236
J. Phys. Chem. 1980, 84, 236-239
Reaction Mechanism of the Homogeneous Thermal Decomposltlon of Acetylene T. Tanzawa and W. C. Gardiner, Jr." Department of Chemistty, University of Texas, Austin, Texas 78712 (Received April 16, 1979; Revised Manuscript Received September 13, 1979) Publication costs assisted by the Robert A. Welch Foundation
A modeling study is reported in which experiments on the rate of and product distribution from CzH2pyrolysis from 625 to 3400 K are described with a single mechanism. The essential primary mechanism at low temperatures proves to consist of an H-atom, vinyl radical chain H + C2H2 C2H3,CzH3 + CzH2 C4H4+ H producing vinyl acetylene at early times. At high temperatures this is replaced by the ethynyl chain H + C2Hz CzH + Hz, CzH + CzHz C4Hz + H producing diacetylene. By considering a variety of studies simultaneously it was possible to assign rate constant expressions to the key elementary reactions. While all of the basic observations on the primary decomposition are accounted for by the final mechanism,uncertainties still remain in the rates of secondary reactions and in the magnitudes of the falloffcorrections required for the unimolecular reactions involved.
-
Introduction The heterogeneous and homogeneous thermal decomposition reactions of C2H2have been frequently studied to obtain basic gas kinetics information and to provide technical data on the process of carbon formation from gaseous hydrocarbons. From the product distribution in the early stages of reaction it is clear that one is dealing with different types of chain reactions at high and low temperatures. At low temperatures, vinyl acetylene, C4H4, is the main primary product, while at high temperatures iacetylene, C4H2, and H2 are the primary products. The m i c mechanism problem is to compile a plausible set of elementary reactions and rate constant expressions which accounts quantitatively for the product formation rates over the wide ranges of temperature and composition which have been studied experimenta1ly.l In connection with a recent shock tube study2 of the reaction we developed such a mechanism and showed that it was capable of describing well the results of all published shock tube experiments as well as our own. By extending the modeling study to lower temperatures it was found that flow experiments down to 1300 K3 could still be described very well, while bulb dissociation experiments (625-745 K)4gave dissociation rates substantially higher than computeda5 This mechanism had two obvious rate constant abnormalities. The abstraction reaction AHoo = 88 kJ H + C2H2 H2 + C2H
-
had to be assigned a non-Arrhenius temperature dependence in apparently flagrant disagreement with the endoergicity in order to permit computation of sufficient extent of reaction for the lower-temperature shock tube studies.6 Similarly, the decomposition reaction H0o = 519 kJ C4Hz M C4H + H M
+
+
+
had to be assigned a rate constant notably faster than the extrapolated directly measured value.' Since the modeling results were quite sensitive to these assignments, it was felt that either some experiments were being misinterpreted or important ingredients were missing from the mechanism. We pursued these questions further by expanding our modeling effort to include optimizing the fit to low-temperature experiments. It appears that at low temperatures the dominant fate of H-atom attack on C2Hzis not abstraction but addition 0022-3654/80/2084-0236$0 1.OO/O
-
-+
H
+ C2H2
+
-
A H o o = -163 kJ
C2H3
and that vinylacetylene is primarily produced by A H o o = 24 kJ CzH3 + CzHz C4H4 H +
+
Incorporating these additions and their further consequences into the mechanism, together with initiation reactions at suitable rates, permitted satisfactory accounting for all results considered, from 625 to 3400 K. There remain unresolved questions about the role of falloff in the unimolecular reactions, and the conflict in the rate constant for C4H2 thermal decomposition. The final mechanism and rate constant set is given in Table I.&"
Met hod Among the numerous lower-temperaturestudies of C2H2 thermal decomposition we chose to model the flow reactor studies of Palmer and Dormish3 (1330-1530 K) and of Munson and Anderson12 (770-1120 K), and the bulb pyrolysis experiments of Silcocks4(625-745 K). It has been pointed out repeatedly1p3J3that an Arrhenius presentation of the apparent second-order decomposition rate constants from all studies indicates good coherence among all results, so that it is to be expected that, once the apparent rate constants of a few typical studies have been matched, all of the other ones will also be matched. The mathematical background of the modeling is described e1~ewhere.l~ It was necessary to make choices for the equations describing, as best approximation, the physical state of the reacting gas. For modeling Silcocks' bulk dissociation experiment^,^ an isothermal constantdensity reaction was assumed. Rate constant adjustments were investigated for a 2-h run at P = 1atm, T = 745 K. For modeling the Palmer and Dormish and Munson and Anderson flow-reactor isothermal reaction at constant inflow velocity was assumed. In course of reaction, the pressure and density then decreased somewhat while the flow velocity increased slightly. Rate constant adjustments were investigated for a 5-s dwell time at P = 1 atm, T = 973 K, and 20% C2H2in He as typical Munson and Anderson conditions, and for a 150-ms dwell time at P = 1 atm, T = 1433 K, and 1% CzH2 in He as typical Palmer and Dormish conditions. For modeling Skinner and Sokoloski's single pulse shock tube experiments, constant density reaction was assumed. Rate constant adjustments were investigated for a 2-ms dwell 0 1980 American Chemical Society
The Journal of Physical Chemistty, Vol. 84, No. 3, 1980 237
Homogeneous Thermal Decomposition of Acetylene
TABLE I: Reaction Mechanism and Rate Constants' reaction kJ 1. C,H, + M 4 C,H + H + M 5 19 2. CzI-ILt C,H, 4 C,H, + H 192 3. H + C,H, C,H, -163 24 4. C,W, + C,H, C,H, t H 5. C,H, t H + M 4 C,H, t M - 330 -88 6. C,H + H, -, C,H, + H 7. C,H + C,H, C,H, t H -75 8. C,H t C,H, C,H, t C,H, -190 9. C,H, + M C,H, + H + M 250 10. C4Hz + M C,H t H + M 519 11. C,H + C,H, C,H, + H - 7!5 12. C,H + C,H, C6H, + H - 7!5 C,H, t H - 80 13. C6H + C,H, 14. C,H + C,H, C,H, t H - 80 15. C4H2t C,H C,H, t H - 80 51!3 16. C6H, + M C,H + H t M 17. C8Hz t M C,H + H + M 51!3 432 18. H, + M -t H t H + M
-
+
-+
-
--f
+
-+ -+ -f
-+
--
-+
-
log A 16.62 12.30 12.74 13.20 15.00 13.53 13.60 13.60 16.00 17.54 13.60 13.60 12.00 12.00 12.00 16.70 16.70 12.35 t 0.5 log
EA,
notes
kJ
10 105 0 0
0 0 250 335 0 0 0 0 0 335 335 387
T
-
ref 8 this work ref 9 this work ref 5 this work and ref 10 ref 2 and 10 this work and ref 1 0 ref 2 ref 2b ref 2 ref 2 ref 2 ref 2 ref 2 ref 2 , est ref 2, est ref 2 and 11
448 192
a Rate constants have mol/cm3 concentration units. The rate constant of reaction 1 0 was obtained from analysis of laser-schlieren data. Unreasonable recombination rate constants are implied if k - ) , [or the analogues of it assigned to reactions 1 6 and 1 7 1 is extrapolated to lower temperatures. I t was found that this has no effect whatever upon any of the loiver-temperature modeling reported here.
l2
C2H2 decay
IO
\f
8
In
'-
z
6
L I
Y, PD -J'--
I
5
\
/
MA
O\
l
r 4 0
m 0
1
2
0
5
4
6
8
10
12
14
16
7 IO'KIT
9
Flgure 2. Comparison with single pulse shock tube data of ref 6. --, experimental; ___, -, computations as indicated. _ _ _ I
IO'KIT
Figure 1. Comparison of computed and experimental rate constants for C2H2decay and product appearance. --and symbolis, experimental results; computed results, Table I mechanism; cornPutd results, reactions 3 and 4 omitted. BK, C2H2decay rate, ref 16; GKMN, C4H2appearance rate, ref 15; AG, ref 17; SS,ref 6; PD, ref 3; TM, ref 18; MA, ref 12; FK, ref 19; CMN, ref 20; S,ref 4. The computed lines for PD and GKMN merge.
-----
time at ID6 = 5 atm, T = 1430 K, and 6% C2Hzin Ar. The other shock tube experiments modeled were those previously considered,2 using laminar boundary layer flow models for incident wave experiments and constant density reaction for reflected wave experiments. The primary comparisons betweein computation and experiment were for the product yiebd6J5and the overall second-order sate constant for CzHz disappearance or C4H2I5or C4H6appearance. For each investigation the extraction of this overall rate constant from computed profiles was done in a manner as close as possible to that employed in analyzing the individual experiments. Adjustments were made to the rate constant expressions for many reactions as the investigation proceeded. At the end, however, the only rate constants which were changed from previously proposed values were those for reactions 2 and
6. A rate constant expression was assigned to reaction 8 by analogy to the experimental result for reaction 7. A rate constant expression for reaction 4 was derived by fitting the k4 values optimdy matching the Silcocks4and Munrion and Anderson12typical results (see above) to an Arrhenius expression. Since previous experiments were far less sensitive to the rate constant expressions for reactions 2 and 6 than those considered here, the final rate constant set amounts to new derived results for kz,k4, and k6, assignment of the rate constant measured for k, also to ks, and acceptance of literature rate constant expressions for all other reactions.
Results The final set of rate constant expressions is given in Table I. The computed overall rate constants are compared to experimental ones in Figure 1.'6'20 A comparison to the product formation rates of Skinner and Sokoloski is shown in Figure 2. Comparison with laser-schlieren observations2 has to be made somewhat indirectly due to the way that these reflect reaction rates. In Figure 3 are shown comparisons of the maximum deflections and times to maximum deflection computed for several C2H2-Ar
238
The Journal of Physlcal Chemistry, Vol. 84, No. 3, 1980 -23
,
I
-
Tanzawa and Gardiner
10
II)
-
9'1
ra E
i6 4.0
3.0 104K/T
Flgure 3. Comparison with experlmental laser-schlieren data. The (open symbols) and D, (filled symbols) are explalned In ref 2. -, Table I mechanism; ----,ref 2 mechanism.
T,
TABLE 11: Sensitivity Spectruma reaction GKMNb PDC 1 2 3 4 5 6
7
8 9
413 313 313 313 73/76 313 31-1 13/17
SHTO' 1 IO
1/4
110 1614 53/23 01- 1 -11-6 l/O 01- 3
1 /4
1 lo 1l o 56/52 117
SLTe
SSf
68/81 68/81
7/36
68/91
For the formation rate of CzH3 in reaction 3 we accepted uncritically the flash photolysis result of Payne and Stief? which is for a unimolecular dissociation reaction at its high-pressure limit. [For references to other measurements of k3 and discussion, see ref 9.1 Clearly this cannot be a correct description of H + CzH2 CzH3 over the entire temperature range considered, as falloff effecta must enter. We assumed that these would be small enough to neglect for that part of the temperature range where the CzH3 chain is important, i.e., up to the Munson and Anderson flow tube experiments. The rate constant expression for the main CzH3reaction in this temperature range, reaction 4, was then adjusted to fit exactly the midrange experiments of Silcocks and Munson and Anderson. While the Arrhenius parameters for k p are reasonable as they stand, they would be affected by changes in the expressions for k3, k6, and k7. Reaction 5 was originally introduced by us in an attempt to account for C4H4production in Silcocks' experiments without the CzH3 chainB6The value given is an estimate only; it does not have a large effect upon any of the computed results. Reaction 6 plays its role in the reverse direction as a chain propagation step at high temperatures. We chose to deal with it in the exothermic direction as CzH H2 in order to emphasize its similarity in nature and rate to reactions 7 and 8; the value of k6 obtained by fitting the laser-schlieren data is seen to confirm that C2H is indeed extremely reactive. While none of the experiments considered here appear to be sensitive to either k7 or ks,the flow reactor experiments of Lange and Wagner clearly require a very large value for the former. We set ks = k7 by analogy. None of the experiments considered was found to have appreciable sensitivity to kg. The value of klo is important for the high temperature conditions, and the laser schlieren data in particular could not be matched with klo values much different from those given by the Table I expression. This is, however, in conflict with a recent direct mea~urement.~ We have no explanation for the discrepancy. The remaining reactions (11-18) are concerned with the observed rates of formation of polyacetylenes16J6at high temperature, and are discussed in ref 2. They are otherwise not required in the mechanism. One reaction omitted from the mechanism is the disproportionationl CzH2 + CzH2 CzH3 + CZH
34/11 1113 54/37 1 /4 -91-9
'
a Entries are 100 X pS as defined in ref 21. The first entry is obtained by multi lying a rate constant by 5, the For C,H,:Ar = 10:90, T,= second by dividing by 5. 2000 K, P,= 5 torr, ref 15; pS for maximum C,H, appearance rate. For C,H,:He= 1:99, T = 1433 K , P = 1 atm, ref 3; pS for average C,H, decay in 150 ms. For C,H,: Ar = 10:90, T = 745 K,P = 1 atm, ref 4 ; pS from average C2H, decay. e Same as d but T = 625 K. f For C,H,:Ar= 6:94, T,= 1430 K, P,= 5 atm, ref 6; pS from C2H2decay.
mixtures. The final sensitivity spectraz1pnfor three typical experiments, at low, mid-range, and high temperatures, are given in Table 11.
Discussion One sees in Figures 1-3 that the match between computation and experiment is good over the entire temperature range studied experimentally. In Table I1 and in Table I1 of ref 2 one sees that the important reactions shift entirely from experiment to experiment, such that in order to match the whole group of results a nearly unique set of rate constant expressions is required. The experimental basis for the set is essentially as follows. The unimolecular dissociation reaction (1) is most uniquely defined by the ARAS experiments of Frank and Just: and their rate constant expression is used. While at higher temperatures than studied by them their expression might be somewhat too fast due to collision efficiency effects, it would be difficult to model their H-atom profiles with any substantially different klexpression. The bimolecular initiation reaction (2) is required to obtain sufficient initiation at lower temperatures, where reaction 1is too slow to compete with chain termination. Its rate constant is not well-defined, since the initiation rate is always coupled to chain propagation rates in defining observed reaction rates. Changes in other rate constant expressions would therefore imply changes in k2.
+
-
As we discussed before,2 taking this reaction as an alternative to reaction 2 could not be made to give agreement with the mass-spectrometric results.16J6We likewise omit as an initiation pathway. consideration of triplet C2Hz23*24 We note finally that our mechanism contains many reactions [(1),(3), (5), (91, (lo), (16), (171, (IS)] which are unimolecular decompositions. For the modeling purposes these were written in simple mass-action forms and given Arrhenius rate constant expressions. In reality, however, all of them except (18) may be expected to have falloff and collision efficiency effects over the range of conditions we considered, and so the mass-action laws and Arrhenius expressions cannot be associated with theoretical limiting forms for elementary reactions. Conclusions The mechanism and rate constant expressions presented here provide satisfactory descriptions of the experimental data on the rate of the homogeneous thermal decomposition reaction of CzHzand of the appearance rates of the
J, Phys. Chem. 1980, 84, 239-246
initial productlq. There remain many uncertaintiles as to the identity and rates of the secondary reactions leading to polymer formation at low temperatures and carbon formation at high temperatures. 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. We thank H. B. Palmer and J. Troe for stimulating discussions, References and Notes (1) M. H. Back, Can. J. Chem., 48, 2199 (1971). (2) T. Tanzawa and W. C. Gardlner, Jr., Seventeenth International Symiposlum on Combustion, Leeds, 1978,‘The Combustion Institute, 1979. (3) H. B Palmer and F. L. Dormlsh, J . Phys. Chem., 68, 1553 (1964). (4) C.G. Silcocks, Proc. R . Soc.London, Ser. A , 242, 411 (1957). (5) T. Tanzawa and W. C. Gardiner, Jr., Dlscusslon contrlbutiori following ref 2.
139
(6) G. B. Skinner and E. M. sdtoloski, J. phys. Chem., 64, 1952 (1960).
(7) P. Frank and Th. Just, Combust. Flame, in press. (8) Th. Just, private communlcation. (9) W. A. Payne and L. J. Stief, J . Chem. Phys., 64, 1150 (1976). (10) W. Lange and H. G. Wagner, Ber. Bunsenges. Phys. Chem., 79,
165 (1975). (11) A. L. Myerson and W. S. Watt, J . Chem. Phys., 48, 425 (1968). (12) M. S. B. Munson and R. C. Anderson, Carbon, 1, 51 (1963). (13) C. F. Cullis and N. H. Franklin, Roc. R . Soc. London, Ser. A , 2180, 139 (1964). (14) W. C. Gardher,Jr., B. F. Waker, and C. 8. W a k e M in “Shock Waves in Chemistry”, A. Lifshitz, Ed., Marcel Dekker, New York, 19BO. (15) I. D. Gay, G. B. Kistiakowsky,J. V. Michael, and H. Niki, J . Chem. Phys., 43, 1720 (1965). (16) J. N. Bradey and G. B. Klstlakowsky, J. Chem. Fbys., 35,264(1961). (17) C. F. Aten and E. F. Greene, Combust. Flame, 5, 55 (1961). (18) G. D. Towell and J. J. Martln, AIChE J., 7, 693 (1961). (19) D. A. Frank-Kamenetzky, Acta fhyslcochim. U.S.S.R., 18, 1148 (1943). (20) C. F. Cullis, 0. J. Minkoff, and M. A. Nettleton, Trans. Faraday S c c , 58, 1117 (1962). (21)W. C. Gardlner, Jr., J. Phys. Chem., 81, 2367 (1977). (22) W. C. Gardiner, Jr., J . Phys. Chem., 83, 37 (1979). (23) C. S. Burton and H. Hunziker, J . Chem. fhys., 57, 339 (1972). (24) I?. W. Wetmore and H. F. Schaefer, 111, J. Chem. phys., 6S, 1648 (1978).
A FT TR Spectroscopic Study of the Ozone-Ethene Reaction Mechanism in 0,-Rich Mixtures Fu Su, Jack G. CaIvert,* and John H. Shaw Departments of Chemistry and Physics, The Ohio State Universw, Columbus, Ohio 43210 (Received August 9, 1979) PubllcatYon costs asslsted by the Environmental Protection Agency
Fourier transform infrared spectroscopy has been employed to study the kinetics and products of the reaction between ozone and ethene (C2H4, cis-CDHCDH, trans-CDHCDH, and C2D4) in the ppm range in gaseous 02--N2 mixtures (18-26 “C, 700 torr). In addition to the expected reaction products, CO, C02,CH20,and HCOt2H, major products of the reaction were indentified as formic acid anhydride and as yet unidentified precursor to (HC0)20. The kinetics of formation of these products were studied in dilute mixtures of Oa and ethene (in 02--N2at 700 torr) with small amounts of CH20,CH3CH0, CO, or SO2 added. These results show tlhat about 38% of the CH202species formed in the Osethene reaction may take part in bimolecular reactions with molecular species present in the ppm range, while about 62% fragment or rearrange in unimolecular steps. Reported here are the second-order rate constants for the 03--ethene primary reaction and the relative rate constants for the first-order fragmentation steps of CH202and the second-order chemical reactions of CH:z02 with CH20,C2H4,CO, and SO2.
Introduction The ozone-alkene gas-phase reactions are of special interest in atmospheric chemistry.l The nature of the products, the mechanism, and the rate constants for the gas-phase reactions have been studied in the extensive, pioneering efforts of CvetanoviE and co-workers2and in more recent years by many other investigators from several laboratocies. The earlier work was interpreted in terms of the Criegee mechanism3 in which ithe original molozonide fragmented to form a carbonyl compound and a reactive species which has become known as the Criegee intermediate. 03
+
- -- \ - 7‘T o *
RzC=CR2
At
Bt
-
P-O-O
RzC
R2C-CR2
R2C=O
CRz ( A t )
(Et)
+ R2COO-
(1)
(2)
(3)
0022-3654,/80/2084-0239$01 .OO/O
O’Neal and Blumstein4 suggested that the diradical intermediate product B+ of (2) could rearrange by internal H-atom transfer together with some subsequent fragmentation steps to produce the excited molecular produicta and free-radicals observed experimentally.6 It was proposed from thermodynamic considerations and unimolecular reaction rate theory that the significant formation of the Criegee intermediate may be unimportant except for the very simplest alkene, C2H4, where fragmentation of the very energy-rich ozonide in (2) and (3) is most likely to occur before collisional equilibration of the initial products. However, Niki et al? showed clearly that a large fraction of the gas-phase ozonolysis of cis-2-butenereaction occurs via the Criegee mechanism as well. The identification of the intermediate diolrirane (CIH2(O),) by Lovas and Suenram’* and Martinez et al.% in the low temperature reaction of ozone with ethene provided convincing evidence for the occurrence of the Criegee mechanism for at least a major fraction of this reaction. Presumably the dioxirane formed from the initial Criegee 0 1980 American Chemlcal Society
