As a general rule, \({\rm{M}}{{\rm{X}}_{\rm{n}}}\) molecules (where \({\rm{M}}\) represents a central atom and \({\rm{X}}\) represents terminal atoms; \({\rm{n = 2 - 5}}\)) are polar if there is one or more lone pairs of electrons on \({\rm{M}}\). \({\rm{N}}{{\rm{H}}_{\rm{3}}}\) (\({\rm{M = N, X = H, n = 3}}\)) is an example. There are two molecular structures with lone pairs that are exceptions to this rule. What are they?

A molecule is a group of two or more atoms bound together by chemical bonds; the term may or may not include ions that meet this condition depending on the context.

Polarity is the property of a molecule with a dipole moment, in which it can be said where the negative charge is denser, moreoften found.

Generally, dipole moment is the sum of all bond dipole moments.

Effect of electron pairs can be modelled as a highly polar bond in the direction of the pair is formed, then we can sum the dipole bondmoment to have dipole moment of the molecule.

Symmetrical molecules do not have a dipole moment, they are not polar.

Symmetry is intuitively understood property, mathematically this would mean that for each type of bond, sum of vectors in direction ofevery such bond is \({\rm{0}}\).

Of the molecular structures on table \({\rm{3}}{\rm{.19}}\) which have free electron pairs

• Linear structure - two bonds, three free electron pairs
• Square planar structure - four bonds, two electron pairsare symmetrical, therefore not polar.

This is additionally explained with the picture below - the arrows on bonds define directions in which the concentrations of negativecharge grow. Orientation is not defined for the bonds, but because of symmetry, in either case sum is \({\rm{0}}\).

Therefore, \({\rm{M}}{{\rm{X}}_{\rm{3}}}\) and \({\rm{M}}{{\rm{X}}_{\rm{4}}}\) are the exceptions obtained.