Worked example 1
Electrolysis of molten NaCl. Write the electrode reactions and identify the sign of each electrode.
Try it first: Notice that with no water present, only Na⁺ and Cl⁻ can react.
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Electrochemistry & Applications of ThermodynamicsElectrochemistryContent level: Core 24 min
Electrolytic Cells and Driven Reactions
What you'll be able to do: Compare electrolytic and galvanic cells, assign electrode signs correctly, and predict the products of an electrolysis.
Best after: Standard Reduction Potentials and Cell Potential
Introduction
A galvanic cell lets a favorable reaction pay you in electricity. An electrolytic cell does the reverse: you pay electricity to force an unfavorable reaction to run. Aluminum, chlorine, and every chrome bumper come from that bargain.
These are recommended, not required. You can start this lesson at any time.
Learning objectives
- Distinguish electrolytic from galvanic cells by sign and spontaneity
- Assign electrode polarity in an electrolytic cell
- Predict electrolysis products, including cases where water competes
- Explain the role of overpotential and molten-salt electrolysis
Lesson
Running a reaction uphill
If E°cell for a reaction is negative, the reverse process is spontaneous, but the forward one can still be forced by supplying electrical energy. That is electrolysis. Charging a rechargeable battery is the same idea: the cell operates galvanically while discharging and electrolytically while charging.
applied voltage > |E°cell|
Signs flip, definitions do not
Oxidation is at the anode and reduction is at the cathode in both cell types. What changes is polarity. In a galvanic cell the anode is negative because it produces electrons; in an electrolytic cell the anode is positive because the power supply is pulling electrons out of it.
Memorize the definitions, then derive the signs from which way the electrons are being pushed.
Overpotential
Thermodynamics gives the minimum voltage; reality demands more. Extra voltage, called overpotential, is needed to overcome the slow kinetics of forming gas bubbles at an electrode surface. Oxygen evolution has a particularly large overpotential, which is a major reason chlorine, not oxygen, is produced when concentrated brine is electrolyzed.
Competing reactions in water
In an aqueous solution the solvent can be reduced (2 H₂O + 2 e⁻ → H₂ + 2 OH⁻) or oxidized (2 H₂O → O₂ + 4 H⁺ + 4 e⁻). The species with the more positive reduction potential is reduced at the cathode, and the species easiest to oxidize reacts at the anode. Since Na⁺ reduction is at -2.71 V, far below water, aqueous sodium chloride yields hydrogen gas, not sodium metal.
2 H₂O + 2 e⁻ → H₂ + 2 OH⁻, E = -0.83 V at pH 7
Molten salts
To make sodium metal you must remove water entirely. In the Downs cell, molten NaCl is electrolyzed: sodium metal forms at the cathode and chlorine gas at the anode. Aluminum is produced the same way from alumina dissolved in molten cryolite, a process that consumes a large fraction of the electricity used by heavy industry.
2 NaCl(l) → 2 Na(l) + Cl₂(g)
Everyday electrolysis
Electroplating deposits a thin metal layer by making the object the cathode in a solution of the plating metal. Anodizing thickens the oxide layer on aluminum by making it the anode. Both are electrolysis with a decorative or protective purpose.
Key ideas
Definition
Electrolytic cell
A cell in which an external source drives a nonspontaneous redox reaction.
Rule
Polarity
Electrolytic: anode positive, cathode negative. Galvanic: the reverse.
Rule
Electrode roles
Oxidation at the anode and reduction at the cathode in both cell types.
Rule
Product selection
The species easiest to reduce plates out; the species easiest to oxidize reacts at the anode.
Definition
Overpotential
Extra voltage beyond the thermodynamic minimum, needed to overcome slow electrode kinetics.
Equation
Minimum applied voltage
V(applied) > |E°cell| + overpotential
- E°cell = negative for the driven reaction
Equation
Water reduction
2 H₂O + 2 e⁻ → H₂(g) + 2 OH⁻
- E = -0.83 V at pH 7
Equation
Water oxidation
2 H₂O → O₂(g) + 4 H⁺ + 4 e⁻
- E = +0.82 V as a reduction at pH 7
Worked examples
Worked example 2
Why does electrolysis of aqueous NaCl produce hydrogen instead of sodium metal at the cathode?
Try it first: Compare the two possible cathode reactions on a potential scale.
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Worked example 3
A reaction has E°cell = -1.23 V. What can be said about the voltage required to drive it?
Try it first: Separate the thermodynamic minimum from the practical requirement.
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Common mistakes
Switching the anode to reduction in an electrolytic cell.
Why it's wrong: The definitions never change; only the polarity does.
Check instead: Anode = oxidation, always.
Calling the electrolytic cathode positive.
Why it's wrong: The power supply forces electrons onto the cathode, making it negative.
Check instead: Follow the electrons from the supply.
Predicting sodium metal from aqueous NaCl.
Why it's wrong: Water is reduced far more easily than Na⁺.
Check instead: Always include water among the candidate reactions in aqueous electrolysis.
Assuming the thermodynamic voltage is enough in practice.
Why it's wrong: Overpotential from slow gas-forming kinetics adds a real extra requirement.
Check instead: Treat |E°cell| as a lower bound.
Thinking electrolysis violates thermodynamics.
Why it's wrong: Energy is supplied from outside, so the overall process still obeys the second law.
Check instead: Account for the external work in ΔG.
Practice this skill
No practice questions are available for this topic yet. You can still practice the whole unit.
What you should now know
An electrolytic cell uses an external power supply to drive a reaction with a negative E°cell and a positive ΔG°. Oxidation still occurs at the anode and reduction still occurs at the cathode, but the electrode signs reverse relative to a galvanic cell: in electrolysis the anode is positive and the cathode is negative, because the power supply pulls electrons away from the anode and pushes them onto the cathode. The applied voltage must exceed the magnitude of E°cell, and in practice an additional overpotential is needed to overcome kinetic barriers, especially for gas-forming reactions. When a solution contains several possible reactions, the species that is easiest to reduce plates at the cathode and the species that is easiest to oxidize reacts at the anode; in aqueous solution, water itself is often the competitor, which is why electrolyzing aqueous NaCl gives hydrogen and chlorine rather than sodium metal. Molten salts remove water from the competition, which is how reactive metals such as sodium and aluminum are produced industrially.
- Electrolysis drives a reaction with negative E°cell using external energy
- Oxidation stays at the anode and reduction at the cathode in both cell types
- In electrolysis the anode is positive and the cathode negative
- Water often outcompetes dissolved ions in aqueous electrolysis
- Molten salts are used to make reactive metals such as sodium and aluminum
Sources and further reading
- Chemistry 2e, Section 17.7: ElectrolysisOpenStax · Rice University · Chapter 17.7View source
Access for free at openstax.org License
- Chemistry 2e, Section 17.5: Batteries and Fuel CellsOpenStax · Rice University · Chapter 17.5View source
Access for free at openstax.org License