Electrolysis harnesses electricity to drive non-spontaneous redox reactions. It's used in various industries for electroplating, metal refining, and chemical production. Understanding the process and products of electrolysis is key to grasping its wide-ranging applications.

Faraday's laws help calculate the amount of substance produced or consumed during electrolysis. These laws, along with concepts like current efficiency and overpotential, are crucial for optimizing electrolytic processes and understanding their real-world limitations.

Electrolysis

Process and applications of electrolysis

• Electrolysis utilizes electrical energy to drive a non-spontaneous redox reaction
  • Requires an external power source such as a battery or electrical outlet
  • Oxidation occurs at the anode which is the positive electrode
  • Reduction occurs at the cathode which is the negative electrode
• Applications of electrolysis encompass various industries and processes
  • Electroplating deposits a thin layer of metal onto a surface (jewelry making, corrosion protection, decorative finishes)
  • Electrolytic refining purifies metals such as copper and aluminum
  • Production of chemical elements and compounds including chlorine and sodium hydroxide
  • Electrochemical machining removes material from a workpiece using electrolysis

Products of electrolysis in solutions

• In molten salts, the cation reduces at the cathode and the anion oxidizes at the anode
  • In molten NaCl, Na+ reduces to Na (s) at the cathode and Cl- oxidizes to Cl2 (g) at the anode
• In aqueous solutions, the products depend on the reactivity of the electrodes and the ions present
  • At the cathode, the order of reduction preference from most to least preferred:
    1. Active metal cations (Al3+, Mg2+)
    2. H+ from H2O to form H2 (g)
    3. Less active metal cations (Cu2+, Ag+)
  • At the anode, the order of oxidation preference from most to least preferred:
    1. Anions of strong acids (Cl-, Br-)
    2. OH- from H2O to form O2 (g)
    3. Less reactive anions (SO42-, NO3-)

Faraday's Laws and Electrolytic Cells

Calculations using Faraday's laws

• Faraday's first law states the mass of a substance produced or consumed at an electrode is directly proportional to the quantity of electricity passed through the cell
  • $m = \frac{Q}{F} \cdot \frac{M}{z}$, where:
    • $m$ represents the mass of the substance (g)
    • $Q$ represents the quantity of electricity (coulombs, C)
    • $F$ represents Faraday's constant (96,485 C/mol)
    • $M$ represents the molar mass of the substance (g/mol)
    • $z$ represents the number of electrons transferred per formula unit
• Faraday's second law states the masses of different substances produced or consumed by the same quantity of electricity are proportional to their equivalent weights which is molar mass divided by the number of electrons transferred

Efficiency and overpotential in electrolytic cells

• Current efficiency is the ratio of the actual amount of product formed to the theoretical amount predicted by Faraday's laws
  • Expressed as a percentage: $\text{Current efficiency} = \frac{\text{Actual yield}}{\text{Theoretical yield}} \times 100%$
  • Inefficiencies can arise from side reactions, back reactions, or physical losses
• Overpotential is the additional potential required beyond the thermodynamic potential to drive an electrolysis reaction at a desired rate
  • Caused by factors such as activation energy barriers, concentration gradients, and resistance in the cell
  • Types of overpotential include activation overpotential, concentration overpotential, and resistance overpotential
  • Higher overpotentials lead to increased energy consumption and decreased current efficiency