Chemical Bonding & Molecular Structure — Complete NEET Guide — Summary

Chemical bonding describes how atoms combine to form molecules with lower energy and greater stability than isolated atoms. For NEET 2026, mastery of three interconnected frameworks — Ionic/Covalent nature, VSEPR, and MOT — is essential as this chapter contributes 3–4 questions per year.

Ionic Bonding

Ionic bonds form by complete electron transfer from electropositive metals to electronegative non-metals, producing oppositely charged ions held together by electrostatic attraction in a crystal lattice. The stability of an ionic compound is measured by its lattice enthalpy (U) — the energy released when gaseous ions combine to form one mole of ionic solid. Lattice enthalpy cannot be measured directly; instead, it is calculated via the Born-Haber cycle:

$\Delta H_f = \Delta H_{sub} + \tfrac{1}{2}\Delta H_{diss} + IE + EA + U$

For NaCl: $\Delta H$ _f = −411, $\Delta H$ _sub(Na) = +108, IE(Na) = +496, ½ $\Delta H$ _diss( $Cl_{2}$ ) = +121, EA(Cl) = −349 kJ/mol → U = −787 kJ/mol.

Fajan's Rules predict the degree of covalent character within ionic bonds. The cation polarizes (distorts) the electron cloud of the anion. Covalent character is greater when: (1) the cation is small and highly charged (high polarizing power), and (2) the anion is large (highly polarizable). Mnemonic: "Small Cat, Big Annie, High Charge = Covalent."

Covalent Bonding and Dipole Moment

Covalent bonds result from shared electron pairs. The dipole moment (μ = q × d, unit: Debye; 1 D = $3.336 \times 10^{-30}$ C·m) measures the polarity of a molecule as a whole. Symmetric molecules — $CO_{2}$ (linear), $BF_{3}$ (trigonal planar), $CCl_{4}$ (tetrahedral), $SF_{6}$ (octahedral) — have μ = 0 because individual bond dipoles cancel by vector addition. Asymmetric molecules ( $NH_{3}$ , $H_{2}O$ , $CHCl_{3}$ ) have non-zero dipole moments due to incomplete cancellation and/or lone-pair contribution.

VSEPR Theory

VSEPR predicts molecular geometry by minimizing electron-pair repulsions around the central atom. The steric number (SN) = number of sigma bonds + lone pairs on the central atom, which directly determines the hybridization and electron geometry. The actual molecular shape is obtained by ignoring lone pairs:

SN	Hybridization	Electron Geometry	Molecular Shape (example)
2	sp	Linear	Linear — Be $Cl_{2}$ , $CO_{2}$
3	$sp^{2}$	Trigonal planar	Trigonal planar ( $BF_{3}$ ) or Bent ( $SO_{2}$ )
4	$sp^{3}$	Tetrahedral	Tetrahedral ( $CH_{4}$ ), Pyramidal ( $NH_{3}$ ), Bent ( $H_{2}O$ )
5	$sp^{3}$ d	Trigonal bipyramidal	TBP ( $PCl_{5}$ ), See-saw ( $SF_{4}$ ), T-shape ( $ClF_{3}$ ), Linear ( $XeF_{2}$ )
6	$sp^{3}$$d^{2}$	Octahedral	Octahedral ( $SF_{6}$ ), Sq. pyramidal ( $BrF_{5}$ ), Sq. planar ( $XeF_{4}$ )

Lone-pair repulsion follows: lp-lp > lp-bp > bp-bp. Each lone pair compresses bond angles: $CH_{4}$ (109.5°) → $NH_{3}$ (107°) → $H_{2}O$ (104.5°).

$XeF_{4}$ worked example: Xe has 8 valence electrons; bonds to 4 F; lone pairs = (8 − 4×1 − 4×1)/2... Using expanded octet: 4 bond pairs + 2 lone pairs = SN 6 → $sp^{3}$$d^{2}$ hybridization. Both lone pairs occupy axial positions to minimize mutual repulsion → square planar shape.

$XeF_{2}$ worked example: Xe bonds to 2 F; 3 lone pairs; SN = 5 → $sp^{3}$ d. Lone pairs in equatorial positions → linear shape.

Valence Bond Theory (VBT)

VBT describes bonds as overlapping atomic/hybrid orbitals. Sigma (σ) bonds form by head-on overlap — they allow free rotation. Pi (π) bonds form by lateral (sideways) overlap — they restrict rotation and are weaker. Single bond = 1σ; double bond = 1σ + 1π; triple bond = 1σ + 2π.

Molecular Orbital Theory (MOT)

MOT constructs molecular orbitals by linear combination of atomic orbitals (LCAO). Each pair of AOs forms one bonding MO (lower energy) and one antibonding MO (higher energy, marked with *).

Bond order = (Nb − Na) / 2, where Nb = electrons in bonding MOs and Na = electrons in antibonding MOs. Higher bond order → shorter bond length → greater bond energy.

Critical distinction — MO filling order depends on atomic number Z:

Z ≤ 7 ( $B_{2}$ , $C_{2}$ , $N_{2}$ ): π2p orbitals fill BEFORE σ2p (due to 2s–2p mixing)
Z > 7 ( $O_{2}$ , $F_{2}$ ): σ2p fills BEFORE π2p (normal order)

Key species from NEET perspective: $O_{2}$ has bond order 2 and is paramagnetic (2 unpaired electrons in degenerate π*2p MOs). VBT incorrectly predicts $O_{2}$ as diamagnetic — a landmark triumph of MOT.

$O_{2}^{-}$ (superoxide): BO = (10 − 7)/2 = 1.5, paramagnetic (1 unpaired $e^{-}$ ) $O_{2}^{2-}$ (peroxide): BO = (10 − 8)/2 = 1, diamagnetic

Resonance and Hydrogen Bonding

Resonance describes electron delocalization when a single Lewis structure is inadequate (e.g., benzene, ozone). The actual molecule is a resonance hybrid of all contributing structures. Delocalization lowers energy and increases stability. Hydrogen bonding (H attached to F, O, or N) abnormally raises boiling points (e.g., $H_{2}O$ vs $H_{2}S$ ) and is responsible for water's unique properties.