OCR A Level Chemistry A: Qualitative analysis (H432)

This study guide teaches section 3.1.4, Qualitative analysis, of the OCR A Level Chemistry A (H432) specification, Version 3.1 (May 2026), for first assessment in 2017. Module 3 (Periodic table and energy) is its home, and it contains one lettered learning outcome, 3.1.4(a), with an anion part (i) and a cation part (ii). Both parts, the additional guidance and the PAG4 link are covered below. H432 is the full A level. According to the specification’s assessment overview, Module 3 is examined in component 01, Periodic table, elements and physical chemistry, and in component 03, Unified chemistry.

For the rest of the course, go to the OCR chemistry hub. Mark what you know on the printable checklist, and spot gaps with the free diagnostics.

What 3.1.4 asks of you

Statement	Skill	Additional guidance (specification)
3.1.4(a)(i)	Identify the anions CO₃²⁻ (reaction with H⁺(aq) to form CO₂ gas), SO₄²⁻ (precipitation with Ba²⁺(aq)) and Cl⁻, Br⁻, I⁻ (see 3.1.3(g)) in an unknown compound, on a test-tube scale	The sequence of tests is carbonate, then sulfate, then halide; the specification notes that BaCO₃ and Ag₂SO₄ are both insoluble. Cross-reference to 2.1.4(c). PAG4; HSW4
3.1.4(a)(ii)	Identify the cation NH₄⁺ by reaction with warm NaOH(aq), forming NH₃	PAG4; HSW4

PAG4 is “Qualitative analysis of ions”. The specification lists its skills as using apparatus for qualitative tests for ions and making and recording qualitative observations. Its suggested activity is identifying the anions and cations in a mixture of Group 2 salts. Later, in 5.3.2, the same tests return alongside tests for transition metal ions.

The outcome asks for “processes and techniques”, so you need more than a list of colours. You need to know which reagent to add, in what order, what you see, what equation explains it, and why the order matters.

Working on a test-tube scale

Each test uses small amounts: a spatula tip of solid dissolved in about 2 cm³ of distilled water, with reagents added a few drops at a time from a dropping pipette.

• Use distilled or deionised water. Tap water can contain chloride and sulfate ions, so it can give a false positive.
• Take a fresh portion of the unknown for each separate test unless the plan says to carry on with the same solution.
• Run a blank (distilled water plus the same reagents) if a result is faint or unclear.
• Warm with a water bath rather than a direct flame when only gentle warming is needed.
• Record what you see in precise words: “white precipitate”, “effervescence”, “colourless gas”, “no visible change”. Write the observation first and the conclusion second.

3.1.4(a)(i): Testing for carbonate ions

Reagent: any dilute acid supplies H⁺(aq). When further tests will follow on the same solution, use dilute nitric acid, for the reasons explained in the sequence section below.

Observation: effervescence (fizzing). Bubble the gas through limewater, a solution of calcium hydroxide, and it turns milky as a white precipitate of calcium carbonate forms. That confirms the gas is carbon dioxide.

Equations:

• Ionic: CO₃²⁻(aq) + 2H⁺(aq) → CO₂(g) + H₂O(l)
• Full, for potassium carbonate: K₂CO₃(aq) + 2HNO₃(aq) → 2KNO₃(aq) + CO₂(g) + H₂O(l)
• Limewater: Ca(OH)₂(aq) + CO₂(g) → CaCO₃(s) + H₂O(l)

This is the neutralisation of a carbonate by an acid from 2.1.4(c), now used as a test. Fizzing alone is not proof. Other reactions also release gases, so the limewater result is what identifies CO₂.

Worked example 1: a carbonate and a gas volume

A student adds excess dilute nitric acid to 0.691 g of potassium carbonate. Calculate the volume of carbon dioxide formed at room temperature and pressure. Use Ar values K = 39.1, C = 12.0, O = 16.0. Take the molar gas volume at RTP as 24.0 dm³ mol⁻¹.

1. Mr(K₂CO₃) = 2(39.1) + 12.0 + 3(16.0) = 138.2
2. n(K₂CO₃) = 0.691 ÷ 138.2 = 0.00500 mol
3. One mole of carbonate gives one mole of CO₂, so n(CO₂) = 0.00500 mol.
4. Volume = 0.00500 × 24.0 = 0.120 dm³ = 120 cm³

3.1.4(a)(i): Testing for sulfate ions

Reagent: aqueous barium ions, Ba²⁺(aq), as barium nitrate or barium chloride solution. Add it to the solution after it has been acidified with dilute nitric acid.

Observation: a white precipitate of barium sulfate.

Equation: Ba²⁺(aq) + SO₄²⁻(aq) → BaSO₄(s)

Why acidify first? Barium carbonate is also a white solid that does not dissolve in water. If carbonate ions are still in the solution, Ba²⁺ precipitates them too and you cannot tell the two precipitates apart. Adding acid first destroys any carbonate as CO₂, so a white precipitate that forms afterwards must be barium sulfate.

Worked example 2: how much sulfate?

Because BaSO₄ is insoluble, the precipitate can be filtered, dried and weighed. A 0.800 g sample of lawn feed contains ammonium sulfate, (NH₄)₂SO₄, and no other sulfate. It is dissolved, acidified and treated with excess barium nitrate. The dry precipitate has a mass of 1.272 g. Calculate the percentage by mass of ammonium sulfate. Use Ar values Ba = 137.3, S = 32.1, O = 16.0, N = 14.0, H = 1.0.

1. Mr(BaSO₄) = 137.3 + 32.1 + 4(16.0) = 233.4
2. n(BaSO₄) = 1.272 ÷ 233.4 = 0.00545 mol
3. One (NH₄)₂SO₄ gives one SO₄²⁻, which gives one BaSO₄. So n((NH₄)₂SO₄) = 0.00545 mol.
4. Mr((NH₄)₂SO₄) = 2(14.0) + 8(1.0) + 32.1 + 4(16.0) = 132.1
5. Mass = 0.00545 × 132.1 = 0.720 g
6. Percentage = 0.720 ÷ 0.800 × 100 = 90.0%

Carry unrounded values through the working. Round only the final answer.

3.1.4(a)(i): Testing for halide ions

This test is set out in 3.1.3(g), and 3.1.4(a) brings it into the analysis sequence.

Reagent: first make the solution acidic with dilute nitric acid; then add AgNO₃(aq), silver nitrate solution.

Observations and follow-up with aqueous ammonia:

Ion in the unknown	Solid formed with Ag⁺(aq)	Its colour	Then dilute NH₃(aq)	Then concentrated NH₃(aq)
chloride	AgCl	white	precipitate disappears	precipitate disappears
bromide	AgBr	cream	precipitate remains	precipitate disappears
iodide	AgI	yellow	precipitate remains	precipitate remains

Ionic equations:

• Ag⁺(aq) + Cl⁻(aq) → AgCl(s)
• Ag⁺(aq) + Br⁻(aq) → AgBr(s)
• Ag⁺(aq) + I⁻(aq) → AgI(s)

White, cream and yellow can look alike in a test tube, especially in small amounts. The ammonia step separates them. The specification says complexes with ammonia are not required other than observations, so you need to know what dissolves, not the formula of what forms.

The nitric acid has the same job as before. Silver carbonate is insoluble, so carbonate ions left in the solution would give a precipitate with Ag⁺ and confuse the result.

3.1.4(a)(ii): Testing for ammonium ions

Method: add aqueous sodium hydroxide to the solid or solution and warm gently. Hold a piece of damp red litmus paper at the mouth of the test tube, without touching the glass.

Observation: the litmus paper turns blue. You may also notice the sharp smell of ammonia, but do not sniff the tube directly.

Equations:

• Ionic: NH₄⁺(aq) + OH⁻(aq) → NH₃(g) + H₂O(l)
• Full, for ammonium chloride: NH₄Cl(aq) + NaOH(aq) → NaCl(aq) + NH₃(g) + H₂O(l)

Two details carry the explanation. Warming matters because ammonia is very soluble in water; at room temperature most of it stays dissolved, and heat drives the gas out. The litmus must be damp because ammonia only shows as an alkali when it dissolves in water and releases OH⁻ ions. Dry paper may not change. If the paper touches the tube, sodium hydroxide on the glass turns it blue and gives a false positive.

The sequence: carbonate, then sulfate, then halide

When one solution may contain several anions, the specification sets the order of tests as carbonate, sulfate, then halide. The reason is that each later reagent also precipitates an ion tested earlier.

Step	Reagent	Positive result	Why it must come at this point
1	Dilute HNO₃(aq), until fizzing stops	Effervescence; gas turns limewater milky	Removes CO₃²⁻, which would otherwise precipitate with both Ba²⁺ and Ag⁺
2	Ba(NO₃)₂(aq), in excess	White precipitate (BaSO₄)	BaCO₃ is insoluble, so carbonate must already be gone. Ag₂SO₄ is insoluble too, so sulfate must be removed before step 3
3	AgNO₃(aq), then NH₃(aq)	White, cream or yellow precipitate	Only halide ions are left to precipitate with Ag⁺

If step 2 gives a precipitate, filter it off and carry on with the filtrate. That removes the sulfate before silver ions are added.

Choice of acid. Hydrochloric acid would add Cl⁻ ions, so step 3 would always show chloride. Sulfuric acid would add SO₄²⁻ ions, so step 2 would always show sulfate. Nitric acid adds only nitrate ions, which none of the reagents precipitates.

Choice of barium salt. Barium chloride would also add Cl⁻ ions and spoil the halide test. Barium nitrate does not. When you only need to test for sulfate and no halide test follows, either barium salt works.

Worked example 3: identifying an unknown salt

A white solid, W, is a single compound. A student records these observations.

Test	Observation
W dissolved in distilled water, then dilute HNO₃(aq) added	no effervescence
Ba(NO₃)₂(aq) added to the acidified solution	no visible change
AgNO₃(aq) added to the same solution	yellow precipitate
Concentrated NH₃(aq) added to the precipitate	precipitate does not dissolve
Fresh W warmed with NaOH(aq), damp red litmus held at the mouth	litmus turns blue

Reasoning, step by step:

1. No effervescence with acid: no carbonate.
2. No precipitate with Ba²⁺ in acid: no sulfate.
3. Yellow precipitate with Ag⁺, insoluble in concentrated ammonia: iodide, I⁻.
4. Alkaline gas on warming with NaOH: ammonia, so the cation is NH₄⁺.
5. W is ammonium iodide, NH₄I.

Equations: Ag⁺(aq) + I⁻(aq) → AgI(s), and NH₄⁺(aq) + OH⁻(aq) → NH₃(g) + H₂O(l).

Writing observations that earn credit

• Name the state and the colour: “white precipitate”, not “goes white” or “goes cloudy”.
• For a gas, describe what you see (effervescence, bubbles) and then the test result, such as “limewater turns milky”.
• “No visible change” is an observation. “Nothing happens” is weaker, because something may happen that you cannot see.
• Keep observation and inference apart. “Cream precipitate” is what you see; “bromide present” is what you conclude.

Common errors

• Using hydrochloric acid to acidify before a halide test, then reporting chloride that came from the acid.
• Using barium chloride for the sulfate test and then testing the same solution for halides.
• Adding barium ions before checking for carbonate, then reading a white BaCO₃ precipitate as sulfate.
• Saying that AgBr dissolves in dilute ammonia, or that AgI dissolves in concentrated ammonia.
• Using dry litmus, blue litmus, or forgetting to warm in the ammonium test.
• Writing BaSO₄ or the silver halides with (aq) instead of (s) in the ionic equation.
• Writing the carbonate ionic equation with only one H⁺.

Next steps

• Condensed tables, method steps and a self-test: Qualitative analysis revision notes
• Original questions with mark-by-mark answers: Qualitative analysis practice questions
• How carbonates neutralise acids: Acids
• For mole, mass and gas-volume methods: Amount of substance
• Recording observations in practical work: Practical skills for the Practical Endorsement

Official syllabus

OCR A Level Chemistry A (H432) specification, Version 3.1 (May 2026), for first assessment in 2017, section 3.1.4 Qualitative analysis, learning outcome (a)(i) and (ii), with its additional guidance and PAG4. Published by OCR.