Forces Between Molecules (Oxford AQA International A Level Chemistry)

Permanent Dipole–Dipole Forces

• There are three types of intermolecular forces

• Temporary dipoles exist in all molecules, but in some molecules there are permanent dipoles

Permanent dipole-dipole interactions

• This is an attraction between a permanent dipole on one molecule and a permanent dipole on another.

• Permanent dipole-dipole bonding usually results in the boiling points of the compounds being slightly higher than expected from temporary dipoles alone

  • It slightly increases the strength of the intermolecular attractions

Comparing butane and propanone

• For small molecules with the same number of electrons, dipole-dipole attractions are stronger than dispersion forces

Diagram to show the structures of butane and propanone

Table to show the difference in intermolecular forces between butane and propanone

Induced Dipole–Dipole Forces

• Induced dipole - dipole forces exist between all atoms or molecules

  • They are also known as van der Waals forces or London dispersion forces

    

• The electron charge cloud in non-polar molecules or atoms are constantly moving

• During this movement, the electron charge cloud can be more on one side of the atom or molecule than the other

  • This causes a temporary dipole to arise

• This temporary dipole can induce a dipole on neighbouring molecules

  • When this happens, the δ+ end of the dipole in one molecule and the δ- end of the dipole in a neighbouring molecule are attracted towards each other

  • Because the electron clouds are moving constantly, the dipoles are only temporary

• The strength of the induced dipole-dipole forces depend on:

  • The number of electrons the molecule has

    • More electrons = stronger induced dipole-dipole forces

  • The relative molecular mass of the molecule

    • Higher mass = stronger induced dipole-dipole forces

  • For example, pentane, C₅H₁₂ has a higher boiling point than propane, C₃H₈ 

Hydrogen Bonding & Water

• Hydrogen bonding is the strongest type of intermolecular force

• For hydrogen bonding to take place the following is needed:

  • A species which has an O, N or F (very electronegative) atom bonded to a hydrogen

• When hydrogen is covalently bonded to an O, N or F, the bond becomes highly polarised

• The H becomes so δ⁺ charged that it can form a hydrogen bond with the lone pair of an O, N or F atom in another molecule

• For example, in water

  • Water can form four hydrogen bonds, because the O has two lone pairs and there are two δ⁺ H atoms

Hydrogen bonding around one water molecule

Hydrogen bonding between water and ammonia

High melting and boiling points

• Hydrogen bonding in water, causes it to have anomalous properties such as high melting and boiling points

  • A lot of energy is therefore required to separate the water molecules and melt or boil it

• The graph below compares the enthalpy of vaporisation (energy required to boil a substance) of different hydrides

• The enthalpy changes increase going from H₂S to H₂Te due to the increased number of electrons in the Group 16 elements

• This causes an increase in the instantaneous dipole - induced dipole forces (dispersion forces) as the molecules become larger

• Based on this, H₂O should have a much lower enthalpy change (around 17 kJ mol^-1)

• However, the enthalpy change of vaporisation is almost 3 times larger which is caused by the hydrogen bonds present in water but not in the other hydrides

Graph to show the enthalpy of vaporisation of different hydrides

Ice

• When water freezes, the water molecules remain fixed in position to form a three-dimensional crystalline structure

• This way of packing the molecules and the relatively long bond lengths of the hydrogen bonds means that the water molecules are slightly further apart than in the liquid form

  • This means that ice is less dense than water, which explains why ice floats in water

The three-dimensional crystalline structure of ice