The GCSE chemistry electrolysis required practical investigates copper sulphate solution using copper electrodes. You measure how the mass of copper deposited on the cathode changes with current or time, linking charge flow to mass transfer between anode and cathode. It is one of the most calculation-rich practicals in the GCSE course.
What is the aim of this required practical?
The electrolysis required practical at GCSE typically asks you to:
- Electrolyse copper(II) sulphate solution using copper electrodes.
- Measure the mass of copper deposited on the cathode over a fixed time.
- Investigate how current (or concentration, or time) affects the mass deposited.
- Plot a graph and draw a conclusion about the relationship.
At a deeper level, the experiment demonstrates Faraday's First Law of Electrolysis — the mass of a substance deposited at an electrode is directly proportional to the charge passed through the electrolyte.
What equipment do you need?
| Equipment | Purpose |
|---|---|
| Copper electrodes (× 2) | Act as anode and cathode |
| Beaker of copper sulphate solution | Electrolyte; provides Cu²⁺ ions |
| Variable resistor (rheostat) | Controls the current |
| Ammeter (in series) | Measures current in amps |
| Power pack or battery | Provides the d.c. supply |
| Stopwatch | Controls the time of electrolysis |
| Electronic balance (to ±0.01 g) | Measures mass of cathode before and after |
| Emery paper (sandpaper) | Cleans electrodes before each run |
| Connecting wires and crocodile clips | Completes the circuit |
Safety: copper sulphate is irritating to skin and eyes. Wear safety goggles and gloves. Do not ingest. Dispose of solution according to school guidelines.
How do you carry out the experiment step by step?
- Clean both copper electrodes with emery paper to remove any oxide layer. Rinse with deionised water, then dry carefully.
- Weigh both electrodes on the electronic balance. Record the initial mass of the cathode (negative electrode) to two decimal places.
- Set up the circuit: connect the power pack, ammeter, variable resistor, and electrodes in series. Immerse both electrodes in the copper sulphate solution.
- Set the current using the variable resistor. For the first run, set the ammeter to (e.g.) 0.2 A. Record the current.
- Start the stopwatch and switch on the power supply simultaneously. Run for exactly 5 minutes (300 seconds).
- Switch off the power supply at exactly 5 minutes.
- Remove the cathode carefully, rinse with deionised water, allow to dry thoroughly (or pat dry with tissue), then weigh it again.
- Calculate the mass of copper deposited: mass deposited = final mass − initial mass.
- Repeat with different current settings (e.g. 0.4 A, 0.6 A, 0.8 A, 1.0 A), cleaning and re-weighing the electrodes each time.
How do you record and process results?
| Current (A) | Time (s) | Initial mass of cathode (g) | Final mass of cathode (g) | Mass of Cu deposited (g) | Charge (C) = I × t |
|---|---|---|---|---|---|
| 0.2 | 300 | 5.42 | 5.62 | 0.20 | 60 |
| 0.4 | 300 | 5.62 | 6.04 | 0.42 | 120 |
| 0.6 | 300 | 6.04 | 6.64 | 0.60 | 180 |
| 0.8 | 300 | 6.64 | 7.44 | 0.80 | 240 |
| 1.0 | 300 | 7.44 | 8.45 | 1.01 | 300 |
Plot: mass deposited (y-axis, g) against charge passed (x-axis, coulombs = A × s). The graph should be a straight line through the origin, showing direct proportionality.
Worked example calculation:
At 0.6 A for 300 s:
- Charge (Q) = I × t = 0.6 × 300 = 180 C
- Mass deposited = 0.60 g
What are the half-equations at each electrode?
At the cathode (negative electrode), copper ions are reduced (gain electrons):
Cu²⁺ (aq) + 2e⁻ → Cu (s)
At the anode (positive electrode), copper atoms are oxidised (lose electrons):
Cu (s) → Cu²⁺ (aq) + 2e⁻
This is why the copper anode gradually loses mass (copper dissolves) and the cathode gains mass (copper deposits) — and why the copper sulphate solution concentration stays approximately constant. The copper anode replenishes the Cu²⁺ ions that are removed at the cathode.
What variables must you control?
To make the investigation a fair test when varying current:
- Time: identical for each run (e.g. 300 seconds)
- Volume and concentration of copper sulphate solution: use the same solution throughout, or fresh solution of the same concentration each time
- Electrode surface area: use electrodes of the same size and ensure the same area is submerged each run
- Temperature: room temperature (work in the same conditions)
- Distance between electrodes: keep constant
Independent variable: current (in amps) Dependent variable: mass of copper deposited (in grams) Controlled variables: time, concentration, electrode area, temperature
Frequently asked questions
Why does the cathode gain mass and the anode lose mass during electrolysis of copper sulphate?
At the cathode, positive copper ions (Cu²⁺) from the solution are attracted and gain electrons (are reduced), forming solid copper metal that deposits on the cathode surface — so the cathode gains mass. At the anode, copper atoms in the electrode lose electrons (are oxidised) and enter the solution as Cu²⁺ ions — so the anode loses mass. The overall effect is a transfer of copper from anode to cathode, which is the basis of commercial electroplating and electrolytic refining of copper.
Why is the graph of mass deposited against charge a straight line through the origin?
Faraday's First Law states that the mass of substance deposited at an electrode is directly proportional to the charge (Q = I × t) passed through the electrolyte. Doubling the charge doubles the number of electrons available, which doubles the number of Cu²⁺ ions reduced to solid copper. The straight line through the origin reflects this direct proportionality. Any deviation from a straight line suggests an uncontrolled variable — most often the current fluctuating or the electrodes not being cleaned properly.
What would happen if you used inert (graphite) electrodes instead of copper?
With graphite electrodes and copper sulphate solution, the cathode reaction is the same (Cu²⁺ + 2e⁻ → Cu). However, at the anode, water is oxidised instead of carbon: 2H₂O → O₂ + 4H⁺ + 4e⁻, producing oxygen gas. Since the anode no longer replenishes Cu²⁺ ions, the blue colour of the copper sulphate solution gradually fades and the solution becomes more acidic as H⁺ ions accumulate. This is a common comparison question in GCSE exams.
Why must you dry the cathode thoroughly before weighing?
Water droplets remaining on the cathode add mass that does not come from deposited copper. If the cathode is weighed while still damp, the recorded final mass will be higher than the true mass of copper deposited, making the calculated mass too large. Thorough drying — either in a warm oven, under a lamp, or by carefully patting dry with tissue — ensures only the mass of the deposited copper is recorded.
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