Solvent Choice (Edexcel International AS Chemistry): Revision Note

Solvent Choice

• 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

Haloalkanes

• Even though haloalkanes contain a polar bond, they are only partially soluble in water as they can not form hydrogen bonds

• This is because there are no H-F, H-O or H-N bonds within the molecule

Due to the difference in electronegativity between the carbon and halogen atoms, the C-X bond is polar (the difference gets smaller from C-F to C-I)

Ionic compounds

• Many ionic compounds will dissolve in polar solvents, e.g. water

• Solubility is dependent on two main factors:

  1. Breaking down the ionic lattice

  2. The polar molecules attract and surround the ions

• Polar molecules, such as water, can break down or disrupt the ionic lattice and surround each ion in solution

• The δ+ end of the polar molecule can surround the negative anion

• The δ- end of the polar molecule can surround the positive cation

• The solubility of an ionic compound depends on the relative strength of the electrostatic forces of attraction within the ionic lattice and the attractions between the ions and the polar molecule

• In general, the greater the ionic charge, the stronger the electrostatic attraction in the lattice and the less soluble an ionic compound tends to be

  • For example, sodium chloride (Na⁺ and Cl^-) dissolves readily in water, whereas magnesium oxide (Mg²⁺ and O^2-) is virtually insoluble 

  • This is a general rule, though, and there are many exceptions