The electronic AP Chemistry book (above) says that O2 has one bond, as evidenced by its being paramagnetic.
I’ve not heard that before, and “Professor Dave” on youtube does not come to that conclusion: Valence Bond Theory, Hybrid Orbitals, and Molecular Orbital Theory - YouTube
Does anyone have a way out of this contradiction? First year AP teacher here, and I never understood this much myself as a student.
So when they get a bond order of 2 themselves in the displayed math, how do they get to say that MO theory shows that O2 is paramagnetic? What math or reasoning goes along with the experimental data? The math they just showed doesn’t.
It looks to me that the electronic AP Chemistry book is saying that the Lewis dot diagram for paramagnetic O[sub]2[/sub] (accounting for MO theory) has a single bond - not that MO theory itself predicts a single bond.
Lewis dot diagrams are are a useful tool, but they do not predict certain things very well and paramagnetism is often one of them. This is something molecular orbital theory does a much better job of. In both cases they are predicting a bond order of 2, but only in MO theory is the paramagnetism correctly predicted. My hope is the AP chemistry book is trying to make this point and pointing out an alternative way to represent the Lewis dot structure so that it correctly predicts paramagnetism. However, as you point out bond order is no longer correctly predicted. Thus we have an important short coming of valence bond theory. Molecular orbital theory in general can be more complicated and harder to wrap your head around, but it also typically gives a clearer picture of the exact bonding in a given molecule.
edit: In regards to your second post, paramagnetism is based on unpaired electrons and MO theory shows their are 2 of them. Valence theory would try to pair these electrons up, making the molecule diamagnetic.
O2 does have a bond and two unpaired electrons, in its lowest energy form. It’s what’s called a triplet. It is exceptional in this regard; in fact, one of its most interesting properties is that, because it’s a triplet and most other molecules are singlets, it’s less reactive than if it was a singlet. It needs to switch or swap to a higher-energy singlet state before reacting.
The all-paired form, the singlet predicted by Lewis’ dots, exists - but in this particular weirdo’s case it isn’t the lowest-energy form.
Interesting substances tend to have unusual behaviors: water does (the extent of its hydrogen bonds), so does oxygen.
That makes no sense.
The paramagnetism and ferromagnetism both arise out of UNPAIRED ELECTRONS.
The outer shell unpaired electrons are paramagnetic as they are unconstrained.
In ferromagnetism, there’s a constraint on the unpaired electron , as its inside…, and it can behave ferromagnetic.
“one bond” isn’t making any sense as a a necessary and sufficient indicator of unpaired electrons… there’s all sorts of electron shell arrangements in molecules that make for paramagnetism.
Here’s a discussion of how to predict molecular stability and magnetic property from orbitals (and where lewis diagrams go wrong.)
http://www.mpcfaculty.net/mark_bishop/molecular_orbital_theory.htm
Well, their overall point is not the single bond, but that the unpaired electrons mean in this case that there can’t be a second bond.
The textbook does the math for bond order and they get 2 for oxygen. This is using 2 unpaired electrons in the pi* 2p MO level, admittedly. If the bond order of 2 does not mean there are 2 bonds, then what does it show? I admit that I must have been incorrect that it means there are definitely 2 bonds, because bond orders like 1.5 are possible.
See if this helps. I don’t know much about MO theory. But I thought I’d at least reply.
And this
A bond order of 1.5 does in fact mean that there are 1.5 bonds. Or, if you prefer, an equal superposition of 1 bond and 2 bond states. It’s a quantum thing.
Oh, yes, the bond of 1.5 is the average over the the different resonance structures. Like if you have a central atom with a single bond on one side, a lone pair, and a double bond on the other side. O3 does this (at least with a simple Lewis structure). But I’m pretty sure you know this. I assume you’re asking how they reconcile the math of getting a bond order of 2 with MO theory but the MO theory is saying it’s a single bond.
Hmm… could the bond strength or bond length indicate a stronger than single bond?
I think you are misinterpreting what they are trying to say. O2 is a classic case where Lewis dot structure does not correctly predict properties and I think all they are trying to demonstrate is this fact. It is a bond order of 2 (as demonstrated in MO theory) and it also has unpaired electrons. These two things can only be observed when you do the full MO treatment, it can’t be easily elucidated from the Lewis Dot structures alone.
The correct bond order is two, representing a double bond. The single bond lewis dot structure is not a valid Lewis dot structure, it violates the octet rule. Again, you can’t use Lewis Dot structures to accurately predict O2 properties. That is the point, nothing else.
Maybe this was already what you were looking at, but this is a description I could generally follow.
http://chemed.chem.purdue.edu/genchem/topicreview/bp/ch8/mo.html
Bond order now is = (bonding electrons - antibonding electrons) / 2 = (8-4)/2 =2. “Although the Lewis structure and molecular orbital models of oxygen yield the same bond order, there is an important difference between these models. The electrons in the Lewis structure are all paired, but there are two unpaired electrons in the molecular orbital description of the molecule. As a result, we can test the predictions of these theories by studying the effect of a magnetic field on oxygen.”
Then use this for the answer for O2, though it seems like people are using slightly different notation for naming their orbitals.
Or maybe Molecular Orbital diagram for the molecule, oxygen, O2. - YouTube for a longer description (and there are others on YouTube).