Willer Select – NEET Special Batch
Chemical Bonding Quick Revision (Lectures 1–5)
Lewis Structures · VSEPR · Hybridization · MOT · Bond Parameters
Lecture 1: Lewis Structures & Chemical Bonding
1. Why Chemical Bonds Form?
- Atoms bond to achieve stability (noble gas-like configuration)
- Driven by energy minimization: more stable = lower energy
2. Octet Rule & Exceptions
| Concept | Explanation | Examples |
|---|---|---|
| Core Principle | Gain/lose/share to get 8 valence electrons | Na+, Cl- |
| Exceptions |
|
BeCl2, SF6, NO |
3. Types of Chemical Bonds
| Type | Formation | Properties | Examples |
|---|---|---|---|
| Ionic | Electron transfer | High MP/BP, soluble in water | NaCl |
| Covalent | Electron sharing | Low MP/BP, bad conductor | CH4, H2O |
| Coordinate | One atom donates both e- | Directional, polar | NH4+ |
4. Lewis Structure Drawing Steps
- Count total valence e- for all atoms
- Identify central atom (least EN, never H/F)
- Connect atoms with single bonds
- Complete octets of terminal atoms, add extras to center
- Form multiple bonds if needed for central atom's octet
- CO2: O═C═O (16e-)
- NH4+: All H–N–H bonds, N (8e-), H (2e- each)
- CO32-: Resonance forms, 24e-
5. Formal Charge (FC)
- FC = (Valence e-) – (Non-bonding e-) – ½(Bonding e-)
- Best Lewis structure has FC values closest to zero
6. Common Mistakes to Avoid
- Forgetting to add/subtract for ions
- Ignoring resonance and expanded octet
Practice: Draw Lewis structures for: (a) SO2 (b) N2 (c) BF3 (d) NO3-
Lecture 2: VSEPR Theory & Molecular Geometry
1. What is VSEPR?
- Electron pairs (bonding/lone) repel to minimize repulsion, shaping molecule
2. Key Terms
| Term | Definition |
|---|---|
| Steric Number (SN) | Total electron domains (bonds + lone pairs) |
| Lone Pair (LP) | Non-bonding pair on central atom |
3. Predicting Geometry: 5 Steps
- Draw Lewis Structure
- Count electron domains
- Use SN to assign Electron Geometry (see table)
- Ignore LPs to get Molecular Geometry
- Adjust for LP-LP > LP-BP > BP-BP repulsion
| Steric No. | Electron Geometry | Angles | Molecular Geometry (with LPs) | Example |
|---|---|---|---|---|
| 2 | Linear | 180° | Linear | BeCl2 |
| 3 | Trigonal Planar | 120° | Bent (1 LP, e.g. SO2) | SO2 |
| 4 | Tetrahedral | 109.5° | Pyramidal/Bent | NH3, H2O |
| 5 | Trigonal Bipyramidal | 90°,120° | See-saw/T-shape/Linear | SF4, XeF2 |
| 6 | Octahedral | 90° | Square pyramidal/planar | XeF4 |
Tip: Memorize all geometries for Steric Number 2-6 and corresponding angles!
VSEPR Flow:
Lewis Structure → Steric Number → Electron Geometry → Lone Pairs → Molecular Geometry
Practice: Predict shape and bond angle:
(a) BF3 (b) SF6 (c) PCl5 (d) I3-
Common errors: Confusing electron/molecular geometry, neglecting LP effects.
Lecture 3: Valence Bond Theory (VBT) & Hybridization
1. VBT Core Idea
- Covalent bonds form by overlap of half-filled atomic orbitals
- Bond directionality depends on axes of overlap (geometry comes from this)
2. Bond Types
| Type | Overlap | Rotation | Strength | Examples |
|---|---|---|---|---|
| Sigma (σ) | Head-on | Free | Strong | H2, Cl2 |
| Pi (π) | Side-on | Restricted | Weak | O2, C2H4 |
3. Hybridization: Why & How
- Explains molecule shape not predictable by VBT
- e.g. CH4 is tetrahedral because of sp3 mixing
4. Hybridization Table
| Type | Orbitals | Geometry | Angle | Examples |
|---|---|---|---|---|
| sp | 1s+1p | Linear | 180° | BeCl2, CO2 |
| sp2 | 1s+2p | Trigonal planar | 120° | BF3 |
| sp3 | 1s+3p | Tetrahedral | 109.5° | CH4 |
| dsp2 | 1d+1s+2p | Square planar | 90° | [Ni(CN)4]2- |
| d2sp3 | 2d+1s+3p | Octahedral | 90° | SF6 |
5. How to Identify Hybridization
- Draw Lewis Structure
- Determine Steric Number = σ bonds + lone pairs
- Match SN: 2→sp, 3→sp2, 4→sp3, 5→dsp3, 6→d2sp3
Practice: Identify hybridization & bond types: (a) C2H2 (b) NH3 (c) XeF4 (d) CO32-
NEET Tip: Map hybridization–shape–angle for fast recall!
Lecture 4: Molecular Orbital Theory (MOT)
1. Core Idea of MOT
- Atomic orbitals form molecular orbitals (MOs) that cover the entire molecule
- Aufbau, Pauli, Hund's rules apply to fill MOs
- Explains paramagnetism of O2, resonance, bond order trends better than VBT
2. Main Rules: LCAO-Method
- Linear Combination of Atomic Orbitals (LCAO) for MO construction
- Bonding MOs (σ/π): lower energy, Antibonding MOs (σ\*/π\*): higher energy (node/star)
3. Typical MO Energy Patterns
| Molecule | MO Order | Bond Order | Magnetism |
|---|---|---|---|
| H2 | σ(1s) < σ\*(1s) | 1 | Dia |
| He2 | σ(1s) < σ\*(1s) | 0 | Dia |
| B2 | π(2p) < σ(2p) | 1 | Para |
| O2 | σ(2p) < π(2p) < π\*(2p) | 2 | Para |
4. Bond Order Formula & Stability
- BO = ½ [bonding e- – antibonding e-], BO > 0 = stable, higher BO = stronger bond
5. Magnetism in MOT
| Type | Spin State | Example |
|---|---|---|
| Paramagnetic | Unpaired | O2, B2 |
| Diamagnetic | All paired | N2, H2 |
6. Heteronuclear Diatomics (CO, NO)
- MO energies closer to more EN atom
- NO: BO=2.5, paramagnetic; CO: BO=3, diamagnetic
Practice:
BO for: (a) O2+ (b) N2+ (c) F2
Magnetism: (a) C2 (b) O2-
Magnetism: (a) C2 (b) O2-
Common mistakes: Using N2 MO order for O2/F2; forgetting antibonding e-
Memorize MO diagrams for O2 & N2; these are NEET favourites!
Lecture 5: Bond Parameters & Polarity
1. Bond Length
| Factor | Effect | Example |
|---|---|---|
| Atom Size | Larger atom → longer bond | C–F < C–I |
| Bond Order | Higher BO → shorter bond | C–C > C=C > C≡C |
| Hybridization | More s-character → shorter bond | C(sp3)–H > C(sp)–H |
2. Bond Angle
- Lone pairs decrease angle: CH4 109.5° > NH3 107° > H2O 104.5°
- Substituent E.N.: More EN = smaller angle (NF3 < NH3)
3. Bond Energy
| Concept | Explanation | Example |
|---|---|---|
| Definition | Energy to break 1 mole of bonds (gas) | H–H: 436 kJ/mol |
| Bond Order | Higher order = stronger bond | C≡C: 839 > C–C: 347 |
| Trend | Decreases down group | H–F > H–Cl > H–Br |
4. Electronegativity (EN)
Pauling: F (4.0) > O (3.5) > Cl (3.0) > C (2.5) > H (2.1)
| ΔEN | Bond Type | Example |
|---|---|---|
| 0–0.4 | Nonpolar covalent | C–H |
| 0.5–1.6 | Polar covalent | O–H |
| ≥1.7 | Ionic | NaCl |
5. Dipole Moment (μ)
- μ = charge × separation, units: Debye
- Symmetrical = μ = 0 (CO2, BF3), asym = Polar (H2O: 1.85 D)
| Molecule | Shape | μ (D) |
|---|---|---|
| CO2 | Linear | 0 |
| H2O | Bent | 1.85 |
| NH3 | Trigonal pyramidal | 1.47 |
6. % Ionic Character
- % Ionic = [1 – e–0.25(ΔEN2)] × 100
- NaCl: 67% ionic; HCl: 19% ionic
7. Resonance & Bond Length
- Resonance equalizes bonds (O3: 128 pm; between single and double)
Practice:
1. Arrange by bond length: (a) C–N, C≡N, C=N (b) C–O, C=O, C≡O
2. Predict polarity: (a) BF3 (b) NF3 (c) CH2Cl2
2. Predict polarity: (a) BF3 (b) NF3 (c) CH2Cl2
Common mistakes:
- Assuming all linear molecules are nonpolar (e.g., HF is polar!)
- Ignoring effect of LPs on bond angles
- Confusing bond energy with dissociation energy
NEET Tip: Memorize ΔEN thresholds and μ (H2O, NH3, CO2)!
Next: Practice & Revision – NCERT questions and NEET-style MCQs.
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