Menu

Chemical Bonding and Molecular Structure

NCERT Class 11 Chemistry Chapter 4

Chemical Bonding for Class 11 Chemistry, JEE and NEET: Lewis structures, ionic and covalent bonds, VSEPR, hybridisation, MO theory and hydrogen bonding.

10 topics Updated 8 Nov 2026

For: Class 11, JEE and NEET, unless a topic says otherwise

In the syllabus: JEE Class 11 syllabus, NEET Class 11 syllabus

With your textbook

NCERT Chemistry Part I (Class 11), Chapter 4 “Chemical Bonding and Molecular Structure”, reprint 2026-27.

Open the official NCERT chapter (PDF, 4.5 MB, ncert.nic.in)

Checked on 19 Oct 2026.

Learn in this order

  1. Kössel–Lewis approach to chemical bonding

    The Kössel–Lewis approach, Lewis symbols, writing Lewis structures, formal charge and the rule’s exceptions.

    NCERT Section 4.1 17 min read 5 solved examples 9 practice questions

  2. Ionic or electrovalent bond

    Electron transfer, factors favouring an ionic (electrovalent) bond, lattice enthalpy and the Born–Haber cycle.

    NCERT Section 4.2 12 min read 5 solved examples 7 practice questions

  3. Bond parameters, resonance and polarity

    Bond length, angle, enthalpy and order, resonance, dipole moment and Fajans’ rules, with solved examples.

    NCERT Section 4.3 17 min read 5 solved examples 10 practice questions

  4. VSEPR theory and the shapes of molecules

    Postulates, shapes of molecules with and without lone pairs and bond angles, with solved examples.

    NCERT Section 4.4 14 min read 5 solved examples 7 practice questions

  5. Valence bond theory: orbital overlap, sigma and pi bonds

    Valence bond theory: how H forms, positive, negative and zero orbital overlap, sigma and pi bonds and their strength, for Class 11, JEE and NEET.

    NCERT Section 4.5 13 min read 4 solved examples 8 practice questions

  6. Hybridisation: sp, sp², sp³, sp³d and sp³d²

    Sp, sp, sp, spd and spd with shapes and bond angles, the steric number shortcut and solved examples.

    NCERT Section 4.6 19 min read 4 solved examples 8 practice questions

  7. Molecular orbital theory, bond order and magnetism

    Molecular orbital theory: LCAO, bonding and antibonding orbitals, the two MO energy orders, bond order and magnetism, for Class 11, JEE and NEET.

    NCERT Section 4.7 16 min read 4 solved examples 8 practice questions

  8. Bonding in homonuclear diatomic molecules: H₂ to Ne₂

    MO configurations, bond orders and magnetism of H to Ne, and why O is paramagnetic.

    NCERT Section 4.8 15 min read 4 solved examples 8 practice questions

  9. Hydrogen bonding: cause, types and effects

    Its cause, intermolecular and intramolecular types with examples, and effects on boiling point and density.

    NCERT Section 4.9 13 min read 4 solved examples 8 practice questions

  10. Metallic bonding: the electron sea model

    Metallic bonding: positive ion cores in a sea of delocalised electrons, what makes it strong, and why metals conduct, shine and bend, for JEE and NEET.

    JEE and NEET only 10 min read 4 solved examples 8 practice questions

Key formulas in this chapter

Kössel–Lewis approach

  • Group valence = dots in the Lewis symbol, or 8 − dots
  • Electrons for a Lewis structure = Σ valence electrons − (charge)
  • Formal charge = V − L − ½B
  • Single bond: 1 shared pair; double: 2; triple: 3

Ionic bond

  • M(g) → M(g) + e: ionisation enthalpy (positive)
  • X(g) + e → X(g): electron gain enthalpy
  • NaCl: 495.8 − 348.7 = 147.1 kJ mol; lattice enthalpy +788 kJ mol (lattice formation −788)
  • Born–Haber: ΔlatticeH = −ΔfH + ΔsubH + ½ΔbondH + ΔiH + ΔegH
  • Lattice enthalpy rises with ionic charge and falls with ionic size

Bond parameters

  • Approximate bond length R ≈ rA + rB
  • Mean bond enthalpy = total dissociation enthalpy ÷ number of bonds
  • μ = Q × r; 1 D = 3.33564 × 10 C m
  • % ionic character = μ(observed) ÷ μ(fully ionic) × 100
  • For comparable bonds, bond order ↑ usually means bond enthalpy ↑ and bond length ↓

VSEPR theory

  • Repulsion order: lp–lp > lp–bp > bp–bp
  • Shortcut domains = ½(V + M − c + a), for the simple cases discussed; verify with a Lewis structure
  • 2, 3, 4, 5, 6 pairs: linear 180°, trigonal planar 120°, tetrahedral 109.5°, trigonal bipyramidal 120°/90°, octahedral 90°
  • CH 109.5° > NH 107° > HO 104.5°

Valence bond theory

  • Single bond = 1σ; double = 1σ + 1π; triple = 1σ + 2π
  • Comparable σ contributions are generally stronger than π contributions between the same atoms
  • H: bond length 74 pm, bond enthalpy 435.8 kJ mol

Hybridisation

  • Steric number = σ bonds + lone pairs on the atom
  • Steric number 2, 3, 4, 5, 6: sp, sp, sp, spd, spd
  • s character: sp 50%, sp 33.3%, sp 25%
  • Angles: sp 180°, sp 120°, sp 109.5°

Molecular orbital theory

  • Bond order = ½ × (bonding electrons − antibonding electrons)
  • Bonding MO: ψA + ψB; antibonding MO: ψA − ψB
  • Up to N: σ1s < σ*1s < σ2s < σ*2s < π2p < σ2p < π*2p < σ*2p
  • O and F: σ1s < σ*1s < σ2s < σ*2s < σ2p < π2p < π*2p < σ*2p

Homonuclear diatomic molecules

  • Bond order = ½ × (bonding electrons − antibonding electrons)
  • Bond orders: H 1, He 0, Li 1, Be 0, B 1, C 2, N 3, O 2, F 1, Ne 0
  • Paramagnetic: B, O (two unpaired electrons each)

Hydrogen bonding

  • Common strong hydrogen bonds: X–H···Y, usually X/Y = F, O or N; broader cases exist
  • Strength: covalent bond > hydrogen bond > van der Waals interaction (typical; ranges overlap)
  • Extent of hydrogen bonding: solid > liquid > gas
  • Hydrogen-bond strength depends on donor, acceptor, charge, distance and geometry

Metallic bonding

  • Elementary trend for comparable metals: more participating electrons and smaller cores often strengthen cohesion
  • Melting points: Na 98 °C < Mg 650 °C < Al 660 °C
  • Group 1 melting points fall: Li 181 °C, Na 98 °C, K 64 °C, Rb 39 °C, Cs 28 °C

More in Chemistry