Chemical Bonding – NEET PG Study Notes
(Chemistry – 10 th‑12 th & Pre‑Clinical)
Goal: Understand the nature, types, and characteristics of chemical bonds.
Why it matters: All medicines, biochemistry, and pharmacology rely on the way atoms hold together.
1. Simple Explanation
A chemical bond is a force of attraction that holds atoms together in a molecule or crystal lattice.
- Electrons (negative) are shared, transferred, or delocalised among atoms.
- The balance of electron density determines whether a bond is ionic (full transfer), covalent (sharing), or metallic (delocalised sea of electrons).
- Additional interactions – hydrogen bonds, van der Waals forces, coordinate (dative) bonds – fine‑tune molecular structure and reactivity.
2. Key Points
| Concept | What it is | How it appears in molecules | Key Take‑Aways |
|---|---|---|---|
| Electronegativity (χ) | Ability of an atom to attract electrons | Determines bond polarity | χ difference > 1.7 → ionic; 0.4–1.7 → polar covalent; < 0.4 → non‑polar covalent |
| Bond Polarity | Unequal sharing of electrons | Creates dipole moments (e.g., H₂O) | Dipole–dipole, hydrogen bonding |
| Octet Rule | Atoms seek 8 valence electrons | Forms ionic/covalent bonds | Exceptions: H (2), d‑block, expanded octet |
| Hybridization | Mixing of atomic orbitals | Predicts molecular shape (sp, sp², sp³) | Key for VSEPR geometry |
| VSEPR Theory | Repulsion of electron pairs dictates shape | Predicts bond angles (e.g., 104.5° in H₂O) | Good for simple molecules |
| Resonance | Delocalised electrons over multiple structures | Stabilises molecules (e.g., benzene) | Resonance energy > 0 |
| Molecular Orbital (MO) Theory | Combines atomic orbitals into MOs (bonding/antibonding) | Explains bond order, magnetism | Bond order = (n_bonding – n_antibonding)/2 |
| Ionic vs. Covalent | Transfer vs. sharing of electrons | NaCl (ionic), H₂O (covalent) | Ionic: high lattice energy; Covalent: bond dissociation energy |
| Metallic Bonding | Delocalised “sea” of electrons | Conductivity, malleability | Bond energy ≈ 1–3 eV per atom |
| Coordinate (Dative) Bond | One atom donates both electrons | Complexes (e.g., NH₃·BH₃) | Often in Lewis acid–base reactions |
| Hydrogen Bond | H–X (X = N, O, F) with lone pair on another atom | Stabilises DNA, water structure | Strength: 5–30 kJ mol⁻¹ |
| Van der Waals Forces | London dispersion, dipole–dipole | Weak, short‑range | Important in polymer physics |
3. Worked Example – Formation of a Water Molecule (H₂O)
-
Electron Count & Octet Rule
- H: 1 electron → needs 1 more (2e⁻ total)
- O: 6 valence electrons → needs 2 more (8e⁻ total)
-
Lewis Structure
H | H : O : H- Two lone pairs on O (4 e⁻).
- Two O–H σ bonds (2 e⁻ each).
-
Bond Polarity
- χ(O)=3.44, χ(H)=2.20 → Δχ≈1.24 → polar covalent
- O has partial negative (δ⁻), H partial positive (δ⁺).
-
Hybridization & Geometry
- O uses sp³ hybrid orbitals → tetrahedral electron geometry.
- Two lone pairs occupy 2 positions → bent molecular shape.
- Bond angle ≈ 104.5° (slightly less than tetrahedral 109.5° due to lone‑pair repulsion).
-
Dipole Moment
- Vector addition of two O–H dipoles → net dipole ≈ 1.85 D (dipole moment of water).
-
Bond Dissociation Energy (BDE)
- O–H BDE ≈ 463 kJ mol⁻¹.
-
Hydrogen Bonding (optional)
- H₂O can donate two H bonds (via H atoms) and accept two (via lone pairs).
Key Take‑Aways from Example
- Octet rule guides Lewis structure.
- Electronegativity difference → bond polarity.
- Hybridization predicts shape and bond angles.
- Dipole moment arises from arrangement of polar bonds.
4. Quick Revision Summary
| Topic | Remember | Quick Check |
|---|---|---|
| Electronegativity | χ = 0 (F) → 4.0 (O) | χ diff > 1.7 → ionic |
| Bond Types | Ionic, covalent (polar/non‑polar), metallic | Identify by χ diff |
| Octet Rule | 8 valence e⁻ |