Chapter 6 – Equilibrium (JEE Chemistry)
1. Introduction to Equilibrium
Equilibrium refers to a state where **two opposite processes occur at equal rates**.
In chemistry, equilibrium is studied mainly in:
- Chemical reactions
- Ionic reactions in solution
2. Types of Equilibrium
- Physical Equilibrium – Phase equilibrium (ice ⇌ water)
- Chemical Equilibrium – Reversible chemical reactions
3. Reversible Reactions
A reaction that can proceed in both forward and backward directions is called a reversible reaction.
$N_2(g) + 3H_2(g) \rightleftharpoons 2NH_3(g)$
4. Law of Mass Action
According to the law of mass action, at constant temperature,
the rate of a chemical reaction is proportional to the product of the active masses of reactants.
5. Equilibrium Constant (Kc)
For reaction:
$aA + bB \rightleftharpoons cC + dD$
$$K_c = \frac{[C]^c[D]^d}{[A]^a[B]^b}$$
Only gaseous or aqueous species are included.
Solids and liquids are omitted.
6. Characteristics of Equilibrium Constant
- Depends only on temperature
- Independent of initial concentrations
- Large K → reaction favours products
- Small K → reaction favours reactants
7. Equilibrium Constant in Terms of Pressure (Kp)
$$K_p = K_c(RT)^{\Delta n}$$
$\Delta n =$ moles of gaseous products − moles of gaseous reactants
8. Reaction Quotient (Q)
Reaction quotient is calculated using the same expression as K but at any moment.
- Q < K → reaction proceeds forward
- Q > K → reaction proceeds backward
- Q = K → system at equilibrium
9. Le Chatelier’s Principle
If a system at equilibrium is subjected to a change,
the equilibrium shifts to counteract the effect of the change.
10. Effect of Concentration
Increasing concentration of reactants shifts equilibrium forward.
11. Effect of Pressure
Increase in pressure shifts equilibrium towards the side with fewer moles of gas.
12. Effect of Temperature
- Endothermic reaction → increase temperature favours forward reaction
- Exothermic reaction → increase temperature favours backward reaction
13. Effect of Catalyst
Catalyst does not change equilibrium position.
It only increases the rate of reaction.
14. Ionic Equilibrium
Equilibrium involving ions in aqueous solution is called ionic equilibrium.
15. Acids and Bases
- Arrhenius concept
- Bronsted–Lowry concept
$HA \rightleftharpoons H^+ + A^-$
16. Ionization Constant of Acid (Ka)
$$K_a = \frac{[H^+][A^-]}{[HA]}$$
17. pH Scale
$$pH = -\log[H^+]$$
Lower pH → stronger acid
Higher pH → stronger base
18. Buffer Solutions
A buffer solution resists change in pH on addition of small amount of acid or base.
$$pH = pK_a + \log\frac{[\text{Salt}]}{[\text{Acid}]}$$
19. Hydrolysis of Salts
Salt of weak acid and strong base produces basic solution.
20. Solubility Product (Ksp)
For $AgCl(s) \rightleftharpoons Ag^+ + Cl^-$
$$K_{sp} = [Ag^+][Cl^-]$$
If ionic product > Ksp → precipitation occurs.
21. Final Revision Checklist
You have mastered equilibrium if you can:
- Calculate Kc, Kp, Q
- Predict equilibrium shift
- Solve pH and buffer problems
- Apply Ksp correctly