Chemical Equilibrium

Introduction to Reversible Reactions

In many chemical reactions, reactants are completely converted into products. These are known as irreversible reactions. However, some reactions can proceed in both the forward and backward directions. These are called reversible reactions, represented by the double arrow (⇌).

Dynamic Equilibrium

Dynamic equilibrium occurs in a closed system when the rate of the forward reaction equals the rate of the reverse reaction. At this point, the concentrations of reactants and products remain constant over time, even though the reactions have not stopped. It is 'dynamic' because molecular activity continues, but 'equilibrium' because there is no net change in the system.

Le Chatelier's Principle

This principle states that if an external change (stress) is applied to a system at equilibrium, the system will shift its equilibrium position in a direction that tends to counteract or oppose that change.

Factors Affecting Equilibrium Position

• Concentration: Increasing the concentration of reactants shifts the equilibrium to the right (towards products). Increasing the concentration of products shifts it to the left (towards reactants).
• Pressure: This affects only gaseous systems where there is a change in the number of moles. An increase in pressure shifts the equilibrium toward the side with the fewer number of gas moles.
• Temperature: For an exothermic reaction (ΔH is negative), increasing temperature shifts the equilibrium to the left. For an endothermic reaction (ΔH is positive), increasing temperature shifts the equilibrium to the right.
• Catalysts: A catalyst increases the rate of both forward and backward reactions equally. Therefore, it does not affect the position of equilibrium; it only helps the system reach equilibrium faster.

The Equilibrium Constant (K)

The equilibrium constant (K_c or K_p) provides a mathematical relationship between the concentrations of products and reactants at equilibrium. Crucially, only temperature can change the numerical value of the equilibrium constant. Changes in concentration, pressure, or the addition of a catalyst do not change the value of K.

Specific Examples

• Action of Steam on Iron: 3Fe(s) + 4H₂O(g) ⇌ Fe₃O₄(s) + 4H₂(g). Since the moles of gas are equal on both sides (4 moles H₂O vs 4 moles H₂), pressure changes do not affect the equilibrium position.
• N₂O₄ / NO₂ System: N₂O₄(g) [Colorless] ⇌ 2NO₂(g) [Brown]. Increasing pressure shifts the equilibrium to the left (1 mole side), making the mixture lighter. Increasing temperature (this is endothermic) shifts it to the right, making it darker brown.