Equilibria

This section covers the concept of chemical equilibrium, including dynamic equilibrium, Le Chatelier's principle, and equilibrium constants. Understanding these concepts is crucial for predicting the outcome of reversible reactions.

Dynamic Equilibrium

A dynamic equilibrium is a state where the rate of the forward reaction is equal to the rate of the reverse reaction. This does not mean the reactions have stopped; rather, they are proceeding at equal rates, resulting in no net change in the concentrations of reactants and products.

Imagine a reaction in a closed container. The forward reaction produces products, and the reverse reaction consumes products and produces reactants. As the forward reaction proceeds, the concentration of reactants decreases, and the concentration of products increases. Simultaneously, the reverse reaction consumes products and produces reactants, counteracting the forward reaction. Eventually, the rates of the forward and reverse reactions become equal, and a dynamic equilibrium is established.

Key Features of Dynamic Equilibrium

• The rates of the forward and reverse reactions are equal.
• The concentrations of reactants and products remain constant (but are not necessarily equal).
• The system is not static; both forward and reverse reactions are still occurring.

Le Chatelier's Principle

Le Chatelier's principle states that if a change of condition is applied to a system in equilibrium, the system will shift in a direction that relieves the stress. 'Stress' refers to any change that disturbs the equilibrium.

Common stresses include:

• Change in Concentration: Adding reactants shifts the equilibrium towards products; adding products shifts the equilibrium towards reactants.
• Change in Pressure: For gaseous reactions, 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.
• Change in Temperature: For exothermic reactions, increasing temperature shifts the equilibrium towards reactants; for endothermic reactions, increasing temperature shifts the equilibrium towards products.
• Addition of an Inert Gas: Adding an inert gas at constant volume does not affect the equilibrium position. Adding an inert gas at constant pressure does not affect the equilibrium position.

Examples of Le Chatelier's Principle

1. Increasing the concentration of reactants: If $A + B \rightleftharpoons C + D$ is at equilibrium, adding more A and B will shift the equilibrium to the right, favouring the formation of C and D.
2. Increasing the pressure: If $N_2(g) + 3H_2(g) \rightleftharpoons 2NH_3(g)$ is at equilibrium, increasing the pressure will shift the equilibrium to the right, favouring the formation of ammonia (NH₃) because there are fewer moles of gas on the product side.
3. Increasing the temperature: If $N_2(g) + 3H_2(g) \rightleftharpoons 2NH_3(g)$ is at equilibrium, increasing the temperature will shift the equilibrium to the left, favouring the decomposition of ammonia into nitrogen and hydrogen because the reaction is exothermic.

Equilibrium Constants

An equilibrium constant, denoted by K, is a numerical value that expresses the ratio of products to reactants at equilibrium. It provides a quantitative measure of the extent to which a reaction proceeds to completion.

For the general reversible reaction:

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

The equilibrium constant, K, is defined as:

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

Where [A], [B], [C], and [D] represent the equilibrium concentrations of reactants and products, and a, b, c, and d are their stoichiometric coefficients.

Types of Equilibrium Constants

• K_c: The equilibrium constant expressed in terms of molar concentrations.
• K_p: The equilibrium constant expressed in terms of partial pressures (for gaseous reactions). It is related to K_c by: $$K_p = K_c (RT)^{\Delta n}$$, where Δn is the change in the number of moles of gas (moles of gaseous products - moles of gaseous reactants).

Interpreting the Value of K

Value of K	Extent of Reaction
K > 1	Equilibrium lies to the right (favours products)
K < 1	Equilibrium lies to the left (favours reactants)
K = 1	Equilibrium is at the point where reactants and products are present in roughly equal amounts.

The value of K is temperature-dependent. Changes in temperature will cause a change in the value of K, shifting the equilibrium to favour either the products or the reactants, depending on whether the reaction is exothermic or endothermic.