were N2O-saturated and contained 4 mM NO2-, 250-
500 µM ABTS2- or Fe(CN)64-, and 8 mM PB at pH 6.8. The dose was 2.4
or 6.9 Gy/pulse, and ko is the observed rate of the formation of ABTS‚- or
Fe(CN)6 in the absence of nitroxide, that is, ko ) 2 × 107 [ABTS2-] or
3-
4-
3 × 106 [Fe(CN)64-]. The results in the presence of Fe(CN)6 are within
experimental error identical to those obtained in the presence of ABTS2-
.
We included only a few data points for each nitroxide to avoid overloading
the figure.
Figure 7. The 1H NMR spectrum of 4-oxo-TPO-H (bottom) and the
electrooxidation product of 4-oxo-TPO (top) in D2O (structure given in eq
7). The spectrum of the hydroxylamine consisted of a singlet at 1.48 ppm
(12 protons of the methyls at the 2 and 6 positions) and a singlet at 2.87
ppm (4 protons at the 3 and 5 positions). The spectrum of the electrooxi-
dation product showed a singlet at 1.55 ppm (6 protons of the methyls at
position 1), a singlet at 2.2 ppm (6 protons of the methyls at position 5),
and a singlet at 2.99 ppm (2 protons at position 2).
to 4-oxo-TPO+, which is already charged, but to a novel
uncharged species having a MW of 169. We therefore suggest
that 4-oxo-TPO+ undergoes a rapid â-elimination of proton via
eq 7:
‚
‚
We have previously found that the reaction of OH or N3
with 4-oxo-TPO forms the same species as that produced
electrochemically, but with different yields.36 We have mistak-
enly identified this product as 4-oxo-TPO+, on the basis of the
observation that, after partial oxidation of 3.66 mM 4-oxo-TPO,
the spectrum of the nitroxide in the visible region was almost
fully recovered upon subsequent electroreduction. However, a
closer examination of the UV and visible spectrum after
complete oxidation at low and high concentrations of the
nitroxide suggests that the nitroxide was not restored when the
electrooxidized solution was subjected to electroreduction. Thus,
the reaction of 4-oxo-TPO with ‚OH or ‚N3 forms the unstable
4-oxo-TPO+, which decays in the absence of a potential
reductant into the same species as is formed electrochemically.
To verify our suggestion, we electrooxidized 4-oxo-TPO in
D2O and compared the H NMR spectrum of the product to
1
that of 4-oxo-TPO-H. The spectrum of the hydroxylamine
(Figure 7, bottom), which consisted of a singlet at 1.48 ppm
due to 12 protons of the methyls at the 2 and 6 positions, and
a singlet at 2.87 ppm corresponding to the 4 protons at the 3
and 5 positions, was anticipated to resemble that of 4-oxo-TPO+.
In fact, it greatly differed from that observed for the electrooxi-
dation product (Figure 7, top). The latter showed a singlet at
1.55 ppm due to protons of the 2 methyls bound to the carbon
with the nitroso group (position 1), a singlet at 2.2 ppm
attributable to the protons of the methyls at position 5, and a
singlet at 2.99 ppm due to the 2 protons at position 2. The
expected singlet of the vinylic proton was overshadowed by
the large water peak at 4.8 ppm, which was present in all of
our samples.
Generally, aliphatic C-nitroso compounds can exist as
monomers, colorless dimers, or tautomeric oximes. However,
tertiary nitroso compounds often exist as monomers, which are
colored and exhibit a typical absorption band between 630 and
790 nm (ꢀ ) 1-60 M-1 cm-1).35 This further supports our
conclusion because the electrooxidized product absorbs at 650
nm (ꢀ ) ca. 14 M-1 cm-1).36
Discussion
‚
The reaction of NO2 with nitroxides forms an intermediate
which can oxidize ferrocyanide, ABTS2-, NADH, and nitrite.
In the case of TPO, the intermediate reacted sufficiently slowly
with nitrite for it to be identified spectrophotometrically as
TPO+. In the case of the other nitroxides, the intermediates
decayed faster in the presence of nitrite via a second-order
reaction. The observed second-order rate constant was linearly
-2
dependent on [nitroxide]o at a constant [NO2-]o, and on
-
[NO2
]
2 at a constant [nitroxide]o (Figures 4, 5). These results
o
suggest that the respective oxoammonium cations are formed
via reaction 6 and decomposed slowly via reaction -6 followed
by reactions 4 and 5.
(35) Williams, D. L. H. Nitrosation; Cambridge University Press: Cambridge,
1988.
(36) Samuni, A.; Goldstein, S.; Russo, A.; Mitchell, J. B.; Krishna, M. C.; Neta,
P. J. Am. Chem. Soc. 2002, 124, 8719-8724.
RNO‚ + ‚NO2 h RN+dO + NO2
(6)
-
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8368 J. AM. CHEM. SOC. VOL. 125, NO. 27, 2003

Reaction of Cyclic Nitroxides with Nitrogen Dioxide
A R T I C L E S
Assuming that reaction 6 is a fast approach to equilibrium, we
found that the rate-limiting step for the decay of RN+dO is
the dimerization/hydrolysis of ‚NO2, and rate eq 8 is obtained,
× 107 M-1 s-1 for X ) OCH3 and OCOCH3, respectively.42 In
acetonitrile, the rates are expected to be similar for nitroxides
studied by us. If anything, these rates should be smaller in water,
because in the latter solvent a stronger solvation of the cation
should increase the outer-sphere components of the barrier to
the electron self-exchange. Still, 1 × 1010 M-1 s-1 must surely
be considered a generous upper limit. As seen in Table 1, the
calculated values for k6 are at least 1-2 orders of magnitude
lower than the experimental values (Table 1). Therefore, reaction
6 most probably takes place via an inner-sphere electron-transfer
mechanism,
2k4k5[NO2 ]o [RN+dO]2
-
2
d[RN+dO]
dt
-
)
)
(k-4 + k5)K6 [RNO‚]o
2
2
kobs[RN+dO]2 (8)
where from the literature 2k4k5/(k-4 + k5) ) 1.3 × 108 M-1
s-1 is calculated. Hence, kobs ) 1.3 × 108([NO2-]o/K6[RNO‚
-
]o)2, and plots of kobs versus 1/[RNO‚]o2 or [NO2
]
2 should be
o
RNO‚ + ‚NO2 h adduct
adduct h RN+dO + NO2
(10)
(11)
linear (Figures 4 and 5) with slope1 ) 1.3 × 108([NO2-]o/K6)2
or slope2 ) 1.3 × 108/([RNO‚]/K6)2, respectively, and (slope1/
slope2)1/2 ) [RNO‚]o[NO2-]o. Hence, (slope1/slope2)1/2 ) (2.6
( 0.1) × 10-7 and (3.7 ( 0.1) × 10-7 M2 for 3-CP and 4-oxo-
TPO, respectively, and are in good agreement with the experi-
mental conditions, that is, 2 mM nitrite and 140 µM 3-CP or
150 µM 4-oxo-TPO. The values of K6, which are derived from
the slope of the lines in Figures 4 and 5, are summarized in
Table 1.
-
where k6 ) k10k11/(k-10 + k11).
Petrov and Kozlov27 studied the reaction of nitroxides with
tetranitromethane in the presence of nitrite and measured the
concentrations of the reagents and products at the minium
concentration of the nitroxide under different initial concen-
trations of the reagents. A complex linear relationship yielded
a straight line, and from the slopes of the lines they deter-
mined K6 ) 2.7 × 103 and 7 × 104 for 4-OH-TPO and TPO,
which are in excellent agreement with our determination.
However, from the intercepts of these lines, they derived k6 )
1.2 × 105 and 2 × 105 M-1 s-1 for 4-OH-TPO and TPO,
respectively, which are more than 3 orders of magnitude lower
than those determined directly (Table 1). Most likely, their
approach was not sensitive enough for accurate determination
of k6.
The rate constant for the oxidation of ferrocyanide and
ABTS2- by RN+dO is extremely high, that is, ca. 1 × 1010
M-1 s-1, which explains why the rate-determining step of
‚
the formation of the oxidized product is the reaction of NO2
with the nitroxide. We simulated the formation of NAD+
using reactions 4-6 and 9 for the reaction of 4 µM ‚NO2 with
50 or 100 µM 4-oxo-TPO in the presence of 100 µM NADH
and 4 mM nitrite, where kobs ) (1.4 ( 0.1) × 104 or (2.3 (
0.1) × 104 s-1, respectively, and obtained k9 ) (7 ( 1) × 108
M-1 s-1
.
‚
The present study may explain why a large excess of NO2
4-oxo-TPO+ + NADH f 4-oxo-TPO-H + NAD+ (9)
had no effect on the EPR spectra and signal intensities of
nitronyl nitroxides.14 These nitroxides are expected to react
The reduction potentials of the RN+dO/RNO‚ couples have
been calculated using K6 and E°(‚NO2/NO2-) ) 1.04 V to be
0.84, 0.93, and 0.88 V for 4-OH-TPO, 4-oxo-TPO, and 3-CP,
respectively. The latter value is in excellent agreement with
‚
rapidly with NO2 to form the corresponding nitronyl oxoam-
monium cations, which could be reduced back to the nitroxides
in the presence of nitrite formed mainly via the self-decomposi-
tion of the large excess of ‚NO2 in aqueous solutions (reactions
4-5).
‚
that determined previously using HO2 as an oxidant, that is,
0.89 V.31 The calculated reduction potentials for all nitroxides
are in good agreement with the midpotentials, which have
been previously determined by cyclic voltammetry.37-39 In the
case of TPO, we calculated K6 ) 8 × 104 using the recent value
of 0.75 V for E°(TPO+/TPO).31 The results are compiled in
Table 1.
Conclusions
The rate constants determined in the present study for the
‚
reaction of NO2 with piperidine and pyrrolidine nitroxides
were determined to be (7-8) × 108 M-1 s-1 and are close to
‚
‚
that determined for the reaction of NO with NO2, that is,
As in the case of HO2‚,31 one can use the Marcus equation k6
) (kakbK6 f)1/2 (ka and kb are the self-exchange rate constants
for the reactants, ln f ) (ln K6)2/4 ln(kakb/1022)40 to show that
reaction 6 does not take place via an outer-sphere electron-
1.1 × 109 M-1 s-1 43
.
These are among the highest rate con-
stants measured for NO2 at physiological pH44 and exceed by
far those with glutathione, cysteine, and urate, which are
considered to be the major scavengers of the radical in biological
systems.45 The kinetics results demonstrate that the reaction
proceeds via an inner-sphere electron-transfer mechanism to
form the respective oxoammonium cation, which is reduced back
‚
‚
transfer mechanism. The self-exchange rate constant for NO2 /
NO2 has been determined to be 9.6 M-1 s-1 41
.
The self-
-
exchange rate constant for X-TPO/X-TPO+ in acetonitrile has
been determined by EPR and 1H NMR to be 5.7 × 107 and 4.4
‚
to the nitroxide while oxidizing nitrite to NO2. Hence, the
nitroxide slows down the decomposition of NO2 into nitrite
‚
(37) Fish, J. R.; Swarts, S. G.; Sevilla, M. D.; Malinski, T. J. Phys. Chem. 1988,
92, 3745-3751.
(38) Morris, S.; Sosnovsky, G.; Hui, B.; Huber, C. O.; Rao, N. U. M.; Swartz,
H. M. J. Pharm. Sci. 1991, 80, 149-152.
(39) Krishna, M. C.; Grahame, D. A.; Samuni, A.; Mitchell, J. B.; Russo, A.
Proc. Natl. Acad. Sci. U.S.A. 1992, 89, 5537-5541.
(40) Marcus, R. A. J. Phys. Chem. 1963, 67, 853-857.
(41) Awad, H. H.; Stanbury, D. M. J. Am. Chem. Soc. 1993, 115, 3636-
3642.
(42) Wu, L. M.; Guo, X.; Wang, J.; Guo, Q. X.; Liu, Z. L.; Liu, Y. C. Sci.
China, Ser. B 1999, 42, 138-144.
(43) Gratzel, M.; Taniguch, S.; Henglein, A. Ber. Bunsen-Ges. Phys. Chem.
1970, 74, 488-492.
(44) Huie, R. E. Toxicology 1994, 89, 193-216.
(45) Ford, E.; Hughes, M. N.; Wardman, P. Free Radical Biol. Med. 2002, 32,
1314-1323.
9
J. AM. CHEM. SOC. VOL. 125, NO. 27, 2003 8369

A R T I C L E S
Goldstein et al.
‚
and nitrate and could serve as a reservoir of NO2 radicals.
Moreover, the respective oxoammonium cation could oxidize
biological targets, causing further damage unless it is reduced
back to the nitroxide by common biological reductants, such
as thiols or urate. These subsequent processes are currently under
investigation.
Acknowledgment. This research was supported (A.S.) by a
grant from the Israel Science Foundation of the Israel Academy
of Sciences. We thank Dr. Henry Fales for mass spectroscopic
analysis and Dr. Gabor Merenyi for helpful discussions.
JA035286X
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8370 J. AM. CHEM. SOC. VOL. 125, NO. 27, 2003