While going through molecular orbital theory, my chemistry teacher mentioned that the reason why there is no significant interaction between $\ce {s}$ orbitals on one atom and $\ce {p_x}$ and $\ce {p_y}$ orbitals on another atom is that the $\ce {s}$ orbitals interact with both lobes of the $\ce {p}$ orbitals, having both in phase and anti phase interactions with these $\ce {p}$ orbitals, taking the z axis to be the internuclear axis.

However, I read that orbitals which do not have the right symmetry simply cannot interact, i.e. their overlap integral is necessarily $\ce {0}$. Thus, is the reasoning given by my chemistry teacher actually valid?

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    $\begingroup$ There is no contradiction behind your "however". It is the same in other words. $\endgroup$ – Ivan Neretin Jan 12 '19 at 14:49

According to my chem textbook: "Molecular orbitals obtained from 2px and 2py orbitals are not symmetrical around the bond axis because of the positive lobes above and below the molecular plane." Your instructor might be pointing out to the bond-symmetry thing in MO formation between pz and s orbitals extrapolated in most chemistry textbooks. But orbital overlap is a result of constructive and destructive interference. As s orbitals are spherical, they have infinite possibilities of orbital overlap with other orbital shapes and thus even overlap with d orbitals (as explained by VSEPR: Valence Shell Electron Pair Repulsion Theory in a similar manner). Thus s and px and s and py orbitals do overlap; if you take px to be of lower energy than py and pz, s-px overlapping occurs in boranes (formed by boron, which has one electron in 2px orbital, and hydrogen, which has one electron in 1s orbital), that often dimerize. This is also the case with methane, where px, py-s overlapping occurs, assuming that in sp3-hybridized state px, py orbitals in C bond with H 1s orbitals.


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