9.4 Thermodynamic and Kinetic Control Student

Topic: 9.4 Thermodynamic and Kinetic Control

Learning Objective

  • Objective 9.4.A: Explain, in terms of kinetics, why a thermodynamically favored reaction might not occur at a measurable rate.

Essential Knowledge

  • 9.4.A.1: Many processes that are thermodynamically favored do not occur to any measurable extent or they occur at extremely slow rates.

  • 9.4.A.2: Processes that are thermodynamically favored but do not proceed at a measurable rate are under "kinetic control."

    • High activation energy can often explain why a process is under kinetic control.

    • Lack of noticeable rate does not imply that the system is at equilibrium.

    • If a process is thermodynamically favored but does not occur at a measurable rate, it is reasonable to conclude that it is under kinetic control.

Notes

  • A negative value for Gibb's free energy ( ΔG < 0) indicates a thermodynamically favorable reaction; however, not all such reactions occur.

  • Example 1: Carbon diamond to carbon graphite is thermodynamically favorable at room temperature, yet diamonds do not turn into graphite over a lifetime.

    • This reaction requires breaking strong covalent bonds in a large macromolecular solid structure, needing high activation energy, making it highly unlikely to occur.

  • Reactions that are thermodynamically favorable but do not occur can be described as being under kinetic control.

  • General factors for kinetically controlled processes:

    • High activation energy

    • Unfavorable orientations of reactants during collisions

    • Requirement for high numbers of particles to collide simultaneously to react

Page 2: Reaction Examples

Exothermic Reaction: 2 AB2(g) → A2(g) + 2 B2(g)

  1. ΔS (Entropy Change): + (Positive)

  2. ΔH (Enthalpy Change): - (Negative)

  3. ΔG (Gibb’s Free Energy): Negative under any conditions (Always)

  4. Condition Change to Proceed: Increase temperature for more energy or use a catalyst to lower activation energy.

Reverse Reaction: A2(g) + 2 B2(g) → 2 AB2(g)

  1. ΔS: - (Negative)

  2. ΔH: + (Positive)

  3. ΔG: Never negative (will never occur at room temperature)

  4. Condition Change: No adjustments make this reaction proceed at room temperature.

WE DO: Iron with Oxygen to Form Iron (III) Oxide

  • Reaction:

    • 4Fe(s) + 3 O2(g) → 2Fe2O3(s)

    • ΔH°rxn = -1648 kJ/mol

    • ΔS°rxn = -544 J/K

  • Tasks: A. Calculate ΔG°rxn at 298 K. B. Reason for slow reaction at room temperature, but faster with water added. C. Sketch energy diagram illustrating both scenarios. D. Effect of adding water on ΔG.

Page 3: Further Development - YOU DO

  1. Sketch Energy Diagrams:

  • A) Exothermic reaction; thermodynamically favorable; occurs quickly.

  • B) Exothermic reaction; thermodynamically favorable; does not occur at measurable rate.

  • C) Endothermic reaction; thermodynamically favorable; occurs quickly.

  1. Catalyst Addition Explanation:

  • Catalysts lower activation energy, enabling reactions that do not normally happen at room temperature, despite negative ΔG.

  1. Exothermic Reaction: A(g) + B(g) → AB(g) (K = 1x10^8 at 298 K) A. ΔS: Information deduced about the entropy change. B. ΔH: Information deduced about enthalpy change. C. ΔG: Information deduced about Gibb's Free Energy. D. Temperature Impact Explanation: Why the reaction does not occur at room temperature but does when temperature is increased.