Dynamic Equilibrium & Le Chatelier's Principle

Understanding Equilibrium Constant $K_c$

In a chemical reaction, when the rate of the forward reaction equals the rate of the reverse reaction, the system is in dynamic equilibrium. At this point, the concentrations of reactants and products remain constant over time.

The equilibrium constant $K_c$ is a value that expresses the ratio of the concentrations of products to reactants at equilibrium, each raised to the power of their coefficients in the balanced equation.

For a general reaction:

$$aA + bB \rightleftharpoons cC + dD$$

The equilibrium constant $K_c$ is given by:

$$K_c = \frac{[C]^c[D]^d}{[A]^a[B]^b}$$

Effects on the Position of $K_c$

1. Effect of Pressure

• Changing the pressure affects the position of equilibrium only if gases are involved and the number of moles of gas changes in the reaction.
• Increasing pressure shifts the equilibrium towards the side with fewer moles of gas.
• Decreasing pressure shifts the equilibrium towards the side with more moles of gas.

2. Effect of Temperature

• Temperature changes can affect both the position of equilibrium and the value of $K_c$.
• For an exothermic reaction, increasing temperature shifts the equilibrium to the left (towards reactants), decreasing $K_c$.
• For an endothermic reaction, increasing temperature shifts the equilibrium to the right (towards products), increasing $K_c$.

3. Effect of Concentration

• Changing the concentration of reactants or products shifts the equilibrium to oppose the change.
• Increasing the concentration of reactants shifts the equilibrium towards the products.
• Increasing the concentration of products shifts the equilibrium towards the reactants.

Examples

For the reaction: $$N_2(g) + 3H_2(g) \rightleftharpoons 2NH_3(g)$$

• Pressure: Increasing pressure shifts equilibrium towards the right (fewer moles of gas).
• Temperature: If exothermic, increasing temperature shifts equilibrium to the left.
• Concentration: Increasing $[N_2]$ shifts equilibrium to the right.