Equilibrium Quick Revision | CBSE Class 11 Chemistry | Full Chapter in 1️⃣5️⃣ Mins | Rapid Revision

Key Concepts:

• Equilibrium (Static & Dynamic)
• Physical Equilibrium (Solid-Liquid, Liquid-Vapor, Solid-Vapor, Solid-Solution)
• Chemical Equilibrium (Homogeneous & Heterogeneous)
• Reversible & Irreversible Reactions
• Equilibrium Constant (Kc, Kp)
• Relationship between Kp and Kc
• Characteristics of Equilibrium Constant
• Reaction Quotient (Qc)
• Gibbs Free Energy and Equilibrium Constant Relationship
• Le Chatelier's Principle
• Electrolytes (Strong & Weak)
• Degree of Dissociation (α)
• Arrhenius, Bronsted-Lowry, and Lewis Acid-Base Theories
• Self-ionization of Water (Kw)
• pKa, pKb, pH, pOH
• Polyprotic Acids
• Common Ion Effect
• Buffer Solutions (Acidic, Basic, Salt)
• Solubility and Solubility Product (Ksp)

1. Introduction to Equilibrium:

• Equilibrium is defined as a state in a process where properties like temperature, pressure, and concentration remain constant over time.
• The discussion focuses on dynamic equilibrium, where forward and backward reactions occur at equal rates, maintaining a constant state despite ongoing changes.
• Example: A tap filling a bucket with a hole, where the inflow equals the outflow, is used to illustrate dynamic equilibrium.

2. Types of Equilibrium:

• Physical Equilibrium: Involves changes in physical properties.
  • Solid-Liquid Equilibrium: Rate of melting equals the rate of freezing.
  • Liquid-Vapor Equilibrium: Rate of evaporation equals the rate of condensation.
  • Solid-Vapor Equilibrium: Rate of sublimation equals the rate of condensation.
  • Solid-Solution Equilibrium: Rate of dissolution equals the rate of crystallization; a saturated sugar solution is given as an example.
• Chemical Equilibrium: Classified into homogeneous (same phase for reactants and products) and heterogeneous (different phases).
  • Homogeneous Example: N2(g) + H2(g) ⇌ NH3(g)
  • Heterogeneous Example: H2O(g) ⇌ H2O(l)

3. Reversible and Irreversible Reactions:

• Equilibrium is observed in reversible reactions, where reactants form products and products revert to reactants.
• Irreversible reactions proceed in one direction only and do not establish equilibrium.

4. Graphical Representation of Equilibrium:

• Graphs illustrate the change in concentration of reactants and products over time, converging at equilibrium.
• Example: N2 + H2 forming NH3 (Haber process) shows reactants decreasing and products increasing until equilibrium is reached.

5. Equilibrium Constant (Kc and Kp):

• The equilibrium constant (Kc) is the ratio of product concentrations to reactant concentrations, each raised to the power of their stoichiometric coefficients.
• Kc = [C]^c [D]^d / [A]^a [B]^b, where a, b, c, and d are stoichiometric coefficients.
• Kp is the equilibrium constant expressed in terms of partial pressures.

6. Relationship Between Kp and Kc:

• Kp = Kc(RT)^Δn, where R is the universal gas constant, T is temperature in Kelvin, and Δn is the change in the number of moles of gaseous products minus gaseous reactants.
• If Δn = 0, Kp = Kc. If Δn < 0, Kp < Kc. If Δn > 0, Kp > Kc.

7. Characteristics of Equilibrium Constant:

• Equilibrium constant is independent of initial concentrations.
• Reversing the reaction inverts the value of Kc.
• Multiplying the reaction by a factor raises Kc to that power.
• Kc has a definite value for each chemical reaction at a particular temperature.

8. Applications of Equilibrium Constant:

• The reaction quotient (Qc) is used to predict the direction of a reaction.
• Qc is calculated using the same formula as Kc but with non-equilibrium concentrations.
• If Qc = Kc, the reaction is at equilibrium. If Qc < Kc, the reaction proceeds in the forward direction. If Qc > Kc, the reaction proceeds in the reverse direction.

9. Gibbs Free Energy and Equilibrium:

• The relationship between Gibbs free energy (ΔG°) and the equilibrium constant (K) is given by: ΔG° = -RT ln K or ΔG° = -2.303RT log K.
• If K > 1, ΔG° < 0, the reaction is spontaneous and favors product formation.
• If K < 1, ΔG° > 0, the reaction is non-spontaneous and favors reactant formation.
• If K = 1, ΔG° = 0, the reaction is at equilibrium.

10. Le Chatelier's Principle:

• States that if a change of condition (e.g., concentration, pressure, temperature) is applied to a system in equilibrium, the system will shift in a direction that relieves the stress.
• Changes in concentration, pressure, temperature, and addition of inert gases can affect equilibrium. Catalysts do not shift equilibrium but increase the rate of reaction.

11. Ionic Equilibrium and Electrolytes:

• Electrolytes are substances that dissociate into ions in solution.
• Strong electrolytes dissociate completely (α = 1), while weak electrolytes dissociate partially (α < 1).
• Example: Acetic acid (CH3COOH) is a weak electrolyte.

12. Acid-Base Theories:

• Arrhenius Theory: Acids produce H+ ions, and bases produce OH- ions in water.
• Bronsted-Lowry Theory: Acids are proton (H+) donors, and bases are proton acceptors.
• Lewis Theory: Acids are electron-pair acceptors, and bases are electron-pair donors.

13. Self-ionization of Water:

• Water can act as both an acid and a base, undergoing self-ionization: H2O + H2O ⇌ H3O+ + OH-.
• The ion product of water (Kw) is [H+][OH-] = 1.0 x 10^-14 at 25°C. Kw is temperature dependent.

14. pH and pOH:

• pH = -log[H+], pOH = -log[OH-].
• pH + pOH = 14.
• pH scale ranges from 0 to 14, with 7 being neutral, below 7 acidic, and above 7 basic.

15. Polyprotic Acids:

• Acids with more than one ionizable hydrogen atom.
• Example: H2SO4 dissociates in two steps, each with its own acid dissociation constant (Ka).

16. Common Ion Effect:

• The decrease in the solubility of a sparingly soluble salt when a soluble salt containing a common ion is added.
• Example: Adding NH4Cl to a solution of NH3/H2O shifts the equilibrium due to the common NH4+ ion.

17. Buffer Solutions:

• Solutions that resist changes in pH upon addition of small amounts of acid or base.
• Types: Acidic buffers, basic buffers, and salt buffers.

18. Solubility and Solubility Product (Ksp):

• Solubility is the extent to which a solute dissolves in a solvent.
• Solubility product (Ksp) is the equilibrium constant for the dissolution of a sparingly soluble salt.
• Example: For a salt AB, Ksp = [A+][B+].

19. Conclusion:

• The video provides a rapid revision of chemical and ionic equilibrium, covering key concepts, formulas, and applications. It emphasizes the importance of understanding equilibrium constants, Le Chatelier's principle, and acid-base theories. The speaker recommends referring to the detailed one-shot video for more in-depth explanations and problem-solving.