Electrolysis & Faraday's Laws
Running redox in reverse — pouring electricity in to force reactions, and counting exactly how much product you get.
⏱️ About 18 min
A galvanic cell gives you electricity for free from a spontaneous reaction. Electrolysis does the opposite: plug in a power supply and you can force reactions that would never happen on their own — splitting water into hydrogen and oxygen, electroplating a spoon with silver, refining aluminium. And the amount of product is no mystery: it's set precisely by how much charge you push through.
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The big idea: An electrolytic cell uses an external power source to drive a non-spontaneous redox reaction. Oxidation is still at the anode and reduction still at the cathode — but the anode is now POSITIVE (the sign flips from a galvanic cell). How much substance forms is fixed by Faraday's laws: it's proportional to the charge, Q = I·t, and F = 96485 C per mole of electrons.
Redox, forced to run uphill
A galvanic cell runs because its reaction is spontaneous (E°cell > 0). An electrolytic cell runs a reaction that is not spontaneous (E°cell < 0) by supplying energy from an external power source — a battery or power supply pumping electrons where they don't want to go.
The electrode jobs don't change — oxidation still happens at the anode, reduction at the cathode (An Ox, Red Cat). What changes is the driving force: instead of the reaction pushing electrons, the power supply does.
⚠️ The anode sign flips
This is the classic trap. In a galvanic cell the anode is negative. In an electrolytic cell the anode is positive — the external supply pulls electrons out of it, so the power source's positive terminal connects there. The chemistry label (anode = oxidation) is unchanged; only the electrical sign flips.
🔑 Faraday's laws of electrolysis
The mass of substance produced at an electrode is proportional to the charge passed, and for a given charge it depends on how many electrons each ion needs. The chain is always: current × time → charge → moles of electrons → moles of product → mass.
\[ Q = I \cdot t \qquad n_{e^-} = \frac{Q}{F}, \quad F = 96485\ \text{C/mol e}^- \]
Charge Q (coulombs) = current I (amps) × time t (seconds). Divide by Faraday's constant to get moles of electrons.
✨ Electrons per ion matter
Depositing one Cu²⁺ needs 2 electrons (Cu²⁺ + 2e⁻ → Cu); one Ag⁺ needs only 1 (Ag⁺ + e⁻ → Ag). So the same charge plates out twice as many moles of silver as copper. Always read the electrons per ion straight from the half-reaction before converting to moles of metal.
📝 Worked example: A current of 2.0 A flows for 30 minutes through molten/aqueous Cu²⁺. How many grams of copper are deposited? (Cu²⁺ + 2e⁻ → Cu; molar mass Cu = 63.55 g/mol.)
- Time in seconds: 30 min × 60 = 1800 s.
- Charge: Q = I·t = 2.0 A × 1800 s = 3600 C.
- Moles of electrons: n = Q/F = 3600 / 96485 = 0.0373 mol e⁻.
- Copper needs 2 e⁻ each: mol Cu = 0.0373 / 2 = 0.01866 mol.
- Mass: 0.01866 mol × 63.55 g/mol = 1.19 g.
✓ About 1.19 g of copper deposited.
✏️ Practice: A current of 1.5 A is passed for 20.0 minutes through a silver solution (Ag⁺ + e⁻ → Ag; molar mass Ag = 107.87 g/mol). What mass of silver, in grams, is deposited at the cathode?
g
Solution
- Time in seconds: 20.0 min × 60 = 1200 s.
- Charge: Q = I·t = 1.5 A × 1200 s = 1800 C.
- Moles of electrons: n = Q/F = 1800 / 96485 = 0.01866 mol e⁻.
- Silver needs 1 e⁻ each, so mol Ag = 0.01866 mol.
- Mass = 0.01866 × 107.87 = 2.01 g of silver.
Check your understanding
1. How does an electrolytic cell differ from a galvanic cell?
Electrolysis is driven by an external power source to force a non-spontaneous reaction. Oxidation still occurs at the anode, reduction at the cathode.
2. In an electrolytic cell, the anode is:
In an electrolytic cell the anode is positive — the opposite sign to a galvanic cell — because the power supply pulls electrons out of it. It is still where oxidation occurs.
3. The same charge is passed through Cu²⁺ and Ag⁺ solutions. Compared with copper, the moles of silver deposited are:
Ag⁺ needs 1 e⁻ per atom while Cu²⁺ needs 2, so the same charge deposits twice as many moles of silver as copper.
✅ Key takeaways
- Electrolysis uses an external power source to drive a non-spontaneous redox reaction (E°cell < 0).
- Oxidation is still at the anode and reduction at the cathode — but the anode is now POSITIVE (sign flips from galvanic).
- Faraday's laws: mass of product is proportional to charge passed, Q = I·t.
- Convert charge to moles of electrons with F = 96485 C/mol e⁻, then use electrons-per-ion to get moles of product.
- The full chain: current × time → charge → moles e⁻ → moles product → mass.
➡️ You've now seen redox from every angle — counting electrons, balancing them, harvesting them for power, tracking their voltage, and forcing them with electricity. That completes the electrochemistry module; from here the course turns to the chemistry of carbon itself.
Want to test yourself on this?
Try the Chemistry practice test →