Class 11 Chemistry Chapter – Equilibrium introduces students to the dynamic balance that exists in reversible reactions. It helps explain how forward and backward reactions occur simultaneously and how the rate of these reactions becomes equal over time. This chapter is the foundation for understanding acid-base chemistry, solubility, and ionic equilibria in higher grades.
In real life, equilibrium governs many processes—from the solubility of gases in oceans to the functioning of our lungs. The NCERT Class 11 Equilibrium chapter develops analytical thinking, numerical ability, and reasoning through laws like Le Chatelier’s Principle and the concept of equilibrium constant.
In this detailed explanation, we’ll study the key topics, formulas, tables, and solved examples that make equilibrium easy to master. By the end, you’ll be able to confidently answer both numerical and conceptual questions in your board and Competitive Exams.
Table of Contents
- Types of Chemical Equilibrium
- Law of Mass Action and Equilibrium Constant
- Le Chatelier’s Principle
- Ionic Equilibrium
- FAQs
Types of Chemical Equilibrium
Physical and Chemical Equilibrium
| Type | Definition | Example |
|---|---|---|
| Physical Equilibrium | Balance between two physical states of a substance | Liquid ⇌ Vapor (e.g., \(H_2O(l) \leftrightarrow H_2O(g)\)) |
| Chemical Equilibrium | Rate of forward and backward reactions are equal | \(N_2 + 3H_2 \leftrightarrow 2NH_3\) |
Physical equilibrium involves phase changes like evaporation or dissolution, while chemical equilibrium occurs in reversible reactions. At equilibrium, reactant and product concentrations remain constant, though the reactions continue at equal rates. For instance, in the Haber process, the concentration of \(NH_3\) remains steady because formation and decomposition rates are equal.
Students should remember that equilibrium is dynamic, not static. It does not mean reactions have stopped; instead, they proceed at the same rate in both directions. Recognizing this helps avoid misconceptions during numerical problem-solving.
Law of Mass Action and Equilibrium Constant
Derivation and Applications
| Expression | Parameter | Description |
|---|---|---|
| \(K_c = \frac{[C]^c [D]^d}{[A]^a [B]^b}\) | Equilibrium constant in concentration | Applicable to reactions in solution phase |
| \(K_p = K_c (RT)^{\Delta n}\) | Relation between K_c and K_p | Used for gaseous reactions |
| \(Q_c = \frac{[C]^c [D]^d}{[A]^a [B]^b}\) | Reaction quotient | Compares present state to equilibrium state |
The Law of Mass Action states that at a constant temperature, the rate of a reaction is proportional to the product of concentrations of reactants raised to their respective stoichiometric powers. It allows us to define equilibrium constants K_c and K_p. These constants help determine whether a reaction favors products or reactants at equilibrium.
When Q_c < K_c, the reaction moves forward; when Q_c > K_c, it shifts backward. Understanding these relationships helps students predict equilibrium direction and evaluate chemical feasibility under varying conditions.
Le Chatelier’s Principle
Effect of Changing Conditions on Equilibrium
| Change | Effect on Equilibrium | Example |
|---|---|---|
| Concentration | Equilibrium shifts to oppose concentration change | Increasing [H_2] drives forward reaction in Haber process |
| Pressure | Favors side with fewer gas moles | Higher pressure favors NH_3 formation |
| Temperature | Depends on endo/exothermic nature | Cooling favors exothermic direction |
| Catalyst | No effect on equilibrium position | Only increases rate to reach equilibrium |
Le Chatelier’s Principle states that if a dynamic equilibrium is disturbed by changing concentration, temperature, or pressure, the system responds to counteract the disturbance. For instance, increasing temperature in an exothermic reaction shifts the equilibrium toward reactants to absorb excess heat.
This principle helps chemists optimize industrial reactions. For example, the Haber Process for ammonia synthesis applies high pressure and moderate temperature to maximize yield. Understanding such adjustments enables better control over chemical processes.
Ionic Equilibrium
Acids, Bases, and Solubility Product
| Concept | Expression | Explanation |
|---|---|---|
| Ionization Constant of Weak Acid | \(K_a = \frac{[H^+][A^-]}{[HA]}\) | Represents degree of ionization |
| Ionization Constant of Weak Base | \(K_b = \frac{[BH^+][OH^-]}{[B]}\) | Represents base strength |
| Solubility Product | \(K_{sp} = [M^{n+}]^x [A^{m-}]^y\) | Helps calculate solubility of sparingly soluble salts |
Ionic equilibrium deals with dissociation of acids, bases, and salts in aqueous solutions. Weak acids and bases ionize partially, governed by their ionization constants K_a and K_b. The smaller the value, the weaker the acid or base. Solubility products (K_{sp}) help determine whether a salt will precipitate or remain dissolved.
For instance, the solubility of AgCl in water can be determined using K_{sp}. If ionic product Q_{sp} > K_{sp}, precipitation occurs; otherwise, the solution remains unsaturated. These principles have wide applications in analytical and environmental chemistry.