Required Practical: Electrolysis

This lesson covers the AQA Required Practical for investigating the electrolysis of aqueous solutions using inert electrodes, as specified in AQA GCSE Chemistry (Required Practical 2, specification 4.4.3). You must be able to describe the apparatus, carry out the practical, identify the products at each electrode, and explain the results using knowledge of ions and redox chemistry. This practical is commonly examined and connects theory directly to laboratory skills.

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Aim

To investigate what happens during the electrolysis of different aqueous solutions using inert (unreactive) electrodes, and to identify the products formed at each electrode.

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Apparatus and Materials

Item	Purpose
Beaker (250 cm3)	To hold the electrolyte solution
Two graphite (carbon) electrodes	Inert electrodes — do not react with the products
DC power supply (or battery pack)	Provides direct current for electrolysis
Connecting wires with crocodile clips	To connect the electrodes to the power supply
Test tubes (optional)	To collect gases produced at the electrodes
Wooden splint	To test for hydrogen (squeaky pop test)
Litmus paper (blue, damp)	To test for chlorine (bleaches litmus)
Glowing splint	To test for oxygen (relights a glowing splint)
Aqueous solutions: copper sulfate, sodium chloride, sodium sulfate, copper chloride	Different electrolytes to test

Exam Tip: Graphite (carbon) electrodes are used because they are inert — they do not react with the products of electrolysis or dissolve into the solution. If reactive electrodes (like copper) were used, the anode would dissolve, affecting the results. Always state that graphite electrodes are inert when explaining the choice of electrode.

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Method

Setting Up the Apparatus

1. Pour approximately 100 cm3 of the electrolyte solution into a clean beaker.
2. Place two graphite electrodes into the solution, ensuring they do not touch each other.
3. Connect the electrodes to a DC power supply using connecting wires with crocodile clips.
4. Connect the positive terminal of the power supply to one electrode (this becomes the anode).
5. Connect the negative terminal of the power supply to the other electrode (this becomes the cathode).

graph TD
    A["DC Power Supply<br/>(6V)"] -->|"Positive wire (+)"| B["ANODE (+)<br/>Graphite electrode"]
    A -->|"Negative wire (-)"| C["CATHODE (-)<br/>Graphite electrode"]
    B --> D["Beaker containing<br/>aqueous electrolyte<br/>solution"]
    C --> D
    D --> E["Observe: bubbles,<br/>colour changes,<br/>metal deposits"]
    E --> F["Test gases at<br/>each electrode"]

    style A fill:#4a90d9,color:#fff
    style B fill:#27ae60,color:#fff
    style C fill:#e74c3c,color:#fff
    style D fill:#9b59b6,color:#fff

Running the Electrolysis

6. Turn on the power supply and set it to approximately 6 volts.
7. Observe both electrodes carefully for several minutes.
8. Record any observations at each electrode:
   • Bubbles of gas
   • Colour changes in the solution
   • Deposits forming on the electrode
   • Smell (cautiously waft — do not inhale directly)

Collecting and Testing Gases

9. To collect a gas, place an inverted test tube filled with the electrolyte solution over the electrode. The gas will displace the solution and collect in the tube.
10. Test the collected gas:

Gas	Test	Positive result
Hydrogen	Apply a burning splint to the mouth of the test tube	Squeaky pop
Oxygen	Insert a glowing splint into the test tube	Splint relights
Chlorine	Hold damp blue litmus paper near the electrode	Litmus paper turns red then white (bleached)

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Results

Electrolysis of Copper Sulfate Solution

Electrode	Observation	Product	Explanation
Cathode (-)	Orange-brown solid deposited on electrode	Copper	Cu2+ ions are attracted, gain 2 electrons (reduction): Cu2+ + 2e- → Cu. Copper is less reactive than hydrogen.
Anode (+)	Bubbles of gas; glowing splint relights	Oxygen	OH- ions from water are discharged (no halide present): 4OH- → 2H2O + O2 + 4e- (oxidation)
Solution	Gradually becomes paler blue, eventually colourless	—	Cu2+ ions (which cause the blue colour) are removed from solution as copper metal

Electrolysis of Sodium Chloride Solution (Brine)

Electrode	Observation	Product	Explanation
Cathode (-)	Bubbles of gas; squeaky pop with burning splint	Hydrogen	H+ ions (from water) are discharged because sodium is more reactive than hydrogen: 2H+ + 2e- → H2
Anode (+)	Bubbles of gas; bleaches damp litmus paper (turns white)	Chlorine	Cl- ions are discharged (halide present): 2Cl- → Cl2 + 2e- (oxidation)
Solution	Becomes alkaline (pH increases); solution near cathode turns universal indicator blue/purple	Sodium hydroxide remains	Na+ and OH- ions remain in solution, forming NaOH

Electrolysis of Sodium Sulfate Solution

Electrode	Observation	Product	Explanation
Cathode (-)	Bubbles of gas; squeaky pop	Hydrogen	Na is more reactive than H, so H+ ions are discharged: 2H+ + 2e- → H2
Anode (+)	Bubbles of gas; relights glowing splint	Oxygen	No halide present, so OH- ions are discharged: 4OH- → 2H2O + O2 + 4e-

Exam Tip: When writing up results for this practical, always state what you OBSERVED (e.g. "bubbles of gas", "brown solid formed on electrode") AND what the product IS (e.g. "hydrogen gas", "copper metal"), AND the test you did to confirm it (e.g. "squeaky pop with a burning splint"). Three elements for full marks: observation, identification, and test.

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Safety Precautions

Hazard	Precaution
Chlorine gas — toxic	Work in a well-ventilated room or use a fume cupboard; do not inhale; only use small volumes
Sulfuric acid / sodium hydroxide — corrosive/irritant	Wear safety goggles and avoid skin contact
Electrical equipment near liquids	Keep the power supply away from liquids; do not touch electrodes while the circuit is on; use low voltage (6V)
Hot electrodes	Allow to cool before handling

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Analysing the Results

Why Different Solutions Give Different Products

The products of electrolysis depend on:

1. The ions present in the solution (from the dissolved compound AND from water)
2. The reactivity of the metal ion compared to hydrogen (for the cathode product)
3. Whether a halide ion is present (for the anode product)

Solution	Ions present	Cathode product	Anode product
Copper sulfate	Cu2+, SO4 2-, H+, OH-	Copper (Cu less reactive than H)	Oxygen (no halide)
Sodium chloride	Na+, Cl-, H+, OH-	Hydrogen (Na more reactive than H)	Chlorine (halide present)
Sodium sulfate	Na+, SO4 2-, H+, OH-	Hydrogen (Na more reactive than H)	Oxygen (no halide)
Copper chloride	Cu2+, Cl-, H+, OH-	Copper (Cu less reactive than H)	Chlorine (halide present)
Potassium iodide	K+, I-, H+, OH-	Hydrogen (K more reactive than H)	Iodine (halide present)

Exam Tip: A common exam question provides an unfamiliar electrolyte and asks you to predict the products. Apply the two rules: (1) At the cathode — compare the metal to hydrogen in the reactivity series; if the metal is more reactive, hydrogen is produced. (2) At the anode — if a halide (Cl-, Br-, I-) is present, the halogen is produced; if not, oxygen is produced. These rules work for every aqueous solution.

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Common Exam Questions on this Practical

Q: Explain why copper is deposited at the cathode during electrolysis of copper sulfate solution. [3 marks]