Electrochemistry: Redox, Galvanic Cells & Electrolysis

Electrochemistry sits at the intersection of electricity and chemistry. It covers two complementary processes: galvanic cells (where spontaneous chemical reactions generate electrical energy) and electrolysis (where electrical energy drives non-spontaneous chemical reactions). From the battery powering your phone to the industrial production of chlorine and aluminium, electrochemistry is everywhere in modern life.

Oxidation & Reduction: OIL RIG

  • OIL RIG: Oxidation Is Loss (of electrons); Reduction Is Gain (of electrons).
  • Redox reactions always occur together — when one species loses electrons (is oxidised), another gains those electrons (is reduced).
  • Example: Zn(s) + CuSO₄(aq) → ZnSO₄(aq) + Cu(s). Zinc is oxidised (Zn → Zn²⁺ + 2e⁻); copper ions are reduced (Cu²⁺ + 2e⁻ → Cu).

Oxidation States

  • Pure elements: oxidation state = 0. Oxygen in compounds: usually −2. Hydrogen in compounds: usually +1.
  • Sum of oxidation states in a neutral compound = 0; in a polyatomic ion = charge of the ion.
  • A species is oxidised if its oxidation state increases; reduced if it decreases.

Galvanic (Voltaic) Cells

  • A galvanic cell converts chemical energy into electrical energy using a spontaneous redox reaction.
  • Anode (oxidation, negative electrode): the more reactive metal (e.g. Zn → Zn²⁺ + 2e⁻).
  • Cathode (reduction, positive electrode): the less reactive metal (e.g. Cu²⁺ + 2e⁻ → Cu).
  • E°cell = E°cathode − E°anode.

Electrolysis

  • Electrolysis uses an external electrical supply to force non-spontaneous chemical reactions.
  • At the cathode (−): cations are reduced. At the anode (+): anions are oxidised.
  • Faraday's Laws: mass deposited ∝ charge passed (Q = It) and molar mass ÷ ionic charge.
  • Electrolysis of brine produces chlorine (anode), hydrogen (cathode), and sodium hydroxide solution.

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