Two kinds of electrochemical cell
In a redox reaction, electrons pass from a reducing agent to an oxidizing agent. If the two are put in separate containers joined by a wire, the electrons must travel through the wire: an electric current. Agalvanic (or voltaic) cell uses a reaction that happens by itself to produce electricity. Anelectrolytic cell uses electricity from a power supply to drive a reaction that would not happen by itself. In both, oxidation happens at the anode and reduction at the cathode.
Inside a galvanic cell
The classic example is the Daniell cell: a zinc electrode in zinc sulfate solution and a copper electrode in copper(II) sulfate solution.
Anode (oxidation): Zn → Zn²⁺ + 2e⁻Cathode (reduction): Cu²⁺ + 2e⁻ → CuZinc atoms leave the anode as ions, and the electrons they leave behind make it the negative electrode. The electrons flow through the external circuit to the copper cathode, the positive electrode, where copper ions pick them up and are deposited as copper. Inside the cell, the current is carried by ions. Thesalt bridge, a tube of salt solution or gel, lets negative ions drift towards the anode and positive ions towards the cathode, keeping both solutions neutral. Electrons never travel through the solutions.
Electrode potentials and cell voltage
Each half-cell has a standard reduction potential, E°, measured against a standard hydrogen electrode at 25 °C with 1 mol/L solutions. The more negative E° is, the more easily the metal gives up electrons. The half-cell with the lower potential becomes the anode, and the cell voltage is:
E°cell = E°cathode − E°anode| Half-reaction | E° (V) |
|---|---|
| Mg²⁺ + 2e⁻ → Mg | −2.37 |
| Zn²⁺ + 2e⁻ → Zn | −0.76 |
| Fe²⁺ + 2e⁻ → Fe | −0.44 |
| Pb²⁺ + 2e⁻ → Pb | −0.13 |
| Cu²⁺ + 2e⁻ → Cu | +0.34 |
| Ag⁺ + e⁻ → Ag | +0.80 |
For zinc and copper, E°cell = 0.34 − (−0.76) = 1.10 V. Metals far apart in the list make bigger voltages. Note that E° does not depend on how many electrons are involved: the silver half-reaction is not doubled when it is balanced.
Why batteries run down
The voltage also depends on concentration, through the Nernst equation. At 25 °C, for a metal ion of charge n:
E = E° + (0.0257 / n) × ln[Mⁿ⁺]As the cell works, cathode ions are used up and anode ions build up, so the two electrode potentials move closer and the voltage falls. When the reaction reaches equilibrium, the voltage is zero: the battery is flat. The simulator shows this on the voltage graph.
Electrolysis
In electrolysis, a power supply pushes electrons into the cathode, which is nownegative, and pulls them from the anode, which is positive. Positive ions move to the cathode and are reduced; negative ions or water molecules are oxidized at the anode. The supply must provide at least the cell’s reversed voltage, plus extra “overpotential” to make the reactions go at a useful rate. Water needs at least 1.23 V to split, but in practice about 1.7 V or more.
How much product? Faraday’s law
Every electron passed reduces or oxidizes a fixed amount of substance. One mole of electrons carries 96,485 coulombs, the Faraday constant F.
charge Q = current I × time tmoles of product = Q ÷ (z × F)Here z is the number of electrons per particle of product (2 for Cu²⁺ → Cu or for H₂, 4 for O₂). That is why electrolysis of water gives twice the volume of hydrogen as oxygen.
Galvanic vs electrolytic
| Galvanic cell | Electrolytic cell | |
|---|---|---|
| Energy | Chemical → electrical | Electrical → chemical |
| Reaction | Spontaneous | Non-spontaneous (driven) |
| Anode | Oxidation, negative | Oxidation, positive |
| Cathode | Reduction, positive | Reduction, negative |
| Example | Batteries, fuel cells | Electroplating, splitting water |
Real-world applications
- Batteries, from alkaline cells to the lithium-ion cell in your phone, are galvanic cells; recharging runs them as electrolytic cells.
- Electroplating coats objects in chromium, nickel, silver or gold.
- Copper refining dissolves impure copper anodes and deposits pure copper on the cathode.
- Aluminium is extracted from its molten oxide by electrolysis.
- Hydrogen for fuel can be made by electrolysis of water using renewable electricity.
- Corrosion protection: blocks of magnesium or zinc are bolted to ships and pipelines, so that they corrode instead of the steel.
Common misconceptions
- “The anode is always positive.” It is negative in a galvanic cell and positive in an electrolytic cell; it is always where oxidation happens.
- “Electrons flow through the salt bridge or the solution.” Ions carry the current there; electrons only travel through the wires and electrodes.
- “A battery stores electricity.” It stores chemicals that react to produce electricity.
- “Doubling a half-equation doubles E°.” Electrode potentials do not depend on the amounts written.
What the model simplifies
Standard electrode potentials and the Faraday constant are real values, and the Nernst equation is applied with concentrations in place of activities, which is only approximate at high concentrations. The internal resistance (4 Ω), the bulb (6 Ω), the 100 mL half-cells, the 20 g electrodes and the electrolysis onset voltages are illustrative. Time is speeded up (1 s = 1 hour for the galvanic cell, 1 minute for electrolysis), and the drawn electrode sizes exaggerate mass changes so they can be seen. Real magnesium also reacts slowly with water, chlorine is partly dissolved in the solution, and real cells slow down because of overpotentials and the time ions take to reach the electrodes. The ions and electrons are a handful of symbolic dots, not to scale.
Key takeaways
- Electrochemical cells separate oxidation (at the anode) from reduction (at the cathode).
- In a galvanic cell a spontaneous reaction drives electrons from anode to cathode through the external wire.
- E°cell = E°cathode − E°anode; the voltage falls as the concentrations change.
- Ions, not electrons, carry the charge through solutions and the salt bridge.
- Electrolysis uses electricity to drive non-spontaneous reactions, and the amount of product is proportional to the charge passed.