Galvanic vs Electrolytic Cells: How Electrochemistry Problems Work

A galvanic (voltaic) cell converts chemical energy to electrical energy through a spontaneous redox reaction — this is how batteries work. An electrolytic cell does the opposite: it uses electrical energy to drive a non-spontaneous reaction — this is how electroplating and electrolysis work. In both types, oxidation occurs at the anode and reduction occurs at the cathode (An Ox, Red Cat — always). The key difference is that galvanic cells have a positive cell potential (E°cell > 0, spontaneous), while electrolytic cells require an applied voltage to force a reaction that would not happen on its own (E°cell < 0, non-spontaneous).

A galvanic cell separates a redox reaction into two half-cells connected by a wire and a salt bridge. The oxidation half-reaction occurs at the anode, releasing electrons. Those electrons travel through the external wire (this is the electric current) to the cathode, where they are consumed by the reduction half-reaction. The salt bridge allows ions to flow between the half-cells to maintain electrical neutrality.

Here is a concrete example: the classic Zn-Cu cell. Zinc metal is oxidized: Zn(s) -> Zn2+(aq) + 2e- (this happens at the anode). Copper ions are reduced: Cu2+(aq) + 2e- -> Cu(s) (this happens at the cathode). Zinc loses mass as it dissolves. Copper gains mass as metal deposits on the cathode. Electrons flow from zinc to copper through the wire.

To calculate the standard cell potential: E°cell = E°cathode - E°anode (using standard reduction potentials from a table). For the Zn-Cu cell: E°cathode (Cu2+/Cu) = +0.34 V. E°anode (Zn2+/Zn) = -0.76 V. E°cell = +0.34 - (-0.76) = +1.10 V. The positive value confirms the reaction is spontaneous.

The mnemonic: in a galvanic cell, the anode is negative (it is the source of electrons) and the cathode is positive (it attracts electrons). This is opposite to electrolytic cells, which confuses students. The rule that never changes: oxidation at the anode, reduction at the cathode.

An electrolytic cell uses an external power source (a battery or power supply) to force a non-spontaneous redox reaction to proceed. The most familiar example is electrolysis of water: 2H2O(l) -> 2H2(g) + O2(g). This reaction has a negative E°cell (-1.23 V), meaning it will not happen on its own. You must apply at least 1.23 V of external voltage to drive it.

In electrolytic cells, the anode is positive (connected to the positive terminal of the power supply) and the cathode is negative (connected to the negative terminal). This is the opposite polarity from galvanic cells — but the chemistry is the same: oxidation still happens at the anode and reduction still happens at the cathode. The power supply forces electrons to flow in the non-spontaneous direction.

Electroplating is a practical application. To silver-plate a spoon, you make the spoon the cathode and a silver bar the anode. Silver dissolves from the anode (Ag -> Ag+ + e-) and deposits on the spoon at the cathode (Ag+ + e- -> Ag). The power supply drives the process. The amount of metal deposited depends on the current, time, and the metal's molar mass — this relationship is quantified by Faraday's laws of electrolysis.

The minimum voltage needed to drive electrolysis is the magnitude of the (negative) cell potential. In practice, you need more voltage than the theoretical minimum due to overpotential — extra voltage required to overcome activation energy barriers at the electrode surfaces. Overpotential is especially significant for gas-evolving reactions like water electrolysis.

Standard cell potentials assume all concentrations are 1 M, all gases are at 1 atm, and the temperature is 25°C. Real cells rarely operate under these conditions. The Nernst equation adjusts the cell potential for actual concentrations:

E = E° - (RT/nF) x ln(Q)

At 25°C, this simplifies to: E = E° - (0.0592/n) x log(Q)

where E° is the standard cell potential, n is the number of moles of electrons transferred, and Q is the reaction quotient (products over reactants, each raised to its coefficient).

Worked example: For the Zn-Cu cell with [Zn2+] = 0.10 M and [Cu2+] = 2.0 M: Q = [Zn2+]/[Cu2+] = 0.10/2.0 = 0.050. E = 1.10 - (0.0592/2) x log(0.050) = 1.10 - (0.0296)(-1.30) = 1.10 + 0.038 = 1.14 V. The cell potential is slightly higher than standard because the concentration conditions favor the forward reaction (low product concentration, high reactant concentration).

At equilibrium, the cell potential is zero (E = 0) and Q = K (the equilibrium constant). Plugging in: 0 = E° - (0.0592/n) x log(K), which gives log(K) = nE°/0.0592. For the Zn-Cu cell: log(K) = 2 x 1.10 / 0.0592 = 37.2, so K = 10^37.2 — an astronomically large number, confirming that this reaction goes essentially to completion. ChemistryIQ practice problems walk through Nernst calculations with different concentration scenarios.

Here is the comparison that exam questions test most frequently.

Energy conversion: Galvanic converts chemical energy to electrical. Electrolytic converts electrical energy to chemical. Spontaneity: Galvanic is spontaneous (E°cell > 0, negative delta G). Electrolytic is non-spontaneous (E°cell < 0, positive delta G, requires external power). Anode charge: Galvanic anode is negative. Electrolytic anode is positive. Cathode charge: Galvanic cathode is positive. Electrolytic cathode is negative.

What does NOT change between them: oxidation always occurs at the anode. Reduction always occurs at the cathode. Anions migrate toward the anode. Cations migrate toward the cathode. Electrons flow from anode to cathode in the external circuit.

The relationship between E°cell and Gibbs free energy is: delta G° = -nFE°cell. For galvanic cells, E°cell is positive, so delta G° is negative (spontaneous). For electrolytic cells, E°cell is negative, so delta G° is positive (non-spontaneous). This connects electrochemistry directly to thermodynamics — a favorite cross-topic exam question.

One trick question to watch for: a rechargeable battery is a galvanic cell when discharging (spontaneous, generating electricity) and an electrolytic cell when charging (non-spontaneous, using external electricity). The same physical device switches between both types depending on which direction current flows.