What you'll learn
Qualitative analysis identifies which ions are present in a substance, using characteristic observations — colours of precipitates, gases given off, flame colours and changes on adding excess reagent. These tests are used in water and soil testing, in forensic and quality-control laboratories, and they form an important part of CAPE practical examinations. This guide covers flame tests, tests for cations using aqueous sodium hydroxide and aqueous ammonia, tests for anions (carbonate, sulfate, sulfite, halides, nitrate and others), tests for gases, the chemistry behind each observation, and how to plan a systematic analysis of an unknown salt.
CAPE practical and written papers reward precise observations (exact colours, "white precipitate soluble in excess", "effervescence") and correct inferences, supported by ionic equations. Learning the tests as a connected system rather than as isolated facts makes them easier to remember.
Key terms and definitions
Qualitative analysis — Identifying the components of a substance.
Precipitate — An insoluble solid formed when two solutions react.
Amphoteric hydroxide — A hydroxide that dissolves in both acids and excess strong base.
Complex ion — A central metal ion surrounded by ligands; formation of complexes can cause a precipitate to dissolve.
Effervescence — Bubbling caused by a gas being given off.
Inference — A conclusion about the identity of an ion drawn from an observation.
Ionic equation — An equation showing only the ions and species that take part in a reaction.
Core concepts
Flame tests
A clean nichrome or platinum wire is dipped in concentrated hydrochloric acid, then into the sample, and held in a blue Bunsen flame. Heat promotes electrons to higher energy levels; as they fall back, they emit light of characteristic colours.
| Ion | Flame colour |
|---|---|
| Li⁺ | crimson red |
| Na⁺ | golden yellow (persistent) |
| K⁺ | lilac (view through blue glass to remove sodium contamination) |
| Ca²⁺ | brick red |
| Sr²⁺ | crimson |
| Ba²⁺ | apple green |
| Cu²⁺ | blue-green |
Testing cations with aqueous sodium hydroxide
Add NaOH(aq) dropwise, then in excess. Many metal ions form insoluble hydroxide precipitates:
| Ion | With a few drops of NaOH | With excess NaOH |
|---|---|---|
| Ca²⁺ | white precipitate | insoluble |
| Mg²⁺ | white precipitate | insoluble |
| Al³⁺ | white precipitate | dissolves → colourless solution |
| Zn²⁺ | white precipitate | dissolves → colourless solution |
| Pb²⁺ | white precipitate | dissolves → colourless solution |
| Cu²⁺ | pale blue precipitate | insoluble |
| Fe²⁺ | green precipitate (turns brown on standing in air) | insoluble |
| Fe³⁺ | red-brown precipitate | insoluble |
| NH₄⁺ | no precipitate; on warming, ammonia gas (turns damp red litmus blue) | — |
Ionic equations, e.g.:
- Cu²⁺(aq) + 2OH⁻(aq) → Cu(OH)₂(s)
- Fe³⁺(aq) + 3OH⁻(aq) → Fe(OH)₃(s)
- Al(OH)₃(s) + OH⁻(aq) → [Al(OH)₄]⁻(aq) — dissolves because aluminium hydroxide is amphoteric.
The green Fe(OH)₂ darkens because it is oxidised by air to Fe(OH)₃.
Testing cations with aqueous ammonia
Aqueous ammonia contains a low concentration of OH⁻ ions, so it also precipitates hydroxides, but excess ammonia can form ammine complexes.
| Ion | With a few drops of NH₃ | With excess NH₃ |
|---|---|---|
| Ca²⁺ | no precipitate (or very slight) | — |
| Mg²⁺ | white precipitate | insoluble |
| Al³⁺ | white precipitate | insoluble |
| Zn²⁺ | white precipitate | dissolves → colourless solution ([Zn(NH₃)₄]²⁺) |
| Pb²⁺ | white precipitate | insoluble |
| Cu²⁺ | pale blue precipitate | dissolves → deep blue solution ([Cu(NH₃)₄(H₂O)₂]²⁺) |
| Fe²⁺ | green precipitate | insoluble |
| Fe³⁺ | red-brown precipitate | insoluble |
Distinguishing white precipitates: Al³⁺ and Zn²⁺ both dissolve in excess NaOH, but only Zn²⁺ dissolves in excess NH₃. Al³⁺ and Pb²⁺ both behave similarly with NaOH; Pb²⁺ can be confirmed with potassium iodide, which gives a bright yellow precipitate of PbI₂.
Tests for anions
| Ion | Test | Positive result |
|---|---|---|
| CO₃²⁻ | add dilute acid | effervescence; gas (CO₂) turns limewater milky |
| SO₄²⁻ | add dilute HCl (or HNO₃), then BaCl₂(aq) (or Ba(NO₃)₂) | white precipitate of BaSO₄, insoluble in acid |
| SO₃²⁻ | add dilute acid and warm | SO₂ gas, turns acidified K₂Cr₂O₇ from orange to green; with BaCl₂, white precipitate soluble in dilute HCl |
| Cl⁻ | add dilute HNO₃, then AgNO₃(aq) | white precipitate (AgCl), soluble in dilute NH₃ |
| Br⁻ | add dilute HNO₃, then AgNO₃(aq) | cream precipitate (AgBr), soluble in concentrated NH₃ |
| I⁻ | add dilute HNO₃, then AgNO₃(aq) | yellow precipitate (AgI), insoluble in NH₃ |
| NO₃⁻ | add NaOH(aq) and aluminium foil (or Devarda's alloy), warm | ammonia gas, turns damp red litmus blue |
Why acidify? In the sulfate and halide tests, acid is added first to remove carbonate (and sulfite) ions, which would otherwise form misleading white precipitates with barium or silver ions. Nitric acid (not HCl) is used before silver nitrate because chloride ions from HCl would precipitate.
Ionic equations:
- CO₃²⁻ + 2H⁺ → CO₂ + H₂O
- Ba²⁺ + SO₄²⁻ → BaSO₄
- Ag⁺ + Cl⁻ → AgCl
Tests for gases
| Gas | Test | Result |
|---|---|---|
| Hydrogen | lighted splint | squeaky "pop" |
| Oxygen | glowing splint | relights |
| Carbon dioxide | limewater | turns milky |
| Ammonia | damp red litmus | turns blue |
| Chlorine | damp blue litmus | turns red, then bleached |
| Sulfur dioxide | acidified K₂Cr₂O₇ paper | orange → green |
| Hydrogen chloride | ammonia vapour (glass rod dipped in conc. NH₃) | dense white fumes of NH₄Cl |
| Nitrogen dioxide | appearance / damp blue litmus | brown gas; litmus turns red |
Planning a systematic analysis
- Note the appearance of the solid (colour suggests transition metal ions: blue/green for Cu²⁺, green for Fe²⁺ and Ni²⁺, yellow-brown for Fe³⁺).
- Carry out a flame test.
- Dissolve in water; test separate portions with NaOH(aq) and NH₃(aq), adding each dropwise and then in excess.
- Test for anions: dilute acid (carbonate/sulfite), acidified barium chloride (sulfate), acidified silver nitrate (halides), and the nitrate test.
- Record observations precisely and make inferences, confirming with a further test where possible.
Interfering ions and confirmatory tests
Some tests can give similar results for different ions, so confirmatory tests are important:
- Fe²⁺ can be confirmed with potassium hexacyanoferrate(III) (dark blue precipitate) and Fe³⁺ with potassium thiocyanate (blood-red colour).
- Sodium contamination masks other flame colours; potassium is confirmed by viewing through blue cobalt glass.
Worked examples
Example 1: Identifying a cation
A solution gives a white precipitate with NaOH(aq) that dissolves in excess. With NH₃(aq), it gives a white precipitate that also dissolves in excess. Identify the cation.
Zn²⁺ — the only common cation whose white hydroxide dissolves in excess of both NaOH and NH₃.
Example 2: A full analysis
A green crystalline solid dissolves in water. With NaOH(aq), a green precipitate forms which turns brown at the surface. With acidified BaCl₂(aq), a white precipitate forms. Identify the salt.
The green precipitate darkening in air indicates Fe²⁺; the white precipitate with acidified barium chloride indicates SO₄²⁻. The salt is iron(II) sulfate.
Example 3: Ionic equation
Write the ionic equation for the reaction of aqueous copper(II) ions with sodium hydroxide, including state symbols.
Cu²⁺(aq) + 2OH⁻(aq) → Cu(OH)₂(s)
Example 4: Why the acid?
Explain why dilute nitric acid is added before silver nitrate in the test for chloride ions. (2 marks)
To remove carbonate ions, which would otherwise form a precipitate of silver carbonate and give a false positive (1). Nitric acid is used rather than hydrochloric acid because HCl would add chloride ions (1).
Example 5: A complete observation–inference table
A white solid, Q, was analysed. Complete the inferences.
| Test | Observation | Inference |
|---|---|---|
| Flame test | brick-red flame | Ca²⁺ present |
| Q dissolved in water; NaOH(aq) added dropwise, then in excess | white precipitate, insoluble in excess | Ca²⁺ or Mg²⁺ (confirmed Ca²⁺ by flame test) |
| NH₃(aq) added | no precipitate | consistent with Ca²⁺ (Mg²⁺ would give a white precipitate) |
| Dilute HCl added to solid Q | effervescence; gas turns limewater milky | CO₃²⁻ present (gas is CO₂) |
Conclusion: Q is calcium carbonate, CaCO₃ — the main component of the limestone found widely across the Caribbean.
Example 6: Separating two white precipitates
A solution contains either Al³⁺ or Mg²⁺. Describe one test to decide which.
Add aqueous sodium hydroxide dropwise and then in excess. Both give a white precipitate, but aluminium hydroxide dissolves in excess (it is amphoteric), forming a colourless solution, whereas magnesium hydroxide remains insoluble. Al(OH)₃(s) + OH⁻(aq) → [Al(OH)₄]⁻(aq)
Example 7: Sulfate or sulfite?
A solution gives a white precipitate with barium chloride. Describe how to tell whether the anion is sulfate or sulfite.
Add dilute hydrochloric acid to the precipitate. Barium sulfate is insoluble in acid, so the precipitate remains. Barium sulfite dissolves in acid, releasing sulfur dioxide, which turns acidified potassium dichromate(VI) paper from orange to green.
Example 8: Testing for ammonium ions in a fertiliser
Describe how to show that a fertiliser contains ammonium ions.
Warm a small sample with aqueous sodium hydroxide. A pungent gas is given off that turns damp red litmus paper blue, showing it is ammonia: NH₄⁺(aq) + OH⁻(aq) → NH₃(g) + H₂O(l)
Writing good observations
| Weak | Precise |
|---|---|
| "It went cloudy." | "A white precipitate formed." |
| "It dissolved." | "The precipitate dissolved in excess, forming a colourless solution." |
| "Gas was given off." | "Effervescence; the gas relit a glowing splint (oxygen)." |
| "It turned blue." | "The pale blue precipitate dissolved in excess ammonia, forming a deep blue solution." |
| "Clear solution." | "Colourless solution" (or name the colour if coloured). |
Examiners award marks for precise colours, states and solubility in excess, so practise writing observations in this form.
Common mistakes and how to avoid them
- Imprecise observations — write "white precipitate, soluble in excess, giving a colourless solution", not just "it dissolved".
- Writing "clear" instead of "colourless" — a clear solution can still be coloured.
- Confusing Al³⁺ and Zn²⁺ — use excess ammonia to distinguish them.
- Forgetting to warm when testing for ammonium and nitrate.
- Using HCl before silver nitrate.
- Mixing up observations and inferences — record what you see, then deduce.
How this links to your Internal Assessment
Qualitative analysis is often assessed directly in CAPE practicals, where you analyse unknown salts and record observations and inferences in tables. An SBA could also apply these tests to real samples — testing local well or river water for chloride, sulfate and iron, analysing soil extracts for nitrate and ammonium, or checking fertilisers for their stated ions. Strong work uses small, controlled quantities, records observations precisely, writes ionic equations, includes confirmatory tests, and follows safety procedures for heating and handling reagents.
Exam technique for Qualitative analysis of ions
- Use observation → inference tables, keeping the two columns separate.
- Describe colours precisely and state whether precipitates are soluble or insoluble in excess.
- Name gases and give the test and result.
- Support inferences with ionic equations where asked.
- Suggest a confirmatory test if two ions could give the same result.
Quick revision summary
- Flame tests: Li⁺ crimson, Na⁺ yellow, K⁺ lilac, Ca²⁺ brick red, Ba²⁺ apple green, Cu²⁺ blue-green.
- NaOH(aq): Cu²⁺ pale blue, Fe²⁺ green, Fe³⁺ red-brown; Al³⁺, Zn²⁺, Pb²⁺ white, dissolving in excess; NH₄⁺ gives ammonia on warming.
- NH₃(aq): Zn²⁺ white, dissolves in excess; Cu²⁺ pale blue → deep blue in excess.
- Anions: CO₃²⁻ (effervescence, limewater), SO₄²⁻ (acidified BaCl₂, white), halides (acidified AgNO₃: white, cream, yellow), NO₃⁻ (Al/NaOH → NH₃).
- Gas tests for H₂, O₂, CO₂, NH₃, Cl₂, SO₂, HCl.
- Record precise observations, make inferences, confirm with further tests.