Types of Forces between Molecules (AQA A Level Chemistry): Revision Note

Intramolecular Forces vs. Intermolecular Forces

Intramolecular forces

• Intramolecular forces are forces within a molecule and are usually covalent bonds

• Covalent bonds are formed when the outer electrons of two atoms are shared

• Single, double, triple and coordinate bonds are all types of intramolecular forces

Intermolecular forces

• Molecules also contain weaker intermolecular forces, which are forces between the molecules

• There are three types of intermolecular forces:

  • Induced dipole–dipole forces, also called van der Waals, London or dispersion forces

  • Permanent dipole – dipole forces are the attractive forces between two neighbouring molecules with a permanent dipole

  • Hydrogen Bonding is a special type of permanent dipole-permanent dipole forces

  • Intramolecular forces are stronger than intermolecular forces

    • For example, a hydrogen bond is about one-tenth the strength of a covalent bond

  • The strengths of the types of bond or force are as follows:

The varying strengths of different types of bonds

Polar Bonds

Polarity

• When two atoms in a covalent bond have the same electronegativity, the covalent bond is nonpolar

• When two atoms in a covalent bond have different electronegativities, the covalent bond is polar, and the electrons will be drawn towards the more electronegative atom

• As a result of this:

  • The negative charge centre and the positive charge centre do not coincide with each other

  • This means that the electron distribution is asymmetric

  • The less electronegative atom gets a partial charge of δ+ (delta positive)

  • The more electronegative atom gets a partial charge of δ- (delta negative)

• The greater the difference in electronegativity, the more polar the bond becomes

Permanent dipole - dipole forces:

• Polar molecules have permanent dipoles

• The molecule will always have a negatively charged and a positively charged end

• Forces between two molecules that have permanent dipoles are called permanent dipole-dipole forces 

• 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

Induced Dipole-Dipole Forces

Induced dipole-dipole forces:

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

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

    

• The electron charge cloud in nonpolar molecules or atoms is constantly moving

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

Relative strength

• For small molecules with the same number of electrons, permanent dipoles are stronger than induced dipoles

  • Butane and propanone have the same number of electrons

  • Butane is a nonpolar molecule and will have induced dipole-induced dipole forces

  • Propanone is a polar molecule and will have permanent dipole-dipole forces

  • Therefore, more energy is required to break the intermolecular forces between propanone molecules than between butane molecules

  • So, propanone has a higher boiling point than butane

Hydrogen Bonds

Hydrogen bonding

• Hydrogen bonding is the strongest form of intermolecular bonding

  • Intermolecular bonds are bonds between molecules

  • Hydrogen bonding is a type of permanent dipole–permanent dipole bonding

• For hydrogen bonding to take place, the following condition applies:

  • 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 bond with the lone pair of an O, N or F atom in another molecule

• For example, in water

  • Water can form two hydrogen bonds, because the O has two lone pairs