Intermolecular Forces (OCR AS Chemistry A): Revision Note

Intermolecular Forces

Intramolecular forces

• Intramolecular forces are forces within a molecule

  • These are typically covalent bonds

• Covalent bonds are formed when atoms share outer electrons

• Single, double, triple and co-ordinate bonds are all types of intramolecular forces

Intermolecular forces

• Molecules also experience weaker intermolecular forces, which act between molecules

• There are three types of intermolecular forces:

  • Induced dipole – dipole forces 

    • Also known as London forces or London dispersion forces 

  • Permanent dipole – dipole forces

    • These are the attractive forces between polar molecules with permanent dipoles

  • Hydrogen Bonding 

    • This is a particularly strong type of permanent dipole - permanent dipole force

  • Intramolecular forces are stronger than intermolecular forces

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

Induced dipole-dipole forces:

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

  • They are also known as London 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

• Therefore the greater the number of electrons the molecule has or the greater the relative molecular mass, the stronger the induced dipole-dipole forces

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

Permanent dipole - permanent dipole forces:

• Polar molecules have permanent dipoles

• The molecule will always have a negatively and 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

Relative strength

• Permanent dipole–dipole forces are stronger than induced dipole–dipole forces in small molecules with similar electron counts

• For example, butane and propanone:

  • Butane (34 electrons) and propanone (32 electrons) have a similar number of electrons and the same relative molecular mass (58)

  • Butane is a nonpolar molecule with induced dipole forces

  • Propanone is a polar molecule with permanent dipole forces

  • More energy is required to break the intermolecular forces between propanone molecules than between butane molecules

  • Therefore, propanone has a higher boiling point than butane

Pd-pd forces are stronger than id-id forces in smaller molecules with a similar number of electrons

Hydrogen Bonding & Water

Hydrogen bonding

• Hydrogen bonding is the strongest form of intermolecular bonding

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

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

• For example, in water

  • Each water molecule can form up to four hydrogen bonds: two through the two lone pairs on its O atom and two through its two H atoms

Properties of water

• Hydrogen bonding in water, causes it to have anomalous properties such as high melting and boiling points, high surface tension and anomalous density of ice compared to water

High melting & boiling points

• Water has high melting and boiling points which is caused by the strong intermolecular forces of hydrogen bonding between the molecules

• In ice (solid H₂O) and water (liquid H₂O) the molecules are tightly held together by hydrogen bonds

• A lot of energy is therefore required to overcome the hydrogen bonds between the water molecules in order to melt or boil them

Hydrogen bonds are strong intermolecular forces which are difficult to break causing water to have high melting and boiling points

Effects of Intermolecular Forces

Properties of Molecular substances 

Ice 

• Solids are denser than their liquids as the particles in solids are more closely packed together than in their liquid state

• In ice however, the water molecules are packed in a 3D hydrogen-bonded network in a rigid lattice

• Each oxygen atom is surrounded by hydrogen atoms

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

• Therefore, ice has a lower density than liquid water

The ‘more open’ structure of molecules in ice causes it to have a lower density than liquid water

Iodine

• The molecular lattice of iodine consists of a neat arrangement of molecules in the crystal lattice which is held together by the weak intermolecular forces 

• Being a non-polar molecule, the weak intermolecular bonding is due to induced dipole - dipole interactions, also called London forces

• Iodine tends to sublime at temperatures approaching 114 ^oC due to weak intermolecular force

  • A purple vapour is observed when iodine sublimes

Crystalline structure of Iodine

Solubility

• The general principle is that 'like dissolves like' so non-polar substances mostly dissolve in non-polar solvents, like hydrocarbons and they form dispersion forces between the solvent and the solute

• Polar covalent substances generally dissolve in polar solvents as a result of dipole-dipole interactions or the formation of hydrogen bonds between the solute and the solvent

• A good example of this is seen in organic molecules such as alcohols and water:

Hydrogen bonds form between ethanol and water

• As covalent molecules become larger their solubility can decrease as the polar part of the molecule is only a smaller part of the overall structure

  • This effect is seen in alcohols for example where ethanol, C₂H₅OH, is readily soluble but hexanol, C₆H₁₃OH, is not

• Polar covalent substances are unable to dissolve well in non-polar solvents as their dipole-dipole attractions are unable to interact well with the solvent

• Giant covalent substances generally don't dissolve in any solvents as the energy needed to overcome the strong covalent bonds in the lattice structures is too great

Conductivity

• As covalent substances do not contain any freely moving charged particles, they are unable to conduct electricity in either the solid or liquid state

• However, under certain conditions some polar covalent molecules can ionise and will conduct electricity

• Some giant covalent structures are capable of conducting electricity due to delocalised electrons

Comparing the Properties of Covalent Compounds Table

• Comparison of covalent and ionic substances:

  • Melting and boiling point: non-polar covalent low; polar covalent low; giant covalent very high; ionic very high

  • Volatility: non-polar covalent highest; polar covalent high; giant covalent low; ionic low

  • Solubility in polar solvents: non-polar covalent insoluble; polar covalent some solubility depending on molecular size; giant covalent insoluble; ionic soluble

  • Solubility in non-polar solvents: non-polar covalent soluble; polar covalent some solubility depending on molecular size; giant covalent none; ionic insoluble

  • Electrical conductivity: non-polar covalent none; polar covalent none; giant covalent none (except graphite and graphene); ionic only when molten or aqueous