Mo Diagram For O2

Building the MO Diagram for O2 From Scratch

Most students get tripped up on the O2 molecular orbital diagram because they blindly apply the nitrogen ordering instead of recognizing when the s-p mixing changes. Here is how it actually works in practice and what to watch out for. Oxygen is one of those cases where the 2s and 2pz orbitals are far enough apart in energy that they don't mix the way they do in B2, C2, or N2. This shifts the sigma bonding orbital above the pi bonding orbitals, and if you draw the diagram using the N2 ordering, your bond order comes out wrong and you miss the paramagnetism entirely. That is not a small detail. It is the entire point of the exercise. The correct energy ordering for O2 is:

2s *2s 2pz 2px = 2py *2px = *2py *2pz Note that 2pz drops below the 2p orbitals here. For N2 and lighter diatomics, that order flips. The crossover happens right around O2 and F2. That is why every textbook that just gives you one generic diagram without explaining the exception will get you in trouble on an exam.

Step-by-Step Construction

Each oxygen atom contributes six valence electrons, so the O2 molecule has twelve to place. Fill from the bottom up, following Hund's rule for the degenerate * orbitals. 2s gets two electrons. *2s gets two. 2pz gets two. The 2px and 2py pair together and hold four electrons total. That accounts for ten electrons. You have two left, and they go into the * antibonding orbitals—one in *2px and one in *2py, unpaired. This is why O2 is paramagnetic, which is something you can verify experimentally with a magnet, not just something to memorize for a test. The bond order calculation is straightforward: bonding electrons minus antibonding electrons, divided by two. That gives (8 - 4) / 2 = 2. A double bond. Consistent with the Lewis structure, but the MO picture tells you more than the Lewis structure ever could, specifically the unpaired electrons and the magnetic behavior.

The Mistake That Costs People Points on Exams

I ran into this constantly when I was TAing upper-level chemistry. Students would draw the orbitals below 2pz correctly, then fill the electrons, but they would pair both remaining electrons into a single * orbital instead of distributing them across the degenerate pair. That violates Hund's rule and changes the entire physical interpretation. Once I made them calculate the bond order from that wrong filling, they got 3 instead of 2, which immediately flagged that something was off. It is a useful self-check if you catch yourself before submitting. The standard MO diagram works fine for homonuclear diatomics like O2, N2, and F2. It gets murky fast for heteronuclear species where the atomic orbitals sit at different energies and you need to weight contributions asymmetrically. It also does not account for electron correlation effects, which matter for excited states and certain transition metal complexes. If you need quantitative accuracy beyond bond order and qualitative magnetism, you move into computational chemistry territory—DFT or post-Hartree-Fock methods—and the simple diagram becomes a teaching tool, not a predictive one. For the vast majority of undergraduate and graduate coursework, the diagram as described above is sufficient. Just remember the energy ordering flip at O2 and fill according to Hund's rule. Everything else follows from there.