Master all skills needed for your practical exam — step-by-step, with interactive quizzes. Designed to be simple and clear!
This section explains what the practical exam expects and how marks are awarded.
1 hour 30 minutes
40 marks — worth 20% of your final Chemistry grade
Carry out experiments, record results, draw conclusions, and answer questions
Following instructions, measuring accurately, recording data, and identifying chemicals
| Type | What You Do | Marks |
|---|---|---|
| Acid-Base Titration | Measure volume of acid/alkali to find concentration | ~15 marks |
| Qualitative Analysis | Identify unknown ions using chemical tests | ~12 marks |
| Other Experiments | Rate of reaction, paper chromatography, etc. | ~8 marks |
| Questions on Results | Calculations, conclusions, sources of error | ~5 marks |
Before touching any equipment, read all steps carefully so you know what to expect.
Use the correct equipment for each measurement. Read the scale at eye level!
Write down every reading immediately. Do NOT do the work from memory later.
Write each step clearly. You can still get marks even if your final answer is wrong.
State your answer clearly using the evidence from your results.
Know your equipment. Using the right tool at the right time earns you marks.
| Equipment | Used For | Accuracy |
|---|---|---|
| Burette | Titration — delivering variable volumes | ±0.05 cm³ |
| Pipette (25 cm³) | Delivering a fixed, accurate volume | ±0.06 cm³ |
| Measuring Cylinder | Rough measurements of volume | ±0.5 cm³ |
| Conical Flask | Holding solution during titration | Not for measuring |
Meniscus tip: Water (and most solutions) curve downward — this dip is called the meniscus. ALWAYS read from the BOTTOM of the curve. Reading from the top gives an incorrectly high value!
Always zero (tare) with an empty container first. Record to 2 decimal places (e.g. 2.34 g).
Place weighing boat first, tare to zero, then add your substance. Never weigh directly on the pan.
For most reactions, a –10°C to 110°C thermometer is suitable.
Gently stir the solution to ensure even temperature distribution before you read.
Don't guess smaller increments than the thermometer scale allows.
Always record the units alongside every measurement: cm³, g, °C, mol/dm³, etc.
Your results table must have these features to score full marks:
Include the quantity AND unit, e.g. "Volume of acid / cm³"
If your burette reads to 2 d.p., record ALL values to 2 d.p. (e.g. 24.50 not 24.5)
One rough titration + at least two accurate (concordant) titrations
Draw a table with ruled lines. Separate rough titration from accurate ones.
The most common experiment in Paper 3. Learn each step and the key words.
Clamp the burette vertically. Rinse with the acid you will use. Fill above 0 cm³, then drain to 0.00 cm³. Remove any air bubbles from the tip.
Use a pipette to transfer exactly 25.0 cm³ of alkali (e.g. NaOH) into a clean conical flask. Add 2–3 drops of indicator (e.g. methyl orange or thymolphthalein).
Open the burette tap and add acid quickly while swirling. When colour changes permanently, record the volume. This gives you an approximate endpoint.
Repeat from step 2. Add acid drop by drop near the rough endpoint. Stop when ONE drop causes a PERMANENT colour change. Record to 0.05 cm³.
Use concordant titres (within 0.10 cm³ of each other) to calculate the mean. Do NOT include the rough titration.
| Indicator | In Acid | At Endpoint | In Alkali |
|---|---|---|---|
| Methyl Orange | 🔴 Red | 🟠 Orange | 🟡 Yellow |
| Thymolphthalein | ⚪ Colourless | 🔵 Faint Blue | 🔵 Blue |
| Litmus | 🔴 Red | 🟣 Purple | 🔵 Blue |
| Universal Indicator | 🔴 Red/Orange | 🟢 Green | 🔵 Blue/Violet |
Thymolphthalein is colourless in acid and neutral solutions, and turns blue in alkaline conditions (pH > 9.4). It gives a very sharp colour change and is suitable for strong acid–strong alkali titrations.
Exam tip: Methyl orange and thymolphthalein are preferred for titrations — they give a sharp, clear colour change. Universal indicator and litmus are NOT suitable because their change is too gradual.
You MUST be able to calculate concentration from titration results.
25.0 cm³ of NaOH was titrated with HCl (concentration 0.100 mol/dm³).
Mean titre of HCl = 22.40 cm³
Step 1: Convert cm³ to dm³: 22.40 ÷ 1000 = 0.02240 dm³
Step 2: Moles of HCl = 0.100 × 0.02240 = 0.00224 mol
Step 3: From equation (1:1 ratio), moles of NaOH = 0.00224 mol
Step 4: Concentration of NaOH = 0.00224 ÷ 0.0250 = 0.0896 mol/dm³
Identify unknown ions using chemical tests. Learn the tests, reagents, and observations.
| Ion | Test | Observation |
|---|---|---|
| Cu²⁺ (copper) | Add NaOH solution | Blue precipitate forms |
| Fe²⁺ (iron II) | Add NaOH solution | Green precipitate forms |
| Fe³⁺ (iron III) | Add NaOH solution | Orange-brown precipitate |
| Zn²⁺ (zinc) | Add NaOH (excess) | White ppt; dissolves in excess NaOH |
| Al³⁺ (aluminium) | Add NaOH (excess) | White ppt; dissolves in excess NaOH |
| Ca²⁺ (calcium) | Add NaOH solution | White precipitate forms |
| NH₄⁺ (ammonium) | Add NaOH, warm | Pungent gas; damp red litmus turns blue |
To tell Al³⁺ and Zn²⁺ apart: add excess ammonia solution — zinc precipitate dissolves in excess ammonia, but aluminium precipitate does NOT.
| Ion | Test | Observation |
|---|---|---|
| Cl⁻ (chloride) | Add dilute HNO₃ then AgNO₃ solution | White precipitate of AgCl |
| Br⁻ (bromide) | Add dilute HNO₃ then AgNO₃ solution | Cream precipitate of AgBr |
| I⁻ (iodide) | Add dilute HNO₃ then AgNO₃ solution | Yellow precipitate of AgI |
| SO₄²⁻ (sulfate) | Add dilute HCl then BaCl₂ solution | White precipitate of BaSO₄ |
| CO₃²⁻ (carbonate) | Add dilute HCl | Colourless gas; turns limewater milky |
| NO₃⁻ (nitrate) | Add NaOH + aluminium powder, warm | Gas produced — damp red litmus turns blue |
Halide memory trick — White, Cream, Yellow: Chloride = White, Bromide = Cream, Iodide = Yellow. They get progressively darker going down Group 7!
| Gas | Test | Positive Result |
|---|---|---|
| Hydrogen (H₂) | Burning splint at mouth of tube | "Squeaky pop" sound |
| Oxygen (O₂) | Glowing splint inside tube | Splint re-ignites (relights) |
| Carbon dioxide (CO₂) | Bubble through limewater | Limewater turns milky/cloudy |
| Chlorine (Cl₂) | Damp litmus paper | Paper bleaches/turns white |
| Ammonia (NH₃) | Damp red litmus paper | Paper turns blue |
| Sulfur dioxide (SO₂) | Damp potassium manganate(VII) paper | Paper turns from purple to colourless |
Sodium (Na⁺) — very bright, persistent yellow
Potassium (K⁺) — faint lilac colour
Calcium (Ca²⁺) — brick red/orange-red
Barium (Ba²⁺) — apple green colour
Copper (Cu²⁺) — blue-green, sometimes called "turquoise"
Lithium (Li⁺) — deep crimson red
How to do a flame test: Clean a platinum/nichrome wire in dilute HCl. Dip it in the sample. Hold it in a blue Bunsen flame and observe the colour produced.
Safety is tested in Paper 3. Know these rules — they could also save your life in the lab!
Safety goggles protect your eyes from splashes of acid, alkali, and other chemicals. ALWAYS wear them whenever chemicals are present.
When heating a test tube, always point the open end away from yourself and others. Contents can erupt suddenly.
Never inhale a gas directly. Hold the container away and gently wave your hand to direct a small amount towards your nose.
Even if you wore gloves, wash hands thoroughly before leaving the laboratory.
Loose clothing near a Bunsen burner can catch fire. Tie hair back whenever using a naked flame.
Chemicals can contaminate food or drinks, causing serious illness.
Can cause serious harm or death even in small amounts. Examples: chlorine gas, lead compounds.
Can catch fire easily. Examples: ethanol, methane. Keep away from flames.
Attacks and destroys living tissue (skin, eyes). Examples: concentrated H₂SO₄, NaOH.
Causes redness or inflammation on contact. Less severe than corrosive.
Provides oxygen to cause/worsen fires. Examples: potassium manganate(VII), H₂O₂.
Do not dispose of down drains. Dispose of as instructed by your teacher.
| Common Error | Effect on Result | How to Avoid |
|---|---|---|
| Parallax error in burette | Incorrect volume recorded | Read at eye level, bottom of meniscus |
| Air bubble in burette tip | Volume reading too low | Open tap fully to flush bubble out |
| Not rinsing pipette with solution | Solution becomes diluted | Rinse twice with the solution used |
| Adding too much indicator | Endpoint colour unclear | Use only 2–3 drops of indicator |
| Stopping too late at endpoint | Over-titration; titre too large | Add drop by drop near endpoint |
| Including rough titre in mean | Inaccurate mean titre | Only average concordant titres |
Exothermic and endothermic reactions are commonly tested in Paper 3. Learn how to measure, record, and interpret temperature changes.
Energy is released to the surroundings. The temperature of the mixture increases. Examples: combustion, neutralisation, displacement reactions, hand warmers.
Energy is absorbed from the surroundings. The temperature of the mixture decreases. Examples: thermal decomposition, dissolving ammonium nitrate, cold packs.
ΔT = Tfinal − Tinitial. A positive ΔT means exothermic. A negative ΔT means endothermic.
Acid + alkali (neutralisation), acid + metal, acid + carbonate, dissolving salts, displacement reactions.
Use a thermometer to record the temperature of your solution before adding the reactant. Record to the nearest 0.5°C.
Place the solution in a polystyrene (foam) cup inside a beaker. This acts as insulation to reduce heat loss to the surroundings and gives a more accurate result.
Add the second reactant all at once (e.g. solid into solution). Start a stopwatch if timing is required.
For exothermic reactions record the highest temperature reached. For endothermic reactions record the lowest temperature reached. Continue stirring throughout.
Use ΔT = Tfinal − Tinitial, then calculate the energy change using the formula q = mcΔT.
The heat energy change (q) is calculated using:
Heat energy change in joules (J)
Mass of solution in grams (g) — usually taken as the total volume of solution in cm³ ≈ g
Specific heat capacity of water = 4.18 J g⁻¹ °C⁻¹ (given in exam)
Temperature change in °C = Tfinal − Tinitial
50 cm³ of HCl was mixed with 50 cm³ of NaOH. Temperature rose from 21.0°C to 27.5°C.
Step 1: ΔT = 27.5 − 21.0 = 6.5°C
Step 2: m = 50 + 50 = 100 g (total solution volume)
Step 3: q = 100 × 4.18 × 6.5 = 2717 J = 2.72 kJ
Conclusion: Reaction is exothermic (temperature increased; q is positive).
| Reaction Type | Example | Exo or Endo? |
|---|---|---|
| Neutralisation | HCl + NaOH → NaCl + H₂O | 🔥 Exothermic |
| Acid + Metal | Zn + H₂SO₄ → ZnSO₄ + H₂ | 🔥 Exothermic |
| Acid + Carbonate | CaCO₃ + HCl → CaCl₂ + H₂O + CO₂ | 🔥 Exothermic |
| Displacement | Zn + CuSO₄ → ZnSO₄ + Cu | 🔥 Exothermic |
| Dissolving NH₄NO₃ | NH₄NO₃(s) → NH₄NO₃(aq) | ❄️ Endothermic |
| Thermal Decomposition | CaCO₃ → CaO + CO₂ (heated) | ❄️ Endothermic |
Label both axes with the quantity AND unit, e.g. "Temperature / °C" and "Time / s".
Do NOT just join all dots with a zigzag. Draw a smooth curve or straight line of best fit through the points.
Heat is lost during mixing. Draw the cooling section back to the time of mixing to find the corrected maximum temperature (extrapolation).
ΔT = corrected maximum temperature − initial temperature. Use this in q = mcΔT.
Extrapolation tip: The cooling part of the graph is approximately linear. Draw a straight line through the cooling points and extend it back to the time of mixing. The temperature at this intercept is the corrected Tmax.
| Error | Effect | How to Minimise |
|---|---|---|
| Heat loss to surroundings | ΔT is smaller than true value; q too low | Use a polystyrene cup with a lid; work quickly |
| Not stirring the mixture | Uneven temperature; incorrect Tmax recorded | Stir continuously with the thermometer |
| Thermometer not fully submerged | Reading does not reflect true solution temperature | Ensure bulb is fully immersed in the solution |
| Using a glass beaker instead of polystyrene | More heat lost; lower ΔT recorded | Always use polystyrene/foam cup as calorimeter |
| Assuming density of solution = 1 g/cm³ | Mass used in q = mcΔT may be slightly inaccurate | Accept this assumption unless told otherwise |
Exam tip: When asked to suggest an improvement to a thermochemical experiment, the most common correct answer is: "Use a polystyrene cup with a lid to reduce heat loss."
Sign convention: If q is positive (temperature rose) → reaction is exothermic. If q is negative (temperature fell) → reaction is endothermic. The sign tells you the direction of energy transfer.
Test your knowledge of Practical Paper 3 — 20 questions covering all topics.