In Covalent Bonding- How do the shared and unshared (Lone) valence electron pairs around a central atom influence molecular shape and overall polarity?
There are a set of guidelines to remember the shape of covalent molecules by looking through its central atom and applying VSEPR rules. The steps are:
1) Write the molecule in Lewis form.
2) Count the total number pairs of valence electrons around the central atom, by the formula (V+A±-charge)/2, where v=no. of own valence electrons of the central atom, A = no. of surrounding bonds from it.
3) Subtract the number of pi-electron pairs (if any) to get X,
4) If X is, 2 = linear, 180 degrees, 3 = trigonal planar, 120 degrees, 4 = tetrahedral, 109.5 degrees, 5 = trigonal bipyramidal, 120, 180 & 90 degrees, 6 = Octahedral (Square bipyramidal) 90 & 180 degrees.
5) Look for the presence of lone electron pairs from the difference in count obtained in 2) and the total number of bonds. If present, modify the structure and bond angles, keeping in mind that the lone pairs repel other lone pairs or bond pairs most dominantly. For molecules where lone pair is absent, pure bond angles and geometries stated above shall prevail.
In order to get overall polarity of a molecule one has to derive the resultant of all the covalent bond polarities and lone pair polarities. For example in NH3 molecule, the three N-H bond polarities and the lone pair polarities are in a reinforcing orientation, whereas in NF3, the three N-F bond polarities and the lone pair polarities are in opposing disposition. This makes NH3 molecule more polar than the NF3 molecule, though both of these molecules are of similar shape.