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    Chemical Bonding and Molecular Structure

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    Chemical Bonding and Molecular Structure – Class 11 Notes

    Chemical Bonding & Molecular Structure

    Detailed Class 11 Notes for CBSE | ISC | WBCHSE – Principles, Types, Shapes, Theories, MOT, Practice MCQs

    1. Introduction

    Chemical Bond: A force of attraction that holds two or more atoms together in a molecule or compound, making them behave as a single unit.

    Understanding bond formation explains why atoms combine, molecular shapes, and physical properties of substances.

    2. Types of Chemical Bonds

    (a) Ionic (Electrovalent) Bond

    • Formed by complete transfer of electrons from a metal to a non-metal.
    • Results in ions held by strong electrostatic forces.
      Example: NaCl (Na → Na+ + e–; Cl + e– → Cl–)
    Ionic compounds are usually crystalline solids with high melting points and conduct electricity in molten/solution state.

    (b) Covalent Bond

    • Formed by mutual sharing of electrons between non-metal atoms to complete their octet.
    • Single, double, or triple bonds: H2, O2, N2
    Cl2: Cl· + ·Cl → Cl:Cl (single covalent bond)
    Covalent compounds are generally soft, low melting, and poor conductors.

    (c) Coordinate (Dative) Bond

    • Both bonding electrons are supplied by one atom (donor).
    • Example: NH4+ formation, SO2, O3

    3. Octet Rule and Its Limitations

    • Atoms attain stability by acquiring 8 electrons in their valence shell (duplet for H, He).
    • Failures: Odd electron species (NO), expanded octet (PF5), incomplete octet (BF3)
    [Insert Lewis structures: BF3, PF5, NO Here]

    4. Lewis Structures and Resonance

    • Lewis structure: Shows valence electrons as dots/crosses. Shared pair = line; lone pair = dots.
    • Formal Charge: The charge assigned to an atom in a given Lewis structure
      Formal charge = (Valence e–) – (Lone e–) – ½ (Bonding e–)
    • Resonance: When two or more valid structures are possible; real molecule is a hybrid.
    [CO32- resonance structures]

    5. Theories of Chemical Bonding

    5.1 Valence Bond Theory (VBT)

    • Bonds form when atomic orbitals overlap; electron pairs localized between atoms.
    • Types: Sigma (σ) – head-on overlap; Pi (π) – sidewise overlap.
    [Sigma and Pi bond formation diagrams]

    5.2 Hybridization

    • Mixing of atomic orbitals to form new hybrid orbitals of equal energy.
      sp, sp2, sp3 etc. for linear, trigonal planar, tetrahedral shapes.
    Examples: CH4 – sp3, BF3 – sp2, BeCl2 – sp

    5.3 Molecular Orbital Theory (MOT)

    MOT was developed by Hund and Mulliken and explains bonding through the formation of molecular orbitals extending over the entire molecule, not just between two atoms.
    • Atomic orbitals combine to form Molecular Orbitals (MOs): The number of MOs formed equals the number of atomic orbitals combined.
    • Bonding vs. Antibonding:
      • Bonding MOs (lower energy): Constructive overlap, electron density between nuclei increases stability (denoted σ, π).
      • Antibonding MOs (higher energy): Destructive overlap, electron density between nuclei decreases (denoted σ*, π*).
    • Filling of MOs:
      • MOs are filled in order of increasing energy (Aufbau principle), with each orbital holding 2 electrons with opposite spins.
      • Bonding orbitals are filled before antibonding orbitals.
    • Types of Molecular Orbitals:
      • σ (sigma): Along internuclear axis (head-on overlap of s-s, s-p, or p-pz).
      • π (pi): Side-to-side overlap of p orbitals (p-px, p-py).
    Bond Order (BO) = ½ × [Number of electrons in bonding MOs − Number in antibonding MOs]

    BO = ( Nbonding  − Nantibonding )/2

    Energy Level Diagram for Homonuclear Diatomics

    [Insert Molecular Orbital Energy Level Diagram for B2 to N2 and O2 to F2 here]
    • Order of orbital energies:
      For Z ≤ 7 (Li2, Be2, B2, C2, N2):
      σ1s < σ*1s < σ2s < σ*2s < π2px = π2py < σ2pz < π*2px = π*2py < σ*2pz
    • For Z ≥ 8 (O2, F2, Ne2):
      σ1s < σ*1s < σ2s < σ*2s < σ2pz < π2px = π2py < π*2px = π*2py < σ*2pz

    Applications of MOT

    • Explains bond order, magnetism, and stability of molecules.
    • Paramagnetism of O2: O2 has two unpaired electrons in π*2p MOs, explaining its observed paramagnetism, which cannot be explained by VBT.
    • Predicts existence/non-existence of molecules (e.g., He2 does not exist as BO=0).
    Examples:
    For O2:
    Configuration (Z=8):
    σ1s2 σ*1s2 σ2s2 σ*2s2 σ2pz2 π2px2 π2py2 π*2px1 π*2py1
    Bonding electrons = 8, Antibonding = 4
    BO = (8 − 4)/2 = 2
    Summary of Bond Order and Magnetism using MOT:
    N2 (BO=3) – Diamagnetic
    O2 (BO=2) – Paramagnetic
    O2– (BO=1.5) – Paramagnetic (1 unpaired)
    He2 (BO=0) – Does not exist
    [Insert diagrams showing electron filling in MOs for N2, O2, He2]

    6. Polarity and Dipole Moment

    • Electronegativity: Tendency to attract shared electrons (Pauling scale).
    • Covalent bonds with differing EN create dipole moments.
    • μ (dipole moment) = q × r, q = charge, r = bond length (unit: Debye)
    Water is highly polar due to bent shape; CO2 is nonpolar despite polar bonds (linear geometry cancels dipoles).

    7. Molecular Shapes: VSEPR Theory

    • Electron pairs (bonding and lone) arrange to minimize mutual repulsion.
    • Lone pairs occupy more space than bond pairs.
    • Common geometries:
      • Tetrahedral (sp3): 4 BP, 0 LP – CH4
      • Pyramidal: 3 BP, 1 LP – NH3
      • Angular/Bent: 2 BP, 2 LP – H2O
      • Trigonal Planar (sp2): 3 BP, 0 LP – BF3
      • Linear (sp): 2 BP – BeCl2, CO2
    [VSEPR shapes: CH4, NH3, H2O, BF3, BeCl2]

    8. Formal Charge and Oxidation Number

    • Formal charge: theoretical charge in a Lewis structure.
    • Oxidation number: actual/ionic charge; often differs from formal charge.

    9. Bond Parameters

    • Bond length: Distance between nuclei; shorter for higher bond order.
    • Bond angle: Angle between two bonds at the atom.
    • Bond order: Higher = stronger, shorter bond. Order: single (1), double (2), triple (3).
    • Bond energy: Energy to break one mole of bonds in gas phase.
    Increasing bond order increases bond energy and decreases bond length.

    10. Exceptions and Special Cases

    • Expanded Octet in PF5, SF6
    • Odd Electron Molecule: NO, NO2
    • Incomplete Octet: BeCl2, BF3
    • Back bonding: Seen in BF3, CO
    Always analyze exceptions to predict molecular structures accurately in exams!

    11. Practice MCQs (Answer & Explanation Below)

    MCQOptions
    1. Which of the following forms an ionic bond?a) H2 b) O2 c) NaCl d) Cl2
    2. The number of lone pairs in NH3 is:a) 0 b) 1 c) 2 d) 3
    3. Shape of BF3 molecule is:a) Tetrahedral b) Trigonal planar c) Bent d) Linear
    4. Which is a polar molecule?a) CO2 b) CH4 c) H2 d) H2O
    5. The bond order of O2 molecule is:a) 1 b) 2 c) 3 d) 4
    6. Bond angle in water (H2O) is:a) 90° b) 104.5° c) 120° d) 180°
    7. Which has a coordinate bond?a) H2 b) NH4+ c) O2 d) H2O
    8. Which molecule has an expanded octet?a) BH3 b) PF5 c) CO2 d) CH4
    9. Valence electrons in BeCl2?a) 4 b) 6 c) 8 d) 10
    10. Resonance occurs in:a) CH4 b) SO3 c) NH3 d) BeCl2
    11. Which is not explained by VSEPR?a) Lone pairs effect b) d-p π bonding c) Hybridization d) Bond length
    12. Which is paramagnetic?a) N2 b) O2 c) CO2 d) HCl
    13. Molecule with highest bond order:a) O2 b) N2 c) F2 d) H2
    14. Which hybridization leads to linear shape?a) sp b) sp2 c) sp3 d) d2sp3
    15. Bond angle in CH4 is:a) 90° b) 104.5° c) 120° d) 109.5°
    16. The term ‘Lewis acid’ refers to:a) Proton donor b) Electron donor c) Electron pair acceptor d) None
    17. Example of molecule with bent shape:a) CO2 b) H2O c) BeF2 d) BF3
    18. Which does not obey the octet rule?a) CH4 b) NH3 c) BeCl2 d) H2O
    19. The correct electron-dot structure for NO3–:a) Single N−O bond only b) Double bonds only c) Resonance hybrid d) None
    20. Bond order for C≡N in HCN:a) 1 b) 2 c) 3 d) 4
    21. Presence of two unpaired electrons in O2 is explained by:a) Lewis theory b) VBT c) MOT d) VSEPR
    22. Which has maximum number of lone pairs on central atom?a) NH3 b) H2O c) BeCl2 d) PF5
    23. CO2 is nonpolar because:a) C has no electronegativity b) Shape is linear c) Bonds not polar d) None
    24. A triple bond consists of:a) 3 σ bond b) 2 σ + 1 π c) 1 σ + 2 π d) 3 π
    25. The formal charge on O in O3 (central O):a) +1 b) 0 c) -1 d) +2
    26. The maximum covalency of nitrogen is:a) 3 b) 4 c) 5 d) 6
    27. The energy required to break a bond is:a) Bond energy b) Ionization enthalpy c) Electron affinity d) Lattice energy
    28. Molecule with one lone pair and three bond pairs:a) NH3 b) CH4 c) H2O d) BF3
    29. Which is a planar molecule?a) NH3 b) CH4 c) BeCl2 d) H2O
    30. Back bonding is seen in:a) C2H4 b) BF3 c) H2O d) NH3

    Answers & Explanations to MCQs

    1. c) NaCl
      Because Na (metal) transfers electron to Cl (non-metal); forms strong ionic bond.
    2. b) 1
      NH3 has three bond pairs and one lone pair on N.
    3. b) Trigonal planar
      BF3 is sp2 hybridized (AX3), 120° angle.
    4. d) H2O
      Bent molecule with polar bonds, net dipole; CO2 is linear so overall nonpolar.
    5. b) 2
      O2 bond order from MOT is 2.
    6. b) 104.5°
    7. b) NH4+
      One bond (from N to H+) is coordinate.
    8. b) PF5
      Central P involves 10 valence electrons (expanded octet).
    9. a) 4
      Be(2e) + two Cl(1e each) = 4 valence electrons around Be.
    10. b) SO3
      SO3, like CO32-, exhibits resonance.
    11. b) d-p π bondingVSEPR only deals with electron pair repulsions, not d-p π bonding.
    12. b) O2
      MOT: O2 has unpaired electrons, so paramagnetic.
    13. b) N2
      N2 bond order is 3 (triple bond).
    14. a) sp
      sp hybridization yields 180° bond angle; linear shape (e.g., BeCl2, CO2).
    15. d) 109.5°
    16. c) Electron pair acceptor
      Lewis acid accepts an electron pair.
    17. b) H2O
      Bent shape; two lone pairs on O.
    18. c) BeCl2
      Be has only 4 electrons in its valence shell (violates octet).
    19. c) Resonance hybrid
      NO3– has 3 equivalent resonance structures.
    20. c) 3
      C≡N bond is a triple bond.
    21. c) MOT
      Molecular Orbital Theory explains the paramagnetism of O2.
    22. b) H2O
      2 bond pairs, 2 lone pairs on O.
    23. b) Shape is linear
      CO2 has polar bonds but is linear, dipoles cancel.
    24. c) 1 σ + 2 π
      Triple bond includes 1 sigma and 2 pi bonds (as in N2 or C≡N).
    25. a) +1
      For central O in O3; see formal charge calculation.
    26. c) 5
      Maximum covalency (e.g., in NH4+, N2O5 etc.)
    27. a) Bond energy
    28. a) NH3
      NH3: 3 bond pairs + 1 lone pair (on N).
    29. c) BeCl2
      BeCl2 is linear and planar.
    30. b) BF3
      B in BF3 gets electron density from F by back bonding.
    © 2025 Chemistry SmartNotes | Designed for Class 11 (CBSE, ISC, WBCHSE)

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