J . Phys. Chem. 1990, 94, 3277-3283
3277
Kinetics and Thermochemistry of the Equilibrium CCI, -t 0,
CCI,O,
J. J. Russell, J. A. Seetula, D. Cutman,* Department of Chemistry, Catholic University of America, Washington, D.C. 20064
F. Danis, F. Caralp, P. D. Lightfoot, R. Lesclaux,* Laboratoire de Photophysique et Photochimie MolBculaire. UniversitP de Bordeaux I , 33405 Talence Cedex, France
C. F. Melius, Combustion Research Facility, Sandia National Laboratory, Livermore. California 94550
and S. M. Senkan Department of Chemical Engineering, Illinois Institute of Technology, Chicago, Illinois 60616 (Received: September 29, 1989; I n Final Form: December 16, 1989) The equilibrium CC13 + O2 z CC1302has been studied experimentally between 351 and 461 K, the temperature range in which it could be clearly observed and characterized. Equilibrium constants were measured as a function of temperature at two laboratories using different techniques and different experimental conditions. In both investigations,CC13was produced by pulsed UV photolysis, and CCI3relaxation to equilibrium, in the presence of 02,was monitored in time-resolved experiments, using either photoionization mass spectrometry or UV absorption spectrometry. The experimental results were combined with calculated entropies of CCI, (74 f 1 cal mol-' K-I) and CCI3O2(83 f 3 cal mo1-l K-I), obtained from theoretical determinations of the structures and internal motions of these radicals, to obtain the enthalpy change of this equilibrium (AH0298= -19.9 (*I .O)kcal mol-'). The results indicate a considerable weakening of the CH3-02 bond caused by chlorine substitution on the CH3 group, from 32.4 kcal mol-' [DHo2,,(CH3-O2)] to 19.9 kcal mol-' [DH02g8(CC13-02)].This effect of chlorine substitution on bond strengths is compared with that found for other compounds of the type CH3-X. The implication of weaker R-02 bond strengths caused by chlorine substitution in the oxidation chemistry of chlorinated fuels is briefly discussed. A heat of formation of CC1302was obtained from the experimental results (AHf0298 = -3.3 ( i 2 . 2 ) kcal mol-').
Introduction Reactions of chlorine-containing free radicals are receiving renewed attention because of their involvement in a wide range of important processes of current interest including the chemistry of stratospheric ozone production and destr~ction,'-~ the incineration of municipal wastes with chlorine-containing polymeric material~,4*~ and the production of commercially useful C2 hydrocarbons from methane by controlled oxy~hlorination.~-~ In all these processes, oxidation of polyatomic free radicals (R) by molecular oxygen is a key elementary step. The kinetics of R + O2 reactions involving chlorine-bearing free radicals have not been widely studied, particularly at elevated temperatures. However, knowledge of the behavior of these elementary reactions, including the anticipated change in reaction mechanism at elevated temperatures, is essential for developing useful quantitative kinetic models of the important processes described above. The addition mechanism R +0 2 RO2 (1) which is the dominant reaction path at low temperatures has been characterized for some chlorine-containing radicals at ambient temperature. They include the CCI3 + 0: and CFCl2 + 02 reactions. None of these reactions have been isolated (or nearly isolated) for direct investigation at elevated temperatures. However, there
-
( I ) Atmospheric Ozone 1985. WMO Global Ozone Research and Monitoring Project, Report No. 16; World Meterological Organization: Geneva, 1986. (2) Singh, H . B.; Kasting, J. F. J . A m " . Chem. 1988, 7. 261. (3) Rodriguez, J. M.; KO,M. K. W.; Sze, N. D. Geophys. Res. Left. 1986, 13. 1292. (4) Senkan, S.M. In Deroxicafion of Hazardous Waste; Exner. J. H., Ed.; Ann Arbor Science: Ann Arbor. MI. 1982: Chanter 3. (5) Miller, D. L.; Senser, D. W.; Cundy, V. i.; Matula, R. A. Hazard. Wasfe 1984. 1 . I . (6) Senkan, S. M . Chem. Eng. f r o g . 1987, 12, 58. ( 7 ) Granada. A.: Karra, S.B.; Senkan, S.M. Ind. Eng. Chem. Res. 1988, 27. 1163. ( 8 ) Ryan, K. R.; Plumb, I . C. In!. J . Chem. Kinef. 1984, 16, 591. (9) Caralp, F.; Lesclaux, R . Chem. fhys. L e f f .1983, 102, 54.
0022-3654/90/2094-3277$02.50/0
are indications that substitution of chlorine for hydrogen in CH3 reduces the R-O2 bond energy significantly. If this is the case, then the anticipated mechanism change in R O2 reactions (caused by the thermal instability of R 0 2 with respect to dissociation to R + 0,) will occur at much lower temperatures for R + O2reactions involving chlorinated free radicals than is observed in the oxidation of alkyl radicals. In an investigation of chlorine atom sensitized oxidation of chloroform, Olbregts reports observations that indicate that the activation energy of kT2is 22 kcal mol-'.'o This would suggest a CC13-O2 bond energy (DHO,,,) CCI, + 0 2 (M) s CC1302 (M) (2)
+
+
+
of =25 kcal mol-' since reaction 2 is significantly in the unimolecular falloff region at the pressures used in Olbregts' investigation." Melius has calculated heats of formation of R and R 0 2 for R = CCI,, CCI2H, and CClH2 using the BAC-MP4 method.I2 From these results one obtains a CCI3-O2 bond energy of 17 kcal mo1-l." These values, while differing significantly, are both much lower than the R-02 bond energy of the parent alkylperoxy radical (CH3-02) which is 32.4 kcal m01-I.~~ We have begun a comprehensive investigation of reaction 2 in order to fully understand its kinetic behavior and to determine the thermochemistry of CCI3O2. The latter information, which includes determinations of the CCI3O2entropy and heat of formation, is needed to identify the plausible reactive routes of this peroxy radical that are likely to be important in atmospheric processes and under combustion conditions. We report here a cooperative study of reaction 2 between 351 and 461 K, the temperature range in which the equilibrium between reactants and products in reaction 2 could be clearly observed and characterized. Equilibrium constants were measured (IO) Olbregts, J . J . fhorochem. 1980, 14, 19. ( 1 I ) Caralp, F.; Danis, F.; Lesclaux, R. To be published.
( I 2) (a) Ho, P.; Coltrin, M. E.; Binkley, J. S.; Melius, C. F. J. fhys. Chem. 1985, 89, 4647. (b) Melius, C. F.; Binkley, J. S. Symp. (Inr.) Combusr., [Proc.] 1985, 20, 5 7 5 .
(13) Melius, C. F. Unpublished results. (14) Slagle, I . R.; Gutman, D. J . Am. Chem. SOC.1985, 107, 5342.
0 1990 American Chemical Society
3278 The Journal of Physical Chemistry, Vol. 94, No. 8, 1990
Russell et al.
TABLE I: Conditions and Results of Experiments To Measure the Equilibrium Constant for the Reaction CC13 Photoionization Mass Spectrometry To Monitor CCls
[O,] x 10-16, T, K
molecules cm-)
[MIX molecules cm-j
k r , s-I
m , , s-I
388 388 388 388 38gb 398 398‘ 408 408 408d 418 418 428 428 439
1.22 2.00 1.44 1.81 1.81 3.28 1.71 4.41 4.44 4.44 4.45 3.47 3.91 5.13 4.05
17.7 17.6 17.2 17.6 17.6 5.77* 17.6 5.76; 5.79 5.79 5.77; 11.9* 5.78* 5.76* 5.77;
2.2 2.9 2.7 6.6 6.6 9.9 0.7 9.1 5.7 5.7 5.9 9.5 3.8 5.5 10.8
210 360 316 322 343 323 347 390 386 369 507 530 392 41 I 369
m,. s-I
4.1 2.9 16.2 11.8 5.6 11.0 11.8 14.9 11.3 11.3 15.5 22.8 14.2 12.0 20.6
+ O2 2 CC1,02 Using AIB 2.55 4.77 5.32 4.64 4.45 5.71 2.12 3.77 3.67 3.41 2.23 I .67 1.10 I .37 0.7 1
K,, atm-’ 3452 4132 3868 3328 3955 2104 1863 1086 1144 I063 726 674 416 398 24 1
‘Asterisk indicates N 2 bath gas (He used otherwise). bExperimentusing half the laser intensity that was used in the previous one. CDatafrom this experiment shown in insert in Figure I . dExperiment with half the precursor concentration that was used in the previous one.
as a function of temperature at two laboratories using different techniques. The results obtained were combined with a calculated entropy change of reaction 2 (obtained from theoretical determinations of the structures and internal motions of CCI, and CC1302)to obtain the enthalpy change of reaction 2. The details and results of the two experimental investigations and the theoretical study of this equilibrium are reported here. Experimental Section
The two independent experimental determinations of the temperature dependence of the equilibrium constant of reaction 2 both used time-resolved techniques to monitor the relaxation of CCI, to equilibrium in the presence of 02. In one study CCI, was detected by photoionization mass spectrometry, and in the other investigation its temporal behavior was followed by UV absorption spectrometry. These two studies are discussed separately in this section. A. Study Using Photoionization Mass Spectrometric Detection oJCCI,. The experimental apparatus and technique used to obtain equilibrium constants in this portion of the study have been described previously in connection with prior investigations of other R + O2 .- R 0 2 equilibria (ones involving hydrocarbon free radical^).'^-'^ Hence only a summary is presented here. The reaction occurred in a heatable 1.05-cm-i.d. uncoated quartz tubular reactor coupled to a photoionization mass spectrometer. Gas flowing through the reactor (at 4m s-l) contained CC13Br (the CCI, source), O2in varying amounts, and the carrier gas, which was usually in great excess (He or N2). Pulsed (5 Hz) unfocused 248-nm radiation from a Lambda Physik EMG 201 MSC excimer laser was collimated and then directed along the axis of the tube. CCI, was produced uniformly along the reactor by the photolysis CCI,Br
248 nm
CCI,
+ Br
(3)
CC13Br concentrations were selected to yield low CCI3 concentrations (6 X IO” molecules cm-, (at room temperature). This suggests a large value for k9 (2 ( f l ) X cm3 molecule-’ s-l). A comparable value has been reported for a similar reaction, CI + CF2C102.23 The absorption cross section of CCI3O2was calibrated relative to that of CCI, by generating both radicals under nearly the same conditions, the only difference being the presence or absence of an excess of 02.The recorded spectra are shown in Figure 2. All (19) Clyne, M . A.; Walker, R. F. J . Chem. Soc., Faraday Trans. I 1973, 69, 1547. (There is a mistake in the rate constant expression of k6 in the abstract of this reference. Log A is -10.840, not -10.480.) (20) Sanhueza, E. J . Photochem. 1977, 7, 325. (21) Hauteclocque, S . J . Photochem. 1980, 14, 157. (22) Lesclaux, R.; Dognon, A. M.; Caralp, F. J . Photochem. Photobiol. A 1987, 4 / , 1 . (23) (a) Carr, R. W.; Peterson, D. G.; Smith, F. K. J . Phys. Chem. 1986, 90, 607. (b) Dagaut, P.; Wallington, T. J.; Kurylo, M. J . J . Photochem. Photobiol. A 1989, 48, 187.
Russell et al.
3280 The Journal of Physical Chemistry, Vol. 94, No, 8, 1990 3
0.14
Y
.u z d
m
'd
8
-m
d
tlms
tlms 10
0
20
'
0
1
0.14
io
F
20
I
t Ills i
10
"0
20
Figure 3. Dccay traces record& by UV absorption spectrometry at 220 nm; 760-Torr total pressure of N2; [Cl,] = 3 X 10l6molecules cm? [CHCI3] = 5 X IO" molecules cm4. (a) T = 371 K, [O,] = 0; (b) T = 371 K, [O,] = 1 .O X IO" molecules cm-); (c) T = 371 K, [O,] = 2.2 X I O l 5 molecules cm-3; (d) T = 461 K. [O,] = 1.36 X IOl7 molecules Solid lines are calculated.
kinetic determinations reported below were performed at 220 nm (for CCI,) and 260 nm (for CC1,02). These wavelengths provided the best compromise between selectivity and sensitivity. 2. Kinetic Analysis of Experimental Results. The absorption profiles were analyzed taking into account the reactions occurring during the experiment. Two reactions were treated as occurring instantly: the conversion of CI into CCI, (reaction 6) and the decomposition of CCI,O (reaction 8). (A more rigorous and lengthy analysis using the actual rates of these reactions yielded the same results, hence justifying these approximations.) In these experiments, the reaction of CCI, with CI2 CCI,
+ c12
-
CCl4 + CI
(10)
did not interfere significantly. The CCI4 produced is a potential contributor to the absorption measured at 220 nm. However, k , , is relatively low in the temperature range of this in~estigation?~ and it was shown in our prior investigation of the recombination react ion CCI,
+ CCI3 + ( M )
4
CzC16 + ( M )
(11)
that this potential interference is not important. The rate constant for the CCI, + 0, recombination reaction, k,, is required for the data analysis to obtain K , ( k 2 / k Z ) . We have investigated the kinetics of reaction 2 at lower temperatures (253-333 K) where the reverse reaction is not important and obtained the result (for 760 Torr) k2 = 1.44 X 1 0-l2( T / 2 9 8 p 6 cm3 molecule-1 s-,. (Details of this investigation, including the pressure and temperature dependences of k,. will be published separately.'!) This expression for k 2 was extrapolated through the temperature range of the current investigation (353-461 K ) to obtain the values needed in the data analysis. While this extrapolation introduces some extra uncertainty in k2 in the higher temperature range, the effect of this on the determination of K 2 was slight because of the low sensitivity of the determination of K 2 to this rate constant. (24) Timonen, R. S.;Russell, J. J.; Gutman, D.h i . J . Chem. Kinel. 1986, 18. 1193.
In these experiments, reaction 2 generally reached equilibrium rapidly. Slower subsequent reactions can appear to shift the equilibrium, since they provide paths to convert CC1302back into CCI,. They include the slow conversion of CCI,02 to CCl,O by reactions 7 or 12, followed by the subsequent rapid decomposition CCI, + CC1,02 2CCI30 (12)
-
of CC130 (reaction 8), which releases chlorine atoms that regenerate CCI, in reaction 6. These cycles were of negligible importance at the highest temperatures of this investigation where equilibrium was established very rapidly, but they had to be taken into account at the lowest temperatures where relaxation to equilibrium was slower (because of the need to use low O2concentrations). To take the above cycles into account in the data analysis, values of k , and k i 2 were also required. k7 at 760 Torr [ 1.6 (f0.6)X T/298)-3.0(*l.o)]was obtained from observations of CCl2O production (monitored by its absorption in the UV over a period of = I s) when O2was in large excess. Under these conditions, CCI, is rapidly converted into CCI3O2and CCI2O is produced by the chain oxidation of CHCI,, reactions 6 , 2 , 7, and 8. The chain length is long since there is no "efficient" termination reaction and reaction 7 is rate determining. The rate constant of reaction I2 could not be obtained by an independent experiment. The value chosen (l.Oz,, X 10-l2) was that which resulted in the best agreement between experiment and simulation at the lowest temperature. At higher temperatures, this reaction rate constant did not affect the simulated decays. 3. Determination of Equilibrium Constants. Values of K 2 ( k , / ! ~ were ~ ) obtained by fitting experimental concentration profiles to simulated ones using the kinetic mechanism and the rate constants discussed above. kz was the only adjustable parameter (since k , was obtained from the extrapolations of lower temperature experiments). Since equilibrium was generally reached rapidly in these experiments, the fitting procedure, while being very sensitive to K2, is relatively insensitive to the absolute values of k 2 and k-2. So in the process of obtaining the correct value of K,, an error in k-2 results that matches any error in the extrapolated value of k , . Therefore, individual values of k2 and
Kinetics and Thermochemistry of CCI,
+ O2 z CCI3O2
The Journal of Physical Chemistry, Vol. 94, No. 8, 1990 3281
TABLE 11: Conditions and Results of Experiments To Measure the Equilibrium Constant for the Reaction CCI3 + O2 ; iCClj02 Using UV Absorption Spectroscopy To Monitor CC13 [Cl,] x 10-16, [CHCIJ X IO-”, [O,] x 10-16, [CCI,] x 10-13, k-2 X K2 X IO-’, T. K molecules molecules cm-, molecules cm-) molecules cm-, 10-2, s-I atm-’ 35 I
360
37 1
b
42 1
46 I
3. I 3. I 3. I 3. I 3.1 3.1 3. I 3. I 3. I 3. I 3.1 2.6 2.6 2.6 2.6 2.6 2.7 2.8 2.8 2.7 2.7 2.6 2.8 2.6 2.7 2.4 2.6 2.6 2.6 2.6 2.5 2.4
4.1 4. I 4.1 4.2 4.2 4.9 5.1 4.7 4.8
5.0 5.0 4.1 4.1 4.1 4.1 4.1
1 .oo
5.2 5.3
1.05 0.18 0.37 0.75
5.1
1.10
5.2
1.90 2.60 0.58 0.95 1.36 1.90 1.90 2.04 2.80 4.70 6.50 10.10
5.1
5.3 4.3 4.2 4. I 4. I 4.2 4.3 4.3 4.1 4.1 4.1
k-2 obtaincd in this manner in the temperature range of this study are of uncertain accuracy and should be used with appropriate caution. In Figure 3a-c, absorption profiles at 220 nm recorded in three different experiments at 371 K are displayed. Only [02] was different in each experiment. In Figure 3a, the CC13 decay in the absence of O2 is shown. This second-order decay is due to CCI, recombination, reaction 1 I . Such a profile was recorded under each set of experimental conditions to establish the initial CCI, concentration. By contrast, in the presence of a large excess of 02,CCI, is instantly converted to CCI3O2(see Figure 3c). The residual absorption shown in Figure 3c is due to CC1,02. Due to the radical chain discussed above, the CCI3O2concentration does not change significantly during this observation time. The slight increase in absorbance at long times is due to the formation of CClzO produced by reaction 8. These observations of the essentially stoichiometric conversion of CCI, to CC1302establish the relative absorption coefficient of CCI, and CCI3O2 at this wavelength under the exact conditions of the experiment and hence provide the accurate calibrations needed to quantitatively relate the observcd temporal behavior of the absorbance with that of the chemical system under study. Finally, in Figure 3b, the intermediate case is shown. Enough O2is present to observe the establishment of the equilibrium in reaction 2 during a time period that is short compared to that of the CCI, CCI, recombination process. (The slower decay in absorbance following the establishment of equilibrium is due primarily to this recombination.) The initial concentration of CCI, varied by less than 5% from one experiment to the next. A typical dccay trace obtained with O2 present at the highest temperature of this study (461 K) is shown in Figure 3d. As shown by the calculated trace (full line), the establishment of equilibrium is almost instantaneous due to the large O2concentration that is required to obscrvc equilibrium at this temperature. The role of secondary and tcrtiary reactions is far less important in the equilibrium constant determination, since K2 is obtained from the total absorbance during the decay and from the decay rate (due mainly to CCI, + CCI,) which is essentially controlled by the CCI,
+
0.1 I 0.22 0.33 0.44 0.66 0.1 1 0.22 0.32 0.43 0.65 1.10 0.21 0.42 0.63
10.1 10.0 10.0 9.9 9.9 9.4 10.0 9.8 9.6 10.2 9.2 8.9 8.7 8.6 9.3 9. I 9.2 9. I 9.0 8.9 8.8
8.8 10.3 10.0 9.9 9.9 9.8 10.5 9.8 9.7 9.7 9.6
3.1 1 2.06 3.34 4.20 3.36 3.32 3.97 6.75 5.93 6.48 6.05 8.40 10.90 11.40 11.40 8.20 126 165 I42 I52 171 I42 672 587 648 65 1 864 668 660 829 68 1 679
5370 81 IO 5000 3980 4910 4480 3750 2200 2510 2290 2460 1540 1 I90 1 I40 1 I40 1580 57 44
51 48 42 51 7.1 8.1 7.4 7.3 5.5 7.1 7.2 5.8 7.0 7.0
concentration at equilibrium. This more favorable condition for determining K 2 was present at the three highest temperatures of this investigation. At each of the five temperatures used in this investigation, equilibrium constants were determined by using different O2 concentrations, typically covering over a factor of 10 in [O,]. As expected, the values of K2 did not depend on [O,]. The conditions of all the experiments and the results obtained are given in Table 11. Equilibrium constants are also plotted in Figure 1. C. Accuracy of Equilibrium Constants. As can be seen in Figure I , there is excellent agreement between the results of the two experimental investigations in the overlapping temperature range of the two studies. Agreement is perfect in the middle of this temperature range, and deviations between K2values, which are highest at the extremes of the mutually covered temperature range, are under 10%. Individual assessments of the accuracy of the determinations of K2 are presented here. The exact agreement between K2values obtained by using two very different experimental methods and very different experimental conditions is verification that the experimental techniques, including the assumptions and calibrations used to reduce and interpret the data, are at least as accurate as presumed. In the photoionization mass spectrometry study, the estimated ( I u) uncertainty in K2 varies from k25% in the middle of the temperature range to *SO% at the high- and low-temperature extremes. The high- and low-temperature limits of this investigation were determined by different factors. At higher temperatures than were used here (>440 K), not enough O2could be added to drive the reaction forward to an observable equilibrium. Below 380 K such low O2concentrations were required that the reaction was too slow to reach equilibrium in the observable reaction time (=20 ms). In the U V absorption experiments, the uncertainty in K2 is f25% at the three highest temperatures (371, 421, and 461 K) where equilibrium was established very rapidly and the determination of K2 was relatively insensitive to knowledge of the rate
3282 The Journal of Physical Chemistry, Vol. 94, No. 8, I990 constants of secondary and tertiary reactions. At the lower temperatures, where the approach to equilibrium was slower and K2 is more sensitive to the accuracy of this information, the uncertainty is greater, f50%. The exact agreement in K 2 values near 408 K, the midpoint of the overlapping temperature range of the two investigations, is taken as a further indication that the values of K 2 near this temperature are more accurately known ( f 2 0 % ) than the individual estimates of uncertainty indicate. Thermochemistry of Reaction 2 The enthalpy change of reaction 2 at 298 K was obtained by using both a Second and Third Law analysis of the data. The procedure used has been described before.I4-I6 Initially, A G O T values were obtained directly from the equilibrium constants. Each AGO7 was corrected slightly (
