Summary
Equilibria involves reversible reactions where products can reform reactants, and the position of equilibrium can be influenced by changing conditions like pressure, concentration, and temperature. The Haber Process exemplifies the balance between yield and rate in industrial reactions.
- Reversible Reactions — reactions where products can revert to reactants. Example: N2 (g) + 3H2 (g) ⇌ 2NH3 (g)
- Equilibrium — the state where the forward and reverse reactions occur at the same rate. Example: No net change in concentrations of reactants and products.
- Le Chatelier’s Principle — principle stating that a system at equilibrium will adjust to counteract changes in conditions. Example: Increasing pressure shifts equilibrium towards fewer gas molecules.
- Yield-Rate Compromise — balancing reaction conditions to optimize both the rate and yield of a product. Example: The Haber Process uses 450°C and high pressure for ammonia production.
Exam Tips
Key Definitions to Remember
- Reversible Reactions
- Equilibrium
- Le Chatelier’s Principle
- Yield-Rate Compromise
Common Confusions
- Confusing the direction of equilibrium shift with changes in pressure or concentration.
- Misunderstanding the effect of temperature changes on exothermic and endothermic reactions.
Typical Exam Questions
- What happens to the equilibrium position when pressure is increased in a gaseous reaction? It shifts towards the side with fewer gas molecules.
- How does increasing temperature affect an exothermic reaction at equilibrium? It shifts the equilibrium position towards the endothermic direction.
- Why is a compromise temperature used in the Haber Process? To balance between a reasonable yield and an acceptable reaction rate.
What Examiners Usually Test
- Understanding of how equilibrium shifts with changes in conditions.
- Application of Le Chatelier’s Principle to different scenarios.
- Explanation of the yield-rate compromise in industrial processes like the Haber Process.
