Chemguide: Core Chemistry 14 - 16

Intermolecular forces

This page introduces the intermolecular forces which exist between individual molecules and help them stick together in liquids and solids.

I am assuming that you have read the page on electronegativity and polar bonds.

What is an intermolecular force?

Intermolecular forces are also known as intermolecular attractions. I tend to use the two terms interchangeably - they mean the same thing.

Intermolecular forces are forces between molecules, in the same way that an intercontinental missile can fly between continents, or an interaction is something happening between, for example, two or more people.

That is quite different from the forces which hold molecules together. If you heat water, H2O, and turn it into steam, you are breaking the forces between the water molecules, but you aren't breaking the covalent bonds within each water molecule.

If you heat solid iodine, I2, it turns to iodine vapour - so you have overcome all the intermolecular forces holding the molecules close to each other, but you still have I2 molecules in the vapour. You haven't done anything to the covalent bonds holding the atoms together.

This is really important - intermolecular forces are forces between one molecule and its neighbour(s). The covalent bonds within the molecule are a quite separate issue.

The origin of intermolecular forces

Intermolecular attractions in polar molecules

Suppose you have a simple molecule like hydrogen chloride, HCl. The hydrogen and chlorine are held together by a covalent bond, but chlorine is more electronegative than hydrogen, so the bonding pair is pulled slightly towards the chlorine end of the bond.

The molecule is polar - one end is slightly positive and the other is slightly negative because of the uneven distribution of the electrons.

If two HCl molecules come close enough together, there is an attraction between the positive end of one and the negative end of the other. There is an intermolecular attraction.

These are usually described as van der Waals dipole-dipole attractions - often just as dipole-dipole attractions.

At ordinary temperatures, when HCl is a gas, the molecules are moving around fast enough that this weak attraction isn't strong enough to hold them together.

But at lower temperatures, where the molecules aren't moving so fast, eventually the intermolecular forces will be strong enough to hold the particles in, first of all, a liquid and, at even lower temperatures, in a solid.

Intermolecular attractions in non-polar molecules - temporary fluctuating dipoles

Note:  This next bit is copied directly from Chemguide's advanced chemistry pages. You probably won't need to know this in any detail at 14-16 year old level. But it isn't very difficult, and is interesting.

Attractions are electrical in nature. In a symmetrical molecule like hydrogen, however, there doesn't seem to be any electrical distortion to produce positive or negative parts. But that's only true on average.

The lozenge-shaped diagram represents a small symmetrical molecule - H2, perhaps, or Br2. The even shading shows that on average there is no electrical distortion.

But the electrons are mobile, and at any one instant they might find themselves towards one end of the molecule, making that end δ-. The other end will be temporarily short of electrons and so becomes δ+.

An instant later the electrons may well have moved up to the other end, reversing the polarity of the molecule.

This constant "sloshing around" of the electrons in the molecule causes rapidly fluctuating dipoles even in the most symmetrical molecule. It even happens in noble gases, like helium, which consist of single uncombined atoms.

If both the helium electrons happen to be on one side of the atom at the same time, the nucleus is no longer properly covered by electrons for that instant.

How temporary dipoles give rise to intermolecular attractions

I'm going to use the same lozenge-shaped diagram now to represent any molecule which could, in fact, be a much more complicated shape. Shape does matter, but keeping the shape simple makes it a lot easier to both draw the diagrams and understand what is going on.

Imagine a molecule which has a temporary polarity being approached by one which happens to be entirely non-polar just at that moment. (A pretty unlikely event, but it makes the diagrams much easier to draw! In reality, one of the molecules is likely to have a greater polarity than the other at that time - and so will be the dominant one.)

As the right hand molecule approaches, its electrons will tend to be attracted by the slightly positive end of the left hand one.

This sets up an induced dipole in the approaching molecule, which is orientated in such a way that the δ+ end of one is attracted to the δ- end of the other.

An instant later the electrons in the left hand molecule may well have moved up the other end. In doing so, they will repel the electrons in the right hand one.

The polarity of both molecules reverses, but you still have δ+ attracting δ-. As long as the molecules stay close to each other the polarities will continue to fluctuate in synchronisation so that the attraction is always maintained.

There is no reason why this has to be restricted to two molecules. As long as the molecules are close together this synchronised movement of the electrons can occur over huge numbers of molecules.

This diagram shows how a whole lattice of molecules could be held together in a solid using these temporary fluctuating dipoles. An instant later, of course, you would have to draw a quite different arrangement of the distribution of the electrons as they shifted around - but always in synchronisation.

These attractions due to temporary fluctuating dipoles are properly known as van der Waals dispersion forces or just as dispersion forces, and that is what I will call them as a matter of routine.

Hydrogen bonding

You can't do a survey of intermolecular forces without mentioning hydrogen bonds. These are really important attractions, although you won't meet them much at this level. They are, for example, the attractions which hold the two strands of DNA together.

Hydrogen bonds occur in compounds where a hydrogen atom is covalently bound to one of the very electronegative atoms, nitrogen, oxygen and fluorine. Two simple examples are water, H2O, and ammonia NH3.

The oxygen and nitrogen attract the bonding pairs quite closely to themselves, and so there is quite a strong dipole on each bond, especially in the water case. Oxygen is more electronegative than nitrogen.

But this isn't just another attraction between polar molecules - it is stronger than that. The hydrogen atoms are so strongly attracted to the lone pairs on nearby water molecules, that there is almost the beginning of a covalent bond being formed.

It isn't really a covalent bond, but it is nevertheless quite a powerful attraction. This attraction between a fairly positive hydrogen atom and a lone pair on a very electronegative nearby atom is called a hydrogen bond.

Summing it all up

There are various forms of intermolecular attractions:

  • Attractions involving temporary fluctuating dipoles.

  • Attractions involving permanent dipoles.

  • Attractions involving hydrogen bonds.

You mustn't assume that these are alternatives. More than one of these maybe present in the same substance, and they just reinforce each other.

For example, in water, there are fluctuating dipoles as well as permanent dipoles as well as hydrogen bonds. This all adds up to amazingly strong intermolecular attractions given the size of the molecule.

On the other hand, butane, a gas used as a fuel, hasn't got an overall dipole on the molecule, and doesn't have any hydrogen atoms attached to a very electronegative atom. The only attractive forces between the molecules are fluctuating dipoles.

As a result, butane is a gas with a boiling point of -0.5°C, but water, with its larger intermolecular attractions is a liquid with a boiling point of 100°C.


As far as exams at this 14-16 year old level are concerned, you can almost certainly forget about most of this page. There are one or two cases where you will need information from this in another context. I will point you back to here when that happens.

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© Jim Clark 2019