Bonding, Shapes and Intermolecular Forces
Types of bonding
- Ionic: electrostatic attraction between oppositely charged ions (metal + non-metal); giant lattice.
- Covalent: a shared pair of electrons (non-metals). A dative (coordinate) bond is a shared pair where both electrons come from one atom.
- Metallic: positive ions in a sea of delocalised electrons.
Electronegativity and bond polarity
Electronegativity is an atom's ability to attract the bonding electrons. It increases across a period and up a group (fluorine is the highest).
- If two bonded atoms differ in electronegativity, the bond is polar (δ+ and δ−).
- A molecule is polar overall only if the polar bonds don't cancel by symmetry (e.g. H₂O is polar; CO₂ is not, because it's linear and symmetrical).
Shapes of molecules — VSEPR
Electron pairs repel and arrange themselves as far apart as possible. Count the bonding pairs and lone pairs around the central atom:
| Electron pairs | Shape | Bond angle |
|---|---|---|
| 2 | Linear | 180° |
| 3 | Trigonal planar | 120° |
| 4 | Tetrahedral | 109.5° |
| 5 | Trigonal bipyramidal | 90°/120° |
| 6 | Octahedral | 90° |
Lone pairs repel more than bonding pairs, reducing the angle by ~2.5° each. E.g. NH₃ (one lone pair) = 107°; H₂O (two lone pairs) = 104.5°.
Intermolecular forces (weakest to strongest)
1. Van der Waals (London/dispersion) forces — temporary induced dipoles; present in all molecules; stronger with more electrons.
2. Permanent dipole–dipole — between polar molecules.
3. Hydrogen bonding — the strongest; occurs when H is bonded to N, O or F. Explains water's high boiling point, ice being less dense than water, and high boiling points of alcohols.
Intermolecular forces (not the covalent bonds) are what break when a simple molecular substance melts or boils.
Worked example
Explain why H₂O has a much higher boiling point than H₂S, despite S being larger.
- Water molecules form hydrogen bonds (H bonded to O), which are much stronger than the permanent dipole–dipole and van der Waals forces in H₂S, so more energy is needed to separate water molecules. ✓
Common mistakes
- Forgetting lone pairs when predicting shape and bond angle.
- Saying a molecule with polar bonds is always polar — check the symmetry (CO₂ is non-polar).
- Confusing the strength of hydrogen bonding with covalent bonds (H-bonds are intermolecular, weaker).
Exam tips
- Use VSEPR: count bonding + lone pairs, then give the shape and angle (subtract for lone pairs).
- Explain boiling points using the type and strength of intermolecular forces.
- Know hydrogen bonding requires H–N, H–O or H–F.
Key facts to remember
- Bonding: ionic, covalent (incl. dative), metallic; electronegativity difference → polar bonds (polar overall depends on symmetry).
- VSEPR: electron pairs repel to give shapes (linear 180°, trigonal planar 120°, tetrahedral 109.5°…); lone pairs repel more.
- Intermolecular forces: van der Waals < dipole–dipole < hydrogen bonding (H–N/O/F); these break on melting/boiling of simple molecules.