Solubility of Fenofibrate in Different Binary Solvents: Experimental

Jul 29, 2016 - Solubility of fenofibrate in ethanol–water and acetone–water mixtures was investigated using dynamic laser monitoring at atmospheri...
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Solubility of Fenofibrate in Different Binary Solvents: Experimental Data and Results of Thermodynamic Modeling Hua Sun,†,‡ Baoshu Liu,*,†,‡ Peihua Liu,† Junli Zhang,§ and Yongli Wang∥ †

College of Chemical and Pharmaceutical Engineering, Hebei University of Science and Technology, Shijiazhuang, Hebei 050018, People’s Republic of China ‡ Hebei Research Center of Pharmaceutical and Chemical Engineering, Shijiazhuang, Hebei 050018, People’s Republic of China § NCPC Hebei Huamin Pharmaceutical Co., Ltd., Hebei 052160, People’s Republic of China ∥ School of Chemical Engineering, Tianjin University, Tianjin 200273, People’s Republic of China

ABSTRACT: Solubility of fenofibrate in ethanol−water and acetone−water mixtures was investigated using dynamic laser monitoring at atmospheric pressure and temperatures between 278.15 and 308.15 K. Solubility increased as temperature and the fraction of organics increased. The solubility data was correlated with the modified Apelblat, NRTL, combined nearly ideal binary solvent/Redlich−Kister (CNIBS/R−K), and Jouyban−Acree (J−A) models. Solvent composition and temperature correlations were conducted, and the apparent dissolution enthalpy, entropy, and Gibbs free energy change of the fenofibrate were calculated using the modified Apelblat model and the van’t Hoff equation. The results shoLw that the dissolution of fenofibrate is endothermic and nonspontaneous, and that the main contributor in the binary solvents studied is enthalpy.

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INTRODUCTION Fenofibrate (CAS No.: 49562-28-9), a lipid lowering agent, is mainly used to reduce cholesterol levels in patients.1 It can be used alone or in combination with statins in the treatment of hypercholesterolemia and hypertriglyceridemia. However, fenofibrate has poor water solubility and low bioavailability.2−5 In industrial processes, the final step in the purification of fenofibrate from binary solvents after synthesis is crystallization. This step directly affects the quality of the final product, and the thermodynamic properties (especially the solubility of the solute) are therefore the key to optimal purification and crystallization. At present, the most common binary solvents for fenofibrate crystallization are ethanol−water and acetone−water. The solid solubilities of three antilipemic agents of clofibric acid, fenofibrate, and gemfibrozil in supercritical carbon dioxide were reported by Chen et al.6 in 2010, and the mobility of amorphous fenofibrate via the differential scanning calorimetry (DSC) and thermally stimulated depolarization currents (TSDC) were reported by Diogo et al.7 in 2016. The solubility of fenofibrate in pure ethanol and in acetone has been reported,8 but no data on solubility in binary ethanol−water or water−acetone are available in the literature. The study of fenofibrate solubility and © XXXX American Chemical Society

the thermodynamic properties of these binary systems have therefore been receiving considerable attention. In this research, solubility was measured using a dynamic laser method. The experimental data on the solubility of fenofibrate in pure ethanol and pure acetone are very similar to data measured by the gravimetric method.8 This suggested that the dynamic laser method was a reliable way to measure the solubility of fenofibrate. The data obtained were correlated using four different models: the modified Apelblat, NRTL, combined nearly ideal binary solvent/Redlich−Kister (CNIBS/ R−K), and Jouyban−Acree (J−A). The thermodynamic properties were calculated by regression of the solubility data using the modified Apelblat model and the van’t Hoff equation.

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EXPERIMENTAL SECTION Materials and Reagents. Fenofibrate is a white crystalline powder; its molecular structure is shown in Figure 1. In this research, its purity was determined using HPLC.9 Differential

Received: March 28, 2016 Accepted: July 19, 2016

A

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uncertainty of ±0.0001 g, were placed in a 100 mL jacketed vessel. The fenofibrate was rapidly mixed using a magnetic stirrer at a constant temperature controlled by a water bath (Shanghai Yiheng MP-501A, uncertainty of ±0.05 K). Additional solvent was subsequently introduced dropwise into the dissolver at a rate of 10−15 g/h using an injector. The equilibrium point was gauged by the intensity change of the laser beam entering the vessel. The experiment was repeated three times under identical conditions. The estimated relative standard uncertainty of the solubility data was less than 2%. The mole fraction solubility of the fenofibrate (xA) and the mole fraction of organic solvents (ethanol or acetone) in the binary solvents (f b) were given by the following equations:

Figure 1. Molecular structure of fenofibrate.

scanning calorimetry (DSC 3, Mettler Toledo) was used to obtain the melting point and the enthalpy of fusion at a heating rate of 10 K·min−1 from 303.15 to 423.15 K. The thermal analysis results are shown in Figure 2. Thermogravimetric

xA =

mA /MA (mA /MA ) + (mB /MB) + (mC /MC)

(1)

fb =

mB /MB (mB /MB) + (mC /MC)

(2)

where mA, mB, and mC represent the masses of fenofibrate, organic solvent, and water, respectively. MA, MB, and MC are the molecular weights of fenofibrate, organic solvent, and water, respectively.

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RESULTS AND DISCUSSION Melting Point and Enthalpy of Fusion. The thermal analysis results are shown in Figure 2. The melting point of fenofibrate was 353.62 K and the enthalpy of fusion of fenofibrate was 31.2 kJ·mol−1, the values were in good agreement with literatures.6−8 Solubility Data. The measured fenofibrate solubility of ethanol−water and acetone−water at different temperatures were given in Table 2 and Figure 4. The solubility data of fenofibrate were compared with the reported data,8 and given in Figure 5 and Figure 6, the results shown that our data were very concordant with the literature. The data showed that solubility increased with temperature and the mole fraction of organic solvents. From Figure 4, it can be seen that the solubility of fenofibrate was higher in acetone−water than in ethanol−water. This follows the empirical but useful “like dissolves like” rule. From Figure 1, we can conclude that the fenofibrate contains the aldehyde group, which is similar in structure to acetone. In intermolecular terms, fenofibrate is a medium-sized aprotic hydrophobic molecule whose intermolecular forces consist of London forces and Debye interactions.8 The dipole moment of fenofibrate rises from 3.92 to 4.87 D when going from εr = 1 to εr = 78.39.15 The ethanol provides the instantaneous dipole moment, since the solubility of fenofibrate is higher in ethanol than in water. In addition, ethanol is a polar protic solvent, while acetone is a dipolar aprotic solvent. The fenofibrate interacts more readily with the acetone molecules, forming dipole−dipole interactions.16 Correlation of Solubility Data Using Thermodynamic Models. In order to quantitatively describe the solid−liquid equilibrium of fenofibrate, modified Apelblat, CNIBS/R−K, J−A, and NRTL models were applied to correlate with the experimental solubility data. Correlation Using the Modified Apelblat Model. The modified Apelblat equation was the classic model for simulation of the solubility data and was used in some literature17−20 B ln xA = A + + C ln T (3) T

Figure 2. Differential scanning calorimeter of fenofibrate. The uncertainty for the temperature of fusion was u(T) = 0.5 K. The uncertainty for the enthalpy of fusion was ur(ΔH) = 0.005.

analysis (DSC Q600, TA Instruments) was used to measure the thermal stability at a heating rate of 10 K·min−1 from 303.15 to 973.15 K. The results are shown in Figure 3. Descriptions of

Figure 3. Thermogravimetric analysis of fenofibrate.

the chemicals used in the experiments are provided in Table 1. Deionized water was used in all experiments. Apparatus and Procedure. A laser monitoring observation technique was applied to measure the solubility of the fenofibrate.10−13 The apparatus is the same as that described in a previous study.14 Predetermined excess amounts of fenofibrate and the solvent of known mass, weighed by an analytical balance (Mettler Toledo AB204-N, Switzerland) with an B

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Table 1. Commercial Sources and Important Physical Properties of Chemicals Used in This Studya chemical name

CAS registry no.

source

formula weight

fenofibrate ethanol acetone water

[49562-28-9] [64-17-5] [67-64-1] [7732-18-5]

North China Pharmaceutical Co. Tianjin Chemical Reagent Co. Tianjin Chemical Reagent Co.

360.84 46.07 58.08 18.02

a b

boiling point (K)

dipole moment

351.45c 329.15c 373.15c

2.88 Dc 1.84 Dc

mass purity

purification method

≥99.8% ≥99.7% ≥99.5%

recrystallization method Ib method Ib distilled

The chemical properties and experiments were measured 100 ± 2 kPa, and all the studied experiments as follows were measured under 100 ± 2 kPa. Analytical grade reagents dehydrated with molecular sieves. cThese data were from Lange’s Handbook of Chemistry, 16th edition, 2005.

Table 2. Solubility Data of Fenofibrate in Ethanol−Water between 278.15 and 308.15 K and Solubility Data of Fenofibrate in Acetone−Water between 278.15 and 303.15 K under 100 ± 2 kPaa f bb

278.15 K

283.15 K

0.1992 0.3820 0.6013 0.7883 1.0000

0.2013 0.7868 4.6114 15.2961 27.4596

0.3122 1.3494 7.6133 21.0211 38.3741

0.1981 0.4211 0.6124 0.8057 1.0000

0.0674 1.3217 9.4582 29.1766 62.7294

0.0874 2.0673 12.1784 37.9691 77.0991

288.15 K

293.15 K

Ethanol−Water (104 xAc) 0.4264 0.5679 2.0344 2.9295 11.7567 15.9809 27.9653 38.9721 51.6506 68.0760 Acetone−Water (103 xAc) 0.1142 0.1376 2.8497 4.0183 17.3381 23.9186 49.9791 67.8086 97.1571 120.4610

298.15 K

303.15 K

308.15 K

0.7234 4.0860 22.0104 54.1592 93.3989

0.8811 5.8155 30.2964 74.2803 131.0742

1.0584 8.3466 44.6021 108.8920 201.4386

0.1668 5.7038 36.5280 87.3905 147.4556

0.2053 8.4598 52.3658 113.7514 180.2957

a

u(T) = 0.05 K; u( f b) = 0.0001; ur(xA) = 0.03; u(p) = 2 kPa. bThe mole fraction of organic solvent in organic−water binary mixed solvents (solute-free). cExperimental mole fraction solubility of fenofibrate.

Figure 4. Experimental mole fraction of fenofibrate in different binary solvents between different temperatures and mole fractions of organic. Correlation results in (a) ethanol−water and (b) acetone−water mixtures.

Figure 5. Experimental and literature8 mole fraction of fenofibrate in ethanol.

Figure 6. Experimental and literature8 mole fraction of fenofibrate in acetone. C

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Table 3. Regression Parameters of the Modified Apelblat Model for Fenofibrate Solubility in Binary Solvent Mixturesa

a

f bb

10−2 Ac

10−4 Bc

0.1992 0.3820 0.6013 0.7883 1.0000

8.5893 4.0735 3.6087 −4.9331 −6.4231

−4.2026 −2.3787 −2.1411 1.6651 2.3231

0.1981 0.4211 0.6124 0.8057 1.0000

2.9018 −1.3264 −9.1730 −1.3687 −0.4247

−1.6176 0.0333 3.4767 0.1895 −0.1297

Cc Ethanol−Water −127.6856 −58.8556 −51.8004 75.8630 98.2429 Acetone−Water −42.9315 22.1812 139.9461 22.4790 7.8822

R2d

104 RMSDe

MDf

0.9993 0.9984 0.9970 0.9997 0.9981

0.0073 0.1230 0.9972 0.7867 3.1832

1.1465 2.8145 3.7417 1.0344 2.6815

0.9988 0.9982 0.9990 0.9995 0.9998

0.0186 1.0207 5.7559 6.3298 4.8851

1.1061 2.2457 1.7522 0.8771 0.4384

u(T) = 0.05 K; u(f b) = 0.0001; u(p) = 2 kPa. bThe mole fraction of organic solvent in organic−water binary mixed solvents (solute-free). cRegression

parameters of the modified Apelblat equation. dSquared correlation coefficients. eRMSD =

∑iN= 1(xiexp − xical)2 N

; fMD = 100

∑iN= 1 |xiexp − xical| / xiexp N

.

Table 4. Regression Parameters of the NRTL Model for Fenofibrate Solubility in Binary Solvent Mixturesa ethanol−water acetone−water a

10−4 Δg12b

10−4 Δg21b

10−4 Δg13b

10−4 Δg31b

10−4 Δg23b

10−4 Δg32b

R2c

103 RMSDd

1.0897 −0.9827

−1.6411 1.6937

0.4799 0.7368

3.3542 2.0326

37.1522 0.9127

6.6047 4.7889

0.9414 0.9702

0.0519 3.8344

u(T) = 0.05 K; u( f b) = 0.0001; u(p) = 2 kPa.

d

RMSD =

b

Regression parameters of the NRTL model.

Squared correlation coefficients.

N

x3 equals (1 − x2), and

A, B, and C are empirical parameters. From eq 3, we can conclude that the values of A and B represent the variation of solution activity coefficients, whereas B and C represent the temperature effect of the enthalpy of fusion. The regressed empirical parameters are listed in Table 3. The corresponding root-mean-square deviation (RMSD) and mean deviation (MD) were used to judge the goodness of the fit. The expressions were defined as follows, and the values obtained are given in Table 3

Gij = exp( −αijτij)

τij =

∑i = 1 (xiexp − xical)2 N

(4)

N

MD = 100

∑i = 1 |xiexp − xical| /xiexp N

Δgij RT

(7)

i , j = 1, 2, 3; i ≠ j

(8)

Here, αij is a parameter related to the nonrandomness of the solution. It normally ranges from 0.20 to 0.47, and was 0.30 in the current study. Δgij represents the cross interaction energy parameter (gij − gjj) for the NRTL model. The solubility data fitting results are given in Table 4. Correlation Using the CNIBS/R−K Model. The CNIBS/ R−K model, originally proposed by Acree et al.,23,24 is a common correlate function used to address the solid−liquid equilibrium of binary solvent systems.25−27 The model relates the solubility of the solute with the mole fraction of the binary solvent mixture. It can be described as follows:

N

RMSD =

(5)

where xexp is the experimental value and xcal is the value i i calculated by the model. N represents the number of experimental data points measured in each system. In this study, N = 7 in ethanol−water and N = 6 in acetone−water. Correlation Using the NRTL Model. The activity coefficient of the NRTL model for a ternary system can be expressed as follows:21,22 ln γA =

c

∑iN= 1(xiexp − xical)2

n

ln xA = fa ln xa + fb ln xb + fa fb

∑ Si(fa

− fb )i

i=0

(9)

where fa and f b are the initial mole fractions of the solvents on a solute-free basis. In this study, fa represented the initial mole fractions of water, f b was the initial mole fraction of organic solvent, xa and xb represented the initial mole fraction solubility of the solutes in each pure solvent, Si was the model constant, and n represented the number of solvents (in this study, n = 2). Substituting (1 − f b) for n = 1 yielded the following:

τ21G21x 2 + τ31G31x3 τ G x + τ31G31x3 − 21 21 2 xA + G21x 2 + G31x3 xA + G21x 2 + G31x3 τ G x + τ32G32x3 xA × + τ12 − 12 12 A xA + G21x 2 + G31x3 x 2 + G12xA + G32x3 τ G x + τ23G23x 2 G12x 2 × + τ13 − 13 13 A x 2 + G12xA + G32x3 x3 + G13xA + G23x 2 G13x3 × x3 + G13xA + G23x 2 (6)

ln xA = ln xa + (ln xb − ln xa + S0 + S1)fb + ( −S0 + 3S1)f b2 + 2S1f b3

(10)

After simplifying, we obtained eq 11.

where x2 represents the mole fraction of organic in the binary solvent, x3 is the mole fraction of water in mixed solvent,

ln xA = B0 + B1fb + B2 f b2 + B3f b3 D

(11)

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Table 5. Regression Parameters of the CNIBS/R−K Model for Fenofibrate Solubility in Binary Solvent Mixturesa

a

Tb/K

B0c

B1c

B2 c

278.15 283.15 288.15 293.15 298.15 303.15 308.15

−11.7970 −11.8819 −12.0058 −11.8552 −11.8085 −12.0116 −12.2682

2.9323 6.7955 9.8921 10.8833 12.2471 15.1184 18.1688

278.15 283.15 288.15 293.15 298.15 303.15

−15.4916 −15.7125 −15.5994 −15.7005 −15.9615 −16.2440

20.2903 23.9251 24.9394 26.8710 29.9279 33.3817

B3 c

Ethanol−Water 11.8066 −8.8404 4.7341 −5.2165 −0.5261 −2.6379 −2.0535 −1.9711 −4.0934 −1.0230 −8.7977 1.3551 −13.4518 3.6461 Acetone−Water −11.6978 1.8153 −18.2101 5.1322 −19.5591 5.5813 −22.1018 6.5119 −26.4963 8.3044 −31.9223 10.7604

R2d

104 RMSDe

MDf

0.9999 0.9998 0.9993 0.9998 0.9999 0.9999 0.9999

0.0284 0.2982 0.8215 0.6445 0.5777 0.2243 0.0705

0.2760 2.0344 4.0606 2.3222 1.5059 0.4261 0.0909

0.9997 0.9999 0.9999 0.9999 0.9998 0.9998

0.7978 0.0118 0.4199 0.0576 1.6995 2.1800

3.7666 0.0438 1.1746 0.1204 2.6583 2.5737

u(T) = 0.05 K; u(f b) = 0.0001; u(p) = 2 kPa. bAbsolute temperature. cRegression parameters of the CNIBS/R−K model. dSquared correlation

coefficients. eRMSD =

∑iN= 1(xiexp − xical)2 N

; fMD = 100

∑iN= 1 |xiexp − xical| / xiexp N

Table 6. Regression Parameters of the J−A Model for Fenofibrate Solubility in Binary Solvent Mixtures between 278.15 and 313.15 Ka

a e

10−3 A0b

A1b

A2b

10−3 A3b

−5.6863

8.0023

7.5431

−0.5298

−6.1089

5.9727

5.7065

4.0493

10−3 A4b

10−3 A5b

Ethanol−Water 4.1883 −6.4450 Acetone−Water −0.1648 −5.5621

10−3 A6b

R2 c

104 RMSDd

MDe

2.4718

0.9971

3.1427

10.1514

3.0763

0.9973

7.1224

u(T) = 0.05 K; u(p) = 2 kPa. bRegression parameters of the J−A model. cSquared correlation coefficients. dRMSD =

MD = 100

13.4538 ∑iN= 1(xiexp − xical)2 N

.

∑iN= 1 |xiexp − xical| / xiexp N

where B0, B1, B2, and B3 are the model parameters obtained by nonlinear regression. The fitting results are given in Table 5. Correlation Using the J−A Model. To simultaneously fit temperature and mixed solvent composition,28−30 the J−A model was derived from the CNIBS/R−K model of Jouyban et al.31 as follows: n

ln xA = fa ln xa + fb ln xb + fa fb

∑ i=0

The modified Apelblat model gave better correlation results than the other models. Thermodynamic Properties of Fenofibrate Solutions. The thermodynamic properties for fenofibrate in ethanol− water and acetone−water were calculated in terms of their apparent standard dissolution enthalpy (ΔHsol), entropy (ΔSsol), and Gibbs free energy (ΔGsol) using the modified Apelblat model and the van’t Hoff equation32,33 The functions were expressed as follows:

Ji (fa − fb )i T

(12)

where Ji represents the constant of the model. When n = 2 and fa = (1 − f b), the model can be rewritten as eq 13. T ln xA = A 0 + A1T + A 2 Tfb + A3fb + A4 f b2 + A5f b3 + A 6f b4

(13)

where A0, A1, A2, A3, A4, A5, and A6 represent the model parameters. The solubility data on fenofibrate fitting are given in Table 6. From Tables 3−6, it can be seen that the R2 values were close to 1, and the RMSD and MD values were low, indicating good agreement between the measured and calculated solubility. The modified Apelblat model provided a regression relating ln xA to T, and was used to predict the solubility of fenofibrate with a fixed binary solvent composition at different temperatures. The CNIBS/R−K model was used to predict the solubility for different binary solvent compositions at a fixed temperature, whereas the J−A model was used to predict the solubility of different binary solvent compositions at different temperatures.

⎛ B⎞ ΔHsol = RT ⎜C − ⎟ ⎝ T⎠

(14)

ΔSsol = R(A + C + C ln T )

(15)

⎛ ⎞ B ΔGsol = −RT ⎜A + + C ln T ⎟ ⎝ ⎠ T

(16)

where A, B, and C represent the parameters obtained by regression using the modified Apelblat model (Table 3) and R is the universal ideal gas constant (8.314 J mol−1 K−1). To minimize the error propagation, the mean harmonic temperature T was defined by eq 17.

T=

N ∑ (1/Ti )

(17)

where N is the number of experimental temperatures studied. In our ethanol−water system, N was 7, Ti ranged from 278.15 to 308.15 K, and T was 292.81 K. In the acetone−water system, N was 6 and T was 290.40 K. E

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dissolution enthalpy, entropy, and Gibbs free energy changes were calculated from the solubility data using the modified Apelblat model and the van’t Hoff equation. The results demonstrate that fenofibrate dissolution in these binary solvents is an endothermic, entropy-driven process.

The relative contributions of enthalpy (%ζH) and entropy (%ζTS) to the dissolution process were compared by using eqs 18 and 19, respectively34,35 %ζH = 100 ×

|ΔHsol| |ΔHsol| + |T ΔSsol|

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(18)

|T ΔSsol| %ζTS = 100 × |ΔHsol| + |T ΔSsol|

AUTHOR INFORMATION

Corresponding Author

(19)

*E-mail: [email protected]. Fax: 0086-311-88632183.

The thermodynamic parameter results are given in Table 7. It can be seen that the ΔHsol values were all positive during the

Notes

Table 7. Thermodynamic Properties for the Dissolution of Fenofibrate at 292.81 K

Funding

f ba

ΔHsolb (kJ mol−1)

0.1992 0.3820 0.6013 0.7883 1.0000

38.56 54.49 51.91 46.25 46.02

0.1981 0.4211 0.6124 0.8057 1.0000

30.83 50.78 48.83 38.52 29.81

ΔSsolc (J mol−1 K−1) Ethanol−Water 50.30 118.27 123.64 111.60 115.57 Acetone−Water 31.35 127.50 135.62 108.85 84.06

ΔGsold (kJ mol−1)

%ζHe

%ζTSf

23.83 19.86 15.71 13.57 12.18

72.36 61.14 58.91 58.60 57.63

27.64 38.86 41.09 41.40 42.37

21.73 13.76 9.45 6.91 5.40

77.21 57.83 55.36 54.92 54.98

22.80 42.17 44.64 45.08 45.02

The authors declare no competing financial interest. The authors thank the support of the Hebei Food and Drug Administration (ZD2015026), the Natural Science Foundation of Hebei Province (No. B2015206108), and the National Natural Science Foundation of China (No. 21406050).

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NOTATIONS

xA

mole fraction of fenofibrate in the solution fa initial mole fractions of water fb initial mole fraction of organic solvents in the solvents △Hfus enthalpy of fusion △Cp,m disparity of heat capacities between liquid and solid at the melting temperature R gas constant T absolute temperature Tm melting temperature (K) a, b the empirical constants of modified Apelblat model A, B, and C empirical parameters RMSD root-mean-square deviation MD mean deviation xexp experimental value i xcal calculated value of model i N number of experimental data Δgij cross interaction energy parameter for the NRTLmodel (gij − gjj) (J/mol) Si model constant n means the number of solvents B0, B1, B2, B3, B4 model parameters of CNIBS/R−K model A0, A1, A2, A3, A4, A5, A6 model parameters of J−A model ΔHsol dissolution enthalpy change (kJ/mol) ΔSsol dissolution entropy change (J/mol) ΔGsol dissolution free Gibbs energy change (kJ/mol)

a

The mole fraction of organic solvent in organic−water binary mixed solvents (solute-free). bThe solution enthalpy of fenofibrate. cThe solution entropy of fenofibrate. dThe apparent free Gibbs energy for the solution process of fenofibrate. eThe relative contributions by enthalpy toward the solution process. fThe relative contributions by entropy toward the solution process.

dissolution process, suggesting that the process is endothermic for both binary solvents. ΔGsol represents the minimum energy that the fenofibrate dissolution process needs to overcome under the experimental conditions. Because the ΔGsol value decreased with the increase in organic mole fractions, it could be deduced that the solute−solvent interactions grew and the fenofibrate solubility was enhanced with the addition of extra ethanol or acetone. The ΔGsol value was lower for acetone− water than for ethanol−water, suggesting that fenofibrate dissolves more readily in the former than in the latter. This was consistent with the solubility trend drawn from the relationship between temperature, mixed solvent composition, and solubility. In addition, the values of ΔSsol were positive, suggesting that the dissolution process was entropy-driven for the two binary solvents. Comparing %ζH and %ζTS, it could be concluded that the enthalpy is the main contributor to the standard Gibbs free energy in fenofibrate dissolution.

Greek Letters

■

γA αij %ζH %ζTS

CONCLUSIONS In this study, the solubility of fenofibrate in ethanol−water and acetone−water mixtures was measured under atmospheric pressure and at temperatures between 278.15 and 308.15 K using dynamic laser monitoring. Solubility was found to increase as the temperature and the organic content of the binary solvents increased. Acetone is a suitable solvent, whereas water is a good antisolvent in the industrial crystallization of fenofibrate. The solubility data fit well with the modified Apelblat, NRTL, CNIBS/R−K, and J−A models. The apparent

activity coefficient nonrandomness parameter relative enthalpy contribution relative entropy contribution

Subscripts

A fus m sol H TS F

fenofibrate fusion melting dissolution property enthalpy entropy DOI: 10.1021/acs.jced.6b00268 J. Chem. Eng. Data XXXX, XXX, XXX−XXX

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DOI: 10.1021/acs.jced.6b00268 J. Chem. Eng. Data XXXX, XXX, XXX−XXX