Chapter 17: Chemical Equilibrium

Le Chatelier's Principle

If a system at equilibrium is disturbed, the system will adjust to counteract the disturbance and restore equilibrium.

Types of Disturbances

Concentration Changes

• Increasing reactant concentration shifts equilibrium toward the products (forward reaction).
• Increasing product concentration shifts equilibrium toward the reactants (reverse reaction).

Temperature Changes

• Exothermic Reaction: Adding heat shifts equilibrium toward the reactants.
• Endothermic Reaction: Adding heat shifts equilibrium toward the products.

Pressure Changes

Affects reactions involving gases:

• Increasing pressure shifts equilibrium toward the side with fewer moles of gas.
• Decreasing pressure shifts equilibrium toward the side with more moles of gas.

The Effect of Catalysts

• Catalysts increase the rate of both the forward and reverse reactions.
• They help achieve equilibrium faster but do not alter the equilibrium position.

Applications of Le Chatelier's Principle

• Optimizing industrial chemical processes like the Haber process for ammonia production.
• Predicting changes in equilibrium when external conditions change.
• Controlling chemical reactions in laboratory and real-world settings.

Examples

Haber Process

$3H{2}_{\mathrm{g}}+N{2}_{\mathrm{g}}\rightleftharpoons 2\mathrm{NH}{3}_{\mathrm{g}}$

• High pressure shifts equilibrium to the side with fewer gas molecules (ammonia).
• Lower temperatures favor ammonia formation but reduce reaction rate.

Questions for Students

1. State Le Chatelier's Principle and explain its significance in chemical equilibrium.
2. How does increasing the concentration of a reactant affect equilibrium?
3. Describe the effect of temperature on an exothermic reaction at equilibrium.
4. Why do catalysts not affect the equilibrium position?
5. Explain how Le Chatelier's Principle is applied in the Haber process.