Chemical Bonding and Molecular Structure

Chemical bonding explains how atoms combine to form molecules and compounds. This is one of the most important topics in JEE Chemistry.

Overview

graph TD
    A[Chemical Bonding] --> B[Ionic Bonding]
    A --> C[Covalent Bonding]
    A --> D[Metallic Bonding]
    C --> C1[Lewis Structures]
    C --> C2[VSEPR Theory]
    C --> C3[Valence Bond Theory]
    C --> C4[Molecular Orbital Theory]
    C3 --> C3a[Hybridization]

Kossel-Lewis Approach

Octet Rule

Atoms tend to achieve 8 electrons in their valence shell (noble gas configuration).

Lewis Dot Structures

• Count total valence electrons
• Arrange atoms (central atom usually least electronegative)
• Distribute electrons as bonding and lone pairs
• Check octet completion

Formal Charge

$$\text{Formal Charge} = \text{Valence e}^- - \text{Lone pair e}^- - \frac{1}{2}(\text{Bonding e}^-)$$

Ionic Bonding

Formation of Ionic Bonds

graph LR
    A[Metal atom] -->|loses e⁻| B[Cation]
    C[Non-metal] -->|gains e⁻| D[Anion]
    B --> E[Ionic Compound]
    D --> E

Lattice Enthalpy

Energy released when gaseous ions combine to form ionic solid:

$$M^+(g) + X^-(g) \rightarrow MX(s) + \text{Lattice Energy}$$

Factors affecting lattice energy:

• Higher charge → Higher lattice energy
• Smaller size → Higher lattice energy

Born-Haber Cycle

Used to calculate lattice enthalpy indirectly:

$$\Delta H_f = \Delta H_{sub} + IE + \frac{1}{2}D + EA + U$$

where U is lattice energy.

Covalent Bonding

Bond Parameters

Parameter	Definition
Bond Length	Distance between nuclei of bonded atoms
Bond Angle	Angle between two adjacent bonds
Bond Order	Number of bonds between two atoms
Bond Energy	Energy required to break one mole of bonds

Electronegativity

Tendency of an atom to attract shared electron pair.

Pauling Scale: F (4.0) > O (3.5) > Cl (3.0) > N (3.0) > Br (2.8) > C (2.5) > H (2.1)

Fajan’s Rules

Covalent character increases when:

1. Cation is small with high charge
2. Anion is large with high charge
3. Cation has pseudo noble gas configuration

VSEPR Theory

Valence Shell Electron Pair Repulsion theory predicts molecular shapes.

Principles

1. Electron pairs around central atom repel each other
2. They arrange to maximize separation
3. Lone pairs occupy more space than bonding pairs

Repulsion Order

$$\text{LP-LP} > \text{LP-BP} > \text{BP-BP}$$

Molecular Geometries

Total e⁻ pairs	Bonding pairs	Lone pairs	Geometry	Example
2	2	0	Linear	BeCl₂
3	3	0	Trigonal planar	BF₃
3	2	1	Bent	SnCl₂
4	4	0	Tetrahedral	CH₄
4	3	1	Trigonal pyramidal	NH₃
4	2	2	Bent	H₂O
5	5	0	Trigonal bipyramidal	PCl₅
5	4	1	See-saw	SF₄
5	3	2	T-shaped	ClF₃
5	2	3	Linear	XeF₂
6	6	0	Octahedral	SF₆
6	5	1	Square pyramidal	BrF₅
6	4	2	Square planar	XeF₄

Valence Bond Theory

Concept

A covalent bond forms when:

1. Two half-filled orbitals overlap
2. Electrons have opposite spins
3. Maximum overlap gives maximum bond strength

Types of Overlap

• σ bond: Head-on overlap (s-s, s-p, p-p axial)
• π bond: Lateral/sideways overlap (p-p lateral)

Hybridization

Concept

Mixing of atomic orbitals to form new hybrid orbitals with different shapes and energies.

Types of Hybridization

graph TD
    A[Hybridization] --> B[sp - Linear]
    A --> C[sp² - Trigonal planar]
    A --> D[sp³ - Tetrahedral]
    A --> E[sp³d - TBP]
    A --> F[sp³d² - Octahedral]

Hybridization	Geometry	Bond Angle	Example
sp	Linear	180°	BeCl₂, C₂H₂
sp²	Trigonal planar	120°	BF₃, C₂H₄
sp³	Tetrahedral	109.5°	CH₄, NH₄⁺
sp³d	Trigonal bipyramidal	90°, 120°	PCl₅
sp³d²	Octahedral	90°	SF₆

Calculating Hybridization

$$\text{Hybridization} = \frac{1}{2}[V + M - C + A]$$

where:

• V = Valence electrons of central atom
• M = Number of monovalent atoms
• C = Charge on cation
• A = Charge on anion

Molecular Orbital Theory (MOT)

Principles

1. Atomic orbitals combine to form molecular orbitals
2. Number of MOs = Number of AOs combined
3. Bonding MO: lower energy, stabilizing
4. Antibonding MO: higher energy, destabilizing

Bond Order

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

where $N_b$ = electrons in bonding MOs, $N_a$ = electrons in antibonding MOs

MO Energy Order

For O₂, F₂ (Z > 7):

$$\sigma_{1s} < \sigma^*_{1s} < \sigma_{2s} < \sigma^*_{2s} < \sigma_{2p_z} < \pi_{2p_x} = \pi_{2p_y} < \pi^*_{2p_x} = \pi^*_{2p_y} < \sigma^*_{2p_z}$$

For N₂, C₂, B₂ (Z ≤ 7):

$$\sigma_{1s} < \sigma^*_{1s} < \sigma_{2s} < \sigma^*_{2s} < \pi_{2p_x} = \pi_{2p_y} < \sigma_{2p_z} < \pi^*_{2p_x} = \pi^*_{2p_y} < \sigma^*_{2p_z}$$

Properties from MOT

Molecule	Bond Order	Magnetic Nature
H₂	1	Diamagnetic
N₂	3	Diamagnetic
O₂	2	Paramagnetic
O₂⁻	1.5	Paramagnetic
O₂²⁻	1	Diamagnetic

Hydrogen Bonding

Conditions

1. H must be bonded to highly electronegative atom (F, O, N)
2. Lone pair on electronegative atom should be available

Types

• Intermolecular: Between different molecules (H₂O, HF)
• Intramolecular: Within same molecule (o-nitrophenol)

Effects

• Higher boiling points
• Solubility in water
• Anomalous properties of water

Practice Problems

1. Predict the shape of XeF₄ using VSEPR theory.

2. Find the hybridization of central atom in SF₄ and explain its shape.

3. Write the MO configuration of O₂⁺ and find its bond order.

4. Explain why H₂O has a bond angle of 104.5° instead of 109.5°.

Further Reading

• Atomic Structure
• Coordination Compounds