Van der Waals and London Dispersion Forces
Van der Waals forces are a type of intermolecular force that occurs because of dipole-dipole interactions. London dispersion force is a sub-type of the Van der Waals force that is predominant in non-polar molecules. An intermolecular force is a force occurring between two different molecules.
All these forces are weaker than both ionic and covalent bonds. They also act on a short distance and are weaker the further molecules are from each other.
Below we will dive into both Van der Waals forces in general and London dispersion forces.
What are Van der Waals Forces?
Van der Waals forces are weak forces between molecules that occur due to either temporary or permanent dipoles. They are also called dispersion forces. The dipoles (either temporary or permanent) cause attraction between the molecules.
Van der Waals forces include London dispersion forces (discussed below), dipole-dipole forces, and ion-dipole. The strongest force of these is ion-dipole followed by dipole-dipole. London dispersion forces are the weakest force.
All molecules experience London dispersion forces. However, it is the only intermolecular force (IMF) present in non-polar molecules. In polar molecules, other IMFs dominate so although London dispersion forces are present they may not be mentioned.
Van der Waals interactions are names after Johannes Diderik van der Waals. Van der Waals was the first to propose intermolecular forces as the reason for the properties of certain molecules.
London Dispersion Force
The London dispersion force occurs between two molecules due to a temporary dipole. The force arises due to the uneven distribution of electrons which causes a temporary dipole. On average, electrons are evenly distributed, but in any instance, they may be clustered on one side of a molecule. This clustering causes an instantaneous dipole. The instantaneous dipole can cause a temporary dipole to form on surrounding molecules. The dipoles on each molecule then exert a force on each other.
This is the weakest type of intermolecular force.
The larger the electron cloud around a molecule, the stronger the instantaneous dipole and, therefore, the attractive force is. This can be seen in the trend for the melting point and boiling point for halogens.
| Halogen | Melting Point (K) |
| F2 | 53.5 |
| Cl2 | 171.6 |
| Br2 | 265.9 |
| I2 | 386.8 |
As the molecule’s electron cloud size increases (moving down rows of the periodic table), the melting point increases. The increased melting point indicates stronger forces that need to be broken between molecules.
The attraction will also be stronger on a longer more spread-out molecule than a compact molecule. The spread-out molecule has more polarizability and contact with surrounding molecules.
The red molecule will have stronger London dispersion forces because of its length compared to the compact blue molecule with the same number of atoms. The longer length means there is more distance for an electron to move and create a stronger dipole.
We know these forces must exist because there are non-polar molecules that are liquid at room temperature. If there are no intermolecular forces between them, they would be gasses.
Example Problems
Which molecule would have larger IMF forces: Bromine or Fluorine?
The correct answer is bromine. Both molecules are non-polar so the main IMF is London dispersion forces. Bromine will have more polarizability due to having more electrons than fluorine.
Which molecule will have larger London dispersion forces: Methane (CH4) or hexane (C6H14)?
Hexane is the correct answer. Hexane will have larger London dispersion forces because the molecule is larger (in terms of atoms) and more spread out (can become more polarized).
Van der Waals Equation
The Van der Waals equation modifies the ideal gas law to account for weak intermolecular forces. Intermolecular forces are corrected by adding an2/V2 to the pressure term. The other correction term is accounting for the volume of molecules
- P = Pressure
- V = Volume
- n = Moles of gas
- R= Ideal gas constant
- T = Temperature
- a = gas constant correction factor (specific for each gas)
- b= gas constant correction factor for size
London Dispersion Forces History
London dispersion forces are named after German physicist Fritz London (1900-1954). Although originally from Germany, he left in 1933 and went to Oxford University. Then in 1939 accepted a position at Duke University. There he also contributed heavily to the development of quantum mechanics.