IB Chemistry SLTopic 6 — Electron TransferPaper 1 & 2Trends~13 min read
Relative Ease of Redox
Any half-equation can be written down. Whether the reaction actually happens is a different question, and it comes down to a tug of war: which species wants the electrons more.
📚 What you need to know
The more reactive a metal, the more easily it is oxidised, and the better a reducing agent it is.
Ease of oxidation increases down group 1 and group 2.
The oxidising power of the halogens decreases down group 17: F2 is the strongest.
A halogen displaces the halide of any halogen below it in the group.
A more reactive metal displaces a less reactive metal from a solution of its ions.
Displacement experiments are how the reactivity series is worked out in the first place.
Aqueous colours: chlorine very pale green, bromine orange, iodine brown.
Metals: how readily they let go
A metal reacts by losing electrons, so “reactive metal” and “easily oxidised” mean the same thing. Lithium reacts steadily with water; sodium fizzes about the surface; potassium ignites. Going down group 1, the outer electron is further from the nucleus and better shielded, so it leaves more easily.
A group 1 metal with water
2Li(s) + 2H2O(l) → 2LiOH(aq) + H2(g) Li: 0 → +1
Group 2 shows the same pattern one step down in vigour. Magnesium reacts only very slowly with cold water, while calcium reacts quickly and visibly.
The two ladders run in opposite senses. At the top of the metals is the substance most desperate to give electrons away; at the top of the halogens, the one most desperate to take them.
Halogens: how eagerly they take
A halogen reacts by gaining an electron to reach a 1– charge, so a reactive halogen is a strong oxidising agent. Going down group 17 the outer shell is further out and better shielded, so the incoming electron is held less strongly and the oxidising power falls.
Oxidising power
F2 > Cl2 > Br2 > I2
Put a more reactive halogen into a solution containing the halide of a less reactive one and it simply takes the electrons: a displacement reaction, visible as a colour change.
The whole pattern is one diagonal. Every reaction sits above the leading diagonal, because a halogen can only displace something less reactive than itself.
Both trends have the same underlying cause, which is worth saying in an answer. Going down either group, the outer shell is further from the nucleus and screened by more inner electrons. That makes it easier for a metal to lose an electron and harder for a halogen to gain one — so metals get more reactive down a group and halogens get less.
Metals displacing metals
The same tug of war explains the classic school demonstration. A more reactive metal pushes its electrons onto the ions of a less reactive one, and the two swap places.
Two observations, one explanation: the blue fades because Cu2+ is being removed from solution, and the brown coating is the copper metal it has become.
This is how the reactivity series is built, not just illustrated. Test every metal against every other metal’s solution, note which combinations react, and the order falls out. A reaction means the added metal is the more reactive one; no reaction means it is not.
WORKED EXAMPLE
State, with a reason, whether each mixture reacts. (a) bromine water added to potassium iodide solution (b) iodine solution added to potassium chloride solution (c) chlorine water added to potassium fluoride solution
(a) Br₂ with I⁻Bromine is above iodine, so it is the stronger oxidising agent and takes the electrons.reacts — the solution turns brownBr₂ + 2I⁻ → 2Br⁻ + I₂(b) I₂ with Cl⁻Iodine is below chlorine, so it is the weaker oxidising agent and cannot take electrons from chloride.no reaction(c) Cl₂ with F⁻Fluorine is above chlorine, so fluoride holds its electrons even more tightly than chloride does.no reaction
WORKED EXAMPLE
Use these results to place the four metals in order of decreasing reactivity. Fe(s) + Cu2+(aq) → Fe2+(aq) + Cu(s) Cu(s) + Mg2+(aq) → no reaction Mg(s) + Zn2+(aq) → Mg2+(aq) + Zn(s) Zn(s) + Fe2+(aq) → Zn2+(aq) + Fe(s)
Take each result in turnA reaction means the solid metal is the more reactive of the pair.Fe > CuMg > Cu (no reaction the other way round)Mg > ZnZn > FeChain them togetherMg > Zn > Fe > CuA “no reaction” line is just as informative as a reaction — it tells you the order the other way round.
WORKED EXAMPLE
Zinc is added to silver nitrate solution. Write the two half-equations and the overall ionic equation, and identify the reducing agent.
Step 1 — which way round?Zinc is above silver in the reactivity series, so zinc gives electrons to Ag⁺.Step 2 — the halvesZn → Zn²⁺ + 2e⁻Ag⁺ + e⁻ → AgStep 3 — match the electrons and addZn(s) + 2Ag⁺(aq) → Zn²⁺(aq) + 2Ag(s)zinc is the reducing agentYou would see the zinc dissolve and a grey deposit of silver form on what remains of it.
💡 Exam tip
Say “more reactive metals are more easily oxidised” — it links reactivity to electrons, which is what the question wants.
Explain group trends with atomic radius and shielding, not just “it is more reactive”.
For a displacement, state which species is oxidised, which is reduced, and give both half-equations.
Quote the colour of the solution before and after. That is where the observation marks are.
Remember the trends run opposite ways: metals get more reactive down a group, halogens get less.
A “no reaction” result is evidence too — use it in ordering questions.
⚠️ Common mix-up
Applying the metal trend to the halogens. Reactivity increases down group 1 but up group 17.
Thinking a halide can displace a halogen. Iodide will never displace chlorine.
Confusing the halogen with its halide in an equation: Cl2 is the oxidising agent, Cl– is not.
Saying “the metal is more reactive so it displaces” without saying it loses electrons more easily.
Forgetting to balance the electrons when the charges differ, as with Zn and Ag+.
Up next: Acids with Reactive Metals — hydrogen sits in the middle of the metals ladder for a reason. Which side of it a metal falls on decides whether an acid will touch it at all.
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