Equilibrium

Le Chateliers Principle Examples

Le Chateliers Principle states that if a change of condition is applied to a system in equilibrium, the system will shift in a direction that relieves the stress. This can involve changes in concentration, pressure, volume, or temperature.

Types of Stress and System Response:

• Change in Concentration:

  • Increasing the concentration of a reactant shifts the equilibrium to the right (towards products).
  • Increasing the concentration of a product shifts the equilibrium to the left (towards reactants).
  • $ K_{eq} $ remains unchanged (unless temperature changes). Example: $ N_2(g) + 3H_2(g) \rightleftharpoons 2NH_3(g) $ Adding more $ N_2 $ will shift the equilibrium to the right, producing more $ NH_3 $ .

• Change in Pressure/Volume: (Applies only to gaseous reactions)

  • Increasing pressure (decreasing volume) favors the side with fewer moles of gas.
  • Decreasing pressure (increasing volume) favors the side with more moles of gas.
  • $ K_{eq} $ remains unchanged (unless temperature changes). Example: $ N_2(g) + 3H_2(g) \rightleftharpoons 2NH_3(g) $ . Increasing pressure shifts equilibrium to the right (fewer moles of gas on the product side).

• Change in Temperature:

  • Exothermic reactions ( $ ΔH < 0 $ ): Increasing temperature shifts the equilibrium to the left (towards reactants), decreasing temperature shifts to the right.
  • Endothermic reactions ( $ ΔH > 0 $ ): Increasing temperature shifts the equilibrium to the right (towards products), decreasing temperature shifts to the left.
  • $ K_{eq} $ changes with temperature. Equilibrium Constant

Examples:

• Haber Process ( $ N_2 + 3H_2 \rightleftharpoons 2NH_3 + heat $ ): This is an exothermic reaction. To maximize $ NH_3 $ production, high pressure, low temperature, and continuously removing $ NH_3 $ are employed. However, low temperatures slow the reaction rate, hence a compromise is needed.

• Dissolution of a slightly soluble salt: $ AgCl(s) \rightleftharpoons Ag^+(aq) + Cl^-(aq) $ Adding $ Cl^- $ ions will shift the equilibrium to the left, decreasing the solubility of $ AgCl $ . This is the common ion effect. Common Ion Effect

Further Notes:

• Equilibrium Constant
• Reaction Quotient
• Gibbs Free Energy and Equilibrium

Practice Problems:

1. Predict the effect of increasing the pressure on the following equilibrium: $ 2SO_2(g) + O_2(g) \rightleftharpoons 2SO_3(g) $

2. Predict the effect of adding more $ H_2 $ to the following equilibrium: $ H_2(g) + I_2(g) \rightleftharpoons 2HI(g) $

3. Predict the effect of increasing the temperature on the following equilibrium: $ N_2(g) + O_2(g) \rightleftharpoons 2NO(g) \quad ΔH > 0 $