IB Chemistry SLTopic 3 — Classifying the ElementsPaper 1 & 2Trends~12 min read
Group 17 Elements and Halide Ions
The halogens run the group 1 trend backwards. These elements react by gaining an electron, so anything that pulls electrons in harder makes them more reactive — and that means the small ones at the top of the group are the dangerous ones. Displacement reactions let you see the order for yourself, in colour.
📘 What you need to know
The halogens are the group 17 non-metals: fluorine, chlorine, bromine, iodine and astatine. They are toxic and reactive.
They exist as diatomic molecules and have an outer configuration of ns2np5 — one electron short of a full shell.
They react by gaining one electron to form 1− halide ions.
Colour deepens and melting and boiling points rise down the group, because bigger molecules have more electrons and stronger London forces between them.
Reactivity decreases down the group, because the incoming electron is added further from the nucleus and is more shielded.
A more reactive halogen will displace a less reactive one from a solution of its halide.
Displacement is a redox reaction: the halogen is reduced, the halide ion is oxidised.
Physical properties down the group
All four halogens you will meet are non-polar diatomic molecules, so the only forces between molecules are London (dispersion) forces. Those forces get stronger as molecules get larger, because a bigger electron cloud is easier to distort. Going down the group, molecular size and electron count both rise, so the forces strengthen and the melting and boiling points climb.
Notice that reactivity and boiling point trend in opposite directions. Boiling point is about forces between molecules; reactivity is about what happens to one atom’s electrons. Different questions, different answers.
Colours in solution matter too. Chlorine water is very pale green, bromine water is orange, and iodine solution is brown. These are the colours you use to identify what has formed in a displacement reaction.
Why reactivity decreases down the group
A halogen reacts by pulling in one extra electron. The incoming electron is attracted by the nucleus, so the question is how strong that attraction is at the point where the electron arrives — which is the outer shell.
Going down the group, the outer shell is further from the nucleus and screened by more inner shells. Fluorine’s incoming electron lands in the second shell, close and barely shielded; iodine’s lands in the fifth, far out and heavily shielded. So fluorine grabs electrons far more effectively, and reactivity falls as you go down.
Two groups, two opposite trends, one explanation. Group 1 wants to lose an electron, so a weak hold helps and reactivity rises downwards. Group 17 wants to gain one, so a weak pull hurts and reactivity falls downwards.
Displacement reactions
Put a halogen into a solution containing the ions of a less reactive halogen and it will take their electrons. That is a displacement reaction, and the colour change tells you it happened.
The blank diagonal is not a gap in the data — adding chlorine to chloride is adding a substance to its own ion, so nothing can happen.
Displacement is redox
Take chlorine added to potassium bromide solution. Chlorine takes electrons from the bromide ions, so bromine is released and chloride ions are left behind.
Full and ionic equations
Cl2(aq) + 2KBr(aq) → 2KCl(aq) + Br2(aq)
Cl2(aq) + 2Br−(aq) → 2Cl−(aq) + Br2(aq)
Split it into half-equations and the electron transfer becomes obvious:
Oxidation: 2Br− → Br2 + 2e− — the halide ion loses electrons
So the more reactive halogen is the better oxidising agent, and oxidising power falls down the group in step with reactivity.
Worked examples
WORKED EXAMPLE
Bromine water is added to colourless sodium iodide solution. Predict what you would see and write the ionic equation.
Step 1: check the order
Bromine is above iodine in group 17, so it is more reactive and will displace it.
Step 2: predict the observation
Iodine is formed, and iodine in solution is brown.
Step 3: write the ionic equationBr₂(aq) + 2I⁻(aq) → 2Br⁻(aq) + I₂(aq)The orange solution turns dark brown as iodine is displacedsodium ions appear on both sides, so they are left out of the ionic equation
WORKED EXAMPLE
Identify what is oxidised and what is reduced in the reaction Cl2 + 2I− → 2Cl− + I2, and name the oxidising agent.
Step 1: track the chlorine
Cl goes from 0 in Cl2 to −1 in Cl−, so it gains electrons.
reductionStep 2: track the iodine
I goes from −1 in I− to 0 in I2, so it loses electrons.
oxidationStep 3: name the agents
The species that causes oxidation is the one being reduced.
Chlorine is reduced and is the oxidising agent; iodide is oxidised and is the reducing agentthe agent labels always feel backwards at first — the oxidising agent is the one that gets reduced
WORKED EXAMPLE
Which pair would react most vigorously: caesium with iodine, or caesium with fluorine? Explain.
Step 1: what each element wants
Caesium must lose an electron; the halogen must gain one.
Step 2: apply both trends
Group 1 reactivity increases down, so caesium is already the best choice available.
Group 17 reactivity decreases down, so fluorine beats iodine.
Step 3: combine
Easiest electron to give, strongest pull to take it.
Caesium with fluorine — lowest ionisation energy meets the most electronegative elementthe two trends run in opposite directions, so the most violent pair is bottom-left with top-right
💡 Exam tip
Learn the three solution colours: chlorine very pale green, bromine orange, iodine brown. Most displacement marks are colour marks.
Write the ionic equation when asked to show what is really happening, and cancel the spectator ions.
Say the halogen is reduced and the halide ion is oxidised. Naming the agents as well usually picks up an extra mark.
Keep the two explanations separate: boiling point is about forces between molecules; reactivity is about attraction for an incoming electron.
Astatine is radioactive and barely studied — predict its properties by extending the trend rather than quoting data.
Halogens are diatomic. Writing Cl instead of Cl2 in an equation costs marks.
⚠ Common mix-up
Assuming reactivity increases down group 17 because it does in group 1. It is the opposite.
Saying covalent bonds break when iodine sublimes. Only the London forces between I2 molecules are overcome.
Mixing up halogen and halide. The halogen is the neutral element; the halide is the 1− ion.
Predicting a reaction with the wrong direction. A halogen never displaces one above it in the group.
Quoting the colour of the halogen instead of the solution. Iodine solid is grey-black but its solution is brown.
Forgetting to balance charge in ionic equations. One Cl2 needs two Br−.
Up next: Metallic and Non-Metallic Oxides — we cross a whole period instead of walking down a group, and watch the oxides change from strongly alkaline to strongly acidic as the bonding changes underneath.
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