Required practicals · Chemistry
Chemistry required practicals
AQA sets 8 required practicals for Chemistry. This page covers the AQA specification only — not Edexcel, OCR or any other exam board — so check with your teacher which board you’re actually sitting before relying on it. (MarkMatch as a whole is AQA-only for now, too.)
Every practical below has a real method to know, but not everything is equally exam-critical — the greenbox under each one is what actually costs marks if you don’t know it; the grey box is useful context you don’t need to stress about memorising.
Preparing a soluble salt
Prepare a pure, dry sample of a soluble salt from an insoluble base/metal and a dilute acid.
Equipment
Dilute acid (e.g. sulfuric acid), Insoluble base (e.g. copper oxide) or a reactive metal, Bunsen burner, filter funnel + paper, evaporating basin, Glass rod, water bath
Method
- Warm the acid gently (do not boil).
- Add the insoluble base/metal a little at a time, stirring, until it is in excess (no more dissolves) — this ensures all the acid has reacted.
- Filter the mixture to remove the unreacted excess solid, keeping the filtrate (salt solution).
- Gently heat/evaporate the filtrate (a water bath, not direct flame, avoids spitting/decomposition) until crystals start to form at the edge — this is saturation point.
- Leave to cool and crystallise; filter and pat dry, or leave crystals to dry naturally.
You definitely need to know
- The solid (base/metal) is added in EXCESS specifically so all the acid reacts — leftover solid is filtered off, leftover acid would contaminate the product.
- Order: react → filter (removes excess solid) → evaporate/crystallise → dry.
- Crystallisation stops once you see crystals forming at the surface/edge of the solution — over-heating past this can decompose the salt.
Safety
Heat gently — boiling dry too fast can cause spitting or decompose the salt.
Common exam question types
- Explain why the base/metal is added in excess.
- Describe how to obtain pure, dry crystals from the salt solution.
- Explain why the evaporating basin is heated gently rather than boiled dry directly.
- Name the salt produced from a given acid + base/metal.
Titration
Use titration to determine the reacting volumes of an acid and an alkali, and find the concentration of one from the other.
Equipment
Burette, pipette + pipette filler, Conical flask, Indicator (e.g. phenolphthalein or methyl orange), White tile
Method
- Use a pipette to measure a fixed volume of alkali into a conical flask; add a few drops of indicator.
- Fill the burette with acid, recording the starting volume.
- Add acid to the alkali, swirling constantly, until the indicator just changes colour (the end point).
- Record the volume of acid used (the titre).
- Repeat until you get concordant results — titres within 0.10 cm³ of each other — and calculate a mean using only the concordant ones.
You definitely need to know
- A white tile is placed under the flask to see the colour change clearly.
- "Concordant results" = titres agreeing within 0.10 cm³ — you must be able to identify which repeats are/aren’t concordant from a table.
- The mean titre is calculated from CONCORDANT results only, not all results, and not including a rough/trial run.
- A rough (fast) titration is normally done first to find an approximate end point, then repeated slowly and precisely near that volume.
Common exam question types
- Identify concordant results from a table and calculate the mean titre.
- Explain why a rough titration is carried out first.
- Explain why the flask is swirled during the titration.
- Calculate concentration (mol/dm³) using the titre and a given equation (moles = concentration × volume).
Electrolysis
Investigate what happens when aqueous solutions are electrolysed using inert electrodes.
Equipment
Electrolysis cell/beaker, Inert electrodes (carbon/graphite or platinum), Power supply, Aqueous solution (e.g. copper chloride, sodium chloride), Test tubes to collect gas
Method
- Set up the cell with inert electrodes dipped in the solution, connected to a DC power supply.
- Switch on and observe each electrode.
- Test any gas produced (e.g. a lit splint for hydrogen — "squeaky pop"; a glowing splint for oxygen — relights).
- For a solution like copper chloride, note the colour change/deposit at each electrode.
You definitely need to know
- At the CATHODE (negative): if the metal is more reactive than hydrogen (e.g. Na, K, Ca), hydrogen gas is produced instead of the metal; less reactive metals (e.g. copper) are deposited.
- At the ANODE (positive): if a halide ion is present, the halogen is produced; otherwise oxygen is produced (from OH⁻ ions in water).
- Gas tests: hydrogen — lit splint gives a squeaky pop; oxygen — relights a glowing splint; chlorine — bleaches damp litmus paper.
- Electrodes must be INERT (carbon/platinum) so they don’t react themselves and interfere with the products.
Safety
Chlorine gas is toxic — only produced in small amounts and tested carefully, never inhaled directly.
Common exam question types
- Predict/explain the product at each electrode for a named solution.
- Describe a test to confirm a named gas.
- Explain why inert electrodes are used.
- Write a half-equation for a named electrode reaction (Higher tier).
Temperature changes in reactions
Measure the temperature change of reacting solutions to compare the energy released/absorbed by different reactions.
Equipment
Polystyrene cup (insulated) inside a beaker for support, Thermometer, Lid with a hole for the thermometer, Two reacting solutions (e.g. acid + metal, or acid + alkali)
Method
- Measure a set volume of one solution into the insulated cup; record its starting temperature.
- Add the second reactant (or a fixed mass of solid), stirring, and record the highest (or lowest) temperature reached.
- Calculate the temperature change: final − initial.
- Repeat for different variables (e.g. different concentrations, or different metals) keeping everything else the same.
| Independent | The reactant variable being compared (e.g. concentration, type of metal/acid) — varies by investigation |
| Dependent | Temperature change (final − initial) |
| Control | Volume of solution; Starting temperature; Insulation/lid used |
You definitely need to know
- A polystyrene cup is used because it INSULATES — it minimises energy transfer to/from the surroundings, so the measured change reflects the reaction, not heat loss.
- An EXOTHERMIC reaction gives a temperature RISE (energy released to surroundings); an ENDOTHERMIC reaction gives a temperature FALL (energy absorbed from surroundings).
- A lid with a small hole for the thermometer further reduces heat loss to the air — expect this as an "improvement" answer.
Common exam question types
- Calculate temperature change from given start/end readings.
- Explain why a polystyrene cup (and lid) is used.
- Identify whether a reaction is exothermic or endothermic from the temperature change and explain what that means for energy transfer.
- Suggest an improvement to reduce heat loss.
Rate of reaction
Investigate how changing a variable (concentration, temperature, surface area, or use of a catalyst) affects the rate of a reaction.
Equipment
Reacting solutions (e.g. sodium thiosulfate + hydrochloric acid, OR magnesium + acid with a gas syringe), Conical flask, paper with a black cross (for the "disappearing cross" method) OR gas syringe/measuring cylinder, Stopwatch, water bath
Method
- "Disappearing cross" method (sodium thiosulfate + HCl): place the flask over a paper cross, start the stopwatch when acid is added, and time how long until the cross is no longer visible through the cloudy precipitate.
- "Gas collection" method (e.g. magnesium + acid): collect the gas produced in a gas syringe, recording the volume at regular time intervals (e.g. every 10 s) until the reaction finishes.
- Repeat, changing ONE variable (e.g. concentration of acid) each time, keeping everything else constant.
- For the gas method, plot volume of gas against time; the initial gradient shows the rate.
| Independent | The factor being tested (e.g. concentration, temperature, surface area, catalyst presence) |
| Dependent | Time taken (disappearing cross) or volume of gas over time (gas collection) |
| Control | Volume/mass of reactants; Temperature (unless it’s the variable being tested); The person judging when the cross disappears, ideally the same person throughout |
You definitely need to know
- Rate = 1 ÷ time is the standard way to express results from the disappearing-cross method.
- For the gas method, the STEEPEST part of a volume-time graph (usually the start) shows the FASTEST rate — reaction slows as reactants are used up, shown by the graph levelling off.
- A limitation of the disappearing-cross method is it’s subjective (different people judge "disappeared" differently) — a very common evaluate question.
- Higher concentration/temperature/surface area, or adding a catalyst, all increase rate — be ready to explain WHY using collision theory (more frequent/more energetic collisions).
Common exam question types
- Calculate rate as 1/time from given data.
- Explain the shape of a volume-time graph, including where it levels off and why.
- Evaluate the disappearing-cross method (reliability, subjectivity).
- Explain, using collision theory, why changing the tested variable changes the rate.
Chromatography
Use paper chromatography to separate and identify substances in a mixture, and calculate Rf values.
Equipment
Chromatography paper, Pencil (not pen — ink would run too), Capillary tube for spotting, Solvent, beaker/tank with a lid
Method
- Draw a pencil line near the bottom of the paper (the "baseline") and mark on it where each sample will be spotted.
- Spot small amounts of each sample onto the baseline using a capillary tube, letting each spot dry before adding more (keeps spots small and concentrated).
- Lower the paper into a shallow layer of solvent — the solvent level must be BELOW the baseline, or the samples would dissolve straight into the solvent instead of running up the paper.
- Put a lid on and leave until the solvent has risen most of the way up the paper (the "solvent front").
- Remove, mark the solvent front immediately, let dry, and measure distances to calculate Rf values.
You definitely need to know
- Rf = distance travelled by the spot ÷ distance travelled by the solvent front — both measured from the baseline.
- The baseline must be drawn in PENCIL and kept ABOVE the solvent level — both are classic "explain why" questions.
- A pure substance gives ONE spot; a mixture gives multiple spots — this is how chromatography distinguishes pure substances from mixtures.
- Matching Rf values (under the same conditions) to a known reference substance identifies an unknown component.
Common exam question types
- Calculate an Rf value from given distances.
- Explain why the baseline is drawn in pencil.
- Explain why the solvent level must start below the baseline.
- Use Rf values (or spot pattern) to identify substances in a mixture / distinguish a pure substance from a mixture.
Identifying ions (flame tests)
Use chemical tests to identify metal ions, carbonate, halide and sulfate ions in unknown single ionic compounds, and simple tests to identify gases (4.8.3.1–4.8.3.5) — a single required practical spanning all five test types, confirmed against AQA's own specification, not just flame tests.
Equipment
Nichrome wire loop, Dilute hydrochloric acid (to clean the loop, and for the carbonate test), Bunsen burner, Sample compounds (solid and in solution), Sodium hydroxide solution, Dilute nitric acid + silver nitrate solution, Dilute hydrochloric acid + barium chloride solution, Test tubes for gas tests, limewater
Method
- FLAME TEST: dip the nichrome wire loop in hydrochloric acid, then hold it in the blue Bunsen flame to clean it (repeat until no colour shows). Dip the clean loop into the sample, then hold it back in the flame. Observe and record the flame colour.
- METAL HYDROXIDE TEST: add a few drops of sodium hydroxide solution to a solution of the sample. Observe the colour of any precipitate formed, and (for a white precipitate) add EXCESS sodium hydroxide to see whether it redissolves.
- CARBONATE TEST: add dilute acid to the sample. If it fizzes, bubble the gas produced through limewater and observe.
- HALIDE TEST: add a few drops of dilute nitric acid, then a few drops of silver nitrate solution, to a solution of the sample. Observe the colour of any precipitate formed.
- SULFATE TEST: add a few drops of dilute hydrochloric acid, then a few drops of barium chloride solution, to a solution of the sample. Observe whether a precipitate forms.
- GAS TESTS: hydrogen — lit splint gives a squeaky pop; oxygen — relights a glowing splint; carbon dioxide — turns limewater cloudy; chlorine — bleaches damp litmus paper (turns it white).
You definitely need to know
- Flame colours: lithium — crimson red; sodium — yellow/orange; potassium — lilac; calcium — orange-red; copper — blue-green/green.
- Metal hydroxide precipitate colours: calcium — white; copper(II) — blue; iron(II) — green; iron(III) — brown/orange-brown; aluminium — white. Aluminium hydroxide (unlike calcium hydroxide) DISSOLVES in EXCESS sodium hydroxide to give a colourless solution — the key way to distinguish Al³⁺ from Ca²⁺, since both start as a white precipitate.
- Carbonate test: fizzing with acid PLUS the gas turning limewater cloudy confirms carbon dioxide — fizzing alone is not enough, since other reactions can also produce gas.
- Halide precipitate colours (with silver nitrate, after acidifying with nitric acid first): chloride — white; bromide — cream; iodide — yellow.
- Sulfate test: a white precipitate (barium sulfate) with barium chloride, after acidifying with hydrochloric acid first, confirms a sulfate ion. Acidifying first is essential — it dissolves any carbonate that might otherwise also give a false-positive white precipitate.
- The loop MUST be cleaned with acid and flame between every flame-test sample, or contamination gives a false colour.
- A mixture of metal ions can mask individual flame colours — this limits the flame test’s reliability for mixtures (instrumental methods like flame photometry are used instead, Higher tier).
Good to know — less essential
- This is Chemistry-only (not Combined Science Trilogy).
Common exam question types
- Identify a metal ion from a stated flame colour, hydroxide precipitate colour/behaviour in excess, or vice versa.
- Given a set of test results (flame colour, precipitate colours, gas test), identify the compound present.
- Describe the correct method for carrying out a named test (flame test, hydroxide test, halide test, sulfate test).
- Explain why nitric acid is added before the silver nitrate halide test, or why hydrochloric acid is added before the barium chloride sulfate test (both remove/avoid a possible carbonate false-positive).
- Explain why the wire loop is cleaned before each flame test.
- Explain a limitation of flame tests for identifying ions in a mixture.
- Describe how to distinguish aluminium and calcium ions using sodium hydroxide (aluminium hydroxide redissolves in excess, calcium hydroxide does not).
Water purification
Produce potable water from an impure source (e.g. seawater) using distillation and filtration.
Equipment
Impure/salty water sample, Distillation apparatus (flask, condenser, thermometer), Filter funnel + paper, Heat source, universal indicator/pH probe
Method
- Filter the impure water first to remove any insoluble solids.
- Set up simple distillation: heat the filtered water in a flask; water evaporates and passes into a condenser, where it cools back to liquid and is collected.
- Dissolved salts stay behind in the original flask, since they don’t evaporate at water’s boiling point.
- Test the distilled water’s purity/quality: check pH is close to 7, and that it evaporates to leave no solid residue.
You definitely need to know
- Distillation works because water has a much lower boiling point than the dissolved salts — the water evaporates and is collected, the salts are left behind.
- "Potable" means safe to drink — NOT the same as "pure" water (chemically pure H₂O). Potable water still contains some dissolved minerals/ions, just at safe levels.
- Filtration removes insoluble solids; distillation is needed to remove dissolved solids (salts).
- A condenser works by cooling the water vapour so it turns back into a liquid — usually shown with water flowing counter-current around it.
Common exam question types
- Describe the process of producing potable water from seawater/an impure source.
- Explain why filtration alone isn’t enough to remove dissolved salts.
- Explain the difference between "pure" and "potable" water.
- Explain how the condenser cools the water vapour.