Board exam focus — Chemical Bonding and Molecular Structure (CBSE & HBSE)

This chapter explains WHY and HOW atoms combine to form molecules. Starting from the Kossel-Lewis approach and the octet rule, it builds up to ionic and covalent bonding, bond parameters, VSEPR-predicted shapes, valence bond theory with hybridization, molecular orbital theory (which finally explains paramagnetism of O2), and the unique role of hydrogen bonding. It is one of the most concept-dense and exam-heavy chapters in Class 11 Chemistry.

Kossel-Lewis Approach, Ionic & Covalent Bonds, Bond Parameters

Why Do Atoms Combine?

Atoms combine to attain a stable, lower-energy configuration — usually the nearest noble-gas electronic arrangement. A chemical bond is the attractive force that holds constituent atoms together in a molecule.

Kossel-Lewis Approach (1916)

• Lewis introduced the idea that atoms achieve stability with 8 electrons in the valence shell (octet) and represented bonding/lone pairs as dots (Lewis symbols).
• Kossel emphasized electron transfer between a highly electropositive (Group 1/2) and a highly electronegative (Group 16/17) atom, forming oppositely charged ions held by electrostatic force.

The number of valence electrons available for bonding equals the dots in a Lewis symbol. For carbon (group 14) there are 4 dots.

The Octet Rule

Atoms combine by gaining, losing, or sharing electrons so that each attains a stable octet (8 e⁻) in its outer shell.

Limitations of the Octet Rule

1. Incomplete octet — LiCl, BeCl₂, BCl₃ (central atom has < 8 e⁻).
2. Expanded octet — PCl₅ (10 e⁻), SF₆ (12 e⁻); possible from period 3 onward (vacant d-orbitals).
3. Odd-electron molecules — NO, NO₂ cannot satisfy the octet.
4. Gives no information about shape or energy of molecules.

Ionic (Electrovalent) Bond

Formed by transfer of electrons. Favoured by:

• Low ionization enthalpy of the metal,
• High electron gain enthalpy (more negative) of the non-metal,
• High lattice enthalpy of the resulting crystal.

Lattice enthalpy = energy released when one mole of an ionic solid forms from its gaseous ions. Higher lattice enthalpy → more stable ionic compound. It increases with higher charge and smaller ionic size (per Coulomb's law).

Bond Parameters

• Bond length — equilibrium internuclear distance. Decreases with bond multiplicity: C–C (154 pm) > C=C (134 pm) > C≡C (120 pm).
• Bond angle — angle between two adjacent bonds at the central atom.
• Bond enthalpy — energy required to break one mole of bonds in gaseous state. Triple > double > single.
• Bond order — number of bonds between two atoms (1, 2, 3). Higher bond order → shorter length, greater enthalpy.
• Resonance — when a single Lewis structure is inadequate, the real structure is a resonance hybrid of canonical forms (e.g., O₃, CO₃²⁻). Resonance lowers energy (resonance stabilization); it does NOT mean the molecule oscillates between structures.

Dipole Moment

μ = q × d (unit: Debye, D; 1 D = 3.33 × 10⁻³⁰ C·m). It is a vector. CO₂ is non-polar (μ = 0, linear, dipoles cancel) while H₂O is polar (μ = 1.85 D, bent).

Trap: BeF₂ and CO₂ are non-polar despite polar bonds because their symmetry makes the vector sum zero.

VSEPR Theory, Valence Bond Theory & Hybridization

VSEPR Theory (Valence Shell Electron Pair Repulsion)

Electron pairs (bonding + lone) around a central atom arrange themselves to minimise repulsion, fixing the molecular geometry.

Repulsion order: lone pair–lone pair (lp–lp) > lone pair–bond pair (lp–bp) > bond pair–bond pair (bp–bp).

Effect of lone pairs on angle (lone pairs compress bond angles):

• CH₄ (0 lp): 109.5°
• NH₃ (1 lp): 107°
• H₂O (2 lp): 104.5°

Trap: As lone pairs on the central atom increase, the bond angle DECREASES because lp–bp repulsion squeezes the bonding pairs.

Valence Bond Theory (VBT)

A covalent bond forms by overlap of half-filled atomic orbitals of two atoms. Greater overlap → stronger bond.

• Sigma (σ) bond: head-on (axial) overlap (s–s, s–p, p–p). Strong; allows rotation.
• Pi (π) bond: sideways (lateral) overlap of p-orbitals. Weaker; no free rotation. A double bond = 1σ + 1π; a triple bond = 1σ + 2π.

Hybridization

Mixing of atomic orbitals of comparable energy to form an equal number of equivalent hybrid orbitals with definite geometry.

Quick formula for steric number (hybridization): Steric number = (number of σ-bonded atoms) + (number of lone pairs). SN 2 → sp, 3 → sp², 4 → sp³, 5 → sp³d, 6 → sp³d².

Only sigma bonds and lone pairs count for hybridization; pi bonds do not.

Molecular Orbital Theory, Bond Order & Hydrogen Bonding

Molecular Orbital Theory (MOT)

Atomic orbitals combine to form molecular orbitals (MOs) that belong to the whole molecule. Two AOs give two MOs:

• Bonding MO (σ, π) — lower energy, electron density between nuclei.
• **Antibonding MO (σ, π)** — higher energy, node between nuclei.

Energy order (for O₂, F₂ and heavier): σ1s < σ1s < σ2s < σ2s < σ2pₓ < (π2pᵧ = π2p_z) < (π2pᵧ = π2p_z) < σ*2pₓ

Energy order (for B₂, C₂, N₂): the π2p orbitals lie BELOW σ2pₓ (due to s–p mixing): σ1s < σ1s < σ2s < σ2s < (π2pᵧ = π2p_z) < σ2pₓ < (π2pᵧ = π2p_z) < σ*2pₓ

Bond Order

Bond order = ½ (Nb − Na), where Nb = electrons in bonding MOs, Na = electrons in antibonding MOs.

• Bond order > 0 ⇒ molecule exists; = 0 ⇒ does not exist (e.g., He₂).
• Higher bond order ⇒ greater stability, higher bond enthalpy, shorter bond length.

Magnetic Behaviour

• Paramagnetic — has unpaired electron(s) (attracted by magnetic field).
• Diamagnetic — all electrons paired (weakly repelled).

MOT's triumph: O₂ has 2 unpaired electrons in its π* orbitals ⇒ paramagnetic. VBT/Lewis structures predicted it diamagnetic. This is the classic exam point.

Hydrogen Bonding

An electrostatic attraction between an H atom covalently bonded to a highly electronegative atom (F, O, N) and a lone pair on another electronegative atom.

• Intermolecular H-bonding (between molecules): raises boiling point — H₂O, HF, NH₃, alcohols. Explains why H₂O is liquid while H₂S is gas; why HF boils higher than HCl.
• Intramolecular H-bonding (within one molecule): e.g., o-nitrophenol — lowers boiling point relative to the para isomer.

Ice is less dense than water because H-bonding forms an open cage-like (hexagonal) structure — vital for aquatic life.

Frequently asked questions

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Do these notes follow CBSE and HBSE?

Yes. The Chemical Bonding and Molecular Structure notes are NCERT-aligned and include guidance for both CBSE and Haryana Board (HBSE), with important questions and MCQs for revision.

What does the Chemical Bonding and Molecular Structure chapter cover?

Concept explanations, key formulas and definitions, fully solved examples and board-pattern practice questions for Chemical Bonding and Molecular Structure.