Le Chatelier's Principle

Definition of Equilibrium

Equilibrium: A condition in a chemical reaction where

• The rate of the forward reaction equals that of the reverse reaction.
• The concentrations of products and reactants remain stable over time.

Equilibrium is a dynamic process, indicating that even though concentrations remain constant, molecular activities persist. This differs from a static state where no alterations occur.

Characteristics of Dynamic Equilibrium

• Steady concentrations are maintained.
• Rates of forward and reverse reactions are equal.
• Functions within a closed system with no exchange of matter with the surroundings.

Graphical Representation

• X-axis: Time
• Y-axis: Concentration

Recognising patterns of stabilisation is essential to understanding dynamic equilibrium.

Equilibrium Constant (K)

Equilibrium Constant (K): Denotes the ratio of the concentrations of products to reactants at equilibrium:

$K = \frac{[\text{Products}]}{[\text{Reactants}]}$

• K varies based on the type of reaction and is crucial for predicting the equilibrium position.
• Temperature impacts K. Alterations in temperature can shift equilibrium.

Introduction to Le Chatelier's Principle

Le Chatelier's Principle: "If a system at equilibrium experiences changes in concentration, temperature, or pressure, it will adjust to oppose the changes."

• Provides insight into how equilibrium adapts to changes.
• Critical for comprehending industrial applications like ammonia production.

Effects of Temperature on Equilibrium

• Endothermic Reactions: Absorb heat. An increase in temperature shifts equilibrium towards products.
  • Example: Baking bread, which involves heat absorption.
• Exothermic Reactions: Release heat. An increase in temperature shifts equilibrium towards reactants.
  • Example: Respiration, which involves heat release.

Effects of Volume and Pressure on Equilibrium

• Pressure Increase: Shifts the equilibrium toward fewer gas moles.
• Volume Decrease: Similar to increased pressure effect, shifting towards fewer moles.
• Pressure Decrease: Shifts equilibrium towards more gas moles.

Nitrogen Dioxide/Dinitrogen Tetroxide Interaction

• Chemical Equation: $\text{2NO}_2(g) \rightleftharpoons \text{N}_2\text{O}_4(g)$

ICE Tables

ICE Tables: Monitor initial, change, and equilibrium concentrations:

	A	B	C
Initial	1	1	0
Change	-x	-x	+x
Equilibrium	1-x	1-x	x

Worked Example: For a reaction A + B ⇌ C with initial concentrations of A = 1 mol/L, B = 1 mol/L, and C = 0 mol/L, and an equilibrium constant K = 4:

1. Set up the ICE table as shown above
2. At equilibrium: K = [C]/([A][B]) = x/((1-x)(1-x)) = 4
3. Simplify: x/(1-x)² = 4
4. Solve: (1-x)² = x/4
5. Expand: 1-2x+x² = x/4
6. Multiply both sides by 4: 4-8x+4x² = x
7. Rearrange: 4x²-9x+4 = 0
8. Using the quadratic formula: x = (9 ± √(81-64))/8 = (9 ± √17)/8
9. Since x must be less than 1 (as it's a concentration change): x ≈ 0.61 mol/L

Therefore, at equilibrium [A] = [B] = 0.39 mol/L and [C] = 0.61 mol/L

Graphical Analysis

These graphs illustrate how reactant and product concentrations evolve over time until equilibrium is achieved.

• Components:
  • X-axis: Time
  • Y-axis: Concentration

Safety Considerations

• Always use gloves, goggles, and a lab coat.
• Ensure proper ventilation in the laboratory.

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