Chemical Bonding & Molecular Structure
(IIT‑JEE / Competitive Exam – Chemistry)
These notes are crafted to give you a deep, exam‑oriented understanding of the key concepts, derivations, and subtle points that frequently appear in IIT‑JEE, JEE‑Advanced, and other competitive exams.
1. Overview
| Category | Key Concept | Representative Models | Typical Exam Tricky Points |
|---|---|---|---|
| Bond Types | Ionic, Covalent, Metallic, Hydrogen, Van der Waals | Lewis, VSEPR, Hybridisation, MO Theory, FMO | Distinguishing polar‑covalent vs. ionic character; “partial” charges. |
| Structural Prediction | Geometry, Bond Order, Dipole Moment | VSEPR, Hybridisation, MO, FMO | 3‑centre 4‑electron bonds; d‑orbital participation; pseudo‑hypervalency. |
| Bond Energies | Dissociation, Enthalpy, Free energy | Bond strength tables, O‑K plot | Extrapolating from heteronuclear diatomics. |
2. Ionic Bonding
| Feature | Details |
|---|---|
| Electrostatic interaction | (U = -\frac{Z_+ Z_- e^2}{4\pi\varepsilon_0 r}) |
| Lattice energy | Born–Landé equation: (U = -\frac{N_A M z^+ z^- e^2}{4\pi\varepsilon_0 r_0}\left(1-\frac{1}{n}\right)) |
| Parameters | (M) (Madelung constant), (n) (Born exponent), (r_0) (nearest‑neighbour distance) |
| Tricky Point | Even “ionic” compounds like NaCl can exhibit covalent character when the electronegativity difference is < 1.7. |
3. Covalent Bonding
| Concept | Key Equations / Derivations |
|---|---|
| Bond Polarity | (\Delta \chi = \chi_A - \chi_B) 0–0.4 → non‑polar, 0.4–1.7 → polar covalent, >1.7 → ionic |
| Bond Dipole | (\mu = \Delta q , d) where (\Delta q = \Delta \chi \times 0.3,e) (approx.) |
| Bond Dissociation Energy (BDE) | (D_0) = energy required to dissociate a molecule to ground‑state atoms. |
| Resonance Energy | Empirical: (E_{\text{res}} = \frac{E_{\text{sum of structures}} - E_{\text{actual}}}{\text{# resonance structures}}) |
| Tricky Point | Resonance structures are not real molecules. They are a bookkeeping device to represent delocalised electrons. |
4. Hybridisation & VSEPR
4.1 Basic Hybridisation
| Hybrid Orbital | Composition | Resulting Geometry |
|---|---|---|
| (sp^3) | (1s + 3p) | Tetrahedral (109.5°) |
| (sp^2) | (1s + 2p) | Trigonal planar (120°) |
| (sp) | (1s + 1p) | Linear (180°) |
Derivation of bond angles
Using dot product of hybrid orbitals, e.g. for (sp^3):
[ \cos\theta = \frac{\mathbf{h}_i \cdot \mathbf{h}_j}{|\mathbf{h}_i||\mathbf{h}_j|} = -\frac{1}{3};\Rightarrow;\theta = 109.5^\circ ]
4.2 VSEPR (Valence Shell Electron Pair Repulsion)
| Electron Geometry | Molecular Geometry | Examples |
|---|---|---|
| 4 pairs | Tetrahedral | CH₄ |
| 3 pairs | Trigonal planar | BF₃ |
| 3 pairs + 1 lone pair | Trigonal pyramidal | NH₃ |
| 4 pairs + 1 lone pair | Tetra‑hedral with a lone pair | H₂O |
Tricky Points
- Lone pairs repel more strongly than bonding pairs → distortion of ideal angles.
- 3‑centre 4‑electron (3c‑4e) bonds in hypervalent molecules (e.g., (\text{Cl}_3^-)) – VSEPR still works but the geometry is not purely tetrahedral.
- D‑orbital participation in (sp^3d) or (sp^2d) hybridisation – often over‑used; real bonding may involve (\pi) overlap instead.
5. Molecular Orbital (MO) Theory
5.1 Construction of MO Diagrams
- Atomic Orbitals (AOs) combine to form MOs (bonding ↓, antibonding ↑).
- Symmetry: SALCs for poly