What a learner can do afterwards
- Writes the reaction quotient for a cell reaction and substitutes it into the Nernst equation with the right electron count
- Calculates the potential of a concentration cell and explains why it is not zero
- Converts a standard cell potential into an equilibrium constant and states the link to Gibbs energy
- Explains why the quotient should strictly contain activities and when the concentration form is good enough
1 · Read
A table potential assumes standard conditions, about 1 M for every solute, but your cell rarely has them. Correct it with E = E standard - (R T / n F) ln Q. Write Q from the balanced cell reaction, each species raised to its stoichiometric power, and read n from the balanced half reactions.
In a concentration cell both electrodes are the same metal, so E standard is zero, yet a voltage appears. Oxidation runs in the dilute half cell and reduction in the concentrated one. For copper at 0.010 M against 1.0 M, Q = 0.010 and n = 2, so at 25 C E = -(0.0257 / 2) x ln 0.010 = +0.059 V. As Q nears 1 the voltage fades, and more electrons give a smaller voltage for the same ratio.
Since minus n F E standard equals minus R T ln K, a standard potential gives an equilibrium constant. For E standard = 0.10 V and n = 2 at 25 C, ln K = 2 x 0.10 / 0.0257 = 7.8, so K is about 2.4e3. Positive E standard means K above 1, so products dominate. As a cell runs down, Q approaches K and E falls to zero.
Strictly, Q holds activities. In dilute solutions concentrations are close enough; in concentrated brines they are not, so use activities there.
You correct the standard potential with Q and n, read concentration cells from Q alone, link E standard to K, and use activities when solutions are concentrated.
Rewritten readings were written for this dataset. The remaining reading summary '17.4: The Nernst Equation' is adapted from OpenStax Chemistry (CC BY 4.0), openstax.org.
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