2005
J. Phys. Chem. 1980, 84, 2005-2009
ARTICLES Thermal Decomposition of Propane 1. Kolke and W. C. Gardiner, Jr." Department of Chemlstry, University of Texas, Austin, Texas 78712 (Received: January 7, 1980) Publlcation costs assisted by the Petroleum Research Fund and the US. Army Research Office
The thermal decomposition of propane in reflected shock waves with temperatures in the range 1300 C Ts< 1700 K and pressures near atmospheric was studied by using IR laser absorption kinetic spectroscopy. Com.puter modeling with a 44-reaction mechanism was used to interpret the absorption profiles. The primary decomposition step proved to be about three times faster than had been deduced from an earlier single-pulseshock-tube study. Possible explanations for the discrepancy are discussed.
Introduction Although the thermal decomposition of propane has been investigated several times by conventional chemical kinetic techniques, only one study of the reaction in the high-temperature range has been reported1 and subjected to interpretation by computer modeling.2 This work provided stable product distributions as functions of temperature from 1100 to 1700 K, composition from 0.4 to 1.6% C3H8 in Ar, and density with starting pressures ranging from 50 to 200 torr. The computer modeling study employed to interpret these results initially included 29 elementary reactions; only 11eventually proved to be important. By concentrating on three typical sets of shock conditions, it was possible to deduce the modifications of literature rate-constant expressions required to match the experimental data satisfactorily. The key reactions proved to be the primary decomposition step C3H8 CH3 C2H5 and the two abstraction reactions CH3 C3H8 CHI + n-C3H7or i-C3H7,for which rate-constant expressions were deduced. The present study was undertaken for two reasons. First, a time-resolved measurement of the C3H8decay rate was sought to provide more direct information on the net (initial plus chain) decomposition rates than is possible with single-pulse shock-tube experiments. Second, a more extensive reaction mechanism, including recent measurements of several important rate constants, was to be investigated to see whether any pathways were overlooked in the earlier work. In this connection a complete sensitivity study3 was indicated. Our experimental and computational results agree in general with the earlier work, except that the rate of the primary decomposition step was found to be about three times faster than reported previously. Possible reasons for this difference are discussed.
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Experimental Section The experiments were done in a 4.3% C3H8-95.7% Ar mixture made from Matheson ultrahigh-purity grade Ar (99.999%) and Matheson instrument grade C3H8 (99.5%, principal impurities i-c4H10 (0.35%) and n-C4H10 (0.05%)). The shock-tube apparatus used in this investigation has been described in detail previo~sly.'~*~ Briefly, attenuation of a 3.39-pm He-Ne beam by hydrocarbons is monitored 0022-36541a012oa4-2005$0 1 .ooio
as a function of time after shock heating. For the experiments reported here, reflected shock waves were used, the laser beam being about 12 mm from the end plate. Shock-front absorptions were used to find absorptivities as functions of temperature for all stable hydrocarbons which are expected to participate (C3H8,C3H6,C2H6,C2H4, C2H2,CHI), and it was shown that CH3 does not a b ~ o r b . ~ J The absorption profiles were interpreted by computer modeling under assumption of a constant-density reaction starting from the reflected shock conditions calculated from the incident shock speed under the assumption of complete vibrational relaxation but no chemical reaction? For species where standard thermochemical properties are not available, rigid-rotor, harmonic oscillator formuhi were used together with the bond dissociation energies of BensonlO to derive thermochemical properties. Results Representative oscilloscopetrace photographs are shown in Figure 1. One sees that compression in the incident shock wave requires about 3 KS, the traversal time of the shock wave across the laser beam. It is followed first by a steep schlieren spike as the reflected shock wave passes the observation point and then by loss of absorption as the C3H8decomposes. One notes that the absorption does not decrease to zero in this temperature range, as decomposition products (primarily CH4 and C2H6)also absorb the laser line. To characterize the rates of reaction we proceeded as follows. The oscilloscope trace photographs were digitized by reading the signal voltage at suitable intervals after the center point of the schlieren spike, which defined time zero. These voltages V(t)were converted into absorbance values A ( t ) = log ( Vo/Vt)) and fit to a spline function from %which the times to reach 98 and 90% of the shock-front absorbance could be found by interpolation. Here V , i s the voltage corresponding to full laser intensity with no absorption. A measure of the reaction rate at the time of reaching 98% absorption was obtained from this spline function also by computing k98/s-l
=
[dA(t)/dt10.98A(0)([C3H815~(C3HS)d~-1
and similarly a kw value for 90% absorption. Here the concentrations are those computed for the reflected shock 0 1980 American Chemical Society
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The Journal of Physical Chemistry, Vol. 84, No. 16, 1980
Koike and Gardiner
(a )
-I
\ v)
W
-
m
Y
W
0 -1
3 -
\ \
I 7
6
IO k K / T
> E
21
Flgure 2. Arrhenius presentation of experimental and computed k,, values: (-) Table I mechanism and rate-constant expressions; (- - -) ref 2 mechanism and rate-constant expressions.
Figure 1. Laser absorption profiles for reflected shock waves in 4.3%
C,H, in Ar, p i = 20 torr: (a) T , = 1370 K, Vo = 262 mV, L = 20 mV; (b) T, = 1730 K, V, = 262 mV, L = 22 mV. Note the schlieren
c
spike on the low-temperature profile, which is not present on incident shock wave laser-absorption records (ref 6).
wave from the starting pressure and the incident shock speed, a(C3Hg)is the absorptivity of C3Hgat the reflected shock temperature, and d = 9.6 cm is the optical path length through the shock tube. The sense of this expression is that, if the only reaction occurring were the second-order decomposition of C3Hgand C3Hgwere the only absorber, the kg8so computed would be the true rate constant. Values of k g g and kw determined this way are shown as functions of temperature in Figures 2 and 3 together with calculated values of the same quantities. \
Interpretation I t was clear from the results of earlier studies of C3H8 pyrolysis that the reaction under our experimental conditions has a complicated chain mechanism. We therefore set out to model the reaction by using as complete a set of elementary reactions as possible and the most reasonable set of rate-constant expressions we could assemble. This information is set forth in Table I. The computer modeling was done on the assumption of an adiabatic, constant-density reaction.8 Values for kg8 and kgOwere calculated in accordance with the definition given in the previous section taking all absorbers of the laser radiation into consideration. The only parameter which was varied during the interpretation research was the rate constant of reaction 1. The final value selected is a factor of 3 faster than the result of Lifshitz, Scheller, and Burcat.' For comparison, values of k98 and kgocomputed by using the mechanism and rate-constant expressions of Lifshitz and Frenklach2 are also shown in Figures 2 and 3. A sensitivity spectrum was calculated for hgg and hgOat 1450 and 1650 K by using the Table I rate constant set for
7
6
IO k K / T
Figure 3. Arrhenius presentation of k g ovalues; lines as in Figure 2.
reference and multiplication and division by 5 to determine the sensitivity. The results are shown in Figures 4 and 5. Discussion The 11-reaction scheme investigated by Lifshitz and Frenklach was intended to describe C3H8thermal decomposition at 1250 K, at which temperature the olefins C3H6 and CzH4as well as CzH6and CH4 are stable products. Their 11 reactions thus comprise the unimolecular decomposition to CH3 and C2H5,the decomposition of C2H6 to C2H4+ H, the radical attacks of H and CH3upon C3H8 to form the two C,H,? four decomposition reactions of C3H7,and the recombination of CH3 to C2HP One first inquires whether it is reasonable to expect this set to be sufficient up to 1700 K. Of all the product species, C2H6
The Journal of Physical Chemisfty, Vol. 84, No. 16, 1980 2007
Thermal Decomposition of Propane
1650 K
1450 K
I
I
2 3
4 5 8
1450 K
2
13 I
_-
4
E?
5
c
8 9
IO II
9
1 10
c
II
24 27
d
-
39
C
40
c I
24 27 39
40 k 98
SCALE
m -20
I
+20
Figure 4. Sensltlvlty spectrum for kgsat 1450 and 1650 K. The numbers refer to the reactions listed in Table I. The abscissa is 100 times the pS value defined in ref 3. The upper bars were calculated by multiplying the Table I rate-constant expressions by 5, the lower bars by multiplying by 1/5. Reaction 20 also shows 100 X pS = -7 for k,, at 1450 K. Positive values correspond to reactions which accelerate the decomposition.
is the least stable; however, its dissociation is to CH3 radicals, which are present in the system anyway, and so one inquires instead whether CH3pathways may have been overlooked. Direct studies of C3H6thermal decomposition14*30 show that this does occur on the shock-tube time scale at temperatures over 1500 K, and the product shows that C3H6deanalysis for C3H8 composition by radical attack already becomes noticeable a t 1300 K. This iiitroduces the consideration of the radicals C3H, and C2H3 and the stable species C3H4. The 44-reaction scheme of Table I includes the reactions that have been discussed in the literature for these species plus some additional recombination steps (see below). In general conformation of the findings of Lifshitz and Frenklach, C3H6does then get consumed at the higher temperatures of our experiments and computations. However, it then turns out that these C3H6reactions have no effect upon the computed kg8 and kw values. The reason is twofold. First, our experiments provide information about the earliest part of the decomposition reaction, before C3H6has built up in concentration. Second, even if C3H6 does get consumed during our observation times, much of this loss is to CH4 formation; since C3H6 and CH4 have comparable absorptivity for the laser wavelength, the experiments are not very sensitive to this rea~tion.~~~ Having additional radicals in the reaction mechanism also permits us to test another aspect of the chemistry. Though chain initiation is entirely through reaction -1, a variety of possibilities exist for chain termination and for
I
k 90
SCALE
m -20 +20
Figure 5. Sensitivity spectrum for kgO.
interchanging the main chain carriers H and CH3. We therefore tested whether additional chain termination routes would affect the computed parameters. This indeed proved to be the case. In Figures 4 and 5 one notes the decelerating effects of reactions 24, 27, 39, and 40. Once a basic understanding was reached that the 44reaction mechanism functions essentially the same way as does the ll-reaction mechanism of Lifshitz and Frenklach, we were able to adjust the rate constants to fit the experimental kss and kN results in an understanding manner. It was soon found that the only change required over the literature rate constants adopted originally was to increae the Lifshitz and Frenklach value of kl by a factor of 3. This resulted in the match to the data shown in Figures 2 and 3. These authors also found it possible to match the product distribution and rate constant at 1250 K with kl = 1012.72 cm3 mo1-ls-l as contrasted to their final result, 1012.40 cm3 mo1-l $-1, and our result, 1012.sscm3 mol-l s-l, and so the agreement of the Table I mechanism with their data may actually be much closer than its agreement with their value of kl, considering the offsetting effect of the additional termination reactions included here. We note in this connection that the fate of CH3 radicals at the higher temperatures considered here is certainly quite different from that found by Lifshitz and Frenklach to prevail at 1250 K. In particular, self-destruction by reactions -24 and 27 is now known, by direct measurements of H and CH3 profiles in simpler sy~stems,3~J~ to be important at the higher temperatures of our study. (The rate constant we assumed for reaction 27 is the same as that reported in ref 31 at 1700 K.) The temperature is not yet high enough that thermal dissociation of CzH4olr CH3 need be c o n ~ i d e r e d .Incorporation ~~~~~ of more recent expressions for the known reactions of CH3 would clearly have some effect upon our conclusions, as would removal
2008
The Journal of Physical Chemistty, Vol. 84, No. 16, 1980
Koike and Gardiner
TABLE I: Reaction Mechanism and Rate Constant Parametersa no. 1 2 3 4 5 6 7 8 9
10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 a
reaction CH, t C,H, = C,H, CH, t C,H, = CH, t i-C,H, CH, t C,H, = CH, t n-C,H, H t C,H, = H, t i-C,H, H t C,H, = H, t n-C,H, C,H, t C,H, = C,H, t i-C,H, C,H, t C,H, = C,H, t n-C,H, i-C,H, = C,H, t CH, n-C,H, = C,H, t CH, i-C,H, = H t C,H, n-C,H, = H t C,H, C,H,- CH, t C,H, C,H, = C,H, t H CH, t C,H, = CH, t C,H, C,H, t C,H, = C,H, t C,H, C,H, t C,H, = C,H, t C,H, H t C,H, = H, t C,H, C,H,= C,H, t CH, C,H, t H = C,H, C,H, t M = CH, t CH, t M CH, t C,H,F CH, t C,H, H t C,H, = H, t C,H, C,H, t M = C,H, t H t M H t C,H, = CH,.t CH, CH, t CH, = C,H, t H CH, t CH, = C,H, t H, CH, t CH, = C,H, t H, C,H, t H = C,H, t H, CH,tM=CH,tHtM CH, t H = CH, t H, C,H, t M = C,H, t H, t M C,H, t M = C,H, t H t M C , H , t C , H , = C , H , t C,H, C,H, t H = C,H, t H, C,H, t C,H, = C,H, t H C,H, t H = C,H, C,H, t H = C,H, t H, H, t M = H t H t M CH, t C,H,= CH, t C,H, C,H, t C,H, = C,H, t C,H, CH, t C,H,= CH, t C,H, C,H, t C,H, = C,H, t C,H, C,H, t C,H, = C,H, t C,H, C,H, t C,H, = C,H, t C,H,
AH" o , k J
-352 -4 1 - 27 -41 - 27 -15 -2 116
102 158 144 394 327 105 - 108 - 80 - 105 230 - 266 367 - 25 - 25 155 - 40 65 40 - 237 - 277 432 0 166 435 280 3 0 - 163 - 269 432 -6 1 - 252 3 28 - 272 - 244
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log A 12.86 12.55 12.55 13.80 13.80 10.70 10.70 12.00 13.80 14.30 13.80 16.10
15.00 10.30 11.00 11.00 12.40 13.48 11.38 111.29 - 25.26 log T -0.26 t 4.0 log T 14.12 15.31 13.57 12.90 12.00 12.00 12.27 17.46 14.86 17.41 17.58 14.82 15.70 12.00 12.74 13.47 12.35 - 0.5 log T 13.00 13.00
11.00 12.00 13.00 13.00
Concentration units are mol ern-,.
of the speculative disproportionation steps 37, 39,40,43, and 44. It would appear unreasonable, however, to undertake further adjustments of rate constants for the important reactions of C3H8decomposition based upon the available data and the type of modeling used by Lifshitz and Frenklach and by us. The reason for this is that the mechanism includes a number of unimolecular decomposition steps, some written ashnimolecular and some as bimolecular steps, but all actually being somewhere in the falloff region for the densities and temperatures considered. Although simple Lindemann-Hinshelwood falloff corrections can readily be incorporated into modeling calculation^,^^ a correction that comes close to fully parameterized RRKM falloff behavior requires more substantial programming efforts.35 Until all unimolecular steps have been subjected to falloff considerations, further refinement of the rate-constant expressions for C3H8 thermal decomposition is not justified. I t also appears unreasonable to compare our result for kl critically with theoretical estimates or rate constants for similar recombination reactions. In this temperature range the high-pressure limit rate constant for CH3 reomb bin at ion^^ is near 1013.5*0.2 cm3 mol- s-l, from the results of many experiments and in good accord with theory. For C2H5 recombination, however, independent methods
EA,
kJ
ref
this work
0 43 43 33.5 33.5 43.5 43.5 144 139 173 159 3 59 368 30 29 34 4.6 152 0 668 35 39 126 0 167 96 0
2 2
11 11 12 12 13 13 13 13 14
13 16 17 17 18 19 20 4 21 22 4 22 23 23
see text
0
22 23 24 25 25 26 25 27 28
379 63 332 411 268 96 31
10 0
see text
388 0 0 42 46 0
29
see text see text 20 17
see text
see text
9
give rate constants ranging from to 1013.6cm3 mol-' s-' with no agreement either on whether an activation energy is evident.3740 Our result, 1012.86, and Lifshitz and Frenklachs result, 1012.4, for CH3 C2H6recombination would suggest that the higher values are not correct, but otherwise the available rate constant data are too discordant to admit detailed analyses. The sensitivity to the rate of reaction -29 noted by Kao and Yeh41for the Lifshiftz and Frenklach conditions was not apparent for our experiments as based upon the Table I mechanism. Acknowledgment. Acknowledgment is made to the donors of the Petroleum Research Fund, administered by the American Chemical Society, for partial support of this research. This research was also supported by the Robert A. Welch Foundation and the U S . Army Research Office.
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References and Notes (1) A. Lifshitz, K. Scheller, and A. Burcat, Roc. Int. Shock Tube Symp., 7th, 7969, 690 (1973). (2) A. Lifshitz and M. Frenklach, J. Phys. Chem., 79, 686 (1975). (3) W.C. Gardiner, Jr., J . Phys. Chem., 81, 2367 (1977). (4) J. H. Owen and W.C. Gardiner, Jr., Proc. Int. Shock Tube Symp., 9th, 1973, 720 (1973). (5) D. B. Olson, T. Tanzawa, and W.C. Gardiner, Jr., Int. J. Chem. Klnet., 11, 23 (1979). (6) D. B. Olson, W.G. Mallard, and W.C. Gardiner, Jr., Appl. Spectrosc., 32,489 (1978).
J. Phys. Chem. 1980, 84, 2009-2015
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(22) P. Camilleri, R. M. Marshall, and J. H. Purnell, J. Chem. Soc.,Faracky Trans. 7, 70, 1434 (1974). (23) W. C. Gardlner, Jr., T. Koike, and T. Tanzawa, J. Phys. Chem., in preparation. (24) Th. Just and P. Roth, Ber. Bunsenges. Phys. Chem., 79,682 (1975. (25) Th. Just, P. Roth, and R. Damm, Symp. (Int.) Combust., [Pmc.], 16, 961 (1977); P. Roth and Th. Just, Ber. Bunsenges. Phys. Chem., 77, 114 (1973). (26) M. L. Boyd, T. M. Wu, andM. H. Back, Can. J. Chem., 46, 2415 (1968). (27) T. Tanzawa and W. C. Gardiner, Jr., Combust. Flame, in press. (28) W. A. Payne and L. J. Stief, J. Chem. Phys., 84, 1150 (1976). (29) A. L. Myerson and W. S. Watt, J. Chem. Phys., 49, 425 (1968). (30) A. Burcat, Fuel, 54, 87 (1975). (31) T. Tsuboi, Jpn. J. Appl. Phys., 17, 709 (1978). (32) K. A. Bhaskaran, P. Frank, and Th. Just, Int. Shock Tube Symp., 12th, 7979, in press. (33) P. Roth, U. Earner,and R. LGhr, Int. Shock Tube Symp., 12th, 1979, in press. (341 J. M. Whiie and W. C. Gardiner. Jr.. J. Phvs. Chem.. 83. 562 11979). ( 3 4 W. C. Gardiner, Jr. and J. Troe, in preparation; J: Troe, J.‘Phys. Chem., 83, 114 (1979). (36) D. B. Olson and W. C. Gardlner, Jr., J. Phys. Chem., 83, 922 (1979). (37) A. SheDD and K. 0. Kutschke, J. Chem. Phvs.. 26, 1020 (1956). . . (38j R. W. Fessenden, J . Phys. Chem., 68, 1508 (1964). (39) R. Hlatt and S. W. Benson, J. Am. Chem. Soc., 94, 25 (1972). (40) R. Hiatt and S. W. Benson, J. Am. Chem. Soc., 94, 6886 (1972). (41) W. Kao and C. Yeh, J. Phys. Chem., 81, 2304 (1977).
(7) T. Koike arid W. C. Gardiner, Jr., Appl. Spectrosc., in press. (8) W. C. Gardiner, Jr., B. F. Walker, and C. B. Wakefield in “Shock Waves in Chemistry”, A. Litshltz, Ed., Marcel Dekker, New York, 1980, Chapter 1. (9) D. R. Stull and H. Prophet, Nafl. Stand. Ref. Data Ser. ( U . S . ,Natl. Bur. Stand.), No. 37 (1971); J. Chao, R. C. Wilhoit, and B. J. Zwolinski, J. Phys. Chem. Ref. Data 2 , 427 (1973); R. Main, NASA Rep., CR-72560, Appendix B. (10) S. W. Benson, “Thermochemical Kinetics”, Wlley, New York, 1968. (11) A. F. Trotman-Dickenson and G. S. Milne, Natl. Stand. Ref. Data Ser. (U.S., Natl. Bur. Stand.), No. 9, 108 (1967). (12) D. L. Allara and D. Edelson, Int. J. Chem. Kinet., 7, 479 (1975). (13) S. W. Benson and H. E. O’Neil, Natl. Stand. Ref. Data Ser. ( U . S . , Nafl. Bur. Stand.), No. 21 (1971). (14) G. A. Chappell and H. Shaw, J. Phys. Chem., 72, 4672 (1968). (15) T. Kunugi, K. Soma, and T. Sakal, Ind. Eng. Chem. Fundam., 9, 319 (1970). (16) R. .I. Cvetanovic and R. S. Irwin, J. Chem. Phys., 46, 1694 (1967) (17) T. Kunugi, T. Sakai, K. Soma, and U. Sasaki, Ind. Eng. Chem. Fundam., 8, 374 (1969). (18) V. V. Voevcdsky and Y. N. Kondratlev, Pmg. React. Kimf., 1, Chapter 2 (1961). (19) K. M. Sundaram and G. F. Froment, Ind. Eng. Chem. Fundam., 17, 174 (1978). (20) A. F. Trotman-Dickenson and E. W. R. Steacie, J. Chem. Phys., 19, 169 (1951). (21) T. C. Clark and J. E. Dove, Can. J. Chem., 51, 2147 (1973). See also C. J. Chen, M. H. Back, and R. A. Back, ibid., 54, 3175 (1976).
Temperature-Dependent Rate Constants for the Reaction of Ground-State Chlorine with Simple Alkanes R. S. Lewis,+ S. P. Sander, S. Wagner, and
R. T.
Watson”
Jet Propulslon Laboratory, California Institute of Technology, Pasadena, California 9 1 103 (Received: September 4, 1979)
The low-pressure discharge-flow-resonance fluorescence technique has been utilized to study the rates and temperature dependences of three chlorine atom-alkane reactions. The reactions have been studied by using a wide range of experimental conditions to ensure the absence of complicating secondary processes. The Arrhenius expression for each bimolecular reaction is expressed in units of cm3 molecule-’ s-’: (1)C1 C2H6 C2H5 + HC1, Izl = (9.01 f 0.48) X exp(-(133 f 15)/T, 220-604 K; (2) C1+ C3H8 C3H7+ HC1, k2 = (1.36 f 0.13) X exp(+(44 f 2 5 ) / T ) , 220-607 K (3) C1+ n-C4Hlo C4H9 + HCl, Iz3 = (2.15 f 0.10) X exp(S(l2 f 26)/2‘), 298-598 K. In addition, the following reaction was studied at 298 K: (4) C1 + i-C4H10 C4H9+ HCl, k4 = (1.46 f 0.06) X lO-’O. The present results are compared with earlier absolute and relative rate constant measurements.
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Introduction Reactions involving abstraction of a hydrogen atom from an RH molecule by atomic chlorine have been of interest to kineticists for many years. In addition to the need for accurate experimental rate data in order to test current kinetic theories, recent interest in these reactions stems from the now well-known stratospheric chlorine-ozone problem.’ This renewed interest in reactions between atomic chlorine and hydrogen-containing molecules (C1+ RH HC1 t R) is due to the role of HC1 as a reservoir for active chlorine in the stratosphere. This has led to numerous studies of the kinetic behavior of atomic chlorine with molecular hydrogen2* and m e t h a r ~ e ~ fover i J ~a~wide range of temperature and pressure using direct kinetic techniques. However, except for a single temperaturedependence study with ethane,14 the reactions of higher alkanes with atomic chlorine have received relatively little recent attention. The introduction of gas chromatography allowed competitive chlorination experiments between molecular hy-+
NASA/NRC Resident Research Associate, 1976-1978. 0022-3654/80/2084-2009$01 .OO/O
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drogen, alkanes, and chlorinated alkanes to be studied by product analysis,21i22 thus providing more accurate rewlts than those available using the consumption m e t h ~ d . ~ ~ J ~ The early work is reviewed by Fettis and K n o ~ In . ~ad~ dition to activation energy differences and Arrhenius A factor ratios determined in the competitive chlorination studies, Fettis and Knox derived absolute Arrhenius parameters for the competitors by employing the absolute exp(-2750/7‘) cm3 Arrhenius expression of 1.38 X molecule-’ s-.l for the C1 H2 reaction as the primary standard. However, as this Arrhenius expression is now thought to be incorrect,26-28 the competitive chlorination results should be combined with the best available Arrhenius expression for the C1 + H2 reaction to yield reevaluated absolute Arrhenius parameters for the competitors. The present study utilized the discharge-flow-resonance fluorescence technique to investigate the kinetic behavior of ground-state atomic chlorine with ethane, propane, and n-butane over a range of temperature and with isobutane at room temperature. The results obtained from the present study will be compared to the reevaluated com-
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0 1980 American Chemical Soclety
