Half-Cell
Definition and meaning of Half-Cell in chemistry.
A half-cell is an electrode immersed in an electrolyte solution, representing one half of an electrochemical cell, in which either an oxidation or a reduction half-reaction takes place.
In more detail
Two half-cells joined by a salt bridge (or porous membrane) and an external wire make a complete galvanic or electrolytic cell. Electrons flow through the external circuit from the half-cell undergoing oxidation (the anode) to the one undergoing reduction (the cathode), while ions migrate through the salt bridge to preserve electrical neutrality in each compartment. Each half-cell is characterized by a standard reduction potential, E°, measured relative to the standard hydrogen electrode (SHE), which is arbitrarily assigned E° = 0 V. Combining the two half-cell potentials gives the overall cell potential, E°cell = E°cathode − E°anode.
Key facts
| Field | Physical Chemistry |
|---|---|
| Reference electrode | Standard hydrogen electrode (SHE), E° = 0 V |
| Key quantity | Standard reduction potential, E° (volts) |
| Example half-reaction | Cu²⁺(aq) + 2e⁻ → Cu(s), E° = +0.34 V |
In the Daniell cell, a zinc strip in ZnSO4 solution (Zn²⁺/Zn half-cell, E° = −0.76 V) is connected via a salt bridge to a copper strip in CuSO4 solution (Cu²⁺/Cu half-cell, E° = +0.34 V), giving a cell potential of 1.10 V as zinc oxidizes and copper ions are reduced.
Frequently asked questions
Can a half-cell generate current on its own?
No. A single half-cell cannot sustain current by itself; it must be paired with another half-cell through a salt bridge and external circuit to complete the redox reaction and allow charge to flow.
How is a half-cell's potential measured?
Its potential cannot be measured in isolation, so it is measured relative to the standard hydrogen electrode, which is defined as 0 V under standard conditions.