The situation with octahedral shaped molecule with lone pairs is more straight- forward than that
with trigonal bipyramidal molecules. The reason for this should be fairly clear, all positions in
an octahedral molecule are equivalent, the bond angles between
all the atoms in the molecule are 900.
Whereas in a trigonal bypramidal molecule we had to consider two different locations, the axial and equatorial positions
this is not the case with octahedral shaped molecules. The image below shows the
shapes of several octahedral molecules that
contain lone pairs of electrons. Recall that we do not consider the
lone of electrons when
determing the final shape. Molecules
with a single lone pair will always form square pyramidal molecules,
since all positions in
an octahedral molecule are equivalent, it does not matter where the
lone pair goes.
Molecules with two lone pairs can reduce the repulsion between the lone pairs by placing them 1800 apart. This will result in the formation of a square planar molecule, as shown below:
The 3 molecules shown below all have shapes based on a tetrahedral arrangement around the central atom. However as we have seen that lone pairs of electrons require more space than bonding electron pairs and this results in the reduction of the bond angles between the bonding pairs as shown below.
Molecules with multiple bonds, that is double or triple bonds between the atoms show a similar effect. Double bonds for example contain a higher electron density than single bonds and so will repel other bonding electron pairs in a manner similar to that of lone pairs. In the example below are shown 2 molecule which have a trigonal planar structure. In a trigonal planar molecule we might expect bond angles of 1200, however as shown below you can see that the double bond requires more space than the single bonds in trigonal planar molecules and the bond angles will be not be the expected 1200.