J . Phys. Chem. 1986, 90, 4148-4156
4148
Kinetics and Mechanism of OH Reactions with Organic Sulfides A. J. Hynes, P. H. Wine,* and D. H. Semmes' Molecular Sciences Branch, Georgia Tech Research Institute, Georgia Institute of Technology, Atlanta, Georgia 30332 (Received: January 16, 1986)
The conventional flash photolysis-resonance fluorescence technique was employed to study reactions of OH with CH3SCH3 ( l ) , CD3SCD3(2), CH3SC2Hs(3), and C2HsSC2Hs(4) in argon buffer gas. Reactivity trends, temperature dependencies, and isotope effects suggest that hydrogen abstraction is the dominant observed reaction pathway under these conditions. A pulsed-laser photolysis-pulsed-laser-induced fluorescence technique was employed to study reactions 1 and 2 in N2,air, and O2 buffer gases. Complex kinetics were observed in the presence of 02.A four-step mechanism involving hydrogen abstraction, reversible addition to the sulfur atom, and scavenging of the (thermalized) adduct by O2 is required to explain all experimental observations. In 1 atm of air, the effective bimolecular rate constant for reaction 1 decreases monotonically from 1.58 X lo-" to 5.2 X cm3 molecule-' s-I over the lower tropospheric temperature range 250-310 K. Over the same temperature range the branching ratio for hydrogen abstraction increases monotonically from 0.24 to 0.87. At 261 K, the rate constant for unimolecular decomposition of the CD3S(OH)CD3adduct is (3.5 f 2.0) X lo6 s-I and the rate constant cm3 molecule-I s-l. for the adduct reaction with O2 is (4.2 f 2.2) X
Introduction The atmospheric sulfur cycle has been the subject of intensive investigation in recent years because of the need to assess the contribution of anthropogenically p r o d u d sulfur to such problems as acid rain, visibility reduction, and climate modification. In heavily industrialized regions such as the eastern United States and western Europe, anthropogenic sulfur emissions exceed natural emissions by about an order of magnitude.'" On a global scale, however, natural sulfur emissions are thought to approximately equal those from anthropogenic source^.^-^ An understanding of the natural sulfur cycle is thus required in oder to establish a "base line" with which anthropogenic perturbations can be compared. Biological activity is generally believed to be a major natural source of reduced Measurements indicate that the predominant reduced sulfur compound entering the atmosphere from the oceans is CH3SCH3 (dimethyl sulfide, DMS).'-13 Hence, DMS plays an important role in the global sulfur cycle. Majot atmospheric sinks for DMS are thought to be reaction with OH and NO3. In marine environments, NO3 levels are typically low and, as a result, DMS is expected to be destroyed primarily by OH. A number of kinetics studies of reaction 1 are reported in the literature.'"22 In addition, several steady-state photolysis-end product analysis studies have recently been O H + CH3SCH3 CH3SCH2 + H2O (la)
-
M +
CH,S(OH)CH3
reported where no rate constant information was obtained but conclusions were drawn concerning the relative importance of hydrogen abstraction and addition to the sulfur atom as reaction pathway^.^^-^^ Despite the rather large data base, neither the rate constant nor the branching ratio for reaction 1 is welldefined. Values for k l have been measured directly with both flash photolysis14*15J7~22 and discharge techniques, with reported 298 K rate constants ranging from 3.2 X 1OI2to 9.8 X cm3 molecule-I s-I and reported activation energies ranging from -352 to +274 cal mol-'. All direct measurements were carried out in the absence of the potentially reactive gas 02.Three competitive kinetics s t u d i e ~ , all ~ ~of ~which ' ~ ~ employed ~~ 1 atm of N, + 0, as the buffer gas, report 298 K rate constants similar to the higher values reported in the direct studies. While there seems to be general agreement that the branching ratio for channel l a is significant, the contribution from channel 1b remains poorly defined. Present address: Department of Chemistry, California Institute of Technology, Pasadena, CA 91 125.
0022-3654/86/2090-4148$01.50/0
In this paper we present the results of an extensive series of experiments aimed at elucidating the kinetics and mechanism of reaction 1 both in the absence and presence of 02. The conventional flash photolysis-resonance fluorescence technique was employed to study reactions 1-4 in argon buffer gas. Observed
-
+ CD3SCD3 products O H + CH3SC2H5 products O H + C2HSSC2H5 products OH
-
(2) (3)
(4)
temperature dependencies, reactivity trends, and isotope effects form a consistent picture that points toward hydrogen abstraction as the dominant observed reaction pathway for all reactions investigated. Also, a pulsed-laser photolysis-pulsed-laser-induced
(1) Galloway, J. N.; Whelpdale, D. M . Atmos. Enuiron. 1980, 14, 409. (2) Moller, D. Atmos. Enuiron. 1984, 18, 19. (3) Moller, D. Atmos. Enuiron. 1984, 18, 29. (4) Cullis, D. F.; Hirschler, M. M. Atmos. Enuiron. 1980, 14, 1263. (5) Granat, L.; Rodhe, H.; Hallberg, R. 0. SCOPE 1976, 7, 89-134. (6) Rodhe, H.; Isaksen, I. J. Geophys. Res. 1980, 65, 7401. (7) Bremner, J. M.; Stele, C. G. Adu. Microb. Ecol. 1983, 2, 155. Watkins, W. E.; Bingemer, H.; (8) Barnard, W. R.; Andreae, M. 0.; Georgii, H.-W. J. Geophys. Res. 1982, 87, 8787. (9) Andreae, M. 0.; Barnard, W. R.; Ammons, J. M. Ecol. Bull. 1983, 35, 167-177. (10) Andreae, M. 0.; Raemdonck, H. Science (Washington, D.C.) 1983, 221, 744. (11) Cline, J. D.; Bates, T. S. Geophys. Res. Lett. 1983, 10, 949. (12) Turner, S. M.; Liss, P. S. J. A m o s . Chem. 1985, 2, 223. (13) Andreae, M. 0.; Ferek, R. J.; Bermond, F.; Byrd, K. P.; Engstrom, R. T.; Hardin, S.; Houmere, P. D.; LeMarrec, F.; Raemdonck, H.; Chatfield, R. B. J . Geophys. Res. 1985, 90, 12891. (14) Atkinson, R.; Perry, R. A.; Pitts, J. N., Jr. Chem. Phys. Lett. 1978, 54, 14.
(15) Kurylo, M. J. Chem. Phys. Lett. 1978, 58, 233. (16) Cox, R. A,; Sheppard, D. Nature (London) 1980, 284, 330. (17) Wine, P. H.; Kreutter, N. M.; Gump, C. A,; Ravishankara, A. R. J . Phys. Chem. 1981, 85, 2660. (18) MacLeod, H.; Poulet, G.; LeBras, G. J. Chim. Phys. 1983, 80, 287. (19) Atkinson, R.; Pitts, J. N., Jr.; Aschmann, S. M. J . Phys. Chem. 1984, 88, 1584. (20) Barnes, I.; Bastian, K. H.; Becker, K. H.; Fink, E. H. Phys.-Chem. Behau. Atmos. Pollut., Proc. Eur. Symp., 3rd 1984, 149-157. (21) Martin, D.; Jourdain, J. L.; LeBras, G. Int. J. Chem. Kinet. 1985, 17, 1247. (22) Wallington, T. J.; Atkinson, R.; Tuazon, E. C.; Ashmann, S. M. Int. J. Chem. Kinet., in press. (23) Niki, H.; Maker, P. D.; Savage, C. M.; Breitenbach, L. P. Int. J . Chem. Kinet. 1983, 15, 647. (24) Hatakeyama, S.; Akimoto, H. J . Phys. Chem. 1983, 87, 2387. (25) Grosjean, D. Enuiron. Sci. Technol. 1984, 18, 460.
0 1986 American Chemical Society
The Journal of Physical Chemistry, Vol. 90, No. 17, 1986 4149
O H Reactions with Organic Sulfides fluorescence technique was employed to study reactions 1 and 2 in nitrogen, air, and oxygen buffer gases. W e find that in the presence of 02,a four-step mechanism involving abstraction, reversible addition, and reaction of the adduct with O2is required to explain all experimental observations. Effective rate constants and branching ratios for use in atmospheric models are determined as a function of temperature, and estimates of rate constants for elementary steps in the complex mechanism are obtained.
8
Experimental Section
Two separate series of experiments were carried out. The first employed a conventional flash photolysis-resonance fluorescence apparatus (FP-RF) while the second employed a pulsed-laser photolysis-pulsed-laser-induced fluorescence apparatus (PLP-PLIF). The two experimental setups are discussed separately below. FP-RF Apparatus. The FP-RF apparatus is described elsewhere.26 A brief review of its operation along with details pertinent to this investigation is given below. A jacketed Pyrex reaction cell with an internal volume of 150 cm3 was used in all experiments. The cell was maintained a t a constant temperature by circulating ethylene glycol from a thermostated bath through the outer jacket. A copperanstantan thermocouple with a stainless steel jacket was inserted into the reaction zone through a vacuum seal, thus allowing measurement of the gas temperature under the precise pressure and flow conditions of the experiment. OH radicals were produced by flash photolysis of H 2 0 at wavelengths between the onset of absorption at 185 nm and the Suprasil cutoff at 165 nm. An OH resonance lamp situated perpendicular to the flashlamp excited fluorescence in the 0-0 band of the A21;+-X211 system. Fluorescence was detected perpendicular to both the flash lamp and resonance lamp by a photomultiplier fitted with an interference filter (309.5-nm peak transmission, 10-nm full width at half maximum). Signals were obtained by photon counting in conjunction with multichannel scaling. For each decay rate measured, sufficient flashes were averaged to obtain a well-defined temporal profile over at least two and usually three l / e times. The flash duration was -50 ps while measured O H lifetimes ranged from 1.25 to 31.9 ms. In order to avoid accumulation of reaction or photolysis products, all experiments were carried out under “slow-flow” conditions. Sulfide reactant was flowed from a 12-L bulb containing a dilute sulfide/Ar gas mixture. An H20/Ar mixture was generated by bubbling argon through high-purity water at 298 K and a pressure of 900 Torr. The sulfide/Ar mixture, H 2 0 / A r mixture, and additional Ar were mixed before entering the reaction cell. The concentration of each component in the reaction mixture was determined from measurements of the appropriate mass flow rates and the total pressure. The fraction of sulfide in the sulfide/Ar gas mixtures was checked frequently by simultaneous measurement of the total pressure of the mixture and UV absorption by the sulfide at 228.8 nm. To carry out these measurements, a Cd penray lamp was used as the light source in conjunction with a 76-cm absorption cell and a band-pass filter-photomultiplier detection system. Required absorption cross sections were measured during the course of the investigation and are as follows (units are cm2): CH3SCH3, 11.6; CD3SCD3, 5.16; CH3SC2H5, 8.45; CH3CH2SCH2CH3, 9.64. PLP-PLZF Apparatus. The PLP-PLIF apparatus was employed to examine OH sulfide reactions under conditions approaching those found in the atmosphere. Pulsed-laser photolysis allows OH to be produced cleanly while completely avoiding photolysis of 02,and P L I F detection is sufficiently sensitive to allow detection of 1Olo OH per cm3 under kinetic conditions in 1 atm of O2 (where the OH fluorescence quantum yield is -0.0005). Another advantage of the PLP-PLIF apparatus over our conventional FP-RF system is a factor of 100 improvement in time resolution. Hence, the PLP-PLIF technique offers a better opportunity to observe the effects of short-lived intermediates.
A schematic of the PLP-PLIF apparatus is shown in Figure 1. The main body of the reaction cell was constructed of brass with a diameter of 3.8 cm and a length of 21 cm. Four Pyrex extensions were fitted to the main body of the cell to physically isolate the laser entrance and exit windows from the reaction zone, thus minimizing detection of laser scattered light. The main body of the cell was constructed with channels in its walls, through which heating or cooling fluids from a thermostated bath could be circulated. To measure temperature, the probe beam entrance window was replaced by a plexiglass plate fitted with a cajon fitting through which a jacketed copper constantan thermocouple could be injected (along the direction of flow) into the reaction zone; this technique allowed measurement of the gas temperature under the precise pressure and flow rate conditions of the experiment. To prevent heterogeneous decomposition of H202,the interior of the brass cell was coated with FEP Teflon and then overcoated with halocarbon wax. As in the FP-RF system, all experiments were carried out under “slow-flow” conditions. The linear flow rate through the reactor was (typically) 15 cm s-l and the laser repetition rate was 10 Hz. Since photolysis was always across the direction of flow, a “fresh” reaction mixture was available for each pulse (the probe beam entered the reactor along the direction of flow but its energy was sufficiently small that no significant buildup of photoproducts occurred during transport from the gas entrance port to the reaction zone). Sulfide concentrations employed in the PLP-PLIF experiments were much larger than in the FP-RF experiments and, as a result, could be measured directly in the slow-flow system by absorption at 228.8 nm. A 2-m absorption cell was employed for the concentration measurements. Measured rate constants were found to be independent of whether the reaction mixture flowed through the absorption cell before entering the reactor or after leaving the reactor. OH was produced by pulsed-laser photolysis of H202. Two different photolysis lasers were employed. One was a KrF excimer laser (Lambda Physik Model EMGZOOE) that could deliver up to 800 mJ per pulse at 248 nm, and the other was a frequencyquadrupled Nd:YAG laser (Quanta Ray Model DCR-2) that could deliver up to 40 mJ per pulse at 266 nm. OH was detected by pulsed-laser-induced fluorescence. The fluorescence excitation source was a frequency-doubled Nd:YAG laser (Quanta Ray Model DCR-2A). Three overlapping rotational lines2’ (Qll, Q1l’, R23)of the OH (A21;+-XZII) 1-0 band were excited simultaneously at 281.9145 nm. A line-narrowing etalon was employed in the dye laser to reduce the output band width to -0.001 nm;
( 2 6 ) Wine, P. H.; Kreutter, N. M.; Ravishankara, A. R. J . Phys. G e m . 1979.83, 3191.
(27) Dieke, G. H.; Crosswhite, H. M. J . Quanr. Spectrosc. Rad. Trans. 1962, 2, 97.
+
-
-
Figure 1. Schematic of the PLP-PLIF Apparatus. AC, absorption cell; BPF, band-pass filter; CdL,cadmium lamp; CM, capacitance manometer; D, frequency doubler; DG, three-channel delay generator; DL, dye laser; EM, energy monitor; GI, gas inlet; HS, harmonic separator; HV, high voltage; PA, picoammeter; PD, photodiode; PM, photomultiplier; PL, photolysis laser; RC, reaction cell; SA, signal averager; T, throttle; YL, Nd:YAG laser; 7-54F,Corning 7-54glass filter.
4150
The Journal of Physical Chemistry, Vol. 90, No. 1 7 , 1986
the probe laser power was typically 1 mJ per pulse. Fluorescence in the 0-0 and 1-1 bands was detected by an RCA 8850 photomultiplier after passing through a pyridine filter, two band-pass filters (310-nm peak transmission, 5-nm fwhm), and a Corning 7-54 filter. The filter system effectively discriminated against Rayleigh-scattered laser light and Raman scattering from N2 and 02.The photomultiplier output was fed to a waveform analyzer (Data Precision Model DATA 6000). The measured pulse height (voltage) when both lasers fired minus the measured pulse height when only the probe laser fired was proportional to the O H concentration in the reactor. The probe laser was triggered at time t (0-10 ms) after the photolysis laser and the signal was collected for (typically) 100 laser shots. Then the delay time was changed and the signal was collected again. Typically, about 10 different delay times were sampled to map out an OH temporal profile over two to three l / e times. The pulse widths of the photolysis lasers were 20 ns (excimer laser) and 6 ns (Nd:YAG laser) while measured OH lifetimes ranged from 5 to 15 000 ws. Chemicals. The pure gases used in this study had the following stated minimum purities: N2, 99.999%; 02,99.99%. Air was Matheson zero grade with total hydrocarbons 5 1 ppm. All buffer gases were used as supplied. H202(90% by weight) was obtained from FMC corporation. It was further concentrated by bubbling buffer gas through it before experiments were undertaken and constantly (24 h per day) during the course of the experiments. H 2 0was Baker HPLC grade and was used as supplied. Sulfides were obtained from Aldrich and had the following stated purities: CH3SCH3, 99+%; CD3SCD3,99+% chemical purity and 99.9 atom% D; CH3SC2H5,99%; C2H5SC2H5, 98%. Samples of each sulfide were transferred under dry nitrogen into Pyrex tubes fitted with high-vacuum stopcocks, degassed at 77 K, and purified by trap-to-trap distillation. The samples were subjected to additional degassing immediately before use. Absorption measurements around 280 nm indicated that impurity concentrations of highly reactive disulfides were negligible.
Results and Discussion FP-RF Experiments. To measure the bimolecular rate constants of interest, it is desirable to establish experimental conditions where the OH temporal profile is governed by the following processes: HzO 2 O H
OH
OH
-+
+ sulfide
-
+H
(5)
k,
products
(i = 1-4)
(i)
loss by diffusion from the detector field of view and reaction with background impurities (6)
Then, if [sulfide] >> [OH] (pseudo-first-order conditions), simple first-order kinetics would be obeyed In ([OH],/[OH],) = (k,[sulfide]
+ k6)t = k’t
(I)
The bimolecular rate constant, ki, is determined from the slope of a k’vs. [sulfide] plot. Observation of OH temporal profiles that are exponential (Le., obey eq I), a linear dependence of k’ on [sulfide], and invariance of k, to variations in flash intensity serve as strong evidence that the only processes that affect the OH time history are reactions i, 5, and 6, although reactions of O H with impurities in the sulfide/buffer gas mixture cannot be ruled out by the above set of observations if [impurity] >> [OH]. Under typical operating conditions (flash energy ca. 60 J, H 2 0 partial pressure ca. 150 mTorr) about 1 X 10” O H radicals are produced by the photoflash.26 At these low radical concentrations, reactions of OH with itself or with products of reactions i cannot contribute significantly to the observed temporal profiles. To investigate the possibility that reactive free radicals generated by photodissociation of the sulfide reactants could contribute to the OH temporal behavior, experiments were carried out for most reaction; where the flash intensity was varied by (typically) a factor of 3. No dependence of ki on flash intensity was observed for any of the reactions investigated. Most experiments were
Hynes et al. carried out at argon pressures of 30-40 Torr although we have previously studied reaction 1 over the range 50-200 Torr.I7 More complete pressure dependence studies of reactions 1 and 2 were carried out with the PLP-PLIF technique and are discussed below. In the FP-RF studies, both k l and k2 were observed to be independent of pressure. As predicted by eq I, exponential decays and linear dependencies of k’ on sulfide concentration were observed for all reactions investigated. The quality of the pseudo-first-order decay plots were similar to those published e l ~ e w h e r e ’ for ~ ~ ~O*H organosulfur reactions. The experimental results are summarized in Table I. Errors quoted for individual k, determinations are 2a and refer only to the precision of the k’vs. [sulfide] data. The absolute accuracy of the results is limited by precision, uncertainties in the determination of reactant concentrations, and other unidentified systematic errors that, of course, we believe to be negligible. We estimate that the absolute accuracy of a typical k, determination is *20%. For the three reactions where temperature-dependent kinetics data were obtained, the data are adequately described by the Arrhenius equation
+
k(T) = A , exp(-E,/RT)
(11)
Arrhenius parameters have been obtained from least-squares I data. Best fit Arrhenius parameters analyses of the In ki(T) vs. T along with 298 K rate constants are summarized in Table 11. Quoted errors in Arrhenius parameters are 2a and represent precision only (uA = Auld). The data for reaction 1 are in good agreement with our previously published results,I7 although the activation energy obtained from the new data is a little larger. Two sets of Arrhenius parameters for reaction 1 are given in Table 11. One set is derived from the data in Table I of this paper while the other is derived from the data in Table I combined with the data in Table I1 of ref 17. Since there is no reason to accept one data set over the other, we prefer the Arrhenius parameters obtained from the combined data sets. Our data show that the activation energy for reaction 4 is very close to zero. We feel that the temperature-independent rate constant k4 = (1.55 f 0.22) x 1O-” cm3 molecule-’ s-’ (error is 2u, precision only) is as good a representation of the data as the Arrhenius parameters given in Table 11. Important observations concerning the FP-RF results are (1) enhanced reactivity for sulfides with secondary hydrogen atoms, (2) a significant kinetic isotope effect for reactions 1 and 2, and (3) small but positive activation energies for reactions 1 and 2. These observations all point toward hydrogen abstraction as the dominant observed reaction pathway
+ H20 (la) CD3SCD2 + HDO (2a) OH + CH3SCH2CH3 CH3SCHCH3+ H 2 0 (major) (3a) CHzSCHzCH3+ H 2 0 (minor) (3b) CH3SCH2CH2+ HzO (minor) (3c) O H + CH3CHzSCH2CH3 CH3CH2SCHCH3(major) + CH3SCH3 OH + CD3SCD3 OH
-
-
-+
CH3SCH2
-
(4a)
CH3CH2SCH2CH2(minor)
(4b)
The expected major abstraction pathways for reactions 3 and 4 are indicated above. The 298 K rate constant for the minor channel (3b) can be estimated to be 0.5kl, or -2.2 X 10-l2 cm3 molecule-’ s-I. Hence, k3,,/k3 is probably -0.26. Rate constants for the minor channels (3c) and (4b) can be “guesstimated” by assuming that primary hydrogens in the /3 position have the same reactivity as primary hydrogens in n-propane. Using the group rate constants summarized by A t k i n ~ o nwe , ~ ~obtain the guesst(28) Wine, P. H.; Thompson, R. J.; Semmes, D. H. In?. J . Chern. Kinet. 1984, 16, 1623.
The Journal of Physical Chemistry, Vol. 90, No. 17, 1986 4151
O H Reactions with Organic Sulfides TABLE I: Summary of FP-RF Rate Constant Data" sulfide CH3SCH3
CD3SCD3
CHPSC2HS
C2H5SCZH5
T, K
P,Torr
flash energy, J
no. exptsb
276 298 298 300 359 374 374 397 253 299 299 299 360 360 418 299 299 299 255 255 269 299 299 338 370
40 40 30 40 40 40 40 40 30 30 30 300 30 30 30 40 40 40 40 40 40 40 40 40 40
60 60 60 60 60 60 60 60 60 60 60 60 34 110 55 60 75 25 60 30 48 58 60 48 60
5 6 5 5 4 4 6 5 6 4 5 5 5 4 5 7 5 5 6 6 4 5 5 5 6
range of k',
+
s-l
lo%, 2a: cm3 molecule-'
32-402 32-356 41-375 32-420 31-383 34-356 32-377 37-374 43-370 45-302 48-330 65-345 42-348 42-338 49-347 38-647 37-595 34-543 43-644 37-798 49-692 37-578 41-590 44-555 35-469
s-l
4.17 f 0.87 4.09 f 1.16 4.44 f 0.23 4.75 f 0.71 5.45 f 0.89 5.97 f 0.07 5.46 f 0.52 5.69 f 0.46 1.46 f 0.14 1.95 f 0.13 1.87 f 0.16 1.98 f 0.18 2.53 f 0.19 2.72 f 0.21 3.11 f 0.18 8.54 f 0.81 8.65 f 1.12 8.31 f 0.53 14.2 f 1.8 17.6 f 2.5 15.4 f 1.6 14.5 i 1.2 16.1 i 2.1 15.4 f 2.3 15.1 f 2.2
-
'Argon was used as the buffer gas in all experiments. H20(0.15 torr) was used as the photolyte in all experiments. The linear flow rate through the reactor was -2 cm/s in all experiments. The photolysis repetition rate was 1 Hz in all experiments. bExperiment is the determination of one pseudo-first-order rate constant. CErrorsrepresent precision only.
TABLE 11: Arrbenius Parameters and 298 K Rate Constants Obtained from the FP-RF Data" sulfide Ab EIR. K k (298 K ) b notes CH3SCHp 13.6 f 4.0 332 i 96 4.46 c, e 9.6 f 2.1 234 f 66 4.38 d, e CD3SCDP 10.3 f 1.7 498 f 51 1.94 c, e CH3SC2HS C2HSSC2HS
13.9 f 6.3 15.5 f 2.2
-31 0
* 132
8.50 15.4 15.5
c, e
f
'Errors are 20. b u n i t s are cm3 molecule-' s-l. cArrhenius parameters calculated with data in Table I. dArrhenius parameters calculated with data in Table I plus data in Table I1 of ref 17. C k(298 K) calculated from Arrhenius parameters. funweighted average of rate constants given in Table I; k4 assumed independent of temperature.
-
-
imates k,/k3 0.022 and k4,,/k4 0.024. PLP-PLIF Experiments. The PLP-PLIF apparatus was employed to investigate OH reactions with CH3SCH3and CD3SCD3 in N,, air, and 0,buffer gases. The relevant reaction scheme was expected to be
+ hv
H202
OH
+ sulfide OH
OH
-+
kl
20H
products
H202
HO2
(7) (j = 1, 2)
+ H20
6) (8)
loss by diffusion from the detector field of view and reaction with background impurities (6)
All experiments were carried out under pseudo-first-order conditions with the sulfide in large excess over OH. As in the FP-RF experiments, exponential OH decays were observed in all experiments with the decay rates being independent of large variations in photolysis laser power. In most experiments the initial OH concentration was between 1 X 10" and 1 X lo1, molecule ~m-~. Typical OH temporal profiles observed in 1 atm of air are shown in Figure 2, while typical plots of k'vs. sulfide concentration are (29) Atkinson, R. Chem. Reu. 1986, 86, 69.
0.011 0
'
'
"
1
0.5
'
"
'
I
I.o
"
'
L
tlme (ms)
Figure 2. Typical O H temporal profiles observed in the PLP-PLIF
-
+
experiments. Reaction: OH CD3SCDI. Experimental conditions: 298 K, 700 Torr air, [OH], 7 X 10" ~ m - ~CD3SCDP . concentrations were (a) 0, (b) 8.43 X l O I 4 ~ m - ~(c), 4.84 X 1015 cm-'. Solid lines were obtained from linear least-squares analyses and give the following pseudo-first-order rate constants: (a) 670 f 110 s-', (b) 3960 f 240 s-I, and (c) 17200 f 1400 s-I (errors are 20 and represent precision only).
shown in Figure 3. The data in Figure 3 clearly demonstrate that the observed bimolecular rate constant (koM the slope of the k'vs. [sulfide] plot) depends on the O2concentration! This result suggests a complex reaction mechanism involving an intermediate that reacts with O2and seems to be in conflict with the conclusion from the FP-RF experiments that the predominant pathway is hydrogen abstraction. Observed pressure dependencies of koM in air at 298 K for reactions 1 and 2 are plotted in Figure 4. The temperature dependence of kobd for reaction 2 in 700 Torr air and 700 Torr O2 is plotted in Arrhenius form in Figure 5; for comparison the Arrhenius line derived from the FP-RF data is also plotted. All measured values of kobsdfor reactions 1 and 2 are summarized in Table 111. Errors quoted in Table I11 are 2a and represent precision only. We estimate the absolute accuracy of a typical koW determination to be *20%. Reaction Mechanism. All experimental observations are
4152
The Journal of Physical Chemistry, Vol. 90, No. 17. 1986
Hynes et al. TABLE I11 Observed Bimolecular Rate Constants as a Function of Temperature, Pressure, and O2Concentration for PLP-PLIF Experiments p, range of kobd f 2a," sulfide T. K Torr M k'. s-I cm3 molecule-' s-l CD3SCD3 261 700 air 1080-50500 11.6 f 1.1 266 700 02 13.5 f 1.2 854-48500
CCDISCD.l(lO~~cm-~)
CH3SCH3
Figure 3. Typical k' vs. reactant concentration plots observed in the PLP-PLIF experiments. Reaction: OH CD3SCD3at 298 K. Buffer gas: (a) 450 Torr N2,(b) 700 Torr Air, (c) 700 Torr 02.Solid lines were obtained from linear least-squares analyses and give the following bimolecular rate constants in units of cm3/molecule-' s-*: (a) 1.82 & 0.11, (b) 3.40 f 0.13, (c) 6.50 f 0.72 (errors are 2a and represent precision only).
+
0
0
a
275 276 287 287 298 298 298 298 298 298 317 321 340 340 361 261 262 279 298 298 298 298 298 298 298 321
606-54200 1650-47300 777-20100 593-23100 1520-18900 193-19800 336-1 7300 804-1 1700 672-1 8900 1290-2 1200 ' 817-16500 620-13600 1030-11470 02 547-7880 air 1110-15200 N2 1060-208000 air 498-24900 372-24200 air 160-13700 N2 53-7500 SF6 15 1-86 10 air air 1960-21800 air 310-28900 air 596-56100 air 1850-65700 air 420-22300 O2
700 700 700 700 450 100 300 500 700 700 700 700 700 700 700 700 700 700 40 500 50 130 340 590 750 700
air air 02 N2 air air air air 02 air 02 air
11.9 f 2.0 9.63 f 0.63 5.29 f 0.44 6.99 f 0.53 1.82 f 0.11 2.10 f 0.15 2.68 f 0.09 2.97 f 0.13 3.400 f 0.13 6.50 f 0.72 3.02 f 0.18 3.72 f 0.27 2.32 f 0.1 1 2.30 f 0.28 2.66 f 0.11 4.29 f 0.48 12.5 f 1.7 9.53 f 0.28 4.80 f 0.1 1 4.75 f 0.15 4.68 f 0.08 5.04 f 0.14 5.18 f 0.34 5.80 f 0.16 6.28 f 0.10 5.43 h 0.30
Errors are 2a and represent precision only.
CHSSCHI
consistent with the following mechanism (written for CD3SCD3 but identical for CH3SCH3): OH + CD3SCD3 CD3SCD2 HDO (2a)
+
+ CDsSCD3
+ OH + CD3SCD3 + M CD3S(OH)CD3 + M CD3S(OH)CD3+ O2 products OH + H202 HO2 + H2O
-
OH
'i O
O400
*
P(Torr)
Figure 4. Plots of kobd vs. pressure in air at 298 K. The data pints at zero pressure are the PLP-PLIF results obtained in the absence of 02.
-
+
-
(2b, -2b)
(9)
(8) loss by diffusion from the detector field of view and reaction with background impurities ( 6 )
In the above mechanism, CD3S(OH)CD3 is a thermalized adduct. In the absence of O2all observed OH removal (by DMS) appears to be via the abstraction route, i.e., reaction 2a. Apparently, reaction -2b is very fast compared to the time scale of our experiments. However, the adduct lifetime must be long enough that it can be scavenged by O2 in competition with decomposition back to reactants. The dependence of k,, on temperature (Figure 5) is qualitatively consistent with the above mechanism. The activation energy for reaction -2b is expected to be much larger than the activation energy for reaction 9. Therefore, the fraction of adduct molecules scavenged by O2 can increase dramatically over a relatively small temperature range. At high O2levels, the adduct can be assumed to be in steady state. Applying the steady-state approximation to the proposed mechanism leads to the expression
where X(T) 2.6
3.0
3.4
3.0
1000/T(K)
Figure 5. Arrhenius plots for the OH + CD3SCD3reaction in 700 Torr air (0)and 700 Torr O2 ( 0 ) . The dashed line is obtained from the FP-RF results (Table I) and represents the Arrhenius behavior in the absence of 02.
I
k,(T)/k-Zb(T)
(IV)
In the above equations k2b and k-2b are bimolecular rate constants that are understood to depend on the identity and concentration of the buffer gas M. Under our experimental conditions, reactions 2b and -2b are probably in the "fall-off" region between the low-pressure third-order limit and the high-pressure second-order limit.
The Journal of Physical Chemistry, Vol. 90, No. 17, 1986 4153
O H Reactions with Organic Sulfides TABLE IV Results Obtained from Simulating the Dependence of k- on [O,] and Temperature for the OH + CD9CD3 Reaction' A. Comparison of Measured and Simulated Values for koMbs simulated kAM -0.7 kcal/mT E2. = 0 M measd kobd 11.5 11.8 air 11.6 f 1.1 26 1
z=
266 275 276 287 287 298 298 317 321 340 340 361
0 2
0 2
air air 0 2
air 0 2
air 0 2
air 0 2
air
13.5 f 1.2 11.9 f 2.0 9.63 dZ 0.63 5.29 f 0.44 6.99 0.53 3.40 f 0.13 6.50 0.72 3.02 dZ 0.18 3.72 f 0.27 2.32 0.1 1 2.30 dZ 0.28 2.66 f 0.11
* *
*
12.4 11.4 8.26 5.38 9.20 3.54 6.69 2.50 3.33 2.45 2.69 2.61
B. Best Fit Parameters E,h = -0.7 kcal/mol cm3 molecule-' s-' 3.04 A ~ iw3' , cm3 molecule-' 5.53 Ex/R
7460
K
20
i
/ O
12.3 11.6 8.21 5.30 9.52 3.48 6.81 2.48 3.29 2.44 2.67 2.61
E2b= 0 11.3 1.64 7760
'All bimolecular rate constants were measured at a total presume of 700 Torr. bunits for koW are loi2cm3molecule-' 8'. CErrorsare 2a,
precision only. We have taken the 13 bimolecular rate constants for reaction
2 that were measured in 700 Torr air or 700 Torr O2 and fit koW(T,[02]) to eq I11 using a trial and error procedure with a least-squares-fitting criterion. It is assumed that over the limited temperature range 260-360 K all elementary reaction rate constants can be expressed in Arrhenius form, Le., eq 11. Arrhenius parameters for reaction 2a were obtained from the Fl-RF results (Table 11). Values for A2br Ax ('AglA-~b), and Ex (EE-2b - E9) were taken as independent variables. By analogy with OH addition to CH3SH,2* CH3SD,28and CH3SSCH3,"EZbwas fixed at -0.7 kcal/mol. To test the sensitivity of the fit to the choice of E2b, the procedure was repeated with klb taken to be temperature independent. The results are summarized in Table IV. The quality of the fit is relatively insensitive to the choice of E2b. Rather different combinations of 4,and Ax give the best fit for the different choices of E 2 b but very similar values for Ex are obtained in both cases. Ex is the parameter that is determined most accurately from the above analysis. Equation I11 does quantitatively reproduce the experimental dependence of koW on temperature and oxygen concentration. We therefore conclude that the proposed mechanism does include all important reactions. As discussed below, Arrhenius parameters obtained from the above analysis can be used in conjunction with other experimental information to obtain rate constants for each elementary reaction in the mechanism. Estimation of Elementary Reaction Rate Constants. The rate equations for the proposed mechanism can be solved analytically
Caulfldol (lO'*crn-')
Figure 6. Dependence of k'on [sulfide] for reactions 1 and 2 at 261 K
in 700 Torr N2.Dashed line is obtained from a linear least-squares analysis of the k'vs. [CH3SCH3]data and gives the rate constant kl = (4.29 f 0.48) X an3molecule-' s-I, where the uncertainty is 20 and represents precision only. Solid lines represent the dependence of k:', on [CD3SCD3]calculated from eq V for various choices of the adduct lifetime (given in the figure), assuming kZ = 1.72 X lo-'' cm3/molecule-' s-' and kfb = 1.16 X lo-'' cm3 molecule-' s-l (see text for details).
in equilibrium. Clearly, if the approach to equilibrium could be observed in real time, then the observed temporal profiles could be modeled to obtain k2band k-2b. Attempts were made to observe the approach to equilibrium in 700-Torr SF6and Ar at both 298 and 260 K. CD3SCD3was used as the sulfide reactant in these experiments because, since ka < kl,, the approach to equilibrium should be easier to observe for the deuterated analogue. The time resolution was 100 ns and meaningful data could be obtained at delay times (after the photolysis laser fired) of 1 2 0 0 ns. The "approach to equilibrium" experiments were not successfulsingle-exponential OH decays were observed under all experimental conditions investigated. This negative result suggests that the adduct lifetime is very short. Simulations of OH temporal profiles with [O,] set equal to zero show that at low sulfide concentration, where the equilibrium concentration of OH is much greater than that of adduct, the slow component of the double-exponential decay directly measures the abstraction rate constant (i.e., kilow= k,,[sulfide] + k6 + k8[H202],j = 1 or 2); this condition was met in the FP-RF experiments. At sufficiently high sulfide concentration, however, k'h becomes less than kb [sulfide] + k6 k8[H20z]. The physical interpretation of this nonlinearity in the kilo, vs. sulfide concentration dependence is quite clear. At high sulfide concentration, [OHIO - ( K + XI) exp(X1t) - ( K + X2) exp(X2t) -a significant fraction of OH is tied up in the form of the (un(V) [OH], X I - A2 reactive) adduct. Hence, the apparent rate of OH removal per unit concentration of sulfide decreases with increasing sulfide where concentration. If kilowcould be measured a t very high sulfide (VI) K = k-2b k,[O2] concentrations, then the kih vs. sulfide concentration dependence be modeled to obtain kIb/k-,%,i.e., the equilibrium constant could X1 = 0.5((a2 - 4/3)'/* - a] (VW for adduct formation and decomposition. Since kfi and k,/k+ are independently determined from the rate constant measure(VIII) X2 = -0.5((az- 4j3)'/' a) ments in O2and air and kjais independently determined from the a = K + kg[H202] k6 + (kza + k2b)[CD$CDs] (IX) FP-RF experiments, rate constants for all elementary reactions in the mechanism would thus be obtained. B = ( k 8 W A I + kcJK + (k2aK ~ ~ ~ ~ ~ [ O ~ I I [ C D(X) ~ S C D B IAttempts to observe nonlinearity in the k i l , vs. [CD3SCD3] dependence were unsuccessful at 298 K but successful a t 261 K. A double-exponential OH decay is predicted. In the absence of The data, obtained in 700 Torr N2buffer gas, are shown in Figure 02,the fast component (kIfasJ represents the approach to equi6, along with simulations that employ previously determined values librium between photolytically produced OH and adduct while for kza and k2b along with different guesses for k-2b. It should the slow component ( k i h ) represents the decay of OH and adduct
+
+
+
+
4154 The Journal of Physical Chemistry, Vol, 90, No. 17, 1986
Hynes et al.
TABLE V Summary of Reported Rate Constants for OH Reactions with CH$CH3 and C2H5SC2H5in the Absence of O2 sulfide technique' I O i 2 k (298 K),b cm3 molecule-I s-' EIRE,K range T, K CH3SCH3 FP-RF 10.0 -179 & 151 300-427 FP-RF 9.1 -131 f 430 273-400 DF-EPR 10.4 -176 h 200d 313-513 DF-EPR 3.2 FP-RF 3.9 -130 f 102 297-400 CK-GC' 5.3 FP-RF 4.4 +234 f 66 248-397 PLP-PLIF 4.8 C2HX2HS DF-EPR 12 FP-RP 15.4 -31 f 132 255-370
,
ref 14 15 18 21 22 22 11, t d td 21 t d
'FP, flash photolysis; DF, discharge flow; CK, competitive kinetics; PLP, pulsed-laser photolysis; RF, resonance fluorescence; EPR, electron paramagnetic resonance; GC, gas chromatography; PLIF, pulsed-laser-induced fluorescence. Calculated from Arrhenius expression when temperature-dependent data were obtained. Errors are 2a. Calculated from reported rate constants at two temperatures. Errors reflect reported uncertainties in individual rate constants. 'OH produced from N2H4+ O3reaction; rate constant measured relative to an assumed rate constant of cm3 molecule-l SKI for the O H + cyclohexane reaction. ftw means this work. 7.3 1 X be noted that decay rates in excess of lo5 s-' had to be measured in order to observe nonlinearity. We find that the value of k-2b that best fits the experimental data is sensitive to the choice of k,,. From the FP-RF results, we obtain kZa= 1.53 X lo-', cm3 molecule-' s-I; using this value in the simulations gives k-2b = 3.6 X lo6 s-'. On the other hand, the data at low CD3SCD3in Figure 6 are most accurately reproduced by taking k,, = 1.72 X lo-', cm3 molecule-' s-'; using this value in the simulations gives k-2b = 2.5 X lo6 s-', Considering all uncertainties, we feel that our data supports a value of (3.5 f 2.0) X lo6 s-' for k-2bat 261 K. The temperature dependence studies of kobd in 700-Torr air and cm3 molecule-' 700-Torr 0, support a value of (1.4 f 0.3) X for k9/k-2bat 261 K (Table IV). Hence, we obtain for k, at 261 K the value (4.2 f 2.2) X lo-', cm3 molecule-' s-'. In summary, the following elementary reaction rate constants have been determined for the O H CD3SCD3reaction at 261 K:
+
kza = (1.6 f 0.2) k2b
X
10-I2 cm3 molecule-' s-l
= (1.15 f 0.20) x lo-'' cm3 molecule-'
s-l
(700 Torr N2
+ 0,)
k-2b = (3.5 f 2.0) x lo6 S-' k9 = (4.2 f 2.2)
X
lo-', cm3 molecule-' s-'
Experiments similar to those described above were also carried out for the OH CH3SCH3reaction. In this case, however, the k',,, vs. [CH3SCH3]dependence was linear for k:low5 2 X lo5 s-* (see Figure 6). Since the abstraction rate for reaction 1 is significantly faster than for reaction 2, we expect that nonlinearity in the k',,,, vs. [CH3SCH3]dependence would be observed only at very large values for k's,ow.Our experimental observations do not imply that k-lbmust be larger than k-2b. While some isotope effect would be expected for the adduct decomposition step, the data in Figure 4 demonstrate that the competition between adduct decomposition and adduct reaction with 0, shows no isotope dependence. Adduct Binding Energy. From the results summarized in Table 15 f 2 kcal/mol. Since reaction IV we obtain E , = E-2b-.E9 9 is quite fast, its activation energy must be small--E9 = 0.5 f 1.0 kcal/mol. This leads to the estimate E-2b = 14.5 f 3 kcal/mol. Another estimate for E-2b can be obtained by using our rate constant k-2b= (3.5 f 2.0) X lo6 s-' a t 261 K and recognizing s-'.~' that A factors for simple bond fissions are typically Substituting the rate constant and A factor information into eq 11 leads to the estimate E-zb = 11.3 f 1.7 kcal/mol. Hence, our results support a value of 13 f 3 kcal/mol for E-2b. Possible Secondary Chemistry Complications. One kinetic complication that would not be detected through the variations in experimental parameters described earlier in the paper is regeneration of O H via sulfide reactions with HO, HOz + CH3SCH3 OH + CH,S(O)CH,, AH = -21 kcal/mol (10)
+
-
+
(30) Benson, S. W. Thermochemical Kinetics; Wiley-Interscience: New York, 1976. (31) Niki, H., private communication.
HO, + CD3SCD3 OH + CD,S(O)CD, (1 1) In the absence of oxygen, H 0 2 can be produced in our experiments via reaction 8. Our experimental conditions were such that only a small fraction of initially produced OH (typically a few percent) reacted with H 2 0 2to produce HO,. Therefore, even if reactions 10 and 11 were gas kinetic, they would have a negligible effect on our rate constant determinations with [O,] = 0. In the presence of oxygen, however, HO, can be produced via the adduct O2 reaction; then if reactions 10 and 11 were sufficiently rapid, (faster than cm3 molecule-' s-') we would underestimate kow: Le., the 'oxygen effect" would be even larger than we have observed! There is no kinetic data for reactions 10 and 11 reported in the literature. However, Niki)' reports unpublished results that suggest that klo I1 X cm3 molecule-' s-I. Also, the similar reaction of HOz with SO, HOz SO, O H SO,, AH = -17 kcal/mol (12) -+
+
+
-
+
is known to be very s I o w ~ * - with ~ ~ measured rate constants ranging cm3 molecule-' s-I. We therefore from 8.7 X to i l X consider it extremely unlikely that reactions 10 and 11 contributed to our observed OH temporal profiles. Comparison with Previous Work. For comparison purposes, it is useful to divide the available data base into two subsetsresults obtained in the presence and absence of 0,. In Table V our results for reactions 1 and 4 are compared with other results obtained in the absence of 0,. There are no other data with which to compare our results for reactions 2 and 3. The three earliest 02-free studies of k , other than our 1981 study," all reported relatively fast rate c o n ~ t a n t s . ' ~ ~ 'However, ~J* two of these studies have now been superceded by more careful studies from the same laboratories.zo~21 It now seems clear that the earlier DF-EPR work of MacLeod et a1.I8 overestimated kl because a heterogeneous reaction between OH and CH3SCH3on the walls of the flow tube contributed significantly to O H removal. The flash photolysis results of Atkinson et al.I4 and KuryloI5 may have been influenced by the presence of reactive impurities in the CH3SCH3samples; CH3SSCH3,for example, reacts with O H 50 times more rapidly than does CH3SCH3.16,17The two recent studies by Martin et aL2' and Wallington et aL2, report 298 K rate constants for reaction 1 that are lower by -25% than the values we obtain. Also, Wallington et al. observe a small negative activation energy for reaction 1 while we observe a small positive activation energy. Our experiments were carried out over a long period of time and employed many different samples of CH3SCH3. In some cases, UV absorption measurements around 280 nm were carried out to check for the presence of the reactive impurities CH3SSCH3 and CH3SH. No evidence for the presence of these impurities was observed. We also found that fractional distillation of CH3SCH3did not affect the observed kinetics. Hence, we do not (32) Payne, W. A.; Stief, L. J.; Davis, D. D. J . Am. Chem. Soc. 1973, 95, 7614. (33) Burrows, J. P.; Cliff, D. I.; Harris, G. W.; Thrush, B. A,; Wilkinson, J. P. T. Proc. R. SOC.London, A 1979, 368, 463. (34) Graham, R. A,; Winer, A. M.; Atkins, R.; Pitts, J. N., Jr. J . Phys. Chem. 1979, 83, 1563.
The Journal of Physical Chemistry, Vol. 90, No. 17, 1986 4155
OH Reactions with Organic Sulfides TABLE VI: Summary of Reported Rate Constants for the OH C H W H 3 Reaction in Air at 298 K and 730 30 Torr Total
*
Pressure
techniaue’ CK-GC CK-GC CK-GC CK-GC CK-FTIR CK-GC CK-GC PLP-PLIF
k,b cm3 molecule-’ s-I
ref reactant
OH prodn scheme
+
TABLE VII: Recommended Rate Constants and Branching Ratios for Reaction 1 in 1 atm of Air
ref
~~
9.1 f 1.4 10f 1 42 f 4 45 f 6 47 f 7 9.3 f 0.7 8.5 f 0.2 6.3 f 0.3
ethylene n-hexane
16 19 20 20 20 22 22 twh
c
d
propylene propy 1ene propylene n- hexane cyclohexane
e
d d e
f g
’CK, competitive kinetics; GC, gas chromatography; FTIR, Fourier transform infrared spectroscopy; PLP, pulsed-laser photolysis; PLIF, pulsed-laser-inducedfluorescence. * Errors are 2a and represent precision only. “ON0 + hu OH. “CH30N0 + hv OH. CNO, + RH + hu OH. JN,H, + 0 9 OH. pH202 + hu OH. * tw means this work.
-
-
4
“Units for koW are
k0Ma
Babs
Badd
15.8 15.2 14.5 13.6 12.5 11.2 9.8 8.5 7.4 6.5 5.9 5.5 5.2
0.24 0.25 0.27 0.29 0.32 0.37 0.42 0.50 0.58 0.67 0.75 0.82 0.87
0.76 0.75 0.73 0.71 0.68 0.63 0.58 0.50 0.42 0.33 0.25 0.18 0.13
cm3 molecule-I s-l.
4
believe that the discrepancy between our values for k l in the absence of O2and those reported by Martin et al. and Wallington et al. can be attributed to the presence of reactive impurities in our CH3SCH3 samples. Another argument against impurity reactions as an explanation for the differences in reported rate constants is that, since OH reactions with CH3SSCH3and CH3SH both display negative activation energies,” the presence of these impurities in our samples would have resulted in observation of a lower activation energy for reaction 1; instead, we observe a higher activation energy than Wallington et al. Other than the results reported in this paper, all studies of reaction 1 in the presence of O2 have been restricted to 297 f 3 K and all have been relative rate measurements. Reported rate constants are summarized in Table VI. Our directly measured rate constant is significantly lower than all values obtained via competitive kinetics techniques. The probable reason for the high rate constants obtained in the competitive kinetics studies is the existence of a secondary reactant (other than OH) that removes CH3SCH3more rapidly than it removes the reference organic. Wallington et alez2have carried out relative rate measurements as a function of O2 concentration using two different schemes for producing OH (see Table VI). They obtained significantly higher values for kl when methyl nitrite photolysis in the presence of O2 and NO was used as an OH source than when the dark reaction of hydrazine with ozone was employed to generate OH. In both cases, they did observe that the rate constant increased with increasing O2concentration. On the basis of the observation that changing the OH source led to a change in the measured rate constant, Wallington et al. have postulated the existence of a secondary pathway for CH3SCH3removal and further speculated that this pathway involves the CH3S radical. Identification of the interfering reaction(s), however, must await further experimentation. Barnes et a1.20 also report relative rate data that indicate that the OH CH3SCH3reaction rate increases with increasing O2concentration. However, the rate constants reported by these authors are much larger than those measured in other laboratories at all O2levels investigated (2-300 Torr 02,740 Torr total pressure). Barnes et al. acknowledge that their results should be considered preliminary since experiments were not carried out over a wide enough range of conditions to completely exclude interferences from species such as O(3P), 03,NO3, and walls. It is interesting to compare our rate constant for reaction 9, cm3 molecule-’ s-I a t 261 K, with the rate (4.2 f 2.2) X constant for reaction of the OH + SO2 adduct with O2 HOS02 + O2 H 0 2 SO3 (13)
+
+
T, K 250 255 260 265 270 275 280 285 290 295 300 305 310
+
-
M a r g i t a r ~has ~ ~recently reported that k13 = (4 f 2) X cm3 molecule-’ s-’ at 298 K i.e., k9 10k13. As discussed in the next section, reaction 9 could involve a hydrogen transfer similar to reaction 13 but may also proceed via a methyl transfer route. Even (35) Margitan, J. J. J . Phys. Chem. 1984, 88, 3314.
if both reactions proceed via the same (hydrogen-transfer) mechanism, the observed difference in reactivity is not unexpected. This is because reaction 9 is about 20 kcal/mol exothermic while reaction 13 is nearly thermoneutral and may, in fact, be slightly endothermic. Implications for Atmospheric Chemistry. The results reported in this paper demonstrate that both the effective rate constant (kobsd)and the branching ratio (addition vs. abstraction) for reaction 1 change dramatically as a function of temperature over the relevant range of lower tropospheric temperatures 250-3 10 K. It should be kept in mind that, for purposes of atmospheric modeling, addition followed by decomposition back to OH + CH3SCH3is treated as no reaction. The “effective” addition pathway represents only those adduct molecules that are scavenged by 0 2 . A majority of our PLP-PLIF experiments were carried out with CD3SCD3as the sulfide reactant because more information about elementary reaction rates could be obtained in this manner. However, enough experiments were carried out with CH3SCH3 to demonstrate that, within experimental uncertainty, koa values for the two compounds differ only by the difference in the abstraction rates. The data in Figure 4 strongly support the validity of this approximation. Plugging the appropriate Arrhenius parameters (obtained from Tables I1 and IV) into eq 111, we obtain the following expression for the temperature dependence of kow for reaction 1 in 760 Torr air (units are cm3 molecule-’ s-l):
+
+
kobsd = [Texp(-234/n 8.46 x exp(7230/T) 2.68 X exp(7810/T)]/ [1.04 X 10”T + 88.1 exp(7460/T)] (XI) Values for kobd at 5-deg intervals have been calculated from eq XI and are tabulated in Table VII. Also tabulated are branching ratios for abstraction (Babs)and addition (Bad,,). The branching ratios were calculated from the relationships
(XIII) Under atmospheric conditions, the abstraction route is thought to result in production of CH3S H2C0via the following reaction seq~ence:~~?~~
+
OH
+ CH3SCH3
--+
H20
+ CH3SCH2
+ 0 2 + M CH3SCH202 + M CH3SCH202 + NO CH3SCH20 + NO2 CH3SCH2O + M CH3S + CH2O + M
CH3SCH2
+
+
--+
(la) (14) (15) (16)
The ultimate fate of CH3S is unknown, although Balla et report direct kinetic evidence that this radical reacts very rapidly with NO and NO2 but negligibly slowly with 02.Possible routes (36) Balla, R. J.; Nelson, H.H.; McDonald, J. R. Chem. Phys., submitted.
J. Phys. Chem. 1986, 90, 4156-4158
4156
+
for the adduct 0, reaction include the following: C H 3 S C H 3 0 H+ O 2 CH3S(0)CH3+ HOz (MezSO) CH3O2 CH3SOH
-
+
-+
(1 7a) (1 7b)
CH3SOH is probably converted to CH3S03H (methanesulfonic acid) by reaction with Oz while the atmospheric fate of M e 2 S 0 is unclear. MezSO has a very low vapor pressure and may be rapidly removed via heterogeneous processes. At 298 K our results demonstrate that reaction 1 in 1 atm of air proceeds 70% via abstraction and 30% via (irreversible) addition. Photooxidation studies have been reported by Niki et al.23 and Hatakeyama and A k i m o t ~where , ~ ~ SO2 yields from OHinitiated oxidation of CH3SCH3were reported to be 22% and 21%, respectively. Large yields of methanesulfonic acid were observed
in both studies. At oresent, there is insufficient information to allow SOzproductioi to be associated with either the abstraction route or the addition route. However, it should be noted that our results suggest that abstraction is the dominant reaction pathway for T > 300 K while addition is the dominant pathway for T < 270 K. Hence, temperature-dependent product analysis studies should shed some light on the detailed pathways for SO2 and C H 3 S 0 3 Hproduction.
Acknowledgment. We thank R. J. Thompson for helping with some of the FP-RF experiments, J. K. Lawson for assistance in the data analysis, and J. M. Nicovich for helpful discussions and assistance in the data analysis. We also thank H. Niki for communicating to us his unpublished results on the HOZ CH3SCH3 reaction. This work was supported by the National Science Foundation through Grant No. ATM-82-17232.
+
Substituent Effects on Rates of One-Electron Oxidation of Phenols by the Radicals C102, NO2, and SO3Zeev B. Alfassi: Robert E. Huie, and P. Neta* Chemical Kinetics Division, Center for Chemical Physics, National Bureau of Standards, Gaithersburg, Maryland 20899 (Received: January 28, 1986)
Rate constants for the reactions of CIOz, NOz, and SO3-radicals with several substituted phenoxide ions have been measured by pulse radiolysis. They vary from the immeasurably slow (lo9 M-' s-l) and depend on the redox potentials of the phenoxide ions and the inorganic radicals. With the weak oxidant SO; reverse reactions were observed in certain cases; i.e., the phenoxyl radical oxidizes sulfite ions. An attempt is made to correlate the rate constants with Hammett's substituent constants and the results are compared with those obtained previously for the reactions of various inorganic radicals with phenols and phenoxide ions.
Introduction
One-electron oxidation of phenoxide ions to phenoxyl by the azide radical takes place very rapidly XC,jHdO-
+ N3
+
XC6H40'
+ N3-
(1)
with k , = 4 X lo9 M-' s-I for a variety of substituted phenols, with little selectivity.] On the other hand, reaction of N3 with the neutral phenols is much more selective XC6H40H
+ N3
+
XC6H40'
H+
N3-
(2)
and its rate constant varies from 4 X lo7 M-' s-l for p-cyanophenol to 4 X lo9 M-' s-' for hydroquinone at pH 5.8.' This difference in selectivity is obviously owing to changes in the one-electron redox potential of the phenols associated with their acid-base equilibria (eq 3). At high pH the phenoxide ions examined have X C 6 H 4 0 H G XC6H40- + H+
(3)
redox potentials in the range of 0-1 V,2-3while in neutral solutions the redox potentials are higher by -0.5 V. The potential for the azide radical was estimated to be E(N3/N3-) = 1.9 V.4 This is high enough to make its reaction with any of the phenoxide ions practically diffusion controlled; only in neutral solutions does selectivity become apparent. To achieve high selectivity in the oxidation of phenoxide ions it is necessary to use radicals which are much weaker oxidants than N3. One such radical is SO3-, which was found to oxidize hydroquinone very rapidly at high pH but not p h e n ~ l . ~In , ~fact, with phenol it enters into the equilibrium $03-
C6H40- G
so?- + C6H40'
(4)
Visiting scientist from Ben-Gurion University, Beer Sheva, Israel.
0022-3654/86/2090-4156$01.50/0
with k4 = 6 X lo5 and k4 = 1 X lo7 M-' s - ' . ~ This indicates that SO3-is a weak oxidant (E(SO,-/SO?-) = 0.63 V vs. which will not attack many phenols. The radicals C102and N O 2 are known to be slightly stronger oxidants (E(C102/C102-) = 0.94 V and E(NO2/NO2-) = 1.03 V vs. NHE)6s7 that can oxidize phenoxide ions relatively slowly and p-methoxyphenoxide more rapidly.8 Therefore, it appears that radicals having a redox potential in the range of 1 V are good candidates for exploring the effect of substituents on rates of oxidation of phenoxide ions. We have examined the reactivity of CIOz and NOz with a variety of substituted phenols, attempted to measure the rate constants for SO3- with several additional phenols, and compared the results with those obtained previously with N3*and with other radicals. Experimental Sectiong
Sodium chlorite was obtained from Eastman, sodium nitrite was a Fisher Certified ACS reagent, and sodium sulfite was a Mallinckrodt Analytical reagent. The phenols were obtained from Aldrich. Water was purified by a Millipore Milli-Q system. Fresh solutions were prepared before each experiment. They were (1) Alfassi, Z. B.; Schuler, R. H. J . Phys. Chem. 1985, 89, 3359. (2) Steenken, S.; Neta, P. J . Phys. Chem. 1982, 86, 3661. (3) Huie, R. E.; Neta, P. J. Phys. Chem. 1984, 88, 5665. (4) Endiwtt, J. F. In Concepts of Inorganic Photochemistry, Adamson, A. W., Fleischauer, P. D., Eds.; Wiley-Interscience: New York, 1975; p 81. (5) Huie, R. E.; Neta, P. J. Phys. Chem. 1985,89, 3918. ( 6 ) Troitskaya, N. V.; Mishchenko, K.P.; Flis, I. E. Russ. J. Phys. Chem. 1959, 33, 77. (7) Wilmarth, W. K.; Stanbury, D. M.; Byrd, J. E.; Po,H. N.; Chua, C. P. Coord. Chem. Rev. 1983, 51, 155. (8) Huie, R. E.; Neta, P. J . Phys. Chem. 1986, 90, 1193. (9) Certain commercial equipment, instruments, or materials are identified in this paper in order to specify adequately the experimental procedure. Such identification does not imply recognition or endorsement by the National Bureau of Standards, nor does it imply that the material or equipment identified are necessarily the best available for the purpose.
0 1986 American Chemical Society
