Bonding Forces Between Molecules
Intermolecular forces are the forces of attraction between individual molecules of a substance, that affect properties such as melting and boiling points.
These forces also affect how strongly molecules of different substances are attracted to each other, so they have an important impact on reactivity.
And because we often use the word “bonding” it is not uncommon for students to confuse intermolecular forces with more “formal” bonds.
So it’s worth us being very clear and stating that we are talking about forces of attraction that don’t involve sharing or exchanging electrons. We focus on the intermolecular forces that occur between covalent molecules.
As you might have guessed, that means that intermolecular forces are weaker than formal bonds between atoms or ions within molecules.
Three Types of Intermolecular Forces
Intermolecular forces vary considerably, and we categorise them into three groups that we will discuss in this article.
We will also consider the differences in the categories and the relative strength of those forces.
You may have already heard of the three types:
van der Waals forces
van der Waals Forces
All atoms are, of course, made up of positively and negatively charged particles even though they are overall neutral. It’s the same with molecules.
Consider that electrons do not sit still within the centre of their orbitals. They move around because they have energy, and they might be located anywhere at any point in time.
So it follows that there will be moments when an atom’s charge will not be evenly distributed, meaning there are moments when an atom has more positively charged and more negatively charged regions. They have a δ+ region and a δ- region, so they have a temporary dipole.
The atom’s neighbours also undergo this constant redistribution of charge meaning that there will be some electrostatic attraction between atoms – even between atoms of noble gas elements.
That electrostatic attraction is a van der Waals force.
Consider too that an atom that has a temporary dipole may induce a dipole in its neighbour atom. If the δ- region is towards the neighbour atom (and close enough to interact) then electrons in the neighbour atom may be slightly repelled. This would cause the neighbour atom to have a temporary dipole with δ+ nearest to the original atom.
The δ- and δ+ regions of those adjacent atoms will naturally exhibit electrostatic attraction toward each other. This attraction is another presentation of van der Waals force.
van der Waals forces are not limited to individual atoms, but also commonly occur between molecules. The molecules may be elements or covalent compounds.
In fact, van der Waals forces are stronger with molecules than with monoatomic elements. The constant redistribution of electrons within each atom of each molecule means there are numerous temporary (or “instantaneous”) dipoles in each molecule.
Not that van der Waals forces are sometimes referred to using other names such as London dispersion forces, induced dipole-dipole or instantaneous dipole-dipole bonds/forces. These are not to be confused with when we refer to dipole-dipole forces which usually means a permanent dipole-dipole attraction
Dipole-dipole forces are possible when covalent molecules have one or more polar bond within their structure. A polar bond has a δ+ end and a δ- end, so it has a permanent dipole. You will sometimes see this type of attraction referred to as a permanent dipole-dipole.
Dipole-dipole is a stronger force of attraction than van der Waals forces.
Consider a polar bond within a covalent molecule.
A polar bond is a bond between atoms of different elements with dissimilar electronegativity. The electrons being shared as part of the covalent bond between the atoms will be more attracted towards the more electronegative atom. This causes there to be a δ- end and a δ+ end to the bond (i.e. a permanent dipole).
The dipole means there is an electrostatic attraction between the molecule and another molecule with a dipole. The ‘other’ molecule could be a molecule of the same substance or another substance.
If the ‘other’ molecule is the same substance this will result in melting and boiling point being higher than could be explained by other properties.
Hydrogen bonds are a dipole-dipole attraction due to a very polar covalent bond between hydrogen (of course) and a very electronegative atom. Hydrogen bonding is the strongest of the intermolecular forces, stronger than a standard dipole-dipole attraction.
We only consider the H-O, H-F and H-N bonds to lead to hydrogen bonding.
H-O, H-F or H-N bonds are particularly polar due to the nature of the hydrogen atom. That atom has only one electron, and that atom is shared in the covalent bond with the other atom. With any of these three elements the electrons in the covalent bond are most likely to be located farther from the hydrogen atom and nearer to the very electronegative atom. This results in the nucleus of the hydrogen atom, and its positive charge, being very exposed – there are no further electrons shielding the positive charge.
This means there is a big charge difference between the δ- end and the δ+ ends of the bond, and that leads to strong attraction and strong intermolecular forces between adjacent molecules.
Relative Strength of Intermolecular Forces
The three types of intermolecular forces are all caused by attraction between areas of charge separation in adjacent molecules. But they don’t all have similar strength.
The strength of the attraction varies primarily due to the magnitude of the charge separation present.
We know that the charge separation is greatest in molecules with H-O, H-N or H-F bonds, and the charge separation is smallest in molecules that undergo only temporary or induced dipoles.
That means that hydrogen bonds are stronger forces than dipole-dipole attractions, and that dipole-dipole attractions are greater than van der Waals forces.
Why are Intermolecular Forces Important?
The amount of attraction between adjacent molecules of covalent substances makes a big impact on the physical properties of the substance.
In particular the melting point and boiling point of a substance are higher if that substance has stronger intermolecular forces.
Let’s compare the melting point and billing points of two comparable substances: water (H2O) and hydrogen sulfide (H2S).
Oxygen and sulfur are both elements in the same group, making this a fair comparison.
The hydrogen sulfide molecule is larger and heavier than a water molecule. Greater size and mass tend to cause a higher melting and boiling point.
Yet the melting point of hydrogen sulfide is around -85°C. The melting point of water is, of course, much higher at 0°C.
And the boiling point of hydrogen sulfide is -60°C (which is why we know H2S as a gas), compared to water’s 100°C.
The much higher melting point and boiling point for water is due to the presence of hydrogen bonding in water. Water has H-O bonds which lead to hydrogen bonding between adjacent water molecules. More energy is required to overcome the attraction between molecules, hence a higher melting point and boiling point.
Hydrogen sulfide has H-S bonds which lead to dipole-dipole attraction, and this attraction is significantly weaker than the hydrogen bonds.