Chemical Bonding: VSEPR, VBT & MOT

Chemical bonding determines molecular geometry, polarity, and reactivity. JEE Main tests three complementary theories: VSEPR (geometry prediction), VBT (orbital overlap and hybridisation), and MOT (bond order and magnetic behaviour).

Ionic vs Covalent Continuum: Fajans' rules quantify covalent character in ionic bonds. High covalent character arises from: small cation size, large anion size, high cation charge, pseudo-noble-gas configuration (e.g., Cu+, Ag+). LiI is more covalent than LiF; BeCl₂ is more covalent than MgCl₂.

VSEPR Theory: Electron pairs (bonding + lone pairs) around a central atom arrange to minimise repulsion. Repulsion order: lp-lp > lp-bp > bp-bp. The steric number (SN = bonding pairs + lone pairs on central atom) determines electron geometry. Molecular geometry depends on how many of those pairs are lone pairs.

SN = 2: linear (180 degrees). SN = 3: trigonal planar (120 degrees). SN = 4: tetrahedral (109.5 degrees). SN = 5: trigonal bipyramidal ($\frac{90}{120}$ degrees). SN = 6: octahedral (90 degrees). SN = 7: pentagonal bipyramidal ($\frac{72}{90}$ degrees).

Key VSEPR examples — linear vs bent geometry:

CO₂ — linear (SN=2, no lone pairs on C)

H₂O — bent geometry (SN=4, 2 lone pairs compress bond angle to 104.5 degrees)

Lone pairs occupy equatorial positions in TBP (they need more room). In octahedral arrangements with 2 lone pairs, they go trans to each other (square planar, e.g., XeF₄).

Valence Bond Theory (VBT): Covalent bonds form by overlap of half-filled (or empty) atomic orbitals. Hybridisation is the mathematical mixing of orbitals to produce equivalent hybrid orbitals directed in space. Key hybridisations: sp (linear, BeCl₂, C₂H₂), sp2 (trigonal planar, BF₃, C₂H₄), sp3 (tetrahedral, CH₄, NH₃, H₂O), sp3d (TBP, PCl₅), sp3d2 (octahedral, SF₆), sp3d3 (PBP, IF₇).

Hybridisation examples — sp vs sp2 vs sp3:

Acetylene (C₂H₂) — sp hybridised carbons, linear, 180 degrees

Ethylene (C₂H₄) — sp2 hybridised carbons, trigonal planar, 120 degrees

Methane (CH₄) — sp3 hybridised carbon, tetrahedral, 109.5 degrees

$\Sigma$ bonds form by head-on overlap; $\pi$ bonds by lateral overlap. In a double bond: 1 $\sigma$ + 1 $\pi$. In a triple bond: 1 $\sigma$ + 2 $\pi$.
Hybridisation of an atom = $\sigma$ bonds + lone pairs on that atom (the steric number).

Molecular Orbital Theory (MOT): Atomic orbitals combine to form molecular orbitals (MOs). Number of MOs = number of AOs combined. Bonding MOs are lower in energy; antibonding MOs are higher. Bond order = (Nb - Na)/2, where Nb = electrons in bonding MOs, Na = electrons in antibonding MOs. If bond order = 0, the molecule is unstable.

For homonuclear diatomics up to N₂ (Z <= 7), the $\sigma$(2p) MO is HIGHER in energy than $\pi$(2p) MOs due to s-p mixing.
For O₂ and beyond (Z >= 8), $\sigma$(2p) is LOWER than $\pi$(2p). This is the critical ordering switch that explains O₂'s paramagnetism.

Dipole Moment and Polarity: $\mu$ = q * d (Debye units). Molecular polarity depends on both bond polarity AND geometry. Symmetric molecules (BF₃, CCl₄, XeF₄) have zero net dipole despite polar bonds. Water (bent, 104.5 degrees) and NH₃ (pyramidal, 107 degrees) have non-zero dipole moments.

Hydrogen Bonding: Strongest intermolecular force (5-40 kJ/mol). Requires H bonded to F, O, or N (highly electronegative, small atoms). Intramolecular H-bonding (e.g., o-nitrophenol) reduces boiling point relative to intermolecular H-bonding (p-nitrophenol).

Intramolecular vs intermolecular hydrogen bonding:

o-Nitrophenol — intramolecular H-bond (lower boiling point)

p-Nitrophenol — intermolecular H-bond (higher boiling point)

Coordinate (Dative) Bonds: Both electrons in the shared pair come from one atom (donor). Examples: NH₄+ (N donates to H+), BF₃.NH₃, CO (C donates lone pair to form triple bond), O₃ (central O donates to terminal O).