(I once showed a class a beaker of water and asked, "Can you light a bulb with this?" They laughed. Then I added a pinch of salt, connected a battery, and the bulb glowed. Their faces changed. That is the power of this chapter. Electricity can turn insulators into conductors, break compounds apart into their elements, and coat one metal onto another. It is one of the most useful tools in all of chemistry — and in this chapter, you will learn how it works.) Think about everything you plug in. Your phone charger, your laptop, your kettle, your lights. Electricity powers the modern world. But electricity does not just power devices — it can also change substances. When electricity passes through certain liquids or solutions, it can break them apart into their elements, deposit metals onto surfaces, and produce useful chemicals. This process is called electrolysis, and it is used everywhere: to extract aluminium from its ore, to chrome-plate car parts, to refine copper, to make chlorine and hydrogen, and to power the batteries in your phone. In this chapter, we are going to learn what happens when electricity meets chemistry. We will look at which substances conduct electricity and which do not, how electrolysis works, what happens at the electrodes, and how it is used in industry. By the end, you will understand one of the most important industrial processes in the world. Let us begin.
Not all substances conduct electricity. Those that do are called conductors, and those that do not are called non-conductors (or insulators).
| Type | Description | Examples |
|---|---|---|
| Metallic conductors | Conduct electricity by the movement of delocalised electrons. Do not undergo chemical change. | Copper, aluminium, silver, graphite |
| Electrolytes | Conduct electricity when molten or dissolved in water. Conduct by the movement of ions. Undergo chemical change. | Sodium chloride solution, dilute sulphuric acid, copper(II) sulphate solution |
| Non-electrolytes | Do not conduct electricity in any state. | Sugar, ethanol, pure water (very poor conductor) |
(Here is the key difference: metallic conductors conduct by electrons, and they do not change chemically. Electrolytes conduct by ions, and they do change chemically — they break down. That is why electrolysis is possible.)
When an electrolyte is dissolved in water or melted, it dissociates into ions. These ions are free to move and carry charge. When electricity is passed through the electrolyte, the ions migrate to the electrodes of opposite charge:
At the electrodes, the ions gain or lose electrons and become neutral atoms or molecules. This is electrolysis.
Key definitions:
When a molten ionic compound is electrolysed, the ions are free to move. Let us look at the electrolysis of molten lead(II) bromide (PbBr₂).
What happens:
Overall equation: PbBr₂ → Pb + Br₂
When an aqueous solution is electrolysed, water is also present. This means there are additional ions (H⁺ and OH⁻) that can compete at the electrodes. The ion that is discharged depends on its position in the electrochemical series and its concentration.
At the cathode: H⁺ ions are discharged: 2H⁺ + 2e⁻ → H₂. Hydrogen gas is released.
At the anode: OH⁻ ions are discharged: 4OH⁻ → 2H₂O + O₂ + 4e⁻. Oxygen gas is released.
Overall: Water is decomposed into hydrogen and oxygen: 2H₂O → 2H₂ + O₂.
At the cathode: Cu²⁺ ions are discharged: Cu²⁺ + 2e⁻ → Cu. Copper is deposited.
At the anode: Copper dissolves: Cu → Cu²⁺ + 2e⁻. The anode loses mass.
Net effect: Copper is transferred from the anode to the cathode. This is the basis of copper refining.
At the cathode: Cu²⁺ ions are discharged: Cu²⁺ + 2e⁻ → Cu. Copper is deposited.
At the anode: OH⁻ ions are discharged: 4OH⁻ → 2H₂O + O₂ + 4e⁻. Oxygen gas is released.
When more than one type of ion is present, the one that is discharged depends on:
| Ion | Ease of Discharge |
|---|---|
| K⁺, Na⁺, Ca²⁺, Mg²⁺, Al³⁺ | Very difficult — remain in solution |
| Zn²⁺, Fe²⁺, Pb²⁺ | Moderate |
| Cu²⁺, Ag⁺, Au⁺ | Easy — discharged in preference to H⁺ |
| Cl⁻, Br⁻, I⁻ | Easy — discharged in preference to OH⁻ at high concentration |
| OH⁻, SO₄²⁻, NO₃⁻ | OH⁻ is discharged in preference to SO₄²⁻ and NO₃⁻ |
Electrolysis is used in many industries. Here are the most important applications.
| Application | What Happens | Why It Matters |
|---|---|---|
| Extraction of reactive metals | Metals like aluminium, sodium, and calcium are extracted from their molten ores by electrolysis. | These metals cannot be extracted by carbon reduction because they are too reactive. |
| Electroplating | A thin layer of one metal is deposited onto another by electrolysis. | Used to make objects look better, resist corrosion, or conduct electricity better. |
| Refining of copper | Impure copper is used as the anode, pure copper as the cathode. Copper transfers from anode to cathode. | Produces 99.99% pure copper for electrical wiring. |
| Manufacture of chlorine and hydrogen | Electrolysis of brine (concentrated sodium chloride solution) produces chlorine, hydrogen, and sodium hydroxide. | These are essential industrial chemicals. |
| Anodising aluminium | Aluminium is made the anode in an electrolytic cell, forming a thick oxide layer. | Protects against corrosion and allows dyeing. |
| Electrolytic capacitors | Used in electronic circuits. | Store and release electrical energy. |
(Electrolysis is one of those rare things that is both simple and incredibly powerful. A battery, two electrodes, and a solution — and you can extract metals, purify them, and coat them. It is the quiet engine of modern industry.)
Brine is a concentrated solution of sodium chloride. When it is electrolysed, three useful products are formed:
Equations:
This process is the basis of the chlor-alkali industry, which produces chlorine, hydrogen, and sodium hydroxide for use in water treatment, soap making, PVC production, and many other industries.
Michael Faraday discovered two important laws of electrolysis:
These laws allow chemists to calculate exactly how much product will be formed during electrolysis. They are essential in industrial electrochemistry.
Key formula: m = (Q × M) ÷ (n × F)
Where:
Question: In the electrolysis of molten lead(II) bromide, state what is formed at the cathode and at the anode, and write the equations.
Solution:
• At the cathode (negative electrode): Pb²⁺ ions gain electrons → lead metal: Pb²⁺ + 2e⁻ → Pb.
• At the anode (positive electrode): Br⁻ ions lose electrons → bromine gas: 2Br⁻ → Br₂ + 2e⁻.
Question: Describe what happens at each electrode when copper(II) sulphate solution is electrolysed using copper electrodes.
Solution:
• At the cathode: Cu²⁺ ions are discharged and copper is deposited: Cu²⁺ + 2e⁻ → Cu.
• At the anode: Copper dissolves from the anode: Cu → Cu²⁺ + 2e⁻.
• Overall: Copper is transferred from the anode to the cathode. This is used in copper refining.
Question: Calculate the mass of copper deposited when a current of 2 A flows for 30 minutes through copper(II) sulphate solution. (Molar mass of Cu = 63.5 g/mol, F = 96500 C/mol.)
Solution:
• Charge Q = current × time = 2 × (30 × 60) = 2 × 1800 = 3600 C.
• Moles of electrons = Q / F = 3600 / 96500 = 0.0373 mol.
• Cu²⁺ + 2e⁻ → Cu, so 2 moles of electrons deposit 1 mole of Cu.
• Moles of Cu = 0.0373 / 2 = 0.01865 mol.
• Mass of Cu = 0.01865 × 63.5 = 1.18 g.
1. What is an electrolyte? Give one example.
An electrolyte is a substance that conducts electricity when molten or dissolved in water and is decomposed by it. Example: sodium chloride solution.
2. Define electrolysis.
Electrolysis is the decomposition of an electrolyte by the passage of an electric current.
3. What is the charge on the cathode and the anode?
Cathode: negative. Anode: positive.
4. What happens to cations and anions during electrolysis?
Cations (positive ions) move to the cathode and gain electrons. Anions (negative ions) move to the anode and lose electrons.
5. State two applications of electrolysis.
Any two of: extraction of reactive metals, electroplating, refining of copper, manufacture of chlorine and hydrogen, anodising aluminium.
6. Describe what happens at the cathode and anode during the electrolysis of molten lead(II) bromide.
Cathode: Pb²⁺ + 2e⁻ → Pb (lead metal deposited). Anode: 2Br⁻ → Br₂ + 2e⁻ (bromine gas released).
7. What are the products of the electrolysis of brine?
Hydrogen gas at the cathode, chlorine gas at the anode, and sodium hydroxide solution remaining.
Question 1 [10 marks]
(a) Distinguish between metallic conductors and electrolytes. [4]
(b) Describe the ionic theory of electrolysis. [3]
(c) Explain what happens at the cathode and anode during the electrolysis of molten lead(II) bromide. [3]
(a) Metallic conductors conduct electricity by the movement of delocalised electrons and do not undergo chemical change. Electrolytes conduct electricity by the movement of ions when molten or in aqueous solution and are decomposed by the current. (4 marks: 2 for metallic conductors, 2 for electrolytes.)
(b) When an electrolyte is dissolved or melted, it dissociates into ions. These ions are free to move. When a current is passed, cations move to the cathode and anions move to the anode. At the electrodes, ions gain or lose electrons and become neutral. (3 marks: 1 for dissociation, 1 for ion movement, 1 for discharge at electrodes.)
(c) At the cathode (negative electrode), Pb²⁺ ions gain electrons and lead metal is deposited: Pb²⁺ + 2e⁻ → Pb. At the anode (positive electrode), Br⁻ ions lose electrons and bromine gas is released: 2Br⁻ → Br₂ + 2e⁻. (3 marks: 1 for cathode, 1 for anode, 1 for equations.)
Question 2 [10 marks]
(a) Describe what happens during the electrolysis of copper(II) sulphate solution using: (i) copper electrodes, (ii) inert electrodes. [6]
(b) Explain why the products are different in the two cases. [4]
(a) (i) With copper electrodes: at the cathode, Cu²⁺ + 2e⁻ → Cu (copper deposited). At the anode, Cu → Cu²⁺ + 2e⁻ (copper dissolves). Copper is transferred from anode to cathode. (ii) With inert electrodes: at the cathode, Cu²⁺ + 2e⁻ → Cu (copper deposited). At the anode, OH⁻ ions are discharged: 4OH⁻ → 2H₂O + O₂ + 4e⁻ (oxygen gas released). (6 marks: 3 for each case.)
(b) The products are different because the anode behaves differently. With copper electrodes, the anode is active — it dissolves and contributes Cu²⁺ ions to the solution. With inert electrodes, the anode does not dissolve, so OH⁻ ions from water are discharged instead, producing oxygen gas. The nature of the electrode determines which ions are discharged at the anode. (4 marks: 1 for active anode, 1 for inert anode, 1 for ion contribution, 1 for conclusion.)
Question 3 [10 marks]
(a) Describe the electrolysis of brine (concentrated sodium chloride solution), stating the products at each electrode and in the solution. [5]
(b) State two uses of chlorine and two uses of sodium hydroxide. [4]
(c) Explain why chlorine is discharged at the anode instead of oxygen in the electrolysis of concentrated brine. [1]
(a) At the cathode: hydrogen gas is released (2H⁺ + 2e⁻ → H₂). At the anode: chlorine gas is released (2Cl⁻ → Cl₂ + 2e⁻). In the solution: sodium hydroxide remains. Overall: 2NaCl + 2H₂O → 2NaOH + H₂ + Cl₂. (5 marks: 1 for cathode, 1 for anode, 1 for solution, 1 for equations, 1 for overall.)
(b) Chlorine: water treatment (killing bacteria), making bleach, making PVC. Sodium hydroxide: making soap, making paper, water treatment, making other chemicals. (4 marks: 1 for each use.)
(c) In concentrated brine, the concentration of Cl⁻ ions is high, so they are discharged in preference to OH⁻ ions. (1 mark.)
Challenge 1: Explain why pure water is a very poor conductor of electricity, but adding a pinch of salt makes it a good conductor.
Pure water contains very few ions — only a tiny amount of H⁺ and OH⁻ from the self-ionisation of water. With so few charge carriers, the conductivity is very low. When salt (sodium chloride) is added, it dissolves and dissociates into Na⁺ and Cl⁻ ions. These ions are free to move and carry charge through the solution. The more ions present, the greater the conductivity. So adding salt increases the number of charge carriers, turning pure water from a poor conductor into a good conductor.
Challenge 2: Explain why copper is refined by electrolysis using copper electrodes, rather than by simply melting and re-casting it.
Copper is refined by electrolysis because it is the only practical way to obtain very high purity (99.99%) needed for electrical wiring. Simply melting and re-casting copper does not remove the impurities (such as silver, gold, and other metals) that reduce its conductivity. In electrolysis with copper electrodes, the impure copper is used as the anode, and pure copper is deposited at the cathode. The impurities either remain in the solution or fall to the bottom as "anode sludge." The copper that transfers to the cathode is extremely pure. This level of purity is essential for electrical applications, where even small amounts of impurity can significantly reduce conductivity.
Challenge 3: During the electrolysis of copper(II) sulphate solution with copper electrodes, the blue colour of the solution remains constant. Explain why, using the reactions at the electrodes.
The blue colour of copper(II) sulphate solution is due to the presence of Cu²⁺ ions. During electrolysis with copper electrodes, Cu²⁺ ions are discharged at the cathode (Cu²⁺ + 2e⁻ → Cu), which would normally reduce the concentration of Cu²⁺ ions and fade the blue colour. However, at the same time, copper dissolves from the anode (Cu → Cu²⁺ + 2e⁻), releasing new Cu²⁺ ions into the solution. The rate at which Cu²⁺ ions are removed at the cathode equals the rate at which they are added at the anode. So the concentration of Cu²⁺ ions in the solution remains constant, and the blue colour stays the same. The net effect is the transfer of copper from the anode to the cathode.
"Electricity can power a light, but it can also break a molecule, purify a metal, and coat a surface. Learn the chemistry behind the current, and you can make it work for you."
