IB Chemistry HLTopic 6 — Electron TransferPaper 1 & 2Trends~10 min read
Relative Ease of Redox
Not every redox reaction you can write on paper will actually happen. Whether it does comes down to a simple contest: is one species more desperate to give electrons away than the other is to hold on to them? Ranking species by that eagerness is what the reactivity series really is.
📚 What you need to know
Metals are reducing agents — they give electrons away. Reactivity increases down Groups 1 and 2.
Halogens are oxidising agents — they take electrons. Oxidising power increases up Group 17.
A more reactive metal displaces a less reactive metal from solution.
A more reactive halogen displaces a less reactive halide from solution.
“No reaction” is a real result and tells you which species is more reactive.
Hydrogen sits in the metal list as a reference point: metals above it react with acids, metals below it do not.
Two ladders, opposite directions
The voltages are the same standard electrode potentials you meet later — the reactivity series is just this list with the numbers hidden.
Metals: down the group, easier to oxidise
Going down Group 1 or Group 2, the outer electron is further from the nucleus and better shielded, so it is held less tightly and lost more easily. That is the whole explanation, and examiners want those words.
Magnesium reacts very slowly with cold water.
Calcium, one row lower, reacts steadily and visibly.
Down Group 1 the same pattern goes from vigorous to genuinely dangerous.
Group 2 with water
Ca(s) + 2H2O(l) → Ca(OH)2(aq) + H2(g)
Do not just write “calcium is more reactive”. Say why: the outer electrons are in a higher energy level, further from the nucleus, with more shielding, so less energy is needed to remove them. Three marks instead of one.
Metal displacement
Put a more reactive metal into a solution of a less reactive metal’s ions and it pushes the ions out of solution by handing them electrons.
Mg → Mg2+ + 2e− (oxidation — magnesium is the reducing agent)
Cu2+ + 2e− → Cu (reduction — copper ions are the oxidising agent)
The sulfate ions never change, so they are spectators and can be left out of the ionic equation entirely.
Halogen displacement
Halogens work the opposite way round: the small ones at the top of the group pull electrons in most strongly, so they are the strongest oxidising agents.
Mix
Does it react?
What you see
Equation
Chlorine + potassium bromide
Yes
Colourless to orange
Cl2 + 2KBr → 2KCl + Br2
Bromine + sodium iodide
Yes
Orange to brown
Br2 + 2NaI → 2NaBr + I2
Iodine + potassium chloride
No
No colour change
Iodine is below chlorine
Chlorine + potassium fluoride
No
No colour change
Chlorine is below fluorine
Colours to know: chlorine water is very pale green and often looks colourless; bromine water is orange, yellow when dilute; iodine solution is brown. The colour change is the evidence a displacement happened.
Worked examples
WORKED EXAMPLE
Zinc is added to tin(II) sulfate and a reaction occurs. Copper is added to zinc sulfate and nothing happens. Place the three metals in order of decreasing reactivity.
Step 1: Read the first resultZinc displaced tin, so zinc is above tin.Zn > SnStep 2: Read the second resultCopper could not displace zinc, so copper is below zinc.Zn > CuStep 3: Place tin and copperTin is above copper in the reactivity series, and no result here contradicts that.Zn > Sn > Cua “no reaction” result is data, not a failed experiment
WORKED EXAMPLE
Write the ionic equation, with half-equations, for chlorine reacting with potassium bromide solution.
Step 1: Identify who does whatChlorine is higher in the group, so chlorine takes the electrons.Step 2: Reduction halfCl2(aq) + 2e− → 2Cl−(aq)Step 3: Oxidation half2Br−(aq) → Br2(aq) + 2e−Step 4: Add, leaving out the potassium spectatorsCl2(aq) + 2Br−(aq) → 2Cl−(aq) + Br2(aq)the solution turning orange is the bromine being formed
WORKED EXAMPLE
Explain why iodine cannot displace chlorine from sodium chloride solution.
Step 1: Compare oxidising powerIodine is at the bottom of the group, chlorine higher up.Step 2: Say what that means for electronsChlorine attracts electrons more strongly than iodine does, because its atoms are smaller with less shielding.Step 3: ConcludeIodine cannot pull electrons off Cl−, so there is no reactiondisplacement only ever runs one way — downhill in reactivity
💡 Exam tip
Explain trends with atomic radius, shielding and nuclear attraction — not just “it is more reactive”.
Give the colour change as your observation in halogen questions. That is what a student in the lab would actually see.
Leave spectator ions out of ionic equations, and say which ones they are if asked.
Treat “no reaction” as useful evidence when ranking metals.
Remember the two directions: metals get more reactive down a group, halogens get less oxidising down a group.
If asked to prove an order, quote the specific experiment that shows each pairwise comparison.
⚠ Common mix-up
Applying the metal trend to halogens. They run in opposite directions down the periodic table.
Saying a halogen “is reduced because it gains oxygen”. Use electrons.
Including spectator ions in an ionic equation.
Forgetting halogens are diatomic. It is Cl2 + 2Br−, so the 2 in front of the bromide is required.
Assuming a more reactive metal always reacts faster. Aluminium is reactive but its oxide layer makes it look sluggish.
Confusing reactivity with the ease of extracting a metal. The more reactive the metal, the harder it is to extract.
Up next: Acids with Reactive Metals — hydrogen sits in the middle of the metal ladder, and where a metal sits relative to it decides whether an acid will attack it at all.
Want this explained one-to-one?
Book a free session with an experienced IB Chemistry tutor and get your trickiest topics made simple.
