Effect of Phase Behavior on the Ethenolysis of Ethyl Oleate in

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J. Phys. Chem. B 2009, 113, 2810–2814

Effect of Phase Behavior on the Ethenolysis of Ethyl Oleate in Compressed CO2 Jiyuan Song, Minqiang Hou, Gang Liu, Jianling Zhang, Buxing Han,* and Guanying Yang Beijing National Laboratory for Molecular Sciences, Institute of Chemistry, Chinese Academy of Sciences, Beijing 100190, China ReceiVed: December 4, 2008; ReVised Manuscript ReceiVed: January 11, 2009

How to enhance the reaction efficiency using greener methods is an important topic. In this work, the phase behavior of the reaction system of metathesis of ethyl oleate with ethene in compressed CO2 was studied at 308.15 and 323.15 K using the Peng-Robinson equation of state. The effect of the phase behavior on the reaction rate and equilibrium conversion was studied. It was demonstrated that addition of CO2 in the reaction system could increase the reaction rate and equilibrium conversion considerably at suitable conditions where the solubility of the reactant in the vapor phase was low, while the solubility of the products was very high. However, at the condition where the solubility of the reactant and products were all high, the reaction rate was much slower. The mechanism for this interesting phenomenon was discussed in detail. Introduction The availability of fossil organic resources is gradually decreasing. Therefore, it is important to look for alternative feedstocks. Production of raw materials from renewable resources for the chemical industry is an interesting topic in terms of its contribution toward sustainable development. Metathesis of unsaturated fatty acids is one of the important routes due to the wide availability of unsaturated fatty acids from various plant oils and animal fats.1 Olefin metathesis is a very important organic reaction, in which olefins are converted into new olefin products via the rupture and re-formation of carbon-carbon double bonds in the presence of catalyst. Ethenolysis of alkyl oleate (1) can produce 1-decene (2) and alkyl 9-decenoate (3) (Figure 1); both are very useful intermediates in the chemical industry.2 It is not surprising that many catalysts for these kinds of reactions have been designed and the reactions have been studied extensively.3 However, how to enhance the reaction efficiency is still an interesting topic. It is well-known that many supercritical fluid (SCF) technologies are environmentally more acceptable. SCFs have been used in many fields, such as extraction and fractionation,4 chemical reactions,5 material processing,6 and microelectronics.7 Chemical reactions in SCFs or under supercritical condition have many advantages.5 For example, reaction rates, yields, and selectivity can be adjusted by varying temperature and pressure. Supercritical (SC) CO2 is the most attractive among SCFs because it is nontoxic, inexpensive, and nonflammable and it has mild critical temperature and critical pressure.8 Many chemical reactions have been conducted in SC CO2, such as hydrogenation,9 hydroformylation,10 oxidation,11 esterification12 and transesterification,13 dehydration of alcohols,14 Friedel-Crafts alkylation,15 etherealization,16 Diels-Alder reaction,17 conversion of CO2 into voluble chemicals,18 and reaction of polymers.19 The unusual properties of the chemical reactions in SCFs or under supercritical state are not well-understood. One of the main reasons is that the phase behavior of the reaction systems are complex and usually are not known, and the phase behavior of the reaction systems can influence the reaction significantly.20 * To whom correspondence should be addressed. Fax:86-10-62562821. E-mail: [email protected].

Figure 1. Ethenolysis of ethyl oleate.

Figure 2. Structure of the first-generation Grubbs catalyst benzylidene-bis(tricyclohexylphosphine)dichlororuthenium.

Therefore, combination of the study of the reaction properties and the phase behavior of the reaction systems is an effective way to disclose the reason for the unusual properties of the reactions in SCFs. In this work, we first studied the phase behavior of the reaction system for the ethenolysis of alkyl oleate in the presence of CO2. Then the reaction was conducted in different phase regions. It was demonstrated that CO2 at suitable pressure could enhance the reaction rate and equilibrium conversion considerably, which depended strongly on the phase behavior of the reaction system. The mechanism for this interesting phenomenon was discussed in detail. As far as we know, this is the first work to study the effect of SC CO2 on the ethenolysis reaction of ethyl oleate. We believe that optimization of the reaction efficiency by controlling the phase behavior of reaction systems using CO2 is also applicable to some other important reactions. In addition, enhancing reaction efficiency by this method is greener because CO2 is nontoxic and widely available. Experimental Section Materials. CO2 and ethene with a purity of (>99.95%) was supplied by Beijing Analytical Instrument Factory. Ethyl oleate was produced by Beijing Chemical Agent Factory with a purity > 99% and used without further purification. The first-genera-

10.1021/jp810672e CCC: $40.75  2009 American Chemical Society Published on Web 02/09/2009

Phase Behavior on Ethenolysis of Ethyl Oleate

J. Phys. Chem. B, Vol. 113, No. 9, 2009 2811

TABLE 1: Critical Parameters and Acentric Factors of CO2, Ethyl Oleate, 1-Decene, and Ethyl 9-Decenoate critical parameters

Tc/K

Pc/MPa

ω

ref

CO2 ethene ethyl oleate 1-decene ethyl decanoate

304.2 282.4 782.0 616.4 680.82

7.38 5.031 1.21 2.218 1.788

0.225 0.089 0.992 0.487 0.742

22 22 23 24 25

focused on how the phase behavior of the reaction system affects the reaction rate and the equilibrium conversion in the presence of CO2. Phase Behavior of the Reaction System. To study the effect of the phase behavior of the reaction system on the reaction rate and conversion of the reaction, we first studied the phase behavior of the reaction system at some typical conditions using Peng-Robinson equation of state (PR-EOS).21

TABLE 2: kij and lij of Ethyl Oleate-Carbon Dioxide, Ethyl 9-Decene-Carbon Dioxide, and 1-Decene-Carbon Dioxide Binary Mixtures 308.15 K

P)

RT aR(T) V-b V(V + b) + b(V - b)

(1)

323.15 K

binary mixtures

kij

lij

kij

lij

carbon dioxide-ethyl oleate carbon dioxide-ethyl 9-decenoate carbon dioxide-1-decene

0.049 0.044 0.107

-0.010 0.010 -0.002

0.046 0.031 0.097

0.005 0.097 -0.008

R2T2c a ) 0.45724 Pc b ) 0.0778

tion Grubbs catalyst benzylidene-bis(tricyclohexylphosphine)dichlororuthenium (Figure 2) was purchased from SigmaAldrich, Inc. Apparatus. The apparatus used for the reaction was similar to that reported previously.20 It consisted mainly of a CO2 cylinder, a high-pressure pump (DB-80), a magnetic stirrer, a constant-temperature water bath, a high-pressure volumevariable view cell, a sample bomb of 10 mL, a temperature controller, and a pressure gauge. The high-pressure view cell was composed of a stainless steel body, a stainless steel piston, and two borosilicate glass windows. The two windows were installed on opposite sides of the cell over the whole height so that the state of the reaction system could be observed clearly. The volume of the view cell could be changed in the range of 20-50 mL by moving the piston. The view cell was immersed in the constant-temperature water bath controlled by a HaakeD3 temperature controller, and the temperature was measured by an accurate mercury thermometer with the accuracy of better than 0.05 K. The pressure gauge was composed of a pressure transducer (FOXBORO/ICT Model 93) and an indicator, and its accuracy was 0.025 MPa in the pressure range of 0-20 MPa. Procedures. In a typical experiment, 1.025 g of ethyl oleate with 1 mol % catalyst was charged into the reactor, and the reactor was placed in the constant-temperature water bath. After thermal equilibrium was reached, the air in the reactor was first replaced by ethene, and then ethene was charged up to 1.5 MPa. At this pressure, the molar ratio of ethene to ethyl oleate was about 4. CO2 was then compressed into the reactor to suitable pressure, and the stirrer in the reactor was started. The pressure was kept constant during the reaction by changing the volume of the reactor. After a certain reaction time the gases in the reactor were released, passing a cold trap with dimethylformamide (DMF) as the absorbent. The solution in the cold trap and the reaction mixture left in the reactor were combined and analyzed using gas chromatography (Agilent 4890D, Agilent Technologies Inc.) with hexadecane as the internal standard. Results and Discussion In the presence of ethane, ethyl oleate (1) can react with ethene (cross-metathesis) to produce 1-decene (2) and ethyl 9-decenoate (3), as illustrated in Figure 1. The other possible way is the self-metathesis and the products are 9-hexadecene and diethyl 9-octadecene-1,18-dioate. However, the self-metathesis reaction is negligible at our experimental conditions, and other byproducts were not considerable. In this work, we

(2)

RTc Pc

(3)

R(T) ) [1 + m(1 - Tr0.5)]2 m ) 0.37464 + 1.54226ω - 0.26992ω2,

(4) when ω e 0.491 (5a)

or

m ) 0.379642 + 1.48503ω - 0.164423ω2 + 0.016666ω3, when ω > 0.491 (5b) where P, V, R, Tc, Pc, ω, and Tr denote the pressure, mole volume, gas constant, critical temperature, critical pressure, acentric factor, and reduced temperature, respectively. Mixing rules should be introduced for multicomponent systems in order to use PR-EOS. Parameters aR and b are expressed according to van der Waals mixing rules as N

N

aR ) Σ Σ xixj(aR)ij, i)1 j)1

N

N

b ) Σ Σ xixjbij i)1 j)1

where (aR)ij and bij represent the cross-coefficients. The classical one-parameter van der Waals model for the cross-coefficient term (aR)ij is calculated from (aR)ij ) [(aR)i(aR)j]1/2(1 - kij). A further adjustable binary interaction parameter lij is introduced in the cross-coefficient term bij, asbij ) 0.5(bi + bj)(1 - lij). The critical properties of CO2, ethene, ethyl oleate, 1-decene, and ethyl 9-decenoate are listed in Table 1. Using the parameters in Table 1 and the vapor-liquid equilibrium data of carbon dioxide-ethyl oleate,26 carbon dioxide-ethyl decanoate, and carbon dioxide-decane27 binary systems, the kij and lij of the corresponding binary mixtures at 308.15 and 323.15 K were obtained in this work and the values are listed in Table 2. Due to the lack of the parameters of ethyl 9-decenoate and its similarity with ethyl decanoate, the parameters of ethyl decanoate were used for ethyl 9-decenoate.25 There are reactants, products, and CO2 in the reaction system. At lower pressures there exist a vapor phase and a liquid phase because the solvent power of CO2 is small. The composition of the two phases changes with temperature and pressure. With the pressure being high enough, the reaction system can exist as a single fluid phase. The solvent power of CO2 for the

2812 J. Phys. Chem. B, Vol. 113, No. 9, 2009

Figure 3. Dependence of solubility of 1-decene, ethyl 9-decenoate, and ethyl oleate in ethene + CO2 mixture at 308.15 K (the partial pressure of ethene is 1.5 MPa).

Figure 4. Dependence of solubility of 1-decene, ethyl 9-decenoate, and ethyl oleate in ethene + CO2 mixture at 323.15 K (the partial pressure of ethene is 1.5 MPa).

reactants and products, which is related directly to the solubility of the components in CO2, is crucial to govern the phase behavior of the reaction system. Therefore, we calculated the solubility of ethyl oleate, ethyl 9-decene, and 1-decene in CO2 + ethene using PR-EOS with parameters in Tables 1 and 2 at 308.15 and 323.15 K. The results calculated are shown in Figures 3 and 4. In the calculation, the kij and lij of ethene-ethyl oleate, ethene-ethyl 9-decenoate, ethene-1-decene, and ethene-carbon dioxide binary systems were set to zero. This should not result in considerable error because the amount of ethene is small, and this approximation is commonly used in this case.28 The calculation simulated the reaction system, i.e., the partial pressure of ethene was 1.5 MPa, and the pressure shown in the figure is the total pressure after adding CO2. It is known that some ethene was consumed during the reaction. The decrease of ethene partial pressure was at most 0.4 MPa in our reaction system, which was compensated for by that of CO2 because our experiment was conducted at constant pressure by changing the volume of the reactor during the reaction. To know how the change of vapor-phase composition affected the solubility of the reactant ethyl oleate and products, we also calculated the solubility of the reactant and the products with ethene partial pressure of 1.1 MPa at the same total pressure. The results indicated that the difference in the solubility at these two ethene partial pressures was not considerable. It is known that CO2 is soluble in liquids, and the solubility depends on temperature and pressure.29 Dissolution of CO2 can expand the liquid and change the physical properties such as viscosity.30 The solubility of CO2 in ethyl oleate at 308.15 and 323.15 K were also calculated, and the results were presented in Figure 5. We also determined the solubility data using the apparatus and procedures reported previously,20a and the results are also presented in Figure 5 for comparison. Figure 5 shows that the calculated data agree very well with the calculated values, verifying the reliability of the calculation method.

Song et al.

Figure 5. Solubility (mole fraction) of CO2 in ethyl oleate at 308.15 and 323.15 K. The points were the data determined, and lines were the calculated values using PR-EOS.

Figure 6. Effect of reaction time and the total pressure on conversion of ethyl oleate at 308.15 K.

Effect of Phase Behavior on the Reaction. On the basis of the phase behavior study, we conducted the reaction at 308.15 K under four typical pressures, without CO2 and with CO2 at total pressure of 5.0, 8.2, and 12.0 MPa, and the results are shown in Figure 6. The figure demonstrates that in the absence of CO2 the reaction could reach equilibrium in about 1 h, and the equilibrium conversion was about 80%. At 5.0 MPa, the reaction rate was faster at the beginning than that without CO2, and the maximum conversion was higher than that in the absence of CO2. At 8.2 MPa, the reaction was faster than that at 5.0 MPa and the maximum conversion was also considerably larger. As the pressure was increased to 12.0 MPa, however, the reaction rate was much slower. This interesting phenomenon can be explained on the basis of the variation of the phase behavior of the reaction system discussed above. At 5.0 MPa, the concentrations of the reactant ethyl oleate and the products in the vapor phase were very low (Figure 3). Therefore, the change of the reaction rate and equilibrium conversion resulted mainly from the dissolution of CO2 in the liquid phase where the catalyst existed and the reaction took place. For simplicity, we discuss qualitatively the effect of CO2 dissolution in the liquid phase on the reaction using the solubility data of CO2 in the reactant shown in Figure 5. At 5.0 MPa, the partial pressure of CO2 was about 3.5 MPa, and the mole fraction of CO2 in carbon dioxide-ethyl oleate was about 0.5. It is well-known that the dissolution of CO2 can reduce the viscosity of liquids.30 Therefore, dissolution of CO2 in the reaction mixture could reduce the viscosity of the liquid phase of the reaction mixture and enhance the mass transfer, and then the reaction rate became faster. At the same time, dissolution of CO2 in the liquid phase also affects the equilibrium conversion, which is discussed as follow. For a reversible reaction, A + B ) C + D, the thermodynamic equilibrium constant can be expressed as

Phase Behavior on Ethenolysis of Ethyl Oleate

Ka )

acad aaab

J. Phys. Chem. B, Vol. 113, No. 9, 2009 2813

(6)

where ai stands for the activity of component i. Ka is a constant at fixed temperature, which can also be expressed as follows.

Ka ) KxKγ

(7)

Kx and Kγ are defined by the following equations.

Kx )

Kγ )

Figure 7. Effect of reaction time and total pressure on conversion of ethyl oleate at 323.15 K.

XcXd XaXb

(8)

γcγd γaγb

(9)

where γa, γb, γc, and γd are activity coefficients of A, B, C, and D, respectively. Kγ should be changed after adding CO2 because CO2 affects the activity coefficients of the components, while Ka depends only on temperature. Therefore, Kx, which is directly related with the equilibrium conversion, is changed after adding CO2. At 8.2 MPa, the solubility of the reactant in the vapor phase was not considerable, while the solubility of the two products was high, as can be known from Figure 3. Therefore, the most products formed were extracted into the vapor phase during the reaction. In our experiment, we observed that the volume of the liquid phase reduced continuously with reaction time, while, at 5.0 MPa, the change of the volume with time was not noticeable. This confirmed the usability and the reliability of the phase behavior data calculated in this work (Figure 3). At this pressure, existence of CO2 was favorable to the reaction both kinetically and thermodynamically. On one hand, the solubility of CO2 in the liquid phase is high, which reduced the viscosity of the reaction mixture, and the reaction rate was faster. On the other hand, the products formed were extracted into the vapor phase, which shifted the reaction equilibrium. Therefore, equilibrium conversion was higher. Figure 3 shows that all the solubilities of the reactant and products in CO2 are very high at 12.0 MPa. Our experiments also showed that all the reactant was extracted into the vapor phase after charging CO2 to this pressure. In other words, there was no liquid phase in the reaction system. This further verifies the reliability of the results in Figure 3. Figure 6 demonstrates that at this pressure the reaction rate was even much slower than that of the reaction without CO2. The reaction equilibrium could not be reached even though the reaction time was prolonged to 3 h, as can be seen from Figure 6. We observed through the windows of the reactor that the catalyst existed at the bottom of the reactor during the reaction at 12 MPa; i.e., the catalyst did not dissolve in the vapor phase because the reaction mixture contained a large amount of CO2, which acted as antisolvent for the catalyst. This is the main reason why the reaction rate was much slower at 12 MPa. In the absence of CO2 and at 5.0 and 8.2 MPa, however, the catalyst could dissolve in the liquid phase in which the reaction took place, which was known by direct observation through the windows of the reactor, and the reaction rate was much faster. This suggests that existence of the liquid phase is necessary for the high reaction rate. Therefore, we can conclude that the reaction efficiency can be optimized by controlling the phase behavior

of the reaction system. The efficiency is highest at 8.2 MPa where the solubility of the reaction products is very high, while that of the reactant ethyl oleate is not considerable. We also conducted the reaction at 323.15 K. On the basis of the phase behavior study at this temperature (Figure 4), we carried out the reaction in the absence of CO2 and at total pressure of 10.4 MPa. The conversion as a function of reaction time is shown in Figure 7. Similarly, the CO2 could enhance the reaction rate and the equilibrium conversion effectively. As can be known from Figure 4, the solubility of the reactant in the vapor phase was very low at 10.4 MPa, while the solubility of the two products was high. We also observed in our experiments that the volume of the liquid phase in the reaction system reduced with reaction time, similar to that at 308.15 K and 8.2 MPa, as discussed above. Similarly, the main reasons for the larger equilibrium conversion and reaction rate were that CO2 extracted the products into the vapor phase and reduced the viscosity of the liquid phase. Conclusion The metathesis of ethyl oleate with ethene was studied at 308.15 and 323.15 K in the presence of CO2. It was demonstrated that at 308.15 K, the reaction rate and equilibrium conversion of the reaction at 5.0 and 8.2 MPa are considerably higher than that in the absence of the CO2. Specifically, at 8.2 MPa where the solubility of the reactant in the vapor phase was low, while the solubility of the products was very high, the reaction efficiency is the highest. The main reasons are that CO2 in the liquid phase can reduce the viscosity, and the products can be extracted into the vapor phase. However, at 12.0 MPa where the solubility of the reactant and products were all high, the reaction rate was much slower because the liquid phase disappears and the catalyst cannot be well-dispersed. The existence of the liquid phase is necessary for the high reaction rate. Similarly, at 323.15 K, the reaction efficiency at 10.4 MPa is considerably higher than that in the absence of CO2. We believe that the reaction efficiency of some other important reactions can also be optimized by controlling the phase behavior of their reaction systems using CO2. Acknowledgment. This work was supported by the National Key Basic Research Project of China (Grant 2006CB202504), the Chinese Academy of Sciences (Grant KJCX2.YW.H16), and the National Natural Science Foundation of China (Grant 20733010). References and Notes (1) Corma, A.; Lborra, S.; Velty, A. Chem. ReV. 2007, 107, 2411. (2) Burdett, K. A.; Harris, L. D.; Margl, P.; Maughon, B. R.; MokhtarZadeh, T.; Saucier, P. C.; Wasserman, E. P. Organometallics 2004, 23, 2027.

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