Cambridge IGCSE • Year 11 • Chemistry
Electrochemistry
Electrolysis, ionic movement, electrode products and electrochemical applications.
Chapter 4
Verified Curriculum Topic
What is Electrochemistry?
Electrolysis, ionic movement, electrode products and electrochemical applications.
Electrochemistry matters because it links chemical ideas, reactions, and reasoning patterns that recur throughout the syllabus. At Year 11 level, students are often expected to define terms accurately, explain processes clearly, and connect theory to reactions, observations, or applications.
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Summary
The One Thing
Electrolysis uses electrical energy to drive a non-spontaneous reaction in an ionic substance, causing ions to move and discharge at electrodes. Products are determined by the ions present, the electrolyte conditions, and the electrode materials; oxidation occurs at the anode and reduction at the cathode.
Reactions, Processes and Experiments
| What happens | Equation or process | What you observe | Type |
|---|---|---|---|
| Positive copper ions gain electrons at the cathode and form copper. | Cu2+ + 2e- -> Cu. | Copper is deposited at the cathode. | Reduction half-equation |
| Negative chloride ions lose electrons at the anode and form chlorine. | 2Cl- -> Cl2 + 2e-. | Chlorine gas is produced at the anode. | Oxidation half-equation |
| Solid ionic compounds are heated until molten so their ions can move. | An ionic compound conducts when molten because its ions can move freely; in the solid state, the ions are fixed in a lattice. | The molten compound conducts electricity; the solid compound does not. | Ionic movement and conductivity |
| Molten lead(II) bromide is decomposed into lead and bromine. | PbBr2(l) -> Pb(l) + Br2(g). | Lead forms at the cathode and bromine forms at the anode. | Electrolysis of a molten ionic compound |
| Molten aluminium oxide is electrolysed to extract aluminium. | Al3+ + 3e- -> Al. | Aluminium is produced at the cathode. | Reduction during metal extraction |
| In aqueous solutions, hydrogen ions may be discharged instead of metal ions. | Hydrogen ions gain electrons at the cathode when the metal is more reactive than hydrogen. | Hydrogen gas is produced at the cathode. | Cathode competition in an aqueous electrolyte |
| Halide ions are discharged at an inert anode in aqueous solution. | Halide ions such as chloride, bromide, and iodide are discharged to form the corresponding halogen. | A halogen is produced at the anode. | Anode discharge of halide ions |
| Hydroxide ions are discharged at an inert anode when halide ions are not preferentially discharged. | Hydroxide ions are discharged and oxygen is produced. | Oxygen gas is produced at the anode. | Anode discharge in an aqueous electrolyte |
| Concentrated aqueous sodium chloride is electrolysed. | 2NaCl + 2H2O -> H2 + Cl2 + 2NaOH. | Hydrogen forms at the cathode, chlorine forms at the anode, and sodium hydroxide remains in solution. | Electrolysis of concentrated aqueous sodium chloride |
| Dilute aqueous sulfuric acid is electrolysed using inert electrodes. | 2H2O -> 2H2 + O2. | Hydrogen forms at the cathode and oxygen forms at the anode. | Electrolysis of dilute aqueous sulfuric acid |
| Copper(II) sulfate solution is electrolysed using copper electrodes. | Copper is transferred from the anode to the cathode. | The anode loses mass and the cathode gains mass. | Electrorefining |
| An object is coated with copper by making it the cathode. | The object is the cathode, a copper anode supplies copper ions, and a copper salt solution is the electrolyte. | A layer of copper forms on the object. | Electroplating |
| The object is prepared before electroplating. | The object is cleaned first so that the coating adheres well. | The coating adheres more effectively to the object. | Preparation for electroplating |
| The current is controlled during electroplating. | The current is controlled to produce an even layer. | A more even metal coating is formed. | Control of electroplating |
| Aluminium oxide is electrolysed industrially at high temperature. | Aluminium oxide must be molten; cryolite lowers its melting point and reduces energy consumption. | Aluminium is produced while less energy is required than would otherwise be necessary. | Industrial extraction of aluminium |
| Carbon anodes react with oxygen during aluminium extraction. | Carbon anodes react with oxygen and are gradually consumed, producing carbon dioxide. | The carbon anodes become smaller and carbon dioxide is produced. | Electrode reaction in aluminium extraction |
| The amount of substance discharged is related to the electrical charge passed. | Q = It, where Q is charge in coulombs, I is current in amperes, and t is time in seconds. | Increasing the current or time increases the amount of substance discharged, provided other conditions remain constant. | Quantitative electrolysis |
| Electrolysis extracts very reactive metals. | Electrolysis is used to extract metals such as aluminium that cannot be obtained by reducing their compounds with carbon. | A reactive metal is produced from its compound. | Metal extraction |
| Electrorefining purifies a metal. | The impure metal is used as the anode and pure metal is deposited at the cathode. | Pure metal forms at the cathode while the anode dissolves. | Electrorefining |
| An electrolytic cell uses electricity to cause a chemical reaction. | An electrolytic cell uses electricity to cause a reaction, while a simple cell produces electricity from a reaction. | Electrical energy is converted into chemical change in an electrolytic cell. | Electrochemical cell |
Key Terms
- Electrolysis: The chemical decomposition of an ionic compound in the molten state or dissolved in water by passing an electric current through it.
- Electrolyte: A molten ionic compound or aqueous solution containing mobile ions that conducts electricity and is decomposed during electrolysis.
- Electrode: A conductor placed in an electrolyte through which electric current enters or leaves the electrolyte.
- Cathode: The negative electrode in electrolysis, where positive ions gain electrons and reduction occurs.
- Anode: The positive electrode in electrolysis, where negative ions lose electrons and oxidation occurs.
- Cation: A positively charged ion that moves towards the cathode.
- Anion: A negatively charged ion that moves towards the anode.
- Oxidation: Loss of electrons, occurring at the anode during electrolysis.
- Reduction: Gain of electrons, occurring at the cathode during electrolysis.
- Inert electrode: An electrode, commonly made of platinum or carbon, that conducts electricity but does not usually react with the electrolyte.
- Ionic movement: Cations move through the electrolyte towards the cathode, while anions move towards the anode; electrons move through the external wires.
- Molten electrolyte: An ionic compound heated until it melts, allowing its ions to move freely without water being present.
- Aqueous electrolyte: An ionic compound dissolved in water; the solution contains ions from the compound and hydrogen ions and hydroxide ions from water.
- Discharge: The process in which an ion gains or loses electrons at an electrode to become an atom or molecule.
- Electroplating: Coating an object with a thin layer of metal by making the object the cathode in an electrolytic cell.
- Electrorefining: Purifying a metal by using the impure metal as the anode and depositing pure metal at the cathode.
- Electrochemical cell: A system in which chemical reactions and electrical energy are linked; an electrolytic cell uses electricity to cause a reaction, while a simple cell produces electricity from a reaction.
Easily Confused
- Cathode and anode: The cathode is where reduction and electron gain occur; the anode is where oxidation and electron loss occur.
- Cation and anion: Cations are positive and move towards the cathode; anions are negative and move towards the anode.
- Molten and aqueous electrolytes: A molten electrolyte contains ions from the ionic compound only, whereas an aqueous electrolyte also contains hydrogen ions and hydroxide ions from water.
- Electrolysis and a simple cell: An electrolytic cell uses electricity to cause a reaction; a simple cell produces electricity from a chemical reaction.
- Electroplating and electrorefining: Electroplating coats an object with a metal, whereas electrorefining purifies a metal by depositing pure metal at the cathode.
- Inert and carbon anodes: An inert electrode does not usually react with the electrolyte, whereas carbon anodes in aluminium extraction react with oxygen and are gradually consumed.
- Ion movement and electron movement: Ions move through the electrolyte, while electrons move through the external wires.
- Halide discharge and hydroxide discharge: In aqueous solutions, halide ions such as chloride, bromide, and iodide may form halogens at the anode; otherwise hydroxide ions form oxygen.
- Metal-ion discharge and hydrogen-ion discharge: In aqueous solutions, hydrogen may form at the cathode instead of the metal when the metal is more reactive than hydrogen.
- Electrolysis and spontaneous reactions: Electrolysis requires electrical energy to drive a non-spontaneous reaction, whereas a spontaneous reaction can produce electrical energy in a simple cell.
What Gets Asked
- Identify the products at the cathode and anode. Marks are lost by reversing the electrodes or confusing oxidation with reduction.
- Write half-equations for electrode reactions. The equation must conserve both atoms and charge, including the correct number of electrons.
- Explain why a solid ionic compound does not conduct electricity but a molten or dissolved compound does. The required distinction is fixed ions in a lattice versus mobile ions.
- Predict products for molten and aqueous electrolytes. Water competes with dissolved ions in aqueous solutions, so the products may differ from those of the corresponding molten compound.
- Describe concentrated aqueous sodium chloride and dilute aqueous sulfuric acid. The specific products must be stated: hydrogen and chlorine with sodium hydroxide remaining for concentrated sodium chloride; hydrogen and oxygen for dilute sulfuric acid.
- Calculate charge or compare the amount discharged using Q = It. The units must be consistent: coulombs, amperes, and seconds; increasing current or time increases the amount discharged when other conditions remain constant.
Flashcards
Quick quiz
What is electrolysis?
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Sign up free — save & unlock everythingLearning objectives
- 4.1Define electrolysis as the decomposition of an ionic compound, molten or in aqueous solution, by an electric current.
- 4.2Describe the electrode products for the electrolysis of molten lead(II) bromide.
- 4.3Predict the products of electrolysis of aqueous solutions, including the effect of ion concentration and reactivity.extended
- 4.4Describe electroplating and explain one reason for electroplating an object, such as preventing corrosion.
- 4.5Explain, using half-equations, the reactions occurring at the anode and cathode during electrolysis.extended
- 4.6Describe the process of purifying copper by electrolysis using a copper anode and cathode.extended
Practice questions
Q1. During the electrolysis of molten lead(II) bromide, which product forms at the cathode (negative electrode)?1 mark · core
- A. Bromine gas
- B. Lead metal
- C. Oxygen gas
- D. Hydrogen gas
Answer: B
- • 1 mark for selecting B
Positive lead ions (Pb2+) are attracted to and gain electrons at the negative cathode, forming lead metal. Negative bromide ions move to the positive anode and are oxidised to bromine gas.
Q2. Write half-equations for the reactions occurring at the cathode and anode during the electrolysis of molten lead(II) bromide.4 marks · extended
Answer: Cathode (reduction): Pb2+ + 2e− → Pb. Anode (oxidation): 2Br− → Br2 + 2e−.
- • 1 mark: correct cathode half-equation with correct species (Pb2+ and Pb)
- • 1 mark: correct balancing of charge/electrons at the cathode
- • 1 mark: correct anode half-equation with correct species (Br− and Br2)
- • 1 mark: correct balancing of charge/electrons at the anode
Q3. Give one reason why an iron object might be electroplated with chromium.2 marks · core
Answer: To prevent the iron from corroding/rusting, and/or to improve its appearance by giving it a shiny surface.
- • 1 mark: prevents corrosion/rusting
- • 1 mark: improves appearance (shiny, decorative finish) — accept either valid reason for full credit across two reasons
Q4. Predict the products formed at each electrode during the electrolysis of concentrated aqueous sodium chloride solution, and explain your answer for the cathode product.3 marks · extended
Answer: Chlorine gas is produced at the anode and hydrogen gas is produced at the cathode. At the cathode, both H+ ions (from water) and Na+ ions are present, but H+ ions are discharged in preference to the more reactive Na+ ions, producing hydrogen gas.
- • 1 mark: chlorine gas at the anode
- • 1 mark: hydrogen gas at the cathode
- • 1 mark: correct explanation that H+ is preferentially discharged over the more reactive Na+ ion
Key ideas to master
- Learn the precise terms, laws, and reaction patterns associated with Electrochemistry.
- Understand why each step or change happens instead of memorising the result only.
- Practise writing balanced equations, comparisons, or structured explanations where relevant.
- Revise common exceptions, observations, and applications that examiners often test.
Common exam prompts
- Define the main idea in Electrochemistry using correct chemical terminology.
- Write or interpret the reactions, observations, or comparisons that belong to this topic.
- Explain why a process happens, not just what happens.
- Summarise the high-yield facts and exceptions examiners often choose from this chapter.
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What is Electrochemistry in Cambridge IGCSE Year 11 Chemistry?
Electrolysis, ionic movement, electrode products and electrochemical applications.
How should I study Electrochemistry effectively?
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