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 inter-particular 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.

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Intermolecular forces between 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 (London dispersion 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)

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Dispersion forces

Induced dipole - dipole 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 dispersion force depends on the relative molecular mass, high mass produces a larger force.

That dispersion 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, dispersion 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.

Dispersion Forces

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Trends in organic chemistry

The homologous series that characterise organic chemistry allow us to see clearly the relationship between relative mass and dispersion forces. A graph plotted of mass against boiling point gives a (very nearly) straight line.

Boiling points of some organic molecules

Notice that methylpropane, CH₃CH(CH₃)CH₃, has a lower boiling point than its straight chain isomer, propane, CH₃CH₂CH₂CH₃. This is due to the 'fine-tuning' of dispersion forces by the surface area to volume ratio difference between the two molecules.

butane	methylpropane

Geometrically speaking, the shape with the lowest surface area to volume ratio is a sphere.

Dispersion forces act at the surface of molecules, so the greater the surface area, the stronger the force (once again, all other things being equal).

A branched chain is a closer approximation to a sphere than a straight chain. Hence, straight chains have a higher boiling point than branched chain molecules, they have a larger surface area at which attractions occur.

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Permanent dipole - dipole interactions

In a body of polar molecules the dipoles can attract their oppisite 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.

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Comparison with dispersion forces

You should remember at this point that ALL molecules have dispersion 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 dispersion forces.

We can see from the table that whereas the relative mass is the same between alkanes and alkanals (aldehydes), the boiling point of the alkanals is much higher. This can only be due to some other form of intermolecular force, other than dispersion force. It is of course due to dipole - dipole forces acting between the relatively negative oxygen on the carbonyl group of the alkanal and the relatively positive carbon of the carbonyl group.

Hydrogen bonding

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 dispersion forces or dipole-dipole interactions, 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.

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The group 15, 16 and 17 hydrides

The effect of hydrogen bonding on intermolecular forces can be demonstrated very well by studying the boiling points of the group 16 hydrides. As expected, the general trend is increased boiling point with increased relative molecular mass (as the dispersion 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 16 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 dispersion forces.

Hence there is a gradual but regular increase from H₂S to H₂Se to H₂Te

The group 15 hydrides

The group 17 hydrides

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Comparing molecules

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 dispersion 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 dispersion influence, in which case any difference in boiling point is due to either permanent dipole-dipole interactions, or hydrogen bonding.

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Dimerisation

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.

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Solubility

The process of dissolution involves the solvent forming bonds with the solute molecules. This bond formation is exothermic, the energy being used to break the interparticular forces holding the structure together. There is also an increase in entropy when a solid dissolves, which contributes to the solubility according to Gibbs Free Energy for the process (section 4.62). If the overall Gibbs Free Energy change is negative, the substance is soluble. (remember: ∆G = ∆H - T∆S)

Substances which hydrogen bond usually dissolve in water. This is because the water molecules themselves can hydrogen bond to the solute. This bond formation makes the process favourable.

Ethanol is an alcohol which is miscible in all proportions with water. This means that it forms hydrogen bonds easily with the water. The solubility of the alcohols decreases as the relative molecular mass increases. As we ascend the homologous series of alcohols the relative mass increases by adding -CH₂- units to the hydrocarbon chain.

The increasing mass means that dispersion forces become more important and the alcohol molecules begin to form intermolecular bonds with each other more easily than with the water molecules.

Intermolecular forces summary

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