Standard electrode potentials
Standard electrode potentials
- The electrode potential $E$ shows how easily a half-cell is reduced.
- We measure it against a fixed reference under standard conditions.
- Combining two half-cells gives a cell potential.
The reference and measuring E°
- The standard hydrogen electrode is the zero reference: $\text{H}_2$ at 1 atm over platinum in $1\ \dfrac{\text{mol}}{\text{dm}^3}$ $\text{H}^+$, defined as 0.00 V.
- To measure a half-cell's $E^{\ominus}$, connect it to the hydrogen electrode and read the voltage (always written as a reduction).
Practice
The standard hydrogen electrode is defined as having a potential of:
It is the reference half-cell, set at exactly 0.00 V; all other E° values are measured against it.
Combining half-cells
$$E^{\ominus}_{\text{cell}} = E^{\ominus}(\text{more positive}) - E^{\ominus}(\text{less positive})$$
- Polarity: the more negative electrode is the negative terminal; electrons flow from it through the circuit.
- Reactivity: a more positive $E^{\ominus}$ → better oxidising agent; a more negative $E^{\ominus}$ → better reducing agent.
Practice
The standard cell potential is:
Subtract the less positive electrode potential from the more positive one.
Practice
A half-cell with a more positive E° is:
A more positive E° means the species is readily reduced — a stronger oxidising agent.
Practice
In a cell, electrons flow through the external circuit from the:
The more negative electrode is the negative terminal; electrons flow from it to the more positive electrode.
You've got it
Key idea
- $E^{\ominus}$ shows how easily a half-cell is reduced (written as a reduction)
- the standard hydrogen electrode = 0.00 V reference
- $E^{\ominus}_{\text{cell}} = E^{\ominus}(\text{more positive}) - E^{\ominus}(\text{less positive})$
- more positive $E^{\ominus}$ = stronger oxidising agent; more negative = stronger reducing agent