Dissolution Thermodynamics of the Solubility of Sulfamethazine in (Acetonitrile + 1-Propanol) Mixtures

Abstract

Background: Solubility is one of the most important parameters in the research and development processes of the pharmaceutical industry. In this context, cosolubility is one of the most used strategies to improve the solubility of poorly soluble drugs, besides allowing to identify some factors involved in the dissolution process. The aim of this research is to evaluate the solubility of sulfamethazine in acetotinitrile + 1-propanol cosolvent mixtures at 9 temperatures (278.15, 283.15, 288.15, 293.15, 298.15, 303.15, 308.15, 313.15, and 318.15 K); a drug used in human and veterinary therapy and two solvents of great chemical-pharmaceutical interest. Methods: The determination was carried out by the shaking flask method and the drug was quantified by UV/Vis spectrophotometry. Results: The solubility of sulfamethazine increases from pure 1-propanol (solvent in which it reaches its lowest solubility at 278.15 K) to pure acetonitrile (solvent in which it reaches its maximum solubility at 318.15 K), behaving in a logarithmic-linear fashion. Conclusions: The increase in solubility is related to the acid/base character of the cosolvent mixtures and not to the solubility parameter of the mixtures. The dissolution process is endothermic and favored by the solution entropy, and also shows a strong entropic compensation.

Keywords: 1-propanol; acetonitrile; cosolvency; solubility; solution thermodynamics; sulfamethazine.

Figure 1

Molecular structure of the 4-amino-N-(4,6-dimethylpyrimidin-2-yl)benzenesulfonamide [9].

Figure 3

Mole fraction solubility of SD (3, •) [24], SMR (3, ∘) [25], and SMT (3, ▲) [this work] in {MeCN (1) + 1-PrOH (2)} mixtures at T/K = 298.15.

Figure 4

Mole fraction solubility of SMT (3) as a function of mixtures’ composition (top) and as a function of the Hildebrand solubility parameter of the mixtures (bottom) in some {MeCN (1) + alcohol (2)} mixtures at T/K = 298.15. ∘: {MeCN (1) + MeOH (1)} [22]; •: {MeCN (1) + 1-PrOH (2)} [this work].

Figure 2

(A) Logarithmic mole fraction solubility of SMT ($x_3$) as function of the mass fraction of MeCN (1) in {MeCN (1) + 1-PrOH (2)} mixtures at different temperatures (in Kelvin). •: 278.15; ∘: 283.15; ▲: 288.15; △: 293.15; ■: 298.15; □: 303.15; ⧫: 308.15; ◊: 313.15; ★: 318.15. (B) Logarithmic mole fraction solubility of SMT ($x_3$) in different {MeCN (1) + 1-PrOH (2)} mixtures compositions ($w_1)$ as function of temperature. □: 0.00; : 0.05; ■: 0.10; : 0.15; △: 0.20; : 0.25; ▲: 0.30; : 0.55; ◊: 0.20; : 0.25; ⧫: 0.30; : 0.55; ∘: 0.60; : 0.65; •: 0.70; : 0.75; ★: 0.80; : 0.85; ⊗: 0.90; : 0.95; ⊕: 1.00.

Figure 5

DSC thermograms of SMT as original sample and phases in equilibrium with saturated solutions. With $T_{\mathrm{fus}}$/K and ${\Delta }_{\mathrm{fus}}{H}^{\circ }$/kJ·mol⁻¹ values and standard uncertainties: Original untreated sample ($T_{\mathrm{fus}}$/K = 468.5 ± 0.5, ${\Delta }_{\mathrm{fus}}{H}^{\circ }$/kJ·mol⁻¹ = 33.57 ± 0.3), 1-PrOH-saturated sample ($T_{\mathrm{fus}}$/K = 468.5 ± 0.5, ${\Delta }_{\mathrm{fus}}{H}^{\circ }$/kJ·mol⁻¹ = 33.8 ± 0.3), MeCN-saturated sample ($T_{\mathrm{fus}}$/K = 468.7 ± 0.5, ${\Delta }_{\mathrm{fus}}{H}^{\circ }$/kJ·mol⁻¹ = 34.27 ± 0.3), and {MeCN (1) + 1-PrOH (2)} ($w_1$ = 0.5) saturated sample ($T_{\mathrm{fus}}$/K = 468.3 ± 0.5, ${\Delta }_{\mathrm{fus}}{H}^{\circ }$/kJ·mol⁻¹ = 34.1 ± 0.3).

Figure 6

Van ’t Hoff plot of the mole fraction solubility of SMT ($x_3$) in different MeCN (1) + 1-PrOH (2) mixtures compositions. ★: $w_1$ = 0.00 (neat 1-PrOH); : 0.05; •: 0.10; : 0.15; ∘: 0.20 : 0.25; ▲: 0.30; : 0.35; △: 0.40; : 0.45; □: 0.50; : 0.55; ■: 0.60; : 0.65; ◊: 0.70; : 0.75; ⧫: 0.80; : 0.85; ⊗: 0.90; : 0.95; ⊕: 1.00 (neat MeCN).

Figure 7

Enthalpy–entropy compensation plots for the solubility of SMT (3) in {MeCN (1) + 1-PrOH (2)} mixtures at $T_{\mathrm{hm}}$/K = 297.6 K. The points represent the mass fraction of MeCN (1) in the {MeCN (1) + 1-PrOH (2)} mixtures in the absence of SMT (3).