4.3: Heteronuclear Diatomic Molecules

Diatomic molecules with two non-identical atoms are called heteronuclear diatomic molecules. When atoms are not identical, the molecule forms by combining atomic orbitals of unequal energies. The result is a polar bond in which atomic orbitals contribute unevenly to each molecular orbital. 

The application of molecular orbital theory to heteronuclear diatomic molecules is similar to the case of homonuclear diatomics, except that the atomic orbitals from each atom have different energies and contribute unequally to molecular orbitals. Recall that atomic orbitals must have compatible symmetry and similar energy to combine into molecular orbitals. In the case where atomic orbitals of like symmetry have different energies, they combine less favorably than orbitals that are closer to one another in energy. As a general rule, orbitals that have energy differences of greater than 10-14 eV do not combine favorably. In the molecular orbital diagram, the closer a molecular orbital is to an atomic orbital, the more that atomic orbital contributes to the molecular orbital. This last point is helpful for back-of-the napkin estimations of what the molecular orbitals "look" like.

In this section, you should learn how to generate molecular orbital diagrams of heteronuclear diatomic molecules. To approach such a problem, we must start with a knowledge of the relative energies of electrons in different atomic orbitals. In other words, we need knowledge of the orbital potential energies (or orbital ionization energies).

Orbital Ionization Energies

There are two approaches you can use to "know" or estimate the atomic orbital energy levels.

1. Use a table of atomic orbital ionization energies, like those found in Table \(\PageIndex{1}\).

2. When you do not have access to a table of values like the one below, use periodic trends in electronegativity and/or ionization energies as your guide to approximate relative values for different atoms.

Table \(\PageIndex{1}\):These are one-electron ionization energies of the valence orbitals calculated by average energies of both the ground-state and ionized-state configurations. (From Harry Gray, “Electrons and Chemical Bonding,” Benjamin, 1964, Appendix)
\[ \begin{array} {|c|ccccccc|c|ccc|}\hline Atom & 1s & 2s & 2p & 3s & 3p & 4s & 4p & Atom & 3d & 4s & 4p \\   \hline H & -13.64 &  &  &  &  &  & &Sc & -4.71 & -5.70 & -3.22 \\ He & -24.55 &  &  &  &  &  & & Ti & -5.58 & -6.08 & -3.35 \\  Li &  & -5.46 &  &  &  &  &  &  V & -6.32 & -6.32 & -3.47\\  Be &  & -9.30 &  &  &  &  & & Cr & -7.19 & -6.57 & -3.47 \\  B &  & -14.01 & -8.31 &  &  &  & &  Mn & -7.93 & -6.82 & -3.60  \\  C &  & -19.47 & -10.66 &  &  &  &  & Fe & -8.68 & -7.07 & -3.72\\  N &  & -25.54 & -13.14 &  &  &  &  & Ni & -10.04 & -7.56 & -3.84\\  O &  & -32.36 & -15.87 &  &  &  & & Cu & -10.66 & -7.69 & -3.97  \\ F &  & -46.37 & -19.72 &  &  &  &  \\  Ne &  & -48.48 & -21.57 &  &  &  &  \\  Na &  &  &  & -5.21 &  &  &  \\  Mg &  &  &  & -7.69 &  &  &  \\  Al &  &  &  & -11.28 & -5.95 &  &  \\  Si &  &  &  & -15.00 & -7.81 &  &  \\ P &  &  &  & -18.72 & -10.17 &  &  \\  S &  &  &  & -20.71 & -11.65 &  &  \\  Cl &  &  &  & -25.29 & -13.76 &  &  \\  Ar &  &  &  & -29.26 & -15.87 &  &  \\  K &  &  &  &  &  & -4.34 &  \\  Ca &  &  &  &  &  & -6.08 &  \\  Zn &  &  &  &  &  & -9.42 &  \\  Ga &  &  &  &  &  & -12.65 & -5.95 \\  Ge &  &  &  &  &  & -15.62 & -7.56 \\  As &  &  &  &  &  & -17.61 & -9.05 \\  Se &  &  &  &  &  & -20.83 & -10.79 \\  Br &  &  &  &  &  & -24.05 & -12.52 \\  Kr &  &  &  &  &  & -27.52 & -14.26 \\ \hline  \end{array} \nonumber \]

 

• 4.3.1: Polar bonds

• 4.3.2: Ionic Compounds and Molecular Orbitals

 

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Sources:

• Gray, Harry.  Electrons and Chemical Bonding, Benjamin, 1964.
• Miessler, Gary L, and Donald A. Tarr. Inorganic Chemistry. Upper Saddle River, N.J: Pearson Education, 2014. Print.

Curated or created by Kathryn Haas