A salt in chemistry is defined as an ionic compound, typically a crystalline solid, composed of a positively charged ion (cation) and a negatively charged ion (anion). These compounds are generally the product of a neutralization reaction between an acid and a base. Solubility refers to the maximum amount of this solute, the salt, that can be dissolved in a given volume of a solvent, such as water. Water serves as an effective solvent because its molecules are polar.
How Ionic Compounds Interact with Water
When an ionic compound is introduced to water, the solvent’s polarity allows it to interact strongly with the salt’s crystal lattice structure. The small, polar water molecules orient their slightly negative oxygen end toward the salt’s positive cations. Simultaneously, the water molecules orient their slightly positive hydrogen ends toward the salt’s negative anions.
This strong attraction overcomes the electrostatic forces, known as lattice energy, holding the ions together in the solid structure. The ions then separate, a process called dissociation, and become surrounded by water molecules. This surrounding of ions by solvent molecules is known as hydration. A salt dissolves when the energy released during hydration is greater than the energy required to break the ionic bonds, resulting in a more stable state for the dissolved ions.
The General Rules for Soluble Salts
Predicting whether a salt will dissolve in water begins by knowing which ions almost always form soluble compounds. Salts containing Group 1 metal cations, such as sodium (\(Na^+\)) and potassium (\(K^+\)), are reliably soluble with virtually no exceptions. Similarly, any salt containing the ammonium ion (\(NH_4^+\)) will readily dissolve in water.
The presence of the nitrate ion (\(NO_3^-\)) in a compound is another strong indicator of solubility. Acetates (\(C_2H_3O_2^-\)) also form compounds that are nearly always soluble in water. These four groups—Group 1 metals, ammonium, nitrates, and acetates—form the foundation for identifying highly soluble salts.
Other anion groups are generally soluble but feature important exceptions. The halide ions, including chloride (\(Cl^-\)), bromide (\(Br^-\)), and iodide (\(I^-\)), typically form soluble salts. Sulfate salts (\(SO_4^{2-}\)), too, are generally soluble. The specific ions that cause these otherwise soluble groups to become insoluble are the key to predicting which salts will not dissolve.
Predicting Insoluble Compounds
Salts that are insoluble are identified by looking for specific combinations of cations and anions that do not dissociate in water. Certain anion groups are inherently likely to form insoluble compounds unless they are paired with one of the “always soluble” cations.
For instance, most salts containing the carbonate (\(CO_3^{2-}\)), phosphate (\(PO_4^{3-}\)), or sulfide (\(S^{2-}\)) ions are insoluble. The exceptions to this rule are only when these anions are paired with Group 1 metals or the ammonium ion.
The majority of hydroxide salts (\(OH^-\)) are also classified as insoluble. The only common exceptions are the hydroxides of Group 1 metals and barium (\(Ba^{2+}\)). Calcium hydroxide (\(Ca(OH)_2\)) is considered sparingly or slightly soluble.
Insoluble salts can also be formed when specific cations combine with the generally soluble halide and sulfate ions. While most chlorides, bromides, and iodides dissolve, they become insoluble when paired with silver (\(Ag^+\)), lead (\(Pb^{2+}\)), or mercury(I) (\(Hg_2^{2+}\)) ions. For example, silver chloride (\(AgCl\)) will precipitate out of a solution, making it an insoluble salt.
Similarly, the generally soluble sulfate ion forms insoluble salts with lead (\(Pb^{2+}\)) and barium (\(Ba^{2+}\)). Strontium (\(Sr^{2+}\)) and calcium (\(Ca^{2+}\)) sulfates are only slightly soluble. Lead sulfate (\(PbSO_4\)) is a well-known example of an insoluble salt.