Introduction
In chemical equilibrium we studied reaction involving molecules only but in ionic equilibrium we will study reversible reactions involving formation of ions in water. When solute is polar covalent compound then it reacts with water to form ions.
At moderate concentrations, there exists an equilibrium between the ions and undissociated molecules. This equilibrium state is called ionic equilibrium.
Electrical Conductors & Electrolytes
Substances, which allow electric current to pass through them, are known as conductors or electrical conductors.
Types of Conductors
| Type | Characteristics | Examples |
|---|---|---|
| Metallic/Electronic Conductors | Conduct electricity without undergoing any chemical change | Metals |
| Electrolytic Conductors | Undergo decomposition (chemical change) when electric current passed | Electrolytes |
Types of Electrolytes
| Type | Degree of Ionisation (α) | Examples |
|---|---|---|
| Strong | α ≈ 1 | HCl, H₂SO₄, NaCl, HNO₃, KOH, NaOH, AgNO₃, CuSO₄, all strong acids/bases/salts |
| Weak | α << 1 | CH₃COOH, NH₄OH, HCN, H₂O, Liq. SO₂, HCOOH, weak acids/bases |
Arrhenius Theory of Electrolytic Dissociation
Postulates: (i) In aqueous solution, molecules of electrolyte undergo spontaneous dissociation to form positive and negative ions. (ii) Degree of ionization (α) = Number of dissociated molecules / Total number of molecules before dissociation. (iii) At moderate concentrations, equilibrium exists between ions and undissociated molecules. (iv) Each ion behaves osmotically as a molecule.
Factors Affecting Degree of Ionisation (α)
Strong electrolytes: α ≈ 1; Weak electrolytes: α << 1
Higher dielectric constant → higher ionising power; Water has highest dielectric constant
α ∝ 1/Concentration ∝ 1/Weight of solution ∝ Dilution ∝ Amount of solvent
Degree of ionisation increases with rise in temperature
Degree of ionisation decreases in presence of strong electrolyte having common ion
Ostwald's Dilution Law
The strength of an acid or base is measured by its dissociation constant. For acetic acid:
Kₐ = [CH₃COO⁻][H₃O⁺] / [CH₃COOH]
For weak electrolytes: Kₐ = Cα²/(1-α) ≈ Cα²
α = √(Kₐ/C) = √(KₐV)
Similarly for weak base NH₄OH: α = √(K_b/C) = √(K_bV)
Ostwald's Dilution Law: For a weak electrolyte, the degree of ionisation is inversely proportional to the square root of molar concentration or directly proportional to the square root of volume containing one mole of the solute.
Dissociation Constants (K_a, K_b)
For Weak Acid HA
Kₐ = [H⁺][A⁻] / [HA]
For Polybasic Acids
Stepwise dissociation with K₁ > K₂ > K₃
H₂PO₄⁻ ⇌ H⁺ + HPO₄²⁻ (K₂)
HPO₄²⁻ ⇌ H⁺ + PO₄³⁻ (K₃)
Overall K = K₁ × K₂ × K₃
For Weak Base BOH
K_b = [NH₄⁺][OH⁻] / [NH₄OH]
Common Ion Effect
The degree of dissociation of an electrolyte (weak) is suppressed by the addition of another electrolyte (strong) containing a common ion.
Add CH₃COONa → CH₃COONa → CH₃COO⁻ + Na⁺
[CH₃COO⁻] increases, equilibrium shifts left
Applications: Qualitative analysis to adjust S²⁻ concentration in second group and OH⁻ concentration in third group.
Isohydric Solutions
Solutions with same common ion concentration show no change in degree of dissociation when mixed.
Solubility Product (K_sp)
Product of concentrations of ions raised to a power equal to the number of times the ions occur in the equation representing dissociation of electrolyte at given temperature when solution is saturated.
K_sp = [Aʸ⁺]ˣ [Bˣ⁻]ʸ
Difference: Ionic Product vs Solubility Product
| Ionic Product | Solubility Product |
|---|---|
| Applicable to all solutions (unsaturated/saturated) | Applied only to saturated solutions |
| Varies with concentration | Constant at given temperature |
| Broad meaning | Specific to saturated equilibrium |
Expressions for Different Salt Types
| Salt Type | Example | K_sp Expression | Solubility (x) |
|---|---|---|---|
| AB (1:1) | AgCl, BaSO₄ | K_sp = x² | x = √K_sp |
| AB₂ (1:2) | PbCl₂, CaF₂ | K_sp = 4x³ | x = ³√(K_sp/4) |
| A₂B (2:1) | Ag₂CrO₄ | K_sp = 4x³ | x = ³√(K_sp/4) |
| A₂B₃ (2:3) | As₂S₃, Sb₂S₃ | K_sp = 108x⁵ | x = ⁵√(K_sp/108) |
| AB₃ (1:3) | AlCl₃, Fe(OH)₃ | K_sp = 27x⁴ | x = ⁴√(K_sp/27) |
Precipitation Criteria
| Case | Condition | Result |
|---|---|---|
| I | IP < K_sp | Unsaturated, no precipitation |
| II | IP = K_sp | Saturated, no precipitation |
| III | IP > K_sp | Supersaturated, precipitation occurs |
Applications of Solubility Product
IP > K_sp → precipitation
Knowing K_sp to find solubility
HCl gas through NaCl solution precipitates pure NaCl
Concentrated NaCl precipitates soap
Group separation based on K_sp and common ion effect
After precipitation if ion in excess
Two electrolytes with common ion in same solution
Acids and Bases
Arrhenius Concept
Acids give H⁺ ions in water; Bases give OH⁻ ions in water.
Limitations: Requires water, doesn't explain non-aqueous behavior, limited neutralization concept, can't explain acidic character of salts like AlCl₃.
Bronsted-Lowry Concept
Acid: Proton donor; Base: Proton acceptor.
Acid Base Conjugate Conjugate
acid base
Conjugate pairs: HCl/Cl⁻ and H₃O⁺/H₂O
Utility: Includes ionic species, explains basic character of Na₂CO₃, NH₃, explains non-aqueous reactions.
Limitations: Can't explain reactions in non-protonic solvents, acid-base oxide reactions without proton transfer, behavior of BF₃, AlCl₃ as acids.
Lewis Concept
Acid: Electron pair acceptor; Base: Electron pair donor.
Types of Lewis Acids: Molecules with incomplete octet (BF₃, AlCl₃), all cations, molecules with empty d-orbitals (SiF₄, SnCl₄), molecules with multiple bond between atoms of dissimilar electronegativity (CO₂, SO₂).
Types of Lewis Bases: Neutral species with lone pair (NH₃, R-OH), negatively charged species/anions.
Hard and Soft Acids and Bases (HSAB)
| Type | Characteristics | Examples |
|---|---|---|
| Hard Acids | Small size, high positive charge, not easily polarized | Li⁺, Na⁺, Be²⁺, Mg²⁺, Al³⁺, BF₃, SO₃ |
| Soft Acids | Large size, low positive charge, easily polarized | Pb²⁺, Cd²⁺, Pt²⁺, Hg²⁺, I₂ |
| Hard Bases | Hold electrons strongly | OH⁻, F⁻, H₂O, NH₃, CH₃OCH₃ |
| Soft Bases | Electrons easily polarized/removed | I⁻, CO, CH₃S⁻, (CH₃)₃P |
Principle: Hard acids prefer hard bases (ionic bonding); Soft acids prefer soft bases (covalent bonding).
Relative Strength of Acids and Bases
Relative Strength Formulas
For weak bases BOH₁ and BOH₂: Strength ratio = √(K_b₁/K_b₂)
Inorganic Acids
| Type | Order | Basis |
|---|---|---|
| Hydrides | HF > H₂O > NH₃ > CH₄ HCl > H₂S > PH₃ > SiH₄ | Electronegativity |
| Hydrides | HF < HCl < HBr < HI H₂O < H₂S < H₂Se < H₂Te | Atomic size |
| Oxyacids | HOI < HOBr < HOCl HIO₄ < HBrO₄ < HClO₄ | Electronegativity |
| Oxyacids (same element) | HOCl < HClO₂ < HClO₃ < HClO₄ H₂SO₃ < H₂SO₄ HNO₂ < HNO₃ | Oxidation number |
| Oxyacids (across period) | H₄SiO₄ < H₃PO₄ < H₂SO₄ < HClO₄ | Left to right |
| Oxyacids (same oxidation state) | HNO₃ > HPO₃; H₃PO₄ > H₃AsO₄; HClO₄ > HBrO₄ > HIO₄ | Size of central atom |
Organic Acids
Compound acidic if conjugate base stabilized through resonance. Phenol acidic (C₆H₅O⁻ stabilized), ethanol neutral (C₂H₅O⁻ not stabilized).
(sp > sp² > sp³ hybridization)
Inorganic Bases
| Factor | Effect | Example |
|---|---|---|
| Electronegativity | Basicity decreases with increase | NH₃ > H₂O > HF |
| Atomic size | Larger size → lesser availability | F⁻ > Cl⁻ > Br⁻ > I⁻; O²⁻ > S²⁻ |
| Charge | Negative charge increases basicity; Positive decreases | OH⁻ > H₂O > H₃O⁺ |
| Alkali hydroxides | Increases down group | LiOH < NaOH < KOH < RbOH < CsOH |
| Alkaline earth hydroxides | Increases down group | Be(OH)₂ < Mg(OH)₂ < Ca(OH)₂ < Sr(OH)₂ < Ba(OH)₂ |
| Group 15 hydrides | Decreases down group | NH₃ > PH₃ > AsH₃ > SbH₃ > BiH₃ |
Organic Bases
Higher electron density on nitrogen → more basic. Guanidine strong base due to resonance-stabilized conjugate acid.
Acid-Base Neutralisation & Salts
Acid + Base → Salt + Water. Nature of resulting solution depends on particular acid and base.
Classification of Salts
| Type | Characteristics | Examples |
|---|---|---|
| Simple Salts | Formed by acid-base interaction | - |
| Normal Salts | No replaceable H or OH | NaCl, NaNO₃, K₂SO₄, Ca₃(PO₄)₂ |
| Acidic Salts | Incomplete neutralization of polybasic acids | NaHCO₃, NaHSO₄, NaH₂PO₄ |
| Basic Salts | Incomplete neutralization of polyacidic bases | Zn(OH)Cl, Mg(OH)Cl, Fe(OH)₂Cl |
| Double Salts | Combination of two simple salts | Ferrous ammonium sulphate, Potash alum |
| Complex Salts | Combination of simple salts/molecular compounds | K₄[Fe(CN)₆] |
| Mixed Salts | Furnishes more than one cation/anion | Ca(OCl)Cl, NaNH₄HPO₄ |
Salt Hydrolysis
Reaction of cation or anion or both ions of salt with water to produce acidic or basic solution. Reverse of neutralization.
| Salt Type | Hydrolysis | Solution | K_h Expression | h Expression | pH Expression |
|---|---|---|---|---|---|
| Weak acid + Strong base | Anionic | Basic | K_h = K_w/K_a | h = √(K_h/C) | pH = 7 + ½pK_a + ½logC |
| Strong acid + Weak base | Cationic | Acidic | K_h = K_w/K_b | h = √(K_h/C) | pH = 7 - ½pK_b - ½logC |
| Weak acid + Weak base | Both | Depends on K_a, K_b | K_h = K_w/(K_aK_b) | h = √(K_h) | pH = 7 + ½pK_a - ½pK_b |
| Strong acid + Strong base | None | Neutral | - | - | 7 |
Degree of hydrolysis: h = (moles hydrolysed) / (total moles taken)
Ionic Product of Water
K_w = [H⁺][OH⁻] = 10⁻¹⁴ at 25°C
[H⁺] = [OH⁻] = 10⁻⁷ M in pure water
ΔH = +57.3 kJ mol⁻¹ (endothermic)
K_w increases with temperature. At 25°C: Neutral [H⁺] = [OH⁻] = 10⁻⁷; Acidic [H⁺] > [OH⁻]; Basic [H⁺] < [OH⁻].
pH Scale
pOH = -log[OH⁻]
pH + pOH = pK_w = 14
| Solution | [H⁺] | [OH⁻] | pH | pOH |
|---|---|---|---|---|
| Acidic | >10⁻⁷ | <10⁻⁷ | <7 | >7 |
| Neutral | 10⁻⁷ | 10⁻⁷ | 7 | 7 |
| Basic | <10⁻⁷ | >10⁻⁷ | >7 | <7 |
pH of Common Materials
| Material | pH | Material | pH |
|---|---|---|---|
| Gastric juice | 1.4 | Rain water | 6.5 |
| Lemon juice | 2.1 | Pure water | 7.0 |
| Vinegar | 2.9 | Human saliva | 7.0 |
| Soft drinks | 3.0 | Blood plasma | 7.4 |
| Beer | 4.5 | Tears | 7.4 |
| Black coffee | 5.0 | Egg | 7.8 |
| Cow's milk | 6.5 | Household ammonia | 11.9 |
Limitations of pH Scale
- Doesn't give immediate idea of relative strengths
- pH can be zero or negative for concentrated solutions
- Very dilute acid (10⁻⁸ M) cannot have pH > 7
pK Value
For conjugate pair: Kₐ × K_b = K_w; pKₐ + pK_b = pK_w = 14
Weak acids have higher pKₐ; Weak bases have higher pK_b.
pH of Very Dilute Solutions
For 10⁻⁸ M NaOH: [OH⁻] = 10⁻⁸ + 10⁻⁷ = 1.1×10⁻⁷; pOH = 6.96; pH = 7.04
Buffer Solutions
Solution whose pH is not altered greatly by addition of small quantities of strong acid or base. Solution of reserve acidity or alkalinity that resists pH change.
Types of Buffer Solutions
| Type | Composition | Examples |
|---|---|---|
| Single Substance | Salt of weak acid and weak base | CH₃COONH₄, NH₄CN |
| Acidic Buffer | Weak acid + Salt with strong base | CH₃COOH + CH₃COONa |
| Basic Buffer | Weak base + Salt with strong acid | NH₄OH + NH₄Cl |
Buffer Action
Reserved base neutralizes added H⁺ ions; Reserved acid neutralizes added OH⁻ ions.
Examples of Buffer Systems
- Phthalic acid + Potassium hydrogen phthalate
- Citric acid + Sodium citrate
- Boric acid + Borax
- Carbonic acid + Sodium hydrogen carbonate (blood, pH 7.4)
- NaH₂PO₄ + Na₃PO₄ or Na₂HPO₄
- Glycerine + HCl
- Gastric juice (pH 1.6-1.7)
Henderson-Hasselbalch Equation
Basic buffers: pOH = pK_b + log([salt]/[base])
When [salt]/[acid] = 10: pH = pKₐ + 1; When [salt]/[acid] = 1/10: pH = pKₐ - 1
Weak acid can prepare buffers with pH range pKₐ ± 1.
pH doesn't change with dilution but varies with temperature due to K_w change.
Buffer Capacity
Maximum when [salt] = [acid] or pH = pKₐ or pOH = pK_b.
Significance of Buffer Solutions
- Colorimetric comparison of [H⁺]
- Removal of phosphate radical in qualitative analysis
- Precipitation of hydroxides in third group analysis
- Industrial applications: fermentation (pH 5-6.5), tanning, electroplating, sugar/paper manufacturing
- Bacteriological research
- Biological systems (blood pH 7.4)
Indicators
Substances used to determine end point in titration. Weak acids/bases that change color in certain pH range.
| Indicator | pH Range | Acid Color | Base Color |
|---|---|---|---|
| Cresol red | 1.2-1.8 | Red | Yellow |
| Thymol blue | 1.2-2.8 | Red | Yellow |
| Methyl yellow | 2.9-4.0 | Red | Yellow |
| Methyl orange | 3.1-4.4 | Pink | Yellow |
| Methyl red | 4.2-6.3 | Red | Yellow |
| Litmus | 5.0-8.0 | Red | Blue |
| Bromothymol blue | 6.0-7.6 | Yellow | Blue |
| Phenol red | 6.4-8.2 | Yellow | Red |
| Thymol blue | 8.1-9.6 | Yellow | Blue |
| Phenolphthalein | 8.3-10.0 | Colorless | Pink |
| Thymolphthalein | 8.3-10.5 | Colorless | Blue |
| Alizarin yellow R | 10.1-12.0 | Blue | Yellow |
| Nitramine | 10.8-13.0 | Colorless | Orange/Brown |
Theories of Indicators
Ostwald's Theory and Quinonoid Theory
Selection of Suitable Indicator
| Titration Type | Suitable Indicators |
|---|---|
| Strong acid vs Strong base | Phenolphthalein, Methyl red, Methyl orange |
| Weak acid vs Strong base | Phenolphthalein |
| Strong acid vs Weak base | Methyl red, Methyl orange |
| Weak acid vs Weak base | No suitable indicator |
Color Change Mechanism
K_in = [H⁺][In⁻] / [HIn]
[HIn]/[In⁻] = [H⁺]/K_in
Color visible when ratio [HIn]/[In⁻]:
- = 10: Acid color (pH = pK_in - 1)
- = 1: Intermediate (pH = pK_in)
- = 0.1: Base color (pH = pK_in + 1)
Effective pH range: pK_in ± 1
Important Points for JEE Main / NEET
Key Points
- Strong electrolytes: α≈1; Weak: α<<1; All salts are strong electrolytes
- Ostwald's Law: α = √(Kₐ/C) for weak acids
- Common ion effect suppresses dissociation of weak electrolytes
- Solubility Product: IP > K_sp → precipitation occurs
- Acid strength increases with oxidation number and electronegativity; decreases down group
- Hydrolysis: Weak acid + Strong base → basic solution; Strong acid + Weak base → acidic solution
- K_w = 10⁻¹⁴ at 25°C; pH + pOH = 14
- Buffers resist pH change; effective range pH = pKₐ ± 1
- Indicators change color at pH = pK_in ± 1
- pH of boiling water = 6.5625 (still neutral)
- At 37°C, pH of neutral solution = 6.8
- Buffer capacity maximum when [salt] = [acid]
- Buffers cannot withstand large amounts of acid/base (~0.1 mol/L maximum)
JEE Main & NEET Weightage
High weightage: Ionic equilibrium + acids/bases + buffers + hydrolysis + K_sp → 4-6 questions.