Polarity and intermolecular forces
When two atoms share electrons unequally, the bond is polar. The more electronegative atom gets a partial negative charge (δ-) and the other a partial positive charge (δ+).
The difference in electronegativity gives a rough guide:
| Difference | Bond type |
|---|---|
| Less than about 0.5 | Nonpolar covalent |
| About 0.5 to 1.7 | Polar covalent |
| More than about 1.7 | Ionic |
The cutoffs are approximate. Bonding changes gradually, not in steps.
A molecule is polar if it has polar bonds that do not cancel. Shape decides this. CO2 has two polar bonds pointing in opposite directions, so it is nonpolar overall. H2O is bent, so its polar bonds add up and the molecule is polar.
Intermolecular forces are attractions between molecules. They are much weaker than bonds, but they set melting and boiling points.
- London dispersion forces act between all molecules and grow with molecular size.
- Dipole-dipole forces act between polar molecules.
- Hydrogen bonds are strong dipole attractions involving H bonded to N, O or F.
Stronger forces mean a higher boiling point.
Worked example: why does water boil at 100 °C but methane at -162 °C?
Water molecules are polar and form hydrogen bonds. Methane is nonpolar and has only weak dispersion forces, so far less energy is needed to separate its molecules.
Common mistake
Thinking boiling breaks the bonds inside molecules. Boiling water separates whole H2O molecules from each other. The bonds between O and H stay intact.