Bond Polarity & Dipole Moments (Cambridge (CIE) A Level Chemistry): Revision Note

Bond Polarity & Dipole Moments

• Electronegativity is the ability of an atom to draw a pair of electrons towards itself in a covalent bond

• Electronegativity increases across a Period and decreases going down a Group

Polarity

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

Bonding electrons in a chlorine molecule

• 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 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

Bonding electrons in a hydrogen chloride molecule

Dipole moment

• The dipole moment is a measure of how polar a bond is

• The direction of the dipole moment is shown by the following sign in which the arrow points to the partially negatively charged end of the dipole:

Representing dipoles

Assigning polarity to molecules

• To determine whether a molecule with more than two atoms is polar, the following things have to be taken into consideration:

  • The polarity of each bond

  • How the bonds are arranged in the molecule

• Some molecules have polar bonds but are overall not polar because the polar bonds in the molecule are arranged in such a way that the individual dipole moments cancel each other out

Polarity in chloromethane

Polarity in tetrachloromethane