1. Physical Trends Down Group 7
The halogens exist as diatomic covalent non-metal molecules (X₂):
- Fluorine (F₂): Pale yellow gas. Most reactive and most electronegative element.
- Chlorine (Cl₂): Pale green gas with a pungent, choking odour.
- Bromine (Br₂): Dense red-brown liquid giving off dense orange-brown vapour.
- Iodine (I₂): Shiny grey-black crystalline solid that sublimes on gentle warming to form a purple vapour.
Boiling points increase down the group: As molecular size increases, the electron count rises (18 → 34 → 70 → 106). Larger electron clouds produce stronger London dispersion forces between molecules, requiring more thermal energy to separate them.
Trends in Electronegativity and Boiling Point
| Halogen | F₂ | Cl₂ | Br₂ | I₂ |
|---|---|---|---|---|
| Electronegativity of the atom | 4.0 | 3.0 | 2.8 | 2.5 |
| Boiling point / °C | -188 | -34 | 59 | 184 |
- Electronegativity decreases down the group. The atoms get larger. The bonding pair of electrons is further from the nucleus and more shielded, so it is attracted less strongly.
- Boiling point increases down the group. Larger molecules have more electrons, so the van der Waals forces between them are stronger and need more energy to overcome.
2. Oxidising Ability and Displacement Reactions
An electron acceptor. Halogen atoms gain electrons to form 1⁻ halide ions: \[ \text{X}_2 + 2e^- \rightarrow 2\text{X}^- \]
Oxidising ability decreases down Group 7: F₂>Cl₂>Br₂>I₂. As atomic radius and electron shielding increase, incoming electrons experience weaker electrostatic attraction to the halogen nucleus.
Displacement Reactions
A more reactive halogen will displace a less reactive halide ion from aqueous solution:
- \(\text{Cl}_2(aq) + 2\text{Br}^-(aq) \rightarrow 2\text{Cl}^-(aq) + \text{Br}_2(aq)\) (Solution turns yellow-orange).
- \(\text{Cl}_2(aq) + 2\text{I}^-(aq) \rightarrow 2\text{Cl}^-(aq) + \text{I}_2(aq)\) (Solution turns brown).
- \(\text{Br}_2(aq) + 2\text{I}^-(aq) \rightarrow 2\text{Br}^-(aq) + \text{I}_2(aq)\) (Solution turns brown).
No reaction happens when a halogen is added to its own halide or to a halide above it in the group, for example Br₂ + Cl⁻ or I₂ + Br⁻.
3. Reducing Ability of Halide Ions with Concentrated Sulfuric Acid
Reducing ability increases down Group 7: \(\text{F}^- < \text{Cl}^- < \text{Br}^- < \text{I}^-\). Iodide ions have the largest ionic radius and greatest shielding, so their outer electrons are most easily lost.
| Solid Halide | Reaction Type | Observations | Products Formed |
|---|---|---|---|
| NaCl | Acid-base only (no redox) | Steamy fumes turning moist blue litmus paper red. | \(\text{HCl}(g) + \text{NaHSO}_4(s)\) |
| NaBr | Acid-base AND Redox | Steamy fumes (HBr), orange-brown fumes (Br₂), choking colourless gas (SO₂). | HBr,Br₂,SO₂ (Sulfur reduced from +6 to +4) |
| NaI | Strong Redox | Steamy fumes (HI), black solid / purple vapour (I₂), choking gas (SO₂), yellow solid sulfur (S), rotten egg smell gas (H₂S). | HI,I₂,SO₂,S,H₂S (Sulfur reduced from +6 down to -2) |
Equations
Sodium chloride (no redox; misty fumes of HCl):
- NaCl + H₂SO₄ → NaHSO₄ + HCl
Sodium bromide (misty fumes of HBr, then orange-brown fumes of Br₂):
- NaBr + H₂SO₄ → NaHSO₄ + HBr
- 2HBr + H₂SO₄ → Br₂ + SO₂ + 2H₂O. Sulfur goes from +6 to +4.
Sodium iodide (purple fumes, a black solid, a yellow solid and a rotten-egg smell):
- NaI + H₂SO₄ → NaHSO₄ + HI
- 2HI + H₂SO₄ → I₂ + SO₂ + 2H₂O. Sulfur goes from +6 to +4.
- 6HI + H₂SO₄ → 3I₂ + S + 4H₂O. Sulfur goes from +6 to 0 (yellow solid).
- 8HI + H₂SO₄ → 4I₂ + H₂S + 4H₂O. Sulfur goes from +6 to -2 (H₂S smells of rotten eggs).
Why does reducing ability increase down the group?
Down the group, the halide ion is larger, and its outer electrons are further from the nucleus and more shielded. So an electron is lost more easily. I⁻ is the strongest reducing agent: it reduces sulfur all the way to -2. For oxidation states and half-equations, see Redox Equations.
4. Qualitative Testing with Acidified Silver Nitrate
To identify unknown halide ions in solution:
- Add dilute nitric acid (HNO₃) to eliminate carbonate (CO₃²⁻) or hydroxide (OH⁻) ions that would produce false precipitates. Never use hydrochloric acid, as it introduces chloride ions!
- Add aqueous silver nitrate (AgNO₃): \[ \text{Ag}^+(aq) + \text{X}^-(aq) \rightarrow \text{AgX}(s) \]
- Confirm the precipitate by adding dilute and concentrated aqueous ammonia.
Explain why
- Why AgNO₃? The silver halides form precipitates of different colours: AgCl white, AgBr cream, AgI yellow.
- Why acidify with HNO₃? It removes carbonate and hydroxide ions, which would also form precipitates with Ag⁺. Do not use HCl, because it adds Cl⁻ ions.
- Why add ammonia? The precipitate colours are hard to tell apart. Their different solubilities in ammonia confirm the halide: AgCl dissolves in dilute NH₃, AgBr dissolves in concentrated NH₃, and AgI does not dissolve in concentrated NH₃.
This test is part of RP3: Qualitative Inorganic Analysis.
5. Disproportionation Reactions of Chlorine
A redox reaction in which the same element is simultaneously oxidised and reduced.
Household Bleach Production
Chlorine reacts with cold dilute aqueous sodium hydroxide:
\[ \text{Cl}_2(g) + 2\text{NaOH}(aq) \rightarrow \text{NaCl}(aq) + \text{NaClO}(aq) + \text{H}_2\text{O}(l) \]
Chlorine is simultaneously reduced to -1 in NaCl and oxidised to +1 in sodium chlorate(I), NaClO (active ingredient in bleach).
The solution of sodium chlorate(I) is bleach. It is used to disinfect (kill bacteria) and to bleach.
Water Purification
Chlorine disinfects drinking water by killing bacteria:
\[ \text{Cl}_2(aq) + \text{H}_2\text{O}(l) \rightleftharpoons \text{HCl}(aq) + \text{HClO}(aq) \]
The products contain chloride ions and chlorate(I) ions. This is also disproportionation: chlorine goes from 0 to -1 in HCl and to +1 in HClO.
Chloric(I) acid (HClO) acts as a powerful oxidising disinfectant. In sunlight, chlorine reacts with water to form chloride ions and oxygen:
\[ 2\text{Cl}_2(aq) + 2\text{H}_2\text{O}(l) \rightarrow 4\text{HCl}(aq) + \text{O}_2(g) \]
Chlorine in water treatment: benefits and risks
- Benefit: chlorine kills bacteria in drinking water and prevents diseases spread by water.
- Risks: chlorine is toxic. It can also react with organic matter in the water to form chlorinated organic compounds.
- Decision: society decides that the health benefits outweigh the risks, so small, controlled amounts are added.
Go deeper
These UK AQA pages cover the same ideas in more detail. They may include content that is not in 9620.
6. Practice Questions
Show Answer & Mark Scheme
Procedure & Observations:
- Add dilute nitric acid (HNO₃), followed by aqueous silver nitrate (AgNO₃). (1 mark)
- A cream precipitate of silver bromide (AgBr) forms. (1 mark)
- Add dilute ammonia solution: the precipitate does not dissolve. (1 mark)
- Add concentrated ammonia solution: the precipitate dissolves to produce a colourless solution. (1 mark)
Show Answer & Mark Scheme
Mark Scheme:
- 8HI + H₂SO₄ → 4I₂ + H₂S + 4H₂O. (1 mark)
- Sulfur goes from +6 to -2. (1 mark)
- Any two from: purple fumes; black solid (I₂); yellow solid (S); smell of rotten eggs (H₂S); misty fumes (HI). (2 marks)
Show Answer & Mark Scheme
Mark Scheme:
- Nitric acid removes carbonate and hydroxide ions, which would also form precipitates with Ag⁺. (1 mark)
- Hydrochloric acid would add Cl⁻ ions, giving a white precipitate whatever the halide. (1 mark)
- The precipitate colours are hard to tell apart, so the different solubilities in ammonia confirm the halide. (1 mark)
Show Answer & Mark Scheme
Mark Scheme:
- Cl₂ + H₂O ⇌ HCl + HClO. (1 mark)
- Benefit: kills bacteria and prevents disease. (1 mark)
- Risk: chlorine is toxic, or it forms chlorinated organic compounds. (1 mark)
- The health benefits outweigh the risks. (1 mark)