The sigma bonds I've seen in the general chem class showed overlapping s (or hybrid) orbitals where electron density was mostly located between both
atoms. So H2 looks egg-shaped.
In this year's organic chem book, it shows a sigma bond between two hydrogen atoms with antibonding 1s orbitals. In that diagram, there's a nodal
plane equidistant from both atoms and the orbitals are kinda pushed away from it.
So where did that antibonding 1s orbital come from? When we look at p orbitals, there's a positive and negative half of each orbital. My
understanding is that the wavefunction outputs positive values for one node and negative for the other. I think I remember hearing that relates to
electron spin. So how does that work for the 1s orbital? If there were a negative sphere overlapping the positive sphere, wouldn't the electron
densities cancel?blogfast25 - 15-8-2015 at 11:48
Like all orbitals (wavefunctions, ψ) 1s orbitals can take on a + or a - form (+ψ or -ψ are both solutions of the Schrödinger equation). When the
signs of two 1s orbitals are the same, they are said to be 'in phase', when the signs are opposite they are said to be 'out of phase'.
When two s orbitals that are in phase meet they reinforce and form a σ bonding molecular orbital.
But when two s orbitals that are out phase meet they 'annihilate' and form a σ<sup>*</sup> anti-bonding molecular
orbital.
H<sub>2</sub> isn't really pear shaped, though. The electron density is symmetrical along the bonding axis.
[Edited on 15-8-2015 by blogfast25]kt5000 - 15-8-2015 at 16:51
That helps, thanks. I have a feeling they're not going to expect us to know 'why' those orbitals form the way they do, but I didn't expect to see
that and I have to understand why
Do the actual wavefunctions get covered in depth in undergrad Chemistry, or does that wait for graduate courses?blogfast25 - 15-8-2015 at 18:18
