Electrochemistry

Spontaneous redox → electrical energy.

• Anode: oxidation (−), electrons leave
• Cathode: reduction (+), electrons arrive
• Salt bridge: maintains neutrality
• Electron flow: anode → wire → cathode

$E^\circ_{cell} = E^\circ_{cathode} - E^\circ_{anode}$

Positive $E^\circ_{cell}$ → spontaneous.

• More positive $E^\circ$ → stronger oxidising agent (better at being reduced)
• More negative $E^\circ$ → stronger reducing agent (more easily oxidised)
• SHE: $E^\circ = 0.00$ V (reference)

$\Delta G^\circ = -nFE^\circ_{cell}$

where $n$ = moles of electrons, $F = 96485$ C/mol

$E^\circ_{cell} = \frac{RT}{nF}\ln K \quad \text{or} \quad \ln K = \frac{nFE^\circ}{RT}$

$E_{cell} = E^\circ_{cell} - \frac{RT}{nF}\ln Q$

At 298 K: $E_{cell} = E^\circ_{cell} - \frac{0.0592}{n}\log Q$

Non-spontaneous redox driven by external voltage.

• Anode still oxidation, cathode still reduction
• External power source reverses natural electron flow
• Examples: electroplating, electrolysis of water, aluminium extraction

Faraday's Laws

$n = \frac{It}{nF} \quad \text{or} \quad m = \frac{ItM}{nF}$

where $I$ = current (A), $t$ = time (s), $n$ = electrons per ion

• $E^\circ_{cell} = E^\circ_{cathode} - E^\circ_{anode}$; positive = spontaneous
• $\Delta G^\circ = -nFE^\circ$
• Nernst: $E = E^\circ - (0.0592/n)\log Q$
• Electrolysis: $m = ItM/(nF)$
• Galvanic: spontaneous; electrolytic: non-spontaneous (needs external voltage)