Laboratory and theoretical study of the oxy radicals in the OH- and Cl-initiated oxidation of ethene

Article abstract of DOI:10.1021/jp981937d

The products of the OH-initiated oxidation mechanism of ethene have been studied as a function of temperature (between 250 and 325 K) in an environmental chamber, using Fourier transform infrared spectroscopy for end product analysis. The oxidation proceeds via formation of a peroxy radical, HOCH₂CH₂O₂. Reaction of this peroxy radical with NO is exothermic and produces chemically activated HOCH₂CH₂O radicals, of which about 25% decompose to CH₂OH and CH₂O on a time scale that is rapid compared to collisions, independent of temperature. The remainder of the HOCH₂CH₂O radicals are thermalized and undergo competition between decomposition, HOCH₂CH₂O → CH₂OH + CH₂O (6), and reaction with O₂, HOCH₂CH₂O + O₂ → HOCH₂-CHO + HO₂ (7). The rate constant ratio, k₆/k₇, for the thermalized radicals was found to be (2.0 ± 0.2) × 10²⁵ exp[-(4200 ± 600)/T] molecule cm^-3 over the temperature range 250-325 K. With the assumption of an activation energy of 1-2 kcal mol^-1 for reaction 7, the barrier to decomposition of the HOCH₂CH₂O radical is found to be 10-11 kcal mol^-1. A study of the Cl-atom-initiated oxidation of ethene was also carried out; the main product observed under conditions relevant to the atmosphere was chloroacetaldehyde, ClCH₂CHO Theoretical studies of the thermal and "prompt" decomposition of the oxy radicals were based on a recent ab initio characterization that highlighted the role of intramolecular H bonding in HOCH₂CH₂O. Thermal decomposition is described by transition state and the Troe theories. To quantify the prompt decomposition of chemically activated nascent oxy radicals, the energy partitioning in the initially formed radicals was described by separate statistical ensemble theory, and the fraction of activated radicals dissociating before collisional stabilization was obtained by master equation analysis using RRKM theory. The barrier to HOCH₂CH₂O decomposition is inferred independently as being 10-11 kcal mol^-1, by matching both of the theoretical HOCH₂CH₂O decomposition rates at 298 K with the experimental results. The data are discussed in terms of the atmospheric fate of ethene.

Full text of DOI:10.1021/jp981937d

8116
J. Phys. Chem. A 1998, 102, 8116-8123
Laboratory and Theoretical Study of the Oxy Radicals in the OH- and Cl-Initiated
Oxidation of Ethene
John J. Orlando* and Geoffrey S. Tyndall*
Atmospheric Chemistry DiVision, National Center for Atmospheric Research, Boulder, Colorado 80303
Merete Bilde
Plant Biology and Biogeochemistry Department, Risø National Laboratory, DK-4000 Roskilde, Denmark
Corinne Ferronato
Groupe de Recherche sur l’EnVironnement et la Chimie Applique´e (GRECA), UniVersite´ Joseph Fourier,
F-38000 Grenoble, France
Timothy J. Wallington
Ford Research Laboratory, SRL-3083, Ford Motor Company, Dearborn, Michigan 48121-2053
Luc Vereecken and Jozef Peeters*
Department of Chemistry, UniVersity of LeuVen, B-3001 LeuVen, Belgium
ReceiVed: April 21, 1998; In Final Form: August 12, 1998
The products of the OH-initiated oxidation mechanism of ethene have been studied as a function of temperature
(between 250 and 325 K) in an environmental chamber, using Fourier transform infrared spectroscopy for
end product analysis. The oxidation proceeds via formation of a peroxy radical, HOCH2CH2O2. Reaction of
this peroxy radical with NO is exothermic and produces chemically activated HOCH2CH2O radicals, of which
about 25% decompose to CH2OH and CH2O on a time scale that is rapid compared to collisions, independent
of temperature. The remainder of the HOCH2CH2O radicals are thermalized and undergo competition between
decomposition, HOCH2CH2O f CH2OH + CH2O (6), and reaction with O2, HOCH2CH2O + O2 f HOCH2-
CHO + HO2 (7). The rate constant ratio, k6/k7, for the thermalized radicals was found to be (2.0 ( 0.2) ×
1025 exp[-(4200 ( 600)/T] molecule cm-3 over the temperature range 250-325 K. With the assumption of
an activation energy of 1-2 kcal mol-1 for reaction 7, the barrier to decomposition of the HOCH2CH2O
radical is found to be 10-11 kcal mol-1. A study of the Cl-atom-initiated oxidation of ethene was also
carried out; the main product observed under conditions relevant to the atmosphere was chloroacetaldehyde,
ClCH2CHO. Theoretical studies of the thermal and “prompt” decomposition of the oxy radicals were based
on a recent ab initio characterization that highlighted the role of intramolecular H bonding in HOCH2CH2O.
Thermal decomposition is described by transition state and the Troe theories. To quantify the prompt
decomposition of chemically activated nascent oxy radicals, the energy partitioning in the initially formed
radicals was described by separate statistical ensemble theory, and the fraction of activated radicals dissociating
before collisional stabilization was obtained by master equation analysis using RRKM theory. The barrier to
HOCH2CH2O decomposition is inferred independently as being 10-11 kcal mol-1, by matching both of the
theoretical HOCH2CH2O decomposition rates at 298 K with the experimental results. The data are discussed
in terms of the atmospheric fate of ethene.
Introduction
atmospheric lifetime (a day or two; see below), leads to a wide
range of observed ethene mixing ratios (tens of parts per
trillion over the oceans to tens of parts per billion in urban
regions).2-4 Despite its short lifetime, high concentrations of
ethene are occasionally observed in the upper troposphere as
the result of convection4 or the lofting of biomass burning
plumes.5,6
The oxidation of ethene in the atmosphere, as is the case for
most hydrocarbons, is initiated predominantly by reaction with
the OH radical.7
The oxidation of non-methane hydrocarbons (NMHCs) in the
troposphere leads to the generation of ozone and other secondary
pollutants such as peroxyacyl nitrates (PANs), other organic
nitrates, and carbonyl species. Ethene is one of the more
abundant NMHCs with a total source strength of about 18-45
Tg/yr.1,2 Various natural sources of ethene to the atmosphere
have been identified, including emissions from vegetation, soils,
and the oceans. Ethene is also an important component of
biomass burning and is produced in the consumption of fossil
fuels.1 The disparate nature of its sources, coupled with its short
OH + C2H4 + M f HOCH2CH2 + M
(1)
S1089-5639(98)01937-9 CCC: $15.00 © 1998 American Chemical Society
Published on Web 09/26/1998

OH- and Cl-Initiated Oxidation of Ethene
J. Phys. Chem. A, Vol. 102, No. 42, 1998 8117
The recommended value for the rate coefficient for reaction 1,
k1 ) 8.2 × 10-12 cm3 molecule-1 s-1, at 1 atm total pressure
and 298 K,8 leads to an atmospheric lifetime for ethene of about
1-2 days. The hydroxyethyl radical rapidly adds O2 to form a
â-hydroxyalkylperoxy radical,9,10 the subsequent reactions of
which lead in large part to the generation of a â-hydroxyalkoxy
radical, HOCH2CH2O.10-14
decomposition rates of the oxy radicals from the separate
statistical ensemble (SSE) theory. The relevance of the data to
the atmospheric fate of ethene is discussed.
Experimental Section
The work described here was carried out in an environmental
chamber, which has been described previously.20,21 The ap-
paratus consists of a 47 L stainless steel cell, 2 m in length,
which is interfaced to a Fourier transform spectrometer (Bomem
DA3.01) operating in the infrared spectral region. Multipass
optics housed within the chamber provided a total path length
of 32.6 m. Absorption spectra were obtained from the co-
addition of 50-200 interferograms recorded at a resolution of
1 cm-1. The temperature of the cell was controlled by flowing
ethanol from a circulating bath (Neslab ULT-80DD) through a
jacket surrounding the cell. The temperature in the chamber,
as monitored using eight thermocouple gauges along its length,
was found to be uniform along the length of the cell to within
(1 K. Photolysis was carried out using a Xe-arc lamp, the
output of which was filtered with a Corning 7-54 filter to give
radiation between 240 and 395 nm.
The OH-initiated oxidation of ethene was studied using the
photolysis of mixtures of CH3ONO (10-58 mTorr), C2H4
(125-360 mTorr), NO (9-90 mTorr), O2 (70-700 Torr), and
N2 (0-630 Torr), at a total pressure of 700 Torr. For studies
of the Cl-initiated oxidation, CH3ONO was replaced in the
mixture by Cl2 (40-145 mTorr) and the O2 pressure was varied
between 2 and 700 Torr. The minor components of the gas
mixtures were swept into the chamber from calibrated volumes
using O2 or N2 carrier gas. Quantification of reactants and
products in these experiments was achieved by comparison to
standard spectra obtained in the environmental chambers at
NCAR or at Ford Motor Company. The absorption feature used
to quantify HOCH2CHO at 1107 cm-1 is strongly overlapped
by that of HCOOH at 1105 cm-1. Therefore, care was taken
to ensure that the NO was not depleted, to avoid buildup of
HCOOH from the reaction of HO2 with CH2O.
HOCH2CH2 + O2 + M f HOCH2CH2O2 + M (2)
HOCH2CH2O2 + NO f HOCH2CH2O + NO2 (3)
HOCH2CH2O2 + RO2 f products
HOCH2CH2O2 + HO2 f products
(4)
(5)
Oxy radicals such as HOCH2CH2O can undergo two basic
types of reaction, reaction with O2 and decomposition via C-C
bond rupture,15,16 and it is this competition that determines which
stable products are obtained in the oxidation process. Envi-
ronmental chamber studies by Niki et al.12,13 using FT-IR
detection of end products revealed the formation of CH2O and
glycolaldehyde (hydroxyacetaldehyde, HOCH2CHO) in the OH-
initiated oxidation of ethene. These data were interpreted in
terms of a competition between unimolecular decomposition
of the oxy radical,
HOCH2CH2O + M f CH2OH + CH2O + M
and its reaction with O2,
HOCH2CH2O + O2 f HOCH2CHO + HO2
(6)
(7)
A rate coefficient ratio k6/k7 ) 1.8 × 1019 molecule cm-3 was
reported for 1 atm total pressure at 298 K. The ratio is expected
to be strongly temperature dependent, owing to the existence
of an energy barrier to the occurrence of reaction 6, but no
studies of the oxidation of ethene at T < 298 K have been
reported to date.
Hydroxyl radicals were generated in the cell from the CW
photolysis of methyl nitrite (CH3ONO) in the presence of added
NO.
The reactions of peroxy radicals with NO are in general
exothermic by about 10-15 kcal mol-1. Recent studies on
halocarbon oxidation conducted in our laboratories17-19 have
shown that a significant fraction of this available energy is
deposited into the internal modes of the oxy radical product,
leading to an enhanced tendency for this radical to undergo
unimolecular decomposition. This chemical activation is a result
of the relatively long lifetime (>1 ps) of the peroxynitrite
intermediate formed in the reaction of the peroxy radical with
NO. Since the barrier to decomposition of the HOCH2CH2O
CH3ONO + hν f CH3O + NO
CH3O + O2 f CH2O + HO2
HO2 + NO f OH + NO2
(8)
(9)
(10)
radical is expected to be only about 12-14 kcal mol-1 16
,
such
Methyl nitrite was prepared from the dropwise addition of
sulfuric acid to a saturated solution of sodium nitrite in
methanol22 and degassed by multiple freeze-pump-thaw cycles
prior to use. For most experiments at 298 K and all experiments
at lower temperatures, 13CH3ONO was used so that formalde-
hyde formed from its photolysis could be differentiated spec-
troscopically from that formed in the ethene oxidation. Other
gases were used as purchased, N2 from the boil-off of a liquid
N2 Dewar, O2 (UHP, U.S. Welding), Cl2 (HP, Matheson), C2H4
(CP grade, Matheson), NO (Linde).
Experiments consisted of between two and five irradiations
of a gas mixture, with an FTIR spectrum recorded after each
irradiation. The overall conversion of ethene was always kept
to less than 5% to avoid significant destruction of the products
(formaldehyde, glycolaldehyde, and chloroacetaldehyde; see
below), which have similar reactivities toward Cl and OH as
ethene.8,23-25
an effect may also be at work in the ethene system and may
influence the observed products. To model the extent of prompt
dissociation, one needs to know both the thermochemistry of
the alkoxy radical decomposition reaction and the energy
partitioning in the nascent products of the RO2 + NO reaction.
To aid in the interpretation of the results, detailed calculations
have been carried out to model the energy distribution in the
products of reaction 3, the fraction of the activated oxy radicals
that decomposes before undergoing collisions, and the rate of
decomposition of thermalized oxy radicals.
The goals of this work were as follows: (1) to determine the
products of the OH- and Cl-initiated oxidation of ethene over
the temperature range encountered in the troposphere; (2) to
investigate whether the extent to which chemical activation of
the oxy radical occurs differs in the two reaction systems; (3)
to interpret the results quantitatively using calculations of the

8118 J. Phys. Chem. A, Vol. 102, No. 42, 1998
Orlando et al.
Results and Discussion
OH-Initiated Oxidation of Ethene: Chemistry of the
HOCH2CH2O Radical. The peroxy radical HOCH2CH2O2 was
generated from the addition of OH to ethene via reactions 1
and 2. On the basis of our previous studies of HFC-134a and
HFC-236cb oxidation,17,18 the following mechanism can be used
to describe the subsequent chemistry involving the formation
of an excited oxy radical.
HOCH2CH2O2 + NO f HOCH2CH2O* + NO2 (3)
HOCH2CH2O* f CH2OH + CH2O
HOCH2CH2O* + M f HOCH2CH2O + M
HOCH2CH2O + M f CH2OH + CH2O + M
HOCH2CH2O + O2 f HOCH2CHO + HO2
CH2OH + O2 f CH2O + HO2
(11)
(12)
(6)
Figure 1. Plot of the inverse yield of glycolaldehyde versus the inverse
oxygen pressure, from the OH-initiated oxidation of ethene at 700 Torr
total pressure. Data shown are for 298 K: filled circles, this work;
open circles, Niki et al.13
(7)
(13)
In this analysis, any oxy radicals that are produced above the
barrier to decomposition are assumed to decompose rapidly on
a time scale that is short relative to collisions, as represented
by reaction 11. Radicals that are produced below the decom-
position barrier suffer numerous collisions and are thus ther-
malized (reaction 12) before undergoing decomposition (reaction
6) or reaction with O2 (reaction 7). If we use the quantity R to
represent the yield of thermalized radicals, then the yield of
glycolaldehyde (YGA) is given as follows:
YGA ) R{k7[O2]/(k7[O2] + k6)}
(A)
and
[YGA]
-1 ) (1/R) [k6/k7[O2] + 1]
(B)
Figure 2. Plot of the inverse yield of glycolaldehyde versus the inverse
oxygen pressure, from the OH-initiated oxidation of ethene at 250 K
and 700 Torr total pressure. The data have been corrected for loss of
HOCH2CH2O radicals by reaction with NO or NO2 as described in eq
E.
Because small conversions of ethene were employed in these
experiments, yields of glycolaldehyde were not obtained relative
to the amount of ethene consumed but instead were obtained
from a consideration of the relative yields of glycolaldehyde
and formaldehyde. In this case, since two molecules of
formaldehyde are formed,13
a slope of 540 ( 80 Torr, from which values of R ) 0.74 (
0.10 and k6/k7 ) (1.32 ( 0.20) × 1019 molecule cm-3 can be
obtained (all uncertainties are 2σ). Also shown in Figure 1 (but
not included in the fit) are data reported by Niki et al.13 under
the same conditions of pressure and temperature. As will be
further discussed below, these data are clearly consistent with
our data and are also indicative of a nonunity intercept in the
1/YGA vs 1/[O2] plot.
The nonunity intercept obtained in Figure 1 is a strong
indication of the occurrence of prompt decomposition (reaction
11) of a significant fraction of the oxy radicals formed in
reaction 3. However, the value of R, and hence the magnitude
of the hot alkoxy radical effect, obtained from Figure 1 is quite
uncertain, owing to a fairly substantial extrapolation. The best
demonstration of the occurrence of prompt decomposition of
the oxy radicals produced in reaction 3 is obtained at low
temperature and high O2 pressures. Under these conditions,
most thermalized radicals should react with O2 (owing to the
strong temperature dependence anticipated for reaction 6) and
significant yields of formaldehyde must originate from prompt
decomposition of chemically activated oxy radicals. Data for
the lowest temperature studied, 250 K, are shown in Figure 2.
Clearly, even at 250 K and 700 Torr of O2, there is a significant
YGA ) [HOCH2CHO]/{0.5[CH2O] + [HOCH2CHO]} (C)
From (B) above, it is clear that values of R and k6/k7 can be
obtained from a plot of the inverse of the glycolaldehyde yield
versus the inverse O2 pressure, provided no other losses of the
thermalized oxy radicals occur. A plot of this type (for 700
Torr total pressure and 298 K) is shown in Figure 1. The data
shown include those obtained using 13CH3ONO as the OH
radical source as well as some initial data obtained using
12CH3ONO. Product yield data obtained using 12CH3ONO had
to be corrected for production of 12CH2O from the decomposi-
tion of 12CH3ONO, reactions 8 and 9:
[12CH2O]t,corr ) [12CH2O]t,obs
-
{[12CH3ONO]o - [12CH3ONO]t}
where [12CH2O]t,corr and [12CH2O]t,obs are the corrected and
observed formaldehyde concentrations at time t and [12CH3-
ONO]o and [12CH3ONO]t are the initial concentration of methyl
nitrite and its concentration at time t. A least-squares fit of the
data shown in Figure 1 yields an intercept of 1.35 ( 0.15 and

OH- and Cl-Initiated Oxidation of Ethene
J. Phys. Chem. A, Vol. 102, No. 42, 1998 8119
yield of formaldehyde and only about a 70% yield of glycol-
aldehyde. As the amount of O2 is decreased, the glycolaldehyde
yield does decrease, but only slightly, owing to a small amount
of decomposition of the thermalized oxy radicals (reaction 6).
At cold temperatures, the possibility exists that the thermalized
oxy radicals react with NO or NO2 in addition to reactions 6
and 7.
HOCH2CH2O + NO (NO2) f
HOCH2CH2-ONO (-ONO2) (14)
In this case, eq A can be modified to take account of the loss
of radicals, assuming similar rate coefficients for the reactions
of NO and NO2 with HOCH2CH2O.
YGA ) R{k7[O2]/(k7[O2] + k6 + k14[NOx])}
(D)
The apparent glycolaldehyde yield, relative to the sum {0.5[CH2O]
+ [HOCH2CHO]}, can then easily be shown to be given by eq
E.
YAPP ) R{k7[O2]/(k7[O2] + k6 + (1 - R)k14[NOx])} (E)
Some experiments were performed using elevated levels of NO
at 250 and 255 K to look for a reduction in the yield of
glycolaldehyde associated with the formation of hydroxynitrites
or nitrates. A small change in the glycolaldehyde yield was
observed (about 10%) when the ratio [NO]/[O2] was increased
to 3000. Analytical fits to the observed data using eq E showed
that the ratio k14/k7 is probably in the range 3000-4000. The
data were thus corrected for loss of the thermalized oxy radicals
using k14/k7 ) 3000 with values for R and k6/k7 estimated from
the experiments at higher temperatures. A plot of 1/YGA versus
1/[O2] for the corrected data is shown in Figure 2 and yields a
ratio k6/k7 ) (1.7 ( 0.5) × 1018 molecule cm-3 and a yield of
stabilized radicals, R, of 0.71 ( 0.10. The reactions of alkoxy
radicals with NO and NO2 are normally at the high-pressure
limit at 700 Torr, with rate coefficients of the order (2-4) ×
Figure 3. Plot of the inverse yield of glycolaldehyde versus the inverse
oxygen pressure, from the OH-initiated oxidation of ethene at 700 Torr
total pressure. Data shown are for 298 K (b), 278 K (3), 273 K (O),
and 265 K (1).
TABLE 1: Summary of Dissociation Rate Coefficients at
700 Torr and Fraction of Stabilized Oxy Radicals Measured
in This Studya
k6/k7
temp K
1018 molecule cm-3
k6 s-1
R
325
315
298
278
273
265
256
250
64 ( 21
39 ( 13
13 ( 2
6.4 ( 1.6
3.7 ( 0.8
2.5 ( 0.5
1.5 ( 0.5
1.7 ( 0.5
8.0 × 105
4.5 × 105
1.3 × 105
5.3 × 104
2.9 × 104
1.9 × 104
1.0 × 104
1.1 × 104
0.74 ( 0.10
0.83 ( 0.15
0.71 ( 0.09
0.79 ( 0.08
0.82 ( 0.10
0.71 ( 0.10
10-11 cm3 molecule-1 s-1 16,26
Thus, the reaction of HOCH2-
.
CH2O with O2 is of the order (6-8) × 10-15 cm3 molecule-1
s-1 at 250 K, similar to the reaction of ethoxy radicals with
O2.16,26
Data were obtained at intermediate temperatures as well
(Figure 3). In all cases, a nonunity intercept in the 1/YGA vs
1/[O2] plot was obtained, with an average value at all temper-
atures studied of 1.34 ( 0.10. Since there was no significant
variation of the intercept with temperature, we report a tem-
perature-independent yield of thermalized oxy radicals of 0.75
( 0.06. Values of the rate coefficient ratio k6/k7 were also
obtained at each temperature studied. The slopes and intercepts
at each temperature are summarized in Table 1. The slopes of
the plots typically had uncertainties of 20%. The data at 256
K were corrected for the loss of RO radicals using eq E. Above
256 K, decomposition of the oxy radical is sufficiently rapid
that the NOx reactions do not have a major effect on the yields,
and no correction was made. The data are plotted in Arrhenius
form in Figure 4.
Product yield experiments were also carried out at 315 and
325 K. Because of the small yields of glycolaldehyde at these
temperatures (due to rapid decomposition of HOCH2CH2O),
experiments were only conducted at 700 Torr of O2. Measured
yields of glycolaldehyde under these conditions were 27 ( 4%
(315 K) and 19 ( 3% (325 K). While accurate values of k6/k7
cannot be obtained in the absence of a full O2 dependence of
the product yields, a rough estimate can be obtained by
a k6 was calculated using k7 ) 6 × 10-14 exp(-550/T) cm3
molecule-1 s-1, the same as for C2H5O + O2.8,26
constructing a plot of 1/YGA versus 1/[O2] using an assumed
intercept of 1.34 (the average value determined above) and the
single data point at 700 Torr. Values of k6/k7 obtained in this
fashion (also shown in Figure 4) have been assigned an
uncertainty of (30%.
A linear least-squares fit to all the data shown in Figure 4
yields k6/k7 ) (2.0 ( 0.2) × 1025 exp[-(4200 ( 600)/T]
molecule cm-3, where the error in the pre-exponential factor
reflects the uncertainty in the room-temperature value. Because
of the large uncertainty in the values obtained at 250 and 256
K (due to a very small variation in the glycolaldehyde with O2
pressure), these points were given a weight of 0.5 in the fit, as
were the single-point determinations at 315 and 325 K. If an
activation energy of 1-2 kcal mol-1 is assumed for the reaction
of HOCH2CH2O with O2, as is the case for the analogous
reactions of CH3O and CH3CH2O,8 then an Arrhenius activation
energy of 9.5-10.5 kcal mol-1 is obtained for the decomposition
of HOCH2CH2O, reaction 6, at 700 Torr. As discussed below,
the reaction is expected to be in its falloff regime such that the
corresponding barrier height Eb can be estimated at 10-11 kcal

8120 J. Phys. Chem. A, Vol. 102, No. 42, 1998
Orlando et al.
mol-1. The values for k6 given in Table 1 have been calculated
using a rate coefficient of 6 × 10-14 exp(-550/T) cm3 molecule-1
s-1 for the reaction with O2, which is the same as that for C2H5O
+ O2.8,26 While considerable uncertainty exists in any individual
determination of R and k6/k7, the theoretical treatment presented
later shows that chemical activation is expected for the barrier
to decomposition measured here; thus, we are justified in
analyzing the data in this way.
Our data can be compared with previous room-temperature
studies of the OH-initiated oxidation of ethene. The product
yield data of Niki et al.13 at room temperature are in very good
agreement with our work (see Figure 1). However, the earlier
study did not account for the occurrence of prompt decomposi-
tion of the HOCH2CH2O* radical. Therefore, despite the
observation of indistinguishable product yields, Niki’s value of
k6/k7 ) 1.8 × 1019 molecule cm-3 at 298 K and 1 atm total
pressure differs from ours.
Barnes et al.14 also conducted chamber studies of the OH-
initiated oxidation of ethene. Their data show an increased yield
of formaldehyde versus glycolaldehyde when NO was added
to their system, consistent with the occurrence of prompt
decomposition of the HOCH2CH2O radical in the presence of
NO. In the absence of NO, the product studies are complicated
by the presence of hydroperoxides and ethene glycol from the
self-reaction, which are very difficult to calibrate. However,
the variation of the yields of formaldehyde and glycolaldehyde
measured by Barnes et al. as a function of O2 pressure is
consistent with the present results.
Cl-Initiated Oxidation of Ethene: Chemistry of the
ClCH2CH2O Radical. A study of the Cl-initiated oxidation
of ethene was also carried out, using photolysis of Cl2 as the
source of Cl atoms. Initial experiments were conducted in the
absence of NO such that the major fate of the ClCH2CH2O2
radical was the self-reaction 17.25,27
Figure 4. Arrhenius plot showing the logarithm of the rate constant
ratio k6/k7 as a function of inverse temperature. Solid line is a least-
squares fit to the data.
in our system.32 The observed product yields can be used to
derive a branching ratio, k17a/(k17a + k17b) ) 0.55, in good
agreement with the data of Yarwood et al.25
The large yield of ClCH2CHO and the absence of CH2O in
the Cl-atom-initiated oxidation of ethene in air clearly indicate
the predominance of reaction 19 over reaction 18 under these
conditions, consistent with previous studies.25,27 An analogous
mechanism for the formation of BrCH2CHO from the reaction
of Br with ethene has also been reported.25,33 Only at very low
O2 concentrations (2-10 Torr) was any formaldehyde observed
in our experiments; these data were consistent with a rate
coefficient ratio of k18/k19 ) (6 ( 3) × 1015 molecule cm-3
.
Cl + C2H4 + M f ClCH2CH2 + M
(15)
Further experiments were carried out in 700 Torr of air with
added NO (10-80 mTorr) to look for evidence of chemical
activation of the oxy radical. Complications arise in this system
due to the generation of OH via reaction 19 followed by reaction
10. The subsequent reaction of OH with ethene then leads to
production of CH2O, which makes unambiguous identification
of reaction 18 difficult. While ClCH2CHO and CH2O were
found in almost comparable yields, no HCOCl was observed.
Simulations of the experiments using a chemical box model of
the system showed that the observed formaldehyde could all
be attributed to OH-initiated chemistry and that very little (if
any) decomposition of the ClCH2CH2O radical was occurring.
Thus, even with NO present, ClCH2CHO is again the dominant
product formed in the Cl-initiated oxidation of ethene, and no
significant chemical activation effect is observed, in contrast
to the OH-initiated oxidation.
Theoretical Prediction of Prompt Decomposition. Reac-
tions of alkylperoxy radicals with NO, forming alkoxy radicals
and NO2, are exoergic by some 10-15 kcal mol-1; moreover,
the reactants can possess significant amounts of thermal internal
energy. Since these processes occur via an intermediate
combination product, a peroxynitrite, much of the excess energy
should end up as vibrational energy of the largest product, the
oxy radical. The term “vibrational” includes here the internal
rotations. The nascent vibrational energy distribution of the
oxy radical, P(Ev), together with the energy-specific rates of
the competing unimolecular dissociation and collisional sta-
bilization processes, will determine the fraction of oxy radicals
that decompose promptly, (1 - R). If the lifetime τ of the
ClCH2CH2 + O2 + M f ClCH2CH2O2 + M (16)
2ClCH2CH2O2 f 2ClCH2CH2O + O2
(17a)
f ClCH2CHO + ClCH2CH2OH + O2
(17b)
ClCH2CH2O f CH2Cl + CH2O
(18)
ClCH2CH2O + O2 f ClCH2CHO + HO2
(19)
ClCH2CH2O2 + HO2 f ClCH2CH2OOH + O2 (20)
Significantly different branching ratios k17a/(k17a + k17b) have
been reported, 0.69 ( 0.02 by Wallington et al.27 and 0.57 (
0.02 by Yarwood et al.25 The CH2Cl radical, if formed, would
be converted predominantly to HCOCl in air and to both CO
and HCOCl at low O2.28,29
Products observed in our system in 700 Torr of synthetic air
were chloroacetaldehyde, ClCH2CHO (molar yield of 86%),
ClCH2CH2OH (yield of 16%), and HCOCl (yield of 3%). The
hydroperoxide product of reaction 20 was not observed in our
system, although modeling studies using the Acuchem program30
with published rate coefficient data31 for reactions 17 and 20
indicate that its yield should be about 35%. The excellent mass
balance observed in our experiments leads us to the conclusion
that the hydroperoxide is converted to ClCH2CHO and H2O on
the chamber surfaces, similar to the behavior of other peroxides

OH- and Cl-Initiated Oxidation of Ethene
J. Phys. Chem. A, Vol. 102, No. 42, 1998 8121
activated intermediate is sufficiently long for random partitioning
of its internal energy over all its vibrational degrees of freedom,
the P(Ev) distribution will also be governed by statistics and
hence can be predicted with an adequate theory. Generally,
the ergodicity condition is met for τ g 10-12 s. By adoption
of relative energies for the reactants, the peroxynitrite intermedi-
ate and the products of reaction 3, as for the 1-butyl analogue,34
one finds an HOCH2CH2OONO* RRKM lifetime τ ≈ 5 × 10-12
s, which meets the ergodicity criterion; at the same time, τ is
short enough to be able to neglect collisional energy loss of the
ROONO* under atmospheric conditions. The dissociation of
peroxynitrites into RO + NO2 occurs on a “type II” potential
energy surface (PES) without a PE maximum in the exit
channel,34 implying a loose and late transition state (TS). Nitrate
formation from the short-lived activated peroxynitrite cannot
compete with the very fast prompt dissociation because the far
more rigid ROONO f RONO2 transition state is entropically
at a disadvantage.
Figure 5. Predicted nascent vibrational energy distribution P(Ev) of
the HOCH2CH2O radical formed in reaction 3 at 298 K; energy grain
size ) 0.239 kcal mol-1. Predicted probability f(Ev) of prompt
dissociation at 760 Torr N2 for Eb(6) ) 11 kcal mol-1
.
The separate statistical ensemble (SSE) theory for statistical
partitioning of the available energy over the various degrees of
freedom (d.f.) of the dissociation products of a vibrationally
activated parent was developed by Wittig et al.35 specifically
for dissociation on a type II PES via a loose TS. Of course,
ergodicity is the primary condition. The SSE theory is a refined
version of phase space theory. SSE recognizes explicitly that
the vibrational energies of the fragments of a dissociation
process are determined solely by the “hot” vibrations of the
activated parent. When a molecule with n vibrational modes
dissociates into two fragments, with n1′ and n2′ vibrational
modes, respectively, the remaining modes evolve into the r )
n - (n1′ + n2′) d.f. of relative translation/rotation motion of
the fragments. For a type II PES, the energy Ed available for
distribution over the above three sets of product d.f. is equal to
the vibrational energy in all n modes of the parent, Ea, less the
dissociation barrier Eo. Wittig et al.35 have shown that, for a
loose, productlike TS, a statistical distribution of Ed over all n
internal d.f. of the dissociating molecule amounts to equal
probabilities for all product quantum states with total energy
Ed. Therefore, the distribution of Ed over the three sets of d.f.
involved, the n1′ and n2′ vibrations of fragments 1 and 2,
respectively, and the r d.f. of relative motion, will be determined
uniquely by the combined densities of quantum states for each
of these sets or, more precisely, by the energy dependences
thereof. The r modes of relative motion are treated as
unhindered free particle d.f., with combined state density
therefore varying as E((r/2)-1). SSE has been validated for a
number of photodissociation processes.35,36 It must be stressed
that the basic tenet of the SSE theory, i.e., that there is no free
flow of energy between the d.f. of overall rotation of a molecule
and its internal modes and that therefore only the internal energy
of the parent will determine the internal energy distribution of
the separating fragments, is not dependent on the way the
activated molecule is prepared, be it by photoexcitation or
otherwise.
in a trans position with respect to the oxy carbon; the lowest of
these (“trans-trans”) lies 1.9 kcal mol-1 above the cis-cis (all
relative energies include zero-point vibration energy differences).
In implementing SSE, the vibration manifolds of all rotamers
were folded together; this amounts to assuming that discrete
vibrational states of different rotamers can arise in reaction 3
with an equal probability, confined by the energy, entirely in
keeping with the basic SSE philosophy. The final P(Ev)
function for the RO resulting from the thermal reaction 3 must
be obtained by integration over the distribution of the Ed
energies. For this reaction without barriers in the entrance or
exit channels,34 the internal energy available to the products,
Ed, is equal to the reaction exoergicity, -∆Er(3), plus the internal
thermal energy of the (variational) entrance transition state, Eth,
which derives from the thermal energy of the reactants and
which therefore shows a distribution. For ∆Er, the value of
-11 kcal mol-1 for the analogous CH3CH2O2 reaction26 was
adopted (an H bridge as in HOCH2CH2O is also found for
HOCH2CH2O2, appearing to be only slightly stronger38). The
Eth distribution of the entrance TS was obtained from the sum
of states function Gq(E)exp(-E/kT),39 using approximate fre-
quencies.38 The average value of Eth at 298 K is 4 kcal mol-1
.
As the peroxynitrite suffers no collisions during its lifetime,
the rotational energy of the exit transition state may differ from
the E(rot) ) (3/2)kT value39 for the entrance transition state, in
accordance with the corresponding moments of inertia. How-
ever, both transition states are variational and similarly loose
such that the rotational energy change should be at most ∼kT;
the correspondingly small effect (<0.6 kcal-1 mol-1) on the
available internal energy Ed was not taken into account. The
resulting nascent P(Ev) distribution function calculated for
HOCH2CH2O formed in reaction 3 at 298 K is shown in Figure
5. The average energy is 9.5 kcal mol-1, i.e., below the
“equipartition” value of 10.5 kcal mol-1; the fwhm width of
the distribution is almost 8 kcal mol-1
.
To implement SSE theory for the products of reaction 3, the
vibrational state densities of the products RO and NO2 were
calculated by exact counts.37 The HOCH2CH2O frequencies
were taken from a theoretical B3LYP-DFT/6-31G** study.38
This study identified several internal rotation conformers. The
lowest-lying, “cis-cis”, with the two O atoms nearly in a cis
arrangement (dihedral angle (50°, two enantiomers), is stabil-
ized by intramolecular hydrogen bonding between the hydroxyl
H and the oxy O. The other rotamers, denoted as “trans”, have
the two O atoms in a trans arrangement and/or the hydroxyl H
Of the activated oxy radicals with internal energy Ev higher
than the dissociation barrier Eb(6), a fraction f(Ev) will decom-
pose promptly, in competition with collisional stabilization.
Since the energy-specific unimolecular dissociation rate constant
k(Ev) increases rapidly with Ev, the fraction f(Ev) tends to unity
for large Ev. The f(Ev) and the overall fraction of oxy radicals
decomposing promptly, (1 - R), which are both pressure
dependent, were obtained by master equation analysis of the
combined unimolecular dissociation/collisional energy transfer
system, with P(Ev) as input. As master equation code we used

8122 J. Phys. Chem. A, Vol. 102, No. 42, 1998
Orlando et al.
our recently developed fast and rigorous direct calculation of
product distribution method,40 which was inspired by the exact
stochastic method of Gillespie.41 The k(Ev) are calculated
according to RRKM theory, based on the (scaled) B3LYP-DFT
vibration frequencies.38 Densities and sums of states needed
for the k(Ev) are obtained by exact counts. Collisional energy-
transfer probabilities were calculated according to the biexpo-
nential Troe model, adopting an average energy transfer per
RO-N2 collision, ∆E , of -100 cm-1 42 and obeying detailed
balance.
demonstrate the sensitivity of the theoretical predictions to the
choice of Eb; moreover, considering that the experimental value
of (1 - R) is 25%, the calculated results endorse an Eb in the
10-11 kcal mol-1 range. The theoretical result is also sensitive
to the value of the exoergicity of the RO2 + NO reaction,
-∆Er(3). Increasing it from 11 to 13 kcal mol-1, for Eb ) 11
kcal mol-1, entails an increase of (1 - R) from 21.5 to 40%.
On the other hand, the calculated values are only moderately
sensitive to the relative stability of the two types of rotamers
and to the vibrational parameters. Figure 5 shows the energy-
specific dissociation probability f(Ev) for Eb ) 11 kcal mol-1
.
The critical parameter here is the energy barrier for the
decomposition of HOCH2CH2O to CH2OH + CH2O, Eb(6).
There are four ways to estimate this barrier. First, from the
experimental data, an Arrhenius activation energy of about 10
kcal mol-1 was derived for the thermal dissociation at 700 Torr.
Second, the modified Evans-Polanyi correlation relating the
activation barrier Eb to the ionization potential IP and the
endothermicity ∆Hd recently proposed by Atkinson16 and shown
The small low-energy tail, below the 11 kcal mol-1 barrier, is
due to collisional activation effects. There will of course be a
pressure effect on the rate of rise of the f(Ev) curve and hence
also on (1 - R); for total pressures of 380, 760, and 1520 Torr
of N2, the calculated (1 - R) values for Eb ) 11 kcal mol-1 are
26.1, 21.5, and 16.9%, respectively. However, such small
changes would be difficult to distinguish in the present
experiments.
in eq F suggests an activation energy of 12.5-14.5 kcal mol-1
depending on the adopted heat of dissociation.
,
Finally, it should be noted that the rate coefficient for the
thermal decomposition of HOCH2CH2O given by Atkinson,16
4.6 × 103 s-1 at 298 K, is much smaller than that measured
here, despite the lower pre-exponential factor deduced in this
work. This is a result of the much larger activation energy used
by Atkinson (14.5 kcal mol-1). Additionally, the present
calculations suggest that the thermal decomposition rate coef-
ficient is still a factor of 3 from the high-pressure limit at 700
Torr, a conclusion that awaits experimental confirmation.
Eb ) (2.4(IP) - 8.1) + 0.36∆Hd
(F)
Third, the above-mentioned B3LYP-DFT study38 identified a
“cis-cis TS”, with the hydroxyl H and oxy O still interacting,
-1
at 9.9 kcal mol
above the cis-cis oxy zero-point level and
also a set of “trans TS” rotamers an additional 1.9-3.2 kcal
mol-1 higher; note that the probable error on Eb for this level
of theory is 3 kcal mol-1 43
Finally, Eb can be estimated from
.
The prompt dissociation of ClCH2CH2O formed in the ClCH2-
CH2O2 + NO reaction is more difficult to quantify theoretically
because experimental data on the barrier to decomposition,
Eb(18), are lacking. Assuming that the difference between the
heats of formation (at 298 K) of ClCH2CH2O and CH3CH2O is
equal to that between C2H5Cl and C2H6, the heat of dissociation
of ClCH2CH2O can be estimated from known data26,46 as 14
kcal mol -1. The corresponding Eb predicted by the recent
Atkinson correlation16 is 17.5 kcal mol-1. Even when an Eb
value of only 15 kcal mol-1 is taken and an RO2 + NO f RO
+ NO2 exoergicity of 12 kcal mol-1 is assumed, the theoretically
calculated prompt dissociation fraction is negligibly small
(<2%), in agreement with the experimental findings. An upper
limit to the decomposition rate of 100 s-1 can be obtained from
the experimental ratio k6/k7 ) 1.3 × 1019 molecule cm-3 at
298 K and 700 Torr. Given that the known k7 values at 298 K
for small C2 and C3 oxy radicals are all close to 1 × 10-14 cm3
molecule-1 s-1 26
one can expect a similar value here. A k7
,
value close to 10-14 is indeed supported by the NO addition
experiments. In this way, one can estimate the thermal rate
coefficient k6 at 1.3 × 105 s-1 at 298 K, with an expected error
of less than a factor of 4. In the thermal dissociation, only the
lower-energy cis-cis rotamers of the oxy radical and of the
transition state are of importance. Using the B3LYP-DFT
vibration frequencies and moments of inertia,38 the pre-
exponential factor of the TS theory high-pressure rate coefficient
was calculated; (kBT/h)(Qq/Q) ) 1.1 × 1013 s-1 at 298 K. On
the basis of the Troe low-pressure nonequilibrium theory44 and
the Troe falloff formalism,45 it was found that the thermal rate
constant k6 at 700 Torr and 298 K is still almost 3 times lower
than its high-pressure limit. Hence, one derives Eb(6) ) 10.2
kcal mol-1, with a possible error of (0.8 kcal mol -1. This
value agrees very well with the experimental T dependence of
k6/k7. We therefore conclude that the barrier is in the 10-11
kcal mol-1 range.
The fraction of HOCH2CH2O* radicals decomposing promptly
was calculated for Eb values of 10 and 11 kcal mol-1. In the
master equation analysis, the (fast) interconversion of the
activated cis-cis and trans rotamers was incorporated, together
with their decomposition via the cis-cis TS and the trans TS’s.
Our master equation codes40 allow for several different activated
species simultaneously, each involved in a large number of
parallel and competing unimolecular reactions and collisional
energy-transfer processes. The results, for 298 K and 760 Torr
of N2 total pressure, can be summarized in terms of the overall
fraction (1 - R). For an Eb value (i.e., the relative energy of
the cis-cis TS with respect to the cis-cis radical ground state)
of 10 kcal mol-1, we calculated 38% prompt dissociation,
whereas for Eb ) 11, we obtained 21.5%. An Eb value of 13
kcal mol-1 results in only 4.8% dissociation. These data
our estimate of k18/k19 ) (6 ( 3) × 1015 molecule cm-3
,
assuming that the rate coefficient for reaction 19 is similar to
that for the reaction of ethoxy with O2. If the A factor for
decomposition is similar to that in reaction 6, an estimate of
the barrier height of 15-16 kcal mol-1 is obtained, in accord
with the experimental and theoretical findings. Thus, we
conclude that the major difference in the behavior of the
â-hydroxy- and â-chloroalkoxy radicals can be attributed to the
difference in the barrier heights to dissociation associated with
the higher ionization potential of the leaving group.
Atmospheric Implications. As in previous studies,12-14 the
OH-initiated oxidation of ethene has been shown to yield CH2O
and glycolaldehyde as the major products. However, we have
shown that the mechanism is more complicated than previously
assumed and occurs via the production of a chemically activated
oxy radical, HOCH2CH2O*. This mechanism must be taken
into account to understand the products obtained in the oxidation
of ethene in different regions of the atmosphere. In urban and
rural regions, the chemistry of the peroxy radical will be
dominated by its reaction with NO, reaction 3, and yields of
formaldehyde (1.6 molecules produced per ethene consumed)
and glycolaldehyde (0.2 molecules produced per ethene con-

OH- and Cl-Initiated Oxidation of Ethene
J. Phys. Chem. A, Vol. 102, No. 42, 1998 8123
sumed) are the same as those reported by Niki et al.13 In the
upper troposphere, NO levels are again sufficiently high to
control the RO2 reactivity and prompt decomposition of the
excited RO radicals will give a formaldehyde yield of about
0.5 molecules per ethene consumed. The remainder of the
radicals will be thermalized and react essentially exclusively
with O2 at the low temperatures present there, giving a
glycolaldehyde yield of about 0.75. In the cleaner regions of
the free troposphere and in the marine boundary layer, where
NOx concentrations are low, other reactions of the peroxy
radicals will occur (i.e., reaction with HO2 and possibly CH3O2
and other peroxy radicals), and the fate of ethene in these regions
will have to be determined by chemical models.
The oxidation of ethene by chlorine atoms is probably only
of significance in the marine boundary layer, particularly during
Arctic sunrise,47 where evidence for elevated Cl atom concen-
trations have been reported.48 The dominant fate of the ClCH2-
CH2O radical produced in the Cl-initiated oxidation of ethene
will be reaction with O2 to produce ClCH2CHO.
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Conclusions
The OH-initiated oxidation of ethene has been shown to yield
formaldehyde and glycolaldehyde as major products. The
oxidation in the presence of NOx has been shown to occur via
the formation of an excited oxy radical, HOCH2CH2O*. About
25% of these excited radicals decompose promptly to yield
formaldehyde, while the remainder are thermalized and undergo
competition between reaction 6 and reaction 7. The rate
coefficient ratio k6/k7 has been determined to be (2.0 ×
1025) exp(-4200/T) molecule cm-3 over the temperature range
250-325 K. The results are in good quantitative agreement
with calculations of the decomposition rates of both chemically
activated and thermalized oxy radicals. Of controlling impor-
tance to these decomposition rates is that the intramolecular
H-bonding stabilization by about 2 kcal mol-1 persists in the
transition state, notwithstanding the greatly increased C-C bond
length.38 The yield of formaldehyde in the oxidation of ethene
in the troposphere will be a strong function of temperature and
will range from a maximum value of 1.6 near the Earth’s surface
to a low value of 0.5 in the upper troposphere. In contrast, the
chlorine initiated oxidation of ethene does not lead to C-C bond
scission, even in the presence of NO.
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Acknowledgment. This work was funded by the Upper
Atmosphere Research Program of the NASA Mission to Planet
Earth. NCAR is sponsored by the National Science Foundation.
L.V. and J.P. are indebted to the European Commission for
financial support. Thanks are due to Frank Flocke of NCAR
for his help in the preparation of the methyl nitrite sample and
to Chris Cantrell and Laura Iraci for helpful comments on the
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