11.2 The Chemical Properties of the Halogen Elements and the Hydrogen Halides

Subject: Cambridge International AS Level Chemistry (9701) Topic: Group 17

2026 📌 Syllabus Objectives

By the end of these notes, you will be able to:

  1. Describe the relative reactivity of the halogen elements as oxidising agents
  2. Describe the reactions of the halogens with hydrogen and explain their relative reactivity in these reactions
  3. Describe the relative thermal stabilities of the hydrogen halides and explain these in terms of bond strengths

📖 Main Content

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Objective 1: The Halogens as Oxidising Agents

What is an oxidising agent?

An oxidising agent is a substance that takes electrons away from another substance. When it does this, the other substance loses electrons and becomes oxidised (its oxidation number goes up). The oxidising agent itself gains electrons and becomes reduced (its oxidation number goes down).

Think of it like this: the halogen is a "electron thief" — it steals electrons from other atoms.

How do halogens act as oxidising agents?

Halogens react with metals by accepting (taking) one electron from the metal atom. When they gain this electron, each halogen atom becomes a negatively charged ion with a 1− charge (called a halide ion).

Example — Calcium reacting with Chlorine:

Ca (s) + Cl2 (g)→Ca2+(Cl−)2 (s)\text{Ca (s) + Cl}_2\text{ (g)} \rightarrow \text{Ca}^{2+}\text{(Cl}^-\text{)}_2\text{ (s)}
  • The calcium loses electrons → it is oxidised (its oxidation number increases)
  • The chlorine gains electrons → it is reduced (its oxidation number decreases)
  • Because chlorine causes the calcium to be oxidised, chlorine is the oxidising agent

How does oxidising power change down Group 17?

The oxidising power of the halogens decreases going down the group. This means:

  • Fluorine (F₂) — the strongest oxidising agent
  • Chlorine (Cl₂) — strong oxidising agent
  • Bromine (Br₂) — moderate oxidising agent
  • Iodine (I₂) — the weakest oxidising agent

In other words: the halogens get less reactive as you go down the group.

Why does oxidising power decrease down the group?

This is explained by electronegativity — which means how strongly an atom attracts electrons towards itself. An atom with high electronegativity pulls electrons strongly; an atom with low electronegativity pulls electrons weakly.

Here is the step-by-step explanation:

Step 1 — Atomic radius increases going down the group. As you go from fluorine to iodine, each element has more electron shells. This means the atom gets bigger. The outer shell (where the new electron would be accepted) gets further and further away from the nucleus (the positively charged centre of the atom).

Step 2 — Shielding increases going down the group. The inner electron shells sit between the nucleus and the outer shell. They act like a shield, blocking the attractive pull of the positive nucleus. The more inner shells there are, the more shielding there is. This is called the shielding effect.

Step 3 — It becomes harder to attract an incoming electron. Because of the larger atomic radius and greater shielding, an incoming electron feels much less attraction towards the nucleus. The halogen therefore has a weaker "pull" on electrons — its electronegativity is lower.

Step 4 — Oxidising power decreases. Since oxidising power depends on the ability to attract and accept electrons, and this ability decreases going down the group, the oxidising power decreases from F to I.

📊 Electronegativity values (approximate):

HalogenElectronegativity
F~4.0 (highest)
Cl~3.0
Br~2.8
I~2.5 (lowest)

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