Intermolecular forces


Intermolecular forces dipole-dipole, London forces, hydrogen bonding versus

Covalent bonds

Dr. Walt Volland, all rights reserved 1998-2011 revised Nov 3, 2011


Attractions between molecules are classified as

  • Dipole-dipole interactions , example: ammonia, NH3
  • London forces also known as van der Waals forces, example: methane, CH4
  • Hydrogen bonding example: water, H2O .



Relative magnitudes of forces

The relative size of these interactions is important so the relative effects are understood.
The relative typical strengths for the different interactions are listed here in descending strength.
Covalent bonds >

Hydrogen bonding >

Dipole-dipole interactions >

London forces

400 kcal >

12-16 kcal >

2-0.5 kcal >

less than 1 kcal

From this we can see that normal covalent bonds are almost 40 times the strength of hydrogen bonds.
Covalent bonds are almost 200 times the strength of dipole-dipole forces, and more than 400 times the size of London forces.


Dipole-dipole interactions
Dipole-dipole interactions exist between molecules that are permanently polar. This requires the presence of polar bonds and a asymmetric molecule. These molecules have a permanent separation of positive and negative charge. In the illustration the H end of HCl is permanently slightly positive charge. The Cl end of HCl has a permanent slight negative charge. the "H" in one molecule is attracted to the "Cl" in a neighbor. The intermolecular force is weak compared to a covalent bond. But this dipole-dipole interaction is one of the stronger intermolecular attractions.

London forces

London forces exist in nonpolar molecules.

These forces result from temporary charge imbalances. The temporary charges exist because the electrons in a molecule or ion move randomly in the structure. The nucleus of one atom attracts electrons form the neighboring atom. At the same time, the electrons in one particle repel the electrons in the neighbor and create a short lived charge imbalance.


These temporary charges in one molecule or atom attract opposite charges in nearby molecules or atoms. A local slight positive charge d+ (lower case Greek delta) in one molecule will be attracted to a temporary instantaneous slight d- negative charge in a neighboring molecule.


London forces in hydrocarbons and organic molecules
The temporary separations of charge produce London force attractions are what associate one nonpolar organic molecule with its neighbors. The possibilities for these interactions go up with increasing molecular size and surface area. A larger surface increases the chances for the "induced" charge separations to interact. If the molecules are linear they have more surface area than if they are folded into a sphere. The linear molecules have higher melting and boiling points because of the increased attractions.

n-pentane linear
Hydrogen bonding

Hydrogen bonding is a unique type of intermolecular molecular attraction. There are two requirements.

First, there must be a covalent bond between a H atom and either F, O, or N. (These are the three most electronegative elements.)

Second, the H atom in the polar bond -O-H,  -N-H, F-H inteact with a lone pair of electrons on a nearby atom of F, O, or N in another molecule or for BIG molecules in another part of the molecule.

liquid water hydrogen bonding model

The normal boiling point for water is 100oC. The graph below shows how high this boiling point is compared to the predicted ( green circles) boiling points.
The observed boiling point of 100oC is almost four times greater than the expected value at about -80oC or -100oC . The predicted boiling point (green circles) from the trend of boiling points for H2Te, H2Se, H2S and H2O is very low. If the trend continued the predicted boiling point would be below -62 oC. The "anomalous" boiling point for water is the result of hydrogen bonding between water molecules.

When can hydrogen bonding exist?

Possible combinations where hydrogen bond can exist.

The first entry shows the covalent bond to the O or N atom. These atoms form two and three covalent bonds. The single covalent bond between O,N,F is shown and the dashed line shows the hydrogen bond. NOTICE the H atom is attracted to a lone pair on the nearby N, O, F atom.

A covalent bond between -O-H ---- :O-

A covalent bond between -N-H----- :O-

A covalent bond between F-H ------ :O-

A covalent bond between -O-H ---- :N-

A covalent bond between -N-H---- :N-

A covalent bond between F-H ----- :N-

A covalent bond between -O-H ----- :F-

A covalent bond between -N-H ---- :F-

A covalent bond between F-H ------ :F-


Hydrogen bonding in an ice crystal model
Summary on hydrogen bonding

Hydrogen bonding is responsible for the expansion of water when it freezes. The water molecules in the solid have tetrahedral spatial arrangement for the two lone pairs and two single bonds radiating out from the oxygen. The lone pairs on the "O" atoms are attracted to nearby water molecules through hydrogen bonds. A cage like structure results. The cage has an hexagon shaped opening.

Click for more on hydrogen bonding in water

Exercise: Which of the following molecules display hydrogen bonding?

 Methane, CH4

 methyl ether, CH3OCH3

 Hydrogen peroxide, H2O2

 methyl alcohol, CH3OH


Answer: The hydrogen peroxide and methyl alcohol have hydrogen bonding between molecules.

The methane lacks highly electronegative atoms bonded to the hydrogen atoms. The methyl ether has an oxygen but the carbons are bonded to the "O" . The hydrogen atoms are not bonded to very electronegative "O" atom.


Dr. Walt Volland, all rights reserved 1998--2011
revised November 3, 2011