9.8 Galvanic (Voltaic) and Electrolytic Cells

An electrochemical cell uses a redox reaction to either produce or consume electrical energy. Galvanic, or voltaic, cells run a thermodynamically favored reaction to generate voltage, while electrolytic cells use an outside power source to drive a thermodynamically unfavored reaction. For AP Chemistry, track where oxidation and reduction happen, which way electrons flow, and whether the cell produces or requires energy.

Galvanic vs Electrolytic Cells Summary

Galvanic cells and electrolytic cells both separate oxidation and reduction into half-cells, but they differ in energy flow. A galvanic cell uses a thermodynamically favored redox reaction to produce electrical energy. An electrolytic cell uses an external power source to drive a thermodynamically unfavored redox reaction.

For AP Chemistry, keep the constant rules separate from the changing features. Oxidation is always at the anode, reduction is always at the cathode, and electrons move through the wire from anode to cathode. What changes is whether the reaction produces voltage on its own or needs outside energy.

Why This Matters for the AP Chemistry Exam

This topic asks you to connect each physical part of a cell to what it does, and to explain how a cell behaves at both the macroscopic and particulate levels. You should be able to read a diagram and identify the anode, cathode, electron flow direction, ion flow through the salt bridge, where mass changes, and where a gas forms. A common skill tied to this topic is explaining how changing the setup of a cell would change the results, so expect questions that test whether you understand why each component is there. Knowing the difference between galvanic and electrolytic cells also sets up later topics that link cell potential to free energy and equilibrium.

Key Takeaways

• Oxidation always occurs at the anode; reduction always occurs at the cathode ("An Ox, Red Cat").
• Electrons travel through the external wire from anode to cathode in both cell types.
• Galvanic (voltaic) cells run a thermodynamically favored reaction and produce a positive cell potential.
• Electrolytic cells run a thermodynamically unfavored reaction and need an external power source to operate.
• A salt bridge supplies spectator ions that keep each half-cell neutral so the reaction can keep going.
• You can describe every operating feature (electron flow, ion flow, electrode mass change, gas evolution) at both the macroscopic and particulate level.

Quick Review of Redox

Electrochemistry is built on redox (oxidation-reduction) reactions, where electrons transfer from the species being oxidized to the species being reduced. Use OIL RIG to keep it straight: Oxidation Is Loss of electrons, Reduction Is Gain of electrons.

Take the reaction:

2AgNO3 + Cu → Cu(NO3)2 + 2Ag

Copper goes from an oxidation number of 0 to +2, so it is oxidized. Silver goes from +1 to 0, so it is reduced. Electrons move from copper to silver. Splitting this into half-reactions:

Oxidation: Cu → Cu2+ + 2e-

Reduction: 2Ag+ + 2e- → 2Ag

(The reduction half-reaction is multiplied by 2 so the electrons balance.)

Galvanic (Voltaic) Cells

A galvanic cell separates the two half-reactions into two half-cells connected by a wire. Because the oxidizing agent gains electrons and the reducing agent loses them, electrons flow through the wire from the anode (oxidation site) to the cathode (reduction site). The driving force pushing those electrons is the cell potential, written as Ecell.

Consider:

Cu + 2Ag+ → 2Ag + Cu2+

Here copper is oxidized (copper anode) and silver ion is reduced (silver cathode). As the cell runs, the copper anode loses mass because Cu is converted to Cu2+, and the silver cathode gains mass because Ag+ is deposited as Ag metal. A voltmeter in the wire measures the cell potential.

The Salt Bridge

The two solutions are connected by a salt bridge, which keeps ions flowing so the cell does not stop. It contains inert ions (spectator ions) that migrate into each half-cell to keep both sides electrically neutral as charge builds up. NaNO3 is a common choice. Without it, charge imbalance would quickly halt the reaction.

In a galvanic cell, the favored reaction gives a positive Ecell, which matches the idea that the reaction is thermodynamically favored.

Electrolytic Cells

An electrolytic cell runs a reaction that is not thermodynamically favored, so it requires an external power source to force it forward. Here Ecell for the spontaneous direction is negative, so a battery or power supply pushes electrons in the nonfavored direction.

Suppose you want to break molten NaCl back into its elements:

2NaCl → 2Na + Cl2

The half-reactions are:

Reduction at cathode: 2Na+ + 2e- → 2Na

Oxidation at anode: 2Cl- → Cl2 + 2e-

This reaction is not favored, so it does not happen on its own. A battery supplies the electromotive force needed to drive the electrons: Cl- is oxidized to Cl2 gas at the anode, and Na+ is reduced to Na metal at the cathode. So you would see gas bubbles forming at the anode and sodium metal building up at the cathode.

Even in an electrolytic cell, the rule holds: oxidation at the anode, reduction at the cathode.

Comparing the Two Cells

Notice that the location of oxidation and reduction does not change between the two cell types. What changes is whether the cell produces energy or needs energy supplied to it.

How to Use This on the AP Chemistry Exam

Reading Cell Diagrams

When you get a diagram, work through it in order: identify which half-cell has oxidation (anode) and which has reduction (cathode), then mark electron flow from anode to cathode in the wire. Track ion movement in the salt bridge: cations move toward the cathode and anions move toward the anode to balance charge. Then connect those to observable changes, like an electrode shrinking, an electrode growing, or gas bubbling off.

Explaining Component Changes

A skill tied to this topic is explaining how changing the setup changes the results. Be ready to reason through questions like what happens if the salt bridge is removed, what gas forms at an electrode, or which electrode gains or loses mass. Tie your answer to the role each component plays rather than just naming parts.

Connecting Macroscopic and Particulate Levels

Strong answers describe both what you would observe (mass change, bubbles, voltmeter reading) and what is happening to the particles (electrons moving through the wire, ions migrating, atoms being deposited or dissolved). Practice writing explanations that move between these two levels.

Common Trap

Do not label electrodes as positive or negative for this topic. The exam will not ask you to assign positive or negative signs to electrodes here, so focus on oxidation/reduction and electron flow instead.

Common Misconceptions

• "Electrons flow through the salt bridge." Electrons travel through the external wire. The salt bridge carries ions, not electrons, to keep each half-cell neutral.
• "The anode is always negative and the cathode is always positive." This topic does not require labeling electrodes as positive or negative, and the assignment can differ between galvanic and electrolytic cells. Stick to anode = oxidation, cathode = reduction.
• "The anode and cathode switch roles in an electrolytic cell." Oxidation is at the anode and reduction is at the cathode in both cell types. What changes is whether the reaction is favored and whether energy is supplied.
• "Electrolytic reactions are impossible because they are unfavored." They are not favored on their own, but an external power source supplies the energy needed to drive them.
• "Multiplying a half-reaction changes its potential." Scaling a half-reaction to balance electrons does not change its reduction potential. Potential is an intensive property.
• "A galvanic cell needs a battery." A galvanic cell produces its own voltage from a favored reaction. The external power source is what an electrolytic cell needs.

Related AP Chemistry Guides

• 9.1 Introduction to Entropy
• 9.3 Gibbs Free Energy and Thermodynamic Favorability
• 9.2 Absolute Entropy and Entropy Change
• 9.4 Thermodynamic and Kinetic Control
• Unit 9 Review: Thermodynamics and Electrochemistry
• 9.5 Free Energy and Equilibrium