Electrochemistry

Q: What is electrolysis and what is an electrolyte?

Electrolysis is the decomposition of an ionic compound by passing an electric current through it when molten or aqueous. An electrolyte is a molten ionic compound or aqueous solution containing free-moving ions that conducts electricity and is decomposed by it.

Q: How do the electrodes and ion movements correspond to charge during electrolysis?

The anode is the positive electrode attracting negative anions , while the cathode is the negative electrode attracting positive cations (remember PANIC : Positive Anode, Negative Is Cathode).

Q: Which processes occur at the cathode and the anode in terms of electron transfer?

Reduction (gain of electrons) occurs at the cathode, and oxidation (loss of electrons) occurs at the anode (remember OIL RIG and An Ox / Red Cat ).

Q: What products are formed during the electrolysis of dilute sulfuric acid, and in what volume ratio?

Hydrogen gas forms at the cathode and oxygen gas forms at the anode in a 2:1 volume ratio of \( \text{H}_2 : \text{O}_2 \).

Chemistry 1110 · CCEA GCSE · 7 min read · Updated: 31 Aug 2026

Welcome to Electrochemistry & Electrolysis!

Have you ever wondered how we extract shiny aluminium metal from rocks in the ground, or how car batteries store power? The answer lies in Electrochemistry. In this chapter, we explore electrolysis—a process where we use an electric current to break chemical compounds apart into their original elements.

Don't worry if this seems tricky at first! Once you learn a few simple memory tricks and rules, predicting the products of electrolysis becomes straightforward.

1. Key Vocabulary: The Language of Electrolysis

Before jumping into experiments, let's break down the essential terms you need to know for your exam:

• Electrolysis: The decomposition (splitting up) of an ionic compound, when molten or dissolved in water, by passing an electric current through it.
• Electrolyte: A molten ionic compound or an aqueous solution containing free-moving (mobile) ions that conducts electricity and is decomposed by it.
• Non-electrolyte: A liquid or solution that does not contain free ions and therefore cannot conduct electricity (e.g., pure water, ethanol, sugar solution, or covalent liquids).
• Electrodes: Rods or plates that carry electrical current into and out of the electrolyte. They are usually made of an unreactive (inert) material like graphite (carbon) or platinum so that they do not react with the products.
• Anode: The positive electrode.
• Cathode: The negative electrode.
• Anion: A negatively charged ion (attracted to the Anode).
• Cation: A positively charged ion (attracted to the Cathode).

Two Super Handy Memory Aids:

• PANIC: Positive Anode, Negative Is Cathode.
• OIL RIG: Oxidation Is Loss (of electrons), Reduction Is Gain (of electrons).

Quick Review: How Do Charges Behave?

Remember the golden rule of electrostatics: Opposites attract!
• Positive cations move towards the negative Cathode.
• Negative anions move towards the positive Anode.

2. How Electrolysis Works: Step-by-Step

Why do ionic solids not conduct electricity? In a solid crystal lattice, ions are locked tightly in fixed positions and cannot move. Electrolysis only works when an ionic compound is:

1. Molten (melted): Heating gives the ions enough energy to break free from the lattice and move freely.
2. Aqueous (dissolved in water): Water molecules separate the lattice so the ions are free to move.

What Happens at the Electrodes?

1. At the Cathode (Negative Electrode):
• Positive ions (cations) gain electrons from the cathode.
• Gaining electrons is reduction.
• For example: \( \text{Pb}^{2+} + 2\text{e}^- \rightarrow \text{Pb} \)

2. At the Anode (Positive Electrode):
• Negative ions (anions) lose electrons to the anode.
• Losing electrons is oxidation.
• For example: \( 2\text{Br}^- \rightarrow \text{Br}_2 + 2\text{e}^- \)

Key Takeaway: Oxidation always happens at the anode; reduction always happens at the cathode (Remember: An Ox and Red Cat).

3. Electrolysis of Molten Compounds

Example 1: Molten Lead(II) Bromide, \( \text{PbBr}_2 \)

This is a classic laboratory demonstration. Solid lead(II) bromide is heated in a crucible until it melts. Two inert graphite electrodes connected to a DC power supply are dipped into the liquid.

Observations:

• At the Cathode: A silvery bead/liquid metal forms (molten lead, \( \text{Pb} \)).
• At the Anode: Bubbles of a pungent, orange-brown gas/vapour are given off (bromine, \( \text{Br}_2 \)).
• Safety Note: This experiment must be done in a fume cupboard because bromine gas is toxic and corrosive.

Half-Equations:

• Cathode (Reduction): \( \text{Pb}^{2+} + 2\text{e}^- \rightarrow \text{Pb} \)
• Anode (Oxidation): \( 2\text{Br}^- \rightarrow \text{Br}_2 + 2\text{e}^- \)
• Overall Equation: \( \text{PbBr}_2\text{ (l)} \rightarrow \text{Pb}\text{ (l)} + \text{Br}_2\text{ (g)} \)

Example 2: Molten Lithium Chloride, \( \text{LiCl} \)

• Cathode: Lithium ions are reduced to silvery lithium metal: \( \text{Li}^+ + \text{e}^- \rightarrow \text{Li} \)
• Anode: Chloride ions are oxidised to pale green/yellow chlorine gas: \( 2\text{Cl}^- \rightarrow \text{Cl}_2 + 2\text{e}^- \)
• Overall Equation: \( 2\text{LiCl}\text{ (l)} \rightarrow 2\text{Li}\text{ (l)} + \text{Cl}_2\text{ (g)} \)

4. Electrolysis of Dilute Sulfuric Acid, \( \text{H}_2\text{SO}_4 \)

When dilute sulfuric acid is electrolysed using inert platinum or graphite electrodes (often in a Hofmann voltameter), the water is split into hydrogen and oxygen gases.

Ions Present:

• From acid and water: \( \text{H}^+ \), \( \text{OH}^- \), and \( \text{SO}_4^{2-} \)

Electrode Reactions and Observations:

• At the Cathode: \( \text{H}^+ \) ions are attracted, gaining electrons to form colourless hydrogen gas.
Half-equation: \( 2\text{H}^+ + 2\text{e}^- \rightarrow \text{H}_2 \)
Test for gas: Squeaky pop test with a burning splint.

• At the Anode: Hydroxide ions (\( \text{OH}^- \)) lose electrons more readily than sulfate ions, forming colourless oxygen gas and water.
Half-equation: \( 4\text{OH}^- \rightarrow \text{O}_2 + 2\text{H}_2\text{O} + 4\text{e}^- \)
Test for gas: Relights a glowing splint.

Volume Ratio:

For every 4 electrons transferred, we make 2 molecules of \( \text{H}_2 \) and only 1 molecule of \( \text{O}_2 \).
Therefore, the volume of hydrogen produced at the cathode is twice (2:1 ratio) the volume of oxygen produced at the anode!

5. Industrial Extraction of Aluminium

Aluminium is very high in the reactivity series (more reactive than carbon), so it cannot be extracted by reduction with carbon. Instead, it must be extracted using electrolysis.

Step 1: Raw Materials and Preparation

• The main ore is bauxite, which is purified to yield white aluminium oxide powder (\( \text{Al}_2\text{O}_3 \)), also called alumina.
• Alumina has a huge melting point of around \( 2070^\circ\text{C} \). Melting it directly would require an enormous amount of expensive energy.
• The Solution: Alumina is dissolved in molten cryolite (\( \text{Na}_3\text{AlF}_6 \)).

Why use cryolite?

1. It lowers the working melting point to around \( 900^\circ\text{C} - 1000^\circ\text{C} \), saving vast amounts of energy and reducing costs.
2. It improves electrical conductivity of the electrolyte mixture.

Step 2: The Electrolytic Cell

• Cathode: The carbon (graphite) lining of the large steel cell.
• Anode: Large carbon (graphite) blocks suspended into the molten electrolyte.

Step 3: Reactions in the Cell

• At the Cathode: Aluminium ions gain electrons and are reduced to molten aluminium metal.
Half-equation: \( \text{Al}^{3+} + 3\text{e}^- \rightarrow \text{Al} \)
Because molten aluminium is denser than the cryolite mixture, it sinks to the bottom and is tapped off periodically.

• At the Anode: Oxide ions lose electrons to form oxygen gas.
Half-equation: \( 2\text{O}^{2-} \rightarrow \text{O}_2 + 4\text{e}^- \)

A Common Exam Question: Why Must the Carbon Anodes Be Regularly Replaced?

At high operating temperatures (\( \sim 950^\circ\text{C} \)), the oxygen gas produced at the anode reacts with the hot carbon (graphite) blocks to produce carbon dioxide gas:
\( \text{C}\text{ (s)} + \text{O}_2\text{ (g)} \rightarrow \text{CO}_2\text{ (g)} \)
This causes the anodes to gradually burn away, meaning they must be replaced periodically.

Economic & Environmental Aspects: Why Recycle Aluminium?

• Huge energy savings: Recycling aluminium uses only about 5% of the energy required to extract it from bauxite (a 95% energy saving!).
• Conserves natural resources: Preserves limited reserves of bauxite ore.
• Reduces greenhouse gases: Lowers the amount of \( \text{CO}_2 \) released into the atmosphere from electricity generation and burning anodes.
• Reduces waste: Prevents bauxite waste (red mud) from polluting landscapes and keeps aluminium cans out of landfills.

6. Common Mistakes to Avoid in Exams

• Confusing the electrodes: Remember PANIC (Positive Anode, Negative Is Cathode).
• Writing diatomic elements as single atoms: Halogens and oxygen always exist as diatomic molecules when formed: write \( \text{Cl}_2 \), \( \text{Br}_2 \), \( \text{I}_2 \), \( \text{O}_2 \), or \( \text{H}_2 \), not \( \text{Cl} \) or \( \text{O} \).
• Forgetting to balance half-equations: Make sure both atoms and charges balance. If you produce \( \text{Br}_2 \), you need \( 2\text{Br}^- \) and \( 2\text{e}^- \).
• Thinking solid salts conduct: Solid ionic compounds do not conduct because ions are fixed in place. Always state that ions are free to move in molten liquid or solution.

7. Summary Checklist

Before moving on to the next chapter, check that you can:
• Define electrolysis, electrolyte, anode, and cathode.
• Use OIL RIG to identify where oxidation and reduction occur.
• Predict products and write balanced half-equations for the electrolysis of molten salts (e.g. \( \text{PbBr}_2 \)).
• Explain the role of cryolite in the extraction of aluminium.
• Explain why carbon anodes must be replaced during aluminium extraction.
• Give key economic and environmental benefits of recycling aluminium.

Quick check

Can you answer these now?

Open each question to check the key ideas from this chapter.

What is electrolysis and what is an electrolyte?

Electrolysis is the decomposition of an ionic compound by passing an electric current through it when molten or aqueous. An electrolyte is a molten ionic compound or aqueous solution containing free-moving ions that conducts electricity and is decomposed by it.

How do the electrodes and ion movements correspond to charge during electrolysis?

The anode is the positive electrode attracting negative anions, while the cathode is the negative electrode attracting positive cations (remember PANIC: Positive Anode, Negative Is Cathode).

Why do ionic substances conduct electricity only when molten or dissolved in water?

In solid ionic compounds, ions are held in fixed positions in the lattice and cannot move; when molten or in solution, the lattice breaks down so ions are free to move and carry charge.

Which processes occur at the cathode and the anode in terms of electron transfer?

Reduction (gain of electrons) occurs at the cathode, and oxidation (loss of electrons) occurs at the anode (remember OIL RIG and An Ox / Red Cat).

What are the observations and half-equations for the electrolysis of molten lead(II) bromide ( \( \text{PbBr}_2 \))?

Cathode: silvery molten lead forms (\( \text{Pb}^{2+} + 2\text{e}^- \rightarrow \text{Pb} \)). Anode: pungent orange-brown bromine gas forms (\( 2\text{Br}^- \rightarrow \text{Br}_2 + 2\text{e}^- \)).

What products are formed during the electrolysis of dilute sulfuric acid, and in what volume ratio?

Hydrogen gas forms at the cathode and oxygen gas forms at the anode in a 2:1 volume ratio of \( \text{H}_2 : \text{O}_2 \).

Why is cryolite added to aluminium oxide (alumina) in the industrial extraction of aluminium?

Cryolite lowers the melting point of alumina from \( \sim 2070^\circ\text{C} \) to \( \sim 900\text{--}1000^\circ\text{C} \) and improves electrical conductivity, saving significant energy and costs.

What reactions take place at the electrodes during the electrolysis of aluminium oxide?

Cathode: \( \text{Al}^{3+} + 3\text{e}^- \rightarrow \text{Al} \) (molten aluminium sinks to the bottom). Anode: \( 2\text{O}^{2-} \rightarrow \text{O}_2 + 4\text{e}^- \) (oxygen gas formed).

Why do the carbon anodes in the aluminium extraction cell need to be replaced periodically?

The oxygen gas produced at the anode reacts with the hot carbon electrodes at \( \sim 950^\circ\text{C} \) to form carbon dioxide gas (\( \text{C} + \text{O}_2 \rightarrow \text{CO}_2 \)), causing them to burn away.

What are the main economic and environmental advantages of recycling aluminium?

Recycling uses only 5% of the energy needed for extraction (a 95% saving), conserves raw bauxite reserves, reduces greenhouse gas emissions, and reduces landfill waste.

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