Colours and states of chlorine, bromine and iodine

· Group 17 elements = halogens; they exist as diatomic molecules, e.g. Cl₂, Br₂, I₂.
· Chlorine, Cl₂ = pale green gas at room temperature.
· Bromine, Br₂ = red-brown liquid with orange-brown vapour at room temperature.
· Iodine, I₂ = grey-black solid with purple vapour when heated.
· Colour trend down the group: halogens become darker from chlorine → bromine → iodine.
· State trend at room temperature: gas → liquid → solid down the group.

This image shows the visible colour differences between halogens. It helps students remember that chlorine is pale green, bromine is red-brown, and iodine is grey-black/purple vapour. Source

Volatility trend down Group 17

· Volatility = how easily a substance evaporates.
· A more volatile substance has a lower boiling point and evaporates more easily.
· Volatility decreases down Group 17: Cl₂ > Br₂ > I₂.
· Chlorine is the most volatile of these three because it is a gas at room temperature.
· Iodine is the least volatile because it is a solid with much stronger intermolecular forces.
· Melting point and boiling point increase down the group, so the elements become less volatile.

Explaining volatility using instantaneous dipole–induced dipole forces

· Halogens are simple molecular substances with weak intermolecular forces between molecules.
· The relevant intermolecular forces are instantaneous dipole–induced dipole forces.
· Down the group, molecules have more electrons and larger electron clouds.
· Larger electron clouds are more easily polarised, so stronger instantaneous dipoles form.
· Stronger instantaneous dipole–induced dipole forces mean more energy is needed to separate molecules.
· Therefore, boiling point increases and volatility decreases down the group.
· Key explanation chain: more electrons → stronger id-id forces → higher boiling point → lower volatility.

Bond strength trend in halogen molecules

· Halogen molecules contain a single covalent bond, e.g. Cl–Cl, Br–Br, I–I.
· Bond strength decreases down the group: Cl–Cl > Br–Br > I–I.
· Down the group, atoms become larger and the bonding pair is further from the nuclei.
· The attraction between the bonding pair of electrons and the two nuclei becomes weaker.
· Therefore, the X–X bond length increases and the X–X bond strength decreases down the group.
· Exam phrase: larger atomic radius → longer bond → weaker attraction to bonding pair → lower bond strength.
· Extra exam note: F–F is anomalously weak because very small fluorine atoms have strong lone pair–lone pair repulsion.

Bond enthalpy and bond strength

· Bond enthalpy = energy needed to break one mole of covalent bonds in the gaseous state.
· A higher bond enthalpy means a stronger bond.
· For the main CIE trend: Cl₂ has a stronger X–X bond than Br₂, and Br₂ has a stronger X–X bond than I₂.
· Bond enthalpy decreases because atomic radius increases and orbital overlap becomes less effective down the group.
· Do not confuse bond strength with intermolecular forces:
· Bond strength = strength of the covalent bond within X₂ molecules.
· Volatility = depends on intermolecular forces between X₂ molecules.

Common exam pitfalls

· Do not write that volatility decreases because covalent bonds get stronger; volatility depends on intermolecular forces, not the X–X covalent bond.
· Do not say iodine is purple solid; iodine is usually described as a grey-black solid that forms purple vapour.
· Do not confuse boiling point trend with bond strength trend: boiling point increases, but X–X bond strength generally decreases down the group.
· Use the phrase instantaneous dipole–induced dipole forces, not just “van der Waals forces”, when a full explanation is required.
· Always link volatility to number of electrons, polarizability, and strength of intermolecular forces.