Determination of Adsorption and Kinetic Parameters for

Apr 14, 2012 - Chemical Engineering Department, Ege University, Bornova, Izmir, Turkey. Ind. Eng. Chem. Res. , 2012, 51 (18), pp 6350–6355. DOI: 10...
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Determination of Adsorption and Kinetic Parameters for Transesterification of Methyl Acetate with Hexanol Catalyzed by Ion Exchange Resin Emine Sert* and Ferhan Sami Atalay Chemical Engineering Department, Ege University, Bornova, Izmir, Turkey S Supporting Information *

ABSTRACT: The transesterification of methyl acetate and 1-hexanol catalyzed by the cation exchange resin Amberlyst-131 was studied to obtain optimum operating parameters, adsorption parameters, and kinetic parameters. The effects of temperature, molar ratio of ester to alcohol, stirrer speed, and catalyst loading on the reaction rate were investigated. The chemical equilibrium constants were obtained from kinetic experiments and theoretically from standard thermodynamic properties at temperatures of 333, 338, 343, and 348 K. The experimental data were tested with the pseudohomogeneous and adsorption based models. The activity coefficients were estimated using UNIQUAC to account for the nonideal thermodynamic behavior of reactants and products for both models. The activation energy for the transesterification reaction was found to be 37.8 kJ mol−1 by the Langmuir−Hinshelwood−Haugen−Watson (LHHW) model, which correlates the experimental data.

1. INTRODUCTION Methyl acetate is a byproduct during the production of PVA (poly(vinyl alcohol)). From 1 ton of PVA, about 1.68 tons of methyl acetate is produced. The industrial application of methyl acetate is low.1 The products of transesterification of methyl acetate with hexanol are methanol, which is a raw material for PVA production, and hexyl acetate, which is an important chemical used as solvent for resins, polymers, fats, and oils. Esters can be produced by various methods, on both laboratory and commercial scales. The simplest route to produce esters with high yields is the direct esterification of acids with alcohol in the presence of homogeneous or heterogeneous catalysts.2 Also transesterification of ester with alcohol in the presence of acidic catalyst is another way to produce esters. The esterification of acids with alcohols and transesterification of esters with alcohols are the most important applications of resin catalyzed reactions and have been studied in numerous papers. Akbay and Altıokka3 reported the production of amyl acetate catalyzed by Amberlyst 15, Amberlyst 36, and Amberlite IR-120 in a batch reactor. The experimental results show that Amberlyst 36 has slightly less catalytic activity than the others. Jimenez et al.4 reported the production of butyl acetate and methanol by reactive and extractive distillation catalyzed by Amberlyst 15. Bozek-Winkler and Gmehling5 converted methyl acetate to butyl acetate using transesterification with Amberlyst 15 as the catalyst. The value of the enthalpy of reaction found is close to 0. Two different kinetic models, pseudohomogeneous (PH) and Langmuir− Hinshelwood (LH), have been applied to describe the reaction kinetics. The PH model provides nearly the same results as the LH model. Steinigeweg and Gmehling6 reported the transesterification process of methyl acetate and 1-butanol by combination of reactive distillation and pervaporation using structured catalytic packings Katapak-S. Their study shows that the combination of reactive distillation with pervaporation is © 2012 American Chemical Society

favorable since conversions close to 100% can be obtained. Xu et al.1 studied the transesterification of methyl acetate and 1butanol catalyzed by the cation exchange resin NKC-9. The effects of temperature, molar ratio of reactants, and catalyst loading on the reaction rate were researched under the condition of eliminating the effect of diffusion. The experimental data were correlated with a kinetic model based on pseudohomogeneous catalysis. The advantage of a heterogeneous catalyst is well-known, and the use of solid ion exchange resin as the catalyst has the following advantages; (a) the catalyst can be easily separated from the reaction products by decantation or filtration; (b) continuous operation in the column is possible; (c) the side reaction can be eliminated and the product purity is high.1 Like esterifications, transesterification reactions are typical equilibrium limited reactions and the determination of thermodynamic equilibrium is required experimentally and theoretically. In general, the probable mechanisms for esterification reactions have been investigated by many researchers. The pseudohomogeneous model (PH) is widely used in esterification systems.7−9 In the PH model, adsorption and desorption of all components are negligible. The Langmuir−Hinshelwood−Haugen−Watson (LHHW) model takes into account the adsorption of all components.10 On the other hand, the Eley−Rideal (ER) model can be applied when reaction between one adsorbed reactant and one nonadsorbed reactant from the bulk liquid phase is assumed to occur.11,12 The ester of 1-hexanol, namely, hexyl acetate, is used in a wide range of industrial applications. n-Hexyl acetate is a fruitysmelling fluid used as flavoring agent or in perfumes. It is generally produced by the reversible, acid catalyzed liquid phase Received: Revised: Accepted: Published: 6350

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esterification of 1-hexanol and acetic acid.13 During this reaction, olefin formation and ether formation were considered as possible side reactions in the literature.14 Schmitt and Hasse14 used batch reactors to study the chemical equilibrium and autocatalyzed reaction kinetics and plug flow reactors to study the heterogeneously catalyzed reaction using Amberlyst CSP2 for the synthesis of n-hexyl acetate using concentrated acetic acid. They also reported the formation of two byproducts, namely, diethyl ether and 1-hexene. In this study, both the recovery of methyl acetate which is produced during PVA production and also the avoidance of above-mentioned side reactions were intended. Although there many studies on esterification and transesterification reactions, any information about the transesterification of methyl acetate with hexanol and the kinetic study of this reaction were not found in the literature. In this study, the transesterification reaction between methyl acetate and hexanol was performed and the effects of temperature, catalyst loading, ester-to-alcohol molar ratio, and stirrer speed were investigated. Also, the chemical equilibrium and the reaction kinetics of this reaction were studied to calculate the activation energy.

3. RESULTS AND DISCUSSION 3.1. Thermodynamic Analysis. 3.1. Calculation of Activities. The nonideality of the liquid phase was corrected by replacing the concentrations of components with activities. The components’ activity coefficients were calculated by the UNIQUAC group contribution method. Chemical Equilibrium. The experimental runs were carried out at temperatures of 333, 338, 343, and 348 K, a molar ratio of ester to alcohol of 1, and a catalyst loading of 4 wt %. Experiments were undertaken to determine the equilibrium mole fractions of methyl acetate, hexanol, hexyl acetate, and methanol. The equilibrium constant was calculated experimentally according to the following formula: K=

x HeAcxMeOH γHeAcγMeOH xMeAcx HeOH γMeAcγHeOH

where xi is the mole fraction of component i at equilibrium and γi is the activity coefficient of component i calculated by the UNIQUAC model. The equilibrium constants were found to be 1.06, 1.11, 1.20, and 1.27 at temperatures of 333, 338, 343, and 348 K, respectively. The heat of reaction can be calculated from the standard enthalpy of formation. The heat of reaction and equilibrium constant at different temperatures were calculated by using standard enthalpies of formation and Gibbs energies for the reactants and products. The heat of reaction was found to be 1.16 kJ mol−1, and the equilibrium constant was weakly dependent on temperature because of the low value of the heat of reaction. Equilibrium constants were also calculated from thermodynamic data at a reference temperature from standard Gibbs energies of reaction:

2. EXPERIMENTAL SETUP AND PROCEDURE 2.1. Materials. Methyl acetate (Merck) and hexanol (Merck) were used as reactants. Amberlyst 131, a cation exchange resin in the H+ form, was obtained from SigmaAldrich. 2.2. Kinetic Runs. The reactor consisted of a two-necked spherical glass flask of 500 mL capacity fitted with a coil condenser to prevent any loss of products. An electrical heater was used to heat the reaction mixture, and the batch content was stirred magnetically using a magnetic stirrer. The temperature of the reaction mixture was controlled using a temperature controller. In a typical run, methyl acetate and catalyst were charged into the reaction vessel in a predetermined ratio. The temperature of the reactor was set to the desired value. After the desired temperature was reached, preheated hexanol was added into the reactor and this was taken as zero time for a run. About 1 mL of liquid sample was withdrawn from the reactor at regular intervals for gas chromatographic analysis. 2.3. Equilibrium Runs. The equilibrium runs were carried out by the same procedure for kinetic runs. Equimolar amounts of methyl acetate and hexanol with a catalyst loading of 4 wt % were allowed to reach equilibrium. The reaction temperatures studied were 333, 338, 343, and 348 K. Equilibrium mole fractions of methyl acetate, hexanol, hexyl acetate, and methanol were calculated to determine the overall equilibrium constant. 2.4. Analysis. A Hewlett-Packard 6890GC gas chromatograph was used to determine the composition of the samples. The gas chromatograph column was HP-FFAP polyethylene glycol connected with flame ionization detector to detect the compounds in the same run. Helium was used as carrier at a flow rate of 6.8 × 10−8 m3 s−1. Injector, detector, and oven temperatures were 473.15, 503.15, and 423.15 K, respectively. The temperature program of gas chromatographic analysis was given as follows: waiting 3 min at 60 °C; heating from 60 to 180 °C at a rate of 5 °C/min; waiting for 3 min at 180 °C.

ln K (T0) = −

ΔG° RT0

(1)

Equilibrium constants at different temperatures were obtained from the integrated form of the van’t Hoff equation: ln K (T ) = ln K (T0) −

ΔH ° ⎛ 1 1⎞ ⎜ − ⎟ R ⎝T T0 ⎠

(2)

Figure 1 shows the variation of the chemical equilibrium constant with temperature, indicating a good agreement between experimental and theoretical values. It can also be

Figure 1. Variation of chemical equilibrium constant with temperature. 6351

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seen that the reaction is slightly endothermic, since the equilibrium constant increases with increasing temperature. 3.2. Absence of Mass Transfer Resistance. To achieve a purely kinetic study, it is necessary to eliminate both external and internal diffusion limitations. In this study, where the reaction of methyl acetate with hexanol over Amberlyst 131 was carried out in a batch reactor, the external and internal mass transfer resistances to transesterification reaction are directly related to stirrer speed and the particle size of the catalyst, respectively. External Mass Transfer Resistance. The experiments were performed at 343 K, initial alcohol-to-ester mole ratio of 1, catalyst loading of 4 wt %, and stirrer speeds of 400, 600, 800, and 1000 rpm to determine the limits at which diffusion limitations will not exist. Figure 2 shows that the conversion of

D lmμm 0.8 =

i=1 i≠l

(5)

where xi is the mole fraction of the component i, Dli is the binary diffusivity of the limiting reactant in component i, and μi and μm are the viscosities of component i and the mixture in cP, respectively. It was found that values of the internal diffusion parameter were significantly less than 1. These results indicate that internal diffusion does not limit the reaction of methyl acetate with hexanol over ion exchange resin for the reaction conditions implemented in this study (Table 1). These results Table 1. Significance of Internal Diffusion after 30 min temp (K)

catal loading (wt %)

HeOH/MeAc

333 338 343 343 343 343 343

4 4 4 6 8 4 4

1 1 1 1 1 0.5 2

Cwp 5.1 5.4 5.8 4.3 3.5 6.1 2.4

× × × × × × ×

10−4 10−4 10−4 10−4 10−4 10−4 10−4

confirm the studies claiming that the effect of external and internal diffusion can be neglected for most of the reactions catalyzed by the ion exchange resins.16 3.3. Optimization of Reaction Conditions. The effects of parameters such as temperature, catalyst loading, and reactant mole ratio on the reaction rate for transesterification of methyl acetate with hexanol were studied. Effect of the Reaction Temperature. The investigation on the effect of temperature is very important because this information is useful in performing the kinetic study and calculating the activation energy. To investigate the effect of the reaction temperature, operations were carried out at temperatures of 333, 338, 343, and 348 K. The initial molar ratio of hexanol to methyl acetate was 1, and the catalyst loading was kept constant at 4 wt %. Figure 3 shows the effect of

Figure 2. Effect of stirrer speed on conversion at 343 K, alcohol-toester ratio of 1, and catalyst loading of 4 wt %.

methyl acetate is the same when the mixture is agitated at speeds of 800 and 1000 rpm. This indicates the absence of external mass transfer limitations above 800 rpm. Therefore, all experiments were conducted at 800 rpm to ensure the external mass transfer limitation does not exist. This observation is in line with the work of Chakrabarti and Sharma,15 wherein it has been established that external diffusion does not generally control the overall rate in ion exchange resin catalyzed processes unless the viscosity of the reactant mixture is very high or the speed of agitation is very low. Internal Mass Transfer Resistance. To investigate the internal diffusion effect on the reaction rate, analysis based on the Weisz−Prater criterion was undertaken. Data from some experiments were fitted to the Weisz−Prater equation in which the experiments were catalyzed by Amberlyst 131:

Cwp =

∑ xiDliμm 0.8

−rA(obs)ρp R c 2 DeC li

(3)

where rA is the reaction rate at a given time, ρp is the catalyst density, Rc is the ratio of the catalyst pellet volume to the catalyst pellet external surface area, Cli is the limiting reactant concentration in the mixture at a given time, and De is the effective diffusivity. The definition of effective diffusivity is given as

De = ξ 2D lm

Figure 3. Effect of temperature on conversion at an alcohol-to-ester ratio of 1 and catalyst loading of 4 wt %.

temperature on the conversion of methyl acetate, which increased rapidly with the increase in the temperature. A similar trend has been observed for other solid catalyzed esterification and transesterification reactions.16−20 Higher temperature yields greater conversion of methyl acetate under identical conditions. Increasing the temperature is apparently favorable for the acceleration of the forward reaction.18

(4)

where ξ is the void fraction and Dlm is the diffusivity of the limiting reactant in the mixture in cm2/s. Dlm was calculated using the Perkin and Geankoplis method: 6352

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Effect of Catalyst Loading. The catalyst loading was expressed as the weight ratio of the catalyst to the total reaction weight. Catalyst loading was varied from 4 to 8 wt % at a temperature of 343 K, feed mole ratio of 1:1, and stirrer speed of 800 rpm. The effect of catalyst loading on the conversion of methyl acetate is given in Figure 4. With the catalyst loading

Figure 6. Effect of catalyst reusability on conversion of methyl acetate at catalyst loading of 4 wt % and temperature of 343 K.

conditions studied, kinetics rather than mass transfer is the ratecontrolling step. The reaction mechanism of transesterification is considered to consist of three reaction steps. The first step is carbocation formation, which results from the combination of a proton with a carbonyl group, which is relatively fast. Thus, it is considered to be always in equilibrium. The second step is the reaction of the carbocation with hexanol, which is slow. The third step is the formation of hexyl acetate and methanol. This step is also very fast and thus can be considered to be in equilibrium. Therefore, the second step is generally believed to be the rate-controlling step.17 Four reaction rate mechanisms that are expressed according to the adsorption status of the reactants were employed to the experimental data to obtain kinetic expression. Pseudohomogeneous Model (PH). The PH model assumes complete swelling of the polymeric catalyst in contact with polar solvents, leading to an easy access of the reactants to the active sites. This model can be derived from the LHHW model considering that none of the components is strongly adsorbed.21 Langmuir−Hinselwood−Haugen−Watson Model (LHHW). Assuming that the process is controlled by the reaction on the catalyst surface, the LHHW model assumes that the reaction takes place between two adsorbed molecules. Eley−Rideal Model (ER(I) or ER(II)). Depending on which of the two reactants is adsorbed, for a single site surface reaction rate-controlling step, the reaction between an adsorbed and a nonadsorbed reactant molecule on the catalyst surface can be represented by the Eley−Rideal model.21 In Table 2, k is the forward reaction rate constant in mol/ g·min and Ki is the adsorption constant for species i. The aim of the kinetic modeling is to minimize the squared differences between the calculated values of the rate and the rate obtained directly from the experimental data.

Figure 4. Effect of catalyst loading on conversion at 343 K and alcohol-to-ester ratio of 1.

increasing, the number of resin functional groups increased and thus the reaction rate increased. The increase of the reaction rate was due to the increase in the total number of acid sites available for the reaction with the increase of catalyst loading. Furthermore, the catalyst loading had no effect on the equilibrium constant.1 Effect of Molar Ratio of Hexanol to Methyl Acetate. The initial molar ratio of the reactants hexanol to methyl acetate was varied between 0.5 and 2 while the other operating variables were kept constant. Figure 5 shows that the conversion of

Figure 5. Effect of alcohol-to-ester ratio on conversion at 343 K and catalyst loading of 4 wt %.

methyl acetate increases with increasing molar ratio of hexanol to methyl acetate. Transesterification of methyl acetate with hexanol is an equilibrium limited chemical reaction because the position of equilibrium controls the amount of product formed.18 Therefore, the use of an excess of hexanol increased the conversion of methyl acetate. Reusability of Catalyst. In this study, after the experiments the catalyst was washed with distilled water and dried at 363 K for reuse. As observed from Figure 6, there is some loss of activity after four successive runs, and the conversion of methyl acetate was decreased by only 6−7% after the fourth run for the transesterification of methyl acetate with hexanol. 3.4. Kinetic Study. The absence of both external and internal diffusion limitations shows that, under the reaction

min φ =

∑

(rcal − rexp)2 (6)

all data samples

The mean relative error percentage between the predicted and experimental mole fractions was calculated by the following formula: mean relative error =

∑all data samples

|xexp − xpred|

nsample

xexp

· 100 (7)

In order to determine the best fit model, the mean relative errors defined by eq 7 were calculated for each model and 6353

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modified LHHW, were applied for correlating the kinetic data at different temperatures. The PH model gave a better correlation between the experimental and model reaction rates compared to adsorption rate models and activation energies. The order of the adsorption constants for the species was found to be KMeOH > KHeOH > KMeAc > KHeAc. The trend in the adsorption constants agree well with the study of Winkler et al.3 while studying the transesterification of methyl acetate with butanol. The trend in the adsorption constant is found to be the same as that of the solubility parameters. The solubility parameters of hexyl acetate, methyl acetate, hexanol, and methanol were found to be 16, 18.7, 23.3, and 29.6 (J/cm3)0.5. The relation between the adsorption constants and the solubility parameters agrees with earlier studies.21

Table 2. Kinetic and Adsorption Parameters kinetic model

temp (K)

k (mol/ g·min)

MRE (%)

KMeAc

KHeOH

KHeAc

KMeOH

PH

333 338 343 348 333 338 343 348 333 338 343 348 333 338 343 348

122.4 157.8 189.7 194.5 334.0 460.2 521.5 615.4 346.7 387.6 423.4 476.7 287.9 345.6 413.2 498.7

9.9 9.1 12.6 10.6 2.3 2.7 2.9 3.1 4.7 4.4 4.8 5.4 4.6 5.8 7.3 8.1

− − − − 3.89 3.78 3.65 3.49 3.56 3.32 3.12 3.01 − − − −

− − − − 5.46 5.29 5.08 4.97 − − − − 5.21 4.89 4.32 4.21

− − − − 2.89 2.78 2.65 2.54 3.01 2.98 2.76 2.54 2.78 2.65 2.41 2.29

− − − − 6.23 6.01 5.89 5.69 6.21 5.98 5.45 5.21 5.98 5.74 5.12 4.97

LHHW

ER(I)

ER(II)

4. CONCLUSION The kinetic behavior of the transesterification of methyl acetate and 1-hexanol, leading to n-hexyl acetate and methanol catalyzed by Amberlyst 131, has been studied experimentally. The significance of the internal mass transfer limitation was quantified by use of the Weisz−Prater criterion for experiments carried out at different operating conditions. It was found that the internal diffusion effect was negligible under the employed transesterification reaction conditions. Experiments were carried out at 800 rpm to ensure eliminating external mass transfer resistance. The heat of reaction was found to be 1.16 kJ mol−1. The transesterification reaction was found to be mildly endothermic. The experimental data were correlated by the LHHW model, and the activation energy was calculated as 37.8 kJ mol−1.

tabulated in Table 2. As shown from Table 2, experimental data were correlated by the LHHW model by a smaller mean relative error, so sorption can have a significant influence on the reaction kinetics. The Arrhenius plot for the transesterification of methyl acetate with hexanol is given in Figure 7. For the

■

ASSOCIATED CONTENT

S Supporting Information *

Tables for the properties of Amberlyst 131 (Table S1); volume parameters, equations for UNIQUAC method, and binary interaction parameters (Tables S2 and S3); standard enthalpies and Gibbs energies of compounds (Table S4) and rate expressions (Table S5). This material is available free of charge via the Internet at http://pubs.acs.org.

Figure 7. ln k vs 1/T.

■

transesterification of methyl acetate with hexanol catalyzed by Amberlyst 131, no value of the activation energy was found in the literature. The activation energy for production of hexyl acetate found in this work is about 37.8 kJ mol−1. The high value of the activation energy confirms the theory that the reaction takes place on the surface of the catalyst and is not diffusion controlled.5 Winkler et al.5 studied the transesterification of methyl acetate with butanol catalyzed by Amberlyst 15, and they found the activation energy to be 37.51 kJ mol−1. Their experimental data correlate the PH and LHHW models with similar deviations. In the study of Cui et al.,17 transesterification of methyl acetate with butanol catalyzed by ionic liquids was investigated and the activation energy was found to be 36.38 kJ mol−1. They found that the sulfate based ionic liquid is more active than both sulfuric acid and Amberlyst 15. Patel and Saha22 studied the esterification of dilute acetic acid with hexanol catalyzed by a macroporous cation exchange resin. The nonideality of each component in the reaction mixture was accounted for by using the activity coefficient with the use of the UNIFAC group contribution method. Both the PH and heterogeneous kinetic models, e.g., LHHW, ER, and

AUTHOR INFORMATION

Corresponding Author

*Tel.: 90 232 311 1493. Fax: 90 232 388 7776. E-mail: emine. [email protected]. Notes

The authors declare no competing financial interest.

■

ACKNOWLEDGMENTS This study was supported by a Ege University Scientific Research Project (10-MUH-027). The kind help of Saman Setoodeh Jahromy during the laboratory studies are also acknowledged.

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ABBREVIATIONS MeAc = methyl acetate HeOH = hexanol HeAc = hexyl acetate MeOH = methanol dx.doi.org/10.1021/ie300350r | Ind. Eng. Chem. Res. 2012, 51, 6350−6355

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REFERENCES

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