Chemistry 30 · Electrochemical cells
Cells that move electrons
A voltaic cell lets a spontaneous redox reaction push electrons through a wire; an electrolytic cell uses electricity to force a reaction that would never happen on its own. Same parts, opposite directions — and one sign flip that catches everyone.
- 1. The parts of a voltaic cell
- 2. Cell notation and voltage
- 3. Electrolytic cells
- 4. Plating, corrosion and calculations
The parts of a voltaic cell
The idea: separate the two half-reactions into different containers and the electrons can only get from one to the other through a wire — which is a current you can use.
| Part | What happens there |
|---|---|
| Anode | Oxidation. The electrode loses mass as metal atoms become ions. |
| Cathode | Reduction. The electrode gains mass as ions plate onto it. |
| Wire | Carries electrons from anode to cathode. |
| Salt bridge | Lets ions move so neither solution builds up charge. Without it the cell stops almost at once. |
| Electrolyte | The solution each electrode sits in, supplying or receiving ions. |
AN OX and RED CAT — ANode OXidation, REDuction CAThode — holds in both kinds of cell. What changes between them is the sign on the electrodes, not the chemistry.
Cell notation and voltage
The idea: the shorthand is always anode on the left, cathode on the right, with a double line for the salt bridge.
Zn(s) | Zn²⁺(aq) || Cu²⁺(aq) | Cu(s)
- A single line is a phase boundary — solid metal meeting its solution.
- The double line is the salt bridge separating the half-cells.
- Reading it: zinc metal is oxidized on the left, copper ions are reduced on the right.
E°cell = E°(cathode) − E°(anode), using reduction potentials for both. For this cell, 0.34 − (−0.76) = +1.10 V. A positive cell potential is what makes it a battery rather than a science project that does nothing.
Watch out: do not multiply a potential when you scale a half-reaction. Voltage is energy per unit charge, so it is the same whether one electron moves or a mole of them.
Electrolytic cells
The idea: connect a power supply and you can force the reaction that would not happen by itself — running the same chemistry backwards.
| Voltaic | Electrolytic | |
|---|---|---|
| Energy | Chemical → electrical | Electrical → chemical |
| E°cell | Positive; spontaneous | Negative; needs a power supply |
| Anode | Negative electrode | Positive electrode |
| Cathode | Positive electrode | Negative electrode |
| Oxidation | At the anode, in both | |
| Examples | Batteries, fuel cells | Electroplating, refining aluminum, electrolysis of water |
The sign swap is the only thing to memorize, and it has a reason: in a voltaic cell the anode pushes electrons out, so it is the negative terminal. In an electrolytic cell the power supply pulls electrons out of the anode, so that electrode is held positive.
Plating, corrosion and calculations
The idea: once you can see which electrode does what, the applications and the arithmetic both follow.
- Electroplating: the object being coated is the cathode, because plating means reducing metal ions onto it. The pure metal bar is the anode, dissolving to replace what plates out.
- Corrosion is oxidation of a metal by its surroundings. Iron rusting is Fe → Fe²⁺ + 2e⁻, with oxygen and water doing the reducing.
- Cathodic protection: attach a metal that is oxidized more readily — zinc or magnesium — and it corrodes instead. It is called a sacrificial anode because it is deliberately spent.
Worked: how much metal plates out
A current of 2.00 A runs for 30.0 minutes through a solution of Cu²⁺.
- Charge: q = It = (2.00 A)(1800 s) = 3600 C.
- Moles of electrons: 3600 C ÷ 96 500 C/mol = 0.0373 mol.
- Moles of copper: Cu²⁺ needs 2 electrons each, so 0.0373 ÷ 2 = 0.0187 mol.
- Mass: 0.0187 mol × 63.55 g/mol = 1.19 g.
The step people skip is the third: the charge on the ion decides how many electrons each atom needs, so Al³⁺ takes three times as much charge per mole as a 1+ ion would.