Understanding Le Chatelier's Principle

This page introduces Le Chatelier's principle in equilibrium reactions. The principle states that when a chemical system at equilibrium is disturbed, it will shift to counteract the change and establish a new equilibrium.

Key points:

• If equilibrium shifts left, the backward reaction is faster, producing more reactants
• If equilibrium shifts right, the forward reaction is faster, producing more products
• Temperature, pressure, and concentration changes can disturb equilibrium
• The system will respond to minimize the effect of the disturbance

Example: For the reaction H₂(g) + I₂(g) ⇌ 2HI(g), a shift left produces more H₂ and I₂, while a shift right produces more HI.

Highlight: Catalysts do not affect equilibrium position, but they do speed up how quickly equilibrium is reached.

Worked Examples of Equilibrium Shifts

This page provides worked examples demonstrating effects of concentration on equilibrium shift for different reaction scenarios.

For the reaction CH₃COOH(l) + C₂H₅OH(l) ⇌ CH₃COOC₂H₅(l) + H₂O(l):

1. Adding more CH₃COOC₂H₅ shifts equilibrium left to produce more ethanoic acid and ethanol, opposing the increase.

2. Removing C₂H₅OH shifts equilibrium left to produce more ethanoic acid and ethanol, counteracting the removal.

For Ce⁴⁺(aq) + Fe²⁺(aq) ⇌ Ce³⁺(aq) + Fe³⁺(aq):

3. Adding water has no effect as it dilutes all ions equally, maintaining the ratio of reactants to products.

Highlight: These example scenarios of equilibrium position changes illustrate how Le Chatelier's principle applies to real chemical systems.