Solubility rules are key in predicting which compounds dissolve in water and which do not. While there are many commonly accepted rules, most have exceptions. These rules are not completely consistent across all textbooks and sources; however, the rules listed below are derived from experimental data and research papers to ensure the best possible accuracy. Before memorizing solubility rules, it is important to understand what makes some compounds more soluble than others.
Solubility is determined largely by lattice energy. Lattice energy is the energy required to separate ions from an ionic crystal. Therefore, the higher the lattice energy, the lower the solubility. Some important trends for lattice energy are that higher ionic charge leads to greater lattice energy, and smaller ions also result in higher lattice energy. Since lattice energy is always endothermic, it may seem confusing why compounds dissolve in water at all. This is explained by hydration energy. Hydration energy is the energy released when water molecules surround an ion and stabilize its charge. The constant competition between lattice energy and hydration energy determines whether a compound is soluble in water.
This is the most well-known solubility rule and is generally very accurate. However, there is an important exception. Lithium forms a very small ion, which leads to unusually high lattice energy. Because of this, compounds such as Li₃PO₄, Li₂CO₃, and LiF are known to be insoluble.
This rule holds true in most cases. Nitrate is unique because it is always soluble. These anions are relatively large and can stabilize charge effectively, resulting in lower lattice energy and higher solubility. Some notable exceptions include silver acetate (AgCH₃COO), mercury(I) acetate [Hg₂(CH₃COO)₂], silver chlorate (AgClO₃), and certain heavy-metal chlorates, which have limited solubility.
The low solubility of carbonate (CO₃²⁻) and phosphate (PO₄³⁻) ions is caused by their large negative charges, which produce high lattice energy. Group 1 metal cations reduce this lattice energy enough to allow solubility. However, lithium remains an exception, so Li₂CO₃ and Li₃PO₄ are still insoluble.
Sulfide ions (S²⁻) form strong ionic lattices with most cations, making these compounds insoluble. Group 1 metal sulfides and ammonium sulfide have lower lattice energies, allowing water to overcome the ionic lattice. Group 2 sulfides are only partially soluble due to their higher lattice energies.
This is one of the most important solubility rules because halide salts appear frequently in chemical reactions. A notable property occurs with lead(II) halides, whose solubility increases significantly with temperature. This behavior is uncommon, as most ionic compounds show little temperature dependence. It is also important to note that aluminum being insoluble with halides does not apply to aluminum fluoride; AlF₃ is highly soluble due to strong hydration of the Al³⁺ ion.
The high solubility of most sulfate salts is due to the large size of the sulfate ion, which lowers lattice energy. Although some textbooks list these exceptions as “slightly soluble,” in laboratory settings all of these cations form visible precipitates in the presence of sulfate.
Most hydroxide salts are insoluble due to strong ionic bonding. Group 1 metal hydroxides are fully soluble, while Ca(OH)₂, Sr(OH)₂, and Ba(OH)₂ show moderate to high solubility.