Thermodynamic Models for Determination of the Solubility of Sulfanilic

Nov 3, 2014 - ABSTRACT: Data on corresponding solid−liquid equilibria of sulfanilic acid in different organic solvents are essential for industrial ...
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Thermodynamic Models for Determination of the Solubility of Sulfanilic Acid in Different Solvents at Temperatures from (278.15 to 328.15) K Yonghong Hu,*,† Qi Zhang,† Ying Shi,‡ Yang Yang,† Limin Cheng,† Cuicui Cao,† and Wenge Yang† †

State Key Laboratory of Materials-Oriented Chemical Engineering, Nanjing Tech University, Nanjing 210009, China Taiyuan Qiaoyou Chemical Industrial Co. Ltd., Taiyuan 030025, China

‡

ABSTRACT: Data on corresponding solid−liquid equilibria of sulfanilic acid in different organic solvents are essential for industrial design and further theoretical studies. In this study, the solubilities of sulfanilic acid were measured in methanol, ethanol, 2-propanol, acetone, acetonitrile, formic acid, and acetic acid with the temperature range of (278.15 to 328.15) K by the analytical stirred-flask method under atmospheric pressure. For the temperature range investigated, the solubility of sulfanilic acid in the solvents increased with increasing temperature. The thermodynamic models, such as the van’t Hoff model, the modified Apelblat model, and the Buchowski−Ksiazaczak λh model were investigated to describe the experimental data. It was found that the modified Apelblat model was the most suitable for predicting the solubility behavior of sulfanilic acid with increasing temperature. The calculated thermodynamic parameters indicated that in each studied solvent the dissolution process of sulfanilic acid is endothermic. Apelblat equation and the Buchowski−Ksiazaczak λh equation have been used to correlate and predict the solubility of sulfanilic acid in pure solvents. Thermodynamic models of C6H7NO3S solubility corresponding to mathematic expressions are developed and employed to correlate with the experimental data. For the thermodynamic parameters of dissolution, the standard enthalpy (ΔH0soln), standard entropy (ΔS0soln), and Gibbs free energy (ΔG0soln) of the solution of sulfanilic acid are determined from the solubility data by the van’t Hoff equation.

1. INTRODUCTION Sulfanilic acid (C6H7NO3S, CASRN: 121-57-3, shown in Figure 1) also called 4-aminobenzenesulfonic acid, is a greyish-white

Figure 1. Chemical structure of sulfanilic acid.

2. EXPERIMENTAL SECTION 2.1. Materials. Sulfanilic acid (C6H7NO3S) with a purity of > 99.8 % was purchased from Sinopharm Chemical Reagent Co., Ltd., China. Its purity was measured by high-performance liquid chromatography (HPLC type DIONEX P680 DIONEX Technologies), and its melting point was measured by a digital melting point system (type WRS-1B, Shanghai Precision & Scientific Instrument Co., Ltd.) as being between (560.85 and 561.30) K (the solid acid of sulfanilic acid exists as a zwitterion and has an unusually high melting point).2 All of the organic solvents for dissolving were supplied by Shanghai Shenbo Chemical Co., Ltd., China. The purities of the solvents were determined in our laboratory by high-performance liquid chromatography, and their mass fraction purities were higher than 0.99. Meanwhile, all chemicals were used as received without further purification. The materials table is depicted in Table 1. 2.2. Apparatus and Procedures. The solubility of sulfanilic acid was investigated, in various organic solvents,

crystal or powder, which is widely used as an important intermediate in the production of pesticides, pharmaceuticals, dyes, chromatographic analytical reagents, and so on.1,2 Sulfanilic acid is a primary aryl amine, which can be obtained by heating aniline with weak fuming sulfuric acid and pouring the reaction product into water. The solubilities of organic compounds in different solvents play an important role for understanding the solid− liquid equilibria (SLE) or phase equilibria in the development of a crystallization process or liquid−liquid equilibria in extraction and extractive or azeotropic distillation processes. More particularly, knowledge of an accurate solubility is needed for the design of separation processes such as extractive crystallization and the safety of operating different processing units such as distillation columns, absorption units, and extraction plants. The solubility of sulfanilic acid can also supply basic and theoretical data for industrial production. To determine proper solvents and to design an optimized production process, it is necessary to know the solubilities of sulfanilic acid in different organic solvents.3 The present work is focused on the study of the solubilities of sulfanilic acid in seven solvents, methanol, ethanol, 2-propanol, acetone, acetonitrile, formic acid, and acetic acid in the temperature range from (278.15 to 328.15) K.4,5 The modified © 2014 American Chemical Society

Received: September 19, 2014 Accepted: October 22, 2014 Published: November 3, 2014 3938

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Table 1. Properties of Solventsa properties

a

solvent

MW/(g·mol−1)

D/(g·mol−1)

source

mole fraction purity

analysis method

methanol ethanol 2-propanol acetone acetonitrile formic acid acetic acid sulfanilic acid

32.04 46.07 60.10 58.08 41.05 46.03 60.05 173.19

0.7915 0.7893 0.7855 0.7848 0.7860 1.2200 1.0492 1.4850

Shenbo Chemicals Shenbo Chemicals Shenbo Chemicals Shenbo Chemicals Shenbo Chemicals Shenbo Chemicals Shenbo Chemicals Sinopharm Chemicals

99.7 99.9 99.7 99.8 99.5 99.5 99.8 99.8

HPLC HPLC HPLC HPLC HPLC HPLC HPLC HPLC

The density values are the standard values (provided by Shanghai Shenbo Chemical Co., Ltd. and Sinopharm Chemical Co., Ltd.).

decrease in the following order: acetic acid > formic acid > 2propanol > methanol > ethanol > acetonitrile. The lines representing the solubilities of sulfanilic acid in acetone intersect the lines representing the solubilities in other solvents. 3.2. Thermodynamic Models. The correlation between temperature and solubility of sulfanilic acid in different solvents can be explained by thermodynamic models. The van’t Hoff model presented in eq 2 was normally used to predict the relation of solubility vs reciprocal temperature (1/T) in solid−liquid equilibrium.

by the analytical stirred-flask dissolution method, and we used a gravimetric method to measure the compositions of the saturated solutions. Saturated solutions of sulfanilic acid were produced by adding 8 mL of the corresponding solvent and excess sulfanilic acid to a jacketed spherical 10 mL Pyrex glass flask. The jacketed flasks were maintained at the desired temperature controlled by a thermostat within ± 0.1 K that was supplied from a constant-temperature water bath (type HWC-52, ShangHai Cany Precision Instrument Co., Ltd.). For each measurement, excess sulfanilic acid was added to a known volume of solvent. Continuous stirring was achieved for fully mixing the suspension using a magnetic stirrer. The stirring continued for about 24 h to ensure solid−liquid equilibrium, and the solution was allowed to settle for 12 h or more before sampling.6−8 The supernatant was taken and filtered (vacuum filtration, pore size of filter membrane is 0.22 μm), and the filtrate was poured into a volumetric flask preweighed by using an analytical balance (Sartorius, CPA225D, Germany) with a resolution of ± 0.01 mg. Finally, 1 mL of solution was transferred into a preweighed covered 5 mL breaker and weighted immediatedly. The breaker was put into a dryer at room temperature and weighed weekly until reaching constant weight. All experiments were repeated three times for reproducibility to obtain the mean values, and the mean values were used to calculate the mole fraction solubility. The saturated mole fraction solubility (x) of the solute in the solvent is calculated by eq 1. x=

m1/M1 m1/M1 + m2 /M 2

ln x = A +

B T

(2)

where x is the mole fraction solubility of sulfanilic acid, T is the corresponding temperature in K, and A and B are constants and are listed in Table 3. To describe solid−liquid equilibria, the temperature dependence of the solubilities of sulfanilic acid in different organic solvents at different temperatures can be described as9 ⎛ 1 ⎞ ΔH ⎛ T ⎞ ΔCp ⎛ Tt ⎞ ΔCp Tt fus ⎜ t ⎜ ln⎜⎜ ⎟⎟ = ln − 1⎟ − − 1⎟ + ⎠ ⎠ RTt ⎝ T R ⎝T R T ⎝ γxx ⎠

(3)

where γx is the activity coefficient of sulfanilic acid on a mole fraction basis, x is the mole fraction solubility of sulfanilic acid, ΔHfus is the enthalpy of fusion of sulfanilic acid, ΔCp is the change of the heat capacity, T is the absolute temperature (K), Tt is the triple-point temperature of sulfanilic acid, and R is the gas constant. For regular solutions, the activity coefficient is obtained by

(1)

ln γx = a +

where m1 and m2 represent the mass of sulfanilic acid and the mass of solvent, respectively. M1 is the molar mass of sulfanilic acid; M2 is the molar mass of solvents.

b T /K

(4)

where a and b are empirical constants. Introducing γx from eq 4 into eq 3 and subsequent rearrangement results in eq 5 which can be written as10,11

3. RESULTS AND DISCUSSION 3.1. Solubility Data. The measured mole fraction solubilities of sulfanilic acid in methanol, ethanol, 2-propanol, acetone, acetonitrile, formic acid, and acetic acid with the temperature range of (278.15 to 328.15) K are presented in Table 2 and graphically shown in Figures 2 and 3 (the solubilities in formic acid and acetic acid are much bigger than for the other solvents, so two graphs were used here). As can be seen from the graphs, it could be found that solubility is a function of temperature, but the trend with temperature of solubility is different. From (278.15 to 328.15) K, the variation of increasing trend with temperature is gentle for acetone, while for the other solvents the behaviors are different, the solubilities of sulfanilic acid increase quickly. The molar solubilities

⎡ ΔH ⎤ ΔCp fus ln x = ⎢ − (1 + ln Tt) − a⎥ R ⎣ RTt ⎦ ⎡ ⎛ ΔHfus ΔCp ⎞ ⎤ 1 ΔCp − ⎢b + ⎜ − ln T ⎟Tt ⎥ + ⎢⎣ R ⎠ ⎥⎦ T R ⎝ RTt

(5)

Because the coefficients are constants, the solubility in the above equation may be written as the modified Apelblat model, which was first used by Apelblat and Manzurola.12,13 The model can give a relatively more accurate correlation with another parameter in eq 2: ln x = A + 3939

B + C ln T T

(6)

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Table 2. Mole Fraction Solubilities of Sulfanilic Acid in Different Organic Solvents with the Temperature Range from (278.15 to 328.15) K under 0.1 MPaa 100 RD

100 RD

T/K

4

10 x

van’t Hoff model (eq 2)

278.15 283.15 288.15 293.15 298.15 303.15 308.15 313.15 318.15 323.15 328.15

2.530 2.70 3.030 3.180 3.480 3.730 4.030 4.290 4.610 4.940 5.250

1.976 −0.741 1.980 −1.572 −0.287 −0.804 −0.248 −0.699 −0.217 0.405 0.571

278.15 283.15 288.15 293.15 298.15 303.15 308.15 313.15 318.15 323.15 328.15

2.470 2.630 2.840 2.990 3.180 3.420 3.620 3.830 4.030 4.320 4.500

1.619 0.380 0.704 −1.003 −1.258 −0.292 −0.552 −0.522 −0.744 1.157 0.444

278.15 283.15 288.15 293.15 298.15 303.15 308.15 313.15 318.15 323.15 328.15

2.130 2.310 2.510 2.640 2.840 3.000 3.230 3.380 3.630 3.850 4.030

0.469 0.866 1.195 −0.758 0.000 −1.333 0.000 −1.479 0.275 0.779 0.248

278.15 283.15 288.15 293.15 298.15

2.910 2.960 3.030 3.10 3.170

1.031 0.000 −0.330 −0.323 −0.315

a

modified Apelblat model (eq 6)

2-Propanol 0.395 −1.481 1.980 −1.257 0.287 −0.268 0.248 −0.233 0.000 0.202 −0.190 Methanol 0.405 −0.380 0.704 −0.669 −0.629 0.585 0.000 0.000 −0.496 0.926 −0.444 Ethanol −0.939 0.000 1.195 −0.379 0.352 −0.667 0.310 −0.888 0.275 0.519 −0.248 Acetone 0.029 0.015 0.010 0.008 0.006

Buchowski− Ksiazaczak λh model (eq 10)

T/K

10 x

van’t Hoff model (eq 2)

303.15 308.15 313.15 318.15 323.15 328.15

3.240 3.320 3.380 3.460 3.530 3.600

−0.617 0.000 −0.296 0.000 0.283 0.556

278.15 283.15 288.15 293.15 298.15 303.15 308.15 313.15 318.15 323.15 328.15

1.320 1.440 1.650 1.780 1.950 2.140 2.350 2.550 2.740 3.010 3.240

278.15 283.15 288.15 293.15 298.15 303.15 308.15 313.15 318.15 323.15 328.15

10.600 11.580 12.900 13.900 15.200 16.600 18.100 19.600 21.300 23.100 24.680

288.15 293.15 298.15 303.15 308.15 313.15 318.15 323.15 328.15

20.370 21.700 22.700 23.980 25.160 26.360 27.690 29.100 30.370

4

modified Apelblat model (eq 6)

Buchowski− Ksiazaczak λh model (eq 10)

Acetone −0.796 −2.222 1.650 −0.943 0.575 0.268 0.744 0.233 0.434 0.607 0.000 −0.405 −0.760 0.352 −0.669 −0.314 0.877 0.552 0.261 −0.248 0.926 −0.444 −1.878 −0.433 1.195 −0.379 0.704 0.000 0.929 −0.592 0.551 0.519 −0.496 −0.344 −0.676 0.000 0.323 0.315

0.005 0.005 0.004 0.004 0.004 0.003 Acetonitrile 2.273 0.000 −0.694 −1.389 2.424 1.818 −0.562 0.000 −1.026 0.000 −0.935 −0.467 −0.426 0.426 −0.392 0.000 −1.095 −1.095 0.664 0.332 0.926 0.000 Formic Acid 2.642 0.283 0.777 −0.432 1.318 1.085 −0.863 −0.504 −1.184 −0.395 −1.145 −0.181 −0.829 0.055 −0.816 −0.102 −0.047 0.235 0.823 0.563 0.527 −0.446 Acetic Acid 0.344 −0.344 0.737 0.553 −0.308 −0.176 −0.167 0.167 −0.397 −0.040 −0.569 −0.228 −0.253 −0.072 0.309 0.241 0.395 −0.066

0.309 0.602 0.296 0.289 0.000 −0.556 −0.758 −2.083 1.818 0.000 0.000 0.000 0.851 0.392 −0.365 0.664 0.000 −0.660 −1.122 0.698 −0.719 −0.395 0.000 0.442 0.306 0.704 1.082 0.041 −1.080 0.230 −0.176 0.375 0.318 0.152 0.289 0.447 −0.033

95 % uncertainties u are u(x) = 2 %, u(T) = 0.05 K, and u(p) = ± 2 KPa.

where T is the absolute temperature, and the A, B, and C are parameters obtained by fitting the experimental solubility data and are shown in Table 4 together with the corresponding rootmean-square deviations (RMSD) between the data and eq 6, in mole fraction units. The values of B and C reflect the variation in the solution activity coefficient and provide an indication of the effect of solution nonidealities on the solute. The agreements between experimental data and data calculated using the thermodynamic models expressed as the relative deviation (RD) defined by eq 3 are listed in Table 2. The parameters of models obtained from linear and nonlinear regressions and their values are listed in Tables 3 to 5. The relative absolute deviation (RAD) and the corresponding

root-mean square deviation (RMSD) as expressed by eqs 7 to 9 are shown in Tables 3 to 5.14,15 x − xci RD = i xi (7) RAD =

1 N

N i=1

⎡1 RMSD = ⎢ ⎢⎣ N 3940

xi − xci xi

∑ N

(8)

∑ (xci − xi) i=1

⎤1/2

2⎥

⎥⎦

(9)

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where x is the mole fraction solubility of sulfanilic acid and T and Tm are the experimental temperature and the standard melting temperature of sulfanilic acid in Kelvin, respectively. λ and h are model parameters determined by the experimental data and are listed in Table 5 together with RMSD values in the system together with the corresponding rmsd’s which are listed in Table 5, respectively. As we can see from Tables 2 to 5, the predicted data of sulfanilic acid in seven pure organic solvents show good agreement with the experimental data with small RMSDs. Taking the solubility data in the selected solvents fitted by the modified Apelblat equation as an illustration, the relative absolute deviations are less than 0.6 %, with individual values (Table 2) not exceeding 2 %. As the Apelblat equation has the best fit of the three equations (Tables 3 to 5), it is proposed as the most suitable for correlating the sulfanilic acid solubility data. The same conclusion can be drawn after analyzing the solubility data and the parameters that fitted by the van’t Hoff equation and the λh equation. Meanwhile, we can see that the 102 RAD values of van’t Hoff equation, the modified Apelblat equation, and the Buchowski−Ksiazaczak λh equation are 0.727, 0.386, and 0.553. This result indicates that the modified Apelblat model proved to be more accurate and suitable for the description of dissolution of sulfanilic acid in the studied solvents at various temperatures. 3.3. Thermodynamic Parameters. The standard molar enthalpy of solution (ΔH0soln) can be related to the temperature and the solubility with the van’t Hoff analysis, as17

Figure 2. Mole fraction solubility (x) of sulfanilic acid versus temperature (T) in the selected organic solvents: ■, 2-propanol; ●, methanol; ▲, ethanol; ★, acetone; □, acetonitrile.

⎛ ∂ ln x ⎞ 0 ΔHsoln = −R·⎜ ⎟ ⎝ ∂(1/T ) ⎠

where R represents the gas constant (8.314 J·mol−1·K−1) and x is the mole fraction of sulfanilic acid at the corresponding absolute temperature T of the solution. ΔH0soln can be obtained from the slope of the solubility curve where ln x is plotted versus 1/T. Figures 4 and 5 show the linearity of the ln x vs (1/T) plots of sulfanilic acid in the selected solvents. In the present work, the mole fraction solubility of sulfanilic acid in the studied solvents at corresponding temperatures was measured. The standard molar Gibbs energy of solution (ΔG0soln) can be calculated according to18

Figure 3. Mole fraction solubility (x) of sulfanilic acid versus temperature (T) in the selected organic solvents: ○, acetic acid; △, formic acid [melting point of acetic acid: (289 to 290) K].

The Buchowski−Ksiazaczak λh equation (eq 10), which was suggested first by Buchowski et al.,16 is another way to describe the solution behavior. The Buchowski−Ksiazaczak λh equation has fitted the experimental data well for many systems with only two parameters λ and h. In this paper, the solubility data were also correlated with the Buchowski−Ksiazaczak λh equation: ⎡1 ⎡ λ(1 − x) ⎤ 1 ⎤ ln⎢1 + ⎥ ⎥ = λh⎢ − ⎣ ⎦ x Tm ⎦ ⎣T

(11)

0 ΔGsoln = −RTmean·intercept

(12)

where the intercept is determined in plots of ln x as a function of (1/T) (Figures 4 and 5). The standard molar entropy of solution (ΔS0soln) is obtained by 0 ΔSsoln =

(10)

0 0 ΔHsoln − ΔGsoln Tmean

(13)

Table 3. Parameters of the van’t Hoff Model for Sulfanilic Acid Solubility in the Selected Organic Solvents at Temperatures from (278.15 to 328.15) K solvents

A

2-propanol methanol ethanol acetone acetonitrile formic acid acetic acid

−6.0 −6.0 −6.0 −7.4 −6.0 −4.0 −4.0

B (−2.04 (−1.69 (−1.79 (−6.00 (−2.50 (−2.38 (−1.45

± ± ± ± ± ± ±

0.07)·105 0.06)·105 0.15)·105 0.03)·105 0.07)·105 0.27)·105 0.04)·105

3941

102 RAD

104 RMSD

0.864 0.789 0.673 0.341 1.038 0.997 0.387 102 RAD = 0.727

0.036 0.029 0.025 0.014 0.023 0.165 0.104

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Table 4. Parameters of the Modified Apelblat Model for Sulfanilic Acid in the Different Organic Solvents at Temperatures from (278.15 to 328.15) K solvent

A

B

2-propanol methanol ethanol acetone acetonitrile formic acid acetic acid

−9.07 ± 0.12 −17.33 ± 0.14 −7.37 ± 0.09 −20.05 ± 0.12 −8.3 ± 0.05 −16.63 ± 0.10 −13.13 ± 0.13

(−1.09 ± 0.04)·10 (−5.21 ± 0.11)·102 (−1.02 ± 0.03)·103 (2.03 ± 0.10)·102 (−1.40 ± 0.22)·103 (−8.65 ± 0.12)·102 (−4.71 ± 0.31)·102

102 RAD

104 RMSD

0.595 0.476 0.525 0.008 0.502 0.389 0.210 102 RAD = 0.386

0.026 0.020 0.018 3.252 0.015 0.077 0.060

C 3

0.84 1.94 0.46 1.99 0.78 2.29 1.51

± ± ± ± ± ± ±

0.05 0.09 0.02 0.10 0.02 0.11 0.07

Table 5. Parameters of the λh Model for Sulfanilic Acid in the Selected Organic Solvents λ

solvent 2-propanol methanol ethanol acetone acetonitrile formic acid acetic acid

0 0 0 0 0 0 0

h (7.17 (1.06 (1.12 (3.17 (8.66 (1.23 (1.95

± ± ± ± ± ± ±

5

0.05)·10 0.03)·106 0.19)·106 0.08)·106 0.22)·105 0.28)·105 0.30)·105

102 RAD

104 RMSD

0.770 0.528 0.698 0.337 0.630 0.561 0.344 102 RAD = 0.553

0.030 0.020 0.023 0.013 0.016 0.112 0.101

Table 6. Thermodynamic Parameters of Dissolution of Sulfanilic Acid in the Selected Organic Solvents at the Mean Temperature ΔH0soln −1

ΔS0soln −1

ΔG0soln −1

solvents

(kJ·mol )

(J·mol ·K )

(kJ·mol−1)

2-propanol methanol ethanol acetone acetonitrile formic acid acetic acid

11.16 ± 0.14 9.20 ± 0.09 9.62 ± 0.10 3.29 ± 0.05 13.64 ± 0.14 12.93 ± 0.09 7.79 ± 0.07

7.99 ± 0.00 −5.672 ± 0.00 −3.89 ± 0.00 −45.06 ± 0.00 19.75 ± 0.00 32.09 ± 0.00 1.22 ± 0.00

8.73 ± 0.14 10.921 ± 0.09 10.80 ± 0.10 16.95 ± 0.05 7.65 ± 0.14 3.20 ± 0.09 7.42 ± 0.07

% ξH

% ξTS

± ± ± ± ± ± ±

17.84 ± 0.18 15.74 ± 0.27 10.91 ± 0.17 80.58 ± 0.60 30.51 ± 0.21 42.95 ± 0.18 4.54 ± 0.03

82.16 84.26 89.09 19.42 69.49 57.05 95.46

0.18 0.27 0.17 0.60 0.21 0.18 0.03

Figure 5. A van’t Hoff plot of the mole fraction solubility (ln x) of sulfanilic acid in the selected organic solvents against 1/T with straight lines to correlate the data: ○, acetic acid; △, formic acid.

Figure 4. A van’t Hoff plot of the mole fraction solubility (ln x) of sulfanilic acid in the selected organic solvents against 1/T with straight lines to correlate the data: ■, 2-propanol; ●, methanol; ▲, ethanol; ★, acetone; □, acetonitrile.

contribution to the standard Gibbs energy by enthalpy and entropy in the solution process,19,20 respectively.

The values of the standard Gibbs energy, enthalpy, and entropy are listed in Table 6 (the uncertainty of the thermodynamic parameters are shown), together with ξH and ξTS. The study of the ξH and ξTS is aimed at comparing the relative

%ξH = 3942

0 |ΔHsoln | 0 0 |ΔHsoln | + |T ΔSsoln |

·100 (14)

dx.doi.org/10.1021/je500867u | J. Chem. Eng. Data 2014, 59, 3938−3943

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0 |T ΔSsoln | 0 |ΔHsoln |

+

0 |T ΔSsoln |

Article

determination of the solubility of 2,5-bis(2-furylmethylidene) cyclopentan-1-one in different solvents at temperatures ranging from 308.15 to 403.15 K. Fluid Phase Equilib. 2014, 367, 57−62. (9) Yang, X. Z.; Wang, J.; Li, G. S. Solubilities of Triadimefon in Acetone+ Water from (278.15 to 333.15) K. J. Chem. Eng. Data 2009, 54, 1409−1411. (10) Ruidiaz, M. A.; Delgado, D. R.; Martínez, F.; Marcus, Y. Solubility and preferential solvation of indomethacin in 1,4-dioxane + water solvent mixtures. Fluid Phase Equilib. 2010, 299, 259−265. (11) Ramírez-Verduzco, L. F.; Rojas-Aguilar, A.; De los Reyes, J. A.; Muñoz-Arroyo, J. A.; Murrieta-Guevara, F. Solid-liquid equilibria of dibenzothiophene and dibenzothiophene sulfone in organic solvents. J. Chem. Eng. Data 2007, 52, 2212−2219. (12) Apelblat, A.; Manzurola, E. Solubilities of L-aspartic, DLaspartic, DL-glutamic, p-hydroxybenzoic, o-anistic, p-anisic, and itaconic acids in water from T = 278 K to T = 345 K. J. Chem. Thermodyn. 1997, 29, 1527−1533. (13) Apelblat, A.; Manzurola, E. Solubilities of o-acetylsalicylic, 4aminosalicylic, 3,5-dinitrosalicylic, and p-toluic acid, and magnesiumDL-aspartate in water from T = (278 to 348) K. J. Chem. Thermodyn. 1999, 31, 85−91. (14) Alves, K. C.; Condotta, R.; Giulietti, M. Solubility of docosane in heptane. J. Chem. Eng. Data 2001, 46, 1516−1519. (15) Delgado, D. R.; Holguín, A. R.; Almanza, O. A.; Martínez, F.; Marcus, Y. Solubility and preferential solvation of meloxicam in ethanol + water mixtures. Fluid Phase Equilib. 2011, 305, 88−95. (16) Buchowski, H.; Ksiazczak, A.; Pietrzyk, S. Solvent activity along a saturation line and solubility of hydrogen-bonding solids. J. Phys. Chem. 1980, 84, 975−979. (17) Zhou, Z.; Qu, Y.; Wang, J.; Wang, S.; Liu, J.; Wu, M. Measurement and correlation of solubilities of (Z)-2-(2-aminothiazol4-yl)-2-methoxyiminoacetic acid in different pure solvents and binary mixtures of water + (ethanol, methanol, or glycol). J. Chem. Eng. Data 2011, 56, 1622−1628. (18) Lei, Z.; Hu, Y.; Yang, W.; Li, L.; Chen, Z.; Yao, J. Solubility of 2(2,4,6-trichlorophenoxy) ethyl bromide in methanol, ethanol, propanol, isopropanol, acetonitrile, n-heptane, and acetone. J. Chem. Eng. Data 2011, 56, 2714−2719. (19) Liu, J. Q.; Wang, Y.; Tang, H.; Wu, S.; Li, Y. Y.; Zhang, L. Y.; Liu, X. Experimental measurements and modeling of the solubility of aceclofenac in six pure solvents from (293.35 to 338.25) K. J. Chem. Eng. Data 2014, 59, 1588−1592. (20) Liu, X.; Hu, Y.; Liang, M.; Li, Y.; Yin, J.; Yang, W. Measurement and correlation of the solubility of maleic anhydride in different organic solvents. Fluid Phase Equilib. 2014, 367, 1−6.

·100 (15)

Table 6 demonstrates that the enthalpy and the standard Gibbs energy of sulfanilic acid are positive in the studied selected organic solvents, indicating that the solution process of sulfanilic acid in the seven solvents is endothermic. Also enthalpy is the main contributor to the standard Gibbs energy of solution during the dissolution, because the values of %ξH are large (except for acetone).

4. CONCLUSIONS The solubility data of sulfanilic acid in seven pure organic solvents within the temperature range between (278.15 and 328.15) K were measured leading to the following conclusions: (1) The solubility of sulfanilic acid in the selected solvents increased with increasing temperature, but the magnitude of the increase varied for different solvents; (2) The solubility data could be successfully correlated using three equations (van’t Hoff, modified Apelblat, and the λh). The modified Apelblat fit the data best; (3) Thermodynamic properties have been evaluated.

■

AUTHOR INFORMATION

Corresponding Author

* Tel.: +86 25 58139208. Fax: +86 25 58139208 E-mail: [email protected]. Funding

This research work was financially supported by Key Topics for State Key Laboratory of Materials-Oriented Chemical Engineering (ZK201304), the Doctoral Fund Project (20133221110010), NSFC (Natural Science Foundation of China) (No. 31471692), and the Joint Innovation and Research Funding-Prospective Joint Research Projects (BY2013005-02). Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS We thank the editor and the anonymous reviewers. REFERENCES

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dx.doi.org/10.1021/je500867u | J. Chem. Eng. Data 2014, 59, 3938−3943