IB Chemistry HLTopic 6 — Electron TransferPaper 1 & 2Core idea~10 min read
Electrolytic Cells
A voltaic cell lets a reaction that wants to happen do useful work on the way. An electrolytic cell does the opposite: it uses electricity to force a reaction that would never happen on its own. That is how we pull sodium out of salt and aluminium out of its ore.
📚 What you need to know
Electrolysis uses electrical energy to drive a non-spontaneous redox reaction.
The compound must be molten or dissolved, so its ions are free to move.
Cations move to the cathode and are reduced; anions move to the anode and are oxidised.
In an electrolytic cell the cathode is negative and the anode is positive — the opposite of a voltaic cell.
From a molten binary compound you always get the metal at the cathode and the non-metal at the anode.
Electrons move through the wires; ions move through the electrolyte. Nothing else carries the charge.
Electrolysing molten lead(II) bromide
This is the simplest possible case. There are only two ions present, so there is no competition and no ambiguity.
Solid lead(II) bromide will not conduct at all. The ions are locked in the lattice, and only melting sets them free to carry charge.
The word gives the game away. “Electrolysis” comes from the Greek for splitting by electricity, and that is exactly what you are watching: a stable compound being pulled apart into the elements it was made from.
Why the charges are the other way round
This is the single most confusing thing in the topic, and it becomes obvious once you ask what is pushing the electrons.
In a voltaic cell, the reaction itself pushes electrons out of the anode, so the anode piles up with electrons and is negative.
In an electrolytic cell, a power supply pulls electrons away from the anode, so the anode is short of electrons and is positive.
What never changes is the chemistry: oxidation happens at the anode and reduction at the cathode, in both types of cell.
A memory hook that works both ways round: an Anode is where Oxidation happens. Learn that pair and you never need the charges memorised.
Why it has to be molten or dissolved: in a solid ionic lattice the ions are fixed in place. They still have charge, but they cannot move, so no current can flow. Melting or dissolving breaks the lattice apart and lets them travel.
What electrolysis is used for
Use
What is electrolysed
Why electrolysis is needed
Extracting aluminium
Molten aluminium oxide in cryolite
Aluminium is too reactive to be displaced by carbon
Extracting sodium
Molten sodium chloride
No cheaper reducing agent is strong enough
Purifying copper
Copper(II) sulfate with copper electrodes
Gives the very high purity that electrical wiring needs
Making chlorine and sodium hydroxide
Concentrated brine
Produces three useful products from cheap salt water
Worked examples
WORKED EXAMPLE
Molten aluminium oxide is electrolysed. Give the electrode half-equations and the overall equation.
Step 1: Identify the ionsAl3+ and O2−. Cations go to the cathode, anions to the anode.Step 2: Cathode, reductionAl3+(l) + 3e− → Al(l)Step 3: Anode, oxidation2O2−(l) → O2(g) + 4e−Step 4: Scale to 12 electrons and add2Al2O3(l) → 4Al(l) + 3O2(g)the carbon anodes burn away in the oxygen, which is why they need replacing
WORKED EXAMPLE
Explain why solid sodium chloride does not conduct electricity but molten sodium chloride does.
Step 1: Say what carries the chargeIn an ionic compound the charge carriers are the ions themselves.Step 2: Describe the solidThe ions are held in fixed positions in the lattice by strong electrostatic attraction, so they cannot move.Step 3: Describe the liquidMelting overcomes those attractions, freeing the ions to move towards the electrodes.Only the molten form has mobile ions, so only it conducts“mobile ions” is the phrase that earns the mark
WORKED EXAMPLE
Molten magnesium chloride is electrolysed. Predict the products at each electrode and state the electrode charges.
Step 1: List the ionsMg2+ and Cl−. There is no water, so no competition.Step 2: Positive ions go to the negative electrodeCathode (−): Mg2+(l) + 2e− → Mg(l)Step 3: Negative ions go to the positive electrodeAnode (+): 2Cl−(l) → Cl2(g) + 2e−Magnesium metal at the negative cathode, chlorine gas at the positive anodemolten binary salt: metal one end, non-metal the other, every time
💡 Exam tip
State the charge and the name of each electrode together: “the cathode, which is negative”.
Include state symbols. For a molten salt the ions are (l), not (aq).
Balance the electrons before writing an overall equation — they must cancel completely.
If asked why a compound must be molten, say the ions must be free to move.
Remember cations go to the cathode. Both words start the same way, which is the easiest hook there is.
For industrial uses, link the method to the reason: aluminium is too reactive for carbon reduction.
⚠ Common mix-up
Using voltaic electrode charges in electrolysis. Here the anode is positive.
Saying electrons move through the electrolyte. Ions do that. Electrons stay in the wires.
Forgetting non-metals are diatomic. 2Br− → Br2 + 2e−, not Br− → Br + e−.
Predicting products for a solution when the question says molten — or the reverse. Water changes everything.
Thinking electrolysis makes electricity. It consumes it.
Writing (aq) for molten compounds. There is no water present at all.
Up next: Oxidation of Alcohols — the topic turns to organic chemistry now, but the idea is identical. Something loses electrons, something else gains them, and the oxidation numbers tell you which.
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