Galvanic (voltaic) and electrolytic cells — AP Chemistry
1. Core Concepts and Cell Notation ★★☆☆☆ ⏱ 3 min
Both galvanic (voltaic) and electrolytic cells are electrochemical cells that separate oxidation and reduction half-reactions to force electron flow through an external wire, rather than direct electron transfer between reactants.
Standard notation for all electrochemical cells follows a universal convention: anode (oxidation half-reaction) is placed on the left, cathode (reduction half-reaction) on the right. A single vertical line | denotes a phase boundary, and a double vertical line || denotes a salt bridge or porous disk that separates half-cells and maintains charge neutrality.
Exam tip: If asked for the direction of anion flow in the salt bridge, remember anions always go to the anode, regardless of cell type — this is a common low-stakes MCQ trap.
2. Electrode Charge and Identification ★★★☆☆ ⏱ 3 min
Only the charge of electrodes reverses between galvanic and electrolytic cells — the reaction at each electrode never changes: oxidation is always at the anode, reduction always at the cathode.
For galvanic cells: the anode produces electrons via oxidation, so it carries a negative charge. Electrons flow from the negative anode through the external wire to the positive cathode, where reduction consumes electrons. The salt bridge maintains charge neutrality: anions flow to the anode to balance positive charge buildup from new cations, and cations flow to the cathode to balance negative charge buildup from consumed cations.
For electrolytic cells: an external battery drives the non-spontaneous reaction, pulling electrons away from the anode and pushing electrons onto the cathode. This gives the anode a positive charge and the cathode a negative charge, but oxidation still occurs at the anode and reduction still at the cathode. Ion flow follows the same rule as galvanic cells: anions flow to the anode, cations flow to the cathode, regardless of cell type.
3. Standard Cell Potential and Spontaneity ★★★☆☆ ⏱ 4 min
Standard reduction potentials ($E^\circ_{\text{red}}$) are tabulated for all common half-reactions, measured relative to the standard hydrogen electrode (SHE), which is assigned $E^\circ_{\text{red}} = 0\ \text{V}$.
E^\circ_{\text{cell}} = E^\circ_{\text{red (cathode)}} - E^\circ_{\text{red (anode)}}
An equivalent form is $E^\circ_{\text{cell}} = E^\circ_{\text{red (cathode)}} + E^\circ_{\text{ox (anode)}}$, where $E^\circ_{\text{ox (anode)}} = -E^\circ_{\text{red (anode)}}$, so both formulas give the same result. A key relationship between cell potential and Gibbs free energy is:
\Delta G^\circ = -nFE^\circ_{\text{cell}}
where $n$ = moles of electrons transferred in the balanced overall reaction, and $F$ = Faraday's constant ($96500\ \text{C/mol }e^-$). From this formula, we get the spontaneity rule: if $E^\circ_{\text{cell}}$ is positive, $\Delta G^\circ$ is negative, and the reaction is spontaneous (galvanic cell). If $E^\circ_{\text{cell}}$ is negative, $\Delta G^\circ$ is positive, and the reaction is non-spontaneous (requires external voltage, so electrolytic cell).
Exam tip: If you struggle with sign errors when calculating $E^\circ_{\text{cell}}$, flip the sign of the anode's reduction potential to get oxidation potential, then add to the cathode's reduction potential. This eliminates subtracting negative numbers.
4. Faraday's Law of Quantitative Electrolysis ★★★★☆ ⏱ 4 min
Electrolytic cells are used to drive non-spontaneous redox reactions for industrial applications. We can relate the amount of product produced to the current passed through the cell and time of electrolysis using Faraday's law of electrolysis.
The core relationship is that total charge ($Q$, measured in coulombs, C) equals current ($I$, measured in amperes, A, where $1\ \text{A} = 1\ \text{C/s}$) multiplied by time ($t$, measured in seconds):
Q = I \times t
Total moles of electrons transferred is then $n_{e^-} = \frac{Q}{F} = \frac{I t}{F}$. We use the stoichiometry of the reduction half-reaction to relate moles of electrons to moles of product, then convert moles to mass or volume as needed. For example, to produce 1 mole of Al from $\text{Al}^{3+}$, you need 3 moles of electrons.
Exam tip: Always convert time to seconds before calculating charge. AP exam questions frequently give time in minutes or hours to test unit conversion recall.
Common Pitfalls
Why: Students memorize charge only for galvanic cells and forget it reverses for electrolytic cells, where the external battery pulls electrons from the anode to make it positive
Why: Students confuse intensive properties ($E^\circ$) with extensive properties ($\Delta G$, enthalpy) that do scale with reaction size
Why: Students mix up ion flow and electron flow when recalling how the circuit is completed
Why: Students rush from moles of electrons straight to mass without writing the half-reaction
Why: Students associate all electrochemical cells with spontaneous galvanic cells that produce voltage
Why: Students confuse the order of notation with charge sign