Precipitation Reactions and Solubility Rules

A precipitation reaction occurs when two aqueous solutions are mixed and an insoluble ionic compound, called a precipitate, forms; solubility rules tell which ionic compounds are soluble in water and which are not. In brief: mixing silver nitrate and sodium chloride solutions gives a white precipitate of silver chloride (Ag⁺(aq) + Cl⁻(aq) → AgCl(s)), because chlorides are soluble except those of silver, lead(II) and mercury(I).

Precipitation reactions are used to identify ions, to remove pollutants from water and to make pigments and other materials. They also occur in nature in the formation of scale, stalactites and kidney stones. This article states the solubility rules with exceptions, shows how to predict a precipitate and write the net ionic equation, lists the colors of common precipitates, works through calculations of yield and limiting reagents, and gives some history.

What are the solubility rules?

The rules are general guidelines for the solubility of ionic compounds in water at room temperature. A compound is called soluble if more than about 1 g dissolves in 100 mL of water. No compound is completely insoluble, and “insoluble” really means very slightly soluble.

When two rules apply to the same compound, the exception takes priority: silver chloride is insoluble even though chlorides are generally soluble.

How do you predict a precipitate?

1. Write the ions present in each solution.
2. Swap the partners to find the possible products, writing correct formulas by balancing charges.
3. Use the solubility rules to decide whether each product is soluble (aq) or insoluble (s).
4. If at least one product is insoluble, a precipitation reaction occurs. Write the balanced molecular equation.
5. Write the complete ionic equation, with soluble strong electrolytes as ions.
6. Cancel the spectator ions to get the net ionic equation.
7. If both products are soluble, there is no reaction.

What are the colors of common precipitates?

Worked examples

1. Silver nitrate and sodium chloride. Possible products: AgCl and NaNO₃. AgCl is insoluble (exception to the chloride rule), NaNO₃ is soluble. Molecular equation: AgNO₃(aq) + NaCl(aq) → AgCl(s) + NaNO₃(aq). Net ionic equation: Ag⁺(aq) + Cl⁻(aq) → AgCl(s). The spectators are Na⁺ and NO₃⁻.
2. Mass of precipitate. 50.0 mL of 0.100 mol/L AgNO₃ and 50.0 mL of 0.100 mol/L NaCl are mixed. Each has 0.0500 L × 0.100 = 5.00 × 10⁻³ mol. The ratio is 1 : 1, so 5.00 × 10⁻³ mol of AgCl forms, which is 5.00 × 10⁻³ × 143.32 = 0.717 g.
3. Limiting reagent. 25.0 mL of 0.200 mol/L Pb(NO₃)₂ (5.00 × 10⁻³ mol of Pb²⁺) is mixed with 25.0 mL of 0.200 mol/L KI (5.00 × 10⁻³ mol of I⁻). Net equation: Pb²⁺ + 2I⁻ → PbI₂. The 5.00 × 10⁻³ mol of I⁻ can form only 2.50 × 10⁻³ mol of PbI₂, so iodide is limiting. The mass of PbI₂ (461.0 g/mol) is 2.50 × 10⁻³ × 461.0 = 1.15 g. The excess lead ions that remain: 5.00 × 10⁻³ − 2.50 × 10⁻³ = 2.50 × 10⁻³ mol.
4. No reaction. KNO₃(aq) + NaCl(aq): possible products KCl and NaNO₃ are both soluble, so there is no precipitate and no reaction.
5. Net ionic equation for copper hydroxide. CuSO₄(aq) + 2NaOH(aq) → Cu(OH)₂(s) + Na₂SO₄(aq). Net: Cu²⁺(aq) + 2OH⁻(aq) → Cu(OH)₂(s). A pale blue precipitate forms.
6. Percent yield. If 0.680 g of AgCl is collected in example 2 (theoretical 0.717 g), the percent yield is 0.680 ÷ 0.717 × 100% = 94.8%.
7. Which ion? A solution gives a white precipitate with barium chloride that does not dissolve in dilute nitric acid. The precipitate is BaSO₄, so the solution contains sulfate ions. (Barium carbonate would also be white, but it dissolves in acid.)

Where are precipitation reactions used?

• Qualitative analysis: identification of ions by the color and solubility of precipitates, such as the silver nitrate test for halides and the barium chloride test for sulfate.
• Water treatment: removal of heavy metals as hydroxides or sulfides, removal of phosphate with iron or aluminum salts and softening of hard water by precipitating calcium as the carbonate.
• Gravimetric analysis: measuring the amount of an ion by forming a precipitate, filtering, drying and weighing it.
• Manufacturing: pigments such as titanium and chromium compounds, and photography with silver halides.
• Medicine: barium sulfate is insoluble and is used as a contrast agent for X-rays.

Where do precipitates form in nature?

Water containing dissolved calcium hydrogencarbonate that drips into a cave loses carbon dioxide and deposits calcium carbonate: Ca²⁺ + 2HCO₃⁻ → CaCO₃(s) + H₂O + CO₂. Over thousands of years this builds stalactites and stalagmites. The same reaction causes limescale in kettles and pipes in hard-water regions. Kidney stones often consist of calcium oxalate or calcium phosphate, which precipitate from urine. Marine organisms precipitate calcium carbonate to build shells and coral.

How did the study of precipitation develop?

The formation of solids by mixing solutions was used for pigment production and for analysis by early chemists. Torbern Bergman systematized qualitative analysis with precipitation reagents in the 1770s, and Jöns Jacob Berzelius and others developed quantitative gravimetric analysis in the early 19th century. Later, Walther Nernst introduced the solubility product in 1889, which gave a quantitative explanation for precipitation: a solid forms when the product of the ion concentrations exceeds the solubility product constant K_sp.

Which mistakes are common?

Mistake 1: Writing the formula of the precipitate without balancing charges

The precipitate from Pb²⁺ and I⁻ is PbI₂, not PbI.

Mistake 2: Including spectator ions in the net ionic equation

Only the ions that form the precipitate appear in the net equation.

Mistake 3: Forgetting the exceptions in the solubility rules

Most chlorides are soluble, but AgCl, PbCl₂ and Hg₂Cl₂ are not.

Mistake 4: Assuming that a precipitate means a complete reaction

A slightly soluble compound still has a small amount of dissolved ions in equilibrium with the solid.