Required Practicals
The AQA GCSE Chemistry specification includes eight required practicals that you must understand thoroughly. In the exam, you will be asked questions about the method, apparatus, variables, results and evaluation of these practicals — even if you did not perform them yourself during the course. This lesson covers every required practical in the detail that AQA expects.
Why Required Practicals Matter
- Required practicals account for approximately 15% of the total marks across both papers.
- Questions can appear on either Paper 1 or Paper 2, depending on the topic.
- You may be asked to describe the method, identify variables, interpret results, suggest improvements or evaluate sources of error.
- Questions often combine practical knowledge with theory — for example, asking you to explain the results using collision theory or equilibrium principles.
Required Practical 1: Making Salts — Preparing a Pure, Dry Sample of a Soluble Salt from an Insoluble Oxide or Carbonate
Paper 1 — Topic 5.4 Chemical Changes
Aim
To prepare a pure, dry sample of a soluble salt (e.g., copper sulfate) by reacting an insoluble base (e.g., copper oxide) with a dilute acid (e.g., sulfuric acid).
Equipment
- Dilute sulfuric acid (approximately 1 mol/dm³)
- Copper oxide powder (insoluble base)
- Beaker (250 cm³)
- Measuring cylinder (50 cm³)
- Glass stirring rod
- Bunsen burner, tripod and gauze
- Filter funnel, filter paper and conical flask
- Evaporating basin
- Spatula
- Eye protection (safety goggles)
Method
- Measure 25 cm³ of dilute sulfuric acid into a beaker using a measuring cylinder.
- Warm the acid gently using a Bunsen burner (do NOT boil).
- Add one spatula of copper oxide powder to the warm acid and stir with a glass rod.
- Continue adding copper oxide powder, one spatula at a time, stirring after each addition, until the copper oxide is in excess (some black powder remains undissolved at the bottom).
- The excess copper oxide ensures that all the acid has been neutralised.
- Filter the mixture through filter paper in a funnel to remove the excess copper oxide. The filtrate (blue liquid) is the copper sulfate solution.
- Pour the filtrate into an evaporating basin.
- Heat the solution gently until about half the water has evaporated (crystals should start to appear at the edges).
- Leave the evaporating basin in a warm place for slow crystallisation to occur over several days. Slow evaporation produces larger, more regular crystals.
- Pat dry the crystals with filter paper.
Variables
| Variable | Detail |
|---|
| Independent Variable (IV) | Not applicable — this is a preparation, not an investigation |
| Dependent Variable (DV) | Mass of salt crystals produced |
| Control Variables | Volume and concentration of acid, type of base, temperature of acid |
Key Equation
CuO + H₂SO₄ → CuSO₄ + H₂O
Copper oxide + sulfuric acid → copper sulfate + water
Expected Results
- The black copper oxide dissolves in the acid, producing a blue solution of copper sulfate.
- When copper oxide is in excess, unreacted black powder remains at the bottom.
- After filtering and evaporating, blue crystals of hydrated copper sulfate form.
Common Errors
- Not adding excess base — this means unreacted acid remains, and the salt is impure.
- Boiling the solution to dryness instead of allowing slow crystallisation — this produces tiny, poor-quality crystals or a powder.
- Not filtering carefully — solid base particles contaminate the salt.
- Forgetting to warm the acid first — the reaction is very slow with cold acid.
Safety
- Wear eye protection at all times.
- Sulfuric acid is corrosive — avoid skin contact and wipe spills immediately.
- Take care when heating — do not boil dry.
How This Is Examined
- Describe the full method for preparing a soluble salt.
- Explain why the base is added in excess.
- Explain why the mixture is filtered.
- Explain why slow evaporation produces better crystals.
- Write a balanced symbol equation for the reaction.
Required Practical 2: Electrolysis — Investigating the Electrolysis of Aqueous Solutions
Paper 1 — Topic 5.4 Chemical Changes
Aim
To investigate what is produced at each electrode during the electrolysis of different aqueous solutions using inert (carbon or platinum) electrodes.
Equipment
- Power supply (DC, low voltage) or battery pack
- Two carbon (graphite) electrodes
- Beaker (250 cm³)
- Connecting leads with crocodile clips
- Solutions: copper sulfate, sodium chloride, sodium sulfate (each approximately 1 mol/dm³)
- Test tubes (to collect gases over water)
- Splint, litmus paper
- Eye protection
Method
- Pour the aqueous solution into a beaker to a depth of about 5 cm.
- Connect two carbon electrodes to a DC power supply using connecting leads.
- Place both electrodes into the solution, ensuring they do not touch each other.
- Turn on the power supply and set it to about 4–6 V.
- Observe what happens at each electrode:
- At the anode (positive electrode): look for bubbles (gas produced) or a deposit.
- At the cathode (negative electrode): look for bubbles or a coating/deposit on the electrode.
- Test any gas produced:
- Collect gas in an inverted test tube over water.
- Test for oxygen using a glowing splint (relights).
- Test for hydrogen using a burning splint (squeaky pop).
- Test for chlorine using damp litmus paper (bleaches it white).
- Repeat with different aqueous solutions.
Variables
| Variable | Detail |
|---|
| Independent Variable (IV) | The aqueous solution used |
| Dependent Variable (DV) | Products formed at each electrode |
| Control Variables | Voltage, size and type of electrodes, volume and concentration of solution, time |
Expected Results
| Solution | Cathode Product | Anode Product | Explanation |
|---|
| Copper sulfate | Copper (pink/brown coating) | Oxygen (bubbles, relights splint) | Cu²⁺ is less reactive than hydrogen, so copper is deposited. SO₄²⁻ cannot be discharged, so water is oxidised to O₂. |
| Sodium chloride | Hydrogen (bubbles, squeaky pop) | Chlorine (bubbles, bleaches litmus) | Na⁺ is more reactive than hydrogen, so H₂ is produced. Cl⁻ is a halide, so Cl₂ is produced (halide rule). |
| Sodium sulfate | Hydrogen (bubbles, squeaky pop) | Oxygen (bubbles, relights splint) | Na⁺ is more reactive than hydrogen, so H₂ is produced. SO₄²⁻ cannot be discharged, so O₂ is produced. |
Common Errors
- Using AC instead of DC — electrolysis requires direct current.
- Electrodes touching — this causes a short circuit.
- Not testing gases correctly — must use correct gas tests.
- Confusing anode and cathode (remember: AN OX — anode is where oxidation occurs; RED CAT — reduction at the cathode).
Safety
- Wear eye protection.
- Chlorine gas is toxic — perform the experiment in a well-ventilated room or fume cupboard.
- Do not touch electrodes while the circuit is on.
Required Practical 3: Temperature Changes — Investigating the Variables That Affect Temperature Changes in Reactions
Paper 1 — Topic 5.5 Energy Changes
Aim
To investigate the temperature change when different variables are changed in a neutralisation reaction (e.g., adding different volumes of acid to a fixed volume of alkali).
Equipment
- Polystyrene cup (insulating container)
- Thermometer (–10 to 110 °C, resolution 0.5 °C or better) or temperature probe
- Measuring cylinders (25 cm³ and 50 cm³)
- Dilute hydrochloric acid (1 mol/dm³)
- Dilute sodium hydroxide solution (1 mol/dm³)
- Lid (with hole for thermometer)
- Eye protection
Method
- Use a measuring cylinder to measure 25 cm³ of sodium hydroxide solution into a polystyrene cup.
- Record the initial temperature of the NaOH.
- Measure 5 cm³ of dilute hydrochloric acid using a measuring cylinder.
- Add the acid to the alkali in the polystyrene cup, stir gently with the thermometer.
- Record the maximum temperature reached.
- Calculate the temperature change (ΔT = final temperature − initial temperature).
- Repeat with 10 cm³, 15 cm³, 20 cm³, 25 cm³, 30 cm³ and 35 cm³ of acid, using a fresh 25 cm³ of NaOH each time.
Variables
| Variable | Detail |
|---|
| Independent Variable (IV) | Volume of acid added |
| Dependent Variable (DV) | Temperature change (ΔT) |
| Control Variables | Volume and concentration of NaOH, concentration of acid, starting temperature, insulation |
Expected Results
- Temperature increases as acid volume increases (exothermic neutralisation).
- Temperature reaches a maximum when the acid and alkali are in the exact stoichiometric ratio (equivalence point, approximately 25 cm³ of acid for 25 cm³ of 1 mol/dm³ NaOH with 1 mol/dm³ HCl).
- Beyond the equivalence point, adding more acid causes the temperature to decrease because the excess cold acid cools the mixture without reacting.
Common Errors
- Not using a polystyrene cup — a glass beaker allows too much heat loss.
- Not stirring — uneven temperature distribution.
- Not recording the starting temperature of the NaOH each time.
- Misreading the thermometer — always read at eye level to the meniscus.
- Not using a lid — heat loss to surroundings reduces the temperature change.
Safety
- Wear eye protection.
- HCl and NaOH are irritants/corrosive — avoid skin contact.
Required Practical 4: Rates of Reaction — Investigating the Effect of a Variable on the Rate of a Chemical Reaction
Paper 2 — Topic 5.6 Rates of Reaction
Aim
To investigate how changing the concentration of a reactant affects the rate of reaction. A common method involves measuring the volume of gas produced over time when magnesium reacts with hydrochloric acid, or measuring the time for a cross to disappear in the sodium thiosulfate–hydrochloric acid reaction.
Method A: Gas Collection (Magnesium and HCl)
Equipment: Conical flask, delivery tube, gas syringe or inverted measuring cylinder in water trough, stopwatch, magnesium ribbon (cut to identical 3 cm lengths), dilute hydrochloric acid at different concentrations, measuring cylinder, balance, eye protection.
- Measure 25 cm³ of HCl (e.g., 2 mol/dm³) into a conical flask.
- Place a 3 cm strip of magnesium ribbon into the flask and immediately connect the gas syringe/delivery tube.
- Start the stopwatch.
- Record the volume of gas collected every 10 seconds until the reaction stops.
- Repeat with HCl at 1.5 mol/dm³, 1.0 mol/dm³, 0.5 mol/dm³ and 0.25 mol/dm³, using a fresh piece of magnesium ribbon each time.
Method B: Disappearing Cross (Sodium Thiosulfate and HCl)
Equipment: Conical flask, white paper with a black cross drawn on it, stopwatch, sodium thiosulfate solution at different concentrations, dilute HCl (fixed volume and concentration), measuring cylinders, eye protection.
- Place a piece of paper with a black cross underneath a conical flask.
- Measure 25 cm³ of sodium thiosulfate solution into the flask.
- Add 5 cm³ of dilute HCl to the flask and start the stopwatch.
- Look down through the top of the flask. Stop the stopwatch when the cross can no longer be seen (due to the sulfur precipitate making the solution opaque).
- Record the time.
- Repeat with different concentrations of sodium thiosulfate, keeping HCl volume and concentration constant.
Variables
| Variable | Detail |
|---|
| Independent Variable (IV) | Concentration of HCl (Method A) or concentration of Na₂S₂O₃ (Method B) |
| Dependent Variable (DV) | Volume of gas produced per unit time (A) or time for cross to disappear (B) |
| Control Variables | Temperature, volume of solutions, mass/length of magnesium (A), volume and concentration of HCl (B) |
Expected Results
- Higher concentration → faster rate of reaction.
- Method A: steeper curve on a volume-time graph for higher concentrations.
- Method B: shorter time for the cross to disappear at higher concentrations.
- Explanation: Higher concentration means more particles per unit volume → more frequent collisions → more successful collisions per second → faster rate.
Common Errors
- Not cleaning magnesium ribbon with sandpaper to remove the oxide layer.
- Gas leaking from the apparatus (poor seal on gas syringe/delivery tube).
- Subjective endpoint in Method B — different people may judge the disappearance of the cross differently.
- Not controlling temperature — carry out all experiments in the same session at the same room temperature.
Safety
- Wear eye protection.
- HCl is an irritant.
- In Method B, sulfur dioxide gas may be produced — ensure good ventilation.
Required Practical 5: Chromatography — Investigating How Paper Chromatography Can Be Used to Separate and Identify Substances
Paper 2 — Topic 5.8 Chemical Analysis
Aim
To use paper chromatography to separate and identify the components of a mixture (e.g., food colourings or inks).
Equipment
- Chromatography paper
- Beaker (250 cm³) with a lid (watch glass or cling film)
- Solvent (water or ethanol, depending on the substances being separated)
- Pencil (NOT pen — ink from a pen would run)
- Ruler
- Capillary tube or fine dropper (for spotting)
- Known reference samples of pure substances
- Eye protection
Method