Electrochemistry

Redox Foundations

Assign oxidation states by the priority rules (free elements zero, oxygen $-2$, hydrogen $+1$ with nonmetals) and track electrons with OIL RIG: oxidation is loss, reduction is gain. The oxidizing agent is itself reduced; the reducing agent is oxidized. Balance redox equations as half-reactions, equalizing electrons before adding.

Galvanic and Electrolytic Cells

A galvanic (voltaic) cell runs a spontaneous reaction to produce current; an electrolytic cell consumes external current to drive a nonspontaneous one. In BOTH, oxidation occurs at the anode and reduction at the cathode (AN OX, RED CAT); the salt bridge maintains neutrality. Cell potential comes from reduction potentials: $E^\circ_{cell} = E^\circ_{cathode} - E^\circ_{anode}$ A positive $E^\circ_{cell}$ means spontaneous, linked to free energy by $\Delta G^\circ = -nFE^\circ_{cell}$

Nonstandard Conditions and Electrolysis

The Nernst equation corrects for concentration: $E = E^\circ - \frac{0.0592}{n}\log Q$ Faraday's law converts charge ($q = It$) into moles of electrons and grams deposited. Biologically, the electron transport chain is a series of redox couples ending at oxygen, and NAD$^+$/NADH shuttles electrons through metabolism.

Key Takeaways

• Anode oxidizes and cathode reduces in every cell type.
• Positive $E^\circ_{cell}$ pairs with negative $\Delta G^\circ$ and $K > 1$.
• Never flip the sign of a reduction potential when using the cathode-minus-anode formula.
• Electrolysis stoichiometry runs through moles of electrons.