Intermolecular Forces

This lesson covers the three types of intermolecular forces: London dispersion forces, permanent dipole-dipole interactions and hydrogen bonding. Understanding these forces is essential for explaining physical properties such as boiling points, solubility and viscosity (AQA specification 3.1.3).

---

What Are Intermolecular Forces?

Key Definition: Intermolecular forces are the forces of attraction between molecules. They are much weaker than covalent bonds (intramolecular forces) but determine the physical properties of covalent substances.

There are three types, in order of increasing strength:

1. London dispersion forces (also called van der Waals forces or induced dipole-dipole forces)
2. Permanent dipole-dipole forces
3. Hydrogen bonding (a special, strong type of permanent dipole interaction)

Important: When a covalent substance boils, it is the intermolecular forces that are broken, not the covalent bonds within the molecules.

---

1. London Dispersion Forces

London forces exist between all molecules (polar and non-polar). They are the only intermolecular force between non-polar molecules.

How They Arise

1. At any instant, the electrons in a molecule are unevenly distributed, creating a temporary (instantaneous) dipole
2. This temporary dipole induces a dipole in a neighbouring molecule
3. The two temporary dipoles attract each other
4. These temporary dipoles are constantly forming, disappearing and re-forming

Instantaneous dipole:     δ+  δ−       δ+  δ−
                        [molecule A] [molecule B]
                              ←attraction→

Factors Affecting Strength of London Forces

Factor	Effect
Number of electrons (Mr)	More electrons → larger, more polarisable electron cloud → stronger London forces
Surface area / chain length	Larger surface area → more points of contact → stronger London forces
Shape	Elongated molecules have more contact area than compact/branched ones

Evidence: Boiling Points of Noble Gases

Noble Gas	Mr	Boiling Point (°C)
He	4	−269
Ne	20	−246
Ar	40	−186
Kr	84	−153
Xe	131	−108

The boiling point increases with increasing number of electrons (Mr), confirming that London forces increase with the number of electrons.

Evidence: Boiling Points of Alkanes

Alkane	Formula	Mr	Boiling Point (°C)
Methane	CH₄	16	−162
Ethane	C₂H₆	30	−89
Propane	C₃H₈	44	−42
Butane	C₄H₁₀	58	0
Pentane	C₅H₁₂	72	36
Hexane	C₆H₁₄	86	69

Boiling points increase with chain length due to increasing London forces.

Effect of Branching

Pentane (bp 36°C) vs 2,2-dimethylpropane (neopentane, bp 10°C):

• Same molecular formula (C₅H₁₂) and Mr
• Neopentane is more compact (spherical) → less surface area for contact
• Weaker London forces → lower boiling point

---

2. Permanent Dipole-Dipole Forces

These exist between polar molecules that have a permanent dipole moment.

How They Arise

Polar molecules have a permanent uneven distribution of charge. The δ+ end of one molecule attracts the δ− end of another.

    δ+  δ−       δ+  δ−       δ+  δ−
    H—Cl    ···   H—Cl    ···   H—Cl

Strength

Permanent dipole-dipole forces are generally stronger than London forces for molecules of similar size. However, London forces still operate between polar molecules in addition to the dipole-dipole forces.

Evidence

Propanone (CH₃COCH₃, Mr = 58, bp = 56°C) vs butane (C₄H₁₀, Mr = 58, bp = 0°C):

• Similar Mr but propanone has a much higher boiling point
• Propanone is polar (C═O dipole) → permanent dipole-dipole forces in addition to London forces
• Butane is non-polar → London forces only

Exam Tip: Even in polar molecules, London forces are usually the largest contribution to the total intermolecular force. Permanent dipole-dipole forces provide an additional attractive force on top of London forces.

---

3. Hydrogen Bonding

Hydrogen bonding is a special, strong form of permanent dipole-dipole interaction.

Key Definition: A hydrogen bond is a strong intermolecular force that occurs when a hydrogen atom bonded to a highly electronegative atom (N, O, or F) is attracted to a lone pair on a nearby N, O, or F atom.

Requirements for Hydrogen Bonding

1. A hydrogen atom covalently bonded to N, O, or F (these are the three most electronegative elements that are small enough)
2. A lone pair on an N, O, or F atom on an adjacent molecule

    δ+        δ−    δ+        δ−
    —O—H ····· :O—H ····· :O—
              hydrogen bond

The hydrogen bond is represented by a dashed or dotted line. It is typically about 1/10th the strength of a covalent bond (typically 10–40 kJ mol⁻¹ compared with 200–500 kJ mol⁻¹ for covalent bonds).

Why N, O, and F?

These elements are:

• Highly electronegative → create a large δ+ on the hydrogen
• Small atoms → their lone pairs are concentrated in a small volume, close to the nucleus
• This combination makes the δ+H···lone pair interaction unusually strong

Common Misconception: Hydrogen bonding does NOT occur with Cl despite chlorine being quite electronegative (3.0). Chlorine's larger atomic radius means its lone pairs are too diffuse to form effective hydrogen bonds. Only N, O, and F form hydrogen bonds.