Experimental and Modeling Studies on the Solubility of d

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Experimental and Modeling Studies on the Solubility of D‑Pantolactone in Four Pure Solvents and Ethanol−Water Mixtures Cui Huang,†,‡ Zhiping Xie,†,‡ Jinchao Xu,†,‡ Yujia Qin,†,‡ Yanni Du,†,‡ Shichao Du,†,‡ and Junbo Gong*,†,‡ †

School of Chemical Engineering and Technology, State Key Laboratory of Chemical Engineering, and ‡The Co-Innovation Center of Chemistry and Chemical Engineering of Tianjin, Tianjin University, Tianjin 300072, People’s Republic of China S Supporting Information *

ABSTRACT: The solubility of D-pantolactone in water, ethanol, methanol, ethyl acetate, and ethanol−water mixtures was determined at temperatures between (278.15 and 318.15) K using a digital densitometer by a static method. The measured solubility data were correlated with Apelblat equation, van’t Hoff equation, nonrandom two liquid model, and Wilson model. The results indicate that the Wilson model is the most suitable model in pure solvents and the Apelblat model is the best model in ethanol−water mixtures.

1. INTRODUCTION D-Pantolactone (dihydro-3-hydroxy-4H-dimethyl-2(3H)-furanone, Figure 1) (CAS Registration No. 599-04-2) is an

difficulty in completely evaporating the solvent from the solution. Laser monitoring was not selected because it is quite time-consuming. The solubility of D -pantolactone was determined by densitometry. For a certain solution, density is a function of solute concentration and temperature. Garside and Mullin proposed the use of an online densitometer to determine solute concentration.8 Qiu and Rasmuson reported successful use of density measurement in measuring concentration of samples extracted from a crystallizer.9 S Rohani, Rawlings, and Miller demonstrated the use of densitometry for accurate online concentration measurements.10−12 In this paper, the D-pantolactone solubility in water, ethanol, methanol, ethyl acetate, and ethanol−water mixtures at temperatures ranging from (278.15 to 318.15) K was measured using a digital densitometer by a balance method. Water is the commonly used solvent, whereas ethyl acetate is the nonignorable impurity during the crystallization procedure. Methanol and ethanol are the most likely solvents to be used during the crystallization process. Because D-pantolactone would convert into plastic crystal at temperatures higher than 335.15 K,7 which is an invalid solid form in the production of vitamins, a useful temperature range from (278.15 to 318.15) K was chosen in the experiment. The experimental data were correlated with the van’t Hoff equation, the modified Apelblat equation, the nonrandom two liquid (NRTL) model, and the Wilson model.

Figure 1. Molecular structure of D-pantolactone.

important drug intermediate in the synthesis of Vitamin B5.1,2 With the expanding vitamin market, demand for D-pantolactone is growing rapidly. At present, the majority of the studies focus on the preparation of D-pantolactone.3−6 Few attention is given to the purification process. Crystallization as the essential step to produce highly purified D-pantolactone must be given special attention. Solid−liquid equilibrium data are essential for the design of a crystallization process. These data are valuable in theoretical study, such as the selection of thermodynamic model, the calculation of yield, and annual output. Up to now, no solubility data of D-pantolactone have been reported in the literature. For this purpose, detailed solubility data of D-pantolactone in different solvent systems are required. There are several approaches for the measurement of solubility. In most cases, it is preferable to determine saturation concentration of the solute using a gravimetric method due to excellent reproducibility. This method is inadvisible for accurate measurement of D -pantolactone solubility because this compoud converts to plastic crystal at temperatures higher than 335.15 K.7 This means when D-pantolactone is heated to 335.15 K, the orientation of the crystal changes and it converts into plastic crystal, which is defined as position ordered and orientation disordered. The thermal instability leads to great © XXXX American Chemical Society

2. EXPERIMENTAL SECTION 2.1. Materials. D-Pantolactone, supplied by Ethyl Chemical Co., Ltd. (Shanghai, China), was purified by recrystallization twice from aqueous solution. The mass fraction of the products, Received: October 17, 2014 Accepted: January 20, 2015

A

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3. THERMODYNAMIC MODELS To quantificationally describe solid−liquid equilibrium, the connection between the solubility of D-pantolactone and temperature was modeled with the following thermodynamic equations. 3.1. van’t Hoff Equation. For real solutions, the logarithm of mole fraction of a solute is linearly related to the reciprocal of the absolute temperature, which can be described as the van’t Hoff equation:15

determined by HPLC (Agilent 1200, American), was higher than 0.996. Methanol, ethanol, and ethyl acetate (analytical grade) were purchased from Jiangtian Chemical Co., Ltd. (Tianjin, China), and their mass fractions were higher than 0.995. All chemicals were used directly without further purification. Deionized water was prepared in our laboratory by a Thermo Scientific Barnstead Pacific TII (Wuzhou Technology, Beijing). 2.2. Experimental Methods. First, a digital densitometer (DMA4500, Anton Paar GmbH) based on vibration tube technology was used to measure the density at 318.15 K for different mole fractions of D-Pantolactone. The thermostatic device inside the instrument can control the sample temperature at a range of (278.15 to 353.15) K in order to eliminate the influence of the temperature fluctuations during the measurement process. The digital densitometer are equipped with three temperature sensors, which can measure environment temperature, heater temperature, and sample cell temperature accurately. The microprocessor system inside the instrument can calculate the real temperature of the sample according to a special formula. The digital densitometer allows setting customized measurement times for the same sample in order to ensure the accuracy of the results. In this experiment, the density of each sample was measured twice. The vibration tube method offers extremely high accuracy, and the precision of the digital densitometer is 0.0001 g/cm3. The experimental data were linear fitted to obtain the correlation between density and the mole fraction of solute. The solubility measurement was performed in a 50 mL conical flask. For each measurement, a certain volume of solvent and an excessive amount of D-pantolactone crystal were added to the conical flask, respectively. The experiment was operated at a constant stirring rate of 150 rpm for 12 h to ensure that the solubility equilibrium was always maintained during the entire measurement process. After settling down for 5 h, 1 mL of supernatant was extracted from the conical flask and filtered with a 0.22 μm PTFE filter. Then the density was determined by the digital densitometer. Some preliminary experiments were conducted to verify that the metastable zone of D-pantolactone is wide and a short sampling process cannot cause precipitation of new crystal. Besides, all the density measurements were operated at 318.15 K to ensure that crystallization and plugging did not occur in the measuring tube. The procedure was repeated three runs for each pair of temperature and solvent composition. Through the average density value, the mole fraction solubility of D-pantolactone in a certain solvent can be obtained from the standard curve of density versus concentration at 318.15 K. All experiments were performed using high-throughput thermostatic cultivation oscillators (HNY-200B, China). The equipment allows eight parallel experiments at the same temperature. Validation of the equipment showed that the temperature of the liquid in the flasket was always 1 K lower than the given temperature. So all the temperatures were revised by a thermometer with an uncertainty of ± 0.1 K. To evaluate the accuracy of the experimental results, a laser monitoring method was also used to measure solubility of Dpantolactone in water. The measurement apparatus of the laser technique is shown in the Supporting Information (Figure S1), and the detailed measuring procedures are the same as mentioned in the literature.13,14

ln x = −

ΔHd ΔSd + RT R

(1)

where x is mole fraction solubility, T is the absolute temperature, R is the gas constant, ΔHd and ΔSd account for dissolution enthalpy and dissolution entropy, respectively. Besides, ΔHd and ΔSd are assumed to be independent of temperature within a certain temperature range. 3.2. Modified Apelblat Equation. The modified Apelblat equation16 is a semiempirical model, which can be expressed as follows: ln x = A +

B + C ln T T

(2)

where A, B, and C are empirical constants. The values of A and B reflect the effect of solution nonidealities on the solubility of solute and the value of C reflects the influence of temperature on the fusion enthalpy.17 3.3. NRTL Model. In the binary system, the activity coefficient can be calculated by the following formula:18 ⎡ τ G 2 τ12G12 2 ⎤ 21 21 ⎥ ln γ1 = x 2 2⎢ + (x 2 + G12x1)2 ⎦ ⎣ (x1 + G21x 2)2

(3)

Here G12 = exp( −α12τ12)

(4)

G21 = exp( −α12τ21)

(5)

τ12 =

τ21 =

(g12 − g22) (RT )

(6)

(g21 − g11) (RT )

(7)

Δg12 (= g12 − g22) and Δg21 (= g21 − g11) are cross interaction energy parameters, independent of temperature and composition. In addition, α12 is a constant that reflects the nonrandomness of the mixture and its value generally varies between 0.20 and 0.47.19 Different values of α12 were chosen to correlate the solubility data of D-pantolactone. It turns out that α12 = 0.40 is the most suitable value because of the smallest relative deviation for the measurement system. 3.4. Wilson Model. The Wilson model20 is another activity coefficient model, which can be expressed as follows: ln γ1 = −ln(x1 + ∧12 x 2) ⎞ ⎛∧ ∧21 + x 2⎜ 12 + ∧12 x 2 − ⎟ x 2 + ∧21 x1 ⎠ ⎝ x1

(8)

Here, B

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∧12 =

⎛ λ − λ11 ⎞ V2 ⎛ Δλ ⎞ V2 ⎟= exp⎜ − 12 exp⎜ − 12 ⎟ ⎝ ⎝ RT ⎠ V1 RT ⎠ V1

∧21 =

⎛ λ − λ 22 ⎞ ⎛ Δλ ⎞ V1 V ⎟ = 1 exp⎜ − 21 ⎟ exp⎜ − 21 ⎝ ⎠ ⎝ RT ⎠ V2 RT V2

Table 2. Experimental Mole Fraction Solubility of DPantolactone in Water by Densitometry and Laser Monitoring at Temperature T and Pressure p = 0.1MPaa

(9)

(10)

where Δλ12 and Δλ21 are the cross interaction energy parameters, and V1 and V2 are the mole volume of solute and solvent, respectively.

4. RESULTS AND DISCUSSION 4.1. Thermodynamic Properties. The melting temperature and enthalpy of fusion of D-pantolactone were determined by differential scanning calorimetry (DSC, Mettler-Toledo, model DSC 1/500, Switzerland) under a nitrogen atmosphere. The DSC curve is exhibited in Figure 2. It is obvious that D-

T/K

102xiden

102xilaser

102RD

288.15 293.15 298.15 303.15 308.15 313.15 318.15

9.30 14.13 23.81 31.37 39.28 46.29 55.28

8.80 13.24 22.73 31.72 40.65 48.94 57.84

5.68 6.72 4.75 1.10 3.37 5.41 4.43

a den xi and xilaser represent the mole fraction solubility measured by densitometry and laser monitoring, respectively. RD reveals the relative deviation of the two methods. The standard uncertainties u are u(T) = 0.1 K, ur(p) = 0.05, ur(x) = 0.03.

good agreement with that by laser monitoring. The mole fraction solubility data of D-pantolactone in pure solvents are displayed in Table 3. The Apelblat equation, the van’t Hoff Table 3. Experimental Mole Fraction Solubility of DPantolactone in Pure Solvents at Temperature T and Pressure p = 0.1MPaa 102RD

Figure 2. DSC curve of D-pantolactone at a heating rate of 10 °C· min−1.

pantolactone has an endothermic peak at 343.46 K and 365.70 K, respectively. As reported in the literature, the endothermic peak at 365.70 K is caused by melting process and the endothermic peak at 343.46 K is due to the converting from the stable form into plastic crystal.7 The melting temperature and enthalpy of fusion of D-pantolactone are listed in Table 1. Table 1. Thermodynamic Properties of D-Pantolactone (0.1MPa)a values

T1/K

ΔH1/kJ·mol−1

T2/K

ΔH2/kJ·mol−1

343.46

14.92

365.70

2.93

a

The standard uncertainties u are u(T) = 0.5 K; relative standard uncertainties u are ur(ΔH) = 0.05 (0.95 level of confidence).

4.2. Standard Curves. The calibration data of density versus D-pantolactone mole fraction are presented in the Supporting Information. Figure S2 and S3 show the standard curve of density versus mole fraction of D-pantolactone in pure solvents and in the mixed (ethanol−water) solvents. As observed in the two diagrams, the density and mole fraction of solute indicate a good linear relationship and the results illustrate that the determination of solubility based on densitometry is dependable. 4.3. Solubility in Pure Solvents. The solubility of Dpantolactone in pure water, ethanol, methanol, and ethyl acetate were measured at temperatures between (288.15 and 318.15) K. Table 2 shows the mole fraction solubility data of Dpantolactone in water by densitometry and laser monitoring. The maximum relative deviation of the two methods is 6.72 %, which proves that the solubility measured by densitometry is in

T/K

102xi

288.15 293.15 298.15 303.15 308.15 313.15 318.15

26.29 30.22 35.53 42.16 47.08 49.58 54.34

288.15 293.15 298.15 303.15 308.15 313.15 318.15

40.42 46.27 50.62 55.25 61.10 66.70 69.40

288.15 293.15 298.15 303.15 308.15 313.15 318.15

29.03 34.52 38.54 44.60 50.12 60.71 66.50

Apelblat

van’t Hoff

Methanol Solvent −1.70 1.37 2.33 1.19 1.66 −0.02 −2.18 −2.17 −2.10 −2.22 1.85 0.56 −0.09 1.50 Ethanol Solvent 0.35 1.27 −1.21 −0.36 0.51 −0.22 1.32 −0.10 −0.30 −0.93 −1.62 −1.23 0.88 1.64 Ethyl Acetate Solvent −1.30 0.63 0.17 −1.12 0.22 0.31 −1.46 0.37 1.80 1.68 −0.69 −1.31 0.54 1.02

Wilson

NRTL

0.48 0.15 −0.32 −0.78 −0.58 0.24 0.81

0.54 0.16 −0.37 −0.86 −0.61 0.27 0.94

0.42 −0.25 −0.23 −0.12 −0.23 −0.19 0.67

0.50 −0.29 −0.29 −0.19 −0.25 −0.12 0.78

0.32 −0.14 0.05 0.01 0.28 −0.22 0.10

0.13 −0.05 0.04 −0.09 0.19 −0.10 −0.15

a

The standard uncertainties u are u(T) = 0.1 K, ur(p) = 0.05, ur(x) = 0.03.

equation, the NRTL model, and the Wilson model were used to correlate the experimental solubility data. Table 4 shows the fitting parameters of different models. As shown in Tables 2 and 3, D-pantolactone solubility increases significantly with the increase of temperature in all of the investigated solvents. At constant temperature, the solubility order meets the following rule: ethanol > ethyl acetate > methanol > water. The dissolving behavior of the solute depends its chemical structure and the polarity of the C

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Table 4. Parameters of the Apelblat, van’t Hoff, NRTL, and the Wilson Model for the Solubility of D-Pantolactone in Pure Solvents model

parameter 4

A/10 B/104 C/104 Δg12/103·J·mol−1 Δg21/103·J·mol−1 Δλ12/103·J·mol−1 Δλ21/103·J·mol−1 ΔH/104·J·mol−1 ΔS/104·J·mol−1·K−1

Apelblat

NRTL Wilson van’t Hoff

water

methanol

ethanol

ethyl acetate

0.0708 −3.4976 −0.0104 3.8138 3.5694 −0.2826 9.2168 2.6996 0.0079

0.0433 −2.1513 −0.0064 4.1405 2.1753 0.2464 7.0926 1.7850 0.0051

0.0201 −1.0493 −0.0029 3.9181 1.0423 −0.0218 5.5414 1.3539 0.0040

−0.0576 2.3795 0.0087 3.7899 2.1295 2.5437 4.3160 2.1972 0.0066

solvent. The reason for the greater solubility of D-pantolactone in alcohols than in water lies in the polarity difference of solvent molecules. The influence of solvent on dissolving capacity can be explained by the empirical rule “like dissolves like”. If the solute molecules and the solvent molecules have similar polarity, solute molecules need to overcome the smaller energy barrier and can combine with solvent molecules more easily. The experimental results agree well with this principle. DPantolactone molecule is an annular structure (Figure 1) and its polarity is small; therefore, D-pantolactone has a relatively high solubility in ethanol and ethyl acetate. When two kinds of solvents have equal polarity, their chemical structure is a key determinant of dissolving capacity. Both D-pantolactone and ethanol contain a hydroxyl group, so solute molecules and solvent molecules are liable to combine in the form of hydrogen bonds, and this contributes to a higher solubility in ethanol. To evaluate the correlation results and select the most suitable model for D-pantolactone solubility in pure solvents, the relative deviation (RD %) and the average relative deviation (ARD %) were calculated. The relative deviation and the average relative deviation are defined as RD % =

102RD

100 N

T/K

10 xi

Apelblat

van’t Hoff

NRTL

0.2

278.15 283.15 288.15 293.15 298.15 303.15 308.15 313.15 278.15 283.15 288.15 293.15 298.15 303.15 308.15 313.15 278.15 283.15 288.15 293.15 298.15 303.15 308.15 313.15 278.15 283.15 288.15 293.15 298.15 303.15 308.15 313.15

11.51 15.71 20.06 25.58 31.40 36.52 42.06 46.48 18.22 22.24 25.85 30.63 35.98 41.87 44.74 52.37 22.51 25.32 31.00 35.15 39.66 45.95 49.11 55.74 25.44 30.05 34.21 39.51 44.06 51.03 56.96 61.16

−0.02 −0.73 1.37 −0.01 −0.94 0.35 −0.35 0.29 0.11 −0.96 1.41 0.47 −0.88 −2.48 3.30 −1.18 −1.53 3.27 −1.75 −0.25 0.63 −2.16 2.11 −0.59 0.08 −0.85 0.79 −0.05 1.65 −1.39 −1.59 1.26

2.33 1.17 0.38 −0.98 −1.98 −1.54 −0.70 2.16 0.88 0.17 0.30 −0.27 −0.91 −1.56 1.39 0.20 0.53 1.20 −0.62 −0.51 −0.33 −1.49 0.96 0.42 0.82 0.08 0.19 −0.42 0.18 −1.17 −0.98 1.46

1.08 0.21 −0.35 −0.80 −0.86 −0.46 0.19 1.14 0.72 0.08 −0.16 −0.48 −0.61 −0.49 0.32 0.73 0.53 0.41 −0.44 −0.52 −0.45 −0.46 0.32 0.73 0.59 −0.01 −0.21 −0.46 −0.26 −0.33 0.03 0.80

0.6

x1,exp i

(11) N

∑ i=1

cal x1,exp i − x1, i

x1,exp i

(12)

xcal 1,i

where and represent the experimental and calculated solubility value, respectively. N account for the amount of experimental dots. Table 5 lists the ARD % of different

0.8

Table 5. ARD% of Different Models in Pure Solvents solvent

water

methanol

ethanol

ethyl acetate

Apelblat NRTL Wilson van’t Hoff

1.5774 2.0444 1.7250 4.7735

1.7013 1.3229 1.1805 3.2775

0.8823 0.6380 0.5548 1.4695

0.8473 0.1992 0.2944 2.0754

2

we

0.4

cal x1,exp i − x1, i

ARD % =

xexp 1,i

Table 6. Experimental Mole Fraction Solubility of DPantolactone in Ethanol−Water Mixtures at Temperature T and Pressure p = 0.1MPaa

a

we is the mass fraction of ethanol in the absence of D-Pantolactone. The standard uncertainties u are u(T) = 0.1 K, ur(p) = 0.05, ur(x) = 0.03.

correlation models. The average relative deviation of the four models are 1.25 % (Apelblat), 1.05 % (NRTL), 0.94 % (Wilson), and 2.90 % (van’t Hoff). Therefore, the Wilson model fits well with the experimental solubility data of Dpantolactone in pure solvents. 4.4. Solubility in Binary Solvent Mixtures. The solubilities of D-pantolactone in ethanol−water mixtures at different temperatures are listed in Table 6. It can be observed

that the solubility in ethanol−water mixtures increases with an increase in temperature and mass fraction of ethanol. The temperature dependence of D-Pantolactone solubility can be described by the modified Apelblat equation (eq 2). The experimental data were fitted with van’t Hoff equation, Apelblat equation, and NRTL model. The parameter values together with the deviation are presented in Tables 7 and 8, respectively. D

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Table 7. Parameters of the Apelblat, the van’t Hoff and the NRTL Model for the Solubility of D-Pantolactone in Ethanol−Water Mixtures model

parameter

we = 0.2

we = 0.4

we = 0.6

we = 0.8

4

0.0620 −3.0360 −0.0091 3.4083 3.6315 2.6014 0.0077

0.0183 −1.0345 −0.0026 3.6421 2.7686 2.0949 0.0062

0.1571 −8.9218 −0.0225 3.6973 2.3192 1.8440 0.0054

0,.1297 −7.6409 −0.0184 3.5286 2.1073 1.7977 0.0054

A/10 B/104 C/104 Δg12/103·J·mol−1 Δg21/103·J·mol−1 ΔH/104·J·mol−1 ΔS/104·J·mol−1·K−1

Apelblat

NRTL van’t Hoff

Table 8. ARD% of Different Models in Ethanol−Water Mixtures Apelblat NRTL van’t Hoff

we = 0.2

we = 0.4

we = 0.6

we = 0.8

0.5071 2.8117 5.6883

1.3484 1.4375 2.1264

1.5356 1.3700 2.0675

0.9566 0.8436 1.1896

The average relative deviation of the three models are 1.70 % (Apelblat), 1.72 % (NRTL), 3.22 % (van’t Hoff). The results indicate that the modified Apelblat equation is the most suitable model for ethanol−water mixtures. On the basis of the modified Apelblat equation, the solubility of D-pantolactone at different temperatures were calculated in mole fraction and plotted versus mass fraction of ethanol. Figure 3 shows good correlation between the experimental data and the fitted solubility curve.

Figure 4. Modified van’t Hoff plots of ln xi versus 1000T‑1 in different pure solvents: ■, water; ●, ethanol; ⧫, ethyl acetate; ▲, methanol.

Table 9. Predicted Values of the Gibbs Free Energy Change water T/K 288.15 293.15 298.15 303.15 308.15 313.15 318.15

ΔGd/KJ·mol 4.23 3.84 3.44 3.05 2.65 2.26 1.86

ΔGd/KJ·mol 3.15 2.90 2.64 2.39 2.13 1.88 1.62

ethanol −1

ethyl acetate −1

ΔGd/KJ·mol 2.01 1.81 1.61 1.41 1.21 1.01 0.81

ΔGd/KJ·mol−1 2.95 2.62 2.29 1.96 1.63 1.30 0.97

ature. The decrease of ΔGd reflects that dissolving process is spontaneous, which coincides with the change law of solubility. Table 4 demonstrates that the ΔHd values of D-pantolactone are positive in the four pure solvents, which illustrates the dissolving process is endothermic and temperature increasing is beneficial to the increase of solubility. Furthermore, the ΔHd values and ΔGd values in water are far higher than that in other solvents. This is because of the difference in energy barrier during dissolution and the lower ΔHd corresponds to larger solubility.

Figure 3. Modified Apelblat plots of ln xi versus walcohol in ethanol− water mixtures: ●, T = 278.15 K; ▶, T = 283.15 K; ▲, T = 288.15 K; ★, T = 293.15 K; ⧫, T = 298.15 K; ▼, T = 303.15 K; ■, T = 308.15 K; ◀, T = 313.15 K.

4.5. Predication of Gibbs Free Energy Change. Figure 4 reveals the logarithm of mole fraction solubility of Dpantolactone versus the reciprocal of absolute temperature in different pure solvents. On the basis of the van’t Hoff equation, the values of dissolution enthalpy and dissolution entropy were calculated from the slope and the intercept of the curve. As shown in Figure 4, the van’t Hoff plot of the logarithm of mole fraction solubility is linear to the reciprocal of the absolute temperature. The Gibbs free energy change during the dissolving process could be expressed by20 ΔGd = ΔHd − T ΔSd

methanol −1

5. CONCLUSION A digital densitometer was successfully applied to measure the solubility of D-pantolactone in four pure solvents and ethanol− water mixtures at temperatures ranging from (278.15 to 318.15) K. Validation of densitometry indicates that the method can provide a powerful tool for off-line concentration measurement. The solubility is positively correlated with temperature according to the experimental data for a given solvent. The solubility values of D-pantolactone in pure solvents satisfy the following order: ethanol > ethyl acetate > methanol > water. This phenomenon is mainly attributed to the difference in the polarity and chemical structure of the solvent. The polarity of ethanol molecules is smaller and it is liable to

(13)

The ΔGd values of D-pantolactone in pure solvents are displayed in Table 9. It is obvious that the Gibbs free energy values are all positive for each temperature and solvent type. In addition, the ΔGd values decrease with the growing temperE

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(12) Miller, S. M.; Rawlings, J. B. Model identification and control strategies for batch cooling crystallizers. AlChE J. 1994, 40, 1312− 1327. (13) Domanska, U.; Kozlowska, M.; Rogalski, M. Solubility of imidazoles in alcohols. J. Chem. Eng. Data 2002, 47, 8−16. (14) Hou, G. Y.; Yin, Q. X.; Yang, Y.; Hu, Y.; Hu, Y.; Zhang, M. J.; Wang, J. K. Solubilities of Adefovir Dipivoxil in Different Binary Solvents at 298.15 K. J. Chem. Eng. Data 2008, 53, 1021−1023. (15) Choi, W. S.; Kim, K. J. Solubility of Forms I and II of clopidogrel hydrogen sulfate in methanol and 2-propanol mixture. J. Chem. Eng. Data 2010, 56, 43−47. (16) Yang, Z. H.; Zeng, Z. X.; Sun, L.; Xue, W. L. Determination and correlation of solubilities of lauric acid in eight alcohols. J. Chem. Eng. Data 2014, 59, 2725−2731. (17) Hojjati, H.; Rohani, S. Measurement and prediction of solubility of paracetamol in water-isopropanol solution. Part 1. Measurement and data analysis. Org. Process Res. Dev. 2006, 10, 1101−1109. (18) Domanska, U.; Marciniak, A. Solubility of 1-alkyl-3-methylimidazolium hexafluorophosphate in hydrocarbons. J. Chem. Eng. Data 2003, 48, 451−456. (19) Wei, D.; Pei, Y. Measurement and correlation of solubility of diphenyl carbonate in alkanols. Ind. Eng. Chem. Res. 2008, 47, 8953− 8956. (20) Wang, S.; Zhang, Y.; Wang, J. Solubility measurement and modeling for betaine in different pure solvents. J. Chem. Eng. Data 2014, 59, 2511−2516.

form hydrogen bonds with D-pantolactone molecules. In ethanol−water mixtures, the solubility increases with increasing temperature and ethanol mass fraction. Different thermodynamic models were selected to correlate the experimental solubility data. The Wilson model and modified Apelblat equation are the most suitable models in pure solvents and ethanol−water mixtures, respectively. These data would be of great importance in the optimization and design of a crystallization process.

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ASSOCIATED CONTENT

S Supporting Information *

Chemical properties (mass purity, molar mass, and molar volume ) of pure materials, experimental solubility, calculated parameters for different models in different pure solvents and ethanol-water mixtures, and predicted values of Gibbs free energy change in pure solvents. This material is available free of charge via the Internet at http://pubs.acs.org.

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AUTHOR INFORMATION

Corresponding Author

*Tel.: 86-22-27405754. Fax: +86-22-27374971. E-mail: junbo_ [email protected]. Funding

We are grateful for the financial support of the National Natural Science Foundation of China (No. NNSFC 21176173), and the National High Technology Research and Development Program (863 Program No.2012AA021202). Notes

The authors declare no competing financial interest.

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REFERENCES

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DOI: 10.1021/je500996n J. Chem. Eng. Data XXXX, XXX, XXX−XXX