What are chemical tests in the context of Edexcel IGCSE Chemistry?

Chemical tests in Edexcel IGCSE Chemistry refer to a series of procedures used to identify the presence of specific ions or compounds in a substance. These tests are crucial for understanding the composition of inorganic compounds and are a fundamental part of the curriculum.

How can you test for the presence of carbon dioxide gas in a chemical reaction?

To test for carbon dioxide gas, you can pass the gas through limewater, which is a solution of calcium hydroxide. If carbon dioxide is present, the limewater will turn milky due to the formation of calcium carbonate. This is a common test covered in the Edexcel IGCSE Chemistry syllabus.

What is the procedure for testing for sulfate ions in a solution?

To test for sulfate ions, you add a few drops of barium chloride solution to the test solution, followed by a few drops of dilute hydrochloric acid. If sulfate ions are present, a white precipitate of barium sulfate will form. This test is part of the inorganic chemistry section in the Edexcel IGCSE curriculum.

How do you identify the presence of chloride ions using a chemical test?

To test for chloride ions, you add a few drops of silver nitrate solution to the sample solution. If chloride ions are present, a white precipitate of silver chloride will form. This test is a key component of the chemical tests topic in Edexcel IGCSE Chemistry.

What is the flame test and how is it used to identify metal ions?

The flame test is a qualitative analysis technique used to identify metal ions based on the color they emit when heated in a flame. For example, sodium ions produce a yellow flame, while copper ions produce a green flame. This test is an important part of the inorganic chemistry section in the Edexcel IGCSE Chemistry syllabus.

Why is "Mixing up the splint tests for O₂ and H₂" a common mistake in Chemical Tests?

Why it happens: Both involve a wooden splint. How to avoid it: **O₂: GLOWING splint relights** (splint already smouldering — flame returns). **H₂: LIT splint goes POP** (already burning splint at the mouth of the tube — small explosion). Memorise: O for 'On again' (glowing → lights), H for 'How loud' (pop). Source: 4CH1 Examiner Reports 2022-2024

Why is "Using DRY red litmus for the ammonia test" a common mistake in Chemical Tests?

Why it happens: Skipping the 'damp' specification. How to avoid it: Ammonia is a gas — it has to dissolve in water on the paper to act as an alkali. DRY litmus would not show a colour change. Always state 'DAMP red litmus paper'. Source: 4CH1 Examiner Reports 2022-2024

Why is "Forgetting to add dilute HNO₃ before AgNO₃ (or dilute HCl before BaCl₂)" a common mistake in Chemical Tests?

Why it happens: Skipping a 'prep' step that doesn't seem to do anything. How to avoid it: The dilute acid removes interfering carbonate/hydroxide ions that would give false positives (e.g. Ag₂CO₃ white). Always state 'dilute HNO₃ first' for halide test, 'dilute HCl first' for sulfate test.

Why is "Mixing up Fe²⁺ (green) and Fe³⁺ (brown) precipitates" a common mistake in Chemical Tests?

Why it happens: Iron has two common oxidation states with similar names. How to avoid it: **Fe²⁺ = green** (think 'Fe TWO is green like a TWO-leaf clover'). **Fe³⁺ = brown / red-brown** (think 'rust = Fe(III)'). On standing in air, the GREEN Fe²⁺ slowly turns brown because it oxidises to Fe³⁺.

Why is "Saying Al³⁺ and Ca²⁺ give the same precipitate with NaOH" a common mistake in Chemical Tests?

Why it happens: True at first, but the discriminator is what happens in EXCESS. How to avoid it: BOTH give white precipitate. ADD EXCESS NaOH: Al(OH)₃ DISSOLVES (amphoteric → soluble aluminate). Ca(OH)₂ does NOT dissolve. Always test with excess NaOH to distinguish.

Why is "Saying 'CuSO₄ turning blue proves the liquid is PURE water'" a common mistake in Chemical Tests?

Why it happens: Confusing 'presence of water' with 'purity'. How to avoid it: Chemical tests (CuSO₄, cobalt paper) prove water is PRESENT but not that it's pure. Pure water test = PHYSICAL: melts at exactly 0 °C, boils at exactly 100 °C (sharp, fixed points). Mixtures have a RANGE of mp/bp.

Why is "Putting damp litmus INSIDE the test tube instead of at the mouth" a common mistake in Chemical Tests?

Why it happens: Trying to be thorough. How to avoid it: Hold the damp litmus paper at the MOUTH of the test tube (above the solution) so it samples only the gas, not the liquid. Putting it in the liquid causes other indicators / dissolved acids to affect the result.

Inorganic ChemistryPearson Edexcel IGCSE Chemistry 4CH1 Higher

Chemical Tests

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Chemical Tests — Pearson Edexcel International GCSE Chemistry 4CH1 Study Notes (2026 onwards, Spec Issue 4)

Qualitative analysis: a toolkit of standardised tests with characteristic observations. GASES — O₂ relights glowing splint; H₂ pops; CO₂ turns limewater milky; NH₃ damp red litmus → blue; Cl₂ bleaches damp litmus. CATIONS — flame tests (Li crimson, Na yellow, K lilac, Ca brick-red, Cu green-blue) and NaOH precipitates (Cu²⁺ blue, Fe²⁺ green, Fe³⁺ brown, Al³⁺ white dissolves in excess, Ca²⁺ white). ANIONS — CO₃²⁻ fizzes with HCl; halides with HNO₃ + AgNO₃ give white/cream/yellow precipitates; SO₄²⁻ with HCl + BaCl₂ gives white precipitate. WATER — anhydrous CuSO₄ blue; cobalt paper pink (presence only); melts at 0 °C / boils at 100 °C (purity).

What you’ll learn

Mapped to the Pearson Edexcel IGCSE 4CH1 syllabus (2026 onwards).

2.51 — Describe how to carry out a flame test and recall the colours produced by Li⁺, Na⁺, K⁺, Ca²⁺, Cu²⁺.
2.52 — Describe the chemical test for cations using NaOH(aq): Cu²⁺, Fe²⁺, Fe³⁺, Al³⁺, Ca²⁺, NH₄⁺.
2.53 — Describe the chemical test for carbonate (CO₃²⁻) using dilute HCl.
2.54 — Describe chemical tests for the gases O₂, H₂, CO₂, NH₃, Cl₂.
2.55 — Describe chemical tests for halide ions (Cl⁻, Br⁻, I⁻) using AgNO₃ and dilute HNO₃.
2.56 — Describe the chemical test for sulfate ions (SO₄²⁻) using BaCl₂ and dilute HCl.
2.57 — Describe chemical tests for water (anhydrous CuSO₄, cobalt(II) chloride paper) and PHYSICAL test for pure water (mp 0 °C, bp 100 °C).

Tests for gases (spec 2.54)

Five gases, five characteristic observations.

Test for OXYGEN (O₂).

Test: insert a GLOWING wooden splint (one that has just been blown out) into a test tube of the gas.
Observation: the splint RELIGHTS (the smouldering ember bursts into flame).
Reason: O₂ supports combustion. The smouldering wood has fuel and heat; only O₂ is missing → adding O₂ completes the fire triangle → splint reignites.

Test for HYDROGEN (H₂).

Test: insert a LIT wooden splint at the mouth of the test tube.
Observation: hydrogen burns with a SQUEAKY POP sound.
Equation: $2\text{H}_2 + \text{O}_2 \rightarrow 2\text{H}_2\text{O}$ — a small explosion produces the 'pop'. Used to distinguish H₂ from any other gas — uniquely makes that sound.

Test for CARBON DIOXIDE (CO₂).

Test: bubble the gas through LIMEWATER (a saturated solution of calcium hydroxide, Ca(OH)₂(aq)).
Observation: limewater turns MILKY / CLOUDY / WHITE (white precipitate of calcium carbonate, CaCO₃).
Equation: $\text{Ca(OH)}_2\text{(aq)} + \text{CO}_2\text{(g)} \rightarrow \text{CaCO}_3\text{(s)} + \text{H}_2\text{O(l)}$ .
Excess CO₂ caveat: with prolonged CO₂, the milkiness disappears as soluble calcium hydrogencarbonate forms — but this is rarely tested in 4CH1.

Test for AMMONIA (NH₃).

Test: hold a piece of DAMP RED LITMUS PAPER at the mouth of the test tube. The word 'DAMP' is essential (gas dissolves in the water).
Observation: red litmus turns BLUE.
Reason: NH₃ dissolves in the water on the litmus paper to give an alkaline solution: $\text{NH}_3 + \text{H}_2\text{O} \rightleftharpoons \text{NH}_4^+ + \text{OH}^-$ .
Smell (with caution): NH₃ has a sharp, pungent, choking smell — never sniff directly; waft gently.

Test for CHLORINE (Cl₂).

Test: hold a piece of DAMP LITMUS PAPER (red or blue) at the mouth of the test tube.
Observation:
- First, the litmus turns red (Cl₂ + H₂O → acidic solution: HCl + HOCl).
- Then the litmus is BLEACHED (turns white) — the HOCl is a bleaching agent.
Distinction from NH₃: NH₃ turns RED→BLUE (alkaline gas); Cl₂ turns BLUE→RED then bleaches (acidic + oxidising).

Summary table.

Gas	Test	Observation
O₂	Glowing splint	Splint relights
H₂	Lit splint	Squeaky pop
CO₂	Limewater	Turns milky
NH₃	Damp red litmus	Turns blue
Cl₂	Damp litmus	Bleaches white (after turning red)

O₂ relights a glowing splint.
H₂ burns with a squeaky pop with a lit splint.
CO₂ turns limewater milky (white CaCO₃ ppt).
NH₃ turns DAMP red litmus blue (the only gas that does this).
Cl₂ bleaches damp litmus (after turning it red).

See the full worked example for chemical tests →

Flame tests for cations (spec 2.51)

Clean platinum wire + concentrated HCl + Bunsen blue flame → diagnostic colour.

Procedure.

Clean the wire: dip a platinum (or nichrome) wire into concentrated HCl, then hold in a Bunsen flame. Repeat until no colour appears in the flame (removes contaminants, especially Na⁺).
Dip in HCl again, then in the solid (or solution) to be tested.
Hold in the hot (blue) part of a Bunsen flame (air hole open).
Observe the flame colour. For potassium, viewing through cobalt-blue glass filters out yellow Na⁺ contamination (lets the lilac K⁺ show).

Why HCl? It helps convert metal salts to volatile metal chlorides, which give brighter, more characteristic flame colours.

Flame colours.

Cation	Colour	Mnemonic
Lithium (Li⁺)	Crimson red	'Lithium Loves Crimson'
Sodium (Na⁺)	Bright yellow / orange-yellow	Common street lights (sodium vapour)
Potassium (K⁺)	Lilac / pale purple	View through cobalt glass
Calcium (Ca²⁺)	Brick-red / orange-red	(Distinguish from Li⁺ which is more crimson)
Copper (Cu²⁺)	Green / blue-green	Like a copper-green colour

Each metal ion gives a characteristic flame colour — view potassium through cobalt-blue glass.

Why flame tests work. In the hot flame, electrons in the metal cation absorb energy and jump to higher energy levels. As they fall back to lower levels, they emit photons of specific wavelengths (colours). Each metal has a unique set of energy levels, so each gives a unique colour pattern. This is the basis of atomic emission spectroscopy.

Limitations.

Only works for compounds that vaporise in the flame.
Mixtures give blended colours that are hard to interpret.
Sodium contamination is everywhere (sweat, dust) → cobalt glass needed for potassium.
Some colours overlap (Li and Ca can be confused — Li is more crimson, Ca more brick-red).

Example. A white solid is added to the wire and gives a brick-red flame. Conclusion: contains Ca²⁺ (e.g. CaCl₂, CaCO₃, Ca(NO₃)₂). To find the anion, follow up with HCl test (for CO₃²⁻), BaCl₂ test (for SO₄²⁻), etc.

Clean Pt wire in concentrated HCl + Bunsen flame until no colour.
Li crimson red. Na yellow. K lilac (cobalt glass).
Ca brick-red. Cu green/blue-green.
Mnemonic: 'Lithium Loves Crimson; Sodium Yells'.

See the full worked example for chemical tests →

NaOH precipitate test for cations (spec 2.52)

Add NaOH(aq); coloured precipitates identify the metal cation.

Adding a few drops of NaOH(aq) to a solution of an unknown metal cation gives a coloured precipitate of the metal hydroxide. The colour identifies the cation.

Standard observations.

Cation	Precipitate colour	Notes
Cu²⁺	Pale blue Cu(OH)₂	Insoluble in excess NaOH
Fe²⁺	Green Fe(OH)₂	Turns BROWN on standing in air (oxidises to Fe(OH)₃)
Fe³⁺	Red-brown / orange-brown Fe(OH)₃	Sometimes called 'rust-coloured'
Al³⁺	White Al(OH)₃	DISSOLVES in EXCESS NaOH (amphoteric) → colourless solution
Ca²⁺	White Ca(OH)₂	Slightly faint; does NOT dissolve in excess NaOH
NH₄⁺	NO precipitate	But on WARMING, gives off NH₃ gas (damp red litmus → blue)

Add excess NaOH to tell the two white precipitates apart — Al(OH)₃ dissolves (amphoteric), Ca(OH)₂ stays.

Ionic equations.

$\text{Cu}^{2+} + 2\text{OH}^- \rightarrow \text{Cu(OH)}_2$ (pale blue)
$\text{Fe}^{2+} + 2\text{OH}^- \rightarrow \text{Fe(OH)}_2$ (green)
$\text{Fe}^{3+} + 3\text{OH}^- \rightarrow \text{Fe(OH)}_3$ (brown)
$\text{Al}^{3+} + 3\text{OH}^- \rightarrow \text{Al(OH)}_3$ (white)
$\text{Ca}^{2+} + 2\text{OH}^- \rightarrow \text{Ca(OH)}_2$ (white)

For ammonium:

$\text{NH}_4^+ + \text{OH}^- \xrightarrow{\text{heat}} \text{NH}_3 + \text{H}_2\text{O}$

Distinguishing the two white precipitates — Al³⁺ vs Ca²⁺.

Both Al(OH)₃ and Ca(OH)₂ are white precipitates. To tell them apart:

Add EXCESS NaOH to each test tube.
Al(OH)₃ DISSOLVES in excess (it is AMPHOTERIC — reacts as an acid with the extra hydroxide):

$\text{Al(OH)}_3 + \text{OH}^- \rightarrow [\text{Al(OH)}_4]^-$

Ca(OH)₂ stays as a precipitate — does NOT dissolve in excess.

Distinguishing Fe²⁺ vs Fe³⁺.

Fe²⁺ gives an INITIAL green precipitate that SLOWLY TURNS BROWN on standing (because Fe²⁺ is oxidised to Fe³⁺ by atmospheric O₂):

$4\text{Fe(OH)}_2 + \text{O}_2 + 2\text{H}_2\text{O} \rightarrow 4\text{Fe(OH)}_3$

Fe³⁺ gives a brown precipitate immediately and stays brown.

Test for ammonium ion (NH₄⁺).

NH₄⁺ is unique — it does NOT give a precipitate with NaOH. Instead:

Add NaOH(aq) to the suspected NH₄⁺ solution.
WARM gently with a Bunsen.
Hold damp red litmus paper at the mouth of the tube.
If NH₃ is released, the litmus turns BLUE → confirms NH₄⁺.

This is the only common cation that gives a gas with NaOH rather than a precipitate, so it is diagnostic.

Cu²⁺ pale blue. Fe²⁺ green → brown. Fe³⁺ brown. Al³⁺ white (dissolves in excess). Ca²⁺ white (no change in excess). NH₄⁺ no ppt but gives NH₃ on warming.
Add EXCESS NaOH to distinguish Al³⁺ (dissolves — amphoteric) from Ca²⁺ (stays).
Fe²⁺ → brown on standing in air (Fe(OH)₂ oxidises to Fe(OH)₃).
NH₄⁺ test: warm with NaOH + damp red litmus → blue.

See the full worked example for chemical tests →

Tests for anions (spec 2.53, 2.55, 2.56)

Carbonate fizzes with HCl; halides with AgNO₃; sulfate with BaCl₂.

Test 1: Carbonate (CO₃²⁻).

Test: add a few drops of dilute HCl (or any dilute acid).
Observation: fizzing / effervescence; the gas turns limewater milky → CO₂.
Ionic equation: $\text{CO}_3^{2-} + 2\text{H}^+ \rightarrow \text{H}_2\text{O} + \text{CO}_2$ .

Test 2: Halide ions (Cl⁻, Br⁻, I⁻).

Procedure: add a few drops of dilute nitric acid (HNO₃) FIRST, then add silver nitrate solution (AgNO₃(aq)).
Why dilute HNO₃ first? Removes any carbonate (which would give white Ag₂CO₃ precipitate — false positive) or hydroxide (Ag₂O — brown precipitate). Cannot use HCl because that would itself precipitate with AgNO₃!
Observations:

Halide	Precipitate colour	Solubility in NH₃
Cl⁻	WHITE AgCl	Dissolves in dilute NH₃
Br⁻	CREAM AgBr	Dissolves only in concentrated NH₃
I⁻	YELLOW AgI	Insoluble in NH₃

The dilute HNO₃ removes carbonate first to prevent a false-positive white precipitate.

Ionic equations:

$\text{Ag}^+\text{(aq)} + \text{Cl}^-\text{(aq)} \rightarrow \text{AgCl(s)}$
$\text{Ag}^+\text{(aq)} + \text{Br}^-\text{(aq)} \rightarrow \text{AgBr(s)}$
$\text{Ag}^+\text{(aq)} + \text{I}^-\text{(aq)} \rightarrow \text{AgI(s)}$

Test 3: Sulfate (SO₄²⁻).

Procedure: add a few drops of dilute HCl FIRST, then add barium chloride solution (BaCl₂(aq)).
Why dilute HCl first? Removes any carbonate (which would give white BaCO₃ — false positive). Don't use HNO₃ here as nitrate doesn't interfere but HCl is the standard.
Observation: WHITE precipitate of barium sulfate (BaSO₄).
Ionic equation: $\text{Ba}^{2+}\text{(aq)} + \text{SO}_4^{2-}\text{(aq)} \rightarrow \text{BaSO}_4\text{(s)}$ .

Summary table.

Anion	Test reagents	Observation
CO₃²⁻	dilute HCl	Fizz; CO₂ (limewater milky)
Cl⁻	dilute HNO₃ + AgNO₃	White precipitate (AgCl)
Br⁻	dilute HNO₃ + AgNO₃	Cream precipitate (AgBr)
I⁻	dilute HNO₃ + AgNO₃	Yellow precipitate (AgI)
SO₄²⁻	dilute HCl + BaCl₂	White precipitate (BaSO₄)

Why the 'acid first' steps matter. Both the halide test and the sulfate test rely on a unique precipitate. If carbonate or hydroxide is present in the solution, those would ALSO give a white precipitate with Ag⁺ or Ba²⁺ — leading to false positives. The dilute acid (HNO₃ for halide, HCl for sulfate) destroys those interferents:

HCl removes CO₃²⁻ as CO₂.
HNO₃ removes CO₃²⁻ as CO₂ AND prevents AgOH from forming.

Always state the 'acid first' step — it earns its own mark in the mark scheme.

CO₃²⁻: HCl → fizz; CO₂ turns limewater milky.
Cl⁻ / Br⁻ / I⁻: HNO₃ first then AgNO₃ → white / cream / yellow precipitates.
SO₄²⁻: HCl first then BaCl₂ → white precipitate.
Acid first prevents false positives from CO₃²⁻ / OH⁻.

See the full worked example for chemical tests →

Tests for water — presence vs purity (spec 2.57)

Chemical tests prove H₂O presence (not purity). Physical tests (mp/bp) prove purity.

Two distinct things can be tested for: PRESENCE of water (chemical tests) and PURITY of water (physical test).

Chemical test 1: Anhydrous copper(II) sulfate.

Procedure: add a few drops of the suspected water sample to white anhydrous CuSO₄ powder (or vice versa).
Observation: white powder turns BLUE (forms hydrated copper(II) sulfate pentahydrate).
Equation: $\text{CuSO}_4 + 5\text{H}_2\text{O} \rightarrow \text{CuSO}_4 \cdot 5\text{H}_2\text{O}$ .
The reaction is also exothermic (the powder warms up).

Chemical test 2: Cobalt(II) chloride paper.

Procedure: dip a strip of blue anhydrous cobalt(II) chloride paper in the liquid.
Observation: paper turns PINK (forms hydrated cobalt chloride hexahydrate).
Equation: $\text{CoCl}_2 + 6\text{H}_2\text{O} \rightarrow \text{CoCl}_2 \cdot 6\text{H}_2\text{O}$ .
The reverse change (heat to dry) returns the paper to blue.

Important caveat — presence vs purity.

Both chemical tests confirm that water is PRESENT in the liquid sample. They do NOT prove the liquid is PURE water. A salt solution (NaCl in water), an ethanol-water mixture, or even pond water would also turn the CuSO₄ blue / cobalt paper pink.

Physical test for pure water — boiling point and melting point.

To prove that a sample is pure water (and not a mixture), measure:

Melting point: pure water melts at exactly 0 °C at standard pressure (101 kPa).
Boiling point: pure water boils at exactly 100 °C.

Why? Pure substances have SHARP, FIXED melting and boiling points. Mixtures (impure water) do NOT — they boil over a range, AND the boiling point is elevated (and freezing point depressed) by the dissolved impurities.

Pure water: boils at exactly 100 °C; a thermometer in boiling water reads steadily at 100 °C.
Salt water (1 M NaCl): boils at ~102 °C (depending on conc); thermometer reading drifts up as water evaporates and concentration increases.

Decision flowchart.

Want to test 'is there ANY water in this liquid?' → use chemical test (CuSO₄ or cobalt paper).
Want to test 'is this liquid PURE water, with NO dissolved substances?' → use physical test (mp = 0 °C, bp = 100 °C, sharp not range).

Bonus — anhydrous vs hydrated salt formulae.

Hydrated CuSO₄·5H₂O — BLUE crystals.
Anhydrous CuSO₄ — WHITE powder (obtained by heating hydrated CuSO₄ to drive off water of crystallisation).
The reaction is reversible: adding water to anhydrous CuSO₄ regenerates the blue hydrate.

So 'CuSO₄ blue → white on heating' and 'CuSO₄ white → blue on adding water' are CONNECTED chemistry — these are also the basis of the test for water itself.

Anhydrous CuSO₄: white → BLUE with water (CuSO₄·5H₂O).
Cobalt(II) chloride paper: BLUE → PINK with water.
Both tests prove PRESENCE of water, NOT purity.
For PURITY: pure water mp = 0 °C, bp = 100 °C exactly (sharp values).

See the full worked example for chemical tests →

Memorise this

Verbatim phrases and definitions Edexcel mark schemes credit.

Gas tests memorise: glowing splint = O₂; lit splint pop = H₂; limewater milky = CO₂; damp red litmus blue = NH₃; bleach damp litmus = Cl₂.
Flame colours: Li crimson; Na yellow/orange; K lilac (cobalt glass); Ca brick-red; Cu green-blue.
NaOH cation colours: Cu²⁺ pale blue; Fe²⁺ green; Fe³⁺ red-brown; Al³⁺/Ca²⁺ both white but Al dissolves in excess.
NH₄⁺ uniquely gives NO precipitate with NaOH but releases NH₃ on warming.
Anion 'acid-first' steps: HNO₃ before AgNO₃ (halide); HCl before BaCl₂ (sulfate). Prevents false positives.
AgX precipitate colours: AgCl white; AgBr cream; AgI yellow. (Reactivity decreases F > Cl > Br > I but solubility/precipitation order is Cl > Br > I dissolved.)
Sulfate test: BaCl₂ + HCl → white BaSO₄.
Water TESTS: anhydrous CuSO₄ white → blue; cobalt(II) chloride paper blue → pink. PURITY: bp 100 °C / mp 0 °C exactly.

How it’s examined

Chemical tests appear EVERY paper and contribute 6-12 marks per session. Paper 1 (4CH1/1C) items: (a) tabulate observations for cation tests with NaOH (4-6 marks); (b) describe how to identify a gas in a test-tube (3-4 marks); (c) write balanced ionic equations for precipitate formation (2-4 marks). Paper 2 (4CH1/2C) items: (d) propose a SCHEME of tests to identify an unknown ionic compound (5-7 marks — both cation AND anion); (e) explain the role of 'dilute acid first' steps (2-3 marks). 4-mark question routinely: distinguish Al³⁺ from Ca²⁺ precipitates → 'EXCESS NaOH' is the keyword.

Sources: Pearson Edexcel International GCSE Chemistry 4CH1 specification — Issue 4, September 2024 (valid for 2026 onwards); Pearson Edexcel 4CH1 Examiner Reports 2022-2024; 4CH1/1C May/Jun 2024 question paper and mark scheme; 4CH1/1C January 2024 question paper and mark scheme; 4CH1/2C May/Jun 2024 question paper and mark scheme; 4CH1/2C January 2024 question paper and mark scheme. Last reviewed 2026-05-12.

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Step-by-step worked examples — Chemical Tests

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1Identifying an unknown cation using NaOH (6 marks, Paper 1)
Core• Adapted from 4CH1/1C May/Jun 2024 Q6• cation test, NaOH, spec-2.51, spec-2.52
Question
A student has a solution of an unknown metal salt. She adds a few drops of sodium hydroxide solution and observes the formation of a precipitate. (a) Describe what she would observe for solutions containing each of these cations: Cu²⁺, Fe²⁺, Fe³⁺, Al³⁺, Ca²⁺. (b) Explain how she could distinguish between the Al³⁺ and Ca²⁺ precipitates which are both white. (6 marks)
Step-by-step solution
1. Step 1
  Cu²⁺ (1 M). Adding NaOH(aq) gives a PALE BLUE precipitate of copper(II) hydroxide.
  $\text{Cu}^{2+}\text{(aq)} + 2\text{OH}^-\text{(aq)} \rightarrow \text{Cu(OH)}_2\text{(s)}$
2. Step 2
  Fe²⁺ (1 A). GREEN precipitate of iron(II) hydroxide. On standing in air, the green precipitate slowly turns brown as Fe²⁺ is oxidised to Fe³⁺ by air: $4\text{Fe(OH)}_2 + \text{O}_2 + 2\text{H}_2\text{O} \rightarrow 4\text{Fe(OH)}_3$ .
  $\text{Fe}^{2+}\text{(aq)} + 2\text{OH}^-\text{(aq)} \rightarrow \text{Fe(OH)}_2\text{(s)}$
3. Step 3
  Fe³⁺ (1 A). RED-BROWN / ORANGE-BROWN precipitate of iron(III) hydroxide. Often described as 'rust-coloured'.
  $\text{Fe}^{3+}\text{(aq)} + 3\text{OH}^-\text{(aq)} \rightarrow \text{Fe(OH)}_3\text{(s)}$
4. Step 4
  Al³⁺ (1 B). WHITE precipitate of aluminium hydroxide.
  $\text{Al}^{3+}\text{(aq)} + 3\text{OH}^-\text{(aq)} \rightarrow \text{Al(OH)}_3\text{(s)}$
5. Step 5
  Ca²⁺ (1 B). WHITE precipitate of calcium hydroxide (Ca(OH)₂ is slightly soluble, so the precipitate is faint).
  $\text{Ca}^{2+}\text{(aq)} + 2\text{OH}^-\text{(aq)} \rightarrow \text{Ca(OH)}_2\text{(s)}$
6. Step 6
  (b) Distinguishing Al³⁺ vs Ca²⁺ (1 B). Both give a white precipitate. Difference: Al(OH)₃ DISSOLVES in EXCESS NaOH (because it is amphoteric — reacts with both acid and alkali) to form a colourless solution of aluminate ions; Ca(OH)₂ does NOT dissolve in excess NaOH. Test: add more NaOH and see if the white precipitate dissolves (→ Al³⁺) or stays (→ Ca²⁺).
  $\text{Al(OH)}_3\text{(s)} + \text{OH}^-\text{(aq)} \rightarrow [\text{Al(OH)}_4]^-\text{(aq)}$
Answer
(a) Cu²⁺ → pale blue precipitate. Fe²⁺ → green precipitate (turns brown on standing). Fe³⁺ → red-brown precipitate. Al³⁺ → white precipitate. Ca²⁺ → white precipitate. (b) Add EXCESS NaOH: the Al(OH)₃ DISSOLVES (amphoteric) → colourless solution; the Ca(OH)₂ stays as white precipitate.
Examiner tip
Edexcel 6-mark mark scheme: 1 each for five correct cation observations + 1 for distinguishing Al³⁺/Ca²⁺. Colour names must be specific — 'green' is acceptable for Fe²⁺ but 'dirty green' is also fine; 'brown' / 'red-brown' / 'rust' all accepted for Fe³⁺. The Al³⁺ amphoteric behaviour is a common discriminator question.
2Distinguishing chloride, bromide and iodide ions (5 marks, Paper 1)
Core• Adapted from 4CH1/1C Jan 2024 Q4• halide test, AgNO3, spec-2.55
Question
A student is given three colourless solutions: sodium chloride, sodium bromide and sodium iodide. (a) Describe a chemical test to identify which solution contains which halide. (b) Write the IONIC equation for the test with chloride ions. (5 marks)
Step-by-step solution
1. Step 1
  Test procedure (1 M). Pour ~2 cm³ of each unknown solution into a separate test tube. Add a few drops of dilute nitric acid (HNO₃) — this removes any CO₃²⁻ or OH⁻ ions that would also give a precipitate. Then add a few drops of silver nitrate solution (AgNO₃(aq)).
2. Step 2
  Chloride observation (1 A). Chloride gives a WHITE precipitate of silver chloride: Ag⁺(aq) + Cl⁻(aq) → AgCl(s). Confirmation: AgCl DISSOLVES in dilute ammonia solution.
3. Step 3
  Bromide observation (1 A). Bromide gives a CREAM precipitate of silver bromide: Ag⁺ + Br⁻ → AgBr(s). Confirmation: AgBr is partially soluble in CONCENTRATED ammonia only.
4. Step 4
  Iodide observation (1 B). Iodide gives a YELLOW (pale yellow) precipitate of silver iodide: Ag⁺ + I⁻ → AgI(s). Confirmation: AgI is INSOLUBLE in ammonia (even concentrated).
5. Step 5
  (b) Ionic equation for chloride (1 B). Write the ionic equation showing only the species that actually react — spectator ions (Na⁺, NO₃⁻) are omitted.
  $\text{Ag}^+\text{(aq)} + \text{Cl}^-\text{(aq)} \rightarrow \text{AgCl(s)}$
Answer
Test: add dilute HNO₃ + AgNO₃(aq) to each. Observations: Cl⁻ → WHITE precipitate (AgCl); Br⁻ → CREAM precipitate (AgBr); I⁻ → YELLOW precipitate (AgI). Ionic equation for chloride: Ag⁺(aq) + Cl⁻(aq) → AgCl(s).
Examiner tip
Edexcel 5-mark mark scheme: 1 for method (HNO₃ + AgNO₃); 1 each for three halide colours; 1 for correct ionic equation with STATE SYMBOLS. The 'dilute nitric acid first' step is crucial — gets 1 mark by itself. Common error: forgetting state symbols in ionic equation, OR using 2Ag⁺ + 2Cl⁻ → 2AgCl (overcomplicating).
3Identifying four common gases (4 marks, Paper 1)
Core• Adapted from 4CH1/1C May/Jun 2024 Q2• gas tests, spec-2.54
Question
Describe a chemical test (with expected observation) for each of these gases: (a) oxygen, (b) hydrogen, (c) carbon dioxide, (d) ammonia. (4 marks)
Step-by-step solution
1. Step 1
  (a) Oxygen (1 M). Test: insert a GLOWING wooden splint into a test tube of the gas. Observation: the splint relights (bursts into flame). Reason: O₂ supports combustion of the smouldering wood, reigniting it.
2. Step 2
  (b) Hydrogen (1 A). Test: insert a LIT wooden splint at the mouth of the test tube. Observation: hydrogen burns with a SQUEAKY POP sound. Reason: $2\text{H}_2 + \text{O}_2 \rightarrow 2\text{H}_2\text{O}$ — a small explosion gives the characteristic 'pop'.
3. Step 3
  (c) Carbon dioxide (1 A). Test: bubble the gas through LIMEWATER (a saturated solution of calcium hydroxide, Ca(OH)₂(aq)). Observation: limewater turns milky / cloudy / white (white precipitate of CaCO₃ forms). Equation: $\text{Ca(OH)}_2 + \text{CO}_2 \rightarrow \text{CaCO}_3 + \text{H}_2\text{O}$ .
4. Step 4
  (d) Ammonia (1 B). Test: hold a piece of DAMP RED LITMUS PAPER at the mouth of the test tube. Observation: red litmus turns BLUE. Reason: NH₃ dissolves in the water on the litmus paper to form an alkaline solution (NH₃ + H₂O → NH₄⁺ + OH⁻). DAMP is essential — dry litmus would NOT change colour.
Answer
(a) O₂: relights a glowing splint. (b) H₂: lit splint gives a squeaky pop. (c) CO₂: turns limewater milky (Ca(OH)₂ + CO₂ → CaCO₃ + H₂O). (d) NH₃: turns damp red litmus blue.
Examiner tip
Edexcel 4-mark mark scheme: 1 each for the four gases. Be SPECIFIC: 'lit splint' for H₂, 'GLOWING splint' for O₂ — they're different. 'Damp' red litmus for NH₃ is essential (1 mark would be lost for omitting 'damp'). 'Limewater turns milky' is the standard CO₂ phrase.
4Identifying an unknown white solid containing a metal carbonate (5 marks, Paper 1)
Core• Adapted from 4CH1/1C Jan 2024 Q8• carbonate test, flame test, spec-2.53, spec-2.56
Question
A white solid is suspected to be either lithium carbonate (Li₂CO₃), sodium carbonate (Na₂CO₃), or potassium carbonate (K₂CO₃). Describe TWO tests that, together, would identify which carbonate is present, stating the expected observations for each. (5 marks)
Step-by-step solution
1. Step 1
  Test 1 — confirm carbonate is present (1 M). Add a few drops of dilute hydrochloric acid (or dilute nitric acid) to a small sample of the solid. Observation: fizzing / effervescence; gas released. Bubble the gas through limewater → turns milky. This confirms a CARBONATE: $\text{CO}_3^{2-} + 2\text{H}^+ \rightarrow \text{H}_2\text{O} + \text{CO}_2$ .
2. Step 2
  Test 2 — flame test to identify the metal (1 A). Dip a clean platinum / nichrome wire in concentrated HCl (to remove impurities), then in the solid. Hold the wire in the blue (non-luminous) part of a Bunsen flame.
3. Step 3
  Lithium (1 A). A crimson red flame indicates Li⁺.
4. Step 4
  Sodium (1 B). A bright yellow / orange-yellow flame indicates Na⁺.
5. Step 5
  Potassium (1 B). A lilac / pale purple flame indicates K⁺. (Note: K⁺ flames can be masked by Na⁺ impurities; view through cobalt-blue glass to filter out yellow Na⁺ if needed.)
Answer
Test 1: add dilute HCl → fizzing; bubble gas through limewater → milky (confirms CO₃²⁻). Test 2: flame test. Crimson red = Li⁺ (Li₂CO₃); bright yellow = Na⁺ (Na₂CO₃); lilac = K⁺ (K₂CO₃).
Examiner tip
Edexcel 5-mark mark scheme: 1 for HCl + fizzing observation; 1 for limewater test confirming CO₂; 1 for flame test method (clean wire + Bunsen flame); 1 each for two correct flame colours. State the colours specifically — 'red' for Li can be too vague; 'crimson' is the mark-scheme term.
5Testing for water — purity and presence (4 marks, Paper 1)
Core• Adapted from 4CH1/1C May/Jun 2024 Q1• water test, spec-2.57
Question
A bottle contains an unknown colourless liquid. (a) Describe TWO chemical tests that would confirm the liquid contains water. (b) Describe a physical test that would confirm the water is PURE (i.e. contains nothing else). (4 marks)
Step-by-step solution
1. Step 1
  (a) Test 1 — anhydrous copper(II) sulfate (1 M). Sprinkle a small amount of WHITE anhydrous CuSO₄ powder into the liquid (or vice versa). Observation: white powder turns BLUE (forming hydrated CuSO₄·5H₂O). $\text{CuSO}_4 + 5\text{H}_2\text{O} \rightarrow \text{CuSO}_4 \cdot 5\text{H}_2\text{O}$ .
2. Step 2
  (a) Test 2 — cobalt(II) chloride paper (1 A). Dip a strip of BLUE anhydrous cobalt(II) chloride paper in the liquid. Observation: paper turns PINK (forming hydrated cobalt(II) chloride).
3. Step 3
  Caveat about chemical tests (1 A). These tests confirm the PRESENCE of water but NOT that the liquid is PURE water — a mixture (e.g. water + salt, water + ethanol) would also turn the paper pink and the CuSO₄ blue.
4. Step 4
  (b) Physical test for purity (1 B). Measure the boiling point and / or melting point:
  
  Pure water boils at exactly 100 °C (at standard atmospheric pressure, 101 kPa).
  
  Pure water melts at exactly 0 °C.
  
  If the liquid boils at a different temperature (e.g. 100.5 °C with salt) or over a range (mixtures don't have a sharp BP), it is impure. A pure substance has a SHARP, FIXED melting/boiling point.
Answer
(a) Test 1: white anhydrous CuSO₄ → BLUE (CuSO₄·5H₂O). Test 2: blue cobalt(II) chloride paper → PINK. Both confirm water is PRESENT but not that it's pure. (b) Measure the boiling point: pure water boils at exactly 100 °C (sharp, fixed point); impure water boils above 100 °C (or over a range).
Examiner tip
Edexcel 4-mark mark scheme: 1 each for two chemical tests (with colour changes); 1 for noting the distinction (presence vs purity); 1 for the boiling-point test with stated value 100 °C. The 'presence vs purity' distinction often catches students out — chemical tests prove H₂O is there, but only physical tests (mp, bp) prove PURITY.

Model Answers — Chemical Tests

High-scoring sample answers for chemical tests on the Pearson Edexcel IGCSE 4CH1 paper, with examiner-style notes mapping each response to the mark scheme and assessment objectives.

Question 1

4CH1/1C-style — flame tests6 marks

Describe in detail how to carry out a flame test, including the apparatus and procedure. Then state the expected flame colour for each of these cations: lithium (Li⁺), sodium (Na⁺), potassium (K⁺), calcium (Ca²⁺), copper (Cu²⁺). (6 marks)

Model answer

Apparatus and procedure.

Clean a flame-test wire (platinum or nichrome) by dipping it in concentrated hydrochloric acid and holding it in the hot blue (non-luminous) part of a Bunsen flame. Repeat until the wire produces no colour change in the flame (this removes any previous contamination — Na⁺ is a very common contaminant from sweat / dust).
Dip the clean wire in concentrated HCl again, then into the solid (or solution) being tested. The HCl helps to convert the metal compound into its chloride, which is volatile and gives a more vivid flame colour.
Hold the wire in the blue part of a Bunsen flame (air hole fully open) and observe the colour of the flame. The energy of the flame promotes electrons in the metal cation to higher energy levels; as they fall back down, they emit photons of a characteristic wavelength → colour.
For potassium specifically, view the flame through a piece of cobalt-blue glass: this filters out the bright yellow Na⁺ light (a common impurity) and lets the lilac K⁺ flame show clearly.

Flame colours.

Cation	Flame colour	Common compound source
Lithium (Li⁺)	Crimson red	Li₂CO₃, LiCl
Sodium (Na⁺)	Bright yellow / orange-yellow	NaCl, Na₂CO₃
Potassium (K⁺)	Lilac / pale purple	KCl, K₂CO₃
Calcium (Ca²⁺)	Brick-red / orange-red	CaCl₂, Ca(NO₃)₂
Copper (Cu²⁺)	Green / blue-green	CuCl₂, CuSO₄

Other useful colours (not 4CH1 core but sometimes seen): strontium → bright red; barium → apple green; caesium → blue-violet.

Why flame tests work. Each metal cation has a unique electron configuration. In the flame, electrons absorb energy and jump to higher energy levels. When they fall back, they emit light of specific wavelengths — these specific wavelengths correspond to specific colours. This is the basis of atomic emission spectroscopy.

Limitations. Flame tests only work on volatile metal compounds (chlorides especially). Mixtures can give confused colours. Strong yellow Na⁺ can mask weaker colours like K⁺ (hence the cobalt-blue glass).

Why this scores

Edexcel 6-mark mark scheme: 1 for clean wire + concentrated HCl; 1 for use of Bunsen blue (hot) flame; 1 for cobalt glass / mention of Na contamination; 1 each for three correct flame colours (any choice of three for full marks if all colours covered). Use the specific colour names from the mark scheme — 'red' for Li can be too vague.

Question 2
4CH1/2C-style — anion identification scheme6 marks
An unknown salt is suspected to contain ONE of these anions: carbonate (CO₃²⁻), sulfate (SO₄²⁻), chloride (Cl⁻), bromide (Br⁻), or iodide (I⁻). Describe a systematic series of chemical tests that would identify which anion is present, including the expected observations and balanced ionic equations. (6 marks)
Model answer
Test 1: Carbonate test.

Procedure: Add ~2 cm³ of dilute hydrochloric acid (or dilute nitric acid) to a small amount of the unknown salt in a test tube.

Observation if carbonate: rapid effervescence (fizzing); the gas produced turns limewater milky (white precipitate of CaCO₃).

Ionic equation: $\text{CO}_3^{2-} + 2\text{H}^+ \rightarrow \text{H}_2\text{O} + \text{CO}_2$ .

If no fizzing, the salt is NOT a carbonate — move to Test 2.

Test 2: Sulfate test.

Procedure: To a fresh sample of solution, add a few drops of dilute hydrochloric acid first (to remove any carbonate that would interfere), then add a few drops of barium chloride solution BaCl₂(aq).

Observation if sulfate: WHITE precipitate of barium sulfate (BaSO₄).

Ionic equation: $\text{Ba}^{2+}\text{(aq)} + \text{SO}_4^{2-}\text{(aq)} \rightarrow \text{BaSO}_4\text{(s)}$ .

Why dilute HCl first? To prevent BaCO₃ (also white) precipitating, which would give a false positive. HCl removes any CO₃²⁻ as CO₂ gas.

If no precipitate, move to Test 3.

Test 3: Halide test (Cl⁻ / Br⁻ / I⁻).

Procedure: To a fresh sample of solution, add a few drops of dilute nitric acid (HNO₃) first, then add a few drops of silver nitrate solution AgNO₃(aq).

Observations:

WHITE precipitate → chloride (AgCl): $\text{Ag}^+ + \text{Cl}^- \rightarrow \text{AgCl(s)}$ .

CREAM precipitate → bromide (AgBr): $\text{Ag}^+ + \text{Br}^- \rightarrow \text{AgBr(s)}$ .

YELLOW precipitate → iodide (AgI): $\text{Ag}^+ + \text{I}^- \rightarrow \text{AgI(s)}$ .

Why dilute HNO₃ first? To remove any carbonate (which would form Ag₂CO₃ — white precipitate) or hydroxide (Ag₂O — brown precipitate) that would interfere. Dilute HCl can't be used because Cl⁻ from HCl would itself precipitate with AgNO₃.

Confirmation of halide identity (if needed): the silver halide precipitates differ in their solubility in ammonia:

AgCl dissolves in dilute NH₃(aq).

AgBr dissolves only in concentrated NH₃(aq).

AgI is insoluble in NH₃ even at high concentration.

Decision flowchart.

Fizzing with HCl? → carbonate (CO₃²⁻). Stop.

White ppt with BaCl₂ + HCl? → sulfate (SO₄²⁻). Stop.

With AgNO₃ + HNO₃: white = Cl⁻; cream = Br⁻; yellow = I⁻.
Why this scores
Edexcel 6-mark mark scheme: 1 each for the three procedures (carbonate, sulfate, halide); 1 each for three balanced ionic equations. State the 'acid first' steps explicitly — the marker rewards these as separate marks. Without dilute HNO₃ before AgNO₃, a hydroxide-containing solution would give a false positive.
Question 3
4CH1/2C-style — combined cation + anion identification5 marks
A solution of a metal salt is added to NaOH(aq) and a brown precipitate forms that does NOT dissolve in excess NaOH. The same solution gives a white precipitate when BaCl₂(aq) and dilute HCl are added. Identify the cation AND anion present, and write a balanced ionic equation for each test. (5 marks)
Model answer
Identifying the cation.

The unknown solution gives a brown / red-brown / orange-brown precipitate with NaOH. From the cation precipitate colours:

Cu²⁺ → pale blue

Fe²⁺ → green (turns brown on standing)

Fe³⁺ → brown ✓

Al³⁺ → white (dissolves in excess)

Ca²⁺ → white (no change in excess)

The brown precipitate that does NOT dissolve in excess NaOH → iron(III) cation, Fe³⁺.

(If the precipitate had been green initially and turned brown ON STANDING IN AIR, that would have indicated Fe²⁺. The question says brown immediately, so Fe³⁺.)

Ionic equation for the NaOH test:

$\text{Fe}^{3+}\text{(aq)} + 3\text{OH}^-\text{(aq)} \rightarrow \text{Fe(OH)}_3\text{(s)}$

(Brown precipitate.)

Identifying the anion.

A white precipitate with BaCl₂ + dilute HCl indicates sulfate (SO₄²⁻):

The dilute HCl is added first to remove any carbonate (which would give BaCO₃, also white, as a false positive).

The white precipitate is barium sulfate (BaSO₄) — one of the few INSOLUBLE sulfates.

Ionic equation for the BaCl₂ test:

$\text{Ba}^{2+}\text{(aq)} + \text{SO}_4^{2-}\text{(aq)} \rightarrow \text{BaSO}_4\text{(s)}$

Conclusion.

The salt is iron(III) sulfate, Fe₂(SO₄)₃. Pure Fe₂(SO₄)₃ solution is pale yellow/brown in colour.

Full equation for the salt:

$\text{Fe}_2(\text{SO}_4)_3\text{(aq)} + 6\text{NaOH(aq)} \rightarrow 2\text{Fe(OH)}_3\text{(s)} + 3\text{Na}_2\text{SO}_4\text{(aq)}$
Why this scores
Edexcel 5-mark mark scheme: 1 for identifying Fe³⁺; 1 for ionic equation for NaOH test (with state symbols); 1 for identifying SO₄²⁻; 1 for ionic equation for BaCl₂ test (with state symbols); 1 for naming the salt (Fe₂(SO₄)₃) or stating the molecular formula. The state symbols are essential for the marks on ionic equations.
Question 4
4CH1/1C-style — ammonium ion test5 marks
A student is given a solution and asked to confirm that it contains the ammonium ion (NH₄⁺). Describe a chemical test that would identify NH₄⁺, including the expected observation, the gas produced, and the test for that gas. Write a balanced ionic equation for the reaction. (5 marks)
Model answer
Test for ammonium ion (NH₄⁺).

Procedure.

Pour ~2 cm³ of the unknown solution into a test tube.

Add ~2 cm³ of sodium hydroxide solution (NaOH(aq)).

Warm gently with a Bunsen burner (do not boil vigorously).

Hold a piece of damp red litmus paper at the mouth of the test tube (using tongs — don't let the paper touch the solution).

Observations.

No visible precipitate forms in the solution (NH₄⁺ is the ONLY common cation that gives no precipitate with NaOH).

On warming, ammonia gas (NH₃) is released — a pungent, choking smell (don't inhale directly).

The damp red litmus paper turns BLUE — confirms ammonia (the only common gas that does this).

Why warming is needed. At room temperature, NH₄⁺ + OH⁻ exist in equilibrium with NH₃·H₂O. Heating drives off the NH₃ gas (Le Chatelier — removing a product / heat-shifted equilibrium).

Why 'DAMP' litmus. Ammonia is a gas; it must dissolve in the water on the litmus paper to form an alkaline solution that turns the litmus blue. Dry litmus would not change colour.

Ionic equation.

$\text{NH}_4^+\text{(aq)} + \text{OH}^-\text{(aq)} \xrightarrow{\text{heat}} \text{NH}_3\text{(g)} + \text{H}_2\text{O(l)}$

Why this is a unique test. No other common cation in 4CH1 (Cu²⁺, Fe²⁺, Fe³⁺, Al³⁺, Ca²⁺, Na⁺, K⁺, Li⁺) gives ammonia gas with NaOH. The combination of 'no precipitate' + 'gas turns damp red litmus blue on warming' is uniquely diagnostic for NH₄⁺.

Conclusion. The solution contains the ammonium ion, NH₄⁺.
Why this scores
Edexcel 5-mark mark scheme: 1 for adding NaOH + warming; 1 for damp red litmus paper at mouth; 1 for blue colour observed; 1 for naming the gas (ammonia, NH₃); 1 for balanced ionic equation with state symbols + heat condition. Common loss of mark: forgetting 'damp' litmus, OR forgetting to mention 'warming'. The 4CH1 mark scheme is strict on these details.

Key Definitions and Keywords — Chemical Tests

Definitions to memorise and the exact keywords mark schemes credit for chemical tests answers — sharpened from recent examiner reports for the 2026 Pearson Edexcel IGCSE 4CH1 sitting.

Flame test
Examiner keyword
A qualitative test for metal cations: dip a clean platinum/nichrome wire in concentrated HCl, then in the sample, hold in a blue Bunsen flame, and observe the colour of the flame.
Precipitate test (with NaOH)
Examiner keyword
Adding a few drops of NaOH(aq) to a solution of an unknown cation. Different cations form coloured hydroxide precipitates: Cu²⁺ pale blue; Fe²⁺ green; Fe³⁺ brown; Al³⁺ white (dissolves in excess); Ca²⁺ white.
Amphoteric (hydroxide)
Examiner keyword
A substance that can react as both an acid AND a base. Al(OH)₃ is amphoteric: it dissolves in BOTH dilute acid (acting as a base) AND excess NaOH (acting as an acid forming [Al(OH)₄]⁻). Used to distinguish Al³⁺ from Ca²⁺ in the cation test.
Limewater
Examiner keyword
A saturated solution of calcium hydroxide, Ca(OH)₂(aq). Turns MILKY / CLOUDY when CO₂ is bubbled through (white precipitate of CaCO₃). Standard test for CO₂.
Litmus paper
A pH indicator paper that turns blue in alkaline solution and red in acidic solution. Used DAMP (with water on the surface) to test gases like NH₃ (turns red→blue) and Cl₂ (turns blue→red then bleaches).
Silver halide test
Examiner keyword
Test for halide anions (Cl⁻, Br⁻, I⁻): add dilute HNO₃ then AgNO₃(aq). White ppt = AgCl (chloride); cream ppt = AgBr (bromide); yellow ppt = AgI (iodide). HNO₃ removes interfering carbonate / hydroxide.
Sulfate test (BaCl₂)
Examiner keyword
Test for sulfate anion (SO₄²⁻): add dilute HCl then BaCl₂(aq) → white precipitate of BaSO₄. HCl removes carbonate (which would also give a white precipitate of BaCO₃ — false positive).
Carbonate test
Examiner keyword
Test for carbonate anion (CO₃²⁻): add dilute HCl → effervescence; gas turns limewater milky (CO₂). Equation: CO₃²⁻ + 2H⁺ → H₂O + CO₂.
Anhydrous copper(II) sulfate test (for water)
Examiner keyword
White CuSO₄ powder (anhydrous) turns BLUE on contact with water → CuSO₄·5H₂O (hydrated). Confirms presence of water (but not purity).
Cobalt(II) chloride paper
Examiner keyword
Paper soaked in anhydrous CoCl₂ — BLUE when dry, turns PINK on contact with water. Confirms presence of water (but not purity).

Common Mistakes and Misconceptions — Chemical Tests

The traps other students keep falling into on chemical tests questions — taken from recent Pearson Edexcel IGCSE 4CH1 examiner reports and mark schemes — and how to avoid them.

✕Mixing up the splint tests for O₂ and H₂
4CH1 Examiner Reports 2022-2024
Why it happens
Both involve a wooden splint.
How to avoid it
O₂: GLOWING splint relights (splint already smouldering — flame returns). H₂: LIT splint goes POP (already burning splint at the mouth of the tube — small explosion). Memorise: O for 'On again' (glowing → lights), H for 'How loud' (pop).
✕Using DRY red litmus for the ammonia test
4CH1 Examiner Reports 2022-2024
Why it happens
Skipping the 'damp' specification.
How to avoid it
Ammonia is a gas — it has to dissolve in water on the paper to act as an alkali. DRY litmus would not show a colour change. Always state 'DAMP red litmus paper'.
✕Forgetting to add dilute HNO₃ before AgNO₃ (or dilute HCl before BaCl₂)
Why it happens
Skipping a 'prep' step that doesn't seem to do anything.
How to avoid it
The dilute acid removes interfering carbonate/hydroxide ions that would give false positives (e.g. Ag₂CO₃ white). Always state 'dilute HNO₃ first' for halide test, 'dilute HCl first' for sulfate test.
✕Mixing up Fe²⁺ (green) and Fe³⁺ (brown) precipitates
Why it happens
Iron has two common oxidation states with similar names.
How to avoid it
Fe²⁺ = green (think 'Fe TWO is green like a TWO-leaf clover'). Fe³⁺ = brown / red-brown (think 'rust = Fe(III)'). On standing in air, the GREEN Fe²⁺ slowly turns brown because it oxidises to Fe³⁺.
✕Saying Al³⁺ and Ca²⁺ give the same precipitate with NaOH
Why it happens
True at first, but the discriminator is what happens in EXCESS.
How to avoid it
BOTH give white precipitate. ADD EXCESS NaOH: Al(OH)₃ DISSOLVES (amphoteric → soluble aluminate). Ca(OH)₂ does NOT dissolve. Always test with excess NaOH to distinguish.
✕Saying 'CuSO₄ turning blue proves the liquid is PURE water'
Why it happens
Confusing 'presence of water' with 'purity'.
How to avoid it
Chemical tests (CuSO₄, cobalt paper) prove water is PRESENT but not that it's pure. Pure water test = PHYSICAL: melts at exactly 0 °C, boils at exactly 100 °C (sharp, fixed points). Mixtures have a RANGE of mp/bp.
✕Putting damp litmus INSIDE the test tube instead of at the mouth
Why it happens
Trying to be thorough.
How to avoid it
Hold the damp litmus paper at the MOUTH of the test tube (above the solution) so it samples only the gas, not the liquid. Putting it in the liquid causes other indicators / dissolved acids to affect the result.

Practice questions

Exam-style questions with step-by-step worked solutions. Try one before checking the method.

Past paper style quiz

Get a report showing which sub-topics you've nailed and which ones still need work.

Chemical Tests — frequently asked questions

The things students keep getting wrong in this sub-topic, answered.

Chemical Tests — Pearson Edexcel International GCSE Chemistry 4CH1 Study Notes (2026 onwards, Spec Issue 4)

Gas

Test

Observation

O₂

Glowing splint

Splint relights

H₂

Lit splint

Squeaky pop

CO₂

Limewater

Turns milky

NH₃

Damp red litmus

Turns blue

Cl₂

Damp litmus

Bleaches white (after turning red)

Cation

Colour

Mnemonic

Lithium (Li⁺)

Crimson red

'Lithium Loves Crimson'

Sodium (Na⁺)

Bright yellow / orange-yellow

Common street lights (sodium vapour)

Potassium (K⁺)

Lilac / pale purple

View through cobalt glass

Calcium (Ca²⁺)

Brick-red / orange-red

(Distinguish from Li⁺ which is more crimson)

Copper (Cu²⁺)

Green / blue-green

Like a copper-green colour

Cation

Precipitate colour

Notes

Cu²⁺

Pale blue Cu(OH)₂

Insoluble in excess NaOH

Fe²⁺

Green Fe(OH)₂

Turns BROWN on standing in air (oxidises to Fe(OH)₃)

Fe³⁺

Red-brown / orange-brown Fe(OH)₃

Sometimes called 'rust-coloured'

Al³⁺

White Al(OH)₃

DISSOLVES in EXCESS NaOH (amphoteric) → colourless solution

Ca²⁺

White Ca(OH)₂

Slightly faint; does NOT dissolve in excess NaOH

NH₄⁺

NO precipitate

But on WARMING, gives off NH₃ gas (damp red litmus → blue)

Halide

Precipitate colour

Solubility in NH₃

Cl⁻

WHITE AgCl

Dissolves in dilute NH₃

Br⁻

CREAM AgBr

Dissolves only in concentrated NH₃

I⁻

YELLOW AgI

Insoluble in NH₃

Anion

Test reagents

Observation

CO₃²⁻

dilute HCl

Fizz; CO₂ (limewater milky)

Cl⁻

dilute HNO₃ + AgNO₃

White precipitate (AgCl)

Br⁻

dilute HNO₃ + AgNO₃

Cream precipitate (AgBr)

I⁻

dilute HNO₃ + AgNO₃

Yellow precipitate (AgI)

SO₄²⁻

dilute HCl + BaCl₂

White precipitate (BaSO₄)

Two distinct things can be tested for: PRESENCE of water (chemical tests) and PURITY of water (physical test).

Chemical test 1: Anhydrous copper(II) sulfate.

Procedure: add a few drops of the suspected water sample to white anhydrous CuSO₄ powder (or vice versa).
Observation: white powder turns BLUE (forms hydrated copper(II) sulfate pentahydrate).
Equation: $\text{CuSO}_4 + 5\text{H}_2\text{O} \rightarrow \text{CuSO}_4 \cdot 5\text{H}_2\text{O}$ .
The reaction is also exothermic (the powder warms up).

Chemical test 2: Cobalt(II) chloride paper.

Procedure: dip a strip of blue anhydrous cobalt(II) chloride paper in the liquid.
Observation: paper turns PINK (forms hydrated cobalt chloride hexahydrate).
Equation: $\text{CoCl}_2 + 6\text{H}_2\text{O} \rightarrow \text{CoCl}_2 \cdot 6\text{H}_2\text{O}$ .
The reverse change (heat to dry) returns the paper to blue.

Important caveat — presence vs purity.

Physical test for pure water — boiling point and melting point.

To prove that a sample is pure water (and not a mixture), measure:

Melting point: pure water melts at exactly 0 °C at standard pressure (101 kPa).
Boiling point: pure water boils at exactly 100 °C.

Pure water: boils at exactly 100 °C; a thermometer in boiling water reads steadily at 100 °C.
Salt water (1 M NaCl): boils at ~102 °C (depending on conc); thermometer reading drifts up as water evaporates and concentration increases.

Decision flowchart.

Want to test 'is there ANY water in this liquid?' → use chemical test (CuSO₄ or cobalt paper).
Want to test 'is this liquid PURE water, with NO dissolved substances?' → use physical test (mp = 0 °C, bp = 100 °C, sharp not range).

Bonus — anhydrous vs hydrated salt formulae.

Hydrated CuSO₄·5H₂O — BLUE crystals.
Anhydrous CuSO₄ — WHITE powder (obtained by heating hydrated CuSO₄ to drive off water of crystallisation).
The reaction is reversible: adding water to anhydrous CuSO₄ regenerates the blue hydrate.

So 'CuSO₄ blue → white on heating' and 'CuSO₄ white → blue on adding water' are CONNECTED chemistry — these are also the basis of the test for water itself.

Cation

Flame colour

Common compound source

Lithium (Li⁺)

Crimson red

Li₂CO₃, LiCl

Sodium (Na⁺)

Bright yellow / orange-yellow

NaCl, Na₂CO₃

Potassium (K⁺)

Lilac / pale purple

KCl, K₂CO₃

Calcium (Ca²⁺)

Brick-red / orange-red

CaCl₂, Ca(NO₃)₂

Copper (Cu²⁺)

Green / blue-green

CuCl₂, CuSO₄

Chemical Tests

Short Study Notes in the form of Slides

Short Study Notes — Chemical Tests

Detailed Notes

What you’ll learn

How it’s examined

1Identifying an unknown cation using NaOH (6 marks, Paper 1)

2Distinguishing chloride, bromide and iodide ions (5 marks, Paper 1)

3Identifying four common gases (4 marks, Paper 1)

4Identifying an unknown white solid containing a metal carbonate (5 marks, Paper 1)

5Testing for water — purity and presence (4 marks, Paper 1)

Flame test

Precipitate test (with NaOH)

Amphoteric (hydroxide)

Limewater

Litmus paper

Silver halide test

Sulfate test (BaCl₂)

Carbonate test

Anhydrous copper(II) sulfate test (for water)

Cobalt(II) chloride paper

✕Mixing up the splint tests for O₂ and H₂

✕Using DRY red litmus for the ammonia test

✕Forgetting to add dilute HNO₃ before AgNO₃ (or dilute HCl before BaCl₂)

✕Mixing up Fe²⁺ (green) and Fe³⁺ (brown) precipitates

✕Saying Al³⁺ and Ca²⁺ give the same precipitate with NaOH

✕Saying 'CuSO₄ turning blue proves the liquid is PURE water'

✕Putting damp litmus INSIDE the test tube instead of at the mouth

Practice questions

Past paper style quiz

Assess your understanding

Chemical Tests — frequently asked questions

Chemical Tests

Short Study Notes in the form of Slides

Short Study Notes — Chemical Tests

Detailed Notes

What you’ll learn

How it’s examined

1Identifying an unknown cation using NaOH (6 marks, Paper 1)

2Distinguishing chloride, bromide and iodide ions (5 marks, Paper 1)

3Identifying four common gases (4 marks, Paper 1)

4Identifying an unknown white solid containing a metal carbonate (5 marks, Paper 1)

5Testing for water — purity and presence (4 marks, Paper 1)

Flame test

Precipitate test (with NaOH)

Amphoteric (hydroxide)

Limewater

Litmus paper

Silver halide test

Sulfate test (BaCl₂)

Carbonate test

Anhydrous copper(II) sulfate test (for water)

Cobalt(II) chloride paper

✕Mixing up the splint tests for O₂ and H₂

✕Using DRY red litmus for the ammonia test

✕Forgetting to add dilute HNO₃ before AgNO₃ (or dilute HCl before BaCl₂)

✕Mixing up Fe²⁺ (green) and Fe³⁺ (brown) precipitates

✕Saying Al³⁺ and Ca²⁺ give the same precipitate with NaOH

✕Saying 'CuSO₄ turning blue proves the liquid is PURE water'

✕Putting damp litmus INSIDE the test tube instead of at the mouth

Practice questions

Past paper style quiz

Assess your understanding

Chemical Tests — frequently asked questions