Galvanic (Voltaic) Cells

Drop a strip of zinc straight into copper(II) sulfate solution and the reaction still happens: zinc gives up electrons to copper ions, Zn turns to Zn²+, and Cu²+ turns to solid copper that plates onto the strip. All the energy of that reaction is wasted as heat, because the electrons jump directly from zinc atom to copper ion with no detour. A galvanic cell keeps the two halves apart in separate beakers, one with a zinc electrode in ZnSO₄ solution, one with a copper electrode in CuSO₄ solution. Now the only path for electrons from zinc to copper is through an external wire, and that wire can run a light bulb, a motor, or a voltmeter.

At the zinc electrode, the anode, zinc atoms lose electrons: Zn → Zn²+ + 2e-. Those electrons travel through the wire to the copper electrode, the cathode, where Cu²+ ions gain them: Cu²+ + 2e- → Cu. Electrons flow anode to cathode through the wire, so conventional current flows the other way. Inside the solutions, a salt bridge filled with an inert electrolyte connects the two beakers, letting ions migrate to balance charge: as Zn²+ builds up on one side and Cu²+ is used up on the other, anions drift toward the zinc side and cations drift toward the copper side. Pull the salt bridge and charge piles up uncompensated on each electrode within seconds, and the circuit stops.

This setup is the physical, measurable twin of the free energy calculations from earlier in the unit: the voltage the cell produces, E_cell, is just another way of reporting the same spontaneity that negative delta G describes, linked by delta G = -nFE_cell.