Measuring Rate Constants for Reactions of the Simplest Criegee

Article pubs.acs.org/JPCA

Measuring Rate Constants for Reactions of the Simplest Criegee Intermediate (CH2OO) by Monitoring the OH Radical Yingdi Liu, Kyle D. Bayes,* and Stanley P. Sander Jet Propulsion Laboratory, California Institute of Technology, 4800 Oak Grove Drive, Pasadena, California 91109-8099, United States ABSTRACT: While generating the CH2OO molecule by reacting CH2I with O2, significant amounts of the OH radical were observed by laser-induced fluorescence. At least two different processes formed OH. A fast process was probably initiated by a reaction of vibrationally hot CH2I radicals. The second process appeared to be associated with the decay of the CH2OO molecule. The addition of molecules known to react with CH2OO increased the observed decay rates of the OH signal. Using the OH signals as a proxy for the CH2OO concentration, the rate constant for the reaction of hexafluoroacetone with CH2OO was determined to be (3.33 ± 0.27) × 10−11 cm3 molecule−1 s−1, in good agreement with the value measured by Taatjes et al.1 The rate constant for the reaction of SO2 with CH2OO, (3.53 ± 0.29) × 10−11 cm3 molecule−1 s−1, showed no pressure dependence over the range of 50−200 Torr and was in agreement with the value at 4 Torr reported by Welz et al.2

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INTRODUCTION The study of reactions of olefins with ozone has a long history. As first suggested by Criegee,3 the initial step appears to be an addition of the O3 across the carbon−carbon double bond, forming a five-member ring. In the gas phase this primary ozonide spontaneously splits apart4 to give two molecules, a carbonyl product and a carbonyl oxide, R1R2COO; the latter is usually referred to as a “Criegee intermediate”. Here R1 and R2 may be hydrogen atoms or carbon fragments, depending on the structure of the original olefin. Criegee intermediates can react in a variety of ways including undergoing an internal rearrangement to form a hot hydroperoxide, which can then decompose to give an OH radical.5−11 The generation of OH from ozone/olefin reactions is an important source of HOx in the atmosphere.5−14 Observed yields of OH range from 10% to nearly 100% depending on the olefin involved. Because the formation of OH can involve energy rich Criegee intermediates, these OH yields tend to be pressure dependent.15,16 In the case of the ozone/ethylene reaction, the simplest Criegee intermediate, formaldehyde oxide or CH2OO, should be formed. Theoretical calculations on this structure suggest that there is a barrier of about 30 kcal in order for CH2OO to rearrange to form OH and CHO.17−19 However, the overall reaction of ethylene with O3 is sufficiently exothermic to form CH2OO + CH2O or OH + CHO + CH2O. Studies at both low pressures and at 1 atm have shown yields of OH of 14−60% for the C2H4/O3 system.15,20 Recent work by Taatjes and co-workers1,2 has shown that it is possible to generate the simplest Criegee intermediate by reacting iodomethyl radicals with molecular oxygen, R.1 (the asterisk indicates excess vibrational energy). CH 2I + O2 → CH 2IOO* © 2014 American Chemical Society

CH 2IOO* → CH 2OO + I

(R.1b)

CH 2IOO* + M → CH 2OO + M

(R.1c)

By measuring ionization appearance potentials, they showed that the product of reaction R.1b was the Criegee intermediate, CH2OO, and not one of its isomers (formic acid or dioxirane). Also, by following the appearance of the iodine atom as a function of pressure, they were able to determine the pressure dependence for the formation of the iodoperoxy radical in reaction R.1c.21,22 They did not report any OH being formed in any of their experiments. In the work described below, we observe OH being formed in the system CH2I + O2 and show how this can be used to follow reactions of CH2OO with other species.

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EXPERIMENTAL DETAILS The chemical reactions were initiated by photolyzing a dilute mixture of CH2I2 and O2 in argon using an excimer laser (Lambda-Physik LPX 120i) operating at 351 nm with a repetition rate of 2−10 Hz. The excimer beam entering the reaction chamber had a diameter of 0.5 cm and a energy of ∼0.6 mJ/pulse. A dye laser beam intersected the excimer laser beam at right angles and was used to detect OH by laserinduced fluorescence.23 The dye laser (Sirah, pumped by Spectra-Physics YHP-40) operated with Rhodamine 6G at 20 kHz. The dye laser beam at the center of the chamber had a diameter of 0.05 cm and an energy of about 200 μW, corresponding to an individual pulse energy of 10 nJ. The dye laser was tuned to 282.0 nm in order to pump the P1(1) line of Received: July 16, 2013 Revised: January 6, 2014 Published: January 8, 2014

(R.1a) 741

dx.doi.org/10.1021/jp407058b | J. Phys. Chem. A 2014, 118, 741−747

The Journal of Physical Chemistry A

Article

the (1,0) band of the OH (A2Σ+ ← X2Π) transition. The OH spectrum was obtained every day before taking data by tuning the dye laser, to ensure that the signal was indeed OH. When off resonance, the background signal was only a few counts s−1. The dye laser was adjusted frequently during the experiment to optimize the OH signal. For the total pressures used, vibrational relaxation of OH(A2Σ+ v = 1) was rapid,24 so OH fluorescence occurred primarily in the (0,0) band at 308 nm. The fluorescence was viewed in a direction perpendicular to both laser beams; a fast quartz lens collected the fluorescence and directed it through a 308 nm interference filter and into a photomultiplier tube (Electron Tubes P25PC-02). Pulses from the PMT were collected by a multichannel scaler (Ortec MCSPCI) as a function of time after the excimer laser pulse. Normally 1000−5000 sweeps were added giving total integration times of 50−250 ms for each 50 μs time bin. Gases flowed through the photolysis cell continuously. Flows of the main carrier gases, argon (Matheson 99.999%) and oxygen (Air Products 99.9995%), were determined by flow controllers (MKS). A fraction of the argon flow was diverted through a bubbler containing liquid CH2I2 (Sigma-Aldrich, 99%). The bubbler was kept at 22 ± 2 °C, at which temperature the vapor pressure of CH2I2 was estimated to be 1.1 Torr.25 Mixtures of other additives, hexafluoroacetone (HFA, Sigma-Aldrich 97%) and SO2 (Matheson 99.98%), were made up in 20 L bulbs diluted ∼1:1000 with argon. All gas flows entered a manifold 50 cm upstream of the photolysis chamber to allow sufficient time for mixing. The pressure in the photolysis cell was measured with a capacitance manometer (MKS Baratron) with an automatic control system (MKS 0248A) used to keep the pressure constant. The temperature near the middle of the photolysis chamber was monitored by a thermocouple and was stable at 22 ± 2 °C. Gas concentrations were calculated using the measured flow rates, the total pressure, and the temperature in the cell.

Figure 1. LIF signals for OH as a function of time after the excimer laser firing for three different oxygen concentrations. From highest to lowest, [O2] = 16.1 × 1014, 8.1 × 1014 and 3.3 × 1014 molecules cm−3. Total pressure was 50 Torr, and [CH2I2] = 3.85 × 1014 molecules cm−3.

Here capital K’s are used to denote various rates, always with units of s−1, while lowercase k’s will be reserved for specific rate constants. The sum of the two terms C1 and C3 gives the amplitude of the fast OH that appears immediately after the flash. The amplitude C2 gives the strength of the slower forming OH signal, which grows at a rate governed by K2. The constant K1 is most evident in the fall of the signal after the maximum, but this exponential decay term applies to both the C1 and C2 signals. The last term in eq 1 was needed to represent the signal at longer times; the origin of this slowly decaying signal is not understood at the present time. The fitted parameters for the three runs of Figure 1 are collected in Table 1. The rate of growth of the slow OH signal is clearly dependent on the oxygen concentration. As can be seen in Figure 1, higher oxygen concentrations increase the rate of growth of the slow OH signal. The fitted constant K2 increases with [O2] at low concentrations, at approximately the rate expected for reaction R.1a ((1.4−1.67) × 10−12 cm3 molecule−1 s−1),26−28 but then increases less rapidly at higher concentrations of oxygen as shown in Figure 2. The difference between the observed values of K2 and an extension of the straight line in Figure 2 must be related to the rate of an intermediate reaction or reactions that are responsible for the formation of the slow OH. The fraction of the OH signal that occurs at very short times, the fast OH, also depends on the oxygen concentration as shown in Figure 3. The derivation of the green line in Figure 3 will be delayed to the Discussion. The dominant loss process for OH radicals formed in this system should be R.2.

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RESULTS When CH2I2/O2 mixtures were photolyzed at 351 nm, the OH radical was formed by at least two different processes, as distinguished by their time dependence. As can be seen in Figure 1, strong signals of OH are evident in the earliest useful channel, from 50 to 100 μs. (The first counting interval of 0− 50 μs picked up scattered light from the excimer flash and so is excluded in Figure 1 and in subsequent analyses.) We suspect that this rapid formation of OH is caused by energy-rich CH2I radicals reacting with O2 (see Discussion). Additional OH is formed in a slower process that results in a maximum OH signal in the region of 1−5 ms, depending on conditions. This slower process is thought to reflect the formation of the Criegee intermediate in the reaction of thermalized CH2I radicals with O2. After the maximum, there is a slow decay of the signal out to 100 ms, at which time the excimer laser fires again and a new sweep is initiated. Changing the repetition rate of the excimer from 10 to 5 or 2 Hz, allowing more time for the slow decay, did not significantly alter the very early OH signals. The empirical eq 1 was used to fit the LIF OH signals as a function of the time after the excimer flash using the method of least-squares.

OH + CH 2I 2 → H 2O + CHI 2

For the conditions of Figure 1 and a rate constant of (R.2) of (4.3 ± 0.5) × 10−12 cm3 molecule−1 s−1,29 the rate of loss of OH should be about 1656 s−1. However, the observed decay rate of OH after the maximum, as measured by K1, is about 120 s−1. This suggests that the observed decay rate represents mainly the decay of the precursor to OH and that the OH is being generated fast enough to almost balance its rapid loss caused by (R.2) (i.e., there is an approximate steady state). If

signal = C1{exp(−K1t )} + C2{exp(−K1t ) − exp(−K 2t )} + C3{exp( −K3t )}

(R.2)

(1) 742



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Table 1. Examples of Parameters Resulting from Fitting Runs Shown in Figures 1 and 4 to Eq 1a Figure Figure Figure Figure Figure Figure a

1 1 1 4 4 4

C1

C2

C3

K1

K2

K3

374 302 179 142 269 105

1688 1749 2018 3088 2707 3073

367 364 308 223 49.4 14.4

121 120 127 189 350 1117

846 605 321 955 1143 2236

8.0 8.4 9.0 21.9 16.5 9.7

[O2] 1.61 8.1 3.3 1.63 1.63 1.63

× × × × × ×

1015 1014 1014 1015 1015 1015

[CH2I2]

[HFA]

× × × × × ×

0 0 0 0 1.19 × 1013 2.78 × 1013

3.85 3.85 3.85 4.26 4.26 4.26

1014 1014 1014 1014 1014 1014

Units: all C, counts/0.05 s; all K, s−1; all [X], molecules cm−3.

The simplest assumption to explain this behavior is that the precursor to the OH is the Criegee intermediate, CH2OO. If this is correct, then the decay rate of the OH signal should be increased as molecules known to react with CH2OO are added. Horie et al.30 were the first to use hexafluoroacetone (HFA) as an effective and clean scavenger for Criegee intermediates. They determined that the intermediate formed during the ozonolysis of ethylene, supposedly CH2OO, reacted 13 times faster with HFA than with acetaldehyde, a known reactant for Criegee intermediates. By directly monitoring the CH2OO, Taatjes et al.1 reported a rate constant for the reaction of HFA with CH2OO of (3.0 ± 0.3) × 10−11 cm3 molecule−1 s−1, confirming the rapid destruction of the simplest Criegee intermediate by HFA. Figure 4 shows the effect of adding HFA in the present experiments for conditions where [CH2I2] and [O2] are Figure 2. Dependence of the rise time constant, K2, on the oxygen concentration for two different days. The straight line shows the expected rate of reactions R.1), based on the average of three rate constants reported in refs 26−28. The [CH2I2] was 3.85 × 1014 molecules cm−3, and the total pressure was 50 Torr.

Figure 4. LIF signals for OH as a function of time for three different concentrations of HFA. The red lines show the least-squares fits of eq 1 to the data. From highest to lowest, [HFA] = 0, 1.19 × 1013, and 2.78 × 1013 molecules cm−3. Total pressure 50 Torr, [O2] = 1.6 × 1015 molecules cm−3, and [CH2I2] = 4.0 × 1014 molecules cm−3.

constant. The decay rate of the OH signal (K1) increases rapidly as HFA is added while the formation rate (K2) stays approximately constant. Separate experiments with other sources of OH showed no effect on its decay rate with much larger additions of HFA. Since OH is not reactive with HFA, the increased decay rate seen in Figure 4 is assumed to be due to the reaction of HFA with the OH precursor, CH2OO. When the fitted values of K1 are plotted against the HFA concentration, straight lines result, as can be seen in Figure 5. Seven runs using two different HFA mixtures were made on four different days. Each run involved determining the decay constant K1 at each of 8−12 HFA concentrations. The slopes of plots of K1 vs [HFA] were taken to represent the rate constant for HFA reacting with CH2OO; these values are collected in Table 2. The average of these rate constants, weighted by the inverse square of their sample standard deviations, is 3.33 × 10−11 cm3 molecule−1 s−1. Using the Student’s t distribution,

Figure 3. Amount of the fast OH signal (present at very short times) as a fraction of the total OH signal as a function of [O2]. The total pressure was 50 Torr. Derivation of the green line is presented in the Discussion section.

this is correct, then the observed rate of decay of the OH signal can be used as a proxy for the rate of decay of the precursor. In order to avoid significant perturbations to the steady state approximation, the decay rates need to be limited to no greater than about half of the expected loss rate of OH due to (R.2). 743



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Figure 5. Fitted decay constant K1 for different concentrations of HFA. The slope of the least-squares line is assumed to be the rate constant for the reaction of CH2OO with HFA. Concentrations in molecules cm−3: [O2] = 1.6 × 1015; [CH2I2] = 3.85 × 1014. Total pressure 50 Torr. Figure 6. Fitted decay constant, K1, for different concentrations of SO2 at a total pressure of 51 Torr. The slope of the least-squares line is assumed to be the rate constant for the reaction of CH2OO with SO2 .

the 95% confidence limits for this average, based on precision only, was ±3.7%. Adding in the uncertainty in the final concentrations of HFA, the combined uncertainty is 8%, or (3.33 ± 0.27) × 10−11 cm3 molecule−1 s−1. This agrees well with the value of (3.0 ± 0.3) × 10−11 reported by Taatjes et al.1 Another molecule known to react with CH2OO is SO2.1,28,31 Theoretical calculations31 suggest that this involves the CH2OO adding to SO2 to give a five-membered ring, as shown in R.3.

Figure 6 shows the increased decay rate of OH as SO2 is added to the system at a total pressure of 51 Torr. The slope of the least-squares line in Figure 6 should give a rate constant for (R.3). Using an error analysis similar to that used for HFA above, the weighted average and combined uncertainty for the rate constant of (R.3) is 8%, or (3.53 ± 0.29) × 10−11 cm3 molecule−1 s−1, in good agreement with a previous measurement of (3.9 ± 0.9) × 10−11 cm3 molecule−1 s−1 made by Welz et al.1 at 4 Torr. The measurements using SO2 were made at four different pressures. Figure 7 shows the present results plotted against total pressure. No pressure dependence is evident over this pressure range. Interference by the OH + SO2 reaction

Figure 7. Measured rate constants for the reaction of CH2OO with SO2 as a function of the total pressure. The black points are from Table 2, and the red point at 4 Torr is from ref 2. The error bars are estimated 95% confidence limits based on the precision of the slopes only.

Table 2. Values of Rate Constants (k) and Their Sample Standard Deviations (s), Both in Units of cm3 molecule−1 s−1, as a Function of Conditionsa

a

reactant

range/1012

[O2]/1014

[CH2I2]/1014

press./Torr

k/10−11

s/10−11

HFA HFA HFA HFA HFA HFA HFA SO2 SO2 SO2 SO2 SO2

0−4 0−4 0−20 0−20 0−16 0−16 0−32 0−9 0−4 0−4 0−4 0−4

8.2 4.1 8.3 16.0 16.0 8.2 8.2 15.0 15.0 15.0 15.0 15.0

4.1 4.1 4.0 4.0 4.0 4.0 4.2 2.2 2.2 2.2 2.2 2.2

50 50 50 50 50 50 50 50 51 100 152 200

3.02 3.38 3.53 3.65 3.25 3.28 3.17 3.62 3.64 3.37 3.43 3.44

0.16 0.21 0.11 0.18 0.10 0.10 0.22 0.09 0.10 0.25 0.09 0.15

Concentrations are given in molecules cm−3. 744



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DISCUSSION The fast formation of OH, as measured by the parameters (C1 + C3), was unexpected. When CH2I2 is photolyzed by a 351 nm photon, there is 30 kcal mol−1 of excess energy over that needed to break the carbon−iodine bond, assuming that the iodine atom is formed in its ground state. The parting fragments (I and CH2I) are both heavy so only a small fraction of this excess will be lost as kinetic energy. Using an impulsive model on a repulsive upper electronic surface, one can calculate that 84% of the excess energy, about 25 kcal/mol, will be retained by the CH2I radical.30 Baughcum and Leone33 and also Lenzer et al.34 have observed such energy-rich CH2I radicals when photolyzing CH2I2 at 248, 303, and 308 nm. So at least some of the CH2I radicals formed in the 351 nm photolysis will be vibrationally hot also. Baughcum and Leone studied the deactivation of the hot CH2I radicals, formed by 248 nm photolysis, by monitoring their infrared emission. They found deactivation of specific vibrational modes of CH2I * by CH2I2 to have rate constants in the neighborhood of 2 × 10−11 cm3 molecule−1 s−1. They estimated that argon was about 14 times slower, which would result in specific relaxation rate constants of about 1.4 × 10−12 cm3 molecule−1 s−1 for argon. The fraction of the observed OH that is fast OH, (C1 + C3)/ (C1 + C2 + C3), is shown in Figure 3 as a function of the oxygen concentration. This can be interpreted as a competition between a very rapid reaction of the vibrationally excited CH2I* with O2 to form CH2OO (R.7) and vibrational relaxation of CH2I* by the bath gas to give thermal CH2I (R.8).

contributes only about 1.5% to the observed loss rates, even at 200 Torr. The yield of OH radicals in the CH2I2/O2 system was estimated by first calibrating the OH detection efficiency using the 351 nm photolysis of molecular chlorine in the presence of methanol (7.8 × 1014 molecules cm−3), oxygen (3 × 1016 molecules cm−3), and nitric oxide (1.2 × 1015 molecules cm−3).32 The following reaction sequence is responsible for generating OH in this system: Cl + CH3OH → CH 2OH + HCl

(R.4)

CH 2OH + O2 → HO2 + CH 2O

(R.5)

HO2 + NO → OH + NO

(R.6)

Article

Modeling shows that almost every Cl atom results in the formation of an OH, but the OH then rapidly reacts with CH3OH and with NO. As a result, the peak OH concentration is only 71−73% of the initial [Cl]. The model predicts that the maximum OH signal should occur at 0.00025 s and that the signal then decays at a rate of about 1000 s−1; both predictions are closely matched by experiment. We estimate that the maximum [OH] in this system was about 1 × 1011 molecules cm−3, which is about a factor of 10 greater than the maximum [OH] in the CH2I2 experiments. Strong LIF signals were recorded as a function of time in the same manner as used for the CH2I2 experiments. On the same day, with the lasers running at the same power, measurements of the OH LIF were made on the CH2I2/O2 system using CH2I2 concentrations that varied by a factor of 5. The total amounts of OH formed in both the Cl2/CH3OH and CH2I2/O2 systems were calculated by multiplying the signal accumulated in each bin by the appropriate OH loss rate times 50 μs and then summing the losses over all bins. The loss rate in the CH2I2/O2 system was taken to be k2[CH2I2]. For the Cl2/CH3OH system the observed exponential decay rate of the LIF signal after its maximum, about 1040 s−1, was used as the loss rate. Then the total OH losses for each system were divided by the expected radical generation for each, σCH2I2[CH2I2]I0 for the Criegee system and 2σCl2[Cl2]I0 for the Cl2/CH3OH system, where I0 represents the 351 nm laser intensity. The resulting normalized losses for the Cl2/CH3OH system were reasonably constant as Cl2 varied by a factor of 4; nine separate measurements had a standard deviation of 12%. The resulting average sensitivity on that particular day was 2.5 × 105 (OH molecules cm−3)/(count s−1). However, a similar calculation for the CH2I2/O2 system gave normalized losses that differed by a factor of 2 as the CH2I2 concentration varied by a factor of 5. The cause of this variation is not known. The ratio of the average of these two normalized OH losses, taken on the same day with the same laser intensities, allows the laser intensities to cancel out and resulted in an average OH yield in the CH2I2/O2 system of 0.21 OH radicals generated per CH2I radical. Using the average of the Cl2/CH3OH runs with the eight individual CH2I2/O2 runs resulted in OH yields from 0.13 to 0.28. However, the uncertainty in all of these yields is probably at least a factor of 1.5 due to the uncertainty in the vapor pressure of liquid CH2I2; a value of 1.1 Torr at 22 °C, based on extrapolation of literature values,24 was used in calculating the above yields.

CH 2I* + O2 → CH 2OO + I

(R.7)

CH 2I* + M → CH 2I + M

(R.8)

The CH2OO formed in (R.7) will probably retain some of the excess energy contributed by the CH2I*, but vibrational relaxation should be rapid for a molecule of this size. The thermal CH2OO thus formed may then proceed to generate OH, the “fast OH”, by the same mechanism that forms the slow OH. Let F be the fraction of the newly formed CH2I* radicals capable of forming fast OH and (1 − F) the fraction of newly formed CH2I that are formed with insufficient energy to give fast OH but that do go on to form slow OH. Then the fraction of the total observed OH that is fast OH would be (C1 + C3)/(C1 + C2 + C3) = Fk 7[O2 ]/(k 7[O2 ] + k 8[M])

(2)

where k7 and k8 are rate constants for (R.7) and (R.8). Inverting both sides of eq 2 gives a linear relationship in [O2]−1: (C1 + C2 + C3)/(C1 + C3) = (1 + (k 8[M]/k 7[O2 ]))/F (3)

The data shown in Figure 3 have been replotted according to eq 3, as shown in Figure 7, and the linear least-squares line gives F = 0.38 and the ratio k8[M]/k7 = 3.1 × 1014 cm3 molecule−1. These fitted parameters have been used to plot the green lines in both Figure 3 and Figure 8. Using the value of [Ar] appropriate for the data in Figure 3, the ratio of k8/k7 is 2 × 10−4. For (R.7) to be 80% complete by 10−4 s after the photolysis pulse, i.e., by the first recorded bin, would require k7 ≥ 5 × 10−11 cm3 molecule−1 s−1 for the lowest 745



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decomposition to dioxirane would still proceed at about 1010 s−1. Since the dioxirane is about 25 kcal/mol more stable than the Criegee and it has just come over a 20 kcal/mol barrier, there should be at least 45 kcal/mol of excess energy in the newly formed dioxirane. Unless removed by vibrational relaxation, this will be more than enough to surmount the barriers to form first the bis-oxy biradical and then formic acid. The newly formed formic acid would have almost 137 kcal/mol of excess energy, more than the 110 kcal/mol needed to form OH + CHO by a simple bond fission. This is a possible path for the formation of the fast OH in the current experiments. The yield of OH from hot formic acid is expected to be less than unity because of several competing reactions, giving H + HCO2, H2 + CO2 or H2O + CO, as well as collisional stabilization to formic acid. The formation of the slow OH may arise from the same reaction sequence as outlined above, but with the energy necessary to surmount the barrier provided by thermal fluctuations. Olzmann et al.39 calculated an average lifetime of 3 s for a thermally equilibrated CH2OO at 298 K. Supposedly this slow decay also involves mounting the 20 kcal barrier to generate dioxirane. The slowest loss rates (K1) observed in the current experiments were 80−100 s−1, corresponding to lifetimes of about 8 ms. Recent experiments that observed CH2OO directly found maximum lifetimes of about 2 ms.1 Both of these observed lifetimes may be influenced by reactions of the Criegee with other gas phase species or on surfaces. Also, the exothermicity available in this reacting system may elevate the vibrational “temperature” and contribute to the short Criegee lifetimes. Work is continuing to identify the specific reactions that are responsible for OH generation in this system. Also, the technique is being used to measure rate constants for reactions of CH2OO with other molecules of atmospheric interest.

Figure 8. Data shown in Figure 3 have been plotted according to eq 3. The least-squares line has an intercept of 2.65 and a slope of 8.2 × 1014 molecules cm−3.

oxygen concentration shown in Figure 3. Using this limit for k7, the rate constant for k8 would be ≥1 × 10−14, about 140 times smaller than was estimated above for the higher energy CH2I * formed by 248 nm photolysis. A slower relaxation rate for the less energetic CH2I* formed by 351 nm photolysis is expected due to the lower density of vibrational states, resulting in less efficient energy transfer to argon. This calculation is not intended to determine a relaxation rate for CH2I*, but just to show that the above model gives plausible values. Early evidence for the existence of a CH2O2 molecule in the gas phase was presented in 1977 by work at the National Bureau of Standards. Lovas and Suenram35 condensed ozone and ethylene at 77 K and then observed the microwave spectrum of dioxirane in the gas phase as the mixture was warmed. Isotope studies confirmed a structure with the two oxygen atoms forming a symmetrical three-member ring with the carbon.35 Parallel work by Martinez et al.,36 using the same generation technique, detected dioxirane mass spectrometrically. It was possible to distinguish the mass 46 dioxirane peak from the mass 46 formic acid peak because they appeared at different times during the warm-up. These early results were interpreted as support for the previous calculations of Wadt and Goddard,37 who had proposed that the decomposition of CH2OO in the gas phase would initially form the more stable dioxirane. Subsequent rearrangements would lead to the bis-oxy biradical, with the oxygen−oxygen bond in dioxirane broken, and then finally to formic acid. Several later ab initio calculations have confirmed this general picture, finding an approximately 20 kcal/mol barrier for CH2OO to form dioxirane and a higher barrier of about 30 kcal/mol for CH2OO to form OH + CHO directly (see for example the recent review by Vereecken and Francisco38). Olzmann et al.39 have used statistical rate theory and a master equation approach to calculate the rates of CH2OO decomposition to either dioxirane or to OH + CHO. If the CH2OO were to retain all of the excess energy of the CH2I* radical, about 25 kcal/mol, they calculate a rate of formation of dioxirane of 1011 s−1. Even if as much as 6 kcal/mol was lost to vibrational relaxation in CH2I* or by the leaving I atom, the

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AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected] (K.D.B.). Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS We thank the two reviewers of this manuscript for their careful and thorough analysis. They have made significant contributions to our understanding of this system. This research was carried out by the Jet Propulsion Laboratory, California Institute of Technology, under contract with the National Aeronautics and Space Administration (NASA). This work was supported by the Upper Atmosphere Research and Tropospheric Chemistry programs and the JPL Posdoctoral Program. The authors also thank Dave Natzic for his technical assistance.

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REFERENCES

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Measuring Rate Constants for Reactions of the Simplest Criegee

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