Experimental Investigation on the Solubility of Oxygen in Toluene and

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Experimental Investigation on the Solubility of Oxygen in Toluene and Acetic Acid Xiankun Wu,†,‡ Zilei Deng,†,‡ Jiujuan Yan,†,‡ Zhongyang Zhang,†,‡ Feng Zhang,*,†,‡ and Zhibing Zhang*,†,‡ †

School of Chemistry and Chemical Engineering, Nanjing University, Nanjing 210093, China Jiangsu Engineering Research Center for Organic Pollution Control and Resources Reuse, Nanjing 210046, China

‡

S Supporting Information *

ABSTRACT: By using the dual vessel equilibrium method, the solubility of oxygen in toluene and acetic acid was experimentally investigated with pressure ranging from 0.1 to 1.0 MPa at temperatures from 293.1 to 383.1 K. The results show that the oxygen solubility either in toluene or in acetic acid increases with a rise in temperature. On the basis of the experimental data, Henry coefficients were derived and expressed as a function of temperature. At the same temperature, Henry coefficient for the oxygen− toluene system is lower than that for the oxygen−acetic acid system. Through analysis of the Gibbs energy (ΔG0), partial molar enthalpy (ΔH0), and the partial molar entropy (ΔS0) of the solvation, we can know that solubilization of oxygen either in toluene or in acetic acid is an endothermic process. To correlate the experimental data, the Peng−Robinson equation of state with the quadratic mixing rule was used for the two systems.

1. INTRODUCTION Catalytic oxidation of liquid toluene in acetic acid system is a significant reaction process. During the oxidation, toluene can be converted into a series of oxidation products, mainly including benzyl alcohol, benzaldehyde, and benzoic acid.1−3 Amomg them, benzaldehyde is the most desirable and valueadded product. Since benzaldehyde is easily overoxidized to benzoic acid, its yield depends mostly on the solubility of oxygen in the liquid phase of the reaction system. Therefore, adequate information about solubility of oxygen is needed to improve the control of the benzaldehyde selectivity and yield in this oxidation.4 Until now, for the system of oxygen−toluene, some experimental data are available in the literature. Field et al.5 reported oxygen solubility in toluene at a normal pressure and room temperature. Fischer et al.6 provided information about the solubility of oxygen up to 10.0 MPa at temperatures between 298.41 and 348.29 K, while Li et al.7 investigated the solubility of oxygen in toluene at higher temperatures ranging from 298.45 to 393.15 K. However, more data on oxygen solubility obtained by different methods in toluene at reaction pressure and temperature are still essential. Acetic acid is a common solvent for toluene oxidation, as well as for oxidation of p-xylene to terephthalic acid which plays an extremely important role in the modern industry.8−11 Lang12 and Rischbieter et al.13 have investigated the solubility of oxygen in aqueous acetic acid. However, few data were reported about oxygen solubility in pure acetic acid, especially under oxidation reaction pressure and temperature. As for oxidation of the liquid substances by oxygen, the mass transfer rate of gas to liquid phase has significant influence on the oxidation rate. In order to increase the reaction rate, a relatively high pressure for improving oxygen solubility in liquid is usually adopted. Oxidation of liquid toluene by oxygen or p-xylene by air is often performed under a higher pressure of 0.5 MPa. However, the pressure has to be controlled under a certain value for © 2014 American Chemical Society

reasons of safety and energy conservation. In addition, the volume fraction of oxygen in gas is usually restricted to less than 7%.7 Therefore, the solubility data of oxygen either in toluene or in acetic acid are essential for optimizing the design and oxidation process. It is good to use a thermodynamic model, such as an equation of state combining with a mixing rule, to predict the gas−liquid equilibrium when no experimental data are available.14 Fischer et al.6 reported that the predictive Soave− Redlich−Kwong equation of state (PSRK EoS) with the Huron−Vidal mixing rule could bring an accurate prediction for oxygen and nitrogen solubility in organic solvents. Li et al.7 also used PSRK EoS with the modified Huron−Vidal mixing rule to adjust the experimental data and showed that the predicted data were in good agreement with experimental one. Windmann et al.15 improved the Peng−Robison equation of state (PR EoS) with the quadratic mixing rule or the Huron−Vidal mixing rule to correlate the experimental data for the solubility of nitrogen in acetone and oxygen in acetone and noted that the results of the quadratic mixing rule were very similar to those from the Huron−Vidal mixing rule. Tenório Neto et al.16 found that the experimental data were satisfactorily represented by the PR EoS with the quadratic mixing rule. The results confirmed that PR EoS with the quadratic mixing rule or the Huron−Vidal mixing rule can be used for good prediction of the oxygen (nitrogen)− organic solvent equilibrium. Since the empirical mixing parameter in the Huron−Vidal mixing rule was obtained at atmospheric pressure, the calculated results at higher pressures are not very satisfactory. Then to extend the present experimental results for further applications, the PR EoS with Received: Revised: Accepted: Published: 9932

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process for the two solvents, respectively. The two vessels were heated to the scheduled temperature, and the vapor pressure (Ps) of the solvent was recorded by the pressure sensor E. Then valve 1 was closed and oxygen was introduced into the storage vessel until the pressure reached a given value (P1). Finally, valve 4 was opened to transfer oxygen from the storage vessel F to absorption vessel J until its pressure reached a stable value (P2). 2.3. Data Processing. The solubilities of oxygen in the different solvents can be expressed in mole fraction,21−25 nG x= nL + nG (1)

the quadratic mixing rule is employed to correlate the experiment data and to obtain the binary interactive parameters.15,17 In this work, a new apparatus was set up to measure the solubility of oxygen both in toluene and in acetic acid. Six isotherms data will be determined for both systems under pressure from 0.1 to 1.0 MPa and temperatures from 293.1 to 383.1 K. The Henry coefficients were calculated on the basis of the measured data by using PR EoS with the quadratic mixing rule.

2. EXPERIMENTAL SECTION 2.1. Materials. Oxygen purity was 99.999%, manufactured by Nanjing Tianze Co. Ltd., China. Toluene and acetic acid, analytical grade and obtained from Nanjing Chemical Reagent Co. Ltd., China, were degassed under vacuum. 2.2. Apparatus and Procedure. Apparatus for measuring oxygen solubility in the toluene or in acetic acid is similar to that used for CO2 absorption in our previous work.18−20 The schematic diagram of the static apparatus is shown in Figure 1.

where nG represents the moles of oxygen dissolved in a solvent and nL the moles of the solvent in the absorption vessel. Since the temperature, initial and equilibrium value of the pressure, and the volume of the both vessels were measured by the above method, nG can be calculated as follows:24,26 nG =

(P − Ps)(V1 + V2 − m /ρ) PV 1 1 − 2 Z1RT Z 2RT

(2)

where V1 and V2 denote the volumes of storage vessel and absorption vessel, respectively; P1 is the initial pressure of oxygen in the storage vessel; P2 is the equilibrium pressure of oxygen in the both vessels; Ps is the saturated vapor pressure of the solvent at temperature T; m and ρ are the mass and the density of the solvent in the storage vessel at temperature T, respectively; Z1 and Z2 are denoted as the compressibility factor of oxygen at the initial pressure P1 and the equilibrium pressure P2 .

3. RESULTS AND DISCUSSION In the present work, the solubility data for oxygen either in toluene or in acetic acid were measured at six temperatures (293.1, 303.1, 323.1, 343.1, 363.1, and 383.1 K) with the equilibrium pressure up to 1.0 MPa. The isothermal data for the partial pressure of oxygen (P), and greater fraction of oxygen in the solution (x), are listed in Tables 1−4 and plotted

Figure 1. Schematic diagram of the static apparatus: A, valve 1; B, valve 2; C, oxygen cylinder; D, valve 3; E, pressure sensor; F, storage vessel; G, valve 4; H, pressure sensor; J, absorption vessel; L, valve 5; M, temperature controller; N, magnetic stirrer; O, oil bath.

Table 1. Experimental P−x Data for Oxygen−Toluene at Different Temperatures

The apparatus mainly consists of two stainless steel (AISI 316L) vessels. The storage vessel (V1 = 168.37 cm3) isolated oxygen before the gas got in touch with toluene or acetic acid in the absorption vessel (V2 = 49.35 cm3), in which the liquid in the vessel can be magnetically stirred. The temperature of both vessels is maintained by an oil bath. The pressure is recorded every 5 s with pressure sensors of different measuring range. The pressure sensor E with a range from 0 to 0.1 MPa was used to monitor the vapor pressure of solvent in the absorption vessel. The pressure sensor H with a range from 0 to 1 MPa was employed to measure the initial pressure (P1) in the storage vessel and equilibrium pressure (P2). The pressure sensors were supplied by Fujing Wideplus Precision Instruments Co., Ltd., with measuring uncertainty of 0.2% in relation to the full scale. In the experiment, first, with all valves except for valve 4 open, a precisely weighed amount of solvent (about 25.10g) was placed into the absorption vessel. Then it was degassed three times by a vacuum pump at a low temperature (271.1 K) and valve 5 was closed. As calculated by the initial mass and final mass of the solvent and after degassing, about 0.10 g of toluene and 0.08 g of acetic acid were lost during the degassing

T = 293.1 K

T = 303.1 K

T = 323.1 K

x

P, MPa

x

P, MPa

x

P, MPa

0.001 20 0.002 18 0.003 32 0.003 64 0.004 69 0.005 79 0.006 93 0.007 92

0.143 0.263 0.395 0.434 0.545 0.696 0.827 0.960

0.001 43 0.002 64 0.003 85 0.004 78 0.005 36 0.006 28 0.007 99 0.008 48

0.153 0.293 0.429 0.526 0.590 0.703 0.813 0.944

0.002 17 0.003 54 0.004 69 0.005 96 0.006 96 0.008 15 0.008 94 0.009 69

0.221 0.364 0.489 0.607 0.705 0.808 0.897 0.974

in Figures 2 and 3, respectively. As can be seen in Tables 1−4, the solubilities of oxygen both in toluene and in acetic acid increase with a rise of temperature. However, toluene has a larger solubility of oxygen than acetic acid does under the same conditions. The experimental data were correlated with PR EoS using the quadratic mixing rule. As shown in Figures 2 and 3, under most of regarded temperature, the experimental data agree well with PR EoS combined with the quadratic mixing rule for both toluene and acetic acid systems. Only slight 9933

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Table 2. Experimental P−x Data for Oxygen−Toluene at Different Temperatures T = 343.1 K

T = 363.1 K

T = 383.1 K

x

P, MPa

x

P, MPa

x

P, MPa

0.001 36 0.002 78 0.004 12 0.005 38 0.006 68 0.007 59 0.008 61 0.009 90

0.128 0.264 0.387 0.512 0.625 0.723 0.827 0.926

0.001 65 0.002 83 0.003 97 0.005 22 0.006 42 0.007 67 0.008 58 0.010 10

0.147 0.258 0.359 0.478 0.584 0.682 0.789 0.903

0.001 31 0.002 56 0.003 88 0.005 28 0.006 50 0.007 76 0.008 85 0.010 10

0.112 0.221 0.338 0.448 0.561 0.667 0.768 0.876

Table 3. Experimental P−x Data for Oxygen−Acetic Acid at Different Temperatures T = 293.1 K

T = 303.1 K

T = 323.1 K

x

P, MPa

x

P, MPa

x

P, MPa

0.000 68 0.001 15 0.001 62 0.002 05 0.002 49 0.002 91 0.003 38 0.004 00

0.162 0.278 0.386 0.497 0.600 0.708 0.816 0.977

0.000 83 0.001 15 0.001 51 0.002 14 0.002 72 0.003 12 0.003 51 0.003 99

0.194 0.267 0.347 0.504 0.629 0.718 0.820 0.923

0.000 70 0.001 21 0.001 59 0.002 25 0.002 86 0.003 27 0.003 87 0.004 40

0.150 0.259 0.344 0.486 0.614 0.715 0.832 0.950

Figure 3. Pressure (P) versus the mole fraction of oxygen (x) in acetic acid.

Table 4. Experimental P−x Data for Oxygen−Acetic Acid at Different Temperatures T = 343.1 K

T = 363.1 K

T = 383.1 K

x

P, MPa

x

P, MPa

x

P, MPa

0.000 71 0.001 29 0.001 71 0.002 28 0.002 85 0.003 40 0.003 95 0.004 49

0.143 0.264 0.350 0.470 0.584 0.700 0.804 0.928

0.000 89 0.001 70 0.002 26 0.002 83 0.003 35 0.003 82 0.004 24 0.004 67

0.175 0.334 0.440 0.554 0.649 0.754 0.824 0.919

0.000 61 0.001 32 0.002 00 0.002 55 0.003 11 0.003 67 0.004 05 0.004 52

0.116 0.252 0.378 0.490 0.591 0.699 0.774 0.870

Figure 2. Pressure (P) versus the mole fraction of oxygen (x) in toluene.

deviations are found for oxygen−toluene system at 383.1 K and oxygen−acetic acid system at 363.1 and 383.1 K. The data of oxygen dissolved in toluene have been reported previously,5−7 and a comparison of this work with the literature has been made in Figure 4. It can be seen that the line of this work is very much similar to that of refs 6 and 7 at almost the same temperature (about 298 K). Because of the smaller temperature deviation and the difference of experimental method and data processing, the slight deviation between our data and literature data was acceptable. This confirmed the correctness of the method used in the present work. Theoretically, for a given temperature, Henry coefficients can be calculated with P and x data:27

Figure 4. Comparison of solubility data of oxygen in toluene.

⎡ f (T , P , y ) ⎤ i ⎥ Hi(T ) = lim ⎢ i P → Ps⎣ xi ⎦

(3)

where fi (T , P , yi ) = pyi φi(T , P , yi ) 9934

(4)

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f i and φi are the fugacity and the fugacity coefficient of oxygen, respectively. xi and yi are the mole fractions of oxygen in the liquid and vapor phase, respectively. It can be seen in Table 5

Generally, the partial molar enthalpy of solvation indicates the strength of interaction between the dissolved gas and the solvent, and partial molar entropy yields information about the level of ordering present in the gas/solvent mixture.34 As shown in Table 7, the partial enthalpies of solvation are positive values for both systems corresponding to an endothermic solubilization, indicating that the solubilization is thus favored at higher temperatures. Moreover, the higher absolute value for the partial molar enthalpy of the oxygen− toluene system suggests a more favorable interaction between the oxygen and the toluene than oxygen with acetic acid. The values for the entropy of the solvation are negative at temperature 293.1 K for both systems, and the more negative value for oxygen−acetic acid entropy reveals a higher ordering degree as oxygen dissolves in acetic acid. However, the deviation is not so significant enough to proceed with further analysis revealing that the energetic effect is more dominant than the entropic effect.35

Table 5. Henry Coefficients for Oxygen in Toluene and Acetic Acid Derived from the Experimental P−x Data Hi, MPa T, K

oxygen−toluene

oxygen−acetic acid

293.1 303.1 323.1 343.1 363.1 383.1

120.6 110.7 98.8 94.5 89.9 86.4

242.5 231.9 216.6 206.3 196.4 192.2

that Henry coefficients of oxygen solubility in toluene or acetic acid dropped with an increase in temperature, meaning that more oxygen can be dissolved in both solvents at a higher temperature. Under a lower pressure, the Henry coefficients can be considered as independent parameter on pressure.28 Then Henry coefficients of oxygen in both solvents can be correlated with a function of temperature and defined as7,29−31 A ln(Hi) = +B (5) T where A and B are equation parameters and |Hi| is the average absolute deviations of Henry coefficient between the experimental data and calculated results with eq 5, as shown in Table 6.

4. PENG−ROBINSON EQUATION OF STATE In this work, the Peng−Robinson equation of state was chosen to correlate the experimental data:

A

B

|Hi|, %

toluene acetic acid

400.426 293.780

3.393 4.478

1.92 0.8

⎛ R2Tc 2 ⎞⎡ ⎟⎢1 + (0.37464 + 1.5422ω a = ⎜0.45724 Pc ⎠⎢⎣ ⎝ ⎛ − 0.26992ω 2)⎜⎜1 − ⎝ b = 0.07780

ΔS 0 =

⎞⎤ ⎟⎟⎥ ⎠⎥⎦

(10)

RTc Pc

(11)

∑ ∑ xixj(aiaj)1/2 (1 − δij)

aM =

i

(12)

j

∑ xibi

bM =

(6)

⎛ ∂ΔG 0 ⎞ 2⎛ ∂ ln Hi ⎞ ⎟ ΔH 0 = −T 2⎜ ⎟ = −RT ⎜ ⎝ ∂T ⎠ P ⎝ ∂T ⎠T

T Tc

where Tc is the critical temperature, Pc is the critical pressure, and ω stands for the acentric factor.37,38 According to quadratic mix rules,39

The influence of temperature on oxygen solubility in toluene or in acetic acid can also be correlated with the Gibbs energy (ΔG0), partial molar enthalpy (ΔH0), and the partial molar entropy (ΔS0) of solvation, which can be calculated from the correlation of Henry coefficient based on the following relationship shown in eqs 6−8:30,32,33 ΔG 0 = RT (ln(Hi))P

(9)

where R is the universal gas constant, V is the molar volume, and a and b are equation of state dependent parameters, defined as follows:36

Table 6. Parameters A and B and the Calculated Results |Hi| with Eq 5 solvent

a(T ) RT − V−b V (V + b) + b(V − b)

P=

(13)

i

where δi j is an adjustable binary interaction parameter to correlate experimental data depending on temperature. The optimal value of δi j can be obtained by nonlinear least squares which is used as the objective function (OF),40,41

(7)

0 ⎛ ∂ ln Hi ⎞ ΔH 0 − ΔG ⎟ − R ln(Hi)P = −RT ⎜ ⎝ ∂T ⎠ P T

N

OF = (8)

∑ [(f1Li

− f 1Gi ) − (f2Li − f 2Gi )]2

(14)

i

Table 7. Gibbs Energy, Partial Molar Enthalpy, and Partial Molar Entropy of Solvation Obtained for Oxygen−Toluene and Oxygen−Acetic Acid solvent

ΔGi|T=293.1K, kJ·mol−1

ΔHi, kJ·mol−1

ΔSi|T=293.1K, J·mol−1·K−1

toluene acetic acid

−11.68 −13.38

3.33 2.44

−28.49 −37.32

9935

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where N is the number of data, and f L and f G are the fugacity of the gas and liquid, respectively.

5. CONCLUSIONS In this work, solubility data (isothermal P−x data) for oxygen in toluene and acetic acid were measured over a wide range of temperature and pressure. The reliability of the experimental data was validated by a comparison with literature data. On the basis of this information, the Henry coefficient was derived for both binary systems. Meanwhile, empirical equations of the Henry coefficient as a function of temperature were obtained for oxygen in both solvents. The observed results showed that the solubility of oxygen in toluene was higher than in acetic acid, while the Henry coefficient of oxygen in toluene or in acetic acid decreased with increasing temperature. The partial molar enthalpy of solvation estimated from the Henry coefficient suggested that the oxygen-dissolving process in both solvents was endothermic. Furthermore, the Peng− Robinson EoS with the quadratic mixing rule was adjusted to obtain experimental data for both systems with a good agreement.

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REFERENCES

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ASSOCIATED CONTENT

S Supporting Information *

Measurement uncertainties of pressure, temperature, volume, and mass; information on the critical properties and acentric factor for oxygen, toluene, and acetic acid; values of binary interaction parameters. This material is available free of charge via the Internet at http://pubs.acs.org.

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ΔS0 = partial molar entropy of solvation, J·mol−1·K−1 f i = fugacity φi = fugacity coefficient A, B = equation parameters R = universal gas constant, ∼8.14 J·mol−1·K−1 V = molar volume, L/mol ai , bi = pure substance parameters aM, bM = mixing substance parameters Tc = critical temperature, K Pc = critical pressure, MPa ω = acentric factor f L = fugacity of liquid f G = fugacity of gas δij = binary interaction parameter

AUTHOR INFORMATION

Corresponding Authors

*F.Z.: tel, 86-25-83596665; fax, 86-25-83593772; e-mail, zf@ nju.edu.cn. *Z.Z.: tel, 86-25-83593772; fax, 86-25-83593772; e-mail, [email protected]. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS This work was financially supported by National Natural Science Foundation of China (Grant 21306078), Natural Science Foundation of Jiangsu Province (Grant BK2011633), and Environmental Protection Scientific Project of Jiangsu Province (Grant 2012017).

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SYMBOLS USED V1 = volume of storage vessel, cm3 V2 = volume of absorption vessel, cm3 P1 = initial pressure, MPa P2 = equilibrium pressure, MPa Ps = vapor pressure, MPa x = mole fraction nG = moles of oxygen dissolved in solvents nL = moles of the solvents Z1 = compressibility factor of oxygen at initial pressure Z2 = compressibility factor of oxygen at equilibrium pressure P = partial pressure of oxygen Hi = Henry coefficient, MPa |Hi| = average absolute deviation of Henry coefficient ΔG0 = Gibbs energy of solvation, KJ·mol−1 ΔH0 = partial molar enthalpy of solvation, KJ·mol−1 9936

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dx.doi.org/10.1021/ie5014772 | Ind. Eng. Chem. Res. 2014, 53, 9932−9937