#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

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


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.