Electrochemical Cells

Electrochemical cells are systems that convert chemical energy into electrical energy or use electrical energy to drive chemical reactions. They consist of two electrodes in contact with an electrolyte, where redox reactions enable electron transfer. Most electrochemical cells include two half-cells connected by an external wire for electron flow and a salt bridge for ion flow. The salt bridge contains an electrolyte solution and maintains charge neutrality by allowing ions—not electrons—to migrate between compartments.

Electrochemical cells are classified into galvanic (voltaic) cells and electrolytic cells. Galvanic cells generate electricity spontaneously through a redox reaction, while electrolytic cells require an external power source to drive a nonspontaneous reaction. Electrolytic cells are widely used in applications such as electroplating metals onto jewelry and industrial metal products.

In any electrochemical cell, oxidation occurs at the anode and reduction occurs at the cathode. In a galvanic cell, the anode is negative, and the cathode is positive because electrons flow from the anode to the cathode through the wire.

A classic example is the Daniell cell, developed by John Daniell in 1836. It consists of a zinc electrode immersed in ZnSO₄ solution and a copper electrode immersed in CuSO₄ solution, separated by a porous barrier or salt bridge. At the anode, zinc is oxidized: Zn → Zn²⁺ + 2e⁻

The released electrons travel through the external circuit to the copper electrode, where reduction occurs: Cu²⁺ + 2e⁻ → Cu

The overall cell reaction is: Zn + Cu²⁺(aq) → Zn²⁺(aq) + Cu

The electromotive force (emf) of the cell is the open-circuit potential difference between its terminals and reflects the driving force of the spontaneous reaction.