Galvanic (Voltaic) Cells

Core Concept

A galvanic (voltaic) cell converts the chemical energy of a spontaneous redox reaction directly into electrical energy. The oxidation and reduction half-reactions are physically separated into two half-cells connected by a wire and a salt bridge. Oxidation always occurs at the anode (negative electrode) and reduction at the cathode (positive electrode). Electrons released at the anode travel through the external wire to the cathode, while ions migrate through the salt bridge to keep the solutions electrically neutral. Because the net reaction is spontaneous, the cell has a positive voltage (E°cell > 0).

Key Ideas

• Anode = oxidation, cathode = reduction. The mnemonic "An Ox, Red Cat" fixes it permanently.
• Electrons flow anode → cathode through the external circuit (conventional current is the opposite way).
• Salt bridge closes the circuit. It lets inert ions migrate so charge does not build up; without it, the cell would stop almost instantly.
• Two half-cells. Each contains an electrode immersed in a solution of its own ions (e.g., Zn in Zn²⁺, Cu in Cu²⁺).
• Spontaneous ⇒ positive voltage. The spontaneous cell is the one whose overall E°cell comes out positive.

Equations and Variables

How It Works

1. In the anode half-cell, the more active metal (e.g., Zn) is oxidized: Zn → Zn²⁺ + 2 e⁻. The electrode loses mass as metal ions dissolve.
2. The freed electrons travel through the external wire toward the cathode.
3. In the cathode half-cell, cations in solution (e.g., Cu²⁺) are reduced at the electrode surface: Cu²⁺ + 2 e⁻ → Cu. The electrode gains mass as metal plates out.
4. To keep each solution neutral, anions migrate toward the anode and cations toward the cathode through the salt bridge.
5. The result is a steady electron current that can do work (light a bulb, run a device) until the cell runs down.

Worked Example

The Daniell cell: Zn(s) | Zn²⁺(aq) || Cu²⁺(aq) | Cu(s).

• Anode (oxidation): Zn(s) → Zn²⁺(aq) + 2 e⁻, E° = +0.76 V (written as oxidation)
• Cathode (reduction): Cu²⁺(aq) + 2 e⁻ → Cu(s), E° = +0.34 V
• Net: Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s)
• Cell potential: E°cell = E°cathode − E°anode = +0.34 − (−0.76) = +1.10 V

Electrons flow from the Zn anode to the Cu cathode; Zn dissolves while Cu plates out. The positive voltage confirms the reaction is spontaneous.

How it works

1. In the anode half-cell, the more active metal (e.g., Zn) is oxidized: Zn → Zn²⁺ + 2 e⁻. The electrode loses mass as metal ions dissolve.
2. The freed electrons travel through the external wire toward the cathode.
3. In the cathode half-cell, cations in solution (e.g., Cu²⁺) are reduced at the electrode surface: Cu²⁺ + 2 e⁻ → Cu. The electrode gains mass as metal plates out.
4. To keep each solution neutral, anions migrate toward the anode and cations toward the cathode through the salt bridge.
5. The result is a steady electron current that can do work (light a bulb, run a device) until the cell runs down.

Common confusions

• "The anode is where reduction happens." — Wrong. Oxidation is at the anode; reduction at the cathode (An Ox, Red Cat).
• "Electrons flow cathode → anode." — Wrong. Electrons flow anode → cathode through the external circuit.
• "The salt bridge carries electrons." — Wrong. It carries ions; electrons travel only through the wire.
• "The cathode always gains mass and the anode always loses it." — Only true for metal/metal-ion half-cells; gas or inert electrodes (e.g., Pt, graphite) don't change mass.
• "A galvanic cell needs an external battery." — Wrong. It produces electricity from a spontaneous reaction; an external source is what an electrolytic cell needs.

Quick review

• Galvanic cell = spontaneous redox → electrical energy.
• Anode (oxidation, −), cathode (reduction, +).
• e⁻ flow anode → cathode; ions flow through the salt bridge.
• Salt bridge prevents charge buildup and completes the circuit.
• E°cell > 0 for a working galvanic cell.