Up to now we have only considered combining two atoms of the same element, which makes things simpler because the same orbitals on each of the two atoms have the same energy. But when the two atoms are different two things change. The fi rst is obvious—the number of electrons contributed by each atom is different. This is easy to accommodate since it just affects the total number of electrons, we need to put into the MO diagram when we fill up the energy levels. So, for example, if you were constructing an MO diagram for NO, the gas nitric oxide (NO, a rather remarkable biological messenger in the human body) rather than N2, you simply put in a total of 15 rather than 14 electrons because O contributes eight electrons and N seven.

Nitric oxide, NO Nitric oxide was for a long time known only as one of the villains of urban air pollution, being formed during the combustion of petroleum and other fossil fuels. In the last 20 years, however, it has become evident that it is much more than that—one unexpected role, which earned its discoverers the Nobel Prize in physiology in 1998, is as a biological messenger, managing the contraction of smooth vessels and hence regulating blood flow.

The second thing that changes when you have two different atoms bonded together is the relative energies of the AOs being combined. It may seem natural to assume that a 2p orbital has the same energy whatever atom it finds itself in, but of course the difference is that an electron in a 2p (or any other) orbital feels an attraction to the nucleus which depends on the nuclear charge. The greater the number of protons in the nucleus, the greater the attraction, and hence the more tightly held, more stable, and lower in energy the electron becomes. This is the origin of electronegativity. The more electronegative an atom is, the more it attracts electrons, the lower in energy are its AOs, and so any electrons in them are held more tightly.

As you move across each row of the periodic table, therefore, electronegativity increases as the energy of each orbital drops. From Li (electronegativity 0.98) across to C (2.55), and on to N (3.04), O (3.44), and F (3.98), the elements steadily become more electronegative and the AOs lower in energy. So our diagram of the orbitals of NO actually looks like this.

We have shown only the 2s and 2p orbitals as the 1s orbitals are much lower in energy, and as you saw in the diagram of N2 on p. 94 their bonding and antibonding interactions cancel each other out. The orbitals on O are lower in energy than the orbitals on N, but they still interact just fi ne. However, there is one interesting consequence: if you look at each bonding orbital, you will see that it is closer in energy to the contributing orbital on O than the contributing orbital on N. Likewise, each antibonding orbital is closer in energy to the contributing orbital on N than the contributing orbital on O. The result is that the MOs are unsymmetrical, and while all the bonding orbitals have a greater contribution from the oxygen AOs, all the antibonding orbitals have a greater contribution from the nitrogen AOs. Overall the diagram shows eight electrons in bonding orbitals and three electrons in antibonding orbitals, so the overall electron distribution is skewed (polarized) towards O, just as you would expect from a comparison of the electronegativities of N and O. The eight electrons in bonding orbitals and three electrons in antibonding orbitals means that NO has a bond order of 2½. It also has an unpaired electron—it is a radical. We can’t easily represent half a bond in valence bond terms, so we usually draw NO with a double bond, representing four bonding electrons. The remaining seven electrons can be shown as three lone pairs and one unpaired electron. Where do we put them? Well, our MO diagram tells us that the unpaired electron occupies an orbital closer in energy to N than to O, so we put that on N. N and O differ only slightly in electronegativity (electronegativity of N 3.04; O 3.44): their orbitals are quite close in energy and form stable covalent bonds. But we also need to consider what happens when two atoms forming a bond differ hugely in electronegativity. We can take sodium (electronegativity 0.93) and chlorine (electronegativity 3.16) as our example. We know from observation that the product of reacting these two elements (don’t try this at home) is the ionic solid Na+Cl−, and the MO energy level diagram tells us why. The AOs we need to consider are the 3s and 3p orbitals of Na (all its lower energy 1s, 2s, and 2p orbitals are filled, so we can ignore those, as we did with the 1s orbitals of N2 and NO above) and the 3s and 3p orbitals of Cl (again, the 1s, 2s, and 2p orbitals are all filled). Here is the diagram, with the Na orbitals much higher in energy than the Cl orbitals.

But these AOs are too far apart in energy to combine to form new MOs and no covalent bond is formed. The orbitals which get filled are simply the 3s and 3p orbitals of the Cl atom. The electrons available to fill these orbitals are the seven provided by Cl plus the one from Na: we end up with Na+ and Cl−. The ionic bonding in NaCl is due simply to the attraction between two oppositely charged ions—there is no orbital overlap. These three different cases where the two combining orbitals differ greatly in energy, only a little, or not at all are summarized below.