Temperature dependence of hydrogen atom reaction with nitrate and nitrite species in aqueous solution

Article abstract of DOI:10.1021/jp970934i

Arrhenius parameters for the reaction of hydrogen atoms with NO₃^-, HNO₂, and NO₂^- in aqueous solution have been determined by the use of pulse radiolysis and electron paramagnetic resonance free induc

Full text of DOI:10.1021/jp970934i

J. Phys. Chem. A 1997, 101, 6233-6237
6233
Temperature Dependence of Hydrogen Atom Reaction with Nitrate and Nitrite Species in
Aqueous Solution†
Stephen P. Mezyk*
Research Chemistry Branch, AECLsWhiteshell Laboratories, Pinawa, Manitoba, R0E 1L0 Canada
David M. Bartels
Chemistry DiVision, Argonne National Laboratory, Argonne, Illinois, 60439
X
ReceiVed: March 13, 1997; In Final Form: June 17, 1997
Arrhenius parameters for the reaction of hydrogen atoms with NO -
, HNO
, and NO
-
in aqueous solution
3
2
2
have been determined by the use of pulse radiolysis and electron paramagnetic resonance free induction
decay attenuation measurements. At 25.0 °C, the calculated rate constants for these compounds are (5.61 (
6
8
9
3
-1 -1
0
.51) × 10 , (3.86 ( 0.09) × 10 , and (1.62 ( 0.05) × 10 dm mol
s
respectively, with corresponding
activation energies of 48.7 ( 1.0 (15.2-84.5 °C), 21.54 ( 0.69 (5.9-62.8 °C) and 15.59 ( 0.36 (6.2-86.8
-
1
°
C) kJ mol . Computer modeling of these systems suggests that ‚H atom reaction with either anion directly
produces hydroxide anion and the corresponding NO radical.
x
Introduction
hydrogen atom with the nitrate and nitrite anions, as well as
nitrous acid. Direct EPR detection of the decay of the hydrogen
The complex decomposition chemistry of nitrate and nitrite
has been investigated for many years, with demonstrated
importance in a variety of fields ranging from atmospheric
atom following pulse radiolysis was the monitoring method of
14,15
choice,
because conventional pulse radiolysis/optical tran-
sient absorption methodology is difficult to use given the weak
absorption of the reactant and product species at very short
1
,2
3
chemistry to high-temperature combustion systems. Recent
studies have especially illustrated the importance of the radia-
tion-induced decomposition of these two compounds in the
nuclear field, where the presence of nitrate, for example, has
been proven to influence the rate of molecular hydrogen
wavelengths. The pulsed EPR-based free induction decay (FID)
16-20
attenuation method
was used because of the pseudo-first-
order scavenging kinetics generally obtained.
4
production in both high level liquid nuclear waste storage and
Experimental Section
upon the start up of nuclear power reactors.5
The procedure used for these EPR experiments has been
Despite over 30 years of investigation, there are still many
uncertainties in the aqueous radiation chemistry of these two
compounds. The hydrated electron has been shown to add to
described in detail in several previous publications,1
6-20
and
thus only a brief description shall be given here. Hydrogen
atoms were generated in aqueous solution within an EPR cavity
by pulse radiolysis, using 3 MeV electrons from a Van de Graaff
accelerator. Stock solutions were prepared by addition of HClO4
6,7
both anions, with established room-temperature rate constants
9
3
-1 -1
9
3
-1 -1
of 9.7 × 10 dm mol
s
and 3.5 × 10 dm mol
s
for
8
nitrate and nitrite, respectively, and the activation energy for
-
3
(
0.10 mol dm , Mallinkrodt A. R. Grade, 69.05%) or borax
-
1
nitrate reaction was determined as 16.0 kJ mol over the
-
2
-3
buffer (1.00 × 10 mol dm , Baker Analyzed) to Millipore-
filtered water. Known volumes of this solution were added to
the recirculating system, and then de-oxygenated using argon
temperature range 283-363 K.9 Nitrite reaction with the
1
0
3
hydroxyl radical occurs with a rate constant of 1.0 × 10 dm
-
1
-1 10
mol s ; however, no reaction of this radical with nitrate
(acid solution) or N2O (borax buffered). Exact acid concentra-
11
has been observed. The many investigations of hydrogen atom
reaction with these two anions have given a wide range of
reported rate constants,8 with all studies performed near room
temperature. The one direct electron paramagnetic resonance
tions were determined by calculation from standardization of
the concentrated acid against 1.029 N HCl (Aldrich, volumetric
standard). The solution was flowed through a flat cell in the
cavity at a rate sufficient to ensure that each cell volume was
completely replaced between pulses. The actual volume irradi-
ated in each pulse was less than 0.10 mL.
,12
(EPR) measurement of hydrogen atom reaction rate constants
1
3
for these two compounds gave a nitrate concentration-
dependent rate constant that was much lower than the competi-
tion-kinetics values.
The approximate average radiation dose for this cell volume
was 1.5, 3.0, and 7.0 krad/pulse for the 12, 25, and 55 ns pulses
used, respectively. For extrapolation to obtain the limiting, zero
dose, rate constants (see later), the relative dose values used
were simply the average beam currents measured on a shutter
positioned before the irradiation cell for the three pulse widths.
These values were checked frequently to monitor any small drift
in the beam. A 35 ns microwave probe pulse was applied to
the sample immediately after irradiation, and the resulting free
In order for the radiation-induced decomposition of nitrate
and nitrite to be properly evaluated under the conditions of
importance, accurate, temperature-dependent, rate constants for
their reaction with all of the primary water radiolysis species
need to be known. In this study, we have measured the
temperature-dependent rate constants for the reaction of the
1
*
To whom correspondence is to be addressed.
Work performed at Argonne under the auspices of the Office of Basic
induction decay of the ‚H atom low-field (mI ) /2) EPR
†
transition recorded on a digital oscilloscope. Typically 500-
Energy Sciences, Division of Chemical Science, U.S. DOE, under Contract
W-31-109-ENG-38.
2000 pulses were averaged to measure each FID, at a repetition
X
Abstract published in AdVance ACS Abstracts, August 1, 1997.
rate of 120 Hz.
S1089-5639(97)00934-1 CCC: $14.00 © 1997 American Chemical Society

6234 J. Phys. Chem. A, Vol. 101, No. 35, 1997
Mezyk and Bartels
Figure 2. Rate constant extrapolations to zero dose for aqueous
hydrogen atom reaction with nitrate at pH 1.0 and 22.3 °C (b) and
Figure 1. Dose dependence of the aqueous hydrogen atom scavenging
rate constant determination for nitrate ion reaction at pH 1.0 and 22.3
6
1.4 °C (9). Error bars shown correspond to one standard deviation
°
C using the Van de Graaff 55 ns (9), 25 ns (b), and 12 ns (2) pulse
obtained from the linear fit to the FID scavenging plots.
widths. Solid lines are linear fits corresponding to calculated rate
6
6
constants of (7.13 ( 0.18) × 10 , (6.04 ( 0.20) × 10 , and (5.13 (
6
3
-1 -1
0
.31) × 10 dm mol
s , respectively.
Scavenging experiments were performed by successive injec-
tions of concentrated, standard solutions of new KNO3 (Aldrich,
9.99%) and NaNO2 (Aldrich, 98.8%) to the pH-adjusted stock
9
water. Accuracy of these concentrations is estimated at better
than 2%.
All ab initio calculations were carried out using the SPAR-
2
1
TAN molecular modelling program of Wavefunction, Inc.
Results
-
‚
H Reaction with NO3 . The general expression for the
effective damping rate of the FID in these experiments is given
1
5,18,20
by
Figure 3. Arrhenius plot of log10 krxn vs 1/T for aqueous hydrogen
1
1
i
ex
)
+ k [S] + k [R ]
(1)
-
-
s
∑
i
atom reaction with NO
3
(9), HNO
2
(b), and NO
2
(2). Solid lines
o
T2(eff)
T2
are weighted linear fits, corresponding to activation energies of (48.7
-
1
(
text).
1.0), (21.54 ( 0.69), and (15.59 ( 0.36) kJ mol , respectively (see
i
ex
where ks[S] is the ‚H atom scavenging rate and ∑k [Ri]
represents the spin-dephasing contribution of second-order spin
exchange and recombination reactions between ‚H atoms and
other free radicals. At the radiation doses typically used in these
experiments the latter term may not be negligible at acid pH’s;
this effect has been observed previously as a slight dose (pulse
TABLE 1: Summary of the Temperature-Dependent Rate
-
Constant Data for Hydrogen Atom Reaction with NO
3
,
-
NO2
2
, and HNO in Aqueous Solution
temp
(°C)
scavenging rate constant
3
-1 -1
species
(dm mol
s )
22-24
width) dependence of the measured scavenging curves.
The
-
6
6
6
7
7
7
7
8
8
8
8
8
9
9
9
9
9
9
9
9
NO
3
15.2
(2.71 ( 0.22) × 10
(4.44 ( 0.10) × 10
(8.70 ( 0.81) × 10
(1.88 ( 0.20) × 10
(3.15 ( 0.37) × 10
(4.85 ( 0.04) × 10
(8.89 ( 3.47) × 10
(1.55 ( 0.13) × 10
(2.12 ( 0.06) × 10
(2.76 ( 0.06) × 10
(3.88 ( 0.07) × 10
(6.49 ( 0.60) × 10
(1.00 ( 0.05) × 10
(1.05 ( 0.09) × 10
(1.27 ( 0.02) × 10
(1.64 ( 0.03) × 10
(2.43 ( 0.05) × 10
(3.21 ( 0.14) × 10
(4.15 ( 0.36) × 10
(4.67 ( 0.16) × 10
overall hydrogen atom reaction rate constant with the nitrate
anion, at pH 1.0 and 22.3 °C,
22.3
30.4
43.7
54.0
61.4
76.5
84.5
5.9
-
‚
H + NO3 f products
(2)
at three different pulse widths (doses) is shown in Figure 1.
Although these scavenging plots exhibit excellent linearity the
HNO
2
6
6
calculated slopes of (5.13 ( 0.31) × 10 , (6.04 ( 0.20) × 10
14.7
25.4
6
3
-1 -1
and (7.13 ( 0.18) × 10 dm mol
s
for the 12, 25, and 55
4
6
3.2
2.8
6.2
4.7
ns pulses, respectively, show that such spin exchange does affect
the measured rate constants for this system.
-
2
NO
To correct the observed rate constants for this dose depen-
dence, limiting values were calculated by extrapolation to zero
dose, as shown in Figure 2. An excellent linear relationship
was obtained, and for the rate constants listed above a limiting
1
24.0
44.5
6
3.4
76.6
6.8
6
3
-1 -1
value of (4.44 ( 0.10) × 10 dm mol
procedure was then repeated over the temperature range 15.2-
4.5 °C (values for 61.4 °C also shown in Figure 2), with all
s
was obtained. This
8
8
than -1 over the entire temperature range of measurement,25
all of the reaction can be attributed to the nitrate anion. From
a linear fit to these values, the temperature-dependent rate
the extrapolated values given in Table 1 and shown in the
Arrhenius plot of Figure 3. Because the pK of HNO3 is less

H Atom Reaction with NOx in Aqueous Solution
constants are well described by the expression
J. Phys. Chem. A, Vol. 101, No. 35, 1997 6235
-
‚H Reaction with NO2 . To ensure complete reaction with
-
-2
only NO2 , measurements were performed in 1.00 × 10 mol
-
3
log k ) (15.28 ( 0.16) - [(48700 ( 1000)/2.303RT] (3)
dm borax buffer, whose temperature dependence has been
1
0
2
25
well characterized. Under these conditions, the hydrogen atom
concentration produced was sufficiently low that there was little
dose effect; therefore, signals were only obtained using the 55
ns pulse of the Van de Graaff accelerator.
with k2 and T in units of dm mol s-1 and K, respectively.
3
-1
-
1
This corresponds to an activation energy of 48.7 ( 1.0 kJ mol .
There have been many previous investigations of the reaction
of the hydrogen atom with the nitrate anion in aqueous solution
at room temperature, with rate constants reported over the range
The rate constant measured at pH 9.2 and 24.0 °C by this
method for the reaction
6
3
-1 -1
(
2.9- 40) × 10 dm mol
s
using a variety of competition
1
,26-35
-
kinetic systems.
Our room temperature (22.3 °C) value
‚H + NO2 f products
(6)
6
3
-1 -1
of (4.44 ( 0.10) × 10 dm mol
s
is at the low end of this
was k ) (1.64 ( 0.03) × 10 dm mol s-1. This value is
9
3
-1
range. There have also been two determinations at lower
temperatures. A direct evaluation at 8 ( 2 °C was performed
by Smaller et al.,1 who obtained a concentration-dependent rate
constant from the decay characteristics of the ‚H atom EPR
signal in nitrate solutions, giving an extrapolated rate constant
6
approximately the average of the reported literature val-
3
26-28,32,36,39,40
9
3
-1
ues,
which range from 0.2-2.8 × 10 dm mol
-
1
s . The EPR measurements were then repeated over the
temperature range 6.2-86.6 °C, with the individual rate
constants given in Table 1 and plotted in Figure 3. The
temperature dependence for this reaction was found to be well
described by the equation
6
3
-1 -1
of k ) 1.4 × 10 dm mol s . This value is slightly lower
6
than our calculated rate constant of k2 ) (1.71 ( 0.18) × 10
3
-1 -1
dm mol s , but there is good agreement within the combined
error of the two studies. An earlier determination by Navon
and Stein27 was based on the measurement of total nitrate loss
in aqueous solution at 4 °C, due to its reaction with hydrogen
radicals created by passing H2 gas through a discharge system.
The rate constants determined by these experiments showed a
very strong pH dependence, with calculated values of (1.15 (
log k ) (11.94 ( 0.06) - [(15590 ( 360)/2.303RT] (7)
1
0
6
-1
corresponding to an activation energy of 15.59 ( 0.36 kJ mol .
Again, these Arrhenius parameters are much lower than those
determined for the nitrate system.
7
6
6
3
0
mol
.22) × 10 , (9.4 ( 0.3) × 10 , and (3.8 ( 1.7) × 10 dm
There has only been one previous determination at lower
-
1
-1
13
s
for pH’s 1-4.7, 7, and 11.6-13.3, respectively. All
temperatures, by Smaller et al. who directly measured a rate
8
3
-1 -1
of these rate constants are much higher than our extrapolated
constant of k ) 7.1 × 10 dm mol
s
at 281 K. This value
6
3
-1 -1
value of (1.27 ( 0.13) × 10 dm mol
s
for this temperature.
is slightly lower than our calculated value of k ) (1.10 ( 0.04)
6
9
3
-1 -1
‚
H Reaction with HNO2. Analogous experiments were
× 10 dm mol
s
at this temperature, but again is within
performed for nitrite. Because the pKa of HNO2 is significantly
higher, pKa ) 3.2 at 25 °C,25 both the acid and anion reactions
could be separately evaluated.
the combined errors of the two studies.
Discussion
Because the dose dependence obtained for nitrate was also
observed for nitrite, rate constants were again extrapolated to
zero dose. The value obtained at pH 1.0 and 25.4 °C by this
method for the reaction
In previous interpretations of radiation and photochemistry
6
,41-44
of nitrite and nitrate solutions
it has generally been
assumed that ‚H atoms would add to the ion, producing the
protonated free radical:
‚
H + HNO f H O + NO
(4)
2
2
H + NO3 f HNO3•-
-
‚
(8)
(9)
was k4 ) (3.88 ( 0.07) × 10 dm mol s-1. This rate constant
is somewhat lower than the single literature value obtained by
ethanol competition kinetics at 26 ( 1 °C of k ) (5.47 ( 0.26)
8
3
-1
‚H + NO2 f HNO2•-
-
8
3
-1 -1 36
×
10 dm mol
s
(allowing for the greater rate constant
Given the reported observation of both of these ion radicals
and determination of their pKa’s by transient absorption
spectroscopy, this was a very reasonable assumption. How-
ever, the activation energy of 48.7 ( 1.0 kJ mol and
37
of hydrogen atom reaction with ethanol at this temperature).
Rate constants were then determined under these conditions
over the temperature range 5.9-62.8 °C, with the individual
extrapolated values again given in Table 1 and plotted in Figure
6,41
-
1
1
5
3
-1 -1
preexponential factor of 1.91 × 10 dm mol
s
found for
3
. Experiments at temperatures greater than 63 °C were not
reaction 2 are both much larger than expected for a simple
hydrogen atom reaction. Similar activation energies and pre-
exponential factors have been observed previously for the
reproducible and gave rate constants that were far higher than
expected from the extrapolation of lower values in the Arrhenius
plot. This behavior is attributed to the known thermal instability
of HNO2 at elevated temperatures.38 From a linear fit to the
measured values, the temperature-dependent rate constant was
found to be well described by the expression
4
5
+ 46
hydrogen atom reaction with H5IO6 and N2H5 . For the
periodic acid complex, the reaction was explained in terms of
a concerted oxygen abstraction, dehydration, and iodine complex
rearrangement on the basis of the known ∆H° and ∆S° of the
periodate hydration equilibria. From a similar analysis the
hydrogen atom reaction with the acid form of hydrazine was
postulated to result in a dissociation of the transition complex
to give ammonia, ‚NH2, and a proton.
In low-level (Hartree-Fock/6-31G* basis) ab initio calcula-
tions, reactions 8 and 9 as written are exothermic by some 53.1
kJ, yielding a stable product. However, addition of polarization
basis functions (p-type functions on ‚H atoms: 6-31G** or
log k ) (12.36 ( 0.12) - [(21540 ( 690)/2.303RT] (5)
1
0
4
-1
corresponding to an activation energy of 21.54 ( 0.69 kJ mol .
Both the activation energy and preexponential factor are much
smaller than determined for nitrate reaction.
There is little doubt that the products of reaction 4 are H2O
and NO as written. In low level ab initio calculations (HF/6-
1G**) a transition state for ‚H atom reaction at oxygen is found
at the large H‚‚‚OHNO distance of 0.13-0.14 nm.
-
3
6-311G**) produce dissociation to OH and NO or NO2. For
the reaction with nitrite, a transition state is found at an

6236 J. Phys. Chem. A, Vol. 101, No. 35, 1997
Mezyk and Bartels
-
ONO‚‚‚H distance of 98 pm. For the reaction with nitrate, the
reaction at oxygen is simply downhill, with no transition state.
It seems unlikely that any transition state would appear at higher
levels of theory, and the high experimental activation energy
must correspond to a qualitatively different reaction due to the
presence of solvent. Calculations required to investigate specific
solvent effects are beyond the scope of this work.
apparently lead to the creation of two OH ions in contact, but
the Grotthus mechanism of proton transfer also allows for
transfer via a bridge of H-bonded water molecules,4 and this
may provide a more reasonable picture. The only obvious
similarity between this proton transfer reaction and the reaction
of ‚H atom with nitrate is the breaking of the N-O bond and
formation of ‚NO2 as a product. The latter aspect of both
reactions may be responsible for the positive activation entro-
pies. At least three strong hydrogen bonds to nitrate are broken,
giving a large entropy increase. The ‚NO2 radical is not
significantly hydrogen bonded, as indicated by its facile
9,40
•
-
•-
If HNO2 and HNO3 do exist in the liquid, it must be due
to solvent stabilization. We should recall in any case that these
radicals and their respective conjugate bases, NO22 and NO3
-
2-
,
are metastable in water, reportedly dissociating according to
4
8
reactions 10-13, with the indicated first-order half-lives at room
diffusion in ice at low temperatures.
6
,41
temperature:
Summary
HNO2•- f NO + OH-
τ1/2 < 12 µs
τ1/2 ) 12 µs (11)
τ1/2 ) 3 µs (12)
(10)
Arrhenius parameters have been established for aqueous
hydrogen atom reaction with the nitrate anion, nitrous acid, and
the nitrite anion by direct experimental measurement as
2
-
-
NO2 + H O f NO + 2OH
2
HNO3•- f NO + OH-
log k ) (15.28 ( 0.16) - [(48700 ( 1000)/2.303RT] (3)
10
2
2
log k ) (12.36 ( 0.12) - [(21540 ( 690)/2.303RT] (5)
2
-
-
10
4
NO3 + H O f NO + 2OH
τ1/2 ) 12.5 µs (13)
2
2
log k ) (11.94 ( 0.06) - [(15590 ( 360)/2.303RT] (7)
1
0
6
Barker et al.43 found by conductivity detection of the OH
-
2-
product that reaction 13, of NO3 with water, was characterized
by 49.4 kJ mol activation energy and a large preexponential
respectively. On the basis of the Arrhenius parameters for
nitrate anion and the results of ab initio computer modeling of
this reaction, it is believed that the previously proposed reaction
mechanism, consisting of simple addition of the hydrogen atom
to produce a metastable radical intermediate, cannot be correct.
-
1
1
3
3
-1 -1
factor of 4.3 × 10 dm mol s . The activation energy of
NO22 hydrolysis was much smaller. The similarity of these
activation parameters to those found for ‚H reaction with nitrate
and nitrite suggests a commonality of mechanism. If the
-
43
-
Rather, the reactions proceed directly to the ultimate OH and
-
immediate products (OH and NO2 or NO) of the two reaction
NOx products.
sets are the same, the transition states and activation energies
might be somewhat similar, even though the reactants appear
to be quite different.
According to semiclassical transition state theory (omitting
tunneling effects), the Arrhenius preexponential factor corre-
sponds to
Acknowledgment. The authors thank David Werst for his
assistance in operating and maintaining the Van de Graaff
accelerator. We acknowledge Rich Marassas and Andy Cook
for performing transient absorption experiments at Argonne
National Laboratory and Dan Meisel and Chuck Jonah for useful
discussions.
q
kT
A ) exp
h
∆S
R
+ 1
(14)
(
)
References and Notes
q
where ∆S is the activation entropy. Based on this formula,
the reaction of the ‚H atom with NO3 must be characterized
by a large increase in entropy (+39.2 J mol K ) as the
transition state is approached. The entropy change is dominated
by changes in low-frequency vibrations and librations and,
especially, changes in the hydrogen bonding.
(1) Løgager, T.; Sehested, K. J. Phys. Chem. 1993, 97, 6664.
(2) Burkholder, J. B.; Mellonki, A.; Ranajit, T.; Ravishankara, A. R.
Int. J. Chem. Kinet. 1992, 24, 711.
-
-
1
-1
(3) Tsang, W.; Herron, J. T. J. Phys. Chem. Ref. Data 1991, 20, 609.
(4) Meisel, D.; Diamond, H.; Horowitz, E. P.; Jonah, C. D.; Matheson,
M. S.; Sauer, M. C., Jr.; Sullivan, J. C.; Barnabas, D.; Cerny, E.; Cheng,
Y. D. Radiolytic Generation of Gases from Synthetic Waste. Technical
Report ANL-91/14; Argonne National Laboratory: Argonne, IL, 1991.
The hydrogen bonding of the nitrate ion has been studied
intensely by NMR relaxation methods and infrared spectros-
(5) Singh, A.; LeBlanc, J. C. Use of Gadolinium as a Neutron Absorber
in Nuclear Reactors: Radiolysis of Gadolinium Nitrate Solutions. Technical
Report AECL-5526; AECLsWhiteshell Laboratories: Pinawa, MB, Canada,
copy.4 Study of NO2 has not been pursued to the same extent.
7
-
1
976.
It is generally acknowledged that the nitrate ion is strongly
hydrogen bonded to water, and this accounts for its relatively
(6) Gr a¨ tzel, M.; Henglein, A.; Taniguchi, S. Ber. Bunsen-Ges. Phys.
Chem. 1970, 74, 292.
(7) Kevan, L. In Radiation Chemistry of Aqueous Systems; Stein, G.,
Ed.; Wiley: Chichester, 1968.
4
7
slow rotational relaxation in the plane of the molecule.
2-
Likewise, strong hydrogen bonding of the NO3 radical to water
has been suggested by its EPR relaxation behavior in irradiated
nitrate/water glasses.48 We can speculate that the direct
energetically “downhill” attack of the ‚H atom on the nitrate
ion oxygen is prevented by the presence of the hydrogen bonds
to surrounding water molecules and that a much higher barrier
(8) Buxton, G. V.; Greenstock, C. L.; Helman, W. P.; Ross, A. B. J.
Phys. Chem. Ref. Data 1988, 17, 513 and references therein.
(9) Chen, R.; Avotinsh, Y.; Freeman, G. R. Can. J. Chem. 1994, 72,
083 and references therein.
1
(10) Values taken as an average from the reported data in: Barker, G.
C.; Fowles, P.; Stringer, B. Trans. Faraday Soc. 1970, 66, 1509. Treinin,
A.; Hayon, E. J. Am. Chem. Soc. 1970, 92, 5821. Buxton, G. V. Trans.
Faraday Soc. 1969, 65, 2150. Adams, G. E.; Boag, J. W.; Michael, B. D.
Trans. Faraday Soc. 1965, 61, 1417. Adams, G. E.; Boag, J. W. Proc.
Chem. Soc. London, 1964, 112.
(11) Broszkiewicz, R. K. Int. J. Appl. Radiat. Isot. 1967, 18, 25.
(12) Anbar, M.; Farhataziz; Ross, A. B. Selected Specific Rates of
Reactions of Transients from Water in Aqueous Solutions: II. Hydrogen
Atom (NSRDS-NBS 51); National Bureau of Standards, U.S. Department
of Commerce, U.S. Government Printing Office: Washington, DC, 1975
and references therein.
(perhaps out-of-plane) attack is entropically favored.
2
-
It is interesting to consider how the reaction of the NO3
radical with water might occur, and what similarity this might
have to the reaction of the ‚H atom with NO3 . On the basis
of the stoichiometry of reaction 13, two hydroxide ions must
be formed. One of these must be formed via proton transfer
from a water molecule to the nitrate ion’s oxygen. This would
-

H Atom Reaction with NOx in Aqueous Solution
J. Phys. Chem. A, Vol. 101, No. 35, 1997 6237
(
13) Smaller, B.; Avery, E. C.; Remko, J. R. J. Chem. Phys. 1971, 55,
(33) Volkert, W. A.; Kuntz, R. R. J. Phys. Chem. 1968, 72, 3394.
(34) Hickel, B., CEA-R-4046, 1970, Saclay, France.
(35) Draganic, Z. D.; Draganic, I. G. J. Phys. Chem. 1972, 76, 2733.
(36) Halpern, J.; Rabani, J. J. Am. Chem. Soc. 1966, 88, 699.
(37) Mezyk, S. P.; Bartels, D. M. J. Phys. Chem. 1997, 101, 1329.
(38) Durrant, P. J.; Durrant, B. Introduction to AdVanced Inorganic
Chemistry, 2nd ed.; J. Wiley and Sons, Inc.: New York, 1970; p 699.
(39) Rabani, J.; Stein, G. J. Chem. Phys. 1962, 37, 1865.
(40) Sutton, J.; Moreau, M. Proceedings of the 2nd Tihany Symposium
on Radiation Chemistry; Dobo, J., Hedrig, P., Eds.; Publishing House of
the Hungarian Academy of Sciences: Budapest, 1967.
(41) Gr a¨ tzel, M.; Henglein, A.; Lilie, J.; Beck, G. Ber. Bunsen-Ges. Phys.
Chem. 1969, 73, 646.
2
414.
(14) Fessenden, R. W.; Varma, N. C. Faraday Discuss. Chem. Soc. 1977,
6
3, 104.
(
15) Bartels, D. M.; Craw, M. T.; Han, P.; Trifunac, A. D. J. Phys.
Chem. 1989, 93, 2412.
(
(
16) Bartels, D. M.; Mezyk, S. P. J. Phys. Chem. 1993, 97, 4101.
17) Roduner, E.; Bartels, D. M. Ber. Bunsen-Ges. Phys. Chem. 1992,
9
6, 1037.
(
(
(
(
18) Han, P.; Bartels, D. M. J. Phys. Chem. 1990, 94, 7294.
19) Han, P.; Bartels, D. M. Chem. Phys. Lett. 1989, 159, 538.
20) Han, P.; Bartels, D. M. J. Phys. Chem. 1992, 96, 4899.
21) Hehre, W. J.; Huang, W. W. Chemistry with Computation: An
introduction to SPARTAN. SPARTAN, version 4.1; Wavefunction, Inc.:
Irvine, CA, 1995.
(42) Gonzalez, M. C.; Braun, A. M. Res. Chem. Intermed. 1995, 21,
837.
(
22) Mezyk, S. P.; Bartels, D. M. J. Chem. Soc., Faraday Trans. 1995,
(43) Barker, G. C.; Fowles, P.; Stringer, B. Trans. Faraday Soc. 1970,
66, 1025.
(44) Daniels, M.; Wigg, E. E. J. Phys. Chem. 1967, 71, 1024.
(45) Mezyk, S. P.; MacFarlane, R.; Bartels, D. M. J. Phys. Chem. 1994,
98, 12594.
9
1, 3127.
(
(
(
23) Mezyk, S. P.; Bartels, D. M. Can. J. Chem. 1994, 72, 2516.
24) Mezyk, S. P.; Bartels, D. M. J. Phys. Chem. 1994, 98, 10578.
25) Perrin, D. D. Ionisation Constants of Inorganic Acids and Bases in
Aqueous Solution, 2nd ed.; IUPAC Chemical Data Series No. 29; Pergamon
Press: New York, 1982.
(46) Mezyk, S. P.; Tateishi, M.; MacFarlane, R.; Bartels, D. M. J. Chem.
Soc., Faraday Trans. 1996, 92, 2541.
(47) Laaksonen, A.; Kovacs, H. Can. J. Chem. 1994, 72, 2278 and
references therein.
(26) Appleby, S.; Scholes, G.; Simic, M. J. Am. Chem. Soc. 1963, 85,
3
891.
(
(
(
(
(
(
27) Navon, G.; Stein, G. J. Phys. Chem. 1964, 69, 1384.
28) Hayon, E.; Moreau, M. J. Chim. Phys. 1965, 62, 391.
29) Sawai, T. Bull Chem. Soc. Jpn, 1966, 39, 955.
30) Chouraqui, M.; Sutton, J. Trans. Faraday Soc. 1966, 62, 2111.
31) Airey, P. L.; Dainton, F. S. Proc. R. Soc. London 1966, A291, 478.
32) Anbar, M.; Neta, P. Int. J. Appl. Radiat. Isotop. 1967, 18, 493.
(48) Symons, M. C. R.; Zimmerman, D. N. J. Chem. Soc., Faraday
Trans. I 1976, 72, 409.
(49) Grunwald, E.; Eustace, D. In Proton Transfer Reactions; Caldin,
E., Gold, V., Eds.; Chapman and Halls Ltd.: London, 1973; p 103.
(50) Bamford, C. H.; Tipper, C. F. H. Proton Transfer. In ComprehensiVe
Chemical Kinetics; Elsevier Scientific: New York, 1977; Vol. 8.