Edexcel A Level Chemistry:复习笔记1.5.2 Intermolecular Forces

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 co-ordinate bonds are all types of intramolecular forces

1.3-Chemical-Bonding-Inter-and-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 or London dispersion forces
    • Permanent dipole – dipole forces are the attractive forces between two neighbouring molecules with a permanent dipole
    • Hydrogen Bonding are 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:

     

1.3-Chemical-Bonding-Strengths-of-different-types-of-bonds

The varying strengths of different types of bonds

 

Induced dipole-dipole forces:

  • Induced dipole - dipole forces exist between all atoms or molecules
    • They are also known as London dispersion forces1.3-Chemical-Bonding-Intermolecular-Forces-2

     

  • 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

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 forces
    • Propanone is a polar molecule and will have permanent 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

     

1.3-Chemical-Bonding-Pd-Pd-vs-Id-Id

Pd-pd forces are stronger than id-id forces in smaller molecules with an equal number of electrons

 

Permanent dipole - dipole forces:

  • Polar molecules have permanent dipoles
  • The molecule will always have a negatively and positively charged end

1.3-Chemical-Bonding-Permanent-Dipole-Permanent-Dipole

  • 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

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

     

1.3-Chemical-Bonding-Water-H-Bonds

Hydrogen bonding in water

 

 

Exam Tip

Make sure to use a dashed, straight line when drawing your intermolecular forces! Hydrogen bonds should start at the lone pair and go right up to the delta positive atom - it must be really clear where your H bond starts and ends.

 

 

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