An electrolytic cell uses an external source of electrical energy to drive a thermodynamically unfavorable (nonspontaneous) redox reaction, one where ΔG° > 0 and E°cell < 0. Oxidation still happens at the anode and reduction at the cathode, but the battery forces the electrons to flow.
An electrolytic cell is the opposite of a battery in action. A galvanic cell runs a spontaneous redox reaction and produces electricity. An electrolytic cell takes electricity from an outside power supply and uses it to push a redox reaction that would never happen on its own. In CED terms (9.7.A.1), an external source of energy makes a thermodynamically unfavorable process occur. Charging a rechargeable battery, splitting molten NaCl into sodium metal and chlorine gas, and electroplating silver onto a brass object are all electrolytic processes.
The redox vocabulary doesn't change. Oxidation still occurs at the anode and reduction still occurs at the cathode (the mnemonic "an ox, red cat" works in both cell types). What flips is the sign convention and the energetics. In an electrolytic cell, the overall reaction has E°cell < 0 and ΔG° > 0, so the applied voltage has to be large enough to overcome that unfavorable free energy change. The power supply is doing thermodynamic work the reaction can't do for itself.
Electrolytic cells live in Topic 9.7 of Unit 9 (Thermodynamics and Electrochemistry) and directly support learning objective 9.7.A, which asks you to explain how external energy sources or coupled reactions drive thermodynamically unfavorable processes. This is the payoff of the whole unit. Everything you learned about ΔG°, spontaneity, and E°cell comes together here. The CED pairs the electrolytic cell with photosynthesis (light driving CO₂ to glucose) and ATP coupling as the three flagship examples of forcing unfavorable chemistry to happen. If you can explain why an electrolytic cell needs a battery while a galvanic cell IS a battery, you understand the thermodynamics-electrochemistry connection the exam is built around.
Keep studying AP Chemistry Unit 9
Galvanic (Voltaic) Cells (Unit 9)
A galvanic cell and an electrolytic cell are the same hardware running in opposite directions. Galvanic cells convert chemical energy to electrical energy spontaneously (E°cell > 0); electrolytic cells consume electrical energy to reverse that flow. Charging a phone battery literally converts a galvanic cell into an electrolytic one, which also swaps which electrode is the anode and which is the cathode.
Gibbs Free Energy and Spontaneity (Unit 9)
ΔG° = -nFE° is the bridge. A negative E°cell means a positive ΔG°, which means the reaction won't go without help. The electrolytic cell is the practical answer to "so what do you do when ΔG° is positive?" You pay for it with electrical energy.
Coupled Reactions (Unit 9)
Topic 9.7 puts electrolytic cells and coupled reactions side by side because they solve the same problem two ways. An electrolytic cell uses an external battery to fund an unfavorable reaction; coupled reactions (like ATP → ADP in biology) fund it by pairing it with a favorable reaction through a shared intermediate. Either way, the total energy budget has to come out favorable.
Oxidation-Reduction Reactions (Unit 4 and Unit 9)
Electrolysis is just a forced redox reaction, so the half-reaction skills from Unit 4 carry straight over. In molten NaCl, Cl⁻ is oxidized to Cl₂ at the anode and Na⁺ is reduced to Na at the cathode. If you can split a reaction into half-reactions and track electrons, you can analyze any electrolytic cell.
Multiple-choice questions usually hand you a setup and ask you to identify what's happening at each electrode or why the cell needs an external voltage. Classic stems include electrolysis of aqueous CuSO₄ with inert electrodes (Cu plates out at the cathode, O₂ bubbles off at the anode), molten NaCl (chloride is oxidized at the anode), and electroplating silver onto an object (the object being plated is the cathode, where Ag⁺ is reduced). Another favorite asks what happens when you reverse the current in a galvanic cell. The answer is that the anode and cathode swap identities, but oxidation still happens at the anode. On free-response questions, electrolytic cells show up in stoichiometry-meets-electrochemistry calculations and in justification prompts where you connect a negative E°cell to a positive ΔG° and explain why an applied voltage is required. Always state WHERE oxidation or reduction occurs and WHY the process is nonspontaneous; vague answers about "electricity making it happen" don't earn points.
A galvanic cell runs a spontaneous reaction (E°cell > 0, ΔG° < 0) and produces electrical energy. An electrolytic cell consumes electrical energy to drive a nonspontaneous reaction (E°cell < 0, ΔG° > 0). In both, oxidation happens at the anode and reduction at the cathode, but the electrode signs flip. In a galvanic cell the anode is negative; in an electrolytic cell the anode is positive because the external battery sets the polarity. Quick test for any exam question: if the cell powers something, it's galvanic; if something powers the cell, it's electrolytic.
An electrolytic cell uses external electrical energy to force a nonspontaneous redox reaction, one with E°cell < 0 and ΔG° > 0.
Oxidation occurs at the anode and reduction occurs at the cathode in BOTH galvanic and electrolytic cells; only the spontaneity and electrode signs differ.
Reversing the current in a galvanic cell turns it into an electrolytic cell and swaps which electrode is the anode and which is the cathode.
In electroplating, the object being coated is the cathode, where metal cations like Ag⁺ are reduced onto its surface.
Charging a battery is an electrolytic process, and the CED groups it with photosynthesis and ATP coupling as examples of external energy driving unfavorable reactions (9.7.A.1).
It's an electrochemical cell that uses an outside power source to drive a redox reaction that wouldn't happen on its own (ΔG° > 0, E°cell < 0). Examples include electrolysis of molten NaCl, electroplating, and charging a rechargeable battery.
Positive. The external battery makes the anode positive in an electrolytic cell, which is the opposite of a galvanic cell, where the anode is negative. In both cell types, oxidation still happens at the anode.
A galvanic cell runs a spontaneous reaction and produces electricity (E°cell > 0); an electrolytic cell consumes electricity to force a nonspontaneous reaction (E°cell < 0). Think discharge versus recharge of the same battery.
Yes. Reduction always happens at the cathode, in every type of electrochemical cell. In electrolysis of aqueous CuSO₄, for example, Cu²⁺ is reduced to Cu metal at the cathode while O₂ gas forms at the anode.
Because the reaction is thermodynamically unfavorable. Since ΔG° = -nFE° and E°cell is negative, ΔG° is positive, so the reaction can't proceed without an energy input. The applied voltage must be large enough to overcome that unfavorable free energy change.
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