Chemistry - Hybridisation of similar energy orbitals

Solution 1:

I would suggest that this entire factory of bogus hybridization stories should be eliminated from general and organic chemistry curricula. The word mixing is worse, as it is a jug where one can add different orbitals and run a blender and here is the mixture. This stuff is good for rote memorization only.

Only after a thorough QM course and mathematical foundation, these ideas should be introduced with an example along with historical development. Neither the textbook author, nor the teacher know what it is. The only real experimental observable is only electron densities via X-ray diffraction or other techniques.

Solution 2:

M. Farooq is correct that hybridization is a concept that seems to lead students astray more often than not, but your intuition within this simple concept is entirely correct.

When s, p and d orbitals are all involved in bonding (I'll avoid the word "hybridization"), it is most often the orbitals that are close in energy, for example, 4s, 3d and 4p rather than 4d with 4s and 4p. This is why d orbital involvement with bonding is most often discussed with respect to transition metals (d-block elements).

For p-block elements whose only empty d orbitals are from the same shell as the p orbitals (eg 4d with 4p), d orbital participation in bonding is now known to be very limited because those d orbitals are so high in energy relative to the p. The same applies to the s orbitals that you propose to involve in bonding. Unfortunately, many textbooks and teachers still refer to this outdated model of hypervalent sulfur and phosphorus compounds, for example, as using those higher energy d orbitals for bonding, which is likely the source of your confusion.

Once you realize that d orbitals from the same shell that are much higher in energy are not generally involved in bonding, then your question becomes moot.