OCR A Level Chemistry A: Qualitative analysis, 5.3.2 (H432)

This guide teaches section 5.3.2, Qualitative analysis, of the OCR A Level Chemistry A (H432) specification, Version 3.1 (May 2026), for first assessment in 2017. The section sits in Module 5 (Physical chemistry and transition elements), inside 5.3 Transition elements. It has one lettered learning outcome, 5.3.2(a): part (i) lists the anions and part (ii) the cations you must identify in an unknown compound. Its additional guidance links it to PAG4 and HSW4, and it points back to 3.1.4(a) and 5.3.1(j). H432 is the full A level. The assessment overview says Module 5 is assessed in component 01 (Periodic table, elements and physical chemistry) and in component 03 (Unified chemistry).

Use the OCR chemistry hub for the rest of the course, tick off each statement on the printable checklist, and find weak spots with the free diagnostics. Afterwards, use the revision notes to condense this page and the practice questions to check it.

What 5.3.2 asks of you

Statement	Skill required	Specification links
5.3.2(a)	Carry out qualitative analysis of ions on a test-tube scale, knowing the processes and techniques that identify the listed ions in an unknown compound	PAG4; HSW4
5.3.2(a)(i)	Identify the anions CO₃²⁻, Cl⁻, Br⁻, I⁻ and SO₄²⁻	see 3.1.4(a)
5.3.2(a)(ii)	Identify the cations NH₄⁺, Cu²⁺, Fe²⁺, Fe³⁺, Mn²⁺ and Cr³⁺	see 3.1.4(a) and 5.3.1(j)

The new content since Module 3 is the set of five transition metal ions. Section 5.3.1(j) supplies the reactions: each ion with aqueous sodium hydroxide and with aqueous ammonia, first a little and then in excess.

Technique: PAG4 and HSW4

PAG4 is “Qualitative analysis of ions”. Its two listed techniques are using apparatus for ion tests and recording qualitative observations. The suggested activity works out which ions are present in a mixture of Group 2 salts. HSW4 is carrying out experimental work with appropriate risk management.

Good practice for a cation test:

1. Dissolve a small amount of the unknown in distilled water. Note the colour of the solution before adding anything; it is your first clue.
2. Add the reagent dropwise, shaking gently after each drop, and record what forms.
3. Then add the same reagent in excess (several cm³ more) and shake again. Record whether the precipitate stays or dissolves, and the colour of any new solution.
4. Use a fresh portion of the unknown for the second reagent.
5. Leave any precipitate to stand for a few minutes and look again. Some precipitates change colour in air.

Risk management: sodium hydroxide solution is corrosive, even when dilute, so wear eye protection. Concentrated ammonia gives off an irritant gas; use the dilute solution where it does the job and keep the stock bottle closed.

5.3.2(a)(i): The anion tests

These are the tests from 3.1.4(a), restated because an unknown here may also contain a transition metal ion.

Anion	Reagent	Positive result	Ionic equation
CO₃²⁻	dilute HNO₃(aq)	effervescence; the gas turns limewater milky	CO₃²⁻ + 2H⁺ → CO₂ + H₂O
SO₄²⁻	Ba²⁺(aq), after acidifying	white precipitate	Ba²⁺(aq) + SO₄²⁻(aq) → BaSO₄(s)
Cl⁻	AgNO₃(aq), after acidifying	white precipitate, soluble in dilute NH₃(aq)	Ag⁺(aq) + Cl⁻(aq) → AgCl(s)
Br⁻	AgNO₃(aq), after acidifying	cream precipitate, soluble in concentrated NH₃(aq) only	Ag⁺(aq) + Br⁻(aq) → AgBr(s)
I⁻	AgNO₃(aq), after acidifying	yellow precipitate, insoluble in concentrated NH₃(aq)	Ag⁺(aq) + I⁻(aq) → AgI(s)

Test for carbonate first, then sulfate, then halide. The specification’s reason is that BaCO₃ and Ag₂SO₄ are both insoluble: carbonate would also precipitate with barium ions, and sulfate with silver ions. Nitric acid removes carbonate without adding a new anion. Barium nitrate is a better choice than barium chloride when a halide test will follow on the same solution.

One extra point arises with transition metals. Their carbonates do not dissolve in water. A coloured solid that will not dissolve in water but fizzes in dilute nitric acid is likely a carbonate. Dissolving it in nitric acid also puts the metal ion into solution, ready for the cation tests.

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

Warm the unknown with aqueous sodium hydroxide. Hold damp red litmus paper in the mouth of the tube without touching the glass. Ammonia gas turns it blue.

NH₄⁺(aq) + OH⁻(aq) → NH₃(g) + H₂O(l)

If the unknown also contains a transition metal ion, a precipitate forms in the tube as well. That does not spoil the test: the litmus responds only to the gas given off. Ammonium salts of transition metals, such as ammonium iron(II) sulfate, therefore give two results from one test.

5.3.2(a)(ii): The transition metal ions

Why both reagents precipitate hydroxides

Sodium hydroxide supplies OH⁻ ions directly. Aqueous ammonia is a weak base and supplies a lower concentration of OH⁻ through an equilibrium:

NH₃(aq) + H₂O(l) ⇌ NH₄⁺(aq) + OH⁻(aq)

That is still enough OH⁻ to precipitate all five metal hydroxides. So the dropwise results with the two reagents look the same. The reagents only differ when you add excess, because excess ammonia supplies NH₃ ligands for ligand substitution, while excess NaOH supplies a high concentration of OH⁻.

The specification’s guidance says only two hydroxides react further:

• with excess NaOH, only Cr(OH)₃ reacts, forming [Cr(OH)₆]³⁻
• with excess NH₃, only Cr(OH)₃ and Cu(OH)₂ react, forming [Cr(NH₃)₆]³⁺ and [Cu(NH₃)₄(H₂O)₂]²⁺

The results table

Ion and solution colour	NaOH(aq), a little	NaOH(aq), excess	NH₃(aq), a little	NH₃(aq), excess
Cu²⁺, pale blue	pale blue precipitate	no further change	pale blue precipitate	dissolves: deep blue solution
Fe²⁺, pale green	green precipitate, turning orange-brown at the surface on standing	no further change	as with NaOH	no further change
Fe³⁺, pale yellow to orange-brown	orange-brown precipitate	no further change	orange-brown precipitate	no further change
Mn²⁺, very pale pink	light brown precipitate, darkening on standing in air	no further change	as with NaOH	no further change
Cr³⁺, violet (often looks green)	grey-green precipitate	dissolves: dark green solution	grey-green precipitate	dissolves: purple solution

Hexaaqua chromium(III), [Cr(H₂O)₆]³⁺, is violet. Chromium(III) salt solutions often look green because ions such as chloride or sulfate replace some water ligands. Either colour points you to the hydroxide test, which settles it.

The Fe²⁺ and Mn²⁺ precipitates darken because oxygen in the air oxidises them. Iron(II) hydroxide becomes iron(III) hydroxide, which is why the surface turns orange-brown.

Equations

The guidance allows either form of formula for precipitation: non-complexed, such as Cu²⁺(aq) and Cu(OH)₂(s), or complexed, such as [Cu(H₂O)₆]²⁺ and Cu(OH)₂(H₂O)₄.

Precipitation (with a little NaOH or NH₃):

• Cu²⁺(aq) + 2OH⁻(aq) → Cu(OH)₂(s)
• Fe²⁺(aq) + 2OH⁻(aq) → Fe(OH)₂(s)
• Fe³⁺(aq) + 3OH⁻(aq) → Fe(OH)₃(s)
• Mn²⁺(aq) + 2OH⁻(aq) → Mn(OH)₂(s)
• Cr³⁺(aq) + 3OH⁻(aq) → Cr(OH)₃(s)

In complexed form, for example: [Fe(H₂O)₆]³⁺(aq) + 3OH⁻(aq) → Fe(OH)₃(H₂O)₃(s) + 3H₂O(l). Count the waters: six on the left, three stay in the solid and three are released.

When the precipitating reagent is ammonia, you can also write NH₃ and H₂O as reactants. For copper: [Cu(H₂O)₆]²⁺(aq) + 2NH₃(aq) → Cu(OH)₂(H₂O)₄(s) + 2NH₄⁺(aq).

Complex formation (with excess reagent):

• Cr(OH)₃(s) + 3OH⁻(aq) → [Cr(OH)₆]³⁻(aq)
• Cr(OH)₃(s) + 6NH₃(aq) → [Cr(NH₃)₆]³⁺(aq) + 3OH⁻(aq)
• Cu(OH)₂(s) + 4NH₃(aq) + 2H₂O(l) → [Cu(NH₃)₄(H₂O)₂]²⁺(aq) + 2OH⁻(aq)

Check atoms and charge. In the copper equation both sides have zero overall charge.

How much is “excess”?

Worked example 1

A student has 10.0 cm³ of 0.0500 mol dm⁻³ copper(II) nitrate solution and 0.400 mol dm⁻³ sodium hydroxide. Calculate the volume of NaOH(aq) needed just to precipitate all the copper(II) ions.

1. n(Cu²⁺) = 0.0500 × (10.0 ÷ 1000) = 5.00 × 10⁻⁴ mol
2. Cu²⁺ + 2OH⁻ → Cu(OH)₂, so n(OH⁻) = 2 × 5.00 × 10⁻⁴ = 1.00 × 10⁻³ mol
3. Volume = 1.00 × 10⁻³ ÷ 0.400 = 2.50 × 10⁻³ dm³ = 2.50 cm³

A typical drop is roughly 0.05 cm³, so this is about fifty drops. The first few drops precipitate only part of the copper. That is why the method says add dropwise, observe, then add several cm³ more to be sure the reagent really is in excess before you decide whether the precipitate dissolves.

Putting it together

Worked example 2: an insoluble pale pink solid

Solid P is pale pink and does not dissolve in water. A student records:

Test	Observation
Dilute HNO₃(aq) added to P; gas bubbled into limewater	effervescence; pale pink solid dissolves to a very pale pink solution; limewater turns milky
Ba(NO₃)₂(aq) added to part of the acidified solution	no visible change
AgNO₃(aq) added to another part	no visible change
NaOH(aq) added dropwise to a third part, then in excess	light brown precipitate, no change in excess; darkens on standing

Reasoning:

1. Insoluble in water, fizzes with acid, gas turns limewater milky: carbonate. MnCO₃(s) + 2H⁺(aq) → Mn²⁺(aq) + CO₂(g) + H₂O(l)
2. No precipitate with Ba²⁺ or Ag⁺: no sulfate, no halide.
3. Light brown precipitate that darkens in air and does not dissolve in excess NaOH: Mn²⁺. Mn²⁺(aq) + 2OH⁻(aq) → Mn(OH)₂(s)
4. P is manganese(II) carbonate, MnCO₃.

Note that the nitric acid was in excess, so the first drops of NaOH neutralise the leftover acid before any precipitate appears.

Worked example 3: two green solutions

Bottle A contains iron(II) sulfate solution; bottle B contains chromium(III) sulfate solution. Both look green. Describe a test that tells them apart.

Add aqueous sodium hydroxide dropwise, then in excess, to a portion of each.

• A gives a green precipitate that does not dissolve in excess and turns orange-brown at the surface on standing: Fe²⁺.
• B gives a grey-green precipitate that dissolves in excess to a dark green solution: Cr³⁺.

The key evidence is what happens in excess, not the shade of green. Excess ammonia would also work: Cr(OH)₃ dissolves to a purple solution, Fe(OH)₂ does not dissolve.

Common errors

• Writing “Fe(OH)₂ dissolves in excess ammonia”. Of the five hydroxides, only Cu(OH)₂ and Cr(OH)₃ dissolve in excess NH₃.
• Saying that Cu(OH)₂ dissolves in excess NaOH. Only Cr(OH)₃ does.
• Describing the Cu²⁺ result with excess ammonia as a “deep blue precipitate”. The precipitate dissolves; the deep blue is a solution.
• Giving Fe³⁺ a “brown solution” result with NaOH. It is an orange-brown precipitate.
• Recording the final colour of an Fe²⁺ test after it has stood, and reporting Fe³⁺. Look straight away and again after standing.
• Writing (aq) on a hydroxide precipitate, or leaving the charge off a complex ion.
• Leaving the water ligands out of the copper–ammonia complex. Use the formula given in the specification, [Cu(NH₃)₄(H₂O)₂]²⁺.
• Forgetting that an excess of acid from the carbonate test must be neutralised before hydroxides can precipitate.

Next steps

• Tables and a self-test to revise from: 5.3.2 revision notes
• Original questions with marked answers: 5.3.2 practice questions
• Oxidation numbers and half-equations for the colour changes in air: Redox
• Concentration and mole calculations: Amount of substance
• Recording observations for PAG work: Practical Endorsement skills

Official syllabus

OCR A Level Chemistry A (H432) specification, Version 3.1 (May 2026), for first assessment in 2017, section 5.3.2 Qualitative analysis, learning outcome (a)(i) and (a)(ii), with its additional guidance (PAG4, HSW4) and the linked statements 3.1.4(a) and 5.3.1(j). Published by OCR.