Intermolecular Forces, Electronegativity and Bond Properties

Specification Reference Physical Chemistry: Chemical bonding 3.6

Quick Notes

Hydrogen bonding 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.
Hydrogen bonding explains water’s high boiling point, surface tension, and why ice is less dense than water.
A polar bond forms when electrons in a covalent bond are shared unequally due to electronegativity difference.
Intermolecular forces exist between molecules and hold molecular substances together, they include:
- Instantaneous dipole–induced dipole (London forces) — in all molecules.
- Permanent dipole–dipole forces — in polar molecules.
- Hydrogen bonds — strongest dipole–dipole force.
Covalent, ionic, and metallic bonds are stronger than intermolecular forces.

Full Notes

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 dipole–dipole 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 from another N, O or F atom to form strong forces of attraction to it.

Hydrogen bonding between molecules can cause higher boiling/melting points than predicted for the size of molecules, high solubility in water, and unique properties in biological molecules.

Examples:

CIE A-Level Chemistry diagram showing hydrogen bonding between molecules such as water, ammonia, and HF.

H₂O (water) → hydrogen bonds increase boiling point significantly.
NH₃ (ammonia) → weaker than water but stronger than Van der Waals.
HF → forms hydrogen bonds, raising boiling points.

Hydrogen bonds are shown as dotted lines
e.g.: H–O···H–O, with partial charges δ⁺ on H and δ⁻ on O.

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.
For example:

CIE A-Level Chemistry diagram showing hydrogen bonding between two water molecules with dotted line, lone pairs, and partial charges.

Hydrogen Bonding and Anomalous Properties of Water

Hydrogen bonding explains several unusual physical properties of water and ice:

High boiling and melting points: strong hydrogen bonds require more energy to overcome, so water boils at 100 °C.
High surface tension: water molecules are strongly attracted at the surface.
Lower density of ice: in ice, water molecules form an open hexagonal lattice held by hydrogen bonds. This creates large gaps between molecules, making ice less dense than liquid water.

CIE A-Level Chemistry diagram showing hexagonal hydrogen-bonded lattice structure in ice with lower density than water.

Electronegativity and Bond Polarity

A polar bond occurs when two atoms in a covalent bond have different electronegativities, so electrons are unequally shared.

The more electronegative atom becomes δ⁻, the less electronegative δ⁺, creating a dipole.

CIE A-Level Chemistry diagram showing unequal electron sharing in a polar covalent bond.

Example:

HCl → H (δ⁺) — Cl (δ⁻)

CIE A-Level Chemistry diagram of HCl showing H δ+ and Cl δ− with bond polarity.

A molecule can be polar or non-polar depending on symmetry of polar bonds:

Non-polar: polar bonds arranged symmetrically so dipoles cancel.

CIE A-Level Chemistry diagram showing symmetrical non-polar molecule with dipoles cancelling out.

Polar: dipoles do not cancel due to asymmetry.

CIE A-Level Chemistry diagram showing examples of polar molecules with permanent dipoles.

Van der Waals’ Forces

Van der Waals’ forces are intermolecular forces between molecules that do not involve chemical bonds. These include:

London dispersion (instantaneous dipole–induced dipole)
Permanent dipole–dipole
Hydrogen bonding

i) London Dispersion Forces

Occur between all molecules (polar and non-polar)
Caused by temporary fluctuations in electron distribution
These create a temporary dipole that induces a dipole in a neighbouring molecule.

CIE A-Level Chemistry diagram showing instantaneous dipole inducing dipole in neighbouring molecule.

Strength increases with number of electrons / size of molecule.
Example CH₄, Cl₂

ii) Permanent Dipole–Dipole Forces

Occur between molecules with permanent dipoles (i.e. polar molecules)
Molecules align with opposite partial charges attracting each other

Example HCl molecules attract each other via permanent dipole-dipole forces

CIE A-Level Chemistry diagram showing permanent dipole-dipole forces between HCl molecules.

iii) Hydrogen Bonding (outlined above)

A strong form of dipole–dipole interaction
Occurs 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.

CIE A-Level Chemistry diagram showing hydrogen bonds as a strong form of dipole-dipole force.

Hydrogen bonding is always stronger than London or permanent dipole forces for comparably sized molecules.

Strength of Bonds and Intermolecular Forces

Order of strength (similar sized molecules):

Ionic, covalent, metallic bonds ≫ intermolecular forces
Hydrogen bonds > Permanent dipole–dipole > London dispersion

Summary

Hydrogen bonding: special strong dipole–dipole between H and N, O, or F.
Explains anomalies in water: high boiling point, surface tension, ice less dense.
Electronegativity difference creates bond polarity: δ⁻/δ⁺.
Molecular polarity depends on shape/symmetry of dipoles.
Van der Waals forces:
- London dispersion → all molecules, temporary dipoles
- Permanent dipole–dipole → polar molecules
- Hydrogen bonding → strongest intermolecular force
Relative strengths: covalent/ionic/metallic ≫ H-bonds > dipole–dipole > London.