Energy & Fuels 1990,4, 277-285
277
Inhibition of Cz Oxidation by Methane under Oxidative Coupling Conditions John C. Mackie* and Julie G. Smith Department of Physical Chemistry, University of Sydney, Sydney, NS W 2006, Australia
Peter F. Nelson and Ralph J. Tyler CSIRO Division of Coal Technology, North Ryde, NSW 2113, Australia Received November 13, 1989. Revised Manuscript Received March 2, 1990 The gas-phase partial oxidation of ethane and ethylene have been studied in a completely stirred tank reactor over the temperature range of 920-1220 K, at total pressures from 40 to 100 Torr, and at residence times between 0.1 and 1 s. Methane was found to have a pronounced inhibiting effect upon the partial oxidation of both C2 species. Ethylene is the principal carbon-containing product of partial oxidation of ethane. The ratio of the selectivities to ethylene and carbon monoxide depends only on the extent of conversion of the ethane and not on the presence of methane in the reactants. Carbon monoxide is the principal product of partial oxidation of ethylene. Selectivity to CO is reduced and selectivity to propene is enhanced when methane is present in the reaction mixtures. A detailed chemical kinetic reaction model has been developed to simulate the partial oxidation of the C2 species and to model the inhibiting effect of the methane on C2 oxidation.
Introduction There is considerable interest in the catalytic conversion of methane, the principal constituent of natural gas, into C2 hydrocarbons, especially ethylene, under partial oxidation conditions at around 800 OC.lv2 It is now generally accepted that the role of the catalyst, usually an alkaline-earth oxide, is to produce methyl radicals and that the reaction to form C2 hydrocarbons takes place essentially completely in the gas The oxidative coupling route from methane to C i s is currently under investigation in several laboratories as a possible commercial process for production of higher hydrocarbons. A potential problem for the utilization of this technique for higher hydrocarbon production is that under the high-temperature oxidative conditions the desired C2 products are considerably more reactive toward oxidation than is methane itself. Thus it is important to ascertain reaction conditions that optimize the yields of the C2 hydrocarbons in the oxidative coupling of methane. Preliminary studies over a lithium-promoted magnesium oxide catalyst indicated that the presence of methane in the product gas inhibited the oxidation of both ethane and ethylene under very fuel-rich conditions at about 800 O C 5 We now report a systematic gas-phase study designed to investigate the inhibition of C2oxidation by methane. The very fuel-rich conditions included reaction between ethane and oxygen and between ethylene and oxygen, in both the presence and absence of methane. Mixture compositions, temperatures, and reactor residence times studied are characteristic of those employed in the catalytic oxidative coupling of methane. There have been few previous studies of the role of additives such as methane in retarding the oxidation of C2 hydrocarbons although there have been some studies on the optimization of C2 production in pyrolysis of binary 1985,107,5062. (2) Jones, C. A.; Leonard, J. J.; Sofranko, J. A. J. &tal. 1987,103,311. (3) Lunsford, J. H.: Campbell, K. D.: Morales, E. J. Am. Chem. SOC. 1987,109,7900. (4)Labinger, J. A.; Ott, K. C. J. Phys. Chem. 1987,91,2682. (5)Tyler, R. J. In Proceedings of Catalysts for Fuels Synthesis and Processtng Workshop; National Energy, Research, Development and Demonstration Council: Melbourne, 1986;p 75.
CH4/H2 mixtures.6 Despite the extensive literature of reaction modeling of combustion of hydrocarbon^,^^^ there have been few attempts to model the oxidation kinetics of C1 or C2 hydrocarbons a t the extremely high fuel/oxidant ratios appropriate to the present oxidative coupling studies. Labinger and Ott4 have presented a simple model for oxidative coupling of methane over a mixed Mn-Mg oxide surface. Their principal interest was in the heterogeneous mechanism, and their gas-phase mechanism omitted many of the free radicals considered relevant to the gas-phase chemistry. Dagaut et al.9 have presented a kinetic model for ethylene oxidation that has been compared with experiment principally for lean and stoichiometric mixtures although some studies of fuel-rich C2H4/O2 mixtures were reported. There does not appear to be any previous attempt to model binary mixtures of C1 and C2 hydrocarbons under partial oxidation conditions. In the present work a detailed kinetic reaction model for the gas-phase C1 and C2 chemistry is developed for partial oxidation conditions and tested experimentally. Experimental Section The hydrocarbon (ethane or ethylene) together with oxygen has been studied in dilute mixtures in argon in both the presence and absence of methane in a completely stirred reactor.1° The vessel comprised a fused-silica outer sphere of net volume 270 cm3. At the center of this sphere was a smaller spherical injector containing 28 holes of approximately 0%" diameter through which reactants were radially injected into the larger sphere. Product streams were taken from three points on the circumference of the outer sphere and,after uniting the streams, passed to the gas analysis system. The reactor was contained in an electrically heated furnace. Temperature measurements indicated (6)Klotz, H. D.; Schulz, G.; Drost, H.; Spangenberg, H. J. 2.Phys. Chem. (Leipzig) 1985,266 (l),101. (7) Warnatz, J. Rate Coefficients in the C/H/O System. In Combustion Chemistry; Gardiner, W. C., Ed.; Springer-Verlag: New York, 1984. (8) Tsang, W.; Hampson, R. F. J. Phys. Chem. Ref. Data 1986,15, 1087. (9)Dagaut, P.;Cathonnet, M.; Boettner, J. C.; Gaillard, F. Combust. Flame 1988,71,295. (10)Mackie, J. C.;Doolan, K. R.; Nelson, P. F. J. Phys. Chem. 1989, 93. 664.
0887-0624 /90/2504-0277$02.50/0 , 0 1990 American Chemical Society I
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Mackie et ai.
278 Energy & Fuels, Vol. 4, No. 3, 1990 ((1)
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Figure 1. (a, b) Temperature dependence of designated species mole fractions in the partial oxidation of ethane. CzHs/Oz/Ar = 4/4/92; total pressure = 50 Torr; t,, = 0.2 s. Symbols are the measured mole fractions; lines are the kinetic model predictions. that the difference between center and wall temperatures of the reactor was less than 1 "C. Temperatures ranged from 920 to 1220 K, total pressures ranged from 40 to 100 Torr, and residence times ranged between 0.1 and 1.0 s. Species concentrations in both reactant and product streams have been measured. C1-C4 hydrocarbons, the carbon oxides, oxygen, hydrogen, and argon have been determined by gas chromatography;yields of water have been measured by means of an electronic hygrometer. Acetaldehyde has been determined by GC and formaldehyde both by the chromotropic acid test and by FTIR spectrometry. Negative tests for hydrogen peroxide were recorded by the triiodide method.
Results and Discussion To ascertain the extent to which methane itself underwent partial oxidation in the temperature region of interest of the ethane and ethylene partial oxidation studies (920-1220 K), some preliminary measurements were made on a mixture of CH4/0,/Ar = 37/4/59 and total pressure 44 Torr. At residence times of 0.1 and 0.3 s no detectable decomposition of methane occurred at temperatures up to 1220 K. At a residence time of 1.1 s the extent of decomposition of CH4 was 0.1% at 1170 K and 0.3% at 1220 K. The only products detected were ethane (first detected at 1170 K) and ethylene (first detected at 1220 K).
0.001
1000
I
1050
1
1100 1150 TEMPERATURE, K
1200
Figure 2. (a, b) Temperature dependence of designated species mole fractions in the partial oxidation of ethane containing added methane. CH,/CzH,/Oz/Ar = 20/4/4/72; total pressure = 50 Torr; t,, = 0.2 s. Symbols are the measured mole fractions; lines are the kinetic model predictions.
Ethane Partial Oxidation. In fuel-rich ethane mixtures (CZH6/O2/Ar= 4/4/92) and pressures between 50 and 100 Torr, the principal products of partial oxidation were C2H4, H,, HzO, and CO. Minor products included CHI and, at the highest temperatures of the studied range, C2H,. Yields of COz were generally low and only trace amounts of C, and C4 hydrocarbons were detected. A t the lowest temperatures at which decomposition was observed the product distribution was more typical of pyrolysis than of partial oxidation with C2H4and Hz predominating in the products. In Figure 1 are shown the mole fractions of all species of significance as a function of temperature in the partial oxidation of ethane at a total pressure of 50 Torr and a reactor residence time of 0.2 s. In this series the reactants were fed in the ratio C2H6/02/Ar= 4/4/92. When methane was added to the reactant mixture, strong inhibition of the decomposition of ethane was observed. In Figure 2 are shown the mole fractions of all species of significance as a function of reactor temperature in the partial oxidation of a reactant mixture of ratio CH4/C2H6/O2/Ar= 20/4/4/72, total pressure 50 Torr, and residence time 0.2 s. Significance reductions in oxygen consumption and in CO and HzO yields are also observed. In order to gauge the effectiveness of methane in inhibiting the partial oxidation of ethane, a series of runs
Energy & Fuels, Vol. 4, No. 3, 1990 279
Inhibition of C2 Oxidation by Methane
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Figure 3. Effect of added methane on the extent of decomposition of C& (solid line) and CzH4 (dashed line). Cz/02 = 4/4. Total pressure = 50 Torr; t,,, = 0.2 s; temperature = 1113 K (constant)for CzH6;temperature = 1178 K (constant)for CzHk
was carried out in which the mole fraction of methane in the reactants was varied. Total pressure and reactor residence time were maintained constant at 50 Torr and 0.2 s, respectively. A constant temperature of 1113 K was chosen for this study. At this temperature, without methane present in the reactants, ethane is 78% decomposed under these reactor conditions. The extent of decomposition of ethane as a function of the mole fraction of methane in the reactants is plotted in Figure 3. As may be seen from this figure, methane has a significant inhibiting effect on partial oxidation of ethane and, at only 10% of the reactant mixture, reduces the decomposition of ethane by a factor of more than 3. Whether methane is present or not in the reactants, the principal carbon-containing products of partial oxidation of ethane are ethylene and carbon monoxide. Kinetic modeling studies, discussed below, support the postulate that, in fuel-rich mixtures with oxygen, dehydrogenation of ethane to ethylene (which also takes place under purely pyrolytic conditions) precedes partial oxidation to carbon monoxide. Two questions, however, may be posed: (i) In the absence of methane in the reactants does the selectivity to C2H4 and CO depend on pressure, residence time, or temperature? (ii) Does the presence of CH4 in the reactants change the mechanism and thereby the selectivity to CzH4 and CO? In Figure 4 selectivities to C2H4 and CO are plotted for all runs irrespective of residence time or pressure as a function of the extent of decomposition of ethane. Runs with and without methane added to the reactants are included in this figure. It may be clearly seen that, irrespective of pressure, reactor residence time, or the presence of CH4 in the reactants, the selectivity to CzH4 or to CO depends solely on the extent of decomposition of the C2HG. This is an important result and implies that there is no change in the reaction mechanism over the studied range of pressures, residence times, or methane concentrations for a given extent of decomposition of ethane. The modeling studies support this conclusion and enable identification of the important rate-determining reactions to be made. Ethylene Partial Oxidation. Fuel-rich ethylene mixtures (C2H4/Oz/Ar= 4/4/92) gave a product distribution characteristic of partial oxidation over the entire
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280 Energy & Fuels, Vol. 4, No. 3, 1990 I
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/ 't C2H6. @ C i s 1 I 1 1000 1020 1040 1060 1080 1100 1120 TEMPERATURE, K
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Figure 5. (a, b) Temperature dependence of designated species mole fractions in the partial oxidation of ethylene. CpH4/0,/Ar = 4/4/92; total pressure = 50 Torr; t,, = 0.2 s. Symbols are the measured mole fractions; lines are the kinetic model predictions. To gauge the extent of inhibition of partial oxidation of ethylene by the addition of CH, to the reactant mixture, a series of runs was carried out at a constant temperature of 1178 K, total pressure of 50 Torr, residence time 0.2 s, and reactant mole percentages of methane ranging from 0 to 40%. At this temperature and these conditions in the absence of CH, in the reactant mixture, ethylene undergoes decomposition to the extent of 63 %. As shown in Figure 3, at only 10 mol % methane in the reactant mixture the extent of decomposition of ethylene is reduced by a factor of more than 3. Thus methane has a comparable inhibiting effect on the partial oxidation of both ethylene and ethane. Reaction Model for Partial Oxidation. A chemical kinetic reaction model that uses the Sandia Laboratories CHEMKIN" and P S R ' ~computer codes has been used to (11) Kee, R. J.; Miller, J. A.; Jefferson, T. H. CHEMKIN. A General-Purpose, Problem-Independent, Transportable Fortran Chemical Kinetics Code Package; SAND 80-8003; Sandia National Laboratories: Albuquerque, NM, 1980.
1.0
I
I
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I
1
5.0
Figure 6. (a) Percent conversion of CpH4 as a function of input mol % of 02. (b) Mol % formaldehyde in products as a function of input mol % of O2 For both (a) and (b), temperature = 954 K (constant),total pressure = 100 Torr, and t,,, = 1.0 s. Inlet CpH4 = 10 mol % (constant). model the reaction products of partial oxidation. The CHEMKIN code provides a symbolic description of a chemical reaction mechanism together with thermochemical properties, rates of chemical production, and sensitivities formulation. The PSR (perfectly stirred reactor) code package, used in conjunction with CHEMKIN, solves the conservation equations and chemical rate equations for a perfectly stirred steady-state tank reactor of specified volume and residence time using a hybrid Newtonltimestepping algorithm. Rate constants in the reaction model have been taken directly from the combustion l i t e r a t ~ r e . ~ , The ~ , ' ~present model incorporates the hydrogen/oxygen flame reactions which produce H, 0, OH, and H 0 2 radicals. Reactions of each of these radicals with all C1 and C2 species present are included. Reactions of methyl, ethyl (C,H,), and vinyl (CzH,) radicals are known to be important in both pyrolysis and oxidation of hydrocarbons and the reactions of these radicals with C1 and Cz species are also included. Although CH2biradicals have not so far been found to play an important role in hydrocarbon oxidation, reactions of this biradical have been included for completeness. At the temperatures of the present study, using the equilibrium constants of Gutman et al.,14 we estimate that the concentrations of the peroxy radicals, ROz, should be suffi(12) Glarborg, P.;Kee, R. J.; Grcar, J. F.; Miller, J. A. PSR: A Fortran Program /or Modeling Well-Stirred Reactors; SAND 86-8209; Sandia National Laboratories: Albuquerque, NM,1986. (13) Glarborg, P.; Kee, R. J.; Miller, J. A. Combust. Flame 1986,65, 177. (14) Slagle, I. R.;Gutman, D. J. Am. Chem. Soc. 1985, 107, 5342.
Energy & Fuels, Vol. 4, No. 3, 1990 281
Inhibition of C2 Oxidation by Methane 1oc
monoxide, carbon dioxide, formaldehyde, acetaldehyde, ketene, hydrogen peroxide, water, hydrogen, propene, and propane. As only traces of C4 hydrocarbons were recorded in the experiments, for tractability in the model, C4'shave been omitted. Also, since C3's are only minor products, to reduce the total number of fundamental reactions in the mechanism, only a rudimentary mechanism involving propene and propane has been incorporated. Indeed, for modeling ethane, C i s could have been omitted because of their unimportance as products. However, the kinetic model was developed to model ethylene partial oxidation reactions as well as those of ethane, and C3 products from C2H4have higher yields. The detailed reaction model is presented in Table I. Reactions 1-18 are the subset of well-known hydrogen/ oxygen flame reactions. Reactions 19-105 have mostly been taken from the combustion literat~re.~,"~ With three exceptions, rate constants for these reactions in the form k = AT" exp (-E/RT) have also been taken directly from the literature. The exceptions are Reaction 42, the unimolecular decomposition of the HCO radical whose rate constant was adjusted by a factor of 1.3 to match the experimental CO profiles: reaction 53, increased by a factor of 2; and Reaction 67, decreased by a factor of 2. The last two adjustments gave better agreement with the observed ethylene profiles. Each of these adjustments was within the recommended uncertainties8 in the respective rate constants. A model containing reactions 1-105 was found suitable for modeling the ethane partial oxidation data but failed to predict the observed low-temperature reactivity of ethylene. Most combustion models for ethylene use the initiation reactions C2H4 + M = C2H2 H2 M (73) and C2H4 + 0 2 = CpH3 + HO2 (105) but under the presence conditions this initiation is too slow to lead to the observed products. Ayranci and Back15 have suggested the initiation reaction C2H4 + C2H4 = C2H3 + C2H5 (106) but the vinyl and ethyl radicals thus formed decompose too slowly at low temperatures. Thus, this reaction, too, is unable to simulate the products observed a t low temperatures. Benson16 suggested the possible direct bimolecular reaction C2H4 + 0 2 = CH2O CH2O
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r
0 01
co2
1060 1100
1120 1140 1160 1160 TEMPERATURE, K
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12x)
Figure 7. (a, b) Temperature dependence of designated species mole fractions in the partial oxidation of ethylene containing added methane. CH4/&H4/O2/Ar = 20/4/4/12;total pressure = 50 Torr; t , = 0.2 s. Symbols are the measured mole fractions; lines are the kinetic model predictions.
ciently low to be neglected relative to the concentrations of R and of HOP. Thus reactions of R 0 2 have not been considered in the model. The following stable molecules have been considered in the model: methane, ethane, ethylene, acetylene, carbon
might initiate low-temperature oxidation of ethylene. The above reaction was thought to take place by initial addition of oxygen across the double bond of ethylene. Experimentally, a t the lowest temperatures a t which decomposition of ethylene can be detected, formaldehyde, carbon monoxide, hydrogen, and water are observed in the products. If, however, the Benson initiation reaction is used, it grossly overpredicts formaldehyde yields and cannot predict the significant yields of CO, H2, and H 2 0 found experimentally since CHzO is known to decompose too s10wly'~at the lowest temperatures of the present study. (15) Ayranci, G.; Back, M. H. Cap. J. Chem. 1981, 46, 2415. (16) Benson, S. W. Thermochemrcal Ktnetrcs, 2nd ed.; Wiley: New York, 1976; p 239. (17) Dean, A. M.; Johnson, R. L.; Steiner, D. C. Combust. Flame 1980, 37, 41. (18) Kerr, J. A.; Parsonage, M. J. Evaluated Kinetic Data on Gas Phase Hydrogen Transfer: Reactions of Methyl Radicals; Butterworths: London, 1976. (19) Payne, W.; Stief, L. J. J. Chem. Phys. 1976, 64, 1150.
Mackie et
282 Energy & Fuels, Vol. 4, No. 3, 1990
Table I. Kinetic Model for Partial Oxidation" forward rate constant no. 1
2 3 4 5 6 7 8 9 10 11
12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60 61 62 63 64 65 66 67 68 69 70 71 72
73 74
75
reaction H2 + Oz = 2 0 H H2 + OH = HzO + H H + 02 = OH + 0 0 + Hz = OH + H H + 02 + M = HOz + M OH + HOz = HzO + 02 H + HOZ = 2 0 H 0 + HOz = 0 2 + OH 2 0 H = 0 + HzO HzOz + M 2 0 H + M HzOz + H = H20 + OH HZ02 + H = HOz + Hz HzOz + OH = Hz0 + HOz 2H02 = HzOz + 0 2 H2 + M = 2H + M Oz + M = 2 0 + M H + OH + M = Hz0 + M H + HOz = Hz + 02 CH4 + M = CH3 + H + M CH4 + 0 2 = CH3 + HO2 CH3 + Hz = CH, + H CHI + 0 = CH3 + OH CH, + OH = CH3 + H20 CH4 + HOz = CH3 + HzO2 CH3 + HOz = CH30 + OH CH, + 0 = CHzO + H CH3 + OH = CHzO + Hz CH3 + 0 2 = CH20 + OH CH, + O2 = CH30 + 0 CH30 + M = CHzO + H + M CH30 + H = CH20 + H2 CH30 + 0 = CHzO + OH CH30 + OH = CHZO + HzO CH30 + Oz = C H z O + H 0 2 CHZO + M = HCO + H + M CHzO + H = HCO + H2 CHzO + 0 = HCO + OH CHzO + OH = HCO + HzO CHZO + 0 2 = HCO + HOz CHZO + HOz = HCO + H20z CHzO + CH3 = CH4 + HCO HCO + M = CO + H + M HCO + H = CO + Hz HCO + 0 = CO + OH HCO + 0 = C 0 2 + H HCO + OH = CO + H2O HCO + Oz = CO + HOz CO + 0 + M = COP + M co + oz= coz+ 0 CO + OH = COZ + H CO + HOz = CO2 + OH C2H6 = 2CH3 C2H5+ H = 2CH3 CzH6 + H = CzH5 + Hz CzH6 + HOz = CzH5 + HZO, CzH6 + 0 = CzH5 + OH CZH, + OH = C2H5 + HzO CzH6 + CH3 = CzHS + CH4 CzH5 = CzH4 + H CzH5 + 02 = CzH4 + HOP CzH5 + H02 = CH3 + CHzO + OH CpH5 + H C2H4 + H2 CzH5 + 0 = CH3CHO + H CZH, + 0 = CHZO + CH3 CzH4 + M = CzH3 + H + M CzH4 + HOz = CH3CHO + OH CpH4 + H = CzH3 + Hz CzH4 + 0 = CH3 + HCO CZH4 + 0 = CHzO + CHZ CzH4 + OH = C2H3 + H20 CPH4 + OH = CH3 + CHzO C2H4 + CH3 = CH4 + CzH3 C2H4 + M = CzHz + Hz + M H + C2Hz = C2H3 CzH3 + H = C2Hz + Hz
log A
n
E
13.23 9.08 17.09 10.26 18.30 13.70 14.40 13.68 8.78 17.08 14.51 12.23 13.00 12.30 12.35 11.27 23.88 13.48 17.00 13.60 2.82 7.08 3.54 11.26 13.30 13.83 12.00 13.72 12.85 14.00 13.30 13.00 13.00 10.80 16.52 13.40 13.26 9.53 13.65 12.30 11.00 14.32 14.08 13.48 13.48 13.48 13.52 13.51 12.40 7.18 13.76 22.50 13.78 2.73 11.48 13.40 9.94 -0.26 13.30 12.51 13.38 12.28 13.70 13.20
0.0 1.3 -0.9 1.0 -1.0 0.0 0.0 0.0 1.3 0.0 0.0 0.0 0.0 0.0 0.5 0.5 -2.6 0.0 0.0 0.0 3.0 2.1 3.1 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 1.2 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 -0.4 0.0 0.0 1.3 0.0 -1.8 0.0 3.5 0.0 0.0 1.0 4.0 0.0 0.0
47.8 3.6 16.6 8.8 0.0 1.0 1.9 1.0 0.0 45.5 8.9 3.8 1.8
0.0
0.0 0.0 0.0 0.0
17.41
9.78 5.81 9.20 13.40 12.68 12.30 0.82 17.41 12.74 13.30
0.0 0.0 0.0 0.0 0.0
2.5 1.2 0.0
0.0 0.0
3.7 0.0 0.0 0.0
0.0
92.6 95.6 0.0
0.7 88.0 56.9 7.7 7.6 2.0 18.6 0.0 0.0 0.0
34.6 25.6 25.0 0.0 0.0 0.0
2.6 81.0 4.0 3.1 -0.4 41.4 11.7 6.1 14.7 0.0 0.0 0.0 0.0 0.0
-4.2 47.7 0.8
22.9 91.1 0.0
5.2 14.9 6.4 1.8 8.3 39.7 5.0
96.6 8.0 12.2 0.7 5.0 1.2 1.0 9.5 79.3 2.4 0.0
reverse rate constant n E 11.70 0.0 29.2 14.28 0.0 21.8 12.75 0.0 -1.8 0.0 8.9 13.38 0.0 44.8 15.03 0.0 73.8 14.78 0.0 40.7 13.25 13.72 0.0 56.3 0.0 20.2 14.33 0.0 -4.5 15.06 0.0 78.6 13.77 0.0 19.2 11.95 13.41 0.0 32.7 0.0 41.9 12.97 0.0 -10.6 13.50 11.57 0.0 -22.1 0.0 114.1 16.13 0.0 58.0 13.86 0.0 -17.4 14.85 0.0 -1.6 11.65 0.0 15.3 14.80 0.0 8.9 12.38 0.0 22.9 13.36 0.0 2.0 9.97 0.0 44.0 13.73 0.0 70.0 15.02 0.0 72.0 13.53 0.0 88.0 13.72 0.0 14.4 13.23 0.0 -12.5 13.37 0.0 66.8 13.26 0.0 64.7 12.61 0.0 82.2 13.65 0.0 12.0 10.38 0.0 -9.3 14.68 0.0 18.0 12.15 0.0 15.0 11.66 0.0 31.5 13.07 0.0 -2.0 12.02 0.0 10.2 11.33 0.0 21.2 11.32 0.0 -0.2 14.42 0.0 89.4 14.76 0.0 87.3 13.81 0.0 110.9 16.12 0.0 104.8 14.85 0.0 31.2 12.43 0.0 121.6 16.05 0.0 54.7 13.52 0.0 25.6 13.99 0.0 85.3 14.92 0.0 -1.4 12.75 11.65 0.0 9.8 0.0 18.6 14.05 0.0 5.4 10.93 0.0 10.2 12.23 0.0 25.4 13.45 0.0 23.9 14.36 0.0 1.7 12.46 0.0 14.0 11.87 0.0 16.3 11.05 0.0 66.3 12.03 0.0 72.6 13.94 0.0 79.8 12.26 0.0 -9.6 14.86 0.0 55.1 9.47 12.62 0.0 15.7 0.0 30.7 11.43 0.0 13.2 11.90 11.41 0.0 14.7 0.0 16.4 11.96 0.0 16.6 13.26 0.0 37.1 15.73 0.0 42.7 12.46 14.17 0.0 64.0
log A
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a b a b a b
a a a b a c b b a c c a a a a b b
a a a a b c
c b c d e c b b c a
a b b b c*
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Inhibition of C2 Oxidation by Methane
no. 76 77 78 79 80 81 82 83 84 85 86 87 88 89 90 91 92 93 94 95 96 97 98 99 100 101 102 103 104 105 106 107 108 109 110 111 112 113
reaction C2H3 + 02 = HCO + CH2O CoH, + 0 = CHXO + H C;H, + OH = CiH2 + H 2 0 CzH3 + OH = CH3CHO CHzCO + M =CH2 + CO + M CHzCO + OH = CHZO + HCO CH2C0 + OH = HCCO + H 2 0 CH2C0 + 0 = HCCO + OH CH2C0 + 0 = CH20 + CO CHZCO + H = HCCO + Hz CHzCO + H = CH3 + CO HCCO + 0 2 = 2CO + OH HCCO + 0 = 2CO + H HCCO + H = CHZ + CO HCCO+OH=HCO+H+CO CHBCHO = CH3 + HCO CHSCHO + 0 2 = CH3CO + HOZ CH3CHO + H = CH3CO + H2 CHBCHO + OH = CH3CO + Hz0 CHBCHO + 0 CH3CO + OH CHSCHO + CH3 = CHSCO + CH4 CH3CH0 + HOz = CH3C0 + H20z CH3C0 = CH3 + CO CzH2+ 0 = HCCO + H CH2 + 0 2 = CH2O + 0 2CH2 = CZHz + Hz CH3 + CHZ = C2Hd + H CH3 + H = CHZ + H2 C,H, + OH = CH, + CO C;H; + o2= c,H~-+ H O ~ 2CzH4 = CZH5 + C2H3 C2Hd + 0 2 = CHZO + HCO + H C2H5 + CzH4 = CzHG + C2H3 C2H3 + CH3 = C3He C2H5+ CH3 = C3H8 i-C3H7 + H = C3H8 CH3 + CzH4 = i-C3H7 C3H6 + H = i-C3H7
Energy & Fuels, Vol. 4 , No. 3, 1990 283 Table I (Continued) forward rate constant log A n E 12.60 13.52 13.48 13.48 15.56 13.45 12.88 13.70 13.30 13.88 13.04 12.18 12.08 14.04 13.00 15.30 13.30 13.60 13.00 12.70 10.93 12.23 13.48 13.72 13.30 13.51 13.60 14.95 12.08 13.62 15.00 13.00 11.56 13.40 12.85 13.30 11.52 12.86
0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.5 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0
-0.3 0.0 0.0 0.0 59.3 0.0 3.0 8.0 0.0 8.0 3.4 2.5 0.0 0.0 0.0 79.1 42.2 4.2 0.0 1.8 6.0 10.7 17.2 3.7 9.0 0.0 0.0 15.1 0.5 57.6 66.0 37.0 16.6 0.0 0.0 0.0 7.7 1.2
reverse rate constant log A n E 12.98 0.0 84.6 15.32 15.04 16.87 13.07 13.21 12.00 11.79 13.40 12.31 11.95 11.50 12.59 13.71 13.18 12.28 14.31 13.26 13.35 12.01 12.15 12.17 11.48 13.07 14.46 15.91 16.29 14.14 12.26 11.27 13.29 10.82 10.41 17.08 15.55 13.93 14.43 14.65
0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0 0.0
85.2 79.4 114.6 -16.7 16.1 15.7 3.3 103.5 5.3 37.0 91.2 105.2 31.0 17.9 -4.1 5.1 23.4 34.6 18.9 26.3 14.4 4.3 22.2 69.7 129.0 61.7 9.9 57.3 -1.7 -2.3 15.6 8.8 95.0 83.8 95.4 34.0 38.7
ref
i
a C C
a a a a a a a a
b a a b a
b b b b a a
k a b
b a a C
1
PW m C
b b d d
OReference 13. *Reference 7. cReference 8. dReference 18. eRate constant X 1.3 that of ref 13. /Rate constant X 2 that of ref 8. #Rate constant X 0.5 that of ref 8. Reference 19. Vinyl thermochemistry from ref 20. 'Reference 21. Rate constant from ref 8 but different products (ref 22). 'Reference 15. "'Reference 23. "Units: A, cm3 mol s-l; E , kcal mol-'. PW, rate constant optimized in present work. *, pressure-dependent rate constant. Value given is k,. Falloff data as cited in reference.
The bimolecular reaction between ethylene and oxygen is very exothermic, and there is sufficient exothermicity to decompose one of the formaldehyde molecules into HCO radicals and H atoms. Furthermore, HCO undergoes rapid reaction into CO and H. Such a process would provide the radicals necessary to decompose the ethylene at low temperatures. Therefore, in an attempt to simulate the experimental observation of significant CO, H2, CH20, and H20 at low temperatures, the following initiation has been postulated: C2H4 + 02 = CH20
+ HCO + H
(107)
Including this reaction in the model enables the low-temperature products of partial oxidation of ethylene to be simulated satisfactorily. Both vinyl and ethyl radicals have been found to be important in these fuel-rich partial oxidation reactions. Such conditions are not usually encountered in combustion studies. Reaction 108 involving reaction of these radicals with the C i s has been included for the present conditions. Since small yields of propene and propane were observed especially in partial oxidation of ethylene, reactions (20) Weissmann, M. A.; Benson, S. W. J. Phys. Chem. 1966,92,4080. (21) Slagle, I. R.; Park, J.-Y.; Heaven, M. C.; Gutman, D.J. Am. Chem. SOC.1964,106, 4356. (22) Harding, L. B. J. Phys. Chem. 1981, 85, 10. (23) Halstead, M. P.; Quinn, C. P. Trans. Faraday SOC.1966,64,103.
109-113 were included. These comprise a minimal set of C3 reactions necessitated to keep the number of reactions within tractable limits. Thus the model can only be expected to give semiquantitative simulations of the C3 yields. The reaction model contains several pressure-dependent reactions. The most important of these are reactions 59 and 74, unimolecular reactions of C2H5 and C2H3, respectively. Their pressure dependences are obtained from falloff calculations given in the references cited in Table I. The model gives good agreement with the experimental ethane reaction products (see Figure 1). The same model has been used to predict the observed inhibition of partial oxidation by methane in a mixture of CH4/C2H6/02/Ar = 20/4/4/72. The results of this modeling are shown in Figure 2. The model correctly predicts the inhibition by methane and reproduces all the major observed experimental species profiles quite well. (The experimental yields of water and carbon monoxide are low throughout this series and have appreciable experimental error.) The kinetic model has also been used to simulate the mole fractions of species present as a function of temperature in the partial oxidation of a mixture of C2H4/ 02/Ar = 4/4/92, total pressure 50 Torr, and residence time 0.2 s. Predicted species profiles are compared with experiment in Figure 5. The model correctly predicts the rate of disappearance of O2 with temperature although it
284 Energy & Fuels, Vol. 4, No. 3, 1990
Muckie et al.
slightly overpredicts the rate of decomposition of C2H4 and H2 a t high temperatures. I t predicts the CO and water yields adequately although underpredicting the observed C02. Most of the hydrocarbon mole fractions are well predicted, especially CHI and C2H6. C3 (propene propane) mole fractions computed by the model agree with experiment as well as can be expected from a minimal C3 set of reactions. The model predicts that a significant concentration of formaldehyde should be present in the products from ethylene even a t the lowest temperatures at which decomposition can be observed. There is considerable experimental error in the measured CH20 mole fractions so that the agreement between experiment and theory can be considered to be quite satisfactory. The only significant discrepancy is in the computed acetylene profiles. The model seriously underpredicts this product, which, however, is never produced in high yield. Comparison between model predictions and experiment are reported in Figure 7 for the partial oxidation of a mixture of CH4/C2H4/02/Ar = 20/4/4/72, total pressure 50 Torr, and residence time 0.2 s. The model correctly predicts the strong inhibition of oxidation of the ethylene by methane and gives a good prediction of the C2H4 and O2profiles with temperature. The model gives an adequate representation of the product mole fractions; however, hydrogen is overestimated at the lower temperatures and water (whose experimental concentrations are significantly lowered over the methane-free runs) is underestimated over the entire range. The model predicts very low yields of carbon dioxide. Indeed, the levels in the product mixtures were too low to measure experimentally. Sensitivity coefficient and rate of production analysis of the kinetic model have been used to identify the ratedetermining reactions in the mechanism. An important question that can be answered by this analysis is: What reactions are important in the inhibition of partial oxidation caused by methane? Consider first the important reactions in the partial oxidation of ethylene as derived from the kinetic model by sensitivity and rate of production analyses. In the absence of methane in the reactants, the most sensitive reaction in the model is the well-known chain-branching reaction
When, however, methane is present in excess in the reactants, by far the most sensitive reaction is that between methyl radicals and ethylene:
C2H4 + CH3 = C2H3 + CH4
+
H
+ 02 = OH + 0
(3)
which leads to rapid acceleration of the oxidation reactions leading to carbon oxides and water. Without methane in the reactants, all of the radicals important in oxidation, OH, H 0 2 , 0, and H, together with O2 and CH3 attack the C2H4 and lead to rapid formation of C2H3 radicals. Most important of these is
C2H4 + OH = C2H3 + H2O
(70)
Overwhelmingly, the most important mode of decomposition of vinyl radicals is by reaction with O2 via
C2H3 + 02 = HCO
+ CH20
(76)
thereby rapidly forming CO. Other important reactions which lead to acceleration in the rate of partial oxidation are the degenerate branching reaction H202 + M = OH + OH + M (10) and the decomposition reactions of the HCO radical HCO + M = CO + H M (42)
+
HCO
+ 0 2 = CO + HO2
both of which generate an active radical.
(47)
(72)
The most important decay route for H atoms now becomes
H
+ CHI = CH, + H2
(-21)
and the chain-branching reaction (3) becomes greatly reduced in importance. Hence there is no rapid acceleration in the rate of decomposition of ethylene. Most flux of decomposing C2H4 now goes via reaction 72. The overall decomposition of C2H4 is slowed down because of reaction -52
2CH3 = C2H6
(-52)
which terminates CH, radicals. There is thus a much slower formation of vinyl radicals and hence oxides of carbon when excess methane is present in the reactants. Increased concentration of CH3 radicals associated with the excess methane leads to a significant increase in propene formation via
CH3 + C2H4 = i-C3H7 i-C3H7 = C3H6 + H
(112) (-113)
It might also be thought that the increased [CH,] when methane is present in excess could bring into prominence reaction 29, the reaction between methyl radicals and O2 to form CH30 and 0 atoms. This reaction, however, has an appreciable activation energy, and even a t elevated methyl concentrations, insufficient reacting flux passes through (29). The above mechanism is in accord with the experimental observation of the effects of methane: the inhibition of partial oxidation and the change in selectivities to lower CO selectivity and to higher C3 selectivity. Sensitivity analyses can also give information about the greater low-temperature reactivity of ethylene when compared with ethane under similar partial oxidation conditions (compare Figures 5 and 1). As mentioned above, it was necessary to postulate reaction 107 to model the initiation of partial oxidation of ethylene a t temperatures around 960-1000 K. This reaction provides an essentially chdin-branchingroute (since HCO readily decomposes into CO + H) a t low temperatures. Although this reaction is important at low extents of conversion of C2H4 (
