AQA Group 7 Halogens Overview
Halogens Overview
- Halogens are diatomic non-metals
- Fluorine (F₂): Pale yellow gas
- Chlorine (Cl₂): Pale green gas
- Bromine (Br₂): Brown-orange liquid
- Iodine (I₂): Grey solid/purple gas
Trends in Group 7
- Boiling Points: Increase down the group due to larger van der Waals forces
- Electronegativity: Decreases down the group, increased shielding = larger atoms attract electrons less effectively
- Atomic Radius: Increases down the group, due to increased shielding.
Oxidising Ability/Reactivity of the Halogens
- Decreases down the group
- More reactive halogens can displace less reactive halide ions in redox reactions
- Cl₂ (Strongest Oxidising Agent) can displace Br⁻ and I⁻ ions
- Br₂ can only displace I⁻ ions
- I₂ cannot displace any halide
Reducing Ability of Halide Ions
- Reducing Power Increases down the group.
- More shells = larger ionic radius = weaker attraction between nucleus and outer electrons
- Halide ions down the group are more easily oxidised
- I⁻ (Strongest Reducing Agent)
- Cl⁻ (Weakest Reducing Agent)
Tests for Halides
Reaction of Halide Salts with Conc H₂SO₄
The increase of reducing power of halide ions down the group is demonstrated by their reactions with H₂SO₄ and how much S is reduced.
Step 1 (Acid-base Reaction):
- NaX (s) + H₂SO₄ (l) → NaHSO₄ (s) + HX (g)
Step 2 (Redox Reaction/s):
- Unbalanced equation: H⁺ + X⁻ + H₂SO₄ → X₂ + (S compound) + H₂O
Fluoride and Chloride Ions (F⁻/Cl⁻)
F⁻ & Cl⁻ ions do not have enough reducing power to reduce S in H₂SO₄, so no redox only acid-base reaction.
- NaF + H₂SO₄ → NaHSO₄ + HF
- NaCl + H₂SO₄ → NaHSO₄ + HCl
Observation: White steamy HF/HCl Fumes
Bromide Ions (Br⁻)
Br⁻ ions are stronger reducing agents than F⁻ & Cl⁻ ions, so they have a redox reaction.
Step 1: NaBr + H₂SO₄ → NaHSO₄ + HBr
Step 2: 2H⁺ + 2Br⁻ + H₂SO₄ → SO₂ + Br₂ + 2H₂O
- Oxidation State (S): +6 → +4
Overall equation: 2NaBr + 3H₂SO₄ → 2NaHSO₄ + Br₂ + SO₂ + 2H₂O
Observations: White steamy HBr fumes, Orange Br₂ fumes and SO₂ (colourless, acidic gas)
Iodide Ions (I⁻)
I⁻ ions are the strongest reducing agents, so they have multiple redox reactions.
Step 1: NaI + H₂SO₄ → NaHSO₄ + HI
Step 2: 2H⁺ + 2I⁻ + H₂SO₄ → SO₂ + I₂ + 2H₂O
- Oxidation State (S): +6 → +4
Step 3: 6H⁺ + 6I⁻ + H₂SO₄ → S + 3I₂ + 4H₂O
- Oxidation State (S): +6 → 0
Step 4: 8H⁺ + 8I⁻ + H₂SO₄ → H₂S + 4I₂ + 4H₂O
- Oxidation State (S): +6 → -2
Overall equation: CBA
Observations: White steamy HI fumes, Black solid I₂ & Purple I₂ fumes, SO₂ (colourless, acidic gas), Solid yellow S and H₂S egg smell
Identifying halide ions with acidified AgNO₃
Precipitate Colors:
- Fluoride (F⁻): Colourless (AgF is soluble)
- Chloride (Cl⁻): White
- Bromide (Br⁻): Cream
- Iodide (I⁻): Yellow
Use ammonia to further distinguish
- AgCl: Dissolves in dilute NH₃
- AgBr: Dissolves in concentrated NH₃
- AgI: Insoluble in concentrated NH₃
- Equation: AgX (s) + 2NH₃ → [Ag(NH₃)₂]⁺ + X⁻
Uses of Chlorine (Cl₂) and Chlorate (ClO⁻)
Chlorine reacts with water in a disproportionation reaction to form Chloric acid and Hydrochloric acid
Water Sterilisation
Chlorate ions (ClO-) kill bacteria, preventing diseases like cholera
- As HClO reacts with bacteria, equilibrium shifts to replace it
Used to disinfect pools and drinking water
- Advantages: Destroys microorganisms, long-lasting, reduces algae growth
- Disadvantages: Chlorine is toxic, potential cancer risk from reaction with organic compounds
- Benefits generally outweigh risks
In sunlight, Cl₂ in water can react to make HCl, so only little is used (& Cl₂ is toxic anyway)
Bleach Production
Disproportionation Reaction: Cl₂ reacts with cold, dilute, aqueous NaOH to form NaClO (bleach)
- Cl₂ + 2NaOH → NaCl + NaClO + H₂O
- Cl₂ is oxidised and reduced
Uses include water treatment, paper bleaching, and cleaning