J . Phys. Chem. 1987, 91, 5140-5143
5740
Rate Constant Measurements of Reactions of C2H with H2, Op, C2H2,and NO Using Color Center Laser Kinetic Spectroscopy J. W. Stephens, Jeffrey L. Hall, H. Soka,+W.-B. Yan, R. F. Curl,* and G. P. Glass* Department of Chemistry and Rice Quantum Institute, Rice University, Houston, Texas 77251 (Received: December 15, 1986; In Final Form: June 10, 1987)
The rate constants for the reactions of the ethynyl radical (C2H) with H2, 02,C2H2,and NO have been measured by following the time decay of an C2H infrared transient absorption line originating from the ground vibronic state using color center laser spectroscopy. For the H2, 02,and NO reactions, the C2H was produced by excimer laser flash photolysis (ArF, 193 nm) of CF3C2H. In the case of the C2H2reaction C2H was produced by flash photolysis of acetylene again using the 193-nm ArF excimer line. Excited states of C2H, which are abundant with 193-nm photolysis, were relaxed by buffering the photolysis cell with -20 Torr of He buffer and 160 mTorr of SFs. Rate constants of 4.2 X lo-", 4.8 X lo-", 1.5 X 10-lo, and 3.5 X lo-" cm3 molecule-' s-l were obtained for the reactions of C2H with O,,H2, C2H2,and NO, respectively.
-
Introduction The ethynyl radical, C2H, is known to be an important intermediate in a number of chemical systems. In combustion it plays a key role in the initial stages of soot formation,' and it has also been detected in interstellar space? and in planetary atmo~pheres.~ Previous absolute rate measurements on C2H have been made primarily by three groups. Laufer and various co-workers have studied reactions with 02,H2, C2H2,and a number of other small hydrocarbons by using flash photolytic production of C2H.4-6 The O2and C2H2reactions were followed by observing the appearance of the reaction products CO and C4H2,respectively, and the rate of reaction with H2 was deduced by measuring the dependence of the rate of butadiyne production (arising from the reaction of ethynyl radicals with the acetylene precursor) on the partial pressure of added H1. Renlund et al.'q8 measured reaction rates with 02, H2, and CH4 by monitoring time-resolved chemiluminescence from CH(A2A), CO(a'3Z+), and C02(X'Zg+,v3>,l)produced in the presence of O2 by the reactions
-
C2H + 0, CH(A2A) + C O , ( R l P ) HCO
(la)
+ CO(~'~Z+)
The C2H was formed by both excimer laser photolysis and IR multiple-photon dissociation of a variety of precursors. In the work of Lange and Wagner, C 2 H was formed in a fast-flow system by microwave discharge of dilute bromoacetylene/helium mixture^.^ The concentrations of C2H ( m / e 25) and various products were followed by mass spectrometry. Because they estimated that mixing was incomplete in their observation zone, the authors only quoted lower limits for the reaction rates of C2H with H,, O,, and acetylene. All of the rate constants previously measured are summarized in Table I. It should be noticed that Laufer and Lange and Wagner appear to be in rough agreement. While the C2H + O2 rate constants of Renlund et al. agree within a factor of 4, their values for the reaction rate of C,H with H2differ by about 2 orders of magnitude from the other two groups. A more recent publicationlo by this group proposes an explanation of these discrepancies on the basis that their earlier measurements involve reactions s f vibrationally or electronically excited fadicals. The A211state is only -4000 cm-' above the X2Z Therefore, if rate constants for the ground-state radical, C2H(X2Z+), are to be measured, it is important either to ensure that no excited state is present or to observe the rate of removal of the ground state directly. Kinetic studies on this reactive radical have been hampered by the fact that no gas-phase absorption spectra of the species are
'
Institut fiir Angewandte Physik, Universitlt Bonn, Wegelerstrasse 8, 53 Bonn, West Germany.
0022-3654/87/2091-5740$01.50/0
TABLE I: Comparison of C2HReaction Rate Constants" reaction C2H C2H C2H C2H
+ H2 + O2 + C2H2 + NO
this work 4.8 (3) X lo-" 4.2 (1) X lo-" 1.5 (1) X lo-'' 3.5 (1) X lo-''
In cm3 molecule-l s-', ref 7. /From ref 8.
Lange and Wagnerb
Laufer et al.
Renlund et al.
1.7 X lo-'' 5.5 X lo-'' 5.0 X lo-''
1.5 X 5.0 X 3.1 X 10-"'
1.2 X 2.1 X lo-"/
* From ref 9.
From ref 4.
From ref 6. From
known in either the visible or UV regions of the spectrum.14 Primarily, kinetic studies on this radical have been carried out by monitoring the rates of formation of various reaction products. Only the work of Lange and Wagner, which employed a mass spectrometric method, directly monitored the C2H concentration. However, one may follow C2H kinetics using the known infrared absorption spectra of C2H.I2J3Recently, several such bands of the C2H radical have been observed in the photolysis of acetylene.15 In the work presented below, a C2Habsorption line of one of these bands, originating from the ground vibronic state of the molecule, has been monitored to measure rate constants for the reaction of C2H(W2Z+(0,0,0))with O,, H2, NO, and HCCH.
Experimental Section The experimental arrangement has been described in detail elsewhere16 and is shown in Figure 1. Briefly, the reagents were (1) (a) Homann, K. H.; Wagner, H. G. Proc. R. SOC.London, A 1967, 307, 141. (b) Warnartz, J. A.; Bwkham, H.; Moser, A,; Wenz, H. W. Symp. ( I n t . ) Combust. [Proc.] 1983, 19, 197. (2) Tucker, K. D.; Kutner, M. L.; Thaddeus, P. Astrophys. J. 1974, 193, L115. (3) Strobel, D. F. Planet. Space Sci. 1982, 30, 839. (4) Laufer, A. H.; Bass, A. M. J. Phys. Chem. 1979, 83, 310. (5) Laufer, A. H. J . Phys. Chem. 1981,85, 3823. (6) Laufer, A. H.; Lechleider, R. J. Phys. Chem. 1984, 88, 66. (7) Renlund, A. M.; Shokoohi, F.; Reisler, H.; Wittig, C. Chem. Phys. Lett. 1981, 84, 293. ( 8 ) Renlund, A. M.; Shokoohi, F.; Reisler, H.; Wittig, C. J. Phys. Chem. 1982,86, 4165. (9) Lange, W.; Wagner, H. G. Ber. Bunsen-Ges. Phys. Chem. 1975, 79, 165. (10) Shokoohi, F.; Watson, T. A.; Reisler, H.; Kong, F.; Renlund, A. M.; Wittig, C. J. Phys. Chem. 1986, 90, 5695. (11) Shih, S . K.; Peyerimhoff, S . D.; Buenker, R. J. J. Mol. Specrrosc. 1979, 74, 124. (12) Carrick, P. G.; Merer, A. J.; Curl, R. F. J. Chem. Phys. 1983, 78, 3652. (13) Curl, R. F.; Carrick, P. G.; Merer, A. J. J. Chem. Phys. 1985, 82, 3479. (14) A UV absorption in matrix isolation tentatively identified as arising from C2H has been reported by: Chang, K. W.; Graham, W. R. M. J . Chem. Phys. 1982, 76, 5238. (15) Yan, W.-B.; Hall, J. L.; Stephens, J. W.; Richnow, M. L.; Curl, R. F. J. Chem. Phys. 1987, 86, 1657.
0 1987 American Chemical Society
The Journal of Physical Chemistry, Vol. 91, No. 22, 1987 5741
Reactions of C2H with H2, 02,C2H2,and NO
A-'-'
TRANSIENT DIGITIZER
COMPUTER
UV BEXCIMER EiM
To
D
A
DIFFERENTIAL AMPLIFIER
PUMP
BU~FERM p :p GAS
Figure 1. Experimental layout. Balance between the two liquid N2 cooled InSb detectors is achieved by rotating the first polarizer. The collinear IR probe and UV photolysis beams counterpropagatethrough the reaction cell: P, polarizer; G, germanium flat; L, lens; I, iris.
photolyzed by 193-nm ArF excimer laser pulses directed along a 1-m-long cylindrical reaction cell, and the C2H concentration was probed with a tunable color-center laser operating between 2.3 and 3.3 pm." The IR signal was detected by liquid N2 mold InSb detectors and recorded by using a transient digitizer interfaced to a PDP 11/23 minicomputer. In order to significantly reduce low-frequency noise, a balanced detector scheme was employed. In this mode the time response of the complete detection system was measured as approximately 1 ps. In following the decay of C2H the laser frequency was locked at an absorption line of interest. Because of its reasonable absorption intensity and the fact that it is known to originate from the ground vibronic state of the molecule, the line chosen for these kinetic studies was the Q l l ( l 1) of the 3600-cm-' band at 3594.393 cm-'.I8 Normally, 1000 decay traces were averaged to obtain a single rate measurement at a particular reactant concentration. For the reactions with 02,H2, and NO, the ethynyl radical was produced by 193-nm photolysis of CF3C2H (SCM Specialty Chemicals), and its temporal behavior was monitored at various pressures of the added reactants. Typically, reaction mixtures contained 35 mTorr of CF3C2H,20 Torr of He, and 160 mTorr of SF,. The helium was added to ensure thermal equilibrium and to moderate the temperature rise following photolysis; the SF6 was added to ensure rapid vibrational relaxation of the C2H. For the reaction with HCCH, acetylene at various pressures was used as the C2H precursor. The flow rate of each reagent was determined by timing the pressure rise that occurred when the gas flowed into the calibrated volume of the reaction cell with all the other gases shut off. All pressures were measured by using a Baratron calibrated against a H g manometer. Partial pressures for individual components were calculated from the relative flow rate and the total pressure. To conserve CF3C2H,the valve between the photolysis cell and the pump was throttled down, resulting in a residence time in the cell of -4 s, during which time 60 excimer flashes occurred. Each laser pulse consists of about 2 X 10I6photons, the UV absorption cross section of CF3C2Hwas measured as less than 2 X cm2, and the cell (tube, 54-mm i.d.) has a cross-sectional area of -23 cm2. Because radial diffusion is rapid compared with the residence time of 4 s, the fraction of precursor decomposed directly by the laser is the product of the total number of photons with the UV absorption cross section divided by the cell cross-sectional area giving N 1% decomposition of the trifluoropropyne leaving the cell, a negligible depletion of the precursor concentration. The substances reacting with C2H are always present in large excess so that their depletion in a single-step process is negligible. A chain mechanism in which the photolysis or reaction products react (16) Hall, J . L.; Zeitz, D.; Stephens, J. W.; Kasper, J. V. V.; Glass, G. P.; Curl, R. F.; Tittel, F. K. J . Phys. Chem. 1986, 90, 2501. (17) Kasper, J. V. V.;Pollock, C. R.; Curl, R. F.; Tittel, F. K. Appl. Opt. 1982, 21, 236.
(18) Yan, W.-B.; Dane, C. B.; Zeitz, D.; Hall, J. L.; Curl, R. F. J . Mol. Spectrosc. 1987, 123, 486.
with additional reactant appears to be ruled out by the overall reaction stoichiometries found by Laufer. Experimentally, when the repetition rate of the photolysis flash was reduced by a factor of 5, no change in the decay curve was observed, confirming that depletion is not important. CF3C2H was used from the cylinder without further purification. Laufer has noted that his sample contained -6% of hydrocarbon impurity which reacted rapidly with C2HsSNOsuch impurities were detected in our sample when analyzed by gas chromatograph/mass spectrometry; however, the presence of any impurities and the unwanted side reactions would not have affected our calculated rate constants because we always subtracted the precursor contribution from all measured decay rates. All other gases were obtained commercially and were of the following purities: He (99.995%), SF6(99.8%), o2(99.6%), H2 (99.99%), NO (99.0%). Some were used from the cylinders; however, NO was passed through a dry ice/acetone trap to remove NO2 impurities. In addition, some experiments with H2 were conducted with a liquid N2 cooled molecular sieve trap placed in the H 2 delivery line. We also made some measurements using higher purity H2 (99.999%). Commercial acetylene was found to contain a significant amount (2%) of acetone impurity, which was removed by two approaches. In some experiments the gas was passed through a stainless steel coil at -78 OC to trap out the acetone. In others the acetylene was passed through an activated charcoal filter (Matheson, Model 454). GC/MS analysis of the gas from the latter method indicated that the residual acetone impurity was less than 1%. Since the reaction of C2H2with C2H is so fast, a 1% impurity can have no effect on the measured rate. The concentration of ethynyl radicals formed in our system can be calculated from the measured CF3C2Habsorption cross section at 193 nm (C2 X cm2) and the excimer laser flux (2 X lo', photons pulse-') to be less than 5 X 10l2 ~ m - ~Thus, . in our experiments, a 1000-fold excess of reactant gas was present even at the lowest partial pressures used, and any contribution from secondary or from radical-radical reactions can be neglected. Wall reactions cannot affect our rate measurements because the time scale of diffusion to the walls is milliseconds at 20 Torr while the reaction time is less than 50 ps. (The cell is 50 mm in diameter with the probing laser beam near its center.) To check for possible systematic errors, well-known rate constants for the OH + C2H6 and N H 2 NO reactions were remeasured in our apparatus under conditions similar to those used and 1.5 X for the study of C2H. Rate constants of 2.7 X lo-'' cm3 molecule-' 8,respectively, were obtained for these and reactions, in close agreement with the values of 2.7 X 1.6 X lo-'' cm3 molecule-' s-' recommended in a recent critical review.lg
+
Results In measurements of the H2, 02, and NO rates, trifluoropropyne was chosen as a C2Hprecursor since it reacts more slowly with C2H than does acetylene. However, even at the lowest pressures that gave adequate C2H absorption signals (typically 35 mTorr), the reaction of the precursor (or its hydrocarbon impurity) with C2H was rapid (decay constant, T = 35 ps). Pressures of added reactants were chosen to give decay rates which were between the limits set by the precursor reaction rate and the detector response time. (No measurement with a decay shorter than 5 times the detector response was accepted.) Some typical C2H absorption-time traces at several pressures of O2are shown in Figure 2. All such traces were found to exhibit single-exponential decays. Decay constants were calculated by using points between 80% and 10% of the maximum signal size and weighting each point according to its magnitude.20 From the scatter of the data points in the plot of decay constants versus reactant pressure, we estimate the random error in the decay (19) Bauch, D. L.; Cox, R. A,; Hampson, R. F., Jr.; Kerr, J. A,; Troe, J.; Watson, R.T. J . Phys. Chem. Re$ Data 1984, 13, 1259. (20) Schmid, G. H.; Csizmadia, V. M.; Mezey, P. G.; Csizmadia, I . G. Can. J . Chem. 1976, 54, 3330.
5742 The Journal of Physical Chemistry, Vol. 91, No. 22, 1987
Stephens et al.
0
1
2
3
4
5
6
7
8
PH? (Torr)
Time ( p s e c s ) Figure 2. Decay traces of C2H absorption (Qll(l 1) of the 3600-cm-l band) with added 02:(a) no 0,; (b) 18.8 mTorr of 02;(c) 111 mTorr of 02.The spike at time = 0 is due to electrical interference from the excimer laser.
100
120
Figure 3. Corrected decay constants for C2Hversus pressure of 0,. The slope of the least-squares fit gives a rate constant of 4.2 X IO-’] cm3 molecule-’ s-’.
constant arising from this procedure as approximately 5%. Because different reaction mixtures contained different amounts of precursor, it was necessary to correct each measured decay rate for the appropriate contribution of the precursor. In order to do this, the apparent rate consant for the reaction of C2H with the precursor was determined, by use of 10 points, as 3.5 X lo-” cm3 s-l. Plots of corrected decay constants versus reactant pressure for O2and H2 are shown in Figures 3 and 4. A summary of the measured rate constants is given in Table I. The uncertainties given in parentheses in terms of the least significant digit reported are one estimated standard deviation of the slope obtained in the least-squares fit of the Stern-Volmer plots. The rates measured here are for C2H in its vibronic ground state as long as the upper state of the infrared transition has relaxed. To test for relaxation, preliminary measureeents were performed on absorptions originating from the C2H(XzZ+(0,0,1))state at 1841 cm-I above the ground state,I5 confirming that this state has a lifetime of less than 1 ps under the experimental conditions employed (limited by our detector response). Thus, vibrational relaxation of this state, which should relax no faster than the upper state of the transition monitored, occurred before ground-state rate measurements were made. Moreover, other studies showed that added SF, had no effect on measured ground-state decay rates. Discussion Rate constants for most of the reactions studied in this paper have been measured previously three times: by Lange and Wagner? by Laufer et aLe and by Renlund et aL8 These rates are listed in Table I. (In this table, the results of Renlund et al. are only given for completeness since they probably refer to reactions of excited stateslo) As can be seen from Table I, the results of Lange and Wagner and of Laufer et al. are generally
Figure 4. Corrected decay constants for C2H versus pressure of H,. The rate constant for a least-squares fit through all the points is 4.8 X lo-” cm3molecule-I s-l: 0,points taken with 99.99% H2; 0 , H2 was passed through a liquid N2 cooled molecular sieve trap; *, 99.999%H, was used.
in good agreement with each other, while our results are strikingly different. In view of this agreement between previous investigations, does it make sense for us to propose yet another set of values for these rate constants? Let us consider first the results of Lange and Wagner and then the relationship of our results to those of Laufer et al. Lange and Wagner followed as we do the concentration of C2H using in their case a fast-flow system and mass spectrometric probing. The C2H was produced by a microwave discharge in a bromoacetylene/helium mixture. Because C2H decayed rapidly in their system, Lange and Wagner were forced to limit their observations to those taken within a total reaction time of less than 0.9 ms, using a total pressure of slightly greater that 4 Torr. Under these conditions, mixing of C2H with the reactant by diffusion is very likely to be. incomplete. (Lange and Wagner estimate that they are in a laminar flow region where mixing is primarily by diffusion.) At a pressure of 4 Torr, we calculate a diffusion length ( ( x 2 ) ’ I 2 )of 6.5 mm for H2 and 3 mm for O2 (or C2H2)in 1 ms. This suggests that diffusive mixing is incomplete at their longest reaction time. Lange and Wagner were very clearly aware of this limitation of their measurements, and for this reason they stated that their rate constants are probably too small and give therefore a lower limit to the reaction rate. It is not clear how close this lower limit will be to the true reaction rate. Laufer and Bass4measured the rate of the C2H C2H2reaction using ultraviolet absorption spectroscopy to follow the appearance of C4H2. They inferred a rate for the C2H H2 reaction by measuring the effect of H2 on the yield of C4H2 and assuming that competition for C2H between H2 and C2H2took place. Laufer and Lechleider6 measured the rate of reaction between C2H and O2 by following the appearance of UV absorption due to CO. In this case, they are quite concerned by the detection by Lange and Wagner of large amounts of C 2 H 0 and attempted to check for a two-step mechanism involving this intermediate by making calculations using their own plus Lange and Wagner’s rate constants. For all three reactions (C2H2, 02,and H2), the rate constants obtained by Laufer and Bass (C2H2and H2) and Laufer and Lechleider (0,)are smaller than the estimated lower bounds of Lange and Wagner. In all cases the rate of appearance of products, not the disappearance of reactant, was followed. There is never a necessity for the rate of appearance of product to be as fast as the rate of disappearance of reactants because the overall reaction can always involve an undetected intermediate. Thus, for the C2H C2H2 reaction, the reaction scheme could be
+
+
+
C,H
-
+ C2H2 kC C4H,’
kD
C4H,’
+H
C4H2
where C 4 H i is an untrackable intermediate isomer of C4H2(such as, for example, a triplet excited state) and kc is the rate constant for creation of the isomer and kD is the rate constant for its destruction.
J. Phys. Chem. 1987,91, 5743-5749 Because Laufer and Bass appeared to measure the rate at only one acetylene pressure (75 mTorr), a reaction scheme involving a C4H3intermediate seems also to be possible since the order of the reaction producing C4H2 with respect to C2H2was not determined. Likewise, in the case of CzH 02,Laufer and Lechleider measured the rate of C O formation for only one O2 pressure so that they did not measure the order of the reaction with respect to 02.In this case a more general multistep reaction such as
+
+ O2 -% H C 2 0 + 0 HC20 + O222CO + OH C2H
is not ruled out by O2 reaction order. For any two-step mechanism the initial slope of the product C.,H2 rise will always be zero, but the time domain of the resulting kink in the concentration versus time curve decreases rapidly as kC[C2H2]/ k Dincreases. Because of interference caused by the photoflash, Laufer and Bass could not make observations during the first 10 ps. As long as the creation rate of the isomer, kc[C2H2],is much greater than its destruction rate, kD,the existence of the two-step mechanism will not be revealed by following only the product C4H2 unless the product is followed at very short
5143
reaction times. Test calculations of the size of this initial effect using our k values for kc and Laufer et al.3 k values for kD indicate that such an initial kink would be undetected by Laufer et al. for both C2H + C2H2and C2H 02.(The time analysis for CzH O2done by Laufer and Lechleider6used a much smaller value of kc than ours and is not relevant here.) We suggest that the mechanism for both these reactions may be more complex than a single step. Unfortunately, we do not have the sensitivity to follow the time evolution of precursor or product concentrations with the current apparatus. The ethynyl radical absorptions, which are easily seen, are especially intense since they invo_lve-vibronictransitions which borrow oscillator strength from the A-X transition. In the future, we plan to incorporate a multipass cell into the present apparatus which should enable us to observe CF3CzH, C2H2,and C4Hz. Also, we hope to monitor H C O and C O by using a diode laser. The results of such investigations should enable some of the questions raised by the present study to be resolved.
+
+
Acknowledgment. This work was supported by the Department of Energy under Grant DE-FG05-85ER 13439 and the Robert A. Welch Foundation under Grant C-07 1. Registry NO.CF,C2H, 661-54-1; CIH, 2122-48-7; H2, 1333-74-0; 02, 7782-44-7; NO, 10102-43-9; HCCH, 74-86-2.
Temperature Dependence of the Rate Constant for the Reaction HS
+ NOP
Niann S. Wang, Edward R. Lovejoy, and Carleton J. Howard* NOAA Aeronomy Laboratory, R/E/AL-2, Boulder, Colorado 80303, and the Department of Chemistry and Biochemistry and CIRES, University of Colorado, Boulder, Colorado 80309 (Received: February 9, 1987; In Final Form: May 19, 1987)
-
-
Two reactions of atmospheric importance involving the HS radical have been studied by using a discharge flow laser magnetic resonance technique: (1) HS + NO2 products; and (2) HS + O2 products. The rate constant for reaction 1 was measured at low pressure (- 1 Torr) between 221 and 415 K and a negative temperature dependence was observed: k l = (2.9 f 0.5) X lo-" exp((240 f 5O)/T) cm3 molecule-' s-'. The room temperature value for kl from this study is higher than three previous values. Secondary chemistry associated with high H2S concentrations in the earlier studies is believed to contribute to the discrepancies. DS was used to investigate the isotope effect for reaction 1 and no evidence has been observed for a primary.isotope effect, k(DS NO2) = (7.3 f 1.1) X lo-'' cm3 molecule-' s-', at 299 K. No reaction was observed with O2 and an upper limit of 1.5 X cm3 molecule-' s-I was assigned to reaction 2.
+
Introduction The increasing interest in acid precipitation problems in recent years has led to greater attention being given to the atmospheric reactions of sulfur compounds. It is becoming clear the the oxidation of reduced sulfur compounds such as H2S, OCS, CS2, and CH3SCH3contributes significantly to the global sulfate and sulfuric acid production. For HzSthe oxidation process is believed to be initiated by the well-characterized reaction with an OH radical which produces HS as a It is necessary to understand the subsequent steps in order to assess the relationship between the sources of reduced sulfur compounds and the sulfate deposition sites. Therefore, it is of interest to study HS kinetics in order to evaluate its role and importance in acid precipitation chemistry. In the present study we focus on two reactions of the HS radical: HS + NOz products (1) HS
+ 0,
--
products
(2)
agreement on the rate constant. Black3 measured k l = (3.5 f 0.4) X lo-" cm3 molecule-' s-' with a flash photolysis laser-induced fluorescence (LIF) method. A discharge flow LIF study by Friedl et aL4 gave kl = (3.0 f 0.8) X lo-'* cm3 molecule-I s-'. Bulatov et a1.5 reported a somewhat lower value of k , = (2.4 f 0.2) X lo-" cm3 molecule-' s-I at 295 K using a less direct method of monitoring the product concentration, [HSO], in a flash photolysis intracavity laser absorption experiment. No reaction between HS and O2 has been observed in three previous studies. The work by Tiee et ale6and Black3 with flash photolysis L I F experiments assigned upper limits for k2 (in cm3 and 4 X respectively. The molecule-' s-') of 3.2 X discharge flow study of reaction 2 by Friedl et aL4 gave k2 5 1 x io-'' cm3 molecule-' s-I. (1) Sze, N. D.; KO, M. K. W. Almos. Environ. 1980, 14, 1223. (2) Leu, M. T.; Smith, R. J. J . Phys. Chem. 1982, 86, 73. (3) Black, G. J . Chem. Phys. 1984,80, 1103. (4) Friedl, R. R.; Brune, W. H.; Anderson, J. G. J. Phys. Chem. 1985,89,
Two previous studies of reaction 1 at 298 K are in very good
5505. ( 5 ) Bulatov, V. P.; Kozliner, M. Z.;Sarkisov, 0. M. Khim. Fir. 1984, 3,
'Author to whom correspondence should be addressed at NOAA R/E/ AL-2, 325 Broadway, Boulder, CO 80303.
1300. (6) Tiee, J. J.; Wampler, F. B.; Oldenborg, R. C.; Rice, W. W. Chem. Phys. Lett. 1981, 82, 80.
0022-3654/87/2091-5743$01.50/0
0 1987 American Chemical Society
