Cells

Rucete ✏ Chemistry In a Nutshell

1. Galvanic (Voltaic) Cells

• Convert chemical energy into electrical energy through spontaneous redox reactions.
• Example reaction: Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s)
• Anode (oxidation): Zn(s) → Zn²⁺ + 2e⁻ (E°ox = +0.76 V)
• Cathode (reduction): Cu²⁺(aq) + 2e⁻ → Cu(s) (E°red = +0.34 V)
• Overall Cell Voltage: E°cell = 0.34 V - (−0.76 V) = 1.10 V
• Electrons flow from anode to cathode.
• A salt bridge maintains electrical neutrality.

Cell notation: Zn ∣ Zn²⁺ ∣∣ Cu²⁺ ∣ Cu

2. Electrolytic Cells

• Use electrical energy to drive non-spontaneous reactions.
• Opposite of galvanic cells in electron flow and polarity:
• Anode is positive, cathode is negative.
• Example: Electroplating or electrolysis of molten salts.

3. Thermodynamics of Electrochemical Cells

• Gibbs Free Energy:
  

  $$\mathrm{\Delta }{G}^{\circ }=-nF{E}_{\text{cell}}^{\circ }$$

• Relationship to Equilibrium Constant: $$E_{\text{cell}}^{\circ }=\frac{0.0591}{n}\log K$$
  $$K={10}^{\frac{n{E}_{\text{cell}}^{\circ }}{0.0591}}$$
• Where:
  • n = number of moles of electrons
  • F = Faraday’s constant
  • K = equilibrium constant