Group 7

At room temperature, the halogens exist in all three states of matter:

Trends in Physical Properties

As you go down Group 7:

All halogens exist as diatomic molecules ($\text{F}_2$, $\text{Cl}_2$, $\text{Br}_2$, $\text{I}_2$) — two atoms bonded by a single covalent bond.

All halogens have seven electrons in their outermost shell:

To achieve a stable full outer shell, halogens gain one electron to form −1 ions (halide ions):

$\text{Cl} + e^- \rightarrow \text{Cl}^-$

The halide ions are: fluoride ($\text{F}^-$), chloride ($\text{Cl}^-$), bromide ($\text{Br}^-$), iodide ($\text{I}^-$).

Reactivity decreases as you go down Group 7:

$\text{F} > \text{Cl} > \text{Br} > \text{I}$

Why Does Reactivity Decrease?

As you go down Group 7:

1. More electron shells → the atom is larger
2. The outer shell is further from the nucleus
3. There is more shielding from inner electrons
4. The nucleus has a weaker attraction for incoming electrons
5. It is harder to gain the extra electron needed
6. Therefore, the element is less reactive

Key contrast: Group 1 metals get MORE reactive going down (easier to lose electrons). Group 7 non-metals get LESS reactive going down (harder to gain electrons).

4.1 Reaction with Metals

Halogens react with metals to form metal halides (ionic compounds):

$2\text{Na} + \text{Cl}_2 \rightarrow 2\text{NaCl}$

$2\text{Fe} + 3\text{Cl}_2 \rightarrow 2\text{FeCl}_3$

The metal atom loses electrons (oxidation) and the halogen atom gains electrons (reduction). This is a redox reaction.

4.2 Reaction with Hydrogen

Halogens react with hydrogen to form hydrogen halides:

$\text{H}_2 + \text{Cl}_2 \rightarrow 2\text{HCl}$

Hydrogen halides dissolve in water to form acids:

• HCl → hydrochloric acid
• HBr → hydrobromic acid
• HI → hydroiodic acid

A more reactive halogen can displace a less reactive halogen from a solution of its salt. This is one of the most important reactions to know for GCSE.

The Rule

A halogen higher in Group 7 will displace a halogen lower in Group 7 from its compound.

Experimental Evidence

Example Equations

Chlorine displacing bromine: $\text{Cl}_2 + 2\text{KBr} \rightarrow 2\text{KCl} + \text{Br}_2$

Chlorine displacing iodine: $\text{Cl}_2 + 2\text{KI} \rightarrow 2\text{KCl} + \text{I}_2$

Bromine displacing iodine: $\text{Br}_2 + 2\text{KI} \rightarrow 2\text{KBr} + \text{I}_2$

Why Displacement Works

In the reaction $\text{Cl}_2 + 2\text{KBr} \rightarrow 2\text{KCl} + \text{Br}_2$:

• Chlorine is more reactive than bromine
• Chlorine atoms gain electrons more easily than bromine atoms
• Chlorine takes the electrons from bromide ions
• Chlorine is reduced (gains electrons): $\text{Cl}_2 + 2e^- \rightarrow 2\text{Cl}^-$
• Bromine is oxidised (loses electrons): $2\text{Br}^- \rightarrow \text{Br}_2 + 2e^-$

• Colour changes in displacement reactions are commonly tested — learn the colours!
• When explaining reactivity, mention distance from nucleus, shielding, and ability to gain an electron
• In displacement equations, the halogen (X₂) reacts with the metal halide (MX) — don't forget the diatomic formula
• Contrast Group 1 and Group 7 trends — opposite patterns with opposite reasons

• Halogens have 7 outer electrons and form −1 ions
• Physical state goes from gas → liquid → solid down the group
• Reactivity decreases down the group (harder to attract electrons)
• More reactive halogens displace less reactive ones from solutions
• Key colours: Cl₂ = green, Br₂ = red-brown, I₂ = grey solid / purple vapour