Inside vs between: bonds and intermolecular forces
A chemical bond (covalent, ionic, metallic) holds atoms together inside a substance's particles. An intermolecular force is an attraction between separate molecules.
Bonds are strong (hundreds of kJ per mole). Intermolecular forces are weak (usually 1–40 kJ per mole). When water boils, the H–O bonds do not break; only the attractions between molecules break. Steam is still H₂O.
All these forces come from electric charges: slightly negative parts (δ−) attract slightly positive parts (δ+).
Polarity: where do δ+ and δ− come from?
Electronegativity is how strongly an atom pulls shared electrons. When two different atoms share electrons, the stronger puller gets δ− and the other gets δ+. That bond is polar.
A whole molecule is polar only if its shape does not cancel the pulls. CO₂ is straight, so its two polar bonds cancel: non-polar. H₂O is bent, so they do not cancel: polar. CCl₄ is symmetric (non-polar); CHCl₃ is not (polar).
Functional groups matter: –OH, –NH₂ and –COOH make a molecule able to hydrogen bond; long C–H chains are non-polar.
The three main types
London dispersion forces
Electrons are always moving. For an instant, more of them are on one side, making a temporary dipole. It pushes the electrons of the neighbour and makes an induced dipole. They attract. All molecules have London forces. They grow with the number of electrons (bigger molecules) and with long, flat shapes that touch more (pentane boils higher than its ball-shaped isomer).
Dipole–dipole forces
Polar molecules have permanent δ+ and δ− ends. They line up + to − (HCl, acetone). For molecules of similar size, these add to London forces, so polar substances boil higher.
Hydrogen bonds
A special, strong dipole attraction. It needs H bonded to N, O or F, and a lone pair on another N, O or F. Examples: water, ammonia, HF, alcohols, DNA base pairs. Hydrogen bonds explain why ice floats: they hold ice molecules in an open pattern, so ice is less dense than liquid water.
Ion–dipole forces (Na⁺ surrounded by water) are stronger still and explain why salt dissolves in water.
'Van der Waals forces' is a common name for London plus dipole–dipole forces.
How the forces set properties
- Melting and boiling point: stronger forces → higher. H₂O 100 °C, H₂S −60 °C (no H-bond), CH₄ −162 °C.
- Group trend: for noble gases and halogens, boiling point rises down the group (more electrons → stronger London forces). I₂ is a solid at room temperature; F₂ is a gas.
- Viscosity and surface tension: glycerol (three –OH groups) is thick; water has high surface tension.
- Vapour pressure and evaporation: weaker forces → evaporates faster (acetone dries quicker than water).
- Solubility: 'like dissolves like'. Polar dissolves in polar (sugar, salt in water); non-polar in non-polar (grease in petrol). Oil and water separate.
Key formulas and definitions
- Strength (similar size): London < dipole–dipole < hydrogen bond < ion–dipole
- Hydrogen bond needs: H–N, H–O or H–F + a lone pair on N, O or F
- More electrons → stronger London forces
- Stronger intermolecular forces → higher m.p., b.p., viscosity, surface tension; lower vapour pressure
- Like dissolves like
Worked examples
1. Which has the higher boiling point, F₂ or Cl₂? Why?
Cl₂. Both are non-polar, so only London forces act. Cl₂ has more electrons, so its London forces are stronger.
2. Name the strongest intermolecular force in (a) CH₄ (b) HCl (c) NH₃.
(a) London only (non-polar). (b) Dipole–dipole (polar, but H is on Cl, not N/O/F). (c) Hydrogen bonding (H on N).
3. Why does ethanol (C₂H₅OH, 78 °C) boil much higher than dimethyl ether (CH₃OCH₃, −24 °C)? Same formula C₂H₆O.
Ethanol has an O–H group, so its molecules hydrogen bond. The ether has no H on O, so only weaker dipole–dipole and London forces act.
4. Why is CO₂ a gas but H₂O a liquid at room temperature, though CO₂ is heavier?
CO₂ is linear, so its bond dipoles cancel: only London forces. H₂O is bent and polar and forms hydrogen bonds, which are much stronger.
5. Does iodine dissolve better in water or in hexane?
Hexane. I₂ is non-polar; it mixes with the non-polar solvent (like dissolves like).
Common mistakes
- Thinking boiling breaks the bonds inside molecules. It only separates molecules; steam is still H₂O.
- Calling any molecule with H a hydrogen bonder. The H must be on N, O or F (CH₄ and HCl do not H-bond).
- Saying non-polar molecules have no attraction at all. All molecules have London forces; big ones like I₂ can even be solid.
- Judging polarity by bonds only. Symmetric shapes (CO₂, CCl₄) cancel polar bonds and make the molecule non-polar.