Equilibrium Changes

This revision note addresses changes in equilibrium with a focus on concentration and partial pressure, which are crucial in chemical reactions and processes, including the Haber process.

Chemical Equilibrium

• Chemical Equilibrium: This state is reached when the rates of the forward and reverse reactions are equal.
  • The concentrations of reactants and products remain steady.
  • It occurs within a closed system.

• Dynamic Nature:
  • This concept indicates that reactions continue to occur, continually adjusting to each other.

Equilibrium Constants

• Equilibrium Constant $K_c$:

  • Defined by the concentrations of reactants and products.
  • Expression: $K_c = \frac{[C][D]}{[A][B]}$

• Equilibrium Constant $K_p$:

  • Pertinent to gaseous reactions, defined in terms of partial pressures.
  • Expression: $K_p = \frac{P_C P_D}{P_A P_B}$

Le Chatelier's Principle

Le Chatelier's Principle: If a system at equilibrium experiences a change, it will respond to counteract the change and restore equilibrium.

• Predicts qualitative shifts when an equilibrium system is altered.

Steps to Apply Le Chatelier's Principle

1. Identify Changes:

   • Concentration: Evaluate whether reactant or product levels have changed.
   • Pressure: Consider how changes relate to the mole number of gases.
   • Temperature: Recognise whether the reaction is endothermic or exothermic.

2. Determine Shift Direction:

   • Concentration Shifts: An increase in reactants drives the equilibrium toward products.
     • Example: Adding reactants shifts the balance towards the products.
   • Pressure Shifts:
     • Example: A reduction in pressure could shift equilibrium to the side with more gas molecules.
   • Temperature Shifts: In endothermic reactions, added heat shifts the equilibrium as heat is absorbed.

3. Predict Outcomes:

   • Diagrammatic representations are helpful for visualising molecular arrangements.

Effect of Concentration

Introduction

• Concentration: The quantity of a substance within a specified volume.

Conceptual Framework

• Equilibrium Shift:
  • Changes in concentration affect the reaction's balance.

Example: Acid-Base Reaction

• Adding hydrochloric acid (HCl) to sodium acetate (CH₃COONa) causes equilibrium to shift right, compensating for increased hydrogen ion concentration.

Predicting Shifts

• An increase in concentration results in a movement toward equilibrium.

Effect of Pressure and Volume

• Pressure Changes:

  • Mole count: Predict movement using the number of moles.
  • Haber Process: Pressure influences ammonia production due to differences in mole numbers between reactants and products.

• Volume Changes:

  • A decrease in volume increases pressure, shifting equilibrium similarly to direct pressure changes.

Applications and Worked Examples

• Example 1: Haber Process

  • Increasing the levels of nitrogen ($N_2$) and hydrogen pushes the equilibrium right, resulting in more ammonia production.
  • This can be represented as: $N_2 + 3H_2 \rightleftharpoons 2NH_3$
  • When we add more $N_2$ or $H_2$, the system counteracts by forming more $NH_3$.

• Example 2: $NO_2$ and $N_2O_4$

  • The reaction is: $2NO_2 \rightleftharpoons N_2O_4$
  • Higher pressure supports the formation of $N_2O_4$, resulting in fewer gas molecules.
  • When pressure increases, the system shifts toward $N_2O_4$ (1 molecule) rather than $NO_2$ (2 molecules).

Practice Problems with Solutions

Problem 1: Simple Concentration Adjustment

• Given: In the reaction $H_2 + I_2 \rightleftharpoons 2HI$, the initial equilibrium has $[H_2] = 0.1M$, $[I_2] = 0.1M$, and $[HI] = 0.4M$. What happens if more $H_2$ is added to increase its concentration to $0.2M$?
• Solution: The equilibrium will shift right to form more HI. The system will use some of the added $H_2$ and some $I_2$ to create more HI until a new equilibrium is established.

Problem 2: Volume and Pressure Change

• Given: For the reaction $N_2(g) + 3H_2(g) \rightleftharpoons 2NH_3(g)$, what happens when the volume is halved?
• Solution: Halving the volume doubles the pressure. Since there are 4 moles of gas on the reactant side (1 + 3) and only 2 moles on the product side, the equilibrium shifts toward the product side (fewer moles of gas), producing more $NH_3$.

Problem 3: Temperature Changes

• Given: For the exothermic reaction $N_2(g) + 3H_2(g) \rightleftharpoons 2NH_3(g) + \text{heat}$, what happens when temperature increases?
• Solution: Since the reaction is exothermic (releases heat), increasing temperature will favour the reverse reaction (which absorbs heat). The equilibrium will shift left, resulting in more $N_2$ and $H_2$, and less $NH_3$.