Required Practicals
AQA A-Level Chemistry has 12 required practicals that you must understand in detail. Although you complete these in the laboratory during the course, they are assessed through written questions in the exams — particularly in Paper 3, Section A. You need to know the method, apparatus, variables, safety precautions, expected results, and common sources of error for each one. Examiners frequently ask you to apply your knowledge of these practicals to unfamiliar contexts.
Required Practical 1: Volumetric Solution and Acid-Base Titration
Aim
Make up a volumetric solution of known concentration and carry out a simple acid-base titration to determine the concentration of an unknown solution.
Method
- Weigh out the required mass of solute using a balance accurate to at least 2 decimal places.
- Dissolve the solute in a small amount of distilled water in a beaker, stirring until fully dissolved.
- Transfer quantitatively to a volumetric flask using a wash bottle to rinse all solution from the beaker.
- Make up to the calibration mark with distilled water, using a dropping pipette for the final drops. The bottom of the meniscus must sit on the line.
- Stopper and invert several times to ensure thorough mixing.
- Using a pipette and pipette filler, transfer a known volume (e.g. 25.0 cm³) of one solution into a conical flask.
- Add a suitable indicator (e.g. methyl orange for strong acid–weak base, phenolphthalein for weak acid–strong base).
- Fill a burette with the other solution, recording the initial reading.
- Add solution from the burette, swirling the conical flask, until the endpoint colour change is observed.
- Record the final burette reading and calculate the titre.
- Repeat until concordant titres are obtained (within 0.10 cm³ of each other).
Key Variables
- Independent: Volume of solution added from the burette
- Dependent: Volume at the endpoint (titre)
- Controlled: Volume of solution in the conical flask, concentration of the standard solution, temperature, indicator
Safety
- Wear safety goggles — acids and alkalis are corrosive or irritant.
- Handle glassware with care — volumetric flasks and burettes are fragile.
- Wash hands after handling chemicals.
Expected Results and Calculations
Use the titration formula: n = c × V (amount in moles = concentration × volume in dm³).
From the balanced equation, find the mole ratio, and hence calculate the unknown concentration.
Common Exam Angles
- Why is a conical flask used rather than a beaker? (Easier to swirl without spillage; shape does not affect the titre.)
- Why use a white tile under the flask? (To see the colour change clearly.)
- Why are rough titres not included in the average? (They are used to find the approximate endpoint; concordant titres give more accurate results.)
- How to improve accuracy: use a more precise balance, ensure the meniscus is read at eye level, use a pipette rather than a measuring cylinder.
- Sources of systematic error: not rinsing the pipette with the solution, using tap water instead of distilled water, reading the meniscus incorrectly.
Required Practical 2: Measurement of an Enthalpy Change (Calorimetry)
Aim
Measure the enthalpy change of a reaction using calorimetry (e.g. neutralisation, displacement, or dissolving).
Method
- Measure a known volume of solution using a measuring cylinder and pour into a polystyrene cup (acts as an insulated calorimeter).
- Record the initial temperature using a thermometer accurate to 0.1 °C.
- Add the second reactant (e.g. acid, metal powder, or solid).
- Stir the mixture and record the temperature at regular intervals (e.g. every 30 seconds) for several minutes.
- Identify the maximum or minimum temperature reached.
- Calculate the temperature change, ΔT.
Calculation
Use q = mcΔT, where:
- q = energy change (J)
- m = mass of solution (g) — assume density of solution is 1.00 g cm⁻³
- c = specific heat capacity (4.18 J g⁻¹ K⁻¹ for dilute aqueous solutions)
- ΔT = temperature change (K or °C — the magnitude is the same)
Then: ΔH = −q / n, where n is the number of moles of the limiting reagent.
Key Variables
- Independent: The type of reaction or the identity of the reactant
- Dependent: The temperature change
- Controlled: Volume and concentration of solution, insulation, stirring method
Safety
- Wear safety goggles. Some reactions are exothermic and solutions may splash.
- If using metals (e.g. zinc), be aware of hydrogen gas production — no naked flames.
- Polystyrene cups can be knocked over — place inside a beaker for stability.
Common Exam Angles
- Why is the value obtained less exothermic than the literature value? (Heat loss to the surroundings; not all energy is transferred to the solution.)
- How to reduce heat loss: use a lid, use better insulation, extrapolate a cooling curve back to the time of mixing.
- Why assume the specific heat capacity of water? (Dilute solutions behave similarly to water; the dissolved solute contributes negligibly to the total mass.)
- Drawing and interpreting a temperature–time graph, including extrapolation to correct for heat loss.
Required Practical 3: Rate of Reaction and Temperature
Aim
Investigate how the rate of reaction changes with temperature, and use the data to calculate the activation energy.
Method
- Set up a reaction whose rate can be measured (e.g. sodium thiosulfate and hydrochloric acid — the cross disappearing method, or the iodine clock reaction).
- Carry out the reaction at several different temperatures (e.g. 20 °C, 30 °C, 40 °C, 50 °C, 60 °C), using a water bath to control temperature.
- Record the time taken for the observable change to occur at each temperature.
- Calculate the rate at each temperature (rate = 1/time if monitoring a fixed change).
Activation Energy Calculation
Using the Arrhenius equation in its logarithmic form:
ln k = ln A − Ea/(RT)
A plot of ln k (or ln(1/t)) against 1/T gives a straight line with gradient = −Ea/R.
Therefore: Ea = −gradient × R, where R = 8.314 J K⁻¹ mol⁻¹.
Key Variables
- Independent: Temperature
- Dependent: Time for the observable change (and hence rate)
- Controlled: Concentrations of solutions, volumes, the criterion for the observable change
Safety
- Sodium thiosulfate and acid produce sulfur dioxide — work in a well-ventilated area or a fume cupboard.
- Take care with hot water baths. Do not heat above 60 °C with a Bunsen burner near the reaction.
Common Exam Angles
- Why does rate approximately double for every 10 °C rise? (A greater proportion of molecules exceeds the activation energy according to the Boltzmann distribution.)
- Why use a water bath rather than direct heating? (Provides more even, controllable temperature.)
- How to plot and interpret an Arrhenius plot (ln k vs 1/T).
- Why is 1/t only proportional to rate if the measured change is always the same?
Required Practical 4: Tests for Cations, Anions, and Gases
Aim
Carry out simple test-tube reactions to identify cations (metal ions), anions, and gases.
Cation Tests (using NaOH(aq) and NH₃(aq))
| Ion | NaOH(aq) added | Excess NaOH(aq) | NH₃(aq) added | Excess NH₃(aq) |
|---|
| Al³⁺ | White ppt | Dissolves (amphoteric) | White ppt | No change |
| Ca²⁺ | White ppt | No change | No ppt (or slight) | — |
| Cu²⁺ | Blue ppt | No change | Blue ppt | Dissolves → deep blue solution |
| Fe²⁺ | Green ppt | No change | Green ppt | No change |
| Fe³⁺ | Brown ppt | No change | Brown ppt | No change |
| Mn²⁺ | Pale pink/buff ppt | No change | Pale pink/buff ppt | No change |
Anion Tests
| Ion | Test | Positive Result |
|---|
| CO₃²⁻ | Add dilute acid | Effervescence; gas turns limewater milky (CO₂) |
| SO₄²⁻ | Add dilute HCl then BaCl₂(aq) | White precipitate (BaSO₄) insoluble in excess HCl |
| Cl⁻ | Add dilute HNO₃ then AgNO₃(aq) | White precipitate (AgCl); soluble in dilute NH₃ |
| Br⁻ | Add dilute HNO₃ then AgNO₃(aq) | Cream precipitate (AgBr); soluble in concentrated NH₃ |
| I⁻ | Add dilute HNO₃ then AgNO₃(aq) | Yellow precipitate (AgI); insoluble in NH₃ |
Gas Tests
| Gas | Test | Positive Result |
|---|
| H₂ | Burning splint | Squeaky pop |
| O₂ | Glowing splint | Relights |
| CO₂ | Bubble through limewater | Turns milky/cloudy |
| Cl₂ | Damp litmus paper | Bleaches (turns white) |
| NH₃ | Damp red litmus paper | Turns blue |
| SO₂ | Bubble through acidified potassium dichromate | Orange to green |
Common Exam Angles
- Why add dilute HCl before BaCl₂ when testing for sulfate? (To remove carbonate and sulfite ions, which would also give white precipitates with Ba²⁺.)
- Why add dilute HNO₃ before AgNO₃ when testing for halides? (To remove carbonate and other ions that form insoluble silver salts.)
- Distinguish between Al³⁺ and Ca²⁺ using excess NaOH (Al(OH)₃ dissolves — amphoteric; Ca(OH)₂ does not).
- Distinguish between Cu²⁺ and Fe²⁺/Fe³⁺ using excess NH₃ (Cu²⁺ gives deep blue [Cu(NH₃)₄]²⁺).
Required Practical 5: Distillation of a Product from a Reaction
Aim
Distil and purify a product from an organic reaction (e.g. preparation of cyclohexene from cyclohexanol by dehydration using an acid catalyst).
Method
- Set up the reaction mixture in a round-bottomed flask with anti-bumping granules.
- Add the acid catalyst (e.g. concentrated phosphoric acid or sulfuric acid).
- Heat gently using a heating mantle or water bath. Attach a distillation apparatus with a condenser and receiving flask.
- Collect the distillate over the correct temperature range (matching the boiling point of the product).
- Purify the distillate by washing with sodium carbonate solution (to remove acidic impurities) in a separating funnel.
- Dry the organic layer using anhydrous calcium chloride or sodium sulfate.
- Redistil to obtain a pure product and record the boiling point range.
Key Variables
- Independent: Not typically varied — this is a preparative practical
- Dependent: Yield, purity (boiling point range)
- Controlled: Masses/volumes of reagents, heating rate
Safety
- Concentrated acids are highly corrosive — wear gloves and goggles.
- Organic liquids are flammable — no naked flames. Use an electric heating mantle.
- Carry out in a fume cupboard if vapours are produced.
Common Exam Angles
- Purpose of anti-bumping granules (prevent superheating and sudden boiling).
- Why use a condenser (to cool vapour and collect the liquid product).
- Why wash with sodium carbonate (to neutralise acidic impurities).
- How to confirm the identity and purity of the product (melting point/boiling point, IR spectroscopy, mass spectrometry).
- Calculating percentage yield and atom economy.
Required Practical 6: Tests for Organic Functional Groups
Aim
Test for the presence of alcohol, aldehyde, alkene, and carboxylic acid functional groups.
Tests