Molecular Polarity (DP IB Chemistry): Revision Note

Molecular polarity

• To determine whether a molecule with more than two atoms is polar, consider:

  • The bond polarity

  • The arrangement of those bonds in space, i.e. the molecular geometry

Molecules with only nonpolar bonds 

• There is no difference in electronegativity between the atoms

  • So, there is no dipole in the bond

  • This means that there is no net dipole for the whole molecule

• Examples include:

  • H₂

  • Cl₂

• It is not possible for a molecule with only nonpolar bonds to be polar

  • If the bonds have no dipole, the molecule cannot have one either

Molecules with polar bonds

• There is a difference in electronegativity between the atoms

  • So, the bonding electrons are shared unevenly creating bond dipoles

• These molecules can be:

  • Nonpolar

  • Polar

Nonpolar overall

• If the bond dipoles are arranged symmetrically, they cancel out

• The molecule has no net dipole moment and is nonpolar

• Examples:

  • CO₂: two C=O dipoles pull in opposite directions

  • CCl₄: four polar C–Cl dipoles arranged symmetrically in a tetrahedral shape

Polar overall

• If the bond dipoles are arranged asymmetrically, they do not cancel

• This creates a net dipole moment, and the molecule is polar

• Examples:

  • H₂O: bent shape causes dipoles to reinforce

  • NH₃: trigonal pyramidal shape creates a net dipole

  • CH₃Cl: tetrahedral shape, but Cl pulls more strongly than H