Intermolecular Forces and Hydrogen Bonding

Full Notes

Types of Intermolecular Forces

Intermolecular forces (IMFs) are weak forces of attraction between molecules.

While they are much weaker than covalent, ionic or metallic bonds, they play a critical role in determining physical properties like boiling point, melting point, solubility and volatility in molecular substances.

There are three types of intermolecular force:

• London forces (also called dispersion or induced dipole forces)
• Permanent dipole–dipole interactions
• Hydrogen bonding

London Forces (Instantaneous Dipole–Induced Dipole)

London Forces occur between all molecules, but are the only type of force between non-polar molecules.

• Caused by temporary fluctuations in electron density, creating instantaneous dipoles.

• Electrons are constantly moving within molecules and unequal electron distribution around a molecule creates an instantaneous dipole that can induce a dipole on a neighbouring molecule.
• Two opposite dipoles from different molecules are attracted to each other. This creates a weak force of attraction between the two molecules.

The strength of London forces increases with number of electrons (larger molecules = stronger forces) and surface contact between molecules.

Examples

• Methane (CH₄) has weak London forces = low boiling point of −161 °C
• Pentane has stronger London forces than methane = higher boiling point of +36 °C

Permanent Dipole–Dipole Forces

Permanent dipole–dipole forces occur between polar molecules, where there is a permanent dipole.

• Stronger than London Forces because the dipoles are permanent, not temporary.
• The greater the partial charges in molecules (higher the polarity), the stronger the permanent dipole–dipole forces of attraction.
• Molecules that can form permanent dipole–dipole forces have higher boiling points compared to similar sized molecules with only London forces between them.

Example Hydrogen chloride (HCl)

δ⁺H attracts δ⁻Cl of neighbouring molecule

Hydrogen Bonding

Hydrogen bonding is a unique type of intermolecular force that only occurs when H is directly bonded to a fluorine, oxygen, or nitrogen atom (all highly electronegative atoms).

• There is a strong attraction force between the H from an N–H, O–H or F–H bond and the lone pair of electrons on another N, O or F atom.
• The proton in the hydrogen atom’s nucleus is left exposed on one side when bonded to N, O or F, and this allows a lone pair of electrons from another N, O or F atom to form strong forces of attraction to it.

Hydrogen bonding causes higher boiling/melting points than otherwise expected, high solubility in water, and unique properties in biological molecules.

Examples

• Water (H₂O): hydrogen bonds between O–H groups
• Ammonia (NH₃): N–H bonds form H-bonds
• Hydrogen fluoride (HF)

Matt’s exam tip

If an exam question asks you to draw hydrogen bonding between two OH groups (such as water molecules), always make sure you draw and label the hydrogen bond with a dotted line that has an angle of 180 degrees between the oxygen, hydrogen and oxygen. Always show the lone pair of electrons on the oxygen and include partial charges.

Anomalous Properties of Water and Ice

High boiling and melting point

H₂O has a melting point of 0°C whereas H₂S, which has a similar structure and shape, has a melting point of only −85.5°C. This is because the strongest type of intermolecular force in H₂O is hydrogen bonding, whereas in H₂S it is only permanent dipole–dipole forces.

Without hydrogen bonding, H₂O would boil below 0°C.

Ice is less dense than liquid water

In ice, hydrogen bonds hold water molecules in an open hexagonal structure, making it less dense than liquid water. When ice melts, molecules have enough energy to overcome some of the hydrogen bonding between them and molecules can move closer, increasing density.

Boiling Point Trends Explained by IMFs

Alkanes:

• Boiling point increases with carbon chain length (more electrons in molecule means stronger London forces)
• Branched isomers have lower boiling points (less surface area of contact between molecules means weaker forces)

Alcohols vs. Alkanes:

• Alcohols (e.g. ethanol) have higher boiling points than alkanes of similar size because of hydrogen bonding between O–H groups.

Hydrogen Halides:

• Boiling point generally increases from HCl → HI (larger molecules = stronger London forces)
• The polarity of the molecules decreases, meaning weaker permanent dipole–dipole forces. However, the size of the molecules increases, meaning stronger induced dipole–dipole forces.

Molecule	Boiling Point (°C)
HF	19.5
HCl	-85
HBr	-66
HI	-35

• HF has a relatively high boiling point because hydrogen bonding exists between molecules.

Comparison of Boiling Points Based on Forces

Summary

• London forces, permanent dipole–dipole interactions, and hydrogen bonds are the three types of IMFs.
• Hydrogen bonding is the strongest and explains anomalies like water’s high boiling point and ice’s low density.
• Boiling point trends depend on molecular size, polarity, and hydrogen bonding.
• Alkanes increase with chain length, alcohols are higher than alkanes, HF unusually high compared to other hydrogen halides.