Halogens: Trends and Oxidising Ability

Spec Mapping — OCR H432 Module 3.1.3 — Group 7 (the halogens), covering the appearance and physical state of F₂, Cl₂, Br₂ and I₂ at room temperature, the trend in boiling point down the group (explained by London dispersion forces), the trend in electronegativity, the trend in oxidising ability (F₂ > Cl₂ > Br₂ > I₂), the use of halogen-halide displacement reactions to demonstrate the oxidising-ability trend, and the cyclohexane extraction technique that distinguishes Cl₂, Br₂ and I₂ in solution (refer to the official OCR H432 specification document for exact wording).

The halogens (Group 7) provide OCR's canonical example of a "down-the-group decrease in reactivity". Where Group 2 metals become more reactive down the group (because they need to lose electrons, and electrons are easier to lose from larger, more shielded atoms), the halogens become less reactive down the group (because they need to gain electrons, and electrons are harder to attract into larger, more shielded outer shells). This mirror-image logic is the conceptual heart of OCR Module 3.1.3. The experimental evidence comes from displacement reactions — a more-reactive halogen displaces a less-reactive halide from solution, with the colour change immediately visible. The cyclohexane extraction step then transfers the displaced halogen into the organic layer where its colour is far more distinctive (especially the unmistakable violet of I₂ in cyclohexane). This lesson develops fluency in writing ionic and half-equations for these displacements, identifying oxidising and reducing agents, and explaining the underlying atomic-structure trend.

Key Definitions:

• Oxidising agent — a species that gains electrons in a redox reaction and is itself reduced.
• Reducing agent — a species that loses electrons and is itself oxidised.
• Displacement reaction — a reaction in which a more-reactive element displaces a less-reactive one from its compound.
• Electronegativity — the ability of an atom in a molecule to attract the bonding pair of electrons in a covalent bond (Pauling scale).

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The halogens — Group 17

The halogens are the Group 17 (older "Group 7") non-metals: F, Cl, Br, I, and the radioactive At (astatine, rarely studied; Ts/tennessine is even rarer). All have outer configuration $ns^2\, np^5$ns2np5 — seven outer electrons — and characteristically form 1− ions ("halides") by gaining one electron to complete a noble-gas-like $ns^2\, np^6$ns2np6 outer shell.

Element	Z	Outer config	State at 25 °C	Colour (gas)	Colour (aqueous)
F₂	9	2s² 2p⁵	Pale yellow gas	Pale yellow	Reacts with water (cannot test in lab)
Cl₂	17	3s² 3p⁵	Yellow-green gas	Yellow-green	Pale green (or pale yellow at low conc.)
Br₂	35	4s² 4p⁵	Dark red liquid	Red-brown vapour	Orange / yellow-brown
I₂	53	5s² 5p⁵	Grey-black solid	Violet vapour	Brown (forms triiodide I₃⁻ in KI)
At₂	85	6s² 6p⁵	Black solid (radioactive)	—	—

Because every halogen atom is one electron short of a noble-gas configuration, they are highly reactive non-metals and the most aggressive oxidising agents in chemistry. F₂ is so reactive that it reacts violently with water and with most metals on contact at room temperature; F₂ is therefore never demonstrated in school labs, and the practical work in this module starts with Cl₂.

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Trend in boiling point

Halogen molecules (X₂) are simple molecular with only London dispersion forces between them. Going down the group, the number of electrons per molecule rises (F₂ has 18, Cl₂ has 34, Br₂ has 70, I₂ has 106), so the polarisability of the electron cloud rises, so the London-force strength rises, so more energy is needed to separate molecules from one another → higher boiling point.

Halogen	Electrons / molecule	Boiling point / °C	State at 25 °C
F₂	18	−188	Gas
Cl₂	34	−34	Gas
Br₂	70	59	Liquid
I₂	106	184	Solid

This is why F₂ and Cl₂ are gases at room temperature, Br₂ is a liquid, and I₂ is a solid. The same principle explains the b.p. trend in the noble gases (He → Rn) and in simple molecular elements generally.

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Trend in electronegativity

Electronegativity (Pauling's measure) is the ability of an atom to attract the bonding pair of electrons in a covalent bond. The halogens are among the most electronegative elements:

Element	Pauling electronegativity
F	4.0 (highest of any element)
Cl	3.0
Br	2.8
I	2.5

Electronegativity decreases down Group 7 because:

• Atomic radius increases down the group → bonding electrons sit further from the nucleus.
• Shielding by inner electrons increases.
• Although the nuclear charge rises, the effects of distance and shielding dominate.
• The nucleus therefore has a weaker pull on the bonding pair → lower electronegativity.

This is the same atomic-structure logic used for ionisation-energy trends — once you internalise the "four factors", every periodic trend in halogen chemistry follows.

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Trend in oxidising ability

An oxidising agent gains electrons (is itself reduced). For a halogen, the operative half-equation is:

$\text{X}_2 + 2\text{e}^- \rightarrow 2\text{X}^-$X2​+2e−→2X−

The halogen is the oxidising agent and the halide ion (X⁻) is its reduced form. Oxidising ability decreases down Group 7 (F₂ > Cl₂ > Br₂ > I₂) because:

• Down the group, the outer shell of the halogen atom is further from the nucleus (larger radius).
• Shielding by inner shells increases.
• An incoming electron is therefore held less strongly (less stabilising energy released on electron capture — the electron affinity falls in magnitude).
• The halogen atom is therefore less ready to gain an electron → weaker oxidising agent.

Fluorine is the strongest oxidising agent in Group 7 (and one of the strongest of any element); iodine is the weakest halogen oxidising agent you will meet at A-Level.

This trend mirrors the Group 2 picture: in Group 2 the metals lose electrons, and electrons are progressively easier to lose down the group → reactivity increases. In Group 7 the halogens gain electrons, and electrons are progressively harder to gain down the group → reactivity decreases. Same physics (atomic radius and shielding), opposite consequence for reactivity depending on which direction the electron moves.

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Halogen-halide displacement reactions

The key experimental evidence for the oxidising-ability trend is displacement: a halogen will displace a less-reactive halide from solution because it is the stronger oxidising agent.

flowchart TD
  A[Cl2 added to KBr aq] --> B[Cl2 oxidises Br- to Br2]
  B --> C[Solution turns orange brown]
  D[Cl2 added to KI aq] --> E[Cl2 oxidises I- to I2]
  E --> F[Solution turns brown/black]
  G[Br2 added to KI aq] --> H[Br2 oxidises I- to I2]
  H --> I[Solution turns brown]
  J[Br2 added to KCl aq] --> K[No reaction Br2 weaker than Cl2]
  L[I2 added to KCl or KBr] --> M[No reaction I2 weakest]

Summary table of observations (aqueous, no organic solvent)

Mixture	Ionic equation	Observation in water
Cl₂(aq) + KBr(aq)	$\text{Cl}_2 + 2\text{Br}^- \rightarrow 2\text{Cl}^- + \text{Br}_2$Cl2​+2Br−→2Cl−+Br2​	Yellow → orange (Br₂ released)
Cl₂(aq) + KI(aq)	$\text{Cl}_2 + 2\text{I}^- \rightarrow 2\text{Cl}^- + \text{I}_2$Cl2​+2I−→2Cl−+I2​	Yellow → brown / black (I₂ released; may see solid)
Br₂(aq) + KI(aq)	$\text{Br}_2 + 2\text{I}^- \rightarrow 2\text{Br}^- + \text{I}_2$Br2​+2I−→2Br−+I2​	Orange → brown (I₂ released)
Br₂(aq) + KCl(aq)	No reaction	No change (Br₂ pale orange retained)
I₂(aq) + KCl(aq)	No reaction	No change
I₂(aq) + KBr(aq)	No reaction	No change

The unambiguous experimental finding: only the more reactive halogen displaces (Cl₂ displaces Br⁻ and I⁻; Br₂ displaces only I⁻; I₂ displaces neither). This empirically establishes the reactivity ranking F > Cl > Br > I and directly confirms the down-the-group decrease in oxidising power.

Cyclohexane extraction — sharpening the colour test

The colours in pure water can be ambiguous because Cl₂, Br₂ and I₂ can all look yellowish in dilute solution. Adding a few cm³ of cyclohexane (a non-polar organic solvent) and shaking the stoppered test tube extracts the dissolved halogen into the upper (less-dense) cyclohexane layer, where its colour is more distinctive:

Halogen	Colour in water	Colour in cyclohexane layer
Cl₂	Pale green / yellow	Pale green / almost colourless
Br₂	Orange / yellow-brown	Orange / yellow-brown
I₂	Brown (in KI: brown due to I₃⁻)	Violet / purple (very distinctive)

The violet I₂ in cyclohexane is the most unambiguous diagnostic in this module — a classic OCR exam image. The colour shift for I₂ (brown in water → violet in cyclohexane) arises because iodine in water forms charge-transfer complexes with water (or with I⁻ to give I₃⁻), whereas in pure non-polar cyclohexane the I₂ molecule retains its "free" violet colour.

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Displacement as redox: half-equations

Take the displacement of Br⁻ by Cl₂: $\text{Cl}_2\text{(aq)} + 2\text{KBr(aq)} \rightarrow 2\text{KCl(aq)} + \text{Br}_2\text{(aq)}$Cl2​(aq)+2KBr(aq)→2KCl(aq)+Br2​(aq)

The ionic equation (potassium is a spectator): $\text{Cl}_2\text{(aq)} + 2\text{Br}^-\text{(aq)} \rightarrow 2\text{Cl}^-\text{(aq)} + \text{Br}_2\text{(aq)}$Cl2​(aq)+2Br−(aq)→2Cl−(aq)+Br2​(aq)

Split into half-equations:

$\text{Cl}_2\text{(aq)} + 2\text{e}^- \rightarrow 2\text{Cl}^-\text{(aq)} \quad \text{(reduction — Cl}_2 \text{ is the oxidising agent)}$Cl2​(aq)+2e−→2Cl−(aq)(reduction — Cl2​ is the oxidising agent)

$2\text{Br}^-\text{(aq)} \rightarrow \text{Br}_2\text{(aq)} + 2\text{e}^- \quad \text{(oxidation — Br}^- \text{ is the reducing agent)}$2Br−(aq)→Br2​(aq)+2e−(oxidation — Br− is the reducing agent)

Oxidation-number changes: Cl: 0 → −1 (reduced), Br: −1 → 0 (oxidised). This is a classic redox reaction in which Cl₂ has been reduced to Cl⁻ by accepting electrons from the Br⁻ ion.

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Worked examples

Example 1 — chlorine and bromide

Q: A student adds chlorine water to a colourless solution of potassium bromide. She then adds a few cm³ of cyclohexane, stoppers and shakes. Describe the observations and write the ionic and half-equations.

Answer: The aqueous layer turns orange/yellow-brown (Br₂ released). On adding cyclohexane and shaking, the orange/yellow-brown Br₂ extracts into the upper cyclohexane layer, giving an orange/yellow-brown organic layer.

Ionic equation: $\text{Cl}_2\text{(aq)} + 2\text{Br}^-\text{(aq)} \rightarrow 2\text{Cl}^-\text{(aq)} + \text{Br}_2\text{(aq)}$Cl2​(aq)+2Br−(aq)→2Cl−(aq)+Br2​(aq)

Half-equations: $\text{Cl}_2 + 2\text{e}^- \rightarrow 2\text{Cl}^- \quad \text{(reduction; Cl}_2 \text{ is the oxidising agent)}$Cl2​+2e−→2Cl−(reduction; Cl2​ is the oxidising agent) $2\text{Br}^- \rightarrow \text{Br}_2 + 2\text{e}^- \quad \text{(oxidation; Br}^- \text{ is the reducing agent)}$2Br−→Br2​+2e−(oxidation; Br− is the reducing agent)

Cl: oxidation number 0 → −1; Br: oxidation number −1 → 0.

Example 2 — chlorine and iodide

Q: Repeat Example 1 with potassium iodide in place of potassium bromide.

Answer: Aqueous layer turns brown/black (I₂ released; if KI is concentrated enough, brown solution due to I₃⁻; solid I₂ may even precipitate as black flakes). After adding cyclohexane and shaking, the upper organic layer shows the distinctive violet/purple colour of I₂.

Ionic equation: $\text{Cl}_2\text{(aq)} + 2\text{I}^-\text{(aq)} \rightarrow 2\text{Cl}^-\text{(aq)} + \text{I}_2\text{(aq)}$Cl2​(aq)+2I−(aq)→2Cl−(aq)+I2​(aq)

Half-equations as above with I in place of Br. Cl: 0 → −1; I: −1 → 0.

Example 3 — bromine and chloride (no reaction)

Q: Bromine water is added to a solution of sodium chloride. State the observations and explain why no reaction occurs.

Answer: No reaction. The bromine water retains its orange/yellow-brown colour; no new colour develops. Reason: Br₂ is a weaker oxidising agent than Cl₂. To produce a reaction, Br₂ would have to oxidise Cl⁻ to Cl₂ — but this requires Br₂ to be the stronger oxidising agent of the two, which it is not. The reaction $\text{Br}_2 + 2\text{Cl}^- \nrightarrow 2\text{Br}^- + \text{Cl}_2$Br2​+2Cl−↛2Br−+Cl2​ does not proceed in either direction; chlorine is "higher up" in the displacement series.