Electrochemical Energy Systems Summary

Electrochemical Energy Systems

Introduction to Electrochemistry

  • Definition: The branch of chemistry that examines phenomena resulting from combined chemical and electrical effects.

Types of Electrochemical Processes

  • Electrolytic Processes: Chemical changes occur upon passage of electrical current.
  • Galvanic (Voltaic) Processes: Chemical reactions produce electrical energy.

Modes of Charge Transport

  • Electrodes: Movement of electrons (or holes).
  • Electrolyte: Movement of ions (positively and negatively charged).

What are Electrochemical Cells?

  • Definition: Devices that convert chemical energy from spontaneous redox reactions into electrical energy.
  • Structure:
    • Two electronic conductors (electrodes).
    • One ionic conductor (electrolyte).

Electrodes in Electrochemical Cells

  • Function: Sites where oxidation or reduction occurs.
    • Anode: Site of oxidation; electrons move from this electrode.
    • Cathode: Site of reduction.
    • Current flow dictates movement of electrons from anode to cathode.

Types of Electrochemical Cells

  • Galvanic Cells: Convert chemical energy to electrical energy (e.g., Daniell Cell).
  • Electrolytic Cells: Convert electrical energy to chemical potential energy (e.g., decomposition of water).

Half-Cells and Reactions

  • Electrochemical cells are composed of half-cells that contain oxidation and reduction half-reactions.
  • Redox Reactions: Involve the transfer of electrons between species. They can be represented as:
    Red{1} + Ox{2} \rightleftharpoons Red{2} + Ox{1}

Standard Electrode Potentials

  • The electrode potential is derived from the change in free energy:
    ext{AG} = - ext{NFE}
  • Standard oxidation and reduction potentials are equal in magnitude but opposite in sign.

Competing Reactions and Oxidation/Reduction

  • The species with lower standard reduction potential gets oxidized:
    • Example: Zinc (Zn) and Copper (Cu) reactions golden rules.

Key Concepts for Cells

  • The concept of RED CAT and AN OX:
    • RED CAT: Reduction occurs at the cathode.
    • AN OX: Oxidation occurs at the anode.

Faraday's Law

  • Definition: Relates quantity of current to the quantity of chemical change.
  • Mathematical Statement: m = \frac{MIt}{nF}
    • Where:
    • m = mass of substance
    • M = molecular weight
    • I = current (A)
    • t = time (s)
    • n = number of electrons transferred
    • F = Faraday constant (96485 C/equiv)

Batteries

Primary Batteries

  • Definition: Non-rechargeable, cell reaction is not reversible (e.g., Daniel Cell, Dry Cell).

Secondary Batteries

  • Definition: Rechargeable batteries, reaction can be reversed (e.g., Lead-acid battery, Ni-Cd battery).

Lead-Acid Battery

  • Invention: Created in 1859 by Gaston Planté.
  • Components:
    • Cathode: Lead dioxide.
    • Anode: Lead.
    • Electrolyte: Diluted sulfuric acid.
  • Reactions:
    • Oxidation at Anode:
      Pb \rightarrow Pb^{2+} + 2e^{-}
    • Reduction at Cathode:
      PbO{2} + 4H^{+} + 2e^{-} \rightarrow PbSO{4} + 2H_{2}O

Nickel-Cadmium (Ni-Cd) Batteries

  • Components: Spongy cadmium (anode) and nickel oxyhydroxide (cathode).
  • Reactions:
    Cd + 2OH^{-} \rightarrow Cd(OH){2} + 2e^{-} 2NiO(OH) + Cd + 2e^{-} \rightarrow 2Ni(OH){2} + 2OH^{-}

Lithium Batteries

  • Properties: High energy density, lightweight, and high power density.
  • Types:
    1. Lithium primary batteries: Non-rechargeable.
    2. Lithium-ion batteries: Rechargeable, uses lithium compounds as anodes.

Fuel Cells

  • Definition: Device that converts chemical energy from fuel directly into electricity.
  • Reaction: Fuel reacts with oxidizing agent (e.g., oxygen) to produce electricity, water being the only byproduct.
  • Types include Alkaline, Direct Methanol, and Polymer Electrolyte Membrane fuel cells.
  • Advantages:
    • Zero emissions.
    • High efficiency and power density.
    • Quiet operation and no need for recharging.