Pushing the reaction uphill
A galvanic cell lets a reaction roll downhill and pays you out in current. Now flip the whole idea: connect an *external* power source — a charger, a power supply — and use it to shove electrons the wrong way, forcing a reaction that would never go on its own. This is electrolysis, and the apparatus that does it is an electrolytic cell. You spend electricity to buy chemistry.
The two kinds of cell are mirror images. In the galvanic cell a spontaneous redox reaction *makes* electricity; in the electrolytic cell, supplied electricity *drives* a non-spontaneous one. The same parts appear — two electrodes, an anode, a cathode, oxidation here and reduction there — only now an outside source, not the chemistry, calls the shots.
What electrolysis builds for us
Run a current through water with a little salt in it and it splits into hydrogen and oxygen gas — a clean way to make hydrogen fuel from nothing but water and electricity. Run it through molten aluminium ore and pure aluminium metal collects at one electrode; almost every aluminium can on Earth was born this way. Dip a steel spoon and a bar of silver in the right solution, switch on the current, and a gleaming silver coat grows on the spoon. That is electroplating.
Notice the common thread: in each case electrolysis lets us *make* something valuable — pure metal, clean fuel, a protective coating — that would not form by itself. Electricity is the lever that pries an unwilling reaction uphill. The price is energy, which is why aluminium smelters are built next to cheap power and why "green hydrogen" only counts as green when the electricity itself is clean.
Faraday's counting law: charge in, grams out
How much silver actually plates onto that spoon? Here is one of the tidiest results in all of chemistry. Each atom that deposits needs a fixed number of electrons — and electrons arrive as electric charge. So the amount of substance you make is simply proportional to the total charge you pass. Double the current or double the time, and you double the metal. This is the heart of Faraday's laws of electrolysis.
To turn that proportion into grams you need one conversion number: the Faraday constant, the total charge carried by one mole of electrons — about 96,500 coulombs. It is just a bridge between two counting systems: charge on one side, moles of electrons on the other. Once you have it, the whole calculation is bookkeeping, no chemistry intuition required.
- Total charge passed = current (in amps) × time (in seconds). This counts how many electrons you delivered.
- Divide by the Faraday constant to convert that charge into moles of electrons.
- Each atom needs a known number of electrons, so dividing again gives the moles of metal — and weighing follows from there.
Corrosion: the battery nobody asked for
Now the dark side of the same physics. Rust is electrochemistry happening where you do not want it. On a damp iron surface, tiny patches act as anode and cathode, the moisture film is the salt solution, and a miniature galvanic cell springs up uninvited. Iron quietly gives up its electrons, dissolves, and crumbles into rust. This unwanted, self-running redox attack on metals is corrosion.
Understanding it as a tiny cell hands us the cure. If iron rusts because it loses electrons, give it a metal that gives up electrons even *more* eagerly — like zinc — bolted on or coated over the iron. The zinc corrodes first and protects the iron, sacrificing itself; this is why ship hulls and underground pipes wear blocks of zinc, and why galvanised nails resist rust. The zinc is, quite literally, a *sacrificial anode*.