Thermodynamics Research Center / ThermoML | Journal of Chemical Thermodynamics

ABSTRACT

Experimental measurement of drug quinocetone solubility in water-co-solvent mixtures of ethanol, isopropanol, dimethyl sulfoxide (DMSO) and N,N-dimethylformamide (DMF) were determined by the static equilibrium method at temperature range of (283.15 328.15) K under ambient condition (101.2 kPa). The solid phase equilibrated with liquor was examined by X-ray power diffraction, resulting in no polymorphic transformation, solvate formation or crystal transition compared to the raw material. Quantitative knowledge about the solvent effect on the solubility was obtained with the help of linear solvation energy relationships in the terms of dipolarity-polarizability, hydrogen-bond donor and acceptor ability, and cavity term of mixtures. The preferential solvation was investigated from their thermodynamic solution properties by means of the inverse Kirkwood-Buff integrals. The preferential solvation parameters (dx1,3) for ethanol, isopropanol, DMSO or DMF were negative in the four co-solvent mixtures with water-rich compositions, which indicated that quinocetone was preferentially solvated by water. Temperature has a little effect on the preferential solvation magnitudes for the ethanol and isopropanol mixtures. The higher solvation by water could be explained in terms of the higher acidic behavior of the solvents interacting with the Lewis basic groups of the quinocetone. In addition, thermodynamic co-solvency models including the Jouyban-Acree model, van t Hoff-Jouyban-Acree model and Modified Apelblat-Jouyban-Acree model were employed to describe the solubility of this drug obtaining average relative deviations lower than 2.14% and root-mean-square deviations lower than 1.75 *10-4 for correlative studies. The standard dissolution enthalpies of quinocetone in the solvent mixtures were positive, demonstrating that the dissolution process of quinocetone was endothermic.