Intermolecular Forces

Full Notes

Intermolecular forces (IMFs) are weak forces of attraction between molecules. They are much weaker than covalent, ionic or metallic bonds.

There are three types of intermolecular force:

• Induced dipole–dipole interactions (London dispersion forces)
• Permanent dipole–dipole interactions
• Hydrogen bonding

Instantaneous Dipole–Induced Dipole (London Dispersion Forces)

Instantaneous Dipole–Induced Dipole Forces (also called London Dispersion Forces) occur between all molecules, but are the only type of force between non-polar molecules (e.g., O₂, CO₂, CH₄).

They are caused by temporary fluctuations in electron density in molecules, 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)
• Surface contact between molecules (for example, branched chain hydrocarbons have lower boiling points than similar sized straight chain hydrocarbons)

Permanent Dipole–Dipole Forces

Permanent dipole–dipole forces occur between polar molecules, where there is a permanent dipole (see electronegativity and bond polarity).

They are stronger than London Forces for similar sized molecules 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 Dispersion forces between them.

The more polar the bond, the stronger the dipole–dipole interactions.

Example Hydrogen chloride

δ⁺H attracts δ⁻Cl of neighbouring HCl molecule.

Hydrogen Bonding

Hydrogen bonding is a unique type of intermolecular force that only occurs when H is directly bonded to F, O, or N 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 means a lone pair of electrons from another N, O or F atom is able to form strong forces of attraction to it.

This causes higher boiling/melting points than expected for a given sized molecule along with a high solubility in water, and unique properties in biological molecules.

Examples Molecules with hydrogen bonding

• 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

Water’s hydrogen bonding causes unusual properties:

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.
• This is why ice floats on water.

Simple Molecular Lattices

Simple molecular lattices are solid structures made up of molecules held together by weak intermolecular forces.

Example Iodine

Iodine (I₂) is a non-polar molecule, held together by induced dipole–dipole interactions (London Dispersion Forces).

It has a low melting and boiling point, does not conduct electricity and readily sublimes from solid to gas.

Properties of Simple Molecular Compounds

• Low melting/boiling points: only weak forces between molecules to overcome.
• No electrical conductivity: no charged particles or ions are free to move.
• Solubility:
  • Non-polar molecules dissolve in non-polar solvents.
  • Polar molecules may dissolve in polar solvents.

Summary

• Intermolecular forces are weaker than covalent, ionic or metallic bonds.
• Three types: London dispersion, permanent dipole–dipole, and hydrogen bonding.
• Hydrogen bonding explains water’s anomalous properties like high melting/boiling points and low density of ice.
• Simple molecular lattices have low melting/boiling points, no conductivity, and solubility depends on polarity.