Identify redox reactions as reactions involving gain and loss of oxygen

6.4 Redox – Core and Supplement (Cambridge IGCSE 0620)

1. What is a Redox Reaction?

A redox (reduction‑oxidation) reaction involves the transfer of electrons between species. In the IGCSE syllabus students are required to identify redox reactions in two ways:

• Oxygen‑transfer (traditional) approach – an element that gains oxygen is reduced, the element that supplies the oxygen is oxidised.
• Oxidation‑number (electron‑transfer) approach – an increase in oxidation number indicates oxidation (loss of electrons); a decrease indicates reduction (gain of electrons).

2. Key Definitions (Core + Supplement)

Oxidation‑Number Rules (Official IGCSE Supplement)

3. Systematic Identification of Redox Reactions

Step‑by‑Step Checklist

1. Write the balanced molecular (or net‑ionic) equation.
2. Oxygen‑transfer test – Does any element gain or lose oxygen atoms?
   
   • Yes → the element that gains O is reduced (oxidising agent); the element that loses O is oxidised (reducing agent). Proceed to 5.
   • No → go to step 3.
3. Oxidation‑number test – Assign oxidation numbers using the rules above.
4. Compare oxidation numbers of each element before and after the reaction:
   • Increase → oxidation (loss of e⁻).
   • Decrease → reduction (gain of e⁻).
5. Identify the oxidising and reducing agents:
   • Oxidising agent = species that is reduced.
   • Reducing agent = species that is oxidised.
6. (Optional, Supplement) Write the half‑equations to confirm electron balance.

Combined Flowchart (simplified)

1. Is there a change in the number of O atoms attached to any element?
   
   • Yes → use the oxygen‑transfer test (step 2).
   • No → calculate oxidation numbers (step 3).
2. From the test, decide which species is oxidised and which is reduced.
3. Write half‑equations (Supplement) to verify electron transfer.

4. Writing Half‑Equations (Supplement)

Half‑equations make the electron flow explicit. Follow the checklist below:

1. Separate the overall reaction into an oxidation half‑equation and a reduction half‑equation.
2. Balance all atoms except H and O.
3. Balance O by adding H₂O.
4. Balance H by adding H⁺ (acidic medium) **or** OH⁻ (basic medium).
5. Balance the charge by adding electrons to the more positive side.
6. If necessary, multiply the half‑equations so that the number of electrons cancels, then add them together.

Worked Example – CuO + CO → Cu + CO₂

1. Overall balanced equation (molecular):
   \(\displaystyle \text{CuO} + \text{CO} \;\longrightarrow\; \text{Cu} + \text{CO}_2\)
2. Oxygen‑transfer test – CuO loses its O atom (Cu is reduced); CO gains O (CO is oxidised).
3. Oxidation‑number test:
   • Cu in CuO: +2 → Cu in Cu: 0 (decrease → reduction).
   • C in CO: +2 → C in CO₂: +4 (increase → oxidation). O remains –2 throughout.
4. Half‑equations (acidic medium):
   • Reduction: \(\displaystyle \text{Cu}^{2+} + 2e^- \rightarrow \text{Cu}\)
   • Oxidation: \(\displaystyle \text{CO} + \text{H}_2\text{O} \rightarrow \text{CO}_2 + 2\text{H}^+ + 2e^-\)
   Adding the two halves gives the overall equation (the \(\text{H}^+\) and \(\text{H}_2\text{O}\) cancel in the net‑ionic form).

5. Common Types of Redox Reactions (Core)

6. Electrolysis (Supplement)

When an aqueous electrolyte is electrolysed, reduction occurs at the cathode and oxidation at the anode.

Typical example – Aqueous copper(II) sulfate

• Cathode (reduction): \(\displaystyle \text{Cu}^{2+} + 2e^- \rightarrow \text{Cu}(s)\)
• Anode (oxidation): \(\displaystyle 2\text{H}_2\text{O} \rightarrow O_2 + 4\text{H}^+ + 4e^-\)
• Overall (multiply the cathode equation by 2 to cancel electrons):
  \(\displaystyle 2\text{Cu}^{2+} + 2\text{H}_2\text{O} \rightarrow 2\text{Cu}(s) + O_2 + 4\text{H}^+\)

Key points for the exam:

• Identify which electrode hosts reduction (cathode) and which hosts oxidation (anode).
• Distinguish between inert electrodes (e.g., Pt, C) and metal electrodes that can themselves be oxidised or reduced.

7. Redox in Metal Extraction (Supplement)

Two classic IGCSE examples are given below.

Iron – Blast furnace

• Overall reaction: \(\displaystyle Fe_2O_3 + 3CO \rightarrow 2Fe + 3CO_2\)
• Oxidation half‑equation (carbon/CO): \(\displaystyle 3CO \rightarrow 3CO_2 + 6e^-\)
• Reduction half‑equation (iron oxide): \(\displaystyle Fe_2O_3 + 6e^- \rightarrow 2Fe + 3O^{2-}\) (often written as \(\displaystyle Fe_2O_3 + 6e^- \rightarrow 2Fe + 3O^{2-}\) and the O²⁻ combines with the CO to give CO₂.)

Aluminium – Hall–Héroult process (electrolysis)

• Overall reaction (molten cryolite mixture): \(\displaystyle 2Al_2O_3 \xrightarrow{\text{electricity}} 4Al + 3O_2\)
• Cathode (reduction): \(\displaystyle Al^{3+} + 3e^- \rightarrow Al\)
• Anode (oxidation): \(\displaystyle 2O^{2-} \rightarrow O_2 + 4e^-\)

8. Detailed Worked Example – Copper(II) Oxide Reduced by Hydrogen

Reaction: \(\displaystyle \text{CuO} + \text{H}_2 \rightarrow \text{Cu} + \text{H}_2\text{O}\)

1. Oxygen‑transfer test – CuO loses its O atom (Cu reduced); H₂ gains O (oxidised).
2. Oxidation‑number test:
   • Cu: +2 → 0 (decrease → reduction).
   • H: 0 → +1 (increase → oxidation).
3. Half‑equations (acidic medium):
   • Reduction: \(\displaystyle \text{Cu}^{2+} + 2e^- \rightarrow \text{Cu}\)
   • Oxidation: \(\displaystyle \text{H}_2 \rightarrow 2\text{H}^+ + 2e^-\)
   Adding gives the overall balanced equation above.
4. Agents – Oxidising agent = CuO (it is reduced); Reducing agent = H₂ (it is oxidised).

9. Practice Questions

1. Identify the oxidising and reducing agents in the thermite reaction:
   \(\displaystyle \text{Fe}_2\text{O}_3 + 2\text{Al} \rightarrow 2\text{Fe} + \text{Al}_2\text{O}_3\)
2. Write the balanced overall equation and the two half‑equations for the reduction of copper(II) oxide by carbon monoxide.
3. State whether the following reaction is a redox process and justify your answer:
   \(\displaystyle \text{NaCl} + \text{AgNO}_3 \rightarrow \text{AgCl} + \text{NaNO}_3\)
4. In acidic solution, electrolyse aqueous copper(II) sulfate. Write the half‑equations occurring at the cathode and anode and identify which electrode is the site of reduction.
5. Explain why potassium permanganate (KMnO₄) acts as a strong oxidising agent in acidic medium, giving the reduction half‑equation for \(\text{MnO}_4^-\).

10. Answers & Mark Scheme

11. Summary Checklist for Identifying Redox Reactions

• Write a balanced molecular or net‑ionic equation.
• Check for oxygen transfer – if present, use the oxygen‑transfer test.
• If no oxygen change, assign oxidation numbers using the official rules.
• Identify any increase (oxidation) or decrease (reduction) in oxidation numbers.
• Label the oxidising agent (species reduced) and the reducing agent (species oxidised).
• Optional: write oxidation and reduction half‑equations to confirm electron balance.
• Confirm that the overall reaction is balanced for both mass and charge.