Effect of N-Methyl-pyrrolidone (NMP) on the Equilibrium Solubility of Meloxicam in Aqueous Media: Correlation, Dissolution Thermodynamics, and Preferential Solvation

Abstract

Meloxicam is an analgesic and anti-inflammatory drug widely prescribed in current therapeutics that exhibits very low solubility in water. Thus, this physicochemical property has been studied in N-methyl-pyrrolidone (NMP)-aqueous mixtures at several temperatures to expand the solubility database about pharmaceutical compounds in aqueous-mixed solvents. The flask-shake method and UV-vis spectrophotometry were used for meloxicam solubility determination as a function of temperature and mixture composition. Several cosolvency models, including the Jouyban-Acree model, were challenged for equilibrium solubility correlation and/or prediction. The van't Hoff and Gibbs equations were employed here to calculate the apparent standard thermodynamic quantities for the dissolution and mixing processes of this drug in these aqueous mixtures. Inverse Kirkwood-Buff integrals were employed to calculate the preferential solvation parameters of meloxicam by NMP in all mixtures. Meloxicam equilibrium solubility increased with increasing temperature, and maximal solubilities were observed in neat NMP at all temperatures. The mole fraction solubility of meloxicam increased from 1.137 × 10^-6 in neat water to 3.639 × 10^-3 in neat NMP at 298.15 K. The Jouyban-Acree model correlated the meloxicam solubility in these mixtures very well. Dissolution processes were endothermic and entropy-driven in all cases, except in neat water, where nonenthalpy- nor entropy-driven was observed. Apparent Gibbs energies of dissolution varied from 34.35 kJ·mol^-1 in pure water to 7.99 kJ·mol^-1 in pure NMP at a mean harmonic temperature of 303.0 K. A nonlinear enthalpy-entropy relationship was observed in the plot of dissolution enthalpy vs dissolution Gibbs energy. Meloxicam is preferentially hydrated in water-rich mixtures but preferentially solvated by NMP in the composition interval of 0.16 < x ₁ < 1.00.

Figure 1

Molecular structure of meloxicam.

Figure 2

Molecular structure of N-methyl-2-pyrrolidone.

Figure 3

Logarithmic mole fraction solubility of meloxicam (ln x₃) as function of the Hildebrand solubility parameter in some {cosolvent (1) + water (2)} mixtures at T = 298.15 K. ●: N-Methyl-pyrrolidone (1) + water (2); ○: dimethyl sulfoxide (1) + water (2); ◊: N,N-dimethylformamide (1) + water (2); and Δ: acetonitrile (1) + water (2).

Figure 4

X-ray diffraction spectra of meloxicam. From top to bottom: crystallized in water, crystallized in the {N-methyl-pyrrolidone (1) + water (2)} (x₁ = 0.50) mixture, crystallized in N-methyl-pyrrolidone, and the original untreated sample.

Figure 5

FTIR spectra of meloxicam. From top to bottom: original untreated sample, crystallized in N-methyl-pyrrolidone, crystallized in the {N-methyl-pyrrolidone (1) + water (2)} (x₁ = 0.50) mixture, and crystallized in water.

Figure 6

van’t Hoff plot of the solubility of meloxicam (3) in {N-methyl-pyrrolidone (1) + water (2)} solvent systems. ○: x₁ = 0.00 (neat water), Δ: x₁ = 0.10, □: x₁ = 0.20, ◊: x₁ = 0.30, ×: x₁ = 0.40, *: x₁ = 0.50, ●: x₁ = 0.60, ▲: x₁ = 0.70, ■: x₁ = 0.80, ◆: x₁ = 0.90, and +: x₁ = 0.10 (neat N-methyl-pyrrolidone).

Figure 7

Enthalpy–entropy compensation plot for the solubility of meloxicam (3) in {N-methyl-pyrrolidone (1) + water (2)} mixtures at T_hm = 303.0 K. The points represent the mole fraction of N-methyl-pyrrolidone (1) in the {N-methyl-pyrrolidone (1) + water (2)} mixtures in the absence of meloxicam (3).

Figure 8

Gibbs energy of transfer of meloxicam (3) from neat water (2) to {N-methyl-pyrrolidone (1) + water (2)} mixtures at T = 298.15 K.

Figure 9

Preferential solvation parameters of meloxicam (3) in some {cosolvent (1) + water (2)} mixtures at T = 298.15 K. ●: N-methyl-pyrrolidone (1) + water (2); ○: dimethyl sulfoxide (1) + water (2); and ◊: N,N-dimethylformamide (1) + water (2).