#H_2O# forms a tetrahedral shape due to the two lone pairs. Why can't this rule be applied to all molecules with lone pairs?

1 Answer
Jun 27, 2016

Answer:

Because we do not describe molecular geometry on the basis of the orientation of the lone pairs.

Explanation:

You are absolutely right that the lone pairs are stereochemically active, and influence geometry, however, the lone pairs are simply along for the ride.

If you consider the chloride ion, #Cl^-#, there are 4 lone pairs around the ion, which are certainly arranged in a tetrahedron, even though we cannot interrogate its shape as the chloride ion.

Sequentially substitute the lone pairs with oxygen:#ClO^(-); ClO_2^(-); ClO_3^(-);ClO_4^-#, the chlorine oxidation state sequentially ascends, as does symmetry, and you finish with perchlorate ion, #ClO_4^-#, a tetrahedron, with a formal #""^(+VII)Cl# centre. I leave it to you to determine the symmetry of the other oxychlorides on the basis of VESPER.