Weak chemical bonds – Intermolecular bonds 1. Van der Waals forces Intermolecular forces are attractions between one molecule and neighboring molecules. All molecules are under the influence of intermolecular attractions, although in some cases those attractions are very weak. These intermolecular interactions are known as van der Waals forces. Even in a gas like hydrogen (H2), if the molecules are slowed down by cooling the gas, the attractions become large enough so the molecules will stick together and form a liquid and then a solid.

The attractions between the H2 molecules are so weak that the molecules have to be cooled to 21 K (-252C) before the attractions are enough to form liquid hydrogen. Helium’ s intermolecular attractions are even weaker – the molecules won’t stick together to form a liquid until the temperature drops to 4 K ( -269 C). Attractions are electrical in nature. In a symmetrical molecule like hydrogen, however, it doesn’t seems to be any electrical distortion to produce positive or negative parts.

H2 symmetrical molecule But that’ s only true in average. In the next figure the symmetrical molecule of H2 is represented. + - H H2 symmetrical molecule 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. This end of the molecule becomes slightly negative (charge -). The other end will be temporarily short of electrons and so becomes slightly positive (+). + + - - H Temporary dipole of H2 An instant later the electrons may move to the other end, reversing the polarity of the temporary dipole of molecule.

This phenomena even happens in monoatomic molecules of rare gases, like helium, which consists of a simple atom. 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. + - + - Temporary dipole of He

How temporary dipoles give intermolecular bonds? A molecule which has a temporary polarity approaches another molecule which happens to be non-polar at that moment. The electrons from the non-polar molecule will be attracted by the slightly positive end of the polar molecule. This is how an induced dipole is forming. + - original temporary dipole non-polar molecule + - + - original temporary dipole induced dipole Dipole-dipole attraction (van der Waals forces)

An instant later the electrons in the left hand molecule can move up to the other end. So, they will repel the electrons in the right hand molecule. The polarity of both molecules reverses, but there is still attraction between - end and + end. As long as the molecules stay close to each other, the polarities will continue to fluctuate in synchronization so that the attraction is always maintained. This phenomena can occur over huge numbers of molecules. The following diagram shows how a whole lattice of molecules could be held together in a solid.

Molecular distribution in a solid + - Molecular distribution in a solid The interactions between temporary dipoles and induced dipoles are known as van der Waals dispersion forces .

Hydrogen bond If we plot the boiling points of the hydrides of the elements of groups 15, 16 and 17 we find that the boiling point of the first elements in each group is abnormally high.

In case of ammonia NH3, water H2O and hydrofluoric acid HF, there must be some additional intermolecular forces of attraction, requiring significantly more heat energy to break them. These relatively powerful intermolecular forces are called hydrogen bonds. Hydrogen bonds are stronger than van der Waal dispersion forces, but weaker than covalent or ionic bonds. Hydrogen bonds can form if: hydrogen is attached directly to one electronegative element (F, O, N) each of the elements to which the hydrogen atom is attached have one “active“ lone pair of electrons.

Let’s consider two water molecules coming close together: The slightly + charge of hydrogen is strongly attracted to the lone pair of electrons; as a result a coordinate bond is formed. This is a hydrogen bond. In liquid water, hydrogen bonds are constantly broken and reformed.

In solid water each water molecule can form hydrogen bonds with other 4 surrounding water molecules, creating a 3-D structure of ice. As a result, the boiling point of H2O is higher than that of NH3 or HF.

In the case of ammonia, the amount of hydrogen bonding is limited by the fact that each nitrogen atom has only one lone pair. As well, in hydrogen fluoride, the number of hydrogen atoms is not enough to form a three- dimensional structure.