Introduction to Group 7 Halogens Revision Notes | GCSE AQA Higher Chemistry

1. Meet the Halogens - The Reactive Seven

Picture the periodic table as a neighbourhood. Group 7 is the street where all the residents have something in common - they're all missing just one thing to feel complete. These elements are called the halogens, and they're some of the most reactive non-metals you'll encounter.

The halogens include fluorine, chlorine, bromine, iodine and astatine. You'll focus on chlorine, bromine and iodine for GCSE. What makes them special? They all have seven electrons in their outer shell, making them desperate to gain one more electron to achieve a stable arrangement.

2. The Halogen Trend - Going Down the Group

As you move down Group 7 from chlorine to iodine, the halogens show a clear pattern. Think of it like a staircase - each step down brings predictable changes in colour, state and reactivity.

Chlorine is a pale green gas at room temperature. Bromine is a red-brown liquid that gives off toxic vapour. Iodine is a dark grey solid that sublimes to purple vapour when heated. Notice the pattern? The colour gets darker and the melting point increases as you go down the group.

But here's the crucial part for your exams: reactivity decreases down the group. Chlorine is more reactive than bromine, which is more reactive than iodine. This happens because the outer electrons get further from the nucleus as atoms get bigger, making it harder to attract that extra electron they need.

3. Halogen Reactions - The Electron Grabbers

Halogens are electron-hungry. When they react, they're trying to gain that one electron they need to fill their outer shell. This determines the type of compounds they form and explains their behaviour in reactions.

When halogens react with metals, they form ionic compounds called halides. The metal loses electrons (becomes a positive ion) and the halogen gains electrons (becomes a negative ion). For example: $Na + Cl_2 \rightarrow 2NaCl$

When halogens react with non-metals, they form covalent compounds by sharing electrons. Think hydrogen chloride: $H_2 + Cl_2 \rightarrow 2HCl$. The atoms share electrons rather than transferring them completely.

Worked example: Writing Halogen Reaction Equations

Write a balanced equation for the reaction between magnesium and bromine.

1. Magnesium metal: $Mg$ Bromine molecules: $Br_2$ — Remember bromine exists as diatomic molecules
2. Magnesium bromide: $MgBr_2$ — Mg loses 2 electrons, each Br gains 1 electron
3. $Mg + Br_2 \rightarrow MgBr_2$ — This shows what reacts and what forms
4. Left: 1 Mg, 2 Br Right: 1 Mg, 2 Br — Already balanced - same number of each atom on both sides

Answer: $Mg + Br_2 \rightarrow MgBr_2$

4. Displacement Reactions - The Halogen Hierarchy

Here's where the reactivity trend becomes really useful. More reactive halogens can displace less reactive ones from their compounds. It's like a chemical version of 'survival of the fittest' - the more reactive halogen wins.

Chlorine can displace bromine from bromide solutions: $Cl_2 + 2KBr \rightarrow 2KCl + Br_2$. The solution changes from colourless to orange-brown as bromine forms. Similarly, chlorine displaces iodine from iodides, and bromine displaces iodine.

But the reverse doesn't work. Iodine cannot displace chlorine from chlorides because iodine is less reactive. Think of it as a one-way street - reactions only go from more reactive to less reactive.

5. Exam Success - Mastering Halogen Questions

Halogen questions often test your understanding of the reactivity trend and your ability to predict reaction products. The key is remembering that everything links back to those seven outer electrons.

When describing halogen properties, always mention the outer shell electrons. When explaining reactivity trends, link it to atomic size and electron attraction. When predicting reactions, use the reactivity order: chlorine > bromine > iodine.