What you'll learn
Qualitative analysis means finding out which substances are present in a sample, as opposed to quantitative analysis, which finds out how much. In CSEC Chemistry it comes down to a set of standard tests: tests for gases, tests for cations using sodium hydroxide and ammonia solution, tests for anions using acid, silver nitrate and barium chloride, flame tests, and the tests for water and purity. This topic also asks you to work through an unknown salt and identify both of its ions from your own observations.
Marks here are among the easiest in the paper to secure, because the answers are fixed and learnable. The catch is that examiners want three things every time — the test, the observation and the conclusion — and an answer giving only two of the three loses marks however correct it is.
Key terms and definitions
Qualitative analysis — finding out which substances or ions are present in a sample.
Quantitative analysis — finding out how much of a substance is present.
Cation — a positively charged ion, such as a metal ion or the ammonium ion.
Anion — a negatively charged ion, such as chloride, sulfate or carbonate.
Precipitate — an insoluble solid formed when two solutions are mixed.
Flame test — identifying a metal ion from the colour it gives to a flame.
Amphoteric hydroxide — a hydroxide, such as those of aluminium and zinc, that dissolves in excess alkali.
Limewater — calcium hydroxide solution, used to test for carbon dioxide.
Anhydrous — describing a substance containing no water.
Instrumental method — a technique such as chromatography or spectroscopy, faster and more sensitive than a chemical test.
Soluble in excess — describing a precipitate that redissolves when more of the reagent is added.
Core concepts
Tests for gases
| Gas | Test | Positive result |
|---|---|---|
| Oxygen | insert a glowing splint | the splint relights |
| Hydrogen | insert a lighted splint | a squeaky pop |
| Carbon dioxide | bubble through limewater | limewater turns milky |
| Ammonia | hold damp red litmus at the mouth of the tube | litmus turns blue, pungent smell |
| Chlorine | hold damp blue litmus | it is bleached white |
| Sulfur dioxide | add acidified potassium dichromate paper | paper turns from orange to green |
Oxygen relights a glowing splint; hydrogen pops a lighted splint. Keep those two apart, because swapping them is the single commonest error in this section. Note too that the litmus must be damp in the ammonia and chlorine tests: the gas has to dissolve before it can affect the indicator, and dry litmus gives no result.
Tests for cations
Two reagents are used: sodium hydroxide solution and ammonia solution. In each case you note the colour of any precipitate and whether it dissolves when more reagent is added.
| Cation | With sodium hydroxide | Note |
|---|---|---|
| Calcium, Ca²⁺ | white precipitate, insoluble in excess | flame test: brick-red |
| Aluminium, Al³⁺ | white precipitate, dissolves in excess | |
| Zinc, Zn²⁺ | white precipitate, dissolves in excess | |
| Iron(II), Fe²⁺ | green precipitate | turns brown on standing |
| Iron(III), Fe³⁺ | red-brown precipitate | |
| Copper(II), Cu²⁺ | blue precipitate | |
| Ammonium, NH₄⁺ | no precipitate, but ammonia gas on warming | turns damp red litmus blue |
The three white precipitates are told apart by the excess step: calcium hydroxide stays insoluble, while aluminium and zinc hydroxides dissolve. Ammonia solution then separates aluminium from zinc, because zinc hydroxide dissolves in excess ammonia and aluminium hydroxide does not. Without the excess step you cannot distinguish them at all, which is why it carries a mark of its own.
Flame tests identify some metals directly: sodium gives yellow, potassium lilac, calcium brick-red, and copper blue-green. The wire is cleaned in acid between tests, because a trace of sodium left behind gives a yellow flame that masks everything else.
Tests for anions
| Anion | Test | Positive result |
|---|---|---|
| Carbonate, CO₃²⁻ | add dilute acid | fizzes, releasing CO₂ which turns limewater milky |
| Chloride, Cl⁻ | add dilute nitric acid, then silver nitrate | white precipitate |
| Sulfate, SO₄²⁻ | add dilute hydrochloric acid, then barium chloride | white precipitate |
| Nitrate, NO₃⁻ | add sodium hydroxide and aluminium foil, warm | ammonia gas released |
For chloride and sulfate the acid is added first, to destroy any carbonate ions that would otherwise produce a misleading white precipitate of their own. That detail is worth a mark and is frequently the difference between full and partial credit.
Why the precipitates form
The cation tests work because metal hydroxides are insoluble. Adding sodium hydroxide supplies hydroxide ions, which combine with the metal ion to give a coloured insoluble solid:
Cu²⁺ + 2OH⁻ → Cu(OH)₂ (blue precipitate) Fe³⁺ + 3OH⁻ → Fe(OH)₃ (red-brown precipitate)
The colours come from the metal ion itself. Transition-metal ions are coloured, which is why iron and copper give colours while calcium, aluminium and zinc all give white. The "soluble in excess" behaviour of aluminium and zinc hydroxides occurs because they are amphoteric — they react with additional alkali to form a soluble compound — whereas calcium hydroxide is not and stays as a white solid.
Understanding why a precipitate forms, rather than memorising a list of colours, makes the whole table far easier to recall under pressure, and it lets you reason about an ion you have not met before.
Writing ionic equations for the tests
Questions often ask for the ionic equation rather than a description, because it shows you understand what is actually reacting. Only the ions that take part are written; the spectator ions are left out.
Ag⁺ + Cl⁻ → AgCl (white precipitate) Ba²⁺ + SO₄²⁻ → BaSO₄ (white precipitate) Zn²⁺ + 2OH⁻ → Zn(OH)₂ (white precipitate)
In the silver nitrate test the nitrate ions and the sodium ions play no part, so they do not appear. Check that the charges balance on both sides: in the barium sulfate equation, 2+ and 2− cancel to give a neutral solid. A common slip is writing the formula of the precipitate with the wrong ratio, so count the charges rather than guessing — iron(III) needs three hydroxide ions, not two.
Testing for water and for purity
Water turns white anhydrous copper(II) sulfate blue, and turns blue cobalt chloride paper pink. Both show that water is present; neither shows that it is pure.
Pure water boils at exactly 100 °C and freezes at 0 °C at normal atmospheric pressure. Dissolved impurities raise the boiling point and lower the freezing point, so a measured boiling point above 100 °C indicates an impure sample. Distinguishing "contains water" from "is pure water" is a distinction examiners test directly.
Instrumental methods
Modern laboratories also use instrumental methods such as chromatography and spectroscopy. Their advantages over simple chemical tests are that they are faster, more sensitive and accurate, and able to work on very small samples. They can also separate and identify several substances in one run. The trade-off is cost and the need for trained operators, which is why the chemical tests above remain standard in a school laboratory.
Working through an unknown salt
A typical practical question gives an unknown salt and asks you to identify both ions. A reliable routine is:
- Note the colour of the solid and of the solution. A blue solution suggests copper(II); a green one suggests iron(II).
- Test for the cation with sodium hydroxide, recording the precipitate colour and its behaviour in excess. Confirm a metal with a flame test where useful.
- Test for the anion: dilute acid first, watching for fizzing that would indicate carbonate; then acidified silver nitrate for a halide, or acidified barium chloride for a sulfate.
- Combine the results. A blue precipitate with sodium hydroxide together with a white precipitate on adding acidified barium chloride means copper(II) sulfate.
Setting the reasoning out step by step, with the observation that led to each conclusion, is exactly what the mark scheme rewards in structured practical questions.
Worked examples
Example 1: Identifying an unknown salt (6 marks). A green solution gives a green precipitate with sodium hydroxide that does not dissolve in excess. On adding dilute hydrochloric acid followed by barium chloride, a white precipitate forms. Identify the salt and justify each conclusion.
The green solution and green precipitate with sodium hydroxide indicate iron(II), since Fe²⁺ gives a green hydroxide that is insoluble in excess. The white precipitate with acidified barium chloride indicates sulfate. The salt is therefore iron(II) sulfate. Each conclusion must be tied to its observation to gain the marks; naming the salt alone would score one at most.
Example 2: Distinguishing two white precipitates (4 marks). Two solutions both give a white precipitate with sodium hydroxide. Describe how you would establish which contains aluminium ions and which contains calcium ions.
Add excess sodium hydroxide to each. The precipitate containing aluminium dissolves in excess, because aluminium hydroxide is amphoteric; the precipitate containing calcium remains insoluble. A flame test confirms the result, since calcium gives a brick-red flame while aluminium gives no characteristic colour.
Example 3: Explaining a false positive (3 marks). A student adds silver nitrate to a solution without first adding dilute nitric acid and records a white precipitate as evidence of chloride. Explain why the conclusion is unsafe.
Carbonate ions present in the solution would react with silver nitrate to give a white precipitate of silver carbonate, which looks identical to silver chloride. Adding dilute nitric acid first destroys any carbonate, so the white precipitate that then forms can only be silver chloride. Without that step the observation does not distinguish chloride from carbonate.
Common mistakes and how to avoid them
Giving only the test or only the result. Always write test, then observation, then conclusion. Most mark schemes in this topic allocate a mark to each.
Confusing the splint tests. A glowing splint relights in oxygen; a lighted splint pops in hydrogen.
Forgetting to acidify before silver nitrate or barium chloride. Carbonate then gives a false positive, and the mark for the acid step is lost as well.
Skipping the "in excess" step. Without it the white precipitates of calcium, aluminium and zinc cannot be told apart.
Using dry litmus paper. The gas must dissolve to affect the indicator, so the paper must be damp.
Saying a test shows water is pure. Anhydrous copper(II) sulfate and cobalt chloride paper show only that water is present. Purity is established from the boiling and freezing points.
Describing a titration when asked to identify an ion. Titration is quantitative. If the question says identify, it wants a qualitative test.
Exam technique for Qualitative Analysis
Lay answers out as a short sequence or small table with the columns test, observation and conclusion. It is quicker to write, easier for the examiner to credit, and it stops you omitting one of the three.
Use the exact reagent names. "Acid" is not enough where the mark scheme specifies dilute nitric acid, and "hydroxide" is not enough where sodium hydroxide solution is meant.
Describe colours in the expected vocabulary — white, green, red-brown, blue, brick-red — rather than approximations such as "orangey" or "dark". Examiners mark against a fixed list.
Where a question supplies observations and asks for the ion, work backwards through the same routine: cation from the hydroxide result, anion from the acid and precipitate results, then combine. Where it supplies the salt and asks for observations, work forwards. Recognising which direction the question runs in saves time.
Finally, read the mark allocation. A six-mark identification question wants both ions and a justification for each; a two-mark question wants a single test and its result.
Quick revision summary
- Qualitative analysis identifies what is present; quantitative analysis measures how much.
- Always give test → observation → conclusion.
- Gases: glowing splint relights (O₂), squeaky pop (H₂), limewater milky (CO₂), damp red litmus blue (NH₃), damp blue litmus bleached (Cl₂).
- Cations with sodium hydroxide: white insoluble in excess (Ca²⁺); white soluble in excess (Al³⁺, Zn²⁺); green (Fe²⁺); red-brown (Fe³⁺); blue (Cu²⁺); ammonia on warming (NH₄⁺).
- Excess ammonia separates zinc (dissolves) from aluminium (does not).
- Flame tests: sodium yellow, potassium lilac, calcium brick-red, copper blue-green.
- Anions: acid for carbonate; acidified silver nitrate for chloride; acidified barium chloride for sulfate; NaOH with aluminium foil for nitrate.
- Acidify first for chloride and sulfate, to remove carbonate.
- Precipitates form because metal hydroxides are insoluble; aluminium and zinc hydroxides are amphoteric and so dissolve in excess alkali.
- Water is detected by anhydrous copper(II) sulfate (white to blue) and cobalt chloride paper (blue to pink); purity is shown by a boiling point of exactly 100 °C.
- Instrumental methods are faster, more sensitive and work on tiny samples.