IB Chemistry SLTopic 3 — Classifying the ElementsPaper 1 & 2Trends~10 min read
Group 17 Elements and Halide Ions
The halogens run the trend backwards. Group 1 metals get more reactive down the group; Group 17 elements get less reactive. That isn’t an exception to remember — it falls straight out of the fact that one group loses electrons and the other gains them.
📚 What you need to know
Group 17 = the halogens: fluorine, chlorine, bromine, iodine, astatine. All non-metals, all toxic.
They exist as diatomic molecules (F2, Cl2, Br2, I2) and have 7 valence electrons, configuration ns2np5.
They react by gaining one electron to form halide ions with a 1– charge.
Colour gets darker and melting and boiling points increase down the group.
Reactivity decreases down the group, because the incoming electron is further out and better shielded.
A more reactive halogen displaces a less reactive one from a solution of its halide.
What the halogens are like
All the halogens have 7 outer electrons, so all of them are one electron short of a full shell. That single fact drives everything on this page: they are all looking to gain one electron.
Because they are simple molecules held together by weak intermolecular forces, their physical properties change steadily down the group. Bigger molecules have more electrons, which means stronger London (dispersion) forces, which means more energy is needed to separate them.
Halogen
State at room temperature
Colour
Colour in solution
Fluorine, F2
Gas
Pale yellow
Too reactive to test safely
Chlorine, Cl2
Gas
Pale green
Very pale green
Bromine, Br2
Liquid
Orange–brown
Orange
Iodine, I2
Solid
Grey–black, sublimes to purple vapour
Brown
Learn the solution colours properly. Displacement questions are usually answered from a colour change, so “the solution turns orange” is often the whole observation mark.
Why reactivity decreases down the group
Here is the mirror image of Group 1. A halogen reacts by pulling in an extra electron, so what helps it is a strong attraction to the nucleus — and that attraction gets weaker as you go down.
🧩 The chain of reasoning
Down the group, each atom has one more electron shell, so the atomic radius increases.
The incoming electron would join a shell that is further from the nucleus…
…and is more shielded by the extra inner shells.
So the attraction for the incoming electron is weaker — electron affinity becomes less negative.
The atom gains an electron less easily → the halogen is less reactive.
So fluorine, the smallest, is the most reactive halogen, and reactivity falls away steadily as you go down to iodine. This is also described as decreasing non-metallic character.
One idea, two directions: in both groups the outer electrons get further out and better shielded going down. That makes losing an electron easier (Group 1 more reactive) and gaining one harder (Group 17 less reactive).
Displacement reactions
Put a halogen into a solution containing halide ions of a less reactive halogen, and it will take their electrons. The more reactive halogen becomes the halide ion, and the less reactive one is set free as an element — usually with an obvious colour change.
The order of reactivity is straightforward: Cl2 > Br2 > I2.
A halogen displaces one below it in the group and never one above. Six of the nine combinations do nothing at all.
Chlorine displaces bromide
Add pale green chlorine solution to colourless potassium bromide solution and the mixture turns orange as bromine is released.
Cl2(aq) + 2KBr(aq) → 2KCl(aq) + Br2(aq)
Bromine displaces iodide
Add orange bromine solution to colourless potassium iodide solution and the mixture turns brown as iodine is released.
Br2(aq) + 2KI(aq) → 2KBr(aq) + I2(aq)
Iodine, being the least reactive of the three, displaces neither — add it to KCl or KBr and nothing happens.
Displacement is redox
Strip out the spectator ions and you can see what is really going on. The potassium ions take no part, so the ionic equation tells the whole story:
Chlorine has gained electrons (reduced) and bromide has lost them (oxidised). The more reactive halogen is always the better oxidising agent — it takes electrons from the other one.
WORKED EXAMPLE
Bromine water is added to (a) potassium chloride solution and (b) potassium iodide solution. Predict what happens in each, with equations where a reaction occurs.
(a) Br₂ with KClChlorine is ABOVE bromine, so it is more reactive. Bromine cannot displace it.no reaction — stays orange(b) Br₂ with KIIodine is BELOW bromine, so bromine displaces it. Solution turns brown.Br₂(aq) + 2KI(aq) → 2KBr(aq) + I₂(aq)
WORKED EXAMPLE
Which pair would react most vigorously: Li and I2, Cs and I2, Li and F2, or Cs and F2?
Group 1: reactivity increases DOWN, so pick the lowestcaesiumGroup 17: reactivity decreases DOWN, so pick the highestfluorineCs and F₂ — the easiest electron to give meets the strongest pull to take it.
💡 Exam tip
Sketch a mini Group 17 down the margin (F, Cl, Br, I) before answering a displacement question. Higher displaces lower — that’s the whole rule.
Quote the colour change when asked for observations, not just “a reaction occurs”.
Practise the ionic equations as well as the full ones. Both are examinable.
Remember the two group trends run in opposite directions, and be ready to explain why.
⚠️ Common mix-up
Halogen vs halide. The halogen is the neutral element (Cl2); the halide is the 1– ion (Cl–).
Reactivity decreases down Group 17 — do not copy the Group 1 direction.
Halogens are diatomic. Writing “Cl” instead of “Cl2” costs equation marks.
A halogen cannot displace one above it. No reaction is a perfectly good answer.
Iodine’s solution is brown, and its vapour is purple. Don’t swap them.
Up next: Metallic and Non-Metallic Oxides — what happens when you burn elements from right across a period, and why the products swing from alkaline to acidic.
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