1. Introduction to Electrochemistry
Electrochemistry deals with the conversion of chemical energy into
electrical energy and vice versa.
It has applications in batteries, fuel cells, electrolysis, and corrosion
prevention.
2. Electrochemical Cells
Galvanic (Voltaic) Cells: Convert chemical energy into electrical energy
through spontaneous redox reactions.
o Example: Daniell Cell (Zn/Cu redox reaction).
o Electrode Potential: The voltage difference between electrodes,
measured using the Nernst equation.
Electrolytic Cells: Use electrical energy to drive non-spontaneous
reactions (e.g., electrolysis of water, metal refining).
3. Nernst Equation and Equilibrium
The Nernst equation relates cell potential to ion concentration.
At equilibrium, the equation links cell potential (E°), Gibbs free energy
(ΔG°), and equilibrium constant (K).
4. Conductance of Electrolytic Solutions
Conductivity (κ) measures how well a solution conducts electricity.
Molar conductivity (Λm) depends on ion concentration and mobility.
Kohlrausch’s Law states that at infinite dilution, ions contribute
independently to conductivity.
5. Electrolysis and Faraday’s Laws
Faraday’s First Law: The mass of a substance deposited at an electrode is
proportional to the charge passed.
Faraday’s Second Law: The amount of different substances liberated by
the same charge is proportional to their equivalent weight.
Used in metal refining (e.g., copper purification) and large-scale metal
production (e.g., Na, Mg, Al).
6. Batteries and Fuel Cells
Primary Batteries (non-rechargeable, e.g., dry cell, mercury cell).
Secondary Batteries (rechargeable, e.g., lead-acid battery, lithium-ion
battery).
Electrochemistry deals with the conversion of chemical energy into
electrical energy and vice versa.
It has applications in batteries, fuel cells, electrolysis, and corrosion
prevention.
2. Electrochemical Cells
Galvanic (Voltaic) Cells: Convert chemical energy into electrical energy
through spontaneous redox reactions.
o Example: Daniell Cell (Zn/Cu redox reaction).
o Electrode Potential: The voltage difference between electrodes,
measured using the Nernst equation.
Electrolytic Cells: Use electrical energy to drive non-spontaneous
reactions (e.g., electrolysis of water, metal refining).
3. Nernst Equation and Equilibrium
The Nernst equation relates cell potential to ion concentration.
At equilibrium, the equation links cell potential (E°), Gibbs free energy
(ΔG°), and equilibrium constant (K).
4. Conductance of Electrolytic Solutions
Conductivity (κ) measures how well a solution conducts electricity.
Molar conductivity (Λm) depends on ion concentration and mobility.
Kohlrausch’s Law states that at infinite dilution, ions contribute
independently to conductivity.
5. Electrolysis and Faraday’s Laws
Faraday’s First Law: The mass of a substance deposited at an electrode is
proportional to the charge passed.
Faraday’s Second Law: The amount of different substances liberated by
the same charge is proportional to their equivalent weight.
Used in metal refining (e.g., copper purification) and large-scale metal
production (e.g., Na, Mg, Al).
6. Batteries and Fuel Cells
Primary Batteries (non-rechargeable, e.g., dry cell, mercury cell).
Secondary Batteries (rechargeable, e.g., lead-acid battery, lithium-ion
battery).
