# Why do some molecules form T shapes instead of trigonal planar shapes [duplicate]

Consider a molecule such as $\ce{ClF_3}$.
Shouldn't the electron clouds (which are more negative than the Chlorine atom) want to repel each other and so the $\ce{ClF_3}$molecule should arrange into a trigonal planar shape where the electron clouds are as far away as possible?

I think at a deeper level than VSEPR, Chlorine is either partially ionized to a hybrid of $\ce{F^- + Cl^+F_2}$ configurations or has some electrons promoted up to d shaped orbitals but I don't know how to calculate the resulting shape. I think the main thing throwing me off is that the $\ce{F^- + Cl^+F_2}$ shapes should have the $\ce{F^-}$ ion repulsed from the lone pair electrons.

• There is a 3-center 4-electron bond between the chlorine and two fluorine atoms. This bond is linear, so the fluorines are necessarily on opposite sides of the chlorine. – f'' Jul 16 '16 at 20:23
• A more general question is $$\mathrm{Why\ does\ VSEPR\ predict\ a\ T-shaped\ geometry\ for AX_3 E_2 \ and\ not\ trigonal\ planar?}$$ We also have to remember that VSEPR was developed to rationalize what we were learning about electrons, bonding, and quantum mechanics withe known geometries. – Ben Norris Jul 16 '16 at 21:08
• – ron Jul 16 '16 at 21:52