Orbital hybridization — the stork

In Grade 11 chemistry, I teach atomic orbitals as part of a bigger lesson on electron configuration. This is super great — students gain insight into why the periodic table looks like it does: s-block, p-block, d-block, f-block.1 Good stuff, since it layers beautifully on Mendeleev’s empirically-based periodic table. In AP Chemistry, I toss ’em a lesson on Quantum Numbers,2 which ices the cake. We go on to review Lewis structures, work through the formal charge concept and resonance, eventually landing in VSEPR-land.

This is all fabulous, but on closer inspection, the atomic orbitals that we know and love — and devoted time to studying — do not mesh with VSEPR Theory. Atomic orbitals alone do not account for the bond angles in real molecules as predicted by VSEPR. Traditionally, this is why we have taught orbital hybridization — to reconcile atomic orbitals with molecular shape. This concept, proposed by Linus Pauling,3 makes a heckuva lot of sense. Students may find it a little weird at first, but once explained, they learn its value. Further, it allows us to understand that all bonds — σ and π — are not equal, which pays dividends in the study of organic chemistry.

But if you dig into the literature, there are problems with the hybridization model.4 The calculated non-equivalence of the lone pairs of electrons on the water molecule and the non-involvement of d-orbitals in hyper-valent compounds are two examples.5 To this end, recent articles urge teachers not to instruct the traditional hybridization model.6

But I do it anyway.

Chemistry, even at the high school level, can be complicated: one size does not fit all. A model such as orbital hybridization has its good points and its shortcomings. My students understand the advantages of orbital hybridization while realizing that it isn’t the final word. Those who proceed to more advanced courses will gain a deeper understanding of chemical bonding.

The orbital hybridization model explains a great deal in a manner that teenaged minds can comprehend. We’re teaching high school students, not chemistry undergraduates. We teach this for the same reason that a mother may tell her child that babies are delivered by the stork or that they “come from mommy’s tummy”, it’s a decent yet imperfect explanation that a youngster can wrap his head around.

Sometimes we walk a fine line: we need to explain things the best way we can. But the explanation has gotta “sell”.


  1. Sounds like a prison for chemists. Think “Heisenberg”.
  2. Recently removed from the AP curriculum, but I teach it.
  3. L. Pauling, Journal of the American Chemical Society, 53 (4),1367–1400, 1931.
  4. See for example https://en.wikipedia.org/wiki/Orbital_hybridisation.
  5. J.M. Galbraith, Journal of Chemical Education, 84, 5–8, 2007.
  6. A. Grushow, Journal of Chemical Education, 88, 860–862, 2011.