Lesson 1 of 12 · 9 min

Galvanic and electrolytic cells

NCERT §2.1

Two rods, two beakers and a strip of salt-soaked gel, and the voltmeter reads 1.13 V. Where does that push come from, and what happens when a stronger supply pushes back?

The story this chapter follows: The science-fair electricity stall

A class runs an electricity stall at the school science fair. On the table: a home-made Daniell cell (zinc in 0.10 M Zn²⁺, copper in 1.0 M Cu²⁺) driving a voltmeter, a conductivity cell with a cell constant of 0.645 cm⁻¹, a tray that copper-plates a brass key at 0.50 A, a beaker of brine under electrolysis, the lab's old scooter battery, a model hydrogen-oxygen fuel cell, and two nails left standing in water, one plain and one galvanised. Their readings carry every lesson.

The lesson in notes

In short

An electrochemical cell couples a redox reaction to an electric circuit. In a galvanic (voltaic) cell a spontaneous reaction produces electrical energy; in an electrolytic cell electrical energy from outside forces a non-spontaneous reaction.

The Daniell cell runs Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s). With both ions at 1 mol dm⁻³ its potential is 1.1 V.

Connect an external source that opposes the cell. While E_ext < 1.1 V, electrons still flow from the Zn rod to the Cu rod (current from Cu to Zn), zinc dissolves and copper deposits.

At E_ext = 1.1 V exactly, no current flows and the reaction stops.

When E_ext > 1.1 V the flow reverses: electrons go from Cu to Zn, zinc is deposited on the zinc rod and copper dissolves. The same hardware is now an electrolytic cell.

Strictly, activities replace concentrations in these equations; in dilute solutions the two are equal.