Chemical Bonding and Molecular Structure
1. Why Do Atoms Combine?
Atoms combine to attain a more stable electronic configuration, usually resembling that of noble gases.
This stability is achieved by:
- Loss of electrons
- Gain of electrons
- Sharing of electrons
2. Kossel–Lewis Concept
Kossel explained ionic bonding while Lewis explained covalent bonding using valence electrons.
Valence electrons are responsible for chemical bonding.
3. Ionic Bond (Electrovalent Bond)
An ionic bond is formed due to complete transfer of electrons from one atom to another.
Example:
$$\text{Na} \rightarrow \text{Na}^+ + e^-$$
$$\text{Cl} + e^- \rightarrow \text{Cl}^-$$
Ionic compounds are hard, crystalline and have high melting points.
4. Covalent Bond
A covalent bond is formed by mutual sharing of electrons between atoms.
Single bond → 1 shared pair
Double bond → 2 shared pairs
Triple bond → 3 shared pairs
5. Lewis Dot Structures
Lewis structures show valence electrons as dots around atomic symbols.
Octet rule: Atoms tend to have 8 electrons in valence shell (exceptions exist).
6. Formal Charge
Formal charge =
$$\text{Valence electrons} - \text{Nonbonding electrons} - \frac{1}{2}(\text{Bonding electrons})$$
Lowest formal charge structure is the most stable.
7. Resonance
Resonance occurs when more than one Lewis structure can be drawn for a molecule.
Actual molecule is a resonance hybrid.
8. VSEPR Theory
VSEPR theory predicts molecular shape by minimizing repulsion between electron pairs.
Order of repulsion:
$$\text{LP–LP} > \text{LP–BP} > \text{BP–BP}$$
9. Molecular Geometry (VSEPR)
| Type | Shape | Example |
|---|---|---|
| AX₂ | Linear | BeCl₂ |
| AX₃ | Trigonal planar | BF₃ |
| AX₄ | Tetrahedral | CH₄ |
| AX₃E | Trigonal pyramidal | NH₃ |
| AX₂E₂ | Bent | H₂O |
10. Hybridisation
Hybridisation is the mixing of atomic orbitals to form equivalent hybrid orbitals.
| Hybridisation | Geometry | Example |
|---|---|---|
| sp | Linear | BeCl₂ |
| sp² | Trigonal planar | BF₃ |
| sp³ | Tetrahedral | CH₄ |
11. Valence Bond Theory (VBT)
According to VBT, bond formation occurs due to overlap of half-filled atomic orbitals.
12. Types of Orbital Overlap
$\sigma$ bond → head-on overlap
$\pi$ bond → sidewise overlap
13. Molecular Orbital Theory (MOT)
Atomic orbitals combine to form molecular orbitals spread over the entire molecule.
Bond Order =
$$\frac{1}{2}(N_b - N_a)$$
If bond order = 0, molecule does not exist.
14. Magnetic Nature
If molecule has unpaired electrons → Paramagnetic
If all electrons are paired → Diamagnetic
15. Bond Length and Bond Energy
Bond length decreases with increase in bond order.
Bond energy increases with bond strength.
16. Polarity of Bonds
Bond polarity depends on electronegativity difference.
Greater electronegativity difference → more polar bond.
17. Dipole Moment
$\mu = q \times r$
Net dipole moment depends on molecular geometry.
18. Hydrogen Bonding
Hydrogen bonding occurs when H is bonded to highly electronegative atoms like F, O, or N.
| Type | Example |
|---|---|
| Intermolecular | H₂O |
| Intramolecular | o-Nitrophenol |
19. Effects of Hydrogen Bonding
- High boiling point
- High viscosity
- Unusual solubility
20. Common Exam Traps
- Confusing shape and hybridisation
- Ignoring lone pair effects
- Wrong bond order calculation
- Forgetting resonance structures
21. Final Revision Checklist
You have mastered this chapter if you can:
- Draw correct Lewis structures
- Predict shape using VSEPR
- Identify hybridisation instantly
- Calculate bond order
- Determine polarity and magnetism