The halogens

Characteristic physical properties of halogens

Halogens exist as diatomic molecules: Cl₂, Br₂, I₂
They are simple molecular substances
Reactivity decreases down Group 7:

Cl₂ > Br₂ > I₂

This is shown by displacement reactions with halide ions.

A more reactive halogen will displace a less reactive halide ion from solution:

Boiling point trend:
Boiling points increase down the group:

Cl₂ < Br₂ < I₂

Reason:

Down the group, molecules have more electrons and larger electron clouds
This causes stronger induced dipole–dipole interactions (London forces) and instaneous dipoles between molecules
More energy is needed to overcome these forces, so boiling points increase.

Halogens

• Halogens have an outer shell electron configuration of ns² np⁵.

• They have 7 outer electrons, so they tend to gain 1 electron in redox reactions to achieve a full outer shell (noble gas configuration).

Trend in reactivity of halogens (Cl₂, Br₂, I₂)

Reactivity decreases down Group 7:
Cl₂ > Br₂ > I₂
This is shown by displacement reactions with halide ions.
A more reactive halogen will displace a less reactive halide ion from solution

Explanation of trend in reactivity of halogens

• Atomic radius increases down the group
  The outer shell is further from the nucleus.

• Electron shielding increases
  More inner electron shells block attraction from the nucleus.

• Attraction between nucleus and incoming electron decreases
  The nucleus has less effective pull on an added electron; lower nuclear attraction

• It becomes less easy to gain an electron and thus less reactive

Disproportionation

Disproportionation is a redox reaction in which the same element is both oxidised and reduced in the same reaction.

Chlorine with water (water treatment)

Cl₂ + H₂O ⇌ HCl + HClO

• Chlorine (0) is:

  • reduced to Cl⁻ in HCl (oxidation state −1)

  • oxidised to Cl in HClO (oxidation state +1)

HClO is a disinfectant used in water treatment (kills bacteria).

Colours of halogens in solutions

Halogen	Aqueous solution	Organic solvent (cyclohexane)
Chlorine, Cl₂	Pale green	Pale green
Bromine, Br₂	Orange	Orange
Iodine, I₂	Brown	Purple/Violet

Bleach production

Cl₂ + 2NaOH → NaCl + NaClO + H₂O

• Chlorine (0) is:

  • reduced to Cl⁻ in NaCl (−1)

  • oxidised to Cl in NaClO (+1)

NaClO is bleach (used for cleaning and disinfection).

Used with cold, dilute aqueous sodium hydroxide

Chlorine in water treatment: benefits vs risks

Benefit

Chlorine is added to drinking water to kill harmful bacteria and pathogens.
It forms hypochlorous acid (HClO) in water, which is a strong disinfectant:

Risk:
Chlorine gas is toxic and dangerous if released during handling or storage.
Chlorine can react with organic matter in water to form chlorinated hydrocarbons, some of which may be harmful or potentially carcinogenic.

Precipitation of halide ions (X-) with aqueous silver ions (Ag⁺)

Supplied via Silver nitrate AgNO3 (acidify nitric acid first)

Ag⁺ + Cl⁻ → AgCl(s) (white precipitate)

• Ag⁺ + Br⁻ → AgBr(s) (cream precipitate)

• Ag⁺ + I⁻ → AgI(s) (yellow precipitate)

Reaction silver halide with aqueous ammonia

• Silver chloride (AgCl):
  AgCl(s) dissolves in dilute NH₃(aq)
  AgCl(s) + 2NH₃(aq) → [Ag(NH₃)₂]⁺(aq) + Cl⁻(aq)

• Silver bromide (AgBr):
  Dissolves only in concentrated NH₃(aq) (slowly)

• Silver iodide (AgI):
  Does not dissolve in ammonia