Unit 2

General and Physical Chemistry

Class 12 Chemistry

Ionic Equilibrium

Class 12 Chemistry – Ionic Equilibrium Notes PDF

On mobile, swipe inside the PDF to read all pages and pinch to zoom.

Introduction

DefinitionIonic equilibrium is the equilibrium established in solutions containing weak electrolytes, sparingly soluble salts or acid–base systems where ions are continuously formed and recombined.

This unit covers acid–base concepts, strength of acids and bases, conjugate pairs, Ostwald’s dilution law, ionic product of water, K_a, K_b, pK values, pH, solubility product, common-ion effect, buffers, indicators, types of salts and hydrolysis of salts.

1. Concepts of Acids and Bases

Arrhenius Concept

Acid

Produces H⁺ (more precisely H₃O⁺) in aqueous solution.

Base

Produces OH⁻ in aqueous solution.

LimitationThe Arrhenius concept is mainly restricted to aqueous systems and cannot conveniently describe all acid–base reactions, such as reactions without H⁺ or OH⁻ generation.

Brønsted–Lowry Concept

Acid: proton donor | Base: proton acceptor.

HA + B ⇌ A⁻ + BH⁺

HA/A⁻ and B/BH⁺ are conjugate acid–base pairs.

Lewis Concept

Lewis acid: electron-pair acceptor. Lewis base: electron-pair donor.

Fig. 1 – Acid–Base Concepts

2. Weak Electrolytes and Ostwald’s Dilution Law

A weak electrolyte ionizes only partially in solution. For a weak monobasic acid HA of initial concentration C and degree of ionization α:

HA ⇌ H⁺ + A⁻

K_a = Cα² / (1 − α)

For a very weak electrolyte, α is small, so:

K_a ≈ Cα² ⇒ α ≈ √(K_a/C)

MeaningDilution increases the degree of ionization of a weak electrolyte.

Fig. 2 – Ionization of a Weak Acid

3. Ionic Product of Water, K_a, K_b and pH

H₂O ⇌ H⁺ + OH⁻

K_w = [H⁺][OH⁻]

At 25°C, K_w is approximately 1.0 × 10⁻¹⁴.

pH = −log[H⁺] | pOH = −log[OH⁻]

pH + pOH = 14 (at 25°C)

pK_a = −log K_a | pK_b = −log K_b

For a conjugate acid–base pair at 25°C:

K_aK_b = K_w and pK_a + pK_b = 14

Solution	Approximate pH at 25°C	Condition
Acidic	< 7	[H⁺] > [OH⁻]
Neutral	7	[H⁺] = [OH⁻]
Basic	> 7	[OH⁻] > [H⁺]

Fig. 3 – pH Scale

pH of Strong Acid/Base

ExampleFor 0.001 M HCl, assuming complete ionization: [H⁺] = 10⁻³ M, so pH = 3.

pH of Weak Acid

[H⁺] ≈ √(K_aC)

ExampleFor a weak acid with K_a = 1.0 × 10⁻⁵ and C = 0.10 M: [H⁺] ≈ √(10⁻⁶) = 10⁻³ M, therefore pH ≈ 3.

4. Solubility Product and Common-Ion Effect

For a sparingly soluble salt AB:

AB(s) ⇌ A⁺ + B⁻

K_sp = [A⁺][B⁻]

For A₂B(s) ⇌ 2A⁺ + B²⁻:

K_sp = [A⁺]²[B²⁻]

Ionic Product and Precipitation

Condition	Meaning
Q < K_sp	Unsaturated; precipitation does not occur.
Q = K_sp	Saturated equilibrium.
Q > K_sp	Supersaturated; precipitation is favored.

Common-Ion Effect

The ionization of a weak electrolyte or the solubility of a sparingly soluble salt is suppressed when a strong electrolyte containing an ion common to that equilibrium is added.

Fig. 4 – Common-Ion Effect on Solubility

5. Buffer Solutions

DefinitionA buffer solution resists large changes in pH when small amounts of acid or base are added.

Acidic Buffer

Weak acid + salt of that acid with a strong base, e.g., CH₃COOH/CH₃COONa.

Basic Buffer

Weak base + salt of that base with a strong acid, e.g., NH₃/NH₄Cl.

pH = pK_a + log([salt]/[acid])

pOH = pK_b + log([salt]/[base])

Fig. 5 – Buffer Action

6. Acid–Base Indicators and Their Selection

An acid–base indicator is a weak organic acid or base whose acidic and basic forms have different colours.

HIn ⇌ H⁺ + In⁻

The suitable indicator is chosen so that its colour-change range overlaps the steep pH change near the titration equivalence region.

Titration	Commonly Suitable Indicator Type
Strong acid + strong base	Several indicators can work because the pH jump is large.
Strong acid + weak base	Indicator changing in acidic range is preferred.
Weak acid + strong base	Indicator changing in basic range is preferred.
Weak acid + weak base	Visual indicators are generally less suitable; instrumental methods may be preferred.

Fig. 6 – Indicator Selection on a Titration Curve

7. Types of Salts and Hydrolysis

Types of Salts

Simple Salt

Produces one type of cation and one type of anion on dissociation, e.g., NaCl.

Acidic Salt

Contains replaceable hydrogen derived from partial neutralization of a polybasic acid, e.g., NaHSO₄.

Basic Salt

Contains hydroxyl or related basic groups due to incomplete neutralization of a polyacidic base.

Complex Salt

Contains complex ions, e.g., K₄[Fe(CN)₆].

Salt Hydrolysis

Salt From	Hydrolysis	Expected Character
Strong acid + strong base	Negligible	Approximately neutral
Weak acid + strong base	Anion hydrolyzes	Basic
Weak base + strong acid	Cation hydrolyzes	Acidic

Fig. 7 – Salt Hydrolysis Summary

8. Worked Numerical Patterns

pH of Strong Base0.010 M NaOH gives [OH⁻] = 10⁻² M. Thus pOH = 2 and, at 25°C, pH = 12.

Ksp and SolubilityFor AgCl(s) ⇌ Ag⁺ + Cl⁻, if molar solubility = s, then K_sp = s². Therefore s = √K_sp.

BufferFor an acidic buffer with equal concentrations of weak acid and its salt, log([salt]/[acid]) = 0, so pH = pK_a.

Weak AcidIf K_a = 1.8 × 10⁻⁵ and C = 0.10 M, then [H⁺] ≈ √(1.8 × 10⁻⁶) ≈ 1.34 × 10⁻³ M; pH ≈ 2.87.

9. Quick Revision & Exam Points

Important Questions

State limitations of Arrhenius acid–base concept.
Define Brønsted–Lowry and Lewis acids and bases with examples.
Explain conjugate acid–base pairs.
Derive Ostwald’s dilution law for a weak electrolyte.
Define K_w, K_a, K_b, pK_a and pK_b.
Calculate pH of strong and weak acids/bases.
Define solubility product and explain precipitation condition.
Explain common-ion effect.
Define buffer and explain its action.
Explain selection of acid–base indicators.
Classify salts and explain salt hydrolysis.

Brønsted acid donates H⁺.
Brønsted base accepts H⁺.
Lewis acid accepts electron pair.
Weak electrolyte ionization increases on dilution.
K_w = [H⁺][OH⁻].
pH = −log[H⁺].
K_sp describes sparingly soluble salt equilibrium.
Common ion lowers ionization/solubility.
Buffers resist pH change.
Salt hydrolysis depends on parent acid/base strength.

Source handling: The original Nepal eNotes PDF linked to this chapter is embedded above. The typed section follows the verified NEB Grade 12 Chemistry syllabus and is written as a searchable, responsive study companion. Because the Drive viewer does not expose the handwritten page text as readable document text here, this section is not presented as a word-for-word transcription.