A precipitation reaction occurs when two soluble ionic solutions are mixed and produce an insoluble solid — the precipitate — which falls out of solution. It is one of the most useful reactions in analytical chemistry because the colour of the precipitate immediately identifies which metal ion or anion was present in the original solution.
What is a precipitation reaction?
When ionic compounds dissolve in water, they dissociate (split) into their constituent ions. If two solutions are mixed and one combination of ions forms an insoluble compound, those ions come together and form a solid that falls out of solution as a precipitate.
The key is solubility: whether a compound dissolves in water. A simple rule set for GCSE:
| Rule | Soluble | Insoluble |
|---|---|---|
| All nitrates | All | — |
| Group 1 salts and ammonium salts | All | — |
| Most chlorides | Most | Silver chloride (AgCl), lead chloride (PbCl₂) |
| Most sulfates | Most | Barium sulfate (BaSO₄), lead sulfate (PbSO₄), calcium sulfate (CaSO₄) |
| Most hydroxides and carbonates | Few (Na, K, NH₄⁺) | Most metal hydroxides and carbonates |
How do you write equations for precipitation reactions?
Step 1 — Full ionic equation:
Write all dissolved species as separate ions (since they are fully dissociated in solution); write the precipitate as a formula with state symbol (s).
Example: mixing silver nitrate solution with sodium chloride solution.
Full ionic equation: Ag⁺(aq) + NO₃⁻(aq) + Na⁺(aq) + Cl⁻(aq) → AgCl(s) + Na⁺(aq) + NO₃⁻(aq)
Step 2 — Net ionic equation (cancel spectator ions):
Na⁺ and NO₃⁻ appear on both sides unchanged — they are spectator ions and are cancelled.
Net ionic equation: Ag⁺(aq) + Cl⁻(aq) → AgCl(s)
The net ionic equation shows only the species that actually change — it applies regardless of which soluble source of Ag⁺ and Cl⁻ was used.
Which metal hydroxide precipitates are used to identify metal ions?
Adding sodium hydroxide solution to a solution of a metal salt produces a metal hydroxide precipitate whose colour identifies the metal ion:
| Metal ion | Precipitate formula | Colour |
|---|---|---|
| Fe²⁺ (iron(II)) | Fe(OH)₂ | Green |
| Fe³⁺ (iron(III)) | Fe(OH)₃ | Brown/rust |
| Cu²⁺ (copper(II)) | Cu(OH)₂ | Blue |
| Al³⁺ (aluminium) | Al(OH)₃ | White (dissolves in excess NaOH — amphoteric) |
| Ca²⁺ (calcium) | Ca(OH)₂ | White (slight — Ca(OH)₂ is slightly soluble) |
| Zn²⁺ (zinc) | Zn(OH)₂ | White (dissolves in excess NaOH) |
Worked example — identifying iron(III) ions:
- Add a few drops of NaOH solution to the unknown solution.
- A rust-brown precipitate forms.
- Net ionic equation: Fe³⁺(aq) + 3OH⁻(aq) → Fe(OH)₃(s)
- Conclusion: iron(III) ions are present.
How are precipitation reactions used to test for anions?
Test for sulfate ions (SO₄²⁻):
- Acidify with dilute hydrochloric acid, then add barium chloride solution.
- A white precipitate of barium sulfate (BaSO₄) confirms sulfate.
- Net ionic equation: Ba²⁺(aq) + SO₄²⁻(aq) → BaSO₄(s)
- The acid removes carbonate ions first (which would also precipitate with barium), ensuring the result is specific to sulfate.
Test for chloride, bromide and iodide ions (halide ions):
- Acidify with dilute nitric acid, then add silver nitrate solution.
| Halide ion | Precipitate | Colour | Solubility in ammonia solution |
|---|---|---|---|
| Cl⁻ | AgCl | White | Dissolves in dilute ammonia |
| Br⁻ | AgBr | Cream | Dissolves in concentrated ammonia only |
| I⁻ | AgI | Yellow | Insoluble in ammonia |
The solubility in ammonia allows you to distinguish the three precipitates if colour alone is ambiguous.
Where are precipitation reactions used outside the laboratory?
| Application | Reaction | Purpose |
|---|---|---|
| Water treatment | Adding lime (Ca(OH)₂) to water containing Mg²⁺ → Mg(OH)₂ precipitate | Softening hard water |
| Photography | AgBr precipitate on film | Light-sensitive layer; AgBr decomposes on exposure to light |
| Kidney stones | Calcium oxalate precipitates in the kidney | A harmful precipitation in the body |
| Antacids | Al(OH)₃ or Mg(OH)₂ forms in the stomach | Neutralises excess acid; these insoluble compounds are swallowed as suspensions |
Frequently asked questions
Why does the precipitate form immediately when the two solutions are mixed?
When ionic solutions are mixed, the ions are already fully dissociated and moving freely throughout the solution. As soon as the two solutions meet, ions of opposite types collide. If the combination is insoluble, the lattice forces between those ions are so strong that they immediately overcome the hydration energy (the attraction of water molecules to the ions) and the ions come together to form a solid lattice. The reaction is essentially instantaneous at the point of mixing — no activation energy is needed because no bonds need to be broken in the ionic species themselves.
What is a spectator ion?
A spectator ion is an ion that is present in the reaction mixture but does not take part in the precipitation reaction. It remains dissolved throughout — it is neither changed nor incorporated into the precipitate. For example, in the reaction of silver nitrate with sodium chloride, Na⁺ and NO₃⁻ are spectator ions: they start dissolved and finish dissolved. Identifying and cancelling spectator ions gives the net ionic equation, which is more fundamental — it shows exactly what is happening at the ionic level, regardless of which soluble salts were used as starting materials.
How do you distinguish between iron(II) and iron(III) using precipitation?
Both iron ions form precipitates with sodium hydroxide solution, but the colours are distinct: Fe²⁺ gives a green precipitate of Fe(OH)₂, while Fe³⁺ gives a rust-brown precipitate of Fe(OH)₃. Green and brown are sufficiently different to distinguish easily. You can also observe that the green Fe(OH)₂ precipitate slowly darkens on standing in air, because the Fe²⁺ is oxidised to Fe³⁺ by oxygen — providing a secondary check that you started with iron(II).
Why must sulfate tests be acidified with hydrochloric acid first?
If carbonate ions (CO₃²⁻) are present in the test solution and you add barium chloride directly, barium carbonate (BaCO₃) also precipitates as a white solid — indistinguishable from barium sulfate (BaSO₄) by eye. Adding dilute hydrochloric acid first dissolves any carbonate ions (CO₃²⁻ + 2H⁺ → CO₂ + H₂O) before the barium chloride is added, ensuring the white precipitate that then forms can only be barium sulfate. The acid must be hydrochloric, not sulfuric (which would introduce SO₄²⁻ ions and give a false positive).
For Socratic GCSE chemistry with Professor Curie — building precipitation logic from the particle model through ionic equations to analytical identification — visit aitutors.me.