On the Gas-Phase Chlorination of Ethane - Industrial & Engineering

Iain W. Sutherland , Neil G. Hamilton , Christopher C. Dudman , Peter Jones , David Lennon , John M. Winfield. Applied Catalysis A: General 2011 399 (...
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Ind. Eng. Chem. Res. 2001, 40, 2226-2235

On the Gas-Phase Chlorination of Ethane Ivar M. Dahl,† Elisabeth M. Myhrvold,† Unni Olsbye,*,†,§ Friedemann Rohr,† Odd A. Rokstad,‡ and Ole Swang† SINTEF Applied Chemistry, Department of Hydrocarbon Process Chemistry, P.O. Box 124 Blindern, N-0314 Oslo, Norway, and Department of Catalysis and Kinetics, Sem Sælandsv. 2A, N-7465 Trondheim, Norway

The gas-phase chlorination of ethane has been studied experimentally and by simulation of the corresponding gas-phase reaction scheme. In addition, some key reactions in ethane chlorination have been studied in detail by quantum chemical methods. Experimental results indicate that the chlorination reactivity of C2H6-xClx decreases with increasing x and, further, that 1,1dichloroethane is the preferred dichlorination product. The preference for 1,1-dichloroethane over 1,2-dichloroethane at higher temperatures has been ascribed to the 1,2-dichloroethane precursor radical decomposing into ethene and a chlorine atom. General agreement was found between experiments, kinetic simulation (with rate constants from the NIST database), and quantum chemical calculations for ethane chlorination. 1. Introduction Chloroethene (vinyl chloride, VCM) is a valuable petrochemical product. Traditional production routes (from ethane) involve steam cracking of ethane, followed by parallel ethene chlorination and oxychlorination, and finally dehydrochlorination of the formed 1,2-dichloroethane (ethyl dichloride) to chloroethene, i.e.

C2H6 f C2H4 + H2 C2H4 + Cl2 f C2H4Cl2

∆H298K ) 32.1 kcal/mol

(i)

∆H298K ) -43.3 kcal/mol (ii)

C2H4 + 2 HCl + 0.5 O2 f C2H4Cl2 + H2O ∆H298K ) -57.8 kcal/mol (iii) C2H4Cl2 f C2H3Cl + HCl ∆H298K ) 16.2 kcal/mol (iv) The ethane cracking reaction (i) is equilibrium-limited and is carried out at temperatures in excess of 900 °C. The reaction is strongly endothermic. Reactions ii and iii are carried out at temperatures below 300 °C. Both reactions are strongly exothermic. Reaction iv is strongly endothermic and is typically carried out at 500°C, because of thermodynamic limitations. The total process

2C2H6 + Cl2 + 1.5O2 f 2C2H3Cl + 3H2O ∆H298K ) -122.2 kcal/mol (v) is exothermal. However, because of the low-temperature exotherms and high-temperature endotherms, the energy efficiency of the conventional process is low. Recently, EVC patented1 an ethane oxychlorination process that gives chloroethene from ethane in one reactor, using HCl, O2, and Cl2 as co-reagents. The reaction is carried out over a CuCl-based catalyst at * Author to whom correspondence should be addressed. † Department of Hydrocarbon Process Chemistry. ‡ Department of Catalysis and Kinetics. § Present address: Nordox Industries, P.O Box 6639 Etterstad, N-0607 Oslo, Norway.

approximately 480-550 °C and atmospheric pressure. Over-chlorinated paraffins are dehydrochlorinated and hydrogenated (together with over-chlorinated olefins) in separate reactors. The hydrogenated product streams are led back to the oxychlorination reactor. The economic potential of this simplified, energy-optimized process has very recently led to the construction of a pilot plant, as well as the announcement of a commercialization-ready process.2 An oxygen-free ethane chlorination process would have significant advantages over an oxychlorination process, from both an environmental and an engineering point of view. The two reactions from ethane to 1,2dichloroethane

C2H6 + Cl2 f C2H5Cl + HCl ∆H298K ) -28.5 kcal/mol (vi) C2H5Cl + Cl2 f ClCH2CH2Cl + HCl ∆H298K ) -26.4 kcal/mol (vii) are both exothermal, which constitutes an advantage compared to the scheme described above. Further, the risk of dioxin emission would be greatly diminished, and the amount of highly corrosive HCl-H2O mixtures would be drastically reduced compared to the oxychlorination process. Equilibrium calculations of ethane chlorination, using HYSYS software,3 indicate that maximum yields of ethene and chloroethene can be obtained from an ethane:chlorine ) 1:1 feed at 500 °C and 1 atm., yielding 83% ethene and 7% chloroethene. The purpose of the present work is to evaluate the potential for direct chlorination of ethane. The reaction network involved is illustrated in Figure 1. Along the lower line in that figure are the directly chlorinesubstituted alkanes. They can be reversibly transferred by dehydrohalogenation to the second line of alkenes. The chloroalkenes can again be reversibly dehydrohalogenated to the alkynes. The arrows pointing down and to the right represent addition reactions of chlorine to the alkenes. Arrows going down and to the left represent hydrogenation reactions of alkenes, which could be a method of returning overchlorinated products

10.1021/ie000850n CCC: $20.00 © 2001 American Chemical Society Published on Web 04/13/2001

Ind. Eng. Chem. Res., Vol. 40, No. 10, 2001 2227

Figure 1. Potential reaction products in ethane chlorination that can be used for monomer (chloroethene or ethene) production. Substances marked with an asterisk can consist of two different isomers.

Figure 3. Reactor and insert.

Figure 2. Ethane chlorination test rig. Heated areas are inside dotted rectangles.

to monomer precursors. The most valuable products in this scheme are, of course, the monomers ethene and chloroethene. Previous work by Zaidman4 indicates that the gasphase chlorination of ethane has a high selectivity toward ethyl chloride at 350-400 °C. However, their work, as well as other investigations in this field,5 was limited in that experiments could be carried out only at full conversion of chlorine. The present experiments were carried out at incomplete chlorine conversion. Simulation of the gas-phase reaction scheme using literature data, as well as quantum chemical modeling, was performed in order to help understand the experimental observations. 2. Experimental Section 2.1. Chlorination Experiments. A schematic view of the ethane chlorination test rig is shown in Figure 2. Ethane, chlorine, and helium were fed through Bronkhorst mass flow controllers. Ethyl chloride (with a boiling point of 12 °C) was fed as a gas by means of a dual-channel peristaltic pump. An Isco pump was used to feed 1,1-dichloroethane and 1,2-dichloroethane in corresponding chlorination tests. To avoid flow pulsation, the evaporator was filled with small glass beads

and kept at a low temperature (40 °C). All gases were mixed before entering the reactor. The gas feed lines after the evaporator were heated to 80 °C in a Carlo Erba furnace. Propene was fed at the reactor outlet to react with chlorine in the reactor effluent. All effluent gas lines were insulated and heated to approximately 100 °C with heating wires to avoid condensation of chlorinated products. Gas analysis was performed using an on-line HP6890 gas chromatograph equipped with a Poraplot Q column (0.32 mm inner diameter, 25 m length) and a HP5973 mass selective detector. Helium was used as the carrier gas. The quartz reactor (30 cm length, 6 mm inner diameter) was placed in a tubular oven. A quartz thermowell with an outer diameter of 3 mm that contained a moveable thermocouple for measurements of the axial temperature profile in the reactor was placed inside the reactor, along the reactor axis (Figure 3). After the reactor was flushed with helium (600 mL/ min) at 350 °C, the temperature difference over the reaction zone was smaller than 5 °C. During the reaction, the temperature difference between the oven and the thermocouple in the center of the reaction zone was less than 20° at reaction temperatures e350 °C. The total gas flow through the reactor was 300 mL/ min (NTP), with 15 mL of ethane/min (or mono- or dichlorinated ethane), 15 or 20 mL of chlorine/min, and 270 mL of helium/min as the standard conditions. Propene (20 mL/min) was added to the reactor effluent as mentioned above. Different conversion levels were achieved by varying the temperature between 100 and 450 °C. Experiments with varying reactor volumes and reactant flows were carried out at 300 and 350 °C. The reactor volume could be reduced by inserting a glass rod in place of the thermocouple well. Analyses were performed with the rod at several different heights in the reactor. The usefulness of propene as a chlorine trapping agent was studied by connecting a mass spectrometer (Fisons Sensorlab) to the effluent gas line. It was observed that residual chlorine in the reactor effluent was completely converted as soon as excess propene was added. GC-MS analysis showed that 1,2-dichloropropane was the dominant reaction product from propene

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Table 1. Rate Constantsa and Arrhenius Parameters for the Elementary Reactions Used in the Simulation Modelb,c rxn no.

reaction

preexponential factor Ad

activation energy Eq (kcal/mol)

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

Cl2 + M f 2Cl* +M Cl2 + H* f HCl + Cl* C2H6 + Cl* f C2H5* + HCl C2H5* + Cl2 f C2H5Cl + Cl* C2H5* f C2H4 + H* C2H5Cl f C2H4 + HCl C2H5Cl + Cl* f CH3CHCl* + HCl C2H5Cl + Cl* f CH2CH2Cl* + HCl CH3CHCl* + Cl2 f CH3CHCl2 + Cl* CH3CHCl* f C2H3Cl + H* CH3CHCl* + HCl f C2H5Cl + Cl* CH2CH2Cl* + Cl2 f CH2ClCH2Cl + Cl* CH2CH2Cl* f C2H3Cl + H* CH2CH2Cl* f C2H4 + Cl* C2H4 + Cl* f CH2CH2Cl* CH3CHCl2 + Cl* f CH2CHCl2* + HCl CH3CHCl2 + Cl* f CH3CCl2* + HCl CH3CCl2* + Cl2 f CH3CCl3 + Cl* CH3CCl2* f CH2CCl2 + H* CH2CHCl2* + Cl2 f CH2ClCHCl2 + Cl* CH2CHCl2* f C2H3Cl + Cl* CH2ClCH2Cl + Cl* f CH2ClCHCl* + HCl CH2ClCH2Cl f C2H3Cl* + HCl CH2ClCHCl* + Cl2 f CH2ClCCl2* + HCl CH2ClCHCl* + Cl2 f CH2ClCHCl2 + Cl* CH2ClCHCl* f C2H3Cl + Cl* 2Cl* + M f Cl2 + M C2H5* + Cl* f C2H5Cl C2H5* + Cl* f C2H4 + HCl 2C2H5* f C2H4 + C2H6 2C2H5* f C4H10 CH3CHCl* + Cl* f CH3CHCl2 CH2CH2Cl* + Cl* f CH2ClCH2Cl CH2CH2Cl* + C2H5* f C2H4 + C2H5Cl 2CH3CHCl* f products

1.00 × 1013 8.59 × 1013 5.43 × 1013 7.59 × 1012 4.31 × 1012 6.92 × 1013 8.61 × 1012 8.13 × 1012 2.63 × 1012 3.40 × 1013 3.16 × 1011 3.95 × 1012 1.40 × 1013 3.90 × 1013 1.87 × 1013 7.05 × 1012 6.14 × 1012 1.00 × 1013 3.20 × 1013 6.31 × 1011 3.20 × 1013 1.33 × 1013 6.46 × 1010 1.00 × 101 2.63 × 10-15 2.00 × 1014 2.24 × 1014 2.74 × 1014 7.23 × 1012 1.45 × 1012 1.15 × 1013 9.5 × 1013 1.00 × 1013 1.00 × 1013 1.15 × 1013

40.60 1.17 0.26 -0.24 37.20 57.80 0.92 1.73 0 43.20 8.00 0 42.10 21.70 0 2.78 1.27 2.99 45.70 0.89 26.10 1.58 46.99 0 -19.78 19.99 -1.80 0 0 0 0 0.24 2.99 2.99 0

a From k ) A exp(-Eq/RT). b Data from the NIST database.6 order), cm3/(s mol) (second-order), and cm6/(s mol2) (third-order).

and reactor effluents, although minor amounts of other chlorinated hydrocarbons were also formed (2-chloropropane from reactions of propene and HCl being most abundant). None of these products interfered with the ethane chlorination products in the GC-MS analysis, thus confirming the reliability of this analytical approach. At temperatures e350 °C, the rates of coke formation were negligible, as no visible coke was formed during a 2-h run. The C2 carbon mass balances were, accordingly, between 95 and 105%. 2.2. Simulations of Gas-Phase Kinetics. The gasphase chlorination of ethane was simulated by solving the coupled differential kinetic equations. The kinetic model employed is listed in Table 1; data are from the NIST database.6 The model was designed to be able to account for the full range of products detected in the experiments. The main criteria for the selection of elementary steps were the result of both qualitative analysis of the reaction products and chemical common sense. Reaction rate expressions for each elementary step were set up using reaction rate constants taken from the NIST database.6 In this database, there are some discrepancies regarding the rate constant for Cl2 dissociation (reaction 1 in Table 1). We selected the data from Schofield et al.7 on the basis that they provided the best fit to our measured results. For each chemical species involved in the reaction network, a mass balance was set up, and the resulting coupled first-order differential equations were solved using the commercial scientific software package MATLAB.8

c

Values for reactions in italics are estimated.

d

Units are s-1 (first-

The simulations were carried out for isobaric and isothermal conditions. This approximation is acceptable given the high dilution of the reactants in inert gas and the relatively shallow experimental temperature profiles for most of the experiments performed at low to moderate conversions. The simulations yield concentrations as functions of residence time for each species involved in the overall reaction. The composition of the feed gas and the reaction temperatures were chosen according to the experimental conditions, and the time span over which the reaction was simulated was chosen to match the experimental residence time. 2.3. Quantum Chemical Computational Details. Calculations based on density functional theory (DFT) were carried out using the DeFT program package.9 Gaussian all-electron basis sets of DZVP quality for hydrogen and TZVP for all other atoms, as well as the gradient-corrected functionals of Becke10 for exchange and of Perdew11 for correlation, were used. The MP2 calculations were done with the SPARTAN package,12 using a 6-31G(d,p) basis set. The Gaussian 98 program package13 was used for the CCSD(T) calculations with cc-pVTZ basis sets.14 The main methodological problem of the quantum chemical part of the present investigation arises from the need to calculate accurate energies for transition states with open-shell electronic structures, in which significant spin density resides in the moiety in which the making and breaking of bonds take place. In a very recent paper,15 Chuang et al. address this problem. In general, they recommend highly correlated ab initio methods for both geometries and

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Figure 4. Product yields as functions of temperature in C2H6-xClx chlorination (x ) 0,1,2). C2H6-xClx:Cl2:He ) 15:15:270 (mL/min, NTP) for x ) 1 or 2, and 15:20:270 for x ) 0. (a) C2H6, (b) C2H5Cl, (c) CH3CHCl2, and (d) CH2ClCH2Cl.

energies. However, for the transition state of the reaction between ethane and a hydroxyl radical, yielding ethyl radical and water, their study showed that the MP2 geometry differs only slightly from the geometries calculated at higher levels of theory. As this reaction closely resembles those under study here, it appears safe to use MP2 geometries in evaluating the CCSD(T) energies. Hence, for the reported DFT and MP2 results, geometries and energies have both been calculated at the same level of theory, whereas the CCSD(T) energies have been calculated for geometries optimized at the MP2 level. For the largest systems, viz. those containing three chlorine atoms or a combination of two chlorine atoms and an open shell, CCSD(T) calculations proved impossible to carry out with our available computational resources. In all cases, species with doublet spin multiplicity were described unrestrictedly, i.e., with separate sets of orbitals for R and β spin. Zero-point vibrational energy corrections were taken from vibrational spectra calculated at the UHF/6-31G(d,p) level and scaled by a factor of 0.89. These spectra also confirm that the right number of imaginary vibrational frequencies are at hand for each stationary point, viz. one for transition states and zero for minima. All degrees of freedom have been included in the geometry optimizations. For all systems, geometries were optimized at the DFT and MP2 levels without any assumptions regarding symmetry. 3. Results and Discussion 3.1. Experimental Results. Temperature Dependence. Product yields obtained during the chlorination of ethane, chloroethane, and 1,1- and 1,2-dichloroethane are shown as functions of temperature in Figure 4. Figure 4a shows that ethyl chloride is the dominant

product from ethane, whereas Figure 4b shows that 1,1dichloroethane is the dominant product from ethyl chloride. These results are in agreement with those obtained during ethane chlorination (Figure 4a), for which 1,1-dichloroethane is the major dichlorination product while only traces of 1,2-dichloroethane are observed. The decrease in conversion for ethane > chloroethane > 1,1-dichloroethane under comparable reaction conditions (Figures 4a-4d) shows that the reactivity decreases in the same order. The preference of 1,1-dichloroethane over (1,2-dichloroethane plus chloroethene) as the product from chloroethane chlorination also shows that this decreased reactivity is most marked at the carbon opposite to that already substituted with chlorine. A similar effect is seen in that 1,1,1trichloroethane is the preferred product from 1,1dichloroethane. Chlorination of the two dichlorinated ethanes (1,1and 1,2-dichloroethane) to trichloroethanes is even slower than chlorination of ethyl chloride. This can be seen from Figure 4c,d. Further, both dichlorinated reactants, and especially 1,2-dichloroethane, yield appreciable amounts of the vinyl chloride monomer (chloroethene) through the dehydrochlorination reaction. This leads to a high conversion of the dichlorinated ethanes in Figure 4, although only part of the conversion represents formation of a trichlorinated product. At 400 °C, the chlorination rate of 1,2-dichloroethane is about one-half of what is seen for 1,1-dichloroethane (Figure 4c,d), whereas the dehydrohalogenation rate is about 5 times higher for 1,2- than for 1,1-dichloroethane. Reactor Volume. From the results in Figure 4, 300350 °C was found to be a suitable temperature range for closer kinetic studies of the reactions. To determine the active volume of the reactor, reactor volume varia-

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Figure 5. Product partial pressure in the reactor effluent vs hydrocarbon partial pressure in the feed at 300 and 350 °C. Filled symbols, 300 °C; open symbols, 350 °C. The Cl2 flow is 15 mL/ min, and the balance up to 300 mL/min is He.

Figure 7. Simulated product yields as functions of reaction temperature.

Figure 6. Comparison of simulated (from the model in Table 1) and experimental ethane conversions as a function of reaction temperature.

tion tests were performed at 300 and 350 °C, using ethane as the reactant with chlorine. The active volume (with insert absent) was estimated to be 3.2 ( 0.3 mL, which is somewhat less than the actual free volume of the reactor. This difference is caused by the fact that the ends of the reactor are at a lower temperature than the middle and the “cold ends” do not contribute to the conversion. Reactant Flow Dependence. Results obtained by varying the (chlorinated) hydrocarbon partial pressures are shown in Figure 5. The chlorination rate appears to be close to first order in all three hydrocarbons. 3.2. Simulation Results. The simulations were carried out including the reactions listed in Table 1 with the parameters from the NIST database.6 Figure 6 shows a comparison of the simulated and experimental ethane conversion data as a function of reactor temperature. The two curves are displaced by some 50 °C but show fair agreement otherwise. This displacement could be caused by the gas temperature being higher than what we measure by the internal thermocouple. However, the displacement is also present at low conversion, where even an adiabatic reaction would not be able to heat the gas as much as 50 °C. We therefore tend to ascribe this discrepancy to the choice of rate constant for Cl2 splitting, where, as mentioned before, we have a wide choice among published rate constants. The bulk of the NIST activation energies range from 40 to 50 kcal/ mol, with a single value reaching as high as 56 kcal/ mol. We selected the values from Schofield et al.7 on the basis that they gave the best fit to our measured results. Whereas the highest of the NIST activation energies is in good agreement with our experimental

results, the corresponding preexponential factor led to reaction rates far below those observed experimentally. The discrepancy between the Schofield activation energy of 40 kcal/mol and our best quantum chemical estimate (vide infra) of 51.3 kcal/mol might appear discouraging. However, the 40 kcal/mol value was measured at elevated temperatures and might be too low because of vibrational excitation of the chlorine molecules. Another possible source of discrepancy is that the apparent experimental activation energy might not be directly comparable to the quantum chemical result. The latter concerns a purely monomolecular reaction, whereas, in the experimental case, the dissociation is triggered (and subsequent reassociation must be assisted) by a collision between Cl2 (Cl•) and some other molecule. The experimental values describe, to some extent, these effects and might, therefore, be more appropriately used in the simulations. It is interesting to note that the spectroscopic value,16 which should be largely independent of collision effects and vibrational excitation, is 57 kcal/ mol. Figure 7 shows the simulated product yields as functions of reaction temperature. The shapes of the curves as well as the relative orders of products is in qualitative agreement with the experimental results (Figure 4). According to both simulations and experiments, 1,1-dichloroethane is the preferred dichloroethane isomer, especially at reaction temperatures above 300 °C. Figure 8 shows simulated selectivities to 1,1-dichloroethane and 1,2-dichloroethane as functions of temperature. The fact that 1,1-dichloroethane is the dominant dichloroethane isomer is rather unexpected from a thermodynamic standpoint because 1,2-dichloroethane is slightly more stable. The difference in free Gibbs energy amounts to 0.3 kcal/mol at 180 °C and 0.6 kcal/mol at 530 °C, which yields an equilibrium constant of K1,1/1,2 )1.5 that is more or less temperatureindependent over the temperature range considered here. Both the 1,1-dichloroethane and 1,2-dichloroethane selectivities are practically constant at a low level up to 300 °C. The selectivity ratio is a constant 2.5, indicating that 1,1-dichloroethane is the dominant dichloroethane isomer in this temperature range. Both

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Figure 8. Selectivities to the 1,1- and 1,2-dichloroethane isomers as functions of temperature.

selectivities start to increase above that threshold, but whereas 1,2-dichloroethane passes through a maximum and decreases again at higher temperatures, the 1,1dichloroethane selectivity increases sharply over the whole temperature range. Closer inspection of the kinetic model provides an explanation for these findings. From the kinetic parameters for the elementary propagation steps in Table 1, we see that the primary attack of ethyl chloride by the chlorine radical is fast and only weakly temperature-dependent. The formation of the 1,1-dichloroethane precursor radical (reaction 7) is nearly a factor 2 faster than the formation of the isomer 1,2-dichloroethane precursor radical (reaction 8). The two different radicals generated in reactions 1 and 2 can form stable products in different ways. Theoretically, they might lose a hydrogen atom to form chloroethene (reactions 10 and 13). However, because of their large activation energies, these reactions do not play any role for the temperatures considered here. The two radicals can alternatively react with Cl2 to form 1,1dichloroethane and 1,2-dichloroethane (reactions 9 and 12, respectively). These reactions are very fast. The rate constant is slightly higher for the formation of 1,2dichloroethane. At lower temperatures, however, reactions 9 and 12 are not rate-determining for the formation of 1,2-dichloroethane and 1,1-dichloroethane. The 1,2-dichloroethane precursor radical can also lose a chlorine radical to form ethene (reaction 14). A corresponding reaction for the 1,1-dichloroethane precursor radical would lead to an unstable CH3CH carbene species and is not very likely. Another possibility for the two radicals to stabilize is the abstraction of atoms from a stable molecule other than Cl2. Compared to the reaction with Cl2, these reactions are relatively slow because the C-H, C-Cl, and C-C bonds (about 100 kcal/mol) are roughly twice as strong as the Cl-Cl bond.

In the case of the 1,1-dichloroethane precursor radical, however, the corresponding reaction with HCl was included because the possibility of losing a chlorine radical according to reaction 14 does not exist. HCl seemed to be a good choice because it has the highest concentration of all species at higher temperatures and conversions. However, compared to reaction 14, reaction 11 is several orders of magnitude slower. This might not be apparent at once by inspection of the listed rate constants and is largely due to the concentrationdependent part of the rate expression. Reaction 14 is a first-order reaction, whereas the radical concentration for reaction 11 is multiplied with a small number (concentration of HCl given in mol/cm3, which for the conditions here is on the order of 10-5). Based on these arguments. the interpretation of the experimental 1,1-dichloroethane/1,2-dichloroethane ratio is as follows: Up to 300 °C, the 1,1-dichloroethane/ 1,2-dichloroethane ratio is nearly constant and solely governed by the ratio of the rates of reactions 7 and 8. Ethene formation by reaction 14 does not play a role at lower temperatures because of its substantial activation energy. As the conditions become increasingly more severe, chlorination of the primary product, ethyl chloride, becomes more important. Hence, the selectivity to both 1,1-dichloroethane and 1,2-dichloroethane increases because of the increasing concentrations of ethyl chloride and chlorine radicals, respectively. At the same time, reaction 14 becomes faster with increasing temperature and competes to an ever greater extent with the less temperature-dependent reaction 12, thus decreasing the 1,2-dichloroethane selectivity at higher temperatures. The selectivity to 1,1-dichloroethane, however, increases over the whole temperature range, as reaction 11 is too slow to compete with the formation of 1,1-dichloroethane. The addition reaction of chlorine radicals to ethene is not listed in Table 1 and has not been considered here so far. This reaction might slow the net ethene formation at higher temperatures because it is the reverse reaction of reaction 14. However, including the addition of chlorine to ethene in the model did not affect the results of the simulations. This is in line with thermodynamic data showing that the equilibrium constant for the dehydrochlorination of ethyl chloride is far on the olefin side at higher temperatures. From the above discussions, the main reaction pathways might be as illustrated in Figure 9. 3.3. Agreement between Kinetic Model and Experimental Data. As stated before, there is an offset of ca. 50 °C between the modeled and experimental results. Two other aspects of the agreement between the model and the experimental data were also investigated: the relative reaction rates of consecutive and parallel reactions. Consecutive Reactions. This comparison examined the relative reaction rates in the reaction sequence ka

kb

kc

ethane 98 chloroethane 98 1,1-dichloroethane 9' 1,1,1-thrichloroethane The global rate constants ka, kb, and kc are determined by the rate constants k3, k7, and k22 in the simulation model. The ratios k7/k3 and k22/k3 are presented as functions of temperature in Figure 10. All of the reactions are first-order in (chloro)hydrocarbon chlorination (see Figure 5). Assuming they have

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Figure 11. Ratio of reaction products from 1-chloroethyl and 2-chloroethyl radicals from simulation and from experiment.

Figure 9. Qualitative picture of the most important carbon reaction pathways in ethane chlorination (from simulation). Arrow thicknesses are not to scale.

Figure 10. Relative chlorination rate constants k7/k3 and k22/k3 (for chlorination of chloroethane and 1,1-dichloroethane relative to chlorination of ethane) from simulation.

identical orders in chlorine concentration, the relative global rates can therefore be described by

rb kb cEC ) ra ka cE

rc kc cDCE ) ra ka cE

where ka, kb and kc are the global rate constants above and cE, cEC, and cDCE are the concentrations of ethane, chloroethane, and 1,1-dichloroethane, respectively. Estimation of these relative constants from the experimental data by global simulation gives

kb ) 0.17 ( 0.05 ka and

kc ) 0.07 ( 0.03 ka These values are in fair agreement with the model,

considering that the rate constants in the NIST database are measured at quite different conditions. Parallel Reactions. The other critical comparison is between the modeled and measured selectivities from chloroethane. In the simulation model, this can be expressed as the ratio of the rates of formation of 1-chloroethyl radical and 2-chloroethyl radical (reactions 7 and 8 in Table 1). Based on the interpretation of experimental findings above, further reactions of the 1-chloroethyl radical lead to the production of 1,1-dichloroethane and 1,1,1trichloroethane, as well as some chloroethene. From Figure 4c, the amount of chloroethene corresponds at all temperatures to 50% of the 1,1,1-trichloroethane production from 1,1-dichloroethane. Further reactions of the 2-chloroetyl radical lead to the formation of ethene, 1,2-dichloroethane, 1,1,2dichloroethane, and the rest of the chloroethene observed. The ratio between these product sums should be comparable to the ratio k7/k8. Such a comparison is made in Figure 11. Again, the agreement between experiment and model is fair, although the experiments show a stronger temperature dependence than the model. 3.4. Quantum Chemical Results. Quantum chemical calculations were performed in an attempt to account for the different reactivities observed in the experiments. Table 2 shows the calculated reaction and activation energies for selected reactions of interest. The discrepancies between the DFT and ab initio results look discouraging, as the transition states for hydrogen capture from ethanes by a chlorine radical disappear at the DFT level. In all cases, DFT calculations predict a metastable intermediate in which the chlorine radical coordinates to the hydrogen atom that is being abstracted. There is reason to suspect the DFT results in this case. Recent literature results17 suggest that the density functional in use, like most other commonly used functionals, severely overestimates the bonding strength of three-electron bonds such as the one found in the intermediate. Comparing the MP2 and CCSD(T) results, we note that the energy differences between analogous reactions are very small. As expected, MP2 systematically overestimates activation energies. The CCSD(T) absolute energies are, however, much closer to the experimental values in the cases where the latter are available. As mentioned above, the dissociation of Cl2 into atoms is reported (in the NIST database6) to require between 40 and 56 kcal/mol, compared to the

Ind. Eng. Chem. Res., Vol. 40, No. 10, 2001 2233 Table 2. Reaction Energies (∆E) and Activation Energies (Eq)a as Predicted by DFT, MP2, and CCSD(T) Calculations rxn no.

reactant(s)

3 7 8 22 16 40 4 9 12 23 20 18 26 41 42 25 43

C2H6 + Cl• CH3CH2Cl + Cl• CH3CH2Cl + Cl• CH2ClCH2Cl + Cl• CHCl2CH3 + Cl• CHCl2CH3 + Cl• C2H5•+ Cl2 CH3CHCl• + Cl2 •CH2CH2Cl + Cl2 CH2ClCHCl• + Cl2 CHCl2CH2• + Cl2 •CCl2CH3 + Cl2 2Cl• CH3CHCl• •CH2CH2Cl CH2ClCHCl• CHCl2CH2•

d

T T T T T T T T T T T T T T T T T

product(s)

∆EDFT

∆EMP2

∆ECCSD(T)

∆ECCSD(T) + ZPEb

EqDFT

EqMP2

EqCCSD(T)

EqCCSD(T) + ZPEb

C2H5• + HCl CH3CHCl• + HCl •CH2CH2Cl + HCl CH2ClCHCl• + HCl CHCl2CH2• + HCl CH3CCl2•+ HCl C2H5Cl + Cl• CH3CHCl2 + Cl• CH2ClCH2Cl + Cl• CH2ClCHCl 2 + Cl• CH2ClCHCl 2 + Cl• CCl3CH3 + Cl• Cl2 •CH2CH2Cl C2H4 + Cl• CH2dCHCl + Cl• CH2dCHCl + Cl•

1.4 -3.1 -0.3 -4.6 1.6 -6.3 -28.9 -22.4 -26.0 -17.8 -24.1 -15.0 -56.4 3.5 22.0 23.8 17.4

9.1 5.3 9.4 5.7 16.7 2.3 -43.5 -42.4 -37.9 -35.9 -41.4 -31.5 -41.8 4.1 36.8 17.4 12.0

4.1 0.4 4.2 n/cc n/c n/c -32.4 -26.7 -31.2 n/c n/c n/c -52.1 3.8 16.4 n/c n/c

-0.9 -4.1 -0.3 n/c n/c n/c -28.2 -23.4 -27.7 n/c n/c n/c -51.3 3.6 16.2 n/c n/c