Chemical Bonding (NEET Chemistry)

Why This Chapter Matters

Chemical bonding is fundamental for NEET Chemistry — 6-10 marks. VSEPR theory, hybridisation, molecular orbital theory, and bond parameters are all tested.

Core Concepts

1. Types of Bonds

Ionic: Metal + Non-metal. Electron transfer. High mp, conducts in solution.

Covalent: Non-metal + Non-metal. Electron sharing. Usually low mp.

Coordinate/Dative: Both electrons from same atom. [NH₄]⁺, H₃O⁺, [Cu(NH₃)₄]²⁺

Metallic: Electron sea. Conductors.

Hydrogen bond: X-H···Y where X,Y = F,O,N. Explains high bp of HF, H₂O, NH₃.

F-H···F strongest hydrogen bond.

2. VSEPR Theory

Electron pairs repel each other. Lone pairs > bond pairs (in terms of repulsion).

Electron pairs	Molecular shape	Example

|---|---|---|

2 bond, 0 lone	Linear	BeCl₂, CO₂
3 bond, 0 lone	Trigonal planar	BF₃
4 bond, 0 lone	Tetrahedral	CH₄
3 bond, 1 lone	Pyramidal	NH₃ (107°)
2 bond, 2 lone	Bent/V-shaped	H₂O (104.5°)
5 bond, 0 lone	Trigonal bipyramidal	PCl₅
4 bond, 1 lone	See-saw	SF₄
6 bond, 0 lone	Octahedral	SF₆
4 bond, 2 lone	Square planar	XeF₄

3. Hybridisation

Hybridisation	Geometry	Bond angle	Examples

|---|---|---|---|

sp	Linear	180°	BeCl₂, C₂H₂, CO₂
sp²	Trigonal planar	120°	BF₃, C₂H₄, benzene, NO₃⁻
sp³	Tetrahedral	109.5°	CH₄, NH₃, H₂O, CCl₄
sp³d	Trigonal bipyramidal	90°,120°	PCl₅
sp³d²	Octahedral	90°	SF₆, [Co(NH₃)₆]³⁺

4. Bond Parameters

Bond order: higher → shorter bond length, higher bond energy

Single > Double > Triple bond length (decreasing)

Triple > Double > Single bond energy (increasing strength)

Bond polarity: difference in electronegativity → dipole moment (μ = q × d)

Molecular polarity: vector sum of all dipole moments

Non-polar despite polar bonds: CO₂ (linear, dipoles cancel), CCl₄ (tetrahedral, cancel), BF₃ (trigonal planar, cancel)

Polar: H₂O, NH₃, CHCl₃ (dipoles don't cancel)

5. Molecular Orbital Theory

Atomic orbitals → Molecular Orbitals

Bonding MO: lower energy (stabilise) | Antibonding MO (σ, π): higher energy (destabilise)

Filling order: σ1s, σ1s, σ2s, σ2s, π2p (= two π2p), σ2p, π2p (= two), σ2p

Bond order = (Nb - Na)/2

Stability: BO > 0 | Paramagnetic if unpaired electrons

Important examples:

H₂: BO = 1, diamagnetic

He₂: BO = 0, doesn't exist

O₂: BO = 2, paramagnetic (2 unpaired e in π*2p)

N₂: BO = 3, most stable diatomic

PYQs (NEET)

NEET 2023: Shape and hybridisation of XeF₄?

Xe has 8 valence electrons. 4 bonds + 2 lone pairs = 6 electron pairs → sp³d² hybridisation

2 lone pairs in axial positions → Square planar molecular shape

NEET 2022: Which molecule has maximum bond angle?

CO₂: linear → 180° (maximum)

NEET 2021: O₂ molecule is paramagnetic. This is best explained by:

Molecular Orbital Theory (2 unpaired electrons in π*2p antibonding orbitals)

Revision Notes

VSEPR: Count total electron pairs (bond + lone) → geometry
Lone pairs cause more repulsion → reduce bond angles

HYBRIDISATION = bond pairs + lone pairs on central atom
sp: linear | sp²: trigonal planar | sp³: tetrahedral
sp³d: trig. bipyramidal | sp³d²: octahedral

NOTABLE SHAPES:
NH₃: pyramidal, 107° (1 LP) | H₂O: bent, 104.5° (2 LP)
PCl₅: sp³d, trig. bipyramidal | XeF₄: sp³d², square planar

MOLECULAR ORBITAL:
BO = (Nb-Na)/2 | Paramagnetic if unpaired electrons
H₂: BO=1 | N₂: BO=3 | O₂: BO=2 (paramagnetic!)
He₂: BO=0 (doesn't exist)

HYDROGEN BONDING:
F>O>N (fluorine forms strongest H-bond)
Explains: high bp of H₂O, HF, NH₃; anomalous properties of water

Chemical Bonding and Molecular Structure