OxfordAQA A-Level Chemistry: Group 7(17), the halogens (9620) – Revision Notes

These notes squeeze section 3.2.3, on the Group 7(17) halogens, of the OxfordAQA International AS and A-level Chemistry (9620) specification, Version 5.3, for International AS and A-level exams May/June 2020 onwards. Covered here: point 3.2.3.1 on trends in properties, point 3.2.3.2 on the uses of chlorine and chlorate(I), and Required practical 3. Every point is International AS, with nothing International A-level only. It belongs to the Unit 1 paper (Inorganic 1 and Physical 1), though the Unit 5 paper (Practical and synoptic) may examine any of it.

Full explanations and worked examples are in the halogens study guide. Afterwards, check your recall with the halogens practice questions. Other 9620 topics are on the course hub, and the printable checklist helps you track progress. A free 10-minute diagnostic shows where to focus. For half-equation technique, revise oxidation, reduction and redox equations.

3.2.3.1 The four trends at a glance

Down the group, F → I	Direction	One-line reason
Electronegativity	falls	larger atom, more shielding, so the bonding pair is held less tightly
Boiling point	rises	more electrons per molecule, so stronger van der Waals forces
Oxidising ability of X₂	falls	an incoming electron is further from the nucleus and more shielded
Reducing ability of X⁻	rises	an outer electron is further out and more shielded, so easier to lose

Add a fifth for the tests: the solubility of AgX in ammonia falls from AgCl to AgI.

Definition: electronegativity is an atom’s power to attract the electron pair in a covalent bond. Learn the specification’s wording exactly.

Data worth knowing: electrons per molecule F₂ 18, Cl₂ 34, Br₂ 70, I₂ 106. Boiling points roughly 85 K, 239 K, 332 K, 457 K. Pauling electronegativities roughly 4.0, 3.2, 3.0, 2.7.

Method: a boiling point explanation in four steps

1. Name the structure: simple molecular, diatomic.
2. Compare electrons per molecule, with numbers.
3. Link more electrons to stronger van der Waals (induced dipole–dipole) forces between molecules.
4. More energy is needed to overcome them.

The X–X covalent bond is not broken on boiling. Writing that it is costs the explanation.

3.2.3.1 Displacement in aqueous solution

Rule: a halogen oxidises the halide ions of any halogen lower in the group.

Halogen added	Cl⁻(aq)	Br⁻(aq)	I⁻(aq)
Cl₂	–	Br₂ forms	I₂ forms
Br₂	none	–	I₂ forms
I₂	none	none	–

General ionic equation, where X sits above Y: X₂ + 2Y⁻ → 2X⁻ + Y₂.

Colour tip: shake with cyclohexane. Iodine turns the upper organic layer violet; in water it looks brown. Bromine gives an orange organic layer.

3.2.3.1 Solid sodium halides with concentrated H₂SO₄

Step 1, all halides (acid–base): NaX + H₂SO₄ → NaHSO₄ + HX. You see steamy fumes.

Step 2, only if X⁻ is a strong enough reducing agent (redox):

Halide	Reduces S to	S oxidation state	Extra signs
F⁻, Cl⁻	–	+6 (no change)	none
Br⁻	SO₂	+4	orange-brown vapour, choking gas
I⁻	SO₂, S, H₂S	+4, 0, −2	purple vapour or black solid, yellow solid, bad-egg smell

Sulfur half-equations (all reductions):

H₂SO₄ + 2H⁺ + 2e⁻ → SO₂ + 2H₂O         S: +6 → +4
H₂SO₄ + 6H⁺ + 6e⁻ → S + 4H₂O           S: +6 → 0
H₂SO₄ + 8H⁺ + 8e⁻ → H₂S + 4H₂O         S: +6 → −2

Pair each with 2X⁻ → X₂ + 2e⁻, scaled so the electrons cancel (×1, ×3 or ×4).

Method: an overall equation that starts from the solid

1. Count electrons: how far does sulfur fall? That fixes the number of halide ions.
2. Write the redox part (halogen, sulfur product).
3. Every Na⁺ ends up in NaHSO₄, so add one H₂SO₄ per sodium.
4. Balance O and H with water, then check every element.

Example, iodide going only as far as SO₂: 2NaI + 3H₂SO₄ → 2NaHSO₄ + I₂ + SO₂ + 2H₂O.

3.2.3.1 Silver nitrate test

Method in steps:

1. Dissolve the sample and add dilute nitric acid.
2. Add silver nitrate solution and note any precipitate.
3. Add dilute ammonia; if the precipitate remains, try concentrated ammonia.

Precipitate	Ion	Dilute NH₃	Conc. NH₃
none	F⁻	–	–
white	Cl⁻	soluble	soluble
cream	Br⁻	insoluble	soluble
yellow	I⁻	insoluble	insoluble

The three “explain why” answers:

Choice	Reason
silver nitrate	precipitates of different colours form with Cl⁻, Br⁻ and I⁻
acidified with HNO₃	removes ions such as CO₃²⁻ and OH⁻, which would also precipitate with Ag⁺; not HCl, which adds Cl⁻
ammonia added	the three colours look alike; solubility in ammonia confirms which halide

Dissolving: AgCl(s) + 2NH₃(aq) → [Ag(NH₃)₂]⁺(aq) + Cl⁻(aq).

3.2.3.2 Chlorine equations

Reaction	Equation	Type
with water	Cl₂ + H₂O ⇌ HClO + HCl	disproportionation (0 → +1 and −1), reversible
with water in sunlight	2Cl₂ + 2H₂O → 4H⁺ + 4Cl⁻ + O₂	redox: Cl reduced, O oxidised
with cold, dilute NaOH	Cl₂ + 2NaOH → NaClO + NaCl + H₂O	disproportionation (0 → +1 and −1)

Names: ClO⁻ is the chlorate(I) ion; NaClO is sodium chlorate(I). Its solution is bleach, also used to disinfect.

Water treatment, in four points: bacteria are killed; water-borne disease is prevented; chlorine is toxic and can produce chlorinated organic compounds; on balance, health gains win over toxic effects. Any chemical added to a water supply (fluoride is another) is judged by weighing pros against cons.

Required practical 3 in one table

Ion	Reagent	Positive result
Mg²⁺ / Ba²⁺ (Group 2)	NaOH(aq), and dilute H₂SO₄	Mg²⁺: white hydroxide precipitate; Ba²⁺: white sulfate precipitate
NH₄⁺	NaOH(aq), warm	gas turns damp red litmus blue
halide	HNO₃, AgNO₃(aq), NH₃(aq)	see the silver nitrate table
OH⁻	red litmus	blue
CO₃²⁻	dilute acid, limewater	fizzing; limewater milky
SO₄²⁻	acidified BaCl₂(aq)	white precipitate

Small worked reminders

Reminder 1, a silver iodide mass. A solution holds 0.00200 mol of iodide ions. Excess acidified silver nitrate gives 0.00200 mol AgI. Mr(AgI) = 107.9 + 126.9 = 234.8, so the mass is 0.00200 × 234.8 = 0.470 g.

Reminder 2, electrons in the NaOH reaction. In Cl₂ + 2OH⁻ → Cl⁻ + ClO⁻ + H₂O, one Cl atom gains one electron (0 → −1) and the other loses one (0 → +1). The totals match, so no electrons appear in the overall equation.

Reminder 3, the mole ratio for bleach. One Cl₂ gives one NaClO. To make 0.0400 mol NaClO you need 0.0400 mol Cl₂, which is 0.0400 × 71.0 = 2.84 g.

Must-know distinctions

• Halogens are oxidising agents (weaker down); halide ions are reducing agents (stronger down).
• Steamy fumes (HX, acid–base) versus coloured vapour (X₂, redox).
• Dilute ammonia dissolves only AgCl; concentrated dissolves AgCl and AgBr.
• Chlorine water in the dark (an equilibrium, disproportionation) versus in sunlight (oxygen released, not disproportionation).
• Nitric acid to acidify silver nitrate; acidified barium chloride for sulfate.

Quick self-test

1. Why is chlorine more electronegative than iodine?
2. Which halogen has the weakest van der Waals forces between its molecules, and why?
3. Bromine water is added to sodium chloride solution. What happens, and why?
4. Write the ionic equation for chlorine displacing iodine from iodide ions.
5. Which sodium halides give only steamy fumes with concentrated sulfuric acid?
6. Name the gas with a bad-egg smell from sodium iodide and concentrated sulfuric acid, and give the oxidation state of sulfur in it.
7. A precipitate from acidified silver nitrate dissolves in concentrated ammonia but not dilute. Name it.
8. Calculate the mass of AgCl formed from 0.00150 mol of chloride ions. (Mr AgCl = 143.4)
9. In the sunlight reaction of chlorine water, which element is oxidised?
10. State one use of the solution made from chlorine and cold, dilute sodium hydroxide.
11. Which reagent tests for sulfate ions in Required practical 3?

Answers

1. Chlorine’s atom is smaller with less shielding, so the bonding pair is closer to the nucleus and attracted more strongly.
2. Fluorine; F₂ has the fewest electrons per molecule (18).
3. No reaction; bromine is a weaker oxidising agent than chlorine, so it cannot oxidise Cl⁻.
4. Cl₂ + 2I⁻ → 2Cl⁻ + I₂.
5. Sodium fluoride and sodium chloride.
6. Hydrogen sulfide, H₂S; −2.
7. Silver bromide, AgBr (cream).
8. 0.00150 × 143.4 = 0.215 g.
9. Oxygen (−2 in water to 0 in O₂).
10. Bleach (or a disinfectant for cleaning).
11. Acidified barium chloride solution.

Where marks are usually lost

• Describing electronegativity with “nuclear charge falls”; it rises, but distance and shielding win.
• Saying covalent bonds break on boiling.
• Comparing boiling points by molecule “size” without electron numbers.
• Mixing the two redox trends: halogen oxidising power falls, halide reducing power rises.
• Writing a redox step for sodium chloride.
• Unbalanced charges in the sulfuric acid half-equations.
• Forgetting the extra H₂SO₄ that forms NaHSO₄ in an equation from the solid salt.
• Quoting “AgBr is insoluble in ammonia” without saying dilute.
• Calling the sunlight reaction disproportionation.

Official syllabus

OxfordAQA, International AS and A-level Chemistry (9620) specification, Version 5.3, for International AS and A-level exams May/June 2020 onwards; within it, section 3.2.3, the halogens of Group 7(17), at International AS: point 3.2.3.1, on trends in properties; point 3.2.3.2, on the uses of chlorine and chlorate(I); and Required practical 3.