Salts are a broad class of ionic compounds formed from the reaction of an acid and a base. These compounds consist of a lattice structure held together by the strong electrostatic attraction between positively charged ions (cations) and negatively charged ions (anions). Solubility refers to the maximum amount of a salt that can dissolve in a specific amount of solvent, typically water, at a given temperature. While common table salt, sodium chloride, dissolves easily, not all salts are soluble. The capacity to dissolve is governed by predictable chemical rules based on the energetic balance between the forces holding the solid together and the forces exerted by the surrounding water molecules.
The Chemistry of Dissolution
For an ionic salt to dissolve in water, the process must overcome two competing forces. The first force is the lattice energy, which represents the intense energy required to break apart the ionic bonds and separate the ions from the rigid, crystalline structure. High lattice energy, common for ions that are small or have a high charge, makes a salt less likely to dissolve.
Working against the lattice energy is the hydration energy, which is the energy released when the separated ions become surrounded by water molecules. Water is a polar solvent, meaning its molecules orient themselves to attract the ions. As the salt’s ions enter the water, this forms a protective layer called a hydration shell around each ion.
Dissolution occurs when the energy released by the hydration of the ions is equal to or greater than the energy required to break the crystal lattice. If the attractive forces between the water molecules and the ions are strong enough, the salt will be soluble. If the lattice energy is too high, the ions remain locked in the solid state, and the salt is considered insoluble, forming a solid precipitate.
General Solubility Guidelines
Chemists use a set of guidelines based on the behavior of specific ions to predict a salt’s solubility. The most reliable guidelines focus on ions that almost always result in a soluble salt, regardless of the ion they are paired with. These ions are often referred to as “always soluble” because their hydration energy is consistently high enough to overcome most lattice structures.
The following ions form salts that are uniformly soluble in water:
- Alkali metal ions (Lithium, Sodium, Potassium, and other Group 1 elements).
- The ammonium ion (\(NH_4^+\)).
- The nitrate ion (\(NO_3^-\)).
- Acetate salts (\(CH_3COO^-\)).
The large size and single negative charge of the nitrate ion, for example, allows it to form soluble compounds with virtually every cation. These categories form the foundation for predicting which salts will dissolve in an aqueous solution.
Key Exceptions to Solubility
Variability in salt solubility arises from specific exceptions to the general guidelines, determining why certain ion combinations are insoluble. Halide ions, including chloride (\(Cl^-\)), bromide (\(Br^-\)), and iodide (\(I^-\)), are generally soluble. However, pairing these halides with the heavy metal ions Silver (\(Ag^+\)), Lead (\(Pb^{2+}\)), or Mercury(I) (\(Hg_2^{2+}\)) results in notably insoluble salts. Silver chloride (\(AgCl\)) is a classic example of an insoluble halide salt that forms a white precipitate.
Sulfates (\(SO_4^{2-}\)) are another group of typically soluble ions that have several important exceptions. The sulfates of Barium (\(Ba^{2+}\)), Strontium (\(Sr^{2+}\)), and Lead (\(Pb^{2+}\)) are insoluble. Calcium sulfate (\(CaSO_4\)) is also an exception, as it is only sparingly soluble in water.
Conversely, some ions are generally insoluble, except when paired with the “always soluble” ions. For instance, most salts containing the hydroxide ion (\(OH^-\)) or the carbonate ion (\(CO_3^{2-}\)) are insoluble. The significant exceptions to this rule are the hydroxides and carbonates of the Group 1 metals and ammonium, which remain soluble due to their strong guideline. This interplay of general rules and specific exceptions allows chemists to accurately predict solubility.