Cell Potential Under Nonstandard Conditions

E° is the voltage only under standard conditions: 1 M solutions, 1 atm gases, 298 K. Real cells rarely sit there for long. As a reaction runs, reactants get used up and products build up, so Q changes, and the actual potential E drifts away from E°. The Nernst equation, E = E° - (0.0592/n) log Q at 298 K, tracks that drift. When Q is small (mostly reactants, far from equilibrium), log Q is very negative, so E is larger than E°. As the reaction proceeds and Q grows, E shrinks. That is why a battery's voltage sags continuously while it drains instead of holding steady until it suddenly dies.

A concentration cell pushes this to the extreme: both half-cells use the same species and the same standard electrode, so E° = 0 exactly. All the voltage comes from the concentration difference itself. The side with higher concentration acts as the cathode, because dilution there is favorable, driving reduction, while the dilute side oxidizes to supply ions into solution. Electrons flow through the wire from the dilute (anode) side to the concentrated (cathode) side, and the cell keeps working until both sides reach the same concentration.

This connects directly to equilibrium: at equilibrium Q equals K, and plugging that into the Nernst equation forces E to exactly zero. That is not a contradiction with potential being 'maximum' earlier — maximum potential occurs at the start, when the system is furthest from equilibrium, and it decreases monotonically to zero as the reaction approaches equilibrium, linking this topic to Delta-G = -nFE and the free energy ideas earlier in the unit.