Intermolecular Forces

Intermolecular Forces

There are three types of intermolecular force to know (in order of ascending strength):


  1. Van der Waals forces (Induced dipole-dipole)
  2. Permanent dipole-dipole forces
  3. Hydrogen bonding


Intermolecular Force Explanation Increases With
Van der Waals Electrons are always moving around in charge clouds, which means that the distribution of charge is always changing; for example, at any one moment one part of the atom could be more negative than the other. This means the atom has a temporary dipole.

This causes temporary dipoles on nearby atoms.

This results in very weak electrostatic attractions between slightly oppositely charged atoms.


·         Atom size

·         Surface area (less branching = more surface area = more points of contact)

Permanent dipole-dipole Caused by electronegativity (see below). In bonds where one atom has a higher electronegativity than the other (eg HCl), the more electronegative atom will attract the bonding pair of electrons towards it, causing a shift in electron density and a polar bond.

The more electronegative atom gets a ‘delta minus’ charge (δ−) and the less electronegative a ‘delta plus’ charge (δ+), which causes weak electrostatic forces between molecules of the compound.

·         Difference in electronegativity of atoms
Hydrogen Bonding Only happens in molecules where hydrogen is bonded to fluorine, nitrogen or oxygen. These atoms are highly electronegative, and so attract the bonding electrons towards them in the covalent bond. The hydrogen atom then has a high charge density (because the atom is so small and positively charged), resulting in it forming strong intermolecular forces with F, N or O atoms on other molecules called hydrogen bonds. (Note: not stronger than a covalent bond).


The hydrogen bonds present in ice are slightly less packed together than in water, meaning that ice is less dense than water and can float on the top of lakes, ponds etc.



Definition of electronegativity: The relative ability to attract the bonding electrons in a covalent bond.