Electrochemical Energy Systems and Corrosion Notes

Electrochemical Potential

  • Describes the potential energy of charge carriers in electrochemical systems.

Electrochemical Cells and Their Types

  • Electrochemical cells: Devices that convert chemical energy into electrical energy (or vice versa).

    • Types of electrochemical cells:

    1. Voltaic (Galvanic) cells: Generate electrical energy from spontaneous redox reactions.

    2. Electrolytic cells: Use electrical energy to drive non-spontaneous reactions.

Batteries

  • Primary batteries: Non-rechargeable, implemented in devices like flashlights.

  • Secondary batteries: Rechargeable, e.g., lithium-ion and lead-acid batteries.

  • Lithium-ion batteries: Widely used due to high energy density and rechargeability.

    • Types: Lithium polymer and lithium titanate.

    • Advantages: Light weight, ability to retain charge well, and higher voltage.

Fuel Cells

  • Convert chemical energy directly into electrical energy through electrochemical reactions.

  • Polymer membrane fuel cells: Use a proton exchange membrane to conduct protons from anode to cathode, and electrons travel through an external circuit.

    • Advantages: High efficiency and lower emissions compared to traditional combustion engines.

Supercapacitors

  • Devices that store electrical energy through electrostatic charge separation; good for quick discharge applications.

Theory of Corrosion

  • Corrosion is the deterioration of materials (primarily metals) due to electrochemical reactions with the environment.

Corrosion Protection Methods

  • Cathodic protection: Protects metals by making them the cathode in electrochemical cells.

  • Sacrificial anodes: More reactive metal used to attract corrosion away from the protected metal.

  • Impressed current: A constant current is applied to counteract the corrosive effects.

Electrochemistry Overview

  • Studies the relationship between electrical energy and chemical reactions, bridging the gap between chemistry, physics, and engineering.

Applications of Electrochemistry

  1. Chloralkali industry: Produces chlorine, caustic soda, and hydrogen

  2. Biomedical applications: Includes electrodialysis and biosensors.

  3. Energy systems: Includes batteries, fuel cells, and photovoltaics.

  4. Corrosion engineering: Understanding and preventing metal degradation.

  5. Environmental remedies: Treatments for pollutants using oxidative strategies.

Fundamental Concepts in Electrochemistry

  • Redox reactions: Transfer of electrons occurs during oxidation-reduction reactions.

    • Oxidation: Loss of electrons (e.g., M
      ightarrow M^{n+} + ne^-).

    • Reduction: Gain of electrons (e.g., M^{n+} + ne^-
      ightarrow M).

  • Interfacial potential difference: Known as electrode potential; developed at the electrode-electrolyte interface due to charge separation.

  • Single electrode potential (E): Tendency of an electrode to gain or lose electrons in a specific solution (half-cell potential).

  • Standard electrode potential (E°): Measured at 1M concentration and 1 atm pressure at 25°C.

Electrochemical Cells: Detailed Discussion

Construction of Galvanic Cell

  • Electrodes: Anode and cathode, typically metals immersed in electrolytes.

  • Electrolytes: Ionic solutions that allow ion flow within the cell.

  • Salt bridge: Prevents charge buildup by allowing ion flow between half-cells.

Daniell Cell Example

  • Components:

    • Zinc anode: Zn (s)
      ightarrow Zn^{2+} (aq) + 2e^-

    • Copper cathode: Cu^{2+} (aq) + 2e^-
      ightarrow Cu (s)

  • Cell potential calculation: Use standard reduction potentials to find overall reaction potential.

Standard Potentials

  • Determining standard potentials using the Standard Hydrogen Electrode (SHE).

  • Electrochemical series helps predict which half-reaction is more favorable based on reduction potentials.

Nernst Equation

  • Relates cell potential under non-standard conditions to concentrations using:
    E = E^ - rac{RT}{nF} ext{ln}(Q) where Q is the reaction quotient.

Faraday's Laws of Electrolysis

  1. The amount of chemical change at an electrode is directly proportional to the quantity of electricity passed.

  2. The different substances will produce amounts of chemical change proportional to their equivalent weights.

Practical Example: Electrolysis

  • Given a current, how to calculate mass deposited using Faraday's laws.

  • Example with magnesium chloride (MgCl2): 1 Faraday = 96500 C to deposit equivalent weight of Mg.

Concepts Around Concentration Cells

  • Concentration cells consist of two half-cells connected but with different ion concentrations.

    • Operation: Ecell is derived from concentration differences.

  • Importance in measurement, such as pH estimations via Nernst equation adjustments based on varying ion concentrations.

Electrolytic Cells

  • Use external electrical energy to drive non-spontaneous chemical reactions, such as decomposing compounds or electroplating.

  • Principle: Anode positive (+), cathode negative (-).

Summary of Differences: Voltaic vs. Electrolytic Cells

  • Spontaneity: Voltaic (spontaneous reactions, G < 0); Electrolytic (non-spontaneous, G > 0).

  • Energy Flow: Voltaic cells generate energy, while electrolytic cells consume energy.