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:
  $Red<em>{1} + Ox</em>{2} \rightleftharpoons Red<em>{2} + Ox</em>{1}$

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:
    $PbO<em>{2} + 4H^{+} + 2e^{-} \rightarrow PbSO</em>{4} + 2H_{2}O$

Nickel-Cadmium (Ni-Cd) Batteries

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