Electrochemistry

What you'll learn

Electrochemistry forms a substantial component of the CIE IGCSE Chemistry syllabus, covering the relationship between electrical energy and chemical reactions. This topic examines electrolysis of molten compounds and aqueous solutions, the extraction and purification of metals, and electrochemical cells. Exam questions frequently test the ability to predict products at electrodes, write half-equations, and explain industrial applications.

Key terms and definitions

Electrolysis — the decomposition of an ionic compound, molten or in aqueous solution, by passing an electric current through it

Electrolyte — a substance that conducts electricity when molten or in aqueous solution, undergoing chemical decomposition in the process

Anode — the positive electrode in electrolysis where oxidation occurs; anions are attracted to this electrode

Cathode — the negative electrode in electrolysis where reduction occurs; cations are attracted to this electrode

Oxidation — the loss of electrons, increase in oxidation state, or gain of oxygen

Reduction — the gain of electrons, decrease in oxidation state, or loss of oxygen

Redox reaction — a chemical reaction in which both oxidation and reduction occur simultaneously

Inert electrode — an electrode that does not participate in the reaction, such as graphite or platinum, used to conduct electricity into and out of the electrolyte

Core concepts

Understanding electrolysis fundamentals

Electrolysis requires three essential components: a source of direct current (DC), an electrolyte, and two electrodes. When an ionic compound melts or dissolves in water, the ions become free to move. Applying an electric current causes positive ions (cations) to migrate toward the cathode, while negative ions (anions) move toward the anode.

At the cathode, cations gain electrons in a reduction reaction. At the anode, anions lose electrons in an oxidation reaction. The mnemonic OIL RIG (Oxidation Is Loss, Reduction Is Gain) helps remember electron transfer, while PANCHO (Positive Anode, Negative Cathode, in electrolysis with HOoks connected) reminds students which electrode is which.

Electrolysis of molten ionic compounds

When a molten ionic compound undergoes electrolysis, only two ions are present — the metal cation and the non-metal anion. The products are predictable:

• At the cathode: the metal cation is reduced to form the metal element
• At the anode: the non-metal anion is oxidized to form the non-metal element

For molten lead(II) bromide, PbBr₂:

Cathode reaction: Pb²⁺ + 2e⁻ → Pb

Anode reaction: 2Br⁻ → Br₂ + 2e⁻

Molten lead forms at the cathode as a silvery liquid, while bromine vapour (brown fumes) is released at the anode. This type of electrolysis is straightforward because water is absent, eliminating competing reactions.

Electrolysis of aqueous solutions

Aqueous electrolysis is more complex because water molecules partially dissociate into H⁺ and OH⁻ ions, creating competing ions at each electrode. The discharge rules determine which ion is preferentially discharged:

At the cathode (reduction occurs):

1. If the metal is less reactive than hydrogen (copper, silver, gold), the metal ions are reduced to metal atoms
2. If the metal is more reactive than hydrogen (sodium, calcium, magnesium, aluminium, zinc, iron), hydrogen gas is produced from H⁺ ions instead

At the anode (oxidation occurs):

1. If a halide ion (Cl⁻, Br⁻, I⁻) is present in reasonable concentration, the halogen is produced
2. If no halide is present, or if sulfate or nitrate ions are the anions, oxygen gas is produced from OH⁻ ions

For aqueous copper(II) sulfate with inert electrodes:

Cathode: Cu²⁺ + 2e⁻ → Cu (copper is less reactive than hydrogen)

Anode: 4OH⁻ → O₂ + 2H₂O + 4e⁻ (sulfate ions remain in solution; hydroxide ions oxidize instead)

Copper metal deposits on the cathode (pink-brown solid), while oxygen gas bubbles form at the anode.

For aqueous sodium chloride (concentrated):

Cathode: 2H⁺ + 2e⁻ → H₂ (sodium is more reactive than hydrogen)

Anode: 2Cl⁻ → Cl₂ + 2e⁻ (chloride ions present in sufficient concentration)

Hydrogen gas evolves at the cathode, and chlorine gas (greenish-yellow) at the anode.

Writing half-equations for electrode reactions

Half-equations show the electron transfer at each electrode. CIE IGCSE Chemistry examiners require precise formatting:

1. Write the ion and the product
2. Balance atoms other than oxygen and hydrogen first
3. Balance oxygen by adding H₂O molecules
4. Balance hydrogen by adding H⁺ ions
5. Balance charge by adding electrons (e⁻)
6. Ensure electrons appear on the left for reduction (cathode) and right for oxidation (anode)

For the discharge of hydroxide ions at the anode:

Step 1: OH⁻ → O₂

Step 2: 4OH⁻ → O₂ (balance oxygen atoms)

Step 3: 4OH⁻ → O₂ + 2H₂O (balance hydrogen atoms)

Step 4: 4OH⁻ → O₂ + 2H₂O + 4e⁻ (balance charge: left side has 4− charge, right side needs 4− from electrons)

Industrial applications of electrolysis

Extraction of aluminium: Aluminium oxide (bauxite ore, purified) is dissolved in molten cryolite to lower the melting point from 2050°C to approximately 950°C. Electrolysis occurs with graphite electrodes:

• Cathode: Al³⁺ + 3e⁻ → Al (molten aluminium sinks to the bottom and is tapped off)
• Anode: 2O²⁻ → O₂ + 4e⁻ (oxygen reacts with the graphite anode, forming CO₂, so anodes require regular replacement)

This process consumes enormous amounts of electricity, making aluminium expensive to produce despite being the most abundant metal in Earth's crust.

Electroplating: A thin layer of one metal is deposited onto another object to improve appearance or prevent corrosion. The object to be plated becomes the cathode, while the anode is made of the plating metal. The electrolyte contains ions of the plating metal.

For silver plating:

• Cathode (object): Ag⁺ + e⁻ → Ag
• Anode (silver bar): Ag → Ag⁺ + e⁻

The anode dissolves at the same rate silver deposits on the cathode, maintaining constant electrolyte concentration.

Purification of copper: Impure copper is the anode, pure copper is the cathode, and copper(II) sulfate solution is the electrolyte:

• Cathode: Cu²⁺ + 2e⁻ → Cu (pure copper deposits)
• Anode: Cu → Cu²⁺ + 2e⁻ (impure copper dissolves)

Impurities fall to the bottom as anode sludge, while pure copper transfers from anode to cathode. This produces the high-purity copper (99.99%) required for electrical wiring.

Electrochemical cells and batteries

Electrochemical cells convert chemical energy into electrical energy — the reverse of electrolysis. A simple cell consists of two different metals (electrodes) placed in an electrolyte. Electrons flow through an external circuit from the more reactive metal (negative terminal) to the less reactive metal (positive terminal).

In a zinc-copper cell with dilute sulfuric acid:

At the zinc electrode: Zn → Zn²⁺ + 2e⁻ (oxidation; zinc dissolves)

At the copper electrode: 2H⁺ + 2e⁻ → H₂ (reduction; hydrogen bubbles form)

Zinc is more reactive than copper, so it preferentially loses electrons, creating a potential difference. The cell produces electricity until the zinc is consumed or the electrolyte is exhausted.

Rechargeable batteries (secondary cells) like the lead-acid car battery undergo reversible reactions. During discharge, chemical energy converts to electrical energy. During charging, electrical energy reverses the reactions, regenerating the reactants.

Non-rechargeable batteries (primary cells) like alkaline batteries contain reactions that cannot be easily reversed. Once the reactants are depleted, the battery is discarded.

Reactivity series and predicting redox reactions

The reactivity series ranks metals by their tendency to lose electrons and form positive ions. More reactive metals are stronger reducing agents (lose electrons more easily), while less reactive metals are weaker reducing agents.

Reactivity series (most to least reactive):

Potassium > Sodium > Calcium > Magnesium > Aluminium > Carbon > Zinc > Iron > Hydrogen > Copper > Silver > Gold

A more reactive metal can displace a less reactive metal from its compound in solution:

Zn(s) + CuSO₄(aq) → ZnSO₄(aq) + Cu(s)

Oxidation: Zn → Zn²⁺ + 2e⁻ (zinc loses electrons)

Reduction: Cu²⁺ + 2e⁻ → Cu (copper ions gain electrons)

Zinc is oxidized (loses electrons), while copper ions are reduced (gain electrons), making this a redox reaction. The blue colour of copper sulfate solution fades as Cu²⁺ ions are removed, and brown copper metal deposits on the zinc.

Worked examples

Example 1: Predict the products when aqueous potassium iodide undergoes electrolysis using inert graphite electrodes. Write the half-equations for the reactions at each electrode. [4 marks]

Solution:

At the cathode: Hydrogen gas is produced

Half-equation: 2H⁺ + 2e⁻ → H₂ [1 mark]

(Potassium is more reactive than hydrogen, so H⁺ ions are discharged instead of K⁺)

At the anode: Iodine is produced

Half-equation: 2I⁻ → I₂ + 2e⁻ [1 mark]

(Iodide is a halide ion, so it is discharged in preference to hydroxide ions)

[1 mark for each correct product, 1 mark for each correct half-equation with state symbols not required but electrons essential]

Example 2: A student electrolyses molten zinc chloride. Describe what the student would observe at each electrode and explain the observations in terms of the reactions occurring. [6 marks]

Solution:

At the cathode: A grey metallic solid (zinc metal) forms/deposits [1 mark]

Zinc ions are reduced: Zn²⁺ + 2e⁻ → Zn [1 mark]

Zinc ions gain electrons [1 mark]

At the anode: A greenish-yellow gas (chlorine) is produced [1 mark]

Chloride ions are oxidized: 2Cl⁻ → Cl₂ + 2e⁻ [1 mark]

Chloride ions lose electrons [1 mark]

Example 3: Copper can be purified by electrolysis. The impure copper is the anode and pure copper is the cathode. The electrolyte is copper(II) sulfate solution.

(a) Write the half-equation for the reaction at the anode. [1 mark]

(b) Explain why the concentration of copper(II) sulfate solution remains constant during this process. [2 marks]

Solution:

(a) Cu → Cu²⁺ + 2e⁻ [1 mark]

(b) Copper dissolves from the anode, producing Cu²⁺ ions [1 mark], while Cu²⁺ ions are removed from solution at the cathode, where they gain electrons to form copper atoms [1 mark]. The rate of production equals the rate of removal, so concentration stays constant.

Common mistakes and how to avoid them

• Confusing anode and cathode charges: Students often think the anode is negative because it attracts anions. In electrolysis, the anode is positive — it attracts negative ions because opposite charges attract. Remember: Anode is Adding electrons (oxidation).

• Writing incomplete half-equations: Omitting electrons or failing to balance charges invalidates the half-equation. Always count the total charge on each side and add electrons to balance. At the cathode, electrons appear on the left (reactants); at the anode, electrons appear on the right (products).

• Forgetting aqueous electrolysis rules: When water is present, H⁺ and OH⁻ ions compete with the dissolved ions. A common error is predicting sodium metal at the cathode during aqueous sodium chloride electrolysis. Sodium is too reactive; hydrogen is produced instead.

• Misidentifying oxidation and reduction: Oxidation is electron loss, not oxygen gain (though oxygen gain is one type of oxidation). Use OIL RIG consistently. In the reaction Cu²⁺ + Zn → Cu + Zn²⁺, zinc loses electrons (oxidized) and copper ions gain electrons (reduced).

• Confusing electrochemical cells with electrolysis cells: Electrolysis uses electrical energy to drive a non-spontaneous reaction. Electrochemical cells (batteries) produce electrical energy from spontaneous chemical reactions. The more reactive metal is negative in an electrochemical cell but would be attracted to the cathode (negative electrode) if it were an electrolyte in electrolysis.

• Stating products without considering electrode material: If the anode is copper (not inert), it participates in the reaction by dissolving: Cu → Cu²⁺ + 2e⁻. This occurs in copper purification and electroplating. Always check whether electrodes are inert (graphite, platinum) or reactive.

Exam technique for Electrochemistry

• Command word 'Predict': State the product formed at each electrode, often followed by 'Write a half-equation'. Award marks depend on correctly identifying products based on the discharge rules. Show working by comparing reactivity or identifying ion types (halide, sulfate, etc.).

• Command word 'Explain': Link observations to underlying chemistry. For example, 'Explain why the anode needs regular replacement in aluminium extraction' requires stating that oxygen is produced, which reacts with the graphite anode to form carbon dioxide, gradually consuming the electrode.

• Half-equation questions typically allocate 1-2 marks: One mark for correct species and electrons, a second mark for balanced atoms and charges. State symbols may earn a mark in some questions but are not always required — check the question carefully.

• Industrial applications attract 3-6 mark questions: Examiners test understanding of real-world electrolysis (aluminium extraction, electroplating, copper purification). Structure answers with: electrode material, electrolyte, half-equations, observations, and economic/practical reasons for the process. Extended response questions may ask about advantages, disadvantages, or environmental considerations.

Quick revision summary

Electrolysis decomposes ionic compounds using electricity. Cations migrate to the negative cathode (reduction occurs), while anions migrate to the positive anode (oxidation occurs). In molten electrolysis, products are simply the metal and non-metal. Aqueous electrolysis involves competing ions; discharge depends on reactivity and ion type. Half-equations show electron transfer, balancing atoms and charge. Industrial applications include aluminium extraction, electroplating, and copper purification. Electrochemical cells convert chemical energy to electrical energy, with the more reactive metal forming the negative terminal.