Electrochemical Cells and Standard Potentials

A cell's voltage comes from the difference between the standard reduction potentials of its two half-cells. The sign of that difference tells you whether the cell reaction is spontaneous (a galvanic cell) or must be driven by a power supply (an electrolytic cell), and it sets the Gibbs energy change. Work through the three stages in order.

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Anode (left): oxidation

Cathode (right): reduction

Remember

The cell is drawn as it is written in cell notation: the anode (oxidation) on the left and the cathode (reduction) on the right.
E°_cell = E°(cathode) − E°(anode), using standard reduction potentials. E° values are not multiplied by the coefficients in the balanced equation.
E°_cell > 0: the reaction as written is spontaneous and the cell is galvanic. E°_cell < 0: the reaction must be driven in an electrolytic cell by a power supply of at least |E°_cell|.
ΔG° = −nFE°_cell, where n is the number of electrons transferred in the reaction as written and F = 96,485 C mol⁻¹. E°_cell does not depend on how the equation is written; ΔG° does.
E°(H⁺/H₂) is defined as exactly 0 V, so a cell with the standard hydrogen electrode (SHE) as its anode reads the cathode's E° directly.
E°_cell is measured on open circuit, when no current flows. Standard conditions: 1 M, 1 bar, 298.15 K.

Standard reduction potentials at 25 °C are those in the Lecture 1 table; Mg²⁺/Mg and Pb²⁺/Pb are from the CRC Handbook, rounded to two decimal places.