Lectures 3-4 Gibbs energy and bond energies

Page 1: Reaction Information

  • Reaction: H2O(l) ↔ H2O(g)

  • Standard Enthalpy of Formation (ΔHӨ):

    • H2O(l): -286 kJ/mol

    • H2O(g): -242 kJ/mol

  • Standard Entropy of Formation (Sө):

    • H2O(l): 70 J/K·mol

    • H2O(g): 189 J/K·mol

  • Question: Is this reaction spontaneous at room temperature?

Page 2: Objectives

  • Objective 1: Define Gibbs free energy.

  • Objective 2: Relate Gibbs free energy (ΔG°) to reaction spontaneity, incorporating enthalpy (ΔH°) and entropy (ΔS°).

  • Objective 3: Show the relationship between ΔG° and equilibrium constant (Keqm).

Page 3: Definition of Gibbs Free Energy

  • Gibbs Free Energy: The amount of "free" or "available" energy a system possesses to do work.

  • It predicts whether a process will occur spontaneously at constant pressure and temperature.

Page 4: Gibbs Equation

  • Thermodynamic equation used to calculate the change in Gibbs energy of a system as a function of temperature.

  • Named after: Josiah Willard Gibbs and Hermann von Helmholtz.

Page 5: Gibbs Equation and Spontaneity

  • Gibbs Equation: ΔG = ΔH – TΔS

    • ΔH: Change in enthalpy

    • T: Temperature in Kelvin

    • ΔS: Change in entropy

  • Relationship: Entropy (ΔS) and enthalpy (ΔH) help determine reaction spontaneity.

Page 6: Interpreting the Gibbs Equation

  • If heat is released (ΔH is negative), some may convert into useful work.

  • Some heat increase order of the system (ΔS is negative).

  • More disorder (ΔS>0) leads to more available energy for work.

Page 7: Interpreting Gibbs Energy

  • If ΔG < 0: Reaction is spontaneous forward/right direction.

  • If ΔG > 0: Non-spontaneous forward direction, spontaneous reverse/left direction.

  • If ΔG = 0: Reaction is at equilibrium (reactants and products equally favored).

Page 8: Interpreting Gibbs for Spontaneity

  • Preferred conditions:

    • Spontaneous: Negative ΔH and positive ΔS.

    • Non-spontaneous: Positive ΔH and negative ΔS.

  • Same signs: Temperature determines spontaneity.

Page 9: Diagram Summary of Gibbs Equation

  • Summary of ΔH, ΔS, ΔG:

    • +ΔH, +ΔS: Reaction spontaneous at high temperature, reverse at low.

    • -ΔH, +ΔS: Spontaneous at all temperatures.

    • +ΔH, -ΔS: Spontaneous reverse at all temperatures.

    • -ΔH, -ΔS: Spontaneous at low temperature, reverse at high temperature.

Page 10: Standard Free Energy

  • Standard Gibbs Free Energy for reaction (ΔGⓇreaction):

    • ΔGⓇreaction = [ΣΔG°f (products) - ΣΔG°f (reactants)]

  • Elements in standard state: ΔG°f = 0.

Page 11: Equilibrium Constant (Keqm)

  • Definition: Ratio of product concentrations to reactant concentrations in a reversible process.

  • Example Reaction: A + B ↔ C + D

    • K = ([C][D])/([A][B])

  • Indicates if products or reactants are favored.

Page 12: ΔG and Keqm Relationship

  • Relationship: ΔGӨ = -RTlnK

  • Interpretation:

    • K > 1: Products favored (ΔG < 0).

    • K = 1: Products and reactants equally favored (ΔG = 0).

    • K < 1: Reactants favored (ΔG > 0).

Page 13: Specific Reaction Evaluation

  • Reaction: H2O(l) ↔ H2O(g)

  • Given:

    • ΔG = -RTlnKc

    • ΔGӨreaction = +8.5 kJ/mol

    • R = 8.314 J/K·mol

  • At 25°C, evaluation:

    • lnK = - [ΔG/(RT)] = - [8500/(8.314*298)] = -3.43

    • K = 0.03 (non-spontaneous).

Page 14: Second Set of Objectives

  • Objective 1: Define bond dissociation energy.

  • Objective 2: Illustrate bond dissociation energy for determining standard enthalpies of reaction and formation.

Page 15: The Chemical Reaction

  • Represented:

    • 1 Methane (CH4) molecule + 2 Oxygen (O2) molecules → 1 Carbon Dioxide (CO2) + 2 Water (H2O).

Page 16: Chemical Reaction Description

  • Chemical reactions involve breaking and forming chemical bonds.

  • Energy needed to break bonds (supplied as heat).

  • Energy released when bonds are formed.

Page 17: Bond Dissociation Energy Definition

  • Bond Dissociation Energy (D): Energy needed to break one mole of bonds in a gaseous covalent substance.

  • Example: H2(g) → H(g) + H(g) with ΔHӨ = 435 kJ/mol.

Page 18: Bond Dissociation Energy in Polyatomic Molecules

  • Bond energy varies; examples given:

    • CH4(g) → CH3(g) + H(g) at ΔHӨ = 427 kJ/mol

    • Average bond enthalpy for C-H = 416 kJ/mol.

Page 19: Multiple Bonds Energy

  • Bond energies for double & triple bonds are not multiples of single bond energies.

  • Strength order: Triple > Double > Single Bonds.

    • Example energies are provided for C-O and C=O bonds.

Page 20: Uses of Bond Energies

  • Applications:

    • Estimate Enthalpy of Reaction

    • Estimate Enthalpy of Formation (for gaseous phase only).

    • Formula: ΔHӨreaction = ΣDreactants - ΣDproducts.

Page 21: ΔHӨreaction Calculation Steps

  1. Write balanced equation for reactants forming products (gaseous).

  2. Draw structural formula of all species.

  3. Identify bonds present.

  4. Input bond energy values.

Page 22: Example Problem Setup

  • Given:

    • DC-I = 343 kJ/mol

    • DH-H = 435 kJ/mol

    • DH-Cl = 432 kJ/mol

  • Task: Calculate ΔHⓇformation of HCl(g).

Page 23: Example Calculation Steps

  1. Balanced Equation: ½ H2(g) + ½Cl2(g) → HCl(g)

  2. Bonds: ½ H-H + ½ Cl-Cl → H-Cl.

  3. Energy Calculation: ½ DH-H + ½ DCl-Cl – DH-Cl = [½(435) + ½(343)] – [432] = -43 kJ.

Page 24: Another Calculation Setup

  • Given:

    • DC-H = 416 kJ/mol

    • DH-H = 435 kJ/mol

    • Enthalpy of atomization of graphite = 717 kJ/mol

  • Task: Calculate ΔHӨformation of CH4(g).