Lecture Notes on Chemical Equilibrium

Disrupting Chemical Equilibrium - CHEM1003 Notes

Overview of Lecture Objectives

• Understand the concept of Le Chatelier’s Principle.
• Describe the factors that affect chemical equilibrium.

Le Chatelier’s Principle

• Definition: Proposed by Henri Le Chatelier in 1884.
• Key Concept: If a system at equilibrium is disturbed, it will shift to counteract the stress.

Reaction Quotient (Q)

• Definition: Ratio of product concentrations raised to their stoichiometric coefficients to reactant concentrations raised to theirs.
• Expression:
  Q = rac{[C]^c imes [D]^d}{[A]^a imes [B]^b}
• Difference between Q and K:
  • Q can be measured at any point of the reaction.
  • K is the equilibrium constant at equilibrium:
  • When Q
    eq K, the system is not at equilibrium.
  • When Q = K, the system is at equilibrium.

Factors Affecting Equilibrium

1. Addition of a Reaction Component

   • Example:
     • aA + bB
       ightleftharpoons cC + dD
   • Effect: Adding more of a reactant shifts equilibrium to the right (towards products).
   • Case Study:
     • N2O4(g)
       ightleftharpoons 2NO_2(g)
     • Adding N2O4 results in more NO_2 being formed, indicated by a darker mixture.

2. Removal of a Reaction Component

   • Effect: Removing a product shifts equilibrium towards the same side (towards reactants).
   • Example: For
     • CH3COOH(aq) + C2H5OH(aq) ightleftharpoons H2O(l) + CH3COOC2H_5(aq)
     • Removing ethyl acetate shifts equilibrium to produce more products.

3. Change in Pressure

   • Only applicable if one or more components are gases.
   • Effect: Increasing pressure shifts equilibrium towards the side with fewer moles of gas.
   • Example:
     • N2O4(g)
       ightleftharpoons 2NO_2(g)

4. Change in Temperature

   • Depends on whether the reaction is exothermic or endothermic.
   • For Exothermic Reactions:
     • Heat is a product. Adding heat shifts equilibrium to the left (towards reactants).
     • Example:
     • 2H2(g) + O2(g)
       ightleftharpoons 2H_2O(l) + ext{HEAT}
   • For Endothermic Reactions:
     • Heat is a reactant. Adding heat shifts equilibrium to the right (towards products).

Change in Temperature and Equilibrium Constant (K)

• The value of K changes with temperature.
• Exothermic Reaction:
  • Adding energy (heat) shifts equilibrium to the left (reactants).
• Endothermic Reaction:
  • Adding energy (heat) shifts equilibrium to the right (products).

Exercise on Equilibrium Changes

• Example Reaction: N2O4(g)
  ightleftharpoons 2NO_2(g), ext{ endothermic, } ext{ΔrHΘ}=+56.9 ext{kJ mol}^{-1}
• Effects of various changes:
  • a) Adding N2O4 increases NO_2.
  • b) Lowering pressure increases NO_2.
  • c) Raising temperature increases NO_2 because it is endothermic.
  • d) A catalyst does not affect the position of equilibrium.

Effect of Catalysts on Equilibrium

• Catalysts:
  • Speed up the rate of reaction without being consumed.
  • Do not affect equilibrium position or the value of K.

The Haber Process

• Reaction: N2(g) + 3H2(g)
  ightleftharpoons 2NH_3(g) + ext{Heat}
• Catalyst Use: Allows reaction to occur faster at 500°C despite being exothermic.

Solubility Equilibria

• Equilibrium Expression:
  • K_{sp} = [M^+] [X^-]
• Example Reaction:
  • Bi2S3(s)
    ightleftharpoons 2Bi^{3+}(aq) + 3S^{2-}(aq)

Quantitative Aspects of Equilibrium Constants

• Example Calculation:
  • For the reaction:
    • 1.00 ext{ mol } SO2(g) + 1.00 ext{ mol } O2(g)
      ightarrow SO3(g), ext{ with } 0.925 ext{ mol } SO3(g)
  • Find K:
  • Q = rac{[SO3]^2}{[SO2]^2 [O_2]} = 2.8 imes 10^2

Summary of Equilibrium Concepts

• K value consistency: Remains constant at a given temperature regardless of reactant/product amounts.
• Equilibrium Positions: May vary, but K remains the same.
• Factors affecting position: Include changes in concentration, pressure, and temperature.
• Solubility Product: Applies to saturated solutions with excess solid.