Intermolecular forces (IMFs) are the attractions between molecules. They determine whether a substance is a gas, liquid, or solid at room temperature, as well as its boiling point, melting point, viscosity, and many other physical properties. Water owes its remarkable properties to unusually strong hydrogen bonds between its molecules.
7.1 Types of Intermolecular Forces
Definition
Intermolecular forces
Intermolecular forces are forces of attraction and repulsion that act between neighbouring molecules (or between molecules and ions). They are significantly weaker than covalent, ionic, or metallic bonds (intramolecular forces). Intermolecular forces determine the physical properties of substances such as melting point, boiling point, and viscosity.
Intermolecular forces are electrostatic attractions between molecules (not between atoms within a molecule — those are INTRAMOLECULAR forces such as covalent bonds). IMFs are MUCH weaker than covalent bonds. It is important to distinguish: breaking a covalent bond requires much more energy than overcoming an intermolecular force.
LONDON DISPERSION FORCES (also called induced dipole–induced dipole forces): These are the weakest IMFs and are present in ALL molecules — polar and non-polar alike. They arise because electrons in a molecule are constantly moving, creating temporary (instantaneous) uneven distributions of charge — temporary dipoles. These temporary dipoles induce dipoles in neighbouring molecules, creating a weak, momentary attraction. London forces increase with: (1) increasing molar mass — more electrons → stronger temporary dipoles; (2) increasing surface area of the molecule.
DIPOLE-DIPOLE FORCES: These occur between POLAR molecules. The partially positive end (δ+) of one polar molecule attracts the partially negative end (δ−) of an adjacent polar molecule. Dipole-dipole forces are stronger than London dispersion forces for molecules of similar size. They are the dominant IMF in polar molecules that lack O–H, N–H, or F–H bonds (e.g. HCl, SO₂, acetone).
HYDROGEN BONDING: This is a special, especially strong type of dipole-dipole force that occurs when a hydrogen atom is covalently bonded to a HIGHLY electronegative atom (N, O, or F — remember 'NOF'). Because H is so small and the bond to N/O/F is so polar, the δ+ on H is very concentrated. It strongly attracts the lone pair on the N, O, or F of an adjacent molecule. Hydrogen bonds are roughly 10× stronger than typical dipole-dipole forces. Examples: H₂O, NH₃, HF, alcohols, DNA base pairs.
ION-DIPOLE FORCES: These occur between an ION (full charge, e.g. Na⁺ or Cl⁻) and a POLAR molecule. The ion attracts the oppositely charged end of the polar molecule's dipole. Ion-dipole forces are the strongest of the IMFs. They are responsible for the dissolution of ionic compounds in polar solvents such as water (e.g. NaCl dissolving in H₂O: Na⁺ attracts the δ− oxygen end of water molecules).
Intermolecular Forces — Weakest to Strongest
| Property | Type of IMF | Occurs between ... / Examples |
|---|---|---|
| 1. London dispersion (weakest) | All molecules (polar and non-polar) | Noble gases, alkanes, Br₂, I₂ |
| 2. Dipole-dipole | Polar molecules (no N–H, O–H, F–H) | HCl, SO₂, acetone (CH₃COCH₃) |
| 3. Hydrogen bonding | Molecules with N–H, O–H, or F–H bonds | H₂O, NH₃, HF, alcohols, DNA |
| 4. Ion-dipole (strongest) | Between ions and polar molecules | NaCl in water (Na⁺ / Cl⁻ + H₂O) |
Watch Out
COMMON MISTAKE — Hydrogen bonding does NOT involve molecular hydrogen (H₂). It is a bond between the δ+ hydrogen on one molecule and a lone pair on N, O, or F of another molecule. Also: hydrogen bonds are INTERMOLECULAR forces, NOT covalent bonds — they are much weaker than covalent bonds.
Worked Example
Identify the strongest type of intermolecular force present in each of the following: (a) CH₄, (b) HCl, (c) H₂O, (d) NaCl dissolved in water.
Given
- (a) CH₄ — non-polar molecule
- (b) HCl — polar molecule, no N–H / O–H / F–H
- (c) H₂O — polar molecule with O–H bonds
- (d) NaCl in H₂O — ions in a polar solvent
Find
Dominant IMF in each case
Solution
- 1(a) CH₄ is non-polar → only London dispersion forces.
- 2(b) HCl is polar, but H is bonded to Cl (not N, O, or F) → dipole-dipole forces.
- 3(c) H₂O has O–H bonds → hydrogen bonding (the O–H qualifies for H-bonding).
- 4(d) Na⁺ and Cl⁻ ions interact with polar H₂O molecules → ion-dipole forces.
Practice Question
Consider the following molecules: N₂, HF, PCl₃, and C₂H₆. (a) For each molecule, identify the strongest type of intermolecular force. (b) Rank the molecules from lowest to highest boiling point and justify your ranking in terms of IMF strength.
(8 marks)
7.2 Anomalous Properties of Water
Water is the most important liquid on Earth and it has several properties that are surprisingly different from what you would predict based on its small molar mass (18 g·mol⁻¹). These ANOMALOUS properties all result from the unusually strong HYDROGEN BONDS between water molecules.
Anomalous properties of water explained by hydrogen bonding
- HIGH BOILING POINT (100°C): H₂O has a much higher bp than H₂S (−60°C) or H₂Se (−41°C) despite lower molar mass — because the many hydrogen bonds require much more energy to break.
- HIGH SPECIFIC HEAT CAPACITY (4 200 J·kg⁻¹·K⁻¹): Water can absorb large amounts of heat energy with only a small temperature rise — hydrogen bonds store energy. This stabilises climate and body temperature.
- HIGH HEAT OF VAPORISATION: A large amount of energy is needed to vaporise water (break the hydrogen bonds) — which is why sweating is so effective at cooling the body.
- ICE IS LESS DENSE THAN LIQUID WATER: In ice, water molecules form a rigid hexagonal lattice with all H-bonds intact, holding molecules slightly further apart than in liquid water. Liquid water is denser than ice — ice floats, insulating aquatic organisms in winter.
- HIGH SURFACE TENSION: The strong hydrogen bonds between surface water molecules create a skin-like effect that allows small insects to walk on water.
Real World
BIOLOGICAL IMPORTANCE: The fact that ice floats is critical for life. In winter, lakes freeze from the surface down. The ice layer insulates the liquid water below, preventing it from freezing solid and allowing aquatic life to survive. If water behaved like most liquids (solid denser than liquid), lakes would freeze from the bottom up and all aquatic life would die.
Practice Question
Explain why water has a much higher boiling point (100°C) than hydrogen sulfide, H₂S (−60°C), even though H₂S has a larger molar mass.
(6 marks)
7.3 Physical Properties and IMF Strength
Definition
Density
A measure of the mass in a unit volume. ρ = m/V
Definition
Specific heat capacity (c)
Specific heat capacity is the amount of thermal energy required to raise the temperature of 1 kg of a substance by 1°C (or 1 K). Q = mcΔT, where m is mass in kg, c is the specific heat capacity in J·kg⁻¹·°C⁻¹, and ΔT is the temperature change. Water has an unusually high specific heat capacity of 4 200 J·kg⁻¹·°C⁻¹.
Definition
Heat of vaporisation
The energy that is needed to change a given quantity of a substance into a gas.
BOILING POINT AND MELTING POINT: These depend directly on the strength of the IMFs. To boil a liquid, the molecules must gain enough kinetic energy to overcome the IMFs holding them together. Stronger IMFs → higher boiling point. Similarly, to melt a solid, the particles must break free of their positions — stronger IMFs (or stronger interionic forces) → higher melting point.
VISCOSITY is the resistance of a liquid to flow. A high-viscosity liquid is 'thick' (e.g. honey, glycerol). Viscosity depends on the strength of IMFs between molecules AND the molecular size and shape. Molecules with strong IMFs flow past each other with more difficulty → higher viscosity. Example: glycerol (many O–H groups, many hydrogen bonds) has much higher viscosity than water.
TRENDS IN BOILING POINT WITHIN A HOMOLOGOUS SERIES: For alkanes (London dispersion forces only), the boiling point increases steadily with chain length (molar mass). Longer chains have more electrons → stronger temporary dipoles → stronger London forces → more energy needed to boil. Branching reduces the boiling point because it reduces the surface area available for London interactions.
Effect of IMF Strength on Physical Properties
| Property | Stronger IMFs | Weaker IMFs |
|---|---|---|
| Boiling point | Higher (more energy to overcome IMFs) | Lower |
| Melting point | Higher | Lower |
| Viscosity | Higher (molecules don't flow past each other easily) | Lower |
| Heat of vaporisation | Higher (more energy to break all IMFs) | Lower |
| Vapour pressure (at same T) | Lower (molecules less likely to escape) | Higher |
Exam Tip
EXAM TIP — When comparing boiling points, always state: (1) what type of IMF each substance has, (2) which is stronger and why, and (3) therefore which has the higher boiling point. A three-step logical argument scores full marks.
Worked Example
The boiling points of propane (C₃H₈, Mr = 44) and ethanol (C₂H₅OH, Mr = 46) are −42°C and +78°C respectively. Explain this large difference in boiling point despite the very similar molar masses.
Given
- Propane: C₃H₈, Mr = 44, bp = −42°C, non-polar molecule
- Ethanol: C₂H₅OH, Mr = 46, bp = +78°C, has O–H group
Find
Explanation of the large difference in boiling point
Solution
- 1Propane is non-polar — it only has London dispersion forces (weakest IMF).
- 2Ethanol has an O–H group — it can form HYDROGEN BONDS (strongest IMF after ion-dipole).
- 3Hydrogen bonds are much stronger than London dispersion forces.
- 4Therefore, much more energy (much higher temperature) is needed to provide ethanol molecules with enough kinetic energy to overcome hydrogen bonds and boil.
- 5Despite similar molar mass, the type of IMF is completely different → large boiling point difference.
Practice Question
The table below shows data for three substances: | Substance | Formula | Boiling point | |---|---|---| | Methane | CH₄ | −161°C | | Propane | C₃H₈ | −42°C | | Pentane | C₅H₁₂ | +36°C | All three are non-polar molecules. (a) What type of IMF is present in all three? (b) Explain the trend in boiling point from methane to pentane. (c) Predict whether hexane (C₆H₁₄) would have a higher or lower boiling point than pentane and justify your answer.
(8 marks)