Ionic Equilibrium

Electrolytes

• Strong electrolytes: Strong acids, strong bases, salts
• Weak electrolytes: Weak acids and weak bases
• Non-electrolytes: Urea, sucrose, etc.

Dissociation of Weak Electrolytes

Weak electrolytes are not completely ionised in solution.

Degree of dissociation (α):

α = (Number of molecules dissociated as ions) / (Total molecules dissolved)

Acid & Base Dissociation Constants

Acid dissociation constant (Ka):

Ka = [H₃O⁺][A⁻] / [HA]

Base dissociation constant (Kb):

Kb = [BH⁺][OH⁻] / [BOH]

Stronger the acid/base → larger the Ka / Kb value

Ostwald’s Dilution Law

α = √(K / C)

Degree of dissociation increases with dilution.

α = √(Ka / C)   ,   α = √(Kb / C)

Acid–Base Concepts

Arrhenius Concept:

• Acid → Produces H⁺ in water
• Base → Produces OH⁻ in water

Brønsted–Lowry Concept:

• Acid → Proton donor
• Base → Proton acceptor

Lewis Concept:

• Acid → Electron pair acceptor
• Base → Electron pair donor

Conjugate Acid–Base Pair

Pairs differing by one proton (H⁺).

pH, pOH & Ionic Product of Water

Kw = [H₃O⁺][OH⁻] = 1 × 10⁻¹⁴ (at 25°C)

pH = −log[H₃O⁺]

pOH = −log[OH⁻]

pH + pOH = 14

Buffer Solutions

Acidic buffer: Weak acid + salt of strong base

Basic buffer: Weak base + salt of strong acid

Henderson Equation:

pH = pKa + log([Salt] / [Acid])

pOH = pKb + log([Salt] / [Base])

Hydrolysis of Salts

Weak acid + strong base:

pH = 7 + ½ (pKa + log C)

Strong acid + weak base:

pH = 7 − ½ (pKb + log C)

Strong acid + strong base: No hydrolysis

Solubility Product (Ksp)

Ksp = [A⁺]ᵐ[B⁻]ⁿ

• Ionic product = Ksp → Saturated
• Ionic product < Ksp → Unsaturated
• Ionic product > Ksp → Precipitation

Relation Between Solubility & Ksp

For AB type:   S = √Ksp

For AB₂ type:  S = (Ksp / 4)¹ᐟ³

For AB₃ type:  S = (Ksp / 27)¹ᐟ⁴

For A₃B₂ type: S = (Ksp / 108)¹ᐟ⁵