Many ionic compounds have solubilities that follow useful empirical patterns. Solubility rules summarize these patterns so that a likely soluble or sparingly soluble salt can be predicted without yet calculating an equilibrium constant.

A practical introductory set of rules is:

• salts containing group 1 cations or $\mathrm{NH_4^+}$ are generally soluble;
• nitrates, $\mathrm{NO_3^-}$, are generally soluble;
• chlorides, bromides and iodides are generally soluble, with important exceptions such as salts of $\mathrm{Ag^+}$, $\mathrm{Pb^{2+}}$ and $\mathrm{Hg_2^{2+}}$;
• sulfates are often soluble, but several salts of heavier or group-2 cations are sparingly soluble;
• carbonates and phosphates are generally sparingly soluble except with group 1 cations or $\mathrm{NH_4^+}$;
• many hydroxides are sparingly soluble, with group 1 hydroxides and some heavier group-2 hydroxides as important exceptions.

For example, classify $\mathrm{Na_2CO_3}$ and $\mathrm{CaCO_3}$. Carbonates are usually sparingly soluble, but sodium is a group 1 cation. Therefore $\mathrm{Na_2CO_3}$ is predicted to be soluble, whereas $\mathrm{CaCO_3}$ is predicted to be sparingly soluble.

Likewise, $\mathrm{AgCl}$ is an exception to the usual solubility of chlorides and is therefore expected to form a solid readily in aqueous reactions.

These rules are qualitative empirical shortcuts, not exact physical laws. Terms such as “soluble” and “insoluble” hide a continuum of finite solubilities, and temperature, composition and complex formation can alter the result. A later treatment using the solubility product $K_{sp}$ quantifies when precipitation actually occurs.

The role of solubility rules at this stage is narrower: identify plausible aqueous ions and likely solid products well enough to write and interpret precipitation reactions.