Cell potential under standard conditions (298.15 K, 1 atm, 1 M concentrations).

Cell potential under non-standard conditions.

A measure of the relative amount of products and reactants present in a reaction at a given time. Used to determine if a reaction is at equilibrium.

The charge of one mole of electrons, approximately 96485 C/mol e-.

Moles of electrons transferred in the balanced redox reaction.

Q = $\frac{[Al^{3+}]^2}{[Mn^{2+}]^3}$

1. Calculate Q. 2. If Q > 1, E_cell < E°_cell. 3. If Q < 1, E_cell > E°_cell.

Set E_cell = 0 and Q = K in the Nernst Equation, then solve for E°_cell.

$E_{cell} = E^circ_{cell} - \frac{RT}{nF}lnQ$ or $E_{cell} = E^circ_{cell} - \frac{0.0592}{n}logQ$

ΔG = -nFE_cell

E°_cell: Cell potential at standard conditions. | E_cell: Cell potential at non-standard conditions.

Q = 1: Standard conditions, all concentrations are 1 M. | Q ≠ 1: Non-standard conditions, concentrations are not 1 M.

At equilibrium: Q = K. | Non-equilibrium: Q ≠ K.

Non-equilibrium: Running galvanic cell, E_cell > 0. | Equilibrium: Dead battery, E_cell = 0.

Q > 1: E_cell < E°_cell. | Q < 1: E_cell > E°_cell.