A novel long path photolysis cell - Application to the reactivity of selected organic compounds toward the nitrate radical (NO3)

Article abstract of DOI:10.1039/b400175c

The atmosphere can be considered as a giant chemical reactor where gases and aerosols play a key role in all chemical and physical processes linked to the Earth's climate. Bimolecular rate coefficients for the reactions of the nitrate radical, NO₃, with methanol, ethanol, acetaldehyde, tert-butyl methyl ether, propionic acid, dimethylmalonate, dimethylsuccinate, dimethyl carbonate, and diethylcarbonate in aqueous solutions were measured using a novel experimental approach. All experiments were conducted at room temperature. The reported rate coefficients for the reactions of NO₃ with methanol, ethanol, and acetaldehyde agreed well with currently recommended values. A major assumption in the reactions was that the NO₃ attack proceeded via an H-abstraction mechanism. The existing data did not support any reliable estimation method, which warrants future work to improve knowledge that is required building reliable structure-activity relationships.

Full text of DOI:10.1039/b400175c

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R E S E A R C H P A P E R
A novel long path photolysis cell—application to the reactivity of
selected organic compounds toward the nitrate radical (NO3)
Davy Rousse and Christian George*
`
Laboratoire d’Application de la Chimie a l ’Environnement (UCBL-CNRS) 43 boulevard du 11
Novembre 1918, F-69622, Villeurbanne, France. E-mail: christian.george@univ-lyon1.fr;
Fax: þ33 (0)4 72 44 8114; Tel: þ33 (0)4 72 43 1489
Received 7th January 2004, Accepted 2nd March 2004
F|rst published as an AdvanceArticle on the web14th April 2004
Bimolecular rate coefficients for the reactions of the nitrate radical, NO3 , with methanol, ethanol, acetaldehyde
tert-butyl methyl ether, propionic acid, dimethylmalonate, dimethylsuccinate, dimethyl carbonate and
diethylcarbonate in aqueous solutions have been measured using a novel experimental approach. This study
was performed using laser flash photolysis (LFP) with a capillary made of Teflon AF2400 as a long path
capillary photolysis cell. Taking benefit of this new material allowing a long optical path length for a very
limited irradiated solution volume, new rate constants were established. All experiments were carried out at
room temperature. Measured rate coefficients for the reaction of the NO3 radical with methanol, ethanol,
acetaldehyde, dimethylcarbonate, diethylcarbonate, dimethylmalonate, dimethyl succinate, propionic acid and
tert-butyl-methyl ether are (units are 105 Mꢀ1
s
ꢀ1): kNO þ ethanol ¼ 4.8 ꢁ 0.5, kNO þ ethanol
¼
3
3
19 ꢁ 3, kNO þ acetaldehyde ¼ 20 ꢁ 3, kNO þ dimethylcarbonate ¼ 0.15 ꢁ 0.04, kNO þ diethylcarbonate ¼ 0.84 ꢁ 0.12,
3
3
3
kNO þ dimethylmalonate ¼ 0.26 ꢁ 0.07, kNO þ dimethylsuccinate ¼ 0.34 ꢁ 0.02, kNO þ propionic acid ¼ 0.77 ꢁ 0.02,
3
3
3
kNO þ tert-butyl-methyl ether ¼ 3.9 ꢁ 1.3. The uncertainties in the above expressions are ꢁ2s and represent precision
3
only. The reported rate coefficients for the reactions of NO3 with methanol, ethanol and acetaldehyde agree well
with currently recommended values. To date, there is no kinetic data reported in the literature for the NO3
radical reaction with dimethylcarbonate, diethylcarbonate, dimethylmalonate, dimethylsuccinate,
propionic acid and tert-butyl-methyl ether. The reaction mechanism is briefly discussed as a function
of bond energies
especially for the reactions of OH with organics containing
oxygenated polyfunctional groups.
Introduction
The atmosphere can be considered as a giant chemical reactor
where gases and aerosols play a key role in all chemical and
physical processes linked to the Earth’s climate. The presence
of condensed matter renders this reactor quite inhomogeneous
with complex exchanges between all phases encountered in the
atmosphere.1
In the atmospheric gas phase, the nitrate radical (NO3) has
been identified as a key oxidant, especially at night when the
importance of the hydroxyl radical (OH) is decreasing.
Accordingly, numerous studies focused on the kinetics and
on the addition and hydrogen-abstraction mechanisms in
which NO3 is involved.2,3
Such interest stems from the fact that the intermediate
length carbon chain and oxygenated hydrocarbon compounds
are expected to be increasingly used as new industrial solvents
to favour water based technology. Currently, C2–C10 dibasic
esters6–8 and carbonates9,10 are being considered to replace
‘‘old’’ additives to obtain ‘‘new’’ environmentally friendly
solvents. However, the atmospheric degradation of dibasic
esters and carbonates will potentially result in the formation
of tropospheric ozone and other secondary pollutants.7 More-
over, the C2–C10 oxygenated organics can potentially form
low vapour pressure secondary reaction compounds that
can be either taken up by clouds or form secondary aero-
sols11,12 (SOA). In turn, these compounds can be removed
from the atmosphere via rainfall, dry or wet deposition and
potentially act as a source of organic pollutants in surface
waters.
However, since the Earth’s atmosphere is far from being
purely gaseous as it contains also aqueous particles, i.e. the
so-called hydrometeors (clouds and fogs). As the nitrate radi-
cal may undergo phase partitioning4 or has direct sources in
the aqueous phase, many of its reactions are also occurring
in water.2 Therefore, the nitrate radical is a key oxidant also
in the aqueous phase, along with the hydroxyl radical (OH).
In our previous paper,5 we discussed the reactivity of the
hydroxyl radical toward some oxygenates, and new rate coeffi-
cients were established for the OH þ dimethylmalonate, OH þ
dimethylsuccinate, OH þ dimethylcarbonate and OH þ diethyl-
To date, the atmospheric fate of C2–C10 dibasic esters and
carbonates remains uncertain. In the atmosphere, these com-
pounds are likely to be involved in multiphase chemistry and
be removed via the attack by OH or NO3 .
In this paper, we present new determinations of the kinetics
of the hydrogen abstraction reactions between NO3 and some
oxygenates, i.e. methanol, ethanol, acetaldehyde tert-butyl
methyl ether, propionic acid, dimethylmalonate, dimethylsuc-
cinate, dimethyl carbonate and diethylcarbonate. This study
has been performed using laser flash photolysis (LFP) in a
capillary made of Teflon AF2400 as a long path capillary
photolysis cell. Taking benefit of this new material allowing
a long optical path length for a very limited irradiated solution
volume, new rate constants were established.
carbonate reactions, followed by
a discussion on the
H-abstraction mechanism. In the present study however, we
focused on the reactivity of NO3 toward a few oxygenates.
Although many studies have dealt with the kinetics of the
reaction between NO3 and a series of organic compounds, very
few studies have been performed for the reactions of NO3 with
the intermediate-sized aliphatic compounds (i.e. C2–C10), and
T h i s j o u r n a l i s Q T h e O w n e r S o c i e t i e s 2 0 0 4
3408
P h y s . C h e m . C h e m . P h y s . , 2 0 0 4 , 6 , 3 4 0 8 – 3 4 1 4

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Details of the experimental procedure that were employed to
study these reactions in aqueous solutions are given below.
beam was subsequently collected by a Tektronix oscilloscope
(type 2430A) via a differentiating circuit, which allows the
change of absorption to be measured with a high sensitivity.
This oscilloscope allows a maximum sampling rate of 100
MHz and has signal-averaging capacities. The full absorption
may then be reconstructed from both continuous and transient
signals once both of them have been transferred to a computer
via IEEE-488 GPIB interface from the oscilloscope and multi-
meter. All devices were synchronized via a delay generator
(DG535, Stanford Research System) which was operated as
the master in our optical chain.
The signal to noise ratio of the measured kinetic curves was
such that in most cases experiments were performed on a single
shot basis. However when necessary (i.e., for low NO3 concen-
trations), the averaging capacities of the oscilloscope were
used. In that case, up to 16 single shot experiments were aver-
aged before transmitting to the computer. In all cases, the
photolysis cell content was fully renewed between each laser
pulse. We ensured that single and averaged signals led to the
same kinetic determination.
Experimental
Laser photolysis apparatus
The NO3 radical was produced by the laser-flash photolysis of
peroxodisulfate or cerium ammonium nitrate precursors.2 The
experimental setup described in this section has been applied
for photolytic generation of sulfate radical from peroxodisul-
fate. Sulfate radicals were then converted into nitrate radical
in the presence of excess nitrate anion:
ꢂ
ꢀ
S2O82ꢀ þ hn ! 2SO4
ðR1Þ
ðR2Þ
ꢂ
ꢂ
SO4 ꢀ þ NO3 ! SO42ꢀ þ NO3
ꢀ
Using reactions (R1) and (R2) implies a certain time to achieve
complete formation of the nitrate radical. This time the gap
depends on laser fluence and concentration of reactants.
Therefore, depending on experimental conditions, the decay
of the nitrate radical may not necessarily be of first order dur-
ing the first microseconds following the laser discharge (as
shown in Fig. 2). Therefore, care must be taken when analys-
ing NO3 decays to be in a time region where the production of
NO3 is complete.
Nitrate radicals were also produced from the photolytic
reaction on cerium nitrate ammonium complex, as an alterna-
tive to the experimental method mentioned above. Solutions
of Ce(IV) complex in nitric acid were used to generate nitrate
radical.
Teflon AF photolysis cell
The centrepiece of the experimental set-up is a liquid core
waveguide (LCW) made of Teflon AF 2400 (BioGeneral,
San Diego, CA). Tubing made from these polymers have
been shown to exhibit excellent optical properties such as
high optical clarity (at l ꢅ 200 nm more than 80% of light
is transmitted through a 220 mm thick film of the polymer
described above) and very low refractive index (i.e.,
n ¼ 1.29 for Teflon AF 2400 grades). The refractive index
observed for the Teflon AF 2400 tube is lower than virtually
all standard temperature and pressure liquids. As a result,
tubing made from the Teflon AF 2400 polymer and filled
with water will conduct light.
The Teflon AF 2400 guide had an inner and outer diameter
of 200 and 800 mm, respectively. The length of the LCW was
varied in the course of the study from 30 to 90 cm. In all cases,
the internal volume of the liquid core waveguide was less than
0.1 ml for a maximum optical path length of 90 cm. Very
small liquid volumes used in the LCW (liquid core waveguide)
is one of the major advantages of using waveguides compared
to a standard White cell design for studying radical liquid
phase reactions. A good White cell can achieve a maximum
path length of a few meters (which often requires expensive
lasers) in a volume of about 50–100 ml. Such a large reaction
volume requires much larger flow rates to replenish the cell
between laser flashes. Under the typical experimental condi-
tions employed, the LCW can achieve much longer optical
path lengths in much smaller volumes without requiring as
powerful light sources as those required in standard White cell
systems.
ꢂ
2ꢀ
½CeivðNO3Þ ꢃ þ H2O þ hv ! ½CeiiiðNO3Þ H2Oꢃ2ꢀ þ NO3
6
5
ðR3Þ
Typical Ce(IV) complex concentration was in the range
5 ꢄ 10ꢀ4–5 ꢄ 10ꢀ3 mol Lꢀ1. Typical experimental nitric acid
concentration was 1.0 mol Lꢀ1
.
A Lambda Physik Excimer Multigas laser (EMG 101), oper-
ated with a Xe/F2 mixture leading to a lasing wavelength at
351 nm was used as the photolysis source. Under optimum
conditions, this laser produces a pulse half-width of 16 ns
and a beam cross-sectional area of 250 mm2 with an output
of 125 mJ per pulse. This energy was periodically checked
using a calibrated Joule-meter (Labmaster E, Coherent Inc.).
Transient species produced by the pulsed laser beam were
monitored by means of time resolved absorption spectroscopy.
A diode laser, with an output wavelength of 635 nm and a
power of 1 mW was used as the analyzing light. The light
escaping the photolysis cell (described below) was fed onto
the active surface of an amplified photodiode. The rise time
(5 to 95%) of the amplified photodiode was always kept below
0.5 ns.
The Teflon photolysis cell used is shown in Fig. 1. The
highly flexible Teflon AF 2400 tubing was loosely coiled (less
than 3 cm diameter) and placed in the excimer laser beam path.
The NO3 radicals were generated within the LCW following
the laser flash on the precursors, as discussed above. Since
The continuous optical signal was first measured before the
laser discharge by a sample-and-hold circuit and the corre-
sponding voltage was read by a multimeter (Keithley 179A).
The transient optical absorption signal produced by the laser
Fig. 1 Schematics of the LCW (liquid core waveguide) micro-flow tube experimental setup.
T h i s j o u r n a l i s Q T h e O w n e r S o c i e t i e s 2 0 0 4
P h y s . C h e m . C h e m . P h y s . , 2 0 0 4 , 6 , 3 4 0 8 – 3 4 1 4
3409

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the waveguide was used as a coil, no concentration gradient
could built up along the waveguide length. However, non-
uniform nitrate radical concentration could be produced if
the laser fluence experienced by the coild was not uniform.
Therefore, in order to avoid any shielding of the front and
back sides of the waveguide, the later was placed in a reactor
(not shown in Fig. 1) the inner walls of which were covered
by an aluminium foil. The reactor had the following dimen-
sions 1 ꢄ 3 ꢄ 3 cm3 (height ꢄ width ꢄ depth). In this way we
ensured an increased uniformity in the irradiation of the wave-
guide. However, we could not measure this uniformity,
we therefore made sure that all kinetics studied in the present
study were unimolecular (after completion of the NO3
production as stated before).
In order to probe the solution contents of the Teflon photo-
lysis cell, the output of the diode laser was focused on the entry
of an unpolished 100-mm diameter fused silica optical fibre.
The fused silica optical fibre was held in the liquid content
of the photolysis cell. Then, the light was allowed to escape
the solid optical fibre and was collected in the liquid core wave-
guide. Again, the Teflon AF tubing conducted the light up to
its end where another fused silica optical (located in the liquid)
collected most of this transmitted light. This second solid opti-
cal then conducted the light to the active surface of the ampli-
fied photodiode. In that way the measured signal is integrated
(or averaged) over the full length of the waveguide. As the
observed decays were first order, this averaging allows the
determination of rate constants even in the unlikely situation
where the NO3 concentration was not uniform along the
waveguide length.
99.5%), dimethyl malonate (Aldrich > 99%), dimethyl succi-
nate (Aldrich > 99%), dimethyl carbonate (Fluka > 99%),
diethyl carbonate (Fluka > 99%) and tert-butyl-methyl ether
(Aldrich > 99.8%). Water was taken from an 18 MO purifica-
tion system (Millipore). In some test experiments, we used oxy-
gen free solutions, which led to the same results as non-
degassed water.
Results
Validation of the technique
In order to check the consistency of the results obtained by this
technique, several well-established rate coefficients have been
re-measured.
NO3 + methanol
The kinetics of the reaction of OH with methanol (R4) has
been studied before (see, for example, Table 1).
NO3 þ CH3OH ! HNO3 þ CH2OH
ðR4Þ
The agreement among the reported values for the rate coeffi-
cient for the reaction of OH with methanol is not very clear as
values are covering more or less an order of magnitude (see
Table 1) As a result, further investigations are warranted to
better define the rate coefficient for the removal of OH by
methanol.
A typical decay for the absorption profiles of the nitrate
radical is shown in Fig. 2, whereas a typical second order plot
for this reaction is shown in Fig. 3. The kinetics were of
pseudo-first order (and treated accordingly) as the methanol
was in large excess (this will be the case of the oxygenates con-
sidered in this study). The methanol concentration was varied
from 1 ꢄ 10ꢀ2 to 4 ꢄ 10ꢀ2 mol Lꢀ1. At 298 K, the measured
The use of a long optical path-length, enabled us to work
with low NO3 concentrations. Accordingly, the UV irradiation
of the nitrate precursor never led to a saturation of the latter as
the nitrate radical formed was much less than 1% of the pre-
cursor concentration. The decay of the nitrate radical was
monitored at a wavelength of 635 nm and was believed to obey
Beer–Lambert’s law.
rate constant is reported to be (4.8 ꢁ 0.5) ꢄ 105 Mꢀ1 sꢀ1
.
The previously reported rate constants for reaction R4 are in
the range from 1.8 to 16 ꢄ 105 Mꢀ1
s
ꢀ1, with the latest deter-
mination by Herrmann and coworkers20 being (5.1 ꢁ 0.3) ꢄ
Determination of rate coefficients
105 Mꢀ1 sꢀ1
.
As this will be discussed in more detail in the results section,
many experimental conditions have been changed (laser vol-
tage discharge, solution composition, temperature) leading to a
total of 200 experiments. In all cases reported here, the decays
were purely first order decays and therefore straightforward
to handle.
It appears that our measurement is in excellent agreement
with the latter study where the experimental conditions were
quite similar to our own conditions. This concerns especially
the pH conditions which differ significantly from those used
by Dogliotti and Hayon,18 Pikaev et al.,17 Neta and Huie,21
Ito et al.14 and Shastri and Huie.16 All these studies were per-
formed under very acidic conditions (i.e., with a pH < 0) result-
ing apparently in most cases in slower reactions rates, by ca.
50% for the studies by Neta and Huie,21 Ito et al.14 and Shastri
and Huie.16
All kinetics were studied at room temperature, i.e. 298 K.
Reagents
Solutions were prepared from the following chemicals (used
without further purification): Na2S2O8 (Prolabo Normapur,
98%), NaNO3 (Aldrich 100%), cerium nitrate ammonium
(Aldrich > 98.5%), nitric acid (Aldrich 1 mol Lꢀ1), perchloric
acid (Aldrich 70%) methanol (Prolabo, Normapur, > 99.8%),
ethanol (Prolabo, Normapur, > 99.8%), acetaldehyde (Aldrich
However, the discrepancy between our studies and those by
Dogliotti and Hayon and Pikaev et al. remains unclear. In fact,
the rate coefficient reported in these studies should have been
comparable to the first three quote at pH < 0 which are lower
than those obtained under milder conditions (as those used
here and by Herrmann and coworkers).
Table 1 Comparison of the room temperature rate coefficients for the reaction of the NO3 radical with methanol obtained in this work with
literature values
Author, year
Precursor
pH
T/K
Method
105 ꢄ k/L molꢀ1 sꢀ1
Literature
13
14
15
16
17
18
Herrmann and Exner, 1994
Ito and Akiho, 1988
Neta and Huie, 1986
Shastri and Huie, 1990
Pikaev and Sibirskaya, 1974
Dogliotti and Hayon, 1967
This work
S2O82ꢀ/K2Ce(NO3)6
K2Ce(NO3)6
HNO3
4.0
<0
<0
<0
<0
7.0
1.0
298
RT
295
295
RT
RT
298
FP
FP
PR
PR
PR
FP
FP
5.1
3.1
2.1
1.8
HNO3
HNO3
16.0
10.0
4.8 ꢁ 0.5
K2Ce(NO3)6
2ꢀ
S2O8
T h i s j o u r n a l i s Q T h e O w n e r S o c i e t i e s 2 0 0 4
3410
P h y s . C h e m . C h e m . P h y s . , 2 0 0 4 , 6 , 3 4 0 8 – 3 4 1 4

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A typical second order plot for this reaction is shown in
Fig. 3. The rate coefficient determined in this study is
(1.9 ꢁ 0.3) ꢄ 106 Mꢀ1 sꢀ1 at 298 K.
Previously reported rate coefficients are ranging from 1.1 to
3.9 ꢄ 106 Mꢀ1 sꢀ1 (see Table 2). Therefore, our determination
lies in the middle of the reported range. However again, we can
notice an excellent agreement with the data reported by
Herrmann and coworkers,20 who measured a rate constant of
(2.4 ꢁ 0.5) ꢄ 106 Mꢀ1 sꢀ1 under conditions directly comparable
to the ones used here.
Again, all lower values presented in Table 2 were obtained in
very acidic conditions (again with pH lower than 0) while the
larger ones were measured at higher pH. The difference is again
in the order of ca. 50%.
It may be speculated that both alcohols, under very acidic
conditions, might exist in their protonated form. However,
taking into consideration the pKa of both alcohols, one can
easily show that this fraction is very small. Therefore, this
speculation has to be rejected as an explanation for the differ-
ence between all experiments, which differ by the reaction
sequence used to produce the radicals. The main difference
between these two approaches (i.e., pulse radiolysis and laser
photolysis) lies in the pH range applied. In fact, H-abstraction
by NO3 produces HNO3 which is undissociated in 6 M HNO3
(conditions employed in pulse radiolysis experiments) but fully
dissociated under the conditions used in the present study. This
additional dissociation process may provide additional energy
to the system and therefore change the overall process. How-
ever, this is still speculative. Accordingly, this difference
remains unresolved.
Fig. 2 Typical NO3 temporal profile of the reaction NO3 þ MeOH in
1M HClO4 at 25 ꢆC, [NO3ꢀ] ¼ 1M, [S2O82ꢀ] ¼ 1M and
[MeOH] ¼ 1.10ꢀ3M.
We can here only underline again that the results obtained
in this work and by Herrmann et al compare quite favourably,
which validates our experimental technique.
NO3 + acetaldehyde
Less room temperature studies of the NO3 þ acetaldehyde
Fig. 3 Plots of the observed pseudo-first-order rate constants of
NO3 radicals for / methanol, L ethanol and K acetaldehyde in 1M
HClO4 at 25 ꢆC, [NO3ꢀ] ¼ 1M and [S2O82ꢀ] ¼ 1M. The ordinates
illustrate the kinetics of the reaction between the nitrate radical and
its precursors or impurities.
reaction kinetics are reported in the literature (see Table 3).
NO3 þ CH3CHðOHÞ2 ! HNO3 þ CH3CðOHÞ2 ðR6Þ
For this aldehyde, or more exactly, for its gem-diol, we mea-
sured a second-order rate coefficient of (2.0 ꢁ 0.3) ꢄ 106 Mꢀ1
s
ꢀ1, as listed in Table 3 while Herrmann et al.20 measured
However, just by considering the excellent agreement with
the study by Herrmann and coworkers, we feel confident that
the rate coefficient determined here is representative of the
reaction between the nitrate radical and methanol under
near-neutral conditions. This agreement, taken as a first step,
can be a validation of our new technique.
(1.9 ꢁ 0.6) ꢄ 106 Mꢀ1
s
ꢀ1. Again, both studies are in excellent
agreement.
On the other hand, Ito et al.22 measured a much higher rate
coefficient (up to 6.2 ꢄ 106 Mꢀ1
s
ꢀ1), which cannot be directly
reconciled with these two measurements. However, Ito et al.22
performed their study at much higher ionic strength. Such a
difference may partly explain the difference between Herrmann
et al.,20 our study and Ito et al., as discussed by Herrmann and
coworkers.
NO3 + ethanol
Many room temperature studies of the NO3 þ ethanol reaction
Finally, with the very good agreement with Herrmann and
co-workers20 (the single study directly comparable to our
experimental conditions), we can argue that the experimental
kinetics are reported (see Table 2).
NO3 þ CH3CH2OH ! HNO3 þ CH3CH2OH ðR5Þ
Table 2 Comparison of the room temperature rate coefficients for the reaction of the NO3 radical with ethanol obtained in this work with
literature values
Author, year
Precursor
pH
T/K
Method
106 ꢄ k/L molꢀ1 sꢀ1
Literature
13
14
15
16
17
18
19
Herrmann and Exner, 1994
Ito and Akiho, 1988
S2O82ꢀ/K2Ce(NO3)6
K2Ce(NO3)6
HNO3
4.0
<0
<0
<0
<0
7.0
<0
1.0
298
RT
295
295
RT
RT
296
298
FP
FP
PR
PR
PR
FP
FP
FP
2.4
1.2
Neta and Huie, 1986
Shastri and Huie, 1990
Pikaev and Sibirskaya, 1974
Dogliotti and Hayon, 1967
Alfassi and Padmaja, 1993
This work
1.4
1.1
HNO3
HNO3
2.2
3.9
K2Ce(NO3)6
K2Ce(NO3)6
1.8
1.5 ꢁ 0.3
2ꢀ
S2O8
T h i s j o u r n a l i s Q T h e O w n e r S o c i e t i e s 2 0 0 4
P h y s . C h e m . C h e m . P h y s . , 2 0 0 4 , 6 , 3 4 0 8 – 3 4 1 4
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Table 3 Comparison of the room temperature rate coefficients for the reaction of the NO3 radical with acetaldehyde obtained in this work with
literature values
Author, year
Precursor
pH
T/K
Method
106 ꢄ k/L molꢀ1 sꢀ1
Literature
2
Herrmann and Zellner, 1998
This work
S2O82ꢀ/K2Ce(NO3)66
4.0
1.0
297
298
FP
FP
1.9
2.0 ꢁ 0.3
2ꢀ
S2O8
approach based on Teflon AF has successfully been validated
and should, therefore, be applied to other studies.
To our best knowledge, there are no other kinetic data for
reactions (R7)–(R12) available in the literature with which to
compare our results.
Determination of new rate constants
Discussion
In the atmosphere, the degradation of the dimethylmalonate,
dimethylsuccinate, dimethylcarbonate, diethylcarbonate, tert-
butyl ether and propionic acid will most likely be governed by
the attack of OH and NO3 radicals [i.e., reactions (R7)–(R12)].
Table 5 compares the reactivity of the different VOCs toward
the nitrate and hydroxyl radicals. It obvious that the latter is
reacting much more efficiently with all compounds. In fact,
the ratio between both rate constants is ranging from 103 to
1.6 ꢄ 104. However, this limited set of data does not exhibit
any specific trend.
A major assumption in the reactions studied above is that
the NO3 attack on the oxygenates proceeds via an H-abstrac-
tion mechanism. If the NO3 attack proceeds via an H-abstrac-
tion, then the observed rate coefficients should be correlated
with the type of bond and bond strength of the ‘‘leaving’’
hydrogen atom. The ‘‘leaving’’ hydrogen atom is probably
the one having the weakest C–H bond strength (it was assumed
that no H-abstraction occurs at the O–H group for alcohols).
Bond dissociation energies (BDE) of the weakest C–H for all
oxygenates are listed in Table 6. Other parameters listed in
Table 6 are: (i) the number of equivalent hydrogen atom nH
(i.e., those having the weakest bond strength); and (ii) the loga-
rithm of the rate constant per H atom kH , which is derived by:
NO3 þ CH3OCðOÞCH2CðOÞOCH3 ! Hþ þ NO3
ꢀ
þ CH3OCðOÞCH2CðOÞOCH2 ðR7Þ
NO3 þ CH3OCðOÞCH2CH2CðOÞOCH3 ! Hþ þ NO3
ꢀ
þ CH3OCðOÞCH2CHCðOÞOCH3 ðR8Þ
NO3 þ CH3OCðOÞOCH3 ! Hþ þ NO3
ꢀ
þ CH3OCðOÞOCH2 ðR9Þ
NO3 þ CH3CH2OCðOÞOCH2CH3 ! Hþ þ NO3
ꢀ
þ CH3CH2OCðOÞOCHCH3 ðR10Þ
NO3 þ ðCH3Þ3COCH3 ! Hþ þ NO3ꢀ þ ðCH3Þ3COCH2
ðR11Þ
NO3 þ CH3CH2COOH ! Hþ þ NO3ꢀ þ CH3CHCOOH
ðR12Þ
kH ¼ kobs=nH
ð1Þ
The experimental approach described above was used to
measure the bimolecular rate coefficients for the reactions of the
NO3 radical with dimethylmalonate (R7), dimethylsuccinate
(R8), dimethylcarbonate (R9), diethylcarbonate (R10), tert-
butyl methyl ether (R11) and propionic acid (R12). The pKa
of the latter acid is 4.87, therefore under our experimental con-
ditions it can safely be assumed that this acid is undissociated.
Typical second order plots for these reactions are shown in
Fig. 4 and all rate coefficients are summarised in Table 4.
Many of the bond dissociation energies have not been
experimentally measured and were, therefore, estimated using
Table 4 Summary of the rate coefficients for the reactions of the NO3
radical with six oxygenates obtained in this work
Compounds
k/L molꢀ1 sꢀ1
NO3 precursor
Diethylcarbonate
Dimethylcarbonate
Dimethylmalonate
Propionic acid
(8.40 ꢁ 1.25) ꢄ 104
(1.5 ꢁ 0.4) ꢄ 104
(2.6 ꢁ 0.7) ꢄ 104
(7.7 ꢁ 0.2) ꢄ 104
(3.9 ꢁ 1.3) ꢄ 105
(3.4 ꢁ 0.2) ꢄ 104
S2O28
S2O28
S2O28
S2O28
tert-Butyl-methyl ether
Dimethylsuccinate
K2Ce(NO3)6
K2Ce(NO3)6
Table 5 Comparison of the reactivity of the oxygenated compounds
toward the nitrate and hydroxyl radicals. The data are from this work
or ref. 5 and references cited therein
kOH
/
kNO3
/
kOH/
kNO3
Reactant
L molꢀ1 sꢀ1
L molꢀ1 sꢀ1
Methanol
Ethanol
(1.3 ꢁ 0.4) ꢄ 109 (4.8 ꢁ 0.5) ꢄ 105
(1.9 ꢁ 0.5) ꢄ 109 (1.9 ꢁ 0.3) ꢄ 106
(2.7 ꢁ 0.9) ꢄ 108 (2.6 ꢁ 0.7) ꢄ 104
(5.3 ꢁ 2.9) ꢄ 108 (3.4 ꢁ 0.2) ꢄ 104
2.7 ꢄ 103
3
1.0 ꢄ
Dimethylmalonate
Dimethylsuccinate
1.0 ꢄ 104
1.6 ꢄ 104
3.4 ꢄ 103
Dimethylcarbonate (5.1 ꢁ 2.7) ꢄ 108 (1.5 ꢁ 0.4) ꢄ 104
Fig. 4 Plots of the observed pseudo-first-order decay rates of NO3
radicals for: / propionic acid in 1M HClO4 at 25 ꢆC, [NO3ꢀ] ¼ 1M
Diethylcarbonate
Acetaldehyde
tert-Butyl-
(7.9 ꢁ 3.2) ꢄ 108 (8.4 ꢁ 1.25) ꢄ 104 9.4 ꢄ 103
3.6 ꢄ 1091
(2.0 ꢁ 0.3) ꢄ 106
1.8 ꢄ 103
and [S2O82ꢀ] ¼ 1M.; L dimethylsuccinate in 1M HNO3 at 25 ꢆC,
[[Ce (NO3)6]2ꢀ];
K
dimethylmalonate in 1M HclO4 at 25 ꢆC,
1.6 ꢄ 1092
(3.9 ꢁ 1.3) ꢄ 105
4.3 ꢄ 103
IV
[NO3ꢀ] ¼ 1M and [S2O82ꢀ] ¼ 1M. The ordinates illustrate the kinetics
of the reaction between the nitrate radical and its precursors or
impurities.
methyl ether
Propionic acid
3.8 ꢄ 1083
(7.7 ꢁ 0.2) ꢄ 104
4.9 ꢄ 103
T h i s j o u r n a l i s Q T h e O w n e r S o c i e t i e s 2 0 0 4
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Table 6 Number of equivalent weak C–H bond (nH) with the bond
dissociation energy (BDE) and the rate constants k
Equival
H-atoms
BDE/
kJ molꢀ1
Compounds
k/L molꢀ1 sꢀ1
Methanol
Ethanol
Acetaldehyde
3
2
3
2
3
2
2
3
4
393
385
396
424
432
444
415
397
424
4.8 ꢁ 0.5 ꢄ 105
1.5 ꢁ 0.3 ꢄ 106
2.0 ꢁ 0.3 ꢄ 106
8.40 ꢁ 1.25 ꢄ 104
1.5 ꢁ 0.4 ꢄ 104
2.6 ꢁ 0.7 ꢄ 104
7.7 ꢁ 0.2 ꢄ 104
3.9 ꢁ 1.3 ꢄ 105
3.4 ꢁ 0.2 ꢄ 104
Diethylcarbonate
Dimethylcarbonate
Dimethylmalonate
Propionic acid
tert-Butyl-methyl ether
Dimethylsuccinate
Fig. 6 Presentation of the logarithm of the rate constants per H
atom abstractable vs. bond dissociation energies for NO3 abstraction
reactions in aqueous solution at room temperature.
semi-empirical group contribution methods.23,24 This method
is independent of the one derived by S. Benson, as the latter
does not allow the estimation of bond dissociation energies
for complex oxygenated molecules. It should be underlined
that these bond energies are, without doubt, carrying some
uncertainties. However in the absence of experimental data,
this estimation procedure is a straightforward solution that
should still illustrate some trends in the rate constants.
It may be assumed that the observed differences of the reac-
tivity of the NO3 radical toward the organic species considered
are due to changes in the activation energies which are
expected to be linked to the bond dissociation energy of the
bond being ruptured.25 If this is the case, there should be a lin-
ear dependence of the logarithm of kH on the bond strength.
A plot of log(kH) versus BDE is shown in Fig. 6. There is indeed
a (weak) correlation between the logarithm of kH and the bond
strength, i.e. a decrease of the rate coefficient is observed
with increasing BDE. As a result, the reaction of the NO3 radi-
cal with oxygenates proceeds via an H-atom abstraction
mechanism.
In summary, we reported in this study on the reactivity of
the nitrate radical NO3 toward several oxygenated com-
pounds. For some of them, we measured the rate coefficient
for the first time. Apparently, the reaction proceeds via an
H-abstraction mechanism as already stated in previous studies.
The existing data do not support any reliable estimation
method, which warrants future work in order to improve
knowledge that is required building reliable structure–activity
relationships.
Conclusions
This work addressed the oxidation of several oxygenated com-
pounds by the nitrate radical using a newly developed experi-
mental technique. The experimental approach described here
takes advantage of the novel Teflon AF 2400 material. The
physical nature of such tubing allows for the construction of
photolysis reaction cells with very low volumes but potentially
very long optical path lengths.
We have successfully applied such photolysis cell to the mea-
surement of the rate coefficients for the liquid phase reaction
of the nitrate radical with methanol, ethanol, acetaldehyde,
tert-butyl methyl ether, propionic acid, dimethylmalonate,
dimethylsuccinate, dimethyl carbonate and diethylcarbonate.
The reported rate coefficients for the reactions of NO3 with
ethanol, methanol and acetaldehyde agree very well with the
literature values. To date, there is no kinetic data reported
for the OH radical reaction with dimethylmalonate, dimethyl-
succinate, dimethylcarbonate, diethylcarbonate, tert-butyl
methyl ether and propionic acid with which to compare our
results.
Acknowledgements
Support of this work by the Programme National de Chimie
´
Atmospherique (PNCA) from the CNRS is gratefully
acknowledged as well as the project PRIMEQUAL by the
French Minister for Environment.
References
1
B. J. Finlayson-Pitts and J. N. Pitts, Chemistry of the Upper and
Lower Atmosphere: Theory, Experiments, and Applications,
Academic Press, San Diego, 2000.
2
3
H. Herrmann and R. Zellner, N-Centered Radicals, 1998,
291–343.
R. E. Huie, in Laboratory Studies of Atmospheric Heterogeneous
Chemistry; Current Problems in Atmospheric Chemistry, Advances
in Physical Chemistry Series, ed. J. R. Barker, World Scientific,
Singapore, 1994, vol. 3, pp. 374–419.
4
5
6
7
T. Imamura, Y. Rudich, R. K. Talukdar, R. W. Fox and A. R.
Ravishankara, J. Phys. Chem. A, 1997, 101, 2316–2322.
C. George, D. Rousse, E. Perraudin and R. Strekowski, Phys.
Chem. Chem. Phys., 2003, 5, 1562–1569.
M. Uusi-Penttila, R. J. Richards, P. Blowers, B. A. Torgerson and
K. A. Berglund, J. Chem. Eng. Data, 1996, 41, 235–238.
F. Cavalli, I. Barnes and K. H. Becker, Int. J. Chem. Kinet., 2001,
33, 431–439.
Fig. 5 Plots of the observed pseudo-first-order decay rates of NO3
radicals for: / diethylcarbonate; L dimethylcarbonate in 1M HclO4
at 25 ꢆC, [NO3ꢀ] ¼ 1M and [S2O82ꢀ] ¼ 1M; K tert-butyl-methyl ether.
in 1M HNO3 at 25 C, [[Ce (NO3)6]2ꢀ] ¼ 5.10ꢀ4 M. The ordinates
ꢆ
IV
illustrate the kinetics of the reaction between the nitrate radical and
its precursors or impurities.
T h i s j o u r n a l i s Q T h e O w n e r S o c i e t i e s 2 0 0 4
P h y s . C h e m . C h e m . P h y s . , 2 0 0 4 , 6 , 3 4 0 8 – 3 4 1 4
3413

View Article Online
8
9
J. Wenger, E. Porter, E. Collins, J. Treacy and H. Sidebottom,
Chemosphere, 1999, 38, 1197–1204.
M. Bilde, T. E. Mo¨gelberg, J. Sehested, O. J. Nielsen, T. J.
Wallington, M. D. Hurley, S. M. Japar, M. Dill, V. L. Orkin,
T. J. Buckley, R. E. Huie and M. J. Kurylo, J. Phys. Chem. A,
1997, 101, 3514–3525.
17 A. K. Pikaev, G. K. Sibirskaya, E. M. Shirshov, P. Y. Glazunov
and V. I. Spitsyn, Dokl. Akad. Nauk SSSR, 1974, 215, 645–648.
18 L. Dogliotti and E. Hayon, J. Phys. Chem., 1967, 71, 3802–3808.
19 Z. B. Alfassi, S. Padmaja, P. Neta and R. E. Huie, J. Phys. Chem.,
1993, 97, 3780–3782.
20 H. Herrmann, Photochemische Bildung, Spektroskopie und Kinetik
freier Radikale in wa¨sseriges Lo¨sung. Habilitationsschrift, Univer-
sita¨t GH Essen, 1997.
21 P. Neta and R. E. Huie, J. Phys. Chem., 1986, 90, 4644–4648.
22 O. Ito, S. Akiho and M. Iino, J. Phys. Chem., 1989, 93, 4079–4083.
23 J. Gasteiger, in Physical Property Prediction in Organic Chemistry,
ed. C. Jochum, M. G. Hicks and J. Sunkel, Springer Verlag,
Heidelberg, 1988, pp 119–138.
10 M. A. Pacheco and C. L. Marshall, Energy Fuels, 1997, 11, 2–29.
11 J. R. Odum, T. P. W. Jungkamp, R. J. Griffin, R. C. Flagan and
J. H. Seinfeld, Science (Washington, D. C.), 1997, 276, 96–99.
12 J. H. Seinfeld, S. N. Pandis, Atmospheric Chemistry and Physics:
From Air Pollution to Climate Change, Wiley, New York, 1998.
13 H. Herrmann, M. Exner and R. Zellner, Geochim. Cosmochim.
Acta, 1994, 58, 3239–3244.
14 O. Ito, S. Akiho and M. Iino, Bull. Chem. Soc. Jpn., 1989, 62,
1606–1611.
15 P. Neta and R. E. Huie, J. Phys. Chem., 1986, 90, 4644–4648.
16 L. V. Shastri and R. E. Huie, Int. J. Chem. Kinet., 1990, 22,
505–512.
24 J. Gasteiger, 2000, PETRA: Parameter Estimation for the
Treatment of Reactivity Applications, http://www2.chemie.
uni-erlangen.de/software/petra/.
25 C. L. Clifton and R. E. Huie, Int. J. Chem. Kinet., 1989, 21,
677–687.
T h i s j o u r n a l i s Q T h e O w n e r S o c i e t i e s 2 0 0 4
3414
P h y s . C h e m . C h e m . P h y s . , 2 0 0 4 , 6 , 3 4 0 8 – 3 4 1 4