14694
J. Phys. Chem. 1996, 100, 14694-14702
Reaction of OH with Dimethyl Sulfide (DMS). 1. Equilibrium Constant for OH + DMS Reaction and the Kinetics of the OH‚DMS + O2 Reaction Stephen B. Barone, Andrew A. Turnipseed, and A. R. Ravishankara* National Oceanic and Atmospheric Administration, EnVironmental Research Laboratories, Aeronomy Laboratory, Boulder, Colorado 80303, and Department of Chemistry and Biochemistry and The CooperatiVe Institute for EnVironmental Sciences, UniVersity of Colorado, Boulder, Colorado 80309 ReceiVed: March 22, 1996; In Final Form: June 18, 1996X
The formation of a weakly bound adduct in the reaction of OH with DMS-d6 was observed between 217 and 240 K using the technique of pulsed laser photolysis/pulsed laser-induced fluorescence. The equilibrium constant for this process, OH + DMS-d6 T OH‚DMS-d6, was measured as a function of temperature. The bond strength of this adduct was determined to be 10.7 ( 2.5 kcal mol-1. The weakly bound adduct was observed to react rapidly with O2. The rate constant for the reaction OH‚DMS-d6 + O2 f products was determined to be (1.00 ( 0.33) × 10-12 cm3 molecules-1 s-1, independent of pressure and temperature. The atmospheric implications of the formation of this adduct and its reaction with O2 to the mechanism of DMS oxidation in the atmosphere are discussed.
Introduction Measurements in the atmosphere suggest that natural sources of organosulfur compounds comprise ∼60% of the total sulfur flux into the marine boundary layer in the southern hemisphere and ∼15% in the northern hemisphere.1 The primary organosulfur compound emitted to the atmosphere is dimethyl sulfide (CH3SCH3, DMS).1 The mechanism of the atmospheric oxidation of DMS is very complex and the branching to final stable end products will depend on the temperature, pressure, and composition of the atmosphere. Quantification of the yields of stable end products is vital in evaluating the role of DMS in climate regulation2 and in understanding the biogeochemical cycling of sulfur. The atmospheric oxidation of DMS has been the subject of many investigations in the past few decades and has been reviewed recently.3-5 The reaction of OH with DMS is currently believed to be the most important process that initiates DMS oxidation in the daytime. Hynes et al.6 observed that the loss rate of OH in the presence of DMS increased significantly upon the addition of O2. They proposed the following mechanism for reaction 1 to account for their observations:
OH + DMS T OH‚DMS OH + DMS f H2O + CH3SCH2
(1f, 1r) (1a)
The symbol “T” indicates the combination of forward and reverse reactions. The enhancement in the OH loss rate in the presence of O2 was explained by the occurrence of a reaction between the OH‚DMS adduct and O2:
OH‚DMS + O2 f products
(2)
In the absence of O2, the adduct would decompose rapidly to re-form OH and, hence, channels 1f and 1r would not be observable at long reaction times (i.e., after equilibration). Chamber studies, by Barnes et al.7 and Wallington et al.8, have confirmed the pronounced dependence of the rate coefficient for removal of DMS on the partial pressure of O2. * Address correspondence to this author at NOAA/ERL, R/E/AL2, 325 Broadway, Boulder, CO 80303. X Abstract published in AdVance ACS Abstracts, August 1, 1996.
S0022-3654(96)00866-0 CCC: $12.00
An electronic ab-initio structure calculation of McKee9 yielded a bound geometry for the OH‚DMS adduct with a bond strength of 6.0 kcal mol-1. Contrary to these results, theoretical calculations and experiments by Gu and Turecek10,11 suggested that OH‚DMS was not a stable intermediate but a transition state in reaction 1. The present work was initiated to investigate the existence of the OH‚DMS species and quantify the rate coefficients involved in its formation, decomposition, and reaction with O2. The accompanying paper12 describes our attempts to identify the products and measure their branching ratios in the OH reaction with DMS. Temperature and O2 concentrations were varied to obtain data relevant to modeling DMS oxidation in the atmosphere. In combination, the two papers present a description of the first several steps in the OH-initiated oxidation of DMS. In this paper we compare our results on the OH + DMS reaction kinetics with those from previous studies and briefly discuss their importance in the atmospheric oxidation of DMS. A more detailed discussion of the atmospheric consequences of these findings will be presented in the companion paper. Experimental Section All experiments utilized pulsed laser photolysis for generating OH and pulsed laser-induced fluorescence (LIF) for detecting it. The apparatus used is identical to that previously employed to detect CH3S13 and CF3O14 and was modified for the detection of OH. The detection of OH via LIF is routine in our laboratory. Therefore, we relate here only the aspects essential to the understanding of the present study. The reaction vessel was a ∼200-cm3 jacketed Pyrex cell with six ports: four were 5-cm long with a 1-in. diameter for transmitting of the probe and photolysis laser beams, one was an inlet for inserting a movable thermocouple, and one was for collecting the laser-induced fluorescence. A series of light baffles were used to minimize scattered probe laser radiation in the cell. The reactor temperature was regulated over the range 217-359 K by flowing either cooled methanol or heated ethylene glycol from a temperature-regulated bath through the jacket of the cell. The temperature of the gases within the reaction volume (defined by the intersection of the photolysis © 1996 American Chemical Society
Reaction of OH with DMS. 1 and probe beams) was measured by inserting a calibrated chromel-alumel thermocouple through a movable injector under gas flow conditions identical to those during the experiment. For experiments using He as the bath gas, we conservatively estimate the uncertainty in the reactor temperature to be (0.5 K because there was no measurable temperature gradient within the reaction volume. However, a measurable temperature gradient was observed when SF6 and N2 were used as bath gases. We conservatively estimate the reaction temperatures to be accurate to ∼(2 K under these conditions. OH radicals were generated by pulsed excimer laser photolysis of either H2O2 at 248 nm (KrF, 5-30 mJ pulse-1 cm-2, pulse width ∼20 ns, 10 Hz) or HONO at 351 nm (XeF, 10-30 mJ pulse-1 cm-2, pulse width ∼20 ns, 10 Hz) on an axis perpendicular to the direction of the gas flow. The photolysis beam was passed through a 5-cm side arm and reduced to a diameter of 0.80 cm by an aperture before traversing through the cell and exiting via another 5-cm side arm. The fluence of the photolysis beam was measured by a disk calorimeter after the beam exited the cell. H2O2 was introduced into the cell by passing a small flow of helium through a glass bubbler containing H2O2. A 90% H2O2 solution was purified further by bubbling He through it for several days prior to use. HONO was prepared, on-line, in a flask by adding 0.1 M NaNO2 solution dropwise into a solution of 10% H2SO4 and swept out of the flask directly into the reactor by a small flow of helium. OH free radical concentrations were monitored by exciting the (A2Σ+, V′ ) 1) r (X2Π, V′′ ) 0) electronic transition at ∼281.9 nm (0.10-1.2 mJ pulse-1 cm-2, pulse width ∼8 ns, 10 Hz) from a dye laser pumped by a pulsed XeCl excimer laser. The excitation beam was orthogonal to the photolysis beam. The fluence of this pulse was measured after the beam traversed through the reactor using a disc calorimeter to monitor its stability. The red-shifted fluorescence, resulting from the (A2Σ+, V′ ) 1) f (X2Π, V′′ ) 1) and (A2Σ+, V′ ) 0) f (X2Π, V′′ ) 0) transitions, was collected on an axis orthogonal to both the photolysis and probe laser beams by a 5-cm focal length UV quartz lens. The fluorescence was passed through an aperture and band-pass filter (λ ) 308.7 nm, with a FWHM of 5.9 nm) before being imaged onto a slit in front of the photomultiplier tube (PMT). The output pulse from the PMT was fed into a gated charge integrator and then to a microcomputer, via an A/D converter, for data analysis. The temporal profiles of the OH radical concentrations were acquired by measuring the LIF signal at various delay times between the photolysis and probe lasers (5 µs-50 ms). The nascent rotational and vibrational energy distributions in the OH radical produced by photolysis of H2O2 and HONO are nonthermal.15,16 The use of an excess of buffer gas rapidly thermalized these rotational and vibrational energy distributions to the ambient temperature (before our earliest observation times, t < 5 µs). The detection limit for OH, defined as S/N ) 1, where S is the time zero signal and N is equal to twice the standard deviation in the mean value of the measured background, was 4 × 108 molecules cm-3 in 100 Torr of He upon integration of 100 laser shots. The photolyte, bath gas, and reactants were slowly flowed (linear flow velocities of 2-10 cm s-1) through the reactor perpendicular to the direction of the photolysis beam. He (Bureau of Mines, g99.9997%), N2 (Scott Gases, g99.9995%), O2 (Scott Gases, g99.99%), and SF6 (Scott Gases, g99.9%) were added to the reactor through calibrated electronic mass flow meters. O2 concentrations were determined from the measured mass flow rates and the cell pressure, measured by a capacitance manometer. Concentrations of DMS-d6 (CD3SCD3, >99.9%) and DMS (99.5%) were monitored on line by UV
J. Phys. Chem., Vol. 100, No. 35, 1996 14695 absorption at 213.9 nm (zinc line) in a 100-cm long absorption cell [σ213.9(DMS-d6) ) 1.23 × 10-18 cm2 molecule-1, measured in this study, and σ213.9(DMS) ) 1.70 × 10-18 cm2 molecule-1 from Hearn et al.17]. Although not essential for the kinetic measurements, knowledge of the photolyte concentration is useful in evaluating the potential role of secondary reactions. H2O2 concentration was monitored on line by absorption at 213.9 nm prior to entering the cell [σ213.9nm(H2O2) ) 3.16 × 10-19 cm2 molecule-1].18 HONO concentrations were not monitored on line directly but were inferred from the measured rate coefficient for OH radical loss due to the OH + HONO reaction, in the absence of DMS-d6 or DMS. Because it enables us to easily separate equilibration of OH with DMS from the irreversible loss of OH, we used DMS-d6, rather than DMS, as the reactant in most of our experiments:
OH + DMS-d6 f products
(3)
In the present experiments, OH loss is governed by both its reversible addition to DMS and its irreversible reaction with DMS. If the time scales for these two processes are different, their rate coefficients can be unambiguously extracted from the measured OH temporal profiles. The rate coefficient for the irreversible reaction of OH with DMS is approximately a factor of 3 faster than the that of analogous reaction with DMS-d6.6 Therefore, we used DMS-d6 in place of DMS. Use of DMS-d6 has an added advantage in that it eliminates any possible production of OH radicals from secondary reactions involving the methyl groups in DMS. Results OH + DMS T OH‚DMS (k1f,r). The reaction of OH with DMS can proceed via several channels:
OH + DMS (+ M) T CH3S(OH)CH3 (+ M) ∆°H ) (determined here) (1f, 1r) OH + DMS f CH3SCH2 + H2O ∆°H ) -25.5 kcal mol-1 (1a) OH + DMS f CH3 + CH3SOH ∆°H ) 0 ( 3 kcal mol-1 (1b) OH + DMS f CH3S + CH3OH ∆°H ) -18.8 kcal mol-1 (1c) The heats of formation of the reactants and products of reactions 1a-1c are taken from the literature.19,18 Reaction 3 (the reaction of OH with the deuterated analog of DMS) has the same possible product channels as reaction 1 with approximately the same reaction enthalpies. Initially, the overall rate coefficients for reaction 1 and its deuterated analog (reaction 3) were measured at 298 K in the absence of O2 under pseudo-first-order conditions in [OH] (i.e., [DMS] g 100[OH]). The postphotolysis temporal profiles of OH at 298 K strictly obeyed the equation
[OH]t ) [OH]0 exp(-k′it)
(4)
In this expression [OH]t and [OH]0 are, respectively, the concentration of OH radicals at times t and 0 (i.e., after it is generated but before any reactive loss occurred). k′i represents the pseudo-first-order decay rate constant for OH loss and is related to the second-order rate coefficient (ki) by the expression
k′i ) ki[DMS] + kd
(5)
14696 J. Phys. Chem., Vol. 100, No. 35, 1996
Barone et al. profiles in the presence of DMS-d6 were not strictly single exponential; they were characterized by an initial rapid decay followed by a slower exponential decay. The rate of the initial decay and the point of inflection were observed to vary with the concentration of DMS-d6. This behavior is indicative of the occurrence of an equilibrium between OH and DMS-d6 in combination with a slower first-order loss process for OH. The observed temporal profiles are consistent with the following series of reactions:
mechanism 1 OH + DMS-d6 T OH‚DMS-d6
(3f, 3r)
OH + DMS-d6 f products other than OH‚DMS-d6 (3bi)
Figure 1. Examples of temporal profiles of [OH] measured at 230 K in 100 Torr (25 Torr of N2 and 75 Torr of He): (A) [DMS-d6] ) 1.70 × 1015 molecule cm-3; (B) [DMS-d6] ) 4.38 × 1015 molecule cm-3. The initial concentrations were the same for both profiles. The lines are the nonlinear least-squares fits of the data to mechanism 1.
Figure 2. Plot of Kp (on a log scale) vs 1000/T (van’t Hoff plot) for reaction 3f,r in different bath gases [He (O), SF6 (4), and N2 (0)]. H2O2 photolysis at 248 nm was used in most cases; experiments with HONO photolysis are indicated by ]. The error bars are the 2σ precision obtained from the nonlinear least-squares fits (see text) of the data. The solid line is the weighted linear least squares fit to the equation ln Kp ) -∆°Hr/RT + ∆°Sr/R (second-law analysis). The dashed line is the weighted least-squares fit of the data using ∆°Sr calculated from statistical mechanics (third-law analysis).
where kd is the first-order OH loss rate coefficient in the absence of DMS (or DMS-d6). By measuring k′i at various DMS (or DMS-d6) concentrations, the rate coefficients for reactions 1 and 3 at 298 K were measured to be k1 ) (4.95 ( 0.35) × 10-12 cm3 molecule-1 s-1 and k3 ) (1.75 ( 0.25) × 10-12 cm3 molecule-1 s-1 at 298 K. The reported uncertainties are the 2σ errors in fitting the measured k′i vs [DMS] data to eq 5 using a weighted linear least-squares routine. The weighting was according to the uncertainties in the measured values of k′i. The obtained values are in good agreement with other recent measurements.3,6 Temporal profiles of [OH] were measured at various temperatures. As shown in Figure 1, at T e 240 K, OH temporal
OH f loss
(6)
OH‚DMS-d6 f loss
(7)
3bi ) 3a + 3b + 3c. We fit the obtained OH temporal profiles to mechanism 1 using a nonlinear least-squares fitting routine, FACSIMILE.20 The fits were constrained by using the measured first-order rate coefficient for the loss of OH in the absence of DMS, k6, and assuming that k7 was 245 K, equilibration was completed before 5 µs. We did not utilize reaction times 260 K), where the bimolecular rate coefficient was directly measured (i.e., under conditions where the adduct formation was negligible). As shown in the figure, a slight (∼20%) deviation from the Arrhenius behavior was observed for temperatures much lower than 260 K. When these points are not included, our data yield k3bi ) (6.5 ( 1.1) × 10-12 exp[-(380 ( 60)/T] cm3 molecule-1 s-1. The error bars reflect 2σ of the standard deviation of the fit to the Arrhenius expression. The values of k3bi measured using HONO as the OH precursor were slightly higher (∼30%). The deviations in the values of k3bi obtained at T < 260 K are likely due to reactions involving the photolytes and/or impurities in the
Figure 3. Plot of k3bi (k3bi ≡ k3a + k3b + k3c) (on a log scale) vs 1/T. OH was generated by 248 nm photolysis of H2O2. The line is the linear least-squares fit of all the data for T > 250. The obtained values of A and Ea/R are given in the text. k3bi values for T < 250 K deviate from the above fit by ∼20%. Note the highly expanded scale for the rate coefficient in this figure.
reactants or photolytes. Possible explanations of these observations are given under Discussion. We prefer to quote the values obtained at T > 260 K. The 248-nm photolysis of HNO3 is another commonly used OH radical source. By photolyzing a mixture of HNO3 [(112) × 1014 molecules cm-3], DMS-d6 [(1-9) × 1015 molecules cm-3], and He (100 Torr) at 298 K, we determined the rate coefficient for reaction 3bi to be (1.72 ( 0.6) × 10-12 cm3 molecule-1 s-1. This value is in excellent agreement with our measurements of k3 using H2O2 and HONO at 298 K. Attempts were made to observe the addition of OH to DMS-d6 using HNO3 photolysis to generate OH radicals. However, the photolysis of a mixture of HNO3 and DMS-d6 at 229 K resulted in very rapid single-exponential decay profiles of OH with time constants of (1.5-7.0) × 104 s-1. In addition, the OH loss rate constant was observed to increase with increasing HNO3 concentrations. In the absence of DMS-d6, the OH temporal
14698 J. Phys. Chem., Vol. 100, No. 35, 1996
Barone et al. to secondary reactions involving the OH + DMS reaction products) was unrelated to the biexponential portion of the OH temporal profiles, we subtracted the small long time regeneration and fit the residual profiles to mechanism 1 to determine values of k1f, k1r and k1bi. The obtained equilibrium constants are summarized in Table 1 and are, within the combined uncertainties of our measurements, the same as those for the DMS-d6 measurements. O2 + OH‚DMS-d6 f Products (k8). Reaction 8 was initially studied at T < 235 K, at which the equilibration between OH and DMS-d6 was observable and a large fraction of the initial OH could be tied up as OH‚DMS-d6. Under these conditions, the long time (t > 50 µs) loss rate coefficients of OH in the presence of an excess of DMS-d6 increased upon the addition of O2. The OH temporal profiles showed an initial rapid decay due to equilibration followed by a second decay, the rate of which increased with increasing concentrations of O2. This behavior is indicative of a reaction between the OH‚DMS-d6 adduct and O2, which leads to an irreversible loss of the adduct and, hence, OH:
OH‚DMS-d6 + O2 f products Figure 4. Plot of k′9 vs [O2] for the O2 + OH‚DMS-d6 reaction at 228 K and [DMS-d6] ) 5.4 × 1015 molecules cm-3 in approximately 30 Torr of N2. The line is the weighted linear least-squares fit. The error bars are the 95% confidence limits derived from the FACSIMILE simulations (see text).
TABLE 3: Rate Coefficients for Reactions 8 and 2 Obtained in the Present Study along with the Pertinent Experimental Conditions temp, K DMS-d6 222 228 228 229 230 230 233 233 234 234 247 258 258 DMS 247 258
bath gas
pressure, Torr
k, 10-13 cm3 molecule-1 s-1
He N2 N2 He N2 He N2 N2 He He N2 N2 N2
100 100 100 100 30 100 75 80 100 100 100 100 200
8.9 ( 1.4 8.5 ( 2.7 9.1 ( 0.6 7.8 ( 3.6 8.6 ( 1.6 8.8 ( 0.9 11.8 ( 3.0 11.5 ( 2.6 10.5 ( 2.2 8.4 (2.4 10.9 ( 0.7 10.1 ( 0.5 9.9 ( 0.3
N2 N2
100 100
10.9 ( 0.5 13.7 ( 2.7
profiles were exponential and the OH loss rate constant was consistent with the known bimolecular rate coefficient for the reaction of OH with HNO3 at 229 K. Our observations indicate that either HNO3 itself or an impurity (such as NO2) in it reacts efficiently with the OH‚DMS-d6 adduct. Due to the existence of this interfering reaction(s), we were unable to determine the equilibrium constant using HNO3 as the OH source. We also attempted to observe equilibration between OH and DMS in a manner similar to that for OH with DMS-d6. The approach to equilibrium was more difficult to separate from the longer time loss, which was controlled by the bimolecular component of reaction 1. However, at high pressures (P g 100 Torr) and low temperatures (T e 230), the equilibration between OH and DMS (nondeuterated) was indeed observed. Under these conditions, the OH temporal profiles were characterized by biexponential decays superimposed upon a small regeneration of OH at long reaction times. We were unable to eliminate the slight OH regeneration in the OH + DMS system. Assuming that the regeneration of OH at long times (which is likely due
(8)
For the [OH] temporal profile to be sensitive to the occurrence of reaction 8, a significant fraction of OH must be tied up in the OH‚DMS-d6 adduct. This requirement limited measurements of meaningful values of k8 to 220-235 K. Temporal profiles of [OH], in an excess of DMS-d6, were monitored while the total O2 pressure was varied [(0.2-2.0) × 1017 molecules cm-3]. The observed [OH] temporal profiles could be accounted for by the following series of reactions:
mechanism 2 OH + DMS-d6 T OH‚DMS-d6
(3f, 3r)
OH + DMS-d6 f products other than OH‚DMS-d6 (3bi) OH f loss
(6)
OH‚DMS-d6 f loss
(9)
Mechanism 2 is nearly identical to mechanism 1, and here reaction 9 includes contributions from reactions of OH‚DMSd6 with O2. Reaction 9 represents the loss of OH‚DMS-d6 expressed as a first-order process. The first-order rate coefficient for this reaction in an excess of O2 (over OH‚DMS-d6) can be expressed as k′9 ) k8[O2] + k9dif, where k9dif is the first-order rate constant for loss of OH‚DMS-d6 in the absence of O2. Reaction 6 represents the OH loss in the absence of DMS-d6 and O2 and is mostly due to reaction with the photolyte and, to a lesser degree, diffusion out of the reaction zone. First, [OH] temporal profiles were measured in an excess of DMS-d6 in the absence of O2. Values for k3f, k3r, and k3bi were obtained by fitting these profiles to mechanism 1 using a nonlinear least-squares routine (FACSIMILE).20 Next, [OH] temporal profiles were monitored in the presence of an excess of DMS-d6 and O2 while the total pressure was kept the same. By using prescribed values of k3bi and k6, the OH temporal profiles were fit to mechanism 2 (using FACSIMILE) to yield values of k3f, k3r, and k′9. The rate coefficient of reaction 6 could be measured directly by monitoring the OH loss rate constant subsequent to its generation via photolysis of H2O2 in the bath gas alone. k8 was calculated from the slope of the plot of k′9 vs [O2]. Figure 4 shows a typical plot of k′9 vs [O2] obtained at 230 K in 100 Torr of He. The obtained values of k3f, k3r, and Kc were in excellent agreement ((5%) with those measured in the absence of O2. Conversely, fixing k3f and k3r
Reaction of OH with DMS. 1
J. Phys. Chem., Vol. 100, No. 35, 1996 14699
to the values obtained in the absence of O2 resulted in a 5 µs. While a constant total pressure (N2 + O2) was maintained, [O2] was varied over the range 1-100 Torr. All [OH] temporal profiles were exponential for at least three lifetimes, and the first-order rate constant (k′obs) was observed to increase systematically with the partial pressure of O2. Figure 5 displays the measured second-order rate constant,
Figure 5. Plots of kobs [kobs ≡ (k'obs - k6)/[DMS-d6 or DMS]] vs [O2] for experiments at 247 K and 100 Torr (N2 + O2) for DMS and DMSd6. The error bars shown are the 95% confidence limits for each kobs. A significant difference in the intercept and asymptote (A * B) is observed for two cases and interpreted to be the kinetic isotope effect in the “bimolecular” channel.
kobs (kobs ≡ (k′obs - k6)/[DMS-d6]), as a function of [O2] at 247 K and 100 Torr of (N2 + O2). The kobs vs [O2] data were fit to eq 13 by varying the magnitude of k8, k3f, and k3r while keeping k3f/k3r ) Kc. Kc values at these temperatures were obtained via extrapolation of the low-temperature equilibrium constant data. A trial and error procedure was used in determining the best values of k8, k3f, and k3r by minimizing the standard deviation from a nonlinear least-squares fit of the data. Table 3 summarizes the obtained values of k8. The error bars are the 95% confidence limits from the fits to eq 13. Experiments at higher temperatures were not possible because the concentrations of O2 required to adequately characterize the enhancement of kobs due to reaction 8 severely degraded the OH detection sensitivity. Our results indicate that over the range 220-258 K, reaction 8 is essentially temperature independent. Analogous experiments were performed at 247 and 258 K using DMS in place of DMS-d6, which yielded similar plots of kobs vs [O2]. Figure 5 also shows one such plot at 247 K for DMS. Assuming that extrapolations of our equilibrium constant data for reaction 3f,r are valid for the OH + DMS system, the data were fit to eq 13, where k3f, k3r, k3bi, and k8 are replaced by k1f, k1r, k1bi, and k2, respectively. The obtained values of k2 are summarized in Table 3. Similar to the case of the DMSd6, experiments at higher temperatures were impossible because of the degradation of the OH signal at large O2 concentrations. The results show that the rate coefficient for O2 + OH‚DMS is independent of the isotopic composition of DMS. Reactions with very slight temperature dependencies, such as reactions 2 and 8, require determinations over a large temperature range to adequately define their activation energies. Due to the limited temperature range over which reactions 2 and 8 could be studied, we prefer to average all measurements and report k8 ) k2 ) (10.0 ( 3.3) × 10-13 cm3 molecule-1 s-1, independent of temperature. The error bars reflect 2σ of the standard deviation of the mean and include estimates of systematic uncertainty in the [O2], [DMS-d6], and [DMS].
14700 J. Phys. Chem., Vol. 100, No. 35, 1996 Discussion OH + DMS-d6 T OH‚DMS-d6. We have directly observed OH temporal profiles at T e 240 K which show that OH radicals add to DMS-d6 to form a stable, thermalized adduct, OH‚DMSd6. There are several experimental observations which suggest that systematic errors did not influence our measurements of Kc. Kc values were the same with two OH sources (H2O2 and HONO) and did not depend on photolysis fluence or [DMSd6]. Although we did not systematically measure k3f and k3r as functions of temperature and pressure, we did measure them at different temperatures and pressures. It is clear that k3r increases with bath gas pressure and temperature while k3f increases with pressures and decreases only slightly with temperature. These behaviors are consistent with the reaction of OH with DMS being an addition process and the decomposition being hindered by a thermal barrier. Even though k3f and k3r changed with pressure and bath gas identity, their ratio, Kc, did not change. This invariance of Kc suggests thermalization of the adduct before it underwent unimolecular decomposition back to reactants. Our low-temperature values of k3bi, which were measured by producing OH from 248-nm photolysis of H2O2, do not extrapolate to the values obtained at T > 258 K, at which adduct formation is minimal. The temperature dependence of reaction 3bi was calculated using data at T g 258 K. k3bi values measured at T < 250 K are approximately 20% larger than the values obtained by extrapolating the higher temperature data (see Figure 3). The values of k3bi determined in experiments in which HONO was photolyzed to generate OH were slightly larger than those measured using H2O2. There are several possible explanations for the observed departure in kbi from Arrhenius behavior at low T. It is possible that we have not accounted for one or more loss pathways for the OH‚DMS-d6 adduct in our model. The exponential nature of our postequilibrium decays in the OH signal suggests that the loss process, if occurring, is first-order with respect to the OH‚DMS-d6 adduct. (However, within the experimental uncertainty of our OH temporal profiles, a small second-order process cannot be ruled out.) One possible loss process for OH‚DMS-d6 adduct is its reaction with the photolytic precursor; however, the value of k3bi did not change significantly with [H2O2] or [HONO], which were varied from 1 × 1013 to 20 × 1013 molecules cm-3. The presence of an NO2 impurity in our HONO source may, at least partly, account for the higher values of k3bi measured using this OH source, if NO2 reacts very rapidly with the OH‚DMS-d6 adduct. Another possible loss process for the OH‚DMS-d6 is its unimolecular decomposition to products other than OH and DMS-d6. In the next section, based on the observed kinetic isotope effect in reaction 1 in large concentrations of O2, we argue against this possibility. A more likely explanation for this discrepancy is the presence of trace amounts of impurities in our DMS-d6 sample. In past investigations of OH + DMS reaction, the presence of reactive impurities in DMS samples has been a significant source of systematic error.21,22 Likely impurities include CH3SH and dimethyl disulfide (DMDS), which react very rapidly with OH (3 × 10-11 and 2 × 10-10 cm3 molecule-1 s-1 at 298 K, respectively). We used high-purity samples of DMS-d6 (>99.9%) and DMS (99.5%). Even if all of the impurities in our samples were DMDS, its reaction with OH would increase the measured rate coefficients for the approach to equilibrium by
