Intermolecular Forces

Intramolecular forces vs intermolecular forces

Let’s look at the bonding in carbon and oxygen. Carbon is a giant covalent structure. This means it contains a large number of atoms held together in a repeating lattice structure by many covalent bonds. Covalent bonds are a type of intramolecular force. In contrast, oxygen is a simple covalent molecule. Two oxygen atoms bond using one covalent bond, but there are no covalent bonds between molecules. Instead there are just weak intermolecular forces. To melt diamond, we need to break these strong covalent bonds, but to melt oxygen we simply need to overcome the intermolecular forces. As you’re about to find out, breaking intermolecular forces is much easier than breaking intramolecular forces. Let’s explore intramolecular and intermolecular forces now.

Intramolecular forces

As we defined above, intramolecular forces are forces within a molecule. They include ionic, metallic, and covalent bonds. You should be familiar with them. (If not, check out Covalent and Dative Bonding, Ionic Bonding, and Metallic Bonding.) These bonds are extremely strong and breaking them requires a lot of energy.

Intermolecular forces

An interaction is an action between two or more people. Something that is international occurs between multiple nations. Likewise, intermolecular forces are forces between molecules. These are weaker than intramolecular forces, and don’t require as much energy to break. They include van der Waals forces (also known as induced dipole forces, London forces or dispersion forces), permanent dipole-dipole forces, and hydrogen bonding. We’ll explore them in just a second, but first we need to revisit bond polarity.

Fig. 1 - A diagram showing the relative strengths of intramolecular and intermolecular forces

Bond polarity

As we mentioned above, there are three main types of intermolecular forces:

• Van der Waals forces.
• Permanent dipole-dipole forces.
• Hydrogen bonding.

How do we know which one a molecule will experience? It all depends on bond polarity. The bonding pair of electrons is not always spaced equally between two atoms joined with a covalent bond (remember Polarity?). Instead, one atom could attract the pair more strongly than the other. This is due to differences in electronegativities.

A more electronegative atom will pull the pair of electrons in the bond towards itself, becoming partially negatively-charged, leaving the second atom partially positively-charged. We say that this has formed a polar bond and the molecule contains a dipole moment.

We can represent this polarity using the delta symbol, δ, or by drawing a cloud of electron density around the bond.

For example, the H-Cl bond shows polarity, as chlorine is much more electronegative than hydrogen.

Fig. 2 - HCl. The chlorine atom attracts the bonding pair of electrons towards itself, increasing its electron density so that it becomes partially negatively charged

However, a molecule with polar bonds may not be polar overall. If all the dipole moments act in opposite directions and cancel each other out, the molecule will be left with no dipole. If we look at carbon dioxide, , we can see that it has two polar C=O bonds. However, because is a linear molecule, the dipoles act in opposite directions and cancel out. is therefore a nonpolar molecule. It has no overall dipole moment.

Fig. 3 - CO2 may contain the polar bond C=O, but it is a symmetrical molecule, so the dipoles cancel out

Types of intermolecular forces

A molecule will experience different types of intermolecular forces depending on its polarity. Let’s explore them each in turn.

Van der Waals forces

Van der Waals forces are the weakest type of intermolecular force. They have lots of different names - for example, London forces, induced dipole forces or dispersion forces. They are found in all molecules, including non-polar ones.

Although we tend to think of electrons as being uniformly distributed throughout a symmetrical molecule, they instead are constantly in motion. This movement is random and results in the electrons being spread unevenly within the molecule. Imagine shaking a container full of ping pong balls. At any moment, there might be a greater number of ping pong balls on one side of the container than on the other. If these ping pong balls are negatively charged, it means the side with more ping pong balls will also have a slight negative charge whilst the side with fewer balls will have a slight positive charge. A small dipole has been created. However, the ping pong balls are constantly moving as you shake the container, and so the dipole keeps on moving too. This is known as a temporary dipole.

If another molecule comes close to this temporary dipole, a dipole will be induced in it as well. For example, if the second molecule draws near to the partially positive side of the first molecule, the second molecule’s electrons will be slightly attracted to the first molecule’s dipole and will all move over to that side. This creates a dipole in the second molecule known as an induced dipole. When the first molecule’s dipole switches direction, so does the second molecule’s. This will happen to all the molecules in a system. This attraction between them is known as van der Waals forces.

Van der Waals forces increase in strength as molecule size increases. This is because larger molecules have more electrons. This creates a stronger temporary dipole.

Fig. 4 - A temporary dipole in one molecule induces a dipole in a second molecule. This spreads throughout all the molecules in a system. These forces are known as van der Waals forces or London dispersion forces

Permanent dipole-dipole forces

As we mentioned above, dispersion forces act between all molecules, even ones that we would consider non-polar. However, polar molecules experience an additional type of intermolecular force. Molecules with dipole moments that do not cancel each other out have something we call a permanent dipole. One part of the molecule is partially negatively-charged, while another is partially positively-charged. Oppositely-charged dipoles in neighbouring molecules attract each other and similarly-charged dipoles repel each other. These forces are stronger than van der Waals forces as the dipoles involved are larger. We call them permanent dipole-dipole forces.

Hydrogen bonding

To illustrate the third type of intermolecular force, let’s take a look at some hydrogen halides. Hydrogen bromide, , boils at -67 °C. However, hydrogen fluoride, , does not boil until temperatures reach 20 °C. To boil a simple covalent substance you must overcome the intermolecular forces between molecules. We know that van der Waals forces increase in strength as molecule size increases. As fluorine is a smaller atom than chlorine, we would expect HF to have a lower boiling point. This clearly isn’t the case. What causes this anomaly?

Looking at the table below, we can see that fluorine has a high electronegativity value on the Pauling scale. It is a lot more electronegative than hydrogen and so the H-F bond is very polar. Hydrogen is a very small atom and so its partial positive charge is concentrated in a small area. When this hydrogen nears a fluorine atom in an adjacent molecule, it is strongly attracted to one of fluorine’s lone pairs of electrons. We call this force a hydrogen bond.

Fig. 5 - Hydrogen bonding between HF molecules. The partially positive hydrogen atom is attracted to one of fluorine’s lone pairs of electrons

Not all elements can form hydrogen bonds. In fact, only three can - fluorine, oxygen and nitrogen. To form a hydrogen bond, you require a hydrogen atom bonded to a very electronegative atom that has a lone pair of electrons, and only these three elements are electronegative enough.

Common molecules that do form hydrogen bonds include water (), ammonia () and hydrogen fluoride. We represent these bonds using a dashed line, as shown below.

Fig. 6 - Hydrogen bonding in water molecules

Hydrogen bonds are a lot stronger than both permanent dipole-dipole forces and dispersion forces. They require more energy to overcome. Going back to our example, we now know that this is why HF has a much higher boiling point than HBr. However, hydrogen bonds are only about 1/10th as strong as covalent bonds. This is why carbon sublimes at such high temperatures - a lot more energy is needed to break the strong covalent bonds between atoms.

Examples of intermolecular forces

Let’s look at some common molecules and predict the intermolecular forces they experience.

Carbon monoxide, , is a polar molecule and so has permanent dipole-dipole forces and van der Waals forces between molecules. On the other hand, carbon dioxide, , only experiences van der Waals forces. Although it contains polar bonds, it is a symmetrical molecule and so the dipole moments cancel each other out.

Fig. 7 - The bond polarity in carbon monoxide, left, and carbon dioxide, right

Methane, , and ammonia, , are similar-sized molecules. They therefore experience similar strength van der Waals forces, which we also know as dispersion forces. However, the boiling point of ammonia is a lot higher than the boiling point of methane. This is because ammonia molecules can hydrogen bond with each other, but methane molecules can’t. In fact, methane does not even have any permanent dipole-dipole forces as its bonds are all non-polar. Hydrogen bonds are a lot stronger than van der Waals forces, so require a lot more energy to overcome and boil the substance.

Fig. 8 - Methane is a non-polar molecule. In contrast, ammonia is a polar molecule and experiences hydrogen bonding between molecules, shown by the dashed line. Note that all the N-H bonds in ammonia are polar, although not all the partial charges are shown