Electrochemistry One-Shot Lecture by Sourav Braina

Introduction

• Electrochemistry deals with the production of electrical energy from chemical reactions and the use of electrical energy to bring about non-spontaneous chemical transformations.
• Key Focus: Spontaneous redox reactions for electricity production and non-spontaneous chemical reactions by electrical energy.

Spontaneous Redox Reactions

• Definition: Involves both oxidation and reduction reactions occurring spontaneously to produce electricity.
• Device: Electrochemical cell (Galvanic or Voltaic cell) where ΔG is negative.
• Example: Daniell Cell.

Daniell Cell

• Components: Zinc strip in ZnSO4 solution and copper strip in CuSO4 solution, connected by a salt bridge (inverted U-tube).
• Process: Zinc undergoes oxidation (Zn to Zn²⁺) at the anode, and copper undergoes reduction (Cu²⁺ to Cu) at the cathode.
• Voltage: 1.1V.
• Electrode Reactions:
  • Anode (Oxidation): Zn → Zn²⁺ + 2e⁻
  • Cathode (Reduction): Cu²⁺ + 2e⁻ → Cu

Non-Spontaneous Redox Reactions

• Definition: Require an external source of electrical energy to occur.
• Device: Electrolytic cell where ΔG is positive.
• Process: Known as electrolysis.

Electrolytic Cell Example

• Structure: Two electrodes connected to a power source, immersed in an electrolyte solution.
• Electrode Reactions:
  • Anode: Oxidation reaction.
  • Cathode: Reduction reaction.

Electrode Potential

• Definition: Tendency of an electrode to gain or lose electrons when it is in contact with its ion solution.
• Standard Electrode Potential (E°): Measured under standard conditions, 1M ion concentration, 298K temperature, and 1 bar pressure for gases.
• Measurement: Using Standard Hydrogen Electrode (SHE) as a reference with an E° of 0V.

Calculation of Cell Potential

• Formula: E°cell = E°cathode - E°anode
• Example Calculation: For a cell with Cu and Ag,
  • E°cell = E°(Ag⁺/Ag) - E°(Cu²⁺/Cu)

Electrochemical Series

• Definition: An arrangement of electrodes in increasing or decreasing order of their standard electrode potentials.
• Applications: Determines oxidizing and reducing strengths of different substances.

Key Points

• Strong Oxidants: Higher E° values, e.g., Fluorine.
• Strong Reductants: Lower E° values, e.g., Lithium.

Nernst Equation

• Used to calculate the electrode potential under non-standard conditions.
• Formula: E = E° - (RT/nF) lnQ
  • At 298K, the simplified form is E = E° - (0.059/n) logQ.

Applications and Calculations

• EMF and ΔG° Relationship: ΔG° = -nFE°, where F is Faraday's constant (96,487 C/mol).
• Example Calculation: For ΔG° of a reaction involving Zn and Cu.

Conductivity and Molar Conductivity

• Definitions:
  • Conductivity (κ): Measure of a solution’s ability to conduct electricity.
  • Molar Conductivity (Λm): Conductivity per unit concentration.
• Formula: Λm = κ / C
  • Units: S·cm²·mol⁻¹.

Faraday's Laws of Electrolysis

• First Law: Amount of substance deposited is proportional to the amount of electric charge passed.
• Second Law: Amounts of different substances deposited by the same quantity of electricity are proportional to their equivalent weights.

Battery Types

• Primary Cells: Cannot be recharged (e.g., Dry cell, Mercury cell).
• Secondary Cells: Reversible, can be recharged (e.g., Lead-acid battery, Ni-Cd battery).

Fuel Cells

• Definition: Convert chemical energy directly into electrical energy.
• Example: Hydrogen-oxygen fuel cell.
• Efficiency: Approximately 70%.

Corrosion

• Definition: Slow conversion of metals into undesirable compounds through oxidation.
• Example: Rusting of iron.
• Prevention Methods: Barrier protection, sacrificial protection, etc.

Acid-Base Titrations

• Concept: Uses the principle of electrochemistry to locate the equivalence point in titrations.

Summary: Electrochemistry covers a wide range of topics from spontaneous and non-spontaneous redox reactions to practical applications in batteries, fuel cells, and corrosion prevention.