**Recent Developments on the Mechanism and Kinetics of Esterification Reaction Promoted by Various Catalysts**

Zuoxiang Zeng, Li Cui, Weilan Xue, Jing Chen and Yu Che *Institute of Chemical Engineering, East China University of Science and Technology, P.R.China* 

#### **1. Introduction**

Esters have played a significant role in daily living and chemical industry, such as plasticizers, fragrance, adhesive and lubricants (Joseph et al., 2005; Mbaraka & Shanks, 2006; Krause et al., 2009; Martínez et al., 2011). The vast majority of esters can be prepared using esterification reaction in the chemical engineering industry. Esterification has acquired further improvement from the engineering side; this mainly depends on the research of esterification kinetics. On the other hand, the need to control chemical reactions at the molecular level, which depends critically on the catalytic mechanism, is rapidly increasing (Salciccioli et al., 2011).

In the recent years, various esterification mechanisms were proposed for homogeneous and heterogeneous systems. Base on the mechanism, numerous kinetic models have been developed to represent the kinetic behaviors of esterification, such as simple orders or the power-law model, the pseudo-homogeneous model, the L-H model, the E-R model, etc.

Herein, we review the mechanisms of esterification catalyzed by inorganic acid, Lewis acid, metallic compounds, solid acids, ion-exchange resin, respectively. Meanwhile, the kinetics for each esterification is also conducted in detail. The study on the mechanism of various catalysts is not only useful in the present applications but it is required to enhance the scope of their applications. Therefore, it is desired to research and develop more efficient catalysts for high yield ester production, under mild reaction conditions. The esterification kinetics is important for simulation and design of a reactor.

A literature survey more than 90 relevant references has been done in this chapter. The studies on esterification are mainly focused on the following aspects: (1) establishing the adequate reaction mechanism according to the research papers, and (2) evaluating the available esterification kinetics.

#### **2. Esterification mechanism and kinetic model**

#### **2.1 Inorganic acids**

Inorganic acids, such as H2SO4, HCl and HI, are still being widely applied in the chemical industry (Liu et al., 2006; Schmitt et al., 2008). Under this condition, they are used as

Recent Developments on the Mechanism and Kinetics

alcohol have a significant effect on the easy of esterification.

OH

HO CH3

Fig. 2. The proton transfers step

reaction is:

carboxylic acids.

H3C C

O

Fig. 3. The simplified mechanism

hypothesis on the rapid step(6):

The reaction rate can therefore be expressed as:

OH

of Esterification Reaction Promoted by Various Catalysts 257

OH

OH

OH2

O CH3

O

CH3OH

<sup>3</sup> <sup>3</sup> 2 2 *A CH OH E* <sup>2</sup> *H O r r kc c k cc* (2)

OCH H3O

H2O H3O H3C C

O CH3

In addition, Streitweiser considered that step (2) is the rate-determining stage. The total

CH3COOH + CH3OH CH3COOCH3 + H2O (1)

The result of the reaction is that the rupture of carboxylic acid acyl-oxygen bond occurs, and the hydroxy of acid is replaced by alkoxy, which is a nucleophilic substitution process of

Esterifications of primary alcohols, secondary alcohols with carboxylic acids comply with this mechanism usually. According to this mechanism, the intermediate with a tetrahedral structure is more crowd than reactants, which makes the structure of carboxylic acid and

The activity of different structures of the alcohol and carboxylic acid in esterification obeys the following order: Alcohol: CH3OH > RCH2OH > R2CHOH; Acid: CH3COOH > RCH2COOH > R2CHCOOH > R3CCOOH (Solomons, 1986; March, 1992; TianQuan 1992; Fei & Zhao, 2009).It shows that straight-chain structure is easier than branched- chain structure

Based on the above mechanism proposed by Streitweiser, Rönnback, et al. developed a esterification kinetic model of carboxylic acid with methanol in the presence of hydrogen iodide though isothermal batch experiments at 30-60℃. The catalyst concentration varied from 0.05 to 10.0 wt%. Because the proton-donation step (1) as well as the subsequent steps

OH H3O H2O H3C C

OH

H3C C +

OH

where *r* is the total rate, *r*2 is the rate of step(2), A is CH3C(OH)2+ and E is CH3COOCH3.

The concentration of the intermediate A is obtained after applying the quasi-equilibrium

to esterification, and the more the branched-chain is, the lower the rate will be.

(3)-(5) is assumed to be rapid, the simplified mechanism shown in Figure 3:

H3C C H2O

OH

H3C C

homogenous catalysts for cheap price, high activity and efficiency (greater density of acid sites per gram) (Sanz et al., 2002), stable performance. There is no doubt that mineral acidcatalyzed esterifications have been the subjects of extensive studies because of their long history (Liu et al., 2006; Vahteristo et al., 2008). For instance, it can be mentioned that Berthelot and Saint-Gilles studied the equilibrium of acetic acid with ethanol in 1862(Huiping & Xiaohua, 2006).

Many researchers have conducted a number of experiments to study the mechanism and kinetics of esterifications promoted by mineral acids (Cardoso et al., 2008). Summarizing recent research and development, esterification is an additive-eliminate course in the presence of mineral acid catalyst, generally complying with the following four typical and mature mechanisms.

#### **2.1.1 Bimolecular reaction**

In 1895, Emil Fischer discovered that esters are formed simply by heating a carboxylic acid in alcohol solution containing a small amount of strong acid catalyst (Zhang et al., 1995; Rönnback et al., 1997; Otera & Nishikido, 2009). The following mechanism, which includes five elementary reactions, was proposed:


The chemical reaction can be expressed as Figure 1.

Fig. 1. Mechanism of bimolecular reaction

Streitwieser explained how the proton transfers specifically in 1985(Streitwieser et al., 1985; Liu et al., 2006; Fei & Zhao, 2009). The above chemical step (3) can be written as Figure 2:

Fig. 2. The proton transfers step

homogenous catalysts for cheap price, high activity and efficiency (greater density of acid sites per gram) (Sanz et al., 2002), stable performance. There is no doubt that mineral acidcatalyzed esterifications have been the subjects of extensive studies because of their long history (Liu et al., 2006; Vahteristo et al., 2008). For instance, it can be mentioned that Berthelot and Saint-Gilles studied the equilibrium of acetic acid with ethanol in

Many researchers have conducted a number of experiments to study the mechanism and kinetics of esterifications promoted by mineral acids (Cardoso et al., 2008). Summarizing recent research and development, esterification is an additive-eliminate course in the presence of mineral acid catalyst, generally complying with the following four typical and

In 1895, Emil Fischer discovered that esters are formed simply by heating a carboxylic acid in alcohol solution containing a small amount of strong acid catalyst (Zhang et al., 1995; Rönnback et al., 1997; Otera & Nishikido, 2009). The following mechanism, which includes

1. Ketonic oxygen is easy to capture hydrogen ion to form protonated carbonyl substrate. 2. The substrate makes carbon atom of carboxyl possess higher electropositive, which is conducive to the attack of nucleophile (RCOH) to generate a tetrahedral intermediate. 3. Transfer of a proton from one oxygen atom to another yields a second tetrahedral

OH H2O H3C C

1 2

OCH H3C C

Streitwieser explained how the proton transfers specifically in 1985(Streitwieser et al., 1985; Liu et al., 2006; Fei & Zhao, 2009). The above chemical step (3) can be written as Figure 2:

OH

OH

OH2

O CH3

OH CH3OH H3C C

OH

OH

H3C C

OH

OH2

O CH3

HO CH3

intermediate and converts the OH group into a good leaving group.

4. The acyl-oxygen bond disconnects and one water molecule lost.

OH

H3C C

OH

5

OH

1862(Huiping & Xiaohua, 2006).

mature mechanisms.

H3C C

H3C C

O

O

**2.1.1 Bimolecular reaction** 

five elementary reactions, was proposed:

5. Loss of a proton gives the ester product.

The chemical reaction can be expressed as Figure 1.

H3O H3C C

H3O OCH H2O

Fig. 1. Mechanism of bimolecular reaction

In addition, Streitweiser considered that step (2) is the rate-determining stage. The total reaction is:

$$\text{CH}\_3\text{COOH} + \text{CH}\_3\text{OH} \xrightleftharpoons \text{CH}\_3\text{COOCH}\_3 + \text{H}\_2\text{O} \tag{1}$$

The result of the reaction is that the rupture of carboxylic acid acyl-oxygen bond occurs, and the hydroxy of acid is replaced by alkoxy, which is a nucleophilic substitution process of carboxylic acids.

Esterifications of primary alcohols, secondary alcohols with carboxylic acids comply with this mechanism usually. According to this mechanism, the intermediate with a tetrahedral structure is more crowd than reactants, which makes the structure of carboxylic acid and alcohol have a significant effect on the easy of esterification.

The activity of different structures of the alcohol and carboxylic acid in esterification obeys the following order: Alcohol: CH3OH > RCH2OH > R2CHOH; Acid: CH3COOH > RCH2COOH > R2CHCOOH > R3CCOOH (Solomons, 1986; March, 1992; TianQuan 1992; Fei & Zhao, 2009).It shows that straight-chain structure is easier than branched- chain structure to esterification, and the more the branched-chain is, the lower the rate will be.

Based on the above mechanism proposed by Streitweiser, Rönnback, et al. developed a esterification kinetic model of carboxylic acid with methanol in the presence of hydrogen iodide though isothermal batch experiments at 30-60℃. The catalyst concentration varied from 0.05 to 10.0 wt%. Because the proton-donation step (1) as well as the subsequent steps (3)-(5) is assumed to be rapid, the simplified mechanism shown in Figure 3:

$$\left\| \underbrace{\cdots \atop \cdots}\_{\mathsf{G}\_{\mathsf{G}\mathsf{H}} +\_{\ell} \mathsf{H} \mathsf{O} \xleftarrow{\cdots} \mathsf{G}\_{\mathsf{G}\mathsf{H}} \xrightarrow{\mathsf{H}^{\mathsf{G}}} \mathsf{H} \mathsf{O}} \underbrace{\mathsf{H}^{\mathsf{G}}}\_{\mathsf{(\(\mathsf{G}\}} \mathsf{H} \xleftarrow{\cdots} \mathsf{G}^{\mathsf{G}}} \|\_{\mathsf{G}^{\mathsf{H}} + \mathsf{O}^{\mathsf{H}} \xleftarrow{\cdots} \mathsf{G}^{\mathsf{G}} \xrightarrow{\mathsf{H}^{\mathsf{G}}} \mathsf{O}} \|\_{\mathsf{G}^{\mathsf{G}}}$$

Fig. 3. The simplified mechanism

The reaction rate can therefore be expressed as:

$$r = r\_2 = k\_2 c\_A c\_{CH\_3OH} - k\_{-2} c\_E c\_{H\_3O^+} \tag{2}$$

where *r* is the total rate, *r*2 is the rate of step(2), A is CH3C(OH)2 + and E is CH3COOCH3.

The concentration of the intermediate A is obtained after applying the quasi-equilibrium hypothesis on the rapid step(6):

Recent Developments on the Mechanism and Kinetics

CH3

H2SO4

CH3

H3C

CH3

Fig. 5. Mechanism of large-steric reaction

C

O

H3C

OH

CH3

C

O

H3C

H2O

H3C

**2.1.4 Sulfuric acid** 

mechanism.

of Esterification Reaction Promoted by Various Catalysts 259

CH3

C

O

CH3

Dhanuka showed that the mechanism of sulfuric acid-catalyzed esterification is the model of alkyl sulfuric intermediate. It's a two-step reaction. R′-O-SO3H firstly is generated via sulfuric acid with alcohol, which is treated as a catalyst in the second step (Lide, 1994):

The alkylation reaction of sulfuric acid with alcohol is irreversible (Aranda et al.,

where *k* is rate constant and *K* is equilibrium constant; *C*A, *C*B, *C*W, *C*E is the concentration of

In 2009, Fei, *et al.* studied on the esterification behaviors of neo-polyhydric alcohols with fatty acids catalyzed by sulfuric acid. It was found that the esterification follows the above

RCOOH + R′OH R'OSO3

2008).Therefore the reaction rate can be expressed as:

carboxylic acid, alcohol, water and ester, respectively.

C

O

CH3

HSO4

CH3OH2

R′OH + H2SO4 R′OSO3 + H2O (6)

RCOOR′ + H2O (7)

*r kC C C C* ( ) *AB W E* /*K* (8)

CH3

H3C

CH3

OH2

CH3

H

OCH3

CH3

CH3

C

O

C

O

H3C

OH2

$$K\_6 = K\_1 = \mathbf{c}\_A \cdot \mathbf{c}\_{H\_2O} \;/\; \mathbf{c}\_{CH\_3COOH} \cdot \mathbf{c}\_{H\_3O^+} \tag{3}$$

The rate equation is rewritten as follows,

$$\mathbf{r} = \mathbf{r}\_2 = K\_1 k\_2 \left( \frac{\mathbf{c}\_{H\_3O^+}}{\mathbf{c}\_{H\_2O}} \right) \left[ \mathbf{c}\_{CH\_3COOH} \cdot \mathbf{c}\_{CH\_3OH} - \frac{\mathbf{c}\_E \cdot \mathbf{c}\_{H\_2O}}{K\_1 \left(k\_2 / k\_{-2}\right)} \right] \tag{4}$$

The product *K*l (*k*2/*k*-2) equals the equilibrium constant of the overall reaction *K*E2. In addition, the product *K*l*k*2 is denoted by a lumped constant *k*<sup>2</sup> '. Thus, we obtain for *r*2,

$$r = r\_2 = k\_2 \left(\frac{c\_{H\_3O^+}}{c\_{H\_2O}}\right) \left(c\_{CH\_3COOH} \cdot c\_{CH\_3OH} - \frac{c\_E \cdot c\_{H\_2O}}{K\_{E2}}\right) \tag{5}$$

The kinetic and equilibrium parameters included in Equation (5) are estimated from experimental data with regression analysis. Simulation of the model with the estimated parameters revealed that Equation (5) can predict the experimental trends in the acidcatalyzed esterification correctly.

#### **2.1.2 Single-molecule reaction**

The tertiary alcohol is prone to generate carbocation under acidic condition. Therefore, the mechanism of tert- esterification is different from that of the primary and secondary alcohol, it follows the single-molecule reaction process(shown in Figure 4):The tertiary alcohol combines with protons to generate protonated alcohol(1); The protonated alcohol is got rid of a molecule water to produce tert-carbocation(2), which is very stable; Then the electrophilic attacking is taken place between the tert-carbocation and oxygen atom of carboxylate, the protonated ester(3) is yielded; The loss of a proton from the protonated ester gives the product(4) (Bart et al., 1994).

$$\mathsf{H}\_{\mathsf{H}} \leftarrow\_{\mathsf{c}\mathsf{RO}\xrightarrow[\mathsf{H}]{\mathsf{H}\_{\mathsf{H}}}\mathsf{R}\_{\mathsf{H}}} \mathsf{C} \underbrace{\mathsf{O}\_{\mathsf{H}\_{2}} \xleftarrow{-\mathsf{H}\_{2}\mathsf{O}} \mathsf{O}\_{\mathsf{H}\_{2}}}\_{\mathsf{H}\_{2}} \xleftarrow{+\mathsf{R}\_{2}\mathsf{O}} \mathsf{R}\_{\mathsf{R}\_{2}} \xleftarrow{+\mathsf{R}\mathsf{CO}\operatorname{OH}} \mathsf{R} \xleftarrow{+\mathsf{R}\mathsf{CO}\operatorname{OH}} \mathsf{R} \xleftarrow{+\mathsf{R}\mathsf{O}} \mathsf{R} \xleftarrow{+\mathsf{R}\mathsf{O}} \mathsf{R} \xleftarrow{+\mathsf{R}\mathsf{O}} \mathsf{R} \mathsf{R}$$

Fig. 4. Mechanism of tert-esterification

#### **2.1.3 Large-steric reaction**

Aromatic acid has serious steric hindrance. If there are both ortho-methyls, just as 2.4.6 trimethyl benzoic acid, the alcohol molecule is so difficult to access to carboxyl that the esterification cannot occur. However, if 2.4.6- trimethyl benzoic acid is dissolved in the 100% sulfuric-acid solution, acylilum ion will be formed as shown in Figure 5. Then the added alcohol with the acylilum ion produced ester. The reaction will conduct smoothly (Streitwieser et al., 1985).

This kind of esterification could occur for the reason that carbon atom of the acylilum ion is sp2 hybridized and coplanar with the benzene ring, Then alcohol molecule can be virtually unhindered to attack the acylilum ion from above or below of the molecular plane. Esterifications conducts with this mechanism are merely few.

2

2

*rr k c c*

*H O*

*r r Kk c c*

2 12

addition, the product *K*l*k*2 is denoted by a lumped constant *k*<sup>2</sup>

'

2 2

The rate equation is rewritten as follows,

catalyzed esterification correctly.

**2.1.2 Single-molecule reaction** 

R3C OH

H

**2.1.3 Large-steric reaction** 

(Streitwieser et al., 1985).

ester gives the product(4) (Bart et al., 1994).

R3C OH2

Fig. 4. Mechanism of tert-esterification

H2O

Esterifications conducts with this mechanism are merely few.

CR3

RCOOH

1 2 3 4

Aromatic acid has serious steric hindrance. If there are both ortho-methyls, just as 2.4.6 trimethyl benzoic acid, the alcohol molecule is so difficult to access to carboxyl that the esterification cannot occur. However, if 2.4.6- trimethyl benzoic acid is dissolved in the 100% sulfuric-acid solution, acylilum ion will be formed as shown in Figure 5. Then the added alcohol with the acylilum ion produced ester. The reaction will conduct smoothly

This kind of esterification could occur for the reason that carbon atom of the acylilum ion is sp2 hybridized and coplanar with the benzene ring, Then alcohol molecule can be virtually unhindered to attack the acylilum ion from above or below of the molecular plane.

*A H O CH COOH H O K K cc c c* 2 3 <sup>3</sup> 6 1 / (3)

12 2

2

H

R CR3 C

C O

O

'. Thus, we obtain for *r*2,

(4)

(5)

O

<sup>R</sup> OCR3 <sup>H</sup>

 <sup>3</sup> <sup>2</sup> 3 3

*H O E HO CH COOH CH OH*

3 2 3 3

 *H O E HO CH COOH CH OH H O E*

*c c c*

*c K* 

*c Kkk*

*c c c*

The product *K*l (*k*2/*k*-2) equals the equilibrium constant of the overall reaction *K*E2. In

The kinetic and equilibrium parameters included in Equation (5) are estimated from experimental data with regression analysis. Simulation of the model with the estimated parameters revealed that Equation (5) can predict the experimental trends in the acid-

The tertiary alcohol is prone to generate carbocation under acidic condition. Therefore, the mechanism of tert- esterification is different from that of the primary and secondary alcohol, it follows the single-molecule reaction process(shown in Figure 4):The tertiary alcohol combines with protons to generate protonated alcohol(1); The protonated alcohol is got rid of a molecule water to produce tert-carbocation(2), which is very stable; Then the electrophilic attacking is taken place between the tert-carbocation and oxygen atom of carboxylate, the protonated ester(3) is yielded; The loss of a proton from the protonated

Fig. 5. Mechanism of large-steric reaction

#### **2.1.4 Sulfuric acid**

Dhanuka showed that the mechanism of sulfuric acid-catalyzed esterification is the model of alkyl sulfuric intermediate. It's a two-step reaction. R′-O-SO3H firstly is generated via sulfuric acid with alcohol, which is treated as a catalyst in the second step (Lide, 1994):

$$\text{R'OH} + \text{H}\_2\text{SO}\_4 \xrightarrow{\text{H}^+} \text{R'OSO} + \text{H}\_2\text{O} \tag{6}$$

$$\text{RCOOH} + \text{R'OH} \xrightarrow{\text{RCOO}\_3} \text{RCOOR'} + \text{H}\_2\text{O} \tag{7}$$

The alkylation reaction of sulfuric acid with alcohol is irreversible (Aranda et al., 2008).Therefore the reaction rate can be expressed as:

$$r = k(\mathbb{C}\_A \mathbb{C}\_B - \mathbb{C}\_W \mathbb{C}\_E \nmid K) \tag{8}$$

where *k* is rate constant and *K* is equilibrium constant; *C*A, *C*B, *C*W, *C*E is the concentration of carboxylic acid, alcohol, water and ester, respectively.

In 2009, Fei, *et al.* studied on the esterification behaviors of neo-polyhydric alcohols with fatty acids catalyzed by sulfuric acid. It was found that the esterification follows the above mechanism.

#### **2.2 Lewis acids**

Recent Developments on the Mechanism and Kinetics

The consumption rate of A (-*r*A ) is expressed by,

respectively.

can be rewritten as

obtained at 185°C,

of Esterification Reaction Promoted by Various Catalysts 261

the progress of the reaction were investigated. The reaction was assumed to be a first order with respect to each reactant. For the irreversible, bimolecular-type second order reaction,

where *C*A, *C*B, *C*C is the concentration of lauric acid, hydroxyl group and catalyst (mol/mL),

The catalyst concentration was constant throughout the reaction; therefore, Equation (11)

where *X*A, *X*B is fractional conversion of the lauric acid and the -OH group, respectively; *C*A0, *C*B0 is initial concentration of the lauric acid and hydroxyl group (mol/mL), respectively;

The esterification reaction of castor oil and lauric acid was carried out using SnCl2•2H2O as catalyst (0.25,0.5,1.0 and 2.0 % w/w catalyst loadings) at 185℃ and the kinetic data were measured. The results showed that there was a good linearity relationship between ln [(1-XB)/(1-XA)] and t. The plot of k′ versus CC was close to a straight line. For the given values of *C*B0, *C*A0 and *C*C, the kinetic model of esterification of castor oil and lauric acid was

It was investigated the esterification of octanoic acid and n-octyl alcohol utilizing metallic chlorides (KCl, CoCl2, MgCl2, ZnCl2, FeCl3 etc.) in a stirred tank reactor (Santos, 1996). The results showed that the best efficiency of the formatted ester (n-octyl octanoate) was obtained with ferric chloride, indicating that the higher electronegativity of the metallic ion and the existence of free d orbitals in the transition metals are responsible for the higher yield of ester. Thus, the existence of free d orbitals in Fe3+ (its configuration: 3d3) gives the possibility of the formation of complexes with OH groups of the reactants, which explains the highest activity found for FeCl3•6H2O in this way. Generally, in the catalytic systems whose bases were constituted by a transition metal, such as Fe, Co, Mn, Zn, the mechanism

The esterification reaction of octanoic acid and n-octyl alcohol was carried out using CoCl2•2H2O as catalyst (0, 0.0385, 0.077 mol/l) at 70℃ and the kinetic data were measured. The experimental curves suggest that the kinetics of esterification between octanoic acid and n-octyl alcohol can be described by an irreversible second order power model, considering the catalyst concentration as a constant in the kinetic model proposed. The activation energy

6

0 0C ln[(1 ) / (1 )] 1.62 10 ( ) *X X B A B A C C Ct* (14)

The integrated form of Equation (12) using fractional conversion is,

of esterification can similarly be ascribed to the above mentioned.

is seen to have a value of 53 kcal/mol (Urteaga et al., 1994).

and k is the second-order rate constant (mL2mol-2min-1).

A+B Product *Catalyst* (10)

<sup>A</sup> CAB *r kC C C* = (11)

A= *<sup>A</sup> <sup>B</sup> r kC C* , *<sup>C</sup> k kC* (12)

0 0 ln[(1 ) / (1 )] ( ) *X X C C kt B A BA* (13)

Lewis acid is a compound or ionic species that can accept an electron pair from a donor compound to form a Lewis adduct, such as ZnCl2, Mg (ClO4)2, Sn[N(SO2-n-C8F17)2]4. It is potentially useful on the production of esters as a catalyst, the reason is that it can be easily separated from the reaction media (Sanchez et al., 1992; Cardoso et al., 2009); moreover, comparing with Bronsted acid, template effects are to be expected as Lewis acid sterically bulkier than a proton. In 1964, Anantakrishnan, et al. investigated the esterification reaction of Ac2O with 50% MeOH, EtOH, PrOH, and iso-PrOH, in 50% Me2CO or Dioxane as solvent, using Lewis acid (ZnCl2) as catalyst.

Along with the scientific and technological progress and social development, more and more researchers have concerned about this issue, including the kinds of catalysts, the mechanism and kinetics of reaction. Thus, the contents containing the reaction mechanism and esterification kinetics have been significantly reviewed according to the type of catalysts as follows.

#### **2.2.1 Metal chloride**

Ethyl oleate was synthesized by the esterification of and ethanol catalyzed by SnCl2•2H2O (Cardoso et al., 2008). Under the circumstance of excess ethanol, the effects of the concentration of the catalyst and oleic acid, and temperature on the reaction rate were investigated. A related esterification mechanism was presented and described as follows: in presence of Sn2+( SnCl2•2H2O) catalyst, the carbonyl of the fatty acid is polarized to activate of substrate, which makes the nucleophilic attack to the molecules by ethanol become more favorable. Cardoso et al. investigated the effect of different carbonic chain of alcohol (methyl alcohol, ethyl alcohol, n-propyl alcohol, n-butyl alcohol) on the conversion of oleic acid into respective ester. The results showed that the conversion rate was down with the increase of carbon chain of alcohol, which indicated that high bulk hindrance occurs on the hydroxyl of the alcohol, and the efficient attack of them to the polarized carbonyl of oleic acid is reduced. However, it is not clear how the carbonyl is polarized by Sn2+.

Kinetic data of esterification of ethanol and oleic acid catalyzed by SnCl2 (nethanol : noleic acid : nSnCl2 = 120 : 1 : 0.01) were measured at the reflux temperature, and kinetic model was obtained as follows,

$$\ln C = -0.00278 \times t - 2.3155 \tag{9}$$

where *C* is the concentration of oleic acid, t is the reaction time.

The effectiveness of the catalyst SnCl2•2H2O has been investigated in a broad range of concentrations, and the results are approximately concomitant with a first order dependence in relation to the catalyst concentration. The effect of the temperature on the initial rate of the esterification was determined, showing that an increase in the reaction temperature was caused a corresponding improvement on the reaction rate, especially at a range of 45-75°C. The value of activation energy for the reaction was determined from the data of the initial rate to be 46.69 kJ·mol-1.

The kinetics of catalytic esterication of castor oil with lauric acid using SnCl2•2H2O was studied (Kulkarni & Sawant, 2003). Effects of the catalyst concentration and temperature on

Lewis acid is a compound or ionic species that can accept an electron pair from a donor compound to form a Lewis adduct, such as ZnCl2, Mg (ClO4)2, Sn[N(SO2-n-C8F17)2]4. It is potentially useful on the production of esters as a catalyst, the reason is that it can be easily separated from the reaction media (Sanchez et al., 1992; Cardoso et al., 2009); moreover, comparing with Bronsted acid, template effects are to be expected as Lewis acid sterically bulkier than a proton. In 1964, Anantakrishnan, et al. investigated the esterification reaction of Ac2O with 50% MeOH, EtOH, PrOH, and iso-PrOH, in 50% Me2CO or Dioxane as solvent,

Along with the scientific and technological progress and social development, more and more researchers have concerned about this issue, including the kinds of catalysts, the mechanism and kinetics of reaction. Thus, the contents containing the reaction mechanism and esterification kinetics have been significantly reviewed according to the type of catalysts

Ethyl oleate was synthesized by the esterification of and ethanol catalyzed by SnCl2•2H2O (Cardoso et al., 2008). Under the circumstance of excess ethanol, the effects of the concentration of the catalyst and oleic acid, and temperature on the reaction rate were investigated. A related esterification mechanism was presented and described as follows: in presence of Sn2+( SnCl2•2H2O) catalyst, the carbonyl of the fatty acid is polarized to activate of substrate, which makes the nucleophilic attack to the molecules by ethanol become more favorable. Cardoso et al. investigated the effect of different carbonic chain of alcohol (methyl alcohol, ethyl alcohol, n-propyl alcohol, n-butyl alcohol) on the conversion of oleic acid into respective ester. The results showed that the conversion rate was down with the increase of carbon chain of alcohol, which indicated that high bulk hindrance occurs on the hydroxyl of the alcohol, and the efficient attack of them to the polarized carbonyl of oleic acid is

Kinetic data of esterification of ethanol and oleic acid catalyzed by SnCl2 (nethanol : noleic acid : nSnCl2 = 120 : 1 : 0.01) were measured at the reflux temperature, and kinetic model was

The effectiveness of the catalyst SnCl2•2H2O has been investigated in a broad range of concentrations, and the results are approximately concomitant with a first order dependence in relation to the catalyst concentration. The effect of the temperature on the initial rate of the esterification was determined, showing that an increase in the reaction temperature was caused a corresponding improvement on the reaction rate, especially at a range of 45-75°C. The value of activation energy for the reaction was determined from the data of the initial

The kinetics of catalytic esterication of castor oil with lauric acid using SnCl2•2H2O was studied (Kulkarni & Sawant, 2003). Effects of the catalyst concentration and temperature on

ln = 0.00278 2.3155 *C t* (9)

reduced. However, it is not clear how the carbonyl is polarized by Sn2+.

where *C* is the concentration of oleic acid, t is the reaction time.

**2.2 Lewis acids** 

as follows.

**2.2.1 Metal chloride** 

obtained as follows,

rate to be 46.69 kJ·mol-1.

using Lewis acid (ZnCl2) as catalyst.

the progress of the reaction were investigated. The reaction was assumed to be a first order with respect to each reactant. For the irreversible, bimolecular-type second order reaction,

$$\text{A} + \text{B} \xrightarrow{\text{Catalyst}} \text{Product} \tag{10}$$

The consumption rate of A (-*r*A ) is expressed by,

$$-r\lambda \equiv \mathrm{kC\_{C}C\_{A}C\_{B}} \tag{11}$$

where *C*A, *C*B, *C*C is the concentration of lauric acid, hydroxyl group and catalyst (mol/mL), respectively.

The catalyst concentration was constant throughout the reaction; therefore, Equation (11) can be rewritten as

$$-r\text{A} \bullet \text{k}' \mathbb{C}\_A \mathbb{C}\_{B'} \text{ }\_{\text{A}'} \text{k}' = k \mathbb{C}\_{\mathbb{C}} \tag{12}$$

The integrated form of Equation (12) using fractional conversion is,

$$\ln[(1 - X\_{\beta}) \ne (1 - X\_{A})] = (\mathbb{C}\_{10} - \mathbb{C}\_{A0})k't \tag{13}$$

where *X*A, *X*B is fractional conversion of the lauric acid and the -OH group, respectively; *C*A0, *C*B0 is initial concentration of the lauric acid and hydroxyl group (mol/mL), respectively; and k is the second-order rate constant (mL2mol-2min-1).

The esterification reaction of castor oil and lauric acid was carried out using SnCl2•2H2O as catalyst (0.25,0.5,1.0 and 2.0 % w/w catalyst loadings) at 185℃ and the kinetic data were measured. The results showed that there was a good linearity relationship between ln [(1-XB)/(1-XA)] and t. The plot of k′ versus CC was close to a straight line. For the given values of *C*B0, *C*A0 and *C*C, the kinetic model of esterification of castor oil and lauric acid was obtained at 185°C,

$$\ln[(1 - X\_{\text{A}}) \; / \; (1 - X\_{\text{A}})] = 1.62 \times 10^{\circ} \text{ (C}\_{\text{B}0} - \text{C}\_{\text{A0}}) \text{C}\_{\text{C}} t \tag{14}$$

It was investigated the esterification of octanoic acid and n-octyl alcohol utilizing metallic chlorides (KCl, CoCl2, MgCl2, ZnCl2, FeCl3 etc.) in a stirred tank reactor (Santos, 1996). The results showed that the best efficiency of the formatted ester (n-octyl octanoate) was obtained with ferric chloride, indicating that the higher electronegativity of the metallic ion and the existence of free d orbitals in the transition metals are responsible for the higher yield of ester. Thus, the existence of free d orbitals in Fe3+ (its configuration: 3d3) gives the possibility of the formation of complexes with OH groups of the reactants, which explains the highest activity found for FeCl3•6H2O in this way. Generally, in the catalytic systems whose bases were constituted by a transition metal, such as Fe, Co, Mn, Zn, the mechanism of esterification can similarly be ascribed to the above mentioned.

The esterification reaction of octanoic acid and n-octyl alcohol was carried out using CoCl2•2H2O as catalyst (0, 0.0385, 0.077 mol/l) at 70℃ and the kinetic data were measured. The experimental curves suggest that the kinetics of esterification between octanoic acid and n-octyl alcohol can be described by an irreversible second order power model, considering the catalyst concentration as a constant in the kinetic model proposed. The activation energy is seen to have a value of 53 kcal/mol (Urteaga et al., 1994).

Recent Developments on the Mechanism and Kinetics

OEt

EtOH CO2

experimental evidence to elucidate the reaction mechanism.

model of related catalysts are reviewed in the following discussions.

Mg(ClO4)2

EtO O

O O

Mg ClO4 ClO4

EtO O

R=aryl,alkyl

O O

(Jousseaume et al., 2003).

**2.3.1 Titanate** 

**2.3 Organometallic compound** 

of Esterification Reaction Promoted by Various Catalysts 263

nucleophilicity. Herein, the reaction mechanism was presented to explain the observed reaction characteristics. In fact, the more acidic and less nucleophilic phenols react faster than aliphatic alcohols. Thus, the release of the alcoholic proton should be involved in the ratedetermining step. A reasonable mechanistic hypothesis is depicted as follows (Figure 6),

OEt

O

I II III

O

OEt

OEt

HO

Due to the ability of coordinating with 1,3-dicarbonyl compounds, Mg(ClO4)2 reacts with diethyl dicarbonate to form complex (II), which can undergo the addition of the alcohol to form intermediate (III). An internal proton shift in (III) can produce intermediate (IV), which can irreversibly decompose to the mixed carbonate (VI) and to the carbonic acid monoester (V). Owing to its high instability, (V) immediately produces EtOH and CO2. The irreversibility of the last two steps drives the overall process towards (III). This explanation accounts for the formation of the mixed carbonate (VI) as the major product of the reaction

Nevertheless, this is a speculative hypothesis; more studies are in progress to find

Organometallic compounds containing bonds between carbons and metals provide a source of nucleophilic carbon atoms which can react with electrophilic carbon to form a new carbon-carbon bond (Banach et al., 2001; Finelli et al., 2004). Organometallics find practical uses in catalytic processes. For example, Titanium tetrabutoxide, Butylhydroxyoxostannane are well-known catalysts for the esterification between diacid and diol in polyester industry (Grzesik et al., 2000; Prabhakarn et al., 2011). The catalytic mechanism and kinetic

Titanium-based catalysts have been known for many years and actually are widely used for the production of poly (butylene terephthalate) (PBT), poly (trimethylene terephthalate)

Fig. 6. Mechanism for preparing alkyl and aryl ethyl carbonates catalyzed by Mg(ClO4)2

V

RO

VI

ROH

EtO

O

O R H

O

O R H

IV

EtO

O

O

Mg ClO4 ClO4

Mg ClO4 ClO4

O

OEt

O

OEt

The reaction rate went on a power law model of second order, one for the alcohol and one for the acid,

$$r = k \text{C}\_{\text{alcohol}} \text{ol} \text{C}\_{\text{acid}} \tag{15}$$

where *k* is kinetic constant (l mol−1min−1)

$$r = 9.30 \times 10^{-4} \,\text{C}\_{\text{alcohol} \diamondsuit \text{C}\_{\text{acid}}} \tag{16}$$


Table 1. Kinetic constant for homogeneous reactions

It was also observed that the homogeneous reactions were fitted much better than the pseudo-homogeneous reactions to proposed kinetics model. The reason for it is that an important influence of the physical steps that could have place in the global mechanism of this reaction.

#### **2.2.2 Perchlorate**

Perchlorates are of great chemical interest and importance. The high electronegativity, together with the relatively low charge density, results in poor complexing ability of the perchlorate ion. Metal perchlorates can therefore act as powerful Lewis acids, with this character mainly being exploited to activate bidentate compounds (Bartoli et al., 2007). Magnesium perchlorate is one of the most active Lewis acids for esterification. By 2003, Gooβen & Döhring synthesized various esters through a decarboxylative esterification of alkyl and aryl carboxylic acids with dialkyl dicarbonates in the presence of 1 mol% of Mg(ClO4)2. However, more in-depth investigations on the mechanism of this reaction are underway.

Until 2006, Bartoli, et al. explored the efficiency of Perchlorates (such as Mg(ClO4)2, LiClO4, Zn(ClO4)2) as catalyst to activate diethyl dicarbonate for preparing alkyl and aryl ethyl carbonates. The results showed that among the above catalysts, Mg(ClO4)2 is the most efficient one. The reactions between diethyl dicarbonate and various alcohols including aliphatic alcohols and aryl alcohols were studied, and the results obtained suggest that the reaction rate depends on the acidity of the starting hydroxy compound, and not on its

The reaction rate went on a power law model of second order, one for the alcohol and one

**Catalyst k[l mol-1 min-1] R2** KCl 0.95×10−4 0.988 CaCl2•6H2O 0.83×10−4 0.980 NiCl2•6H2O 1.86×10−4 0.805 CoCl2•6H2O 1.69×10−4 0.994 MgCl2•6H2O 5.87×10−4 0.986 MnCl2•6H2O 1.20×10−4 0.852 FeCl2•6H2O 1.62×10−4 0.977 BaCl2•2H2O 0.96×10−4 0.955 SnCl2•2H2O 10.00×10−4 0.943 ZnCl2•XH2O 1.54×10−4 0.876 AlCl3•6H2O 6.91×10−4 0.876 SnCl4 9.50×10−4 0.896 FeCl3•6H2O 9.30×10−4 0.845

It was also observed that the homogeneous reactions were fitted much better than the pseudo-homogeneous reactions to proposed kinetics model. The reason for it is that an important influence of the physical steps that could have place in the global mechanism of

Perchlorates are of great chemical interest and importance. The high electronegativity, together with the relatively low charge density, results in poor complexing ability of the perchlorate ion. Metal perchlorates can therefore act as powerful Lewis acids, with this character mainly being exploited to activate bidentate compounds (Bartoli et al., 2007). Magnesium perchlorate is one of the most active Lewis acids for esterification. By 2003, Gooβen & Döhring synthesized various esters through a decarboxylative esterification of alkyl and aryl carboxylic acids with dialkyl dicarbonates in the presence of 1 mol% of Mg(ClO4)2. However, more in-depth

Until 2006, Bartoli, et al. explored the efficiency of Perchlorates (such as Mg(ClO4)2, LiClO4, Zn(ClO4)2) as catalyst to activate diethyl dicarbonate for preparing alkyl and aryl ethyl carbonates. The results showed that among the above catalysts, Mg(ClO4)2 is the most efficient one. The reactions between diethyl dicarbonate and various alcohols including aliphatic alcohols and aryl alcohols were studied, and the results obtained suggest that the reaction rate depends on the acidity of the starting hydroxy compound, and not on its

*r kC C* = alcohol acid (15)

<sup>4</sup> *r CC* =9.30 10 alcohol acid (16)

for the acid,

this reaction.

**2.2.2 Perchlorate** 

where *k* is kinetic constant (l mol−1min−1)

Table 1. Kinetic constant for homogeneous reactions

investigations on the mechanism of this reaction are underway.

nucleophilicity. Herein, the reaction mechanism was presented to explain the observed reaction characteristics. In fact, the more acidic and less nucleophilic phenols react faster than aliphatic alcohols. Thus, the release of the alcoholic proton should be involved in the ratedetermining step. A reasonable mechanistic hypothesis is depicted as follows (Figure 6),

Fig. 6. Mechanism for preparing alkyl and aryl ethyl carbonates catalyzed by Mg(ClO4)2

Due to the ability of coordinating with 1,3-dicarbonyl compounds, Mg(ClO4)2 reacts with diethyl dicarbonate to form complex (II), which can undergo the addition of the alcohol to form intermediate (III). An internal proton shift in (III) can produce intermediate (IV), which can irreversibly decompose to the mixed carbonate (VI) and to the carbonic acid monoester (V). Owing to its high instability, (V) immediately produces EtOH and CO2. The irreversibility of the last two steps drives the overall process towards (III). This explanation accounts for the formation of the mixed carbonate (VI) as the major product of the reaction (Jousseaume et al., 2003).

Nevertheless, this is a speculative hypothesis; more studies are in progress to find experimental evidence to elucidate the reaction mechanism.

#### **2.3 Organometallic compound**

Organometallic compounds containing bonds between carbons and metals provide a source of nucleophilic carbon atoms which can react with electrophilic carbon to form a new carbon-carbon bond (Banach et al., 2001; Finelli et al., 2004). Organometallics find practical uses in catalytic processes. For example, Titanium tetrabutoxide, Butylhydroxyoxostannane are well-known catalysts for the esterification between diacid and diol in polyester industry (Grzesik et al., 2000; Prabhakarn et al., 2011). The catalytic mechanism and kinetic model of related catalysts are reviewed in the following discussions.

#### **2.3.1 Titanate**

Titanium-based catalysts have been known for many years and actually are widely used for the production of poly (butylene terephthalate) (PBT), poly (trimethylene terephthalate)

Recent Developments on the Mechanism and Kinetics

expression in the integrated form becomes,

constant; *C*C is catalyst concentration.

TBT as a catalyst (Tiwari & Sawant, 2005).

**2.3.2 Butylhydroxyoxo-stannane** 

catalyst is also proposed by TIAN, et al.

expressed as follow,

(Equation (21)),

of Esterification Reaction Promoted by Various Catalysts 265

catalyst in the temperature range of 165-185℃.The reaction rate for the catalytic reaction is the sum of the rates of both the uncatalyzed (Equation (20)) and the catalyzed reactions

CH (CH ) COOH+CH (CH ) CH OH 3 2 20 3 2 8 2 CH (CH ) COOCH (CH ) CH +H O 3 2 20 2 28 3 2 (20)

3 2 20 32 2 <sup>8</sup> CH (CH ) COOH+CH CH CH OH *Catalyst* 3 2 20 2 28 3 2 CH (CH ) COOCH (CH ) CH H O + (21)

Excess of the acids is used to get almost complete conversion of the alcohols, and the unreacted acid was easily removed from the ester as a sodium salt. Assuming that it is first order dependence of the reaction rate on each reactant and the catalyst, the overall rate

where *X*A is fractional conversion of behenic acid; *C*A0 is initial concentration of behenic acid; *c c uc k kC k* , *k*c is catalyzed reaction rate constant, *k*uc is uncatalyzed reaction rate

The kinetic data show that there is a reasonably good agreement between experimental points and the ones calculated by Equation (22). The values of activation energy obtained for the uncatalyzed (∆*E*uc) and catalyzed (∆*E*c) reaction of behenic acid with decanol, lauryl alcohol, myristyl alcohol and cetyl alcohol are shown in Table 2.The model Equation (22) is also appropriate for the reaction of erucic acid with cetyl alcohol and oleyl alcohol using

> Alcohol ∆*E*uc(kJ mol-1) ∆*E*c (kJ mol-1) decanol 68.1 86.2 lauryl alcohol 69.2 79.6 myristyl alcohol 64.5 78.6 cetyl alcohol 67.3 87.1

Stannum-based catalysts such as butylhydroxyoxo-stannane (BuSnOOH) are commonly used for the synthesis of poly (ethylene terephthalate) (PET) (Patel et al., 2007), poly (butylene terephthalate) (PBT) and so on. Due to the similarity in coordination number and electronegativity between Ti and Sn, as shown in Figure 8, the mechanism of BuSnOOH

The initial rate method will also be used in this reacting system. The kinetic model of monoesterification reaction between terephthalic acid (TPA) and 1,4-butanediol (BDO) is

Table 2. Activation energy value for uncatalyzed and catalyzed reaction

*X X C kt A A* /(1 ) A0 (22)

0.9 2 *r kC C Sn* B A (23)

(PTT), and so on. Recently, the developments of titanium-based catalysts make a marvelous progress (Liu et al., 2006), such as titanium dioxide based catalyst (C-94) mainly focusing on designing to be stable and having good activity and color.

The mechanism of the mono-esterification between terephthalic acid (TPA) and 1,4-butanediol (BDO) catalyzed by Ti(OBu)4 was proposed (Tian et al., 2010). As shown in Figure 7, the reaction involves the formation of an adduct between a carbonyl group and Ti atom.

Fig. 7. Mechanism of the mono-esterification between BDO and TPA catalyzed by Ti(OBu)4

Although the reaction system of TPA and BDO is heterogeneous, it can be assumed that the esterification occurs only in the liquid phase. The initial rate method is used to predict the reaction rate. The kinetic model of mono-esterification between TPA and BDO catalyzed by Ti(OBu)4 in the temperature of 463-483K was investigated (Bhutada & Pangarkar, 1986). The reaction rate r can be described as,

$$r = k \mathbb{C}\_{\Lambda}^{n} \mathbb{C}\_{\mathbb{B}}^{m} \tag{17}$$

where *C*A, *C*B is the concentrations of TPA and BDO, respectively, *k* is the reaction rate constant, *n* and *m* is the reaction order with respect to A and B.

The reaction order can be obtained from experiments in which the initial rates are measured at a series of initial reactant concentrations. The results show that the reaction order for TPA and BDO is nearly a constant of 0.7 and 0.9, respectively. Therefore, the rate equation of the esterification is written as follows,

$$r = k\_{\rm Ti} \mathbb{C}\_{\rm B}^{0.9} \mathbb{C}\_{\rm A}^{0.7} \tag{18}$$

$$
\ln k\_{1l} = 11.73 - \text{7.67} \times \text{10}^3 \text{ / T (R}^2 = 0.997) \tag{19}
$$

where *k*Ti is the rate constant (L0.6·mol−0.6·min−1), *R*2 is the correlation coefficient.

The reaction of behenic acid with fatty alcohols (decanol, lauryl alcohol, myristyl alcohol and cetyl alcohol) was studied by Tiwari, et al. (Tiwari & Sawant, 2005) using TBT as

(PTT), and so on. Recently, the developments of titanium-based catalysts make a marvelous progress (Liu et al., 2006), such as titanium dioxide based catalyst (C-94) mainly focusing on

The mechanism of the mono-esterification between terephthalic acid (TPA) and 1,4-butanediol (BDO) catalyzed by Ti(OBu)4 was proposed (Tian et al., 2010). As shown in Figure 7, the

C

O

Ti

HO(CH2)4OH

COOH

C

O

O

O

H

HO(CH2)4

*n m r kC C* (17)

*Ti* B A *r kC C* (18)

O H

O

Ti

O

O

O H

O

O

COOH

Fig. 7. Mechanism of the mono-esterification between BDO and TPA catalyzed by Ti(OBu)4

Although the reaction system of TPA and BDO is heterogeneous, it can be assumed that the esterification occurs only in the liquid phase. The initial rate method is used to predict the reaction rate. The kinetic model of mono-esterification between TPA and BDO catalyzed by Ti(OBu)4 in the temperature of 463-483K was investigated (Bhutada & Pangarkar, 1986). The

A B

where *C*A, *C*B is the concentrations of TPA and BDO, respectively, *k* is the reaction rate

The reaction order can be obtained from experiments in which the initial rates are measured at a series of initial reactant concentrations. The results show that the reaction order for TPA and BDO is nearly a constant of 0.7 and 0.9, respectively. Therefore, the rate equation of the

0.9 0.7

3 <sup>2</sup> ln 11.73 7.67 10 / ( 0.997) *Ti k R <sup>T</sup>* (19)

The reaction of behenic acid with fatty alcohols (decanol, lauryl alcohol, myristyl alcohol and cetyl alcohol) was studied by Tiwari, et al. (Tiwari & Sawant, 2005) using TBT as

where *k*Ti is the rate constant (L0.6·mol−0.6·min−1), *R*2 is the correlation coefficient.

reaction involves the formation of an adduct between a carbonyl group and Ti atom.

O

O

designing to be stable and having good activity and color.

COOH

COOH

constant, *n* and *m* is the reaction order with respect to A and B.

HO(CH2)4OOC(C6H4)4COOH H2O

Ti

Ti

O

O

O

O

reaction rate r can be described as,

esterification is written as follows,

O

O

O

O

catalyst in the temperature range of 165-185℃.The reaction rate for the catalytic reaction is the sum of the rates of both the uncatalyzed (Equation (20)) and the catalyzed reactions (Equation (21)),

$$\text{CH}\_3\text{(CH}\_2\text{)}\_2\text{COOH} + \text{CH}\_3\text{(CH}\_2\text{)}\_8\text{CH}\_2\text{OH} \rightleftharpoons \text{CH}\_3\text{(CH}\_2\text{)}\_2\text{COOCH}\_2\text{(CH}\_2\text{)}\_8\text{CH}\_3 + \text{H}\_2\text{O} \quad \text{(20)}$$

$$\text{CH}\_3\text{(CH}\_2\text{)}\_2\text{COOH} + \text{CH}\_3\text{(CH}\_2\text{)}\_8\text{CH}\_2\text{OH} \xrightarrow{\text{Catalyst}} \text{CH}\_3\text{(CH}\_2\text{)}\_2\text{COOCH}\_2\text{(CH}\_2\text{)}\_8\text{CH}\_3\text{+H}\_2\text{O (21)}$$

Excess of the acids is used to get almost complete conversion of the alcohols, and the unreacted acid was easily removed from the ester as a sodium salt. Assuming that it is first order dependence of the reaction rate on each reactant and the catalyst, the overall rate expression in the integrated form becomes,

$$X\_A \nmid (1 - X\_A) = \mathbb{C}\_{A0} kt \tag{22}$$

where *X*A is fractional conversion of behenic acid; *C*A0 is initial concentration of behenic acid; *c c uc k kC k* , *k*c is catalyzed reaction rate constant, *k*uc is uncatalyzed reaction rate constant; *C*C is catalyst concentration.

The kinetic data show that there is a reasonably good agreement between experimental points and the ones calculated by Equation (22). The values of activation energy obtained for the uncatalyzed (∆*E*uc) and catalyzed (∆*E*c) reaction of behenic acid with decanol, lauryl alcohol, myristyl alcohol and cetyl alcohol are shown in Table 2.The model Equation (22) is also appropriate for the reaction of erucic acid with cetyl alcohol and oleyl alcohol using TBT as a catalyst (Tiwari & Sawant, 2005).


Table 2. Activation energy value for uncatalyzed and catalyzed reaction

#### **2.3.2 Butylhydroxyoxo-stannane**

Stannum-based catalysts such as butylhydroxyoxo-stannane (BuSnOOH) are commonly used for the synthesis of poly (ethylene terephthalate) (PET) (Patel et al., 2007), poly (butylene terephthalate) (PBT) and so on. Due to the similarity in coordination number and electronegativity between Ti and Sn, as shown in Figure 8, the mechanism of BuSnOOH catalyst is also proposed by TIAN, et al.

The initial rate method will also be used in this reacting system. The kinetic model of monoesterification reaction between terephthalic acid (TPA) and 1,4-butanediol (BDO) is expressed as follow,

$$r = k\_{Sn} \mathbb{C}\_{\text{B}}^{0.9} \mathbb{C}\_{\text{A}}^{2} \tag{23}$$

Recent Developments on the Mechanism and Kinetics

sulfate-supported metal oxides (SO42−/MxOy), TPA/SnO2.

**2.4 Solid acids** 

**2.4.1 TPA/SnO2**

C15H31 C OH

C15H31 C

measured.

constant.

O

O

of Esterification Reaction Promoted by Various Catalysts 267

The research for solid acids has become active since the early 1970s. In 1971, Isao, et al. investigated the esterification of ethanol with acetic acid on silica-alumina, and a simple kinetic model based on a Langmuir-Hinshelwood mechanism was proposed. Since then, the esterification catalyzed by solid acids is widely studied and largely reported (Jiang et al., 2008; Jothiramalingam & Wang, 2009; Li et al., 2010). Herein, the esterification mechanisms and kinetic models according to the different types of solid acids are reviewed, such as

The esterification of palmitic acid with methanol using 12-Tungstophosphoric acid

15 31 3 15 31 3 2

(A) (B) (C) (D)

OH

The catalyst initiates the esterication reaction by donating a proton to palmitic acid molecule. The palmitic acid is then subjected to nucleophilic attack by the hydroxyl group of

As the large excess of methanol, it could be safely assumed to be a rst-order pseudohomogeneous reaction. The esterification reaction was carried out using 15 wt% TPA/SnO2 as catalyst(wA: wcat=5:1)in the temperature range of 45-65℃and the kinetic data were

where *X*A is the conversation of palmitic acid, *t* is the reaction time, *k* is the reaction rate

<sup>k</sup> C H COOH+CH OH C H COOCH +H O (25)

CH3OH

H2O

OCH3

C15H31 C

C15H31 C

*ln kt* (1-X )= A (26)

OH

OH

OH2

O

H

CH3

OCH3

OH

(TPA)/SnO2 was investigated (Srilatha et al., 2011). The reaction is shown as follows:

C15H31 C OH

C15H31 C

OH

As shown in Figure 9, the mechanism of this esterification is proposed,

H

methanol, and the reaction continues with water elimination.

Fig. 9. Mechanism of acid catalyzed esterification of carboxylic acid

OCH3

The reaction rate can be described as follows:

H

$$
\ln k\_{\rm Su} = 9.58 - 6.42 \times 10^3 \text{ / (R}^2 = 0.999\text{)}\tag{24}
$$

where *C*A and *C*B is the concentrations of TPA and BDO, respectively; *k*Sn is the rate constant (L1.9·mol−1.9·min−1) .

Fig. 8. Mechanism of the mono-esterification between BDO and TPA catalyzed by BuSnOOH

#### **2.3.3 Fluorous metaloxide**

In 2004, Otera investigated the transesterification and esterification of various acids and alcohols using 1, 3-disubstituted tetraalkyldistannoxanes as catalyst, and found that the catalytic mechanism different from that proposed for tetraalkyldistannoxanes. In the latter case, the initial step is substitution of the bridging X group to give an alkoxydistannoxane intermediate, (YR2SnOSnR2OR)2, which works as an alkoxyl donor. In contrast, for 1, 3 disubstituted tetraalkyldistannoxanes, no substitution takes place at the bridging position by an alkoxy group. As described above, the bridging chlorine in 1, 3-disubstituted tetraalkyldistannoxanes is never substituted by isothiocyanate ion, but experiences strong association.

A related mechanism of the esterification is proposed that both alcohol and acid coordinate on the terminal tin atom (Yoshida et al., 2006), on which the interchange between the hydroxyl groups of carboxylic acids and the alkoxy groups of alcohols takes place. Meanwhile, since water, one of the products, is less fluorophilic, the esterification of carboxylic acids with alcohols should proceed efficiently.

Fluorous distannoxanes exhibit unusually high preference for fluorous solvents over common organic solvents thanks to coverage of the molecular surface with fluoroalkyl groups. While the fluorous biphase technology had been high-lighted mainly in terms of facile separation of products and fluorous catalysts, it has now been revealed that the equilibrium is also controllable under fluorous biphase conditions. Thus, the esterification has been completely driven in the desired direction by use of a 1:1 ratio of starting materials without recourse to any dehydration technique.

#### **2.4 Solid acids**

266 Chemical Kinetics

where *C*A and *C*B is the concentrations of TPA and BDO, respectively; *k*Sn is the rate constant

Sn

OH

C4H9 O

C

O

HO(CH2)4OH

COOH

C

O

H

HO(CH2)4

O

H

O

OH

C4H9 Sn O

O H

COOH

In 2004, Otera investigated the transesterification and esterification of various acids and alcohols using 1, 3-disubstituted tetraalkyldistannoxanes as catalyst, and found that the catalytic mechanism different from that proposed for tetraalkyldistannoxanes. In the latter case, the initial step is substitution of the bridging X group to give an alkoxydistannoxane intermediate, (YR2SnOSnR2OR)2, which works as an alkoxyl donor. In contrast, for 1, 3 disubstituted tetraalkyldistannoxanes, no substitution takes place at the bridging position by an alkoxy group. As described above, the bridging chlorine in 1, 3-disubstituted tetraalkyldistannoxanes is never substituted by isothiocyanate ion, but experiences strong

A related mechanism of the esterification is proposed that both alcohol and acid coordinate on the terminal tin atom (Yoshida et al., 2006), on which the interchange between the hydroxyl groups of carboxylic acids and the alkoxy groups of alcohols takes place. Meanwhile, since water, one of the products, is less fluorophilic, the esterification of

Fluorous distannoxanes exhibit unusually high preference for fluorous solvents over common organic solvents thanks to coverage of the molecular surface with fluoroalkyl groups. While the fluorous biphase technology had been high-lighted mainly in terms of facile separation of products and fluorous catalysts, it has now been revealed that the equilibrium is also controllable under fluorous biphase conditions. Thus, the esterification has been completely driven in the desired direction by use of a 1:1 ratio of starting materials

HO(CH2)4OOC(C6H4) H2O 4COOH

Fig. 8. Mechanism of the mono-esterification between BDO and TPA catalyzed by

(L1.9·mol−1.9·min−1) .

OH

C4H9

Sn O

Sn O

OH

**2.3.3 Fluorous metaloxide** 

C4H9

BuSnOOH

association.

COOH

COOH

carboxylic acids with alcohols should proceed efficiently.

without recourse to any dehydration technique.

3 2 ln 9.58 6.42 10 /( 0.999) *Sn k R* (24)

The research for solid acids has become active since the early 1970s. In 1971, Isao, et al. investigated the esterification of ethanol with acetic acid on silica-alumina, and a simple kinetic model based on a Langmuir-Hinshelwood mechanism was proposed. Since then, the esterification catalyzed by solid acids is widely studied and largely reported (Jiang et al., 2008; Jothiramalingam & Wang, 2009; Li et al., 2010). Herein, the esterification mechanisms and kinetic models according to the different types of solid acids are reviewed, such as sulfate-supported metal oxides (SO4 <sup>2</sup>−/MxOy), TPA/SnO2.

#### **2.4.1 TPA/SnO2**

The esterification of palmitic acid with methanol using 12-Tungstophosphoric acid (TPA)/SnO2 was investigated (Srilatha et al., 2011). The reaction is shown as follows:

$$\text{C}\_{13}\text{H}\_{31}\text{COOH} + \text{CH}\_{3}\text{OH} \xrightleftharpoons \text{C}\_{19}\text{H}\_{31}\text{COOCH}\_{3} + \text{H}\_{2}\text{O} \tag{25}$$

$$\text{(A)}\qquad \text{(B)}\qquad \qquad \text{(C)}\qquad \qquad \qquad \text{(D)}$$

As shown in Figure 9, the mechanism of this esterification is proposed,

Fig. 9. Mechanism of acid catalyzed esterification of carboxylic acid

The catalyst initiates the esterication reaction by donating a proton to palmitic acid molecule. The palmitic acid is then subjected to nucleophilic attack by the hydroxyl group of methanol, and the reaction continues with water elimination.

As the large excess of methanol, it could be safely assumed to be a rst-order pseudohomogeneous reaction. The esterification reaction was carried out using 15 wt% TPA/SnO2 as catalyst(wA: wcat=5:1)in the temperature range of 45-65℃and the kinetic data were measured.

The reaction rate can be described as follows:

$$-ln(\mathbf{1} \cdot \mathbf{X}\_{\rm A}) = kt \tag{26}$$

where *X*A is the conversation of palmitic acid, *t* is the reaction time, *k* is the reaction rate constant.

Recent Developments on the Mechanism and Kinetics

written as follows:

rate constant.

(*M*=*C*B0/*C*A0).

The equation can be rearranged to be:

range of 393-443K is given as following equations,

temperature range of 303.15- 323.15K (Otera, 1993).

enhance the catalytic activity.

follows:

of Esterification Reaction Promoted by Various Catalysts 269

The stirring rate is sufficient to overcome the diffusion limitation of reactive species. Therefore, the performance of pseudo-homogeneous model can be considered as satisfactory to correlate the kinetic data for the esterification. The rate equation can be

> <sup>A</sup> <sup>A</sup> 1AB 1EW *dC r kC C k C C*

where *C*A, *C*B, *C*E and *C*W is the concentration of myristic acid, methanol, myristic acid methyl ester and water respectively, *k*1 is the forward rate constant and *k*-1 is the backward

> *dX kC X M X X dt K* <sup>2</sup> A 1 A0 A A A

where *X*A is the conversion of myristic acid,*C*A0 is the initial concentration of myristic acid, *K*e is the equilibrium constant, *M* is the concentration ratio of methanol to myristic acid

The relationship between temperature and the forward rate constant in the temperature

In 2004, Jiang, et al. studied the esterification of n-pentanol with benzoic acid using Alpillared clay (PILC) supported SO42− /TiO2 superacid catalyst (Fang et al., 2010), and find

superacid is necessarily promoted by the sulfur of the metal oxides, the more acid sites formed, the higher catalytic activity exhibited. Therefore, Al-PILC carrier can effectively

In 2010, Zubir, et al. investigated the kinetic behavior of the heterogeneous catalyzed esterification of oleic acid with ethanol using tungstated zirconia as a catalyst in the

 CH CH CH=CH CH COOH+CH CH OH 3 2 7 7 <sup>2</sup> 3 2 CH CH CH=CH CH COOCH CH +H O 3 2 <sup>7</sup> <sup>2</sup> <sup>7</sup> <sup>232</sup> (33) (A) (B) (C) (D)

The kinetics of esterification reaction can be expressed using a pseudo-homogeneous second-order equilibrium model in the absence of any intraparticle diffusional limitation as

1 <sup>1</sup> 1 *Ae <sup>e</sup>*

The linear coefficient is 0.996, and the activation energy is 22.51 kJ /mol.

that it is known that the generation of super acid sites in the system of SO4

*<sup>k</sup> <sup>X</sup> <sup>K</sup>*

 2

*Ae Ae*

*dt* (29)

*e*

<sup>1</sup> [(1 )( ) ] (30)

*k X MX* (31)

<sup>1</sup> *k T* exp 2.88 2707 / (32)

*r kCC CC K <sup>A</sup> AB CD e* / (34)

<sup>2</sup>− /MxOy solid

The plots of-ln (1-*X*A) versus *t* at different temperatures showed that there was a good linearity relationship. Linear fitting to ln*k* - 1/T curve was carried out and the reaction rate constant can be described as follows:

$$k = \exp\left(8.66 - 4369 \,/\, T\right) \tag{27}$$

The activation energy was obtained as 36.33 kJ mol−1.

#### **2.4.2 SO4 2− /MxOy**

The solid superacid (SO4 <sup>2</sup>− /MxOy) is a new type of catalyst used in esterication (Jiang et al., 2004). MxOy are usually some transition metal oxides such as ZrO2 and TiO2.

In 2010, Rattanaphra, et al. investigated the esterication of myristic acid with methanol catalyzed by sulfated zirconia. The mechanism of esterification is probably following Langmuir-Hinshelwood model (Arata, 2009; Reddy & Patil, 2009). It is possible that methanol and myristic acid are preferentially adsorbed on the Bronsted acid sites of sulfated zirconia during esterication. The hydroxyl group of methanol is protonated by Bronsted acid on the catalyst surface while the protonation of myristic acid on an adjacent site leads to the carbocation. Deprotonation of methanol oxygen produces the nucleophile, which attacks the carbocation to generate a tetrahedral intermediate. As shown in Figure 10, the tetrahedral intermediate eliminates water to form ester.

Fig. 10. Mechanism of esterication catalyzed by sulfated zirconia

In 2011, Rattanaphra, et al. also investigated the kinetic model of myristic acid esterification with methanol using sulfated zirconia as catalyst. The reaction is shown as follows:

$$\text{CH}\_3\text{(CH}\_2\text{)}\_{12}\text{COOH} + \text{CH}\_3\text{OH} \overset{k\_i}{\underset{k\_{-1}}{\rightleftharpoons}} \text{CH}\_3\text{(CH}\_2\text{)}\_{12}\text{COOCH}\_3 + \text{H}\_2\text{O} \tag{28}$$

$$\text{(A)}\tag{A}\tag{B}\tag{10.1}$$

$$\text{(A)}\tag{B}\tag{10.2}$$

The plots of-ln (1-*X*A) versus *t* at different temperatures showed that there was a good linearity relationship. Linear fitting to ln*k* - 1/T curve was carried out and the reaction rate

The solid superacid (SO42− /MxOy) is a new type of catalyst used in esterication (Jiang et

In 2010, Rattanaphra, et al. investigated the esterication of myristic acid with methanol catalyzed by sulfated zirconia. The mechanism of esterification is probably following Langmuir-Hinshelwood model (Arata, 2009; Reddy & Patil, 2009). It is possible that methanol and myristic acid are preferentially adsorbed on the Bronsted acid sites of sulfated zirconia during esterication. The hydroxyl group of methanol is protonated by Bronsted acid on the catalyst surface while the protonation of myristic acid on an adjacent site leads to the carbocation. Deprotonation of methanol oxygen produces the nucleophile, which attacks the carbocation to generate a tetrahedral intermediate. As shown in Figure 10, the

al., 2004). MxOy are usually some transition metal oxides such as ZrO2 and TiO2.

*k T* exp 8.66 4369 / (27)

constant can be described as follows:

**2.4.2 SO4**

B+ B<sup>+</sup>

CH3OH

B B

C OH

OCH

R

**2− /MxOy**

The activation energy was obtained as 36.33 kJ mol−1.

tetrahedral intermediate eliminates water to form ester.

B+ B O

CH3

RCOOH

<sup>B</sup> <sup>B</sup> H2O

with methanol using sulfated zirconia as catalyst. The reaction is shown as follows:

In 2011, Rattanaphra, et al. also investigated the kinetic model of myristic acid esterification

1 1 *k k*

(A) (B) (E) (W)

B B O

C R

O

H2O OCH

O

HO

CH3

B B O

<sup>H</sup>

B B O <sup>O</sup> CH3

C R

C R

O

HO

HO

CH (CH ) COOCH +H O 3 2 12 3 2 (28)

CH3

H

H

<sup>H</sup> <sup>B</sup> <sup>B</sup>

R

H

O

Fig. 10. Mechanism of esterication catalyzed by sulfated zirconia

CH (CH ) COOH+CH OH 3 2 12 <sup>3</sup>

C R OCH

The stirring rate is sufficient to overcome the diffusion limitation of reactive species. Therefore, the performance of pseudo-homogeneous model can be considered as satisfactory to correlate the kinetic data for the esterification. The rate equation can be written as follows:

$$-r\_{\rm A} = -\frac{d\mathbf{C\_{A}}}{dt} = k\_{\rm 1}\mathbf{C\_{A}}\mathbf{C\_{B}} - k\_{-\rm 1}\mathbf{C\_{E}}\mathbf{C\_{W}}\tag{29}$$

where *C*A, *C*B, *C*E and *C*W is the concentration of myristic acid, methanol, myristic acid methyl ester and water respectively, *k*1 is the forward rate constant and *k*-1 is the backward rate constant.

The equation can be rearranged to be:

$$\frac{dX\_{\rm A}}{dt} = k\_1 \mathbb{C}\_{\rm A0} [(1 - X\_{\rm A})(M - X\_{\rm A}) - \frac{1}{K\_e} X\_{\rm A}^2] \tag{30}$$

where *X*A is the conversion of myristic acid,*C*A0 is the initial concentration of myristic acid, *K*e is the equilibrium constant, *M* is the concentration ratio of methanol to myristic acid (*M*=*C*B0/*C*A0).

$$K\_e = \frac{k\_1}{k\_{-1}} = \frac{X\_{Ae}^2}{\left(1 - X\_{Ae}\right)\left(M - X\_{Ae}\right)}\tag{31}$$

The relationship between temperature and the forward rate constant in the temperature range of 393-443K is given as following equations,

$$k\_1 = \exp\left(2.88 - 2707 \text{ / T}\right) \tag{32}$$

The linear coefficient is 0.996, and the activation energy is 22.51 kJ /mol.

In 2004, Jiang, et al. studied the esterification of n-pentanol with benzoic acid using Alpillared clay (PILC) supported SO42− /TiO2 superacid catalyst (Fang et al., 2010), and find that it is known that the generation of super acid sites in the system of SO42− /MxOy solid superacid is necessarily promoted by the sulfur of the metal oxides, the more acid sites formed, the higher catalytic activity exhibited. Therefore, Al-PILC carrier can effectively enhance the catalytic activity.

In 2010, Zubir, et al. investigated the kinetic behavior of the heterogeneous catalyzed esterification of oleic acid with ethanol using tungstated zirconia as a catalyst in the temperature range of 303.15- 323.15K (Otera, 1993).

$$\text{CH}\_3\text{(CH}\_2\text{)}\_7\text{CH=CH(CH}\_2\text{)}\_7\text{COOH} \\ \text{CH}\_3\text{CH}\_2\text{OH} \xrightleftharpoons \text{CH}\_3\text{(CH}\_2\text{)}\_7\text{CH=CH(CH}\_2\text{)}\_7\text{COOCH}\_2\text{CH}\_3\text{-H}\_2\text{O} \tag{33}$$

$$\text{(A)}\tag{34} \qquad \text{(B)}\tag{10} \qquad \text{(C)}\tag{11}$$

The kinetics of esterification reaction can be expressed using a pseudo-homogeneous second-order equilibrium model in the absence of any intraparticle diffusional limitation as follows:

$$-r\_A = k \left( \mathbf{C}\_A \mathbf{C}\_B - \mathbf{C}\_C \mathbf{C}\_D \;/\; K\_e \right) \tag{34}$$

Recent Developments on the Mechanism and Kinetics

Saha, 2007; Jeřábek et al., 2010; Erdem & Kara, 2011).

The ER model of this esterification can be described as follow,

*ri*

such as Langmuir-Hinshelwood (LH) model.

1

follows,

rate.

of Esterification Reaction Promoted by Various Catalysts 271

The esterification kinetics of acetic acid with iso-butanol catalyzed by Dowex 50 Wx2 was studied (Izci & Bodur, 2007). It is considered to be reversible reactions and can be described by pseudo-homogeneous (PH) model. The general reaction rate expression can be written as

where subscripts A, B, E and W is acid, alcohol, ester and water, respectively, *k*1 is forward

It is found that this bimolecular type is second order reaction. The values of *k*1 for different temperatures are 0.24×103 (318K), 1.00×103 (333K), 3.07×103 (348K). The activation energy was found to be 1.745 kJ·mol−1 in the presence of Dowex 50 Wx2. However, the PH model does not take into account the resin swelling ratio, adsorption of the components and the non-ideal thermodynamic behavior of reactants and products (Ali &Merchant, 2006; Patel &

The kinetics of Dowex 50 Wx8-catalyzed esterification was studied between acetic acid and benzyl alcohol (Ali & Merchant, 2009). The swelling ratio of Dowex 50 Wx8 in different solvent was measured, and the results show that it decreases in the order of water, benzyl alcohol, acetic acid, benzyl acetate. Water appears to be preferentially adsorbed by the catalyst from a binary solution of acetic acid and water, and hinders the approach of butanol to the protonated acid. Therefore, the water exerts an adverse effect on the esterification

The initial reaction rates of esterification between acetic acid and benzyl alcohol were measured at various conditions, and the Eley-Rideal (ER) model was used to correlate the data and showed a high degree of fit, indicating that the surface reaction between adsorbed alcohol and acid in the bulk is the rate-limiting step during the initial stage of the reaction.

/

*k T <sup>f</sup>* 55770 exp 6661.4 / (40)

(39)

*Ka K a* 

ln 1279 / 6.810 *K T <sup>a</sup>* (41)

where *M*cat is the mass of the catalyst, g; *a*acid, *a*alc, *a*ester, *a*water is the activity of acid, alcohol, ester, water in the liquid phase, respectively; *k*f is forward reaction rate constant for the esterification, mol g-1 s-1. *K*a is activity reaction equilibrium constant. *K*acid, *K*alc, *K*water is

However, as the reaction proceeds, the acid adsorption term might have to be introduced and in such an event the reaction kinetics would be represented by a dual site mechanistic,

adsorption equilibrium constant for the alcohol present in the system, respectively.

The activation energy form the above relationship was found to be 55.4kJ/mol.

*cat f alc acid alc ester water a alc alc water water*

*M kK a a a a K*

reaction rate constant (L mol−1 min−1), *K* is the equilibrium constant of the reaction.

*r kC C C C K* A 1A B E W / (38)

where *C*A ,*C*B, *C*C, *C*D is the concentration of oleic acid, ethanol, ethyl oleate and water, respectively, *k* is the kinetic constants for the forward reaction, *K*e is the equilibrium constant.

$$\mathbf{K\_e = exp}(29.42 - 10536 \,/\,\mathrm{T})\tag{35}$$

$$\mathbf{k} = \exp\left(23.43 - 6242 \,\mathrm{/\,T}\right). \tag{36}$$

The value of activation energy is 51.9 kJ·mol-1. The goodness-of-fit of the experimental data to the proposed model is assessed by comparing the experimental reaction rate with the theoretical prediction, and the experimental data are reproduced with errors not greater than 10%.

#### **2.5 Ion-exchange resin**

Ion-exchange resins, especially the cation-exchange resins such as Dowex, Amberlyst series are manufactured mainly by sulfonation of ethylbenzene first, followed by a cross-link with divinylbenzene (Liu & Tan, 2001; Alexandratos, 2008; Tesser et al., 2010). Because of their selective adsorption of reactants, surface acid site features, and swelling nature, these resins not only catalyze the esterification reaction but also affect the equilibrium conversion. They also show excellent performance such as reusable, mechanical separation, continuous operation as a heterogeneous catalyst in esterication (Yang et al., 2007; JagadeeshBabu et al., 2011; Ju et al., 2011; Toor et al., 2011).

By 1965, Bochner, et al. evaluated the performance of Dowex 50WX-8 as a catalyst for the esterification of salicylic acid with methanol. Through analyzing the experimental results, the Langmuir-isotherm applied to this system to describe the reaction rate. Since then, various mechanisms and kinetics of esterification between acids and alcohols catalyzed by cation- exchange resins were proposed. Here is an overview of the reaction mechanisms and kinetic models according to the type of cation-exchange resins.

#### **2.5.1 Dowex**

Dowex, a common cation-exchange resin, is widely used as catalysts in esterification (Vahteristo et al., 2009). In 2005, Kulawska, et al. investigated the esterification of maleic anhydride with octyl, decyl or dodecyl alcohol over Dowex 50Wx8-100, and the kinetic data were measured in the temperature range of 403-433 K (C0ALC/C0MA =5:1). Based on the data, a first order reaction was found- first order with respect to acid and zero order with respect to alcohol. The reaction rate can be described as follows,

$$r = -\frac{dc\_{\rm M}}{dt} = kc\_{\rm M} \tag{37}$$

where *c*M is the concentration of maleic monoester, t is reaction time.

The values of activation energy are 66.0 (±0.65) kJ/mol, 58.6 (±0.4) kJ/mol, and 66.1(±0.4) kJ/mol for dioctyl, didencyl, and didodecyl maleate formation, respectively. However, the effects of the catalyst (the particle size and the concentration) and the mass transfer resistance on the reaction rate were not taken into account.

where *C*A ,*C*B, *C*C, *C*D is the concentration of oleic acid, ethanol, ethyl oleate and water, respectively, *k* is the kinetic constants for the forward reaction, *K*e is the equilibrium

The value of activation energy is 51.9 kJ·mol-1. The goodness-of-fit of the experimental data to the proposed model is assessed by comparing the experimental reaction rate with the theoretical prediction, and the experimental data are reproduced with errors not greater

Ion-exchange resins, especially the cation-exchange resins such as Dowex, Amberlyst series are manufactured mainly by sulfonation of ethylbenzene first, followed by a cross-link with divinylbenzene (Liu & Tan, 2001; Alexandratos, 2008; Tesser et al., 2010). Because of their selective adsorption of reactants, surface acid site features, and swelling nature, these resins not only catalyze the esterification reaction but also affect the equilibrium conversion. They also show excellent performance such as reusable, mechanical separation, continuous operation as a heterogeneous catalyst in esterication (Yang et al., 2007; JagadeeshBabu et

By 1965, Bochner, et al. evaluated the performance of Dowex 50WX-8 as a catalyst for the esterification of salicylic acid with methanol. Through analyzing the experimental results, the Langmuir-isotherm applied to this system to describe the reaction rate. Since then, various mechanisms and kinetics of esterification between acids and alcohols catalyzed by cation- exchange resins were proposed. Here is an overview of the reaction mechanisms and

Dowex, a common cation-exchange resin, is widely used as catalysts in esterification (Vahteristo et al., 2009). In 2005, Kulawska, et al. investigated the esterification of maleic anhydride with octyl, decyl or dodecyl alcohol over Dowex 50Wx8-100, and the kinetic data were measured in the temperature range of 403-433 K (C0ALC/C0MA =5:1). Based on the data, a first order reaction was found- first order with respect to acid and zero order with respect

M

The values of activation energy are 66.0 (±0.65) kJ/mol, 58.6 (±0.4) kJ/mol, and 66.1(±0.4) kJ/mol for dioctyl, didencyl, and didodecyl maleate formation, respectively. However, the effects of the catalyst (the particle size and the concentration) and the mass transfer

*dc r kc*

M

*dt* (37)

K exp 29.42 10536 / T <sup>e</sup> (35)

k exp 23.43 6242 /T . (36)

constant.

than 10%.

**2.5.1 Dowex** 

**2.5 Ion-exchange resin** 

al., 2011; Ju et al., 2011; Toor et al., 2011).

kinetic models according to the type of cation-exchange resins.

to alcohol. The reaction rate can be described as follows,

resistance on the reaction rate were not taken into account.

where *c*M is the concentration of maleic monoester, t is reaction time.

The esterification kinetics of acetic acid with iso-butanol catalyzed by Dowex 50 Wx2 was studied (Izci & Bodur, 2007). It is considered to be reversible reactions and can be described by pseudo-homogeneous (PH) model. The general reaction rate expression can be written as follows,

$$-r\_{\rm A} = k\_1 \left( \mathbf{C}\_{\rm A} \times \mathbf{C}\_{\rm B} - \mathbf{C}\_{\rm E} \times \mathbf{C}\_{\rm W} \,/\, K \right) \tag{38}$$

where subscripts A, B, E and W is acid, alcohol, ester and water, respectively, *k*1 is forward reaction rate constant (L mol−1 min−1), *K* is the equilibrium constant of the reaction.

It is found that this bimolecular type is second order reaction. The values of *k*1 for different temperatures are 0.24×103 (318K), 1.00×103 (333K), 3.07×103 (348K). The activation energy was found to be 1.745 kJ·mol−1 in the presence of Dowex 50 Wx2. However, the PH model does not take into account the resin swelling ratio, adsorption of the components and the non-ideal thermodynamic behavior of reactants and products (Ali &Merchant, 2006; Patel & Saha, 2007; Jeřábek et al., 2010; Erdem & Kara, 2011).

The kinetics of Dowex 50 Wx8-catalyzed esterification was studied between acetic acid and benzyl alcohol (Ali & Merchant, 2009). The swelling ratio of Dowex 50 Wx8 in different solvent was measured, and the results show that it decreases in the order of water, benzyl alcohol, acetic acid, benzyl acetate. Water appears to be preferentially adsorbed by the catalyst from a binary solution of acetic acid and water, and hinders the approach of butanol to the protonated acid. Therefore, the water exerts an adverse effect on the esterification rate.

The initial reaction rates of esterification between acetic acid and benzyl alcohol were measured at various conditions, and the Eley-Rideal (ER) model was used to correlate the data and showed a high degree of fit, indicating that the surface reaction between adsorbed alcohol and acid in the bulk is the rate-limiting step during the initial stage of the reaction. The ER model of this esterification can be described as follow,

$$\dot{m} = \frac{\left[\boldsymbol{M}\_{cat}\boldsymbol{k}\_f \boldsymbol{K}\_{alc} \left(\boldsymbol{a}\_{acid}\boldsymbol{a}\_{alc} - \left(\boldsymbol{a}\_{cster}\boldsymbol{a}\_{water} \; / \boldsymbol{K}\_a \right) \right) \right]}{\left[\boldsymbol{1} + \boldsymbol{K}\_{alc}\boldsymbol{a}\_{alc} + \boldsymbol{K}\_{water}\boldsymbol{a}\_{water}\right]} \tag{39}$$

$$k\_f = 55770 \times \exp\left(-6661.4 \,/\, T\right) \tag{40}$$

$$
\ln\left(K\_{\mu}\right) = \left(-1279 \,\mathrm{\,\mathrm{\,\,\mathrm{\,\,{\,\,{\,\,{\,\,{\,\,{\,\,{\,\,{\,\,{\,\,{\,\,{\,\,{\,\,{\,\,{\,\,{\,\,{\,\,{\,\,{\,\,\,{\,\,{\,\,{\,\,\,{\,\,\,{\,\,\,{\,\,\,{\,\,\,{\,\,\,{\,\,\,{\,\,\,{\,\,\,\,{\,\,\,\,{\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\}\right)}\right)
$$

where *M*cat is the mass of the catalyst, g; *a*acid, *a*alc, *a*ester, *a*water is the activity of acid, alcohol, ester, water in the liquid phase, respectively; *k*f is forward reaction rate constant for the esterification, mol g-1 s-1. *K*a is activity reaction equilibrium constant. *K*acid, *K*alc, *K*water is adsorption equilibrium constant for the alcohol present in the system, respectively.

The activation energy form the above relationship was found to be 55.4kJ/mol.

However, as the reaction proceeds, the acid adsorption term might have to be introduced and in such an event the reaction kinetics would be represented by a dual site mechanistic, such as Langmuir-Hinshelwood (LH) model.

Recent Developments on the Mechanism and Kinetics

experimental mole fractions of acid is 5.71%.

**2.5.2 Amberlyst** 

Tsai et al., 2011).

of Esterification Reaction Promoted by Various Catalysts 273

propionic acid at the same time intervals decreases with increasing the chain length of the alcohol, indicating that the chain length of alcohol has a retarding effect on the reaction rate. It is ascribing to the steric hindrance. Meanwhile, the LH model is used to correlate the above experimental data, and the total average error between the predicted and

Interestingly, the kinetic study of esterifications between the same acid (acetic acid) and different alcohol (benzyl alcohol, 2-propanol) in the presence of Dowex 50x8 reveal that the ER model and the LH model is the most suitable predictive model, respectively. These observations seem to indicate that the type of acid and alcohol play an important role in

Amberlyst, acidic cation-exchange resin bead, was widely used as excellent heterogeneous acid catalysts for a wide variety of organic reactions owing to its macroreticular and continuous open pore structure (Pöpken et al., 2000; Kolah et al., 2007; Pereira et al., 2008;

In 2011, Tsai, et al. investigated the kinetic behavior of heterogeneous esterification of dibasic acids such as glutaric acid with methanol over Amberlyst 35. As shown in Equation (42) and Equation (43), the esterification includes two reversible reactions in series

> 1 2 *k*

3 4 *k*

The large excess of methanol was taken in this study due to the low solubility of glutaric acid in methanol. The kinetic data which had been determined experimentally for the esterication were correlated with quasi-homogeneous (*QH*) model (Schmitt & Hasse, 2006). Comparing with the experimental values, the QH model represents well the conversions of glutaric acid and the mole fractions of each constituent component varying with the contact time. The values of the rate constants were determined by fitting the kinetic data to the reaction rate expressions simultaneously. The temperature dependent *k*1, *k*2, *k*3, *k*4 were listed in Table 3,

It should be noted that the kinetic constants reported above include the swelling effects of Amberlyst 35 beads and just for this two-stage esterification system. However, the non-ideal

thermodynamic behavior of reactants and products was not taken into account.

(C) (B) (E) (D) where A, B, C, D, E is glutaric acid, methanol, monomethyl glutarate, water, dimethyl glutarate, respectively, *k*1, *k*3 is forward reaction rate constants, and *k*2, *k*4 is backward reaction

(A) (B) (C) (D)

*<sup>k</sup>* HOOC(CH ) COOCH +H O 23 3 2 (42)

*<sup>k</sup>* CH OOC(CH ) COOCH +H O <sup>3</sup> 23 3 2 (43)

*ki i k i* <sup>0</sup> exp / *E RT* 1,2,3,4 (44)

accompanying with an intermediate of monomethyl glutarate (MMG):

HOOC(CH ) COOH+CH OH 2 3 <sup>3</sup>

HOOC(CH ) COOCH +CH OH 23 3 3

rate constants, respectively,(cm6g−1min−1mol−1).

determining the number of sites involved in this heterogeneously catalyzed reaction.

In 2011, Ju, et al. measured the kinetic data for the esterification of butyric acid with nbutanol over Dowex 50Wx8-400, and correlated with various types of kinetic models. Strong resin water affinity was taken into account, and the non-ideality of the system was considered by applying the no-random two liquid (NRTL) model. The comparisons of the conversion of butyric acid with reaction time between experimental data and ones predicted by these kinetic models reveal that the ER model and LH model are the most reasonable fit for describing the mechanism, with the total average error 12.52%. Surface reaction is the rate determining step, and the affinity between resin and water is found to be not strong. Therefore, the most possible esterification reaction mechanisms can be proposed as following lists:


In 2007, Ali, et al. studied the esterification of 1-propanol with propionic acid catalyzed by Dowex 50Wx8-400. The experiments were carried out over a temperature range of 303.15- 333.15K, and the reaction mechanism for the esterification was proposed (Lilja et al., 2002). The reaction is initiated by the transfer of a proton from the catalyst to the carboxylic acid,and the carbonium ion is formed during the reaction. The ion is accessible for a nucleophilic attack by the hydroxyl group from the alcohol. After that, a molecule of water is lost from the ion. Finally, the catalyst is recovered by the transfer of proton from the ion to the catalyst surface. This mechanism is represented by the following scheme (Figure 11):

Fig. 11. Mechanism for the esterification catalyzed by Dowex 50Wx8-400

The donation of a proton is commonly assumed to be a fast step, and the nucleophilic substitution is usually assumed to be the rate determined step.

Based on the above reaction mechanism, external and internal diffusion can be negligible, and the E-R model (Equation (39)) is applied. The activity coefficients of components can be predicted using UNIFAC model. Meanwhile, by taking into account the strong water affinity for Dowex 50Wx8-400, a correction term (*α*) was added to the activity term for water in rate expression. The experimental results for the esterification between 1-propanol and propionic acid were modeled according to the above kinetic model. The total average error between the predicted and experimental mole fractions of acid is 1.65%. The activation energy for the forward reaction was estimated to be 67.3 kJ /mol.

The effect of different alcohols on the conversion of propionic acid was investigated using methanol, ethanol, 1-propanol and 1-butanol. The results reveal that the conversion of propionic acid at the same time intervals decreases with increasing the chain length of the alcohol, indicating that the chain length of alcohol has a retarding effect on the reaction rate. It is ascribing to the steric hindrance. Meanwhile, the LH model is used to correlate the above experimental data, and the total average error between the predicted and experimental mole fractions of acid is 5.71%.

Interestingly, the kinetic study of esterifications between the same acid (acetic acid) and different alcohol (benzyl alcohol, 2-propanol) in the presence of Dowex 50x8 reveal that the ER model and the LH model is the most suitable predictive model, respectively. These observations seem to indicate that the type of acid and alcohol play an important role in determining the number of sites involved in this heterogeneously catalyzed reaction.

#### **2.5.2 Amberlyst**

272 Chemical Kinetics

In 2011, Ju, et al. measured the kinetic data for the esterification of butyric acid with nbutanol over Dowex 50Wx8-400, and correlated with various types of kinetic models. Strong resin water affinity was taken into account, and the non-ideality of the system was considered by applying the no-random two liquid (NRTL) model. The comparisons of the conversion of butyric acid with reaction time between experimental data and ones predicted by these kinetic models reveal that the ER model and LH model are the most reasonable fit for describing the mechanism, with the total average error 12.52%. Surface reaction is the rate determining step, and the affinity between resin and water is found to be not strong. Therefore, the most possible esterification reaction mechanisms can be proposed as

1. Single site mechanism: the adsorbed butyric acid onto the catalysts reacts with non-

2. Single site mechanism: the adsorbed n-butanol onto the catalyst reacts with non-

3. Dual site mechanism: both the reactants adsorbed on the catalyst surface and react

In 2007, Ali, et al. studied the esterification of 1-propanol with propionic acid catalyzed by Dowex 50Wx8-400. The experiments were carried out over a temperature range of 303.15- 333.15K, and the reaction mechanism for the esterification was proposed (Lilja et al., 2002). The reaction is initiated by the transfer of a proton from the catalyst to the carboxylic acid,and the carbonium ion is formed during the reaction. The ion is accessible for a nucleophilic attack by the hydroxyl group from the alcohol. After that, a molecule of water is lost from the ion. Finally, the catalyst is recovered by the transfer of proton from the ion to the catalyst surface. This mechanism is represented by the following scheme (Figure 11):

k1

k-1

<sup>R</sup>′OH

RCOOH <sup>H</sup> RCH(OH)2 RCOOR′ H2O

The donation of a proton is commonly assumed to be a fast step, and the nucleophilic

Based on the above reaction mechanism, external and internal diffusion can be negligible, and the E-R model (Equation (39)) is applied. The activity coefficients of components can be predicted using UNIFAC model. Meanwhile, by taking into account the strong water affinity for Dowex 50Wx8-400, a correction term (*α*) was added to the activity term for water in rate expression. The experimental results for the esterification between 1-propanol and propionic acid were modeled according to the above kinetic model. The total average error between the predicted and experimental mole fractions of acid is 1.65%. The activation

The effect of different alcohols on the conversion of propionic acid was investigated using methanol, ethanol, 1-propanol and 1-butanol. The results reveal that the conversion of

Fig. 11. Mechanism for the esterification catalyzed by Dowex 50Wx8-400

substitution is usually assumed to be the rate determined step.

energy for the forward reaction was estimated to be 67.3 kJ /mol.

following lists:

there.

adsorbed n-butanol in the bulk.

adsorbed butyric acid in the bulk.

k1 k-1 Amberlyst, acidic cation-exchange resin bead, was widely used as excellent heterogeneous acid catalysts for a wide variety of organic reactions owing to its macroreticular and continuous open pore structure (Pöpken et al., 2000; Kolah et al., 2007; Pereira et al., 2008; Tsai et al., 2011).

In 2011, Tsai, et al. investigated the kinetic behavior of heterogeneous esterification of dibasic acids such as glutaric acid with methanol over Amberlyst 35. As shown in Equation (42) and Equation (43), the esterification includes two reversible reactions in series accompanying with an intermediate of monomethyl glutarate (MMG):

HOOC(CH ) COOH+CH OH 2 3 <sup>3</sup> 1 2 *k <sup>k</sup>* HOOC(CH ) COOCH +H O 23 3 2 (42) (A) (B) (C) (D)

$$\text{HCOOC(CH}\_2\text{)}\_3\text{COOCH}\_3 + \text{CH}\_3\text{OH} \xrightleftharpoons[k]{k\_1} \text{CH}\_3\text{OOC(CH}\_2\text{)}\_3\text{COOCH}\_3 + \text{H}\_2\text{O} \tag{43}$$

(C) (B) (E) (D) where A, B, C, D, E is glutaric acid, methanol, monomethyl glutarate, water, dimethyl glutarate, respectively, *k*1, *k*3 is forward reaction rate constants, and *k*2, *k*4 is backward reaction rate constants, respectively,(cm6g−1min−1mol−1).

The large excess of methanol was taken in this study due to the low solubility of glutaric acid in methanol. The kinetic data which had been determined experimentally for the esterication were correlated with quasi-homogeneous (*QH*) model (Schmitt & Hasse, 2006). Comparing with the experimental values, the QH model represents well the conversions of glutaric acid and the mole fractions of each constituent component varying with the contact time. The values of the rate constants were determined by fitting the kinetic data to the reaction rate expressions simultaneously. The temperature dependent *k*1, *k*2, *k*3, *k*4 were listed in Table 3,

$$k\_i = k\_{i0} \exp\left(-\Delta E \;/\,\text{RT}\right) \left(\mathbf{i} = \mathbf{1}, \mathbf{2}, \mathbf{3}, \mathbf{4}\right) \tag{44}$$

It should be noted that the kinetic constants reported above include the swelling effects of Amberlyst 35 beads and just for this two-stage esterification system. However, the non-ideal thermodynamic behavior of reactants and products was not taken into account.

Recent Developments on the Mechanism and Kinetics

<sup>3</sup> 3 2 <sup>3</sup> CH CH(OH)COOH+CH CH OH <sup>1</sup>

behavior of the liquid-solid catalytic esterification:

A 2

373.15 K over entire range of the experimental feed compositions.

shown as follows,

*r*

reaction mechanism is proposed to be:

CH CHCOOH+CH CHOHCH OH 2 32

<sup>2</sup> 2CH CHCOOH <sup>3</sup> *<sup>k</sup>*

<sup>2</sup> 3 6 2CH CHCOOC H OH <sup>4</sup> *<sup>k</sup>*

of Esterification Reaction Promoted by Various Catalysts 275

2.117×10−5 (353.15 K), respectively. The predicted values of lactic acid fractional conversion

The heterogeneous esterification of propionic acid with n-butanol over Amberlyst 35 for the synthesis of n-butyl propionate was investigated (Lee et al., 2002). The esterification is

> 1 *k k*

(A) (B) (C) (D) The kinetics data of the liquid-solid catalytic esterification are correlated with various kinetic models, over wide ranges of temperature and feed composition. The activity coefficients calculated using the NTRL model are utilized to represent the non-ideality behavior of the species in the liquid solutions. Meanwhile, the effects of film diffusion and pore diffusion appear to be negligible at the experimental conditions. The results reveal that the Langmuir-Hinshelwood (LH) model yielded the best representation for the kinetic

The activation energy of reaction is about 6.81kJmol−1. This model is capable of representing the kinetic behavior of the liquid-solid catalytic esterification at temperatures from 353.15 to

The kinetics of esterication between acrylic acid and propylene glycol in the presence of Amberlyst 15 was investigated (Altokka & Ödeş, 2009). Taking into account the general esterication reaction as well as polymerization of acrylic acid and products, the overall

> 1 2 *k*

(A) (P) (T) (W)

(A) (T1) (W)

 (T) (T2) (W) Some of the dimer molecules formed here may transform into polymer. Assuming that each

reaction step is elementary, the corresponding rate expression can be written.

2.9782 exp( 7664 / ) a a 0.01847 exp(819.06 / )a a 1 3.70a

D

CH CH(OH)COO(CH ) CH +H O <sup>3</sup> 23 3 2 (49)

*T T*

A B C D

*<sup>k</sup>* CH CHCOOC H OH+H O <sup>2</sup> 36 2 (51)

C H O +H O 12 18 5 2 (53)

CH CHCOOOCCHCH +H O <sup>2</sup> 2 2 (52)

(50)

are in close agreement with the experimental within the range of experimental error.


Table 3. Each of reaction rate constant for the esterification

In 2008, Schmid, et al. investigated the reaction kinetics of the consecutive, reversible liquidphase esterification of ethylene glycol with acetic acid to ethylene glycol monoacetate and ethylene glycol diacetate using Amberlyst 36 as catalyst in the temperature range from 333.15K to 363.15K. As shown in Equation (45) and Equation (46), the reaction contains two parts: the formation of the mono-acetylated ethyleneglycol and its hydrolysis as backward reaction; the consecutive esterification of the mono-acetylated ethylene glycol to ethylene glycol diacetate and its hydrosis:

$$\text{CH}\_3\text{COOH} + \text{(CH}\_2\text{OH)}\_2 \xrightarrow[k\_{-1}]{k\_{-1}} \text{CH}\_3\text{OOC(CH}\_2\text{)}\_2\text{OH} + \text{H}\_2\text{O} \tag{45}$$

(A) (B) (C) (D)

2 -2 CH COOH+CH COO(CH ) OH CH COO(CH ) OOCCH +H O <sup>3</sup> <sup>3</sup> 2 2 <sup>3</sup> 22 3 2 *k <sup>k</sup>* (46) (A) (C) (E) (D)

The acetic acid initiates swelling of the resin, which results in easy accessible acid groups for the reaction and free mobility of all the components. Therefore, a pseudo-homogeneous (PH) kinetic model is applicable. For the reaction, there is,

$$r\_{\rm m} = \frac{1}{m\_{\rm cat}} \frac{1}{v\_i} \frac{dn\_i}{dt} = k\_1 a\_{\rm A} a\_{\rm B} - k\_{-1} a\_{\rm C} a\_{\rm D} + k\_2 a\_{\rm C} a\_{\rm A} - k\_{-2} a\_{\rm E} a\_{\rm D} \tag{47}$$

where *k*1, *k*2, *k*−1, *k*−2 is the forward and backward rate constants, respectively, (mol g-1 s-1), *a*<sup>i</sup> is the activity of each component, and it is calculated using the UNQUAC model, *m*cat is the mass of catalyst.

The values of *k*1, *k*2, *k*−1, *k*−2 is 539.14, 156.43, 114.41, 95.68, respectively, and the mean squared deviation between calculated and experimental mole fractions is 4.53×10-2.

The kinetic behavior of esterification of lactic acid with isopropanol over Amberlyst 15 was investigated (Toor et al., 2011; Kirbaslar et al., 2001), at different temperatures from 323K to 353K. The ER model was used to describe the reaction mechanism that takes place between adsorbed molecules of isopropanol and the molecules of lactic acid in bulk. The adsorption of ester is reported to be negligible. Hence the rate equation can be described as follows,

$$\frac{dX\_{\Lambda}}{dt} = k\_1 \mathbb{C}\_{\Lambda 0} \left( w \,/\, V \right) \left[ \left( 1 - X\_{\Lambda} \right) \left( M - X\_{\Lambda} \right) - \frac{X\_{\Lambda}^2}{K\_{\text{e}}} \right] \tag{48}$$

where *k*1 is Rate constant for esterification reaction, *M* is molar ratio of isopropanol to lactic acid, *K*e is equilibrium rate constant.

The value of activation energy is found to be 22l J/mol, and at each different temperature, the rate constants is 1.058×10−5 (323.15 K), 1.235×10−5 (333.15 K), 1.588×10−5 (343.15 K),

*k*0 3.346x107±1.01 6.416x103±1.00 3.487x105±1.01 1.958x109±1.01 Δ*E*/*R* 4417.4±1.7 84.5±0.001 2042.2±1.7 5974.6±4.0

In 2008, Schmid, et al. investigated the reaction kinetics of the consecutive, reversible liquidphase esterification of ethylene glycol with acetic acid to ethylene glycol monoacetate and ethylene glycol diacetate using Amberlyst 36 as catalyst in the temperature range from 333.15K to 363.15K. As shown in Equation (45) and Equation (46), the reaction contains two parts: the formation of the mono-acetylated ethyleneglycol and its hydrolysis as backward reaction; the consecutive esterification of the mono-acetylated ethylene glycol to ethylene

> 1 *k k*

(A) (B) (C) (D)

2 -2 CH COOH+CH COO(CH ) OH CH COO(CH ) OOCCH +H O <sup>3</sup> <sup>3</sup> 2 2 <sup>3</sup> 22 3 2 *k*

(A) (C) (E) (D) The acetic acid initiates swelling of the resin, which results in easy accessible acid groups for the reaction and free mobility of all the components. Therefore, a pseudo-homogeneous

m 1AB 1CD 2CA 2ED

where *k*1, *k*2, *k*−1, *k*−2 is the forward and backward rate constants, respectively, (mol g-1 s-1), *a*<sup>i</sup> is the activity of each component, and it is calculated using the UNQUAC model, *m*cat is the

The values of *k*1, *k*2, *k*−1, *k*−2 is 539.14, 156.43, 114.41, 95.68, respectively, and the mean

The kinetic behavior of esterification of lactic acid with isopropanol over Amberlyst 15 was investigated (Toor et al., 2011; Kirbaslar et al., 2001), at different temperatures from 323K to 353K. The ER model was used to describe the reaction mechanism that takes place between adsorbed molecules of isopropanol and the molecules of lactic acid in bulk. The adsorption of ester is reported to be negligible. Hence the rate equation can be described as follows,

> / 1 *dX <sup>X</sup> kC w V X M X dt <sup>K</sup>*

where *k*1 is Rate constant for esterification reaction, *M* is molar ratio of isopropanol to lactic

The value of activation energy is found to be 22l J/mol, and at each different temperature, the rate constants is 1.058×10−5 (323.15 K), 1.235×10−5 (333.15 K), 1.588×10−5 (343.15 K),

 <sup>2</sup> A A 1 A0 A A

e

(48)

squared deviation between calculated and experimental mole fractions is 4.53×10-2.

(47)

*r ka a k a a ka a k aa*

CH OOC(CH ) OH+H O <sup>3</sup> 22 2 (45)

*<sup>k</sup>* (46)

Table 3. Each of reaction rate constant for the esterification

3 2 <sup>2</sup> CH COOH+ CH OH <sup>1</sup>

(PH) kinetic model is applicable. For the reaction, there is,

1 1 *<sup>i</sup> cat i*

*m v dt*

*dn*

glycol diacetate and its hydrosis:

mass of catalyst.

acid, *K*e is equilibrium rate constant.

*k*<sup>1</sup> *k*<sup>2</sup> *k*<sup>3</sup> *k*<sup>4</sup>

2.117×10−5 (353.15 K), respectively. The predicted values of lactic acid fractional conversion are in close agreement with the experimental within the range of experimental error.

The heterogeneous esterification of propionic acid with n-butanol over Amberlyst 35 for the synthesis of n-butyl propionate was investigated (Lee et al., 2002). The esterification is shown as follows,

$$\text{CH}\_3\text{CH(OH)COOH} + \text{CH}\_3\text{(CH}\_2\text{)}\_3\text{OH} \xrightleftharpoons[k\_{\text{-}l}]{k\_{\text{-}l}} \text{CH}\_3\text{CH(OH)CO(CH}\_2\text{)}\_3\text{CH}\_3 + \text{H}\_2\text{O} \tag{49}$$

$$\text{(A)}\tag{\text{B}}\tag{\text{B}}\tag{\text{B}}\tag{\text{B}}\tag{\text{B}}\tag{\text{B}}$$

The kinetics data of the liquid-solid catalytic esterification are correlated with various kinetic models, over wide ranges of temperature and feed composition. The activity coefficients calculated using the NTRL model are utilized to represent the non-ideality behavior of the species in the liquid solutions. Meanwhile, the effects of film diffusion and pore diffusion appear to be negligible at the experimental conditions. The results reveal that the Langmuir-Hinshelwood (LH) model yielded the best representation for the kinetic behavior of the liquid-solid catalytic esterification:

$$-r\_{\rm A} = \frac{2.9782 \times 10^{10} \exp(-7664 \text{ } / \text{T}) \left[ \text{a}\_{\rm A} \text{a}\_{\rm B} - 0.01847 \exp(819.06 \text{ } / \text{T}) \text{a}\_{\rm C} \text{a}\_{\rm D} \right]}{\left[ 1 + 3.70 \text{a}\_{\rm D} \right]^2} \tag{50}$$

The activation energy of reaction is about 6.81kJmol−1. This model is capable of representing the kinetic behavior of the liquid-solid catalytic esterification at temperatures from 353.15 to 373.15 K over entire range of the experimental feed compositions.

The kinetics of esterication between acrylic acid and propylene glycol in the presence of Amberlyst 15 was investigated (Altokka & Ödeş, 2009). Taking into account the general esterication reaction as well as polymerization of acrylic acid and products, the overall reaction mechanism is proposed to be:

$$\text{CH}\_2\text{CHCOOH} + \text{CH}\_3\text{CHOHCH}\_2\text{OH} \xrightleftharpoons[k]{k\_1} \text{CH}\_2\text{CHCOOC}\_3\text{H}\_6\text{OH} + \text{H}\_2\text{O} \tag{51}$$

$$\begin{array}{cc} \text{(A)} \\ \text{(A)} \end{array} \tag{\text{(P)}} \tag{\text{(P)}} \tag{\text{(A)}} \tag{\text{(A)}}$$

$$\text{H}\_2\text{CH}\_2\text{CHCOOH} \xrightleftharpoons[k]{} \text{CH}\_2\text{CHCOOCHCH}\_2\text{+H}\_2\text{O}\tag{52}$$

$$\text{(A)}\tag{1}$$

$$\text{(A)}\tag{1}$$

<sup>2</sup> 3 6 2CH CHCOOC H OH <sup>4</sup> *<sup>k</sup>* C H O +H O 12 18 5 2 (53) (T) (T2) (W)

Some of the dimer molecules formed here may transform into polymer. Assuming that each reaction step is elementary, the corresponding rate expression can be written.

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$$\begin{aligned} \frac{d\mathbf{C\_A}}{dt} &= -k\_1 \mathbf{C\_A} \mathbf{C\_P} + k\_2 \mathbf{C\_T} \mathbf{C\_W} - k\_3 \mathbf{C\_A^2} \\ \frac{d\mathbf{C\_P}}{dt} &= -k\_1 \mathbf{C\_A} \mathbf{C\_P} + k\_2 \mathbf{C\_T} \mathbf{C\_W} \\ \frac{d\mathbf{C\_T}}{dt} &= k\_1 \mathbf{C\_A} \mathbf{C\_P} - k\_2 \mathbf{C\_T} \mathbf{C\_W} - k\_4 \mathbf{C\_T^2} \\ \frac{d\mathbf{C\_W}}{dt} &= k\_1 \mathbf{C\_A} \mathbf{C\_P} - k\_2 \mathbf{C\_T} \mathbf{C\_W} \end{aligned} \tag{54}$$

The reaction rate constants *k*1, *k*2, *k*3 and *k*4, in Equation (54) are determined, and shown in Table 4, where *k*1, *k*2 is forward and backward reaction rate constants in Equation(51), and *k*3, *k*4 is forward reaction rate constants in Equation(52) and Equation(53), respectively,(L min−<sup>1</sup> mol−1), *T* is absolute temperature in K.

Together with the experimental data, the concentration-time curves based on the model were obtained under given reaction conditions; there is a reasonably good agreement between calculated curves and experimental points.


Table 4. Each of reaction rate constant for the esterification

#### **3. Conclusions**

This chapter discussed esterification mechanisms, and evaluated the kinetics objectively and quantitatively, which provided a most effective way to select catalyst and design reactor for different esterification systems. It is discovered that some new catalysts (such as lipases, room temperature ionic liquids) have being used in esterification; nevertheless, there are few research on the case. Herein, it is worthy to be investigated deeply.

#### **4. References**


A 2 1AP 2TW 3A

*dC kC C kC C kC*

*dC kC C kC C*

*dC kC C kC C*

1AP 2TW T 2 1AP 2TW 4T

(54)

1AP 2TW

The reaction rate constants *k*1, *k*2, *k*3 and *k*4, in Equation (54) are determined, and shown in Table 4, where *k*1, *k*2 is forward and backward reaction rate constants in Equation(51), and *k*3, *k*4 is forward reaction rate constants in Equation(52) and Equation(53), respectively,(L min−<sup>1</sup>

Together with the experimental data, the concentration-time curves based on the model were obtained under given reaction conditions; there is a reasonably good agreement

*k*0 4.249x109 1.015x106 1.584x1027 1.871x1013 *ΔE/R* 9667 7797 24710 12340

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*dC kC C kC C kC*

P

*dt*

*dt*

*dt*

*dt*

mol−1), *T* is absolute temperature in K.

**3. Conclusions** 

**4. References** 

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between calculated curves and experimental points.

Table 4. Each of reaction rate constant for the esterification

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**1. Introduction** 

**13** 

*Japan* 

Hiroshi Akatsuka

*Tokyo Institute of Technology,* 

**Progresses in Experimental Study** 

Nitrogen plasmas have been widely applied to material or electronic engineering, for example, metal-surface treatment for super-hard coating, preparation of innovative insulating layers, etc. For these processes, various plasma parameters are of great importance to control the characteristics of the prepared materials. From the engineering point of view, non-intrusive measurement method is desirable. Consequently, optical emission spectroscopy (OES) measurement is one of the best methods to examine various plasma parameters. In addition to applications, the nitrogen plasma plays an important role in many fundamental problems related to environmental issues. To understand the kinetics of radicals in the upper atmosphere or ionosphere including NOx generation or ozone destruction, we should study the kinetics of the electronic or vibrational excited states of

nitrogen discharge (Fridman, 2008; d'Agostino et al., 2008; Guerra & Loureiro, 1997).

Koike et al., 2004; Kobori et al., 2004; Yuji et al, 2007).

In most of industrial applications or geophysical phenomena, the nitrogen plasma is in a state of non-equilibrium. Consequently, the gas temperature is much lower than the electron temperature. In order to estimate approximate value to the gas temperature, rotational temperature measurement is being frequently applied in practical researches and industries (Hrachová et al., 2002). Particularly, the spectrum of the second positive system (2PS), emitted as transitions from C 3u state to B 3g state, is easy to obtain its vibrational and rotational temperatures owing to its simple transition scheme. These temperatures are determined as best theoretical fitting parameters for the spectrum observed experimentally. Observation of this band is also easy from the experimental point of view (Phillips, 1976;

In the meanwhile, the first positive system (1PS), another band in the visible wavelength region, was difficult to analyze because of the complicated selection rule of the corresponding transition from B 3g state to A 3u+ state. In order to obtain many experimental parameters as well as to understand the characteristics of non-equilibrium of each electronically excited state, quantitative measurement of 1PS of several vibrational levels is also significant for the understanding of nitrogen plasma (Sakamoto et al., 2006).

Another important parameter for the practical applications or for fundamental discussion of the nitrogen plasma is the dissociation degree, which is considered to be the one of the most

 **Optical Emission Spectroscopy** 

 **of N2 Plasma Diagnostics by** 

