This lesson explains how acids, bases, and salts behave in aqueous solutions, how to calculate pH, how buffer solutions resist pH change, and how sparingly soluble salts dissolve.

1. Core Concepts (Short Notes)

17.1 Strong vs. Weak Acids/Bases

• Strong acids/bases: Fully ionize in water (HCl, HNO₃, NaOH).

• Weak acids/bases: Partially ionize (CH₃COOH, NH₃).

Strength depends on degree of dissociation, not concentration.

17.2 pH and pOH

• pH = −log[H⁺]

• pOH = −log[OH⁻]

• pH + pOH = 14 (at 25°C)

pH scale:

• 0–7: acidic

• 7: neutral

• 7–14: basic

17.3 Ionic Product of Water (Kw)

Kw = [H⁺][OH⁻] = 1 × 10⁻¹⁴ at 25°C

17.4 Acid Dissociation Constant (Ka)

For weak acid HA: HA ⇌ H⁺ + A⁻ Ka = [H⁺][A⁻] / [HA]

17.5 Base Dissociation Constant (Kb)

For weak base B: B + H₂O ⇌ BH⁺ + OH⁻ Kb = [BH⁺][OH⁻] / [B]

17.6 Relationship Between Ka, Kb, and Kw

Ka × Kb = Kw

17.7 Indicators

Indicators change color depending on pH. Examples:

• Methyl orange: red → yellow

• Phenolphthalein: colorless → pink

Each indicator has a pH range where color changes.

17.8 Buffer Solutions

Buffers resist changes in pH. Two types:

• Acidic buffer: weak acid + salt of acid (e.g., CH₃COOH + CH₃COONa)

• Basic buffer: weak base + salt (e.g., NH₃ + NH₄Cl)

17.9 Henderson–Hasselbalch Equation

For acidic buffer: pH = pKa + log ([salt] / [acid])

For basic buffer: pOH = pKb + log ([salt] / [base])

17.10 Solubility Product Constant (Ksp)

For a sparingly soluble salt: AB(s) ⇌ A⁺ + B⁻ Ksp = [A⁺][B⁻]

If ionic product > Ksp → precipitate forms.

2. Key Formulas to Memorize

• pH = −log[H⁺]

• pOH = −log[OH⁻]

• Kw = [H⁺][OH⁻] = 1×10⁻¹⁴

• Ka = [H⁺][A⁻] / [HA]

• Kb = [BH⁺][OH⁻] / [B]

• Ka × Kb = Kw

• Ksp = product of ion concentrations

• Henderson–Hasselbalch: pH = pKa + log (salt/acid)

3. Tips & Tricks for Exams

• Weak acid pH calculations often assume: [H⁺] ≈ √(Ka × concentration).

• For buffers: If salt = acid, then pH = pKa.

• Temperature affects Kw; pH 7 may not always be neutral.

• Indicators must be chosen based on the pH range of titration.

• Ksp problems require writing dissolution equations carefully.

• If Q > Ksp, precipitation occurs.

• Polyprotic acids dissociate stepwise; consider Ka₁ > Ka₂.

4. Important Points to Remember

• pH measures H⁺ concentration logarithmically.

• Buffers maintain pH in biological and industrial systems.

• Ksp determines solubility; low Ksp = very insoluble salt.

• Indicators show different colors in acidic and basic environments.

• Ka and Kb measure acid/base strength, not concentration.

Ionic equilibrium links directly to titration and solubility problems.