Review: Types of Intermolecular Forces
Intermolecular forces (IMFs) are attractive forces between molecules — not within molecules (those are intramolecular bonds such as covalent bonds). The strength of the IMF present in a substance determines its physical properties: boiling point, viscosity, surface tension, vapour pressure, and more.
| IMF Type |
Found In |
Relative Strength |
Example |
| London Dispersion Forces (LDF) |
All molecules |
Weakest (increases with molar mass) |
Noble gases, CH₄, Cl₂ |
| Dipole-dipole |
Polar molecules |
Medium |
HCl, SO₂, acetone |
| Hydrogen bonding |
Molecules with N–H, O–H, or F–H bonds |
Strongest IMF |
H₂O, NH₃, HF, ethanol |
Effect of IMF Strength on Boiling Point
Boiling is the process of converting a liquid to a gas — it requires enough energy to overcome the IMFs holding molecules together. Therefore: stronger IMFs → higher boiling point. Larger molecules also have more electrons, leading to stronger LDF and higher boiling points even in non-polar substances.
Hydrogen halide boiling point anomaly: The general trend from HCl to HI is HCl < HBr < HI — LDF increases with molar mass. BUT HF has an anomalously high BP (19.5°C) despite having the lowest molar mass in the group, because HF molecules form hydrogen bonds (F–H···F). This much stronger IMF requires far more energy to overcome.
Group 16 hydrides show a similar anomaly:
| Molecule |
Boiling Point (°C) |
Dominant IMF |
| H₂O |
100 |
Hydrogen bonding (anomalously high) |
| H₂Te |
−2 |
LDF + dipole-dipole |
| H₂Se |
−41 |
LDF + dipole-dipole |
| H₂S |
−60 |
LDF + dipole-dipole |
Noble gases show a clean LDF trend — only London dispersion forces, all increasing with molar mass:
| Noble Gas | BP (°C) |
|---|
| He | −269 |
| Ne | −246 |
| Ar | −186 |
| Kr | −153 |
| Xe | −108 |
Hydrides of period 2 elements (H₂O, HF, NH₃) all have anomalously HIGH boiling points due to hydrogen bonding.
Other Physical Properties
Viscosity is the resistance of a liquid to flow. Stronger IMFs hold molecules more tightly, making flow harder — so stronger IMF → higher viscosity. Glycerol has three –OH groups, forming extensive hydrogen bonds, making it very viscous. Water is more viscous than ethanol for the same reason.
Surface tension is the energy required to increase the surface area of a liquid. Molecules at the surface experience a net inward force from the molecules below. Stronger IMF → higher surface tension. Water has very high surface tension due to hydrogen bonding — enough for pond skaters (water striders) to walk on the surface.
Vapour pressure is the pressure of vapour in equilibrium above a liquid in a closed container. If IMFs are weak, molecules escape to the vapour phase more easily, so weaker IMF → higher vapour pressure. Highly volatile liquids such as diethyl ether have high vapour pressure and low boiling points.
Capillary action depends on the balance between adhesion (attraction between liquid and container wall) and cohesion (attraction between liquid molecules). Water in a glass tube: adhesion > cohesion → the meniscus curves upward and water rises. Mercury in a glass tube: cohesion > adhesion → the meniscus curves downward and mercury dips.
Anomalous Properties of Water
Water's unusual properties all arise from extensive hydrogen bonding between H₂O molecules. Each water molecule can form up to four hydrogen bonds (two as donor, two as acceptor).
- High boiling point (100°C) — much higher than expected for such a small molecule (molar mass 18 g·mol⁻¹).
- High specific heat capacity (4 200 J·kg⁻¹·K⁻¹) — water resists temperature change, stabilising Earth's climate and maintaining stable body temperatures in living organisms.
- Ice is less dense than liquid water — in ice, each H₂O molecule forms four H-bonds in a rigid, open hexagonal lattice. This structure has more space between molecules than liquid water, so ice is less dense and floats on water. This is ecologically critical: lakes and ponds freeze from the top down, and the floating ice layer insulates the water below, allowing aquatic life to survive winter.
- High surface tension — pond skaters and other insects can walk on the water surface.
- Universal solvent — the polar water molecule dissolves polar and ionic solutes. Ion-dipole forces form hydration shells around ions (e.g. Na⁺ and Cl⁻ when NaCl dissolves).
Why ice floats: The fact that ice is LESS dense than liquid water is unique among substances. When water freezes, the H-bond lattice forces molecules into fixed positions with more space between them than in the denser liquid. This is why lakes freeze from the surface down — the ice layer insulates the water below, allowing aquatic life to survive winter.
Solubility and "Like Dissolves Like"
The principle like dissolves like states that a solvent dissolves solutes with similar polarity:
- Polar solvent (water) dissolves polar molecules and ionic compounds. NaCl dissociates into Na⁺ and Cl⁻, each surrounded by hydration shells of water molecules via ion-dipole forces.
- Non-polar solvent (hexane, CCl₄) dissolves non-polar molecules such as wax, oils, and iodine — held together by LDF of similar strength.
- Miscible liquids mix in all proportions (water + ethanol — both polar and capable of H-bonding).
- Immiscible liquids do not mix (water + oil — very different polarities; water molecules H-bond with each other, excluding non-polar oil).
Vapour Pressure and Boiling Point Connection
A liquid boils when its vapour pressure equals the external (atmospheric) pressure. This explains:
- Cooking at altitude: lower atmospheric pressure means water's vapour pressure equals it at a lower temperature — water boils below 100°C and food takes longer to cook.
- Pressure cooker: the sealed pot raises the pressure above atmospheric, so water must reach a higher temperature before its vapour pressure equals the external pressure — water boils above 100°C and food cooks faster.
IEB Extension: Colligative Properties & Raoult's Law
Colligative properties depend on the NUMBER of solute particles dissolved, not their identity. Key examples:
- Boiling point elevation: ΔTb = Kb × m, where m = molality (mol·kg⁻¹). Adding solute raises the BP.
- Freezing point depression: ΔTf = Kf × m. Adding solute lowers the freezing point (antifreeze in car radiators).
- Raoult's Law: Psoln = Xsolvent × P°solvent, where X = mole fraction of solvent. Adding a solute reduces the mole fraction of solvent, lowering its vapour pressure — and therefore raising its boiling point.
- Supercritical fluids: Above the critical temperature and critical pressure, a substance exists as neither liquid nor gas — a supercritical fluid. CO₂ supercritical fluid is used as a "green" solvent in the decaffeination of coffee.