Equilibrium

Dynamic Equilibrium

  • In chemical reactions, when the rate of the forward reaction equals the rate of the reverse reaction, a state of dynamic equilibrium is reached.
  • It’s crucial to remember that equilibrium does not mean that the reactions cease. Instead, the concentrations of reactants and products become constant.
  • The system at dynamic equilibrium exhibits certain characteristics. For instance, it can only be achieved in a closed system — no reactants or products are allowed to escape.
  • A closed system could be a sealed container or a solution where precipitation or gas formation reactions are occurring.

Le Chatelier’s Principle

  • Le Chatelier’s principle is a tool that helps in predicting the effect of changes in conditions (pressure, temperature, concentration) on the position of equilibrium.
  • Changes in concentration: When the concentration of a reactant or product changes, the equilibrium shifts to counter this change. If a reactant is added, equilibrium will shift to the right to use up the added reactant. If a product is removed, equilibrium will shift to the right to produce more product.
  • Changes in pressure: Pressure changes only affect gaseous reactions where there is a change in the number of moles. If pressure is increased, the equilibrium will shift towards the side with fewer moles of gas, and vice-versa.
  • Changes in temperature: An increase in temperature will favour the endothermic reaction (absorbs heat). Similarly, a decrease in temperature will favour the exothermic reaction (exudes heat).

Haber Process

  • The Haber process is a practical application of the concepts of equilibrium and Le Chatelier’s principle.
  • This process involves the reaction of nitrogen gas and hydrogen gas to produce ammonia: N2(g) + 3H2(g) ⇌ 2NH3(g)
  • The production of ammonia via the Haber process is exothermic. Therefore, increasing the temperature will shift the equilibrium position to the left, reducing the yield of ammonia.
  • The forward reaction involves a decrease in the number of gas moles (from 4 to 2). Therefore, increasing the pressure will shift the equilibrium to the right, increasing the ammonia yield. However, a balance between the rate of reaction and cost needs to be sought in practice.

Remember, understanding how to manipulate the position of equilibrium is essential for industrial processes and can also have significant economic implications. A thorough understanding of the equilibrium principles is thus fundamental.