Lesson 18: Required Practicals — Chemistry

Chemistry required practicals test your ability to make substances, measure changes, and analyse mixtures — the core skills of a chemist. This lesson covers making salts, electrolysis, temperature changes in reactions, rates of reaction, chromatography, and testing water. Each practical has a standard method examiners expect you to reproduce, specific apparatus (burettes, pipettes, thermometers), and characteristic calculations. Learn the methods as ordered sequences and the "why" behind each step.

Practical 1 — Making a Salt (Copper Sulphate)

Method: warm dilute sulphuric acid in a beaker (water bath or gentle Bunsen heat); add excess black copper oxide powder, stirring; the solution turns blue as copper sulphate forms; filter the mixture to remove unreacted solid; evaporate the filtrate gently until crystallisation point (dip a glass rod — crystals form on cooling); leave to crystallise; pat dry on filter paper. Equation: CuO + H₂SO₄ → CuSO₄ + H₂O.

Why each step: warming speeds the reaction; excess base ensures all acid reacts (no acid contaminates the salt); stirring mixes reactants; filtration removes the insoluble excess; evaporating to crystallisation point then cooling slowly grows large pure crystals. Safety: wear goggles (hot acid), handle hot apparatus with tongs.

Exam questions: write the method in order from jumbled steps; explain why excess solid is used; suggest how to make a different salt (e.g. zinc chloride from zinc oxide and HCl — apply the pattern); calculate the yield or atom economy. Insoluble salts (lead iodide) are made by precipitation instead — mixing lead nitrate and potassium iodide solutions, then filtering.

Practical 2 — Electrolysis of Aqueous Solutions

Method: set up inert (graphite) electrodes in the solution, connected to a DC power supply; pass current for a fixed time; observe and test products. For sodium chloride solution (brine): hydrogen at the cathode (squeaky pop test), chlorine at the anode (bleaches damp litmus paper), leaving sodium hydroxide. For copper sulphate with copper electrodes: copper deposits on the cathode (pink layer, mass increases) — the basis of electroplating and purification.

Predicting products: at the cathode, hydrogen is discharged unless the metal is very reactive (then the metal deposits — e.g. copper); at the anode, oxygen (from hydroxide) is discharged unless a halide is present in concentrated solution (then the halogen forms). Half-equations: 2H⁺ + 2e⁻ → H₂ (reduction); 2Cl⁻ − 2e⁻ → Cl₂ (oxidation). Remember OILRIG.

Exam questions: predict products for given solutions and electrodes; write half-equations; explain observations (e.g. "the cathode gains a pink coating — why?"); describe electroplating (object as cathode in a solution of the plating metal). Variables in an investigation: current, time, concentration — link charge passed (Q = It) to mass deposited for higher-tier calculations.

Practical 3 — Temperature Changes (Neutralisation/Displacement)

Method: measure 25 cm³ of dilute HCl into a polystyrene cup (insulator) with a lid; record its temperature; add a measured volume of NaOH; stir, record the highest temperature reached; repeat with different volumes of alkali; plot temperature rise against volume — the peak shows the neutralisation point (stoichiometric ratio). Alternatively, add metals to acid and measure temperature rises to compare reactivity.

Variables: independent — volume of alkali (or metal type); dependent — temperature change; controls — acid concentration and volume, starting temperature, same cup, stirring. Polystyrene cups minimise heat loss; lids reduce it further. Repeats and means improve reliability. Safety: goggles — acids and alkalis; wipe spills.

Exam questions: interpret temperature–volume graphs (explain the rise then fall — excess alkali dilutes and cools the mixture); calculate energy change (higher-tier: q = mcΔT); explain heat losses and how to reduce them (insulation, lids); link to exothermic (temperature rises — neutralisation, combustion) versus endothermic (temperature falls — thermal decomposition, sports injury packs).

Practical 4 — Rates of Reaction

Two standard methods. (a) Measuring gas: react magnesium with HCl in a conical flask connected to a gas syringe; record gas volume every 15 seconds; plot volume against time — the gradient is the rate; the curve flattens when the reaction finishes. (b) Disappearing cross: mix sodium thiosulphate with HCl in a flask over a cross drawn on paper; time how long until the cross disappears (sulphur precipitate clouds the mixture); repeat at different temperatures or concentrations; rate = 1 ÷ time.

Variables: independent — concentration, temperature, surface area, or catalyst; dependent — rate (gradient or 1/time); controls — everything else. Collision theory explains all results: higher concentration/temperature/surface area → more frequent (and, for temperature, more energetic) collisions → faster rate. Catalysts provide an alternative pathway with lower activation energy — unchanged at the end.

Exam questions: describe either method in full; sketch and interpret rate graphs (steeper = faster; same final volume = same amount of reactant); explain results with collision theory; evaluate (e.g. "the cross method is subjective — how could timing be made more objective?" → use a colorimeter or light sensor). Calculating rates from gradients (volume ÷ time) is a key maths skill.

Practical 5 — Chromatography

Method: draw a pencil baseline near the bottom of chromatography paper (pencil — ink would dissolve and interfere); spot the inks/mixtures on the line; suspend in a beaker with a small depth of solvent (below the baseline); the solvent rises, carrying components at different rates; remove before the solvent reaches the top; mark the solvent front; calculate Rf = distance moved by spot ÷ distance moved by solvent front.

Analysis: each substance has a characteristic Rf in a given solvent; compare spots to known samples to identify components. A pure substance gives one spot; a mixture gives several. Applications: testing the purity of drugs, analysing inks in forensics, checking food dyes. Safety: some solvents are flammable/toxic — keep away from flames, work ventilated.

Exam questions: calculate Rf values from diagrams (measure carefully, same units); identify which ink contains which dyes by matching spots; explain why the baseline is pencil; suggest why two spots have the same Rf (same substance, or coincidence — test in a different solvent). Paper chromatography of amino acids or plant pigments (chlorophyll) are classic contexts.

Practical 6 — Testing Water and Analysis

Tests for gases (also examined here): hydrogen — squeaky pop with a lit splint; oxygen — relights a glowing splint; carbon dioxide — turns limewater milky; chlorine — bleaches damp litmus paper. Tests for ions (higher-tier): flame tests (lithium crimson, sodium yellow, potassium lilac, calcium orange-red, copper green); sodium hydroxide test for metal ions (coloured precipitates — copper blue, iron(II) green, iron(III) brown, aluminium/zinc white, dissolving in excess); carbonate test (add acid — CO₂ turns limewater milky); sulphate test (acidified barium chloride — white precipitate); halides (acidified silver nitrate — white/cream/yellow precipitates).

Water purity analysis: test samples for pH, dissolved solids (evaporate and weigh), and specific ions. Potable water must be safe: filtration removes solids, chlorination/UV kills microbes. Exam scenarios give you a table of test results for unknown solutions and ask you to identify each — work through systematically, using positive results to confirm and negative results to eliminate.

Required practical on water: investigate the quality of water samples (e.g. tap, rain, pond) by measuring pH and evaporating to find dissolved solids. Evaluate: which is purest? Is it safe to drink? (Purity ≠ safety — distilled water is pure but pond water may contain pathogens.) Link to environmental chemistry: testing river water downstream of a factory for pollutants.

Conclusion — Why This Matters

Chemistry required practicals are where theory becomes substance: you make salts, split compounds with electricity, measure energy changes, time reactions, separate mixtures and analyse water. For the exam, each practical must be recallable as an ordered method with apparatus, variables, safety points and data processing — from the crystallisation point in salt preparation to the Rf calculation in chromatography. The calculations woven through them (moles in titrations, q = mcΔT, rates from gradients) are the mathematical backbone of GCSE chemistry — practise them in context until fluent. Evaluation follows the universal template: repeatability, validity, accuracy, specific improvements. Notice how the practicals connect the topics: electrolysis applies bonding and the reactivity series; rates apply collision theory; salt preparation applies acid-base patterns. This interconnectedness is what examiners probe with synoptic questions — the student who sees one chemistry, not six isolated practicals, is the student who earns the top grades. And practically: these are real laboratory skills, the foundation of any future in science, medicine or engineering.

Worked example — rates from the disappearing cross

Sodium thiosulfate reacts with hydrochloric acid, forming sulfur that clouds the mixture. At 20 °C the cross disappears after 96 s; at 40 °C after 24 s. Rates (1 ÷ time) are 0.0104 s⁻¹ and 0.0417 s⁻¹ — the rate has roughly quadrupled for a 20 °C rise, consistent with the rule of thumb that rate doubles for each 10 °C. The explanation must mention both collision frequency and, decisively, the proportion of collisions with energy above the activation energy. Evaluation: judging cloudiness by eye is subjective — a light sensor or colorimeter measuring transmitted light gives objective, continuous data, and the same observer (or sensor) should be used throughout.

Worked example — chromatography Rf values

A food colouring spot travelled 4.2 cm while the solvent front reached 7.0 cm, giving Rf = 4.2 ÷ 7.0 = 0.60. A second spot at 0.60 in the same solvent matches a reference dye run alongside — matching Rf values in the same conditions are evidence (not absolute proof) that the substances are the same.

Common mistakes in this set of practicals

• Drawing the chromatography baseline in pen — pen ink is itself a mixture of dyes and will chromatograph; always use pencil, and keep the baseline above the solvent level.
• Evaporating a salt solution to dryness when making crystals — crystallisation needs slow evaporation to the point of crystallisation, then cooling; boiling dry decomposes hydrated crystals.
• Heating flammable liquids with a naked flame — temperature-change and evaporation experiments with ethanol need a water bath.

Exam Experience

Question 1. Describe the tests for hydrogen and for carbon dioxide, including the positive results. [2 marks]

Question 2. In a chromatography experiment, a dye spot moved 4.8 cm from the baseline and the solvent front moved 6.4 cm. Calculate the Rf value of the dye. [2 marks]

Question 3. A student electrolyses copper sulfate solution using inert (graphite) electrodes. Describe the observations she would make at each electrode, explain the products formed, and state how each product could be tested or identified. [6 marks]

Mark scheme

Q1: hydrogen — a lit splint at the mouth of the tube gives a squeaky pop (1); carbon dioxide — bubble through limewater, which turns milky/cloudy (1).

Q2: Rf = distance moved by dye ÷ distance moved by solvent = 4.8 ÷ 6.4 = 0.75 (1 substitution, 1 answer; no unit).

Q3: Level 3 (5–6 marks): observations and explanations for both electrodes with identification — at the negative electrode (cathode) a pink-brown/orange coating of copper forms, because copper ions (Cu²⁺) are positively charged, are attracted to the cathode, and gain electrons to form copper atoms (reduction); at the positive electrode (anode) bubbles of a colourless gas form — oxygen, because hydroxide ions from the water discharge in preference to sulfate ions; the blue colour of the solution fades as copper is removed. Identification: the gas relights a glowing splint (oxygen); the coating's appearance and the mass increase of the cathode identify the copper. Level 2 (3–4 marks): products correct for both electrodes but explanations (ion movement or discharge) thin, or tests missing. Level 1 (1–2 marks): one product named with no explanation.