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DSE Chemistry Recap – Redox | Megan (2026)

Redox

DSE Chemistry

Megan (2026)

For use on Friday, 4 September 2026

01

Redox Fundamentals

Definitions and Redox Equations

A piece of magnesium ribbon is added to an aqueous solution of copper(II) sulphate. A reddish-brown solid is formed and the blue colour of the solution gradually becomes less intense.

Mg(s) + CuSO4(aq) ⟶ MgSO4(aq) + Cu(s)
  1. Define oxidation and reduction in terms of the transfer of electrons.
  2. Explain why oxidation and reduction must always occur together in a redox reaction.
  3. State the change in oxidation number of magnesium during the reaction.
  4. Write the oxidation half equation for magnesium.
  5. Write the reduction half equation for copper(II) ions.
  6. Hence write the overall ionic equation for the redox reaction.
  7. Identify the oxidizing agent and the reducing agent in the reaction.
10 marks
02

Electrochemistry

A Simple Chemical Cell and a Porous Pot

The diagram shows a simple chemical cell. A copper electrode and copper(II) sulphate solution are placed inside a porous pot. The porous pot is placed in zinc sulphate solution containing a zinc electrode. The electrodes are connected through a voltmeter.

V electron flow Zinc electrode Copper electrode ZnSO4(aq) CuSO4(aq) Porous pot

A zinc–copper chemical cell with the two electrolytes separated by a porous pot.

  1. State which electrode is the negative electrode and which is the positive electrode.
  2. Write the half equation for the reaction occurring at the zinc electrode.
  3. Write the half equation for the reaction occurring at the copper electrode.
  4. State the direction of electron flow through the external circuit.
  5. Write the overall ionic equation for the cell reaction.
  6. Describe one observable change at the zinc electrode and one observable change at the copper electrode.
  7. Explain the purpose of the porous pot.
  8. State the energy conversion occurring when the cell operates.
12 marks
03

Inert Electrodes

A Chemical Cell Using Graphite Electrodes

A chemical cell is constructed using two graphite electrodes. Half-cell X contains acidified iron(III) ions and iron(II) ions. Half-cell Y contains iodine and iodide ions. The two half-cells are joined by a salt bridge.

V Half-cell X Half-cell Y Fe3+(aq), Fe2+(aq) I2(aq), I−(aq) Salt bridge Graphite Graphite

The graphite electrodes provide conducting surfaces but do not take part chemically in the reaction.

When the cell operates, iron(III) ions oxidize iodide ions.

  1. Explain why graphite electrodes are required in both half-cells.
  2. Write the oxidation half equation involving iodide ions.
  3. Write the reduction half equation involving iron(III) ions.
  4. Write the overall ionic equation for the cell reaction.
  5. Identify the oxidizing agent and the reducing agent.
  6. State which half-cell contains the negative electrode.
  7. State the direction of electron flow in the external circuit.
  8. Explain one function of the salt bridge.
11 marks
04

Sulphuric Acid

Dilute and Concentrated Sulphuric Acid

Two experiments involving sulphuric acid are carried out separately.

Experiment Reactants Conditions
1 Zinc and sulphuric acid Dilute sulphuric acid at room temperature
2 Copper and sulphuric acid Hot concentrated sulphuric acid
  1. State two observations when zinc reacts with dilute sulphuric acid.
  2. Write the ionic equation for the reaction between zinc and dilute sulphuric acid.
  3. Identify the species reduced in Experiment 1.
  4. Explain why copper does not normally react with dilute sulphuric acid.
  5. State two observations when copper is heated with concentrated sulphuric acid.
  6. Write the full chemical equation for the reaction between copper and hot concentrated sulphuric acid.
  7. State the change in oxidation number of sulphur in Experiment 2.
  8. Identify the oxidizing agent in Experiment 2.
  9. Explain why the reaction with concentrated sulphuric acid is a redox reaction.
12 marks
05

Cell Comparison

Chemical Cells and Electrolytic Cells

Chemical cells and electrolytic cells both involve oxidation and reduction reactions at electrodes. However, their operation and energy changes are different.

Feature Chemical cell Electrolytic cell
Nature of reaction Spontaneous Non-spontaneous
Energy conversion 1 2
Sign of anode 3 4
Sign of cathode 5 6
Reaction at anode 7
Reaction at cathode 8
  1. Complete entries 1–8 in the table.
  2. State the direction in which electrons travel through the external circuit, using the terms anode and cathode.
  3. Explain why a chemical cell does not require an external power supply during normal operation.
  4. Explain why an electrolytic cell requires an external direct-current power supply.
  5. A student says, “The anode is always negative.” Explain why this statement is incorrect.
  6. State one similarity, other than the use of electrodes, between a chemical cell and an electrolytic cell.
14 marks

Question 1

  1. Oxidation is the loss of electrons. Reduction is the gain of electrons.
  2. The electrons lost by one species must be gained by another species. Therefore, oxidation and reduction occur simultaneously.
  3. The oxidation number of magnesium changes from 0 to +2. Magnesium is oxidized.
  4. Mg(s) ⟶ Mg2+(aq) + 2e−
  5. Cu2+(aq) + 2e− ⟶ Cu(s)
  6. Mg(s) + Cu2+(aq) ⟶ Mg2+(aq) + Cu(s)
  7. Oxidizing agent: Cu2+(aq). It accepts electrons and is reduced.
    Reducing agent: Mg(s). It donates electrons and is oxidized.
OIL RIG: Oxidation Is Loss of electrons; Reduction Is Gain of electrons. An oxidizing agent is itself reduced, while a reducing agent is itself oxidized.

Question 2

  1. The zinc electrode is the negative electrode, and the copper electrode is the positive electrode.
  2. Zn(s) ⟶ Zn2+(aq) + 2e−
  3. Cu2+(aq) + 2e− ⟶ Cu(s)
  4. Electrons flow from the zinc electrode to the copper electrode through the external circuit.
  5. Zn(s) + Cu2+(aq) ⟶ Zn2+(aq) + Cu(s)
  6. At the zinc electrode, the zinc electrode becomes thinner or loses mass because zinc atoms form zinc ions.

    At the copper electrode, reddish-brown copper is deposited and the electrode gains mass. The blue colour of the copper(II) sulphate solution also becomes less intense as copper(II) ions are removed.
  7. The porous pot separates the two electrolytes but allows ions to move between them. This completes the internal circuit, maintains electrical neutrality and prevents the rapid direct mixing of the solutions.
  8. Chemical energy is converted into electrical energy.
In a chemical cell, oxidation occurs at the negative electrode and reduction occurs at the positive electrode. Electrons leave the negative electrode and travel through the external circuit to the positive electrode.

Question 3

  1. All the reacting species in the half-cells are in solution, so there is no conducting solid reactant that can act as an electrode. Graphite provides a conducting surface for electron transfer without taking part chemically in the reaction.
  2. 2I−(aq) ⟶ I2(aq) + 2e−
  3. Fe3+(aq) + e− ⟶ Fe2+(aq)
  4. The iron half equation must be multiplied by two before combining the equations:
    2Fe3+(aq) + 2I−(aq) ⟶ 2Fe2+(aq) + I2(aq)
  5. Oxidizing agent: Fe3+(aq). It accepts electrons and is reduced to Fe2+(aq).

    Reducing agent: I−(aq). It donates electrons and is oxidized to iodine.
  6. Half-cell Y, containing I2 and I−, contains the negative electrode. Oxidation and the release of electrons occur in this half-cell.
  7. Electrons flow from the graphite electrode in half-cell Y to the graphite electrode in half-cell X.
  8. The salt bridge allows ions to move between the half-cells, completes the internal circuit and maintains electrical neutrality. It also reduces direct mixing of the two solutions.
Inert electrodes such as graphite or platinum are used when no conducting solid reactant is present. They conduct electrons and provide a surface for the electrode reaction, but they are not consumed in the overall reaction.

Question 4

  1. Any two:
    • Effervescence or bubbles of a colourless gas are observed.
    • The zinc gradually dissolves or becomes smaller.
    • The reaction mixture may become warmer.
  2. Zn(s) + 2H+(aq) ⟶ Zn2+(aq) + H2(g)
  3. H+(aq) ions are reduced. They gain electrons to form hydrogen gas:
    2H+(aq) + 2e− ⟶ H2(g)
  4. Copper is below hydrogen in the reactivity series. It cannot displace hydrogen from a dilute non-oxidizing acid. Therefore, copper does not normally react with dilute sulphuric acid.
  5. Any two:
    • The copper gradually dissolves.
    • A blue solution containing copper(II) ions forms.
    • A colourless gas with a choking or pungent smell is produced.
  6. Cu(s) + 2H2SO4(conc) ⟶ CuSO4(aq) + SO2(g) + 2H2O(l)
  7. The oxidation number of sulphur changes from +6 in H2SO4 to +4 in SO2. Sulphur is reduced.
  8. Concentrated sulphuric acid is the oxidizing agent. It causes copper to be oxidized and is itself reduced to sulphur dioxide.
  9. Copper changes from oxidation number 0 to +2, so copper is oxidized. Sulphur changes from +6 to +4, so sulphur is reduced. As oxidation and reduction occur together, the reaction is a redox reaction.
Dilute sulphuric acid: shows the usual acidic properties and reacts with metals above hydrogen to produce hydrogen gas.

Hot concentrated sulphuric acid: is an oxidizing agent and can oxidize copper, producing sulphur dioxide instead of hydrogen.

Question 5

  1. Entry Correct answer
    1 Chemical energy to electrical energy
    2 Electrical energy to chemical energy
    3 Negative
    4 Positive
    5 Positive
    6 Negative
    7 Oxidation
    8 Reduction
  2. Electrons travel from the anode to the cathode through the external circuit.
  3. A chemical cell uses a spontaneous redox reaction. The reaction releases chemical energy and converts it into electrical energy, so an external power supply is not required.
  4. The reaction in an electrolytic cell is non-spontaneous. An external direct-current power supply provides electrical energy and forces the redox reaction to occur.
  5. The statement is incorrect because the anode is negative in a chemical cell but positive in an electrolytic cell. However, oxidation always occurs at the anode in both cells.
  6. Any one:
    • Both involve redox reactions.
    • Oxidation occurs at the anode in both cells.
    • Reduction occurs at the cathode in both cells.
    • Electrons travel through the external circuit from the anode to the cathode in both cells.
    • Both require an electrolyte containing mobile ions.
AnOx: oxidation always occurs at the anode.
RedCat: reduction always occurs at the cathode.

The signs of the electrodes depend on the type of cell, but the locations of oxidation and reduction do not change.
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