Le Chatelier’s Principle (OCR A Level Chemistry A): Revision Note

Equilibrium & Reaction Conditions

Position of the equilibrium

• The position of the equilibrium refers to the relative amounts of products and reactants in an equilibrium mixture

• When the position of equilibrium shifts to the left, it means the concentration of reactants increases

• When the position of equilibrium shifts to the right, it means the concentration of products increases

Le Chatelier’s principle

• Le Chatelier’s principle says that if a change is made to a system in dynamic equilibrium, the position of the equilibrium moves to counteract this change

• The principle is used to predict changes to the position of equilibrium when there are changes in temperature, pressure or concentration

Effects of concentration

How the equilibrium shifts with concentration changes

Effects of pressure

• Changes in pressure only affect reactions where the reactants or products are gases

How the equilibrium shifts with pressure changes

Effects of temperature

How the equilibrium shifts with temperature changes

Equilibrium & Catalysts

Effects of catalysts

• A catalyst is a substance that increases the rate of a chemical reaction (they increase the rate of the forward and reverse reaction equally)

• Catalysts only cause a reaction to reach equilibrium faster

• Catalysts therefore have no effect on the position of the equilibrium once this is reached

Investigating changes to the equilibrium position with concentration

• The equilibrium between aqueous chromate ions, CrO₄^2-, and dichromate ions, Cr₂O₇^2- is sensitive to changes in acid concentration

• Solutions of chromate and dichromate ions have different colours so it is easy to see any shift in the equilibrium position

• By adding dilute sulfuric acid, we can increase the concentration of H⁺ (aq) in the solution

• This increases the rate of the forward reaction causing the equilibrium position to shift to minimise the change in H⁺ (aq) concentration

  • This decreases the concentration of the added reactant, H⁺ (aq)

  • Equilibrium shifts to the right, making more products

  • Solution turns orange due to the formation of more Cr₂O₇^2- (aq)

• By adding aqueous sodium hydroxide, we can decrease the concentration of H⁺ (aq) in the solution

• The added OH^-(aq) ions react with H⁺(aq) ions forming water

• H⁺ (aq) + OH^- (aq) → H₂O (l)

• This decreases the rate of the forward reaction causing the equilibrium position to shift to minimise the change in H⁺(aq) concentration

  • This decreases the concentration of reactant that has been removed, H⁺ (aq)

  • Equilibrium shifts to the left, making more H⁺ (aq) reactant

  • Solution turns yellow due to the formation of more CrO₄^2- (aq) 

Investigating changes in equilibrium position with temperature

• Cobalt chloride, CoCl₂, dissolves in water to form a pink solution

• The dissolving process produces an equilibrium between two different coloured cobalt complexes

[Co(H₂O)₆]²⁺ (aq) + 4Cl^- (aq) ⇌ [CoCl₄]^2- (aq) + 6H₂O (l)

• The forward reaction in this process is endothermic and the backward reaction is exothermic

• By heating up the solution we can increase the amount of heat energy in the system

  • This causes the equilibrium to shift to minimise the change

    • Equilibrium shifts to the right favouring the endothermic reaction (ΔH is positive)

    • This allows the system to take heat energy in and minimise the increase in temperature

    • The solution turns blue as more CoCl₄^2- (aq) is formed

• Cooling down the solution removes the heat energy from the system

  • This again causes the equilibrium to shift to minimise the change

    • Equilibrium shifts to the left favouring the exothermic reaction (ΔH is negative)

    • This allows the system to release heat energy and minimise the decrease in temperature

    • The solution turns pink as more Co(H₂O)₆²⁺ (aq) is formed

Operational Conditions

• Equilibrium reactions are involved in some stages of large-scale production of certain chemicals

• An understanding of equilibrium and Le Chatelier’s principle is therefore very important in the chemical industry

Haber process

• The Haber process involves the synthesis of ammonia according to:

N­₂ (g) + 3H₂ (g) ⇌ 2NH₃ (g)      ΔH_r = -92 kJ mol^-1

• Le Chatelier’s principle is used to get the best yield of ammonia

Maximising the ammonia yield

Pressure

• An increase in pressure will result in the equilibrium shifting in the direction of the fewest molecules of gas formed to reduce the pressure

• In this case, the equilibrium shifts towards the right so the yield of ammonia increases

• An increase in pressure will cause the particles to be closer together and therefore increasing the number of successful collisions leading to an increased reaction rate

• Very high pressures are expensive to produce therefore a compromise pressure of 200 atm is chosen

Temperature

• To get the maximum yield of ammonia the position of equilibrium should be shifted as far as possible to the right as possible

• Since the Haber process is an exothermic reaction, according to Le Chatelier’s principle the equilibrium will shift to the right if the temperature is lowered

• A decrease in temperature will decrease the energy of the surroundings so the reaction will go in the direction in which energy is released to counteract this

• Since the reaction is exothermic, the equilibrium shifts to the right

• However, at a low temperature the gases won’t have enough kinetic energy to collide and react and therefore equilibrium would not be reached therefore compromise temperature of 400-450 ^oC is used in the Haber process

• A heat exchanger warms the incoming gas mixture to give molecules more kinetic energy such that the gas molecules collide more frequently increasing the likelihood of a reaction

Catalysts

• In the absence of a catalyst the reaction is so slow that hardly anything happens in a reasonable time!

• Adding an iron catalyst speeds up the rate of reaction

Contact process

• The Contact process involves the synthesis of sulfuric acid according to:

2SO­₂ (g) + O₂ (g) ⇌ 2SO₃ (g)   ΔH_r = -197 kJ mol^-1

• Le Chatelier’s principle is used to get the best yield of sulfuric acid

Maximising the sulfuric acid yield

Pressure

• An increase in pressure will result in the equilibrium shifting in the direction of the fewest molecules of gas formed to reduce the pressure

• In this case, the equilibrium shifts towards the right so the yield of sulfuric acid increases

• In practice, the reaction is carried out at only 1 atm

• This is because Kp for this reaction is already very high meaning that the position of the equilibrium is already far over to the right

• Higher pressures than 1 atm will be unnecessary and expensive

Temperature

• The same principle applies to increasing the temperature in the Contact process as in the Haber process

• A compromise temperature of 450 ^oC is used

Catalysts

• The Contact process uses vanadium(V) oxide as a catalyst to increase the rate of reaction