IB Chemistry SLTopic 6 — Electron TransferPaper 1 & 2Core idea~12 min read
Electrolytic Cells
A voltaic cell lets a reaction run downhill and collects the electricity. An electrolytic cell does the opposite: it spends electricity to drag a reaction uphill, and pulls stable compounds apart in the process.
📚 What you need to know
Electrolysis uses electrical energy to drive a non-spontaneous redox reaction.
The electrolyte must be molten or dissolved, so that 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 anode is positive and the cathode is negative — the opposite of a voltaic cell.
Oxidation is still at the anode and reduction still at the cathode. Only the signs change.
For a molten binary salt, the metal forms at the cathode and the non-metal at the anode.
Uses: extracting reactive metals, purifying copper, electroplating, and making chlorine and sodium hydroxide.
Why the compound has to melt
Solid sodium chloride does not conduct. Its ions are locked in a lattice and cannot go anywhere, so no charge can flow. Melt it, or dissolve it, and those same ions become mobile — and a mobile ion is a charge carrier.
Molten lead(II) bromide is the standard demonstration because both products are obvious: a bead of silvery lead at the cathode and brown bromine vapour at the anode.
The ions move for the reason you would expect: opposite charges attract. Positive ions are pulled to the negative electrode, negative ions to the positive one. That is where the names come from — cations go to the cathode, anions go to the anode — and it works in every cell, whichever way round the signs are.
The polarity flips, the chemistry does not
This catches almost everyone once. Oxidation happens at the anode and reduction at the cathode in both kinds of cell. What changes is which electrode carries which sign, and it changes because the driving force is coming from a different place.
In a voltaic cell electrons pile up on the anode as the metal dissolves, making it negative. In an electrolytic cell the supply pulls electrons off the anode, making it positive.
If you can only hold one fact, hold this one: oxidation is at the anode in every cell. Work out the polarity afterwards by asking who is pushing the electrons — the chemistry, or the power supply.
Predicting the products
For a molten binary ionic compound it is straightforward, because there are only two ions in the melt. The metal cation can only be reduced, and the non-metal anion can only be oxidised.
Identify the ions present in the melt or solution.
Send the positive ions to the cathode and write a gain of electrons.
Send the negative ions to the anode and write a loss of electrons.
Balance the electrons in each half-equation, and remember non-metals such as Cl2 and Br2 are diatomic.
What it is used for
Electrolysis is the only practical way to extract the most reactive metals: nothing is a strong enough reducing agent to displace aluminium from its ore, so electricity does the job instead. It also purifies copper, and it plates a thin layer of one metal onto another.
The anode slowly dissolves as Ag+ ions leave it, and the same ions are deposited on the object. The solution’s concentration barely changes, because silver is being replaced as fast as it is used.
Application
What happens
Why electrolysis
Extracting aluminium
Al3+ reduced to Al at the cathode
no chemical reducing agent is powerful enough
Purifying copper
impure copper anode dissolves, pure copper deposits at the cathode
separates copper from impurities that will not dissolve
Electroplating
a thin metal layer is deposited on the object
gives corrosion resistance or appearance cheaply
Chlor-alkali process
brine gives chlorine, hydrogen and sodium hydroxide
three useful products from one cheap raw material
WORKED EXAMPLE
Molten magnesium chloride is electrolysed using inert electrodes. Predict the product at each electrode, write the half-equations, and give the overall equation.
Step 1 — the ions presentThe melt contains Mg²⁺ and Cl⁻ ions only.Step 2 — cathode, negativeMg²⁺ is attracted there and gains electrons.Mg²⁺ + 2e⁻ → MgStep 3 — anode, positiveCl⁻ is attracted there and loses electrons. Chlorine is diatomic, so two ions are needed.2Cl⁻ → Cl₂ + 2e⁻Step 4 — combineMgCl₂(l) → Mg(l) + Cl₂(g)Molten magnesium at the cathode, green chlorine gas at the anode.
WORKED EXAMPLE
Explain why solid potassium iodide does not conduct electricity, but molten potassium iodide does.
In the solidThe K⁺ and I⁻ ions are held in fixed positions in the ionic lattice.no mobile charge carriers, so no conductionIn the meltMelting overcomes the electrostatic attractions and frees the ions to move through the liquid.the mobile ions carry the chargeSay mobile ions, not “free electrons”. Ionic compounds have no delocalised electrons — that is a metal.
WORKED EXAMPLE
A steel spoon is to be plated with silver. State which electrode the spoon should be connected as and why, and write the half-equation for what happens there.
Which electrode?Silver must be deposited as metal, which means Ag⁺ ions gaining electrons — a reduction.the spoon is the cathode, connected to the negative terminalThe half-equationAg⁺(aq) + e⁻ → Ag(s)The positive ions are attracted to the negative electrode, which is exactly where you want the metal to end up.
💡 Exam tip
State that the electrolyte must be molten or in solution so the ions are mobile.
Work out the products by asking which ion goes where, then whether it gains or loses electrons.
Remember that halogens, oxygen and hydrogen are diatomic when you balance the anode equation.
Say anode positive, cathode negative for an electrolytic cell — and note it is the reverse of a voltaic one.
Never write electrons travelling through the electrolyte. Ions move there.
For extraction questions, explain that no chemical reducing agent is strong enough for the most reactive metals.
⚠️ Common mix-up
Putting oxidation at the cathode because the sign has changed. The sign changes; the reaction does not.
Saying an ionic solid conducts because it contains ions. They have to be free to move.
Writing Cl instead of Cl2 at the anode.
Forgetting to balance the electrons between the two half-equations.
Describing electrolysis as spontaneous. It runs only while the supply is connected.
Up next: Oxidation of Alcohols — electrons do not only move between metals and ions. The rest of this topic applies the same bookkeeping to organic molecules, where oxidation and reduction have a very particular look.
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