Learn & Review: Electrochemistry | One Shot Marathon | Class 12 | Gethu Batch | CBSE

Electrochemistry: A Comprehensive Summary

This session provides a detailed overview of electrochemistry, focusing on concepts relevant for board exams. It covers electrochemical cells, electrolytic cells, redox reactions, cell representation, electrode potentials, the Nernst equation, and Gibbs energy.

1. Introduction to Electrochemistry

• Definition: Electrochemistry deals with the relationship between chemical reactions and electrical energy.
• Focus: The session aims to cover 70-80% of the electrochemistry chapter in a "one-shot" format, prioritizing exam-relevant material.

2. Electrochemical vs. Electrolytic Cells

• Electrochemical Cell (Galvanic/Voltaic Cell):
  • Converts chemical energy to electrical energy.
  • Driven by spontaneous chemical reactions (e.g., Zn + CuSO₄ → ZnSO₄ + Cu).
  • Characterized by a negative change in Gibbs free energy (ΔG < 0).
• Electrolytic Cell:
  • Converts electrical energy to chemical energy.
  • Drives non-spontaneous chemical reactions by supplying external energy.

3. Redox Reactions and Electron Transfer

• Core Concept: All electrochemical processes involve the movement of electrons.
• Definitions:
  • Oxidation: Loss of electrons (OIL - Oxidation Is Loss). Occurs at the Anode.
  • Reduction: Gain of electrons (RIG - Reduction Is Gain). Occurs at the Cathode.
• Mnemonics:
  • OIL RIG
  • Red Cat An Ox (Reduction at Cathode, Oxidation at Anode)
• Agents:
  • Reducing Agent: A substance that gets oxidized (loses electrons) and causes reduction in another substance.
  • Oxidizing Agent: A substance that gets reduced (gains electrons) and causes oxidation in another substance.

4. Types of Electrochemical Cells

• Galvanic/Voltaic Cell: Spontaneous reaction produces electrical energy.
• Electrolytic Cell: Non-spontaneous reaction requires electrical energy input.

5. The Daniell Cell: A Galvanic Cell Example

• Setup:
  • Anode (Left): Zinc electrode immersed in Zinc Sulfate (ZnSO₄) solution. Zinc gets oxidized (Zn → Zn²⁺ + 2e⁻).
  • Cathode (Right): Copper electrode immersed in Copper Sulfate (CuSO₄) solution. Copper ions get reduced (Cu²⁺ + 2e⁻ → Cu).
  • Salt Bridge: Connects the two half-cells.
• Function of Salt Bridge:
  • Completes the electrical circuit.
  • Maintains electrical neutrality in the half-cells by allowing ion migration (e.g., anions move to the anode compartment, cations move to the cathode compartment).
  • Typically made of KCl or KNO₃ jellified with agar-agar.
• Cell Representation:
  • Anode is written first, followed by the salt bridge (double vertical line), then the cathode.
  • Phase changes are indicated by single vertical lines.
  • Example: Zn(s) | ZnSO₄(aq) || CuSO₄(aq) | Cu(s)

6. Electrode Potential and Cell EMF

• Electrode Potential: The potential difference developed at the interface between an electrode and its electrolyte due to charge separation.
• Cell EMF (Electromotive Force): The total potential difference between the two electrodes of a galvanic cell.
  • Calculated as: E_cell = E_cathode - E_anode (using reduction potentials).
  • Alternatively: E_cell = Oxidation Potential (Anode) + Reduction Potential (Cathode).
• Standard Electrode Potential (E°): Electrode potential measured under standard conditions (298 K, 1 atm pressure, 1 M concentration).
• Standard Hydrogen Electrode (SHE):
  • A reference electrode with a defined standard potential of 0 volts.
  • Consists of a platinum electrode coated with platinum black, immersed in a 1 M H⁺ solution, with H₂ gas bubbled at 1 atm pressure.
  • Used to measure the standard electrode potentials of other half-cells.
• Reactivity Series and Potentials:
  • Metals with negative reduction potentials (e.g., Li, K, Na, Zn) are more reactive than hydrogen and readily get oxidized.
  • Metals with positive reduction potentials (e.g., Cu, Ag, F₂) are less reactive than hydrogen and tend to be reduced.
  • Strong oxidizing agents have high reduction potentials (e.g., F₂).
  • Strong reducing agents have low (negative) reduction potentials (e.g., Li).
• EMF of Daniell Cell: E°_cell = E°(Cu²⁺/Cu) - E°(Zn²⁺/Zn) = 0.34 V - (-0.76 V) = 1.1 V.

7. The Nernst Equation

• Purpose: Relates the cell potential (EMF) to non-standard conditions (concentration, temperature).
• Equation: E_cell = E°_cell - (RT / nF) * ln(Q) Where:
  • E_cell: Cell potential under non-standard conditions.
  • E°_cell: Standard cell potential.
  • R: Gas constant (8.314 J/K·mol).
  • T: Temperature in Kelvin.
  • n: Number of moles of electrons transferred in the balanced reaction.
  • F: Faraday's constant (96487 C/mol).
  • ln(Q): Natural logarithm of the reaction quotient.
• Simplified Form at 298 K: E_cell = E°_cell - (0.0591 / n) * log(Q)
• Reaction Quotient (Q): Ratio of product concentrations to reactant concentrations, raised to their stoichiometric coefficients. Solids are excluded (concentration = 1).
  • For a reaction aA + bB → cC + dD, Q = ([C]^c * [D]^d) / ([A]^a * [B]^b).

8. Equilibrium and the Nernst Equation

• At Equilibrium: E_cell = 0 and Q = K_c (Equilibrium constant).
• Relationship: 0 = E°_cell - (RT / nF) * ln(K_c)
  • Rearranging gives: E°_cell = (RT / nF) * ln(K_c)
  • At 298 K: E°_cell = (0.0591 / n) * log(K_c)

9. Gibbs Free Energy (ΔG)

• Relationship with Cell Potential: ΔG = -nFE_cell
• Standard Gibbs Free Energy: ΔG° = -nFE°_cell
• Spontaneity:
  • ΔG < 0 (or E_cell > 0): Spontaneous reaction.
  • ΔG > 0 (or E_cell < 0): Non-spontaneous reaction.
  • ΔG = 0 (or E_cell = 0): Reaction is at equilibrium.
• Relationship with Equilibrium Constant: ΔG° = -RT ln(K_c)

10. Key Takeaways and Applications

• Electrochemical cells harness spontaneous redox reactions to generate electricity.
• Electrolytic cells use electricity to drive non-spontaneous reactions.
• Redox reactions are fundamental, involving electron transfer (oxidation at anode, reduction at cathode).
• The Daniell cell is a classic example illustrating these principles.
• The Nernst equation allows calculation of cell potentials under varying conditions.
• Standard Hydrogen Electrode (SHE) serves as a crucial reference point for measuring electrode potentials.
• Gibbs free energy provides thermodynamic insight into the spontaneity of electrochemical reactions.
• Understanding the reactivity series and standard reduction potentials is key to predicting reaction feasibility.