3.1.3 Bonding - Forces acting between molecules
Forces between particles
Matter tends to stick together. We know this because solids and liquids exist!
This tells us that there are forces of attraction between particles of matter, the so-called interparticular forces. The types of particles vary from one compound to another, but there is only one fundamental force of attraction, which is electrostatic in nature, the force of attraction betweeen partial, or total, positive and negative charges.
Positive charges attract negative charges and vice versa.
Intermolecular forces in covalent molecules
When the particles that we are dealing with are molecules, the forces of attraction between the molecules, the interparticular forces, are called intermolecular forces. Covalent molecules are electronically neutral, so where do the positive and negative charges that give rise to electrostatic attraction come from?
The way forces are created is determined by the polarity of the compound. There are only two basic mechanisms.
- Temporary dipole -dipole interactions (Van der Waals' forces)
- Permanent dipole dipole interactions
However, permanent dipole - dipole interactions may be subdivided into:
- Permanent dipole-dipole interactions
- Hydrogen bonding (a special case of permanent dipole-dipole)
Van der Waal's forces
Van der Waals' forces of attraction (also known as London dispersion forces) exist between ALL particles. It is thought that they are due to vibration of the nucleus within the negative charge cloud, creating polarity of temporary positive and negative charge within molecules.
The vibrations set up sympathetic vibrations in neighbouring molecules that brings opposite partial charges into close proximity. It is this that causes the attraction between molecules.
This temporary state may cause attraction between two molecules, pulling them together. The magnitude of Van der Waals force depends on the relative molecular mass, high mass produces a larger force.
That van der Waals' forces can be strong is often ignored when dealing with intermolecular forces. However, as molecules get larger it becomes the dominant force. The important thing is not to cause confusion. Out of the possible intermolecular forces, all other things being equal, van der Waals' forces are the weakest, but this doesn't mean that they are necessarily weak.
In sulfur and iodine, for example, the force is strong enough to cause them to be solids at room temperature. The relative molecular mass of sulfur is 256 and the relative molecular mass of iodine is 254.
- sulfur Mr 256 m.p. = 113ºC
- Iodine Mr 254 m.p. = 114ºC
Notice the similarity between the relative masses and the melting points of sulfur and iodine. The slight difference can be attributed to the different molecular shapes giving rise to slightly different volume to surface area ratios (see below) in the two types of molecule.
Example: Explain, by reference to the intermolecular forces, why sulfur has a higher melting point than phosphorus.
Both sulfur and phosphorus are non-polar covalent solids. The only force of attraction that exists between molecules of both kinds is van der Waals' force. The strength of this force is dependent primarily on the relative molecular mass.
Sulfur exists as S8 molecules with a relative mass of 256.
Phosphorus exists as P4 molecules with a relative mass of 124.
Sulfur has a much higher relative mass than phosphorus, stronger van der Waals' forces and, because of this, a higher melting point.
Permanent dipole-dipole interactions occur between molecules that have permanent dipoles. As we have seen in the polarity section, a permanent dipole is created in a molecule when there is an electronegative atom, such as a oxygen, nitrogen or a halogen.
The permanent dipole consists of regions of partial positive charge and regions of partial negative charge within the same molecule.
By convention we use the small greek letter delta, d, to represent a partial charge. Hence, d+ means a partial positive charge and d- means a partial negative charge.
Permanent dipole - dipole interactions
In a body of polar molecules the dipoles can attract their opposite charges on neighbouring molecules. The partial positive charges are attracted to the partial negative charges on other molecules.
The consequence is that the molecules require larger forces to pull them apart. This means that the boiling point and the enthalpy of vaporisation is higher.
Comparison with Van der Waals' forces
You should remember at this point that ALL molecules have van der Waals' forces. So any comparison must take this into account.
The best way to compare then is to use molecules that have similar relative mass. The vander Waals' forces must be similar in the two types of compound as they are a function of the relative mass. Hence, any difference can be thought of as due to the permanent dipole-dipole forces.
Alkanes are examples of compounds that have only van der Waals' forces.
|Alkanes||Mr||b.p. /ºC||Aldehydes||Mr||b.p. /ºC|
We can see from the table that whereas the relative mass is the same between alkanes and aldehydes, the boiling point of the aldehydes is much higher. This can only be due to some other form of intermolecular force, other than van der Waals' force.
It is, of course, due to dipole- dipole forces acting between the relatively negative oxygen on the carbonyl group of the aldehyde and the relatively positive carbon of the carbonyl group.
Hydrogen bonding is a special case of dipole-dipole attractions.
When hydrogen is bonded to nitrogen, oxygen or fluorine, a very strong dipole is formed, making the hydrogen very strongly (partially) positive. This hydrogen is then attracted to the lone pairs on other similar molecules (nitrogen, oxygen and fluorine all have lone pairs) forming a hydrogen bond, which is stronger than van der Waal's or dipole-dipole, but weaker than normal covalent bonding.
The reason for the strength of the positive charge on the hydrogen atoms, lies in the fact that it has only one bonding pair of electrons. Usually this pair is shared equally between hydrogen and the other bonded atom. However, in the case of hydrogen attached to oxygen, the oxygen pulls the electrons towards it leaving the hydrogen 'naked'.
As the hydrogen atom has no more electron shells, effectively it is the nucleus that is revealed. This is extremely small compared to an atom, so the developing partial charge has a much higher charge density than normal.
The group 6 hydrides
The effect of hydrogen bonding on intermolecular forces can be demonstrated very well by studying the boiling points of the group 6 hydrides. As expected the general trend is increased boiling point with increased relative molecular mass (as the van der Waals' force increases. However, water is completely anomalous to the trend, as it has a much higher boiling point than expected.
This anomalous boiling point can be explained by the high degree of hydrogen bonding between water molecules. Each molecule has two partially positive hydrogen atoms and it also has two lone pairs of electrons on the oxygen atom.
This means that the water molecules are able to form two hydrogen bonds per molecule with other water molecules.
The other members of the group VI hydrides show the effect of increasing relative molecular mass on boiling point. This is to be expected, as the increase in relative molecular mass causes an increase in the strength of the van der Waals' forces.
Hence there is a gradual but regular increase from H2S to H2Se to H2Te.
Group VI hydrides
So how do we decide which is the most important force within a molecule?
The answer is that we can only compare molecules in which we can 'cancel out' factors. For example, it doesn't make sense to try and compare a very large molecule with only van der Waals' forces, with a small molecule having hydrogen bonding. Each has a high degree of intermolecular interaction, but they arise from different sources, so can't fairly be compared.
We can. however, compare molecules with similar relative molecular mass, 'cancelling out' the van der Waals' influence, in which case any difference in boiling point is due to either permanent dipole-dipole interactions, or hydrogen bonding.
Example: Compare and explain the boiling points of the molecules CH3OCH3 and C2H5OH.
Methoxymethane is a slightly polar molecule, with a partial negative charge on the oxygen. Ethanol is a polar molecule with an O-H group. Both molecules have the same relative mass and so the influence of van der Waals' forces can be discounted in any comparison.
The boiling point of ethanol is much higher than methoxymethane, as ethanol is able to form hydrogen bonds, whereas methoxymethane has rather weaker dipole-dipole interactions.
Some molecules have the ability to 'link' to other similar molecules using hydrogen bonding, in such a way that the molecules behave as two in one, so to speak. These double molecules are called dimers.
Ethanoic acid is able to form two hydrogen bonds with another ethanoic acid molecule, making a structure with double the relative molecular mass.
Ethanoic acid has a boiling point of 118ºC, reflecting the high apparent relative molecular mass of 120.
Ethanoic acid dimer