The Halogens

Section 1 of 10Boiling point trend

Trend

• Trend boiling point increases down the group
• Halogens are non-metal diatomic molecules, such as F₂, with simple molecular structures.

Explanation

• Explanation weak London forces act between the molecules
  • (instantaneous dipole–induced dipole)
• An instantaneous imbalance of electrons gives a dipole that induces a dipole in a neighbouring molecule.
• More electrons in a larger molecule give stronger London forces.
• More energy is needed to separate the molecules, so the boiling point rises.

Section 2 of 10Displacement reactions and observations

• Displacement
  • a more reactive halogen displaces a less reactive halogen from an aqueous solution of its halide
  • Reactivity increases up the group, so chlorine is most reactive and iodine least of these three.

Halogen	Colour in aqueous solution	Colour in organic solvent
Chlorine	Very pale green (often colourless)	Colourless
Bromine	Yellow	Yellow
Iodine	Brown (sometimes black solid)	Purple

• Chlorine with bromide
  • 2KBr(aq) + Cl₂(aq) → 2KCl(aq) + Br₂(aq)
  • Bromine is oxidised from −1 to 0 and chlorine is reduced from 0 to −1; potassium is unchanged.
  • The ionic equation is 2Br⁻(aq) + Cl₂(aq) → 2Cl⁻(aq) + Br₂(aq), with K⁺ as the spectator ion.

Section 3 of 10Reactivity trend and halogens as oxidising agents

Trend

• Trend oxidising power decreases down the group
• The atomic radius increases.
• There are more inner shells, so shielding increases.
• There is less nuclear attraction for an electron from another species.

Redox reactions

• Redox reactions each halogen has 7 outer electrons (s² p⁵) and gains one electron to form a 1− ion
• Ca(s) + Cl₂(g) → CaCl₂(s), made of Ca²⁺ and 2Cl⁻.
• The halogen oxidises the metal, whose oxidation number rises, and is itself reduced as its oxidation number falls.

Section 4 of 10Disproportionation of chlorine

Disproportionation of chlorine

Disproportionation

Disproportionation a reaction in which the same species is both oxidised and reduced

• Chlorine with water and with cold dilute aqueous alkali are examples.

With water

• Cl₂ + H₂O → HCl + HClO
• Chlorine is reduced to HCl and oxidised to HClO (chloric(I) acid).
• HClO sterilises water by killing bacteria and dissociates to H⁺ and ClO⁻, which also sterilises.
• Universal indicator turns red because of the acid, then colourless as HClO bleaches it.

With cold dilute aqueous NaOH

• Cl₂ + 2OH⁻ → Cl⁻ + ClO⁻ + H₂O
• Chlorine is oxidised from 0 to +1 in ClO⁻.
• Chlorine is reduced from 0 to −1 in Cl⁻.
• The mixture of NaCl and NaClO (sodium chlorate(I)) is used as bleach and a disinfectant.

Section 5 of 10Water treatment with chlorine

Method

• Method chlorine added to a water supply forms HClO, which dissociates to ClO⁻, both sterilising agents
• It is used for drinking water and swimming pools.

Risk and benefit

• Risk and benefit the health benefit of killing bacteria outweighs the risks
• The risks are toxic effects and the possible formation of chlorinated hydrocarbons.

Section 6 of 10Testing for halide ions

1

Dissolve the sample and add dilute nitric acid, which prevents carbonate giving a false positive.

2

Add silver nitrate solution dropwise.

3

Note the colour of any precipitate: Ag⁺(aq) + X⁻(aq) → AgX(s).

4

Add dilute then concentrated ammonia to confirm which halide it is.

Halide	Precipitate	Effect of ammonia
Chloride	White AgCl	Dissolves in dilute ammonia
Bromide	Cream AgBr	Insoluble in dilute, dissolves in concentrated
Iodide	Yellow AgI	Insoluble in both

Section 7 of 10Order of tests for anions

Order of tests for anions

Carbonate first, because BaCO₃ and Ag₂CO₃ are insoluble and would interfere with later tests.

Sulfate second, because Ag₂SO₄ is insoluble and would give a false positive in the halide test.

Halide last, once carbonates and sulfates are removed.

• Why it works
  • the HCl in the sulfate test removes remaining carbonate, and the HNO₃ in the halide test removes remaining carbonate or sulfate

Section 8 of 10Tests for carbonate, sulfate and ammonium ions

Carbonate

• Carbonate add dilute hydrochloric acid and pass any gas into limewater
• Effervescence occursCO₃²⁻(aq) + 2H⁺(aq) → CO₂(g) + H₂O(l).
• Limewater turns milky as CO₂ + Ca(OH)₂ → CaCO₃ + H₂O.
• The delivery tube bung must be attached as soon as the solutions are mixed so no gas escapes.

Sulfate

• Sulfate acidify with dilute hydrochloric acid, then add a few drops of aqueous barium chloride (or nitrate)
• A white precipitate of BaSO₄ forms: Ba²⁺(aq) + SO₄²⁻(aq) → BaSO₄(s).

Ammonium

• Ammonium warm with aqueous sodium hydroxide
• NH₄⁺(aq) + OH⁻(aq) → NH₃(g) + H₂O(l).
• Ammonia is very soluble, so bubbles are unlikely;
  • it is identified by its pungent smell or by turning red litmus paper blue.

Section 9 of 10Key terms

Term	Meaning
Displacement reaction	A more reactive halogen displaces a less reactive halogen from its halide solution
Disproportionation	A reaction in which the same species is both oxidised and reduced
Oxidising agent	A species that accepts electrons and is reduced
Chloric(I) acid	HClO, a sterilising agent formed when chlorine reacts with water
Limewater	Aqueous calcium hydroxide, used to test for carbon dioxide
London forces	Weak forces between molecules from instantaneous and induced dipoles

Section 10 of 10Exam technique

1

Saying reactivity of halogens increases down the group — it decreases

2

Describing disproportionation as simple oxidation — the same species is oxidised and reduced

3

Adding the silver nitrate test before removing carbonate and sulfate — do carbonate, sulfate then halide

4

Using hydrochloric acid to acidify the halide test — use nitric acid to avoid adding chloride