J . Phys. Chem. 1990, 94, 2381-2386
2381
Oxidation of CS2 by Reaction with OH. 1. Equilibrium Constant for the Reaction OH 4- CS2 e CS20H and the Kinetics of the CS20H O2 Reaction
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Timothy P. Murrells, Edward R. Lovejoy, and A. R. Ravishankara* Aeronomy Laboratory, National Oceanic and Atmospheric Administration, Boulder, Colorado 80303, the Department of Chemistry and Biochemistry, and the Cooperative Institute for Research in Environmental Sciences, University of Colorado, Boulder, Colorado 80309 (Received: May 8 , 1989; In Final Form: October 6. 1989)
The reaction of OH with CS2 was studied in He (50 Torr), N2 (9-40 Torr), and SF, (30-60 Torr) over the temperature range 249-318 K by using pulsed laser photolysis to generate OH and pulsed-laser-inducedfluorescence to detect OH. The rapid approach of [OH] to equilibrium was observed at each temperature, indicating reversible formation of a CS20H adduct. Analysis of the [OH] temporal profiles provided the rate coefficients for the forward and reverse components of the reaction OH + CS2 + M 9 CS20H + M and, hence, the equilibrium constant. The temperature dependence of the equilibrium = -10.9 f 1.0 kcal mol-’ and the entropy change for the reaction, hS0298= constant yielded the heat of reaction, -24.0 f 4.4 cal K-’ mol-’. No evidence could be found for a direct non-adduct-forming channel for the OH CS2 reaction cm3 molecule-’ s-I. An enhancement by O2 of the OH loss rate in CS2 leading to k(OH + CS2 products) I2 X was observed, confirming the occurrence of a reaction between CS20H and 02.Analysis of the [OH] temporal profiles cm3 molecule-I s-’ which was independent of gave a rate coefficient for the CS,OH 0, reaction of (2.6 & 1.0) X temperature from 249 to 299 K.
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Introduction Carbon disulfide, CS2, is released by plants’ and by biomass burning2 and is present at parts per trillion levels in the Earth’s troposphere.3 The mechanism of the oxidation of CS2 in the atmosphere has been the subject of many investigations over the past decade. Detection of OCS in the troposphere4 and the stratosphereS and Crutzen’s hypothesis6that OCS is a major source of sulfur in the stratospheric aerosol layer led to a search for the origin of OCS. It was proposed that the CSz could be oxidized in the troposphere to form OCS. Photolysis was ruled out as a possible mechanism since absorption by CS2 in the near-UV region does not lead to dissociation.’ The reaction of O(3P) with CS2 is reasonably fast* but cannot be responsible for CS2 removal in the troposphere since the O(3P) concentration is very low. Therefore, the reaction of O H with CS2 was proposed as the major pathway for the degradation of CS2 and formation of OCS. Kurylo9 measured the rate coefficient for reaction 1, kl, to be 1.9 X cm3 molecule-I s-’ at 298 K using pulsed HzO photolysis
vantage of this structure and minimized CIS, photolysis by using a CS,-filled cell to filter out the radiation that dissociates CS2. Iyer and RowlandI3 determined the rate coefficient for OCS formation in reaction 1 to be less than 3 X cm3 molecule-’ s-I, at 298 K. These authors used a competitive method where the rate of formation of OCS was measured relative to the rate of loss of a reference molecule, c o , C3H8, or i-C4HI0. They generated OH via C W photolysis of H 2 0 2at 254 nm. Following irradiation, the product and reactant were separated and measured by radio gas chromatography. The results of Wine et al.12 and Iyer and Rowland13 were confirmed by the work of Leu and Smith14 and Biermann15 et al. Leu and Smith used a discharge flow apparatus and measured O H loss via resonance fluorescence and OCS production via mass spectrometry. The study of Biermann et al. with a discharge flow-photoionization mass spectrometer system was not aimed at measuring k l but did
OH + cs2-L HS + ocs (1) to produce OH which was detected by resonance fluorescence. He encountered problems due to secondary chemistry involving the photodissociation products of CS2 which he attempted to overcome by varying many of the experimental parameters. It was believed that he had minimized secondary chemistry contributions and the value of k l he measured was used in model calculations. Kurylo9 also proposed that HS and OCS, the only energetically allowed products considered at that time, are formed via a complex mechanism, since his measured rate coefficient was too fast for a four-centered reaction. The same year, Atkinson cm3 molecule-’ s-I at 298 K. et aLiOreported that k , < 7 X These authors also used the pulsed photolysis-resonance fluorescence method. However, most model calculations’’ utilized Kurylo’s value and considered reaction 1 as a sink of CS2 and a source of OCS. Wine et a1.lZ also used the pulsed photolysis-resonance fluorescence method and found that kl measured by following the disappearance of OH was less than 1.5 X cm3 molecule-’ s-I when CS2photolysis was greatly suppressed. CS2 has a very large absorption cross section in the 165-185-nm region where H 2 0 is photolyzed to generate OH. The CS2 absorption spectrum, unlike that of H,O, is highly structured.’ Wine et al. took ad-
(1) Goldan, P. D.; Kuster, W. C.; Albritton, D. L.; Fehsenfeld, F. C. J. Atmos. Chem. 1987, 5, 439 and references therein. (2) Andreae, M. 0. The Biogeochemical Cycling oJSulfur and Nitrogen in the Remote Atmosphere; Galloway, J. N., et al., Eds.; Reidel: Amsterdam, 1985, and references therein. (3) For example see: (a) Hanst, P. L.; Spiller, L. L.; Watts, D. M.; Spence, J. W., Miller, M. F. J. Air.Pollut. Control. Assoc. 1975, 25, 1220. (b) Torres, A. L.; Maroulis, M. J.; Goldberg, A. B.; Bandy, A. R. Trans. Am. Geophys. Union 1978, 59, 1082. (c) Aneja, V. P.; Overton, Jr., J. H.; Cupitt, L. T.; Durham, J. L.; Wilson, W. E. Nature 1979, 282, 493. (4) Maroulis, P. J.; Torres, A. L.; Bandy, A. R. Geophys. Res. Lett. 1977, 4, 510 and references therein. ( 5 ) Mankin, W. G.; Coffey, M. T.; Griffith, D. W. T.; Drayson, S. R. Geophys. Res. Lett. 1979, 6, 853. (6) Crutzen, P. J. Geophys. Res. Lett. 1976, 3, 73. (7) Okabe, H. Photochemistry of Small Molecules, Wiley: New York, 1978. (8) DeMore, W. B.; Molina, M. J.; Sander, S. P., Golden, D. M.; Hampson, R. F.; Kurylo, M. J.; Howard, C. J.; Ravishankara, A. R. “Chemical Kinetics and Photochemical Data For Use in Stratospheric Modeling”; Evaluation No. 8, JPL Publication 87-47; Jet Propulsion Laboratory: Pasadena, CA, 1987. (9) Kurylo, M. J . Chem. Phys. Lett. 1978, 58, 2 3 8 . (10) Atkinson, R.; Perry, R. A,; Pitts, Jr., J. N. Chem. Phys. Lett 1978, 54, 14. (11) (a) Sze, N. D.; KO, M. K. W. Nature 1979, 278,731. (b) Logan, J. A.; McElroy, M. B.; Wofsy, S. C.; Prather, M. J. Nature 1979, 281, 185. (12) Wine, P. H.; Shah, R. C.; Ravishankara, A. R. J . Phys. Chem. 1980,
‘Address correspondence to this author at NOAA, R/E/AL2, 325 Broadway, Boulder, C O 80303.
84, 2499.
(13) Iyer, R. S.; Rowland, F. S. Geophys. Res. Lett. 1980, 7, 797. (14) Leu, M. T.; Smith, R. H. J . Phys. Chem. 1982, 86, 958. ( 1 5) Biermann, H. W.; Harris, G. W.; Pitts, Jr., J. N. J . Phys. Chem. 1982, 86, 2958.
0022-3654/90/2094-238 1%02.50/0 0 1990 American Chemical Society
2382 The Journal of Physical Chemistry, Vol. 94, No. 6,1990 generate results in agreement with Leu and Smith. The mixing ratio of CS2 in the troposphere decreases with altitude which indicates that there are efficient sinks for CS2 ( s a , = 10 days).16 The measurements by Wine et al.I2 show that reaction 1 is too slow to account for the CS, loss. The reaction of electronically excited CS, with O2 was proposed.17 Jones et aL'* have measured the yield of OCS in this process and find that, on a global basis, it could account for a third of CS2 loss. More information is needed to confirm this mechanism. Measurements of k , in the presence of O2by Cox et aLI9 yielded a higher value. Cox et at. generated O H via photolysis of HONO and carried out a competitive measurement similar to that of Iyer and Rowland. The suspicion that O2was somehow altering the overall rate coefficient was confirmed by Cox's groupZoand Barnes et aLZ1 using C W photolysis in chambers and by Wine and Ravishankara22in a pulsed photolysis experiment. The observations from these studies clearly showed that k , increased with increased pressure in a constant concentration of 0, and increased with 0, concentration at a fixed total pressure. These results were interpreted as evidence for a complex reaction pathway where O H adds to C S 2 to yield a species C S 2 0 H which establishes an equilibrium with O H and CS,. The 0, enchancement of the O H loss was believed to be due to a reaction between C S 2 0 H and O2 CS,
+ O H s"lr, CSzOH kll
C S 2 0 H + O2
k3
products
(2f, 2r) (3)
where reactions 2f and 2r are dependent on pressure in the system. To check this hypothesis, it is necessary to show that indeed an adduct is formed, an equilibrium is attained, and the adduct reacts with 02.It is important to determine the values of the individual rate coefficients, k2f, k2,, and k3, as functions of temperature and have determine the products of reaction 3. Chamber studies18*20*21 shown that the ultimate products are mostly OCS and SO,; so, reaction 3 should lead to species that generate OCS and SO,. Our results, reported here and in the following paper, clearly prove the correctness of the above hypothesis. In this paper we describe the measurements of K, = k2f/k2rand k3 as a function of temperature. In the companion paper we describe identification and yields of the products. Since this work was initiated, Hynes et al.23and Bulatov et al.24 have carried out some measurements similar to those reported here. Hynes et al. measured the overall rate coefficients for the loss of O H in CS2 and air and equilibrium constants at various temperatures. Bulatov et al. just probed the equilibrium process at 298 K. The results from these studies will be compared with ours in the Discussion section.
Experimental Section The O H + CS, and C S 2 0 H + O2 reactions were studied by generating OH via pulsed laser photolysis and detecting it by pulsed-laser-induced fluorescence. A description of the apparatus used in the present study has been given el~ewhere,,~ so only those (16) Bandy, A. R.; Maroulis, P. J . ; Shalaby, L.; Wilner, L. A. Geophys. Res. Lett. 1981, 8, 1180. (17) Wine, P. H.; Chameides, W. L.; Ravishankara, A. R. Geophys. Res. Lett. 1981, 8, 543, and references therein. (18) Jones, B. M . R.; Cox, R. A.; Penkett, S. A. J . A m " . Chem. 1983, I , 65. (19) Cox, R. A.; Sheppard, D. Nature 1980, 284, 30. (20) Jones, B. M. R.; Burrows, J . P.; Cox, R. A.; Penkett, S. A. Chem. Phys. Lett. 1982, 88, 372. (21) Barnes, I.; Becker, K . H.; Fink, E. H.; Reiner, A,; Zabel, F.; Niki, H . Int. J . Chem. Kinet. 1983, 15, 631. (22) Wine, P. H.; Ravishankara. A . R. Presented at the Second Symposium on Composition of the Nonurban Troposphere, May 25-28. 1982, Williamsburg, VA. (23) Hynes. A . J . ; Wine, P. H.; Nicovich, J. M. J . Phys. Chem. 1988, 92, 3846. (24) Bulatov, V. P.; Cheskis, S. G.; logansen, A. A,; Kulakov, P. V.; Sarkisov, 0. M.; Hassinen, E. Chem. Phys. Lett. 1988, 153, 258
Murrells et al.
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details which are pertinent to this work are given here. The reaction vessel was a 150-cm3 Pyrex cell which was equipped with four optical ports for transmission of the two laser beams through the cell and a fifth port for observing the laser-induced fluorescence. Each laser beam entered and exited the cell through quartz windows mounted at Brewster's angle. The photolysis and probe laser beams intersected at right angles in the center of the cell. Fluorescence from the observation volume where the two laser beams intersect was detected through a fifth port orthogonal to the plane containing the laser beams. The gas flow was also perpendicular to both laser beams. The reaction cell was maintained to within f0.5 K of a desired temperature by circulating a thermostated fluid through its outer jacket. The fluid was ethylene glycol for temperatures higher than 298 K and ethanol for temperatures below 298 K. The temperature of the gas within the cell was measured by a chromelalumel thermocouple positioned close to the observation volume. The cell was housed within a chamber through which dry nitrogen was continuously flowed at a pressure of a few Torr. Hydroxyl radicals were produced by photolysis of H 2 0 2at 248 or 266 nm. A Nd:YAG pumped frequency-doubled dye laser was line of the (1-0) (A2Z+ X2n) transition used to excite the Q1(l) of O H at 281.9 nm. The resultant fluorescence from the (1-1) and (0-0) bands of the (A2Z+ X2n)transition was observed with a photomultiplier tube after it passed through a combination of lenses and a band-pass filter centered at 308.7 nm. The output pulse from the photomultiplier tube was sent either to a fast transient digitizer or to a gated charge integrator. The signal from the integrator, which is proportional to [OH], was transmitted to a computer. In general, the signal from 100 laser shots was averaged. The time dependence of [OH] was determined by changing the delay time between the photolysis laser pulse and the probe laser pulse, typically from 0.2 ps to 10 ms. The delay times were always much longer than the duration of the photolysis and probe laser pulses which were typically 10 ns. Both lasers were operated at a repetition rate of 10 Hz. The gas mixture containing the reagents and bath gas flowed through the reaction cell at a linear velocity of 10-20 cm s-I so that each photolysis laser pulse initiated chemistry in a fresh reaction mixture and the accumulation of products in the observation volume was avoided. The energies of the photolysis laser pulses were monitored with a laser power meter and were constant ( 3 5 % ) over the time it took to record an O H concentration-time profile. The energy of the probe laser was not monitored, but the reproducibility of the measurement of O H signal at any delay time during an experiment was better than 5% so that it was not necessary to normalize the signal to laser energy. In some experiments using 248-nm photolysis, the OH signal was normalized by using the signal at 200 ns which was measured frequently throughout the experiment. The photolysis laser fluence was varied from 5 to 15 mJ cm-2 at 248 nm and from 10 to 30 mJ cm-, at 266 nm giving initial O H concentrations [OH],, in the range (5-20) X 10" molecules ~ m - ~CS2 . absorbs weakly at these wavelengths26 and in these experiments the maximum amount of CS2 photolyzed ranged from 1% to 14% of [OH],. A small amount of OH (, [CS,] = 1.8 X 10I6 molecules c ~ n - ~The . lines are calculated profiles using kzr= 3.8 X cm3 moleculed s-I, k2r = 5200 s-I, and k3 = 2.7 X cm3 molecule-' s-l. (0)[O,] = 0, (0) [O,] = 4.2 X 10l6molecules cm-!
-24.0 f 4.4cal K-' mol-', and AHf0298(CS20H)= 26.4 f 1.1 kcal mol-'. The existence of a non-adduct-forming channel for the OH CS2 reaction was evaluated by examining the dependence of k4 on [CS,]. In these studies [H202]and [MI were held constant. No change in k , was observed at 299 K as [CS,] was increased from 0 to 3.3 X 1OI6 Assuming a minimum detectable variation in k4 of 50 s-I, this observation gives an upper limit for k(OH CS2 products) I 2 X lo-" cm3 molecule-' s-'. 2. The CS20H+ 0,Reaction. The rate of O H decay in the O H + CS2 system at longer times was increased by the addition of 02.This observation is consistent with the removal of the CS20H adduct by reaction with 0,: CS20H + O2 products (3) Figure 3 shows typical O H temporal profiles with and without O2added in 50 Torr of He at 274 K. O H temporal profiles were recorded at different [O,] in 50 Torr of He at 299, 274, and 249 K. Removal of C S 2 0 H by reaction 3 is equivalent to an en-
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The Journal of Physical Chemistry, Vol. 94, No. 6, 1990 2385
Oxidation of CS2 by Reaction with O H
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/
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kl
I
I
profiles are subject to considerable error in studies of this nature. Hence the values are not necessarily accurate and may appear as disagreements in such comparisons. Bulatov et al.24have recently measured k2f and kzr at room temperature in 150 Torr of Ar. They used a time-resolved laser-induced fluorescence experiment and obtained an equilibrium constant (=kzr/k2,) of 2.6 X lo-'' cm3 molecule-' which agrees with our results at 295 K. Bulatov et al. also measured a rate cm3 molecule-' s-' for the reaction coefficient of (9 f 3) X C S 2 0 H 03. Our upper limit rate constant for the direct, non-adduct-forming products, is k I2 X cm3 molecule-' channel, O H CS, s-I. Other upper limits measured with CS2 photolysis suppressed, to 7 X cm3 molecule-' s-l at 298 range from 1.5 X K.I2-l4 All of the studies based their measurements of k, on the rate of O H decay after equilibrium had been e ~ t a b l i s h e d . ' ~ , ' ~ J ~ However, this procedure underestimates the rate coefficient because decomposition of C S 2 0 H constantly replenishes OH. The errors in the reported values are greatest for the studies conducted at low temperature and using high [CS,] because a greater fraction of the initial O H is converted to CS20H. In most of the previous measurements at 298 K, the error is small because only -3% of the initial OH is converted to C S 2 0 H . However, in the study of Wine et al.'* the fraction of the initial O H converted to C S 2 0 H was up to 20% at 297 K and 60% at 251 K. Our measurement is not subject to this source of error. However, the Wine et al. room temperature value is still the lowest upper limit (value corrected for a fraction of OH tied up as CSzOH is 51.8 X cm3 molecule-' s-' at 297 K). Our results indicate that the C S 2 0 H 0, reaction has very little, if any, temperature dependence between 249 and 299 K. This conclusion was also reached by Hynes et al. who used a different approach to determine k3. In our study, less than 8 Torr of 0, was added to the O H CS2 equilibrium mixture. Therefore, k3[02]was usually less than or comparable to kzr and the initial rapid approach of O H to equilibrium could always be observed separately from the subsequent decay of O H due to reaction of CS20H with 0,. In contrast, Hynes et al. added at least 140 Torr of 0, and single-exponential decays of O H were observed because k3[0,] >> k,,. From the dependence of the effective bimolecular rate constant for the O H CS2 reaction on [O,], they extracted an average, temperature-independent (25 1-348 K) value for k3 which is in good agreement with the values of (2.9 f 1.1) X measured in our work. The Hynes et al. values for k3 at 140 and 280 Torr lie within their range of values at 700 Torr. Furthermore, the very similar values for k3 which we obtain in 50 Torr of He suggest that k3 is independent of pressure. There are several exothermic O H producing channels for reaction 3 which are discussed in the following paper. Our measurements of k3 reflect only those channels not producing OH. An OH-producing channel would appear in the fits as an increase in k,, as [O,] increased. No systematic dependence of k2, or K , on [O,] was found, even at low temperature when the fits are most sensitive to changes in k2,. Examination of our data at 249 K reveals that less than 30%of the CS20H + 0, reaction could lead to O H regeneration. Hynes et aLz3 also concluded that O H regeneration is quite small. Chamber studies'8~20~21 have shown that OCS and SO2are the ultimate products of CS2 oxidation by O H in air. However, in order to fully assess the implications that the O H + CS2 and CS20H 0, reactions have to atmospheric models, it is necessary to know how these products are formed, what other intermediate species may be produced, and with what yield. These topics are dealt with in the following paper. In conclusion, this work has clearly demonstrated that an equilibrium is established in the OH + CS, reaction. The equilibrium constant has been measured over the temperature range 249-3 18 K. The temperature dependence of K, yields an enthalpy change of 10.9 f 1 .O kcal mol-' for the CS2-OH bond formation and AHf02ss(CS20H)= 26.4 1 . 1 kcal mol-I. The enhancement in the loss rate of O H in the presence of CS2 by 0, has been confirmed and accurately modeled by assuming a reaction
+
+
0
5
c o2 I (
10
I5
20
1 0 ' molecule ~ cm-3 1
Figure 4. First-order CS20H loss rate constant, k,, as a function of O2 concentration. T = 274 K,P = 50 Torr (He). Line is a weighted linear least-squaresfit to the data. The slope of the line yields k,. Error limits are 2u based on nonlinear least-squares fits for k,, the first-order rate coefficient for the removal of CS20H.
hancement in k5 in eq I provided [O,] >> [OH],. In the present experiments [O,]/[OH], > lo4 at all times. Values of k5 were obtained by fitting the temporal profiles to eq I, with k2f,kzr,and k4 fixed at their values derived from experiments without 0,. The bimolecular rate coefficient for reaction 3, k3, is related to kS by the equation k5 = k3[02] + k' where k'refers to all first-order loss processes for C S 2 0 H other than reaction with 0,. Figure 4 shows a plot of kS versus [O,] at 274 K. Linear least-squares analyses of the k5 versus [O,] data sets at each temperature gave values for k3 of (2.3 f 0.6) X cm3 molecule-' s-l at 299 K, (2.9 f 0.3) X cm3 molecule-' s-I at 274 K, and (2.7 f 0.7) X cm3 molecule-' s-I at 249 K. The errors are 20 precision limits. Measurements of rate coefficients by curve fitting are subject to considerable systematic errors. We believe we have minimized these errors by using a wide range of reactant concentrations and operating under conditions where the two components of O H decay are separated. We estimate that the systematic errors are less than *lo% and cm3 report the following rate coefficients: (2.3 f 0.8) X cm3 molecule-' s-' molecule-' s-' at 299 K, (2.9 f 0.6) X at 274 K, and (2.7 f 1.0) X cm3 molecule-' s-' at 249 K. The intercept on the k5 axis of each plot gives k'which showed no systematic dependence on temperature. The mean value of k'was 200 f 200 s-'. The same values of k3 were obtained when k2fand kzr were variable parameters in the fitting procedure. The resulting values of K , = k2f/k2rwere also in very good agreement with those measured without 0, present.
Discussion Recently, Hynes et ai.23measured slightly lower values of K , at every temperature and obtained a smaller enthalpy change for reaction 2 (9.9 f 1.2 kcal mol-'). The major difference between the two studies is the range of pressures used: Hynes et al. studied the equilibrium in the presence of 65-690 Torr of N, or He, whereas our study used 9-60 Torr of N2,SF6,or He. K, is not dependent on pressure, but the forward and reverse rate constants, k2fand k2r,and hence the rate of approach to equilibrium (k2f[CS2] k,,), increase with increasing pressure. The accuracy of the klf and kzr measurements depends on obtaining a complete O H temporal profile during the approach to equilibrium. At low pressure the equilibrium is slower and is easier to observe. However, at low pressure it is important to verify that the CS20H adduct is thermally equilibrated with the bath gas. That K, does not depend on pressure and on the nature of the bath gas suggests that the C S 2 0 H adduct was thermalized in our experiments. It is not possible to compare the individual values of the forward and reverse rate coefficients, k2f and k2,, measured in the two studies because of the different pressures used. Furthermore, as already pointed out, kZfand kzr determined from fitting the O H
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J . Phys. Chem. 1990, 94, 2386-2393
2386
of the C S 2 0 H adduct with 02.The results yield a temperature cm3 independent (249-299 K ) value of (2.6 f 1.0) X molecule-' s-I for the rate coefficient of the CS20H + O2reaction.
Utter for assistance with the AMPAC calculations. This work was supported by NOAA as part of the National Acid Precipitation Assessment Program.
Acknowledgment. We thank Lester Lambert for carrying out the CS2 absorption cross section measurements and Dr. R. G .
Registry No. O H , 3352-57-6; CS,, 75-15-0; CS,OH, 123132-54-7; 02,7782-44-7.
Oxidation of CS, by Reaction with OH. 2. Yields of HO, and SO, in Oxygen Edward R. Lovejoy, Timothy P. Murrells, A. R. Ravishankara,* and Carleton J. Howard Aeronomy Laboratory, National Oceanic and Atmospheric Administration, Boulder, Colorado 80303, the Department of Chemistry and Biochemistry, and the Cooperative Institute for Research in Environmental Sciences, University of Colorado, Boulder, Colorado 80309 (Received: May 8, 1989; In Final Form: October 6, 1989)
The products of the OH-initiated oxidation of CS2 have been investigated. Analysis of the OH loss and the HS production using a laser magnetic resonance (LMR) discharge-flow apparatus yielded k l C 3 X cm3 molecule-' s-I for the reaction OH + CS2 HS + OCS (1) (340 K and 5 Torr (He)). The oxidation of CS2is enhanced by O2due to the following mechanism: OH + CS2 2 CS20H (2), and CS20H + O2 -,products (3). H 0 2 was identified by LMR detection as a major product from this chemistry. The H 0 2 yield was measured in a pulsed photolysis experiment by modeling OH temporal profiles with NO added to convert H 0 2 to OH. The H 0 2 yield was 95 i 15% of the OH consumed by reactions 2 and 3 (50 Torr (He), 249 and 299 K). Discharge-flow experiments employing chemical ionization mass spectrometric (CIMS) detection of OH and SO2 showed that 90 & 20% of the OH consumed by reactions 2 and 3 leads to SO2 production.
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Introduction This is the second of two papers on the oxidation of CS2 by OH. The first paper presents a study of the kinetics of the initial CS2oxidation steps.' This paper is concerned with the identification and quantification of the products of the CS2 oxidation. Recent work indicates that the metathesis reaction OH + CS2 HS + COS is slow, k I2 X cm3 molecule-' s-I at 298 K,'s2 and that O H reacts with CS2predominantly by forming an adduct which decomposes back to reactants.'q3
-
OH
+ CS2
kX
CS2OH
kb
The radical products and the elementary steps following reaction 3 have not been identified. There are many exothermic pathways for reaction 3 (see Scheme I, where AH:298 values are given after the reaction). SCHEME I CS2OH 02 H02 CS20 ? (3a) H 0 2 + OCS S -58 kcal mol-' (3b)
+
------
(2)
The enthalpy change for reaction 2 is approximately -1 1 kcal mol-' and the lifetime of C S 2 0 H is short (lifetime = 4.5 ps at 298 K and 661 Torr of N2).3 The CS20H adduct reacts with O2 C S 2 0 H + O2
k3
products
(3)
with a rate coefficient of k3 = 3 X cm3 molecule-' s-'.'s3 The reaction sequence, (2) followed by (3), is a major loss process for CS2 in the Earth's atmosphere. The rate coefficient for the loss of O H by reaction with CS2 in O2 has been measured under tropospheric condition^^-^ ( k ( 0 H CS,) = 2 X cm3 molecule-' s-l, 295 K and 700 Torr) and predicts an atmospheric lifetime which is consistent with estimates based on atmospheric measurements of CS, (7 = 10 days).6 Many studies on the products of the O H + CS2 reaction in the presence and the absence of O2 have been performed. Iyer and Rowland' observed OI4CS formation in the OH + I4CS2system but they could not rule out reactions other than OH + CS2 as the source of OCS. Leu and Smith* also detected OCS in the OH CS2 system using an electron impact ionization mass spectrometer coupled to a discharge-flow apparatus. Their experiments were complicated by wall reactions so they could not unambiguously identify the OCS source. HS production has not been observed. Chamber studies have shown that the end products of the oxidation of one CS2 in air are one SO2 and one OCS.435
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*Address correspondence to this author at NOAA, R/E/AL2, 325 Broadway, Boulder, CO 80303.
0022-3654/90/2094-2386$02.50/0
--
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+
+ + C O + S2 -19 kcal mol-' O H + SO2 + C S -23 kcal mol-' O H + SO + OCS -56 kcal mol-] HOCS + SO2 ? H + OCS + SO, -79 kcal mol-' H S 0 2 + OCS -1 13 kcal mol-' HOSO + OCS -154 kcal mol-' HCO + SO2 + S -22 kcal mol-' HSO + CO + SO -52 kcal mol-' HSO + C 0 2 + S -55 kcal mol-' HS + SO2 + CO -89 kcal mol-' HS + SO + C 0 2 -85 kcal mol-' H02
adduct
?
(3c) (3d) (3e) (30 (3g) (3h) (3i) (3j) (3k) (31) (3m) (3n) (30)
(1) Murrells, T. P.; Lovejoy, E. R.; Ravishankara, A. R. J . Phys. Chem., previous paper in this issue. (2) Wine, P. H.; Shah, R. C . ;Ravishankara, A. R. J . Phys. Chem. 1980, 84, 2499. (3) Hynes, A. J.; Wine, P. H.; Nicovich, J. M. J . Phys. Chem. 1988, 92, 3846. (4) Barnes, I.; Becker, K. H.; Fink, E. H.; Reimer, A.; Zabel, F.; Niki, H. Int. J. Chem. Kinet. 1983, 15, 631. (5) Jones, B. M. R.; Cox, R. A,; Penkett, S.A. J . Atmos. Chem. 1983.1, 65 _.
(6) Bandy, A. R.; Maroulis, P. J.; Shalaby, L.; Wilner, L. A. Geophys. Res. Lett. 1981, 8, 1180. (7) Iyer, R. S.; Rowland, F. S. Geophys. Res. Lett. 1980, 7 , 191. (8) Leu, M. T.; Smith, R. H . J . Phys. Chem. 1982, 86, 958.
0 1990 American Chemical Society
