Chemical Bonding and Molecular Structure
'Atoms are held together by the desire to achieve the noble gas configuration.' — Kossel and Lewis
1. Chapter Overview
Why do atoms BOND? And what DETERMINES the shape of molecules? This chapter answers these questions. It covers the LEWIS theory of bonding, IONIC and COVALENT bonds, VSEPR THEORY (for predicting molecular shapes), HYBRIDISATION (mixing of orbitals), MOLECULAR ORBITAL THEORY (MO), and HYDROGEN BONDING. Understanding bonding is ESSENTIAL for predicting chemical reactivity.
2. Kossel-Lewis Approach to Chemical Bonding
- Octet Rule: Atoms combine to achieve an OUTER SHELL of 8 electrons (noble gas configuration)
- Lewis Symbols: Represent valence electrons as dots around the element symbol
- Lewis Structures: Show HOW atoms are bonded and where LONE PAIRS are
Formal Charge
- Formal Charge = Valence e⁻ — [Lone pairs + ½(Bonding e⁻)]
- The MOST STABLE Lewis structure has formal charges CLOSEST to zero
Limitations of Octet Rule
- Incomplete octet (H, Be, B, Al) — BeCl₂, BF₃
- Expanded octet (P, S, Cl, Xe) — PCl₅, SF₆, XeF₄
- Odd-electron molecules — NO₂, NO
3. Ionic (Electrovalent) Bond
- Formation: COMPLETE transfer of electrons from one atom to another
- Conditions: Low IE (metal) + High EG (non-metal) + High lattice energy
- Properties: High melting point, soluble in water, conduct electricity in molten state
Factors Affecting Ionic Bond Strength
- Lattice Energy ∝ (Charge₁ × Charge₂)/(r₁ + r₂)
- Higher charge → Stronger bond (MgO > NaCl)
- Smaller ions → Stronger bond (LiF > CsI)
4. Covalent Bond
- Formation: SHARING of electrons between atoms
- Types:
- Single bond (σ bond): 1 shared pair
- Double bond (σ + π): 2 shared pairs
- Triple bond (σ + 2π): 3 shared pairs
Bond Parameters
| Parameter | Definition | Units |
|---|---|---|
| Bond Length | Distance between bonded nuclei | pm or Å |
| Bond Angle | Angle between two bonds | degrees |
| Bond Enthalpy | Energy required to break 1 mole of bonds | kJ/mol |
| Bond Order | ½(Bonding e⁻ — Antibonding e⁻) | Unitless |
Polarity of Bonds
- Non-polar covalent: Equal sharing (H₂, Cl₂, N₂)
- Polar covalent: Unequal sharing (HCl, H₂O)
- Ionic: Complete transfer (NaCl)
- Electronegativity difference determines bond type:
- ΔEN < 0.4: Non-polar covalent
- 0.4 < ΔEN < 1.7: Polar covalent
- ΔEN > 1.7: Ionic
5. VSEPR Theory (Valence Shell Electron Pair Repulsion)
Basic Principle
- Electron pairs (bonded AND lone pairs) arrange themselves to MINIMISE repulsion
- Lone pair — Lone pair > Lone pair — Bond pair > Bond pair — Bond pair
Molecular Shapes
| Total e⁻ Pairs | Lone Pairs | Geometry | Shape | Example |
|---|---|---|---|---|
| 2 | 0 | Linear | Linear | BeCl₂, CO₂ |
| 3 | 0 | Trigonal planar | Trigonal planar | BF₃ |
| 3 | 1 | Trigonal planar | Bent | SO₂, O₃ |
| 4 | 0 | Tetrahedral | Tetrahedral | CH₄ |
| 4 | 1 | Tetrahedral | Trigonal pyramidal | NH₃ |
| 4 | 2 | Tetrahedral | Bent | H₂O |
| 5 | 0 | Trigonal bipyramidal | Trigonal bipyramidal | PCl₅ |
| 5 | 1 | Trigonal bipyramidal | Seesaw | SF₄ |
| 6 | 0 | Octahedral | Octahedral | SF₆ |
6. Hybridisation
- Definition: Mixing of ATOMIC ORBITALS to form EQUIVALENT HYBRID ORBITALS
Types of Hybridisation
| Type | Orbitals Mixed | Geometry | Bond Angle | Examples |
|---|---|---|---|---|
| sp | 1s + 1p | Linear | 180° | BeCl₂, C₂H₂ |
| sp² | 1s + 2p | Trigonal planar | 120° | BF₃, C₂H₄ |
| sp³ | 1s + 3p | Tetrahedral | 109.5° | CH₄, NH₃, H₂O |
| dsp² | 1d + 1s + 2p | Square planar | 90° | [Ni(CN)₄]²⁻ |
| sp³d | 1s + 3p + 1d | Trigonal bipyramidal | 90°, 120° | PCl₅ |
| sp³d² | 1s + 3p + 2d | Octahedral | 90° | SF₆ |
Worked Problem
Q: Determine hybridisation and shape of NH₃. A: N has 5 valence e⁻, 3 bond pairs (H) + 1 lone pair = 4 electron pairs → sp³ hybridisation. Shape: TRIGONAL PYRAMIDAL (lone pair pushes bonds in).
7. Molecular Orbital (MO) Theory
Key Concepts
- Atomic orbitals combine to form MOLECULAR ORBITALS (bonding σ/π and antibonding σ*/π*)
- Number of MOs = Number of atomic orbitals combined
- Bond Order = (N_b — N_a)/2
- MOs are filled following Aufbau, Pauli, and Hund's rules
Electronic Configuration of Homonuclear Diatomics
- For molecules up to N₂ (2p orbitals): σ1s < σ1s < σ2s < σ2s < π2p_x = π2p_y < σ2p_z < π2p_x = π2p_y < σ*2p_z
- For O₂, F₂: Order of σ2p_z and π2p_x/y SWAPS
Magnetic Behaviour
- All e⁻ paired → DIAMAGNETIC (repelled by magnetic field)
- Unpaired e⁻ → PARAMAGNETIC (attracted by magnetic field)
- O₂ is PARAMAGNETIC (2 unpaired electrons in π* orbitals) — MO theory explains this!
Worked Problem
Q: Calculate bond order of O₂ and N₂. A: N₂ (14 e⁻): σ1s² σ1s² σ2s² σ2s² π2p⁴ σ2p². BO = (10 — 4)/2 = 3. O₂ (16 e⁻): σ1s² σ1s² σ2s² σ2s² σ2p² π2p⁴ π*2p². BO = (10 — 6)/2 = 2.
8. Hydrogen Bonding
- Definition: ATTRACTIVE force between H atom (bonded to high EN atom: F, O, N) and a LONE PAIR on another EN atom
- Intermolecular H-bonding: Between DIFFERENT molecules (H₂O, HF, NH₃)
- Intramolecular H-bonding: Within the SAME molecule (o-nitrophenol)
Effects of H-Bonding
- Higher boiling points (H₂O > H₂S despite same group)
- Ice is LESS dense than water (open structure due to H-bonds)
- Proteins and DNA structure (alpha helices, double helix)
9. Common Mistakes
- Molecules with lone pairs are NOT the same shape as the electron pair geometry: NH₃ is trigonal PYRAMIDAL (not tetrahedral), H₂O is BENT (not tetrahedral)
- Bond order ≠ number of bonds: For O₂, bond order = 2, but O₂ has a DOUBLE bond with 2 unpaired electrons
- Odd-electron species cannot be represented by Lewis structures properly: NO, NO₂
- Intramolecular H-bonding LOWERS boiling point (no inter-molecular association)
- d-orbitals participate in hybridisation ONLY for heavier elements (n ≥ 3)
10. CBSE Exam Focus
- Lewis dot structures of molecules (3-mark)
- VSEPR theory — predicting shapes (5-mark)
- Hybridisation of central atom in molecules (3/5-mark)
- Molecular orbital theory — bond order calculations (5-mark)
- Magnetic behaviour using MO theory (3-mark)
- Hydrogen bonding — types and effects (3-mark)
11. Key Formulas
- Bond Order = (N_b — N_a)/2
- Dipole Moment μ = q × d (Debye units)
- % Ionic character = (μ_observed / μ_calculated) × 100%
- Formal Charge = V — (L + B/2)
12. Self-Test (5+ Q&A)
Q1: Determine hybridisation and shape of SF₄. A: S has 6 valence e⁻, 4 bond pairs + 1 lone pair = 5 pairs → sp³d hybridisation. Shape: SEE-SAW (not trigonal bipyramidal!).
Q2: Calculate bond order and predict magnetic behaviour of O₂²⁻ (peroxide ion). A: O₂²⁻ has 18 e⁻. BO = (10 — 8)/2 = 1. All e⁻ paired → DIAMAGNETIC.
Q3: Why is H₂O bent but CO₂ linear? A: H₂O: O has 2 bond pairs + 2 lone pairs → sp³ → BENT. CO₂: C has 2 double bonds, NO lone pairs → sp → LINEAR.
Q4: What is hydrogen bond strength compared to covalent bond? A: H-bonds (10-40 kJ/mol) are MUCH WEAKER than covalent bonds (200-400 kJ/mol) but STRONGER than van der Waals forces.
Q5: Explain why O₂ is paramagnetic using MO theory. A: O₂ has 16 e⁻. The π*2p orbitals contain 2 e⁻ in DIFFERENT orbitals with PARALLEL SPINS (Hund's rule). These unpaired electrons cause paramagnetism.
13. Conclusion
Chemical bonding EXPLAINS why compounds form and what shapes they take. Lewis theory provides a STARTING POINT. VSEPR predicts molecular shapes from electron pair repulsion. Hybridisation explains EQUIVALENT bonds in molecules like CH₄. MO theory gives the MOST COMPLETE picture — explaining bond order, magnetic behaviour, and spectroscopy. Hydrogen bonding, while WEAK, has enormous consequences for water's properties and biological molecules. Bonding concepts are ESSENTIAL for understanding organic chemistry and chemical reactions.
