Required practical: tests for gases, ions and water

What you'll learn

This required practical focuses on qualitative analysis — identifying unknown substances using characteristic chemical tests. You must know the specific tests for common gases (hydrogen, oxygen, carbon dioxide, chlorine), positive ions (metal cations), negative ions (anions), and water. These tests appear frequently in both Paper 1 and Paper 2, and understanding the underlying chemistry is essential for achieving top grades.

Key terms and definitions

Qualitative analysis — laboratory techniques used to identify which chemical species are present in a sample, without measuring quantity

Precipitate — an insoluble solid formed when two solutions react together, appearing as a cloudiness or solid settling in the liquid

Carbonate — a compound containing the CO₃²⁻ ion, which reacts with acids to produce carbon dioxide gas

Flame test — a procedure where a metal compound is held in a Bunsen flame to produce a characteristic colour, used to identify metal ions

Spectator ions — ions present in a reaction mixture that do not participate in the chemical change and remain unchanged in solution

Effervescence — the vigorous formation of bubbles of gas, often described as fizzing

Lithium hydroxide — LiOH, a strong alkali that reacts with acids and dissolves in water to produce alkaline solutions (used in flame tests)

Ionic equation — a chemical equation showing only the ions that participate in the reaction, omitting spectator ions

Core concepts

Tests for gases

Hydrogen (H₂)

When testing for hydrogen gas, use the squeaky pop test:

• Collect the gas in a test tube
• Hold a lit splint to the mouth of the test tube
• Positive result: A squeaky pop sound indicates hydrogen is present
• The hydrogen burns rapidly with oxygen in the air: 2H₂ + O₂ → 2H₂O

This test is commonly used when testing the reactivity of metals with acids, as hydrogen is a product of these reactions.

Oxygen (O₂)

The test for oxygen uses a glowing splint:

• Insert a glowing (not flaming) splint into the test tube
• Positive result: The splint relights, bursting into flame
• Oxygen supports combustion but does not burn itself

This test is essential when investigating decomposition reactions, such as the thermal decomposition of metal carbonates or the catalytic decomposition of hydrogen peroxide.

Carbon dioxide (CO₂)

Carbon dioxide is tested using limewater (aqueous calcium hydroxide):

• Bubble the gas through limewater
• Positive result: The limewater turns from colourless to cloudy/milky white
• A white precipitate of calcium carbonate forms: Ca(OH)₂(aq) + CO₂(g) → CaCO₃(s) + H₂O(l)

Note: If you continue bubbling CO₂ through the cloudy limewater, it eventually becomes clear again as the calcium carbonate reacts to form soluble calcium hydrogencarbonate.

Chlorine (Cl₂)

Chlorine gas is identified using damp litmus paper:

• Hold damp blue or red litmus paper in the gas
• Positive result: The litmus paper is bleached white
• Chlorine is a toxic, pale green gas with a distinctive choking smell
• The bleaching occurs because chlorine is a powerful oxidising agent

Tests for positive ions (cations)

Flame tests for metal ions

Flame tests identify metal cations by the characteristic colour they produce when heated in a Bunsen flame:

Method:

1. Clean a nichrome wire loop by dipping it in dilute hydrochloric acid, then heating it in a blue Bunsen flame until no colour is seen
2. Dip the clean wire loop into the acid again, then into the solid sample
3. Hold the wire loop in the edge of the blue Bunsen flame
4. Observe the colour produced

Results:

• Lithium (Li⁺): Crimson/red flame
• Sodium (Na⁺): Yellow/orange flame (very intense)
• Potassium (K⁺): Lilac/purple flame (can be difficult to see through yellow from sodium contamination)
• Calcium (Ca²⁺): Orange-red flame
• Copper (Cu²⁺): Blue-green flame

The sodium flame colour is particularly intense and can mask other colours, so thorough cleaning between tests is essential.

Sodium hydroxide tests for metal ions

Add a few drops of sodium hydroxide solution to a solution of the unknown metal ion. Many metal ions form coloured precipitates:

Results:

• Calcium (Ca²⁺): White precipitate
• Copper (Cu²⁺): Blue precipitate
• Iron(II) (Fe²⁺): Green precipitate
• Iron(III) (Fe³⁺): Brown/orange precipitate
• Aluminium (Al³⁺): White precipitate, dissolves in excess sodium hydroxide
• Magnesium (Mg²⁺): White precipitate

The ionic equation for copper(II) is: Cu²⁺(aq) + 2OH⁻(aq) → Cu(OH)₂(s)

Distinguishing white precipitates: Only aluminium hydroxide dissolves when excess sodium hydroxide is added. Calcium hydroxide is slightly soluble (appears as a faint white precipitate), whereas magnesium hydroxide is insoluble (appears as a thick white precipitate).

Ammonium ions (NH₄⁺)

To test for ammonium ions:

1. Add sodium hydroxide solution to the sample
2. Warm the mixture gently
3. Test any gas produced with damp red litmus paper
4. Positive result: The litmus paper turns blue, indicating ammonia gas (NH₃)

The reaction is: NH₄⁺(aq) + OH⁻(aq) → NH₃(g) + H₂O(l)

You may also detect ammonia by its characteristic pungent smell, though this should not be your primary test method in exams.

Tests for negative ions (anions)

Carbonate ions (CO₃²⁻)

Add dilute acid (hydrochloric or nitric) to the sample:

• Observation: Effervescence (fizzing) occurs
• Test the gas produced using limewater
• Positive result: Limewater turns cloudy, confirming carbon dioxide

The general equation is: Carbonate + Acid → Salt + Water + Carbon dioxide

For calcium carbonate: CaCO₃(s) + 2HCl(aq) → CaCl₂(aq) + H₂O(l) + CO₂(g)

Sulfate ions (SO₄²⁻)

1. Add dilute hydrochloric acid (to remove any carbonate ions that might interfere)
2. Add barium chloride solution (BaCl₂)
3. Positive result: A white precipitate of barium sulfate forms

The ionic equation is: Ba²⁺(aq) + SO₄²⁻(aq) → BaSO₄(s)

The hydrochloric acid is essential to prevent false positives from carbonate ions, which would also form a white precipitate with barium chloride.

Halide ions (Cl⁻, Br⁻, I⁻)

1. Add dilute nitric acid (to remove interfering ions like carbonates)
2. Add silver nitrate solution (AgNO₃)
3. Observe the colour of the precipitate formed

Results:

• Chloride (Cl⁻): White precipitate of silver chloride
• Bromide (Br⁻): Cream precipitate of silver bromide
• Iodide (I⁻): Yellow precipitate of silver iodide

The ionic equation for chloride is: Ag⁺(aq) + Cl⁻(aq) → AgCl(s)

Further test: The precipitates have different solubilities in ammonia solution, which can help distinguish them, though this is beyond GCSE requirements.

Tests for water

Test for the presence of water

Use anhydrous copper(II) sulfate:

• Add the white anhydrous powder to the liquid
• Positive result: The powder turns blue
• The reaction is: CuSO₄(s) + 5H₂O(l) → CuSO₄·5H₂O(s)
• The blue compound formed is hydrated copper(II) sulfate

Alternatively, use anhydrous cobalt(II) chloride paper:

• The paper is blue when dry
• Positive result: Turns pink in the presence of water

Important: These tests only confirm water is present; they do not prove the sample is pure water.

Test for pure water

To determine if water is pure, measure its boiling point:

• Pure water boils at exactly 100°C (at standard atmospheric pressure of 1 atm/101 kPa)
• Impurities raise the boiling point
• Also measure freezing point: pure water freezes at exactly 0°C

For complete purity testing, you could also check the water has a pH of 7.

Worked examples

Example 1: Identifying an unknown gas

Question: A student adds dilute hydrochloric acid to a white solid. A gas is produced. When this gas is bubbled through limewater, the limewater turns cloudy. Name the gas produced and identify the negative ion present in the white solid. [3 marks]

Answer:

• Gas produced: carbon dioxide [1 mark]
• The limewater test is positive for carbon dioxide [1 mark]
• Negative ion present: carbonate (CO₃²⁻) [1 mark]

Examiner note: You must name both the gas and the ion. Simply stating "carbonate" without writing the ion formula may not gain full marks at higher tier.

Example 2: Metal ion identification

Question: A student performs tests on an unknown metal compound to identify the metal ion present.

When the compound is held in a Bunsen flame using a nichrome wire, an orange-red flame is observed.

When sodium hydroxide solution is added to a solution of the compound, a white precipitate forms.

The precipitate does not dissolve when excess sodium hydroxide is added.

Identify the metal ion present and explain your reasoning. [4 marks]

Answer:

• Metal ion: calcium (Ca²⁺) [1 mark]
• Orange-red flame is characteristic of calcium [1 mark]
• White precipitate with sodium hydroxide is consistent with calcium [1 mark]
• Calcium hydroxide does not dissolve in excess sodium hydroxide, unlike aluminium hydroxide, confirming it is calcium not aluminium [1 mark]

Examiner note: Both tests are necessary to conclusively identify calcium, as several metal ions form white precipitates.

Example 3: Halide ion testing

Question: Describe how you would test a solution to determine whether it contains chloride, bromide or iodide ions. Include the expected observations for each ion. [4 marks]

Answer:

• Add dilute nitric acid to the solution [1 mark]
• Then add silver nitrate solution [1 mark]
• Chloride produces a white precipitate, bromide produces a cream precipitate, iodide produces a yellow precipitate [2 marks — 1 mark for two correct, 2 marks for all three correct]

Alternative acceptable answer: Students may state "add acidified silver nitrate" for the first two marks.

Common mistakes and how to avoid them

• Confusing the gas tests: Students often mix up which test goes with which gas. Remember: squeaky pop for hydrogen, relighting splint for oxygen, limewater for carbon dioxide, bleaching for chlorine. Create a mnemonic or flashcards to memorise these.

• Forgetting to add acid before anion tests: Always add dilute hydrochloric acid before testing for sulfates (then add barium chloride) or dilute nitric acid before testing for halides (then add silver nitrate). This removes interfering ions that could give false positives.

• Not distinguishing between white precipitates: Calcium, magnesium and aluminium all give white precipitates with sodium hydroxide. Remember only aluminium hydroxide dissolves in excess sodium hydroxide. Use flame tests to distinguish calcium from magnesium.

• Thinking the water test proves purity: Anhydrous copper(II) sulfate turning blue proves water is present but does not prove the sample is pure water. Only measuring the boiling point (100°C) or freezing point (0°C) can confirm purity.

• Poor technique in flame tests: Not cleaning the wire loop properly between tests leads to sodium contamination (producing a persistent yellow flame). Always clean thoroughly with acid and heat until colourless.

• Writing incomplete ionic equations: When asked for ionic equations, omit spectator ions and include state symbols. For example, Cu²⁺(aq) + 2OH⁻(aq) → Cu(OH)₂(s), not CuSO₄ + 2NaOH → Cu(OH)₂ + Na₂SO₄.

Exam technique for "Required practical: tests for gases, ions and water"

• Learn the exact observations: Examiners want precise descriptions. Write "blue precipitate" not "blue substance" for copper(II) hydroxide. Write "squeaky pop" not "loud noise" for hydrogen. Vague answers lose marks.

• Command words matter: "Describe a test" requires both the method (what you add) and the observation (what you see). "Identify" just requires naming the substance. "Explain" requires reasoning linked to chemical reactions.

• Method questions need logical steps: When describing a test procedure, write numbered steps in order. State what you add first, what you add second, and what you observe. Don't forget safety precautions when asked.

• Extended response questions: For 4-6 mark questions involving multiple tests, structure your answer clearly. One test per paragraph. State the reagent, the observation, and what it proves. Link observations to conclusions using "therefore" or "this indicates."

Quick revision summary

Master the characteristic tests: squeaky pop (hydrogen), relighting splint (oxygen), limewater cloudy (carbon dioxide), bleached litmus (chlorine). For cations, learn flame test colours and hydroxide precipitate colours. For anions, remember carbonates fizz with acid, sulfates give white precipitate with barium chloride (after adding HCl), halides give coloured precipitates with silver nitrate (after adding HNO₃). Anhydrous copper sulfate turns blue when water is present; pure water boils at exactly 100°C. Practice writing ionic equations and precise observations for maximum marks.