Galvanic vs Electrolytic Cells Cheat Sheet

Key Facts

• Oxidation is loss of electrons, reduction is gain of electrons, and this can be remembered as OIL RIG.
• Oxidation always occurs at the anode, and reduction always occurs at the cathode in both galvanic and electrolytic cells.
• In a galvanic cell, the reaction is spontaneous, so $E^{\circ}_{\text{cell}} > 0$ and $\Delta G^{\circ} < 0$.
• In an electrolytic cell, the reaction is nonspontaneous, so an external power source is required and $E^{\circ}_{\text{cell}} < 0$ for the reaction as written.
• The standard cell potential is calculated with $E^{\circ}_{\text{cell}} = E^{\circ}_{\text{cathode}} - E^{\circ}_{\text{anode}}$ using reduction potentials.
• The relationship between cell potential and free energy is $\Delta G^{\circ} = -nFE^{\circ}_{\text{cell}}$, where $n$ is moles of electrons and $F = 96485\ \text{C/mol e}^{-}$.
• The Nernst equation at $25^{\circ}\text{C}$ is $E_{\text{cell}} = E^{\circ}_{\text{cell}} - \frac{0.0592}{n}\log Q$.
• In line notation, the anode is written on the left, the cathode is written on the right, and the salt bridge is shown with $||$.

Vocabulary

Galvanic cell

A cell that converts chemical energy into electrical energy using a spontaneous redox reaction.

Electrolytic cell

A cell that uses electrical energy from an outside source to force a nonspontaneous redox reaction.

Anode

The electrode where oxidation occurs in any electrochemical cell.

Cathode

The electrode where reduction occurs in any electrochemical cell.

Cell potential

The voltage of an electrochemical cell, calculated under standard conditions by $E^{\circ}_{\text{cell}} = E^{\circ}_{\text{cathode}} - E^{\circ}_{\text{anode}}$.

Salt bridge

A connection that allows ions to move between half-cells so charge does not build up.