Kinetics of a Criegee intermediate that would survive high humidity and may oxidize atmospheric SO2

Article abstract of DOI:10.1073/pnas.1513149112

Criegee intermediates are thought to play a role in atmospheric chemistry, in particular, the oxidation of SO₂, which produces SO₃ and subsequently H₂SO0₄, an important constituent of aerosols and acid rain. However, the impact of such oxidation reactions is affected by the reactions of Criegee intermediates with water vapor, because of high water concentrations in the troposphere. In this work, the kinetics of the reactions of dimethyl substituted Criegee intermediate (CH₃)₂ COO with water vapor and with SO₂ were directly measured via UV absorption of (CH₃)₂COO under near-atmospheric conditions. The results indicate that (i) the water reaction with (CH₃)₂ COO is not fast enough (k_H2O < 1.5 × 10^-16 cm³s^-1) to consume atmospheric (CH₃)₂COO significantly and (ii) (CH₃)₂COO reacts with SO₂ at a near-gas-kinetic-limit rate (k_SO2 = 1.3 × 10^-10 cm³s^-1). These observations imply a significant fraction of atmospheric (CH₃)₂ COO may survive under humid conditions and react with SO₂, very different from the case of the simplest Criegee intermediate CH₂OO, in which the reaction with water dimer predominates in the CH₂OO decay under typical tropospheric conditions. In addition, a significant pressure dependence was observed for the reaction of (CH₃)₂COO with SO₂, suggesting the use of low pressure rate may underestimate the impact of this reaction. This work demonstrates that the reactivity of a Criegee intermediate toward water vapor strongly depends on its structure, which will influence the main decay pathways and steady-state concentrations for various Criegee intermediates in the atmosphere.

Full text of DOI:10.1073/pnas.1513149112

Kinetics of a Criegee intermediate that would survive
high humidity and may oxidize atmospheric SO2
Hao-Li Huang (黃皓立)a, Wen Chao (趙彣)a,b, and Jim Jr-Min Lin (林志民)a,b,1
aInstitute of Atomic and Molecular Sciences, Academia Sinica, Taipei 10617, Taiwan; and bDepartment of Chemistry, National Taiwan University, Taipei
10617, Taiwan
Edited by John H. Seinfeld, California Institute of Technology, Pasadena, CA, and approved July 27, 2015 (received for review July 4, 2015)
Typical water concentration in the troposphere (1.3 × 1017 to
Criegee intermediates are thought to play a role in atmospheric
chemistry, in particular, the oxidation of SO2, which produces SO3
8.3 × 1017 cm−3 at the dew point of 0–27 °C) is orders of magnitude
higher than those of atmospheric trace gases like SO2, NO2, and
volatile organic compounds (VOC) (on the order of 1012 cm−3 or
less). Although it has been shown that CIs may react very fast with
SO2, NO2, and organic acids (2, 3, 14), the reactions of CIs with
atmospheric water vapor would still strongly influence the fates and
concentrations of atmospheric CIs (see Fig. 1 for a simplified
schematic). As expected, the reactivity of CIs toward water vapor
would govern the modeling results of atmospheric H2SO4 forma-
tion from CIs (8, 15, 16).
However, there had been discrepancies about the reactivity of
CIs toward water. Whereas studies (17–20) using C2H4 ozonolysis
as a CH2OO source show substantial reactivity of CH2OO toward
water vapor, despite a large scatter (10−17 to 10−12 cm3s−1) in the
reported rate coefficient, other studies (2, 10, 21) using the CH2I+O2
reaction as a CH2OO source reported negative observation for the
CH2OO reaction with water vapor.
More recently, Chao et al. (22) and Berndt et al. (23) in-
vestigated the reaction of CH2OO with water vapor using the
CH2I+O2 reaction and the C2H4 ozonolysis as their CH2OO
sources, respectively. Both groups observed clear second-order
kinetics with respect to the concentration of water and concluded
that reaction with water dimer predominates in the decay of
CH2OO under atmospheric conditions and that previous studies
may require some reinterpretations. The reported rate coefficient
of the CH2OO reaction with water dimer is large, about 7 × 10−12
cm3s−1 (22), leading to extremely fast decay rate of CH2OO under
typical tropospheric conditions (Table 1).
and subsequently H2SO4, an important constituent of aerosols and
acid rain. However, the impact of such oxidation reactions is af-
fected by the reactions of Criegee intermediates with water vapor,
because of high water concentrations in the troposphere. In this
work, the kinetics of the reactions of dimethyl substituted Criegee
intermediate (CH3)2COO with water vapor and with SO2 were di-
rectly measured via UV absorption of (CH3)2COO under near-atmo-
spheric conditions. The results indicate that (i) the water reaction
with (CH3)2COO is not fast enough (kH2O < 1.5 × 10−16 cm3s−1) to
consume atmospheric (CH3)2COO significantly and (ii) (CH3)2COO
reacts with SO2 at a near–gas-kinetic-limit rate (kSO2 = 1.3 × 10−10
cm3s−1). These observations imply a significant fraction of atmo-
spheric (CH3)2COO may survive under humid conditions and react
with SO2, very different from the case of the simplest Criegee
intermediate CH2OO, in which the reaction with water dimer pre-
dominates in the CH2OO decay under typical tropospheric conditions.
In addition, a significant pressure dependence was observed for the
reaction of (CH3)2COO with SO2, suggesting the use of low pressure
rate may underestimate the impact of this reaction. This work dem-
onstrates that the reactivity of a Criegee intermediate toward water
vapor strongly depends on its structure, which will influence the
main decay pathways and steady-state concentrations for various
Criegee intermediates in the atmosphere.
atmospheric chemistry Criegee intermediate SO2 oxidation
|
|
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chemical kinetics
Taatjes et al. (3) and Sheps et al. (24) have reported that the anti-
form of methyl-substituted CI (CH3CHOO, R1 = H in Fig. 1) reacts
with water vapor much faster than the syn- form (R1 = CH3 in
Fig. 1) does. Quantum-chemical investigations (25–27) as well as
studies of alkene ozonolysis (20, 28) also indicate that the structure
of a CI strongly influences its reactivity toward water vapor. If one
type of CI reacts slowly with water vapor but reacts quickly with
nsaturated hydrocarbons are emitted into the atmosphere in
large quantities from either human or natural sources.
U
Ozonolysis of unsaturated hydrocarbons produces highly reactive
Criegee intermediates (CIs) (1), which may (i) decompose to
radical species like OH radicals or (ii) react with a number of at-
mospheric species, for example, with SO2 to form SO3 and with
NO2 to form NO3 (2, 3). The SO2 oxidation by CIs has gained
special attentions because the SO3 product would be converted
into H2SO4, an important constituent of aerosols and acid rain (4–
8). For example, Mauldin et al. (4) have speculated that Criegee
intermediate reactions with SO2 may account for the discrepancy
between the observed and modeled concentrations of H2SO4 in a
boreal forest region, where various alkenes are emitted by trees.
Recently, Welz et al. (2) demonstrated an efficient method to
prepare a CI in a laboratory by the reaction of iodoalkyl radical
with O2 (for example, CH2I + O2 → CH2OO + I). This method
can produce a CI of high enough concentration that allows direct
detection. With photoionization mass spectrometry (PIMS) de-
tection, Welz et al. (2) measured the rate coefficients of the sim-
plest CI (CH2OO) reactions with SO2 and NO2. Notably, these
new rate coefficients, confirmed by a few later investigations (9–
11), are orders of magnitude larger than those previously used (12,
13) in atmospheric models (e.g., MCM v3.3, available at mcm.leeds.
ac.uk/MCM/browse.htt?species=CH2OO), suggesting a greater role
of CIs in atmospheric chemistry. This result also indicates previous
ozonolysis analyses may be affected by complicated and partly un-
known side reactions and may contain errors in some of the reported
rate coefficients.
Significance
Ozonolysis of alkenes produces highly reactive Criegee in-
termediates. Whereas water dimer efficiently scavenges the
simplest Criegee intermediate CH2OO in the troposphere, this
study clear demonstrates that water vapor does not react
with dimethyl substituted Criegee intermediate (CH3)2COO, at
least not fast enough to significantly consume (CH3)2COO in
the troposphere. On the other hand, (CH3)2COO reacts with
SO2 three times faster than CH2OO does, indicating Criegee
intermediates of a structure similar to (CH3)2COO are poten-
tial candidates for an efficient oxidant in the atmospheric
SO2 oxidation.
Author contributions: J.J.-M.L. designed research; H.-L.H. and W.C. performed research;
H.-L.H. and W.C. analyzed data; and H.-L.H. and J.J.-M.L. wrote the paper.
The authors declare no conflict of interest.
This article is a PNAS Direct Submission.
1To whom correspondence should be addressed. Email: jimlin@gate.sinica.edu.tw.
This article contains supporting information online at www.pnas.org/lookup/suppl/doi:10.
1073/pnas.1513149112/-/DCSupplemental.
www.pnas.org/cgi/doi/10.1073/pnas.1513149112
PNAS
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R3
R2
R4
R1
the case, syn-CIs may have higher steady-state concentrations in the
troposphere and may still play an important role in the SO2 oxida-
tion; otherwise the steady-state concentrations of syn-CIs would still
be low due to their fast consumption by reactions with water vapor
(unless their sources are significantly larger than current estimation)
and we might need to find another candidate for the unknown ox-
idant [oxidant X in the work by Mauldin et al. (4)] in the SO2
atmospheric chemistry.
To shed some light on this important issue, we performed direct
kinetic measurements of the reactions of dimethyl-substituted CI,
(CH3)2COO, with water vapor and with SO2 under near-atmo-
spheric conditions. By introducing the water reactant at high
concentrations, the rate coefficient of (CH3)2COO reaction with
water can be better constrained. In contrast with the fast reaction
of CH2OO with water dimer, this result shows that the relative
probabilities of (CH3)2COO reactions with SO2 and with water
vapor are comparable in the troposphere, so water alone would
not completely scavenge (CH3)2COO, suggesting CIs of similar
structures may play a more important role in the atmospheric
oxidation of SO2.
O
O
SO2
O
+
SO3
O3
+
R2
R1
R2
H2O
Products
R1
O
R3
R4
(H2O)2
H2O
H2SO4
Fig. 1. Reaction scheme showing competitions for CIs between reactions
with water (monomer and dimer) and with SO2.
SO2, these CIs may accumulate to higher concentrations and have
higher probability to oxidize atmospheric SO2. Table 1 shows se-
lected rate coefficients for relevant CI reactions and the effective
first-order decay rate coefficients (keff) of small CIs under an at-
mospheric condition.
As will be discussed in detail in Discussion and Conclusions, the
steady-state concentration of a particular CI would depend on its
formation rate and effective decay rate coefficient; its impact on
the SO2 oxidation would further depend on its concentration and
reaction rate coefficient with SO2. Experimental results (Table 1)
show that CH2OO and anti-CH3CHOO react with water vapor
very quickly (3, 22–24). Thus, their steady-state concentrations
would be too low to have a significant impact in SO2 oxidation
under typical atmospheric conditions, as shown in modeling results
(15, 16). On the other hand, previous experimental data for syn-
CH3CHOO (3, 24) are not precise enough to determine its main
decay pathways in the atmosphere.
Quantum-chemistry (25–27) calculations predicted that the anti-
form of CIs (CIs with R1 = H in Fig. 1, including CH2OO) react
with water vapor very quickly and that the syn- form of CIs (CIs
with R1 ≠ H in Fig. 1, including dialkyl-substituted CIs) react
slowly with water vapor. Here, steric hindrance of the alkyl group
may account for the structure dependence in the reactivity. How-
ever, due to uncertainty in the calculated rate coefficients, it is
unclear about the main decay channels of the syn-CIs in the
atmosphere. For example, some theoretical investigation (29) shows
that the reactions of water vapor with syn-CIs may still be fast
enough (with a large uncertainty) to efficiently scavenge atmospheric
syn-CIs, whereas some other calculations (25) suggest that these
reactions are too slow to consume syn-CIs significantly. If the latter is
Results
Fig. 2A shows time-resolved difference absorption spectra recor-
ded in the (CH3)2CI2/O2 pulsed photolysis system, in which the
laser pulse defines delay time t = 0. The photodissociation of the
(CH3)2CI2 precursor produces (CH3)2CI radicals; (CH3)2COO was
formed through the reaction of (CH3)2CI + O2 → (CH3)2COO + I
(30). At t = 50 μs, a strong absorption band peaked at 330 nm was
observed, which decays with the delay time. At very long delay time
(e.g., 2,000 μs), a negative difference absorption signal peaked at
296 nm was observed, which corresponds to the depletion of the
(CH3)2CI2 precursor. Fig. 2B shows the spectra recorded at a very
similar condition but adding SO2 gas to scavenge CIs. The ab-
sorption signal in Fig. 2B consists mainly of the depletion of
(CH3)2CI2 and a small amount of IO [a byproduct, likely from I +
(CH3)2COO → IO + (CH3)2CO]. Other possible species in the
absorption cell are SO3 and acetone, which absorb rather weakly in
the studied wavelength region [the absorption cross-sections σ for
the relevant species are σ(SO3) < 1 × 10−21 cm2, σ(acetone) < 5 ×
10−20 cm2, and σ ∼ 10−17 cm2 for a CI] (30, 31). Following the
established method of spectral analysis in similar systems (22, 31–
33), the absorption band of (CH3)2COO can be obtained from the
difference between the spectra in Fig. 2 A and B and is shown in
Fig. 2C. The resultant spectrum of (CH3)2COO is slightly broader
Table 1. Reported bimolecular rate coefficients k and effective first-order rate coefficients (keff
k[Coreactant]) for simple CI reactions with H2O, (H2O)2, and SO2 at a given atmospheric
=
condition
CI
Coreactant
[Coreactant]/cm−3
k/cm3s−1
keff/s−1
Reference
CH2OO
H2O
(H2O)2
SO2
5.4 × 1017
6.0 × 1014
1.2 × 1012
5.4 × 1017
<1.5 × 10−15
6.5 × 10−12
3.9 × 10−11
1.0 × 10−14
2.4 × 10−14
6.7 × 10−11
2.2 × 10−10
<4 × 10−15
<2 × 10−16
2.4 × 10−11
2.9 × 10−11
<1.5 × 10−16
<1.3 × 10−13
1.3 × 10−10
<810
3,900
47
5,400
13,000
80
(22)
(22)
(2)
(3)
(24)
(3)
(24)
(3)
(24)
anti-CH3CHOO
H2O
SO2
H2O
SO2
1.2 × 1012
5.4 × 1017
1.2 × 1012
260
syn-CH3CHOO
<2,200
<110
29
(3)
(24)
35
(CH3)2COO
H2O
(H2O)2
SO2
5.4 × 1017
6.0 × 1014
1.2 × 1012
*
*
<81*
<78*
160
This work
This work
This work
The assumed concentrations of H2O and SO2 correspond to a relative humidity RH = 70% and a SO2 mixing
ratio of 50 ppb at 298 K and 1 atm. Only data from direct kinetic measurements are selected.
*The rate constant of the H2O reaction with (CH3)2COO is obtained by assuming the rate constant of the (H2O)2
reaction with (CH3)2COO is zero and vice versa (SI Appendix, Table S2); thus, these two effective decay rates should
not be added together.
10858
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Huang et al.

reaction between two (CH3)2COO molecules did not dominate
in the decay of (CH3)2COO, leading to the observed first-order–
like decay.
The decay of the (CH3)2COO signals can be well described by
Eqs. 1 and 2. Under the high O2 pressure (∼10 torr) used in this
study, the formation time of (CH3)2COO is within 1 μs. This for-
mation time is relatively short in comparison with the decay time of
(CH3)2COO. Thus, the fitting for the decay typically started at t =
10 μs to decouple the formation kinetics.
0.03
0.02
0.01
0.00
-0.01
A [SO2] < 2x1012 cm−3
delay time:
50 μs
150 μs
250 μs
500 μs
2000 μs
-0.02
0.03
Â
Ã
Â
Ã
delay time:
50 μs
B [SO2] = 5.3x1014 cm−3
ðCH3Þ2COO = ðCH3Þ2COO 0e−k′t
,
[1]
[2]
0.02
0.01
150 μs
250 μs
500 μs
2000 μs
IO formation
Â
Ã
Â
Ã
d ðCH3Þ2COO
−
= k′ ðCH3Þ2COO
= ðk0 + kSO2½SO2ꢀÞ ðCH3Þ2COO .
dt
0.00
Â
Ã
-0.01
-0.02
Fig. 4 shows the observed pseudo–first-order rate coefficient (k′− k0)
as a function of SO2 concentration, where the k0 term (Eq. 2)
accounts for the (CH3)2COO decay when no SO2 was added
(see SI Appendix, Table S1 for the values of k0). A linear relation-
ship between k′ and [SO2] is found with the slope corresponding
to the rate coefficient kSO2 for the (CH3)2COO reaction with
SO2. In Fig. 4, it is obvious that the slope is different at differ-
ent total pressure.
(CH3)2CI2 depletion
CDifference spectra between A and B
delay time:
50 μs
0.03
0.02
0.01
0.00
150 μs
250 μs
500 μs
2000 μs
Fig. 5 plots the reaction rate coefficient kSO2 as a function of the
total pressure with N2 as the buffer gas. At pressure higher than
100 torr ([M] = 3.2 × 1018 cm−3), the rate coefficient is leveled.
This high-pressure-limit rate coefficient is 1.3 × 10−10 cm3s−1, about
3 times the reported rate coefficient of CH2OO reaction with SO2
[kCH2OO+SO2 = (3.9 0.7) × 10−11 cm3s−1] (2). At low pressure, the
rate coefficient drops. A simple model in Fig. 6 is used to analyze
the experimental data.
Quantum-chemistry studies (29) suggested that CI reaction
with SO2 would first go through a barrierless formation of a
cyclic intermediate, INT* (R1). In addition, because the for-
mation of the cyclic intermediate is quite exothermic, its sta-
bilization requires collision with a third body (R2). Fur-
thermore, the energized intermediate INT* may dissociate
back to the reactants or form products directly (likely SO3 +
carbonyl); (R4) connects the stabilized intermediate INT to
the final products.
280 300 320 340 360 380 400 420 440 460 480 500
Wavelength / nm
Fig. 2. Typical transient absorption spectra recorded in the (CH3)2CI2/O2
photolysis system at selected delay times after the photolysis laser pulse
(t = 0). Total pressure was 100.1 torr (PN2 = 89.9 torr, PO2 = 10.2 torr,
[(CH3)2CI2]0 = 1.1 × 1014 cm−3). (A) Spectra recorded without adding SO2.
(B) Spectra recorded at [SO2] = 5.3 × 1014 cm−3; (CH3)2COO was consumed
by its reaction with SO2. (C) Difference spectra between the two sets of
spectra in A and B. The main spectral carrier is assigned to (CH3)2COO be-
cause the reaction of (CH3)2COO with SO2 is fast. The residue contributions
of the IO byproduct and the (CH3)2CI2 precursor are minor and have been
removed in C.
than but still consistent with the jet-cooled spectrum reported by
Liu et al. (30). See SI Appendix, Fig. S2.
2.5
3
[SO2] < 2.0x1012 cm−
Exp #3, Ptotal = 50.0 Torr
In the kinetic measurements, we chose a detection window of
335–345 nm to monitor the change in (CH3)2COO concentrations.
In this detection window, the absorption signal of IO is negligible
[σ(IO) = 2.7 × 10−19 cm2 << σ((CH3)2COO) ∼ 10−17 cm2] (22); the
absorption of the precursor (CH3)2CI2 is small but not negligible.
Fortunately, Fig. 2B shows that the depletion of the precursor is a
constant after the photolysis laser pulse, which would not affect our
kinetic analysis for (CH3)2COO.
3
3
3
[SO2] = 2.5x1013 cm−
[SO2] = 4.8x1013 cm−
[SO2] = 9.9x1013 cm−
2.0
1.5
1.0
0.5
0.0
-0.5
Representative difference transient absorption traces re-
corded at various SO2 concentrations are shown in Fig. 3; a
more complete set of the experimental data can be found in
SI Appendix, Figs. S3–S6 and Table S1. As mentioned above,
(CH3)2COO is the main spectral carrier for the absorption at
340 nm. The rapid rise of the signal after the photolysis laser
pulse was due to (CH3)2COO formation. When SO2 is added, the
reaction of (CH3)2COO with SO2 dominates in the observed
decay of (CH3)2COO. When no SO2 is added, the (CH3)2COO
decay is due mainly to reactions of (CH3)2COO with radical
species, including I atoms, OH radicals [possibly from decom-
position of (CH3)2COO (30)], and (CH3)2COO itself, similar to
the case of CH2OO (33). Because the concentration of (CH3)2COO
in our experiment was low (on the order of 1011 cm−3), the self-
0
200
400
600
800
1000
Time / μs
Fig. 3. Typical temporal profiles of the absorbance change at 340 nm. The
change in absorbance was mainly due to (CH3)2COO and was monitored in
real time with a balanced photodiode detector at various SO2 concentra-
tions. The smooth lines are single exponential fit to the data.
Huang et al.
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[H2O], even for high water concentration up to 7 × 1017 cm−3
(RH = 90% at 298 K). Fig. 7 shows the first-order rate coefficients
k′ of (CH3)2COO decay as a function of [H2O]. We have varied
the experimental conditions including the total pressure, pre-
cursor concentration, photolysis laser power, etc. (SI Appendix,
Table S2) and the results show that the effect of water in the decay
of (CH3)2COO is extremely weak. More detailed analysis indicates
the decay rate coefficient of (CH3)2COO depends mainly on the
radical concentration (SI Appendix, Fig. S12). We estimate an
upper limit of 1.5 × 10−16 cm3s−1 for the rate coefficient of the
(CH3)2COO reaction with H2O, based on the fluctuation of the
experimental data (SI Appendix, Table S2).
300 Torr, Exp #10
100 Torr, Exp #7
50 Torr, Exp #4
50 Torr, Exp #3
20 Torr, Exp #1
14000
12000
10000
8000
6000
4000
2000
0
Discussion and Conclusions
So far, there is no method available to detect the concentration
of any CI in the atmosphere. The estimation of the CI concen-
tration can only be done by knowing its formation rate and
consumption rate. The impact of a given CI (e.g., to the SO2
oxidation) depends on its concentration and the rate coefficient
of its reaction with SO2, as shown in the following (Eqs. 5–9):
0
2
4
6
8
10
12
[SO2] / 1013cm−
3
Fig. 4. Pseudo–first-order rate coefficient of (CH3)2COO reaction with SO2
plotted as a function of [SO2]. k0 is the rate coefficient at zero [SO2]. The
measurements were performed at 298 K with N2 as the buffer gas at the
indicated total pressure. The error bars indicate the SDs; the number of in-
dependent data can be found in SI Appendix, Table S1 for each experiment.
Solid lines are linear fit to the data at the corresponding pressure. Data at
higher and lower pressure ranges are shown in SI Appendix, Figs. S7 and S8.
O3 + alkene → CI kform
,
[5]
[6]
CI → Products kdecay
,
ÃÂ
Ã
kform½O3 alkene
½CIꢀss =
,
[7]
[8]
Applying steady-state approximation to INT* leads to the
following:
kdecay
Â
Ã
k
decay = ktherm + kH2O½H2Oꢀ + kw2 ðH2OÞ2
d½CIꢀ
−
= k1½CIꢀ½SO2ꢀ − k−1½INTpꢀ = kSO2½CIꢀ½SO2ꢀ,
ꢀ ꢀ ꢀ ꢀ ꢀ ꢀ ꢀ ꢀ k1ðk3 + k2½MꢀÞꢀ ꢀ ꢀ ꢀ ꢀ ꢀ ꢀ ꢀ
[3]
[4]
+ kSO2½SO2ꢀ + kNO2½NO2ꢀ + . . . ,
dt
ÃÂ
Ã
Â
Ã
Â
Ã
kform½O3 alkene
d½SO2ꢀ
kSO2
=
.
= −kSO2½CIꢀss SO2 = −kSO2
SO2
,
ꢀ k−1 + k3 + k2½Mꢀꢀ
dt
kdecay
[9]
Theoretical analysis (29) suggests that (R2) is the main pathway
(97%) for the (CH3)2COO reaction with SO2 under atmospheric pres-
sure. Thus, we may simplify the mechanism by assuming k2[M] >>
k3, unless the pressure is low. (R4) may be a slow reaction, but if
there are no other competing processes, the yield of R4 may still
be high, as indicated by previous ozonolysis studies (20, 28, 34).
The best fit of Eq. 4 to the data of Fig. 5 is shown as the red
line (the best fit gives k3 =0). At the high-pressure limit, kSO2 = k1 =
1.32 × 10−10 cm3s−1. Nonzero k3 leads to a nonzero low-pres-
sure-limit rate constant k1k3/(k-1+k3) but a value larger than 6 ×
10−11 cm3s−1 for this term cannot fit the data satisfactorily (green
line in Fig. 5). The error in k1 mostly comes from the un-
certainty in the absolute concentrations of SO2, which is less than
where [CI]ss is the steady-state concentration of the CI; kform is the
formation rate constant of the CI from its precursors (e.g., ozone
Total Pressure / Torr
0
100
200
300
400
500
600
700
800
14
12
10
8
k1k3
k−1 + k3
= 6×10−11cm3s−1
10%. Thus, we report k1 = (1.32 0.13)×10−10 cm3s−1
.
To further elucidate the role of the buffer gas, we changed the
buffer gas from N2 to CO2 or Ne and performed similar experiments
at 50 torr. SI Appendix, Fig. S9 shows that CO2 is a more efficient
collider than N2 and Ne is a less efficient collider, as one may expect.
Furthermore, we also investigated the pressure effect in the reaction
of CH2OO with SO2. However, the reaction rate coefficient of
CH2OO with SO2 does not exhibit significant pressure dependence
(see SI Appendix, Figs. S13–S15 and Table S3 and ref. 10).
The pressure dependence of the reaction of (CH3)2COO with
SO2 has implications in atmospheric chemistry. Applying low-
pressure rates in a model may underestimate the impact of the
(CH3)2COO reaction with SO2. Similarly, the pressure depen-
dence of the reactions of SO2 with other CIs needs to be inves-
tigated because quite a number of rate coefficients of CI reactions
are measured under low-pressure conditions [∼4 torr for PIMS
experiments (2, 3, 14)].
6
k1[M ]
(1.32 0.02)×10−10[M ]
(4.88 0.32)×1017 +[M ]
kSO2
=
=
k−1
k2
4
+[M ]
2
0
0
5
10
15
20
25
Total Number Density / 1018cm−3
Fig. 5. Observed second-order rate coefficient of (CH3)2COO reaction with SO2
plotted as a function of the total pressure and number density. The fitting is
plotted as solid lines. The error bars of the data indicate the SDs. The errors in the
fitted equation are the SDs resulting from the fitting only, not including any
possible systematic uncertainty yet. There are 18 independent measurements.
The kinetics of the (CH3)2COO reaction with water vapor were
also measured. In sharp contrast with the fast reaction of CH2OO
with water vapor, the decay of (CH3)2COO barely depends on
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Huang et al.

discrepancy may originate from the complexity of the ozonolysis
experiments, in which it is difficult to fully quantify the side re-
actions during the long reaction time.
*
O
k1
O
k3
O
S
(CH3)2COO + SO2
dimethyl CI
Products
Products
O
(R3)
k-1
(R1)
It is important to note that the reactivity of a particular CI
toward water vapor strongly depends on its structure. Table 1
indicates that a methyl group substitution for R1 may alter the
rate coefficient by orders of magnitude. Although ozonolysis
experiments (20, 28, 34) have also shown a trend for the struc-
ture dependence, the magnitude of the reactivity difference is
much smaller than what is observed in the direct kinetic mea-
surements (Table 1). In addition, the ozonolysis experiments
cannot distinguish the anti- and syn- conformers of CIs, which
have quite different reactivity.
INT*
M
M
(R2)
k2 [M]
O
k4
O
O
S
O
INT
(R4)
Fig. 6. Proposed mechanism for CI reaction with SO2.
The results summarized in Table 1 give direct evidence that the
reaction of water with (CH3)2COO is greatly hindered by the
methyl group at the R1 position. Interestingly, methyl substitution
seems to enhance the reactivity of CIs toward SO2. One can
imagine that CIs with more complicated substitution groups may
also react with water slowly but react with SO2 quickly, similar to
(CH3)2COO. Such CIs may be the candidate for the oxidant X in
the SO2 oxidation, an important issue raised recently (4).
Considering that much of the VOC emissions consist of a large
variety of alkenes, ranging from simple alkenes like C2H4 and C3H6
to bigger alkenes like isoprene, monoterpenes, sesquiterpenes, etc.,
various CIs are expected to form in the atmospheric ozonolysis
reactions. To assess the impact of the CI+SO2 reaction class on the
atmosphere, it is critical to know the atmospheric concentration of
CIs. As mentioned above, it requires the relevant rate coefficients to
estimate the atmospheric concentration of a CI. Water is the third
most abundant molecule in the air. The reactions of CIs with water
are crucial in determining the concentrations and fate of the CIs.
The slow rates of water reactions with (CH3)2COO or with similar
CIs are very difficult to measure, but very important in estimating
the concentrations of those CIs and thus the oxidizing capacity of
the atmosphere.
and an alkene); kdecay would be the summation of the effective
decay rate constants keff of all possible decay pathways; ktherm is
the thermal decomposition rate coefficient; kH2O, kw2, kSO2
,
and kNO2 are the bimolecular rate coefficients of CI reactions
with H2O, (H2O)2, SO2, and NO2, respectively. The ktherm for
(CH3)2COO has been reported recently to be 3 0.4 s−1 at 293
K (34) with a significant temperature dependence (28).
Considering the possible concentrations of H2O and SO2 in the
troposphere, we can evaluate the effective decay rate constants of
simple CIs by water reaction and by SO2 reaction. As shown in
Table 1, the consumptions of CH2OO and anti-CH3CHOO by
water (including monomer and dimer) are extremely fast, >103 s−1
,
predominating in their kdecay. The very large kdecay would result in
very low steady-state concentrations for these CIs, limiting their
roles in oxidizing other atmospheric species like SO2 and NO2 (see
Eqs. 7–9).
On the other hand, the reaction of (CH3)2COO with H2O is much
slower, such that the water reaction would not limit the (CH3)2COO
concentration in the troposphere anymore. In addition, the rate
coefficient of (CH3)2COO reaction with SO2 is larger than those of
other known CI reactions with SO2, indicating a greater role of
(CH3)2COO in the atmospheric SO2 oxidation. In other words, there
is a strong structure dependence in the CI oxidation of SO2. For
(CH3)2COO, its kdecay is smaller and its kSO2 is larger, such that the
SO2 oxidation rate by (CH3)2COO is faster in comparison with those
for CH2OO and anti-CH3CHOO, assuming similar formation rates
(see Eq. 9). In addition, for a CI with small kH2O and kw2, its ktherm
would be another important factor that also influences its [CI]ss and
oxidation capacity. A related issue is that the thermal decomposition
of syn-CIs [including (CH3)2COO] may form OH radicals through a
1,4 H-migration process (30), which may be an important non-
photolytic OH source in the troposphere. Again, structure de-
pendence of ktherm for various CIs would need further investigation.
Previous studies of ozonolysis of 2,3-dimethyl-2-butene (tet-
ramethyl ethylene, R1 = R2 = R3 = R4 = CH3) (20, 28, 34) have
provided some information about the (CH3)2COO reactions
with H2O and with SO2. But, the results are far from consistency.
First, the reported kSO2 are on the order of 10−13 cm3s−1 (34),
much smaller than the value determined from the direct kinetic
measurement (Table 1). A similar situation also happens for
smaller CIs like CH2OO and CH3CHOO (12, 13, 35), suggesting
there might be some systematic issues in the ozonolysis experi-
ments. Very recently, Newland et al. (20) measured the removal of
SO2 in the presence of alkene–ozone systems and concluded that
the SO2 removal displays a clear dependence on relative humidity
for all four alkene ozonolysis systems [ethene (to form CH2OO),
cis-2-butene (to form CH3CHOO), trans-2-butene (to form
CH3CHOO), and 2,3-dimethyl-2-butene (to form (CH3)2COO)],
confirming a significant reaction for stabilized CIs with H2O.
However, Berndt et al. (28) investigated the H2SO4 formation as a
function of [H2O] in similar ozonolysis reactions and reported
quite different values. The results for (CH3)2COO are summarized
in Table 2. The result by Newland et al. is not consistent with
the results of Berndt et al. and this work. The reason for this
Materials and Methods
The experimental setup has been described in detail elsewhere (22, 32). In
brief, (CH3)2COO was generated from photolysis of a gaseous mixture con-
sisting of (CH3)2CI2, O2, and buffer gas (N2) at 248 nm (KrF excimer laser)
via the established preparation method: (CH3)2CI2 + hν → (CH3)2CI + I;
(CH3)2CI +O2 → (CH3)2COO + I (2, 30). (CH3)2COO was monitored via its strong
UV absorption (30). Continuous probe light went through the photolysis reactor
(25 or 20 mm inner diameter, 76 cm long) six or eight times to enhance the
absorption signal.
1500
1250
1000
750
500
Exp #W1, 200 Torr
Exp #W2, 200 Torr
Exp #W3, 200 Torr
Exp #W4, 200 Torr
Exp #W5, 200 Torr
Exp #W6, 200 Torr
Exp #W7, 400 Torr
Exp #W8, 500 Torr
250
0
0
1
2
3
4
5
3
6
7
[H2O] / 1017cm−
Fig. 7. Pseudo–first-order rate coefficient of (CH3)2COO as a function of
[H2O]. Solid lines are linear fit to the data of each experiment. There are 223
data points (see SI Appendix, Table S2 for the experimental conditions).
Huang et al.
PNAS
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vol. 112
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Table 2. Ratio of the rate coefficients of (CH3)2COO reactions with H2O and with SO2, kH2O/kSO2
kH2O/kSO2
Method
Reference
<1.2 × 10−6
Absolute direct kinetic measurements of (CH3)2COO
Relative measurements of the H2SO4 product
Relative measurements of the SO2 and O3 consumptions
This work
Berndt et al. (28)
Newland et al. (20)
<4 × 10−6
(8.7 2.5) × 10−5
Spectral Measurements. Transient absorption spectra of the reaction system
were measured with a continuous broadband light source (Energetiq, EQ-99)
and a time-gated iCCD spectrometer (Andor, SR303i and DH320T-18F-03). The
light source was projected to the entrance of the absorption cell by an
achromatic lens (Thorlabs ACA254-100-UV). To enhance the absorption sig-
nal, the probe light was reflected eight times through the photolysis reactor
by a spherical mirror (R = 1 m, Thorlabs, CM750-500-F01) and a SiO2 prism.
The probe beam and the photolysis beam were overlapped collinearly in the
photolysis reactor. For the iCCD measurement, the reference spectrum was
recorded 200 μs before the photolysis pulse. Transient spectra at delay times
50, 100, 150, 200, 250, 300, 500, 1,000, and 2,000 μs were recorded.
5800E) and monitored by its UV absorption (190–330 nm) with a D2 lamp
(Ocean Optics, D-2000) and a spectrometer (Ocean Optics, Maya2000 Pro).
Relative Humidity Measurements. The relative humidity was adjusted by
controlling the mixing ratio of dry and moisturized buffer gases with mass
flow controllers (Brooks, 5850E or 5800E) and monitored with a humidity
sensor (Rotronic, HC2-S). The temperature for measuring the water reaction
was controlled at 298.2 0.5 K.
Precursor Preparation. The precursor, (CH3)2CI2, was synthesized following a
reported method (36). In brief, acetone was added to hydrazine monohydrate
(80 °C, >1 h) to obtain acetone hydrazone (CH3)2C=NNH2. Saturated solution of
iodine in ethyl ether was added to acetone hydrazone (at room temperature),
which is mixed with ethyl ether and triethylamine, to obtain the final product.
The structure of the synthesized (CH3)2CI2 was checked with H-NMR spectros-
copy [(CH3)2CI2: 3.00 ppm (6H, s, Me) in CDCl3] (37).
Kinetics Measurements. Absorption signal at 340 nm was measured in real
time by using a balanced photodiode detector (Thorlabs, PDB450A) and a
bandpass filter (Edmund Optics, 65129, 10-nm OD4 band pass filter) and the
same light source. A time-dependent transmittance change (<1%) was ob-
served after the photolysis pulse even without adding any sample. This
background did not depend on [H2O], or on [SO2]. It can be subtracted by
performing a background run under the same experimental condition except
adding (CH3)2CI2. The presented data are after the background subtraction.
ACKNOWLEDGMENTS. The authors thank Prof. Jim-Min Fang, Ms. Ling-Wei Li,
and Ms. Che-Hsuan Chang for help in organic synthesis and NMR measurement;
Ms. Liang-Chun Lin and Mr. Chun-Hung Chang for help in experiments; and
Prof. Yuan-Tseh Lee for discussion. This work was supported by Academia Sinica
and Ministry of Science and Technology, MOST 103-2113-M-001-019-
MY3, Taiwan.
[SO2] Measurements. The SO2 concentration was adjusted by controlling the
amount of SO2 in buffer gas with mass flow controllers (Brooks, 5850E or
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