OxfordAQA International GCSE Chemistry: Identification of ions (9202) – Revision Notes

These revision notes condense section 3.4.3 Identification of ions of the OxfordAQA International GCSE Chemistry (9202) specification, Version 6.2, for exams May/June 2018 onwards. They cover statements a to f and the flame test required practical. There are no tiers, so you need every line. The full teaching, with worked examples, is in the identification of ions study guide.

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The whole unit on one table

Ion	Reagent(s)	Positive result
Li⁺	flame test	crimson flame
Na⁺	flame test	yellow flame
K⁺	flame test	lilac flame
Ca²⁺	flame test	red flame
Ba²⁺	flame test	green flame
Al³⁺	NaOH(aq), then excess	white precipitate, dissolves in excess
Ca²⁺	NaOH(aq), then excess	white precipitate, stays
Mg²⁺	NaOH(aq), then excess	white precipitate, stays
Cu²⁺	NaOH(aq)	blue precipitate
Fe²⁺	NaOH(aq)	green precipitate
Fe³⁺	NaOH(aq)	brown precipitate
CO₃²⁻	dilute acid; gas into limewater	fizzing; limewater turns cloudy white
Cl⁻	dilute HNO₃, then AgNO₃(aq)	white precipitate (silver chloride)
Br⁻	dilute HNO₃, then AgNO₃(aq)	cream precipitate (silver bromide)
I⁻	dilute HNO₃, then AgNO₃(aq)	yellow precipitate (silver iodide)
SO₄²⁻	dilute HCl, then BaCl₂(aq)	white precipitate (barium sulfate)

Memory hook for the flames, in table order: “Crimson Yachts Like Red and Green” (Li, Na, K, Ca, Ba).

Flame tests (3.4.3 a and the required practical)

Method in steps

1. Blue, roaring Bunsen flame (air hole open).
2. Clean a nichrome wire loop: dip in dilute hydrochloric acid, heat, repeat until the flame shows no colour.
3. Pick up a little solid on the moistened loop.
4. Hold it at the edge of the flame; record the colour at once.
5. Clean the wire (or use a new loop) between samples.

Risks and controls: hot wire and burner (handle by the holder, cool on a heat-proof mat); acid can spit (eye protection); barium compounds harmful (wash hands).

Evaluation points

• Sodium contamination gives a strong yellow that can hide other colours, especially lilac.
• A mixture of metal ions may show one colour only, so a flame test suits a single compound.
• A sooty yellow flame (air hole closed) makes every result look yellow.
• Repeat with a freshly cleaned wire to check a result.

Sodium hydroxide tests (3.4.3 b and c)

Method in steps

1. Put about 2 cm³ of the unknown solution in a test tube.
2. Add sodium hydroxide solution a few drops at a time. Note the colour of any precipitate.
3. For a white precipitate, keep adding sodium hydroxide until it is in excess. Note whether the precipitate dissolves.

Must-know distinctions

• White three: Al³⁺, Ca²⁺, Mg²⁺. Only aluminium hydroxide dissolves in excess.
• Ca²⁺ or Mg²⁺? NaOH cannot tell them apart. Use a flame test: red for calcium, no characteristic colour for magnesium.
• Iron(II) or iron(III)? Green precipitate for iron(II), brown for iron(III). A green precipitate left in air slowly turns brown, so record the first colour.
• Blue is copper(II) only.

Sample equation, aluminium sulfate with a little sodium hydroxide:

Al₂(SO₄)₃(aq) + 6NaOH(aq) → 2Al(OH)₃(s) + 3Na₂SO₄(aq)

Check the balance: 2 Al, 3 SO₄ and 6 Na on each side; 6 OH⁻ go into 2 Al(OH)₃.

Carbonates (3.4.3 d)

Method in steps

1. Add dilute acid (hydrochloric is fine) to the unknown.
2. Look for fizzing.
3. Pass the gas into limewater.
4. Limewater goes cloudy white: a white precipitate of calcium carbonate forms, so the gas is carbon dioxide and the unknown is a carbonate.

General word equation: carbonate + acid → salt + water + carbon dioxide.

The gas tests themselves (section 3.4.2) are in the purity, chromatography and gas tests guide.

Halides (3.4.3 e)

Method in steps

1. Acidify the solution with dilute nitric acid.
2. Add silver nitrate solution.
3. Colour of precipitate: white (Cl⁻), cream (Br⁻), yellow (I⁻).

Why each reagent?

Choice	Reason
Acid added first	removes carbonate ions, which would also give a precipitate with silver ions
Nitric, not hydrochloric	hydrochloric acid adds chloride ions, giving a false white precipitate
Silver nitrate	silver ions form insoluble silver halides with halide ions

Sulfates (3.4.3 f)

Method in steps

1. Add dilute hydrochloric acid to the solution.
2. Add barium chloride solution.
3. White precipitate of barium sulfate = sulfate ions present.

Carbonate ions left in the sample would form barium carbonate, also white, so the acid is there to destroy them first. Sulfuric acid must not be used: it contains sulfate ions.

Must-know distinctions

Easily confused	How to keep them apart
Green flame vs green precipitate	Green flame = barium ion. Green precipitate with NaOH = iron(II) ion
Crimson vs red flame	Crimson = lithium. Red = calcium
Acid for the halide test vs the sulfate test	Nitric acid with silver nitrate; hydrochloric acid with barium chloride
White precipitate with AgNO₃ vs with BaCl₂	White with silver nitrate = chloride. White with barium chloride = sulfate
Fizzing vs proof of carbonate	Fizzing alone could be hydrogen from a metal; limewater turning cloudy confirms CO₂

Planning tests for an unknown salt

1. Dissolve the solid in distilled water.
2. Divide the solution into separate portions, one per test.
3. Positive ion: flame test (Group 1 metals, calcium, barium) or NaOH (aluminium, calcium, magnesium, copper(II), iron(II), iron(III)).
4. Negative ion: acid and limewater (carbonate), acidified silver nitrate (halides), acidified barium chloride (sulfate).
5. Combine the two ions into a formula using their charges.

Never run the halide test on a portion that has already had barium chloride or hydrochloric acid added. Both bring in chloride ions.

Small worked reminders

Reminder 1: brown precipitate plus sulfate. A yellow-brown solution gives a brown precipitate with sodium hydroxide solution. Another portion, acidified with dilute hydrochloric acid, gives a white precipitate with barium chloride solution.

• Brown hydroxide: iron(III), Fe³⁺.
• White precipitate with acidified barium chloride: sulfate, SO₄²⁻.
• Charges 3+ and 2−: two Fe³⁺ (total 6+) balance three SO₄²⁻ (total 6−). The compound is iron(III) sulfate, Fe₂(SO₄)₃.

Reminder 2: lilac flame plus fizzing. A powder colours the flame lilac. It effervesces in dilute hydrochloric acid, and the gas turns limewater cloudy white.

• Lilac flame: potassium, K⁺.
• Carbon dioxide from an acid: carbonate, CO₃²⁻.
• Two K⁺ balance one CO₃²⁻. The compound is potassium carbonate, K₂CO₃.

Reminder 3: showing a precipitate in an equation. If you are asked for an ionic equation for the chloride test, show only the ions that form the solid, with state symbols:

Ag⁺(aq) + Cl⁻(aq) → AgCl(s)

The nitrate and the other positive ion stay dissolved, so they are left out.

Quick self-test

1. State the flame colour of a potassium compound.
2. A solution gives a blue precipitate with sodium hydroxide solution. Name the metal ion.
3. A white hydroxide precipitate dissolves in excess sodium hydroxide. Name the metal ion.
4. Name the reagent added before silver nitrate in the halide test.
5. Give the colour of silver iodide.
6. Explain why dilute sulfuric acid is not used in the sulfate test.
7. A solid fizzes with dilute acid but the gas has no effect on limewater. Is the solid a carbonate?
8. A solid colours a flame red, and silver nitrate (with dilute nitric acid) turns its solution yellow and cloudy. Name the compound and give its formula.
9. Two white solids, one a calcium salt and one a magnesium salt, both give white hydroxide precipitates that stay in excess. How would you tell them apart?
10. Calculate the relative formula mass of silver chloride, AgCl. (Ar: Ag = 108, Cl = 35.5)
11. Iron(II) sulfate solution is tested with sodium hydroxide. Give the colour of the precipitate and the colour it may become after standing in air.

Answers

1. Lilac.
2. Copper(II), Cu²⁺.
3. Aluminium, Al³⁺.
4. Dilute nitric acid.
5. Yellow.
6. It contains sulfate ions, so it would give a white precipitate even with no sulfate in the sample.
7. No. Carbon dioxide would turn limewater cloudy white; the gas is something else, such as hydrogen from a metal.
8. Calcium iodide, CaI₂ (Ca²⁺ from the red flame; I⁻ from the yellow silver halide; two I⁻ balance one Ca²⁺).
9. Flame test: the calcium salt gives a red flame; the magnesium salt gives no characteristic colour.
10. 108 + 35.5 = 143.5.
11. Green at first; it slowly turns brown as iron(II) is oxidised to iron(III).

Where marks are usually lost

• Giving a flame colour for the wrong metal, most often swapping crimson (lithium) and red (calcium).
• Writing only “white precipitate” for aluminium without saying it dissolves in excess sodium hydroxide.
• Claiming sodium hydroxide alone can distinguish calcium from magnesium.
• Describing the halide test with hydrochloric acid, or the sulfate test with sulfuric acid.
• Leaving out the acid entirely, so the carbonate interference is not removed.
• Writing “goes cloudy” for the carbonate test without saying the gas is passed into limewater.
• Naming the ion when the question asks for the compound, or writing an unbalanced formula such as AlSO₄.
• In the flame test practical, giving no way to stop contamination between samples.

Official syllabus

OxfordAQA International GCSE Chemistry (9202) specification, Version 6.2, for exams May/June 2018 onwards, published by OxfordAQA. Scope of these notes: section 3.4.3 Identification of ions, with the flame test required practical from section 6.2.