Chemical Bonding and Molecular Structure — Class 11 MCQs with Answers

Chemical Bonding and Molecular Structure — important questions & MCQs with answers (Class 11 Chemistry)

6 solved questions from this chapter's difficulty levels, each with its answer and explanation. Tap an option to check yourself.

0 of 6 answered

1. Q1 of 6Easy

   NaCl has which type of bond?

   AIonicBCovalentCMetallicDHydrogen
   Show answerHide answer

   Answer: Ionic

   Explanation: Bond type follows the electronegativity gap: Na (0.9) and Cl (3.0) differ by ~2.1, too large for sharing, so Na transfers an electron to Cl, forming Na⁺Cl⁻ — an ionic bond, not covalent.

2. Q2 of 6Easy

   Hybridisation in BF₃:

   AspBsp²Csp³Dsp³d
   Show answerHide answer

   Answer: sp²

   Explanation: Boron in BF₃ has 3 bond pairs and no lone pair, so VSEPR gives sp² hybridisation and a trigonal planar shape with 120° angles. sp³ would need a 4th group or a lone pair, which BF₃ has neither.

3. Q3 of 6Medium

   Bond order of O₂ from MOT:

   A3B1C2D2.5
   Show answerHide answer

   Answer: 2

   Explanation: MO theory gives O₂ 8 bonding electrons (σ2s, σ2p, π2p) and 4 antibonding (σ*2s, π*2p), so bond order = (8−4)/2 = 2, matching the O=O double bond drawn in its Lewis structure.

4. Q4 of 6Medium

   Shape of CH₄:

   ATetrahedralBSquare planarCLinearDOctahedral
   Show answerHide answer

   Answer: Tetrahedral

   Explanation: Carbon's 4 bond pairs in CH₄ with no lone pair give sp³ hybridisation and a tetrahedral shape with 109.5° angles — the ideal geometry for 4 equivalent electron domains and zero lone-pair distortion.

5. Q5 of 6Hard

   Polar molecule among these:

   AH₂OBCO₂CCH₄DBF₃
   Show answerHide answer

   Answer: H₂O

   Explanation: H₂O is bent (2 lone pairs on O), so its 2 O–H bond dipoles don't cancel, giving a net dipole. CO₂, CH₄, and BF₃ are symmetric (linear, tetrahedral, trigonal planar) with identical bonds, so theirs cancel to zero.

6. Q6 of 6Hard

   Hybridisation of carbon in C₂H₄ (ethene):

   AspBsp²Csp³Dsp³d
   Show answerHide answer

   Answer: sp²

   Explanation: Each carbon in H₂C=CH₂ forms 3 σ bonds (2 to H, 1 to C) plus 1 π bond from an unhybridised p-orbital, needing only 3 hybrid orbitals — sp². The leftover p-orbital handles the π bond of the C=C double bond.

Chemical Bonding and Molecular Structure notes: key concepts (Class 11 Chemistry)

This chapter explains why atoms combine and what shape the result takes. Starting from Kössel–Lewis and the octet rule, it builds ionic and covalent bonding, Fajans' rules, VSEPR geometry, valence bond theory with hybridisation, and molecular orbital theory, then closes with bond parameters, dipole moment and hydrogen bonding.

Lewis structures and the octet

Atoms combine to attain a noble-gas octet by sharing or transferring valence electrons; Lewis dot structures show bonding pairs and lone pairs around each atom.

Ionic bond

Formed by electron transfer from a low-IE metal to a high-EA non-metal; its stability is set by lattice enthalpy, rising with charge and smaller ions.

Covalent bond and parameters

Shared pairs bind atoms; bond length (internuclear distance) shortens and bond enthalpy grows as bond order rises, so triple bonds are shortest and strongest.

Coordinate bond

A covalent bond where both shared electrons come from one atom (donor); once formed it is identical to a normal covalent bond, as in NH₄⁺ and O₃.

Formal charge

Hypothetical charge on an atom = valence e⁻ − lone-pair e⁻ − ½ bonding e⁻; the Lewis structure with the smallest formal charges is the most stable.

Exceptions to the octet rule

The octet fails for incomplete octets (BF₃), odd-electron species (NO), and expanded octets using d-orbitals (SF₆, PCl₅ with 12 and 10 electrons).

Fajans' rules

Small, highly charged cations with high polarising power distort large anions, adding covalent character to an ionic bond — why AlCl₃ is largely covalent.

VSEPR theory

Electron pairs around the central atom orient to minimise repulsion (lp–lp > lp–bp > bp–bp); lone pairs bend shapes, as in bent H₂O and pyramidal NH₃.

Valence bond theory

A bond forms by overlap of half-filled atomic orbitals; greater overlap gives a stronger bond, and the shared pair has opposite spins localised between nuclei.

Hybridisation

Mixing of atomic orbitals of similar energy gives equivalent hybrids: sp (linear, 180°), sp² (trigonal, 120°), sp³ (tetrahedral, 109.5°).

Sigma and pi bonds

A σ bond has head-on axial overlap and is strong; a π bond has sideways overlap and is weaker; single = 1σ, double = 1σ+1π, triple = 1σ+2π.

Molecular orbital theory

Atomic orbitals combine into bonding and antibonding molecular orbitals filled by Aufbau, Pauli and Hund; it explains O₂'s paramagnetism, which VBT cannot.

Dipole moment and polarity

μ = q × d, a vector in debye; it measures polarity. CO₂ is non-polar (bonds cancel) while bent H₂O is polar, so symmetry decides the net moment.

Hydrogen bonding

An attraction between H bonded to N, O or F and a lone pair on another such atom; it explains water's high boiling point and ice being less dense than water.