AQA GCSE Chemistry required practicals: All 8 explained

There are 8 required practicals in the AQA GCSE Chemistry specification (8462), and questions on them can appear on either Paper 1 or Paper 2. You will not actually carry out an experiment in the exam hall, but you will be expected to describe the method, identify variables, interpret results and evaluate the technique as if you had done it yourself.

Practical-skills questions are worth at least 15% of the total marks across both papers. AQA tests whether you understand the chemistry behind the experiment, not whether you have memorised a recipe. This guide lists all 8 required practicals with what you need to know for each one, then sets out the question types AQA returns to year after year.

The table below lists every required practical on the AQA GCSE Chemistry specification, grouped by the paper they fall under. Paper 1 covers Topics 1 to 5 (atomic structure, bonding, quantitative chemistry, chemical changes, energy changes). Paper 2 covers Topics 6 to 10 (rates of reaction, organic chemistry, chemical analysis, the atmosphere, using resources).

For each required practical, aim to be able to answer four things from memory. The method (step by step), the variables (independent, dependent and at least two control variables), the expected results, and the conclusion you can draw. Below is a breakdown of each one.

You make a pure, dry sample of a soluble salt by reacting an insoluble base or carbonate with a warm acid. For example, copper sulfate is made by adding excess copper oxide to warm dilute sulfuric acid, stirring until no more reacts (the acid becomes neutral). You then filter to remove the excess solid, evaporate some water using a water bath or evaporating basin, and leave to crystallise. Finally, you pat the crystals dry between filter paper. Common exam questions: Why add excess base, why filter, why crystallise rather than boil dry (gentle evaporation preserves crystal shape and prevents decomposition), and how to identify which acid produces which salt.

Both Foundation and Higher tier students do the titration practical itself. What is restricted to Higher tier is the calculation work-up: Concentrations in mol/dm³ and g/dm³, and moles from volumes. Foundation papers can still examine the apparatus, indicator and end point.

The method: A pipette is used to measure 25 cm³ of alkali (e.g. sodium hydroxide) into a conical flask with an indicator (methyl orange or phenolphthalein). Acid is added from a burette until the indicator changes colour, marking the end point. The first titration is a rough run; you then repeat to get concordant results (within 0.10 cm³ of each other) and calculate a mean titre. On Higher tier, the unknown concentration is found using moles = concentration × volume. Common exam questions: Why use a pipette for the alkali but a burette for the acid (more accurate volume control), why repeat for concordant results, and (Higher only) calculations of unknown concentration from the mean titre.

You set up an electrolysis cell using two inert (carbon or platinum) electrodes dipped in a salt solution, connected to a low-voltage power supply. Gases produced at each electrode are collected in inverted test tubes and tested. At the cathode, hydrogen is produced unless the metal is less reactive than hydrogen (e.g. copper from copper sulfate). At the anode, oxygen is produced from sulfates and nitrates; chlorine is produced from concentrated halide solutions. The standard gas tests are a lit splint that pops for hydrogen, a glowing splint that relights for oxygen, and damp litmus paper that bleaches white for chlorine. Common exam questions: Predict products for a given solution, identify the gas from a description, and explain why hydrogen rather than sodium is produced at the cathode from sodium chloride solution.

You investigate the temperature change of a reaction (typically a neutralisation or displacement) in an insulated polystyrene cup. A known volume of acid is measured into the cup and its temperature recorded. Alkali is added and the maximum temperature recorded. The independent variable might be the concentration or volume of one reactant, and the dependent variable is the temperature change Δθ. The polystyrene cup with a lid minimises heat loss. Common exam questions: Why use a polystyrene cup (insulation), why use a lid (reduces heat loss to evaporation), calculate Δθ from initial and final temperatures, and identify whether the reaction is exothermic or endothermic.

Two methods, depending on the reaction. For sodium thiosulfate and hydrochloric acid (a colour change reaction), you place the flask on a paper with a cross drawn on it and time how long the cross takes to disappear behind the cloudy sulfur precipitate. Faster reaction = shorter time. For magnesium and dilute hydrochloric acid (a gas evolution reaction), you collect the hydrogen produced in a gas syringe or inverted measuring cylinder and record volume against time. Independent variables can be concentration, temperature, surface area or use of a catalyst. Common exam questions: Sketch a curve of volume against time, explain how rate changes with concentration using collision theory, and identify when the reaction has finished from the graph.

You investigate which dyes are in a mixture (e.g. food colourings or inks). Spots of mixture and reference dyes are placed on a pencil line near the bottom of chromatography paper. The paper is suspended in a beaker with the solvent below the pencil line. As the solvent moves up the paper, it carries the dyes with it at different rates. Each dye's Rf value is calculated as distance moved by the dye divided by distance moved by the solvent. Same Rf in the same solvent means the same substance. Common exam questions: Why use pencil (ink would dissolve), why place spots above the solvent (otherwise they would dissolve into it), calculate Rf, and identify which dyes are present by matching Rf values.

You identify which positive and negative ions are present in an unknown solution using a set of standard tests. Flame tests identify some cations: Lithium gives crimson, sodium yellow, potassium lilac, calcium orange-red, copper green. Sodium hydroxide solution gives coloured precipitates – copper(II) blue, iron(II) green, iron(III) brown, aluminium and calcium white (aluminium dissolves in excess). Halide ions give silver halide precipitates with silver nitrate and dilute nitric acid – chloride white, bromide cream, iodide yellow. Sulfate ions give a white precipitate with barium chloride and dilute hydrochloric acid. Carbonate ions fizz with dilute acid and the gas turns limewater milky. Common exam questions: Given a set of test results, identify the compound, and describe a step-by-step plan to identify an unknown.

You compare different water samples (e.g. tap water, distilled water, seawater, mineral water) by measuring pH with universal indicator and finding the mass of dissolved solids. To find dissolved solids, an evaporating basin is weighed empty, then with 25 cm³ of the water sample, and finally after evaporating the water in a water bath. The mass of solids = (final mass with solids) − (mass of empty basin). Pure water leaves no residue. Common exam questions: Explain why tap water leaves a residue (contains dissolved minerals), describe how to make water potable (filtration and sterilisation), and explain the difference between potable water and pure water (potable = safe to drink but contains dissolved substances; pure = only H₂O molecules).

AQA tends to recycle a small set of question styles for required practicals. Knowing the patterns means you can prepare your answers in advance.

Describe the method questions ask you to list the steps in order. Marks come from specifying equipment, quantities and exactly how measurements are taken. Vague answers like 'add some acid' tend to score little – say what concentration, what volume, and into what apparatus.

Identify the variables questions want the independent variable (what you change), the dependent variable (what you measure) and at least two control variables (what you keep the same). State how each one is controlled.

Explain the results questions need you to use chemistry – not just describe what happened. For example, the rate of reaction increases with concentration because there are more particles per unit volume, leading to more frequent collisions.

Evaluate the method questions ask you to suggest improvements. The standard categories are more precise measuring equipment, more repeats and a mean, better-controlled variables, and reducing systematic errors like heat loss to the surroundings.

Six-mark extended response questions sometimes focus entirely on one required practical. These need a logical structure – describe the method, explain the chemistry, and link the answer to any data shown.

Reading through the method is rarely enough. AQA tests whether you can think like a chemist, not whether you can recite a textbook.

For each practical, write the method out from memory and then check it against your notes or a Cognito video. Pay attention to the steps you forgot – those are the marks you would have dropped in the exam.

Learn the colours cold. Flame test colours, hydroxide precipitate colours, halide silver salt colours – repeated drilling tends to be the most reliable way to remember them. Make flashcards or a single A4 colour sheet and review it once a day in the run-up to the exam.

Drill the chemistry behind each practical. Why does aluminium hydroxide dissolve in excess sodium hydroxide but calcium hydroxide does not? Why does sodium chloride solution produce hydrogen at the cathode rather than sodium? These are the kinds of follow-up questions AQA loves.