Chemistry- Equilibria

• Reversible reaction


  A reaction that can go either direction depending on the conditions of the reaction
• Reversible reaction


  1. Forward reaction (reactants to products)

  2. Reverse reaction (products to reactants)
• Examples of reversible reactions


  • Ammonium chloride dissociation
  • Hydrated copper(II) sulfate dehydration
• Static equilibrium


  Both forward and backward reactions stopped at equilibrium
• Dynamic equilibrium


  Reactants changing to products and products changing to reactants at equal rates

  
• Equilibrium can only be achieved if none of the reactants or products can escape
• Le Chatelier's Principle


  If any change is made to the conditions of an equilibrium, the equilibrium will shift to oppose that change

  
• Increasing concentration of a substance in an equilibrium

  Equilibrium shifts to counteract the change

  
• Adding dilute sodium hydroxide to bromine water equilibrium

  Equilibrium shifts to decrease H+ ion concentration, making the solution colourless

  
• Adding dilute sulfuric acid to bromine water equilibrium


  Equilibrium shifts to increase H+ ion concentration, making the solution yellow-brown
• Adding acid to chromate/dichromate equilibrium


  Equilibrium shifts to favour dichromate, making the solution orange
• Adding hydroxide ions to chromate/dichromate equilibrium


  Equilibrium shifts to favour chromate, making the solution yellow
• Increasing temperature of N2O4 ⇌ 2NO2 equilibrium


  Equilibrium shifts to favour endothermic forward reaction, making the solution brown

  
• Decreasing temperature of N2O4 ⇌ 2NO2 equilibrium


  Equilibrium shifts to favour exothermic reverse reaction, making the solution colourless

  
• Enthalpy change (ΔH)


  The heat absorbed or released during a chemical reaction

  
• Equilibrium between colourless dinitrogen tetroxide (N2O4) and brown nitrogen dioxide (NO2)

  1. N2O4 (g) ⇌ 2NO2 (g)

  2. ΔH = +ve

  
• An increase in temperature


  Favours the endothermic reaction
• In the N2O4 ⇌ 2NO2 equilibrium


  The reaction is endothermic
• Increasing the temperature


  Shifts the equilibrium to the right, resulting in more NO2 being formed and the reaction mixture becoming brown

  
• A decrease in temperature


  Favours the exothermic reaction
• Decreasing the temperature in the N2O4 ⇌ 2NO2 equilibrium


  Shifts the equilibrium to the left, resulting in more N2O4 being formed and the reaction mixture becoming colourless

  
• Exothermic reaction


  A reaction accompanied by the evolution of heat
• Endothermic reaction


  A reaction accompanied by the absorption of heat
• If the temperature of the reaction vessel is increased

  The exothermic reaction is favoured so that the temperature is lowered
• If the temperature of the reaction vessel is decreased

  The endothermic reaction is favoured so that the temperature is increased
• Le Chatelier's Principle


  If a system at equilibrium is subjected to a change in one of the conditions (concentration, temperature, pressure), the system will shift to counteract the change and establish a new equilibrium

  
• If the pressure of a system is increased


  The equilibrium will shift to favour the side with the fewer number of gaseous molecules
• If the pressure of a system is decreased


  The equilibrium will shift to favour the side with the greater number of gaseous molecules
• Equilibrium: 2SO2 (g) + O2 (g) ⇌ 2SO3 (g)

  1. 3 moles of gaseous reactants on the left-hand side, 2 moles of gaseous products on the right-hand side

  2. An increase in pressure would move the equilibrium to the right and result in more SO3 being formed
• Pressure can only affect the position of equilibrium if there is a change in the total gas volume
• Adding a catalyst to a reaction at equilibrium has no effect on the position of equilibrium
• Catalyst


  Speeds up the rate at which a reaction reaches equilibrium (rate of attainment of equilibrium), but does not affect the concentration of the products (yield) or reactants
• Haber Process - Industrial Manufacture of Ammonia

  1. N2 (g) + 3H2 (g) ⇌ 2NH3 (g)

  2. The forward reaction is exothermic

  
• To favour the forward reaction (producing more ammonia) in the Haber Process:

  Use high pressure, high temperature, and high concentrations of reactants
• The catalyst in the Haber Process speeds up the rate at which equilibrium is reached, but does not increase the yield of ammonia
• The forward reaction in the Contact Process


  Is exothermic, so lowering the temperature should cause the equilibrium to move to the right, producing more SO3
• The forward reaction in the Contact Process

  Brings about a decrease in pressure, so increasing the pressure should favour the forward reaction
• The catalyst in the Contact Process speeds up the rate at which equilibrium is reached, but does not increase the yield of SO3
• Conversion of SO3 to sulfuric acid

  1. SO3 (g) + H2SO4 (l) → H2S2O7 (l) (oleum)

  2. H2S2O7 (l) + H2O (l) → 2H2SO4 (l)