Group 7: The Halogens

The halogens — fluorine (F₂), chlorine (Cl₂), bromine (Br₂), iodine (I₂), and astatine (At₂) — are reactive non-metals in Group 7 of the Periodic Table. At AQA A-Level (3.2.3), you need to understand trends in physical properties, oxidising ability, displacement reactions, reactions with NaOH, hydrogen halides, and the silver nitrate test.

---

Physical Properties

Halogen	Formula	Colour	State at room temperature	Boiling point / °C
Fluorine	F₂	Pale yellow	Gas	−188
Chlorine	Cl₂	Yellow-green	Gas	−34
Bromine	Br₂	Red-brown	Liquid	59
Iodine	I₂	Dark grey / purple vapour	Solid	184

Trend in Boiling Points

Boiling points increase down Group 7 because:

• All halogens exist as diatomic molecules (X₂) with weak London (dispersion) forces between them.
• Going down the group, the molecules have more electrons, making them more polarisable.
• This increases the strength of the London forces, so more energy is needed to separate the molecules.

---

Electronegativity and Oxidising Ability

Electronegativity

Electronegativity decreases down Group 7: F (4.0) > Cl (3.2) > Br (3.0) > I (2.7).

The atomic radius increases and shielding increases, so the nucleus is less effective at attracting bonding electrons.

Oxidising Ability

The halogens are oxidising agents — they gain electrons: X₂ + 2e⁻ → 2X⁻

Oxidising ability decreases down the group because:

1. Atomic radius increases, so the incoming electron is further from the nucleus.
2. Shielding by inner electrons increases.
3. Both factors mean the nucleus attracts the incoming electron less strongly.
4. The electron affinity becomes less exothermic (less energy released on gaining an electron).

Key Definition: An oxidising agent is a substance that gains electrons (is itself reduced) and causes another substance to be oxidised.

---

Displacement Reactions

A more reactive (stronger oxidising) halogen will displace a less reactive halide ion from solution. This demonstrates the trend in oxidising ability.

	KCl(aq)	KBr(aq)	KI(aq)
Cl₂(aq)	No reaction	Orange/brown solution (Br₂ formed)	Brown solution (I₂ formed)
Br₂(aq)	No reaction	No reaction	Brown solution (I₂ formed)
I₂(aq)	No reaction	No reaction	No reaction

Equations

Chlorine displaces bromide: Cl₂(aq) + 2KBr(aq) → 2KCl(aq) + Br₂(aq)

Ionic equation: Cl₂(aq) + 2Br⁻(aq) → 2Cl⁻(aq) + Br₂(aq)

• Chlorine is reduced (0 → −1): oxidising agent.
• Bromide is oxidised (−1 → 0).

Chlorine displaces iodide: Cl₂(aq) + 2KI(aq) → 2KCl(aq) + I₂(aq)

Ionic equation: Cl₂(aq) + 2I⁻(aq) → 2Cl⁻(aq) + I₂(aq)

Bromine displaces iodide: Br₂(aq) + 2KI(aq) → 2KBr(aq) + I₂(aq)

Ionic equation: Br₂(aq) + 2I⁻(aq) → 2Br⁻(aq) + I₂(aq)

Exam Tip: In displacement reactions, always identify which species is oxidised and which is reduced. The halogen molecule is always the oxidising agent (it gains electrons).

---

Reactions of Halogens with NaOH

Chlorine with Cold Dilute NaOH (Disproportionation)

Cl₂(g) + 2NaOH(aq) → NaCl(aq) + NaClO(aq) + H₂O(l)

• Chlorine is simultaneously oxidised (0 → +1 in NaClO) and reduced (0 → −1 in NaCl).
• This is a disproportionation reaction.
• NaClO (sodium chlorate(I), bleach) is used in water treatment and as a disinfectant.

Chlorine with Hot Concentrated NaOH

3Cl₂(g) + 6NaOH(aq) → 5NaCl(aq) + NaClO₃(aq) + 3H₂O(l)