AQA GCSE Chemistry 8462: Chemical analysis – Revision Notes

These notes condense Topic 8, Chemical analysis (sections 4.8.1 to 4.8.3), of the AQA GCSE Chemistry (8462) specification, for teaching from September 2016 with exams from June 2018 (version 1.1). The topic is examined on Paper 2 at Foundation and Higher tier, and none of it is Higher tier only. It includes Required practical 6 (chromatography) and Required practical 7 (identifying ions).

For full explanations and worked examples, use the Chemical analysis study guide. Then test yourself with the Chemical analysis practice questions. The AQA GCSE Chemistry hub and the printable checklist show where this topic sits in the course.

Key definitions

• Pure substance (chemistry): a single element or compound, not mixed with any other substance.
• Pure (everyday): nothing added; unadulterated, in its natural state (e.g. “pure milk”).
• Formulation: a mixture designed as a useful product, with components in carefully measured quantities, each with a particular purpose.
• Stationary phase: the part that does not move (the paper).
• Mobile phase: the part that moves (the solvent).
• Rf value: distance moved by substance ÷ distance moved by solvent.
• Precipitate: an insoluble solid formed when two solutions react.
• Instrumental method: analysis by machine; accurate, sensitive and rapid.

Purity and formulations

Evidence	Pure	Impure (mixture)
Melting point	sharp, at the known value	lower, over a range
Boiling point	sharp, at the known value	over a range
Chromatogram	one spot in every solvent	may give two or more spots

Formulations named in the spec: fuels, cleaning agents, paints, medicines, alloys, fertilisers, foods.

How to spot one in a question: it is a mixture, each part has a job, and the amounts are measured and fixed.

Chromatography (Required practical 6)

Method in steps:

1. Pencil origin line about 1–2 cm from the bottom of the paper.
2. Small, concentrated spots of each sample on the line.
3. Solvent level below the origin line; lid on the container.
4. Remove before the solvent reaches the top; mark the solvent front straight away.
5. Measure both distances from the origin; spot distance to the centre of the spot.
6. Rf = substance distance ÷ solvent distance.

Worked reminder: solvent moved 6.0 cm; spot centre 4.2 cm from the origin. Rf = 4.2 ÷ 6.0 = 0.70.

• Same Rf as a reference in the same solvent → likely the same compound.
• Different solvent → different Rf values.
• Separation depends on the distribution of each substance between the mobile and stationary phases.
• Rf is always less than 1 and has no units. Quote it to a sensible number of significant figures.

Reading a chromatogram, in steps:

1. Count the spots for each sample. One spot suggests a single substance; two or more spots mean a mixture.
2. For identification, compare horizontally: a spot at the same height as a reference spot (same Rf, same solvent) matches that reference.
3. A spot that matches no reference is an unknown substance not in the reference set.
4. A spot left on the origin line did not dissolve in that solvent – try another solvent.
5. To claim a substance is pure, it should give one spot in every solvent tried, not just one.

Tests for gases

Gas	Test	Result
H2	burning splint at the mouth of the tube	squeaky pop
O2	glowing splint in the gas	relights
CO2	limewater (calcium hydroxide solution)	turns milky
Cl2	damp litmus paper	bleached white

Flame tests

Ion	Colour
Li⁺	crimson
Na⁺	yellow
K⁺	lilac
Ca²⁺	orange-red
Cu²⁺	green

Mixtures: one colour can mask another. Flame colours of other metals are not needed.

Hydroxide precipitates (add NaOH solution)

Ion	Precipitate	Excess NaOH
Al³⁺	white	dissolves
Ca²⁺	white	no change
Mg²⁺	white	no change
Cu²⁺	blue	no change
Fe²⁺	green	no change
Fe³⁺	brown	no change

Writing hydroxide equations, in steps:

1. Charge on the metal ion = number of OH⁻ needed: Cu²⁺ → Cu(OH)2, Fe³⁺ → Fe(OH)3.
2. Put that number in front of NaOH.
3. Balance the sodium salt on the right.

CuSO4 + 2NaOH → Cu(OH)2 + Na2SO4
MgCl2 + 2NaOH → Mg(OH)2 + 2NaCl
Fe²⁺ + 2OH⁻ → Fe(OH)2

Anion tests

Ion	Add	Result
CO3²⁻	dilute acid	fizzing; gas turns limewater milky
Cl⁻	dilute nitric acid, then silver nitrate	white precipitate
Br⁻	dilute nitric acid, then silver nitrate	cream precipitate
I⁻	dilute nitric acid, then silver nitrate	yellow precipitate
SO4²⁻	dilute hydrochloric acid, then barium chloride	white precipitate

Why acidify first? Carbonate ions also react with silver nitrate and barium chloride to form precipitates. Adding the acid first removes them as carbon dioxide, so a precipitate can only come from the ion being tested. Nitric acid is used for halides because hydrochloric acid would add chloride ions and give a false positive for chloride.

Worked reminder: a carbonate fizzes with acid because carbon dioxide forms, e.g. CaCO3 + 2HCl → CaCl2 + H2O + CO2. The gas turns limewater milky.

Identifying an unknown single compound (Required practical 7), in steps:

1. Flame test → Li⁺, Na⁺, K⁺, Ca²⁺ or Cu²⁺?
2. NaOH solution → which hydroxide colour? If white, add excess to check for Al³⁺.
3. Dilute acid → fizzing and milky limewater means carbonate.
4. New sample: nitric acid + silver nitrate → halide?
5. New sample: hydrochloric acid + barium chloride → sulfate?
6. Combine the cation and anion to name the compound.

Instrumental methods

• Advantages over chemical tests: accurate, sensitive, rapid.
• Flame emission spectroscopy: sample in a flame → light through a spectroscope → line spectrum.
• Match the lines to a reference set to identify the metal ions – works for mixtures.
• Line intensity gives concentration, read from a calibration graph or table of standards.

Interpreting a spectrum, in steps:

1. Line up the sample spectrum with each reference spectrum on the same scale.
2. A metal ion is present only if all of its reference lines appear in the sample.
3. Lines left over after matching belong to another ion – look for a second match.
4. For concentration, read the sample’s intensity against the standards (same line, same conditions) and scale any dilution back up.

Worked reminder: a sample shows lines at positions that match every line of reference metal ion P and every line of reference Q, plus nothing else. The sample contains P and Q ions only.

Must-know distinctions

• Burning vs glowing splint: hydrogen pops vs oxygen relights.
• Pure (chemistry) vs pure (everyday): single substance vs nothing added.
• White precipitates: Al³⁺, Ca²⁺, Mg²⁺ with NaOH; AgCl with silver nitrate; BaSO4 with barium chloride. Read which reagent was used.
• Cream vs yellow: silver bromide vs silver iodide.
• Flame test vs flame emission spectroscopy: colour by eye, single ions, masked in mixtures vs line spectrum, identifies ions in mixtures and measures concentration.

Quick self-test

1. Define a pure substance in chemistry.
2. Sample Z melts between 76 °C and 81 °C. Is it pure? Give a reason.
3. The solvent moved 7.5 cm and a spot moved 3.0 cm. Calculate Rf.
4. In another run the solvent moved 6.0 cm and a spot’s Rf was 0.45. How far did the spot move?
5. Why is the origin line drawn in pencil?
6. Describe the test for oxygen and its result.
7. What flame colour does a potassium compound give?
8. A solution gives a white precipitate with NaOH that dissolves in excess. Name the ion.
9. Write a balanced equation for copper(II) sulfate reacting with sodium hydroxide.
10. Which reagents test for a sulfate ion, and what is seen?
11. Standard potassium solutions of 0, 4, 8 and 12 mg/dm³ give intensities 0, 18, 36 and 54. A sample gives 45. Find its concentration.
12. Give two advantages of instrumental methods over chemical tests.

Answers

1. A single element or compound, not mixed with any other substance.
2. No. It melts over a range, not at one specific temperature.
3. 3.0 ÷ 7.5 = 0.40.
4. 0.45 × 6.0 = 2.7 cm.
5. Pencil does not dissolve in the solvent; ink would separate and move up the paper.
6. Insert a glowing splint into the gas; it relights.
7. Lilac.
8. Aluminium, Al³⁺.
9. CuSO4 + 2NaOH → Cu(OH)2 + Na2SO4.
10. Dilute hydrochloric acid, then barium chloride solution; white precipitate.
11. 54 ÷ 12 = 4.5 units per mg/dm³; 45 ÷ 4.5 = 10 mg/dm³.
12. Any two: more accurate, more sensitive (small samples), more rapid.

Where marks are usually lost

• Measuring Rf from the bottom of the paper instead of the origin line.
• Giving Rf with units, or as a value above 1 (the fraction was upside down).
• Writing the solvent level above the origin line in a method.
• “Glowing splint pops” – mixing up the hydrogen and oxygen tests.
• Using dry litmus for chlorine, or saying it turns red.
• Stating “a precipitate forms” without its colour.
• Forgetting that aluminium hydroxide is the only one that dissolves in excess NaOH.
• Adding hydrochloric acid before silver nitrate in a halide test.
• Naming a compound from one test only – you need both the cation and the anion.

Official syllabus

AQA GCSE Chemistry (8462) specification, for teaching from September 2016, GCSE exams June 2018 onwards, version 1.1, published by AQA – section 4.8 Chemical analysis.