Solubility is affected by the identity and properties of the solute, the identity and properties of the solvent, temperature, pressure, pH, the presence of other dissolved substances and the physical form of the solute. There is no single factor that universally increases or decreases solubility because different substances respond differently to the same environmental change.
At the molecular level, solubility reflects the competition between interactions within the pure solute or solvent and interactions formed between solute and solvent, together with entropy and the overall Gibbs free energy change. For ionic solids, lattice energy and hydration energy are particularly important. For molecular compounds, polarity, hydrogen bonding, dispersion forces and crystal packing can dominate. A complete solubility assessment therefore considers both thermodynamic and chemical factors rather than relying on simple rules such as heat makes solids dissolve or like dissolves like.
Temperature is one of the most familiar factors affecting solubility. Many solid solutes become more soluble when heated, but the direction and magnitude of the change depend on the enthalpy and entropy of dissolution and on changes in the solid phase. If dissolution is endothermic, increasing temperature often favours dissolution according to thermodynamic principles. If dissolution is exothermic, increasing temperature can reduce equilibrium solubility. Some substances show relatively small temperature dependence, while others show dramatic changes.
Gases commonly behave differently: their solubility in liquids often decreases as temperature rises because dissolved gas molecules become more likely to escape into the gas phase. This explains why warm water generally contains less dissolved oxygen than cold water. Temperature therefore affects both condensed solutes and gases, but often in different ways.
Pressure strongly affects the solubility of gases but generally has relatively little effect on most solids and liquids under ordinary laboratory conditions. Henry's law states that, under appropriate conditions, the concentration of a dissolved gas is proportional to its partial pressure above the solution. Increasing carbon dioxide pressure increases carbon dioxide solubility, while releasing the pressure allows gas to leave solution.
Divers experience an important practical consequence of gas solubility because increased ambient pressure can increase the amount of nitrogen dissolved in body tissues, and rapid decompression can cause bubbles to form. Carbonated drinks provide a familiar example outside medicine: carbon dioxide is dissolved under elevated pressure and escapes as bubbles when pressure is reduced. These effects arise from gas liquid equilibrium and demonstrate why pressure must be considered when studying dissolved gases.
The chemical nature of the solute and solvent is fundamental. Polar solutes generally dissolve more readily in polar solvents, while nonpolar solutes generally dissolve more readily in nonpolar solvents, because favourable interactions must compensate for the disruption of existing interactions. Hydrogen bonding can strongly affect solubility. Ethanol is miscible with water because its hydroxyl group interacts strongly with water while its carbon chain is small enough that the polar interaction remains dominant.
Long hydrocarbon chains become progressively less compatible with water. Ionic compounds may dissolve when hydration sufficiently compensates for lattice energy, but a highly stable ionic lattice can make even an ionic substance sparingly soluble. Molecular shape, functional groups, polarizability, crystal packing and intermolecular forces therefore all contribute to solubility.
pH can dramatically affect the solubility of acids, bases and substances containing ionisable functional groups. A weak acid may be relatively poorly soluble in its neutral molecular form but become much more soluble when converted into a charged conjugate base under alkaline conditions. Similarly, a weak base can become more soluble in acidic conditions when protonated to form a charged conjugate acid. This principle is widely used in pharmaceutical formulation and acid base extraction.
For an acid HA, increasing pH can shift the equilibrium towards A minus, potentially increasing apparent aqueous solubility. For a base B, decreasing pH can shift the equilibrium towards BH plus. The total observed solubility can therefore include both intrinsic solubility of the neutral species and additional solubility caused by ionisation. This is one reason why pH control is an important method for manipulating solubility.
Other dissolved substances can alter solubility through common ion effects, complex formation, ionic strength and chemical reactions. If a sparingly soluble salt establishes an equilibrium involving a particular ion, adding that same ion can shift the equilibrium towards the solid and decrease solubility. This is the common ion effect. In contrast, a ligand that binds a dissolved metal ion can lower the free concentration of that ion and potentially increase the apparent solubility of the original solid.
EDTA provides a familiar example of complex formation affecting metal ion chemistry. Ionic strength can also alter activities, meaning that simple concentration based equilibrium calculations become less accurate at higher concentrations. Chemical reactions can remove dissolved species and thereby pull additional solid into solution. Solubility is consequently connected to the entire network of equilibria present rather than being an isolated property of one compound.
The physical form of a substance can also affect apparent solubility. Different crystal polymorphs can have different lattice energies and therefore different equilibrium solubilities. Amorphous forms may have higher apparent solubility than stable crystalline forms because their internal structure is less ordered and often higher in free energy. Hydrates and solvates can have different dissolution properties from anhydrous forms.
Particle size usually has a major effect on dissolution rate because smaller particles provide greater surface area, but it does not necessarily change the equilibrium solubility of a sufficiently large bulk phase. At very small particle sizes, surface energy can influence thermodynamic solubility, so the distinction is not absolutely rigid. Pharmaceutical scientists therefore consider polymorphism, crystallinity, particle size, formulation and solid state when controlling drug solubility and dissolution behaviour.
Solubility is also affected by the physical and chemical environment in practical systems. Mixing and stirring generally increase the rate at which equilibrium is reached by reducing concentration gradients around particles, but they do not normally change the final equilibrium solubility. Grinding a solid increases surface area and therefore speeds dissolution, again without ordinarily changing the bulk equilibrium value. Surfactants can increase apparent solubility by forming micelles that incorporate poorly water soluble molecules.
Cosolvents can alter solvent polarity and provide a more favourable environment for otherwise poorly soluble compounds. Biological fluids contain proteins, lipids, salts and other substances that can bind solutes or alter their distribution. Consequently, real world solubility can differ substantially from values measured in pure water.
