A-Level Chemistry — Iodine and its Properties (the glamorous purple one)

A-Level Chemistry — Iodine and its Properties (the glamorous purple one)

If chlorine is the loud, attention-seeking halogen and bromine is the moody one, iodine is the dramatic artist: it sits there looking like a dull grey solid… then quietly produces a purple vapour that makes everyone in the lab suddenly pay attention.

This post is a tidy A-Level tour of iodine’s key physical and chemical properties, plus the bits exam questions love to poke.

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1) Where iodine sits and what that tells you

Iodine is a Group 17 (halogen) element. Like the others, it exists as diatomic molecules: I₂.

Trends down Group 17 (F₂ → Cl₂ → Br₂ → I₂):

• Melting/boiling points increase (bigger molecules → stronger intermolecular forces)

• Colour gets darker

• Reactivity decreases (harder to gain an electron as atoms get larger and shielding increases)

So iodine is less reactive than chlorine and bromine, but it still does plenty of chemistry.

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2) Physical properties you can actually see

Appearance and state

• Grey/black crystalline solid at room temperature

• Produces a purple vapour when warmed

Sublimation (a favourite classroom moment)

Iodine can sublime: solid → gas without becoming a liquid first (under normal lab conditions).
That purple vapour is iodine gas (still I₂ molecules).

Why does iodine sublime easily?
Inside each I₂ molecule, the I–I covalent bond is strong. But between molecules the attractions are only London dispersion forces—and warming supplies enough energy to overcome those.

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3) Solubility: iodine is fussy

Iodine is non-polar overall, so:

• Low solubility in water

• Much more soluble in non-polar solvents (e.g. cyclohexane / hexane)

Colour clue (classic practical / exam):

• In water: brown/yellow-brown (a mixture of I₂ and some I₃⁻ if iodide is present)

• In cyclohexane: vivid purple (I₂ really showing off)

If you add iodide ions (from KI), iodine forms triiodide:

$$\mathrm{I_2(aq) + I^-(aq) \rightleftharpoons I_3^-(aq)}$$

This helps “pull” iodine into aqueous solution.

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4) Simple redox behaviour (the exam engine)

Iodine is an oxidising agent (it accepts electrons), but weaker than chlorine or bromine.

Reduction half-equation:

$$\mathrm{I_2 + 2e^- \rightarrow 2I^-}$$

That’s the backbone of loads of questions: titrations, displacement, and redox calculations.

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5) Displacement reactions: who can bully whom?

A more reactive halogen displaces a less reactive one from its halide.

So:

• Chlorine displaces iodide:

$$\mathrm{Cl_2 + 2I^- \rightarrow 2Cl^- + I_2}$$

• Bromine displaces iodide:

$$\mathrm{Br_2 + 2I^- \rightarrow 2Br^- + I_2}$$

• But iodine does not displace bromide or chloride.

Observation: formation of iodine gives brown solution (or purple in organic layer).

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6) Iodine as a test reagent (two big ones)

(A) Starch test (iconic)

Iodine forms a deep blue-black complex with starch (specifically amylose helices).
This is used to detect iodine, and also in clock reactions where iodine appears/disappears.

(B) Iodometry / iodimetry (A-Level titration territory)

• Iodine can be titrated with thiosulfate:

$$\mathrm{I_2 + 2S_2O_3^{2-} \rightarrow 2I^- + S_4O_6^{2-}}$$

Starch indicator is added near the end point (when solution is pale) for a sharp finish.

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7) A quick word on safety and good lab habits

• Iodine vapour is irritant — use a fume cupboard for heating/sublimation demos.

• Avoid skin contact (stains and irritates).

• Use small quantities: iodine is “spectacular per gram”.

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Mini “exam-style” check questions (with answers)

1. Why are iodine’s melting and boiling points higher than chlorine’s?
   Because iodine molecules are larger with more electrons, so London dispersion forces are stronger.

2. Write the ionic equation for chlorine reacting with potassium iodide solution.

$$\mathrm{Cl_2 + 2I^- \rightarrow 2Cl^- + I_2}$$

3. Why does iodine appear purple in cyclohexane but brown in water?
   Cyclohexane dissolves molecular I₂ well (purple). In water iodine is poorly soluble and can form I₃⁻ in the presence of I⁻, giving brown shades.

4. State the half-equation for iodine being reduced.

$$\mathrm{I_2 + 2e^- \rightarrow 2I^-}$$