OK, so that title was just unfair click-bait. It is an acronym for Anti-Bonding Is More Anti-Bonding Than Bonding Is Bonding. It refers to the way we draw molecular orbital diagrams. We usually draw them with the bonding orbital going down below the atomic orbitals of which it is comprised by, say, 10 cm on the whiteboard, and we make the antibonding orbital go *up* by 10 cm. 

The problem is, this is wrong. The splitting of orbitals is asymmetric. The antibonding orbital should go *up* more than the bonding orbital goes *down*. OK, so that's cool, but it's a small correction, right?

Well, it depends. Specifically, it depends on the overlap. Without going too far into the weeds, the amount of overlap (S) can range from 0-1, with 0 meaning no overlap (orthogonal orbitals, no bonding *or* antibonding) to 1 (the orbitals are 100% on top of each other, and you just fused the nuclei). 

So how asymmetric are these bonding and antibonding orbitals? How much should you change your whiteboard orbitals? Here's a quick chart. Assuming your bonding orbital goes down by 10 cm in all drawings, the antibonding orbital should go up by:

S: 0.01, AB up by 10.2 cm (this is a super-wimpy bond....think "elemental iodine"

S: 0.10, AB up by 12.2 cm (this is, to the best of my recollection, typical for a C-C single bond)

S: 0.20, AB up by 15 cm

S: 0.40, AB up by 23 cm (this is, to the best of my recollection, the overlap in H2, the first MO diagram our students ever see). 

So why should we care? Well, first of all, it's a True Fact About the Universe. And i kinda think we should teach those. But also, it means that when you stick antibonding electrons in orbitals with great overlap, that these electrons really weaken the bond! I mean, even more than you'd expect by doing Bond Order = bonding pairs-antibonding pairs. So where does this happen? With the p* electrons in the main group to the right hand side of the p-block. N-N, O-O, and F-F bonds are crazy weak because you've got antibonding electrons in  orbitals with great overlap. We know that hydrazine, peroxide, and fluorine are scary, and this is why!

But the ultimate reason is dioxygen. This is an O=O double bond as weak as an O-H single bond. And if we write the combustion of glucose, if you use the right isomer of glucose, you find that you are making and breaking the same number of bonds. So the incredible exothermicity of sugar metabolism is purely due to the fact that you are making stronger bonds, and breaking weaker ones. 

And twelve of those weaker ones are these wimpy ABIMABTBIB-hobbled bonds between oxygen. How much does this matter? If we compare the BDE of O=O and a C=C bond (an analog where we don't have to worry about the fact that ABIMABTBIB), we can see that this effect is worth about 25% of the combustion energy of glucose. Which is about the energy you need to run your brain.

So you can think about the fact that ABIMABTBIB as the feature of chemistry that pays for our beautiful minds, and our ability to study chemistry. What students and professors can't appreciate that?