Chemical Bonding: MDCAT Notes and MCQs
VSEPR theory
Electron pairs around the central atom repel and arrange as far apart as possible. Repulsion order: lone pair–lone pair > lone pair–bond pair > bond pair–bond pair, so lone pairs compress bond angles.
| Bond pairs + lone pairs | Shape | Example (angle) |
|---|---|---|
| 2 + 0 | Linear | BeCl2, CO2 (180°) |
| 3 + 0 | Trigonal planar | BF3 (120°) |
| 2 + 1 | Bent / angular | SnCl2, SO2 (< 120°) |
| 4 + 0 | Tetrahedral | CH4 (109.5°) |
| 3 + 1 | Trigonal pyramidal | NH3 (107.5°) |
| 2 + 2 | Bent | H2O (104.5°) |
| 5 + 0 | Trigonal bipyramidal | PCl5 (90°, 120°) |
| 6 + 0 | Octahedral | SF6 (90°) |
Sigma and pi bonds
- σ bond: head-on (axial) overlap; electron density along the internuclear axis; stronger; free rotation possible; formed first. Every single bond is σ.
- π bond: sideways (lateral) overlap of parallel p orbitals; electron density above and below the axis; weaker, more reactive; restricts rotation. Double bond = 1σ + 1π; triple = 1σ + 2π.
Hybridization
| Type | Geometry | Angle | Examples |
|---|---|---|---|
| sp3 | Tetrahedral | 109.5° | CH4, C2H6, NH3, H2O |
| sp2 | Trigonal planar | 120° | BF3, C2H4, benzene |
| sp | Linear | 180° | BeCl2, C2H2 |
Polarity & dipole moment
- Dipole moment μ = q × r; unit Debye (1 D = 3.336 × 10−30 C m). Vector quantity.
- Symmetrical molecules are non-polar even with polar bonds: CO2, BF3, CH4, CCl4 (μ = 0). Polar: H2O (1.85 D), NH3 (1.46 D), HCl (1.03 D), CHCl3.
- Uses: predicting shape (H2O must be bent since μ ≠ 0; CO2 linear since μ = 0), solubility ("like dissolves like"), boiling points, reactivity of polar bonds.
- Ionic character: unequal sharing in a polar covalent bond gives partial ionic character; electronegativity difference ≈ 1.7 → ≈ 50% ionic. % ionic character = (μobserved / μionic) × 100.
Bond energy
Energy required to break one mole of bonds in gaseous molecules into atoms (kJ/mol). Higher bond energy → stronger, shorter bond. C≡C (837) > C=C (611) > C–C (348). H–F > H–Cl > H–Br > H–I. Increases with bond order and polarity; decreases with atomic size.
PMDC learning outcomes for this unit
- VSEPR, hybridization, polarity, bond energy: 10.1 Use VSEPR theory to describe the shapes of the molecules.
- VSEPR, hybridization, polarity, bond energy: 10.2 Describe the features of sigma and pi bonds.
- VSEPR, hybridization, polarity, bond energy: 10.3 Describe the shapes of simple molecules using orbital hybridization.
- VSEPR, hybridization, polarity, bond energy: 10.4 Determine the shapes of some molecules from the number of the bonded pairs.
- VSEPR, hybridization, polarity, bond energy: 10.5 Predict the molecular polarity from the shapes of molecules.
- VSEPR, hybridization, polarity, bond energy: 10.6 Explain what is meant by the term ionic character of the covalent bond.
- VSEPR, hybridization, polarity, bond energy: 10.7 Describe how knowledge of molecular polarity can be used to explain some physical and chemical properties of the molecules.
- VSEPR, hybridization, polarity, bond energy: 10.8 Define bond energies and explain how they can be used to compare bond strength of different chemical bonds.