Intermolecular Forces

Learning Objectives (OCR Spec 2.2.2)

By the end of this lesson you should be able to:

• Describe the three types of intermolecular force required by OCR: London (induced dipole-dipole), permanent dipole-dipole, and hydrogen bonding
• Explain the origin of each type of force
• Rank intermolecular forces by strength
• Use intermolecular forces to predict and explain trends in boiling point, solubility and viscosity
• Identify molecules that can form hydrogen bonds

Why Intermolecular Forces Matter

Covalent bonds hold atoms together within a molecule (intramolecular). Intermolecular forces act between separate molecules. When a substance melts or boils, you break intermolecular forces, not covalent bonds. That is why water boils at 100 degC but breaking an O-H bond needs ~460 kJ mol-1.

The Three Main Types

graph TD
  A[Intermolecular Forces] --> B[London dispersion forces]
  A --> C[Permanent dipole-dipole]
  A --> D[Hydrogen bonding]
  B -.->|Present in all substances| E[Weakest]
  C -.->|Polar molecules only| F[Medium]
  D -.->|N-H, O-H, F-H only| G[Strongest type of IMF]

Relative strengths (typical values)

Type	Strength range (kJ mol-1)
London dispersion forces	1-10 (can be much higher for large molecules)
Permanent dipole-dipole	5-25
Hydrogen bonding	10-40
Covalent bond (for comparison)	150-1000+

All intermolecular forces are much weaker than covalent bonds.

1. London Dispersion Forces

Also called induced dipole-dipole forces, van der Waals forces, or dispersion forces.

Origin

At any instant the electrons in a molecule are not perfectly symmetrically distributed - there is a random, transient instantaneous dipole. This tiny dipole induces an opposite dipole in a neighbouring molecule, and the two are attracted.

These fluctuating dipoles are continually appearing, disappearing and shifting, so on average they provide a weak attractive force between all molecules.

Key Facts

• Present in all molecules (polar or non-polar)
• The only intermolecular force between non-polar molecules
• Strength increases with:
  • Number of electrons (larger molecules, greater molar mass)
  • Surface area of contact (straight-chain alkanes have stronger London forces than branched isomers because the molecules can pack closer together)

Trend Example: Alkanes

Alkane	Mr	Boiling point (degC)
CH4	16	-162
C4H10	58	-0.5
C8H18	114	126
C16H34	226	287

Larger alkanes have more electrons and therefore stronger London forces, requiring more energy to separate them from neighbours.

Trend Example: Halogens

Halogen	State at r.t.	Boiling point
F2	gas	-188 degC
Cl2	gas	-34 degC
Br2	liquid	59 degC
I2	solid	184 degC

All non-polar, but increasing numbers of electrons give stronger London forces down the group.

2. Permanent Dipole-Dipole Forces

Origin

Polar molecules have a permanent dipole (delta+ and delta- charges). The delta+ end of one molecule attracts the delta- end of a neighbouring molecule.

   delta+   delta-     delta+   delta-
     H --- Cl            H --- Cl
                          \----^
                       dipole-dipole
                       attraction

Key Facts

• Present only in polar molecules (those with a net dipole moment)
• Operate in addition to London forces
• Typically 5-25 kJ mol-1

Example: HCl vs F2

Both have similar molar masses (36.5 vs 38), but HCl (polar) boils at -85 degC while F2 (non-polar) boils at -188 degC. The additional dipole-dipole interactions in HCl raise the boiling point significantly.

3. Hydrogen Bonding

Origin

Hydrogen bonding is a particularly strong type of dipole-dipole interaction that arises when H is covalently bonded to a small, highly electronegative atom with a lone pair: specifically N, O or F. The H is very delta+ because:

• Its only electron is pulled strongly toward the electronegative atom
• It has no inner shells to shield the nucleus
• It is very small (so it can approach the lone pair closely)