Ionic Equilibrium - Lecture Notes

Introduction

• Important chapter: Ionic Equilibrium
• Comparison to Chemical Equilibrium

Concepts Covered

• Understanding one concept at a time, similar to chemical equilibrium
• Important for exams like JEE Mains and JEE Advanced
• Multiple concepts related to acid and base

Electrolytes and Non-Electrolytes

• Electrolytes: Conduct electricity in molten/aquatic states
  • Strong Electrolytes: Fully ionize (e.g., strong acids, bases, salts)
  • Weak Electrolytes: Partially ionize (e.g., weak acids, weak bases, sparingly soluble salts)
• Non-Electrolytes: Do not conduct electricity (e.g., glucose, urea)

Ionic Equilibrium

• Fundamentals of Ionic Equilibrium: Establishment of dynamic equilibrium between ions and un-ionized molecules in weak electrolytes
• Notations: Use of standard notation for dissociation constants and concentrations
• Degree of Dissociation (α)
• Acid-Base Theories: Arrhenius, Bronsted-Lowry, and Lewis
  • Arrhenius: Acid produces H+, Base produces OH-
  • Bronsted-Lowry: Acid donates proton (H+), Base accepts proton
  • Lewis: Acid accepts electron pair, Base donates electron pair
    • Examples: NH₃ + BF₃ → NH₃BF₃

Water's Ionic Product (KW)

• Self-ionization of water: 2H₂O ↔ H₃O⁺ + OH⁻
• KW at 25°C: 10⁻¹⁴, changes with temperature

Degree of Ionization & Relationship with Ionic Product

• Calculating degree of ionization
• Understanding weak electrolytes’ limited ionization
• Relation of KW with temperature variations
• PH Calculation: Using -log [H⁺]
• Notions of Pure Water, acidic, and basic solutions

Buffer Solutions

• Definition: Solutions that resist changes in pH upon the addition of small amounts of acid/base
• Types: Acidic Buffers, Basic Buffers
  • Acidic Buffers: Weak acid + salt of weak acid (e.g., CH₃COOH + CH₃COONa)
  • Basic Buffers: Weak base + salt of weak base (e.g., NH₄OH + NH₄Cl)
• Buffer Capacity: Ability of buffer to neutralize additions of acid/base
• Buffer Range: Effective pH range of buffer
• Henderson-Hasselbalch Equation: Used for calculating the pH of buffer solutions

Solubility Product (KSP)

• Definition: Product of ion concentrations of saturated solution raised to power of their stoichiometric coefficients
• Calculations: Relationship between solubility and KSP for different salts
  • Examples: Ca(OH)₂, AG₂CO₃
• Common Ion Effect on Solubility: Adding common ions decreases solubility
• Ionic Product (IP): Determines if solution is saturated, undersaturated, or leads to precipitation
• Comparisons: IP < KSP (undersaturated), IP = KSP (saturated), IP > KSP (precipitation)
• Complexation: Formation of complex ions increases solubility (e.g., NH₃ with Ag⁺)

Conclusion

• The chapter extensively covers ionic equilibrium concepts
• Importance in solving numerical problems for competitive exams
• Summary of key points and preparation for advanced topics ahead

Exercises and Practice Questions

• Examples provided during lecture for better understanding
• Homework problems for practice
• Emphasis on understanding and applying concepts