Ch 04: Chemical Bonds & Reactions

⚗️ Chapter At a Glance

📘 Chemical Bond

Bond formation can lower energy. The octet rule describes many main-group compounds, with exceptions such as the hydrogen duet and electron-deficient species.

Major Bond Types

Types of Chemical Reactions

Berzelius

Coined “catalysis”

Fe + promoters

Haber's Process Catalyst

📊 Bond Types — Quick Comparison

Bond Type	Formation	Example	Water Solubility	Conductor?
Ionic	Electron Transfer	NaCl, MgO	Often soluble; not universal	Bad (solid); Good (aqueous)
Covalent	Electron Sharing	O₂, H₂O	Variable with molecular polarity	Poor (exception: Graphite)
Hydrogen	N/O/F with H	H₂O, NH₃, DNA	—	—
Coordinate	Sharing by same atom	NH₄⁺, H₃O⁺	—	—

⚡ Ionic Bond (Electrovalent)

📘 Formation

By electron transfer from one atom to another. One atom loses e⁻ (becomes cation +) and other gains e⁻ (becomes anion −).

Na → gives 1 e⁻ → Na⁺; Cl → takes 1 e⁻ → Cl⁻ → NaCl
Mg → gives 2 e⁻; O → takes 2 e⁻ → MgO

Properties

Many ionic compounds dissolve in water; exceptions depend on lattice and hydration energies
High melting point
Solid state → Bad conductor
Aqueous/molten state → Good conductor
Ionic nature

🔗 Covalent Bond

📘 Formation

By sharing of electron pairs between atoms (usually non-metals).

O₂ → Non-polar covalent (same atoms)
H₂O → Polar covalent (different atoms)

Properties

Solubility depends on polarity and intermolecular forces; many non-polar compounds are poorly soluble
Soluble in organic solvents
Low melting & boiling point (Exception: Diamond)
Poor conductor (Exception: Graphite)

💧 Hydrogen Bond

📘 Formation

An attraction involving H already covalently bonded to an electronegative atom and another electronegative atom with an available electron pair

Intermolecular H-Bond

Between different molecules. Examples: H₂O (water), NH₃ (ammonia). Responsible for high BP of water.

Intramolecular H-Bond

Within one molecule; standard examples: o-nitrophenol and salicylaldehyde. DNA base pairing between strands is intermolecular.

🔀 Coordinate Covalent Bond

📘 Formation

Sharing of an electron pair where both electrons come from the same atom (donor). Also called dative bond.

Examples of coordinate-bond formation: NH₄⁺ and H₃O⁺
Found in Hydronium ion (H₃O⁺)
One atom donates both electrons of the shared pair

🌐 Van der Waals' Force

Weak force of attraction or repulsion between molecules
Different from covalent and ionic bonds
Weaker than chemical bonds
Responsible for properties of noble gases in liquid/solid state

⚠️ Exam Trap — Bond Properties

Ionic bond: Solid = bad conductor; Aqueous = good conductor
Covalent bond exceptions: Diamond (high MP), Graphite (good conductor)
Strong hydrogen bonds most commonly involve N, O and F at this introductory level
DNA double helix held by Hydrogen bonds
Water's high boiling point due to intermolecular H-bonds

⚡ Sigma (σ) and Pi (π) Bonds

—

Single Bond

1 Sigma (σ) bond

H–H (H₂)

Double Bond

1 σ + 1 π bond

O=O (O₂)

≡

Triple Bond

1 σ + 2 π bonds

N≡N (N₂)

Bond	Composition	Example	π bonds per indicated multiple bond
Single (–)	1 σ only	H₂, CH₄, H₂O	0
Double (=)	1 σ + 1 π	O₂, CO₂, C₂H₄	1
Triple (≡)	1 σ + 2 π	N₂, C₂H₂	2

⚠️ Exam Trap — Sigma & Pi

Each ordinary single, double or triple covalent link contains one σ bond; extra components are π bonds
Pi bonds only present in double and triple bonds
Triple bond = 1σ + 2π → 2 Pi bonds (NOT 3)
N₂ molecule has a triple bond → 1σ + 2π → very stable

🧬 Molecular Shapes & Structures

H₂O

Water

Angular / V-shaped

Polar Covalent Bond

NH₃

Ammonia

Pyramidal (Trigonal Pyramidal)

Polar Covalent Bond

CH₄

Methane

Tetrahedral

Non-polar Covalent Bond

CO₂

Carbon Dioxide

Linear

Non-polar Covalent Bond

⚠️ Exam Trap — Molecular Shapes

H₂O = Angular/V-shaped (NOT linear; O has 2 lone pairs)
NH₃ = Pyramidal (NOT tetrahedral; N has 1 lone pair)
CH₄ = Tetrahedral (no lone pairs on C)
CO₂ is linear: central carbon has no lone pairs; oxygen atoms do have lone pairs
H₂O is polar; CO₂ is non-polar even though it has polar bonds (symmetric)

🔢 Valency

Valency = combining ability of an atom (no. of electrons used/shared/transferred)
Fluorine: valency 1; oxidation state −1 in its compounds and 0 in F₂
IA and IIA groups: valency = group number (Frankland's rule)
Cl (2,8,7) and Na (2,8,1): usual valency 1; chloride charge is −1, sodium-ion charge is +1
Fe, Cu, S, Hg, Sn — have variable valency (more than one valency)

Cross-Multiply Valencies → SubscriptsAlCl₃: Al(+3), Cl(−1) → Al₁Cl₃ → AlCl₃

🔢 Oxidation Numbers

📘 Definition

Oxidation state is a formal electron-bookkeeping charge, not necessarily an actual ionic charge. Fractional values can be averages in mixed-valence systems.

Element / Species	Oxidation Number	Notes
H (Hydrogen)	+1	−1 in metal hydrides (NaH)
O (Oxygen)	−2	−1 in peroxides (H₂O₂); +2 in OF₂
F (Fluorine)	−1	−1 in compounds; 0 in elemental F₂
Neutral molecule	0	Sum of oxidation numbers = 0
Na, K	+1	Alkali metals: +1 in ordinary compounds, 0 as elements
Mg, Ca, Sr	+2	Alkaline-earth metals: +2 in ordinary compounds, 0 as elements

⚠️ Exam Trap — Oxidation Numbers

O = −2 (usually); −1 in H₂O₂; 0 in O₂
H = +1 (usually); −1 in metal hydrides (NaH, CaH₂)
F has oxidation state −1 in compounds and 0 in elemental F₂
In neutral molecule, sum of all oxidation numbers = 0
In ions, sum = charge on the ion

⚗️ Types of Chemical Reactions

① Combination Reaction

A + B → C

2H₂ + O₂ → 2H₂O (Water formation)

② Decomposition Reaction

AB → A + B

CaCO₃ →(heat)→ CaO + CO₂↑

③ Displacement Reaction

AB + C → AC + B

Fe + CuSO₄ → FeSO₄ + Cu (Fe more reactive than Cu)

④ Double Displacement

AB + CD → AD + CB

Na₂SO₄ + BaCl₂ → BaSO₄↓ + 2NaCl (ion exchange)

⑤ Exothermic Reaction

Reactants → Products + Energy (heat released)

CH₄ + 2O₂ → CO₂ + 2H₂O + Heat | Also: Combustion, Respiration

⑥ Endothermic Reaction

Reactants + Energy → Products (heat absorbed)

Thermal decomposition of CaCO₃ absorbs heat; photosynthesis requires energy input

⑦ Reversible Reaction

A + B ⇌ AB

Proceeds in both forward and backward directions; reaches equilibrium

⑧ Irreversible Reaction

A + B → AB

Proceeds only in one direction; goes to completion

⑨ Redox Reaction

Oxidation + Reduction occur simultaneously

Fe + CuSO₄ → FeSO₄ + Cu (Fe oxidised → Fe²⁺; Cu²⁺ reduced → Cu)

⚠️ Exam Trap — Reactions

Exothermic = energy released; Endothermic = energy absorbed
Respiration = exothermic (releases energy); Photosynthesis = endothermic
Displacement: more reactive metal displaces less reactive from solution
Double displacement = ion exchange → often gives precipitate (↓)
CaCO₃ → CaO + CO₂ = Decomposition (NOT combination)

🔄 Oxidation & Reduction

🔴 Oxidation

Combination with O or electronegative element (F, Cl, N)
Loss of H or electropositive element
Loss of electrons (e⁻)
Oxidation number Increases ↑

e⁻ → away ↑ ON

🟢 Reduction

Combination with H or electropositive element (Na, Mg)
Loss of O or electronegative element
Gain of electrons (e⁻)
Oxidation number Decreases ↓

e⁻ → gained ↓ ON

Oxidising Agents (Oxidisers)

Substances that cause oxidation (get reduced themselves)

H₂O₂ (Hydrogen peroxide)
O₂ (Oxygen)
F₂ (Fluorine)
KMnO₄ (Potassium permanganate)

Reducing Agents (Reducers)

Substances that cause reduction (get oxidised themselves)

Li (Lithium)
Na (Sodium)
NaBH₄ (Sodium borohydride)
H₂ (Hydrogen gas)

📘 Memory Trick — OIL RIG

Oxidation Is Loss (of electrons) | Reduction Is Gain (of electrons)

⚠️ Exam Trap — Redox

Oxidising agent gets reduced; Reducing agent gets oxidised
Oxidation = loss of e⁻ = ON increases | Reduction = gain of e⁻ = ON decreases
In Fe + CuSO₄: Fe is oxidised (0 → +2); Cu²⁺ is reduced (+2 → 0)
H₂O₂ can be both oxidising AND reducing agent (dual redox behaviour; not an acid–base amphoterism label)

⚙️ Catalyst

📘 Definition

A catalyst increases reaction rate through a catalytic pathway and is regenerated overall. Berzelius coined “catalysis” in the 1830s.

Homogeneous Catalysis

Catalyst and reacting system occupy a single phase

2SO₂(g) + O₂(g) —NO(g)→ 2SO₃(g)Both reactants and catalyst in gaseous state

Heterogeneous Catalysis

Catalysis involves more than one phase, typically at an interface

N₂(g) + 3H₂(g) →(Fe solid)→ 2NH₃(g)Gaseous reactants, solid Fe catalyst — Haber's Process

Positive Catalyst

Increases reaction rate

Inhibitor (traditional “negative catalyst”)

Decreases reaction rate

Catalysis terminology

Berzelius

📋 Important Processes & Their Catalysts

Process	Catalyst / enzyme / microorganism
Ghee from vegetable oil (hydrogenation)	Nickel (Ni)
Contact method — Sulphuric acid (H₂SO₄)	V₂O₅ in modern Contact process; Pt used historically
Deacon process — Chlorine (Cl₂) formation	Cupric Chloride (CuCl₂)
Small peptides → amino acids (intestinal digestion)	Intestinal peptidases (historically “erepsin”)
Vinegar: oxidation of ethanol to acetic acid	Acetobacter aceti (microorganism)
Conversion of milk into curd	Lactic-acid bacteria such as Lactobacillus; lactose is the substrate
Haber's Process — Ammonia (NH₃)	Iron-based catalyst; potassium compounds as promoter, with Al₂O₃ in conventional formulations

⚠️ Exam Trap — Catalysts

Haber–Bosch: N₂ + 3H₂ ⇌ 2NH₃; iron catalyst, commonly potassium-compound promoter
Ghee from oil: catalyst = Nickel (Ni)
H₂SO₄ (Contact method): catalyst = V₂O₅ (most asked)
Milk-to-curd fermentation: lactic-acid bacteria; their enzymes catalyse the reactions
Berzelius coined catalysis; enzymes are biological catalysts

🎯 High-Yield Exam Points

Chemical bonding lowers energy; the octet is a useful main-group rule with important exceptions
Ionic bonding: electrostatic attraction of ions formed by electron transfer; water solubility varies; mobile ions conduct in a melt or suitable solution
Covalent bond: electron-pair sharing. Water solubility varies; many molecular compounds conduct poorly. Graphite is a conducting network solid.
Diamond = covalent bond but high melting point (exception)
Graphite = covalent but good conductor (exception)
H-bond: N, O, F with H; DNA held by H-bonds
Coordinate-bond formation examples: NH₄⁺ and H₃O⁺
Single bond = 1σ | Double = 1σ+1π | Triple = 1σ+2π
H₂O = Angular (V-shape) | NH₃ = Pyramidal | CH₄ = Tetrahedral | CO₂ = Linear
F valency = 1. Oxidation states in compounds: F = −1; usually O = −2 and H = +1
Combination: A+B→C | Decomposition: AB→A+B | Displacement: more reactive displaces less
Respiration = Exothermic | Photosynthesis = Endothermic
Oxidation = loss of e⁻, ON increases | Reduction = gain of e⁻, ON decreases
OIL RIG: Oxidation Is Loss | Reduction Is Gain
Berzelius: coined “catalysis” in the 1830s
Haber–Bosch: N₂ + 3H₂ ⇌ 2NH₃; iron catalyst with potassium-based promoters
H₂SO₄ (Contact): catalyst = V₂O₅
Ghee from oil: catalyst = Nickel (Ni)
Curd fermentation: lactic-acid bacteria such as Lactobacillus

📋 Quick Reference

Bond/Reaction	Key Fact
Ionic (Electrovalent)	e⁻ transfer; NaCl
Covalent	e⁻ sharing; O₂, H₂O
H-Bond	N/O/F with H; DNA
Coordinate	Same atom donates; H₃O⁺
Single bond	1σ only
Double bond	1σ + 1π
Triple bond	1σ + 2π
Exothermic	Energy released
Endothermic	Energy absorbed
Oxidation	Loss of e⁻; ON ↑
Reduction	Gain of e⁻; ON ↓

⚠️ Most Common Exam Traps

Ionic solid = bad conductor; ionic solution = good conductor
Diamond = covalent bond but hard & high MP
Graphite = covalent bond but conducts electricity
H₂O = V-shaped (NOT linear)
CO₂ = linear, non-polar (NOT bent like H₂O)
Triple bond = 1σ + 2π (NOT 3 sigma)
Oxidation states: F = −1 in compounds, 0 in F₂; O usually −2, but −1 in H₂O₂
Respiration = exothermic (NOT endothermic)
Reducing agent gets oxidised; oxidising agent gets reduced
Haber's = Fe catalyst (NOT V₂O₅); Contact (H₂SO₄) = V₂O₅

⚗️ Chemical Bonds & Reactions