Chemical Bonding B.Tech. 1st Year

Chapter 1: Chemical Bonding

Complete Detailed Theory with Examples, Diagrams & Formulas

1. Introduction to Chemical Bonding

Chemical bonding is the force of attraction between atoms that enables the formation of chemical compounds. It occurs due to the tendency of atoms to attain stable electronic configuration (octet rule).

Reason for bonding: To achieve noble gas configuration (8 electrons in valence shell).

Example: Na (2,8,1) loses 1 electron → Na⁺ (2,8)
Cl (2,8,7) gains 1 electron → Cl⁻ (2,8,8)

2. Types of Chemical Bonds

• Ionic Bond: Transfer of electrons → Electrostatic attraction between cations and anions.
• Covalent Bond: Sharing of electrons (single, double, triple).
• Coordinate Bond: Both electrons from one atom (donor → acceptor).
• Metallic Bond: Delocalized electrons in metal lattice.
• Hydrogen Bond: Between H and highly electronegative atom (F, O, N).

3. Lewis Dot Structures & Formal Charge

Representation of valence electrons as dots.

Formal Charge = Valence e⁻ − (Non-bonding e⁻ + ½ Bonding e⁻)

CO₂: O=C=O (carbon has 2 double bonds, formal charge = 0)
NH₄⁺: 4 bonds on N → Formal charge on N = +1

4. VSEPR Theory – Molecular Geometry

Valence Shell Electron Pair Repulsion Theory: Electron pairs repel each other and stay as far apart as possible.

Electron Pairs	Bond Pairs	Lone Pairs	Geometry	Example
2	2	0	Linear	BeCl₂, CO₂
3	3	0	Trigonal Planar	BF₃
4	4	0	Tetrahedral	CH₄
4	3	1	Trigonal Pyramidal	NH₃
4	2	2	Bent/V-shaped	H₂O

5. Valence Bond Theory (VBT) – Hybridization

Mixing of atomic orbitals to form new hybrid orbitals.

sp → 2 orbitals (linear, 180°)
sp² → 3 orbitals (trigonal, 120°)
sp³ → 4 orbitals (tetrahedral, 109.5°)

CH₄: Carbon uses sp³ hybridization
C₂H₄: Carbon uses sp² (one p orbital for π bond)
C₂H₂: Carbon uses sp (two p orbitals for two π bonds)

6. Molecular Orbital Theory (MOT)

Atomic orbitals combine to form molecular orbitals.

• Bonding MO (lower energy)
• Antibonding MO (higher energy)

Bond Order = \frac{1}{2} (N_b - N_a)

O₂: (σ1s)²(σ*1s)²(σ2s)²(σ*2s)²(σ2p_z)²(π2p_x)²(π2p_y)²(π*2p_x)¹(π*2p_y)¹
→ Bond Order = ½(10-6) = 2 (Paramagnetic)

7. Intermolecular Forces

• Van der Waals: London dispersion, dipole-induced dipole
• Dipole-Dipole: Between polar molecules
• Hydrogen Bonding: F-H…F, O-H…O, N-H…N (strongest)

H-bonding explains high boiling point of H₂O, NH₃, HF.

8. Fajans' Rules – Polarization

Small cation + large anion + high charge → More covalent character.

• High charge on cation → More polarizing power
• Large anion → More polarizable
• Pseudo-inert gas configuration (d¹⁰) → More polarization (e.g., Cu⁺, Ag⁺)

AlCl₃ → Covalent (small Al³⁺, high charge)
NaCl → Ionic (large Na⁺, low charge)

9. Applications in Engineering Materials

• Ionic → Ceramics, glass
• Covalent → Diamond, SiC (hardness)
• Metallic → Conductors, ductility
• H-bonding → Polymers, adhesives, DNA structure

© 2025 Engineering Chemistry | Detailed Notes by Experts | For B.Tech First Year

Engineering Chemistry
Chemical Bonding - Comprehensive Notes

1. Introduction to Chemical Bonding

Chemical bonding is the attractive force that holds atoms together in molecules, ions, or crystals. It occurs to achieve stable electronic configuration (octet rule or duplet rule).

• Reason for bonding: To attain noble gas configuration (8 electrons in valence shell).
• Types of bonds formed: By transfer or sharing of electrons.
• Bond energy: Energy required to break the bond.
• Bond length: Equilibrium distance between nuclei.
• Bond angle: Angle between two adjacent bonds.

2. Types of Bonds

Ionic Bond

Electrostatic attraction between oppositely charged ions formed by complete transfer of electrons.

• High melting point, soluble in water, conduct electricity in molten/aqueous state.
• Example: NaCl → Na⁺ + Cl⁻

Covalent Bond

Sharing of electrons between atoms (usually non-metals).

• Single (1 pair), Double (2 pairs), Triple (3 pairs).
• Polar covalent: Unequal sharing (e.g., HCl).
• Non-polar: Equal sharing (e.g., Cl₂).

Coordinate (Dative) Bond

Both electrons shared come from one atom.

• Donor: Lone pair provider (Lewis base).
• Acceptor: Electron deficient (Lewis acid).
• Example: NH₄⁺, H₃O⁺, [Co(NH₃)₆]³⁺

Metallic Bond

Attraction between metal cations and delocalized valence electrons ("electron sea").

• Explains conductivity, malleability, ductility, luster.
• Strength increases with more valence electrons and smaller atomic size.

Hydrogen Bonding

Strong dipole-dipole attraction between H attached to N, O, F and another N, O, F.

• Intramolecular: Within molecule (e.g., o-nitrophenol).
• Intermolecular: Between molecules (e.g., H₂O, NH₃, HF).
• Causes high boiling point of H₂O, DNA structure, protein folding.

3. Lewis Dot Structures and Formal Charge

Lewis structure shows valence electrons as dots and bonds as lines.

Formal charge = Valence e⁻ – (Non-bonding e⁻ + ½ Bonding e⁻)

• Best structure: Lowest formal charges, negative charge on electronegative atom.
• Resonance: Delocalization of electrons (e.g., CO₃²⁻, benzene).

4. VSEPR Theory – Molecular Geometry and Shapes

Valence Shell Electron Pair Repulsion Theory: Electron pairs repel each other and stay as far apart as possible.

Electron Pairs	Steric No.	Geometry	Bond Angle	Examples
2	2	Linear	180°	BeCl₂, CO₂
3	3	Trigonal planar	120°	BF₃, SO₃
4	4	Tetrahedral	109.5°	CH₄, NH₄⁺
3+1	4	Trigonal pyramidal	<109.5°	NH₃
2+2	4	Bent	<109.5°	H₂O
5	5	Trigonal bipyramidal	90°, 120°	PCl₅
6	6	Octahedral	90°	SF₆, [Co(NH₃)₆]³⁺

Lone pairs repel more than bond pairs → distortion in shape.

5. Valence Bond Theory (VBT) – Sigma and Pi Bonds, Hybridization

Overlap of atomic orbitals forms bonds.

• Sigma (σ) bond: Head-on overlap (single bond).
• Pi (π) bond: Sideways overlap (double/triple bonds).

Hybridization Types

Type	Orbitals	Geometry	Example
sp	1s + 1p	Linear	BeCl₂, C₂H₂
sp²	1s + 2p	Trigonal planar	BF₃, C₂H₄
sp³	1s + 3p	Tetrahedral	CH₄, NH₃
sp³d	1s + 3p + 1d	Trigonal bipyramidal	PCl₅
sp³d²	1s + 3p + 2d	Octahedral	SF₆

6. Molecular Orbital Theory (MOT)

Linear Combination of Atomic Orbitals (LCAO) forms molecular orbitals.

• Bonding MO (σ, π): Lower energy, constructive interference.
• Antibonding MO (σ*, π*): Higher energy, destructive interference.

Bond Order = ½ (N₆ – Nₐ)

Molecule	Electron Config	Bond Order	Magnetic
H₂	(σ1s)²	1	Diamagnetic
O₂	(σ2s)²(σ*2s)²(σ2p)²(π2p)⁴(π*2p)²	2	Paramagnetic
N₂	(σ2s)²(σ*2s)²(π2p)⁴(σ2p)²	3	Diamagnetic

7. Comparison of VBT and MOT

Feature	VBT	MOT
Concept	Overlap of half-filled orbitals	Combination of atomic orbitals
Explains paramagnetism	No	Yes (O₂)
Bond order	Not explained	Explained
Resonance	Needed	Not needed

8. Intermolecular Forces

• Van der Waals (London dispersion): Temporary dipoles (all molecules).
• Dipole-Dipole: Between polar molecules.
• Hydrogen bonding: Strongest (already covered).

Strength: Hydrogen bond > Dipole-dipole > London forces

9. Fajans' Rules – Polarization and Covalency

Small cation + large anion + high charge → high polarizing power → more covalent character.

1. Small size of cation → high polarization.
2. Large size of anion → easily polarized.
3. High charge on cation/anion → more covalent.
4. Pseudo inert gas configuration (d¹⁰) > inert gas (e.g., Cu⁺ > Na⁺).

Example: AlCl₃ (covalent) > MgCl₂ > NaCl (ionic)

10. Applications in Engineering Materials

• Ionic solids: Ceramics (Al₂O₃, ZrO₂) – high hardness, refractory.
• Covalent network: Diamond, SiC, Si₃N₄ – superhard abrasives.
• Metallic bonds: Alloys (steel, brass) – conductivity, strength.
• Hydrogen bonding: Polymers (nylon, Kevlar), adhesives.
• Van der Waals: Lubricants (graphite), gecko tape.
• Coordinate bonds: Chelating agents in corrosion inhibitors, catalysts.
• MOT in semiconductors: Band gap in Si, Ge, GaAs.

Remember for Exams:
• Draw Lewis structures with formal charges.
• Predict shape using VSEPR (mention lone pairs).
• Calculate bond order and magnetism using MOT.
• Apply Fajans' rules for % covalent character.
• Hybridization ↔ Geometry link is crucial.