What are the elements in Group 7 of the periodic table?

Group 7 of the periodic table, also known as the halogens, includes the elements fluorine (F), chlorine (Cl), bromine (Br), iodine (I), and astatine (At). These elements are known for their high reactivity and are commonly studied in the Edexcel IGCSE Chemistry curriculum.

Why are Group 7 elements called halogens?

Group 7 elements are called halogens, which means 'salt-formers', because they readily react with metals to form salts. For example, sodium chloride (table salt) is formed when chlorine reacts with sodium. This characteristic is a key focus in Inorganic Chemistry studies at the IGCSE level.

How does reactivity change as you move down Group 7?

In Group 7, reactivity decreases as you move down the group from fluorine to astatine. This is because the atomic size increases, making it harder for the nucleus to attract additional electrons needed for reactions. This trend is an important concept in Edexcel IGCSE Chemistry.

What are the physical states of Group 7 elements at room temperature?

At room temperature, the physical states of Group 7 elements vary: fluorine is a pale yellow gas, chlorine is a greenish gas, bromine is a reddish-brown liquid, and iodine is a dark purple solid. Astatine is rarely encountered but is expected to be a solid. Understanding these states is crucial for IGCSE Chemistry students.

What are some common uses of Group 7 elements?

Group 7 elements have various applications: chlorine is used in water purification and disinfectants, bromine in flame retardants, iodine in antiseptics and nutritional supplements, and fluorine in toothpaste and Teflon production. These uses highlight the practical importance of halogens in everyday life, a topic covered in the Edexcel IGCSE Chemistry syllabus.

Why is "Stating that Group 7 reactivity INCREASES going down (like Group 1)" a common mistake in Group 7?

Why it happens: Generalising from Group 1. How to avoid it: Group 1 LOSES electrons → bigger atoms easier to lose → more reactive down. Group 7 GAINS electrons → bigger atoms harder to attract → LESS reactive down. The trends go in OPPOSITE directions. Memorise: 'gaining = harder when shielded'. Source: 4CH1 Examiner Reports 2022-2024

Why is "Saying the iodide displacement reaction turns the solution PURPLE" a common mistake in Group 7?

Why it happens: Confusing the vapour colour with the solution colour. How to avoid it: I₂ in solution (especially aqueous) is BROWN. I₂ as a VAPOUR is purple. Iodine dissolved in non-polar solvents (cyclohexane) is purple. State 'brown solution' for an aqueous displacement.

Why is "Saying 'bonds get stronger down Group 7, so mp/bp increase'" a common mistake in Group 7?

Why it happens: Conflating covalent bonds with intermolecular forces. How to avoid it: Within X₂ molecules, the covalent bonds actually WEAKEN going down (longer bonds). But mp/bp depends on INTERMOLECULAR forces between molecules — these increase down the group because larger molecules have more electrons → stronger London dispersion forces. Source: 4CH1 Examiner Reports 2022-2024

Why is "Writing I₂ + KCl → KI + Cl₂ (i.e. I₂ displacing Cl)" a common mistake in Group 7?

Why it happens: Trying to balance any pair of halogen + halide. How to avoid it: Only a MORE REACTIVE halogen can displace a less reactive one. Reactivity order: Cl > Br > I. So Cl₂ + KBr/KI works; Br₂ + KI works; I₂ + KCl / KBr does NOT (no reaction).

Why is "Confusing the chlorine test (bleaches damp litmus) with the ammonia test (damp red litmus turns blue)" a common mistake in Group 7?

Why it happens: Both involve damp litmus paper. How to avoid it: **Cl₂**: turns damp BLUE litmus RED, then BLEACHES it (red → white). **NH₃**: turns damp RED litmus BLUE. Cl₂ is acidic + an oxidiser (bleaches); NH₃ is alkaline.

Why is "Writing Cl₂ + H₂O → HCl + HOCl with a single arrow (forward only)" a common mistake in Group 7?

Why it happens: Treating it as a normal reaction. How to avoid it: The reaction of chlorine with water is REVERSIBLE — use the equilibrium arrows ⇌: Cl_2 + H_2O leftharpoons HCl + HOCl. The presence of HOCl is what kills bacteria.

Inorganic ChemistryPearson Edexcel IGCSE Chemistry 4CH1 Higher

Group 7

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Group 7 / The Halogens — Pearson Edexcel International GCSE Chemistry 4CH1 Study Notes (2026 onwards, Spec Issue 4)

The halogens F, Cl, Br, I, At — Group 7 elements with 7 outer-shell electrons that gain ONE to form X⁻ ions. Physical states change going down the group: F₂/Cl₂ gases → Br₂ liquid → I₂ solid (sublimes to purple vapour). Melting and boiling points INCREASE down the group due to stronger intermolecular (London dispersion) forces between larger diatomic molecules. REACTIVITY DECREASES down the group (F > Cl > Br > I) because the outer shell is further from the nucleus and shielded by more inner electrons — harder to attract an incoming electron. Displacement reactions: a more reactive halogen displaces a less reactive halide from its salt. Uses: water chlorination, bleach manufacture, brominated flame retardants, iodine antiseptic and salt iodisation — with benefit/risk discussions.

At a glance

Group 7 = F, Cl, Br, I, At — diatomic molecules X₂; 7 outer-shell electrons; gain 1 to form X⁻.
States at 20 °C: F₂ very pale yellow gas; Cl₂ pale green-yellow gas; Br₂ red-brown liquid; I₂ dark grey solid (sublimes to purple vapour).
mp/bp INCREASE down the group — larger molecules → stronger London dispersion forces between molecules.
Reactivity DECREASES down the group (F > Cl > Br > I) — atomic radius increases + shielding increases → weaker attraction for incoming electron.
Displacement: more reactive halogen displaces less reactive halide. Cl₂ + 2KBr → 2KCl + Br₂; Cl₂ + 2KI → 2KCl + I₂; Br₂ + 2KI → 2KBr + I₂.
Reactions with H₂: H₂ + Cl₂ → 2HCl (explosive in sunlight); H₂ + Br₂ → 2HBr (slow); H₂ + I₂ ⇌ 2HI (reversible, slow).
Uses: Cl₂ — water disinfection, bleach, PVC. Br₂ — flame retardants, photographic film. I₂ — antiseptic, iodised salt (prevents goitre).

What you’ll learn

Mapped to the Cambridge IGCSE 4CH1 syllabus (2026 onwards).

2.58 — Recall the physical state, colour and electron arrangement of fluorine, chlorine, bromine and iodine at room temperature.
2.59 — Explain the trend in physical properties (mp/bp) down Group 7 in terms of molecular size and intermolecular forces.
2.60 — Describe the trend in reactivity going down Group 7 (decreasing).
2.61 — Explain the reactivity trend in terms of electron arrangement, atomic radius, and electron shielding.
2.62 — Recall the reactions of the halogens with hydrogen to form hydrogen halides.
2.63 — Predict and describe halogen displacement reactions (e.g. Cl₂ + KBr).
2.64 — Write balanced symbol and ionic equations for displacement reactions and identify the oxidising / reducing species.
2.65 — Describe the use of chlorine in water treatment, bleach and PVC; the use of bromine in flame retardants; the use of iodine as antiseptic + dietary supplement.
2.66 — Discuss the benefits and risks of halogen use in society (e.g. chlorination of water).

Physical states, colours and mp/bp trend (spec 2.58, 2.59)

States change down: gas → liquid → solid. mp/bp increase down because of stronger London dispersion forces.

Atomic structure of the halogens.

Halogen	Symbol	At. no.	Electron config	Outer-shell electrons
Fluorine	F	9	2,7	7
Chlorine	Cl	17	2,8,7	7
Bromine	Br	35	2,8,18,7	7
Iodine	I	53	2,8,18,18,7	7
Astatine	At	85	2,8,18,32,18,7	7 (radioactive, very rare)

All have 7 outer-shell electrons → all behave similarly chemically (Group 7). All exist as diatomic molecules (F₂, Cl₂, Br₂, I₂) — two halogen atoms sharing one electron pair, achieving a stable octet.

Physical states at room temperature (20 °C).

Halogen	Colour	State	mp / °C	bp / °C
F₂	Very pale yellow	GAS	−220	−188
Cl₂	Pale green-yellow	GAS	−101	−34
Br₂	Red-brown	LIQUID (the only liquid Group 7)	−7	+59
I₂	Dark grey / black	SOLID (sublimes to purple vapour)	+114	+184

Trend in colour. Darkens going down: pale yellow (F₂) → green-yellow (Cl₂) → red-brown (Br₂) → dark grey/purple (I₂). The vapour follows the same pattern.

Gas → liquid → solid down the group as London dispersion forces between molecules grow stronger.

Trend in mp/bp. Both melting AND boiling points INCREASE substantially going down the group — bp range is over 370 °C.

Why mp/bp increase down the group.

Halogens form simple molecular substances — diatomic X₂ molecules with weak INTERMOLECULAR FORCES (London dispersion / Van der Waals) holding the molecules together in the solid/liquid state.

When you melt or boil a halogen, you have to overcome these INTERMOLECULAR forces — NOT the covalent bonds within each X₂ molecule. The covalent bonds stay intact.

Going down the group:

Each molecule contains more ELECTRONS (F₂: 18 e⁻ → I₂: 106 e⁻).
More electrons → more 'electron cloud' that can deform → larger temporary dipoles.
Larger temporary dipoles → stronger London dispersion forces between adjacent molecules.
Stronger intermolecular forces → MORE energy needed to overcome them → HIGHER mp and bp.

Key mark-scheme phrase: 'as the molecule has more electrons, the LONDON DISPERSION / VAN DER WAALS forces between molecules become stronger'.

Don't confuse intermolecular forces with covalent bonds. The covalent bond strength WITHIN X₂ molecules actually DECREASES going down (longer bonds, weaker overlap), but this doesn't affect mp/bp.

Iodine sublimes — a famous demo. Solid I₂ is dark grey and crystalline. On gentle heating, it goes DIRECTLY from solid to a beautiful PURPLE vapour without melting (skipping the liquid phase) — this is sublimation. The vapour can recondense on a cool surface as iodine crystals again.

All Group 7 atoms have 7 outer-shell electrons; exist as X₂ diatomic molecules.
States at 20 °C: F₂/Cl₂ gases; Br₂ liquid; I₂ solid (sublimes to purple vapour).
mp/bp increase down the group.
Reason: more electrons → stronger London dispersion (intermolecular) forces.

See the full worked example for group 7 →

Reactivity trend down Group 7 (spec 2.60, 2.61)

DECREASES down (F most reactive; I least). Larger atom + shielding → weaker attraction for incoming electron.

The trend. Reactivity DECREASES going DOWN Group 7. The order is F > Cl > Br > I. Fluorine is the most reactive halogen (and one of the most reactive elements overall — reacts with almost anything including glass and noble gases). Iodine is the least reactive of the stable halogens.

How halogens react. Halogens react by GAINING ONE electron to form a stable −1 ion (a halide), achieving a complete outer shell (noble-gas configuration):

$\text{X}_2 + 2e^- \rightarrow 2\text{X}^-$

So 'reactivity of a halogen' = how easily it can GAIN an electron.

Why the trend goes this direction — three factors.

(1) Atomic radius INCREASES down the group.

Going down: F (smallest) → Cl → Br → I (largest). Each successive halogen has ONE MORE complete electron shell. So the outer-shell electrons (and the place where the 8th electron must join) are FURTHER from the nucleus.

(2) Electron shielding INCREASES.

More inner shells of electrons (between the nucleus and the outer shell) = more SHIELDING / SCREENING of the nuclear charge. For F (one inner shell, 1s²) shielding is small. For I (four inner shells: 1s² 2s²2p⁶ 3s²3p⁶3d¹⁰ 4s²4p⁶4d¹⁰) shielding is much greater.

(3) Effective nuclear charge on incoming electron DECREASES.

The combined effect: the EFFECTIVE attractive force from the nucleus on an incoming electron is REDUCED as you go down the group. The nucleus is further away AND its pull is screened by more inner electrons.

Conclusion. It is HARDER for larger halogen atoms to attract an incoming electron → harder to react → LESS REACTIVE going down. F is most reactive (small + little shielding); I is least reactive.

Comparison with Group 1 — OPPOSITE trend.

In Group 1 (alkali metals), reactivity INCREASES going down because Group 1 reacts by LOSING an outer electron — and the outer electron is more easily lost when it's further from the nucleus and more shielded.

Group	Reaction type	Bigger atom + more shielding →	Trend down
1	LOSE 1 outer electron	EASIER to lose	More reactive
7	GAIN 1 outer electron	HARDER to attract	LESS reactive

Memory tip. 'Gaining gets harder when shielded'. 'Losing gets easier when shielded'.

Evidence — reactions with hydrogen.

$\text{H}_2 + \text{F}_2 \rightarrow 2\text{HF}$ — explosive even in DARK at low temperature.

$\text{H}_2 + \text{Cl}_2 \rightarrow 2\text{HCl}$ — explosive in BRIGHT LIGHT (e.g. UV) or with a spark.

$\text{H}_2 + \text{Br}_2 \rightarrow 2\text{HBr}$ — slow at high temperature.

$\text{H}_2 + \text{I}_2 \rightleftharpoons 2\text{HI}$ — very slow, reversible (equilibrium).

This sequence is a clear demonstration of reactivity DECREASING down the group.

Reactivity DECREASES down Group 7 (F most reactive, I least).
Halogens react by GAINING 1 electron → X⁻.
Going down: atomic radius increases AND shielding increases.
Effective attraction of nucleus on incoming electron DECREASES → less reactive.
OPPOSITE trend to Group 1 (which loses an electron, so trend reverses).

See the full worked example for group 7 →

Halogen displacement reactions (spec 2.63, 2.64)

More reactive halogen displaces less reactive halide from solution. Redox: more reactive is reduced; halide is oxidised.

The rule. A more reactive halogen (higher in Group 7) will DISPLACE a less reactive halogen from a solution of its halide salt.

Reactivity order (top to bottom): Cl > Br > I. So:

Cl₂ can displace Br⁻ AND I⁻ from their salts.
Br₂ can displace only I⁻ (not Cl⁻, because Cl is MORE reactive than Br).
I₂ cannot displace any other halide (it's the least reactive).

Examples — balanced equations and observations.

(1) Chlorine + potassium bromide solution.

$\text{Cl}_2\text{(aq)} + 2\text{KBr(aq)} \rightarrow 2\text{KCl(aq)} + \text{Br}_2\text{(aq)}$

Ionic: $\text{Cl}_2 + 2\text{Br}^- \rightarrow 2\text{Cl}^- + \text{Br}_2$ .

Observation: colourless KBr solution turns ORANGE/BROWN (Br₂ formed).

(2) Chlorine + potassium iodide solution.

$\text{Cl}_2\text{(aq)} + 2\text{KI(aq)} \rightarrow 2\text{KCl(aq)} + \text{I}_2\text{(aq)}$

Ionic: $\text{Cl}_2 + 2\text{I}^- \rightarrow 2\text{Cl}^- + \text{I}_2$ .

Observation: colourless KI solution turns BROWN (I₂ formed); purple vapour may be visible above the solution.

(3) Bromine + potassium iodide solution.

$\text{Br}_2\text{(aq)} + 2\text{KI(aq)} \rightarrow 2\text{KBr(aq)} + \text{I}_2\text{(aq)}$

Ionic: $\text{Br}_2 + 2\text{I}^- \rightarrow 2\text{Br}^- + \text{I}_2$ .

Observation: orange Br₂ solution becomes DARKER BROWN (I₂ forming).

The full displacement matrix.

A more reactive halogen (higher in the group) displaces a less reactive halide from its salt.

Add halogen →	KCl(aq)	KBr(aq)	KI(aq)
Cl₂	No rxn	YES — brown (Br₂)	YES — brown (I₂)
Br₂	No rxn	No rxn	YES — brown (I₂)
I₂	No rxn	No rxn	No rxn

Redox analysis. Each displacement is a REDOX reaction:

The more reactive halogen is REDUCED (gains electrons → X₂ becomes X⁻).
The halide ion of the less reactive halogen is OXIDISED (loses electrons → X⁻ becomes X₂).
The more reactive halogen acts as the OXIDISING AGENT; the halide acts as the REDUCING AGENT.

In Cl₂ + 2I⁻ → 2Cl⁻ + I₂:

Cl₂ gains 2 e⁻ (one per Cl atom) → reduced.
2I⁻ each loses 1 e⁻ → oxidised.

Practical use. Displacement reactions are a way to confirm the reactivity order of halogens experimentally — they were used in the 19th century to establish the periodic trends.

More reactive halogen displaces less reactive halide.
Cl₂ displaces Br⁻ and I⁻; Br₂ displaces I⁻; I₂ displaces neither.
Solution colour: Br₂ brown/orange; I₂ brown (purple vapour).
Redox: more reactive halogen REDUCED (gains e⁻); halide OXIDISED (loses e⁻).

See the full worked example for group 7 →

Uses of halogens — benefits and risks (spec 2.65, 2.66)

Cl₂ water/bleach/PVC; Br₂ flame retardants; I₂ antiseptic/iodised salt. All with benefit/risk balance.

Chlorine — uses.

(1) Drinking water and swimming pool sterilisation (most important).

Chlorine kills bacteria, viruses and other pathogens. ~0.5-2 mg/L added to drinking water:

$\text{Cl}_2\text{(g)} + \text{H}_2\text{O(l)} \rightleftharpoons \text{HCl(aq)} + \text{HOCl(aq)}$

HOCl (chloric(I) acid / hypochlorous acid) is the active germicide.

Benefit: prevents water-borne diseases (cholera, typhoid, dysentery) — saves millions of lives per year. Provides 'residual disinfection' during distribution.
Risk: toxic at high concentrations (respiratory irritant — Cl₂ used as a WW1 chemical weapon). Reacts with organic matter to form trihalomethanes (e.g. chloroform) — suspected carcinogens. Industrial handling hazards.

(2) Bleach manufacture (sodium chlorate(I), NaClO).

$\text{Cl}_2 + 2\text{NaOH} \rightarrow \text{NaCl} + \text{NaClO} + \text{H}_2\text{O}$

Used in household bleach, laundry detergents, swimming pool sanitisers.

Benefit: whitens fabrics, cleans/disinfects surfaces.
Risk: corrosive; reacts dangerously with acidic cleaners → releases toxic Cl₂.

(3) PVC (polyvinyl chloride) plastic manufacture. Used for pipes, flooring, insulation. Risk: PVC manufacture and disposal can release dioxins; alternative plastics being explored.

(4) HCl manufacture, pharmaceutical synthesis.

Bromine — uses.

(1) Flame retardants (PBDEs — polybrominated diphenyl ethers added to fabrics, plastics, foams, electronics).

Benefit: reduces fire deaths and injuries; saves thousands of lives per year.
Risk: many brominated retardants are PERSISTENT, BIOACCUMULATIVE, may disrupt hormone function. Several banned in EU.

(2) Silver bromide (AgBr) in photographic film.

Benefit: black-and-white photography (largely historical now).
Risk: silver mining environmentally damaging.

Iodine — uses.

(1) Antiseptic (tincture of iodine).

Iodine dissolved in alcohol/water; applied to wounds and skin before surgery to kill bacteria.

Benefit: prevents wound infection.
Risk: skin staining; allergic reactions in some people.

(2) Iodised salt (KI added to NaCl).

Provides dietary iodine to prevent goitre (enlargement of thyroid) and cretinism (mental retardation in children of iodine-deficient mothers).

Benefit: universal salt iodisation could prevent millions of cognitive-development cases globally per WHO estimates.
Risk: excess iodine can cause iodine-induced hyperthyroidism (rare).

Fluorine — uses (mentioned briefly).

Fluoride in drinking water and toothpaste — strengthens tooth enamel (forms harder fluoroapatite). Risk: fluorosis (mottled teeth) at excessive doses.
PTFE (Teflon) — non-stick coating, made from fluorine-containing monomers.
CFCs (now banned) — used as refrigerants until damage to ozone layer was discovered.

Benefit/risk framework.

All halogens are toxic in high concentrations; the value to society depends on CONTROLLED, low-dose application. The 'right' level balances:

Benefit: lives saved, public health, fire safety, industrial uses.
Risk: toxicity, environmental persistence, allergic reactions, by-products (e.g. trihalomethanes from chlorination).

This is a key exam theme — examiners reward students who give a BALANCED discussion, not just a one-sided list of benefits or risks.

Cl₂: drinking water + pool disinfection (saves millions of lives); bleach manufacture; PVC.
Br₂: flame retardants (saves lives but environmentally persistent); photographic film (historical).
I₂: antiseptic (tincture) + iodised salt (prevents goitre).
F (compounds): fluoride in toothpaste / water; Teflon.
All halogens have benefit + risk — exam wants BALANCED discussion.

See the full worked example for group 7 →

Memorise this

Verbatim phrases and definitions Cambridge mark schemes credit.

Group 7 atoms: 7 outer electrons → gain 1 to form X⁻.
States at 20 °C: F₂ gas (pale yellow); Cl₂ gas (green-yellow); Br₂ liquid (red-brown); I₂ solid (sublimes to PURPLE vapour).
mp/bp INCREASE down → London dispersion / Van der Waals forces between molecules get stronger.
Reactivity DECREASES down: F > Cl > Br > I. Reason: larger atom + more shielding → weaker attraction for incoming electron.
Group 7 trend is OPPOSITE to Group 1 (gain vs lose electron).
Displacement: Cl₂ displaces Br⁻ and I⁻; Br₂ displaces I⁻; I₂ displaces none.
Cl₂ + 2KBr → 2KCl + Br₂ (orange/brown solution).
Cl₂ + H₂O ⇌ HCl + HOCl (water chlorination — HOCl kills bacteria).
Halogen tests with AgNO₃: AgCl white; AgBr cream; AgI yellow.

How it’s examined

Group 7 / halogens contribute 6-10 marks per session across the two papers. Paper 1 (4CH1/1C) items: (a) physical states and colours of halogens (3-4 marks); (b) explain reactivity trend in terms of atomic radius + shielding (4-5 marks); (c) predict and write equation for displacement reactions (4-5 marks). Paper 2 (4CH1/2C) items: (d) explain in detail why reactivity decreases down the group (5-6 marks — must mention BOTH atomic radius AND shielding); (e) discuss uses of halogens with benefit/risk balance (5-6 marks). Comparison with Group 1 reactivity (opposite trend) is a popular synoptic question. Mark-scheme phrase: 'as atomic radius increases AND shielding increases, attraction for an incoming electron decreases'.

Sources: Pearson Edexcel International GCSE Chemistry 4CH1 specification — Issue 4, September 2024 (valid for 2026 onwards); Pearson Edexcel 4CH1 Examiner Reports 2022-2024; 4CH1/1C May/Jun 2024 question paper and mark scheme; 4CH1/1C January 2024 question paper and mark scheme; 4CH1/2C May/Jun 2024 question paper and mark scheme; 4CH1/2C January 2024 question paper and mark scheme. Last reviewed 2026-05-12.

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Step-by-step worked examples — Group 7

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1Physical states and colours of the halogens (5 marks, Paper 1)
Core• Adapted from 4CH1/1C Jan 2024 Q7• physical properties, halogens, spec-2.58, spec-2.59
Question
State the colour and physical state at room temperature (20 °C) of each of these halogens: fluorine (F₂), chlorine (Cl₂), bromine (Br₂), iodine (I₂). State the trend in melting/boiling points down the group and explain it in terms of molecular structure. (5 marks)
Step-by-step solution
1. Step 1
  Fluorine F₂ (1 M). Very pale yellow GAS at room temperature. Extremely reactive — too reactive to handle outside specialist labs.
2. Step 2
  Chlorine Cl₂ (1 A). Pale green-yellow / yellow-green GAS at room temperature. Pungent, choking smell. The colour of the gas is visible in a gas jar.
3. Step 3
  Bromine Br₂ (1 A). Red-brown / orange-brown LIQUID at room temperature. The only Group 7 element that is liquid at 20 °C. Evaporates easily to give a brown vapour (volatile). Toxic — handle in a fume cupboard.
4. Step 4
  Iodine I₂ (1 B). Dark grey / black SOLID at room temperature; shiny, almost metallic in appearance. Sublimes on warming (changes directly from solid to a beautiful PURPLE vapour without melting). Less reactive than Cl₂ and Br₂.
5. Step 5
  Trend in mp / bp (1 B). Melting points AND boiling points INCREASE down the group. F₂ (bp −188 °C) < Cl₂ (bp −34 °C) < Br₂ (bp 59 °C) < I₂ (bp 184 °C). Explanation: halogens form diatomic MOLECULES (F₂, Cl₂, Br₂, I₂). Going down the group, the molecules contain more electrons → STRONGER INTERMOLECULAR FORCES (London dispersion / Van der Waals forces) — more energy needed to separate molecules. So mp and bp increase. Note: the COVALENT bond strength within X₂ molecules DECREASES down the group, but it's the intermolecular forces that matter for mp/bp.
Answer
F₂: very pale yellow GAS. Cl₂: pale green-yellow GAS. Br₂: red-brown LIQUID. I₂: dark grey SOLID (sublimes to purple vapour). Trend: mp and bp INCREASE down the group because halogens form diatomic molecules with stronger intermolecular forces as molecules get larger / more electrons.
Examiner tip
Edexcel 5-mark mark scheme: 1 each for four colours/states + 1 for mp/bp trend with explanation. Specific colour terms expected: 'green-yellow' for Cl₂, 'red-brown' for Br₂, 'dark grey' or 'purple' (in vapour form) for I₂. Vague 'yellow' or 'brown' may lose marks. The explanation MUST mention intermolecular forces — not just 'molecules bigger'.
2Explaining the reactivity trend in Group 7 (5 marks, Paper 1)
Core• Adapted from 4CH1/1C May/Jun 2024 Q8• reactivity trend, shielding, spec-2.60, spec-2.61
Question
State and explain the trend in reactivity of the halogens going DOWN Group 7. Refer to atomic structure (atomic radius, shielding) in your answer. (5 marks)
Step-by-step solution
1. Step 1
  State the trend (1 M). Reactivity DECREASES going down Group 7. F is the most reactive halogen; I is the least reactive (of the stable halogens). F > Cl > Br > I.
2. Step 2
  Halogen reactions involve gaining electrons (1 A). Halogens have 7 electrons in their outer shell. They react by GAINING ONE ELECTRON to form a stable −1 ion (X⁻), achieving a full outer shell (like a noble gas).
  $\text{X}_2 + 2e^- \rightarrow 2\text{X}^-$
3. Step 3
  Atomic radius increases (1 A). Going down the group, each halogen has ONE MORE complete electron shell than the one above it. So the atomic radius INCREASES: F < Cl < Br < I. The outer (7th-electron-containing) shell is FURTHER from the nucleus.
4. Step 4
  Shielding increases (1 B). With more inner shells of electrons (between the nucleus and the outer shell), there is more electron SHIELDING / screening of the nuclear charge from the outer shell. The effective nuclear charge experienced by the outer-shell electron (and any incoming electron) is reduced.
5. Step 5
  Conclusion — weaker attraction (1 B). Combination of (a) larger atomic radius (outer shell further from nucleus) and (b) more shielding (inner electrons block the nuclear pull) means that the attraction of the nucleus for an INCOMING electron is WEAKER. So it is HARDER to gain an electron → LESS REACTIVE going down. F is most reactive (small, little shielding); I is least reactive (large, more shielding).
Answer
Trend: reactivity DECREASES down Group 7 (F > Cl > Br > I). Reason: halogens react by gaining ONE electron to form X⁻. Going down the group: atomic radius INCREASES (more electron shells); shielding INCREASES (more inner electrons block nuclear attraction). So the attraction of the nucleus for an incoming electron is WEAKER → harder to gain an electron → less reactive.
Examiner tip
Edexcel 5-mark mark scheme: 1 for stating decreasing reactivity; 1 for halogen reactions involve gaining 1 electron; 1 for atomic radius increases (more shells); 1 for shielding increases; 1 for explaining why this makes gaining electron harder. Don't just say 'electrons further away' — must mention SHIELDING explicitly. Note this is OPPOSITE to the Group 1 trend (Group 1 reactivity increases down).
3Predicting and explaining a halogen displacement reaction (5 marks, Paper 1)
Core• Adapted from 4CH1/1C Jan 2024 Q9• displacement, halogens, spec-2.63, spec-2.64
Question
Chlorine gas is bubbled through a colourless solution of potassium iodide (KI). Predict and explain what would be observed. Write a BALANCED chemical equation and identify which species has been oxidised and which has been reduced. (5 marks)
Step-by-step solution
1. Step 1
  Prediction — Cl₂ is more reactive (1 M). Chlorine is HIGHER in Group 7 than iodine → MORE reactive (smaller atom, less shielding → gains electrons more easily). A more reactive halogen DISPLACES a less reactive halogen from its salt solution. So Cl₂ will displace I from KI.
2. Step 2
  Observation (1 A). The colourless solution turns BROWN / orange-brown as iodine (I₂) is formed and dissolves (iodine solutions are brown). On standing or warming, a purple vapour of I₂ may be visible above the solution. A black precipitate of I₂ may form if the iodide is concentrated.
3. Step 3
  Balanced equation (1 A). Cl₂ takes the place of I in the salt; I₂ is released.
  $\text{Cl}_2\text{(g)} + 2\text{KI(aq)} \rightarrow 2\text{KCl(aq)} + \text{I}_2\text{(aq)}$
4. Step 4
  Ionic equation (1 B). Spectator ions (K⁺) cancel.
  $\text{Cl}_2 + 2\text{I}^- \rightarrow 2\text{Cl}^- + \text{I}_2$
5. Step 5
  Oxidation / reduction (1 B). Each I⁻ LOSES an electron to become I (in I₂) — I⁻ is OXIDISED. Each Cl in Cl₂ GAINS an electron to become Cl⁻ — Cl₂ is REDUCED. This is a REDOX reaction: Cl₂ acts as the OXIDISING AGENT (it oxidises I⁻); I⁻ acts as the REDUCING AGENT.
Answer
Prediction: Cl₂ is more reactive than I (smaller atom, less shielding), so Cl₂ DISPLACES I from KI. Observation: colourless solution turns BROWN (I₂ formed); may see purple vapour. Equation: Cl₂ + 2KI → 2KCl + I₂. Ionic: Cl₂ + 2I⁻ → 2Cl⁻ + I₂. I⁻ oxidised (loses e⁻); Cl₂ reduced (gains e⁻).
Examiner tip
Edexcel 5-mark mark scheme: 1 for stating Cl more reactive (with reason); 1 for observation (brown colour from I₂); 1 for balanced symbol equation (with state symbols); 1 for ionic equation; 1 for correct redox labels (Cl₂ reduced, I⁻ oxidised). 'Brown' is the standard term — 'orange-brown' or 'red-brown' also accepted. Common error: saying the solution turns purple — purple is the VAPOUR, the solution is brown.
4Uses of chlorine — water treatment and bleach (6 marks, Paper 2)
Extended• Adapted from 4CH1/2C May/Jun 2024 Q3• uses of halogens, spec-2.65, spec-2.66
Question
Chlorine is added to drinking water in many countries. (a) State why chlorine is added to drinking water and write an equation for its reaction with water. (b) Discuss TWO benefits and TWO risks of chlorinating drinking water. (c) State ONE other major industrial use of chlorine, giving a balanced equation for the reaction. (6 marks)
Step-by-step solution
1. Step 1
  (a) Why chlorine in water (1 M). Chlorine is added as a DISINFECTANT — it kills bacteria, viruses and other pathogenic micro-organisms that cause water-borne diseases (cholera, typhoid, dysentery). Even tiny amounts (~0.5 mg/L) provide a 'residual' chlorine in pipes that prevents re-infection during distribution.
2. Step 2
  (a) Equation (1 A). Chlorine reacts with water in an equilibrium to form HCl and HOCl (chloric(I) acid, also called hypochlorous acid). The HOCl is the active germicide.
  $\text{Cl}_2\text{(g)} + \text{H}_2\text{O(l)} \rightleftharpoons \text{HCl(aq)} + \text{HOCl(aq)}$
3. Step 3
  (b) Two benefits (1 A). (i) Kills disease-causing bacteria and viruses → prevents water-borne diseases like cholera, typhoid, dysentery → saves MILLIONS of lives per year globally. (ii) Provides 'residual disinfection' — small amount remains in pipes throughout the water-distribution network so re-contamination during distribution is prevented. (Alternatives accepted: cheap; well-understood technology.)
4. Step 4
  (b) Two risks (1 B). (i) Chlorine is TOXIC in high concentrations — must be carefully dosed (~0.5-2 mg/L; higher levels can cause respiratory irritation or worse). (ii) Chlorine can react with organic compounds in water to form trihalomethanes (THMs), e.g. chloroform CHCl₃, which are SUSPECTED to be carcinogenic. (iii) Some people don't like the chemical taste or smell. (iv) Cl₂ gas itself is dangerous to handle for water-treatment workers.
5. Step 5
  (c) Other use — chlorine in bleach (1 B). Chlorine + cold dilute sodium hydroxide → sodium chlorate(I) (the active ingredient in household bleach):
  $\text{Cl}_2\text{(g)} + 2\text{NaOH(aq)} \rightarrow \text{NaCl(aq)} + \text{NaClO(aq)} + \text{H}_2\text{O(l)}$
6. Step 6
  (c) Other uses summary (1 B). Other valid uses (one accepted): (i) manufacture of HCl (hydrogen chloride); (ii) manufacture of PVC (polyvinyl chloride plastic); (iii) sterilising swimming pool water; (iv) production of polymers and pharmaceuticals. Each should come with a brief reason.
Answer
(a) Chlorine kills bacteria and viruses — disinfects drinking water. Cl₂ + H₂O ⇌ HCl + HOCl. (b) Benefits: prevents water-borne disease (cholera, typhoid); residual disinfection during distribution. Risks: Cl₂ toxic at high doses; can form trihalomethanes (suspected carcinogens). (c) Used to make bleach: Cl₂ + 2NaOH → NaCl + NaClO + H₂O. (Alternative valid uses: making HCl; making PVC; pool sterilisation.)
Examiner tip
Edexcel 6-mark mark scheme: 1 for disinfectant + reason; 1 for water equation; 1 each for two benefits + two risks; 1 for second industrial use with equation. Must give NUANCED benefit/risk analysis — examiners want a 'discuss' not just 'state'. The trihalomethanes / chloroform point is a high-tariff observation that distinguishes top answers.

Model Answers — Group 7

High-scoring sample answers for group 7 on the Pearson Edexcel IGCSE 4CH1 paper, with examiner-style notes mapping each response to the mark scheme and assessment objectives.

Question 1

4CH1/1C-style — physical properties of halogens6 marks

Describe and explain the trends in physical properties of the halogens (F₂, Cl₂, Br₂, I₂) going DOWN Group 7. Include: colour, physical state at room temperature, melting point, boiling point. Explain the trend in melting / boiling points in terms of structure and bonding. (6 marks)

Model answer

Physical properties of the halogens (at 20 °C, 1 atm).

Halogen	Formula	Colour	State (20 °C)	Melting point / °C	Boiling point / °C
Fluorine	F₂	Very pale yellow	GAS	−220	−188
Chlorine	Cl₂	Pale green-yellow	GAS	−101	−34
Bromine	Br₂	Red-brown	LIQUID	−7	59
Iodine	I₂	Dark grey (purple vapour)	SOLID	114	184
Astatine	At	Very dark / unknown	(Solid; radioactive, very rare)	~302	~337

Observations.

Colour darkens going down the group. F₂ is barely visible (very pale yellow); Cl₂ is pale green-yellow; Br₂ is red-brown; I₂ is dark grey/black solid. The vapour colours follow the same trend: F₂ very pale yellow, Cl₂ yellow-green, Br₂ orange-brown, I₂ purple.
Physical state changes from gas → liquid → solid down the group. F₂ and Cl₂ are gases (very low mp/bp). Br₂ is the only Group 7 element that is liquid at room temperature. I₂ is a solid that sublimes when warmed — i.e. goes directly from solid to purple vapour without melting (this is an iconic chemistry demonstration).
mp and bp INCREASE down the group. F₂ bp = −188 °C → I₂ bp = +184 °C — that's a range of over 370 °C.

Explanation — why mp and bp increase.

All Group 7 elements exist as SIMPLE MOLECULAR (covalent) substances of formula X₂. Two halogen atoms share one electron pair (a single covalent bond) to give a diatomic molecule. The X₂ molecules are held together in the bulk substance by weak INTERMOLECULAR FORCES (London dispersion forces / Van der Waals forces) — NOT by the strong covalent bonds (those are inside each molecule).

To melt or boil the substance, you must overcome the INTERMOLECULAR forces (not the covalent bonds — they stay intact).

Going down the group, each successive halogen molecule contains MORE ELECTRONS (F₂ = 18, Cl₂ = 34, Br₂ = 70, I₂ = 106 electrons). More electrons → stronger London dispersion forces because there is more electron cloud that can deform → larger temporary dipoles → stronger attractive force between neighbouring molecules.

So stronger intermolecular forces → more energy needed to overcome them → higher mp and higher bp. The covalent bond STRENGTH within X₂ molecules actually DECREASES down the group, but that does NOT affect mp/bp (only affects bond energies).

Why this matters. It's why F₂ and Cl₂ are gases (easily separated molecules) but I₂ is a solid at room temperature.

Why this scores

Edexcel 6-mark mark scheme: 1 each for four states/colours; 1 for trend in mp/bp; 1 for explanation in terms of intermolecular forces / more electrons. Always state 'INTERMOLECULAR forces (Van der Waals / London dispersion)' — DON'T say 'the bonds get stronger' (the COVALENT bonds within molecules actually weaken). Common loss-of-mark error.

Question 2
4CH1/2C-style — Group 7 reactivity trend6 marks
Explain in detail why fluorine is more reactive than chlorine, which is more reactive than bromine, which is more reactive than iodine. Refer to electron configuration, atomic radius, electron shielding and the strength of attraction for an incoming electron. (6 marks)
Model answer
Electron configurations.

Halogen Electron config Outer-shell electrons
F 2,7 7 in shell 2
Cl 2,8,7 7 in shell 3
Br 2,8,18,7 7 in shell 4
I 2,8,18,18,7 7 in shell 5

All halogens have 7 outer-shell electrons. They react by GAINING ONE more electron to fill the outer shell (achieving the noble gas configuration → very stable). The product is a halide anion X⁻ with a complete outer octet.

$\text{X}_2 + 2e^- \rightarrow 2\text{X}^-$

The 'reactivity' of a halogen is essentially its ability to gain an electron. The more easily it gains an electron, the more reactive it is.

Trend: reactivity DECREASES going down the group.

F is the most reactive halogen — it will steal electrons from almost ANYTHING (even noble gases, glass, and water in some conditions).

I is the least reactive of the stable halogens — it reacts slowly with hydrogen, slowly with metals, and the reactions are easily reversed.

Explanation — three factors working together.

(1) Atomic radius increases. Going down the group, each successive halogen has ONE MORE COMPLETE ELECTRON SHELL than the one above. So the atomic radius increases: F < Cl < Br < I. The outermost shell (where the incoming 8th electron must go) is FURTHER FROM THE NUCLEUS. Electrostatic attraction follows an inverse-square law (Coulomb's law) — distance matters a lot.

(2) Shielding increases. The inner-shell electrons (between the nucleus and the outer shell) SHIELD / SCREEN the outer shell from the full nuclear charge. The more inner shells there are, the more shielding. For F, there is only one inner shell (1s²). For I, there are four inner shells (1s² 2s²2p⁶ 3s²3p⁶3d¹⁰ 4s²4p⁶4d¹⁰).

(3) Effective nuclear charge experienced by the incoming electron. The combined effect of (1) and (2) is that the EFFECTIVE NUCLEAR CHARGE experienced by an incoming electron decreases going down the group. The 'electron cloud' between the nucleus and the bonding region is thicker; the pull from the nucleus is more attenuated. So the incoming electron is LESS STRONGLY ATTRACTED.

Conclusion.

For F: small atom, little shielding → STRONG attraction for an incoming electron → highly reactive.

For I: large atom, lots of shielding → WEAK attraction for an incoming electron → less reactive.

This explains why F₂ reacts violently with hydrogen even at low temperatures and in the dark, while I₂ reacts very slowly and reversibly (H₂ + I₂ ⇌ 2HI requires high temperatures and gives an equilibrium mixture).

Comparison with Group 1. Group 1 reactivity INCREASES down the group because Group 1 metals react by LOSING an outer electron — the further away that electron is and the more shielding there is, the EASIER it is to lose. Group 7 is the OPPOSITE: it's about GAINING an electron, so a smaller atom (less shielding) is more reactive.
Why this scores
Edexcel 6-mark mark scheme: 1 for stating decreasing reactivity; 1 for halogens gain 1 electron to form X⁻; 1 for atomic radius increase; 1 for shielding increase; 1 for outer-shell electron experiences weaker attraction; 1 for comparison/contrast with Group 1 reactivity trend (or alternative full mark). Must use the TERM 'shielding' explicitly — describing as 'inner electrons get in the way' is informal but acceptable. Don't say 'nuclear charge gets weaker' (it actually INCREASES) — the EFFECTIVE nuclear charge decreases due to shielding.

Halogen	Electron config	Outer-shell electrons
F	2,7	7 in shell 2
Cl	2,8,7	7 in shell 3
Br	2,8,18,7	7 in shell 4
I	2,8,18,18,7	7 in shell 5

Question 3

4CH1/2C-style — halogen displacement matrix6 marks

Complete the following table showing the results of mixing each halogen with each potassium halide solution. Where a reaction occurs, write 'reaction' and state the colour change; where no reaction occurs, write 'no reaction'. Then write balanced equations for ALL the reactions that occur. (6 marks)

Model answer

Reactivity order: Cl₂ > Br₂ > I₂. A halogen will displace ONLY the halides BELOW it (less reactive).

Displacement table.

	KCl(aq) (colourless)	KBr(aq) (colourless)	KI(aq) (colourless)
Cl₂(aq) (pale yellow)	No reaction	Reaction — solution turns orange/brown (Br₂ formed)	Reaction — solution turns brown (I₂ formed)
Br₂(aq) (orange)	No reaction	No reaction	Reaction — solution turns brown (I₂ formed)
I₂(aq) (brown)	No reaction	No reaction	No reaction

Logic. Cl₂ is the most reactive — it displaces Br⁻ (giving Br₂) and I⁻ (giving I₂). Br₂ is intermediate — it displaces only I⁻ (not Cl⁻, because Cl is more reactive than Br, so Br can't take Cl's place). I₂ is the least reactive — it can't displace any of them.

Balanced equations (three reactions occur):

(1) Chlorine + potassium bromide:

$\text{Cl}_2\text{(g)} + 2\text{KBr(aq)} \rightarrow 2\text{KCl(aq)} + \text{Br}_2\text{(aq)}$

Ionic: $\text{Cl}_2 + 2\text{Br}^- \rightarrow 2\text{Cl}^- + \text{Br}_2$ .

Observation: colourless KBr solution turns ORANGE / RED-BROWN as Br₂ is formed.

(2) Chlorine + potassium iodide:

$\text{Cl}_2\text{(g)} + 2\text{KI(aq)} \rightarrow 2\text{KCl(aq)} + \text{I}_2\text{(aq)}$

Ionic: $\text{Cl}_2 + 2\text{I}^- \rightarrow 2\text{Cl}^- + \text{I}_2$ .

Observation: colourless KI solution turns BROWN as I₂ is formed; may see PURPLE vapour above the solution.

(3) Bromine + potassium iodide:

$\text{Br}_2\text{(aq)} + 2\text{KI(aq)} \rightarrow 2\text{KBr(aq)} + \text{I}_2\text{(aq)}$

Ionic: $\text{Br}_2 + 2\text{I}^- \rightarrow 2\text{Br}^- + \text{I}_2$ .

Observation: orange Br₂ solution turns DARKER BROWN as I₂ is formed (already dark colour, so change is subtle but visible).

Redox interpretation. In each displacement, the more reactive halogen is REDUCED (gains e⁻ → goes from X₂ to X⁻); the halide ion of the less reactive halogen is OXIDISED (loses e⁻ → goes from X⁻ to X₂). The more reactive halogen acts as the OXIDISING AGENT.

Why this scores

Edexcel 6-mark mark scheme: 3 for correctly identifying which combinations react vs no reaction (with colour changes); 3 for three balanced equations (ionic acceptable, with state symbols). Common error: predicting I₂ + KCl gives a reaction (it doesn't — I is less reactive than Cl). The reactivity order Cl > Br > I is the simple rule.

Question 4
4CH1/2C-style — uses of halogens with benefit/risk6 marks
Discuss the use of halogens in modern society. Describe TWO uses of chlorine, ONE use of bromine and ONE use of iodine. For EACH use, identify ONE benefit AND ONE risk. (6 marks)
Model answer
Chlorine — Use 1: Drinking water and swimming pool sterilisation.

Process: chlorine is added to water at ~0.5-2 mg/L. It reacts with water to form HOCl (hypochlorous acid), which kills bacteria and viruses.

Benefit: Prevents water-borne diseases like cholera, typhoid, dysentery — historically responsible for saving MILLIONS of lives. Provides residual disinfection during distribution.

Risk: At high doses, chlorine is TOXIC (respiratory irritant; the gas was used as a chemical weapon in WW1). Can react with organic matter in water to form trihalomethanes (e.g. CHCl₃), which are suspected carcinogens. Workers handling Cl₂ must follow strict safety precautions.

Chlorine — Use 2: Manufacture of bleach (sodium chlorate(I), NaClO).

Process: Cl₂ + 2NaOH → NaCl + NaClO + H₂O. The NaClO is the active bleaching agent used in household bleach, laundry detergent and pool sanitisers.

Benefit: Whitens fabrics, cleans surfaces, kills bacteria on surfaces (hygiene). Used in paper manufacture and textile bleaching industry.

Risk: NaClO is corrosive (eye and skin damage). When mixed with acidic cleaners (e.g. toilet cleaner), it releases toxic Cl₂ gas. Bleach overuse releases chlorine compounds into waterways → environmental damage.

Bromine — Use: Flame retardants for fabrics, electronics, foam.

Process: Brominated organic compounds (e.g. PBDEs — polybrominated diphenyl ethers) are incorporated into plastics, fabrics, foam mattresses, electronic circuit boards. The bromine atoms interfere with combustion reactions → less easily catches fire.

Benefit: Reduces fire-related deaths and injuries. Allows TVs, computers and furniture to meet safety standards. Estimated to save thousands of lives per year.

Risk: Many brominated flame retardants are PERSISTENT, BIOACCUMULATIVE TOXIC chemicals — they don't break down in the environment, build up in animal tissue, and may disrupt thyroid function and child development. Many PBDEs have been BANNED in the EU and elsewhere.

(Alternative bromine use: silver bromide AgBr in photographic film/paper — light-sensitive, makes black-and-white photography possible. Benefit: photography records history; Risk: silver mining is environmentally damaging. With digital photography this use has declined sharply.)

Iodine — Use: Antiseptic and dietary supplement.

Process: Iodine solution (tincture of iodine — I₂ dissolved in alcohol) or iodide salts (KI) applied to skin/wounds to kill micro-organisms. Iodised table salt: small amounts of KI added to NaCl to ensure dietary iodine intake.

Benefit: Prevents wound infection (antiseptic). Iodised salt prevents iodine-deficiency disorders — particularly goitre (enlargement of the thyroid gland) and cretinism (mental retardation in children). The WHO estimates that universal salt iodisation could prevent millions of cases of cognitive impairment globally.

Risk: Excess iodine can cause iodine-induced hyperthyroidism (too much thyroid hormone). Iodine solutions stain skin and fabric. Some individuals are allergic to iodine (anaphylaxis risk).

Summary — benefit/risk analysis. All halogens are toxic in high concentrations; their usefulness depends on controlled, low-dose application. Society generally accepts the small risks (cancer, environmental persistence) because the benefits (public health, fire safety) are large. Continual research seeks safer alternatives where possible (e.g. ozone or UV disinfection of water; non-brominated flame retardants).
Why this scores
Edexcel 6-mark mark scheme: 1 each for four valid uses + 1 for at least four benefit/risk pairs (i.e. balanced discussion). Students often state uses without addressing the RISK side. The 'tincture of iodine' / 'iodised salt' answer for iodine is classic.

Question 5

4CH1/2C-style — Group 1 vs Group 7 comparison5 marks

Compare and contrast the reactivity trends in Group 1 (alkali metals) and Group 7 (halogens). Explain why the trends go in OPPOSITE directions, referring to the type of reaction each group undergoes and the role of atomic radius / electron shielding. (5 marks)

Model answer

Trends — opposite directions.

Group 1 (Li, Na, K, Rb, Cs): Reactivity INCREASES going DOWN the group. Li (slow with water) → Na (vigorous) → K (very vigorous, ignites H₂) → Rb / Cs (explosive). The most reactive Group 1 metals are at the BOTTOM.
Group 7 (F, Cl, Br, I): Reactivity DECREASES going DOWN the group. F (extremely reactive) → Cl (very reactive) → Br (less reactive) → I (least reactive of the stable halogens). The most reactive Group 7 element is at the TOP.

Reason — different type of reaction.

The two groups react by OPPOSITE electron-transfer mechanisms:

Group 1: each atom has ONE outer-shell electron. To react, it must LOSE that electron to form a +1 cation (Na → Na⁺ + e⁻). The EASIER it is to lose the electron, the MORE REACTIVE the element.
Group 7: each atom has SEVEN outer-shell electrons. To react, it must GAIN ONE more electron to form a −1 anion (½Cl₂ + e⁻ → Cl⁻). The EASIER it is to gain an electron, the MORE REACTIVE the element.

Atomic radius and shielding — same physics, opposite effect.

In BOTH groups, going down:

Atomic radius INCREASES (more shells).
Inner-shell shielding INCREASES.
The effective attraction of the nucleus on the OUTER ELECTRON SHELL DECREASES.

But the EFFECT of decreased attraction is OPPOSITE for the two groups:

For Group 1 (losing an electron): Weaker nuclear attraction means the outer electron is EASIER TO LOSE. So bigger atom + more shielding = more reactive (going down → more reactive).

For Group 7 (gaining an electron): Weaker nuclear attraction means an INCOMING electron is LESS STRONGLY PULLED IN. So bigger atom + more shielding = HARDER to gain an electron = LESS reactive (going down → less reactive).

Memory tip.

Group 1: "LOSING" — bigger + further → easier → more reactive down.

Group 7: "GAINING" — bigger + further → harder → LESS reactive down.

Summary table.

Feature	Group 1 (Alkali metals)	Group 7 (Halogens)
Outer-shell electrons	1	7
Reaction type	LOSE 1 electron	GAIN 1 electron
Ion formed	M⁺	X⁻
Reactivity trend down group	INCREASES	DECREASES
Most reactive	Cs (at bottom)	F (at top)
Reason	Outer electron easier to lose with more shielding	Incoming electron harder to attract with more shielding

Why this scores

Edexcel 5-mark mark scheme: 1 for stating opposite directions; 1 for Group 1 reaction (lose electron); 1 for Group 7 reaction (gain electron); 1 for explaining shielding effect for Group 1; 1 for explaining shielding effect for Group 7 (opposite outcome). This is a classic 'synoptic' question linking Topic 1 (Periodic Table) with Topic 2 (Groups 1 and 7).

Key Definitions and Keywords — Group 7

Definitions to memorise and the exact keywords mark schemes credit for group 7 answers — sharpened from recent examiner reports for the 2026 Pearson Edexcel IGCSE 4CH1 sitting.

Halogen
Examiner keyword
An element in Group 7 of the periodic table: F, Cl, Br, I, At (and the synthetic Ts). All have 7 outer-shell electrons; all exist as diatomic molecules X₂; all form X⁻ ions in salts.
Halide / halide ion
Examiner keyword
A negative ion X⁻ formed when a halogen gains one electron. Examples: F⁻ (fluoride), Cl⁻ (chloride), Br⁻ (bromide), I⁻ (iodide). Found in ionic salts like NaCl, KBr.
Displacement reaction (Group 7)
Examiner keyword
A reaction in which a more reactive halogen takes the place of a less reactive halogen in a salt. e.g. Cl₂ + 2KBr → 2KCl + Br₂. The halogen higher in Group 7 is more reactive; it displaces the less reactive halogen from its halide.
Sublimation (iodine)
Examiner keyword
Change of state directly from solid to gas (or vapour) without passing through the liquid state. Iodine sublimes on gentle heating: solid I₂ (grey) → purple vapour. Used as a classic chemistry demo.
Diatomic molecule
Examiner keyword
A molecule containing two atoms. Halogens exist as diatomic molecules (F₂, Cl₂, Br₂, I₂) because they share one pair of electrons in a single covalent bond, achieving a stable octet.
Shielding (electron screening)
Examiner keyword
The reduction in the attraction of the nucleus on outer electrons (or incoming electrons) caused by inner electron shells. More inner shells → more shielding → weaker effective nuclear attraction.
Redox reaction (Group 7 displacement)
Examiner keyword
A reaction involving both reduction and oxidation. In halogen displacement (e.g. Cl₂ + 2I⁻ → 2Cl⁻ + I₂): Cl₂ is REDUCED (gains e⁻); I⁻ is OXIDISED (loses e⁻). The more reactive halogen is the OXIDISING AGENT.
Disinfectant (chlorine in water)
A substance that kills micro-organisms (bacteria, viruses, fungi). Chlorine is the most widely used disinfectant for drinking water and swimming pools, providing both immediate kill and residual protection.
Flame retardant (brominated)
A chemical added to materials (fabrics, plastics, foams, electronics) to slow or prevent ignition and reduce fire spread. Brominated flame retardants are common; some are persistent in the environment.
Tincture of iodine
A solution of iodine (I₂) and potassium iodide (KI) in alcohol/water. Used as a skin antiseptic to disinfect wounds — kills bacteria on contact.

Common Mistakes and Misconceptions — Group 7

The traps other students keep falling into on group 7 questions — taken from recent Pearson Edexcel IGCSE 4CH1 examiner reports and mark schemes — and how to avoid them.

✕Stating that Group 7 reactivity INCREASES going down (like Group 1)
4CH1 Examiner Reports 2022-2024
Why it happens
Generalising from Group 1.
How to avoid it
Group 1 LOSES electrons → bigger atoms easier to lose → more reactive down. Group 7 GAINS electrons → bigger atoms harder to attract → LESS reactive down. The trends go in OPPOSITE directions. Memorise: 'gaining = harder when shielded'.
✕Saying the iodide displacement reaction turns the solution PURPLE
Why it happens
Confusing the vapour colour with the solution colour.
How to avoid it
I₂ in solution (especially aqueous) is BROWN. I₂ as a VAPOUR is purple. Iodine dissolved in non-polar solvents (cyclohexane) is purple. State 'brown solution' for an aqueous displacement.
✕Saying 'bonds get stronger down Group 7, so mp/bp increase'
4CH1 Examiner Reports 2022-2024
Why it happens
Conflating covalent bonds with intermolecular forces.
How to avoid it
Within X₂ molecules, the covalent bonds actually WEAKEN going down (longer bonds). But mp/bp depends on INTERMOLECULAR forces between molecules — these increase down the group because larger molecules have more electrons → stronger London dispersion forces.
✕Writing I₂ + KCl → KI + Cl₂ (i.e. I₂ displacing Cl)
Why it happens
Trying to balance any pair of halogen + halide.
How to avoid it
Only a MORE REACTIVE halogen can displace a less reactive one. Reactivity order: Cl > Br > I. So Cl₂ + KBr/KI works; Br₂ + KI works; I₂ + KCl / KBr does NOT (no reaction).
✕Confusing the chlorine test (bleaches damp litmus) with the ammonia test (damp red litmus turns blue)
Why it happens
Both involve damp litmus paper.
How to avoid it
Cl₂: turns damp BLUE litmus RED, then BLEACHES it (red → white). NH₃: turns damp RED litmus BLUE. Cl₂ is acidic + an oxidiser (bleaches); NH₃ is alkaline.
✕Writing Cl₂ + H₂O → HCl + HOCl with a single arrow (forward only)
Why it happens
Treating it as a normal reaction.
How to avoid it
The reaction of chlorine with water is REVERSIBLE — use the equilibrium arrows ⇌: $\text{Cl}_2 + \text{H}_2\text{O} \rightleftharpoons \text{HCl} + \text{HOCl}$ . The presence of HOCl is what kills bacteria.

Practice questions

Exam-style questions with step-by-step worked solutions. Try one before checking the method.

Past paper style quiz

Get a report showing which sub-topics you've nailed and which ones still need work.

Group 7 — frequently asked questions

The things students keep getting wrong in this sub-topic, answered.

Group 7

Short Study Notes in the form of Slides

Short Study Notes — Group 7

Detailed Notes

What you’ll learn

How it’s examined

1Physical states and colours of the halogens (5 marks, Paper 1)

2Explaining the reactivity trend in Group 7 (5 marks, Paper 1)

3Predicting and explaining a halogen displacement reaction (5 marks, Paper 1)

4Uses of chlorine — water treatment and bleach (6 marks, Paper 2)

Halogen

Halide / halide ion

Displacement reaction (Group 7)

Sublimation (iodine)

Diatomic molecule

Shielding (electron screening)

Redox reaction (Group 7 displacement)

Disinfectant (chlorine in water)

Flame retardant (brominated)

Tincture of iodine

✕Stating that Group 7 reactivity INCREASES going down (like Group 1)

✕Saying the iodide displacement reaction turns the solution PURPLE

✕Saying 'bonds get stronger down Group 7, so mp/bp increase'

✕Writing I₂ + KCl → KI + Cl₂ (i.e. I₂ displacing Cl)

✕Confusing the chlorine test (bleaches damp litmus) with the ammonia test (damp red litmus turns blue)

✕Writing Cl₂ + H₂O → HCl + HOCl with a single arrow (forward only)

Practice questions

Past paper style quiz

Assess your understanding

Group 7 — frequently asked questions