Ind. Eng. Chem. Res. 1992,31, 75-80 Chem. Prod. Res. Dev. 1977,16,171-177. Hiatt, R. R. Twenty Years of Peroxide Chemistry. In Frontiers of Free Radical Chemistry; Pryor, W. A,, Ed.; Academic Press: New York, 1980; pp 225-230. Hombek, R.; Heenan, D. F.; Januszkiewicz, K. R.; Sulek, H. H. Oxidation of Aluminum Cold Rolling Base Oils. Lubr. Eng. 1989,45, 56-64. Hsu, S.M.; Ku, C. S.; Pei, P. T. Oxidative Degradation Mechanisms of Lubricants. In Aspects of Lubricant Oxidation; Stadtmiller, W. H., Smith, A. N., Eds.; ASTM Special Technical Publication 916;American Society for Testing and Materials: Philadelphia, 1986;pp 27-48. Jensen, R. K., Ford Motor Co. Scientific Research Laboratory. Personal communication, 1989. Jensen, R. K.; Korcek, S.; Mahoney, L. R.; Zinbo, M. Liquid-Phase Autoxidation of Organic Compounds at Elevated Temperatures. 1. The Stirred Flow Reactor Technique and Analysis of Primary Producta from n-Hexadecane Autoxidation at 120-180 OC. J. Am. Chem. SOC.1979,101,7574-7584. Jensen, R. K.; Korcek, S.; Mahoney, L. R.; Zinbo, M. Liquid-Phase Autoxidation of Organic Compounds at Elevated Temperatures. 2. Kinetics and Mechanisms of the Formation of Cleavage Products in n-Hexadecane Autoxidation. J. Am. Chem. SOC.1981,103, 1742-1749. Jensen, R. K.; Korcek, S.; Zinbo, M.; Johnson, M. D. Initiation in Hydrocarbon Autoxidation at Elevated Temperatures. Znt. J. Chem. Kinet. 1990,22,1095-1107. Jette, S. J.; Shaffer, D. L. Polymerization of Sunflower Oil Diesel Fuel: Copper Catalysis in Contaminated Lubrication Oil. Znd. Eng. Chem. Res. 1988,27,47-50. Jones, W. R. A Review of Liquid Lubricant ThermallOxidative Degradation; NASA Technical Memorandum 83465; NASA: Washington, DC, 1983. Korcek, S.; Jensen, R. K. Relation between Base Oil Composition and Oxidation Stabilitv at Increased Temrmatures. ASLE Trans. 1976,19,83-94. Korcek, S.; Jensen, R. K., Ford Motor Co. Scientific Research Laboratorv. Unpublished results, 1990. Korcek, S.; Cheiier, J. H. B.; Howard, J. A,; Ingold, K. U. Absolute Rate Constants for Hydrocarbon Autoxidation. XXI. Activation Energies for Propagation and the Correlation of Propagation Rate
75
Constants with Carbon-HydrogenBond Strengths. Can. J.Chem. 1972,50,2285-2297. Korcek, S.; Johnson, M. D.; Jensen, R. K.; Zinbo, M. Determination of the High-Temperature Antioxidant Capability of Lubricants and Lubricant Components. Znd. Eng. Chem. Prod. Res. Deu. 1986,25,621-627. Lee, K. W.; Choi, M. J.; Kim, S. B.; Choi, C. S. Liquid-Phase Oxidation of n-Dodecane in the Presence of Boric Acid. Znd. Eng. Chem. Res. 1987,26,1951-1955. Mill, T.; Mayo, F.; Richardson, H.; Irwin, K.; Allara, D. L. Gas-and Liquid-Phase Oxidations of n-Butane. J. Am. Chem. SOC.1972, 94,6802-6811. Naidu, S. K.; Klaus, E. E.; Duda, J. L. Evaluation of Liquid Phase Oxidation Products of Ester and Mineral Oil Lubricants. Znd. Eng. Chem. Prod. Res. Dev. 1984,23,613-619. Naidu, S. K.; Klaus, E. E.; Duda, J. L. Kinetic Model for HighTemperature Oxidation of Lubricants. Znd. Eng. Chem. Prod. Res. Dev. 1986,25,596-603. Reddy, K. T.; Cernansky, N. P.; Cohen, R. S. Modified Reaction Mechanism of Aerated n-Dodecane Liquid Flowing Over Heated Metal Tubes. Energy Fuels 1988,2,205-213. Spearot, J. A. Viscosity of Severely Oxidized Engine Oil. Znd. Eng. Chem. Prod. Res. Dev. 1974,13,259-267. Steiner, E. C.;Blau, G. E.; Agin, G. L. Introductory Guide to SimuSolv Modeling and Simulation Software; The Dow Chemical Co.: Midland, MI, 1986. Suresh, A. K.; Sridhar, T.; Potter, 0. E. Autocatalytic Oxidation of Cyclohexane-Modeling Reaction Kinetics. AZChE J. 1988,34, 69-80. Tseregounis, S. I.; Spearot, J. A.; Kite, D. J. Formation of Deposits from Thin Films of Mineral Oil Base Stocks on Cast Iron. Znd. Eng. Chem. Res. 1987,26,886-894. Van Sickle, D. E. Oxidation of 2,4,6-Trimethylheptane. J. Org. Chem. 1972,37,755-760. Van Sickle, D. E.; Mill, T.; Mayo, F. R.; Richardson, H.; Gould, C. W. Intramolecular Propagation in the Oxidation of n-Alkanes. Autoxidation of n-Pentane and n-Octane. J. Org. Chem. 1973,38, 4435-4440. Received for review June 21, 1991 Accepted September 19, 1991
Experimental Study of the Pyrolysis and Oxidative Pyrolysis of C2HsCl Ramazan Yildirim and Selim M. Senkan* Department of Chemical Engineering, University of California, Los Angeles, California 90024
The gas-phase pyrolysis and oxidative pyrolysis of chloroethane (C2H5C1)were studied experimentally in a 2.1-cm4.d. quartz flow reactor operating at 1-atm pressure, at temperatures in the range 605-630 O C , and at reaction times in the range 0.7-1.2 s. Pyrolysis products in the absence of O2were primarily C2H4 and HC1, with trace levels of C2H3Cland C2He In the presence of 02,however, the selectivity toward C2H3Clformation increased considerably and reached levels as high as 30% based on the amount of C2H5C1reacted. The presence of O2also increased the rate of decomposition of C2H5C1.
Introduction The chlorine-catalyzed oxidative-pyrolysis (CCOP) process was recently developed as a practical method to convert methane, the major component in natural gas, into more valuable C2 and higher molecular weight hydrocarbons (Senkan, 1987a). According to the CCOP process, CHI is chlorinated first to form chlorinated methanes, primarily monochloromethane (CM), followed by the oxidative pyolysis of CM to form Cz products such as CzH4, C2H2,C2&, CzH3C1,synthesis gas (CO and H2),and HC1 in the second step (Senkan, 1987b). The HC1 produced can either be converted back to chlorine via the well-known Deacon reaction and recycled, or can be used to oxy-
* To whom correspondence should be addressed.
chlorinate methane to form CMs, thus completing the catalytic cycle for chlorine (Senkan, 1987b). A detailed chemical kinetic mechanism for the oxidative pyrolysis of CH&l was also developed (Karra and Senkan, 1988). Since natural gas frequently contains substantial quantities of C,H, (typicdy 530% by volume), its impact on the CCOP process is of practical interest. In particular, the development of a better understanding of the oxidative pyrolysis of CzH&l, which forms as the primary chlorination product of C2H6,is important to better assess the nature of the product distributions in the CCOP process utilizing natural gas directly without separation. It should be noted that, although the product HC1 is corrosive, it can also be easily removed from the reaction products by water scrubbhg, leaving behind a gaseous mixture that can be treated by conventional separation equipment. That
0888-588519212631-00~5$03.00/0 0 1992 American Chemical Society
76 Ind. Eng. Chem. Res., Vol. 31, No. 1, 1992 Sampling Probe
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Gas Chromatograph Gas Cylinders
Figure 1. Schematics of the experimental system.
is, only the chemical reactor has to withstand corrosive gases at high temperature. In prior investigations,the pyrolysis of C2H5C1has been studied in static and flow reactors and in shock tubes (see, for example, Barton and Howlett (1949), Tsang (1964), Holbrook and Marsh (1967),and Gorshkov et al. (1986)). In all these studies, the stoichiometryof the overall process was describable by the following unimolecular decomposition reaction: C2H5C1= CzH4 + HC1 (AHr = 18 kcal/mol) (1) Oxidative pyrolysis of C2H5C1was studied first by Gorshkov et al. (1983a,b),who reported the formation of substantial levels of CzHBCltogether with CzH4 However, these investigators did not provide any quantitative details; in particular the effects of operating conditions on conversion and selectivity were not reported. Recently, Fisher et al. (1989) studied the combustion of C2H5Clby injecting it into the postflame region of a turbulent flow propaneair combustor under fuel-lean conditions. They also observed the formation of significant levels of C2H3Cltogether with C2H4, consistent with the earlier results of Gorshkov et al. (1983a). Fisher et al. also extended the reaction mechanism developed in our laboratories for CH3Cl (Karra et al., 1988) to CzH5C1combustion, and had a limited success in accounting for their experimental measurements. The results provided by Gorshkov et al. (1983a) and Fisher et al. (1989) are intriguing as they suggest that the following overall dehydrogenation reaction CzH5Cl = CzH&l+ H2 (AHr = 32 kcal/mol) (2) is also important in the oxidative decompition of C2H,Cl. Since the formation of CzH3Clhas consequences for the
mechanism of oxidation of chlorinated hydrocarbons and on the CCOP process, the effects of O2 on the pyrolysis of C2H5C1must be better understood and quantified. In this paper, we report on the results of a preliminary experimental study on the oxidative pyrolysis of CzH5C1 conducted in a flow reactor. The effects of temperature, mixture composition, and reaction time on reactant conversion and product selectivity were determined. A detailed chemical kinetic mechanism for the oxidation and pyrolysis of C2H6C1is presently under development, and the results of this investigation will be communicated at a later date.
Experimental System The fast flow reactor facility used to study the pyrolysis and oxidative pyrolysis of CzH5Clis shown in Figure 1. The experiments were conducted under atmospheric pressure in a 2.1-cm-i.d. quartz reactor which was about 100 cm long. At this diameter and at the temperatures and pressures involved, the reaction proce%ses are expected to be dominated by gas-phase kinetics, with minor contributions from surface induced reactions. The reactor was placed inside a three-zone Lindbergh furnace. The CzH5Cl and CZH5C1/Ozmixtures were injected directly into preheated nitrogen carrier gas using an air-cooled probe through radially directed injection holes. The upstream 15 cm of the furnace was used to preheat the nitrogen carrier gas. Gas flows were established by the combined use of rotameters and needle valves. The needle valves were operated under critical flow conditions in order to establish uncoupled flow rates. Mean gas flow velocities in the
Ind. Eng. Chem. Res., Vol. 31, No. 1, 1992 77 Table I. Experimental Conditions Investigated mixture composition, mol % expt C2HSCl O2 N2 temp, "C A 1.02 98.98 606-630 B 1.02 4.09 94.89 606-630 C 1.02 0.0-7.73 98.98-91.25 620 D 1.02 4.09 94.89 617 E 1.02 98.98 617
residence time, s 0.93" 0.93" 0.93 0.70-1.17 0.70-1.17
" Mean residence time (see text). reactor were in the range 1-5 m/s during the experiments suggesting laminar flow conditions were present. However, these conditions only result in an uncertainty of about 510% in the concentration profile measurements based on plug flow considerations (Cathonnet et al., 1981). This is due to high dilutions of the reactants and high temperatures present in the system which promotes diffusion, as well as due to the high aspect ratio of the reactor, which was in excess of 15 in our studies. Gas sampling was achieved by withdrawing gases through a cooled quartz sampling probe positioned centrally at the downstream end of the reaction zone. Gases withdrawn passed through the sample loop of a gas chromatograph (GC) (Hewlett-Packard 5880A) which was equipped with two packed columns (Porapak N, 6 ft, and molecular sieve 5A, 6 ft) and a thermal conductivity detector operated with helium carrier gas. n e samples were introduced into the GC by computer-controlled valve injection. Calibration gases were acquired from Matheson Co. (Joliet, IL), and standard procedures of analysis using GC were employed throughout the experiments to convert measurements into absolute mole fractions. We estimate that the absolute mole fractions determined in the present studies should be accurate within *5%. Temperatures were measured by Chromel/Alumel thermocouples. These measurements indicated that the reactor was nearly isothermal, i.e., within f 3 OC, in the central portion of the reactor in all experiments.
Results and Discussion Three sets of experiments were conducted to ascertain the effects of temperature, oxygen concentration, and residence time on the rate of decomposition of CzH5Cland product selectivities. In Table I the specific experimental conditions investigated are summarized. In all experiments, the concentration of CzH5C1was kept at about 1% to preserve the near-isothermal conditions in the system. Studies using higher CzH5C1concentrations lead to substantial cooling of the reaction mixture due to the endothermicity of reaction 1. Nitrogen was used as the diluent, and when Oz was introduced an equivalent amount of Nz was removed from the mixture to keep the CzH5Clmole fraction constant at about 1%. Carbon balances were better than 98%, with no apparent formation of soot or coke under all conditions investigated. Experiment A was conducted to establish the effects of temperature on the rate of CzH5C1pyrolysis and on product distribution, and to better assess the impact of Oz in subsequent experiments. In Figure 2, the mole percents for CzH5Cl,CZH4, CzH3C1,and C,H, are presented as a function of reador temperature measured at a nominal reaction time of about 0.93 s. In this and subsequent figures lines have been drawn through experimental data points (indicated by symbols) to show trends. The data in Figure 2 were acquired by sampling gases when both the injection and sampling probes were spatially fixed and by changing the temperature of the furnace. In addition, the feed gas composition and flow rates were maintained
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Figure 2. Effect of temperature on C2H5Cldecomposition and product distribution in the absence of O2at a nominal residence time of 0.93 s.
the same. Consequently, the data points shown in Figure 2 do not precisely correspond to identical residence times because of differences in gas density and number of moles caused by differences in temperature and different extents of reaction, respectively. However, since differences in temperatures between data points were at most 30 K (630-605 "C), reaction times would be different by at most 3.5% because of changes in gas density, Le., (30 K/878 K) X 100 = 3.5%). Similarly, the increase in gas flow rate caused by the increase in number of moles would be at most 2% based on the complete conversion of CzH5C1into CzH4, CzH3C1,Hz, and HC1. Both these errors are well within the normal experimental error limits of f10% expected in these types of studies. As evident from Figure 2, the major decomposition product of CzH5C1is CzH4 (selectivity >95%). In addition, some CzH3Cl, with CzH3C1/CzH4ratios in the range 0.03-0.05, and trace quantities of CzH6 also did form. These results are in complete agreement with those reported by earlier investigators (Barton and Howlett, 1949; Holbrook and Marsh, 1967; Tsang, 1964; Gorshkov et al., 1986) and support that the unimolecular decomposition of CzH5C1(i.e., reaction 1) is the primary reaction route in the absence of O2 The data for CzH5C1decomposition were also analyzed by using the following separate rate expressions to ascertain the presence of secondary reactions: -d[ CZH&1] /dt = k3 [CZH5C11 (3) d[CzH,I/dt = ~ ~ ( [ C Z H ~ - C[CzH41) ~IO (4) where the brackets represent concentrations and [C2H5ClIo is the initial concentration of CzH5C1. The rate expressions 3 and 4 describe the total rate of disappearance of C2H5C1 and the rate of formation of CzH4, respectively. In the absence of other reaction pathways, i.e., reaction 1being the only reaction, the rate expressions 3 and 4 would have been equivalent. However, this clearly was not the case as evident by the formation of products other than C2H4 shown in Figure 2. In Figure 3 the Arrhenius plots of two-parameter rate coefficients k3 (k3 = k30 exp(-E,/RT)) and k4 (k4 = k40 exp(-E,/RT)) are presented. From this figure activation energies can be determined to be 44 and 61 kcal/mol for
78 Ind. Eng. Chem. Res., Vol. 31, No. 1, 1992 -0.700
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1
2
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6
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Figure 5. Effect of O2concentration on C2H6C1decomposition and product distribution at 620 OC and a nominal residence time of 0.93 9.
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620
Temperature
625
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Figure 4. Effect of temperature on CIHBCldecomposition and product distribution in the presence of 4.09% O2at a nominal residence time of 0.93 s.
k3 and k4, respectively. The activation energy of 61 kcal/mol for k4 is in reasonable agreement with the value of 56 kcal/mol reported for reaction 1 (Holbrook and Marsh, 1967; Tsang, 1964), indicating that the unimolecular decomposition of C2H5C1was the major pathway for CzH4 production in our experiments. The activation energy of 44 kcal/mol for k3,on the other hand, is indicative of the presence of additional reaction pathways for CzH5C1 decomposition, since carbon balances were excellent in all the experiments. These additional reactions are likely to involve the attack of free radicals, characterized by lower activation energies, thereby lowering the apparent activation energy for the rate of disappearance of C2H5C1. The formation of trace products, such as C2H3C1and CzH6, support the presence of free-radical reactions. In Figure 4, the effect of temperature on C2H5Clpyrolysis is shown when 4.09% O2 was present in the mixture (experiment set B). As seen from this figure, the presence of O2 in the system significantly affects the rate and
mechanism of CzH5C1pyrolysis. Most importantly, O2 accelerates C2H5C1pyrolysis and results in a dramatic (10-fold) increase in the production of C2H3C1. These results support the findings of Gorshkov et al. (1983a,b) and Fisher et al. (1989). Additional producta include trace levels of C&, which was also observed in pure C2H5Cl pyrolysis, as well as some CO and C02,all of which increase with increasing temperature. It is particularly significant to note that in this and most other experiments very little consumption of O2occurs, well below the limits of accuracy of gas chromatographic measurements. Consequently, O2 concentrations are not explicitly presented in any of the figures. These results, however, are consistent with the presence of large concentrations of O2 in the system and the low levels of CO and C02formed in the experiments. As seen in Figure 4, C02 concentration exponentially increases with increasing temperature, indicative of the onset of flame reactions. Further increases in temperature indeed have shown the establishment of flames in the reactor, concomitant with the complete conversion of all the hydrocarbons to CO and C02together with a significant decrease in O2concentration. Flame formation clearly establishes the upper limit in operating temperature in the oxidative pyrolysis of C2H5C1. From Figure 4, the C2H3C1/C2H4 ratio can be shown to increase from 0.26 to 0.48 when the temperature is increased from 605 to 630 "C. This result is totidy expected in view of the increasing importance of free-radical reactions with increasing temperature. In Figure 5, the effects of O2 concentration on C2H5C1 conversion and C2H4 and C2H,C1production are presented at a reaction temperature of about 620 OC and at a nominal residence time of about 0.93 s (experiment set C). From this figure it can be seen that the CzH3C1/C2H4ratio decreases from 0.39 to 0.21 upon increasing the O2 concentration from 1.6 to 7.73%. These results indicate an order of magnitude increase in CzH3Cl concentrations when compared to the levels observed in pure C2H5C1pyrolysis. As seen in Figure 5, the specific effect of O2 in promoting the formation of C2H3C1is highest at low O2 concentrations. This is an expected result since O2is consumed very little in our studies. In the present set of experiments an increase in Oz concentration beyond 2% had no effect on C2H3C1production. These results, therefore, suggest the
Ind. Eng. Chem. Res., Vol. 31, No. 1, 1992 79 1 .o
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Figure 6. Effect of residence time on C2H5Cl conversion and product distribution in the presence of O2 at 617 OC.
Figure 7. Effect of residence time on C2H5C1conversion and product distribution in the absence of O2 at 617 OC.
use of lowest O2 concentrations consistent with the specific mixture composition to maximize C2H3Clproduction. Low O2 levels would be desirable to minimize the dilution of
increase in C2H3Clproduction when compared to the pyrolysis of pure C2H5C1. Although C2H4was the major product in both the absence and presence of 02, C2H3Cl/C2H4ratios reached levels as high as 0.39 in the latter experiments. The specific effect of O2in promoting C2H3C1production was largest at low concentrations, consistent with the low conversions of this reactant. These results suggest that oxidative pyrolysis of chlorinated hydrocarbons may be exploited as a manufacturing process to synthesize a broader range of products from abundant natural feedstocks.
products and the potential for the formation of destructive flames. As shown in Figure 5, the effects of O2 on CO and C02 are also interesting. The concentration profiles of these species exhibit first an increasing and then a decreasing trend with increase in O2 levels. Such a behavior for CO is expected since it forms as an intermediate in C2H5Cl combustion. The increase in C02production with increase in O2 concentration is also expected since it is produced from the oxidation of CO. However, the subsequent decrease in C02 levels with increasing O2 concentration is surprising. High O2 concentrations apparently decrease C2H,C1 conversion and thus adversely affects CO production, thereby reducing C02production. Uncertainties in the experimental data can also be the reason for these nonuniform trends. High O2 concentrations also had an adverse effect on C2HG.However, this result is totally expected on the basis of our current understanding of the chemical kinetics of combustion of C2H6and other hydrocarbons (Westbrook and Dryer, 1984). In Figure 6, the effects of residence time on C2H5C1 conversion and product distributions are presented for a prereaction composition of 1.02% C2H5C1and 4.09% O2 at a temperature of 617 OC (experiment set D). A similar plot for the pyrolysis of C2H5Clalone is presented in Figure 7. As evident from Figure 6, the C2H5C1concentration decreases and those for the products increase more rapidly with increasing residence time, indicative of the build-up of free radicals in the system. It is particularly significant to note the expontential increase in both CO and C 0 2 concentrations, which suggest the onset of rapid combustion reactions. In contrast, the concentrations change linearly with residence time in pure C2H5C1pyrolysis (see Figure 7 again),and this is in agreement with the primarily nonchain character of C2H5C1decomposition via reaction 1. In conclusion, the presence of O2was shown to accelerate the rate of decomposition of C2H5C1and led to a 10-fold
Acknowledgment This research was supported, in part, by funds from the U S . Environmental Protection Agency, Grant No. R815136-01-0. We also thank S. Sethuraman for his help during the experiments.
Literature Cited Barton, D. H. R.; Howlett, K. E. The Kinetics of the Dehydrochlorination of Substituted Hydrocarbons. Part III. The Mechanism of the Thermal Decompositions of Ethyl Chloride and of 1:l-Dichloroethane. J . Chem. SOC.1949,165. Cathonnet, M.; Boettner, J. C.; James, H. Experimental Study and Numerical Modeling of High Temperature Oxidation of Propane and n-Butane. Eighteenth Symposium (International) on Combustion; The Combustion Institute Pittsburgh, 1981;p 903. Fisher, E. M.; Koshland, C. P.; Hall, M. J.; Sawyer, R. F.; Lucae, D. Experimental and Numerical Study of the Thermal Destruction of C2H5Cl. Twenty-Third Symposium (International) on Combustion; 1990; p 895. Gorahkov, S. V.; Kolbanovskii,Y. A.; Razovskii, A. Y.; Chernyak, N. Y. Chain Dehydrochlorination of Chloroethane. Kinet. Katal. 1983a,24 (l), 253. Gorshkov, S. V.; Kolbanovskii, Y. A.; Razovsku, A. Y.; Chernyak, N. Y. High-Temperature Transformations of Chloroethane in the Presence of Oxygen. Kinet. Katal. 1983b,24 (l),254. Gorahkov, S. V.;Kolbanovskii, Y. A.; Razovskii, A. Y.;Chernyak, N. Y. High-Temperature Conversions of Alkyl Halides. Dehydrochlorination of Chloroethane. Kinet. Katal. 1986,27 (21,263. Holbrook, K. A.; Marsh, A. R. W.Unimolecular Gas-Phase Pyrolysis of Ethyl Chloride. Trans. Faraday SOC.1967,63,643. Karra, S. B.; Senkan, S. M. A Detailed Chemical Kinetic Mechanism for the Pyrolysis and Oxidation of CH3C1. Ind. Eng. Chem. Res. 1988,27,1163.
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Karra, S. B.; Gutman, D.; Senkan, S. M. A Detailed Chemical Kinetic Mechanism for the Combustion of CH3Cl. Combust. Sci. Technol. 1988,60,45. Park, J. Y.; Heaven, M. C.; Gutman, D. Kinetics and Mechanism of the Reactions of Vinyl Radical with Molecular Oxygen. Chem. Phys. Lett. 1984, 104, 469. Senkan, S. M. Production of Higher Molecular Weight Hydrocarbons from Methane. U.S.Patent 4,714,796, 1987a. Senkan, S. M. Conversion of methane into higher molecular weight
hydrocarbons by the Chlorine-Catalyzed Oxidative-Pyrolysis (CCOP) process. Chem. Eng. Prog. 1987b, 12, 58. Tsang, W. Thermal Decomposition of Some Alkyl-Halides by a Shock-Tube Method. J . Chem. Phys. 1964,41,2487. Westbrook, C. K.; Dryer, F. A. Chemical Kinetic Modeling of Hydrocarbon Combustion. Prog. Energy Combust. Sci. 1984,10,1.
Received for review June 14, 1991 Accepted September 9, 1991
Pretreatment of a Vanadia-Titania Catalyst for Partial Oxidation of o -Xylene under Industrial Conditions Valentin A. Nikolov and Asen I. Anastasov* Institute of Chemical Engineering, Bulgarian Academy of Sciences, Acad. G. Bontcheu str., El. 103, Sofia 1113, Bulgaria
The pretreatment of a vanadia-titania catalyst for oxidation of o-xylene realized during the first 50 days of its life is investigated. An industrial multitubular fixed bed reactor is studied. The aim is to reveal the meaning of the pretreatment, which requires well-defined technological conditions to be maintained. A hypothesis with respect to the pretreatment is assumed and verified experimentally. According to it, the high temperatures in the front zone of the bed during the pretreatment lead to an artificial partial deactivation of the catalyst in this zone. The hypothesis is proved also by additional experiments carried out in a pilot nonisothermal, nonadiabatic integral reactor identical with a single tube of the industrial apparatus and by some physicochemical analyses. Temperature profiles obtained in the industrial reactor are simulated by a two-dimensional heterogeneous model taking into consideration the reduced catalyst activity in the front zone. 1. Introduction Phthalic anhydride is one of the most important products in modern industrial organic synthesis. Owing to its high reactivity, it is widely used in the chemical industry and its production is increasing continuously (Hoffmanand Riddle, 1988). The vapor-phase oxidation in air of o-xylene over a fixed bed of V205-Ti02catalyst promoted mainly by antimony, potassium, rubidium, and phosphorus oxides is used in more than 70% of the total production of phthalic anhydride (Wainwright and Foster, 1979). As the process is highly exothermic, it is carried out mainly in multitubular reactors (De Virgilis and Gerunda, 1982; Cybulski, 1981). A conception considering nonuniform catalyst activity along the bed as more suitable for the reactor performance is widespread in the literature. In order to achieve an appropriate temperature profile, Caldwell and Calderbank (1969) suggest to reduce the activity in the front zone of the bed and propose the catalyst in this region to be diluted by inert pellets. To accomplish an activity profiling of the bed, Pirkle and Wachs (1987),simulating the proceas of o-xylene oxidation by a two-dimensional homogeneous model, use an activity factor F. It is a function of the bed length and takes values less than 1.0 in the front part of the reactor. However, it is well-known that an industrial reactor is always loaded with a catalyst of uniform activity. Regardless of the many papers devoted to o-xylene oxidation (Smith and Carberry, 1975; Amouroux et al., 1976; Wolfahrt and Hoffman, 1979; Soria Lopez et al., 1981; Kerschenbaum and Lopez-Isunza, 1982; Bobrov et al., 19861, there is no available information on how the catalyst activity changes along the bed during the pretreatment of the catalyst. It is also known that neither the catalyst nor the reaction components along the reactor length can be sampled and analyzed during the operation of an industrial
* To whom correspondence should be addressed.
Table I. Catalyst Properties specific surf., m2/g specific vol of the pores (Hg), (m3/kg) X bed density, (kg/m3) X vanadium/titanium ratio by w t thickness of catal layer, mm support: diam of porcelain spheres, mm
lo6
10.0 269.0 1.5 0.055 0.1 6.0
multitubular reactor. On the other hand, it is possible to evaluate approximately the state of the reaction and catalyst by measuring the temperature profiles in the tubes. The reactor performance depends considerably on the character of these profiles, which are a function of the inlet o-xylene concentration, flow rate, and coolant temperature. A t the same time, the parameters mentioned change significantly in the first 2 months of the catalyst life. It seems that these variations are especially important for the activation of the V205-Ti02catalyst and achievement of a suitable activity profile reflecting a long We of the catalyst. Both the behavior and state of the catalyst in the first days of its use have not been discussed in the literature till now. Bearing this in mind, it is interesting to study the formation of the catalyst structure during this period, as well as the effect of maintenance-definite values of inlet oxylene concentration, flow rate, and coolant temperature on the character of this structure. An additional goal is to establish the dependence between the catalyst activity and the temperature profiles obtained in the reador during the pretreatment. 2. Experimental Section An industrial reactor containing 8920 contact tubes of 3250-mm length and 25-mm inner diameter was studied (Figure 1). The heat of the reactions is removed by a coolant, an eutectic mixture of KN03 and NaNOz (ratio 57:43). The industrial reactor is charged with a V205-Ti02 catalyst promoted by Pz05. The characteristics of the catalyst are given in Table I.
0888-5885/92/2631-0080$03.~~/0 0 1992 American Chemical Society
