J . Phys. Chem. 1989, 93, 5 184-5 188
5184
Kinetics and Thermochemistry of C2H3 and CH2CCI Radicals: Study of the C2H3 = C2H, CI and CH2CCl HCI = C2H3CI CI Reactions
+
+
+
+ HCI
J. J. Russell, S. M. Senkan,* Department of' Chemical Engineering, Illinois Institute of' Technology, Chicago, Illinois 6061 6
J . A. Seetula, and D. Gutman* Department of Chemistry, The Catholic University of America, Washington, DC 20064 (Receiced: December 19, 1988)
The kinetics of the reactions of C2H3 and CH2CCl radicals with HC1 have been studied in a tubular flow reactor coupled to a photoionization mass spectrometer in the temperature range 298-495 K. Each radical was produced by the pulsed homogeneous photolysis of a suitable precursor using an excimer laser. Radical decay and product growth profiles were monitored in time-resolved experiments. The Arrhenius rate expression obtained from the measured C2H3 + HCI rate constants is 6.6 (f1.0)X exp(670 (k.230) cal mol-'/RT) cm3 molecule-' s-I. The heat of formation (67.1 (10.6) kcal mol-') and entropy (55.9 (f2.6) cal mol-' K-I) of C2H3were determined from this kinetic information on the C2H3 + HC1 reaction together with recently measured rate constants for the CI + C2H4reaction. A Third Law treatment of the same data using a calculated entropy yields essentially the same heat of formation (66.9 (*0.3) kcal mol-') but with reduced uncertainty limits. The rate constant for the CHzCCl + HCI reaction was found to be temperature independent, having an average cm3 molecule-' s-'. This result was used to obtain a lower value over the temperature range covered of 2.2 (f0.3) X limit for the CH2CCl heat of formation of 60 (h0.3) kcal mol-'. The reduced reactivity caused by chlorine substitution at the radical site is briefly discussed.
Introduction Recently we have begun to investigate the kinetics and thermochemistry of polyatomic free radicals in which chlorine atoms are incorporated into one or both reactants.l In addition to its fundamental value, this information is needed for the development of detailed and quantitative chemical kinetic mechanisms for many chemical processes including the pyrolysis and oxidation of chlorinated hydrocarbons (CHC'S),~.~ the pyrolysis of hydrocarbons in the presence of HC1,4 and the laser-initiated reactions of chlorinated e t h a n e ~ . ~These .~ processes have significant commercial manufacturing potential, e.g., for the production of C2 hydrocarbons from natural gas' and for producing chlorinated ethylenes.s,6 This same information is also needed to trace the degradation of chlorine-containing hazardous material undergoing thermal decomposition and oxidation. Such knowledge can be used to develop strategies for disposing of these materials by combust ion .s We have now investigated the kinetics of two related reactions, those of vinyl and I-chlorovinyl radicals with HCI:
+ HCI CHZCCI + HCI C2H3
+
+
+ C1 C2H3C1 + C1
C2H4
-
(1)
(2)
(The CI-atom-transfer reactions are 12 kcal mol-' endothermic and are not competing routes under our experimental conditions.) Both reactions are important elementary steps in the processes discussed above. Vinyl radicals are believed to be among the most important intermediates initiating polymerization processes that lead to soot f o r m a t i ~ n . ~ , 'Knowledge ~ of the reactions that ( I ) Russell, J. J.; Seetula, J. A.; Senkan, S. M.; Gutman, D. In!. J . Chem. Kinet. 1988, 20, 759. (2) Karra, S. 8.; Senkan, S. M. Ind. Eng. Chem. Res. 1988, 27, 1163. (3) Chang, W. D.; Senkan, S. M. Enoiron. Sci. Techno/. in press. ( 4 ) Lyon, R. K.; Mitchell, J . E. Modification of Thermal Alkylation with HCI. Adu. Chem. Ser. 1979, N o . 183, 289. (5) Wolfrum. J . Laser Chem. 1986, 6, 125. (6) Scheneider, M.; Wolfrum, J. Ber. Bunsen-Ges. Phys. Chem. 1986, 90, 1058.
(7) Granada, A,; Karra, S.; Senkan, S. M. Ind. Eng. Chem. Res. 1987, 26, 1901. (8) Senkan, S. M. In Detoxication of Hazardous Waste; Exner, J. H., Ed.; Ann Arbor Science: Ann Arbor, MI, 1982; Chapter 3. (9) Frenklach, M.; Clary, D. W.; Yuan, T.; Gardiner, Jr. W . C.; Stein, S. E. Combust. Sei. Technol. 1986, 50, 79.
0022-3654/89/2093-5 184$01.50/0
compete with these polymerization steps is essential for obtaining a quantitative understanding of the sooting properties of flames and other oxidation processes involving chlorinated fuels and feedstocks. Reactions 1 and 2 are examples of important competing reactions of this sort once HC1 has accumulated during the pyrolysis or oxidation of the chlorine-containing fuels. Both reactions 1 and 2 are nearly thermoneutral. Hence, under most process conditions, the forward and reverse reactions are both important elementary steps. Knowledge of the thermochemistry of these reactions (basically knowledge of the thermochemistry of the vinyl and 1-chlorovinyl radicals since that of the remaining species is well-known) combined with knowledge of the kinetics of these reactions in one direction provides the information needed to describe the behavior of these elementary reactions in any chemical process. The kinetics of reaction 1 has been studied recently while that of reaction 2 has not been investigated before. Krueger and Weitz" measured k , between 278 and 370 K in time-resolved experiments, ones in which the growth of C2H4was monitored by diode laser absorption. Parmar and Benson12 obtained information on the kinetics of reaction 1 using a very low pressure reactor. They determined values of k-' and K 1 (the equilibrium constant) at six temperatures between 263 and 338 K. From the results of the latter study, values of k l can be easily calculated. The values of k , obtained from these two studies are displayed in Figure 1. Over the common temperature range of the two investigations, the values obtained are between a factor of 2.5 and 4 apart. The thermochemistry of C2H3 has been widely studied and recently reviewed.l3-I6 There continue to be significant disparities in the values of the heat of formation of C2H3 derived from different experiments, values which range from 60 to 72 kcal mol-'. The origins of some of these disparities have also been dis(IO) Harris, S. J.; Weiner, A. M.; Blint, R. J. Combust. Flame 1988, 72, 91. ( 1 1 ) Krueger, H.; Weitz, E. J . Chem. Phys. 1988, 88, 1608. (12) Parmar, S. S.; Benson, S . W. J . Phys. Chem. 1988, 92, 2652. (13) Berkowitz, J.; Mayhew, C. A.; Ruscic, B. J . Chem. Phys. 1988, 88, 7396. (14) Shiromaru, H.; Achiba, Y.; Kimura, K.; Lee, Y.T. J . Phys. Chem. 1987, 91, 17. ( I 5) Skinner, G. B.; Rao, V. S. J . Phys. Chem. 1988, 92, 63 13. (16) Sharma. R. B.; Semo, N. M.; Koski, W. S . Int. J . Chem. Kinet. 1985, 17, 831.
0 1989 American Chemical Society
The Journal of Physical Chemistry, Vol. 93, No. 13, 1989 5185
Kinetics and Thermochemistry of CzH3 and CHzCCl
I
T(r0 5:5
300
423
i
1
53;
H
‘Ci-
T
clc 1:
o m 0 e ~ lno a s
Th 5 S.2CY
j
I
1
20
30
lO0Oi T
iIVE t ~ s e c l
I
40
1
1
I
I
0
5
10
15
[HCI]
(K.’)
Figure 1. Arrhenius plot of measured or determined C2H3+ HCI rate constants: open triangles, ref 11; open squares, ref 12. Solid line through current k , values from Arrhenius expression presented in Table I.
-14
x
10
(molecules cm’)
Figure 2. Plot of the CH2CCIdecay constant vs [HCI] for experiments conducted at 342 K. Inserts are reactant and product signal intensity profiles recorded during experiments conducted at [HCI] = 1.53 X l O I 5 molecules cm-) (closed circle in the figure). Lines through the ion signal data are exponential functions from nonlinear least-squares fits. These fits resulted in a CHzCCl decay constant of 440 s-I and a C2H,CI growth constant of 442 s-l.
The “chemical kinetic method”, one based on determining AHo and ASofor a reaction involving the radical of interest from measurements of the temperature dependence of the reaction equilibrium constant, was used by Parmar and Benson.lZ containing the mass spectrometer. As the gas beam transversed This method is free of effects that complicate the interpretation the ion source, a portion was photoionized and then mass selected. of experiments that rely on establishing ionization thresholds, Temporal ion signal profiles were recorded from a short time effects which appear particularly serious in the case of C Z H ~ . ’ ~ ~ ’ ~ before each laser pulse to -25 ms after the pulse with a multiParmar and Benson obtained a heat of formation for C2H3 from channel scaler. Data from 5000 to 30000 repetitions of the a modified chemical kinetic method, a Third Law analysis of their experiment were accumulated before the data were analyzed. results which did not require knowledge of the temperature deGas flow velocities used were near 5 m s-I. This assured that pendence of K l but did require prior knowledge of the C2H3 the gases in the reactor were replaced between laser pulses. entropy. (A Second Law analysis by Parmar and Benson using Initial conditions in these experiments were chosen to essentially K,(7‘) yielded a CzH3 entropy which was anomalously high.) isolate the reaction of interest for kinetic study. The elementary There has been no prior determination of the heat of formation processes consuming the vinyl and chlorovinyl radicals were the of CHZCCI. reaction with HCI (reaction 1 or 2) and a heterogeneous first-order In the current investigation, we report new measurements of wall loss process (characterized by the first-order constant kw). the rate constants of reaction 1 in part to resolve the current By use of extremely low initial free-radical concentrations (typdisparity in measured values but also to provide the kinetic in), and radiically (2-5) X lolo molecules ~ m - ~radical-radical formation needed to directly obtain the CzH3 heat of formation cal-atom reactions were suppressed; Le., they had negligible rates by using the chemical kinetic method without requiring any compared to the reaction under study. This condition was achieved thermochemical information about this radical. The kinetics of and verified by reducing the concentration of the radicals (either reaction 2 was also investigated to learn the effect of chlorine by attenuation of the photolysis laser intensity or by lowering the substitution in the vinyl radical on the rate constant of this reaction. free-radical precursor concentration) until the measured radical The latter study only provided a lower limit for the CHzCCl heat decay constant (in the absence of HCI) no longer depended on of formation because k-2 has not as yet been determined. these variables. Under the pseudo-first-order conditions used in these experiExperimental Section ments (HC1 in excess), both the vinyl and chlorovinyl radicals The details of the experimental apparatus used and the prodecayed exponentially in both the absence and presence of HCI. cedures employed both to identify mechanisms and to measure A rate constant for a particular temperature and gas density was rate constants have been described previously.” Therefore, only determined from the slope of the line fitted through the measured a summary is presented here. Pulsed ( 5 Hz) unfocused 193-nm radical exponential decay constants (k’) vs [HCI] (from [vinyl], radiation from a Lambda Physik EMG lOlE excimer laser was = [vinyl],, exp(-k’t)). According to the model, k’ is a linear directed along the axis of a heatable, 1.05-cm-i.d., tubular reactor function of [HCI]; k’= kl[HCI] + k, (or kz[HCI] + kw). This coupled to a photoionization mass spectrometer. The gas flowing expected linear relationship was observed in all the experiments. through the tube contained the radical precursors in low conA representative ion signal decay profile and decay constant centrations ( 50) and an excess of the carrier gas 2. Also shown in the second insert in this figure is the growth (He, >97.5%). Laser fluence was attenuated (by a factor of 2) profile of the product, CzH3C1. to achieve negligible HCI photolysis. The effect of total gas pressure on both k l and kz was also The gas from the reactor was sampled through a 0.40-mminvestigated. Over the range of conditions covered (3- and 1.3-fold diameter hole in the side of the reactor and formed into a beam changes in pressure for reactions 1 and 2, respectively), no by a conical skimmer before it entered the vacuum chamber measurable effect of pressure on either rate constant was observed. The conditions of experiments and results obtained from all experiments are presented in Table I. (17) Slagle, 1. R.; Gutman, D. J . A m . Chem. SOC.1985, 107, 5342.
5186
The Journal of Physical Chemistry, Vol. 93, No. 13. I989
TABLE I: Conditions and Results of Rate Constant Experiments wall [He] X [HCI] X IO-'), k,, coating k X IOi2, 7,K atoms molecules c ~ n ' ~ s-' material" cm3 molecule-' C2H3 HCI C2H4 + C1 4.34-20.2 101 HW 2.03 297 4.80 HW 2.07 297 4.80 101 4.00-19.2 4.84 5.89-21.6 HW 2.1 1* 101 297 4.80 3.24-1 9.7 99 T 2.02 298 3.49-18.2 109 T 2.05 296 15.0 HW 1.93 4.98 316 3.68-22.9 98 HW 1.76 3.96-20.3 87 4.87 342 3.6 1-20.5 91 T 1.81 4.87 342 1.54 372 3.97-28.7 100 FHW 4.83 93 4.92 T 1.61 372 4.50-26.9 T 1.60 372 3.41-27.0 91 4.91 4.12-26.7 1.39 104 T 4.95 406 T 1.37 450 3.51-27.0 91 4.97 T 1.49 4.77-27.9 101 5.03 450 4.93-27.9 106 T 1.34 495 5.01
+
k , = 6.6 ( f l . O )
297 297 316 342 342 372 406 450 495
9.91 7.61 9.88 9.94 7.91 10.0 8.47 8.46 8.79
X
wavelength proceeds via two routes having about equal importance:21 5-l
C2H3Br
+
exp(670 (f230) cal mol-'/RT) cm3 molecule" s-I
CH2CCI + HCI 29.5-1 98 23.8-174 24.7-171 23.2-178 3 1.2-1 85 25.1-167 26.2- 17 5 29.4-1 69 27.0-190
Russell et al.
-
+
C2H3CI CI 103 H W 88 T 87 HW 97 HW T 96 108 FHW 96 T 100 T 90 T
0.219 0.235 0.218 0.217 0.222 0.218 0.225 0.224 0.225
k2 = 2.2 ( f 0 . 3 ) X IO-" cm3 molecule'' s-' Reactor wall coatings: halocarbon wax (HW), fluorinated halocarbon wax (FHW), and Teflon (T). * N 2 used as carrier gas.
In the study of reaction I , the effect of the bath gas on k , was also tested at 297 K. Helium was replaced by nitrogen. The measured value of k l when N2 was used was the same as that obtained when He was the bath gas. The independence of k , and k 2 on density and on the identity of bath gas used is considered to be the confirmation that radical internal energy relaxation is complete before reaction occurs. Three different materials were used to coat the interior wall of the tubular reactor to reduce the unimolecular heterogeneous loss of C 2 H 3and to test for possible bimolecular heterogeneous loss. The materials were Teflon,'* halocarbon wax,I9 and fluorinated halocarbon wax.2o While each coating material reduced the unimolecular heterogeneous vinyl radical decay constant k,, it did not affect the values of k , or k2 determined in these experiments. The data analysis takes into account the unimolecular heterogeneous process but cannot discriminate between homogeneous and heterogeneous bimolecular reactions unless the rate of the heterogeneous process can be altered selectively. The independence of k , and k2 on the nature of the wall coating (or its absence) is interpreted as indicating that heterogeneous bimolecular reactions were of negligible importance. The gases were obtained from Linde (He, >99.995%) and Matheson (HCI, >99.998%; C2H3Br,>99.5%) and were used as provided. 1 , I -C2H2CI2(purity >99%) was obtained from Kodak and was purified by fractional distillation. Three photoionization energies were used in this investigation. A chlorine resonance lamp (8.9-9.1 eV) was used to detected C2H3 and CH,CCI radicals, a hydrogen lamp (10.2 eV) for the detection of C2H3,C2H3CI,and 1 , l -C2H2C12,and an argon resonance lamp ( 1 1.6-1 1.8 eV) for the detection of C 2 H + A . Study ofthe C2H3 + HCI Reaction. I . Photolytic Source of C2H3. The C2H3 radicals were produced by the 193-nm photolysis of C2H3Br. About 15% of the precursor photodecomposed during each laser pulse. The photodecomposition at this (18) Teflon 851.201. (19) Halocarbon Wax Corp., Hackensack, NJ, Series 1500. (20) Badechhape, R. B.; Kamarchik, P.; Conroy. A. P.; Glass, G. P. Inr. J. Chem. Kinet. 1916. 8, 23.
€
C2H3
+ Br
(3a)
C2H2
+ tBr
(3b)
2. Rate Constant Determinations. Reaction 1 was studied between 297 and 495 K, the upper temperature being determined by the thermal stability of the Teflon wall coating material. The rate constants measured for reaction 1 are presented in Figure 1 along with the values reported by others. The conditions of the experiments are given in Table I. A least-squares fit of the k l values obtained in the present study yields the following Arrhenius expression:
k , = 6.6 ( f l . 0 ) X exp(670 (f230) cal mol-'/R7')
cm3 molecule-' s-I
The indicated error limits in the Arrhenius expressions for k , and k, take into account the accuracies of measurements (gas flow rates and total pressure) and data analysis procedures including the accuracy of the radical decay constant determinations. The most probable uncertainty in each k , and k2 determination is f 15%. 3. Products of the C2H3+ HCI Reaction. The detectable product of the C2H3 HCI reaction, C2H4,was searched for in an experiment conducted at 417 K. It was detected and its temporal behavior recorded. The exponential growth constant for C2H4was essentially the same as the decay constant of C2H3 under the same conditions, confirming that C2H4is a primary product of reaction 1. B. Study of the CH,CCI HCI Reaction. 1 . Photolytic Source of CH2CCI. The CH2CCI radicals were generated by the 193-nm laser photolysis of 1,1-C2H2C12. Nearly 10% of the precursor photodecomposed when unfocused radiation from the photolysis laser was used. Experiments were conducted to identify the major photolysis pathways. The products C2HC1, C2H2,and CH2CC1appeared instantly following the laser pulse, indicating the presence of three photolysis routes:
+
+
I ,I-C,H,CI,
r -t--
C,HC1 C,H,
+
(4)
HCi
+ 2c1(or + C),I
CH,CCI
+CI
(4b) (4c)
No experiments were conducted to establish the relative importance of routes 4a-4c largely because this information was not required to obtain the kinetic information sought in this investigation. All three routes appear to be of comparable significance solely on the basis of the magnitudes of the recorded product ion signals. 2. Rate Constant Determinations. Reaction 2 was studied at seven temperatures between 297 and 495 K. Again, the stability of the Teflon coating determined the upper temperature limit of these experiments. A sample CH2CCI+ion signal decay profile is shown in Figure 2 along with a plot of k'vs [HCI] obtained for a set of experiments performed to measure k,. The conditions of these experiments and the results obtained are also presented in Table I. The values of k2 were essentially temperature independent, having an average value
k 2 = 2.2 (f0.3)
X
cm3 molecule-l s-'
3. Products of the CHICCl + HCl Reaction. The expected product which can be detected, C2H3CI,was monitored in experiments conducted at both 377 and 540 K. It had to be monitored at m / e = 64 (corresponding to C,H3Cl3'), because at m / e = 62 there was interference from the photolysis product C2HC1.37 (21) Slagle, I. R.; Park, J.-Y.; Heaven, M. C.; Gutman, D. J. Am. Chem. SOC.1984. 106. 4356.
The Journal of Physical Chemistry, Vol. 93, No. 13, 1989 5187
Kinetics and Thermochemistry of C2H3 and CH2CCl The C2H3CI+ion signal rose exponentially, again mirroring the observed decay of CH2CCl (see the second insert in Figure 2). The measured exponential growth constant of C2H3CIwas the same as the decay constant of the CH2CCl radical, consistent with the mechanism of reaction 2.
Thermochemistry of Vinyl and I-Chlorovinyl Radicals A. Thermochemistry of C2H3. It is possible to determine both the heat of formation and the entropy of C2H3 from the kl values reported here with the use of the k-, determinations reported by Parmar and Benson.I2 To obtain the heat of formation, either a Second Law or Third Law procedure can be used (the former requiring no significant additional information and the latter requiring the entropies of the reactants and products in reaction 1).22 Both procedures are used here to obtain the C2H3 heat of formation, the latter to demonstrate the consistency of the two kinds of determinations and to obtain a value with reduced limits of uncertainty. 1 . Second Law Determination. Reactions 1 and -1 were not investigated over the same temperature range. The overlapping temperatures of the two studies are centered near 310 K. We associate the difference in the measured activation energies for the forward and reverse reactions with the standard enthalpy change for the reaction at 310 K (the activation energy for reaction -1 was reported to be 2.8 (f0.4) kcal mol-'):12 AH'(310) = -0.67 (f0.23) - 2.8 (f0.4) = -3.47 (f0.63) kcal mol-' Tabulated heat capacity function^^^,^^ for reactants and products in reaction 1 were then used to modify this enthalpy change slightly to yield the following value at 298 K: AH'(298) = -3.45 (f0.63) kcal mol-' The heat of formation of C2H3 was then determined from the well-established standard state heats of formation of C2H4, HCI, and Cl:22*23 AHf0(298) = 67.1 (f0.6) kcal mol-] The entropy of C2H3at 298 K was obtained from the same rate constants. First, the equilibrium constant for reaction 1 ( K , = k l / k - , ) at 298 K (a temperature common to both studies) was obtained from the reported Arrhenius expression for k l and kI: K,(298) = (2.03 (f0.3)
X
10-12)/(4.6 (f0.5)X
= 4.4 ( f l . 1 )
The standard Gibbs free energy change for reaction 1 was then obtained from this equilibrium constant: AG'(298) = -870 (f170) cal mol-, The entropy change of reaction 1 at 298 K was obtained from this free energy change and the previously calculated value of AH'(298): AS'(298) = -8.6 (f2.6) cal mol-' K-' Finally, the absolute entropy of C2H3 was calculated from the well-established standard-state entropies for C2H4, HCI, and c]:23,24 S'(298) = 55.9 (f2.6) cal mol-' K-'
2. Third Law Calculation. There have been ab initio theoretical studies of the structure and internal motions of c2H3.25-27 This (22) Benson, S. W. Thermochemical Kinetics; Wiley: New York, 1976. (23) Burcat, A. In Combustion Chemistry; Gardiner, Jr., W. C., Ed.; Springer-Verlag: New York, 1984. (24) J A N A F Thermochemical Tables, 3rd ed.; J . Phys. Chem. ReJ Dura 1985, 14. (25) Harding, L. B.; Wagner, A. F.; Bowman, J. M.; Schatz, G. C.; Christoffel, K. J . Phys. Chem. 1982, 86, 4312.
information was used to calculate the C2H3 entropy, S'(298) = 55.5 (f0.5)cal mol-, K-I. The relatively low estimated error in the calculated entropy is due largely to the absence of any vibration frequencies in C2H3lower than about 740 C ~ - I . ~ ' A significant uncertainty in such a frequency (f100 cm-I) results in only a 0.1 cal mol-' K-' uncertainty in the entropy at this temperature. The value of ASo for reaction 1 at 298 K obtained by using this entropy combined with the value of AG'(298) calculated from values of k , and k-, at this temperature yields a new value for AH'(298) and the heat of formation of the C2H3 radical: AH,0(298) = 66.9 (h0.3) kcal mol-' The error limits reported in this section for the C2H3 heat of formation and entropy are maximum possible uncertainties. They reflect the cumulative uncertainty from all primary information used to calculate each thermochemical property. B . Thermochemistry of CH2CCI. There is no knowledge of the magnitude or temperature dependence of k2.Therefore, no thermochemical properties of CH2CCl can be obtained from the results of the current investigation. However, since k , has no activation energy, it is reasonable to presume that the reaction is exothermic, Le., AH'(298) C 0. This information establishes a lower limit of 60 (f0.5) kcal mol-, for the heat of formation of the 1-chlorovinyl radical.
Discussion A. Rate Constants of Reaction 1 . As shown in Figure 1, the values of k l obtained in the current investigation lie systematically above those reported by Krueger and Weitzl' by a factor of about 1.3. Both sets of rate constants display the same mild negative Arrhenius temperature dependence. The possible cause of this systematic difference is not apparent. Agreement is still considered good since constant uncertainty limits for the two sets of rate constants generally overlap. There is poor agreement between the current k , values and those that can be derived from the use of k-' and K I values reported by Parmar and Benson.I2 The error bars displayed in Figure 1 reflect the combined error limits for k-, and K l reported by these authors.I2 In a prior investigation, using the same very low pressure reactor, Dobis and Benson28investigated the equilibrium CI CH4 = HCI + CH3. In that investigation, values of the C1 + CH4 rate constant were measured and were shown to be in excellent agreement with prior determinations. However, the equilibrium constants also obtained in the same study, in conjunction with the measurements of rate constants for the C1 + CH4 reaction, resulted in the determination of C H 3 HCI rate constants that were not in good agreement with our direct determinations.' Largely from these observations we have concluded that the current very low molecule rate pressure reactor experiments yield C1 atom constants that are as accurate as indicated, but the more complex experiments required to obtain equilibrium constants yield results that are more uncertain than stated. B. Rate Constant ofReaction 2. Chlorine substitution at the radical site in the vinyl radical reduced the reactivity with HCI. Changes in both the preexponential factor and the activation energy reflect this behavior. Similar behavior has been observed in R CI2 and R H I reaction when R = CH3-xC1x.28*29It appears that the stabilization of polarized R(6'). .H. .X(6-) transition states is reduced when electron-withdrawing chlorine atoms are attached to the radical site in R.3' C. Thermochemistry of the C2H3 Radical. The Second Law (67.1 (f0.6) kcal mol-') and Third Law (66.9 (h0.3) kcal mol-') determinations of the C2H3 heat of formation reported here provide
+
+
+
+
+
(26) Curtiss, L. A.; Pople, J . A. J . Chem. Phys. 1988, 88, 7405. (27) Harding, L. B. Private communication, 1988. (28) Dobis, 0.;Benson, S. W. Int. J . Chem. Kinef. 1987, 19, 691. (29) Timonen, R. S.; Russell, J. J.; Gutman, D. Znr. J . Chem. Kinet. 1986, 18, 1193. (30) Seetula, J . A,; Russell, J . J . ; Gutman, D. Unpublished results. (31) Timonen, R . S.; Russell, J . J.; Sarzynski, D.; Gutman, D. J . Phys. Chem. 1987, 91, 1873.
5188
J . Phys. Chem. 1989, 93, 5188-5195
essentially identical results because the former method also yields an experimentally determined C2H3 entropy (55.9 ( f 2 . 6 ) cal mol-' K - ' ) which is virtually the same as the theoretical value (55.5 (f0.5) cal mol-' K-I). Since the enthalpy of formation and entropy determinations are coupled, high accuracy in the determination of one of these thermochemical quantities is a strong indication that the other is also accurate. In our prior study of the CH, + HCI reaction, a similar kinetic analysis, one which coupled our measured CH3 + HCI rate constants with those for the CI CHI reaction reported by Dobis and Benson and others,28 we also obtained exactly the correct entropy for C H 3 and a heat of formation which was within f 0 . 2 kcal mol-' of the most accurate prior determinations.' Recent studies using neutral reactants and products have obtained very similar values for the heat of formation of the C2H3 radical. Parmar and Bensont2obtained a C2H3 heat of formation within 1 kcal mol-' of our value, in spite of the fact that their values for k l are lower than ours by a factor of 4. Wodtke and Lee have observed the velocity distribution of DF(v=4) from the F + C2D4 D F C2D3reaction.32 Since this information is a very sensitive function of the energetics of the reaction, it was possible to use the results of their experiments to determine a C2H3 heat of formation. They report a value of 65.6 (f0.5) kcal mol-'.
+
-
+
(32) Wcdtke, A. M.; Lee, Y . T. In Advances in Gas Phase Photochemistry and Kinetics, Molecular Photodissociation Dynamics; Ashford, M . N. R., Baggot, J. E., Eds.; Royal Society of Chemistry: London, 1987; p 31.
Recent studies that have used ionization threshold methods continue to yield significantly higher values for the C2H3 heat of formation. For example, Berkowitz et all3 suggest a value between 69 and 72 kcal mol-' while Shiromaru et ai. indicate a rather high upper limit of 78.6 kcal mol-'.'4 (All values are for 298 K.) In the former investigation, data interpretation was complicated by the fact that a significant structural change occurs between the vinyl radical and cation, and, in the latter study, there was interference from background signals and the ionization onset was difficult to establish because of the gradual increase in the ionization cross section above threshold for the process which was under investigation, C2H4 C2H3+ H + e-. D. C-H Bond Energies in C2H4 and C2H3CI. The thermochemistry presented above may also be used to determine the C-H bond enthalpy at 298 K in C2H4(106.5 (f0.3) kcal mol-') and a lower limit for the C-H bond enthalpy in C2H3C1for the hydrogen bonded to the same carbon atom as CI (>103.6 (f0.2) kcal mol-]).
-
+
Acknowledgment. This research was supported by funds from the National Science Foundation (Grant CBT-8603737). W e thank Dr. Irene R. Slagle for her advice, help, and support and Dr. Wing Tsang and Dr. Larry B. Harding for useful information on the structure and properties of C2H3. J.A.S. also thanks the Natural Science Council of the Academy of Finland and the Alfred Kordelin Foundation for fellowships. Registry No. C2H,, 2669-89-8; CH2CCI, 57095-76-8
Energetic Constraints in the 218-nm Photolysis of Acetic Acid Sally S. Hunnicutt, Leslie D. Waits, and Joyce A. Guest* Department of Chemistry, University of Cincinnati, Cincinnati, Ohio 4.5221 (Received: January 12, 1989)
Acetic acid has been photolyzed at 218 nm under both room temperature and jet-cooled conditions. The nascent OH X2II(v"=O) photofragment has been probed by use of laser fluorescence excitation to determine its scalar and vector quantities. Jet-cooled acetic acid leads to OH photofragments with about 30% less rotational excitation than does the room temperature acetic acid. In both cases, only a few percent of the total available energy is partitioned into rotational energy of OH. Fragment OH is produced in u" = 0 only. The Fl(2113/2)and F2(2nl,2jspin states are statistically populated, and A-doublet levels are equally populated. The OH product radical was determined to have no rotational alignment. Doppler spectroscopy shows that a significant fraction, about 4976, of the available energy is imparted to C H 3 C 0 and O H fragment translation. The energy partitioning shows that the dominant path for OH production from 218-nm acetic acid photolysis is CH3COOH C H 3 C 0 + OH, with no subsequent decomposition of the acetyl fragment. An impulsive model predicts the mean fragment translational energies extracted from the experiments. Dissociation to three fragments, CH3 + CO + OH, cannot be ruled out as a minor channel.
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Introduction Electronic excitation with ultraviolet light can be used to deposit selected amounts of energy into gas-phase polyatomic molecules, often initiating photochemical reactions.',2 When the photochemistry leads to atomic or diatomic fragments, the fragment internal state distributions, velocities, and angular quantities can be examined by laser excitation The elucidation of triatomic and tetratomic photofragmentation dynamics has been dramatically advanced by the application of detailed fragment probing techniques. ( I ) Calvert, J . G.;Pitts, J . N. Photochemistry; Wiley: New York, 1966. ( 2 ) Turro, N. J. Modern Molecular Photochemistry; Benjamin/Cummings: Menlo Park, CA, 1978. (3) Molecular Photodissociation Dynamics; Ashfold, M . N. R., Baggott, J . E., Eds.; The Royal Society of Chemistry: London, 1987. (4) Reisler, H.; Wittig, C. Annu. Reu. Phys. Chem. 1986, 37, 307-349. ( 5 ) Leone, S . R. Annu. Reu. Phys. Chem. 1984, 35. 109-136.
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The issues to be addressed in the photolysis of polyatomic molecules like acetic acid are intrinsically different and more complicated. Questions of a more chemical nature can be asked when multiple reaction paths are possible. The same type of microscopic fragment details that help assign dynamics for small molecule fragmentation may reveal information about macroscopic photochemistry of larger polyatomics in favorable cases. Once a photoproduct is identified, we must next discover the reaction paths by which it is produced. Only then can the microscopic details be used to mode1 the dissociation dynamics and thus gain information about the potential surfaces of the reacting species. This paper will describe the photofragmentation of acetic acid via its '(n,r*) absorption excited at 218 nm. The absorption spectrum6 in this band is broad and structureless. Acetic acid (6) Robin, M. B. Higher Excited States of Polyatomic Molecules; Academic: New York, 1975; Vol. 2.
0 1989 American Chemical Society
