galvanic cell
/ gal-VAN-ik sel /
A galvanic cell is what we usually just call a battery: a device that turns a chemical reaction directly into electricity. Two different metals sit in solutions, joined by a wire and a connecting bridge, and a reaction that wants to happen anyway is harnessed to push electrons through the wire — lighting a bulb or spinning a motor along the way.
Here is the layout. The cell is split into two halves. In one half, a metal is oxidized, giving up electrons; in the other, a substance is reduced, taking them in. The electrons cannot leap straight across the solution, so they are forced to travel the long way — out through the external wire, doing useful work, and into the other electrode. Meanwhile a salt bridge lets ions move to keep each side electrically balanced. The whole thing runs because the overall reaction is spontaneous: it releases energy on its own.
Galvanic cells are everywhere outside the lab — the cell in a flashlight, a car battery while it discharges, a fuel cell in a spacecraft. The catch is that a galvanic cell only delivers power while its reactants last; once the metal is used up or the chemicals reach equilibrium, the voltage falls to zero and the cell is "dead."
The classic Daniell cell pairs a zinc strip in zinc sulfate with a copper strip in copper sulfate. Zinc dissolves and gives up electrons; copper ions plate out onto the copper electrode; the wire between them carries about 1.1 volts.
A Daniell cell: zinc dissolves, copper plates out, and the wire carries the current.
A galvanic cell is the opposite of an electrolytic cell: here a spontaneous reaction produces electricity, whereas an electrolytic cell consumes electricity to force a non-spontaneous one. In a galvanic cell the anode is the negative terminal.