Solubility of Salicylic Acid in Some (Ethanol + Water) Mixtures at Different Temperatures: Determination, Correlation, Thermodynamics and Preferential Solvation

Equilibrium mole fraction solubility of salicylic acid in nine aqueous-ethanolic mixtures, as well as in neat water and neat ethanol, was determined at seven temperatures from T = (293.15 to 323.15) K. Salicylic acid solubility in these mixtures was adequately correlated with well-known correlation/prediction methods based on Jouyban-Acree model. Apparent thermodynamic quantities, i.e. Gibbs energy, enthalpy, and entropy, for the dissolution and mixing processes, were computed by means of the van’t Hoff and Gibbs equations. The enthalpy–entropy compensation plot of enthalpy vs. Gibbs energy of dissolution was not linear exhibiting positive slopes from neat water to the mixture of w1 = 0.30 and from the mixture of w1 = 0.50 to neat ethanol indicating enthalpy-driven drug transfer processes but negative in the interval of 0.30 < w1 < 0.50 indicating entropy-driven drug transfer processes from more polar to less polar solvent systems. Moreover, by using the inverse Kirkwood–Buff integrals it is observed that salicylic acid is preferentially solvated by water molecules in water-rich mixtures but preferentially solvated by ethanol molecules in those mixtures of 0.24 < x1 < 1.00.


Introduction
Solubility is one of the most critical parameters in reaching the desired drug concentration to achieve the required pharmacological response [1,2]. In this regard, because solubility is very important in pharmaceutical product design, formulation studies and its future developments, it must be addressed from the early phases of product discovery to avoid the production of candidates with insufficient solubility [3,4]. In addition, the solubility information also plays a crucial role in the design Extended author information available on the last page of the article of separation and purification processes such as crystallization, precipitation, and supercritical fluid extraction in the pharmaceutical industries [1,2,5]. Moreover, the dependence of solubility on temperature allows to carry out the relevant thermodynamic analysis to deeply insight into the molecular mechanisms involved in the dissolution process [6,7].
Salicylic acid (IUPAC name: 2-Hydroxybenzoic acid, molecular structure shown in Fig. 1), a phenolic acid compound, is a metabolite of salicin, one of the oldest pain relievers derived from willow bark [8,9]. At the same time, it is a precursor of a well-known drug, aspirin, and is widely used as an intermediate for production of many industrial compounds [8,[10][11][12][13][14]. In this sense, salicylic acid is one of the active ingredients of cosmetic products. In addition to being used as an antipyretic, antiinflammatory and analgesic, it is an antiseptic and food preservative in toothpaste [8][9][10][11]15]. Salicylic acid is also known as a potent plant hormone as it affects the growth and development of plants by generating a wide range of metabolic and physiological responses [5,10,11,16].
Based on the above reasons, the main purposes of this study are to, (1) determine the effect of the mixtures composition and temperature on the equilibrium solubility of salicylic acid in binary solvent mixtures of {ethanol (1) + water (2)} within the temperature range T = (293.15 to 323.15) K, (2) correlation the experimental solubility data using the selected mathematical models, (3) calculation of the apparent thermodynamic properties of salicylic acid dissolution in the binary solvent systems and (4) evaluate the preferential solvation parameters of salicylic acid by ethanol in solvent mixtures by means of the inverse Kirkwood-Buff integrals.

Materials
Raw salicylic acid was of analytical reagent grade and provided by Sigma-Aldrich Chemical (Steinheim, Germany). Acetonitrile, ethanol and formic acid (more than 121 Page 4 of 27 99.5 % pure) were bought at Merck Chemical (Istanbul, Turkey). Ultra-pure water (18.2 MΩ·cm) was obtained from the MilliPore Milli − Q − Gradient water purification system (Billerica, MA, USA). Detailed descriptions of all the chemicals are given in Table 1.

Solubility Determinations
All {ethanol (1) + water (2)} solvent mixtures were prepared by mass using a Kern ABJ 220-4NM analytical balance (Germany) with sensitivity ± 0.1 mg, in quantities of 50.00 g. The mass fractions of ethanol of the nine mixtures prepared, varied by 0.10 from w 1 = 0.10 to w 1 = 0.90.
Equilibrium solubility of salicylic acid against mass fraction of ethanol in {ethanol (1) + water (2)} binary mixtures in the temperature range of T = (293.15 to 323.15) K as a function of temperature at atmospheric pressure was measured by the analytical shake-flask method [28] and defined as follows. In this method, excess salicylic acid was placed in the sealed flask containing known mass ratios of pure solvents and binary solvent mixtures. Then, the suspensions were allowed to equilibrate in a constant-temperature bath (± 0.1 K) with shaking for 18 h. When the equilibrium was reached, all the saturated mixtures were centrifuged and an aliquot of the supernatant solution rapidly was diluted with the water − acetonitrile mixture (50:50 % v/v) for quantitative determination by chromatographic analysis. The solubility samples were then analyzed by HPLC and described following.
All HPLC analyses were performed in triplicate using an Agilent 1200 HPLC system (Santa Clara, CA, USA) coupled with a reversed-phase column Zorbax SB C18 (4.6 × 150 mm, 5 μm, Santa Clara, CA, USA). The salicylic acid samples were eluted using the mixture of acetonitrile:water:formic acid (20:79:1 % v/v/v) as mobile phase at the flow rate of 0.7 mL·min -1 . The volume of the injected sample was 10 μL and it was detected at 240 nm with a diode array detector. Salicylic acid determination was performed at ambient temperature. Apart from solubility measurement, thermal and spectroscopic analyzes such as differential scanning calorimetry (DSC) and powder X-ray diffraction (XRD) were used to investigate the nature and crystal structure of raw and equilibrated salicylic acid samples before and after solubility determination.

Differential Scanning Calorimetry Analysis
Solid samples of salicylic acid obtained after solubility determination were analyzed by means of a differential scanning calorimetric analyzer to identify the nature of the raw and equilibrated forms. DSC thermograms of samples was obtained using a Mettler Toledo STARe System DSC 3 series (Ohio, ABD). The device was calibrated by Indium standard to determine the accuracy of obtaining melting temperatures and heats of fusion. Solid samples of 4.0 − 6.0 mg were placed in a crimped sealed aluminum crucible. Then, the samples were heated under a dynamic nitrogen atmosphere (40 mL·min −1 ) over the temperature range of t = (25 to 440) °C at a heating rate of 10 °C·min −1 .

X-Ray Diffraction Analysis
To determine the crystal form of salicylic acid, both raw and after equilibration in neat water, the mixture w 1 = 0.50, and neat ethanol, the X-ray powder diffraction analyses were conducted. The powder XRD patterns of the salicylic acid were recorded on Rigaku Smart Lab system (Tokyo, Japan) using CuKα radiation (1.5418 Å). The samples were scanned at 2θ° from ~ 10° to 90°.

Mole Fraction Solubility of Salicylic Acid
Mole fraction equilibrium solubility values of salicylic acid in neat solvents and nine {ethanol (1) + water (2)} mixtures at seven temperatures from T = (293.15 to 323.15) K and atmospheric pressure of 90 kPa are shown in Table 2 and depicted in Fig. 2. Minimum and maximum salicylic acid solubilities are, respectively, observed, in neat water and neat ethanol, at all temperatures studied. Salicylic acid solubility increases with temperature-arising, which imply endothermic dissolution processes. As indicated above, the solubility of this drug has been earlier reported in the literature and Fig. 3 allows the comparison at T = 298.15 K. It is noteworthy that our solubility values are in very good agreement with those reported by Jouyban et al. [23], although some differences are observed regarding those reported by Matsuda et al. [24], Peña et al. [25], Seidell [29] and Halford [30], in particular in mixtures of intermediate composition, however, these differences are lower than 3.0 % in almost all cases.
Moreover, Fig. 4 depicts the salicylic acid solubility as function of the Hildebrand solubility parameters of the {ethanol (1) + water (2)} mixtures (δ 1+2 ). Hildebrand solubility parameter of solvent mixtures is a polarity index widely used in pharmaceutical sciences. δ 1+2 values were calculated from the corresponding δ values of the pure solvents, i.e. δ 1 = 26.5 MPa 1/2 for ethanol and δ 2 = 47.8 MPa 1/2 for water [31,32] and mixtures compositions. Volume fractions (f i ) were considered by assuming additive behavior as described in Eq. 1 [33,34]: Accordingly, organic compounds reach maximum solubilities in solvent systems exhibiting the same or similar Hildebrand solubility parameters [35,36]. Thus, the δ 3 value of salicylic acid would be lower than δ 1 (neat ethanol δ value, 26.5 MPa 1/2 ) at T = 298.15 K, where maximum solubilities are observed at all temperatures. Nevertheless, the calculated Fedors δ 3 value of salicylic acid is 31.3 MPa 1/2 as shown in Table 3 [37], which is higher than the ethanol δ value. This result demonstrates that some other solvent and solutes properties apart of polarity must be involved in drug solubility and dissolution processes.
Otherwise, Fig. 5 allows the comparison of equilibrium solubilities of salicylic acid (2-hydroxybenzoic acid) regarding benzoic acid at T = 298.15 K [38]. As  observed, benzoic acid is more soluble than salicylic acid in the majority of the cases.

The Cosolvency Models Applied to Salicylic Acid Solubility
Among various cosolvency models used to calculate drug solubility in mixed solvents at an ambient and/or various temperatures [39,40], the log-linear model of Yalkowsky is the simplest model [41], which requires only two experimental determinations to predict the solubility at other solvent compositions. The model calculates the solubility at isothermal condition and is: where N is the number of experimental data points.
As mentioned above, Eq. 2 is capable of estimating drug solubility using solubility data in the mono-solvents. It could be combined with the van't Hoff equation as:  [29]; Δ: Jouyban et al. [23]; □: Matsuda et al. [24]; ◊: Peña et al. [25]; × : Halford [30]. Lines are just a visual guide  As an extension to the above mentioned models, the Jouyban-Acree model is presented by adding the two-body and three-body interaction terms and is the most accurate model to describe a drug solubility in mixed solvents at various temperatures and expressed as [23]: where J i terms are the model constants computed using a no intercept least square analysis [40]. The generated solubility of salicylic acid in {ethanol (1) + water (2)} was fitted to Eq. 5 and the trained model is: The F value of Eq. 6 was 751, the correlation and the model constants were significant with p < 0.0005. Equation 6 is valid for calculating the solubility of salicylic acid in {ethanol (1) + water (2)} mixtures at various temperatures by employing the solubility data of salicylic acid in ethanol and water at each T of interest. The obtained MPD for the back-calculated solubility data of salicylic acid using Eq. 6 was 8.0 %.  Fig. 6 depicts the simulated values obtained by using Eq. 7 at all temperatures studied. As observed, deviations between experimental and calculated solubility values are not distributed regularly, being the higher deviations observed at highest and lowest temperatures. In practical applications of Eq. 7, one may train the model using a minimum number of seven experimental data points and then predict the rest of required data in any solvent composition and temperature of interest as has been shown in an earlier work [43]. When the model trained with the solubility data in ethanol and water at T = (293. 15

Solid Phases' Analyses
DSC thermograms of salicylic acid as original sample and after dissolving it in neat water, in the aqueous mixture of w 1 = 0.50, and in neat ethanol (Figs. S1 to S4 shown as supplementary material), exhibit two endothermic peaks corresponding to the melting and thermal degradation of salicylic acid, respectively. The X-ray diffraction spectra for salicylic acid without any treatment and after dissolving it in neat water, in the aqueous mixture of w 1 = 0.50, and in neat ethanol, are shown in Figs. S5-S8 (as supplementary material). As observed the positions of characteristic peaks are comparable in all samples. Thus, salicylic acid did not suffer crystal polymorphic transitions or solvates formation in these experiments.

Ideal Solubility and Activity Coefficients of Salicylic Acid in Mixed Solvents
Ideal solubility of salicylic acid ( x id 3 ) as a function of temperature was calculated by means of Eq. 9: Here, Δ fus H is the molar enthalpy of fusion of the pure salicylic acid (at the melting point: 23.05 kJ·mol -1 [25]), T fus is the absolute melting point (432.5 K), T is the absolute solution temperature, R is the gas constant (8.3145 J·mol -1 ·K -1 ), and ΔC p is the difference between the molar heat capacity of the salicylic acid crystalline form and the molar heat capacity of the hypothetical super-cooled salicylic acid liquid form at the respective dissolution temperature [44]. Owing the difficulty in ΔC p determination, it was considered as the same as the entropy of fusion (Δ fus S = Δ fus H/T fus , i.e. 53.29 J·mol -1 ·K -1 ). It is important to keep in mind that ideal solubility involves the tendency of solute molecules in its solid state to separate into the hypothetical liquid state without considering the solvent nature and thus, it depends only on the melting properties of solute. Table 2 shows that the ideal solubilities of salicylic acid are higher than the experimental solubilities at all the temperatures studied in solvent systems of 0.0 ≤ w 1 ≤ 0.70 but they are lower in ethanolrich mixtures and neat ethanol. Table 4 summarizes the salicylic acid activity coefficients (γ 3 ) in {ethanol (1) + water (2)} mixtures, which were calculated as the quotient: x id 3 x 3 from the experimental and ideal solubilities summarized in Table 2  Here subscript s stands for the solvent system (which corresponds to the neat solvents or aqueous-ethanol binary mixtures), e ss, e 33 and e s3 denote the solvent-solvent, solute-solute and solvent-solute interaction energies, respectively. However, it is necessary to keep in mind that in multicomponent systems like ethanol-watersalicylic acid, some water-cosolvent interactions are also present. These additional interactions could play a significant role in the magnitudes of dissolution and equilibrium solubility of drugs. V 3 is the molar volume of the super-cooled liquid salicylic acid and φ s is the volume fraction of the solvent system in the saturated solutions. In the case of low x 3 values, the term (V 3 φ s 2 /RT) may be considered almost constant regardless the composition of the solvent system. Thus, the γ 3 values would depend mainly on e ss , e 33 and e s3 [45]. Here, e ss and e 33 are unfavorable for salicylic acid dissolution and equilibrium solubility but e s3 favors the respective drug dissolution processes. The contribution of e 33 towards the equilibrium solubility of salicylic acid could be considered as almost constant in all the solvent systems studied.
Thus, a qualitative analysis could be made based on the e ss, e 33 and e s3 energetic terms described in Eq. 10. Thus, e ss is highest in neat water (δ = 47.8 MPa 1/2 ) and lowest in neat ethanol (δ = 26.5 MPa 1/2 ) [31,32]. Neat water and water-rich mixtures, exhibiting γ 3 values higher than 300, would imply high e ss and low e s3 values. Otherwise, in ethanol-rich mixtures, exhibiting γ 3 values lower than 1.00, the e ss values are relatively low and the e s3 values would be comparatively high.

Apparent Thermodynamic Functions of Dissolution
All apparent thermodynamic quantities of dissolution of salicylic acid in {ethanol (1) + water (2)} mixtures were estimated at T = 298.15 K. Thus, the apparent standard enthalpy changes of dissolution (∆ soln H°) were obtained by means of the modified van't Hoff equation, as [46]: The apparent standard Gibbs energy changes for the dissolution processes (∆ soln G°) were calculated by means of Eq. 12: Here, the intercepts used are those obtained in the regressions of ln x 3 vs. (1/T -1/298.15). Figure 7 depicts the solubility van't Hoff plots for the neat solvents water and ethanol and nine {ethanol (1) + water (2)} mixtures. Linear regressions with determination coefficients higher than 0.993 were obtained in all cases [47][48][49]. Standard apparent entropic changes for dissolution process (∆ soln S°) were obtained from the respective ∆ soln H° and ∆ soln G° values using Eq. 13 [46]: Table 5 summarizes the standard apparent molar thermodynamic functions for dissolution of salicylic acid (3) in all the {ethanol (1) + water (2)} solvent systems at T = 298.15 K.
Apparent standard Gibbs energies, enthalpies and entropies of salicylic acid dissolution are positive in all cases, which implies endothermic and entropy-driven dissolution processes. Moreover, Δ soln G° values decrease continuously from neat water to neat ethanol indicating more affinity of salicylic acid by semipolar solvent media. As observed, initially the Δ soln H° values decrease from neat water to the mixture of w 1 = 0.30 and later they increase with the ethanol proportion to reach a new (13)  maximum value in the mixture of w 1 = 0.50. After, they decrease continuously to reach the minimum value in neat ethanol. On the other hand, Δ soln S° values increase from neat water to the mixture of w 1 = 0.60 and later they decrease continuously with the ethanol proportion to reach the minimum value in neat ethanol. Otherwise, the relative contributions by enthalpy (ζ H ) and entropy (ζ TS ) toward the salicylic acid dissolution processes were calculated by means of the following equations [50]: As shown in Table 5 the main contributor to the positive standard molar Gibbs energies of dissolution of salicylic acid in these solvent systems was the positive enthalpy, with values higher than 0.572, which demonstrates the energetic predominance toward all these dissolution processes. It is noteworthy that enthalpy and entropy contributions in the mixture of w 1 = 0.50 and neat ethanol are the same as those observed for ideal dissolution process.

Apparent Thermodynamic Quantities of Mixing
Global dissolution processes of salicylic acid in {ethanol (1) + water (2)} mixtures may be represented by means of the following hypothetical process: Here, the hypothetical stages are as follows: i) the heating and fusion of salicylic acid at T fus = 432.5 K, ii) the cooling of the liquid salicylic acid to the considered temperature (T = 298.15 K), and iii) the subsequent mixing of both the hypothetical super-cooled liquid salicylic acid and the respective solvent system at this temperature [51]. This allowed the calculation of the individual thermodynamic contributions by fusion and mixing toward the overall dissolution process by means of the following equations: where Δ fus H T 298 and Δ fus S T 298 represent the thermodynamic quantities of salicylic acid fusion and its cooling at T = 298.15 K, which in turn, are calculated by means of the following equations [52]: Table 6 summarizes the apparent thermodynamic quantities of mixing of hypothetical super-cooled liquid salicylic acid with all the aqueous-ethanol mixtures and the neat solvents at T = 298.15 K. Gibbs energies of mixing are positive from neat water to the mixture of w 1 = 0.80 because the experimental solubilities of salicylic acid are lower than the ideal solubilities, as indicated above. As observed, the contributions by the mixing thermodynamic quantities, Δ mix H° and Δ mix S° values are positive in neat water and in all the {ethanol (1) + water (2)} mixtures but negative in neat ethanol. In this way, entropy-driving is observed for salicylic acid mixing processes from neat water to the mixture of w 1 = 0.90 (Δ mix H° > 0, Δ mix S° > 0) but enthalpy-driving is observed for mixing in neat ethanol (Δ mix H° < 0, Δ mix S° < 0). Moreover, to compare the relative contributions by enthalpy (ζ H ) and entropy (ζ TS ) to the mixing processes, two equations analogous to Eqs. 14 and 15, were employed. Thus, in almost all cases the main contributor to Gibbs energy of mixing is the enthalpy, whereas, in the mixture of w 1 = 0.90 the mixing entropy is the dominant thermodynamic function.

Enthalpy-Entropy Compensation Analysis
Extra-thermodynamic studies including enthalpy-entropy compensation analysis provide a powerful tool to inquiry the main molecular mechanisms involved in different chemical processes involving organic compounds [53,54]. Non-enthalpy-entropy compensation effects have been reported in the dissolution processes of several drugs in {ethanol (1) + water (2)} mixtures as summarized in the literature [55]. Reported studies were performed to identify the main mechanisms involved in the dissolving cosolvent action of ethanol. Normally, weighted plots of Δ soln H° vs. Δ soln G° have been used for performing such an analysis [56][57][58]. Figure 8 shows that salicylic acid exhibits a non-linear Δ soln H° vs. Δ soln G° trend in {ethanol (1) + water (2)} mixtures with variable positive slopes from neat water to the mixture of w 1 = 0.30 and from the mixture of w 1 = 0.50 to neat ethanol but negative slopes in the interval of 0.30 ≤ w 1 ≤ 0.50. Accordingly, in the first cases the transfer of salicylic from more polar to less polar solvent systems is enthalpy-driven, whereas, in the mixtures of 0.30 ≤ w 1 ≤ 0.50 it is entropy-driven.

Preferential Solvation of Salicylic Acid
The preferential solvation parameters of salicylic acid (compound 3) by ethanol (compound 1) in the {ethanol (1) + water (2)} mixtures (δx 1,3 ) are defined as Eq. 20: where x L 1,3 is the local mole fraction of ethanol in the molecular environment of salicylic acid and x 1 is the bulk mole fraction of ethanol in the initial binary solvent mixture free of salicylic acid. If δx 1,3 is positive salicylic acid is preferentially solvated by ethanol. On the contrary, salicylic acid is preferentially solvated by water if this parameter is negative. The values of δx 1,3 were obtained from the inverse Kirkwood-Buff integrals (IKBI) as described earlier [59][60][61], based on the following expressions: with, Here, κ T is the isothermal compressibility of the {ethanol (1) + water (2)} mixtures. V 1 and V 2 are the partial molar volumes of ethanol and water in the mixtures, and V 3 is the partial molar volume of salicylic acid in the solvent mixtures. The function D, described by Eq. 25, corresponds to the first derivative of the standard molar Gibbs energies of transfer of salicylic acid from neat water to {ethanol (1) + water (2)} mixtures regarding the mole fraction of ethanol in the solvent mixtures free of drug. The function Q, described by Eq. 26, involves the second derivative of the excess molar Gibbs energy of mixing of both solvents ( G Exc 1+2 ) regarding the mole fraction of water in the mixtures [59][60][61]. V cor is the correlation volume and r 3 is the molecular radius of salicylic acid, which in turn is approximately calculated by means of Eq. 27, where N Av is the Avogadro's number.
Definitive V cor values require iteration because they depend closely on the local mole fractions of ethanol and water around the salicylic acid molecules. This iteration process was performed by replacing δx 1,3 and V cor in the Eqs. 20, 21 and 24 to recalculate x L 1,3 until obtaining a non-variant value of V cor . Figure 9 shows the Gibbs energies of transfer of salicylic acid from neat water to {ethanol (1) + water (2)} mixtures at T = 298.15 K. These values were calculated from the mole fraction solubility values reported in Table 2 using Eq. 28:  Table 7 were calculated from the first derivative of Eq. 29 solved in successive steps of x 1 = 0.05. For {ethanol (1) + water (2)} mixtures the Q, RTκ T , V 1 and V 2 values at T = 298.15 K were taken from the literature [62,63]. V 3 was considered as the one calculated by the Fedors method, i.e. 90.9 cm 3 ·mol -1 (Table 3). Table 7 shows that G 1,3 and G 2,3 values are negative in all the solvent systems. Salicylic acid r 3 value was calculated as 0.330 nm. As indicated above,  V cor values reported in Table 7 were obtained after three iterations. Moreover, Table 7 also summarizes the preferential solvation parameters of salicylic acid by ethanol molecules (δx 1,3 ) at T = 298.15 K. Figure 10 shows non-linear variation of salicylic acid δx 1,3 values regarding the ethanol mole fraction in the mixtures free of drug. Initially, the addition of ethanol to water makes negative the δx 1,3 values of salicylic acid in the composition interval 0.00 < x 1 < 0.24. Maximum negative δx 1,3 value is obtained in the mixture x 1 = 0.15 (δx 1,3 = -5.21 × 10 -2 ), which is higher than |1.0 × 10 -2 |. Hence, these results could be considered as a consequence of real preferential solvation effects of this drug by water molecules, rather than a consequence of uncertainties propagation in performed IKBI calculations [64,65].
In the mixtures composition interval of 0.24 < x 1 < 1.00 the local mole fractions of ethanol around salicylic acid molecules are higher than those in the bulk mixtures free of drug. Maximum positive δx 1,3 value is obtained in the mixture x 1 = 0.45 (δx 1,3 = 6.20 × 10 -2 ), which is also higher than |1.0 × 10 -2 |. Thus, these results could be considered as a consequence of preferential solvation effects of this drug by ethanol molecules [64,65]. In solution salicylic acid mainly acts as a Lewis acid due to the hydrogen atoms in its -COOH and -OH groups ( Fig. 1) in order to establish hydrogen bonds with proton-acceptor functional groups in the solvents (free electron pairs in the oxygen atoms of the -OH groups). In addition, this drug could act as a Lewis base due to free electron pairs in oxygen atoms of carboxyl and hydroxyl and groups to interact with acidic hydrogen atoms in both solvents. Thus, it is conjecturable that in mixtures of 0.24 < x 1 < 1.00 salicylic acid is acting as a Lewis acid with ethanol molecules because this cosolvent is more basic than water, as described by the Kamlet-Taft hydrogen bond acceptor parameters, namely β = 0.75 for ethanol and 0.47 for water [66].
Ultimately, Fig. 10 allows the comparison of preferential solvation results of salicylic acid regarding those of structurally related benzoic acid [38]. As observed, maximum negative δx 1,3 value is higher for salicylic acid, which could be a consequence of the intramolecular hydrogen bonding between carboxyl and hydroxyl groups, although it is observed in different mixture compositions. However, maximum positive δx 1,3 value is higher for benzoic acid although it is also observed in different mixture composition.

Conclusions
Solubility and dissolution physicochemical properties of salicylic acid in {ethanol (1) + water (2)} mixtures depend strongly on the cosolvent mixtures composition. Experimental solubility values of salicylic acid were adequately correlated with the classical Jouyban-Acree model and other well-known correlation models. Apparent thermodynamic quantities of dissolution and mixing were calculated based on van't Hoff and Gibbs equations. Non-linear enthalpy-entropy compensation was found for salicylic acid in these mixtures indicating different transfer mechanisms regarding the solvent mixtures composition. Moreover, salicylic acid is preferentially solvated by water in water-rich mixtures but preferentially solvated by ethanol in mixtures of 0.24 < x 1 < 1.00 at T = 298.15 K. Finally, the thermodynamic results presented in this communication could be useful in optimizing different physical and chemical processes involving salicylic acid.