Transesterification of Methyl Acetate and n-Butanol Catalyzed by

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Ind. Eng. Chem. Res. 2006, 45, 6648-6654

Transesterification of Methyl Acetate and n-Butanol Catalyzed by Amberlyst 15 Ewa Boz3 ek-Winkler and Ju1 rgen Gmehling* Department of Industrial Chemistry, UniVersity of Oldenburg, D-26111 Oldenburg, Germany

The reaction kinetics and chemical equilibrium of liquid-phase transesterification of methyl acetate and n-butanol to n-butyl acetate and methanol in the temperature range of 313.15 to 330.15 K were investigated. The reaction has been catalyzed by the acidic ion-exchange resin, Amberlyst 15. The chemical equilibrium constant obtained from kinetic experiments is in qualitative agreement with the one calculated from standard thermodynamic properties. The influence of the catalyst loading, initial reactant molar ratio, and temperature on the kinetics was studied. Two kinetic models, pseudohomogeneous (PH) and Langmuir-Hinshelwood (LH), were used to describe the reaction rate. Independent binary liquid sorption experiments were used to fit the adsorption constants needed for the LH model. The results of the more simple pseudohomogeneous model provide similar good results as the LH model. Introduction Methyl acetate is a byproduct during the production of poly(vinyl)alcohol (PVA). Per ton of PVA, 1.68 tons of methyl acetate are produced.1 Since methyl acetate is of limited industrial importance and methanol is a feedstock for PVA synthesis, the most attractive way would be to convert methyl acetate into methanol and n-butyl acetate, since n-butyl acetate is an important solvent for plastics, resins, gums, and coatings. n-Butyl acetate can also be used as an extracting agent, as an intermediate in organic synthesis, or in the photographic industry.2 One of the possibilities to obtain the desired products would be the transesterification of methyl acetate with n-butanol, which leads to n-butyl acetate and methanol. Because of the two binary maximum pressure azeotropes,3 namely, methanolmethyl acetate and n-butanol-n-butyl acetate, a conventional process with reaction followed by separation would cause a huge separation effort. Additionally, the low chemical equilibrium constant (close to unity) would result in low conversion and high capital costs. An alternative would be the application of reactive distillation or the combination of the reaction step with a membrane-separation process, e.g., pervaporation, to separate the reaction products from the reactants and, therefore, overcome the azeotropes and chemical equilibrium limitations. Besides kinetic information for the chemical reaction, reliable information about the heat and mass transfer, the vapor-liquid equilibrium behavior, the various pure-component properties, and the chemical equilibrium is required. The esterification of acids with alcohols is one of the most important applications of resin catalysis and has been studied in numerous papers. While the reaction rate in the gas phase is described very well by the Langmuir-Hinshelwood (LH) model demonstrating the heterogeneous nature of the catalyst,4 there is no unified opinion about the best kinetic model for reactions in the liquid phase. Several authors preferred the pseudohomogeneous (PH) expression,5 although others reported good results for the LH model.6 Published information about the kinetics of transesterification reactions catalyzed by ion-exchange resins is very limited. The commonly applied model to describe the reaction rate of the transesterification reaction is the PH model.7,8 From the good results obtained with the LH model for the * Corresponding author. Tel.: +49 441 7983831. Fax: +49 441 7983330. E-mail: [email protected].

esterification reaction, one could expect similar good results for transesterification reactions. In this paper, the chemical equilibrium and the reaction kinetics of transesterification of methyl acetate and n-butanol forming n-butyl acetate and methanol were studied. The reaction can be presented as

CH3COOCH3 + CH3(CH2)3OH / CH3COO(CH2)3CH3 + CH3OH A strongly acidic ion-exchange resin, Amberlyst 15, was used as heterogeneous catalyst. The forward and backward reactions were investigated, and two different kinetic models, namely, the pseudohomogeneous and the Langmuir-Hinshelwood, were applied to describe the reaction kinetics. The adsorption constants needed for the LH model were obtained in two ways: as parameters fitted to the independent binary liquid sorption experiments and as additional parameters fitted directly to the kinetic data. Similar kinetic parameters were obtained as given by Steinigeweg and Gmehling.7 Additionally in this work, the effect of the temperature, the initial molar ratio, and the amount of catalyst on the reaction kinetics is shown and a comparison of different kinetic models is presented. This should ensure a more reliable description of reaction kinetics for the transesterification of methyl acetate with n-butanol catalyzed by Amberlyst 15. Experimental Section (a) Chemicals. The chemicals were of analytical grade (99.8%) and were used without further purification except drying over molecular sieve (3 Å). Methanol and n-butanol were supplied by Roth, methyl acetate was supplied by Merck, and n-butyl acetate was supplied by Acros. The water content was verified by Karl Fischer titration (for each chemical below 100 ppm). (b) Catalyst. Amberlyst 15 (supplied by Aldrich), which is a macroreticular ion-exchange resin, was used as a catalyst. To remove impurities, prior to use, the catalyst was washed several times with distilled water until the supernatant liquid was colorless. The catalyst was then dried at 353.15 K under vacuum, until the mass remained constant. Usually, this procedure takes 2 days. (c) Analytics. All samples were analyzed by gas chromatography using an HP 6890 with a thermal conductivity detector

10.1021/ie060536e CCC: $33.50 © 2006 American Chemical Society Published on Web 08/24/2006

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(TCD) (He as carrier gas at 1.682 bar; HP-Innowax 30 m × 0.035 mm; split 10:1; temperature program 313 K hold for 4 min, heat at 65 K/min to 453 K hold for 2.5 min; retention times: methyl acetate 3.3 min, methanol 4.0 min, n-butyl acetate 6.0 min, and n-butanol 6.7 min). (d) Apparatus and Procedure. (d.1) Kinetics. The kinetic experiments were conducted in a stirred thermostated glass reactor with a volume of 500 cm3. The temperature of the mixture was kept constant within (0.1 K. The plate-type stirrer with baffle was set to a speed of 300 rpm. The reflux condenser was used to avoid any loss of volatile compounds. Liquid samples of ∼1 cm3 were taken with a syringe through a porous filter, to avoid catalyst lost. Methyl acetate or, in the case of the reverse reaction, n-butyl acetate was placed in the reactor together with the catalyst and preheated to the desired temperature. n-Butanol, or for the reverse reaction, methanol, was preheated in a separate vessel, and after reaching the reaction temperature, the alcohol was added to the reactor. The stirrer and the time measurement were started immediately. During each experiment (duration 7-24 h), between 15 and 20 samples were taken. Each sample was cooled rapidly to ∼278 K to avoid any further reaction and analyzed by GC. Measurements were performed between 313.15 and 330.15 K. Besides the temperature, the mass of catalyst and the initial molar reactant ratio of BuOH to MeOAc were varied in the ranges 8-18 wt % and 0.56-1.91, correspondingly. For the backward reaction, the initial molar reactant ratio of MeOH to BuOAc was kept constant (1:1). As was shown by Steinigeweg and Gmehling7 for the stirrer speed between 200 and 600 rpm, no external mass-transfer resistance was observed. This is in good agreement with the results published by Po¨pken et al.6 for the methyl acetate synthesis. Therefore, all experimental runs were performed with a stirrer speed of 300 rpm. Steinigeweg and Gmehling7 and Po¨pken et al.6 reported as well that, because of the fact that Amberlyst 15 is composed of small gel-type microspheres with large macropores,9 internal mass transfer can be excluded for the transesterification reaction. (d.2) Sorption Experiments. Sorption experiments were carried out at 298.15 K in a set of thermostated Teflon tanks with a volume of 100 cm3 each, placed on an automatic shaker tray. Into each tank a fixed amount, 5 ( 0.01 g, of the vacuumdried catalyst was placed and the mixture of the binary nonreactive system (∼20 g) was added. After equilibrium was reached, which takes ∼24 h, liquid samples were taken and analyzed by GC. (d.3) Swelling Experiments. Swelling experiments were conducted at 298.15 K in a sealed, graduated cylinder with a volume of 10 ( 0.1 cm3. A known amount of vacuum-dried catalyst (between 1 and 4 g) was placed in the cylinder, and 5 g of solvent was added. The cylinder was placed in an ultrasonic bath, and after 24 h, the catalyst volume was measured. For each component, the experiments were repeated at least three times.

Table 1. Relative van der Waals Volumes, ri, and Surfaces, qi11 component

ri

qi

methanol methyl acetate n-butanol n-butyl acetate

1.4311 2.8042 3.4543 4.8274

1.4320 2.5760 3.0520 4.1960

Table 2. Temperature-Dependent UNIQUAC Interaction Parameters7 i

j

aij (K)

bij

cij (K-1)

1 2 n-butanol methyl acetate 1 2 n-butanol methanol 1 2 n-butyl acetate methanol 1 2 n-butanol n-butyl acetate 1 2 n-butyl acetate methyl acetate 1 2

2 1 2 1 2 1 2 1 2 1 2 1

326.2 62.97 -21.08 339.3 600.3 -636.2 492.7 9.396 -48.26 260.1 -1340 937.1

0.725 -0.720 0.556 -1.174 -4.047 4.550 -0.083 -0.337 0.200 -0.499 4.297 -2.963

-2.365 × 10-3 1.17 × 10-3 0.0 0.0 6.44 × 10-3 -7.25 × 10-3 0.0 0.0 -4.54 × 10-4 2.12 × 10-4 0.0 0.0

component 1

component 2

methyl acetate methanol

equilibrium constant, Ka, from the equilibrium composition, using eq 1, where the nonidealities in the liquid phase were taken into account by the activity coefficients, γi, which were calculated using the UNIQUAC equation.10 A quadratic temperature dependence of the interaction parameters ∆uij ) aij + bijT + cijT 2 was used to describe the activity coefficients. The required van der Waals values ri and qi for the pure components (see Table 1) were taken from the Dortmund Data Bank (DDB), version 2006, which was kindly placed to our disposal by DDBST GmbH, Oldenburg.11 The required interaction parameters, which are presented in Table 2, were fitted previously in our research group and published elsewhere.7 In Figure 1, the equilibrium constant obtained in this way is plotted as the function of the inverse absolute temperature, together with the values calculated using standard thermodynamic properties. When the standard enthalpy of reaction is assumed to be constant, in the small temperature interval investigated, the standard enthalpy of reaction ∆h0r and the standard Gibbs energy of reaction ∆g0r can be obtained using the van’t Hoff equation

ln Ka(T 0) ) -∆g0r /RT 0 ln(Ka) ) ln Ka(T 0) -

( )

∆h0r 1 1 R T T0

The standard thermodynamic properties obtained from the measurements are given in Table 3, together with the values calculated from the tabulated standard thermodynamic properties

Results and Discussion (a) Chemical Equilibrium. The equilibrium constant of the transesterification reaction is defined as indicated by the following equation

Ka )

∏aiν ) Kx × Kγ ) ∏xiν × ∏γiν i

i

i

(1)

Most of the kinetic experiments lasted long enough that chemical equilibrium was reached. This allowed us to calculate the

Figure 1. Chemical equilibrium constants Ka ([) and the best fit of our data (s), compared with an estimation based on tabulated values (see Table 4) for the standard heats of formation and Gibbs energy of formation (- - -).

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Table 3. Standard Enthalpies and Standard Gibbs Energies of Reaction source

∆h0r (kJ‚mol-1)

∆g0r (kJ‚mol-1)

calcd from thermo. properties, Table 4 obtained by linear regression, Figure 1

3.23 2.23

-0.53 0.20

Table 4. Standard Enthalpies and Standard Gibbs Energies of Formation12 component

∆h0f (kJ‚mol-1)

∆g0f (kJ‚mol-1)

n-butyl acetate n-butanol methyl acetate methanol

-485.60 -274.60 -411.90 -200.94

-312.60 -150.30 -324.20 -162.32

products;12

of the individual reactants and see Table 4. However, one should be aware that a small error in ∆g0f and ∆h0f can cause a relatively large error for ∆g0r and ∆h0r , due to the low values of these thermodynamic standard properties. The value of ∆h0r is relatively close to zero; this means that there is nearly no temperature dependence of the equilibrium constant. As can be seen in Figure 1, the slopes of the dashed line, which represents the results obtained from the tabulated standard thermodynamic properties, and the solid line, which is the best fit of the experimental data, are rather similar, indicating a good agreement of the calculated and the experimental standard enthalpies of reaction. It can also be seen that the reaction is slightly endothermic, since the equilibrium constant Ka increases with increasing temperature. The temperature dependence of the equilibrium constant can be expressed by ln Ka ) 0.8158 - 267.9/T, which describes the best fit of the experimentally obtained equilibrium constants. (b) Reaction Kinetics. In the present work, several kinetic experiments for the heterogeneous catalyzed transesterification of methyl acetate with n-butanol, as well as the backward reaction, were carried out. The effects of catalyst loading, reaction temperature, and initial reactant molar ratio on reaction kinetics were determined. According to other authors, e.g., Sanz et al.5 and Cunill et al.,13 the reaction rate, based on the amount of dry catalyst, can be obtained from r0i ) n0/mcat(dX/dt)t)0, where the differential term is calculated as the slope of the fifth-degree polynomial curve fitted to the conversion-versus-time data. (b.1) Effect of Catalyst Loading. The effect of the amount of catalyst (between 8 and 18 wt %) on the conversion of methyl acetate is shown in Figure 2. In all cases, for the chemical equilibrium constant the same value was obtained. Chemical equilibrium was reached faster with higher catalyst loading.

Figure 3. Initial reaction rate versus catalyst loading at 319.15 K (experimental data [, PH model s) and 330.15 K (experimental data 2, PH model - - -).

Figure 4. Influence of the temperature, 313.15 K (experimental data [, s) and 325.15 K (experimental data 2, PH model - - -), on the conversion of MeOAc at 13 wt % of catalyst and initial molar reactant ratio BuOH/ MeOAc ) 1.

Figure 3 shows a linear relationship between the initial reaction rate, r0, and the amount of dry catalyst for two different temperatures and initial molar reactant ratio 1:1. The experimental results are in good agreement with the calculated ones, which proves the validity of the model. (b.2) Effect of the Reaction Temperature. The temperature of the kinetic experiments was varied between 313.15 and 330.15 K. As shown in Figure 4, with increasing of temperature, the reaction rate increases, but nearly the same equilibrium conversion of methyl acetate is obtained in the investigated temperature range. Similar results were published in the literature14,15 for esterification reactions. The reason is that the heat of reaction for esterification and transesterification reactions is small; i.e., the equilibrium constant depends only slightly on temperature.16 The temperature dependence of the rate constants is expressed by Arrhenius’ law:

ki ) k0i exp

Figure 2. Influence of the amount of catalyst, 9.4 wt % (experimental data [, PH model s) and 11.9 wt % (experimental data 2, PH model - - -), on the conversion of MeOAc at 319.15 K and initial molar reactant ratio BuOH/MeOAc ) 1.

( ) -EA,i RT

The activation energy can be calculated from the slope d ln ki /d(1/T) in the Arrhenius diagram, which is the natural logarithm of the rate constant plotted versus the inverse temperature. (b.3) Effect of the Initial Molar Reactant Ratio. The initial molar ratio of the reactants BuOH to MeOAc was varied between 0.56 and 1.91 during the experiments. Figure 5 shows that the equilibrium conversion of MeOAc increases with increasing molar ratio of n-butanol to methyl acetate. (c) Reaction Kinetic Modeling. (c.1) Pseudohomogeneous Model. Many of the resin-catalyzed reactions can be classified as quasi-homogeneous. The idealized homogeneous state requires complete swelling of the resin and total dissociation of the polymer-bound-SO3H groups.17 A pseudohomogeneous

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Figure 5. Influence of the initial molar reactant ratio, BuOH/MeOAc ) 1 (experimental data [, PH model s) and BuOH/MeOAc ) 1.7 (experimental data 2, PH model - - -), on the conversion at 319.15 K for 12 wt % of catalyst.

Figure 6. Arrhenius diagram of the rate constants for forward (experimental data 2, PH model s) and backward (experimental data 9, PH model - - -) reaction.

model can be applied for systems where the mass-transfer resistance is negligible and one of the reactants or solvents is highly polar. This model is based on the Helfferich approach, which treats catalysis by ion-exchange resins as homogeneous catalysis confined within the internal catalyst mass, where the reactants, products, and solvents are in distribution equilibrium with the bulk solution. The swelling of the resin particle in the presence of polar solvents leads to an easy accessibility of the acid groups for the reaction and free mobility of all the components.17,18 Transesterification reactions are known to be reversible second-order reactions. Therefore, a pseudohomogeneous kinetic model can be used:

1 1 dni ) k1aBuOHaMeOAc - k-1aMeOHaBuOAc (2) mcat νi dt

Figure 7. Relative adsorption of methanol from methyl acetate-methanol mixtures on Amberlyst 15 and calculated dependence (s) assuming constant adsorbed mass.

In the investigated reaction, the stoichiometric coefficient νi ) 1. The kinetic parameters can be fitted to the experimental data by numerically integrating the kinetic equations using a fourthorder Runge-Kutta method and minimization of the mean square deviation F ) 1/ns∑ns(xMeOAc,calc - xMeOAc,exp)2, between the calculated and experimental mole fractions of methyl acetate using the Simplex-Nelder-Mead method.19 In the case of a pseudohomogeneous kinetic model, four 0 adjustable parameters, the preexponential factors k01 and k-1 and the activation energies EA,1 and EA,-1, have to be fitted. The results are given in Table 5. In Figure 6, the reaction rate constants for forward and backward reactions calculated from single experiments at a given temperature are plotted in the Arrhenius diagram. The rate constants calculated from the PH model are included as lines. From the plot, it can be seen that the temperature dependence of the rate constants can be described by Arrhenius’ law and that the pseudohomogeneous model is able to reproduce the kinetic data. (c.2) Adsorption-Based Model. To find out whether sorption effects are important for the investigated kinetics, additionally the Langmuir-Hinshelwood (LH) model was applied. Assuming

that the process is controlled by the reaction on the catalyst surface, the LH model assumes that the reaction takes place between two adsorbed molecules. The LH kinetic model can be described by the following relation

r)

r)

1 1 dni

)

k1aBuOHaMeOAc - k-1aMeOHaBuOAc

mcat νi dt

(1 +

∑i

(3) Ki ai)2

where Ki are the adsorption constants for component i. The adsorption constants were obtained by two different procedures. First, the constants were fitted to the experimental data as additional parameters, together with the activation energies and the preexponential factors. In the second procedure, the adsorption constants were fitted using data from the independent sorption experiments. All four nonreactive binary systems, i.e., methyl acetate-methanol, n-butyl acetate-nbutanol, methyl acetate-n-butyl acetate, and methanol-nbutanol, were investigated in sorption experiments (Figures 7-10, respectively). The influence of the mass ratio of solvents

Table 5. Parameters and Objective Function of the Different Kinetic Models forward reaction kinetic model PH model LH model adsorption constants from kinetics LH model adsorption constants from adsorption

backward reaction

objective function

EA,1 (kJ mol-1)

k0-1 (mol g-1 s-1)

EA,-1 (kJ mol-1)

F

38.38 491.94

40.89 37.52

38.64 461.75

40.85 37.32

22.19 × 10-3 21.51 × 10-3

491.65

37.52

461.46

37.31

21.52 × 10-3

(mol

k01 g-1 s-1)

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Figure 8. Relative adsorption of butanol from n-butyl acetate-n-butanol mixtures on Amberlyst 15 and calculated dependence (s) assuming constant adsorbed mass.

Figure 9. Relative adsorption of methyl acetate from n-butyl acetatemethyl acetate mixtures on Amberlyst 15 and calculated dependence (s) assuming constant adsorbed mass.

to catalyst, which varied between 2 and 8, was studied for the system methanol-n-butanol. No significant influence on sorption was found within the investigated range; see Figure 10. The adsorption constants were obtained from the relation for binary systems proposed by Po¨pken et al.,6 which assumes a Langmuir-type adsorption based on the mass6

m0(w01 - wL1 ) mS K1a1wL2 - K2a2wL1 ) mcat mcat 1 + K1a1 + K2a2

(4)

Po¨pken et al.6 showed that, for the methyl acetate hydrolysis reaction, the adsorbed molar amount per unit mass of catalyst varied much more (between 3.58 and 26.5 mmol/g) than the adsorbed mass per unit mass of catalyst (between 0.265 and 0.478 g/g). They concluded that the values based on mass can be assumed to be constant for all the components. To check whether the same simplification can be assumed for the system investigated in this work, swelling experiments were carried out for all four components. As can be seen in Table 6, also in this case the adsorbed mass varies less than the adsorbed molar amount, and therefore, it was assumed to be constant. For all components, only one parameter was fitted (mS/mcat ) 0.99). The required sorption parameters, which are adsorption constants and adsorbed mass, were fitted to the binary

Figure 10. Relative adsorption of methanol from n-butanol-methanol mixtures on Amberlyst 15 and calculated dependence (s) assuming constant adsorbed mass. Mass of solvents to mass of catalyst ratio is equal to 4 ([) or varies between 2 and 8 (4).

sorption data using the Simplex-Nelder-Mead method.19 As shown in Figures 7-10, the calculated sorption results reproduce the data within experimental error. When all adsorption constants are divided by molecular weight, the one for methanol is the highest. From the adsorption constants given in Table 7, it can be seen that sorption is stronger for alcohols. This is in good agreement with the results of Mazzotti et al.20 They found that the affinity of Amberlyst 15 for the acetic acid, ethanol, ethyl acetate, and water follows the sequence of water > alcohol > acid > ester. The comparatively high value obtained for mS/mcat might be explained by the increase in swelling ratio when the liquid phase is a mixture and not a pure component.6 To show which of the components is sorbed the most, the sorption experiments were carried out also for the quaternary system. Because of the possibility of reaction, which would occur in the quaternary system, the experiment was performed at equilibrium composition. A comparison of the weight fractions of each of the components before and after the sorption equilibrium shows (see Table 6) that methanol is the best and n-butyl acetate is the worst sorbed compound of the system. The parameters obtained from the correlation of the experimental kinetic data using the pseudohomogeneous and LH models are listed in Table 5. The high values of activation energies confirm the theory that the reaction takes place on the surface of the catalyst and is not diffusion controlled. This means that the external and internal mass-transfer resistance can be neglected. The results are in good agreement with the ones published by Steinigeweg and Gmehling.7 All the models show similar mean-squared deviation and are able to represent the experimental data within experimental error. The results of the LH model with adsorption constants obtained from the independent sorption experiments are in good agreement with the results obtained from the LH model with adsorption constants fitted directly to the kinetic data. The results are shown in Table 7. The highest adsorption constant for alcohols and the lowest for n-butyl acetate show the trend, which is in good agreement with the results obtained from sorption experiments.

Table 6. Experimental Swelling Ratios, Adsorbed Volumes, Masses, and Moles Per Gram of Dry Amberlyst 15 Obtained for the Pure Components and Equilibrium Quaternary Sorption Data at 298 K pure-component adsorption data

quaternary adsorption data

component

swelling ratio

adsorbed volume (cm3 g-1)

adsorbed mass (g g-1)

adsorbed amount (mol g-1)

initial weight fraction

weight fraction after sorption

n-butyl acetate n-butanol methyl acetate methanol

1.33 1.59 1.42 1.58

0.233 0.417 0.296 0.409

0.205 0.338 0.276 0.324

1.765 × 10-3 4.559 × 10-3 3.729 × 10-3 10.104 × 10-3

0.375 0.265 0.248 0.112

0.409 0.273 0.241 0.078

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Conclusions

Figure 11. Mole fractions of methyl acetate versus time: experimental data ([), PH (s), and LH (- - -). Table 7. Adsorption Constants and Equilibrium Adsorbed Mass Fitted to the Binary Adsorption Data and Kinetic Data adsorption constants fitted to the kinetic data component n-butyl acetate n-butanol methyl acetate methanol

adsorption constants fitted to the adsorption data

original

divided by mol wt

original

divided by mol wt

2.67 5.61 4.01 5.00

0.0229 0.0757 0.0541 0.1561

2.67 5.62 4.01 4.99

0.0229 0.0758 0.0541 0.1557

equilibrium adsorbed mass, mS/mcat

source

Ka

PH model LH model adsorp. constants from kinetics LH model adsorp. constants from adsorption from linear regression (Figure 1) from thermodynamic properties (Table 4)

0.976 0.982 0.980 0.921 1.238

In chemical equilibrium, the reaction rate is equal to zero. Therefore, from both kinetic expressions (eqs 2 and 3), the following simple relation can be obtained,

k1 k-1

Acknowledgment Financial support from the Arbeitsgemeinschaft Industrieller Forschungsvereinigungen (AiF Nr. N11705/03) is gratefully acknowledged. We also would like to thank Dr. J. Rarey and Dr. M. T. Sanz Diez for help and support and E. Hanuscheck for performing some of the kinetic experiments. Nomenclature

0.99

Table 8. Chemical Equilibrium Constant at 298.15 K

Ka )

The kinetic behavior of the transesterification of methyl acetate and n-butanol, leading to n-butyl acetate and methanol catalyzed by Amberlyst 15, has been studied experimentally. The value of enthalpy of reaction found for this reaction is close to zero. This means that the chemical equilibrium is nearly temperature independent. Two different kinetic models, pseudohomogeneous and Langmuir-Hinshelwood, have been applied to describe the reaction kinetics. To obtain reliable results, independent sorption experiments were carried out. The PH model provides nearly the same results as the LH model, although it is simpler and contains less parameters. The PH model can easily be incorporated into commercial simulators to describe the reaction kinetics, e.g., for the design of reactive distillation processes or membrane reactors.

(5)

Using the kinetic parameters listed in Table 5, the chemical equilibrium constant can be calculated. In Table 8, chemical equilibrium constants calculated from the kinetic parameters obtained from the PH and LH models, as well as the constants obtained from Figure 1 and from the thermodynamical properties (Table 4), are given. Good agreement is found between the experimentally obtained equilibrium constant (Figure 1) and the one calculated from eq 5 for different models. Figure 11 shows a comparison of the experimental and the predicted mole fractions using the PH and LH models as a function of time. As can be seen, for both models there is a good agreement with the experimental data. Almost no difference between the models can be observed. The pseudohomogeneous model, often used to describe esterification and transesterification reactions catalyzed by ion-exchange resins,6,8,21 despite less parameters, provides similar results as the LH model. This means the pseudohomogeneous model, which is simpler and consists of less parameters, can be successfully applied for the design of reactive distillation processes or membrane reactors. But one should keep in mind that, in the systems where the components differ in polarity, sorption can have a significant influence on reaction kinetics.

a ) activity aij ) parameter used in eq 2, K bij ) parameter used in eq 2 cij ) parameter used in eq 2, K-1 EA ) activation energy, kJ‚mol-1 F ) objective function ∆g0f ) standard Gibbs energy of formation, kJ‚mol-1 ∆g0r ) standard Gibbs energy of reaction, kJ‚mol-1 ∆h0f ) standard enthalpy of formation, kJ‚mol-1 ∆h0r ) standard enthalpy of reaction, kJ‚mol-1 Ka ) equilibrium constant in terms of activities Kx ) equilibrium constant in terms of mole fractions Kγ ) equilibrium constant in terms of activity coefficients Ki ) adsorption constant k ) rate constant, mol g-1 s-1 k0 ) preexponential factor, mol g-1 s-1 m0 ) total solvent mass, g mS ) total adsorbed mass, g mcat ) mass of catalyst, g n0 ) initial mole number of MeOAc ni ) number of moles nS ) number of experimental data R ) general gas constant, 8.314 33 J‚mol-1‚K-1 r ) reaction rate, mol gcat-1 s-1 T ) absolute temperature, K t ) time, s ∆uij ) UNIQUAC parameters, K w0 ) liquid-phase weight fraction wL ) equilibrium liquid-phase weight fraction X ) conversion x ) mole fraction γ ) activity coefficient ν ) stoichiometric coefficient Subscripts and Superscripts 0 ) standard state/conditions 1, -1 ) forward and backward reaction, respectively calc ) calculated value

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exp ) experimental value i, j ) ith and jth components, respectively BuOAc ) n-butyl acetate BuOH ) n-butanol MeOAc ) methyl acetate MeOH ) methanol Literature Cited (1) Fuchigami, Y. Hydrolysis of Methyl Acetate in Distillation Column Packed with Reactive Packing of Ion-Exchange Resin. J. Chem. Eng. Jpn. 1990, 23, 354. (2) Gangadwala, J.; Mankar, S.; Mahajani S. Esterification of Acetic Acid with Butanol in the Presence of Ion-Exchange Resins as Catalysts. Ind. Eng. Chem. Res. 2003, 42, 2146. (3) Gmehling, J.; Menke, J.; Krafczyk, J.; Fischer, K. Azeotropic Data, 3 parts; VCH-Verlag: Weinheim, Germany, 2004. (4) Shimizu, S.; Hirai, C. Kinetic Study of Liquid-Phase Esterification with Sulfonic Acid Cation-Exchange Resin of the Macroreticular Type. I. Heterogeneous-Pseudohomogeneous Resin Catalysis. Chem. Soc. Jpn. 1986, 59, 7. (5) Sanz, M. T.; Murga, R.; Beltra´n, S.; Cabezas, J. L. Kinetic Study for the Reactive System of Lactic Acid Esterification with Methanol: Methyl Lactate Hydrolysis Reaction. Ind. Eng. Chem. Res. 2004, 43, 2049. (6) Po¨pken, T.; Go¨tze, L.; Gmehling, J. Reaction Kinetics and Chemical Equilibrium of Homogeneously and Heterogeneously Catalyzed Acetic Acid Esterification with Methanol and Methyl Acetate Hydrolysis. Ind. Eng. Chem. Res. 2000, 39, 2601. (7) Steinigeweg, S.; Gmehling, J. Transesterification Processes by Combination of Reactive Distillation and Pervaporation. Chem. Eng. Process. 2004, 43, 447. (8) Jime´nez, L.; Garvı´n, A.; Costa-Lo´pez, J. The Production of Butyl Acetate and Methanol via Reactive and Extractive Distillation. I. Chemical Equilibrium, Kinetics and Mass-Transfer Issues. Ind. Eng. Chem. Res. 2002, 41, 6663. (9) Pitochelli, A. R. Ion Exchange Catalysis and Matrix Effects; Rohm and Haas Co.: Philadelphia, PA, 1980.

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ReceiVed for reView April 27, 2006 ReVised manuscript receiVed July 6, 2006 Accepted July 20, 2006 IE060536E