Chemical Equilibrium and Le Chatelier’s Principle: Study Notes for GOB Chemistry

Chemical Equilibrium

Introduction to Equilibrium

Chemical equilibrium occurs when the rates of the forward and reverse reactions in a chemical system are equal, resulting in constant concentrations of reactants and products. This concept is fundamental in understanding how chemical reactions behave under various conditions.

• Kinetics studies the rate at which reactants are converted into products.

• Thermodynamics deals with the direction and extent to which a chemical reaction at equilibrium will shift.

• Le Chatelier’s Principle states that if a system at equilibrium is disturbed, it will adjust itself to counteract the disturbance and restore equilibrium.

Le Chatelier’s Principle

Definition and Application

Le Chatelier’s Principle helps predict how a system at equilibrium will respond to changes in concentration, temperature, pressure, or the addition of catalysts.

• Disturbances include changes in concentration, temperature, pressure, or the addition/removal of substances.

• The system will shift in the direction that minimizes the disturbance.

Example Reaction

For the endothermic reaction: $4\text{NH}_3 (g) + 3\text{O}_2 (g) \rightleftharpoons 2\text{N}_2 (g) + 6\text{H}_2\text{O} (g)$, $K_c = 6.73 \times 10^4$

Change	Direction of Shift
Addition of a catalyst	Forward (no change in equilibrium position, only rate increases)
Decreasing the volume	Forward (shifts to side with fewer moles of gas)
Removing H2O (g)	Forward (shifts to replace removed product)
Increasing the temperature	Forward (endothermic reaction favors products with heat added)
Addition of NH3 (g)	Forward (shifts to use up added reactant)
Decreasing the pressure	Forward (shifts to side with more moles of gas)
Removing NH3 (g)	Forward (shifts to replace removed reactant)
Addition of a precipitate	Forward (if it removes a product from solution)
Addition of an inert gas at constant volume	Forward (no effect on equilibrium position, but may affect partial pressures)

Factors Affecting Equilibrium

Concentration, Pressure, and Temperature

• Concentration: Adding or removing reactants/products shifts equilibrium to oppose the change.

• Pressure/Volume: Increasing pressure (by decreasing volume) shifts equilibrium toward the side with fewer moles of gas.

• Temperature: For endothermic reactions, increasing temperature favors products; for exothermic, it favors reactants.

• Catalysts: Increase the rate at which equilibrium is reached but do not affect the position of equilibrium.

Practice Example: Shifting Equilibrium

For the reaction: $\text{CH}_4 (g) + \text{F}_2 (g) \rightleftharpoons \text{CF}_4 (g) + \text{HF} (g)$, $\Delta H = +38.1$ kJ/mol

• Adding CF4: Shifts left (increases product concentration)

• Removing F2: Shifts left (decreases reactant concentration)

• Decrease temperature: Shifts left (endothermic reaction, lowering temperature favors reactants)

• Decrease container volume: Shifts left (if fewer moles of gas on left)

• Increase partial pressure of HF: Shifts left (increases product concentration)

• Exception: Decreasing the container volume may not always shift left, depending on the number of moles of gas on each side.

Thermodynamics and Equilibrium Constants

Equilibrium Constant ($K$) and Temperature

• The equilibrium constant $K$ changes with temperature.

• For endothermic reactions, $K$ increases with temperature; for exothermic, $K$ decreases.

• Example: If $K$ increases with temperature, the reaction is endothermic.

Practice Example: Interpreting $K$ Values

Given: $2A (g) + 3B (g) \rightleftharpoons C (g)$ At 25°C, $K = 5.2 \times 10^4$; at 50°C, $K = 1.7 \times 10^7$

• Conclusion: The reaction is endothermic (since $K$ increases with temperature).

Effect of Catalysts on Equilibrium

Role of Catalysts

• Catalysts speed up both the forward and reverse reactions equally.

• They do not change the position of equilibrium, only the rate at which equilibrium is achieved.

Practice Example: Catalyst Addition

For the reaction: $\text{CaCO}_3 (s) \rightleftharpoons \text{CaO} (s) + \text{CO}_2 (g)$, $K_c = 3.2 \times 10^{-23}$

• Adding a catalyst does not shift equilibrium; it only increases the rate of reaching equilibrium.

Visualizing Equilibrium: Particle Diagrams

Interpreting Particle Diagrams

• Particle diagrams can be used to visualize the progress of a reaction toward equilibrium.

• The container with the greatest number of product particles (e.g., AB) indicates the reaction has proceeded further toward completion.

Example:

For the reaction $A_2 (g) + B_2 (g) \rightarrow 2 AB (g)$, the container with more AB molecules (as shown in the diagram) has proceeded more to completion.

Additional info: These notes are based on standard GOB Chemistry curriculum topics related to chemical equilibrium, Le Chatelier’s Principle, and the effect of various changes on equilibrium systems.