Precipitation Reactions

Precipitation reactions are central to qualitative analysis (identifying unknown ions) and to making insoluble salts at IGCSE. Recognising when a precipitate will form requires knowledge of solubility rules.

Definition

A precipitate is an insoluble solid formed when two solutions are mixed. A precipitation reaction produces this insoluble product from soluble reactants.

Solubility rules

You must know these rules to predict whether a precipitate forms:

Soluble compounds

Type	Rule	Exceptions
All sodium, potassium, ammonium compounds	Always soluble	None
All nitrates	Always soluble	None
All chlorides	Soluble	Except AgCl, PbCl2
All sulfates	Soluble	Except BaSO4, CaSO4, PbSO4

Insoluble compounds

Type	Rule	Exceptions
All carbonates	Insoluble	Except Na2CO3, K2CO3, (NH4)2CO3
All hydroxides	Insoluble	Except NaOH, KOH, Ca(OH)2 (slightly soluble)

Common precipitation reactions

Halide tests (using silver nitrate)

Add dilute nitric acid then AgNO3(aq):

Halide	Equation	Precipitate
Cl-	Ag+(aq) + Cl-(aq) -> AgCl(s)	White
Br-	Ag+(aq) + Br-(aq) -> AgBr(s)	Cream
I-	Ag+(aq) + I-(aq) -> AgI(s)	Yellow

See silver nitrate for full details.

Sulfate test (using barium nitrate)

Add dilute nitric acid then Ba(NO3)2(aq):

Ba2+(aq) + SO4^2-(aq) -> BaSO4(s) — white precipitate

The white precipitate of barium sulfate is insoluble in excess. This confirms the presence of sulfate ions.

Cation tests (using sodium hydroxide)

Add NaOH(aq) to the test solution:

Cation	Precipitate colour	Soluble in excess NaOH?
Cu2+	Blue	No
Fe2+	Green	No
Fe3+	Orange-brown	No
Al3+	White	Yes (dissolves)
Zn2+	White	Yes (dissolves)
Ca2+	White (slight)	No
Mg2+	White	No

Ionic equations:

Cu2+(aq) + 2OH-(aq) -> Cu(OH)2(s)

Fe2+(aq) + 2OH-(aq) -> Fe(OH)2(s)

Fe3+(aq) + 3OH-(aq) -> Fe(OH)3(s)

Al3+(aq) + 3OH-(aq) -> Al(OH)3(s) — dissolves in excess: Al(OH)3 + OH- -> Al(OH)4-

Lead(II) iodide

Pb(NO3)2(aq) + 2KI(aq) -> PbI2(s) + 2KNO3(aq)

Ionic: Pb2+(aq) + 2I-(aq) -> PbI2(s)

Observation: bright yellow precipitate.

Making insoluble salts by precipitation

To make an insoluble salt (e.g. BaSO4, AgCl, PbI2):

1. Mix solutions of two soluble salts that contain the required ions
2. Filter to collect the precipitate
3. Wash with distilled water to remove soluble impurities
4. Dry in an oven or between filter papers

Example: to make lead(II) iodide:

• Mix lead(II) nitrate solution with potassium iodide solution
• Filter the yellow precipitate
• Wash and dry

See preparation of salts and making salts decision tree.

Writing ionic equations

Steps:

1. Write the full balanced equation
2. Split all soluble ionic compounds into ions; keep solids, liquids, gases as formulas
3. Cancel spectator ions (ions identical on both sides)
4. What remains is the ionic equation

Example: NaCl(aq) + AgNO3(aq) -> AgCl(s) + NaNO3(aq)

Full ionic: Na+(aq) + Cl-(aq) + Ag+(aq) + NO3-(aq) -> AgCl(s) + Na+(aq) + NO3-(aq)

Cancel Na+ and NO3-: Ag+(aq) + Cl-(aq) -> AgCl(s)

Describe how you would prepare a pure, dry sample of barium sulfate. [4 marks]

• Mix barium chloride solution with sodium sulfate solution (or dilute sulfuric acid) [1]
• A white precipitate of barium sulfate forms [1]
• Filter the mixture to collect the precipitate [1]
• Wash the precipitate with distilled water, then dry in a warm oven or between filter papers [1]

A solution contains either Fe2+ or Fe3+ ions. Describe a test to distinguish between them using sodium hydroxide solution. [3 marks]

• Add sodium hydroxide solution to the test solution [1]
• If Fe2+ is present, a green precipitate of Fe(OH)2 forms [1]
• If Fe3+ is present, an orange-brown (rust-coloured) precipitate of Fe(OH)3 forms [1]