In-Depth Notes on Electrolysis and Electrochemical Cells 36

Overview of Electrolysis and Electrochemical Cells

  • Electrolysis: A process driven by an external voltage that prompts a nonspontaneous chemical reaction.

  • Electrochemical Cells: Two types of electrochemical cells exist:

    • Galvanic Cells: Spontaneous reaction that generates electrical energy.
    • Electrolytic Cells: Nonspontaneous reaction requiring external power.

Key Concepts in Electrochemistry

  • Redox Reactions: Reduction and oxidation occur simultaneously:

    • Oxidation: Loss of electrons.
    • Reduction: Gain of electrons.

  • Electrochemical Potential: The ability of a species to gain or lose electrons, measured in volts (V). Each half-reaction has a standard reduction potential (E°).

Electrochemical Potential and Terminology

  • Standard Electrode Potential: A measurement of the tendency of a chemical species to be reduced, represented by E°.

    • E.g., Cr³⁺ + e⁻ ⇄ Cr²⁺ has E° = -0.424 V.

  • Galvanic Cell Components:

    • Anode: Site of oxidation (negative electrode).
    • Cathode: Site of reduction (positive electrode).

Example Problem Analysis

  • Given System:

    • Cell: Pt(s)|Cr²⁺ (aq, 0.10M), Cr³⁺ (aq, 0.20M)||Cu²⁺ (aq, 0.10M)|Cu(s).

  • Cell Diagram: Visual representation of the electrochemical cell, identifying anode and cathode.

  • Determine Galvanic Nature: Check reaction spontaneity using standard reduction potentials to predict cell behavior.

Understanding Batteries as Galvanic Cells

  • Functionality: Batteries act as galvanic cells by storing and converting chemical energy into electrical energy.

  • Alkaline Dry Cell Example:

    • Anode reaction: Zn (s) + 2OH⁻ (aq) → Zn(OH)₂ (s) + 2e⁻.
    • Cathode reaction: 2MnO₂ (s) + H₂O (l) + 2e⁻ → Mn₂O₃ (s) + 2OH⁻ (aq).
    • Total reaction calculated to confirm 1.5V outputs across different battery sizes (AAA, AA, etc.).

Electrolysis of Water

  • Electrolytic Process: When energy is applied (e.g., 9V), water (H₂O) decomposes into hydrogen (H₂) and oxygen (O₂).

    • Half-Reactions:
    • At cathode: 2H⁺ (aq) + 2e⁻ → H₂ (g).
    • At anode: O₂ + 4H⁺ (aq) + 4e⁻ → 2H₂O.

  • Equilibrium and Reversibility: Understanding the conditions under which electrolysis occurs and how products relate to reactants.

Running an Electrolysis Reaction

  • Impact of Salt Addition: Increasing ion concentration in water improves conductivity, enhancing product yield (H₂ & O₂).

  • Example of Electroplating: Utilizing electrolysis to deposit a metal layer onto a surface, through reactions such as Ag⁺ + e⁻ → Ag (s).

Quantitative Analysis of Electrolysis

  • Faraday’s Law: Interrelates the charge (Q) to the amount of substance deposited:

    • q=nFq = nF, where n is moles of electrons, F is Faraday's constant (96485 C/mol e⁻).

  • Example Problem: Given a current of 8.0 A over 10 seconds, calculate the mass of copper deposited using the electrolysis formula related to moles of electrons.