The electron pair repulsion theory states that the electron pairs in the valence energy level of an atom repel each other, and therefore are arranged as far apart as possible. For example, H2O:
             Due to this theory, different molecules with different amounts of pairs of electrons have different shapes.
             Some common shapes of molecules include linear, trigonal planar, tetrahedral, trigonal pyramidal and v-shaped (bent) molecules. Examples are drawn below-
             a) linear (CO2)
             The electron dot diagrams must be drawn first in order to work out the shape of the molecule.
             The polarity can be determined from the shape of the molecule. In essence, a molecule is polar if there is an overall electronegativity difference in the molecule. This can be determined using some basic rules, which state that:
             If a molecule is of the form AB2 and is linear, it is non-polar.
             If a molecule is of the form AB3 and is trigonal planar, it is non-polar
             If a molecule is of the form AB4 and is tetrahedral, it is non-polar.
             All other molecules are polar. Using these rules, we can determine the polarity of the earlier used examples.
             a) AB2 Ö b) AB4 Ö c) AB2 Ö d) AB3 Ö e) AB3 Ö
             linear Ö tetrahedral Ö bent X trigonal planar Ö trigonal pyramidal Ö
             non-polar non-polar polar non-polar polar
             As we have just seen, it is easy to determine the molecular polarity of a particular molecule after seeing the shape. Knowing the polarity and the nature of the molecul
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