Group 7: The Halogens

The halogens — fluorine (F₂), chlorine (Cl₂), bromine (Br₂), iodine (I₂), and the rare radioactive astatine (At) — are the most reactive non-metals in the Periodic Table. They sit one electron short of a noble-gas configuration (ns²np⁵), and that single electron-deficit dominates every aspect of their chemistry: their intense colours, their oxidising power, the displacement reactions that demonstrate that power, and the silver-nitrate test that lets us identify their halide ions in unknown solutions. This lesson develops a single, organising idea: oxidising power of the halogens DECREASES down the group, while reducing power of the halide ions INCREASES down the group — the exact opposite trend to Group 2, and a beautiful periodic-table inversion that examiners love to test. You will see this idea worked out in physical properties, halogen–halide displacement (often with cyclohexane to lift the halogen into a coloured organic layer), the colour-by-colour AgNO₃ + NH₃ ladder for halide identification, and the spectacularly differentiated redox chemistry of halides with concentrated H₂SO₄. The lesson closes with disproportionation of Cl₂ in NaOH (cold dilute → bleach; hot concentrated → weedkiller) and the public-health bargain of chlorinated drinking water.

Spec mapping (AQA 7405): This lesson maps to §3.2.3 (Group 7, the halogens) in the full AQA A-Level Chemistry specification. Trends in physical properties and oxidising ability build directly on §3.1.1 (atomic structure) and §3.1.3 (bonding/polarisability), and contrast deliberately with the Group 2 trends established in lesson 1 (§3.1.6 / §3.2.2). Redox bookkeeping uses the oxidation-state and half-equation conventions formalised in §3.1.7 (oxidation, reduction and redox equations) and the qualitative electrode-potential reasoning developed in §3.1.11 (electrochemical cells, A2). The halide-with-AgNO₃ test and the cyclohexane displacement procedure are the practical anchors for Required Practical 10 (qualitative tests for ions), with Required Practical 11 (the iodine–thiosulfate titration) drawing on the same I₂/I⁻ redox chemistry. Refer to the official AQA specification document for the exact wording of each section.

Assessment objectives: AO1 items include recall of the colours and physical states of F₂–I₂ at room temperature, the AgNO₃ precipitate colours (white/cream/yellow) and their behaviour with dilute and concentrated ammonia, and the direction of the oxidising-power and reducing-power trends down the group. AO2 items include predicting displacement reactions in colour-coded grids, writing and balancing ionic equations (including disproportionation of Cl₂ in cold/hot NaOH), and identifying an unknown halide from observations. AO3 items demand explanation: rationalising why oxidising power decreases despite F₂ being anomalously reactive, explaining the differential reduction of concentrated H₂SO₄ by Cl⁻/Br⁻/I⁻ in terms of E° values, and constructing a balanced risk–benefit argument for chlorination of drinking water.

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Physical Properties: Colour, State and the London-Force Trend

Halogen	Formula	Colour at RT	State at RT	Melting / °C	Boiling / °C
Fluorine	F₂	Pale yellow	Gas	−220	−188
Chlorine	Cl₂	Green-yellow	Gas	−101	−34
Bromine	Br₂	Red-brown	Liquid	−7	59
Iodine	I₂	Purple-black (lustrous solid; purple vapour)	Solid	114	184

The single trend behind every entry in this table is the strength of London (dispersion) forces between diatomic X₂ molecules. All four halogens are non-polar, so dipole–dipole and hydrogen bonding play no role; the only intermolecular force is the temporary, induced-dipole London force, and that force scales with the polarisability of the electron cloud. Polarisability rises sharply down the group because the outermost electrons sit progressively further from the nuclei (atomic radius increases) and are less tightly held. F₂ has 18 electrons in a tight volume, weakly polarisable → a gas with a very low boiling point. I₂ has 106 electrons spread over a much larger volume, highly polarisable → a lustrous purple-black solid at room temperature, sublimating to a brilliant violet vapour when warmed.

A useful AQA-style framing: when asked "explain why boiling point increases down Group 7", three marking points are usually available — (i) all halogens are non-polar diatomic molecules held only by London forces; (ii) the number of electrons (and the polarisability of the molecule) increases down the group; (iii) stronger London forces require more energy to overcome, so b.p. increases. Avoid loose phrases such as "the bonds get stronger" — the intramolecular X–X covalent bond is not what melts; the intermolecular London force is.

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Oxidising Power: Why It Decreases Down the Group

A halogen acts as an oxidising agent by gaining an electron per atom:

½ X₂ + e⁻ → X⁻ (reduction half-equation)

The strength of the oxidising agent is measured by the standard electrode potential E° for this half-reaction. At A-Level you should know the direction of the trend; the numerical values are useful background for A*-level synoptic answers:

Half-reaction	E° / V
½ F₂ + e⁻ → F⁻	+2.87
½ Cl₂ + e⁻ → Cl⁻	+1.36
½ Br₂ + e⁻ → Br⁻	+1.09
½ I₂ + e⁻ → I⁻	+0.54

Oxidising power therefore decreases in the order F₂ > Cl₂ > Br₂ > I₂, and the F₂/Cl₂ gap is much larger than the gaps lower down. Three factors explain the trend:

1. Atomic radius increases down the group, so the incoming electron ends up further from the nucleus.
2. Shielding by inner-shell electrons increases in proportion to the number of complete shells (F has [He] core, I has [Kr] 4d¹⁰ core).
3. Effective nuclear charge experienced by the incoming electron remains roughly constant across a single group, so radius and shielding dominate: the electron affinity becomes less exothermic, and E° falls.

Examiners specifically reward the wording "the incoming electron is held less tightly because of increased radius and shielding" rather than vague references to "size".

A subtlety — why F is anomalous.* F₂ has the most positive E° in the table, but the electron affinity of F (−328 kJ mol⁻¹) is actually less exothermic than that of Cl (−349 kJ mol⁻¹). Despite this, F₂ is a stronger oxidising agent than Cl₂ because (a) the F–F bond is anomalously weak (only +158 kJ mol⁻¹; cf. Cl–Cl +243 kJ mol⁻¹) due to lone-pair repulsion across the very short F–F bond, and (b) the high charge density of F⁻ leads to a large negative hydration enthalpy (−506 kJ mol⁻¹). The full Born–Haber analysis shows that the bond-enthalpy term and the hydration term combine to drive F₂ reduction strongly, even though the electron-affinity term itself does not. Quoting any of this in an A* answer is high-end synoptic reasoning.

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Reducing Power of Halide Ions: The Mirror Trend

The reduction half-equation, run in reverse, gives the oxidation of a halide ion:

X⁻ → ½ X₂ + e⁻

The thermodynamic strength of X⁻ as a reducing agent is the negative of the corresponding E° — so the trend is exactly reversed: reducing power of X⁻ INCREASES down the group: F⁻ < Cl⁻ < Br⁻ < I⁻. Iodide is the strongest reducing agent of the lot, because the outer 5p electron in I⁻ is held weakly: the radius is large (220 pm vs F⁻ at 133 pm), the effective nuclear charge experienced is moderate, and the electron is straightforward to remove.

The two trends are not independent — they are paired: a strong oxidising halogen necessarily produces a halide ion that is a weak reducing agent, because the same E° governs both directions of the redox couple. Examiners frequently catch students by asking the trends in opposite directions in the same paper. Remember: "strong oxidiser → weak reducer of its conjugate halide" and vice versa.

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Displacement Reactions: Putting the Trend on Show

A more reactive halogen will displace a less reactive halide ion from solution. The reaction is a single-step electron transfer; the more oxidising halogen pulls electrons off the less oxidising halide, regenerating the latter as its diatomic molecule.

Add ↓ to →	KCl(aq)	KBr(aq)	KI(aq)
Cl₂(aq)	No reaction	Orange/brown solution (Br₂)	Brown solution (I₂)
Br₂(aq)	No reaction	No reaction	Brown solution (I₂)
I₂(aq)	No reaction	No reaction	No reaction

Equations

Chlorine displaces bromide:

Cl₂(aq) + 2KBr(aq) → 2KCl(aq) + Br₂(aq)

Ionic equation: Cl₂(aq) + 2Br⁻(aq) → 2Cl⁻(aq) + Br₂(aq)

• Chlorine is reduced (0 → −1): the oxidising agent.
• Bromide is oxidised (−1 → 0): the reducing agent.

Chlorine displaces iodide:

Cl₂(aq) + 2KI(aq) → 2KCl(aq) + I₂(aq)

Ionic equation: Cl₂(aq) + 2I⁻(aq) → 2Cl⁻(aq) + I₂(aq)

Bromine displaces iodide:

Br₂(aq) + 2KI(aq) → 2KBr(aq) + I₂(aq)

Ionic equation: Br₂(aq) + 2I⁻(aq) → 2Br⁻(aq) + I₂(aq)

Using cyclohexane as an observation aid

The aqueous colours of Br₂ (orange) and I₂ (brown) overlap and can be hard to distinguish, especially at low concentrations. A standard trick is to layer cyclohexane (an immiscible, colourless organic solvent) on top of the aqueous mixture and shake. The non-polar halogens partition strongly into the organic layer, where they show clearly different colours:

Halogen in cyclohexane	Colour
Cl₂	Pale green / colourless
Br₂	Orange / yellow-orange
I₂	Violet / purple (diagnostic)

The diagnostic colour is the violet of I₂ in cyclohexane: it is unmistakable, and AQA examiners often phrase questions around "after shaking with cyclohexane, the upper layer turned violet — identify the halide ion originally present". (Older specifications used CCl₄ or 1,1,1-trichloroethane; cyclohexane is the modern non-toxic substitute and is the solvent specified for RP10.)

Exam tip: In any displacement-reaction question, write the ionic equation, then explicitly state the oxidation-state changes and which species is the oxidising agent / reducing agent. Two of the three available marks usually depend on that bookkeeping.

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Reactions of Solid Halides with Concentrated H₂SO₄

This reaction is the single best demonstration in school chemistry that reducing power of halide ions increases down the group. Concentrated sulfuric acid is a strong acid and a potential oxidising agent (S is at +6 in H₂SO₄ and can be reduced to +4, 0, or −2). What happens depends entirely on whether the halide ion is reducing enough to pull sulfur down to lower oxidation states.

NaCl + conc. H₂SO₄ — acid–base only

NaCl(s) + H₂SO₄(l) → NaHSO₄(s) + HCl(g)

• Observations: white solid moistens; steamy white fumes of HCl gas evolve.
• No colour, no SO₂ smell, no redox. Chloride is too weak a reducing agent to reduce sulfur from +6.

NaBr + conc. H₂SO₄ — partial redox

Step 1 (acid–base): NaBr(s) + H₂SO₄(l) → NaHSO₄(s) + HBr(g)

Step 2 (redox): 2HBr(g) + H₂SO₄(l) → Br₂(g) + SO₂(g) + 2H₂O(l)

• Observations: steamy white HBr fumes accompanied by orange/brown fumes of Br₂ and the choking smell of SO₂.
• Oxidation-state changes: Br: −1 → 0; S: +6 → +4. Bromide is reducing enough to drop sulfur by two.

NaI + conc. H₂SO₄ — full redox cascade

Step 1 (acid–base): NaI(s) + H₂SO₄(l) → NaHSO₄(s) + HI(g)

Step 2 (mild redox): 2HI(g) + H₂SO₄(l) → I₂(s) + SO₂(g) + 2H₂O(l)

Step 3 (deeper redox): 6HI(g) + H₂SO₄(l) → 3I₂(s) + S(s) + 4H₂O(l)

Step 4 (deepest): 8HI(g) + H₂SO₄(l) → 4I₂(s) + H₂S(g) + 4H₂O(l)

• Observations: purple/black I₂ vapour and solid; yellow elemental sulfur deposits; the rotten-egg smell of H₂S.
• Oxidation-state changes: I: −1 → 0; S: +6 → +4 (SO₂), 0 (S), and −2 (H₂S). Iodide can take sulfur all the way down by eight oxidation states.

Exam tip: When asked to compare these three reactions, name the gases evolved (HX, then SO₂ / S / H₂S as appropriate), give the colour of any halogen formed, and explicitly state the sulfur oxidation states. Marks are almost always awarded for the redox bookkeeping, not just the observation. The contrast also rules out conc. H₂SO₄ as a laboratory reagent for preparing pure HBr or HI — phosphoric acid (H₃PO₄) is used instead because it is non-oxidising.

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Silver-Nitrate Test for Halide Ions (RP10 anchor)

The AgNO₃ test is the standard inorganic identification for an unknown halide ion in aqueous solution. It is the practical anchor for Required Practical 10 (qualitative tests for ions). The procedure must be precise because two interferences — carbonate and hydroxide — would otherwise give false positives by precipitating Ag₂CO₃ or AgOH.

Procedure:

1. Acidify with dilute nitric acid (HNO₃). This removes carbonate (as CO₂) and hydroxide (as water), so neither can co-precipitate.
2. Add aqueous silver nitrate (AgNO₃). A coloured precipitate forms if a halide is present.
3. Note the precipitate colour (white / cream / yellow).
4. Confirm with ammonia. Add dilute NH₃ first; if no change, add concentrated NH₃.

Halide	Precipitate	Colour	Behaviour with dilute NH₃	Behaviour with concentrated NH₃
Cl⁻	AgCl	White	Dissolves	Dissolves
Br⁻	AgBr	Cream	Insoluble	Dissolves
I⁻	AgI	Yellow	Insoluble	Insoluble

(AgF is fully soluble in water — fluoride gives no precipitate in this test, which is itself diagnostic of fluoride against the other three halides.)

Equations

Ag⁺(aq) + Cl⁻(aq) → AgCl(s) [white precipitate]

Ag⁺(aq) + Br⁻(aq) → AgBr(s) [cream precipitate]

Ag⁺(aq) + I⁻(aq) → AgI(s) [yellow precipitate]

Why ammonia dissolves the precipitates differently

Ammonia is a Lewis base. It coordinates to Ag⁺ to form the linear, colourless diamminesilver(I) complex:

AgX(s) + 2NH₃(aq) → [Ag(NH₃)₂]⁺(aq) + X⁻(aq)

The position of this equilibrium depends on the lattice enthalpy of the AgX solid. AgCl has the smallest lattice enthalpy, so even dilute NH₃ generates enough [Ag(NH₃)₂]⁺ to pull the equilibrium to the right and dissolve the solid. AgBr has a larger lattice enthalpy, so concentrated NH₃ is needed. AgI has the largest lattice enthalpy (because I⁻ is large and polarisable, with strong covalent character in the Ag–I bond), so even concentrated NH₃ cannot shift the equilibrium and AgI stays insoluble.

Exam tip: Always write the acidification step first when describing the test, and always specify "dilute nitric acid". Forgetting either loses an easy mark. A common A-Level error is to acidify with HCl, which would itself precipitate AgCl — a self-defeating choice.

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Disproportionation of Chlorine in NaOH

Chlorine reacts with sodium hydroxide in a classic disproportionation — Cl in Cl₂ (oxidation state 0) is simultaneously oxidised and reduced. The product depends sharply on temperature and concentration:

Cold dilute NaOH (≈ 15 °C) → household bleach

Cl₂(g) + 2NaOH(aq) → NaCl(aq) + NaClO(aq) + H₂O(l)

• Chlorine is reduced (0 → −1 in Cl⁻) and simultaneously oxidised (0 → +1 in ClO⁻, the chlorate(I) or hypochlorite ion).
• Sodium chlorate(I) is the active ingredient in household bleach and pool sanitiser.

Hot concentrated NaOH (≈ 70 °C) → sodium chlorate(V)

3Cl₂(g) + 6NaOH(aq) → 5NaCl(aq) + NaClO₃(aq) + 3H₂O(l)

• Chlorine disproportionates to −1 (in Cl⁻) and +5 (in ClO₃⁻, the chlorate(V) ion).
• Sodium chlorate(V), NaClO₃, is a strong oxidising agent and is widely used as a weedkiller (and historically in matches and pyrotechnics).

Definition. Disproportionation is a redox reaction in which the same element is simultaneously oxidised and reduced. Examiners want this exact wording.

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Chlorine in Water Treatment: Risk and Benefit

Adding small quantities of chlorine (typically 0.5–2 ppm = mg dm⁻³) to drinking water disinfects it. Chlorine reacts with water to set up an equilibrium:

Cl₂(aq) + H₂O(l) ⇌ HCl(aq) + HClO(aq)

The active disinfectant is chloric(I) acid (HClO), which penetrates bacterial cell walls and oxidises essential intracellular biomolecules. The introduction of chlorinated drinking water in industrial cities at the turn of the 20th century is one of the single largest public-health interventions in history, virtually eliminating cholera, typhoid and dysentery from developed-world urban populations.

Risks must also be discussed in any exam answer asking for evaluation:

• Cl₂ is itself a toxic gas; supply and dosing systems must be carefully engineered.
• Chlorine reacts with trace organic matter (humic acids from soil and plant decay) to form chlorinated organic by-products such as chloromethanes (CHCl₃) and haloacetic acids; some of these are suspected carcinogens at long-term exposure.
• In some jurisdictions the risk/benefit balance has been shifted by using chloramines (NH₂Cl) or ozone instead.

Exam tip: When asked to "discuss the use of chlorine in water treatment", structure your answer as benefit (kills pathogens, prevents waterborne disease at very low Cl₂ ppm); risk (toxic gas; chlorinated by-products); evaluation (net benefit dramatically outweighs the residual risk at currently regulated dose limits). A balanced two-sided answer is essential for the AO3 marks.

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Practical Skills Box (RP10 anchor)

The qualitative-test work in RP10 typically combines the AgNO₃ halide test with the displacement/cyclohexane procedure as part of identifying ions in an unknown sample.