The Nernst equation relates cell potential to non-standard conditions: E = $E_{standard}$ - $\frac{RT}{nF}$ln(Q). At 25 degrees C (298 K), this simplifies to E = $E_{standard}$ - $\frac{0.0592}{n}$log(Q). The variable n is the number of electrons transferred in the balanced cell reaction. Q is the reaction quotient with products over reactants. For a concentration cell (same electrodes, different concentrations), $E_{standard}$ = 0, so E = -$\frac{0.0592}{n}$log(Q). The cell drives toward equalising concentrations. At equilibrium, E = 0, giving $E_{standard}$ = $\frac{0.0592}{n}$log(K). This connects electrochemistry to equilibrium constants. Common error: using the wrong value of n or forgetting to balance the equation first.