Voltaic (Galvanic) Cell
Voltaic (Galvanic) Cells
A voltaic cell converts chemical energy → electrical energy via a spontaneous redox reaction. The two half-cells are separated so electrons must travel through an external wire.
Key Components
- Anode (−): Oxidation occurs here. The more reactive metal dissolves, releasing electrons into the wire.
- Cathode (+): Reduction occurs here. Metal ions in solution gain electrons and deposit as solid metal.
- External wire: Electrons flow from anode → cathode through the wire, producing a current.
- Salt bridge: Contains unreactive ions (e.g. KNO₃). Allows ion migration to maintain electrical neutrality. Cations move toward the cathode, anions toward the anode.
- Voltmeter: Measures the potential difference (EMF / voltage) of the cell.
Electrolytic Cells
An electrolytic cell uses electrical energy → chemical energy to force a non-spontaneous reaction to occur. An external power supply drives electrons in the opposite direction.
Example: Electrolysis of Molten NaCl
Cathode (−): Na⁺ + e⁻ → Na (reduction. Sodium metal deposited)
Anode (+): 2Cl⁻ → Cl₂ + 2e⁻ (oxidation. Chlorine gas evolved)
Overall: 2NaCl(l) → 2Na(l) + Cl₂(g)
Voltaic vs Electrolytic. Comparison
| Voltaic | Electrolytic | |
|---|---|---|
| Energy | Chemical → Electrical | Electrical → Chemical |
| Spontaneous? | Yes (ΔG < 0) | No (ΔG > 0) |
| Anode | Negative (−) | Positive (+) |
| Cathode | Positive (+) | Negative (−) |
| Power source | None (self-generating) | External battery/supply |
In BOTH types: oxidation at anode, reduction at cathode. An Ox, Red Cat.
Standard Electrode Potential E° (HL)
The standard electrode potential (E°) measures the tendency of a half-cell to undergo reduction under standard conditions (298 K, 1 mol dm⁻³, 1 atm). All values are measured relative to the Standard Hydrogen Electrode (SHE), which is assigned E° = 0.00 V.
Interpreting E° Values
- More positive E° → stronger tendency to be reduced → stronger oxidising agent
- More negative E° → stronger tendency to be oxidised → stronger reducing agent
Calculating E°cell
E°cell = E°cathode − E°anode
If E°cell > 0 → reaction is spontaneous (voltaic cell works).
If E°cell < 0 → reaction is non-spontaneous (electrolysis required).
Worked Example: Zn–Cu Cell E°cell
Given: E°(Cu²⁺/Cu) = +0.34 V, E°(Zn²⁺/Zn) = −0.76 V
Cu²⁺ has the more positive E° → it is reduced (cathode).
Zn has the more negative E° → it is oxidised (anode).
E°cell = (+0.34) − (−0.76) = +1.10 V
Positive → spontaneous. This cell will produce electricity.
Gibbs Free Energy & Cell Potential (HL)
ΔG° = −nFE°cell
where n = moles of electrons transferred, F = Faraday constant (96 485 C mol⁻¹).
Positive E°cell → negative ΔG° → spontaneous.
Electroplating
Electroplating uses electrolysis to coat an object with a thin layer of metal. The object to be plated is the cathode; the plating metal is the anode. The electrolyte contains ions of the plating metal.
Think About It
Why does the salt bridge contain ions like KNO₃ rather than reactive metal ions?
K⁺ and NO₃⁻ are spectator ions. They don't participate in redox reactions. The salt bridge completes the circuit by allowing ion flow to maintain electrical neutrality in each beaker, without interfering with the cell reaction.