Addition of oxide radical ions (O-) to nitrite and oxide ions (O2-) to nitrogen dioxide [2]

Full text of DOI:10.1021/ja994230t

J. Am. Chem. Soc. 2000, 122, 3773-3774
3773
-
O- to oxygen, CN , SCN ). Whereas the reaction of oxide
-
- 4c
Addition of Oxide Radical Ions (O ) to Nitrite and
2
-
Oxide Ions (O ) to Nitrogen Dioxide
radicals with nitrite ions has been studied before, it was suggested
-
-
to proceed via electron transfer. The transfer of O from CO
radicals to nitrite ions (and many other possible acceptors,
including NO
) was investigated earlier.4b However, this reaction
is too slow to allow detection of the intermediate and NO was
suggested to be the product.
3
Richard W. Fessenden and Dan Meisel*
Notre Dame Radiation Laboratory and
Department of Chemistry and Biochemistry
UniVersity of Notre Dame, Notre Dame, Indiana 46556
2
2
-
In this report we provide evidence that O at least partially
adds to nitrite ions.
Donald M. Camaioni*
Pacific Northwest National Laboratory
Richland, Washington 99352
NO2- + O f NO
-
2-
(4)
3
ReceiVed December 3, 1999
ReVised Manuscript ReceiVed February 22, 2000
We also find indications for the reaction of O2 with NO
reverse reaction -2. There are several reports in the literature
that organic radicals undergo an analogous addition reaction with
-
, the
2
The redox chemistry of nitrogen oxides is complex, involving
both radicals and radical ions. Knowledge of their kinetics and
reaction mechanisms is fundamental to understanding these
reactions in general and especially those relevant to understanding
5
nitrite to yield nitroanions. We estimate that reaction 4 is highly
exoergic (∆G° ) -20 kcal/mol).6
Pulse radiolysis was utilized to generate the radicals and time-
resolved ESR to unequivocally identify them and study their
x
of atmospheric NO chemistry. Renewed interest in radiolytically
induced processes of these species results from their frequent
presence in nuclear materials in various storage facilities. As part
of a study to understand the radiation effects in nitrate and nitrite
7
kinetics. Experiments were initially done to determine if the ESR
2-
lines of NO
2
and NO
3
could be seen in solution. The detection
from the reaction of eaq with nitrate was tried first.
2-
-
of NO
3
2-
solutions, the nitrate dianion, NO
3
, in water has been reinves-
2-
Lines attributable to NO
3
were found in neutral to basic (1 M
1
tigated recently. The one-electron reduction potential of nitrate
in water (eq 1) was determined and the mechanism of its
decomposition in water elucidated (reaction 2).
-
OH ) solutions of nitrate. The ESR parameters are a ) 43.4 G
and g ) 2.00458 with a line width of 0.20 G. These parameters
2-
resemble those of NO
3
in solids, where there is some variability
8
depending on the particular matrix. These lines disappeared when
O was added (to convert eaq to additional OH) so the
identification is not in doubt. The oxidation of nitrite by OH in
O saturated solutions was then studied. No lines attributable
NO3- + e f NO3 (E° ) -0.89 V)
-
2-
(1)
(2)
-
N
2
2-
-
NO3 + H O S 2OH + NO
4c
2
2
N
2
to NO
2
were found in solutions containing 1-10 mM nitrite at
Many acids catalyze the decomposition reaction 3. Protonation
2
near neutral pH. On this basis, the ESR line width of NO in
water must be larger than ∼3 G.
2
-
-
-
NO3 + HA f OH + A + NO
(3)
but
Next, experiments were done with solutions of nitrite saturated
2
2
-
with N
2
O but looking at the expected lines of NO
3
. Figure 1
-
2-
of the radical dianion may proceed via the acid form HNO
its formation was questioned in this study. Analysis of the kinetics
3
shows the time dependence of the central ESR line of NO
3
for
-
various concentrations of OH . Other than the two reactions
of reaction 3 suggested that the proton transfer from the HA is
described above, reactions -2 and 4, it is difficult to suggest any
N-O2 bond. Mezyk and Bartels
-
2-
concerted with scission of the O
2
other pathway for the generation of NO
3
in this system. In
-
proposed a similar mechanism based on analysis of the activation
solutions containing g0.1 M OH , formation of the radical
dianion can be seen (<5 µs after the pulse) and then it decays
via an apparent first-order process (τ > 10 µs). At the higher
base concentrations this decay leads to a nonzero pseudoequi-
librium level that decays at a slower time scale. The maximum
2
parameters of the reaction of the H atom with nitrate. Accord-
ingly, the reaction may be viewed as an O2 transfer reaction to
-
the acid (or water).
From the redox potential of eq 1 and of NO
-
2
(E°(NO
2
/NO
2
)
/
3
-
-
)
1.04 V; all redox potentials are vs NHE) and nitrate (E°(NO
3
observed concentration increases with [OH ], but is less than the
-
-2
-
NO
2
) ) 0.01 V) one calculates K
2
≈ 1 × 10 M for the equi-
total initial yield of OH and e aq. The rate of decay appears to
1,4
-
librium constant of reaction 2. Combined with the rate constant
decrease upon increasing [OH ]. Furthermore, the rate of ap-
3
-1 -1
k
2
) 10 M s , the rate constant of the back-reaction is k-2
≈
pearance of the radical is slower than the time resolution of the
ESR spectrometer. The center-field ESR line was used in all
kinetic analyses to minimize polarization effects. The observed
kinetic features could not be fit with the known reactions for the
system (Table 1), nor could it be fit even by allowing eq 2 to be
5
-2 -1
1
× 10 M s . Thus it is conceivable that at molar concentra-
- 2-
2
tions of OH the oxide ion, O , may react with NO
to produce
. Such a finding is entirely novel; it predicts that conditions
may be accessible for hydroxide to add to NO thereby converting
a strongly oxidizing species to a strongly reducing one.
2-
NO
3
2
2
reversible, regardless of the value selected for K . We suggest
2-
2-
Another novel route to the production of NO
3
may exist at
that reaction 4 at least partially generates NO
3
.
highly basic solutions. The conjugated base of hydroxyl radicals,
(
5) Kornblum, N.; Ackermann, P.; Swiger, R. T. J. Org. Chem. 1980, 45,
-
O (pK
a
) 11.9), rarely adds to organic molecules but its addition
5
294. Kornblum, N.; Cheng, L.; Davies, T. M.; Earl, G. W.; Holy, N. L.;
to small inorganic molecules and ions is known (e.g., addition of
Keber, R. C.; Kestner, M. M.; Manthey, J. W.; Musser, M. T.; Pinnick, H.
W.; Snow, D. H.; Stuchal, F. W.; Swiger, R. T. J. Org. Chem. 1987, 52, 196.
Russell, G. A.; Metcalf, A. R. J. Am. Chem. Soc. 1979, 101, 2359. Russell,
G. A.; Khanna, R. K. J. Am. Chem. Soc. 1985, 107, 1450. Beckwith, A. L. J.;
Norman, R. O. C. J. Chem. Soc. (B) 1969, 403. Zeldes, H.; Livingston, R. J.
Am. Chem. Soc. 1968, 90, 4540.
(
1) Cook, A. R.; Dimitrijevic, N.; Dreyfus, B. W.; Meisel, D.; Curtiss, L.
A.; Camaioni, D. M. J. Phys. Chem. B. Submitted for publication.
(
2) Mezyk, S. P.; Bartels, D. M. J. Phys. Chem. A 1997, 101, 6233.
3) Stanbury, D. M. Reduction Potentials InVolVing Inorganic Free Radicals
(
-
-
-/NO -
) + E°(NO - 2-) ) 1.77 -
/NO
3 3
in Aqueous Solution; Sykes, A. G., Ed.; Academic Press: San Diego, 1989;
Vol. 33, pp 69-138.
(6) E
4
) E°(O /OH ) - E°(NO
3
2
0.01 - 0.89 ) 0.87 V.
(
4) (a) Gratzel, M.; Henglein, A.; Taniguchi, S. Ber. Bunsen-Ges. Phys.
(7) Madden, K. P.; McManus, H. J. D.; Fessenden, R. W. ReV. Sci. Instrum.
1994, 65, 49.
Chem. 1970, 74, 292, 488. (b) Henglein, A. Radiat. Phys. Chem. 1980, 15,
1
2
51. Lilie, J.; Hanrahan, R. J.; Henglein, A. Radiat. Phys. Chem. 1978, 11,
25. (c) Buxton, G. V.; Greenstock, C. L.; Helman, W. P.; Ross, A. B. J.
(8) See, for example: Rustgi, S. N.; Box, H. C. J. Chem. Phys. 1973, 73,
4763 (a ) 49.7, g ) 2.0046). Edo, B.; Iwasaki, M. J. Phys. Chem. 1982, 86,
2084 (a ) 43.0, g ) 2.0040).
Phys. Chem. Ref. Data 1988, 17, 513.
1
0.1021/ja994230t CCC: $19.00 © 2000 American Chemical Society
Published on Web 03/30/2000

3774 J. Am. Chem. Soc., Vol. 122, No. 15, 2000
Communications to the Editor
Table 1. Kinetic Model for Simulating the Pulse Radiolysis of Alkaline Nitrite Solutions
no.
reaction
forward (back) rate constants
refs
-
-
-
9
-1 -1
1
2
3
4
5
6
7
8
9
e
e
aq + N
2
O f N
2
+ O
9.1 × 10 M
s
(a) ref 4c
-
-
9
-1 -1
aq + NO
2
(+ H
2
O) f NO + OH
+ OH
3.5 × 10 M
s
ref 15a
-
-
10
-1 -1
OH + NO
2
f NO
2
1 × 10
M
s
ref 4c
-
-
-
10
-1 -1
(9.4 × 10 s-1)
7
OH + OH S O + H
2
O
1.3 × 10
M
s
ref 15b
-
-
7
-1 -1
H + OH f e aq + H
2
O
2.0 × 10 M
s
ref 15c
-
-
-
+ 2OH-
7
-1 -1 a
O + NO
2
(+ H
f NO
O S NO
2
O) f NO
2
4.7 × 10 M
s
(b)
(b)
-
2-
7
-1 -1
O + NO
2
3
1.8 × 10 M
s
2
-
+ 2OH-
3
-1 -1
(500 M s-1
(6800 s for hydrolysis)
-2
NO
3
+ H
2
2
1 × 10 M
s
s
)
(b) forward rate: refs 1, 4a
ref 15d
ref 15d
ref 4a and 15d
8
-1 -1
-1
2NO
2
S N
(+ H
NO + NO
2
O
4
5 × 10 M
-
-
-
+2H+
-1
1
0
1
N
2
O
4
2
O) f NO
3
+ NO
2
1000 s
-
1 × 10 M s-1
9
-1
1
2
(+ 2OH ) f 2NO
2
a
N O concentrations were estimated to be 0.11, 0.16, 0.22, and 0.23 M for 2, 1, 0.3, and 0.1 M NaOH, respectively, using Henry’s constant for
2
water and the modified Sechenov equation to correct for effects of ionic strength. b Rate constants for reactions 6 and -8 of Table 1 were
adjusted to give the best fit to the curves in Figure 1.
1
6a
16b
We do not know the details of the faster component that leads
2
“directly” to NO . Because nitrite ion is ambivalent, the possibility
that the ET path might occur via attack at the O termini of nitrite
was considered. This mode of attack might also produce an
adduct, e.g., peroxynitrite dianion.13 No resonances assignable to
such an adduct were detected in the ESR spectra. To assess the
potential for forming ONOO2 from addition of O to NO
-
-
-
, the
free energy of such a reaction was calculated using ab initio and
2
14
self-consistent reaction field solvation methods. The calculations
2-
2-
show that NO
Interestingly, the free energy for reaction 2 was calculated to be
kcal/mol by the same methodology, some 25 kcal/mol more
3
is more stable than ONOO by 34 kcal/mol.
4
positive than the experimental value. Evidently, the parametriza-
tion of the solvation model is inaccurate for these dianions. None-
theless, even if the energies of the dianions are scaled downward
by 25 kcal/mol, the peroxynitrite dianion is too high in energy to
be formed. Thus, ONOO2 is not a plausible intermediate.
Finally, we come back to reaction -2. Even though thermo-
dynamically the reaction is feasible, experimental evidence is
weak. Adjustments to several rate constants from their reported
values could eliminate it from the scheme. For example, the sim-
ulated rate of decay is sensitive to the rates of radical termination,
but except for reaction 11 in Table 1, the constants are not well
established in strongly alkaline solutions. Nonetheless, because
the simulations following the scheme in Table 1 do predict the
-
-
2-
3
Figure 1. Effect of [OH ] on the time-dependent ESR signal for NO
(
central line) following pulse radiolysis of 6 mM KNO
2
solutions saturated
-
with N
2
O: (.) 0.1, (0) 0.3, (~) 1.0, and (O) 2.0 M NaOH. Lines are
observed second-order dependence on [OH ], we kept it in our
calculated from simulations using the kinetic model in Table 1.
computation. From the “fitted” rate constants for eq 2 we obtain
K ) 2 M, which is reflected in a more negative redox potential
2
Simulations showed that such a reaction is consistent with the
than experimentally measured (-1.0 V rather than -0.89 V).
-
9
observed dependence of the yields and kinetics on [OH ]. An
acceptable match to the data of Figure 1 is obtained on reacting
In conclusion, we provide experimental results showing that
O- radicals add, at least partially, to NO
oxidizing environment, dominated by OH/O radicals, is tran-
-
2
. Thus, a strongly
-
-
2-
7
O with NO
and to give NO
approximately 25% of the attack yields NO
reproduce well the dependence of the maximum observed
2
to give NO
2
3
with a rate constant of 1.8 × 10
-
7
-1 -1
with rate constant of 4.7 × 10 M s , i.e.,
siently transformed into a strongly reducing environment domi-
2-
3
. These values
2-
nated by the NO
3
in the presence of high concentrations of base.
2-
-
Acknowledgment. Support of this work by the U.S. Department of
Energy, Office of Basic Energy Sciences and by the Environmental Man-
agement Science Program is gratefully acknowledged. This is Contribu-
tion NDRL No. 4184 from the Notre Dame Radiation Laboratory. Pacific
Northwest National Laboratory is operated by Battelle Memorial Institute
for the U.S. Department of Energy under Contract DE-AC06-76RLO
concentrations of NO
3
on [OH ]. However, to obtain good fit
to the decay of the dianion, it was necessary to include the reverse
-
2
-1
reaction in eq 2. Best fit is obtained with k-2 ) 500 M s .
2-
-
The production of NO
3
is not significant below 0.1 M OH
even though the pK of OH radicals is 11.9. The reactivity of
a
-
-
OH toward NO
2
is much greater than that of O and it produces
1830.
10
Therefore, attack on nitrite by OH is competitive with O-
NO .
2
JA994230T
attack at most base concentrations. The sum of rate constants for
-
O attack on nitrite ions measured here is about 20% of the liter-
(14) The electronic energies of the two ions in water were calculated with
the Gaussian-98 program using the COSMO solvation model and B3LYP
density function method. Both 6-31+G* and 6-311+G* basis sets were used
and thermochemical corrections to Gibb’s standard free energies were
calculated and applied. Results for 6-311+G* and 6-31+G* were within 1
kcal/mol of each other.
(15) (a) Elliot, A. J. M.; McMcracken, D. R.; Buxton, G. V.; Wood, N. D.
J. Chem. Soc., Faraday Trans. 1990, 86, 1539. (b) Zehavi, D.; Rabani, J. J.
Phys. Chem. 1971, 75, 1738. (c) Han, P.; Bartels, D. M. J. Phys Chem 1990,
94, 7294. (d) Treinin, A. H., E. J. Am. Chem. Soc. 1970, 92, 5821. Gratzel,
M. H., A.; Lilie, J.; Beck, G. Ber. Bunsen-Ges. Phys. Chem. 1969, 73, 646.
(16) (a) Oxides of Nitrogen, IUPAC Solubility Sereies; Young, C. L., Ed.;
Permagon: Oxford, England, 1981; Vol. 8. (b) Hermann, C.; Dewes, I.;
Schumpe, A. Chem. Eng. Sci. 1995, 50, 1673.
ature rate constant.11 The discrepancy may result from changes
in the activity of water at the high base concentrations. It is, how-
-
12
ever, clear that separating the kinetics of OH from O is difficult.
(
9) The IBM, Almaden Research Center, CA, “Chemical Kinetic Simulator
1
.1” program (Hinsberg, W.; Houle, F.; Allen, F.) was used in the simulations.
-
(
10) Previous work shows that HONO
2
is not a bound intermediate.
- 2-
Therefore, the reaction goes to NO
2
and OH instead of H+ and NO
3
.
(
11) Treinin, A.; Hayon, E. J. Am. Chem. Soc. 1970, 92, 5821. Buxton, G.
V. Trans. Faraday Soc. 1969, 65, 2150.
(
(
12) Even at 1 M OH- ∼48% of the attack on nitrite is by OH.
13) S. Lymar, private communication.