Measurement and Correlation of the Solubility for 4, 4

Jun 29, 2015 - The modified Apelblat, λh, and van,t Hoff equations were employed to correlate experimental solubility data, and each selected correla...
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Measurement and Correlation of the Solubility for 4,4′Diaminodiphenylmethane in Different Solvents Yijie Deng, Li Xu,* Xiaobo Sun, Liang Cheng, and Guoji Liu* School of Chemical Engineering and Energy, Zhengzhou University, Zhengzhou, Henan 450001, People’s Republic of China S Supporting Information *

ABSTRACT: In this experiment, the solubility data of 4,4′-diaminodiphenylmethane in different pure solvents, including methanol, ethanol, 2-propanol, 1butanol, toluene, chloroform, and benzene, were measured by a synthetic method from (293.15 to 333.15) K at atmospheric pressure. It is indicated that the solubility of 4,4′-diaminodiphenylmethane in each solvent increases with increasing temperature. The modified Apelblat, λh, and van’t Hoff equations were employed to correlate experimental solubility data, and each selected correlation equation could give a good fitting result of the relationship between solubility and temperature. The modified Apelblat equation shows the best agreement, in general. Furthermore, the standard molar enthalpy, standard molar Gibbs energy, and standard molar entropy of 4,4′-diaminodiphenylmethane in the dissolution process were calculated based on the van’t Hoff analysis and Gibbs equation.

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INTRODUCTION 4,4′-Diaminodiphenylmethane, also known as 4,4′-methylenedianiline (MDA, C13H14N2, CAS registry No.101-77-9, Figure 1), is usually used as monomer, chain extender, or cross-linking

equation, and van’t Hoff equation. In addition, the thermodynamic properties in the dissolution process for the system, including standard molar enthalpies, standard molar entropies, and the standard molar Gibbs energies, were also calculated by the methods of van’t Hoff analysis and the Gibbs equation.8−10

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EXPERIMENTAL SECTION Materials. 4,4′-Diaminodiphenylmethane was purchased from Aladdin Reagent Co., Ltd., and the purity of MDA was 0.990 in mass fraction. The rude materials (MDA) were purified by recrystallization in ethanol, and the purity of MDA was 0.996 in mass fraction, as determined by gas chromatography (GC) (Shimadzu GC-2010) in our laboratory. All of the organic solvents (obtained from Tianjin Kemel Chemical Reagents Co., Ltd.) were analytical reagent grade with mass fraction purities greater than 0.99, as determined by GC, and were used without any further purification. Apparatus and Procedure. In this study, the solubility of MDA in different solvents was measured using a synthetic method.11−13 The experimental procedure was described in our previous work.14 A jacketed dissolution vessel was employed to dissolve the solute in the experiment, and the working volume of the glass equilibrium cell was about 50 mL. The equilibrium cell was heated at a desired temperature, and a water circulator was used to keep the experimental temperature constant. To prevent the solvent from evaporating, a reflux condenser was directly connected to the vessel. The dissolving process of MDA was observed using the laser detecting system. The temperature of solution was measured

Figure 1. Chemical structural formula of MDA.

agent in the field of polymers.1−3 The pure MDA is a white crystalline solid that has chemical properties similar to those of m-phenylenediamine. In industrial mass production, MDA is usually produced by condensation of aniline and formaldehyde. The crystallization process of MDA is the critical step that determines the quality of the final product.4 It is well-known that the solubility data of the solid compound in different solvents play an important role in the development and operation of crystallization processes.5 Moreover, it is of great importance for obtaining MDA with high quality and a good crystal habit to determine the solid−liquid equilibrium solubility of MDA in different solvents. Therefore, in order to provide essential fundamental data for industrial crystallization and production, detailed measurements for the solubility data of MDA are required. However, until to now, no experimental solubility data of MDA in pure organic solvents have been reported in the literature. In the present work, the solubility of MDA in commonly used solvents was determined from (293.15 to 333.15) K at atmospheric pressure by using a synthetic method with a laser monitoring observation technique.6,7 The experimental data were correlated with the modified Apelblat equation, λh © 2015 American Chemical Society

Received: December 7, 2014 Accepted: June 10, 2015 Published: June 29, 2015 2028

DOI: 10.1021/je501111d J. Chem. Eng. Data 2015, 60, 2028−2034

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Table 1. Experimental and Correlated Mole Fraction Solubility (x1) of MDA in Different Pure Solvents at Different Temperatures and Pressures (p = 0.1 MPa)a λh equation

Apelblat equation solvent

methanol

ethanol

2-propanol

1-butanol

benzene

toluene

chloroform

van’t Hoff equation

T (K)

100x1exp

100x1cal

100RD

100x1cal

100RD

100x1cal

100RD

293.15 298.15 303.15 308.15 313.15 318.15 323.15 328.15 293.15 298.15 303.15 308.15 313.15 318.15 323.15 328.15 293.15 298.15 303.15 308.15 313.15 318.15 323.15 328.15 333.15 293.15 298.15 303.15 308.15 313.15 318.15 323.15 328.15 333.15 293.15 298.15 303.15 308.15 313.15 318.15 323.15 293.15 298.15 303.15 308.15 313.15 318.15 323.15 328.15 333.15 293.15 298.15 303.15 308.15 313.15 318.15 323.15

12.72 15.72 18.70 21.91 25.85 30.04 34.72 40.25 3.35 4.61 6.37 8.42 11.04 15.02 19.79 26.74 1.10 1.36 1.78 2.36 3.12 4.11 5.36 6.96 8.98 1.87 2.31 2.96 3.71 4.76 5.82 7.13 8.85 10.78 2.07 2.99 3.90 4.99 6.60 8.65 11.42 1.56 1.95 2.44 3.00 3.80 4.87 5.99 7.50 9.32 12.35 14.23 16.52 19.70 23.01 26.85 30.64

12.93 15.56 18.56 21.96 25.80 30.09 34.87 40.14 3.47 4.65 6.22 8.33 11.15 14.92 19.96 26.68 1.04 1.37 1.81 2.39 3.13 4.10 5.34 6.95 9.00 1.84 2.35 2.97 3.74 4.68 5.82 7.19 8.83 10.78 2.19 2.89 3.82 5.03 6.61 8.69 11.39 1.54 1.94 2.43 3.06 3.83 4.80 6.00 7.48 9.33 12.14 14.32 16.82 19.67 22.93 26.62 30.79

−1.65 1.02 0.75 −0.23 0.19 −0.17 −0.43 0.27 −3.58 −0.87 2.35 1.07 −1.00 0.67 −0.86 0.22 5.45 −0.74 −1.69 −1.27 −0.32 0.24 0.37 0.14 −0.22 1.60 −1.73 −0.34 −0.81 1.68 0.00 −0.84 0.23 0.00 −5.80 3.34 2.05 −0.80 −0.15 −0.46 0.26 1.28 0.51 0.41 −2.00 −0.79 1.44 −0.17 0.27 −0.11 1.70 −0.63 −1.82 0.15 0.35 0.86 −0.49

13.09 15.62 18.54 21.88 25.68 29.99 34.85 40.31 3.06 4.34 6.08 8.39 11.41 15.29 20.19 26.23 1.07 1.40 1.84 2.40 3.13 4.07 5.29 6.90 9.06 2.00 2.47 3.04 3.74 4.60 5.67 7.02 8.74 11.00 2.13 2.87 3.82 5.06 6.66 8.71 11.35 1.60 1.99 2.47 3.07 3.81 4.74 5.92 7.44 9.43 12.19 14.34 16.81 19.64 22.88 26.60 30.85

−2.91 0.64 0.86 0.14 0.66 0.17 −0.37 −0.15 8.66 5.86 4.55 0.36 −3.35 −1.80 −2.02 1.91 2.73 −2.94 −3.37 −1.69 −0.32 0.97 1.31 0.86 −0.89 −6.95 −6.93 −2.70 −0.81 3.36 2.58 1.54 1.24 −2.04 −2.90 4.01 2.05 −1.40 −0.91 −0.69 0.61 −2.56 −2.05 −1.23 −2.33 −0.26 2.67 1.17 0.80 −1.18 1.30 −0.77 −1.76 0.30 0.56 0.93 −0.69

13.11 15.64 18.55 21.88 25.68 29.98 34.84 40.29 3.18 4.44 6.12 8.36 11.31 15.14 20.10 26.45 0.97 1.32 1.78 2.38 3.15 4.14 5.40 6.97 8.94 1.82 2.33 2.97 3.74 4.69 5.83 7.20 8.83 10.77 2.09 2.84 3.82 5.07 6.69 8.74 11.32 1.44 1.87 2.40 3.06 3.88 4.87 6.07 7.52 9.25 12.08 14.30 16.84 19.71 22.97 26.64 30.74

−3.07 0.51 0.80 0.14 0.66 0.20 −0.35 −0.10 5.07 3.69 3.92 0.71 −2.45 −0.80 −1.57 1.08 11.82 2.94 0.00 −0.85 −0.96 −0.73 −0.75 −0.14 0.45 2.67 −0.87 −0.34 −0.81 1.47 −0.17 −0.98 0.22 0.09 −0.97 5.02 2.05 −1.60 −1.36 −1.04 0.88 7.69 4.10 1.64 −2.00 −2.11 0.00 −1.34 −0.27 0.75 2.19 −0.49 −1.94 −0.05 0.17 0.78 −0.33

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Table 1. continued a

Standard uncertainty of experimental temperature u is u(T) = 0.05 K, and relative standard uncertainties are ur(x1) = 0.02 and ur(p) = 0.05. The relative deviation (RD) = (x1exp − x1cal)/x1exp.

Figure 2. T−x1 curves for different solvents: △, 1-butanol; ▲, methanol; ●, ethanol; □, chloroform; ■, toluene; ★, benzene; ○, 2-propanol; (A)  , fitting results of the modified Apelblat equation; (B) ---, fitting results of the λh equation; (C) ···, fitting results of the van’t Hoff equation.

with a precision mercury thermometer inserted into the solution of the vessel. The accuracy of the thermometer was ± 0.05 K. An analytical balance (type TG328B, Shanghai Balance Instrument Works Co., Ltd., China) with an uncertainty of ± 0.0001 g was employed to weigh the solute and solvent. During the experimental measurement, predetermined excess amounts of solute and solvent of known mass were added into the dissolution vessel. After being continuously stirred for 2 h, a small amount of solid particles did not dissolve in the solution and were totally suspended in the dissolution vessel. An additional solvent of known mass was added to the solution in the vessel. This process was repeated several times until solid particles were completely dissolved. The intensity of the laser beam penetrating through the solution gradually increased along with the increasing amount of dissolving solute. When the last particle of solute disappeared, the solution was clear, and the intensity of the laser beam reached the highest value. The corresponding temperature and the solvent mass

consumed in the measurement could be obtained. Each measurement was conducted three times with reasonable consequences, and the mean value of the measurements was considered to calculate saturated mole solubility. The standard uncertainty of the measured solubility is estimated as ur(x) = 0.02. The mole fraction solubility (x1) of MDA in the studied solvents could be calculated from the following equation:15 x1 =

m1/M1 m1/M1 + m2 /M 2

(1)

where x1 is the mole fraction solubility of MDA in the solvent; m1 and m2 represent the mass of the solute and solvent, respectively; and M1 and M2 represent the molecular mass of the solute and solvent, respectively. The reliability of the experimental apparatus was verified by comparing results between experimental values and literature values of the solubility of urea in water. It is found from Table 1S 2030

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Table 2. Correlation Results of the Modified Apelblat Equation for MDA in Different Pure Solventsa solvent

A

B

C

102RAD

103RMSD

methanol ethanol 2-propanol 1-butanol chloroform toluene benzene

62.201 −221.839 −127.974 −9.4930 −20.5558 −138.481 −177.187

−5580.08 5325.96 1268.04 3383.81 −1637.82 2510.81 3529.45

−7.9588 35.2617 20.9620 2.9995 4.2308 22.1348 28.3990

0.59 1.33 1.16 0.80 0.86 0.78 1.84 ∑(102RAD) = 7.36

1.28 1.12 0.27 0.39 1.79 0.34 0.71 ∑(103RMSD) = 5.90

A, B, and C are parameters of the modified Apelblat equation. The relative average deviation (RAD) = (1/N)∑Ni=1|(x1exp − x1cal)/x1exp|. The root mean square deviation (RMSD) = [(1/N)∑Ni=1(x1exp − x1cal)2]1/2.

a

λh Equation. The Buchowski−Ksiazczak λh (eq 3) could be used to describe the relation between solubility, activity, and temperature, which was originally developed by Buchowski et al.19,20 It is shown as follows:

(Supporting Information) that the deviation of solubility data of urea in our experiment was no more than 2%. Results indicated that experimental data are in agreement with the reported data.16

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⎡ ⎛ 1 λ(1 − x1) ⎤ 1 ⎞ ln⎢1 + − ⎥ = λh⎜ ⎟ x1 Tm/K ⎠ ⎦ ⎣ ⎝ T /K

RESULTS AND DISCUSSION Solubility Data. The experimental solubility data of MDA in a variety of solvents are presented in Table 1 and Figure 2. It is shown that the solubility of MDA in the seven pure solvents increases smoothly with increasing temperature, and the solubility of MDA in alcohols increases with the increasing polarity index of alcohols. The solubility of the MDA in seven different organic solvents from (293.15 to 333.15) K is in the following order: methanol > chloroform > ethanol > benzene > 1-butanol > toluene > 2-propanol, which is not consistent with the polarity order of solvents [polarity order: methanol > ethanol > 1-butanol > 2-propanol > chloroform > benzene > toluene].17 The results indicate that the polarity of the solvents is not the only factor that determines the solubility of MDA in the studied solvents. Some similar situations were also observed in other solvent property parameters for selected solvents, such as dipole moment, δ Hildebrand solubility parameter, and dielectric constant.17 It is indicated that the solubility order of MDA in the studied solvents was the result of net contributions of these parameters. From Figure 2, ethanol is considered to be a suitable solvent for the separation and purification of MDA in industry. The reasons why ethanol is a suitable recrystallization solvent are as follows: (1) the increasing rate of solubility in selected alcohols is the largest in ethanol, and ethanol shows the most suitable dissolution property; (2) as a kind of relatively safe organic solvent, the toxicity of ethanol in selected solvents is the smallest; (3) with a lower boiling point, ethanol is more easily removed and recycled; and (4) from the point of view of the industrial cost, ethanol is much cheaper than other organic solvents.23 Correlation with Various Models. Modified Apelblat Equation. The temperature dependence of the mole fraction solubility of solute in various solvents could be well-correlated by the modified Apelblat equation derived from the Williamson equation according to the previous work.18 The modified Apelblat equation could usually give a relatively accurate correlation with three parameters and is shown as eq 2: ln(x1) = A +

B + C ln(T /K ) T /K

(3)

where x1 is the mole fraction solubility of MDA and T is absolute equilibrium temperature in K. From Figure 3, Tm

Figure 3. DSC plot of the MDA.

(absolute melting temperature of MDA in K) and molar fusion heat, which were determined by means of differential scanning calorimetric (DSC, NETZSCH, STA409PC), are 363.3 K and 18.7 kJ/mol,21,24 respectively. The λ and h are adjusted parameters that could be obtained by fitting solubility data and are presented in Table 3. van’t Hoff Equation. According to the solid−liquid phase equilibrium theory, the relationship between temperature and mole fraction solubility can be correlated by the van’t Hoff equation,22 which is denoted as ln(x1) = a +

b T /K

(4)

where x1 is the mole fraction solubility of MDA, T is absolute equilibrium temperature in K, and a and b are adjusted parameters of eq 4 and are listed in Table 4. The relative deviation (RD) between the calculated and the experimental values is defined as

(2)

where x1 represents the mole fraction solubility of MDA and T represents the corresponding temperature in K. A, B, and C are adjusted parameters in the equation, which are listed in Table 2. 2031

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Table 3. Correlation Results of the λh Equation for MDA in Different Pure Solventsa solvent

λ

h

102RAD

103RMSD

methanol ethanol 2-propanol 1-butanol chloroform toluene benzene

1.0089 1.8713 0.2154 0.1552 0.6635 0.1426 0.6047

3071.79 3325.01 21457.9 21046.7 4014.98 24277.5 8435.52

0.74 3.56 1.68 3.13 0.90 1.58 1.80 ∑(102RAD) = 13.39

1.68 3.30 0.52 1.38 1.89 0.72 0.77 ∑(103RMSD) = 10.26

λ and h are parameters of the λh equation. The relative average deviation (RAD) = (1/N)∑Ni=1|(x1exp − x1cal)/x1exp|. The root mean square deviation (RMSD) = [(1/N)∑Ni=1(x1exp − x1cal)2]1/2. a

Table 4. Correlation Results of the van’t Hoff Equation for MDA in Different Pure Solventsa solvent

a

b

102RAD

103RMSD

methanol ethanol 2-propanol 1-butanol chloroform toluene benzene

8.4979 16.4115 13.8939 10.7934 7.9482 11.2419 14.3127

−3086.86 −5821.88 −5433.26 −4338.28 −2949.61 −4538.32 −5329.04

0.73 2.41 2.07 0.85 0.86 2.21 1.85 ∑(102RAD) = 10.98

1.70 2.21 0.52 0.40 1.84 0.70 0.94 ∑(103RMSD) = 8.31

a and b are parameters of the van’t Hoff equation. The relative average deviation (RAD) = (1/N)∑Ni=1|(x1exp − x1cal)/x1exp|. The root mean square deviation (RMSD) = [(1/N)∑Ni=1(x1exp − x1cal)2]1/2.

a

RD =

x1exp − x1cal x1exp

standard molar Gibbs free energy ΔGsol0 (kJ·mol−1) of solution of MDA in different solvents could be calculated by eqs 8−11, and these equations are shown as follows:24,25

(5)

⎛ ∂ ln x1, T ⎞ ΔHsol 0 = −R ⎜ ⎟ ⎝ ∂(1/T ) ⎠ P

The relative average deviation (RAD) is employed to evaluate the error of different models and is defined as RAD =

1 N

N

x1exp − x1cal x1exp

∑ i=1

⎛ ⎞ ∂ ln x1, T = −R ⎜ ⎟ ⎝ ∂[(1/T ) − [(1/Tmean)]] ⎠ P

(6)

Furthermore, the root mean square deviation (RMSD) is used to evaluate the accuracy and predictability of the three models and is expressed as ⎡1 RMSD = ⎢ ⎢⎣ N exp

N

∑ (x1exp i=1

⎤1/2 cal 2 ⎥ − x1 ) ⎥⎦

Tmean =

N N 1 Ti

∑i

ΔGsol 0 = −RTmean × intercept (7)

ΔSsol 0 =

cal

where x1 is the experimental mole solubility values, x1 is the calculated mole solubility values, and N is the number of experimental points. The values of the RD, RAD, and RMSD are listed in Tables 2−4, respectively. As we can see from Tables 2−4, the calculated solubility data are in good agreement with the experimental data due to small RSMDs and RADs. It could be observed that the total RAD obtained by the modified Apelblat equation, λh equation, and van’t Hoff equation are 7.36, 13.39, 10.98, respectively. The best fitting result of the relationship between equilibrium solubility and temperature is given by the modified Apelblat equation. Therefore, we could conclude that the modified Apelblat equation is more suitable for solvent selection and model research in the process of crystallization of MDA. Thermodynamic Properties of the Solution. Thermodynamic properties of the dissolution process for the MDA in different solvents were investigated. On the basis of van’t Hoff analysis,23 the standard molar dissolution enthalpy ΔHsol0 (kJ· mol−1), standard molar entropy ΔSsol0 (J·mol−1 K−1), and

(8)

ΔHsol 0 − ΔGsol 0 Tmean

(9) (10)

(11)

where x1 is the mole fraction solubility of MDA, Tmean is the mean harmonic temperature, T is absolute equilibrium temperature in K, R is universal gas constant and its value is 8.314 J·K−1 mol−1. The ΔHsol0, ΔGsol0, and ΔSsol0, respectively, represent the standard molar enthalpy, Gibbs free energy, and entropy of dissolving process.25 The standard molar enthalpy (ΔHsol0) can be obtained from the slope, where ln(x1) is plotted with 1/T. In eq 10, the intercept can be acquired by plotting ln(x1) as a function of (1/T − 1/Tmean) and is shown in Figure 4. The calculated dissolution standard molar enthalpy ΔHsol0 (kJ·mol−1), standard molar entropy ΔSsol (J·mol−1 K−1), and standard molar Gibbs free energy ΔGsol (kJ·mol−1) are presented in Table 5. The comparisons between the relative contributions of the standard molar enthalpy and standard molar entropy to the standard molar Gibbs free energy25 by eqs 12 and 13 are listed in Table 5. 2032

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tally determined by a synthetic method at temperatures ranging from 293.15 to 333.15 K, and the following conclusions were obtained: (1) the solubility of MDA in all selected solvents increases with increasing temperature, and the solubility of MDA increases in seven different solvents as below: 2-propanol < toluene < 1-butanol < benzene < ethanol < chloroform < methanol; (2) the mole fraction solubility of MDA in methanol is far greater than that of the other six solvents in this work, and the increments with temperature in ethanol are the greatest; (3) the experimental solubility data were well correlated by the modified Apelblat, λh, and van’t Hoff equations; the modified Apelblat equation gives the best description of the relationship between equilibrium solubility and temperature; (4) the dissolution process of MDA in all studied solvents is endothermic and entropy-driven. Moreover, the major contributor to the standard molar Gibbs free energy is enthalpy, which is due to all values of %ξH > 53 in the dissolution process of the MDA in the selected solvents.

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Figure 4. A van’t Hoff diagram of the mole fraction solubility (ln x1) of the MDA versus (1/T − 1/Tmean) in the studied solvents: △, 1butanol; ▲, methanol; ●, ethanol; □, chloroform; ■, toluene; ★, benzene; ○, 2-propanol.

%ξH =

%ξS =

|ΔHsol 0| |ΔHsol 0| + |TmeanΔSsol 0| |TmeanΔSsol 0| |ΔHsol 0| + |TmeanΔSsol 0|

ASSOCIATED CONTENT

S Supporting Information *

Comparisons of the literature and experimental solubility of urea as a function of temperature in water. The Supporting Information is available free of charge on the ACS Publications website at DOI: 10.1021/je501111d.

× 100

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(12)

× 100

AUTHOR INFORMATION

Corresponding Authors

*Tel.: +86 0371 67781713. Fax: +86 0371 67781713. E-mail: [email protected] (L.X.). *Tel.: +86 0371 67781713. Fax: +86 0371 67781713. E-mail: [email protected] (G.L.).

(13)

where Tmean represents the mean harmonic temperature, and its meaning is the same as that of eq 9. From Table 5, the standard molar enthalpy and entropy of the dissolving process for all solvents are positive. It is demonstrated that the dissolution process of MDA in seven pure organic solvents is entropydriven and endothermic. The dissolution process of MDA is endothermic because the intermolecular interaction force between the solute (MDA) molecules and solvent molecules is greater than that between the solvent molecules. The positive values of standard molar Gibbs free energy (ΔGsol0) in dissolution process indicate that the process was not spontaneous. When the calculated standard molar Gibbs energy is compared in all selected solvents, it is shown that the lower the value of ΔGsol0, the larger the solubility of MDA. In addition, the values of %ξH are far greater than those of %ξS, which indicates that the major contributing factor to the standard molar Gibbs free energy (ΔGsol0) is the enthalpy in the dissolution process of the MDA in the studied solvents.

Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS We appreciate the editors and the anonymous reviewers for their valuable suggestions. REFERENCES

(1) Liu, Y. F.; Liao, C. Y.; Hao, Z. Z. The polymerization behavior and thermal properties of benzoxazine based on o-allylphenol and 4,4′diaminodiphenyl methane. React. Funct. Polym. 2014, 75, 9−15. (2) Asundaria, S. T.; Patel, K. C. Synthesis, characterization, and antimicrobial studies of bissydnones based on 4,4′-diaminodiphenyl methane. Synth. Commun. 2010, 40, 1899−1906. (3) Men, W. W.; Lu, Z. J. Synthesis and characterization of 4,4′diaminodiphenyl methane based benzoxazines and their polymers. J. Appl. Polym. Sci. 2007, 106, 2769−2774. (4) Adel, N.; Chokri, J.; Arbi, M.; Manef, A. Solubility of gallic acid in liquid mixtures of (ethanol + water) from (293.15 to 318.15) K. J. Chem. Thermodyn. 2012, 55, 75−78.

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CONCLUSIONS The solubility data of MDA in methanol, ethanol, 2-propanol, 1-butanol, benzene, toluene, and chloroform were experimen-

Table 5. Thermodynamic Function Values of the Dissolving Process for MDA in Different Solvents at Mean Temperatures solvent

ΔHsol0 (kJ·mol−1)

ΔSsol0 (J·mol−1 K−1)

ΔGsol0 (kJ·mol−1)

%ξH

%ξS

methanol ethanol 2-propanol 1-butanol chloroform toluene benzene

25.96 46.99 43.44 35.88 24.34 36.45 43.76

71.59 131.99 110.09 89.14 65.50 89.46 117.25

3.75 6.04 9.02 8.01 4.18 8.49 7.66

53.89 53.44 55.79 56.28 54.70 56.59 54.80

46.11 46.56 44.21 43.72 45.30 43.41 45.20

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2034

DOI: 10.1021/je501111d J. Chem. Eng. Data 2015, 60, 2028−2034