Topic 9.8: Galvanic (Voltaic) and Electrolytic Cells

Galvanic and electrolytic cells: Liberty’s hidden battery

This is the inside of a statue you’ve seen a thousand times, and in the early 1980s, when engineers climbed in to take a close look at her, she was slowly destroying herself. The iron bars that hold her copper skin in shape had rusted so badly that they’d swollen up, and all that rust was pushing on the copper from the inside.

That was the Statue of Liberty from the inside, and her skin is copper, thinner than two pennies stacked together, hung on an iron skeleton designed by Gustave Eiffel, the same engineer who went on to build the Eiffel Tower. Her builders knew something about these two metals, which is that when iron touches copper in wet, salty air, the iron gets eaten, so they put a layer of shellac-coated asbestos between the two to keep them apart.

But she spent nearly a century standing in the salt air of New York Harbor, and over that time the shellac fell apart and the asbestos turned dry and porous, which let salty moisture from the air get in and reach both metals. Metallurgists at the National Bureau of Standards studied pieces of her skeleton, and they found that the two metals, back in contact, were making electric currents like a battery, and that current was eating the iron.

So the layer that was supposed to keep the two metals apart ended up letting them touch again. And it let in the salt water that completed the circuit. By the end of this lesson, you’ll know which metal gets eaten. And you’ll know why.

And then you’ll decide exactly where you’d cut Liberty’s circuit to save her. So let’s go down into that narrow gap between the metals. I want you to watch what the iron is actually doing. Here’s the iron on one side, and the copper on the other. Between them is a thin film of salt water.

Every so often an iron atom lets go of the metal and drifts off into the water as an iron ion, with a charge of two plus. It leaves two electrons behind in the metal, so we write it as iron going to iron two plus plus two electrons, and losing electrons like that is called oxidation.

Those electrons don’t stay in the iron, though, because the iron is touching the copper, so they run straight through the metal into the copper. At the copper’s surface, oxygen molecules from the air and water molecules pick them up and become hydroxide ions, and gaining electrons like that is called reduction. This is all sped up a lot. Each particle stands in for a huge number of real ones.

Now here’s something worth thinking about before I tell you, because it’s a mistake AP graders see a lot. The electrons travel through the metal, so what carries the charge through the salt water in between?

And when iron ions and hydroxide ions meet in the water, they form iron two hydroxide, and oxygen oxidizes that further into orange-brown rust, layer after layer. So there’s a complete circuit here with two halves, electrons carrying the charge through the metal and ions carrying it through the water, and if you break either one of them, the whole thing stops.

The place where oxidation happens is called the anode, so the iron is the anode, and it’s the metal being eaten, and the place where reduction happens is called the cathode, which here is the copper’s surface. All of this is hidden in a crack you could barely slide a card into. So let’s set the same idea up on a bench. There, we can watch every piece of it.

Inside a galvanic cell: anode, cathode and salt bridge

Here’s our bench, with a different pair of metals. On the left, a strip of zinc sits in a solution of zinc nitrate. Next to it is a strip of copper in copper nitrate. It’s blue because of the copper ions in it.

And this exact picture, two metals in two separate beakers, is a classic AP free-response setup. On the 2025 exam it was zinc and aluminum, and the very first thing they asked about was what happens at the anode.

Let’s connect the zinc to the copper with a wire, through a meter that measures how hard the cell pushes the electrons, and the needle doesn’t move at all, because we’ve only built one of the two paths, the one for the electrons. Now I’ll add the second path, a glass tube full of potassium nitrate solution called a salt bridge, and the needle jumps to one point one zero volts.

Now let’s swap the meter for a little motor. The reaction actually runs, and the motor spins. At the zinc, atoms leave the metal as zinc ions and leave two electrons each behind, so the zinc is being oxidized, and that makes the zinc the anode. Those electrons travel through the wire and the motor over to the copper, where copper ions from the blue solution grab two each and turn into copper metal, so that’s reduction, and the copper is the cathode.

Over time the zinc strip thins and the copper strip thickens, so you could follow the reaction just by weighing the electrodes. So what’s the bridge actually doing? The zinc beaker keeps gaining positive zinc ions, while the copper beaker keeps losing positive copper ions. The bridge balances that by sending negative nitrate ions into the zinc side and positive potassium ions into the copper side, which keeps both beakers neutral. And not a single electron goes through the bridge.

Now I want you to predict something. If I lift the salt bridge out while the motor’s running, what happens to the motor? Now let’s put that cell on the board. We’ll draw it the same way AP questions do.

On the left, zinc metal is oxidized to zinc ions, so the zinc electrode is the anode, and on the right, copper ions are reduced to copper metal, so the copper electrode is the cathode. Add the two halves, and the two electrons cancel, which leaves zinc plus copper two plus going to zinc two plus plus copper, the reaction this cell runs. Electrons flow through the wire from the anode to the cathode, and in the salt bridge, negative ions drift toward the anode side and positive ions drift toward the cathode side.

And there’s one rule that holds in every electrochemical cell you’ll ever see, which is that oxidation happens at the anode and reduction happens at the cathode. Plenty of students remember it as an ox, red cat. Some people learn that the anode is the one on the left. But that’s just how this picture happens to be drawn, and the anode is wherever oxidation happens.

Here’s the same cell drawn the other way round, with the copper on the left. Which electrode is the anode, and which way do the electrons go through the wire?

Line the bench up with the statue and almost every part has a match, zinc and iron are the anodes, copper is the cathode in both, the wire matches the spot where her metals touch, and the bridge matches the salt water in the old insulation. The statue just doesn’t have a meter or a motor, so none of that energy does anything useful, and all the reaction does is eat the iron. A cell like either of these, where the reaction runs all by itself and gives off electrical energy, is called a galvanic cell, or a voltaic cell, and chemists call a reaction that runs on its own like that thermodynamically favored.

Electrolytic cells: forcing a reaction backward

Now let’s make this cell do something it would never do on its own, which is run backward, so I’m swapping the motor for a power supply that can push electrons either way. The cell itself pushes electrons from the zinc to the copper with one point one zero volts, so to send them the other way, the supply has to push harder than that. Go ahead and turn the supply up. Keep going until the cell starts running backward.

Now zinc ions are taking electrons and turning back into zinc metal, so the zinc strip grows, and the copper is giving up electrons and dissolving, so the blue around it gets deeper. A cell where an outside power source forces a reaction to happen is called an electrolytic cell, and the reaction inside it is thermodynamically unfavored, which is exactly why it needs the push. So in this electrolytic cell, which electrode is the anode now?

Now let’s swap the copper for a strip of platinum, a metal that won’t give up its own electrons here, and turn the supply up further, because the new reaction needs a bigger push, and bubbles start streaming off the platinum. Something still has to be oxidized at the anode, so water gives up the electrons, and those bubbles are oxygen gas, while the platinum itself doesn’t lose any mass at all.

Let’s put the two kinds of cell side by side. A galvanic cell runs a favored reaction and gives you electrical energy, and an electrolytic cell runs an unfavored reaction because a power supply forces it to. But look at the bottom two rows, because in both kinds of cell, oxidation happens at the anode and reduction happens at the cathode.

There’s one thing we still haven’t explained, though. The supply had to beat one point one zero volts, exactly what the meter read, so where does that number come from, and could we have predicted it?