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Kinetic of myristic acid esterification with methanol in the presence of triglycerides over sulfated zirconia
Renewable Energy 36 (2011) 2679e2686
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Renewable Energy journal homepage: www.elsevier.com/locate/renene
Kinetic of myristic acid esterification with methanol in the presence of triglycerides over sulfated zirconia Dussadee Rattanaphra a, Adam P. Harvey c, Anusith Thanapimmetha a, b, Penjit Srinophakun a, b, * a
Bioprocess Laboratory, Department of Chemical Engineering, Faculty of Engineering, Kasetsart University, Bangkok 10900, Thailand Center of Excellent for Petroleum, Petrochemicals, and Advanced Materials, Department of Chemical Engineering, Faculty of Engineering, Kasetsart University, Bangkok, Thailand c School of Chemical Engineering and Advanced Materials, Newcastle University, Newcastle upon Tyne, NE1 7RU, UK b
a r t i c l e i n f o
a b s t r a c t
Article history: Received 16 November 2010 Accepted 22 February 2011 Available online 26 March 2011
The esterification of myristic acid with methanol in presence of triglycerides using sulfated zirconia prepared by solvent-free method as the heterogeneous catalyst was studied. The effects of reaction temperature (393e443 K), catalyst loading (1e3 %wt) and molar ratio of oil to methanol (1:4e1:20) on the conversion of myristic acid were investigated. The experimental data was interpreted with a second order pseudo-homogeneous kinetic model. The kinetic parameters were obtained. A good agreement between the experimental data and the model was observed. Sulfated zirconia prepared by solvent-free method exhibited high catalytic activity for this reaction. Low activation energy of 22.51 kJ mol1 was obtained in the range of temperature of 393e443 K. Ó 2011 Elsevier Ltd. All rights reserved.
Keywords: Esterification Myristic acid Triglyceride Sulfated zirconia Biodiesel
1. Introduction Biodiesel represents a promising alternative fuel for use in diesel engines due to its many advantages: it is renewable fuel, it creates low emissions of CO, SOx, NOx, it is biodegradable and non-toxic [1,2]. Biodiesel can be produced through transesterification process. Normally, refined vegetable oils such as rapeseed, soybean, sunflower and palm oils are used as biodiesel feedstock [3]. Even though these oils yield a high quality biodiesel, they are very expensive. This causes a higher cost of biodiesel as compared to the traditional petroleum diesel [4]. Moreover, an availability of these feedstocks is limited as they are used in the food industry [5]. The use of low quality oils like non-edible oil, waste grease, fats can result in a lower production cost of biodiesel [6]. However, these oils usually have high free fatty acids (FFA) concentrations for example used cooking oil and greases have free fatty acid levels greater than 6% [7]. The presence of large amounts of FFA strongly influences the performance of homogeneous base-catalyzed transesterification. These components cause the formation of soap through ponification
* Corresponding author. Bioprocess Laboratory, Department of Chemical Engineering, Faculty of Engineering, Kasetsart University, Bangkok 10900, Thailand. Tel.: þ662 942 8555; fax: þ662 561 4621. E-mail address: [email protected] (P. Srinophakun). 0960-1481/$ e see front matter Ó 2011 Elsevier Ltd. All rights reserved. doi:10.1016/j.renene.2011.02.018
which makes a serious problem of product separation [8]. Homogeneous acid catalysts have potential to simultaneously catalyze both esterification and transesterification processes of high FFA feedstock [9]. However, the use of these catalysts are non-practical for biodiesel production due to slow reaction rate, requirement of high temperature, high molar ratio of oil to alcohol, the separation of catalyst, serious environmental and corrosion related problems [10]. Recently, a twostep process is proposed to prepare biodiesel from high FFA oils. The first step is esterification of FFA with methanol which is catalyzed by homogeneous acid catalysts (generally sulfuric acid). When FFA content in oil is reduced to lower than 0.5e1 %wt, transesterification of triglyceride with methanol, the second step, is then carried out by homogeneous base catalysts (normally sodium or potassium hydroxide). This process increases the production cost of biodiesel since it requires extra separation steps to remove the catalyst in both stages [8,10,11]. Heterogeneous acid catalysts are developed to solve the problems associated with homogeneous catalysts. The main advantages of these catalysts are easy separation, reusability, low corrosion, and more environmentally friendly [12]. The heterogeneous base catalysts show high catalytic activity, however, most of them are very sensitive to trace amount of FFA and/or water in raw material [13]. Therefore, the use of heterogeneous acid catalysts is potentially attractive because of they can catalyze the transesterification of vegetable oils or animal fats with high content of FFA and water, such as deep-flying oils from restaurants and food processing
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without the formation of undesired products. These catalysts can also promote simultaneously transesterification of triglycerides and esterification of FFA that causes cheaper product cost [7]. Usage of heterogeneous acid catalysts in both transesterification and esterification reactions have been reported in several literatures. Heteropolyacids are typical strong Bronsted acids and have great potential to catalyze biodiesel synthesis. However, a major drawback of these catalysts is their relatively low thermal stability, low surface area and solubility in polar media [14]. Even though the supported heteropolyacid catalysts were developed to improve qualities, they still have problem associated with the continuous leaching of heteropolyacid from support into the reaction medium [15]. Cs-salt heteropolyacids exhibits high catalytic activity and high stability to leaching [15,16]. However, because of their insolubility, makes it is more difficult to prepare by conventional aqueous impregnation method [17]. Suflonic ion-exchange resins like Amberlyst 15 and Nafion NR50 showed higher catalytic activity in esterification than transesterification reaction [11]. These resins are expensive and have low thermal stability [18e20]. Sulfated zirconia (SO42/ZrO2) is classified as a heterogeneous superacid catalyst. It has recently received considerable attention as a promising catalyst for industrial processes due to its strong acid properties [21]. This catalyst presents both active Bronsted and Lewis acidic sites [22]. It is also higher acid strength than heteropolyacid, suflonic ion-exchange resins and other heterogeneous acid catalysts [23]. Sulfated zirconia was found to be a potential catalyst in the esterification of dodecanoic acid with 2-ethylhexanol, 1-propanol and methanol at 130e180 C, compared to the various heterogeneous catalysts such as zeolites ion-exchange resins and mixed metal oxides [18]. This catalyst has been successfully used as heterogeneous catalyst for esterification of palmitic acid in sunflower oil with methanol at 60 C [24] and esterification of high FFA in refined soybean oil with methanol at 35e60 C [25]. Six solid catalysts (ZrO2, ZnO, SO42/SnO2, SO42/ ZrO2, KNO3/KL zeolite and KNO3/ZrO2) were carried out for transesterification of crude palm kernel oil and crude coconut oil with methanol at 200 C and sulfated zirconia gave highest catalytic activity [26]. In addition, sulfated zirconia exhibited good catalytic activity and good stability in several simultaneous transesterification of triglycerides and esterification of FFA [27e29]. However, a major concern of the use of this catalyst in the reactions is the leaching of sulfate groups by hydrolysis. It causes the deactivation of catalyst [18,27e29]. Previously, simultaneous transesterification of rapeseed oil and esterification of myristic acid with methanol using sulfated zirconia prepared by solvent-free method as a catalyst was studied [29]. The esterification of myristic acid seems to be a promising reaction for biodiesel production. However, no researcher has reported the kinetic of this reaction. There are several reports on the kinetic of FFA esterification in presence of triglyceride using different heterogeneous acid catalysts. Various models were proposed for this system either to explain the phenomenon of reaction or determine the corresponding kinetic parameters. Eley-Ridal (ER) model as represented the adsorption phenomenon, was a superior model for esterification of high FFA (approximately 50%) in the presence of soybean oil [30]. LangmuireHinshewoodeHougeneWatson (LHHW) model as characterized by increasing number of parameters was reasonable model to describe esterification of FFA in waste cooking oil [31]. A pseudo-homogeneous second order reversible (PH) model was realizable and applied for several esterification reactions of FFA in the presence of triglyceride [3,30,32,33]. In this work, the kinetic model of myristic acid esterification with methanol in presence of triglycerides using sulfated zirconia as catalyst was investigated. The effect of operating parameters namely reaction temperature, catalyst loading and molar ratio of oil
to methanol over this kinetic were also studied. The corresponding kinetic parameters were obtained.
2. Materials and methods 2.1. Materials Refined rapeseed oil was purchased from Henry Colbeck Ltd, England. The water content and free fatty acid content of this oil were determined which are 0.01 %wt and 0.3 %wt respectively. Myristic acid (purity 98%) and silver nitrate (purity 99.5%) were supplied from Fluka. Zirconium (IV) oxide chloride octahydrate (purity 99.5%), ammonium sulfate (purity 99%) and Karl e Fischer reagent were obtained from Riedel de Haen. Sulfuric acid (purity 97.5%) and heptane for HPLC analysis were purchased from SigmaeAldrich. Methanol (HPLC grade), 2- propanol (purity 99.99%), 35% ammonia solution (laboratory reagent grade) were obtained from Fisher Scientific. Barium chloride dehydrate was supplied from Acros Organics. Methyl heptadecanoate (purity 99.5%) for GC analysis from Fluka was used as an internal standard. Supelco standard 37 Component FAME mix from SigmaeAldrich was used to identify fatty acid methyl ester peaks.
2.2. Catalyst preparation of sulfated zirconia Sulfated zirconia was prepared following the solvent-free method by Sun et al. [22]. ZrOCl2.8H2O and (NH4)2SO4 in a molar ratio of 1:6 were ground in an agate mortar for 20 min at room temperature (25 C). After standing for 18 h at room temperature in air, the sample was calcined at 600 C for 5 h.
2.3. Equipment A 100 ml mini bench top reactor, model 4560 (Parr Instrument Company, Illinois, USA) was used for esterification of myristic acid with methanol in presence of triglycerides. The reactor vessel made of Type 316 stainless steel. The reactor was equipped with a magnetic stirrer drive, a four-blade impeller and thermocouple. The head assembly of reactor was also equipped with a gas inlet valve for charging N2 gas into the reactor and a gas release valve for releasing pressure. A HPLC pump K-120 was provided for injection of material into the reactor vessel with flow rate 10 ml/min. The pressure in the reactor was used to force out the liquid samples through a dip tube which was fitted with a fine filter through the sampling tube. The process conditions including the heating mantle, reaction temperature, pressure, feed rate, and impeller speed were controlled by WinISO software from H.E.L (H.E.L Inc., Lawrenceville, NJ).
2.4. Reaction procedure Initially, 30 g of rapeseed oil with 10% by weight (3.611 g) of myristic acid and catalyst were charged into the reactor. After the reactor was closed, nitrogen was used in order to purge the air left in the reactor. When the reactant and catalyst were heated to the desired reaction temperature with stirring rate at 600 rpm, nitrogen was introduced into the system to ensure that at the desired reaction temperature, the reactants are the liquid phase. The reaction was started when methanol was added into the reactor. Liquid samples were withdrawn through a sampling line at different time during a run.
D. Rattanaphra et al. / Renewable Energy 36 (2011) 2679e2686
2.5. Analysis
Table 1 List of experimental runs and operating conditions.
The withdrawn samples were allowed to settle down for overnight to separate into two layers. Myristic acid methyl ester in oil phase were analyzed by a Hewlett Packard 5890 Series II gas chromatography (GC) equipped with flame ionization detector (FID) and a fused silica capillary column (30 m 0.32 mm 0.25 mm). The GC temperature condition was oven temperature of 210 C using helium as a carrier gas, flame ionization detector temperature of 250 C and injector temperature of 250 C. A 10 mg/ml of methyl heptadecanoate solution was used as an internal standard and myristic acid methyl ester content expressed as a mass fraction in percent were calculated by comparing the peak area between the sample and methyl heptadecanoate. The peak identification was made by comparing the retention time between the sample and the standard compound. In addition, the oil phase was also determined free fatty acids content by titration of the sample in 2-propanol with 0.025 N sodium hydroxide (NaOH) using phenolphthalein as indicator. The acidity (AD) of the sample was defined as follows:
CNaOH VNaOH MMA 100ð%wtÞ m 1000
(1)
where CNaOH is the concentration of NaOH, VNaOH is the volume of NaOH, m is the sample weight, MMA is the molecular weight of myristic acid, 228.38 g mol1 and after that, the conversion of FFA can be calculated by Eq. (2):
X ¼
1
AD AD0
100
(2)
where X is the conversion of FFA and AD0 is the initial acidity before the reaction The water content was determined in both upper aqueous and lower oil phases by Karl e Fischer titration. The operating conditions of experimental runs are summarized in Table 1.
Temp. ( C)
Oil (g)
Myristic acid (g)
Methanol (g)
Catalyst (g)
Initial acidity (%wt)
Molar ratio oil/methanol
1 2 3 4 5 6 7
120 150 170 170 170 170 170
30.24 30.34 30.18 30.48 30.54 50.64 40.66
3.6158 3.6147 3.6667 3.6175 3.6151 6.0283 4.8182
21.7662 21.7662 21.7662 21.7662 21.7662 7.2554 14.5125
0.9027 0.9039 0.9124 0.3064 0.6014 1.5093 1.2078
10.013 10.010 10.019 10.018 10.011 10.017 10.006
1:20 1:20 1:20 1:20 1:20 1:4 1:10
From the relationship of CA and XA, dCA =dt ¼ CA0 dXA =dt as well as the correlations of the changes in B, E and W to A CA ¼ CA0 ð1 XA Þ, CB ¼ CA0 ðM XA Þ and CE ¼ CW ¼ CA0 XA therefore, equation (4) becomes:
h i dXA ¼ CA0 k1 ð1 XA ÞðM XA Þ k1 XA2 dt
(5)
where XA is the conversion of myristic acid, CA0 is the initial concentration of myristic acid, CB0 is the initial concentration of methanol and M is the concentration ratio of methanol to myristic C acid M ¼ B0 . CA0 At equilibrium state, dCA =dt ¼ 0 and thus the equilibrium constant ðKe Þ can be calculated from:
Ke ¼
2 XAe k1 ¼ k1 ð1 XAe ÞðM XAe Þ
(6)
XAe is the conversion of myristic acid at equilibrium stage, Ke is the equilibrium constant. Equation (6) can be rearranged to be:
dXA 1 ¼ k1 CA0 ð1 XA ÞðM XA Þ XA2 dt Ke
(7)
Solving equation (7), the integrated form for the reaction appears as following equation:
3. Mathematical model For this experimental system, the stirring rate is sufficient to overcome the diffusion limitation of reactive species e catalyst [10,12,34]. Therefore, the performance of pseudo-homogeneous model can be considered as satisfactory to correlate the esterification of myristic acid with methanol in the presence of triglycerides using sulfated zirconia as a catalyst. The esterification of myristic acid in rapeseed oil and methanol using sulfated zirconia as a catalyst is represented as follows:
k1 A + B
E + W
(3)
k-1
where, A is myristic acid, B is methanol, E is myristic acid methyl ester and W is water. This reaction can be considered to be a second order reversible reaction. Therefore, the rate equation can be written as follows:
rA
Run
dCA ¼ ¼ k1 CA CB k1 CE CW dt
2a1 XA M 1 a2 M 1 þ a2 ¼ a2 k1 CA0 t In 2a1 XA M 1 þ a2 M 1 a2
0.7 0.6 0.5 0.4 0.3 0.2 0.1 0.0 0
(4)
where CA, CB, CE and CW are the concentration of myristic acid, methanol, myristic acid methyl ester and water respectively, k1 is the forward rate constant and k1 is the backward rate constant.
(8)
where
Concentration (mol/L)
AD ¼
2681
20
40
60
80 100 120 140 160 180 200 Time (min)
Fig. 1. Change of compositions during the esterification of myristic acid with methanol in presence of triglycerides: reaction condition; reaction temperature of 443 K, catalyst loading of 3 %wt, molar ratio of oil to methanol of 1:20 and stirring rate 600 rpm: ) myristic acid methyl ester; ( ) FFA; ( ) water. (
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a1 ¼
1
1 Ke
(9)
h
a2 ¼ ðM þ 1Þ2 4a1 M
To consider the effect of catalyst loading on the forward rate constant, the following equation was applied:
k1 ¼ kcat Wcat
i1=2
(10)
M1þa2 1 XA M1a2 The result of plotting of In½ð2a 2a1 XA M1þa2 ÞðM1a2 Þ vs. t as being
a straight line through the original gives the rate constants.
(11)
where, kcat is the catalyst loading coefficient and Wcat is the ratio of catalyst weight to myristic acid. Arrhenius equation is applied to consider the effect of reaction temperature on the reaction rate that shows as the following equation:
k1 ¼ AeE=RT 100
a
(12)
where, A is the pre-exponential factor, E is the activation energy, R is the idea gas constant and T is the reaction temperature (K).
80 Conversion (%)
4. Results and discussion 60
4.1. Experimental results
40
393 K 20
423 K 443 K
0 0
100
5
10 15 Time (min)
20
25
b
Conversion (%)
80
60
40
1 %wt
20
2 %wt 3 %wt 0 0
100
5
10 15 Time (min)
20
25
0.5 Concentration of myristic acid methyl ester (mol/L)
c
Conversion (%)
80
60
40 1:4
20
1:10 1:20
0 0
The variation of compositions during the esterification of myristic acid with methanol in presence of triglycerides using sulfated zirconia as catalyst is shown in Fig. 1. It can be seen that the concentration of myristic acid methyl ester increased rapidly that related to a dramatic decrease in the FFA concentration during the first 20 min. However, there was the increasing of FFA concentration afterward. The excess water was promoted during the course of reaction. The side-reactions such as thermo/catalytic cracking, hydrolysis of triglyceride and dehydration of methanol might occur that cause the formation of extra FFA and water [29]. From the results, at the reaction time of 20 min was considered to be the equilibrium of esterification. Conversion profiles of the esterification of myristic acid at different levels of reaction temperature, catalyst loading and molar ratio of oil to methanol are represented in Fig. 2. The reaction was varied from 393 to 443 K. The reaction conditions were catalyst loading of 3 %wt, molar ratio of oil to methanol of 1:20 and stirring rate of 600 rpm. The reaction pressure was adjusted corresponding to reaction temperature to ensure that the reactants were in the liquid form. For example, at the desired reaction temperature of 443 K, pressure was controlled at 22 bars. As shown in Fig. 2a, the reaction rates increased with an increase in reaction temperature. However, the final conversion, at 20 min, was independent of reaction temperature.
5
10 15 Time (min)
20
25
Fig. 2. Conversion of FFA versus time for myristic acid methyl ester formation: (a) Effect of reaction temperature (b) Effect of catalyst loading (c) Effect of molar ratio of oil to methanol.
0.4
0.3
0.2
0.1
0.0 1
2
3 4 Number of runs
5
Fig. 3. Reusability of the sulfated zirconia in the esterification of myristic acid with methanol in presence of triglycerides.
D. Rattanaphra et al. / Renewable Energy 36 (2011) 2679e2686
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4.2. Estimation of the equilibrium constants
Table 2 Estimated equilibrium constants at different levels of reaction temperature. Run
Reaction temperature (K)
Initial AD (%wt)
Final AD (%wt)
Ke (e)
1 2 3
393 423 443
10.013 10.010 10.019
0.539 0.535 0.535
3.95 101 3.98 101 3.99 101
The equilibrium constants can be calculated from the final acidity at equilibrium (equation (6)). These values at different levels of reaction temperature, catalyst loading and molar ratio of oil to methanol are summarized in Table 2e4, respectively. The reaction temperature and catalyst loading did not affect the equilibrium constant. The slightly higher equilibrium constant was seen when the more of reaction temperature and catalyst loading were
Table 3 Estimated equilibrium constants at different levels of catalyst loading. Run 4 5 3
Catalyst loading (%wt) 1 2 3
Initial AD (%wt) 10.018 10.011 10.019
Final AD (%wt) 0.543 0.538 0.535
a
Ke (e) 1
3.92 10 3.96 101 3.99 101
The esterification of myristic acid was also carried out at catalyst loading of 1%, 2% and 3 %wt to evaluate the influence of catalyst loading. The reaction parameters were reaction temperature of 443 K, reaction pressure of 22 bars, molar ratio of oil to methanol of 1:20 and stirring rate of 600 rpm. The results are illustrated in Fig. 2b. A significant increase in the reaction rate was observed when catalyst loading doubled from 1 to 2 %wt. However, the increasing of catalyst loading further to 3 %wt did not improve the reaction rate. The variation of catalyst loading did not show any significant effect on the final conversion. The effect of molar ratio of oil to methanol on esterification of myristic acid was studied by varying the molar ratio of oil to methanol from 1:4 to 1:20. The condition parameters were reaction temperature of 443 K, reaction pressure of 22 bars, catalyst loading of 3 %wt and stirring rate of 600 rpm. The reaction rate and final conversion increased with the increasing of molar ratio of oil to methanol (Fig. 2c). It can also be noted that a rise in the molar ratio of oil to methanol from 1:4 to 1:10 resulted in significant increases in the reaction rate and finial conversion. In this work, the sulfated zirconia were reused several times under the condition parameters: reaction temperature of 443 K, reaction pressure of 22 bars, catalyst loading of 3 %wt, molar ratio of oil to methanol of 1:20, stirring rate of 600 rpm and reaction time of 180 min. The sulfated zirconia showed a minor loss of catalytic activity with the second run, and thereafter no significant loss (Fig. 3). Since this reaction system promoted significant water production through the esterification of myristic acid and the side reactions as mention above, it can cause the leaching of sulfate groups into the reaction mixture. Therefore, the leaching of sulfate groups was examined. The result showed that sulfate groups were leached out from the fresh catalyst. For spent catalysts, there was no significant leaching of sulfate groups [29]. It can be concluded that the leaching of sulfate groups lead to deactivation of fresh catalyst. After the catalyst was used several times, the coverage of the catalyst surface with organic hydrophobic components such as triglyceride, intermediate glycerides, FFA and esters might protect acidic sites from leaching [35], so the catalyst became more stable.
b
c
Table 4 Estimated equilibrium constants at different levels of molar ratio of oil to methanol. Run
Molar ratio of oil to methanol
Initial AD (%wt)
Final AD (%wt)
Ke (e)
6 7 3
1:4 1:10 1:20
10.017 10.006 10.019
1.648 0.820 0.535
5.46 101 4.99 101 3.99 101
Fig. 4. Second order reversible model for calculation the reaction rate constants at different levels of (a) reaction temperature (b) catalyst loading (c) molar ratio of oil to methanol.
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Table 5 Estimated forward and backward rate constants at different levels of reaction temperature.
Table 7 Estimated forward and backward rate constants with the different levels of molar ratio of oil to methanol.
Run
Reaction temperature (K)
Forward rate constant (k1)a
Backward rate constant (k1)a
Run
Molar ratio of oil to methanol
Forward rate constant (k1)a
Backward rate constant (k1)a
1 2 3
393 423 443
2.13 102 3.33 102 4.66 102
5.38 102 8.36 102 1.17 101
6 7 3
1:4 1:10 1:20
5.81 102 5.43 102 4.66 102
1.06 10-1 1.09 101 1.17 101
employed. The molar ratios of oil to methanol showed a significant effect on the equilibrium constant. The increasing of molar ratio of oil to methanol with the decreasing of the equilibrium constant was observed. It is probably because a higher molar ratio of oil to methanol could promote transesterification of triglyceride that can compete with the esterification of myristic acid. This might slow down the forward rate of esterification that causes a lower equilibrium constant. 4.3. Estimation of the reaction rate constants Equation (8) was applied to the experimental data at different levels of reaction temperature, catalyst loading and molar ratio of oil to methanol in order to calculate the reaction rate constant. The plots of the left-hand side of equation (8) against time are showed in Fig. 4. As can be seen from Fig. 4, the straight lines as passing through the original were obtained. All straight lines gave a good linear correlation in the range of 0.952e0.990. It is clearly that the proposed kinetic model appropriates for this reaction. The forward rate constants (k1) can be calculated from a slope of each straight line. The backward rate constants (k1) can be calculated by using the correlation of the forward rate constant and the equilibrium constant as in the equation (6). All forward and backward rate constants at the different ranges of reaction temperature, catalyst loading and molar ratio of oil to methanol are summarized in Table 5e7 respectively. The forward rate constant scaled up with the raised reaction temperature and catalyst loading (Tables 5 and 6). The increasing of forward rate constant with decreasing molar ratio of oil to methanol was observed (Table 7). This reason is likely as mentioned in the effect of molar ratio of oil to methanol on the equilibrium constant. It can be noted that all backward rate constants was higher than their forward rate constants. Moreover, higher backward rate constant was also obtained at higher reaction temperature, higher catalyst loading and higher molar ratio of oil to methanol. From previous experimental data, a rich of ester obtained from the transesterification of triglyceride and substantial water were generated simultaneously with the esterification of myristic acid. These components might promote the hydrolysis of ester. This caused the higher backward rate constant of myristic acid esterification than that of forward rate constant. In addition, when the reaction was carried out at higher reaction temperature,
Table 6 Estimated forward and backward rate constants at different levels of catalyst loading. Run
Catalyst loading (%wt)
Forward rate constant (k1)a
Backward rate constant (k1)a
4 5 3
1 2 3
2.02 102 3.58 102 4.66 102
5.14 102 9.05 102 1.17 101
a
Units of the forward rate constant and backward rate constant are L mol1 min1.
a Units of the forward rate constant and backward rate constant are L mol1 min1.
catalyst loading and molar ratio of oil to methanol, the reaction of triglyceride proceeded faster. Therefore, the higher backward rate constant was obtained.
4.4. Estimation of catalyst loading coefficient The coefficient number for describing the kinetic effect of catalyst loading for the forward rate constant (kcat) was estimated by a plot of the forward rate constant against the ratio of the catalyst weight to myristic acid. As shown in Fig. 5, the forward rate constant increased linearity with catalyst amount (equation (11)). Linear correlation coefficient of 0.991 indicated reliable linearity. The slope of this line represented kcat that was 0.162 L mol1 min1. 4.5. Activation energy The dependence of the forward rate constant on reaction temperature is described by the Arrhenius law, as given in equation (12). Fig. 6 illustrates Arrhenius-Van’t Hoff plot of the forward rate constant in the temperature range of 393e443 K. A satisfactory linear coefficient of 0.996 was obtained. Slope of the straight line in Fig. 6 can be applied to calculate the activation energy. The activation energy was 22.51 kJ mol1. The pre-exponential factors can be obtained from the intercept of y-axis in Fig. 6 that was 17.78 L mol1 min1. The relationship between temperature and the forward rate constant is given as following equations.
2707 k1 ¼ exp 2:88 T
(13)
Table 8 shows the comparison of the activation energy of this experiment and other literatures. It was found that the activation energy of this experiment was lower than that of several literatures. 0.06
0.05
0.04
k1
a Units of the forward rate constant and backward rate constant are L mol1 min1.
0.03
0.02
0.01 0.04
0.08
0.12
0.16
0.20
0.24
0.28
Wcat Fig. 5. A plot of the forward rate constant against the ratio of catalyst weight to myristic acid: reaction condition; reaction temperature of 443 K, reaction pressure of 22 bars, molar ratio of oil to methanol of 1:20 and stirring rate of 600 rpm.
D. Rattanaphra et al. / Renewable Energy 36 (2011) 2679e2686
-2.5
In k1
-3.0
-3.5
-4.0
-4.5 0.00220
0.00230
0.00240
0.00250
0.00260
1/T (K-1) Fig. 6. Arrhenius-Van’t Hoff plot for the forward rate constant in the temperature range 393e443 K: reaction condition; catalyst loading of 3 %wt, molar ratio of oil to methanol of 1:20 and stirring rate of 600 rpm.
Table 8 Comparison of the activation energy of this experiment and other literatures. Esterification of FFA
Catalyst
Temperature Activation Reference range (K) Energy (Ea) (kJ mol1)
50% oleic acid in soybean oil with methanol 56.8% oleic acid in soybean oil with methanol 38% FFA in vegetable oil with methanol 100% FFA obtained from hydrolysis of soybean oil with methanol Dodecanoic acid with 2ethylhexanol 10% myristic acid in rapeseed oil with methanol
Sulphonic acid resin
353e393
66.1
[20]
Acid ion-exchange polymeric resin
323e373
58.6
[22]
Polystyrensulphonic acid
363e393
70.3
[23]
Cation-exchange resin 333e353
59.4
[24]
Sulfated zirconia by 333e443 conventional method
55.5
[26]
Sulfated zirconia by solvent-free method
22.5
This experiment
393e443
High catalytic activity of this catalyst can greatly lower the activation energy. 5. Conclusions The kinetic of esterification of myristic acid with methanol in presence of triglycerides using sulfated zirconia prepared by solvent-free method as a catalyst was studied under different reaction temperature, catalyst loading and molar ratio of oil to methanol. A pseudo-homogeneous second order reversible model was demonstrated to justify the experimental data. The kinetic parameters were determined. This model was adequate to provide a satisfactory interpretation of the experiment results. High catalytic activity of sulfated zirconia resulted in a very low activation energy. Notations
A AD
pre-exponential factor, [L mol1 min1] acidity, [%wt]
AD0 CA CA0 CB CB0 CE CNaOH CW E k1 k1 kcat Ke m M MMA R t T VNaOH Wcat X XA XAe
2685
initial acidity before the reaction, [%wt] concentration of myristic acid, [mol L1] initial concentration of myristic acid, [mol L1] concentration of methanol, [mol L1] initial concentration of methanol, [mol L1] concentration of myristic acid methyl ester, [mol L1] concentration of NaOH, [mol L1] concentration of water, [mol L1] activation energy, [kJ mol1] forward rate constant, [L mol1 min1] backward rate constant, [L mol1 min1] catalyst loading coefficient, [L mol1 min1] equilibrium constant, [e] sample weight, [g] concentration ratio of methanol to myristic acid, [e] molecular weight of myristic acid, [g mol1] ideal gas constant, [J mol1 K1] reaction time, [min] reaction temperature, [K] volume of NaOH, [ml] ratio of catalyst weight to myristic acid, [e] conversion of FFA, [e] conversion of myristic acid, [e] conversion of myristic acid at equilibrium stage, [e]
Acknowledgements The authors would like to thank the Ministry of Science and Technology, Thailand for financial support, and Professor Allen Wright and Ms Julie Parker (both of CEAM, Newcastle University, UK) for use and training on their high temperature Parr reactor.
References [1] Kulkarni MG, Dalai AK. Waste cooking oil-an economical source for biodiesel: a review. Ind Eng Chem Res 2006;45:2901e13. [2] Bajpai D, Tyagi VK. Biodiesel: source, production, composition, properties and its benefits. J Oleo Sci 2006;55:487e502. [3] Pasias S, Barakos N, Alexopoulos C, Papayannakos N. Heterogeneously catalyzed esterification of FFAs in vegetable oils. Chem Eng Technol 2006;29:1365e71. [4] Zhang Y, Dube MA, Mclean DD, Kates M. Biodiesel production from waste cooking oil: 2. Economic assessment and sensitivity analysis. Bioresour Technol 2003;90:229e40. [5] Behzadi S, Farid MM. Review: examining the use of different feed stock for the production of biodiesel, Asia-Pac. J Chem Eng 2007;2:480e6. [6] Lopez DE, Suwannakarn K, Goodwin Jr JG, Bruce DA. Reaction kinetics and mechanism for the gas-and liquid-phase esterification of acetic acid with methanol on tungstated zirconia. Ind Eng Chem Res 2008;47:2221e30. [7] Vyas AP, Verma JL, Subrahmanyam N. A review on FAME production processes. Fuel 2010;89:1e9. [8] Peng BX, Shu Q, Wang JF, Wang GR, Wang DZ, Han MH. Biodiesel production from waste oil feedstocks by solid acid catalysis. Process Safety Environ Protect 2008;86:441e7. [9] Lopez DE, Suwannakarn K, Bruce DA, Goodwin Jr JG. Esterification and transesterification on tungstated zirconia: effect of calcination temperature. J Catal 2007;247:43e50. [10] Jacobson K, Gopinath R, Meher LC, Dalai AK. Solid acid catalyzed biodiesel production from waste cooking oil. Appl Catal B: Environ 2008;85:86e91. [11] Lam MK, Lee KT, Mohamed AR. Homogeneous, heterogeneous and enzymatic catalysis for transesterification of high free fatty acid oil (waste cooking oil) to biodiesel: a review. Biotechnol Adv B: Environ 2010;28:500e18. [12] Kulkarni MG, Gopinath R, Meher LC, Dalai AK. Solid acid catalyzed biodiesel production by simultaneous esterification and transesterification. Green Chem 2006;8:1056e62. [13] Katada N, Hatanaka T, Ota M, Yamada K, Okumura K, Niwa M. Biodiesel production using heteropoly acid-derived solid acid catalyst H4PNbW11O40/ WO3-Nb2O5. Appl Catal A: Gen 2009;363:164e8. [14] Srilatha K, Lingaiah N, Prabhavathi Devi BLA, Prasad RBN, Venkateswar S, Sai Prasad PS. Esterification of free fatty acids for biodiesel production over heteropoly tungstate supported on niobia catalysts. Appl Catal A: Gen 2009;365:28e33.
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[15] Alsalma A, Kozhevnikova EF, Kozhevnikov IV. Heteropoly acids as catalysts for liquid-phase esterification and transesterification. Appl Catal A: Gen 2008;349:170e6. [16] Chai F, Cao F, Zhai F, Chen Y, Wang X, Su Z. Transesterification of vegetable oil to biodiesel using a hetero-polyacid solid catalyst. Adv Synth Catal 2007;349:1057e65. [17] Fumin Z, Jun W, Chaoshu Y, Xiaoqian R. Catlytic performance of heteropoly compounds supported on dealuminated ultra-stable Y zeolite for liquid-phase esterification. Sci China B: Chem 2006;49:140e7. [18] Kiss AA, Dimian AC, Rothenberg G. Solid acid catalysts for biodiesel productionetowards sustainable energy. Adv Synth Catal 2006;348:75e81. [19] Helwani Z, Othman MR, Aziz N, Fernando WJN, Kim J. Technologies for production of biodiesel focusing on green catalytic techniques: a review. Fuel Process Technol 2009;90:1502e15014. [20] Lou WY, Zong MH, Duan ZQ. Efficient production of biodiesel from high free fatty acid-containing waste oils using various carbohydrate-derived solid acid catalysts. Bioresour Technol 2008;99:8752e8. [21] Yadav GD, Nair JJ. Sufated zirconia and its modified versions as promising catalysts for industrial processes. Micropor Mesopor Mater 1999;33:1e48. [22] Sun Y, Ma S, Du Y, Yuan L, Wang S, Yang J, et al. Solvent-free preparation of nanosized sulfated zirconia with Brønsted acidic sites from a simple calcination. J Phys Chem B 2005;109:2567e72. [23] Brei VV. Superacids based on zirconium dioxide. Theor Exp Chem 2005;41:165e75. [24] Ni J, Meunier FC. Esterification of free fatty acids in sunflower oil over solid acid catalysts using batch and fixed bed-reactors. Appl Catal A: Gen 2007;333:122e30. [25] R.K. Balasubramanian, J.P. Obbard. Esterification of free fatty acids using sulfated zirconia. In: A. Sayigh, editor. World Renewable Energy Congress; 2008. p. 239e42.
[26] Jitputti J, Kitiyanan B, Rangsunvigit P, Bunyakiat K, Attanatho L, Jenvanitpanjakul P. Transesterification of crude palm kernel oil and crude coconut oil by different solid catalysts. Chem Eng J 2006;116:61e6. [27] Garcia CM, Teixeira S, Marciniuk LL, Schuchardt U. Transesterification of soybean oil catalyzed by sulfated zirconia. Bioresour Technol 2008;99:6608e13. [28] Fu B, Gao L, Niu L, Wei R, Xiao G. Biodiesel from waste cooking oil via heterogeneous superacid catalyst SO42-/ZrO2. Energy Fuels 2009;23:569e72. [29] Rattanaphra D, Harvey A, Srinophakun P. Simultaneous conversion of triglyceride/free fatty acid mixtures into biodiesel using sulfated zirconia. Top Catal 2010;53:773e82. [30] Tesser R, Casale L, Verde D, Serio MD, Santacesaria E. Kinetics of free fatty acids esterification: batch and loop reactor modeling. Chem Eng J 2009;154:25e33. [31] Meng-yu G, Deng P, Li M, En Y, Jianbing H. The kinetics of the esterification of free fatty acids in waste cooking oil using Fe2(SO4)3/C catalyst. Chinese J Chem Eng 2009;17:83e7. [32] Tesser R, Serio MD, Guida M, Nastasi M, Santacesaria E. Kinetics of oleic acid esterification with methanol in the presence of triglycerides. Ind Eng Chem Res 2005;44:7978e82. [33] Su CH, Fu CC, Gomas J, Chu IM, Wu WT. A heterogeneous acid-catalyzed process for biodiesel production from enzyme hydrolyzed fatty acids. AIchE J 2008;54:327e36. [34] Srilatha K, Issariyakul T, Lingaiah N, Sai Prasad PS, Kozinski J, Dalai AK. Efficient esterification and transesterification of used cooking oil using catalyst. Energy Fuels 12-tungstophosphoric acid (TPA)/Nb2O5 2010;24:4748e55. [35] Omota F, Dimian AC, Bliek A. Fatty acid esterification by reactive distillation: Part 2- kinetics-based design for sulphated zirconia catalysts. Chem Eng Sci 2003;58:3175e85.
Contents lists available at ScienceDirect
Renewable Energy journal homepage: www.elsevier.com/locate/renene
Kinetic of myristic acid esterification with methanol in the presence of triglycerides over sulfated zirconia Dussadee Rattanaphra a, Adam P. Harvey c, Anusith Thanapimmetha a, b, Penjit Srinophakun a, b, * a
Bioprocess Laboratory, Department of Chemical Engineering, Faculty of Engineering, Kasetsart University, Bangkok 10900, Thailand Center of Excellent for Petroleum, Petrochemicals, and Advanced Materials, Department of Chemical Engineering, Faculty of Engineering, Kasetsart University, Bangkok, Thailand c School of Chemical Engineering and Advanced Materials, Newcastle University, Newcastle upon Tyne, NE1 7RU, UK b
a r t i c l e i n f o
a b s t r a c t
Article history: Received 16 November 2010 Accepted 22 February 2011 Available online 26 March 2011
The esterification of myristic acid with methanol in presence of triglycerides using sulfated zirconia prepared by solvent-free method as the heterogeneous catalyst was studied. The effects of reaction temperature (393e443 K), catalyst loading (1e3 %wt) and molar ratio of oil to methanol (1:4e1:20) on the conversion of myristic acid were investigated. The experimental data was interpreted with a second order pseudo-homogeneous kinetic model. The kinetic parameters were obtained. A good agreement between the experimental data and the model was observed. Sulfated zirconia prepared by solvent-free method exhibited high catalytic activity for this reaction. Low activation energy of 22.51 kJ mol1 was obtained in the range of temperature of 393e443 K. Ó 2011 Elsevier Ltd. All rights reserved.
Keywords: Esterification Myristic acid Triglyceride Sulfated zirconia Biodiesel
1. Introduction Biodiesel represents a promising alternative fuel for use in diesel engines due to its many advantages: it is renewable fuel, it creates low emissions of CO, SOx, NOx, it is biodegradable and non-toxic [1,2]. Biodiesel can be produced through transesterification process. Normally, refined vegetable oils such as rapeseed, soybean, sunflower and palm oils are used as biodiesel feedstock [3]. Even though these oils yield a high quality biodiesel, they are very expensive. This causes a higher cost of biodiesel as compared to the traditional petroleum diesel [4]. Moreover, an availability of these feedstocks is limited as they are used in the food industry [5]. The use of low quality oils like non-edible oil, waste grease, fats can result in a lower production cost of biodiesel [6]. However, these oils usually have high free fatty acids (FFA) concentrations for example used cooking oil and greases have free fatty acid levels greater than 6% [7]. The presence of large amounts of FFA strongly influences the performance of homogeneous base-catalyzed transesterification. These components cause the formation of soap through ponification
* Corresponding author. Bioprocess Laboratory, Department of Chemical Engineering, Faculty of Engineering, Kasetsart University, Bangkok 10900, Thailand. Tel.: þ662 942 8555; fax: þ662 561 4621. E-mail address: [email protected] (P. Srinophakun). 0960-1481/$ e see front matter Ó 2011 Elsevier Ltd. All rights reserved. doi:10.1016/j.renene.2011.02.018
which makes a serious problem of product separation [8]. Homogeneous acid catalysts have potential to simultaneously catalyze both esterification and transesterification processes of high FFA feedstock [9]. However, the use of these catalysts are non-practical for biodiesel production due to slow reaction rate, requirement of high temperature, high molar ratio of oil to alcohol, the separation of catalyst, serious environmental and corrosion related problems [10]. Recently, a twostep process is proposed to prepare biodiesel from high FFA oils. The first step is esterification of FFA with methanol which is catalyzed by homogeneous acid catalysts (generally sulfuric acid). When FFA content in oil is reduced to lower than 0.5e1 %wt, transesterification of triglyceride with methanol, the second step, is then carried out by homogeneous base catalysts (normally sodium or potassium hydroxide). This process increases the production cost of biodiesel since it requires extra separation steps to remove the catalyst in both stages [8,10,11]. Heterogeneous acid catalysts are developed to solve the problems associated with homogeneous catalysts. The main advantages of these catalysts are easy separation, reusability, low corrosion, and more environmentally friendly [12]. The heterogeneous base catalysts show high catalytic activity, however, most of them are very sensitive to trace amount of FFA and/or water in raw material [13]. Therefore, the use of heterogeneous acid catalysts is potentially attractive because of they can catalyze the transesterification of vegetable oils or animal fats with high content of FFA and water, such as deep-flying oils from restaurants and food processing
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without the formation of undesired products. These catalysts can also promote simultaneously transesterification of triglycerides and esterification of FFA that causes cheaper product cost [7]. Usage of heterogeneous acid catalysts in both transesterification and esterification reactions have been reported in several literatures. Heteropolyacids are typical strong Bronsted acids and have great potential to catalyze biodiesel synthesis. However, a major drawback of these catalysts is their relatively low thermal stability, low surface area and solubility in polar media [14]. Even though the supported heteropolyacid catalysts were developed to improve qualities, they still have problem associated with the continuous leaching of heteropolyacid from support into the reaction medium [15]. Cs-salt heteropolyacids exhibits high catalytic activity and high stability to leaching [15,16]. However, because of their insolubility, makes it is more difficult to prepare by conventional aqueous impregnation method [17]. Suflonic ion-exchange resins like Amberlyst 15 and Nafion NR50 showed higher catalytic activity in esterification than transesterification reaction [11]. These resins are expensive and have low thermal stability [18e20]. Sulfated zirconia (SO42/ZrO2) is classified as a heterogeneous superacid catalyst. It has recently received considerable attention as a promising catalyst for industrial processes due to its strong acid properties [21]. This catalyst presents both active Bronsted and Lewis acidic sites [22]. It is also higher acid strength than heteropolyacid, suflonic ion-exchange resins and other heterogeneous acid catalysts [23]. Sulfated zirconia was found to be a potential catalyst in the esterification of dodecanoic acid with 2-ethylhexanol, 1-propanol and methanol at 130e180 C, compared to the various heterogeneous catalysts such as zeolites ion-exchange resins and mixed metal oxides [18]. This catalyst has been successfully used as heterogeneous catalyst for esterification of palmitic acid in sunflower oil with methanol at 60 C [24] and esterification of high FFA in refined soybean oil with methanol at 35e60 C [25]. Six solid catalysts (ZrO2, ZnO, SO42/SnO2, SO42/ ZrO2, KNO3/KL zeolite and KNO3/ZrO2) were carried out for transesterification of crude palm kernel oil and crude coconut oil with methanol at 200 C and sulfated zirconia gave highest catalytic activity [26]. In addition, sulfated zirconia exhibited good catalytic activity and good stability in several simultaneous transesterification of triglycerides and esterification of FFA [27e29]. However, a major concern of the use of this catalyst in the reactions is the leaching of sulfate groups by hydrolysis. It causes the deactivation of catalyst [18,27e29]. Previously, simultaneous transesterification of rapeseed oil and esterification of myristic acid with methanol using sulfated zirconia prepared by solvent-free method as a catalyst was studied [29]. The esterification of myristic acid seems to be a promising reaction for biodiesel production. However, no researcher has reported the kinetic of this reaction. There are several reports on the kinetic of FFA esterification in presence of triglyceride using different heterogeneous acid catalysts. Various models were proposed for this system either to explain the phenomenon of reaction or determine the corresponding kinetic parameters. Eley-Ridal (ER) model as represented the adsorption phenomenon, was a superior model for esterification of high FFA (approximately 50%) in the presence of soybean oil [30]. LangmuireHinshewoodeHougeneWatson (LHHW) model as characterized by increasing number of parameters was reasonable model to describe esterification of FFA in waste cooking oil [31]. A pseudo-homogeneous second order reversible (PH) model was realizable and applied for several esterification reactions of FFA in the presence of triglyceride [3,30,32,33]. In this work, the kinetic model of myristic acid esterification with methanol in presence of triglycerides using sulfated zirconia as catalyst was investigated. The effect of operating parameters namely reaction temperature, catalyst loading and molar ratio of oil
to methanol over this kinetic were also studied. The corresponding kinetic parameters were obtained.
2. Materials and methods 2.1. Materials Refined rapeseed oil was purchased from Henry Colbeck Ltd, England. The water content and free fatty acid content of this oil were determined which are 0.01 %wt and 0.3 %wt respectively. Myristic acid (purity 98%) and silver nitrate (purity 99.5%) were supplied from Fluka. Zirconium (IV) oxide chloride octahydrate (purity 99.5%), ammonium sulfate (purity 99%) and Karl e Fischer reagent were obtained from Riedel de Haen. Sulfuric acid (purity 97.5%) and heptane for HPLC analysis were purchased from SigmaeAldrich. Methanol (HPLC grade), 2- propanol (purity 99.99%), 35% ammonia solution (laboratory reagent grade) were obtained from Fisher Scientific. Barium chloride dehydrate was supplied from Acros Organics. Methyl heptadecanoate (purity 99.5%) for GC analysis from Fluka was used as an internal standard. Supelco standard 37 Component FAME mix from SigmaeAldrich was used to identify fatty acid methyl ester peaks.
2.2. Catalyst preparation of sulfated zirconia Sulfated zirconia was prepared following the solvent-free method by Sun et al. [22]. ZrOCl2.8H2O and (NH4)2SO4 in a molar ratio of 1:6 were ground in an agate mortar for 20 min at room temperature (25 C). After standing for 18 h at room temperature in air, the sample was calcined at 600 C for 5 h.
2.3. Equipment A 100 ml mini bench top reactor, model 4560 (Parr Instrument Company, Illinois, USA) was used for esterification of myristic acid with methanol in presence of triglycerides. The reactor vessel made of Type 316 stainless steel. The reactor was equipped with a magnetic stirrer drive, a four-blade impeller and thermocouple. The head assembly of reactor was also equipped with a gas inlet valve for charging N2 gas into the reactor and a gas release valve for releasing pressure. A HPLC pump K-120 was provided for injection of material into the reactor vessel with flow rate 10 ml/min. The pressure in the reactor was used to force out the liquid samples through a dip tube which was fitted with a fine filter through the sampling tube. The process conditions including the heating mantle, reaction temperature, pressure, feed rate, and impeller speed were controlled by WinISO software from H.E.L (H.E.L Inc., Lawrenceville, NJ).
2.4. Reaction procedure Initially, 30 g of rapeseed oil with 10% by weight (3.611 g) of myristic acid and catalyst were charged into the reactor. After the reactor was closed, nitrogen was used in order to purge the air left in the reactor. When the reactant and catalyst were heated to the desired reaction temperature with stirring rate at 600 rpm, nitrogen was introduced into the system to ensure that at the desired reaction temperature, the reactants are the liquid phase. The reaction was started when methanol was added into the reactor. Liquid samples were withdrawn through a sampling line at different time during a run.
D. Rattanaphra et al. / Renewable Energy 36 (2011) 2679e2686
2.5. Analysis
Table 1 List of experimental runs and operating conditions.
The withdrawn samples were allowed to settle down for overnight to separate into two layers. Myristic acid methyl ester in oil phase were analyzed by a Hewlett Packard 5890 Series II gas chromatography (GC) equipped with flame ionization detector (FID) and a fused silica capillary column (30 m 0.32 mm 0.25 mm). The GC temperature condition was oven temperature of 210 C using helium as a carrier gas, flame ionization detector temperature of 250 C and injector temperature of 250 C. A 10 mg/ml of methyl heptadecanoate solution was used as an internal standard and myristic acid methyl ester content expressed as a mass fraction in percent were calculated by comparing the peak area between the sample and methyl heptadecanoate. The peak identification was made by comparing the retention time between the sample and the standard compound. In addition, the oil phase was also determined free fatty acids content by titration of the sample in 2-propanol with 0.025 N sodium hydroxide (NaOH) using phenolphthalein as indicator. The acidity (AD) of the sample was defined as follows:
CNaOH VNaOH MMA 100ð%wtÞ m 1000
(1)
where CNaOH is the concentration of NaOH, VNaOH is the volume of NaOH, m is the sample weight, MMA is the molecular weight of myristic acid, 228.38 g mol1 and after that, the conversion of FFA can be calculated by Eq. (2):
X ¼
1
AD AD0
100
(2)
where X is the conversion of FFA and AD0 is the initial acidity before the reaction The water content was determined in both upper aqueous and lower oil phases by Karl e Fischer titration. The operating conditions of experimental runs are summarized in Table 1.
Temp. ( C)
Oil (g)
Myristic acid (g)
Methanol (g)
Catalyst (g)
Initial acidity (%wt)
Molar ratio oil/methanol
1 2 3 4 5 6 7
120 150 170 170 170 170 170
30.24 30.34 30.18 30.48 30.54 50.64 40.66
3.6158 3.6147 3.6667 3.6175 3.6151 6.0283 4.8182
21.7662 21.7662 21.7662 21.7662 21.7662 7.2554 14.5125
0.9027 0.9039 0.9124 0.3064 0.6014 1.5093 1.2078
10.013 10.010 10.019 10.018 10.011 10.017 10.006
1:20 1:20 1:20 1:20 1:20 1:4 1:10
From the relationship of CA and XA, dCA =dt ¼ CA0 dXA =dt as well as the correlations of the changes in B, E and W to A CA ¼ CA0 ð1 XA Þ, CB ¼ CA0 ðM XA Þ and CE ¼ CW ¼ CA0 XA therefore, equation (4) becomes:
h i dXA ¼ CA0 k1 ð1 XA ÞðM XA Þ k1 XA2 dt
(5)
where XA is the conversion of myristic acid, CA0 is the initial concentration of myristic acid, CB0 is the initial concentration of methanol and M is the concentration ratio of methanol to myristic C acid M ¼ B0 . CA0 At equilibrium state, dCA =dt ¼ 0 and thus the equilibrium constant ðKe Þ can be calculated from:
Ke ¼
2 XAe k1 ¼ k1 ð1 XAe ÞðM XAe Þ
(6)
XAe is the conversion of myristic acid at equilibrium stage, Ke is the equilibrium constant. Equation (6) can be rearranged to be:
dXA 1 ¼ k1 CA0 ð1 XA ÞðM XA Þ XA2 dt Ke
(7)
Solving equation (7), the integrated form for the reaction appears as following equation:
3. Mathematical model For this experimental system, the stirring rate is sufficient to overcome the diffusion limitation of reactive species e catalyst [10,12,34]. Therefore, the performance of pseudo-homogeneous model can be considered as satisfactory to correlate the esterification of myristic acid with methanol in the presence of triglycerides using sulfated zirconia as a catalyst. The esterification of myristic acid in rapeseed oil and methanol using sulfated zirconia as a catalyst is represented as follows:
k1 A + B
E + W
(3)
k-1
where, A is myristic acid, B is methanol, E is myristic acid methyl ester and W is water. This reaction can be considered to be a second order reversible reaction. Therefore, the rate equation can be written as follows:
rA
Run
dCA ¼ ¼ k1 CA CB k1 CE CW dt
2a1 XA M 1 a2 M 1 þ a2 ¼ a2 k1 CA0 t In 2a1 XA M 1 þ a2 M 1 a2
0.7 0.6 0.5 0.4 0.3 0.2 0.1 0.0 0
(4)
where CA, CB, CE and CW are the concentration of myristic acid, methanol, myristic acid methyl ester and water respectively, k1 is the forward rate constant and k1 is the backward rate constant.
(8)
where
Concentration (mol/L)
AD ¼
2681
20
40
60
80 100 120 140 160 180 200 Time (min)
Fig. 1. Change of compositions during the esterification of myristic acid with methanol in presence of triglycerides: reaction condition; reaction temperature of 443 K, catalyst loading of 3 %wt, molar ratio of oil to methanol of 1:20 and stirring rate 600 rpm: ) myristic acid methyl ester; ( ) FFA; ( ) water. (
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a1 ¼
1
1 Ke
(9)
h
a2 ¼ ðM þ 1Þ2 4a1 M
To consider the effect of catalyst loading on the forward rate constant, the following equation was applied:
k1 ¼ kcat Wcat
i1=2
(10)
M1þa2 1 XA M1a2 The result of plotting of In½ð2a 2a1 XA M1þa2 ÞðM1a2 Þ vs. t as being
a straight line through the original gives the rate constants.
(11)
where, kcat is the catalyst loading coefficient and Wcat is the ratio of catalyst weight to myristic acid. Arrhenius equation is applied to consider the effect of reaction temperature on the reaction rate that shows as the following equation:
k1 ¼ AeE=RT 100
a
(12)
where, A is the pre-exponential factor, E is the activation energy, R is the idea gas constant and T is the reaction temperature (K).
80 Conversion (%)
4. Results and discussion 60
4.1. Experimental results
40
393 K 20
423 K 443 K
0 0
100
5
10 15 Time (min)
20
25
b
Conversion (%)
80
60
40
1 %wt
20
2 %wt 3 %wt 0 0
100
5
10 15 Time (min)
20
25
0.5 Concentration of myristic acid methyl ester (mol/L)
c
Conversion (%)
80
60
40 1:4
20
1:10 1:20
0 0
The variation of compositions during the esterification of myristic acid with methanol in presence of triglycerides using sulfated zirconia as catalyst is shown in Fig. 1. It can be seen that the concentration of myristic acid methyl ester increased rapidly that related to a dramatic decrease in the FFA concentration during the first 20 min. However, there was the increasing of FFA concentration afterward. The excess water was promoted during the course of reaction. The side-reactions such as thermo/catalytic cracking, hydrolysis of triglyceride and dehydration of methanol might occur that cause the formation of extra FFA and water [29]. From the results, at the reaction time of 20 min was considered to be the equilibrium of esterification. Conversion profiles of the esterification of myristic acid at different levels of reaction temperature, catalyst loading and molar ratio of oil to methanol are represented in Fig. 2. The reaction was varied from 393 to 443 K. The reaction conditions were catalyst loading of 3 %wt, molar ratio of oil to methanol of 1:20 and stirring rate of 600 rpm. The reaction pressure was adjusted corresponding to reaction temperature to ensure that the reactants were in the liquid form. For example, at the desired reaction temperature of 443 K, pressure was controlled at 22 bars. As shown in Fig. 2a, the reaction rates increased with an increase in reaction temperature. However, the final conversion, at 20 min, was independent of reaction temperature.
5
10 15 Time (min)
20
25
Fig. 2. Conversion of FFA versus time for myristic acid methyl ester formation: (a) Effect of reaction temperature (b) Effect of catalyst loading (c) Effect of molar ratio of oil to methanol.
0.4
0.3
0.2
0.1
0.0 1
2
3 4 Number of runs
5
Fig. 3. Reusability of the sulfated zirconia in the esterification of myristic acid with methanol in presence of triglycerides.
D. Rattanaphra et al. / Renewable Energy 36 (2011) 2679e2686
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4.2. Estimation of the equilibrium constants
Table 2 Estimated equilibrium constants at different levels of reaction temperature. Run
Reaction temperature (K)
Initial AD (%wt)
Final AD (%wt)
Ke (e)
1 2 3
393 423 443
10.013 10.010 10.019
0.539 0.535 0.535
3.95 101 3.98 101 3.99 101
The equilibrium constants can be calculated from the final acidity at equilibrium (equation (6)). These values at different levels of reaction temperature, catalyst loading and molar ratio of oil to methanol are summarized in Table 2e4, respectively. The reaction temperature and catalyst loading did not affect the equilibrium constant. The slightly higher equilibrium constant was seen when the more of reaction temperature and catalyst loading were
Table 3 Estimated equilibrium constants at different levels of catalyst loading. Run 4 5 3
Catalyst loading (%wt) 1 2 3
Initial AD (%wt) 10.018 10.011 10.019
Final AD (%wt) 0.543 0.538 0.535
a
Ke (e) 1
3.92 10 3.96 101 3.99 101
The esterification of myristic acid was also carried out at catalyst loading of 1%, 2% and 3 %wt to evaluate the influence of catalyst loading. The reaction parameters were reaction temperature of 443 K, reaction pressure of 22 bars, molar ratio of oil to methanol of 1:20 and stirring rate of 600 rpm. The results are illustrated in Fig. 2b. A significant increase in the reaction rate was observed when catalyst loading doubled from 1 to 2 %wt. However, the increasing of catalyst loading further to 3 %wt did not improve the reaction rate. The variation of catalyst loading did not show any significant effect on the final conversion. The effect of molar ratio of oil to methanol on esterification of myristic acid was studied by varying the molar ratio of oil to methanol from 1:4 to 1:20. The condition parameters were reaction temperature of 443 K, reaction pressure of 22 bars, catalyst loading of 3 %wt and stirring rate of 600 rpm. The reaction rate and final conversion increased with the increasing of molar ratio of oil to methanol (Fig. 2c). It can also be noted that a rise in the molar ratio of oil to methanol from 1:4 to 1:10 resulted in significant increases in the reaction rate and finial conversion. In this work, the sulfated zirconia were reused several times under the condition parameters: reaction temperature of 443 K, reaction pressure of 22 bars, catalyst loading of 3 %wt, molar ratio of oil to methanol of 1:20, stirring rate of 600 rpm and reaction time of 180 min. The sulfated zirconia showed a minor loss of catalytic activity with the second run, and thereafter no significant loss (Fig. 3). Since this reaction system promoted significant water production through the esterification of myristic acid and the side reactions as mention above, it can cause the leaching of sulfate groups into the reaction mixture. Therefore, the leaching of sulfate groups was examined. The result showed that sulfate groups were leached out from the fresh catalyst. For spent catalysts, there was no significant leaching of sulfate groups [29]. It can be concluded that the leaching of sulfate groups lead to deactivation of fresh catalyst. After the catalyst was used several times, the coverage of the catalyst surface with organic hydrophobic components such as triglyceride, intermediate glycerides, FFA and esters might protect acidic sites from leaching [35], so the catalyst became more stable.
b
c
Table 4 Estimated equilibrium constants at different levels of molar ratio of oil to methanol. Run
Molar ratio of oil to methanol
Initial AD (%wt)
Final AD (%wt)
Ke (e)
6 7 3
1:4 1:10 1:20
10.017 10.006 10.019
1.648 0.820 0.535
5.46 101 4.99 101 3.99 101
Fig. 4. Second order reversible model for calculation the reaction rate constants at different levels of (a) reaction temperature (b) catalyst loading (c) molar ratio of oil to methanol.
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Table 5 Estimated forward and backward rate constants at different levels of reaction temperature.
Table 7 Estimated forward and backward rate constants with the different levels of molar ratio of oil to methanol.
Run
Reaction temperature (K)
Forward rate constant (k1)a
Backward rate constant (k1)a
Run
Molar ratio of oil to methanol
Forward rate constant (k1)a
Backward rate constant (k1)a
1 2 3
393 423 443
2.13 102 3.33 102 4.66 102
5.38 102 8.36 102 1.17 101
6 7 3
1:4 1:10 1:20
5.81 102 5.43 102 4.66 102
1.06 10-1 1.09 101 1.17 101
employed. The molar ratios of oil to methanol showed a significant effect on the equilibrium constant. The increasing of molar ratio of oil to methanol with the decreasing of the equilibrium constant was observed. It is probably because a higher molar ratio of oil to methanol could promote transesterification of triglyceride that can compete with the esterification of myristic acid. This might slow down the forward rate of esterification that causes a lower equilibrium constant. 4.3. Estimation of the reaction rate constants Equation (8) was applied to the experimental data at different levels of reaction temperature, catalyst loading and molar ratio of oil to methanol in order to calculate the reaction rate constant. The plots of the left-hand side of equation (8) against time are showed in Fig. 4. As can be seen from Fig. 4, the straight lines as passing through the original were obtained. All straight lines gave a good linear correlation in the range of 0.952e0.990. It is clearly that the proposed kinetic model appropriates for this reaction. The forward rate constants (k1) can be calculated from a slope of each straight line. The backward rate constants (k1) can be calculated by using the correlation of the forward rate constant and the equilibrium constant as in the equation (6). All forward and backward rate constants at the different ranges of reaction temperature, catalyst loading and molar ratio of oil to methanol are summarized in Table 5e7 respectively. The forward rate constant scaled up with the raised reaction temperature and catalyst loading (Tables 5 and 6). The increasing of forward rate constant with decreasing molar ratio of oil to methanol was observed (Table 7). This reason is likely as mentioned in the effect of molar ratio of oil to methanol on the equilibrium constant. It can be noted that all backward rate constants was higher than their forward rate constants. Moreover, higher backward rate constant was also obtained at higher reaction temperature, higher catalyst loading and higher molar ratio of oil to methanol. From previous experimental data, a rich of ester obtained from the transesterification of triglyceride and substantial water were generated simultaneously with the esterification of myristic acid. These components might promote the hydrolysis of ester. This caused the higher backward rate constant of myristic acid esterification than that of forward rate constant. In addition, when the reaction was carried out at higher reaction temperature,
Table 6 Estimated forward and backward rate constants at different levels of catalyst loading. Run
Catalyst loading (%wt)
Forward rate constant (k1)a
Backward rate constant (k1)a
4 5 3
1 2 3
2.02 102 3.58 102 4.66 102
5.14 102 9.05 102 1.17 101
a
Units of the forward rate constant and backward rate constant are L mol1 min1.
a Units of the forward rate constant and backward rate constant are L mol1 min1.
catalyst loading and molar ratio of oil to methanol, the reaction of triglyceride proceeded faster. Therefore, the higher backward rate constant was obtained.
4.4. Estimation of catalyst loading coefficient The coefficient number for describing the kinetic effect of catalyst loading for the forward rate constant (kcat) was estimated by a plot of the forward rate constant against the ratio of the catalyst weight to myristic acid. As shown in Fig. 5, the forward rate constant increased linearity with catalyst amount (equation (11)). Linear correlation coefficient of 0.991 indicated reliable linearity. The slope of this line represented kcat that was 0.162 L mol1 min1. 4.5. Activation energy The dependence of the forward rate constant on reaction temperature is described by the Arrhenius law, as given in equation (12). Fig. 6 illustrates Arrhenius-Van’t Hoff plot of the forward rate constant in the temperature range of 393e443 K. A satisfactory linear coefficient of 0.996 was obtained. Slope of the straight line in Fig. 6 can be applied to calculate the activation energy. The activation energy was 22.51 kJ mol1. The pre-exponential factors can be obtained from the intercept of y-axis in Fig. 6 that was 17.78 L mol1 min1. The relationship between temperature and the forward rate constant is given as following equations.
2707 k1 ¼ exp 2:88 T
(13)
Table 8 shows the comparison of the activation energy of this experiment and other literatures. It was found that the activation energy of this experiment was lower than that of several literatures. 0.06
0.05
0.04
k1
a Units of the forward rate constant and backward rate constant are L mol1 min1.
0.03
0.02
0.01 0.04
0.08
0.12
0.16
0.20
0.24
0.28
Wcat Fig. 5. A plot of the forward rate constant against the ratio of catalyst weight to myristic acid: reaction condition; reaction temperature of 443 K, reaction pressure of 22 bars, molar ratio of oil to methanol of 1:20 and stirring rate of 600 rpm.
D. Rattanaphra et al. / Renewable Energy 36 (2011) 2679e2686
-2.5
In k1
-3.0
-3.5
-4.0
-4.5 0.00220
0.00230
0.00240
0.00250
0.00260
1/T (K-1) Fig. 6. Arrhenius-Van’t Hoff plot for the forward rate constant in the temperature range 393e443 K: reaction condition; catalyst loading of 3 %wt, molar ratio of oil to methanol of 1:20 and stirring rate of 600 rpm.
Table 8 Comparison of the activation energy of this experiment and other literatures. Esterification of FFA
Catalyst
Temperature Activation Reference range (K) Energy (Ea) (kJ mol1)
50% oleic acid in soybean oil with methanol 56.8% oleic acid in soybean oil with methanol 38% FFA in vegetable oil with methanol 100% FFA obtained from hydrolysis of soybean oil with methanol Dodecanoic acid with 2ethylhexanol 10% myristic acid in rapeseed oil with methanol
Sulphonic acid resin
353e393
66.1
[20]
Acid ion-exchange polymeric resin
323e373
58.6
[22]
Polystyrensulphonic acid
363e393
70.3
[23]
Cation-exchange resin 333e353
59.4
[24]
Sulfated zirconia by 333e443 conventional method
55.5
[26]
Sulfated zirconia by solvent-free method
22.5
This experiment
393e443
High catalytic activity of this catalyst can greatly lower the activation energy. 5. Conclusions The kinetic of esterification of myristic acid with methanol in presence of triglycerides using sulfated zirconia prepared by solvent-free method as a catalyst was studied under different reaction temperature, catalyst loading and molar ratio of oil to methanol. A pseudo-homogeneous second order reversible model was demonstrated to justify the experimental data. The kinetic parameters were determined. This model was adequate to provide a satisfactory interpretation of the experiment results. High catalytic activity of sulfated zirconia resulted in a very low activation energy. Notations
A AD
pre-exponential factor, [L mol1 min1] acidity, [%wt]
AD0 CA CA0 CB CB0 CE CNaOH CW E k1 k1 kcat Ke m M MMA R t T VNaOH Wcat X XA XAe
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initial acidity before the reaction, [%wt] concentration of myristic acid, [mol L1] initial concentration of myristic acid, [mol L1] concentration of methanol, [mol L1] initial concentration of methanol, [mol L1] concentration of myristic acid methyl ester, [mol L1] concentration of NaOH, [mol L1] concentration of water, [mol L1] activation energy, [kJ mol1] forward rate constant, [L mol1 min1] backward rate constant, [L mol1 min1] catalyst loading coefficient, [L mol1 min1] equilibrium constant, [e] sample weight, [g] concentration ratio of methanol to myristic acid, [e] molecular weight of myristic acid, [g mol1] ideal gas constant, [J mol1 K1] reaction time, [min] reaction temperature, [K] volume of NaOH, [ml] ratio of catalyst weight to myristic acid, [e] conversion of FFA, [e] conversion of myristic acid, [e] conversion of myristic acid at equilibrium stage, [e]
Acknowledgements The authors would like to thank the Ministry of Science and Technology, Thailand for financial support, and Professor Allen Wright and Ms Julie Parker (both of CEAM, Newcastle University, UK) for use and training on their high temperature Parr reactor.
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