How Solubility Rules Predict Precipitation Reactions

A precipitate is a low-solubility product

A precipitation reaction occurs when ions already dissolved in water recombine to form a compound with low solubility. The visible solid is the precipitate. Introductory problems use qualitative solubility rules to make this prediction without concentration or equilibrium calculations.

Use a three-step prediction method

1. Separate each aqueous ionic compound into its cation and anion.
2. Pair each cation with the other compound's anion to identify the possible double-replacement products.
3. Apply the solubility rules to both products. If either is insoluble, predict a precipitate and write that solid's neutral empirical formula. If both are soluble, report no precipitate.

Always balance ionic charges when writing the empirical formula. A $2+$ cation and a $2-$ anion combine in a 1:1 ratio; a $2+$ cation with a $1-$ anion requires two anions.

A fresh example

Mixing aqueous sodium carbonate and calcium chloride gives possible products sodium chloride and calcium carbonate. Sodium salts and most chlorides are soluble, while calcium carbonate is insoluble under the standard guidelines. The precipitate is therefore $\mathrm{CaCO_3(s)}$:

$$\mathrm{Ca^{2+}(aq)+CO_3^{2-}(aq)\rightarrow CaCO_3(s)}.$$

Know when qualitative rules are not enough

Solubility rules are appropriate for the ordinary aqueous mixtures used in introductory exercises. Near a solubility threshold, actual precipitation depends on ion concentrations and the solubility product $K_{sp}$. Temperature, complex-ion formation, and other specified conditions can also change the result, so use equilibrium data when a problem supplies them.