Le Chatelier’s Principle and Yield Optimisation

Introduction to Chemical Equilibrium

• Many chemical reactions are reversible, meaning they can proceed in both forward and reverse directions.
• A dynamic equilibrium is established when the rate of the forward reaction equals the rate of the reverse reaction. The concentrations of reactants and products remain constant at equilibrium, but the reactions continue to occur.
• Extent of reaction refers to how far a reaction proceeds towards completion.
• Yield is the amount of product obtained in a chemical reaction.

KEY TAKEAWAY: Equilibrium is a dynamic state where forward and reverse reaction rates are equal, not a static state where reactions stop.

Le Chatelier’s Principle

• Le Chatelier’s Principle: If a change of condition (stress) is applied to a system in equilibrium, the system will shift in a direction that relieves the stress.
• Stresses include:
  • Change in concentration
  • Change in pressure (for gaseous systems)
  • Change in temperature

Effect of Concentration Changes

• Adding reactant: Shifts the equilibrium to the right (towards product formation) to consume the added reactant.
• Adding product: Shifts the equilibrium to the left (towards reactant formation) to consume the added product.
• Removing reactant: Shifts the equilibrium to the left (towards reactant formation) to replenish the removed reactant.
• Removing product: Shifts the equilibrium to the right (towards product formation) to replenish the removed product.

Effect of Pressure Changes (Gaseous Systems)

• Pressure changes primarily affect gaseous equilibria.
• Increasing pressure: Shifts the equilibrium towards the side with fewer moles of gas to reduce the pressure.
• Decreasing pressure: Shifts the equilibrium towards the side with more moles of gas to increase the pressure.
• Adding an inert gas: No change in equilibrium position, provided the volume of the container remains constant (as the partial pressures of the reactants and products remain unchanged).

Effect of Temperature Changes

• The effect of temperature depends on whether the reaction is exothermic or endothermic.
• Exothermic reaction: Releases heat ($ΔH < 0$). Treat heat as a product.
  • Increasing temperature: Shifts the equilibrium to the left (towards reactants) to consume the added heat.
  • Decreasing temperature: Shifts the equilibrium to the right (towards products) to generate heat.
• Endothermic reaction: Absorbs heat ($ΔH > 0$). Treat heat as a reactant.
  • Increasing temperature: Shifts the equilibrium to the right (towards products) to consume the added heat.
  • Decreasing temperature: Shifts the equilibrium to the left (towards reactants) to generate heat.

Effect of Catalysts

• Catalysts increase the rate of both forward and reverse reactions equally.
• Catalysts do not affect the equilibrium position or the yield of the reaction. They only help the reaction reach equilibrium faster.

EXAM TIP: When applying Le Chatelier’s principle, always state the stress applied, the direction of the shift, and why the equilibrium shifts in that direction.

Optimising Reaction Conditions for Yield

• Optimising yield involves considering both rate and equilibrium factors.
• A compromise may be necessary to achieve both a reasonable rate and a high yield.

Factors Affecting Rate and Yield

Example: Haber Process

• The Haber process is the synthesis of ammonia ($NH_3$) from nitrogen ($N_2$) and hydrogen ($H_2$):

  $N_2(g) + 3H_2(g) \rightleftharpoons 2NH_3(g)$ $ΔH = -92 kJ/mol$ (exothermic)

• Optimisation:

  • Temperature: Lower temperatures favour ammonia production (exothermic), but the reaction rate is slow. A compromise temperature of around 400-450°C is used.
  • Pressure: High pressure favours ammonia production (fewer moles of gas on the product side). High pressures (200-300 atm) are used.
  • Concentration: Excess hydrogen can be used.
  • Catalyst: Iron catalyst is used to increase the rate of reaction.
  • Product Removal: Ammonia is liquefied and removed to shift the equilibrium towards product formation.

Using Graphical and Thermochemical Information

• To determine optimal conditions, consider:
  • Thermochemical equation: Indicates whether the reaction is exothermic or endothermic.
  • Equilibrium constant (K): Indicates the relative amounts of reactants and products at equilibrium. A large K indicates that the equilibrium lies towards the products.
  • Graph of yield vs. temperature: Shows the relationship between temperature and yield.

STUDY HINT: Practice applying Le Chatelier’s principle to different reactions and scenarios. Draw concentration-time graphs to visualize the changes in concentration as equilibrium shifts.

Limitations of Le Chatelier’s Principle

• Le Chatelier’s principle provides a qualitative prediction of the direction of equilibrium shift.
• It does not provide quantitative information about the extent of the shift.
• It assumes ideal conditions and may not be accurate for complex systems.

COMMON MISTAKE: Forgetting to consider the sign of ΔH when predicting the effect of temperature changes on equilibrium. Remember exothermic reactions release heat (treat heat as a product), and endothermic reactions absorb heat (treat heat as a reactant).

Summary of Key Concepts

• Reversible reactions and dynamic equilibrium.
• Le Chatelier’s principle: Predicting the effect of changes in concentration, pressure, and temperature on equilibrium position.
• Optimising reaction conditions: Balancing rate and yield considerations.
• Understanding the role of catalysts.

VCAA FOCUS: VCAA often presents equilibrium scenarios and asks students to predict the effect of different changes on the equilibrium position and to justify their answers using Le Chatelier’s principle. Expect questions involving the Haber process and other industrial applications.