J. Phys. Chem. 1985,89, 5505-5510
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Kinetics of SH with NO2, Oar Op, and H202 Randall R. Friedl,? William H. Brune, and James G. Anderson* Department of Chemistry and Center for Earth and Planetary Physics, Harvard University, Cambridge, Massachusetts 021 38 (Received: May 6, 1985) A low pressure (1-8 torr of He) discharge flow reactor with (a) LIF detection of SH using a high repetition rate (20 kHz) metal atom laser to circumvent severe predissociation of the A22+ state and (b) resonance fluorescence detection of OH has been used to examine the kinetics of the title reactions at 298 K: S H + NOz HSO + NO, kl = (3.0 0.8) X lo-’’ cm3 s-I; S H O3 HSO 02,k2 = (3.2 f 1.0) X cm3 s-l; S H + O2 SO + OH, k3 < 1 X lo-’’ cm3 s-l; SH + HzOz HZS HO,, k4a; SH + H202 HSOH + OH, k4b; SH H2Oz HSO + HZO, k,. k4 = kha + k 4 b + k , < 5x cm3 s-l. Regeneration of SH in reaction 2 by HSO + O3is observed and is used to infer a rate constant for HSO O3 products of (1.0 0.4) X cm3 s-I. Absence of OH production in that reaction implies that the primary products are SH 202. Isotope experiments with H2S replaced by DzS substantiate that conclusion and yield a reaction rate constant for DSO O3 SD 20, of 9 X cm3 s-l. Brief discussions of SH reactivity compared with OH and Br are offered, as well as a summary of the atmospheric chemistry of SH.
-+
+
-
+ +
-
+
+
-
+
+ NO2 SH + O3
-
-+
SH
+0 2 -
SH
+ Hz02
HSO
+ NO
HSO + 02
(2)
SO + OH
(3)
products
(4)
-
*
+
+
Introduction The identification of acid deposition as one of the major environmental concerns of the present time has led to an increased interest in chemical processes involving sulfur-containing compounds. Despite this interest, current progress toward an understanding of earth’s atmospheric “sulfur cycle” is restricted, in part, by a lack of kinetic information for several radical species, most notably S H , HSO, and HOSOZ. The SH radical, in particular, is known to be formed in the troposphere by the oxidation of H2Sby OH’J however, its subsequent fate is virtually unknown. For modeling purposes, SH reaction rates have been estimated by analogy with OH radical kinetic^,^ a method which is only suggestive and, from a detailed molecular viewpoint, highly inadequate. In the present study, we have used laser-induced fluorescence detection of SH, resonance fluorescence detection of OH, and the discharge-flow technique to obtain room temperature kinetic data for various SH reactions of potential atmospheric importance, namely
SH
--
Although reactions 1-4 have all received attention in the literature, this attention has mainly succeeded only in an identification of the products of reactions 1 and 24,5and in a provocative discussion of the potential atmospheric importance of SH react i o n ~ .In ~ terms of kinetic data, an absolute rate constant has been reported only for reaction 1 [(3.5 f 0.4) X lo-” cm3 molecule-’ s-I] .4 Two recent flash photolytic studies conducted by Black4 and Tiee et aL6 have detected no reaction between SH and 02. However, the upper limits derived for k3 from these experiments, 4 X lo-’’ and 3.2 X cm3 molecule-’ s-l, are insufficient to preclude Ozas an important SH oxidant. An indirect experiment conducted by Becker et al.’ on HzS/03mixtures indicates that reaction of SH with O3 is relatively rapid. However, a recent attempt by Black4 to directly observe this reaction using the flash photolytic technique was unsuccessful due to an apparent heterogeneous reaction between H2S and O3 which produced SH. In this paper we present kinetic measurements for reactions 1-4 as well as observations relevant to the reaction between HSO and 03.In addition, the absolute rate constants derived in this work are used to speculate on the fate of atmospheric SH and to demonstrate a similarity between the reactivities of SH and Present address: Jet Propulsion Laboratory, 4800 Oak Grove Drive, Pasadena, CA 91 109.
0022-3654/85/2089-5505$01.50/0
Br, a finding which correlates with other thermochemical and electronic similarities between these two species.8
Experimental Section Experimental Apparatus and Chemical Sources. The apparatus used in these experiments is shown in Figure 1 and is similar to one described previo~sly.~Reactions of interest were initiated in a Teflon-coated, 60-cm-long, 2.5-cm-i.d. Pyrex “reaction zone” section by addition of the excess reactants, Oz,03,H202,and NO2, through a sliding Pyrex injector tube to flowing mixtures of trace SH (< 3 X 10” ~ m - ~and ) excess He ((3-25) X 10I6 ~ m - ~ ) . Reaction times were varied by movement of the injector and were calculated from a knowledge of the distance from injector tip to detection axis and the bulk flow velocity (500-1300 cm s-l) produced by the mechanical pump (Welch Model 1397). SH was monitored by laser-induced fluorescence (LIF) using an in-house designed and built copper-vapor pumped doubled dye laser system for excitation of the XzII A2Z bandhead at 323.7 nm. Average UV powers as high as 8 m W were obtained with this laser system but, due to the high laser repetition frequency (- 17 kHz), peak powers were only about 20 W and saturation of the SH absorption profile was of no concern even at the highest average powers. A second detection axis was used for resonance fluorescence (RF) detection of OH. Calibration of this axis was accomplished by addition of excess H atoms to measured amounts (counts of NO,. R F sensitivity was approximately 3 X s-’)/(molecule ~ m - with ~ ) a scattered light signal of -25 counts S-I, giving an OH detection limit for integration times over 10 s of 5 X 10" ~ m - and ~ , [O,] = 8 X 10l6~ m - no ~ ,OH was observed ). based upon the absence of O H ([OH], < 1 X lo8 ~ m - ~ Thus, and upon eq 9, we obtain an upper limit for reaction 3 of k3 < 1 X cm3 molecule-' s-l which is in agreement with the results of the previous investigations by Black4 and Tiee et al.s and decreases by a factor of four the lowest upper limit previously obtained. No further improvement in the determination of k3 was afforded by lowering [H2S], since small quantities of O H were produced by the sequence H + O2 + M H 0 2 M, H H02 2 0 H , as confirmed both experimentally and by a computer model which included these two reactions along with the relevant sulfur reactions. In contrast to reaction 3, the lack of SH decay in the presence of H202cannot be attributed to any potential regeneration of SH and does in fact indicate the absence of reaction. Among the exothermic channels available to this reaction SH H202 H2S H 0 2 AH = -1.7 kcal/mol (4a)
-
+
--
+ HSOH + OH HSO + H 2 0
+
+
AH = -20 kcal/mol AH = -45 kcal/mol
-
(4b) (4c)
the only possible SH regeneration sequence, reactions 4b and 7, can be dismissed because of the preferential depletion of OH by reaction with H202(at least at large [H202]). As a check of this conclusion, an attempt to detect OH was made, and, as expected, no OH was detected. Thus, based upon the experimental observations, we obtain as an upper limit for the sum of the possible reaction channels cm3 molecule-' s-l k4 = (k4a klb k k ) < 5 X
+
+
This upper limit is orders of magnitude smaller than the reported rate constant for the reaction of OH with H202but similar to the upper limit reported for the analogous Br rection. A lower limit to the activation energy for reaction 4a may be estimated by taking the experimentally determined A factor for the analogous OH reaction ( A = 3.0 X 10-l2cm3 s-I).I6 This results in an activation energy of >3.8 kcal. The fact that both SH and Br do not abstract an H atom from H202as readily as does OH can be conveniently rationalized in (16) NASA/JPL 'Chemical Kinetic and Photochemical Data for Use in Stratospheric Modeling, Evaluation No. 6"; Jet Propulsion Laboratory: Pasadena, CA, 1983; JPL 83-62.
Figure 7. A schematic diagram of SH chemistry relevant to the atmosphere. Solid and broken lines denote dominant and minor reaction pathways, respectively.
terms of product bond strengths (Do(HS-H) = 90.2 kcal/mol, Do(Br-H) = 86.5 kcal/mol, and Do(HO-H) = 118 kcal/mol). Correlations between product bond strengths (or reaction enthalpies) and activation energy have been found for a variety of other hydrogen-abstraction reaction sets; for example, OH + RH4, 0 RH,, and CH3 RH4.I7 One relatively successful empirical relation between bond energy and reaction barrier has been expressed in the bond energy-bond order (BEBO) model developed by Johnston and Parr.I8*l9 In Figure 6 , BEBO predictions based upon eq 10, where D i B is the product bond energy, E A C T is the
+
+
EACT
(1
=
+
%)( [ + 1-
1
(1
+
31'""1" (10)
energy of activation, AH,,, is the enthalpy of reaction, and p is an empirical parameter which is greater than 1, are compared with experimentally determined activation energies for reactions of various atomic and diatomic radicals with H202. As seen in Figure 6 , the BEBO model is capable of producing the observed trend in activation energy, especially if the experimental value of E A C T for 0 H202is halved as suggested by Mayer and SchielerZ0to account for the extra valence electron on the oxygen atom. While the success of the simple BEBO theory for H 2 0 2 reactions is encouraging, its applicability is apparently restricted only to H-atom abstractions. A comparison of the BEBO model with a set of experimental data for radical-ozone reactions gives virtually no agreement.
+
Summary and Conclusions In this paper, room temperature rate constants have been presented for the reactions of SH with O3and NOz. Both reactions are facile and presumably proceed with little activation energy. In addition, upper limits have been determined for the reaction rates of SH with O2 and HzOz. Both of these reactions are immeasurably slow i n the present apparatus. With respect to H 2 0 2and 03,the reactivity of SH has been found to be similar to that of Br and not to that of OH. Coupling this observation with the fact that SH and Br possess similar ~~
(17) Krech, R. H.; McFadden, D. L. J. Am. Chem. SOC.1977, 99,8402. (18) Johnston, H. S.; Parr, C. J . Am. Chem. SOC.1963, 85, 2544. (19) Johnston, H. S. "Gas Phase Reaction Rate Theory"; Ronald Press: New York, 1966. (20) Mayer, S. W.; Schieler, L. J . Phys. Chem. 1968, 72, 2628.
J. Phys. Chem. 1985,89, 5510-5513
5510
polarizabilities, heats of formation, and electronegativities suggests that a simple correlation exists between reactivity and macroscopic radical properties. For the H,Oz case in particular, a simple rationale for the kinetic behavior appears to be related to the strength of the product hydride bonds. However, no convenient rationale is discernable for the O3 data. Clearly, a larger body of SH and Br kinetic data is required in order to pursue further this avenue of investigation. In terms of the atmospheric chemistry of SH, which is illustrated in Figure 7, the rate constant values obtained from this work for k l and kz indicate that conversion of SH to HSO in the troposphere occurs within seconds. However, owing to reaction 6a, HSO interconverts rapidly with SH so that the ultimate fate of tropospheric SH is inextricably linked to the fate of HSO. Removal of H S O can proceed via several channels: HSO 0 2 SO HOz (10)
+
-+
+ HO2 HSO + NO2 + M HSO
-+
+ SO + H202
+
HSONO,
+M
(11) (12)
Reaction 11 is likely to be faster than H0, self-reaction since the bond strength of H-SO is significantly less than that of H-00. If this is the case, then irreversible oxidation of HSO will take place on the order of 1 h, only slightly longer than the time required for reconversion of HSO to SH and cycling of sulfur through the reaction 2 and 6a couplet will w u r only a few times. However, should HSO be unusually stable, a steady-state S H population will persist, cycling through the ozone couplet will occur repeatedly, and SH disproportionation reactions will acquire increased significance despite the small reactant concentrations of OH and SH. Obviously, kinetic data on HSO is required in order to resolve these possibilities.
Acknowledgment. We gratefully acknowledge L. Lapson for his design and construction of the copper-vapor laser and J. Demusz for his design of the laser trigger pulse amplifier circuit. This work was supported by the National Science Foundation under Grant ATM-8311695. Registry NO. SH, 13940-21-1; NO2, 10102-44-0; 7782-44-7; H202, 7722-84-1.
03,
10028-15-6; 02,
Pressure Dependence of the Rate and Stolchlometry of Water Photolysis over Platfnlzed TiO, Catalysts Keiiti Yamaguti and Shinri Sato* Research Institute for Catalysis, Hokkaido University, Sapporo 060, Japan (Received: May 29, 1985)
Water photolysis over platinized TiOz (anatase and rutile) catalysts was carried out in the presence of inert gas (He) to examine the effect of ambient pressure on the reaction rate and stoichiometry. The evolution rates of hydrogen and oxygen were independent of the He pressure up to 1 atm over anatase and rutile catalysts. A thick water layer on the catalysts, however, gave a remarkable hindrance effect on the rate of water photolysis even under reduced pressure. When the photolysis is very slow, a nonstoichiometric ratio of hydrogen to oxygen (H2/02> 2) was observed due to the slow adsorption of oxygen on TiOz. The photoadsorption of oxygen on the catalysts was trivial under the present experimental conditions.
Introduction Although a number of works on the photocatalytic water decomposition on semiconductor photocatalysts have been reported,'-I6the oxygen produced is often much less than stoichiometric. (1) Bulatov, A. V.; Khidekel, M. L. Izu. Akad. Nauk SSSR Ser. Khim. 1976, 1902.
(2) Yoneyama, H.; Koizumi, M.; Tamura, H. Bull. Chem. SOC.Jpn. 1979, 52, 3449. (3) Schrauzer, G. N.; Guth, T. D. J . A m . Chem. SOC.1977, 99, 7189. (4) Van Damme, H.; Hall, W. K. J . A m . Chem. SOC.1979, 101, 4373. (5) Kawai, T.; Sakata, T. Chem. Phys. Letr. 1980, 72, 87. (6) (a) Sato, S . ; White, J. M. Chem. Phys. Letr. 1980, 72, 83. (b) Sato, S . ; White, J. M. J . Catul. 1981, 69, 128. (7) Wagner, F. T.; Somorjai, G. A. J . A m . Chem. SOC.1980,102, 5459. (8) (a) Domen, K.; Naito, S.;Soma, M.; Ohnishi, T.; Tamaru, K. J . Chem. SOC.,Chem. Commun. 1980, 543. (b) Domen, K.; Naito, S.; Ohnishi, T.; Tamaru, K.; Soma, S. J . Phys. Chem. 1982, 86, 3657. (c) Domen, K.; Naito, S.; Ohnishi, T.; Tamaru, K. Chem. Phys. Lett. 1982, 92, 433. (9) (a) Lehn, J. M.; Sauvage, J. P.; Ziessel, R. Nouv. J . Chim. 1980, 4 , 623. (b) Lehn, J. M.; Sauvage, J. P.; Ziessel, R.; Hilaire, L. Isr. J . Chem. 1982, 22, 168. (10) (a) Duonghong, D.; Borgarello, E.; Gratzel, M. J . Am. Chem. SOC. 1981, 103,4685. (b) Borgarello, E.; Kiwi, J.; Pelizzetti, E.; Visca, M.; Gratzel, M. J . Am. Chem. SOC.1981, 103, 6324. (c) Borgarello, E.; Kiwi, J.; Gratzel, M.; Pelizzetti, E.; Visca, M. J. Am. Chem. SOC.1982, 104, 2996. (1 1) (a) Kalyanasundaram, K.; Borgarello, E.; Gratzel, M. Helv. Chim. Acta 1981, 64, 362. (b) Kalyanasundaram, K.; Borgarello, E.; Duonghong, D.; Gratzel, M. Angew. Chem. 1981, 93, 1012. (12) Mills, A.; Porter, G. J. Chem. SOC.,Faraday Trans. I 1982, 78, 3659.
0022-3654/85/2089-5510$01.50/0
For example, Gratzel and his co-workers reported efficient water photolysis under atmospheric pressure using colloidal TiO, loaded with both Pt and RuO,, but the 0, evolved was much less than stoichiometric.10 Mills and Porter, on the other hand, reported the UV irradiation of an aqueous suspension of Pt/Ti02 catalysts under atmospheric pressure produced H2 but not 0, in the gas phase.I2 To our knowledge, stoichiometric photolysis of liquid water on powdered TiO, catalysts has not been observed under atmospheric pressure. The nonstoichiometric formation of 0, has been explained in terms of (1) the photooxidation of organic compounds from vacuum valve grease,s (2) the high solubility of oxygen in water (1.4 wmol/mL under 1 atm at 20 OC),I4 (3) the photoadsorption of oxygen onto TiO, particles (1 nmol per 1.0-1.2 g),lo~Iz or (4) the formation of peroxides.lS We have found that stoichiometric water photolysis proceeds over metallized TiOz and SrTiO, catalysts under reduced pressure with high quantum efficiency up to 29%.6*'6In our experiments, (13) Stevens, C. G.; Wessman, N . J.; Bowman, J. E.; Ramsey, W. J. Chem. Phys. Lett. 1982, 91, 335. (14) Turner, J. E.; Hendewerk, M.; Somorjai, G. A. Chem. Phys. Lett. 1984, 105, 581. (15) (a) Kiwi, J.; Gratzel, M. J . Phys. Chem. 1984, 88, 1302. (b) Kiwi, J.; Gratzeel, M. Proc. Int. Congr. Catul., 8th 1984, 545. (c) Kiwi, J.; Morrison, C. J. Phys. Chem. 1984, 88, 6146. (16) (a) Yamaguti, K.; Sato, S . Nippon Kaguku Kaishi 1984, 258. (b) Yamaguti, K.; Sato, S.Proc. Int. Congr. Caral., 8th 1984,477. (c) Yamaguti, K.; Sato, S . J. Chem. SOC.,Faraday Trans. 1 1985, 81, 1237.
0 1985 American Chemical Society
