Thermal decomposition of ethylene

COMBUSTION AND FLAME 39:241-253 (1980)

241

Thermal Decomposition of Ethylene T. TANZAWA and W. C. GARDINER, JR. Universi~_ of Texas, Austin, Texas 78712

Laser-schlieren profiles o f incident shock waves in 2.5, 5, and 10% C 2 H 4 in A r were recorded for the ranges o f shock front conditions 2 0 0 0 < T 2 < 2540°K, 1.8 x 10 - 6 < p < 5.4 × 1 0 - 6 mol/cm 3. Data analysis was a c c o m p l i s h e d by c o m p u t e r m o d e l i n g using a 14-reaction mechanism. Most, but not all, previous observations could be accounted for with the final rate constant set. F o r C 2 H 4 + M ~ C 2 H 2 + H 2 + M the expression log (k I/cm 3 mol - I s - 1) = 17.47 - 340 k J / R T , for C 2 H 4 + M ~ C 2 H 3 + H + M the expression log ( k 2 / c m 3 mol - 1 s - 1) = 17.49 - 4 0 0 kJ/RT, and for C 2 H 3 + H ~ C 2 H 2 + H 2 the rate constant k 6 = 1013 cm 3 mol - I s - 1 were obtained.

INTRODUCTION The mechanism of ethylene thermal decomposition is of importance in elucidating high temperature mechanisms of hydrocarbon pyrolysis and oxidation in general, because ethylene rises to high concentrations in the course of these reactions, Until recently, the study of ethylene pyrolysis was still at the stage where the choice between radical and molecular decomposition channels Call4 + M = C2H 2 + Ha + M

AH~ = 166 kJ

(1)

C2H4 + M = C2H 3 + H + M

AH~ = 435 kJ

(2)

was unclear. Previous shock tube experiments carried out using single-pulse shock tube (SPST), [1, 2, 3], TOF mass spectrometry [4], and ir emission techniques [5, 6] did provide important facts, SPST experiments for 1300 < T < 3000°K indicared that the final products included 1, 3-butadiene [1] and methane and ethane [2] in addition to the main products acetylene and hydrogen. Gay er al. [4] and Bauer and Jeffers [7] found that the ethylene decay rate is first order in the ethylene concentration and half order in inert gas concentration. Homer and Kistiakowsky [5] and Roth and Just [6] measured pyrolysis rates over the temperature range 1675 to 2210°K in agreement with those of Gay et al. by following C2H4 and

CzH z through their thermal ir emissions. A definite mechanism that could account for the observations had not been established by these investigations. Recently, Just et al. [8] (JRD) measured the hydrogen atom concentration in dissociating ethylene as a function of time using atomic resonance absorption spectrophotometry (ARAS) in very dilute mixtures, in which contributions from secondary reactions were strongly suppressed. They concluded that ethylene decomposes through two different channels simultaneously, that is, Reactions (1) and (2), and determined a rate constant expression for Reaction (2), namely, k 2 (JRD) = 1017.41 exp (-404kJ/RT) cm a tool- 1 s- 1 . The CzH 4 disappearance rate constants measured previously [6], being much faster than this expression, could then be assigned to reaction (1), that is, kl(RJ ) = 1017.32 exp ( - 3 3 ) kJ/RT) cm a mo1-1 s- x . Tanzawa [9] reported a rate constant expression for Reaction (1) obtained from laserschlieren data assuming that Reaction (1) is dominant over Reaction (2), and subsequent chain reactions, immediately after the arrival of shock front. The agreement between Ref. [6] and Ref. [9] rate constant expressions for Reaction (1) is excellent. A complete report on our laser-schlieren experiments is given here. It is also of interest to compare the whole available set of experimental

Copyright © 1980 by The Combustion Institute P u b l i s h e d b y Elsevier N o r t h H o l l a n d , Inc. 52 Vanderbilt Avenue, New York, NY 10017

0010-2180/80/090241+13501.75

242 observations on ethylene decomposition with predictions from a comprehensive reaction mechanism. We report a reaction mechanism for ethylene pyrolysis that includes a large set of secondary reactions and is able to account quantitatively for many previous experimental conditions-but not for all. EXPERIMENTAL The shock tube apparatus and procedures for data acquisition and analysis have been described in detail previously [10]. Experiments were done over the following ranges of conditions: 2000 < T 2 < 2540°K, 1.8 × 10- 6 < P2 < 5.4 × 10--6 mol c m - a ; test gas compositions 2.5, 5, and 10% C2H4 diluted in Ar. The C2H4, Phillips research grade, specified to be 99.96% pure, was used without further purification. The Ar was Matheson ultra high purity grade, specified to be 99.99% pure. Mixtures were prepared manometricaUy in 41 dm 3 stainless steel bulbs, Data interpretation was performed with the aid of computer modeling. The flow models used were variable-area reactive flow [11 ] for incident shock waves and constant density [12] for reflected shock waves. Thermochemical properties were computed from polynomial fits to JANAF [13] or other published data [14] except for C2H4 [10]. The reaction mechanism adopted for describing the chemistry, based on the cited and other [1519] previous studies, is shown in Table 1. The computation of laser beam deflections was carried out during most of this research by a procedure that we subsequently recognized as incorrect in principle and possibly significantly inaccurate. It was based on using the computed density gradient and an incorrectly averaged local specific refractivity to obtain computed deflections, Specific refractivities were available only for several of the stable molecules. These specific refractivity values were C2H4, 0.304 [20] ; C2H2, 0.531 [20], Ar, 0.159 [20], Ar, 0.159 [21] ; H, 2.59 [22] ; and H 2, 1.54 [21] cma/g. Most of these values refer to 589 nm light. The total local specific refractivity of the reacting m i x t u r e was computed from the incorrect formula K t o t a 1 = ~,Kici/~,ci, where Ki is the specific refractivity of

T. TANZAWA and W. C. GARDINER, JR. component i and ci is its local concentration; the summation was done for the species whose K i values are known on the assumption that neglecting the other species would lead to negligible error in K t o t a 1. After it was recognized that the averaging formula was incorrect as well as approximate, it was replaced by a complete calculation of the local refractivity from molar refractivities of each species at the laser wavelength 633 nm. Where experlmental values were available at several wavelengths, molar refractivities at 633 nm were calculated through a nonlinear least squares analysis to fit the Cauchy dispersion equation. Where experimental refractivity values were unavailable, a summation of atomic refractivities was performed and the molar refractivity at 633 nm then derived in the same manner. (A complete report on this procedure is in preparation.) A complete rerun of the modeling calculations was then done with the laser deflection being computed from the local refractive index gradient. It turned out that the differences between the results of the two series of computations were so slight that no modifications of any of the conclusions were indicated. ANALYSIS AND RESULTS Examples of laser-schlieren oscilloscope records are shown in Fig. 1. The range of conditions for which such records could be utilized was determined by the following considerations. We noticed small, but measurable deflections in pure argon shock waves. These are due to wall boundary layer effects. Such deflections, when extrapolated to the arrival of the shock front, were found to be almost independent of shock temperature for the present experimental conditions. The location of the shock front on laser-schlieren signals recorded using a quadrant photodiode has been described elsewhere [18]. The gradient calculated from the deflection due to the wall boundary layer effect was subtracted from the total computed initial density gradient due to chemical reaction. At the lowest end of the temperature range, the ratio (dp/dx)t=o,eh~ie~/ (dp/dx)t=O,boundarylayer became roughly ~1, SO that no experiments were performed at still lower temperatures. Above 2600°K, the profiles did not

THERMAL DECOMPOSITION OF ETHYLENE

243 TABLE 1

Assumed Mechanism and Final Rate Constant Expressions*. Reaction No. 1 2 3 4 5 6 7 8 9 10 11 12 13 14

Reaction C2H 4 + M C2H4 + M C2H 4 + C 2 H 4 C2H 4 + H C2H 2 + C 2 H 4 C2H 3 + H C2H3 + M C2H5 + M C2H 2 + M C2H2+C2H2 C4Ha + M C2H 2 + H C2H2 + CzH C4H 2 + M

= = = = = = = = = = = = = =

C2H 2 + H 2 + M C2Ha+H+M C2H 5 + C 2 H 3 C2H 3 + H 2 C4H 6 + H C2H 2 + H 2 C2H2+H+M C2H4+H+M C2H+H+M C4Ha+H C4H2+ H +M C2H + H 2 C4H2 + H C4H+H+M

A H ~/kJ

Log A

n

Ea/kJmol--1

Footnote

165 435 275 3 35 -269 163 160 520 190 253 88 -77 395

17.47 17.49 14.26 14.85 12.00 13.00 14.90 15.31 16.62 13.00 16.00 .89 13.60 17.54

0 0 0 0 0 0 0 0 0 0 0 3.2 0 0

340 400 270 60 30 0 130 130 450 190 250 2 0 340

a b c d e f g h i / k l m n

Rate constants are in the form o f k = A T n exp (-Ea/RT), withA in cm 3 mol--1 s--1 K--n units. This work, optimization of fit to initial deflections, see text. Ref. [8], using ARAS, multiplied by 1.2 to improve fit to present laser-schlieren data. Ref. [15], using conventional hot bulb techniques at around 600°K coupled with gas chromatography. Ref. [6]. Ref. [16], multiplied by 2.0 in course of optimizing computed deflection profiles at higher temperatures. For our estimate, see text. Ref. [16], RRK calculation. Ref. [10], RRK calculation. Ref. [17], using ARAS. Ref. [18], assumed. Ref. [18], estimate. l Ref. [18]. m Ref. [19], multiplied by 1.3, microwave flow discharge with mass spectrometry at 320°K. n Ref. [18].

* a b c d e f g h i -/

show a density gradient p l a t e a u , b u t instead showed nearly e x p o n e n t i a l decay. This m e a n t that we had to confine our e x p e r i m e n t s to the temperature range 2 0 0 0 to 2 5 4 0 ° K . F o r the analysis o f the initial laser b e a m deflection, we first assume that R e a c t i o n (1) is d o m i n a n t over the o t h e r possible initial Reactions (2) and (3), for our e x p e r i m e n t a l conditions, that it is kin-

negligible reverse rate, thermal equilibrium e x c e p t for dissociation, ideal shock flow, and very small e x t e n t o f reaction. The bimolecular rate constant expression t h e y derive is

etically second order [ 6 ] , and that the radical chain m e c h a n i s m does n o t influence the initial laser b e a m deflection. These assumptions will be justified later. Breshears et al. [23] derived a relation b e t w e e n the dissociation reaction rate o f d i a t o m i c molecules and the p o s t s h o c k density gradient associated w i t h that rate under the assumptions o f

where M 0 is the m e a n initial molecular weight o f the test gas, Po and Pi the pre- and p o s t s h o c k densities, respectively, U s the shock velocity, (dp/dx)i the density gradient inferred f r o m the b e a m deflection and the specific refractivity o f the test gas, ct the e x t e n t o f reaction, and (da/dp)i can be evaluated b y solving the Rayleigh flow equations

Mop o Us(dp/dx)i(dot/dp) i kl -

pi 2

244

T. TANZAWA and W. C. GARDINER, JR. 40[I ~ /~

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A reaction mechanism for ethylene pyrolysis derived from preliminary consideration of our and previous workers' experiments is shown in Table 1. Table 2 shows the corresponding sensitivity spectrum for the final rate constant set [24]. (The sensitivity or pS value is the slope ofalog-log graph of a computed observable versus a rate constant; in Table 2 it was estimated by multiplying the Table 1 rate constants by 5. For pS = - 1 , one has a quantity that is inversely proportional to the assumed rate constant value; for pS = +1, linear proportionality exists; for pS = 0, the computed quantity is insensitive to the assumed value of the rate constant.) Table 2 confirms that Reaction (1) dominates at early stages of the reaction, before the chain mechanism initiated by Reaction (2) takes over, and that later parts of the profiles provide information about other elementary reaction rates. We modeled a number of our observed deflection profiles over the first 6 /as laboratory time. Comparisons o f experimental with computed profiles are shown in Fig. 3. (Deflection is defined here as the angular deflection of the laser beam in

Fig. 1. Representative laser-schlieren oscilloscope records: (a) T 2 = 2540°K, P2 = 3.85 gmol cm - 3 , 2.5% C2H 4 in Ar; (b) T 2 = 2300°K, P2 = 2.06/~mol cm- 3 , 5% C2H 4 in Ar; (c) T 2 = 2170°K, P2 = 2.44 #mol cm - 3 , 10% C2H 4 in Ar.

~

]0

for small values of a. This was done essentially by adding Ah = 165 kJ X o o c [ C 2 H 4 ] 2 / p 2 t o t h e specific enthalpy and solving the constant-area steady inviscid flow equations for Ap. Setting (da/dp)i = Aa/Ap gave constant derivative values for Aa up to 10 - 3 . An Arrhenius plot for Reaction (1) obtained by using this kl equation and the (dot/dx)i values calculated from the observed beam deflections is shown in Fig. 2, together with the JRD ka expression. The rate expression obtained from a least-squares fit (in the T - 1 , In k plane) to the initial deflection data for all three mixtures is

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k 1 = 3.00 X 1017 exp [ - 3 4 0 kJ/RT] X cm a mol-1 s-1.

Fig. 2. Arrhenius plot for k 1. - - - - R e f . [8]; A 2.5%, C2H4; • 5% C2H4; • 10% C2H4. Reduction of initial deflection data to k 1 values is described in the text.

THERMAL DECOMPOSITION OF ETHYLENE

245 TABLE 2

Final pS Spectra a for Typical Shock Tube Experiments on Ethylene Pyrolysis. This Work

JRDb

GKKNc

SSd

T= 2440°K

T= 2200°K

T= 1700°K

T= 1168°K

Reaction No.

(do/dx)o.5/~s

(dp/dx) 5#s

[H] 60~s

k3/2

k3/2

1 2 3 4 5 6 7 8 9 10 11 12 13

0.42 0.26 0.02 0.05 0 0 0.12 0 0 0 0 0 0

-0.02 0.12 0.02 0.16 0 -0.16 0.20 0.05 0 0 0 0 0

-0.78 0.81 0 -0.10 0 0 0 0 0 0 0 0 0

0.34 0.32 0.32 0.04 0.19 -0.11 0.35 0 0 0 0 0 0

0.15 0.03 0.19 0.23 0.43 -0.07 0.17 0 0 0 0 0 0

a See Ref. b See Ref. c See Ref. d See Ref.

[241. [8]. [4]. [1].

radians, assuming that the beam is deflected as a unit, divided by the optical path length through the shock tube, 7.62 cm [10]. Good agreement is found at the early stages o f reaction. The calculation does not, however, account for the small hump o f the deflection profile for the 10% C2H 4 mixture. It cannot be attributed readily to an aerodynamic disturbance, because the deflection proffrie behind a shock wave in pure argon always was flat over a substantially longer time interval. For these experimental conditions, that is, T < 2 5 4 0 ° K , the small hump could mean acceleration of reaction due to formation of 1,3-butadiene in Reaction (5). While the thermal decomposition o f 1, 3-butadiene has not yet been studied at shock tube temperatures, Skinner and Sokoloski [1] (SS)did report 1,3-butadiene as a final product in their low temperature study of ethylene pyrolysis, and concluded that it is stable enough that the el'fect due to 1,3-butadiene pyrolysis must be minor, The main decomposition reactions they suggested were: C4H6 = C2H4 + C2H2

AH~ = 164 kJ

(15)

C2H 6 = C2H 2 ÷ C2H 2 + H 2

AH~ = 330 kJ.

(16)

However, we excluded them from consideration for our conditions after preliminary computer simulations showed their effects to be negligible. At higher temperatures the small deflection " h u m p " observed in the 10% C2H4 mixture gradually became flatter, and at the highest temperatures studied it even became a dip, as seen in Fig. (a). The appearance of refractive index gradient "humps" in the latter stages o f reaction at lower temperatures indicates acceleration o f the reaction associated with endothermicity. This observation could be rationalized as 1, 3-butadiene formation, Reaction (5), competing with chain termination, Reaction (6). As there are no experimental data available, we estimated rate constants for Reactions (5) and (6) by optimizing the latter stage of the deflection profiles at the high and low temperatures. Reaction (5) was indeed influential for the dip of deflection at high temperatures and Reaction (6) for the hump of deflection at low temperatures. The calculations could never be made to reproduce a comparable hump in the lower temp-

246

T. TANZAWA and W. C. GARDINER, JR. 3

C2H 6 + M = C H 3 + C H 3 + M

(a)

2

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00

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Fig. 3. Comparison of experimental deflections (solid line) with simulations (broken line): (a) T2 = 2440°K, P2 = 1.95tamolcm--a, 10%C2H 4 in Ar;(b) T2 = 2330°K, p2 = 2.44/~mol cm-3, 5% C2H4 in Ar; (c) T2 = 2060°K, P2 = 3.21 jamol cm--3, 10% C2H4 in Ar.

erature region; however, Reaction (6) was deftnitely necessary to maintain a long, flat density gradient, Asaba et al. [2] (AYH) report, in opposition to SS, that they did not detect 1,3-butadiene as a final product. Methane and ethane, which are thermodynamically favored, were found instead in high concentrations in addition to acetylene, which was the most abundant pyrolytic product, Methane and ethane would have to be produced by reactions such as CH 4 + M = CH a + H + M

AH~ = 432 kJ

(17)

CH a + H = CH a + Hg.

AH~ = --1 kJ

(18)

AH~=367kJ

(19)

C2HB + H = CzHB + H2

AH~ = _ 30 kJ

(20)

C2H 5 + H = CHa + CHa

AH~ = --272 kJ

(21)

We calculated the concentrations of methane and ethane at the end of a 1 ms dwell time for the AYH conditions by adding these reactions to the mechanism in Table 1, using the rate constant expressions of Ref. [10]. As shown in Table 3, methane and ethane were computed to be two orders of magnitude lower in concentration than 1,3-butadiene at the lower temperature, while at the higher temperatures, the formation of methane, ethane, and 1,3-butadiene were replaced by diacetylene formation. We therefore excluded the reaction routes forming methane and ethane from the reaction mechanism. It does not appear likely that species with odd numbers of carbon atoms, or ethane, can be produced by pyrolysis o f ethylene. A broader comparison o f experiment with computation can be accomplished by selecting deflections at characteristic times in order to overview the comparison over a wide temperature range. We chose the deflections at 0.5, 3, and 6/as. The comparison is shown in Fig. 4. It is seen that over the experimental temperature range and over the observation period of the density gradients, the computation is in good agreement with the experiments. The next computational task was to reproduce the three-halves order overall rate constants. Since SS [1], AYG [2], and Skinner, Sweet, and Davis [3] (SSD) report overall rate constants that are first order in the ethylene concentration, we converted them into three-halves order by dividing the reported first-order rate constants by the square root o f the appropriate inert gas concentrations. The three-halves order rate constants from all cited studies are shown in Fig. 5. It is seen that the cornputation fails to reproduce the 0.466% C2H4 mixture results of SS, the 0.1% C2H4 mixture results o f SSD and the 10% Calla mixture results of AYH. Otherwise, excellent agreement is found over the temperature range from 1200 to 2200°K. The discrepancies for the SS, SSD, and AYH data do not admit o f a definite explanation. Both are single-pulse shock tube experiments in which

THERMAL DECOMPOSITION OF ETHYLENE

247 TABLE 3

Calculated Fractional Conversions to Pyrolytic Products in 10% C2H 4 in Ar for the Conditions o f AYHa: P5 = 10 atm, Dwell Time = 1 ms

C2H4 1500°K 2000°K 2500°K

0.78 0.17 7.8 × 10 - 3

C2H2

H2

0.13 0.77 0.85

0.17 0,84 1.1

C4H6 4.0 × 10 - 2 5.6 × 10 - 3 1.3 × 10 - 5

C4H2

CH4

C2H6

2.0 × 10 --4 2.2 × 10 - 2 6.8 × 10 - 2

1.9 × 10 --4 4.8 X 10 - 3 1.5 × 10 - 3

2.2 × 10 --4 3.1 × 10 --4 5.2 × 10 - 7

a See Ref. [2].

soot formation may occur before cooling by a rarefaction wave. This might account for the discrepancy. Alternatively, the assumed mechanism may simply not be complete enough, We also compared the experimental hydrogen atom profiles obtained by JRD [8] with com-

puted profiles, as shown in Fig. 6. Excellent agreement is found at the early stage of the reaction; however, deviations of up to 50% are observed at the last stages of observation, as shown in Fig. 6(b). We were unable to devise any mechanistic explanation for this discrepancy.

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Fig. 4. Deflection data for 2.5% C2H 4 in Ar. *, m, and • are taken at 0.5, 3.0, and 6.0 ps laboratory time.

248

T. TANZAWA and W. C. GARDINER, JR. 8

4

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0 I 4

I 5

I

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104 K/T Fig. 5. Simulation of three-halvesrate constants obtained by previous workers. • Homer and Kistiakowsky using ir emission; ..... GKKN using TOF mass spectrometry; • (.446% C2H~) and • (6% C2H4) SS with SPST; - - - AYH with SPST: . . . . . . . . . . SSD with SPST. Solid lines were computed for each experiment as indicated.

An overview of the species and reaction rate profiles implied by the Table 1 mechanism and rate constant expressions for a typical high temperature and a typical low temperature experiment is given in the appendix,

DISCUSSION The main question concerning the identity of the initial reaction in ethylene pyrolysis was answered by JRD [8] and is further verified by the present study. The unimolecular decomposition of ethylene occurs by two competing channels, that is, Reactions (1) and (2). Both Reactions (1) and (2) are required in the mechanism in order to account for the experimentally observed long, flat deflection profiles. This mechanism for ethylene pyrolysis also accounts for many, but not all, results from previous shock tube studies and has enough

radical chain character to account, qualitatively, for the rapid appearance of HD in pyrolysis of CzD4-CzH4 mixtures [4]. The mechanism even reproduced the slight change in activation energy observed around 1500°K. In the case of the AYH results [2], the computations could not confirm ethane and methane as major products, while 1, 3-butadiene did rise to high concentrations at the lower temperature, as reported by SS, [1 ], and so it appears likely that the experimental conditions of AYH (P5 ~" 10 arm) may have made analysis complex. We verified in a pS calculation that in the framework of the Table 1 mechanism, 1,3-butadiene formation through Reaction (5) is important in determining deflection profile humps. (Although not in determining the parameters shown in Table 2.) Alternative acceleration reactions may be sought, for example, in the mechanism of acetylane pyrolysis. We recently completed an extensive study of acetylene pyrolysis [18] in which we de-

THERMAL DECOMPOSITION OF ETHYLENE

249

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B

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u 20

•

.=. lO

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I 200

I

I 400

(a)

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TIME / ps

12

10

•

8

¢o ! tJ O

E

6

~-~ 4

2

O0

I

I 200

I

(b)

I 400

TIME / ps

Fig. 6. Simulation of the hydrogen atom concentration profiles obtained by JRD [8]. (a) T 5 = 1970°K, P5 = 11.9 gmol cm--3 50 ppm C2H 4 in Ar;(b) T 5 = 2060°K, P5 = 10.4 ~mol cm - 3 , 20 ppm C2H 4 in Ar.

250

T. TANZAWA and W. C. GARDINER, JR.

veloped a mechanism with which we successfully accounted for previous workers' data and our own laser-schlieren experiments on C2H2 pyrolysis over a temperature range from 1200 to 3300°K. With that complete set of acetylene pyrolysis mechanism reactions, but excluding the present Reaction (5), the small hump of the C2H 4 deflection profiles, previously mentioned, could not be accounted for. An acceleration by means of the acetylene pyrolysis mechanism, achieved by adjusting the rate constants of suitable reactions, could be made to generate a deflection hump, but it inevitably appeared much too soon in comparison with the experimental data. The following reactions may be considered as further alternatives, C2H4 + C2 H = C4H4 + H C4H 4 + M = C 4 H a + H + M

AH~ = - 6 2 kJ

We infer that the small deflection hump in our lower temperature experiments is indeed a result of 1, 3-butadiene formation. The C2H4 decomposition rate itself is so fast at higher temperatures that the contribution of Reaction (22) cannot be substantial there either.

CONCLUSIONS The two-channel initiation reaction mechanism established by .IRD was successfully verified by laser-schlieren experiments. A second-order rate constant for Reaction (1) was measured up to 2540°K. A proposed ethylene pyrolysis mechanism was shown to be capable of quantitatively explaining many previous observations taken over a wide range of experimental conditions.

(22)

AH~) = 3 3 0 k J

(23)

However, they are unlikely because in the low temperature region where only low concentrations of C2H are expected, Reaction (22) cannot compete with Reaction (5).

APPENDIX The course of reaction in shock-initiated C2H4 pyrolysis, as described by the Table 1 mechanism,

-6 -

C2H4

-8

-~

'7,

H2 C2H2 C4H6

-]0

C2H3 H

-12

-14

I 0

I 400

I

I 800

I

I ]200

I

I 1600

I

I 2000

TIME / ps F~. 7. Computed concentration profiles for an incidentshock wave in 10% C2H4 in Ar, T2= 2500°K, P2 = 3.85 pmol/cm 3.

THERMAL DECOMPOSITION

OF ETHYLENE

251

--4

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-10

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400

l

800

t

1200

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6 2

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1600 2000 TIRE / ]as

Fig. 8. Computed reaction rate profiles for the shock wave of Fig. 7. Numbers refer to elementary reactions as listed in Table 1.

-6" .... H2 C2H2 ~'~

-8 --

'-"

H

~

~

C4H2

~

~__

C2H4

-I0

C2H C4H6 C2H3

-12 ~ 0

~ 2

4

6

8 TIME / Us

Fig. 9. Computed concentration profiles for a reflected shock wave in 5% C2H 4 in Ar, T 5 = 1300°K, P5 = 35.4/~mol/cm 3.

252

T. TANZAWA and W. C. GARDINER, JR.

-2

-3

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-4

7

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4 -

r.D O ..-I

6

2

0

2

4

6

8 TIME / ]as

Fig. 10. Computed reaction rate profiles for the shock wave of Fig. 10. Numbers refer to elementary reactions as listed in Table 1.

is best understood by considering the computed profiles of species concentrations and reaction rates. These are shown in Figs. 7 and 8 for a high temperature experiment and in Figs. 9 and 10 for a low temperature experiment. The reaction conditions for Figs. 7 and 9 are typical of our experiments. In Fig. 7 we note the rapid drop of [C2H4] and its essentially quantitative replacement by CzH2 and Hz. H and C2H a are the important intermediates at short times, with only the former persisting. The further reaction products achieve only minor concentrations, and for describing the main fate of C2H4 one could ignore them. They must be important, of course, in the eventual production of soot. In Fig. 8 one notes that the primary decomposition rates, of Reactions (1) and (2), fall even more

rapidly than [C2H4], due to the temperature drop. Reaction (1) dominates the reaction only for the first microsecond, the chain sequence (4 + 7) later accounting for most of the decomposition. All other steps are of minor importance and could be omitted as far as accounting for C2H4 decomposition under these conditions is concerned. The reaction conditions for Figs. 9 and 10 are typical of the high concentration experiments of Ref. [1]. In Fig. 9 we see that [CzH4] remains nearly constant throughout the experiment. The chain carriers H and CzHs rise slowly in concentration, and the fate of the C2H4 is that about half goes to CzHz + Hz and half goes to C4H6 + Hz. In Fig. 10 these facts are clearly seen to be consequences of the chain reactions (4), (5), and (7). Chain initiation by Reactions (2) and (3) is now at

THERMAL DECOMPOSITION OF ETHYLENE a constant rate throughout the experiment, which accounts for the steady increase in rate of the chain reactions. Tile relative product yield is determined by the relative rates of Reactions (5) and (7), while the absolute yield, and hence the loss rate of C2H4, is determined by the rates of Reactions (3), (4), (5), and (7). (see Table 2). All other reactions are of minor importance as far as C2H4 decomposition under these conditions is concerned, The authors wish to thank the U.S. A r m y Research Office and the Robert A. Welch Foundation for grants that made this research possible.

REFERENCES 1. Skinner, G. B., and Sokoloski, E. M.,J. Phys. C h e m . 64:1028 (1960). 2. Asaba,T., Yoneda, K., and Hikita, T., Kogyo Kagaku Zasshi65:1811 (1962). 3. Skinner, G. B., Sweet, R. C., and Davis, S. K., J. Phys. Chem. 75:1 (1971). 4. Gay, I. D., Kern, R. D., Kistiakowsky, G. B., and Niki, H.,J. Chem. Phys. 45:2371 (1966). 5. Homer, J. B., and Kistiakowsky, G. B., J. Chem. Phys. 47:5290 (1967). 6. Roth, P., and Just, Th., Ber. Bunsenges, Physik. Chem. 77:1114 (1973). 7. Bauer, S. H.,Eleventh Symposium (International) on Combustion, The Combustion Institute, Pittsburgh, 1967, p. 105. 8. Just, Th., Roth, P., and Damm, R., Sixteenth Symposium [International) on Combustion, The Combustion Institute, Pittsburgh, 1977, p. 961.

253 9. Tanzawa, T., Sixteenth Symposium {International) on Combustion, The Combustion Institute, Pittsburth, 1977, p. 969. 10. Olson, D. B., Tanzawa, T., and Gardiner, Jr., W. C., Int. J. Chem. Kinet. 11:23 (1979). 11. Bertin, J. J., Ehrhardt, E. S., Gardiner, Jr., W. C., and Tanzawa, T., Proceedings of the Tenth Shock Tube Symposium, Kyoto, Japan, 1975, p. 595. 12. Lifshitz, A., Frenklach, M., Schechner, P., and Carroll, H. F.,Int. J. Chem. Kinet. 7:753 (1975). 13. Stull, D. R., and Prophet, H.,JANAFThermochemical Properties, 2nd ed., NSRDS-NBS 37, National Bureau of Standards, Washington, D.C., 1971. 14. Duff, R. S., and Bauer, S. H., J. Chem. Phys. 36: 1754 (1962). 15. Boyd, M. L., Wu, T. M., and Back, M. H., Can. J. Chem. 46: 2415 ( 1968). 16. Benson, S. W., and Haugen, G. R., J. Phys. Chem. 71:1735 (1967). 17. Just, Th., Private communication, 1977. 18. Tanzawa, T., and Gardiner, Jr., W. C., Seventeenth Symposium (International) on Combustion, Leeds, England, 1978, The Combustion Institute, 1979, p. 563 ; See also T. Tanzawa and W. C. Gardiner, Jr., J. Phys. Chem. 84:236 (1980). 19. Lange, W., and Wagner, H. G., Ber. Bunsenges, Physik. Chem. 79:165 (1975). 20. Hanbook of Chemistry and Physics, 50th ed., Chemical Rubber, Cleveland, Ohio, 1963. 21. Wright,J. K., Shock Tubes, Methuen, London, 1961. 22. Podolsky, B., Pro. Natl. Acad. Sci. U.S. 14:253 (1928). 23. Breshears, W. D., Bird, P. F., and Keifer, 1. H., J. Chem. Phys. 55:4017 (1971). 24. Gardiner, Jr.,W. C.,J. Phys. Chem. 81:2367 (1977).

Received 20 November 1978; revised 26 December 1979