G11-Chemistry-STB-2023-web

Chemical Equilibrium and Catalysts

• Equilibrium and Catalysts:

  • Addition of a catalyst does not shift the equilibrium position of a chemical reaction.

  • A catalyst increases the rates of both forward and reverse reactions by the same factor, leaving the equilibrium composition unchanged.

  • It hastens the time taken to reach equilibrium without altering the equilibrium constant or concentrations.

Le Chatelier’s Principle and Industrial Applications

• Le Chatelier’s Principle:

  • Industrial processes often utilize Le Chatelier’s principle to maximize product yield under varying conditions.

  • Example: Contact Process for Sulfuric Acid Production:

    • Reaction: 2SO2 (g) + O2 (g) ⇌ 2SO3 (g)

      • ΔH = -196 kJ

    • The equilibrium position is adjusted by controlling temperature and pressure.

    • Higher pressures favor the formation of SO3 due to reduced volume.

    • Optimum pressure is typically around 200 - 400 atm.

Haber Process Overview

• Haber Process for Ammonia Production:

  • Established by Fritz Haber, optimized by Carl Bosch.

  • Reaction: N2 (g) + 3H2 (g) ⇌ 2NH3 (g)

    • Enthalpy change: ΔH = -92 kJ/mol.

  • The production of ammonia is favored by:

    • High pressure

    • Low temperature (though low temp slows the reaction rate).

  • Catalyst: Iron is commonly used to expedite the reaction by lowering activation energy.

  • Yield nearly 100% at 200 °C and above 750 atm.

Temperature Effects on Equilibrium Constant

• Temperature and Equilibrium Constant:

  • As temperature increases, the equilibrium constant (K) generally decreases for exothermic reactions such as the Haber process:

    • Example data:

      • At 25 °C, K ≈ 6.4x10²

      • At 500 °C, K ≈ 1.5x10⁻⁵.

Ammonia Production Optimization

• Continuous Removal of Ammonia:

  • In industrial processes, ammonia is continuously removed to drive the equilibrium further towards the production of ammonia.

  • Uniquely, the production process is designed such that the reaction does not reach equilibrium, favoring efficiency.

• Heat exchange systems are implemented to recover and use heat produced in the reaction to preheat reactants.

Summary of Equilibrium Constants

• Key Equilibrium Concepts:

  • Dynamic Equilibrium: A state where concentrations of reactants and products remain constant.

  • The Law of Mass Action: The rate of a reaction at equilibrium is proportional to the product of the concentrations of the reactants raised to the power of their coefficients.

  • Equilibrium Expression:

    • For a general reaction: aA + bB ⇌ mM + nN

    • The equilibrium constant (K_C) is:

      • K_C = [M]^m[N]^n / [A]^a[B]^b

  • If K_C > 1, products are favored; if K_C < 1, reactants are favored.

Practical Applications and Exercises

• Exercises and Topics Related to Equilibrium:

  • Understand how temperature, pressure, and concentration shifts affect equilibrium as per Le Chatelier’s principle.

  • Evaluate products formed under various conditions and predict changes in equilibrium based on experimental data.

• The significance of catalysts in industry and their effects on reaching equilibrium should also be explored through practical applications and real-world examples.