Draw the Lewis Structure
Determine the Shape
Determine the Polarity
First, count the number of substituents around the central atom, remembering
that a non-bonding pair (lone pair) of electrons counts as a single
substituent, as does any type of bonded atom.
Determine the geometry according to the VSEPR model (use Table 6.2 as a guide).
Decide on the shape of the molecule by placing substituents
(atoms or lone pairs) into the geometry, and considering that the lone
pairs are "invisible". Note that in linear, triangular planar,
and tetrahedral geometries, all of the substituent locations are
First, look up the electronegativity value for each atom in the molecule (see Fig 6. ). Remember that this value gives a relative measure of how strongly a particular kind of atom pulls on the electrons in a bond.
Next, take the difference between each pair of electronegativity values in order to determine the polarity of each individual bond in the molecule. The difference reflects how much harder one atom pulls on the electrons than the other. If the difference is very small, then the electrons will be equally balanced between the two atoms (i.e., a nonpolar bond). If the difference is large, then the electrons will be much closer to the atom with the higher electronegativity, and therefore that end of the bond will be negatively charged.
Finally, try to assess how all of the bond polarities in the
molecule will "add up," while trying to visualize the 3-dimensional
shape of the molecule. Sometimes the polarities of the bonds may
cancel each other altogether, yielding a nonpolar molecule (e.g., CO2).
Other times, polarities will partially cancel, but will still result in
a net polarity (e.g., H2O and H2CO). Note
that in H2CO, the C-H bonds have very little polarity (and
much of that cancels out), but that the polar C=O bond causes the
molecule to be polar.