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By using the shake-flask method, the acetoguanamine solubilities in water, acetonitrile, methanol, n-butanol, isopropanol, isobutanol, ethanol, n-propanol, N,N-dimethylformamide (DMF), 1,4-dioxane, ethyl acetate, and ethylene glycol (EG) were acquired over the temperatures from 283.15 to 323.15 K at ambient pressure (p= 101.2 kPa). The mole fraction of acetoguanamine in equilibrated liquor increased as the temperature increased and followed the order in different solvents: DMF > EG > ethyl acetate > 1,4-dioxane > n-butanol > isobutanol > n-propanol > isopropanol > ethanol > methanol > water > acetonitrile. The acquired solubility was correlated via the Apelblat equation. The largest root-mean-square deviation value and largest relative average deviation were, respectively, 28.12 × 10–6and 2.38 × 10–2.

The solubilities of d-tryptophan in neat n-propanol, ethyl acetate, ethylene glycol (EG), n-butanol, acetonitrile, 1-octanol, and dimethyl sulfoxide and of l-tyrosine in ethyl acetate, EG, isopropanol, acetonitrile, n-butanol, and N,N-dimethylformamide (DMF) were determined via a shake-flask method at temperatures from 283.15 to 328.15 K under about 101.2 kPa. At the fixed temperature, the mole fraction solubility of d-tryptophan was the largest in EG and the lowest in 1-octanol. However, for l-tyrosine, the maximum solubility was observed in the ethyl acetate and minimum in n-butanol. X-ray powder diffraction analysis was carried out to analyze the solids equilibrated with the corresponding liquor, indicating that no polymorphic transformation or solvate formation took place in the studied systems. The solubility correlation was performed via the Apelblat equation. The obtained maximum root-mean-square deviation and relative average deviation values were 13.65 × 10–5and 4.32%, respectively. In addition, the solvent–solvent and solute–solvent interactions were investigated on the basis of the approach of linear solvation energy relationship.

The o-nitrophenylacetonitrile solubilities in some pure solvents such as n-propanol, methanol, toluene, ethanol, isopropanol, acetonitrile, acetone, cyclohexane, ethyl acetate, n-butanol, 2-butanone, isobutanol, acetic acid, water, 1,4-dioxane, and ethylbenzene were determined via a shake-flask method at temperatures from 278.15 to 333.15 K under about 101.2 kPa. Good dissolution ability was found for o-nitrophenylacetonitrile in the organic solvents. The order of solubility magnitude from high to low, except for the solvents of ethyl acetate and 1,4-dioxane, was as follows: acetone > 2-butanone > acetonitrile > toluene > (acetic acid, ethylbenzene) > methanol > ethanol > n-propanol > isopropanol > n-butanol > isobutanol > cyclohexane > water. X-ray powder diffraction analysis was carried out to analyze the solids equilibrated with the corresponding liquor, indicating that no polymorphic transformation or solvate formation took place in the 16 solvents. The correlation of the determined solubility was performed via the Wilson model, Apelblat equation, nonrandom two-liquid model, and λhequation. Maximum root-mean-square deviation and relative average deviation were 49.66 × 10–4and 4.97%, respectively. The Apelblat equation presented the best correlating results among these models. In addition, the solvent–solvent and solute–solvent interactions were investigated on the basis of the approach of the linear solvation energy relationship.

The determination of 2-amino-6-chloropurine solubility in water and 11 organic solvents including dimethyl sulfoxide (DMSO), N,N-dimethylformamide (DMF), 2-butanone, ethylene glycol (EG), n-butanol, n-propanol, isobutanol, isopropanol, 1,4-dioxane, ethanol, and methanol was performed by using the shake flask technique. The experiments were conducted covering the temperatures from 278.15 to 333.15 K at ambient pressure p= 101.2 kPa. The values of solubility for 2-amino-6-chloropurine in selected solvents increased as the temperature rised. The maximum solubility value of 2-amino-6-chloropurine was recorded in DMSO, and the minimum, in water. They ranked as DMSO > DMF > 2-butanone > EG > n-butanol > n-propanol > isobutanol > isopropanol >1,4-dioxane > ethanol > methanol > water. The solubility data obtained were mathematically described through the Apelblat equation. The largest percentage of average relative deviation was 1.75 × 10–2, and the maximum value of root-mean-square deviation, 1.14 × 10–4. Moreover, using the Kamlet and Taft linear solvation energy relationship analysis of the solvent effect, the type and degree of solvent, solute, and solvent–solvent interactions were recognized.

The saturated thermodynamic solubilities of micoflavin in four mixtures of dimethyl sulfoxide (DMSO, 1) + water (2), isopropanol (1) + water (2), propylene glycol (PG, 1) + water (2), and ethanol (1) + water (2) are reported in this contribution. All experiments were carried out via the saturation shake-flask technique at temperatures from 278.15/293.15 to 318.15 K. The solubility data were maximum in neat DMSO (isopropanol, PG, or ethanol) for the four mixtures studied. The micoflavin solubility was mathematically well described through the Jouyban–Acree model, attaining root-mean-square deviation values lower than 3.68 × 10–4and relative average deviation values lower than 6.27%. The local mole fractions of DMSO (isopropanol, ethanol, PG and water) around micoflavin were quantitatively studied via the method of inverse Kirkwood–Buff integrals on the basis of the solubility data determined. The solvation of micoflavin was preferential by water for these mixtures in the water-rich compositions, whereas in the cosolvent-rich and intermediate compositions, the solvation of micoflavin was preferential by DMSO (isopropanol, PG, or ethanol) in the DMSO (isopropanol, PG, or ethanol) + water mixtures. In addition, the cosolvency effect was reorganized by considering the Kamlet and Taft linear solvation energy relationships, and the relative importance of the solvent–solvent and solvent–solute interactions was distinguished. Results indicated that the micoflavin solubility variation was governed by the cavity term in the DMSO + water solution and the cavity and α terms in the aqueous solutions of PG, isopropanol, and ethanol.

The mutual solubilities for ternary systems of p-nitroacetophenone (1) + m-nitroacetophenone (2) + methanol (3), p-nitroacetophenone (1) + m-nitroacetophenone (2) + toluene (3), and p-nitroacetophenone (1) + m-nitroacetophenone (2) + ethyl acetate (3) were investigated through a steady method at 283.15, 298.15, and 313.15 K, respectively, under a local ambient pressure of 101.2 kPa. The isothermal phase diagrams of the three systems were constructed in terms of the obtained mutual solubilities. At a fixed temperature, there were two neat solids formed, namely neat p-nitroacetophenone and neat m-nitroacetophenone, which were recognized through Schreinemakers’ wet residue method. The crystallization regions of neat p-nitroacetophenone and neat m-nitroacetophenone increased with decreasing temperature. The mutual solubility was correlated through the Wilson and NRTL equations. The relative average deviations for p-nitroacetophenone + m-nitroacetophenone + methanol/toluene/ethyl acetate systems were, respectively, 4.44 × 10–2, 3.86 × 10–2, and 3.35 × 10–2for the NRTL equation and 4.33 × 10–2, 2.51 × 10–2, and 2.29 × 10–2for the Wilson equation. The acquired solubility data and the ternary phase diagrams for the three systems were very significant for designing and optimizing the purification process of p-nitroacetophenone and m-nitroacetophenone mixtures.

The stable solid–liquid phase equilibrium for ternary systems of terephthaldialdehyde + 4-hydroxybenzaldehyde + ethyl acetate and terephthaldialdehyde + 4-hydroxybenzaldehyde + acetone was determined using the isothermal saturation method at 293.15, 303.15, and 313.15 K under ambient pressure of 101.2 kPa. In terms of the obtained mutual solubility data, the isothermal phase diagrams of the two systems were constructed. At a given temperature, there were two pure solids formed, pure 4-hydroxybenzaldehyde and pure terephthaldialdehyde, which were identified by Schreinemakers’ wet residue method and X-ray powder diffraction analysis. The crystallization zones of pure 4-hydroxybenzaldehyde and pure terephthaldialdehyde increased with rising temperature. The crystalline zone of terephthaldialdehyde was larger than that of 4-hydroxybenzaldehyde at the same temperature. The mutual solubility data of solid–liquid phase equilibrium were correlated with nonrandom two-liquid (NRTL) and Wilson models. The calculated solubility data with NRTL model agreed well with the determined values. The mutual solubility data and ternary phase diagrams for the two systems were of great significance for optimizing the purification procedure of 4-hydroxybenzaldehyde and terephthaldialdehyde.

The solubility of trimetazidine hydrochloride in 12 neat solvents namely ethanol, methanol, isopropyl alcohol, n-propanol, acetone, toluene, 2-butanone, ethyl acetate, n-butanol, acetonitrile, N,N-dimethylformamide (DMF), and 1,4-dioxane and binary liquid mixtures methanol + (ethanol, DMF, or ethyl acetate) was measured using the isothermal method over the temperature range from 278.15 to 323.15 K under local atmospheric pressure of 101.2 kPa. The observed mole fraction solubility was the largest in methanol and smallest in toluene. The experimental data in monosolvents was correlated through the Apelblat equation; while in mixed solvents, by the van’t Hoff–Jouyban–Acree model, Jouyban–Acree model and Apelblat–Jouyban–Acree model. Consequently, for the selected monosolvents, the largest relative average and root-mean-square deviations were 0.79% and 0.227 × 10–4, respectively; and for the binary solvent mixtures, 3.00% and 2.22 × 10–4. Solvent–solvent and solute–solvent interactions were analyzed in the 12 net solvents with linear solvation energy relationship analysis. The solubility in monosolvents depends significantly upon the Hildebrand parameter and nonspecific dipolarity/polarizability interactions. The values of log x12deviated from the mean values of log x12in the monosolvents. Results were interpreted based on different modes of the solvation interactions.

The solubility of terephthaldialdehyde in N-methyl-2-pyrrolidone (NMP) at temperatures ranging from T= 293.15 to 318.15 K and the mutual solubility for ternary systems of terephthalic acid + terephthaldialdehyde + N,N-dimethylformamide (DMF) and terephthalic acid + terephthaldialdehyde + NMP at T= 298.15, 308.15, and 318.15 K were measured with the technique of isothermal saturation under 101.2 kPa. Six phase diagrams were built on the basis of the determined solubility. There were two solids formed in each ternary system, which were terephthalic acid and terephthaldialdehyde. Each diagram consisted of one cosaturated point, three crystalline areas, and two cosaturated curves. Furthermore, the nonrandom two-liquid mode and Wilson model were used in correlating and calculating the solubility of terephthaldialdehyde in NMP and the mutual solubility data for terephthalic acid + terephthaldialdehyde + DMF/NMP systems. The maximum root-mean-square deviation was 3.58 × 10–3, and maximum value of relative average deviation was 1.74%. The liquid–solid phase diagram and the thermodynamic analysis could provide the basis for obtaining terephthalic acid with high purity.

The thiamphenicol solubility in aqueous cosolvent solutions of ethanol (1), methanol (1), 1,4-dioxane (1), and N,N-dimethylformamide (DMF, 1) was measured via the isothermal dissolution equilibrium method at temperatures ranging from 278.15 to 318.15 K under local pressure (101.1 kPa). At fixed composition of ethanol (methanol, 1,4-dioxane, or DMF) and temperature, the solubility of thiamphenicol was larger in DMF + water mixtures than in the ethanol/methanol/1,4-dioxane mixtures. The local solvent proportions were acquired with the method of inverse Kirkwood–Buff integrals. The absolute value of these preferential solvation parameters were all lower than 1.0 × 10–2for ethanol (1) + water (2) and 1,4-dioxane (1) + water (2) solutions in water-rich compositions and for methanol (1) + water (2) solutions in whole compositions. In the former two cosolvent mixtures in intermediate compositions and cosolvent-rich regions, thiamphenicol was preferentially solvated by cosolvent. However, for the DMF (1) + water (2) solutions, water solvated preferentially thiamphenicol in water-rich compositions and by DMF in intermediate and DMF-rich compositions. This case by cosolvent might be illustrated based on higher basic behavior of water, which interacted with Lewis acidic groups of thiamphenicol. In addition, the solubility of thiamphenicol was described with the van’t Hoff–Jouyban–Acree, Jouyban–Acree, and Apelblat–Jouyban–Acree models. The obtained average relative deviations were no greater than 1.85%. Furthermore, the solvent effect treatment through the KAT-LSER model indicated that the solubility variation was significantly affected by the cavity term.