Electrochemical Cell Lab — Galvanic vs Electrolytic Cells with Nernst Equation

When AP Chemistry, USNCO, and IChO exams ask you to compare galvanic and electrolytic cells, the fastest path to the right answer is a clear side-by-side mental model — which is exactly what this comparison-focused electrochemical cell lab is built for, rather than walking through every electrochemistry topic in sequence.

The defining difference is spontaneity: a galvanic cell converts a spontaneous redox reaction (ΔG < 0, E°cell > 0) directly into electrical energy, while an electrolytic cell uses an external electrical energy source to force a non-spontaneous reaction (ΔG > 0, E°cell < 0) to occur. Every other difference between the two cell types follows from this one fact.

Electrode polarity is the single most commonly confused point: in a galvanic cell, the anode is negative and the cathode is positive, because electrons are being pushed out of the anode by the spontaneous reaction itself. In an electrolytic cell, the external power source forces electrons in the opposite direction, making the anode positive and the cathode negative. In both cell types, though, oxidation always occurs at the anode and reduction always occurs at the cathode — that identification never changes, only the sign.

The Nernst equation, E = E° - (RT/nF)ln Q, extends standard-condition calculations to real, non-standard concentrations, and applies identically to both cell types. As a galvanic cell discharges, Q increases and E decreases toward zero (equilibrium); running the same reaction electrolytically in reverse, sufficient applied voltage can force the reaction against its natural direction entirely.

This lab displays both cell types side by side with synchronized concentration sliders, letting you watch cell potential respond to the Nernst equation in real time while directly comparing the wiring, electrode signs, and spontaneity of each — purpose-built for the classic "compare and contrast galvanic and electrolytic cells" exam question.