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You’re handed a small tube of white or coloured powder and told it’s a single ionic compound. Your task: name it. Because every salt is made of a cation (positive ion) and an anion (negative ion), the job splits neatly into two halves. This guide gives a practical order of tests for each half, the observations to expect, and the reasons behind them, so you can follow the logic rather than memorise a table.
The scheme covers the ions most commonly met in school and first-year university courses. For the bigger picture of how analysts approach unknowns, see how chemists identify an unknown compound.
Before you start
- Wear eye protection. Sodium hydroxide and concentrated acids are corrosive, and silver nitrate stains skin.
- Use small amounts: a spatula tip of solid dissolved in about 2 cm³ of distilled water is plenty.
- Use distilled or deionised water throughout. Tap water contains chloride, sulfate and calcium ions that can give false positives.
- Record every observation, including “no change”. A negative result is still evidence.
Stage 1: Look at the sample
Colour is a free first clue.
| Colour of solid or solution | Possible ion |
|---|---|
| blue | Cu²⁺ |
| pale green | Fe²⁺ or Ni²⁺ |
| yellow-brown or orange-brown | Fe³⁺ |
| pink | Co²⁺, or Mn²⁺ (very pale) |
| white solid, colourless solution | s-block ions, NH₄⁺, Al³⁺, Zn²⁺, Pb²⁺ and most common anions |
Note whether the solid dissolves in water. Many carbonates, and most hydroxides of transition metals, don’t. If it’s insoluble, try dilute nitric acid (and watch for fizzing, which already hints at carbonate). The solubility rules help here.
Stage 2: Identify the cation
Test A: Flame test (for metal ions)
Dip a clean nichrome or platinum wire loop, moistened with dilute hydrochloric acid, into the solid and hold it at the edge of a blue Bunsen flame.
| Flame colour | Ion |
|---|---|
| yellow-orange (strong, persistent) | Na⁺ |
| lilac | K⁺ |
| crimson red | Li⁺ |
| orange-red (brick red) | Ca²⁺ |
| red | Sr²⁺ |
| apple green | Ba²⁺ |
| blue-green | Cu²⁺ |
The colours come from electrons dropping back to lower energy levels and emitting light of specific wavelengths. Sodium is a common contaminant and its yellow can mask lilac potassium; viewing through blue cobalt glass filters out the yellow. See flame test colours for the full explanation.
Test B: Add sodium hydroxide solution
To about 2 cm³ of the solution, add dilute sodium hydroxide drop by drop, then add excess. Many metal hydroxides are insoluble and have characteristic colours.
| Observation with a few drops | With excess NaOH | Ion |
|---|---|---|
| blue precipitate | stays | Cu²⁺ (Cu(OH)₂) |
| green precipitate, turns orange-brown on standing in air | stays | Fe²⁺ (Fe(OH)₂ oxidises to Fe(OH)₃) |
| orange-brown (rust-coloured) precipitate | stays | Fe³⁺ (Fe(OH)₃) |
| white precipitate | dissolves to a colourless solution | Al³⁺, Zn²⁺ or Pb²⁺ (amphoteric hydroxides) |
| white precipitate | stays | Ca²⁺ or Mg²⁺ |
| no precipitate | — | Na⁺, K⁺, NH₄⁺ (or Ba²⁺ in dilute solution) |
Why do some precipitates dissolve in excess? Aluminium, zinc and lead hydroxides are amphoteric: they react with extra hydroxide to form soluble complex ions such as [Al(OH)₄]⁻. See amphoteric substances.
Test C: Separating the white precipitates
If Test B gave a white precipitate, a second reagent narrows it down.
Dissolved in excess NaOH (Al³⁺, Zn²⁺ or Pb²⁺)? Repeat with aqueous ammonia:
- Zn²⁺: white precipitate that dissolves in excess ammonia (forms [Zn(NH₃)₄]²⁺).
- Al³⁺: white precipitate that does not dissolve in excess ammonia.
- Pb²⁺: white precipitate, insoluble in excess ammonia. Confirm with potassium iodide, which gives a bright yellow precipitate of PbI₂.
Stayed in excess NaOH (Ca²⁺ or Mg²⁺)? Use the flame test: Ca²⁺ gives orange-red; Mg²⁺ gives no colour. Or add aqueous ammonia: Mg²⁺ gives a white precipitate, while Ca²⁺ gives little or none because calcium hydroxide is more soluble.
Test D: The ammonium ion
If there was no precipitate with NaOH, test for NH₄⁺. Warm the solution with sodium hydroxide and hold damp red litmus paper at the mouth of the tube (not touching it).
NH₄⁺ + OH⁻ → NH₃ + H₂O
Ammonia gas turns damp red litmus blue, and it has a sharp, characteristic smell (don’t sniff directly). See litmus paper.
Stage 3: Identify the anion
The order of anion tests matters, because some reagents would interfere with later tests. A sensible order is: carbonate → sulfate → halide → nitrate.
Test 1: Carbonate (CO₃²⁻)
Add dilute hydrochloric or nitric acid to the solid or solution. Fizzing indicates a gas. Bubble the gas through limewater; if it turns milky, the gas is carbon dioxide and the anion is carbonate (or hydrogencarbonate).
CO₃²⁻ + 2H⁺ → CO₂ + H₂O Ca(OH)₂ + CO₂ → CaCO₃(s) + H₂O
Why first? Carbonate reacts with barium and silver ions to form precipitates too, which would give false positives in the next tests. Acidifying first removes it.
Test 2: Sulfate (SO₄²⁻)
To a fresh sample, add dilute hydrochloric acid, then barium chloride (or barium nitrate) solution. A white precipitate of barium sulfate means sulfate is present.
Ba²⁺ + SO₄²⁻ → BaSO₄(s)
The acid is there to destroy any carbonate or sulfite, whose barium salts are also white and insoluble.
Test 3: Halides (Cl⁻, Br⁻, I⁻)
To a fresh sample, add dilute nitric acid, then silver nitrate solution.
| Precipitate | Halide | Solubility in ammonia |
|---|---|---|
| white | Cl⁻ (AgCl) | dissolves in dilute ammonia |
| cream | Br⁻ (AgBr) | dissolves only in concentrated ammonia |
| yellow | I⁻ (AgI) | insoluble even in concentrated ammonia |
Use nitric acid, not hydrochloric acid: hydrochloric acid would add chloride ions and always give a white precipitate. Because the colours can be hard to tell apart, the ammonia step is the confirmatory test. Fluoride gives no precipitate, because silver fluoride is soluble. See halogens.
Test 4: Nitrate (NO₃⁻)
Nitrates give no precipitate in any of the tests above, since nearly all nitrates are soluble. To test, add sodium hydroxide solution and a little aluminium powder (Devarda’s alloy works even better), then warm carefully. The nitrate is reduced to ammonia, which turns damp red litmus blue.
This test only works if ammonium ions have already been ruled out (otherwise they’d give ammonia too).
The whole scheme at a glance
Cation:
- Colour of the solid or solution.
- Flame test.
- NaOH, a few drops then excess: coloured precipitate → Cu²⁺, Fe²⁺ or Fe³⁺; white and dissolves → Al³⁺, Zn²⁺, Pb²⁺; white and stays → Ca²⁺, Mg²⁺.
- Ammonia or KI to separate the white precipitates.
- No precipitate → warm with NaOH and test for NH₃.
Anion:
- Dilute acid + limewater → carbonate.
- HCl then BaCl₂ → white precipitate = sulfate.
- HNO₃ then AgNO₃ → white/cream/yellow = Cl⁻/Br⁻/I⁻; confirm with ammonia.
- NaOH + Al, warm → NH₃ = nitrate.
Worked example
A green solid dissolves in water to give a pale green solution.
- Colour: pale green, so Fe²⁺ or Ni²⁺.
- NaOH: green precipitate that slowly turns orange-brown at the surface. That’s Fe(OH)₂ oxidising in air, so the cation is Fe²⁺.
- Dilute HCl: no fizzing, so no carbonate.
- HCl then BaCl₂: white precipitate, so sulfate is present.
The compound is iron(II) sulfate, FeSO₄ (commonly the hydrate FeSO₄·7H₂O, see hydrates).
Common mistakes
- Adding a whole test tube of NaOH at once, so you never see the precipitate form and then dissolve.
- Using HCl before the silver nitrate test.
- Skipping the carbonate test and then misreading a barium carbonate precipitate as sulfate.
- Forgetting that tap water and dirty glassware can contain chloride.
Key takeaways
- Identify the cation and anion separately.
- Flame tests and NaOH precipitates identify most cations; excess NaOH and ammonia separate the white ones.
- Test anions in order: carbonate, sulfate, halide, nitrate, so earlier ions don’t interfere.
- Choose the acid carefully: HCl before barium chloride, HNO₃ before silver nitrate.
- Test yourself with the qualitative analysis practice questions.
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