AQA GCSE Chemistry lists eight required practical activities, with titration, identifying ions and water purification on the separate Chemistry specification only. Many Chemistry practical questions combine technique with calculation, so practise the maths as well as the method.
Preparing a pure, dry sample of a soluble salt from an insoluble oxide or carbonate, heating dilute acid with a Bunsen burner and evaporating with a water bath or electric heater.
Aim
Produce a pure, dry sample of a soluble salt by neutralising an acid with an excess of an insoluble base.
Method
Warm about 25 cm³ of dilute acid in a beaker using a Bunsen burner, then stop heating.
Add the insoluble oxide or carbonate a spatula at a time, stirring, until no more reacts and solid remains in excess.
Filter the mixture to remove the unreacted solid, keeping the filtrate.
Pour the filtrate into an evaporating basin and heat gently with a water bath or electric heater until crystals start to form.
Leave the solution to cool so crystals form slowly.
Pat the crystals dry between filter paper.
Variables
Independent
The acid and base combination chosen, which determines the salt made
Dependent
Mass and purity of the dry crystals produced
Control
Volume and concentration of acid
Excess of base added
Same drying method
Equipment
Dilute acid, Insoluble metal oxide or carbonate, Beaker and stirring rod, Filter funnel and paper, Evaporating basin, Bunsen burner or electric heater.
Results and observations
The reaction mixture fizzes with a carbonate as carbon dioxide is released. Once the reaction is complete, solid remains undissolved and the filtrate is a clear salt solution that yields crystals on evaporation.
Calculations
Percentage yield = (mass of salt actually made ÷ maximum theoretical mass) × 100
Theoretical mass comes from a balanced equation and relative formula masses
Graphs and data
No graph is needed; results are recorded as observations and masses.
Evaluation
Adding the base in excess ensures all the acid reacts, so the final solution is not acidic.
Filtering removes the excess solid, which is essential for purity.
Evaporating to dryness over a flame can decompose the salt or cause spitting; slow crystallisation gives better crystals.
Yield is usually below 100% because some solution is lost in filtering and transferring.
Common mistakes
Adding just enough base rather than an excess.
Boiling the solution dry instead of leaving it to crystallise.
Forgetting that the excess solid must be filtered off before evaporation.
Quick self-test
Why is the insoluble base added in excess?+
To make sure all of the acid reacts, so no acid remains in the final solution.
How do you know the reaction is complete?+
Solid stops dissolving and remains at the bottom of the beaker, and with a carbonate the fizzing stops.
Why is the solution left to crystallise rather than evaporated to dryness?+
Slow crystallisation gives larger, purer crystals and avoids decomposing the salt or losing it through spitting.
Ready to test yourself? Load this practical into Recall and revise it actively.
Required practical 2 · separate chemistry only
Titration
Determining the reacting volumes of a strong acid and a strong alkali by titration, and (Higher Tier) calculating the concentration of one solution in mol/dm³ and g/dm³.
Aim
Find accurately what volume of acid reacts with a known volume of alkali, and use that to find an unknown concentration.
Method
Rinse the pipette with the alkali and use it to transfer a fixed volume, typically 25.0 cm³, into a conical flask.
Add a few drops of indicator such as phenolphthalein or methyl orange.
Rinse the burette with the acid, fill it and record the initial reading to the nearest 0.05 cm³.
Run acid in quickly to get a rough titre, noting the approximate end point.
Repeat, adding dropwise near the end point, and record the volume at the first permanent colour change.
Repeat until you have concordant titres within 0.10 cm³ and calculate a mean of those.
Variables
Independent
Volume of acid added
Dependent
Volume of acid required to neutralise the alkali (the titre)
Control
Same volume and concentration of alkali
Same indicator and number of drops
Same end-point judgement
Equipment
Burette and stand, Pipette and filler, Conical flask, White tile, Indicator, Acid and alkali solutions.
Results and observations
Phenolphthalein turns from pink to colourless when titrating acid into alkali; methyl orange turns from yellow to red. The titre is the volume at the first permanent colour change.
Calculations
Mean titre uses concordant results only
moles = concentration (mol/dm³) × volume (dm³), where volume in dm³ = cm³ ÷ 1000
Use the balanced equation's ratio to find moles of the unknown, then concentration = moles ÷ volume
Concentration in g/dm³ = concentration in mol/dm³ × relative formula mass
Graphs and data
No graph is required; results go in a titration table showing initial, final and titre values to two decimal places.
Evaluation
Only concordant titres are averaged because the rough titre overshoots the end point.
Rinsing apparatus with the solution it will hold prevents dilution errors.
A white tile makes the colour change easier to see, reducing end-point error.
Overshooting even one drop can change the titre measurably.
Common mistakes
Including the rough titre in the mean.
Recording burette readings to the nearest whole cm³ instead of 0.05 cm³.
Forgetting to convert cm³ to dm³ in the moles calculation.
Quick self-test
Why is the pipette rinsed with the alkali before use?+
Any water left inside would dilute the alkali and change the number of moles transferred.
25.0 cm^3 of alkali needs 20.0 cm^3 of 0.100 mol/dm^3 acid in a 1:1 reaction. Calculate the alkali concentration.+
Moles of acid = 0.100 x 0.0200 = 0.00200 mol, so the alkali is 0.00200 / 0.0250 = 0.0800 mol/dm^3.
Ready to test yourself? Load this practical into Recall and revise it actively.
Required practical 3
Electrolysis
Investigating what happens when aqueous solutions are electrolysed using inert electrodes, developing and testing a hypothesis.
Aim
Predict and identify the products formed at each electrode when different aqueous solutions are electrolysed.
Method
Half fill an electrolysis cell or beaker with the solution to be tested.
Place two inert graphite electrodes in the solution, not touching, and connect to a d.c. power supply.
Position an inverted test tube over each electrode to collect any gas.
Switch on and observe both electrodes, recording bubbling, colour changes or deposits.
Test the gases collected: a lit splint pops for hydrogen; a glowing splint relights in oxygen; damp litmus paper is bleached by chlorine.
Repeat with the other solutions and compare with your predictions.
Variables
Independent
The aqueous solution being electrolysed
Dependent
The products formed at the cathode and the anode
Control
Same voltage and time
Same inert electrodes
Same volume and concentration of solution
Equipment
d.c. power supply, Graphite (inert) electrodes, Electrolysis cell or beaker, Test tubes, Solutions such as copper chloride, sodium chloride and sodium sulfate, Splints and litmus paper.
Results and observations
At the cathode, hydrogen is produced unless the metal is less reactive than hydrogen, in which case the metal is deposited. At the anode, a halogen is produced if a halide is present; otherwise oxygen is produced from hydroxide ions.
Calculations
No routine calculation, though relative volumes of gas can be compared
Half equations describe what happens at each electrode
Graphs and data
Results are recorded as a table of solution against products at each electrode.
Evaluation
Inert electrodes are used so that the electrodes themselves do not react and change the products.
Small gas volumes can make tests unreliable, so allow enough time to collect gas.
Chlorine is toxic and dissolves in water, so yields appear lower than expected.
Common mistakes
Mixing up the electrodes: reduction happens at the cathode (negative), oxidation at the anode (positive).
Forgetting that water provides H⁺ and OH⁻ ions in aqueous solutions.
Writing half equations that do not balance for charge.
Quick self-test
Why are inert electrodes used?+
So the electrodes do not take part in the reaction and change the products formed.
Predict the products of electrolysing aqueous sodium chloride.+
Hydrogen at the cathode, because sodium is more reactive than hydrogen, and chlorine at the anode because a halide is present.
Describe the test for the gas produced at the cathode in that reaction.+
Hold a lit splint at the mouth of the tube; hydrogen burns with a squeaky pop.
Ready to test yourself? Load this practical into Recall and revise it actively.
Required practical 4
Temperature changes
Investigating the variables that affect temperature changes in reacting solutions, such as acid plus metals, acid plus carbonates, neutralisation and displacement.
Aim
Measure the temperature change of a reaction and find how a chosen variable affects it.
Method
Place a polystyrene cup inside a beaker for support and add a fixed volume of the first solution.
Record the starting temperature with a thermometer.
Add a measured volume or mass of the second reactant, stir and put a lid on to reduce heat loss.
Record the highest (or lowest) temperature reached.
Repeat with different volumes or concentrations of the reactant being varied.
Repeat each combination and calculate mean temperature changes.
Variables
Independent
The variable investigated, for example volume or concentration of one reactant
Dependent
Temperature change of the mixture
Control
Total volume of solution
Starting temperature
Same insulation and lid
Same stirring
Equipment
Polystyrene cup with lid, Beaker, Thermometer, Measuring cylinders, Acid and alkali (or metal/carbonate), Stirrer.
Results and observations
Exothermic reactions show a temperature rise; endothermic reactions show a fall. In a neutralisation, the temperature rise peaks when the acid and alkali react in exactly the right proportions and falls again as excess reactant cools the mixture.
Calculations
Temperature change = highest temperature − starting temperature
Mean temperature change from repeats
Graphs and data
Plot temperature change against volume of reactant added; two straight lines that intersect identify the point of maximum change.
Evaluation
A polystyrene cup and lid reduce heat loss to the surroundings, the main source of error.
The thermometer resolution (often 1 °C) limits precision; a digital probe is better.
Repeats identify anomalies caused by inconsistent stirring or delayed reading.
Common mistakes
Recording the final temperature after the mixture has started cooling rather than the maximum.
Changing the total volume between runs, which changes the mass being heated.
Describing a temperature rise as endothermic.
Quick self-test
Why is a polystyrene cup used instead of a glass beaker?+
Polystyrene is a good insulator, so less energy is transferred to the surroundings and the temperature change measured is closer to the true value.
The temperature falls from 21 to 15 degrees C. What type of reaction is this?+
Endothermic, because energy is taken in from the surroundings.
Why should the total volume be kept the same in every run?+
A different total volume would change the mass being heated, so the temperature change would not be comparable.
Ready to test yourself? Load this practical into Recall and revise it actively.
Required practical 5
Rates of reaction
Investigating how changes in concentration affect rates of reaction using two methods: measuring the volume of gas produced, and observing a change in colour or turbidity.
Aim
Show how concentration affects reaction rate, measured both by gas production and by a change in cloudiness.
Method
Gas method: add a measured volume of hydrochloric acid to marble chips or magnesium in a conical flask fitted with a bung and gas syringe.
Record the volume of gas collected at regular intervals, for example every 10 seconds.
Repeat with different acid concentrations, keeping everything else the same.
Turbidity method: mix sodium thiosulfate solution with hydrochloric acid in a flask placed over a printed cross.
Time how long the cross takes to disappear as sulfur forms.
Repeat at different thiosulfate concentrations and calculate a mean time for each.
Variables
Independent
Concentration of the reactant solution
Dependent
Volume of gas produced in a set time, or time for the cross to disappear
Control
Temperature
Volume of solutions
Surface area and mass of solid
Same observer and same cross for turbidity
Equipment
Conical flask and bung, Gas syringe, Measuring cylinders, Stopwatch, Sodium thiosulfate and hydrochloric acid, Printed cross card.
Results and observations
Higher concentration gives more gas in the same time and a shorter time for the cross to disappear. Gas volume curves are steepest at the start and flatten when a reactant runs out.
Calculations
Mean rate = quantity of product formed ÷ time taken (for example cm³/s)
Rate from the turbidity method = 1 ÷ time (s⁻¹)
Rate at a given moment = gradient of the tangent to the curve (Higher Tier)
Graphs and data
Plot volume of gas against time for each concentration; a steeper initial gradient means a faster rate. For turbidity, plot 1/time against concentration.
Evaluation
Gas can escape while the bung is being fitted, lowering early readings.
Judging when the cross disappears is subjective, so the same observer should judge every run.
Temperature must be controlled because it strongly affects rate.
Repeats allow a mean and expose anomalies.
Common mistakes
Plotting time on the y-axis when the question asks for volume against time.
Reading the gradient at the plateau instead of at the start when asked for initial rate.
Forgetting that surface area, not just mass, affects the rate with solids.
Quick self-test
Why does the gas volume curve flatten out?+
One of the reactants has been used up, so no more product is formed.
Explain why higher concentration increases rate.+
There are more reactant particles in the same volume, so collisions are more frequent.
The cross disappears in 25 s. Calculate the rate.+
Rate = 1 / 25 = 0.04 s^-1.
Why should the same person judge every cross?+
Judging when the cross disappears is subjective, so using one observer keeps the end point consistent.
Ready to test yourself? Load this practical into Recall and revise it actively.
Required practical 6
Chromatography
Investigating how paper chromatography separates and distinguishes between coloured substances, including calculating Rf values.
Aim
Separate the components of a coloured mixture and identify them using Rf values.
Method
Draw a pencil start line about 1 cm from the bottom of the chromatography paper.
Spot small samples of each ink or dye on the line and let them dry, repeating to concentrate the spot.
Pour solvent into a beaker so the level is below the pencil line.
Hang the paper so the bottom edge dips into the solvent without submerging the spots.
Cover the beaker and let the solvent rise until it is near the top.
Remove the paper, immediately mark the solvent front in pencil and let it dry.
Variables
Independent
The substance or mixture spotted onto the paper
Dependent
Distance travelled by each spot, and the Rf value calculated from it
Control
Same solvent and depth
Same paper type and size
Same start line position
Same run time or solvent front distance
Equipment
Chromatography paper, Beaker and lid or watch glass, Solvent (water or ethanol), Pencil and ruler, Capillary tubes or dropper.
Results and observations
A pure substance produces one spot; a mixture separates into several. Substances more attracted to the solvent than to the paper travel further.
Calculations
Rf = distance travelled by the spot ÷ distance travelled by the solvent front
Both distances are measured from the pencil start line to the centre of the spot
Rf values are always between 0 and 1 and have no units
Graphs and data
No graph; results are recorded as measured distances and calculated Rf values in a table.
Evaluation
The start line must be pencil because ink would dissolve and run.
If the solvent level is above the start line the spots dissolve into the solvent instead of moving up the paper.
Covering the beaker prevents solvent evaporation, which would change how far the front travels.
Rf values are only comparable when the same solvent and paper are used.
Common mistakes
Measuring to the edge of a spot instead of its centre.
Forgetting to mark the solvent front before it evaporates.
Quoting Rf with units or as a value greater than 1.
Quick self-test
Why must the start line be drawn in pencil?+
Pencil is insoluble, whereas ink would dissolve in the solvent and travel up the paper.
A spot moves 4.5 cm and the solvent front moves 9.0 cm. Calculate Rf.+
Rf = 4.5 / 9.0 = 0.50.
How can you tell a substance is pure from a chromatogram?+
A pure substance produces a single spot in every solvent.
Ready to test yourself? Load this practical into Recall and revise it actively.
Required practical 7 · separate chemistry only
Identifying ions
Using chemical tests to identify the ions in unknown single ionic compounds, covering flame tests through to sulfates.
Aim
Identify the positive and negative ion in an unknown ionic compound using qualitative tests.
Method
Flame test: dip a clean nichrome wire loop in acid, then in the solid, and hold it in a blue Bunsen flame; record the flame colour.
Metal hydroxide test: add sodium hydroxide solution dropwise to a solution of the compound and record any precipitate colour, including whether it dissolves in excess.
Carbonates: add dilute acid and bubble any gas through limewater.
Halides: acidify with dilute nitric acid, then add silver nitrate solution and record the precipitate colour.
Sulfates: acidify with dilute hydrochloric acid, then add barium chloride solution.
Combine the positive and negative ion results to name the compound.
Variables
Independent
The unknown compound being tested
Dependent
The observations: flame colour, precipitate colour and gas test results
Control
Clean apparatus between tests
Same volumes of reagents
Known samples tested alongside for comparison
Equipment
Nichrome wire loop, Bunsen burner, Sodium hydroxide solution, Dilute nitric and hydrochloric acid, Silver nitrate and barium chloride solutions, Limewater, Test tubes.
Results and observations
Flame colours indicate the metal ion. Sodium hydroxide gives coloured precipitates for transition metal ions and a white precipitate for several others. Effervescence with acid that turns limewater cloudy indicates a carbonate. Silver nitrate gives white, cream or yellow precipitates for chloride, bromide and iodide. Barium chloride gives a white precipitate with sulfates.
Calculations
No calculation; the practical is qualitative
Write balanced ionic equations for the precipitation reactions
Graphs and data
Record all observations in a systematic results table linking test, observation and conclusion.
Evaluation
The wire loop must be cleaned between flame tests or a previous ion's colour carries over.
Sodium contamination easily masks other flame colours because its yellow flame is intense.
Acidifying before the silver nitrate or barium chloride test removes carbonate ions that would otherwise give a false positive.
Flame colours can be hard to distinguish by eye; flame emission spectroscopy is more reliable.
Common mistakes
Skipping the acidification step before the halide or sulfate test.
Describing a precipitate without naming the colour.
Reporting a colour change rather than a precipitate for the sodium hydroxide test.
Quick self-test
Why is the sample acidified before adding silver nitrate?+
To remove carbonate ions, which would otherwise form a white precipitate and give a false positive.
A white precipitate forms with barium chloride. What ion is present?+
Sulfate ions.
Why must the flame test wire be cleaned between samples?+
Residue from the previous sample would produce its own flame colour and mask the result.
Ready to test yourself? Load this practical into Recall and revise it actively.
Required practical 8 · separate chemistry only
Water purification
Analysing and purifying water samples from different sources, including pH, dissolved solids and distillation.
Aim
Test water samples for pH and dissolved solids, then purify a sample by distillation and check the result.
Method
Test the pH of each water sample with universal indicator paper or a pH probe.
Weigh a clean, dry evaporating basin.
Add a measured volume of the sample, evaporate the water and reweigh to find the mass of dissolved solids.
Set up a distillation apparatus with a condenser and collect the distillate from a fresh sample of the same water.
Test the pH of the distillate and repeat the evaporation test on it.
Compare the results before and after distillation.
Variables
Independent
The source of the water sample, and whether it has been distilled
Dependent
pH and mass of dissolved solids per unit volume
Control
Same volume of each sample
Same evaporation method and drying time
Same pH measurement method
Equipment
Water samples, Universal indicator paper or pH probe, Evaporating basin and balance, Bunsen burner and tripod, Distillation apparatus with condenser, Thermometer.
Results and observations
Untreated samples leave a solid residue after evaporation and may not be neutral. The distillate leaves little or no residue and has a pH close to 7, showing dissolved solids have been removed.
Calculations
Mass of dissolved solids = mass of basin plus residue − mass of empty basin
Concentration of dissolved solids in g/dm³ = mass of solids ÷ volume of sample in dm³
Graphs and data
Compare samples with a bar chart of dissolved solids per dm³ alongside a table of pH values.
Evaluation
The basin must be completely dry before both weighings or the residue mass is overestimated.
Distillation removes dissolved solids but does not remove substances with a similar boiling point to water.
Universal indicator gives pH to the nearest whole number; a pH probe is more precise.
Distillation uses a lot of energy, which is why it is not the usual method for large-scale drinking water.
Common mistakes
Assuming a clear sample is pure because dissolved solids are invisible.
Weighing the basin while it is still warm or damp.
Confusing distillation with filtration, which does not remove dissolved substances.
Quick self-test
Why does a clear water sample still leave a residue after evaporation?+
Dissolved solids are invisible in solution but remain behind when the water evaporates.
A 50 cm^3 sample leaves 0.20 g of residue. Calculate the concentration in g/dm^3.+
0.20 / 0.050 = 4.0 g/dm^3.
Give one limitation of distillation for producing drinking water.+
It requires a large amount of energy, which makes it expensive on a large scale.
Ready to test yourself? Load this practical into Recall and revise it actively.
Practical names follow the required practical activities listed in the official AQA GCSE Chemistry 8462 specification. Always check the current specification for the authoritative list.
How to answer GCSE required practical questions
Practical questions rarely ask you to recite a method. They usually give data, a diagram or an unfamiliar setup and test whether you can think like an experimenter. The exact skills depend on the practical and on the specification you are sitting, so use this as a checklist rather than a prediction.
Identifying variablesPick out what was changed, what was measured and what had to stay the same.
Choosing control variablesName the specific variables that would otherwise affect the dependent variable.
Describing a valid methodWrite steps in order, with quantities and a clear measurement, so someone else could repeat it.
Selecting measurementsChoose apparatus with a sensible resolution and range for what you are measuring.
Recording resultsUse a table with headings, units and a consistent number of decimal places.
Calculating valuesMeans, percentage change, rate, gradient, density and other calculations set by the practical.
Plotting and reading graphsSensible scales, labelled axes with units, plotted points and a line of best fit.
Identifying anomaliesSpot results that do not fit the pattern, exclude them from means and suggest a cause.
Evaluating reliabilityExplain how repeats, means and controlled variables strengthen a conclusion.
Sources of errorSeparate random error from systematic error and say which affects your data.
Suggesting improvementsChange one specific thing and say what it would improve, rather than 'be more careful'.
Drawing conclusionsState the relationship the data supports and refer back to the values you collected.
Applying to new contextsUse the same reasoning on an unfamiliar experiment you have never carried out.
GCSE Required Practical Exam Questions
These are original practice questions written for this page, not past-paper questions. They cover the skills practical questions tend to test: variables, calculations, anomalies, graphs and improvements.
ChemistryA student's rough titre is 25.90 cm³ and their next three titres are 24.85, 24.90 and 25.40 cm³. Which values should they average, and why?+
Average 24.85 and 24.90 cm³ because these are concordant, within 0.10 cm³ of each other. The rough titre and the 25.40 cm³ result are excluded as they are less accurate.
ChemistryIn a rates investigation, 48 cm³ of gas is produced in 40 seconds. Calculate the mean rate of reaction.+
48 ÷ 40 = 1.2 cm³/s.
ChemistryA spot travels 3.6 cm while the solvent front travels 8.0 cm. Calculate the Rf value.+
Rf = 3.6 ÷ 8.0 = 0.45. Rf has no units and is always between 0 and 1.
ChemistrySuggest why the temperature rise measured in a polystyrene cup is smaller than the true value.+
Some energy is transferred to the surroundings, to the cup and to the thermometer, so less energy warms the solution than the reaction actually releases.
Practical revision suits active recall because most of it is small, precise detail: which reagent, which variable, which unit. Reading a method again feels productive but rarely shows you what you cannot remember. Testing yourself does.
Press a button on any practical above to load it here, or paste your own notes.
Loaded: AQA GCSE Chemistry practicals. Edit the notes if you want to narrow the focus, then generate.
17 words100 to 500 words works bestTry a sample:
Free to try, no card needed
Paste your notes, choose a mode and press Generate, or load a sample above.
AQA GCSE Required Practicals FAQs
What are GCSE required practicals?+
They are the practical activities AQA states every student must carry out during the course. There is no separate practical exam at GCSE: instead, the written papers include questions that draw on the knowledge, skills and understanding you gained from doing them.
How many AQA GCSE Biology required practicals are there?+
The AQA GCSE Biology (8461) specification lists ten required practical activities. Seven of them are shared with GCSE Combined Science, while practicals 2, 8 and 10 are on the separate Biology specification only.
How many AQA GCSE Chemistry required practicals are there?+
The AQA GCSE Chemistry (8462) specification lists eight required practical activities. Titration, identifying ions and water purification appear on the separate Chemistry specification rather than Combined Science.
How many AQA GCSE Physics required practicals are there?+
The AQA GCSE Physics (8463) specification lists ten required practical activities, of which thermal insulation and the light practical are on the separate Physics specification only.
Do required practicals appear in GCSE exams?+
AQA states that written papers include questions requiring knowledge gained from carrying out the specified practicals. The exact questions vary from paper to paper, so no one can tell you which practical will come up or how many marks it will carry.
How should I revise required practicals?+
For each practical, be able to state the aim, outline the method in order, identify the independent, dependent and control variables, describe what the results show, do any calculation involved and suggest one improvement. Then self-test rather than re-reading, which is where flashcards and exam-style questions help.
Are required practicals the same for every exam board?+
No. Every board has to cover the same underlying apparatus and techniques requirements, but the named practical activities and their wording differ between AQA, Edexcel and OCR. Use the list that matches the specification you are entered for.
Where can I check the official list?+
AQA publishes the required practical activities in the practical assessment section of each specification, and provides a required practical handbook with suggested methods. Always treat AQA as the authoritative source for the current list.