Intermolecular Forces: London Dispersion, Dipole-Dipole, and Hydrogen Bonding

What Are Intermolecular Forces?

Intermolecular forces (IMFs) are the attractive forces that exist between molecules. They are much weaker than the covalent bonds within molecules, but they determine crucial physical properties like boiling point, melting point, viscosity, and solubility.

Key distinction: Intramolecular forces (bonds) hold atoms together inside a molecule. Intermolecular forces hold molecules together with each other. When a substance boils, you break intermolecular forces — not covalent bonds.

Learning Goals: By the end of this guide, you should be able to:

Identify and describe the three main types of intermolecular forces.

Predict the dominant IMF in a given molecule.

Explain boiling point trends using IMF strength.

Compare IMFs in a table format suitable for exams.

Avoid the most common exam mistakes about forces between molecules.

The Three Types of Intermolecular Forces

1. London Dispersion Forces (LDF)

Also called van der Waals forces or instantaneous dipole-induced dipole forces, these are the weakest type of IMF — but they exist in all molecules.

How they form:

At any moment, the electrons in a molecule may be unevenly distributed, creating a temporary (instantaneous) dipole.
This temporary dipole induces a dipole in a neighbouring molecule.
The two temporary dipoles attract each other.

Strength depends on:

Number of electrons (more electrons → larger electron cloud → easier to polarise → stronger LDF)
Surface area (longer chain molecules have more contact area → stronger LDF)
Shape (branched molecules have less surface contact → weaker LDF)

Example: Why does pentane ( $C_5H_{12}$ ) boil at 36°C but neopentane ( $C(CH_3)_4$ ) boils at 10°C?

Both have the same molecular formula, but pentane is a straight chain with more surface contact area, resulting in stronger London forces and a higher boiling point. Neopentane is spherical with less surface area.

2. Dipole-Dipole Forces (Permanent Dipole)

These occur between polar molecules — molecules with a permanent dipole due to differences in electronegativity.

How they form: The $\delta^+$ end of one polar molecule is attracted to the $\delta^-$ end of a neighbouring polar molecule.

Requirements:

The molecule must have polar bonds
The molecular shape must be asymmetric so the dipoles don't cancel out

Polar (has dipole-dipole)	Non-polar (no dipole-dipole)
$HCl$ (linear, one polar bond)	$CO_2$ (linear, dipoles cancel)
$CHCl_3$ (asymmetric)	$CCl_4$ (tetrahedral, dipoles cancel)
$H_2O$ (bent, dipoles don't cancel)	$CH_4$ (tetrahedral, non-polar C–H)

Note: Polar molecules also have London forces. The dipole-dipole forces are in addition to the London forces.

3. Hydrogen Bonding

The strongest type of intermolecular force. It is a special case of dipole-dipole interaction that occurs when hydrogen is bonded to one of three very electronegative atoms: N, O, or F.

Why it's special:

H is tiny (no inner electron shells), so the $\delta^+$ charge is concentrated.
N, O, and F are very electronegative, creating a large $\delta^-$ .
The result: an unusually strong electrostatic attraction to a lone pair on a nearby N, O, or F.

Consequences of hydrogen bonding:

Water ( $H_2O$ ) has an anomalously high boiling point (100°C) compared to $H_2S$ (–60°C), despite being a smaller molecule.
Ice is less dense than liquid water (the open lattice structure with hydrogen bonds creates more space).
Hydrogen bonding is essential for DNA base pairing and protein folding in biology.

Comparison Table: All Three IMFs

Feature	London Dispersion	Dipole-Dipole	Hydrogen Bonding
Also called	Van der Waals, LDF	Permanent dipole	H-bond
Strength	Weakest	Moderate	Strongest IMF
Present in	ALL molecules	Polar molecules only	N–H, O–H, or F–H only
Cause	Temporary fluctuations in electron cloud	Permanent charge separation	H bonded to N/O/F attracted to lone pair
Depends on	Number of electrons, surface area, shape	Dipole moment of molecule	Electronegativity of N/O/F
Example	$Ar$ , $CH_4$ , $I_2$	$HCl$ , $CHCl_3$	$H_2O$ , $NH_3$ , $HF$

Intermolecular Forces Explorer

Visualise London forces, dipole-dipole, and hydrogen bonding in 3D. Compare how different molecules interact and see why boiling points vary.

Explore IMFs in 3D

Predicting and Comparing Boiling Points

The strength and type of IMFs directly determine boiling points. Here's a systematic approach:

Step-by-Step Method

Identify the IMFs present in each molecule.
Compare the strongest IMF — hydrogen bonding > dipole-dipole > London forces.
If the strongest IMF is the same, compare molecular size (more electrons → stronger London forces).

Example: Rank in Order of Increasing Boiling Point

$CH_4$ , $NH_3$ , $SiH_4$ , $PH_3$

Molecule	Strongest IMF	Relative Size	Predicted BP
$CH_4$	London only	10 electrons	Lowest
$SiH_4$	London only	18 electrons	Low (bigger than $CH_4$ )
$PH_3$	Dipole-dipole + London	18 electrons	Medium
$NH_3$	Hydrogen bonding	10 electrons	Highest

Actual: $CH_4$ (–161°C) < $SiH_4$ (–112°C) < $PH_3$ (–87°C) < $NH_3$ (–33°C) ✓

Worked Examples

Example 1: Identify the IMFs in Ethanol ( $C_2H_5OH$ )

Solution:

Ethanol has O–H bonds → hydrogen bonding is present.
The molecule is polar (asymmetric, O is electronegative) → dipole-dipole forces also present.
All molecules have London dispersion forces.

All three types are present, but hydrogen bonding is the dominant (strongest) IMF.

Example 2: Why does $HF$ have a lower boiling point than $H_2O$ ?

Solution: Both have hydrogen bonding, so we need to look deeper:

Water can form 2 hydrogen bonds per molecule (2 lone pairs + 2 O–H bonds).
HF can form only 1 hydrogen bond per molecule (despite F being more electronegative than O).
More hydrogen bonds per molecule → stronger overall intermolecular attraction → higher boiling point for water.

Example 3: Why is $I_2$ a solid but $Cl_2$ is a gas at room temperature?

Solution: Both are non-polar, so only London dispersion forces act.

$I_2$ has 106 electrons → very large, polarisable electron cloud → strong London forces.
$Cl_2$ has 34 electrons → smaller electron cloud → weaker London forces.
Stronger LDF in $I_2$ means a higher melting point, making it a solid at 25°C.

Common Mistakes

"Breaking bonds when boiling" — Boiling breaks intermolecular forces, NOT covalent bonds. When water boils, the O–H bonds remain intact; only the hydrogen bonds between $H_2O$ molecules break.
"London forces only exist in non-polar molecules" — London forces exist in all molecules, including polar ones and those with hydrogen bonding. They are just usually overshadowed by stronger IMFs.
"All molecules with H bonds have hydrogen bonding" — Hydrogen bonding requires H bonded to N, O, or F specifically. $CH_4$ has hydrogen atoms but no hydrogen bonding because C is not electronegative enough.
"Bigger = higher boiling point, always" — Size matters for London forces, but IMF type matters more. $H_2O$ (18 g/mol) boils at 100°C, while $C_4H_{10}$ (58 g/mol) boils at -1°C, because hydrogen bonding in water is much stronger than London forces in butane.
Calling hydrogen bonds "bonds" — Hydrogen bonds are intermolecular forces, not true chemical bonds. They are about 10× weaker than covalent bonds.

Exam Tips (A-Level / AP / IB)

Always state the specific type of IMF — never just say "van der Waals". Specify: London dispersion, dipole-dipole, or hydrogen bonding.
When explaining boiling points, use the format: "[molecule] has [type of IMF] which are [strong/weak] because [reason], therefore [more/less] energy is needed to overcome them."
For "explain" questions (not "state"), you need to describe how the force arises — e.g., "temporary fluctuations in the electron cloud create instantaneous dipoles."
If two molecules have the same IMF type, compare them using size (number of electrons) or surface area (chain length).

Frequently Asked Questions

What is the strongest intermolecular force?

Hydrogen bonding is the strongest intermolecular force. It occurs when H is bonded to N, O, or F and is attracted to a lone pair on a nearby electronegative atom.

Are van der Waals forces and London forces the same thing?

In many textbooks, "van der Waals forces" is used as an umbrella term for all IMFs, or sometimes specifically for London dispersion forces. To avoid confusion, it's better to use the specific name: London dispersion forces, dipole-dipole, or hydrogen bonding.

Why does ice float on water?

In ice, water molecules form a crystal lattice held by hydrogen bonds. This lattice has an open hexagonal structure with more space between molecules than in liquid water, making ice less dense. This is why ice floats.

How do intermolecular forces affect solubility?

"Like dissolves like" — polar solutes dissolve in polar solvents (both have dipole-dipole or H-bonding interactions). Non-polar solutes dissolve in non-polar solvents (London forces). Oil doesn't dissolve in water because London forces cannot compete with water's strong hydrogen bonds.

Can a molecule have all three types of IMF?

Yes! A molecule like ethanol ( $C_2H_5OH$ ) has all three: London dispersion forces (present in all molecules), dipole-dipole forces (it's polar), and hydrogen bonding (it has an O–H group).

Hydrogen Bonds — A deeper dive into hydrogen bonding, including its role in water and biological systems.
Chemical Bonds — Understand the intramolecular forces (ionic, covalent, metallic) before studying intermolecular forces.
VSEPR Theory — Molecular shape determines polarity, which determines IMF type.