3846
J. Phys. Chem. 1988, 92, 3846-3852
Kinetics and Mechanism of the Reaction of OH with CS, under Atmospheric Conditions A. J. Hynes,* P. H. Wine,* and J. M. Nicovich Molecular Sciences Branch, Georgia Tech Research Institute, Georgia Institute of Technology, Atlanta, Georgia 30332 (Received: October 1 , 1987; In Final Form: January 7 , 1988)
A pulsed laser photolysis-pulsed laser induced fluorescence technique has been employed to study the reaction of OH with CS, in the presence of He, N,, air, and 0, buffer gases and over the temperature range 250-350 K. In the absence of O,, evidence for rapid, reversible formation of a CS20H adduct is observed. Analysis of observed OH temporal profiles gives the forward and reverse rates of adduct formation and hence the equilibrium constant. A heat of reaction of -9.9 f 1.2 kcal mol-' is obtained from the temperature dependence of the equilibrium constant. The temperature dependence of the bimolecular rate coefficient for the forward addition reaction in 680 Torr of N2 + O2 is well-represented by the Arrhenius expression 6.9 X exp( 1150/T) cm3 molecule-l s-'. A rapid reaction of OH with CS, is observed in the presence of 02,confirming the observations reported in previous competitive rate studies. The observed bimolecular rate constant (kobsd) for the OH + CS2 reaction at 295 K in 700 Torr of air is found to be (1.5 f 0.1) X lo-', cm3 molecule-' s-I; koM increases dramatically with decreasing temperature. All experimental observations are consistent with a simple three-step reaction mechanism: adduct formation followed by adduct decomposition in competition with an adduct + 0,reaction. Analysis of our data using a steady-state approximation based on the above three-step mechanism leads to a rate coefficient of (2.9 f 1.1) X cm3 molecule-' s-I for the adduct + O2 reaction.
Introduction A detailed understanding of the mechanism of CS, oxidation is important in order to assess the role of this species in the global sulfur cycle. It is also of considerable practical importance since two of the potential oxidation products, COS and SO,, have a major environmental impact via their roles as greenhouse and pollutant gases. Reported atmospheric concentrations of CS2are highly variable,' ranging from 0 to 20 parts per trillion (ppt) in the free troposphere to 100-200 ppt in polluted air; the CS2vertical profile suggests a relatively short atmospheric lifetime. Laboratory studies of the reaction O H + CS, products (1)
-
present a somewhat contradictory picture of the role of the hydroxyl radical in CS, oxidation. All direct studies of reaction 1 have been performed in the absence of the potentially reactive gas 02. Both flash p h o t ~ l y s i s ~and - ~ discharge flow studies5v6 indicate that k, is extremely slow. Similar results are obtained in competitive studies in the absence of 02.'However, in 1 atm of air, competitive rate studiess-I0 report the Occurrence of a rapid reaction with an apparent bimolecular rate constant of =2 X lo-', cm3 molecule-' s-l. Adduct formation followed by adduct reaction with O2 in competition with adduct decomposition has been invoked to explain these observations.8-10 OH CS2 M s C S 2 0 H M ( l a , -la)
+
+ C S 2 0 H + O2
-
+
products (2) In the direct time-resolved studies no evidence was obtained for adduct formation on the millisecond time scale of the experiments.," Using techniques similar to those employed in a recent study of the O H + dimethyl sulfide reaction," we have studied reaction 1 in O,, 02/N2,N,, and H e buffer gases as a function of pressure and temperature. We have been able to directly confirm the observations of the competitive studiess-I0 on the "02rate enhancement". We have also been able to directly observe rapid, reversible adduct formation between OH and CS, and to determine the forward and reverse rates and, hence, the equilibrium constant as a function of temperature. The observed behavior in the presence and absence of oxygen represents a totally consistent picture and enables us to obtain rate constants for the elementary processes of adduct formation, decay, and reaction with O2and also to determine effective rate constants for reaction 1 under atmospheric conditions for use in modeling calculations. *Authors to whom correspondence may be addressed.
0022-3654/88/2092-3846$01.50/0
Experimental Section The PLP-PLIF apparatus has been described in detail elsewhere." Modifications and a brief review of its operation are given below. The reaction cell was constructed of Pyrex with an internal diameter of 2.5 cm and a length of 40 cm. Four mutually perpendicular side arms, 2.5 cm in diameter and 3 cm in length, were attached to the center of the cell. The photolysis laser passed through two of the side arms across the direction of gas flow while the probe laser passed along the length of the cell. Fluorescence was detected through a third side arm, perpendicular to the photolysis and probe beams. The central 25-cm length of the cell was jacketed to permit the flow of heating or cooling fluid from a thermostated bath. A copper-constantan thermocouple with a stainless steel jacket was inserted into the reaction zone through a vacuum seal, allowing measurement of the gas temperature under the precise pressure and flow conditions of the experiment. O H was produced by pulsed excimer laser photolysis of H202 at 248 nm; the typical photolysis energy was 30 mJ cm-,. OH was detected by pulsed laser induced fluorescence using a Nd:YAG pumped, frequency doubled dye laser. The dye laser output (bandwidth 0.001 nm, energy 1 mJ pulse-]) excited the Ql 1, Q1l', and R,3 lines of the (A2Z+-X211) 1-0 band at ~ 2 8 2nm. Fluorescence in the 0-0 and 1-1 bands was detected by an EM1 98 13QB photomultiplier after passing through collection optics and filters to discriminate against scattered light. In some experiments, the photomultiplier was gated off for the duration of the excimer laser pulse by using an EM1 GBlOOlA pulser. (1) Khalil, M. A.; Rasmussen, R. A. Atmos. Enuiron. 1984, 18, 1805. (2) Kurylo, M. J. Chem. Phys. Lett. 1978, 58, 239. (3) Atkinson, R.; Perry, R. A.; Pitts, J. N., Jr. Chem. Phys. Left. 1978, 54, 14. (4) Wine, P. H.; Shah, R. C.; Ravishankara, A. R. J . Phys. Chem. 1980, 84, 2499. ( 5 ) Leu, M. T.; Smith, R. H. J . Phys. Chem. 1982,86, 958. ( 6 ) Biermann, H. W.; Harris, G. W.; Pitts, J. N., Jr. J . Pbys. Chem. 1982, 86, 2958. (7) Iyer, R. S.; Rowland, F. S. Geophys. Res. Lett. 1980, 7 , 797. (8) Jones,B. M. R.; Burrows, J. P.; Cox, R. A,; Penkett, S. A. Chem. Phys. Lett. 1982, 88, 372. The authors have asked us to point out that ref 8 and 9 contain typographical errors in the labeling of k' in Table I (ref 8) and Table 3 (ref 9). The units of k' should be s-'. (9) Jones, B. M. R.; Cox, R. A.; Penkett, S. A. J. Atmos. Chem. 1983,1, 6.. 5.
(10) Barnes, I.; Becker, K. H.; Fink, E. H.; Reiner, A,; Zabel, F.; Niki, H. Int. J. Chem. Kinet. 1983, 15, 631. (1 1 ) Hynes, A. J.; Wine, P. H.; Semmes, D. H. J . Phys. Chem. 1986,90, 4148.
0 1988 American Chemical Society
Kinetics of the OH
+ CS2 Reaction
The Journal of Physical Chemistry, Vol. 92, No. 13, 1988 3841
The photomultiplier output was appropriately terminated (50-5000 R depending on required time resolution) and fed to a 100-MHz waveform analyzer to obtain the average peak voltage for (typically) 100 laser shots. The voltage when both lasers fired minus the voltage when only the probe laser fired was proportional to the OH concentration in the reaction zone. In order to minimize jitter between the photolysis and probe lasers, the excimer laser, Nd:YAG laser flashlamps, and Nd:YAG laser Q-switch were triggered by the three channels of a digital delay generator (California Avionics 103DR) which was itself triggered at 10 Hz. In this configuration, the jitter between the photolysis laser (20-11s pulse width) and probe laser ( 6 4 s pulse width) was less than 10 ns. The minimum increment in delay was 100 ns. Signal was collected for a range of delays varying from 0 to 40 ms. In kinetic runs typically 10-15 delays were sampled to map out an OH profile over two to three l / e times, while in equilibration runs typically 30-40 delays were sampled in order to accurately profile both the fast approach to equilibrium and the subsequent slow decay of the OH signal. All experiments were carried out under “slow-flow” conditions. The linear flow rate through the reactor was (typically) 10 cm s-l, and the laser repetition rate was 10 Hz. Since photolysis was across the direction of flow, a “fresh” reaction mixture was available for each pulse. CS2 concentrations were measured directly both before entering the reactor (using either a 4- or 17.3-cm absorption cell, the 213.9-nm Zn resonance line, and a monochromator/photomultiplier combination) and after exiting the reactor (using a 60.3-cm absorption cell, the 313.3-nm Hg line, and an interference filter/photomultiplier). Required absorption cross sections were measured as part of this study and are (in units of cm2) 36.0 at 213.9 nm and 0.70 at 313.3 nm. The pure gases used in this study had the following stated 99.99%. Air was zero grade minimum purities: N,, 99.999%; 02, O, and air were with total hydrocarbons less than 1 ppm. NZ, used as supplied. CS, (Aldrich Gold Label) had a stated purity of 99+%; before use, it was purified by trap-to-trap distillation (210-77 K) followed by repeated degassing at 77 K. H202was 90% pure by weight (10% H 2 0 ) ; it was further concentrated by bubbling N, through it for several days before experiments were undertaken and constantly during the course of the experiments.
Results and Discussion Observation of the Approach to Equilibrium. If the initial step in the OH CS2 reaction is the reversible formation of an adduct, then direct observation of the equilibration process should be possible in the absence of 02.Analysis of OH temporal profiles, then, allows the forward and reverse rates of adduct formation to be determined directly. The rate equations for the reaction sequence (la), (-la), (4), (5) can be solved analytically as long as pseudo-first-order conditions are obeyed (i.e., [OH]
