J . Phys. Chem. 1991, 95,2629-2636
2629
High-Resolution Infrared Flash Kinetic Spectroscopy of OH Radicals Aram Schiffman,* David D. Nelson, Jr., Marin S. Robinson, and David J. Nesbittt Joint Institute for Laboratory Astrophysics, University of Colorado and National Institute of Standards and Technology, and Department of Chemistry and Biochemistry, University of Colorado, Boulder, Colorado 80309-0440 (Received: May 1, 1990; In Final Form: December 4, 1990)
A high-resolution infrared flash kinetic spectrometer is used for time- and frequency-resolved studies of the OH radical. OH is produced by 193-nm excimer laser photolysis of HNOJbuffer gas mixtures in a 100-cm flow tube and is probed via weak fractional absorption of light from a widely tunable (2.35-3.59 pm) single-mode (Av I2 MHz) color center laser. The IR absorption technique allows fast ( [OH] by a factor of 500-100000, [R] is essentially constant and the OH decays according to [OH] = [OH]oe-(F)‘ = [OH]&/rob
(12) where k’= ki[RH] and robsis the observed lifetime. The signals are fitted to single-exponential decays with a nonlinear leastsquares program to determine robsfor a series of alkane concentrations. Figure Sa,b shows a typical signal from the reaction with 2.0 X 1015/cm3n-butane, along with the fitted decay and corresponding semilog plots, demonstrating the quality of singleexponential decay observed over 4 lifetimes. Finally, the observed rates l / r h are plotted vs [RH]. The rate constants are obtained from the slopes of these linear plots. IV. Systematic Checks Thermalization of the radicals is anticipated to be rapid with the large buffer excesses used. However, it is conceivable that unexpectedly slow relaxation channels could exist. The 193-nm photolysis of HNO, is known from LIF studiesa to produce a substantial fraction of the O H in rotational states as high as N 18 where R-T energy transfer is expected to be less efficient. These studies also indicate that the vibrational states are formed with 195% in v = 0 and 15% in u = 1, but to our knowledge there are no studies quantifying the population in the remaining energetically accessible v = 2 to v = 5 levels. Indeed, Lester and c o - ~ o r k e r srecently ~~ have observed OH(v= 1,2) from 193-nm
-
(40) Jacobs, A.; Kleinemanns, K.; Kuge, H.; Wolfrum, J. J. Chem. Phys.
1983, 79, 3162.
I
-5t -7
0
100
200
300
400
500 600 700
TIME ( p ~ ) Figure 5. (a) OH decay due to reaction with 2.0 X 10’5/cm3n-butane. The solid line is the fitted single-exponential decay. The residuals from the fit are shown below. The signal displays single-exponential behavior for over 4 lifetimes. (b) Semilog plot of the same data.
TABLE 11: Measured Rate Constant for OH + n-C4Hlo (in lW1* cm3 moiecuie-I s-9 k3 conditions 2.35 f 0.08 9 Torr of Ar, P(2.5)l-” 2.28 f 0.10 5 Torr of Ar, P(2.5)l2.22 f 0.12 9 Torr of N,, P(2.5)l2.26 f 0.1 1 9 Torr of SF6, P(2.5)l2.26 f 0.11 9 Torr of Ar, P(2.5)1+ 2.32 f 0.10 9 Torr of Ar, P(4.5)I2.28 f 0.11 9 Torr of Ar, P(2.5)2+ “Notation indicates the IR transition. P indicates AJ = - 1 , where the number in parentheses is J”, 1 or 2 indicates the spin-orbit state (Q = 3/2 or Q = 1/2, respectively), and + or - indicates the parity of the A-doublet.
photolysis of HNO, in a supersonic jet using LIF detection. Using direct IR absorption methods, we have also observed population
Schiffman et al.
2634 The Journal of Physical Chemistry, Vol. 95, NO.7, 1991 0.15
0.12c
1
0.06
r
C 0.04 I-
\
0.02
I
0
1 7
[ETHANE]
( I 0l6
1
I
J
14
21
28
3
"0
9
6
IS
12
[PROPANE] ( IOi6 molecu l e s / c m 3 )
mo Ie c u I e s / c m 3
Figure 6. Plot of I / r vs ethane concentration. All points represent
measurements under the standard conditionsof 9 Torr of argon buffer, monitoring the P(2.5)l- transition. in OH(v= 1,2), but careful power-dependent studies suggest that these species do not arise from a single-photon dissociation process. With these concerns in mind, we perform two independent checks to verify that complete rotational and vibrational thermalization of the OH radicals occurs on time scales which are rapid with respect to the reaction times. In the first, the identity and concentration of the buffer gas are varied, and a rate constant is determined. If relaxation is occurring on the time scales of the reactions, contributions to the observed decays will appear which will vary with the energy-transfer efficiencies of the various buffer gases. Table I1 lists the measured rate constant for reaction 3 as determined with -9 Torr of each of the species argon, N2, and SF6and -5 Torr of argon. In all cases, the value is constant to 15%. In the second test, decays are measured with the laser tuned to a series of transitions monitoring (1) several J rotational states, (2) both spin-orbit states, and (3) both *A-doublets. If equilibrium is established rapidly with respect to reactive processes, the same decay rate will be observed from all OH probe transitions. Table I1 lists the results from measurements of J = 2.5 and 4.5 in F = 1 and on J = 3.5 in F = 2. Here again, in all cases the rate constants was determined to be the same to better than 5%. In order to verify that the populations of any vibrationally excited O H are low, we have tuned the color center laser to transitions from u = 1 and v = 2. We have, in fact, observed OH in these states; the populations, which vary quadratically with UV flux, are made to constitute only 1 5 % of the total population with the low UV pulse energies used. Thus,contributions from vibrationally excited OH, even if relaxed slowly, will have negligible effect on the measured rates. The presence of small concentrations of reactive impurities in the H N 0 3 , though quite unlikely after the multiple distillations and methods of synthesis described earlier, would tend to cause enhancement of the observed decay rates. However, since [HNO,] even a 10% impurity which reacts is maintained at 16 X 1013/cm3, a t a gas kinetic rate would lead to a decay time of only 1 ms. This is on the same time scale as diffusion loss and would have a virtually unobservable effect. Furthermore, since the H N 0 3 concentration is kept constant for all runs, such a contaminant would contribute a zero offset to our plots of 1/rObvs alkane concentration but would not affect the slope. The determinations of the alkane concentrations depend on accurate knowledge of the cell pressure as measured on the capacitance manometer and measurements of the flow rates via timing pressure rises into a known volume. The pressure meter has been checked against a mercury manometer as a standard reference, as well as with several other calibrated pressure meters in the group. Note that only relative flow rates are required, and since all flowmeters are calibrated with the same volume, any small errors would influence all flow proportionally and not affect the final result.
-
(41) Beck, K.M.;Berry, M.T.;Brustein, M.R.;Lester, M.I. Chem. Phys. Leu. 1989, 162,203.
Figure 7. Plot of 1/r vs propane concentration. All points represent measurements under the standard conditionsof 9 Torr of argon buffer,
monitoring the P(2.5)1-transition. 0.25
v)
$ 0.15
\
C 2
0
[n
- BU TA N E]
4
6
10
8
( I O l 6 m o Ie c u I e s / c m
)
Figure 8. Plot of 1/r vs n-butane concentration. All points represent measurements under the standard conditions of 9 Torr of argon buffer, monitoring the P(2.5)1- transition unless otherwise noted: 0 , standard conditions; 0,P(2.5)1+;A, P(1.5)2-;+, P(2.5)2-; X, P(4.5)1-; 0 , 9 Torr of SF6 buffer, V, 9 Torr of N, buffer.
r * t -- Ic o'20
In
t
0.16 0.121
0.08
0.04
t
I
I
I
I
1.5
3.0
4.5
6.0
[ISOBUTANE]
molecules/cm 3 1
Figure 9. Plot of I/' vs isobutane concentration. All points represent measurements under the standard conditions of 9 Torr of argon buffer, monitoring the P(2.5)1- transition.
The reactions under study do exhibit a temperature dependence, so considerable care is taken to minimize cell heating. Under our conditions of p(HN0,) = ZmTorr, with the UV fluxes employed in these experiments, the local temperature rise in the illuminated column is calculated to be 52 K, i.e., comparable to day-to-day fluctuations in the room temperature. From temperature-dependent studies2*this magnitude of temperature rise contributes a t most a 2% change in the observed rates. V. Results and Discussion Reaction rates are measured for each of the reactions (1)-(4). Each time trace is signal averaged for 5000 pulses and fitted to a single-exponential decay. The concentration of each alkane is varied Over more than an order of magnitude to establish accurate,
The Journal of Physical Chemistry, Vol. 95, NO. 7. 1991 2635
Flash Kinetic Spectroscopy of OH Radicals TABLE 111: Rate Constants for OH + A b n e Reactions (in
cm3 mok~uie-~ s-9 reagent this work ethane 0.243 f 0.012
recommended
previous values
0.28'*
0.232 f 0.01633 0.264 f 0.01731 0.231 f 0.0432 0.259 f 0.02142 0.230 f 0.26" 0.214 f 0.3u 0.31 i 0.02153 0.298 0.02154 0.251" 0.25645 1.128 1.05f 0.0442 1.10 f 0.0446 1 .2645 1.20 f 0. 1855 0.83 f 0.17)O 2.02 A 0.1031 2.55 f 0.7329 2.56" 2.35 f 0.3547 2.45" 2.67 f 0.22*' 2.72 f 0.2752 2.66 f 0.8529 2.1" 2.2049 2.52 f O.OSm
*
propane
1.02 f 0.05
butane
2.32 f 0.08
isobutane
2.1 1 f 0.09
linear plots of l / i o b vs [RH]. The plots of l/iObsvs alkane concentration for each alkane are shown in Figure 6-9, along with the best fit line through the data. The rate constants for (1)-(4) are determined from the best fit slopes to the data in Figures 6-9 and are listed in Table 111; the stated error bounds are 24 from the linear fits. The reaction cm3 molecule-' of O H with ethane, k l = (2.43 f 0.12) X s-I, is the slowest of the four reactions under study, since ethane contains only primary hydrogens. The rate constant for the propane reaction (2) is considerably larger, consistent with the weaker bond strengths of the two secondary hydrogens, with k2 cm3 molecule-' s-' The fastest reactions = ( 1 -02 f 0.05) X (3) and (4) with the butanes are found to have rates k3 = (2.32 f 0.08) X cm3 molecule-' s-I for n-butane and k4 = (2.1 1 f 0.09) X cm3 molecule-' s-I for isobutane. It is useful to compare the results obtained in our flash kinetic spectrometer with previous values. Table 111contains results from previous ex riments, as well as recommended values from recent reviews.28*2In general, good agreement is achieved between our results and those of the LMR and flash photolysis/LIF probe experiments. We also note that, in the few cases of noticeable disagreement with some of the older results, and our values tend to be lower. Our value for the reaction with ethane (1) is in reasonable agreement with the flash photolysis/resonance fluorescence studies of Tully et a1.,42 of Lee and Tang,3z and of Wallington et ale4, and falls within experimental uncertainties of the flash photolysis/LIF studies of Tully and cu-workers" and the discharge flow work of Bourmada et alSM It also agrees closely with revised values from the discharge flow/laser magnetic resonance measurements of Howard and E v e n ~ n . 3The ~ flash photolysis/kinetic spectroscopy result reported by Greiner4s after correction for the effects of an observed secondary reaction is -25% higher than our value. However, a plot from data presented in ref 45 of observed decay rate vs ethane concentration reveals a nonzero intercept (see ref 33) and a slope (Table 111) which is in good agreement with the rate constant from the present study. The value from the flash photolysis/resonance absorption studies of Overend et ale3'also agrees within experimental error. The rate
ge
(42) Tully, F. P.; Ravishankara, A. R.; Carr, K. I. Inr. J . Chem. Kiner. 1983, 15, I 1 1 1. (43) Wallington,T. J.; Neuman, D. M.; Kurylo. M. J. Inr. J. Chem. Kiner. 1987, 19. 725. (44) Bourmada, N.; LaFage, C.; Devolder, P. Chem. Phys. Lerr. 1987,136, 209. (45) Greiner, N. R. J . Chem. Phys. 1970, 53, 1070.
constant presently used in atmosphere modeling28 is an average taken from these studies and thus includes the high values reported by Howard and Evenson and by Greiner. The resulting recommended rate constant of 2.8 X lo-', cm3 molecule-' s-' is -15% higher than that reported here. We suggest, then, that the recommended OH ethane rate constants should be. reevaluated in light of the discrepancies and the lower value consistent with the present study ( k , = 2.43 X lo-', cm3 molecule-' s-I) should be adopted. The values reported for the propane reaction (2) have been more consistent. Our value is in good agreement with most of the previous data, particularly the flash photolysis/resonance fluorescence42and flash photolysis/LIF46 experiments. The present study concurs that the current recommended value for OH propane ( k , = 1.1 X cm3 molecule-' s-I) is reasonable. There are somewhat fewer values available for the reactions of n-butane (3) and isobutane (4). However, our values are markedly lower than the recommended values in both cases. We agree to well within our 4%error limits with the flash photolysis/resonance fluorescence measurements of S t u h P for n-butane. We are in close agreement with flash photolysis LIF studies of Droege and Tully4*for n-butane and Tully et al. for isobutane. The ratio of n-butane to isobutane reaction rates reported by Darnall et aLsowas determined in a competitive reaction scheme with a number of reagents, using gas chromatography to determing branching ratios. The resulting value for k3/k4is 1.08, which agrees closely with the ratio k3/k4= 1 . 1 1 obtained in the present study. Pareskevopoulos and NipS' performed flash photolysis/ resonance absorption experiments are reported k3 15% higher than our value. A close examination of their experimental conditions reveals that they used O H concentrations which were -3-6 times greater and n-butane concentrations more than 10 times smaller than those employed here. Due to the possibility of extremely fast radical-radical reaction channels, this experimental regime of [OH]/[n-butane] concentrations can lead to appreciable perturbations of the OH kinetic lifetimes. Indeed, in duplicating those experimental conditions we observe systematic increases in the apparent decay rates, attributable to the secondary reaction of OH with butyl radical (see eq lo), which could easily account for this discrepancy. Perry et al.s2 report a similarly high number for k,; they do not provide an estimate of their O H concentrations, and the possibility of similar secondary reactions is not ruled out. The rate constants currently accepted for atmospheric modeling, averages which include these higher values, are 10%larger for n-butane, and -25% larger for isobutane, than the rate constants reported here. We again suggest that the recommended rate constants be revised and the lower values consistent with the present and several previous studies ( k , = 2.32 X cm3 molecule-' s-I, k4 = 2.1 1 X cm3 molecule-' s-') be adopted. The alkanes studied in this paper are at sufficiently low concentrations in the atmosphere that the above rate constant revisions would not lead to qualitatively different modeling of atmospheric chemical phenomena. However, their low concentrations make these alkanes excellent naturally occurring tracer molecules which reflect the chemistry of their local environments. Recently, concentrations of atmospheric ethane, propane, n-butane, and isobutane, measured in the field periods of several weeks, have been employed in models of atmospheric global air In
+
+
I
-
-
(46) Droege, A. T.; Tully, F. P. J. Phys. Chem. 1986, 90, 1949. (47) Stuhl, F. Z . Naturforsch., A 1973, 28, 1383. (48) Droege, A. T.; Tully, F. P. J . Phys. Chem. 1986, 90, 5937. (49) Tully, F. P.; Goldsmith, J. E. M.; Droege, A. T. J. Phys. Chem. 1986, 90, 5932, (50) Darnall, K. R.; Atkinson, R.; Pitts, J. N., Jr. J . Phys. Chem. 1978, 82, 1581. (51) Pareskevopoulos, G.; Nip, W. S. Can. J . Chem. 1980, 58, 2146. (52) Perry, R. A.; Atkinson, R.; Pitts, J. N., Jr. J . Chem. Phys. 1976,61, 5314. (53) Jeong, K.-M.; Hsu,K.-J.;Jeffries, J. B.;Kaufman, F. J. Phys. Chem. 1984, 88, 1222. (54) Nielson, 0. J.; Munk, J.; Pagsberg, P.; Sillesen, A. Chem. Phys. Lett. 1986, 128, 168. (55) Baulch, D. L.; Campbell, I. M.; Saudders, S. M. J . Chem. Soc., Faraday Trans. I 1985, 81, 259.
2636
J . Phys. Chem. 1991, 95, 2636-2644
the simplest reasonable model that the predominant alkane loss channel is reaction with OH, the changing concentrations of these four species are used to infer the age and origin of low-lying air masses. However, predictions of the variations in relative alkane concentrations, based on previous rate constants and the recent field observations, conflict with the observed ratios. Accurate determinations of the rate constants for these reactions are therefore important to implicate either the model itself or the rate constants used in it'as the source of these discrepancies. The lower values of these rates found in the present study lead to increased disagreement between predictions and observations, indicating that the models may neglect important second-order effects and imply that assumptions regarding the spatial uniformity of the alkane tracer molecules or the effects of dilution on the traveling air masses may be incorrect.*' VI. Conclusions High-resolution infrared flash kinetic spectroscopy is used with a flow tube arrangement for time- and frequency-resolved studies of the O H radical. The OH is produced from photolysis by an excimer laser pulse and monitored via direct infrared absorption of light from a high-resolution color center laser which is tuned to the peak of or scanned over individual transitions from a specific rovibrational, spin-orbit and A-doublet quantum state. The collinear geometry of the excimer and color center lasers, along with the absolute infrared cross sections which have been measured
previously in this spectrometer, allows direct determination of OH number densities from the fractional IR absorption. The performance of this spectrometer is demonstrated by measurements of reaction rates of OH with ethane, propane, n-butane, and isobutane. These studies are facilitated by the high sensitivity and fast response of the real-time IR absorption signals. Complications due to wall reactions are eliminated, as the radicals are probed at the center of the flow cell. Extensive checks are performed to assure that the measurements are not complicated by competing radical-radical reactions, energy-transfer mechanisms, or temperature effects. Determinations of the four OH hydrocarbon rate constants are made which indicate that a significant number of previous studies have yielded excessively high values for the ethane, n-butane, and isobutane reactions. It is recommended that the rate constants currently used in atmospheric modeling be revised to reflect the lower values from the present study and that the assumptions underlying current atmospheric air flow models be reexamined in light of discrepancies between observations and predictions based on these rate constants.
+
Acknowledgment. This work has been supported by grants from the Air Force Office of Scientific Research. The authors thank Dr. A. R. Ravishankara and Dr. C. J. Howard for many valuable discussions. Registry No. HNO,,7697-37-2; OH, 3352-57-6; ethane, 74-84-0; propane, 74-98-6; butane, 106-97-8; isobutane, 75-28-5.
Structural Dependence of HF Vfbratlonal Red Shms in Ar,HF, n = 1-4, via Hlgh-Resolutfon Slit Jet Infrared Spectroscopy Andrew McIlroy, Robert Lascola, Christopher M. Lovejoy, and David J. Nesbitt*lt Joint institute for Laboratory Astrophysics, University of Colorado and National institute of Standards and Technology, and Department of Chemistry and Biochemistry, University of Colorado, Boulder, Colorado 80309-440 (Received: May 10, 1990)
-
The rotationally resolved v = 1 0 HF stretching spectra of Ar,,HF, n = 1-4, have been observed by using a slit jet, difference frequency infrared laser spectrometer. The red shift of the HF vibrational frequency is seen to be sensitively dependent on the placement of the Ar with respect to the projection of the HF dipole moment; the largest incremental red shift is observed for the ArHF linear geometry. The n = 1-3 red shifts account for almost half (9.65-19.26 cm-') of the shift observed for H F in an Ar matrix (42.4 cm-I), suggesting that only the nearest neighbors contribute significantly to the perturbation of the H F vibrational frequency. The geometry of the experimentally observed isomer of Ar4HFplaces the fourth Ar in what would be the second coordination layer from the HF where it has little effect (
