CHM 101 ELECTROCHEMISTRY Module 3 Powerpoint

CHM 101 - Electrochemistry

Module Overview

• Course Title: CHM 101 Electrochemistry

• Instructor: Prof. O.O. Soriyan

• Focus: Concentration effects, Nernst equation, Redox reactions, Oxidation potentials.

Key Objectives

• Understand the relationship between cell potential and Gibbs free energy (∆G).

• Examine the relationship between standard cell potential and equilibrium constant (K).

• Apply the Nernst equation to calculate the emf of galvanic cells and define the standard potential of a cell.

Thermochemistry and Electrochemistry

Spontaneity of Reactions

• For a reaction to be spontaneous, free energy (∆G) must be negative:

  • ∆G = - ve

• The cell potential (E° cell) must be positive:

  • E° cell = + ve

Free Energy Change Calculation

• The energy change for an electrochemical reaction is given by:

  • ∆G° = -nFE° cell

  • Where:

    • n = number of moles of electrons transferred in the reaction (from balanced equation).

    • F = Faraday constant (96500 C/mole).

Conditions for Cell Reaction feasibility

• a) E° cell positive: Reaction feasible.

• b) E° cell negative: Reaction not feasible.

• c) E° cell zero: Reaction in equilibrium.

Example Calculation: Free Energy of Daniel Cell

Reaction

• Zn + Cu²⁺ → Zn²⁺ + Cu

  • E° cell = +1.10 V

Half Reactions

• Anode: Zn → Zn²⁺ + 2e⁻ (E° = +0.76 V)

• Cathode: Cu²⁺ + 2e⁻ → Cu (E° = +0.34 V)

Overall Reverse Reaction

• Overall: Zn + Cu²⁺ → Zn²⁺ + Cu

  • E° cell = 0.76 V + 0.34 V = 1.10 V

Free Energy Calculation

• n = 2 (moles of electrons)

• ∆G = -nFE° cell

• ∆G = -2 (96500 C/mol)(1.10 V) = -2.10 × 10⁵ C·V/mol = -210 kJ/mol

Assignment

• Calculate free energy for the reaction: N₂(g) + 3H₂(g) → 2NH₃(g) with E° cell = 0.057 V.

Electrochemical Series

• Arrangement of various electrode systems according to their standard reduction potentials.

Key Data: Standard Electrode Potentials at 298 K

• Table listing various half-reactions and their standard reduction potentials.

  • Example:

    • F₂(g) + 2e⁻ → 2F (E° = 2.87 V)

    • Zn²⁺ + 2e⁻ → Zn (E° = -0.76 V)

Trends in Reactivity

• Reactivity of metals decreases down the series.

• Metals above hydrogen can displace hydrogen from dilute acids.

• Electropositive character of metals decreases down the series.

• In general:

  • A redox reaction is feasible when the substance with a higher reduction potential gets reduced.

Questions

• Q7: Is acidified permanganate a stronger oxidizing agent than acidified dichromate? Provide chemical equation and standard potential.

• Q8: Can KMnO₄ oxidize Iron(II) to Iron(III) in acidic aqueous solution?

Free Energy Change at Equilibrium

• ∆G° = -RT ln K, where K = equilibrium constant, R = gas constant.

• Relating equations:

  • -RT ln K = -nFE° cell

  • E° cell = RT/nF ln K

Equilibrium Constants Calculations

Example: Equilibrium Constant for Daniel Cell

• Reaction: Zn + Cu²⁺ → Zn²⁺ + Cu

• E° cell = +1.10 V

• ln K = nE° cell / (0.025693 V)

• K = 1.5 × 10³⁷ (moving to products)

Summary of Reaction Parameters at Standard State

Effects of Concentration on Cell Potential

• Electrochemical cells function at concentrations less than 1M; potential decreases as concentration varies.

• Normal starting conditions (e.g., Daniel Cell E° cell = 1.10 V).

• As the battery operates, ion concentrations change, reducing cell potential until equilibrium is reached (∆G = 0, E° cell = 0 V).

Nernst Equation

• Formula: E° cell = E° cell o − (RT/nF) ln Q, where Q is the reaction quotient.

• E° cell = E° cell o − (0.059 V/n) log Q at 298 K.

Example: Calculate Potential of Daniel Cell

Given

• Zn²⁺ = 0.10 mol/L, Cu²⁺ = 0.0010 mol/L

• E° cell = +1.10 V

Calculation

• E° cell = E° cell o − 0.059 V/n log Q

• Q = [Zn²⁺]/[Cu²⁺]= 0.10/0.0010 = 100

• E° cell = 1.10 V − 0.059 V (2) log(100)

• E° cell = 1.10 V − 0.059 V (4) = 1.10 V − 0.236 V = +1.04 V

Equilibrium and Cell Potential

• At equilibrium, E° cell = 0, which implies Q = K.

• 0 = E° cell o − 0.059 n log K.

• At 298K, log K = nE° cell/0.059.