Describe and explain the displacement reactions of halogens with other halide ions

The Periodic Table – Group VII (Halogen) Properties

Learning Objective

Describe and explain the displacement reactions of halogens with other halide ions, including the redox nature of the processes and the observable colour changes.

1. Halogen Group Overview

• Group VII (Group 17) elements: fluorine (F), chlorine (Cl), bromine (Br), iodine (I) and astatine (At).
• All exist as di‑atomic molecules (e.g., _Cl2, _Br2) and are highly reactive non‑metals.
• Reactivity decreases down the group: F > Cl > Br > I > At.
• A more reactive halogen can oxidise the halide ion of a less reactive halogen, producing the elemental form of the latter and its own halide ion.

2. Physical Properties of the Halogens

3. Reactivity Series (Most to Least Reactive)

4. Redox Terminology

• Oxidising agent: the more reactive halogen (e.g., Cl₂, Br₂) – it gains electrons (is reduced).
• Reducing agent: the halide ion of the less reactive halogen (e.g., Br⁻, I⁻) – it loses electrons (is oxidised).
• Standard reduction potentials increase down the series, which explains the observed displacement order.

5. General Form of a Halogen Displacement Reaction

When a more reactive halogen X₂ is added to a solution containing the halide ion of a less reactive halogen Y⁻:

Overall equation

X₂ + 2 Y⁻ → 2 X⁻ + Y₂

Half‑equations

• Reduction (oxidising agent): X₂ + 2e⁻ → 2 X⁻
• Oxidation (reducing agent): 2 Y⁻ → Y₂ + 2e⁻

6. Specific Displacement Reactions (with Half‑Equations)

1. Fluorine displaces all other halogens
   • F₂ + 2Cl⁻ → 2F⁻ + Cl₂
     Reduction: F₂ + 2e⁻ → 2F⁻  Oxidation: 2Cl⁻ → Cl₂ + 2e⁻
   • F₂ + 2Br⁻ → 2F⁻ + Br₂
     Reduction: F₂ + 2e⁻ → 2F⁻  Oxidation: 2Br⁻ → Br₂ + 2e⁻
   • F₂ + 2I⁻ → 2F⁻ + I₂
     Reduction: F₂ + 2e⁻ → 2F⁻  Oxidation: 2I⁻ → I₂ + 2e⁻
2. Chlorine displaces bromide and iodide
   • Cl₂ + 2Br⁻ → 2Cl⁻ + Br₂
     Reduction: Cl₂ + 2e⁻ → 2Cl⁻  Oxidation: 2Br⁻ → Br₂ + 2e⁻
   • Cl₂ + 2I⁻ → 2Cl⁻ + I₂
     Reduction: Cl₂ + 2e⁻ → 2Cl⁻  Oxidation: 2I⁻ → I₂ + 2e⁻
3. Bromine displaces iodide
   • Br₂ + 2I⁻ → 2Br⁻ + I₂
     Reduction: Br₂ + 2e⁻ → 2Br⁻  Oxidation: 2I⁻ → I₂ + 2e⁻
4. Astatine (theoretical)
   • At₂ + 2I⁻ → 2At⁻ + I₂ (predicted, follows the same pattern)

7. Observable Colour Changes (Diagnostic Tool)

• Cl₂ (colourless) added to KBr → brown/red‑brown Br₂ appears.
• Br₂ (red‑brown) added to KI → violet I₂ vapour is liberated.
• F₂ reactions are usually not observed visually because both F₂ and the displaced halogen are colourless gases.

8. Factors Influencing the Reaction

• Concentration of the more reactive halogen: Higher concentration drives the equilibrium to the right (Le Chatelier’s principle).
• Temperature: Most displacements occur readily at room temperature; gentle heating can speed slower reactions (e.g., Br₂ + 2I⁻).
• Physical state of the halogen: Gases must dissolve in water to react; liquids (Br₂) can be added directly.

9. Practical Example – Displacement of Bromide by Chlorine

1. Add a few drops of chlorine water to a potassium bromide solution.
2. Cl₂ dissolves, giving colourless Cl₂(aq).
3. Reaction: Cl₂ + 2Br⁻ → 2Cl⁻ + Br₂
4. Observation: a brown/red‑brown colour appears, confirming the formation of bromine.

10. Summary

Halogen displacement reactions are classic redox processes that illustrate the relative reactivity of Group VII elements. The more reactive halogen acts as the oxidising agent, gaining electrons to form its halide ion, while the less reactive halide ion is oxidised to the elemental halogen. Recognising the colour changes and writing both overall and half‑equations enables accurate prediction of laboratory outcomes and aligns with the Cambridge IGCSE Chemistry 0620 syllabus requirements.