Equilibrium

• Reversible reaction

  • Some reactions are able to be reversed, but many are not
• Non-reversible reaction

  • Mg(s) + 2HCl(aq) → MgCl2(aq) + H2(g)
• Reversible reactions

  • NH3 (g) + HCl (g) ⇋ NH4Cl (aq)
• Dynamic Equilibrium

  1. If a reversible reaction is carried out in a closed vessel/container and at constant temperature, then an equilibrium is established

  2. It is ongoing and both reactions are happening

  3. An example of a physical change in equilibrium
• Chemical Equilibrium

  1. Reached when the rate of the forward reaction equals the rate of the reverse reaction

  2. Concentration of the reactants and products does not change with time
• Le Chatelier’s principle states that if a dynamic equilibrium is disturbed by changing the conditions, the position of equilibrium moves to counteract (lessen) the change
• Factors affecting the position of equilibrium are concentration, pressure, and temperature
• Position of equilibrium is a way of expressing the direction in which the reaction is progressing
• If the conditions of the experiment change, the composition of the equilibrium mixture will also change
• Effect of changing concentration

  • When the concentration of reactants is increased, the equilibrium must shift towards the products to use up the added reactants
  • When the concentration of products is increased, the equilibrium must shift to the left to break down the extra products
• When the concentration of reactants is increased

  The equilibrium must shift towards the products to use up the added reactants
• When the concentration of products is increased

  The equilibrium must shift towards the reactants to use up the added product
• Increasing pressure (or decreasing volume)

  Increases concentration
• If the pressure increases

  The reaction will shift to the side with the fewest moles of gas to relieve the stress on the system
• If the pressure decreases

  The reaction will shift to the side with the most moles of gas to relieve the stress on the system
• Example

  • CO(g) + 2H2(g) ⇋ CH3OH(g)
  • Low volume (high pressure) → products
  • High volume (low pressure) → reactants
• Endothermic processes

  Are favoured when temperature increases (to absorb the added heat)
• Exothermic processes

  Are favoured when temperature decreases (to release more heat)
• Example

  • In the first step for synthesis of nitric acid, ammonia is oxidized to nitric oxide by the reaction: 4NH3(g) + 5O2(g) ⇋ 4NO(g) + H2O(g) ΔH = –905.6 kJmol-1
  • If T ↑, endothermic reaction is favoured and equilibrium will move left
  • If T ↓, exothermic reaction is favoured and equilibrium will move right
• A catalyst on an equilibrium reaction

  • Speeds up the forward reaction and the back reaction equally as it lowers the Ea for both the forward and back reactions by the same amount
  • The position of equilibrium is unaffected
  • The equilibrium will, however, be reached earlier when a catalyst is used
• The equation for the Haber Process is: N2(g) + 3H2(g) ⇋ 2NH3(g)
• The Haber Process is exothermic, so ∆H is negative
• The catalyst for the Haber Process is iron, Fe
• The ideal conditions for maximum yield in the Haber Process are pressure and temperature
• The actual conditions for the Haber Process are pressure and temperature
• The equation for the Contact Process is: 2SO2(g) + O2(g) ⇋ 2SO3(g)
• The Contact Process is exothermic, so ∆H is negative
• The catalyst for the Contact Process is vanadium pentoxide, V2O5
• The ideal conditions for maximum yield in the Contact Process are pressure and temperature
• The actual conditions for the Contact Process are pressure and temperature
• temperature change

  A) to the left

  B) to the right

  C) to the right

  D) to the left

  
• pressure/concentration

  A) to the right

  B) to the left

  C) no change

  
• concentration

  A) to the right

  B) to the left

  C) to the left

  D) to the right