Physical chemistry
Nernst equation
The Nernst equation extends electrode potentials beyond standard conditions by accounting for activities; it underlies concentration cells and electrochemical pH measurement.
Potential depends on the reaction quotient Q. For an overall reaction transferring n electrons, increasing product activity shifts cell potential logarithmically with Q.
Model and quantities
Read the relation together with assumptions about state, experimental conditions, and sign conventions. Keep units consistent and check dimensions before interpreting a result.
Definition: Nernst equation
E° is the standard potential, R the gas constant, T absolute temperature, F Faraday’s constant, n the balanced electron count, and Q the activity quotient. At 298 K, RT/F is about 0.0257 V.
Quantities in the relation are defined for the reaction or system at hand. In particular, distinguish standard-state quantities from actual conditions and do not infer a mechanism from a general expression alone.
Example: Worked example
At 25 °C, a cell has E°=0.80 V, n=1, and Q=10. Calculate E using RT/F≈0.0257 V.
Solution
E=0.80−0.0257 ln(10)≈0.741 V. Since Q>1, the correction lowers the potential relative to E°.
| Concept | Description | Unit / note |
|---|---|---|
| Key relation | Use under stated conditions | Check units and sign convention |
| Measured quantity | Relates a state or process | Compare data with model |
| Scope | Model specific conditions | Check assumptions first |
With an H⁺-selective electrode, potential depends on H⁺ activity, allowing a calibrated pH meter to convert voltage into pH. The ideal slope is about 59 mV per pH unit at 25 °C.
In the worked example, which result follows from the given data?
Which statement is consistent with this lesson?
References
- Peter Atkins, Julio de Paula (2014). Physical Chemistry