NEET Notes: Chemical Equilibrium

Introduction

Chemical equilibrium is a state of balance in a chemical reaction where the rates of the forward and reverse reactions are equal. This results in the concentration of reactants and products remaining constant over time. Chemical equilibrium plays a crucial role in understanding dynamic processes in chemical reactions and is essential for NEET aspirants to master.

Dynamic Nature of Chemical Equilibrium

Chemical equilibrium is dynamic, meaning that even though the concentrations of reactants and products remain constant, the reactions are still occurring. The forward and reverse reactions continue at the same rate, and this balance is what defines equilibrium. At equilibrium, there is no net change in the concentration of substances involved in the reaction.

Characteristics of Chemical Equilibrium

• The rate of the forward reaction is equal to the rate of the reverse reaction.
• The concentration of reactants and products remains constant but not necessarily equal.
• The equilibrium can be achieved from either direction: starting with reactants or products.
• At equilibrium, the system appears stable and unchanging, though it is dynamic on the molecular level.

Law of Chemical Equilibrium

The law of chemical equilibrium, also known as the Equilibrium Law, states that for a reversible reaction at a constant temperature, the ratio of the product of the concentrations of the products to the product of the concentrations of the reactants, each raised to the power of their respective stoichiometric coefficients, is constant. This constant is called the equilibrium constant (K_c).

Equilibrium Constant (K_c)

The equilibrium constant is a measure of the position of equilibrium for a reaction. It is defined for a general reversible reaction:

aA + bB ⇌ cC + dD

The equilibrium constant expression is:

Kc = [C]c [D]d / [A]a [B]b

Where the concentrations of the reactants (A, B) and products (C, D) are expressed in moles per liter (mol/L), and the exponents are the stoichiometric coefficients of the substances in the balanced chemical equation.

Types of Equilibrium

Homogeneous Equilibrium

In a homogeneous equilibrium, all the reactants and products are in the same phase (solid, liquid, or gas). An example is the equilibrium established in a gas-phase reaction.

N2 (g) + 3H2 (g) ⇌ 2NH3 (g)

Heterogeneous Equilibrium

In heterogeneous equilibrium, reactants and products are present in different phases. The equilibrium expression for such reactions is written considering only the concentrations of the gaseous and aqueous components, as solids and liquids do not appear in the expression.

CaCO3 (s) ⇌ CaO (s) + CO2 (g)

Le Chatelier's Principle

Le Chatelier’s Principle states that if a system at equilibrium is disturbed by a change in temperature, pressure, or concentration of one of the reactants or products, the system will shift in the direction that tends to counteract the disturbance. This principle helps predict the direction of shift when equilibrium is disturbed.

Effect of Concentration

Increasing the concentration of reactants will cause the equilibrium to shift towards the products to restore balance. Conversely, increasing the concentration of products will shift the equilibrium towards the reactants.

Effect of Pressure

In gaseous reactions, increasing the pressure by decreasing the volume will shift the equilibrium towards the side with fewer moles of gas. Similarly, decreasing the pressure will shift the equilibrium towards the side with more moles of gas.

Effect of Temperature

If the reaction is exothermic (releases heat), increasing the temperature will shift the equilibrium towards the reactants (endothermic direction). Conversely, decreasing the temperature will shift the equilibrium towards the products. For endothermic reactions, the opposite occurs.

Equilibrium Constant and Temperature

The equilibrium constant (K_c) is temperature-dependent. For exothermic reactions, K_c decreases with an increase in temperature, while for endothermic reactions, K_c increases with temperature. This relationship is described by the Van’t Hoff equation:

ln(K2/K1) = -ΔH/R (1/T2 - 1/T1)

Importance of Chemical Equilibrium

Chemical equilibrium has significant applications in various industries and biological systems. Understanding the factors that affect equilibrium helps in controlling reactions to achieve desired products and efficiencies. For example, the production of ammonia in the Haber process, the synthesis of sulfuric acid, and even metabolic pathways in the human body are all influenced by chemical equilibrium principles.

Applications of Chemical Equilibrium

Industrial Applications

• Haber Process: The synthesis of ammonia from nitrogen and hydrogen gas.
• Contact Process: The production of sulfuric acid from sulfur dioxide.
• Fertilizer Industry: Ammonia and phosphate fertilizers are produced by manipulating equilibrium conditions.

Biological Systems

• Respiratory Equilibrium: The exchange of oxygen and carbon dioxide in the lungs.
• Buffer Systems: Biological systems maintain pH stability by regulating equilibrium reactions in blood and cells.

Solubility Equilibria

In solubility equilibria, a substance dissolves to form a solution until the maximum solubility is reached. At this point, the rate of dissolution equals the rate of crystallization, and equilibrium is established.

The solubility product constant (K_sp) is used to describe the equilibrium position of a sparingly soluble salt in water:

Ksp = [A]m [B]n

Common Ion Effect

The common ion effect occurs when an ion already present in a solution reduces the solubility of a salt that contains that ion. This is due to Le Chatelier’s Principle, where the equilibrium shifts to reduce the concentration of the common ion.

Conclusion

Chemical equilibrium is a crucial concept in chemistry that explains the behavior of reversible reactions. By understanding the principles of equilibrium, such as the equilibrium constant, Le Chatelier’s Principle, and the factors that affect equilibrium, NEET aspirants can better grasp how reactions proceed and are controlled in various contexts, from industrial applications to biological systems.