Chemical Bonding and Molecular Structure – Class 11 Notes

Introduction

Chemical bonding is a fundamental concept in chemistry that explains how atoms combine to form molecules and compounds. Understanding chemical bonding helps us comprehend the structures, shapes, and properties of various substances in nature.

Atoms bond to attain a stable electronic configuration, often the noble gas configuration, which is associated with minimum energy and maximum stability. This process of bond formation involves redistribution of electrons between atoms.

Why Do Atoms Bond?

• Atoms bond to attain stability.
• Most atoms follow the Octet Rule: atoms tend to gain, lose, or share electrons to have 8 electrons in their valence shell.
• Bond formation lowers the potential energy of the system.

Types of Chemical Bonds

There are mainly three types of chemical bonds:

1. Ionic Bond (Electrovalent Bond)

• Formed by transfer of electrons from one atom to another.
• Usually occurs between metals (which lose electrons) and non-metals (which gain electrons).
• The atom that loses electrons becomes a cation and the one that gains becomes an anion.
• These oppositely charged ions are held together by electrostatic attraction.

Example:
Na (sodium) + Cl (chlorine) → Na⁺ + Cl⁻ → NaCl

Properties of Ionic Compounds:

• High melting and boiling points
• Conduct electricity in molten or aqueous state
• Generally soluble in water

2. Covalent Bond

• Formed by sharing of electrons between atoms.
• Usually occurs between non-metals.
• The shared electrons belong to both atoms, resulting in a stable molecule.

Types of Covalent Bonds:

• Single bond – sharing of 1 electron pair (e.g., H₂, Cl₂)
• Double bond – sharing of 2 electron pairs (e.g., O₂)
• Triple bond – sharing of 3 electron pairs (e.g., N₂)

Properties:

• Lower melting and boiling points than ionic compounds
• Do not conduct electricity
• Often exist as gases or liquids

3. Coordinate (Dative) Covalent Bond

• A type of covalent bond where the shared pair of electrons comes from only one atom.
• Once formed, it is indistinguishable from a normal covalent bond.

Example:
NH₃ + H⁺ → NH₄⁺
(The lone pair on N is donated to H⁺)

Octet Rule

Atoms tend to gain, lose, or share electrons so as to have 8 electrons in their outermost shell.

Limitations of Octet Rule:

• Incomplete octet: e.g., H (2e⁻), Be (4e⁻), B (6e⁻)
• Expanded octet: e.g., PCl₅, SF₆
• Molecules with odd number of electrons: e.g., NO, NO₂

Lewis Structures

• Represent bonding in molecules using dots and lines.
• Dots represent valence electrons.
• Lines represent shared pairs (bonds).

Steps to draw:

1. Count total valence electrons.
2. Identify central atom (usually least electronegative).
3. Form single bonds between central and surrounding atoms.
4. Distribute remaining electrons to complete octets.
5. Form double or triple bonds if needed.

Example:
CO₂: O=C=O

Formal Charge

Used to determine the most stable Lewis structure.

Formula:
Formal Charge (FC) = [Valence electrons] – [Non-bonding electrons] – ½[Bonding electrons]

The structure with the lowest formal charges on atoms is more stable.

Resonance

When more than one valid Lewis structure is possible, the molecule shows resonance.

• Actual structure is a resonance hybrid, a blend of all structures.
• Resonance increases stability.

Example:
O₃ (ozone), NO₃⁻ (nitrate ion)

VSEPR Theory (Valence Shell Electron Pair Repulsion)

Used to predict the shape of molecules.

Basic idea: Electron pairs (bonding and lone pairs) around the central atom repel each other and try to stay as far apart as possible.

Electron Pairs	Shape	Bond Angle	Example
2	Linear	180°	BeCl₂
3	Trigonal planar	120°	BF₃
4	Tetrahedral	109.5°	CH₄
5	Trigonal bipyramidal	90°, 120°	PCl₅
6	Octahedral	90°	SF₆

Effect of lone pairs:

• Lone pairs cause greater repulsion.
• They reduce bond angles.

Example:
NH₃ is trigonal pyramidal (107°), H₂O is bent (104.5°)

Hybridization

Concept: Atomic orbitals mix to form new equivalent orbitals called hybrid orbitals.

Types of Hybridization:

Type	Orbitals involved	Shape	Example
sp	1 s + 1 p	Linear	BeCl₂
sp²	1 s + 2 p	Trigonal planar	BF₃
sp³	1 s + 3 p	Tetrahedral	CH₄
sp³d	1 s + 3 p + 1 d	Trigonal bipyramidal	PCl₅
sp³d²	1 s + 3 p + 2 d	Octahedral	SF₆

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Molecular Orbital Theory (MOT)

Explains bonding using wave nature of electrons. Atomic orbitals combine to form molecular orbitals (MO).

Types of MOs:

• Bonding orbital (σ, π) – lower energy, stable
• Antibonding orbital (σ, π)** – higher energy, unstable

Molecular Orbital Energy Diagram (for 1st and 2nd period elements)

• For O₂, F₂: σ2s < σ2s < σ2p < π2p < π2p < σ*2p
• For B₂, C₂, N₂: σ2s < σ2s < π2p < σ2p < π2p < σ*2p

Bond Order = ½ (No. of electrons in bonding MOs − in antibonding MOs)

• Bond order > 0 → stable
• Bond order = 0 → unstable

Example:
O₂: 8 bonding electrons, 4 antibonding → B.O. = ½(8−4) = 2

O₂ is paramagnetic (has unpaired electrons) – proved by MOT.

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Hydrogen Bonding

A special type of dipole-dipole attraction between a hydrogen atom and a highly electronegative atom like N, O, or F.

Types:

• Intermolecular: Between molecules (e.g., water)
• Intramolecular: Within the same molecule (e.g., o-nitrophenol)

Effects of H-bonding:

• Higher boiling points (e.g., H₂O > H₂S)
• Ice is less dense than water due to H-bonding

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Polarity of Bonds and Molecules

• Non-polar bond: Electrons shared equally (e.g., Cl₂)
• Polar covalent bond: Electrons shared unequally due to electronegativity difference (e.g., HCl)

Dipole Moment (μ):

μ = Q × r
(Q = charge, r = distance between charges)

• Vector quantity
• Measured in Debye (D)
• Indicates polarity of molecules

Examples:

• HCl is polar, has dipole moment.
• CO₂ is non-polar (dipoles cancel out due to linear shape).

Comparison of Different Bonds

Property	Ionic Bond	Covalent Bond	Hydrogen Bond
Electron Involvement	Transfer	Sharing	Attraction
Strength	Strong	Moderate	Weak
Conductivity	In solution	Poor	No
Boiling Point	High	Variable	High (if H-bond)
Examples	NaCl	H₂O, CH₄	H₂O, NH₃