solubilities

Solubilities are the extent to which solutes dissolve in a solvent to form a homogeneous solution at a defined temperature and pressure. They are often expressed as molar solubility (moles per liter) or as the maximum amount of solute that can dissolve in one liter of solvent. The solubility of most solids in water increases with temperature, while the solubility of gases generally decreases as temperature rises. Pressure mainly affects gas solubility.

Solubilities depend on the chemical compatibility of solute and solvent, their polarity, and the presence of other substances in solution. The common ion effect lowers solubility when ions from the solvent share a component with the dissolved solid. Acid-base reactions, complex formation, and pH changes can alter apparent solubility by creating more soluble species. Solubility can also vary with crystal form and particle size.

For sparingly soluble salts, solubility is described by the solubility product constant, Ksp, which expresses the equilibrium between the solid and its dissolved ions: Ksp = [A^+]^m [B]^n. A larger Ksp corresponds to greater solubility under the same conditions. When a solution already contains one of the ions, the solubility often decreases further due to the common ion effect.

Practical implications include precipitation and crystallization during solution processing, as well as the dissolution rate and bioavailability of drugs. Supersaturated solutions can form under certain conditions but are typically unstable and may crystallize.

Solubilities can be increased by heating, converting the solute into a more soluble form (such as a complex), using a cosolvent or surfactant, adjusting pH, or reducing particle size to enhance dissolution rate even if equilibrium solubility is unchanged.