91.CHEMICAL BONDING (38)- Covalent Bonding(37) – Molecular Orbital Theory(12) – Heteronuclear Diatomic molecules(1)

With this post we start discussing MOT theory for diatomic heteronuclear molecules – molecules formed by two different atoms.

The MO diagrams for heteronuclear atoms are slightly different than the homonuclear ones.This is because in heteronuclear molecules, the contribution from two different atoms is different.

It is imperative that for two different atoms to interact, they must have similar energies and should have same symmetry characteristics w.r.t the internuclear axis.If these conditions are not met then there can be no overlap between them.

The relative energy of the bonding orbitals determines the magnitude of the bond energy of the molecule, ΔE.The more electronegative element AOs  have lower energy than the other atom. 

As seen in the figure below, a more electronegative atom has lower energy AO .The bonding MO formed, is just a little more lower in energy than this AO and so the net decrease in energy( ΔE) is less.


The MOs of heteronuclear atoms are not symmetrical as in the case of homonuclear molecules.The MOs are polarized .The BMOs are shifted  towards the more electronegative atom i.e the more electronegative atom attracts the bonding MOs more towards itself.Thus,

the more electronegative atom(nucleophile) contributes more to the BMO→ Mathematically, MOs have larger coefficient(c) on more electronegative atom. 

the less electronegative element(electrophile) contributes more to the ABMO → MOs have smaller coefficients on less electronegative atom. 

Thus, in the above figure , the BMO is drawn closer to  the more electronegative element and the ABMO is drawn closer to the less electronegative element.

The bond between homonuclear atoms will always be stronger than the heteronuclear ones owing to perfectly matched AOs.The strength depends on the relative energies of the atomic orbitals participating in bond formation. AOs which are closer in energy will overlap and those with huge energy difference won’t.

Type of molecule


Bond strength


ΔE = 0

Very strong bond

(as AOs are perfectly matched).

Heteronuclear with less energy difference between AOs

ΔE is less

Strong bond

Heteronuclear with more energy difference between AOs

ΔE is more but ΔE < 12eV

Weak bond

Heteronuclear with energy difference more than 12 eV

ΔE > 12eV

No interaction between orbitals

The following figure shows us relative energies for AOs of different atoms.The data is collected using a technique called photoelectron spectroscopy (We will study this technique in detail in a later post).


In the above figure, the potential energy(in eV) of an valence orbital is plotted on the Y axis and the atomic number is plotted on the X axis.
( We consider kinetic energy for moving objects and potential energy for stationary things.So, we talk about kinetic energy of an electron and potential energy of an orbital).

We start with 1s orbital for H and He , as they are the valence orbitals for these atoms.Each curve represents a different  atomic orbital.As seen in the above figure, the orbital energy of H atom is – 13.6 eV. This means that it takes -13.6 eV energy to remove the valence (1s) electron out of that orbital.

If the energy difference between two orbitals is greater than 12 eV , the AOs will not interact.Thus, looking at the figure above, we can predict which orbitals can interact with each other and the ones that cannot. As highlighted in the figure, the 1s orbital of H CANNOT react with the 2s of fluorine atom as –

Energy of 1s of H atom = -13.6ev &
Energy of 2s of F atom = – 40.2 eV.
The energy difference is huge(26.6 eV).

Energy of 2p of F atom = -18.6 eV.
Energy difference = 18.6 -13.6 = 5eV
∴1s of H can interact with 2p of F atom.
(NOTE – the negative sign is not considered while calculating the energy difference ).

Based on this information ,we shall begin drawing MO diagrams for heteronuclear species in the next few posts. Till then ,

Be a perpetual student of life and keep learning…

Good day !

Image source –


References and Further Reading –


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