A salt is any ionic compound composed of positively charged ions (cations) and negatively charged ions (anions) held together by electrostatic forces. When a salt is placed in water, it undergoes dissolution, separating into its constituent ions. Solubility refers to the maximum amount of salt that can dissolve in a specific volume of solvent at a set temperature. This process exists in a dynamic equilibrium, where the solid salt is continuously dissolving while the dissolved ions are simultaneously re-forming the solid structure. The point at which the solution becomes saturated represents this equilibrium, and several conditions can shift this balance to make a salt less soluble.
Temperature’s Dual Role in Dissolution
The simple addition of heat energy to a saturated solution does not always guarantee an increase in solubility; the effect of temperature depends entirely on the nature of the salt’s dissolution process. For the majority of solid salts, dissolving in water is an endothermic process, meaning it absorbs thermal energy from the surroundings. When the system is heated, the added energy drives the dissolution process forward, allowing more of the salt to dissolve and thus increasing its solubility. A common example of this is the dissolution of ammonium nitrate, which is used in instant cold packs and causes the surrounding water to feel colder as it dissolves.
Some salts exhibit an exothermic dissolution process, meaning dissolving releases heat into the solution. For these salts, increasing the temperature introduces a stress on the equilibrium by adding more heat. To relieve this stress, the system shifts its balance away from the dissolved ions and back toward the solid salt. This results in a decrease in solubility as the temperature rises, forcing some dissolved ions to re-form the solid compound. Calcium hydroxide is one such compound whose solubility decreases noticeably as the water temperature increases.
The Effect of a Shared Ion
One of the most direct ways to reduce a salt’s solubility is by introducing an ion already present in the solution, a phenomenon known as the Common Ion Effect. When a sparingly soluble salt dissolves, it establishes a specific ratio between the solid form and the dissolved ions, quantified by the Solubility Product Constant (\(K_{sp}\)).
Consider the example of silver chloride (\(AgCl\)), which dissolves to form silver ions (\(Ag^{+}\)) and chloride ions (\(Cl^{-}\)). If a second salt, such as sodium chloride (\(NaCl\)), is added to this saturated solution, it introduces a large quantity of chloride ions, which are “common” to the original equilibrium. This sudden increase in the concentration of chloride ions disrupts the established \(K_{sp}\) ratio.
To re-establish the equilibrium, the system must consume the excess chloride ions by combining them with the dissolved silver ions. This reaction forces the equilibrium to shift back toward the solid silver chloride. The result of this shift is the precipitation of more solid \(AgCl\), causing the overall amount of silver chloride dissolved in the water to decrease significantly.
How pH Changes Impact Solubility
Changes in the acidity or basicity of the water, measured by the pH scale, can also reduce a salt’s solubility, but this effect is specific to salts whose ions can react with hydrogen ions (\(H^{+}\)) or hydroxide ions (\(OH^{-}\)). The solubility of a salt is not affected by pH if its ions are derived from a strong acid and a strong base, such as sodium chloride, because these ions do not react with water. However, if a salt contains an anion that is the conjugate base of a weak acid, its solubility becomes highly sensitive to the surrounding pH.
Lowering the pH by adding acid generally increases the solubility of these salts because the added hydrogen ions react with and effectively remove the salt’s anion from the solution. Conversely, the solubility of these salts is reduced when the pH is raised, meaning the solution becomes more basic. When the concentration of hydrogen ions is reduced in a basic solution, the salt’s anion is less likely to be protonated.
This leaves a higher concentration of the free anion in the water, which acts like a common ion and shifts the dissolution equilibrium back toward the solid salt. For instance, salts containing the fluoride ion (\(F^{-}\)), the conjugate base of the weak acid hydrofluoric acid (\(HF\)), become less soluble when the pH is increased. The higher concentration of hydroxide ions in a basic solution leaves more \(F^{-}\) ions to combine with the cation and form the solid salt again.