Bonding notes IB syllabus > bonding (sl) > 4.3 

These notes were written for the old IB syllabus (2009). The new IB syllabus for first examinations 2016 can be accessed by clicking the link below.

IB syllabus for first examinations 2016

4.3 - Intermolecular forces


4.3.1: Describe the types of intermolecular forces (attractions between molecules that have temporary dipoles, permanent dipoles or hydrogen bonding) and explain how they arise from the structural features of molecules. The term van der Waals' forces can be used to describe the interaction between non-polar molecules.


Van der Waal's forces

All particles exhibit Van der Waal's forces. These are induced dipole - dipole attractions, meaning that the dipole that exists is only temporary in that it is caused by vibrations of the nucleus within the negative charge cloud.

Vibrations of the nucleus within the electron charge cloud creates a temporary dipole, shown here as delta + and delta - charges.

This causes sympathetic vibrations of the nucleus in neighbouring atoms, so that the delta + and delta - charges become arranged opposite one another, plus near minus.

There is now attraction between the opposite delta charges on neighbouring atoms, creating Van der Waal's forces.

Eventually all of the atoms or molecules vibrate in unison, making the partial charges enough to hold the structure together.

When enough energy is given to a substance held together by only Van der Waal's forces, they are overcome and the compound melts.

Factors affecting Van der Waal's forces

The factors that affect dipole -dipole interactions are illustrated very well by consideration of the melting and boiling points of the elements phosphorus, sulphur and iodine. All three substances are obviously made up of their own atoms and can possess no permanent dipoles. The forces between the molecules are created by induced dipole-dipole attractions. The greater the force the higher the melting and boiling points (more energy is needed to separate the molecules)

molecule relative molecular mass melting point /ºC boiling point /ºC
I2 254 114 184
P4 108 44 280
S8 256 113 445

As can be seen, there is a clear correlation between relative molecular mass and melting point. The higher the Mr the higher the melting point. Sulphur and iodine have almost identical relative masses (256 and 254 respectively) and they also have almost identical melting points of 113 and 114ºC.

Phosphorus, with a much lower relative mass, 108, has a much lower melting point, 44ºC

Phosphorus molecule Mr=108 Iodine molecule Mr=254 Sulphur molecule Mr=256

Boiling point discrepancy

However, the melting point only partially overcomes the intermolecular forces. To separate the molecules completely we must arrive at the boiling point (the point where the vapour pressure is equal to the atmospheric pressure).

Inspection of the boiling points of sulphur and iodine, which are very different, reveals that they seem to be experiencing very different forces between molecules than before. How can this be?

The answer lies in the second factor affecting Van der Waal's forces, the shape of the molecules.

Molecular shape

Molecules that are more sperical in shape have a lower surface area to volume ratio. They effectively have less surface area on which the Van der Waal's attractions can act.

This can be shown using simple maths by considering the surface area to volume ratio of a sphere and a cube (less sperical)

Shape Surface area Volume Surface area/ volume ratio
sphere (radius r) 4πr2 4/3πr3 0.484
cube (side =d) 6d2 d3 0.600

Where did these figures come from?

Something seems to happen to the sulphur molecules between melting and boiling points.

Careful investigation of molten sulphur during heating reveals several stages. At 114ºC the sulphur melts to form a light yellow mobile (runny) liquid. At this point the S8 molecules are separating.

On further heating the sulphur starts to darken, firstly to orange and then to red and becomes much more viscous (sticky). There is clearly an increase in intermolecular forces over this range of temperatures. This can be explained by the crown shaped molecules of S8 sulphur breaking one of the bonds in the crown and opening to become linear chains of S8.

Sulphur S8 crowns (yellow) heat Sulphur S8 chains (dark red)

The crown-shaped S8 molecules have a more spherical shape than the S8 long chain molecules. The longer chain-shaped molecules experience greater Van der Waal's induced dipole - dipole attractions, so a liquid made of S8 chains gets redder and more viscous (stickier) as the ratio of straight chain molecules to crown molecules increases.

If heating is continued the red molten S8 chain sulphur gets a little less viscous (sticky) as the energy available to overcome the Van der Waal's forces increases. There is a tendency of the sulphur to vaporise as S2 molecules near the boiling point.

Plastic sulphur

When molten (melted) sulphur is poured rapidly into cold water a form of sulphur rather like rubber is produced - called plastic sulphur. This only happens with the red molten sulphur that contains the S8 chains. If the light yellow molten sulphur (containing crown shaped molecules) is poured into water it simply solidifies as crystalline sulphur.

The S8 chain sulphur, however, cools too rapidly to allow the chains to reform into crowns and they lose energy while parallel to one another. This creates a structure with the chains all aligned but held only by the Van der Waal's forces (and some cross linking). The structure can be stretched like rubber.

Plastic sulphur is unstable at room temperature, as the S8 chains slowly return to S8 crown molecules. The rubber like structure gets more brittle as the crystals are slowly formed that use the S8 crowns as their molecular bulding bricks.


4.3.2: Describe and explain how intermolecular forces affect the boiling points of substances. The hydrogen bond can be illustrated by comparing physical properties of: H2O and H2S, HF and HCl, NH3 and PH3, CH3OCH3 and C2H5OH, C3H8, CH3CHO and C2H5OH.


 


Resources

Solubility demo - "like dissolves like" (University of Washington)


Useful links