I would add that they are even more unique when represented in an abstract space called momentum space, where the geometric differences translate into actual distances between the electron pairs. So the answer to your question in general terms is that electron pairs settle into molecule-sized orbitals (molecular orbitals) that are physically spread out over the entire ring, but which nonetheless have unique geometric features that keep them distinguished in real space. What that means is that your initial intuition is essentially correct if all the electron pairs shared exactly the same molecular orbital. However, I will nonetheless add a simpler approach to answering your question, which is this: Pauli exclusion holds as long as there is something in space or in spin (spin being the pairs of electrons) that firmly distinguishes two "stable" states. I very much like the first answer, especially the superb visuals! (Note that Pauli exclusion must also hold in atoms where this latter statement is true!) Nor is there any need to introduce new fictitious quantum numbers. ![]() This is not because of Pauli exclusion, but because the Hamiltonian, $L_z$ and $L^2$ no longer form a complete set of commuting observables. Therefore, molecular orbitals are not indexable by the same set of quantum numbers as can be done for atomic orbitals. Angular momentum operators $L_z$ and $L^2$ commute with atomic Hamiltonians because of rotational and Runge-Lenz invariance however, they do not commute with molecular Hamiltonians because molecules have neither of these symmetries. Your discussion of orbital and spin angular momenta is completely irrelevant. ![]() While on average the electrons look like they are all in the same region of space, if you measure the positions at any time, no two electrons will ever occupy exactly the same point in space. There is no conflict with Pauli exclusion because you can see (and calculate!) that the orbitals are orthogonal by virtue of having an appropriate number of nodes. However, their corresponding probability densities can be superposed onto similar regions of real space. The six pi electrons occupy three different molecular orbitals of very different shapes, as can be seen here, for example, and in Richard Terrett's answer. What tends to be confusing in these pictures is it sometimes unclear that the pictures are sometimes of the molecular orbitals and sometimes of the electron density, which is an observable quantity and quite another thing altogether. ![]() The short answer to your question is that no, they don't all occupy the same orbital but yes, they occupy very similar regions in space. and Schulten, K., 'VMD - Visual Molecular Dynamics', J. ORCA – an ab initio, Density Functional and Semiempirical program package, Version 2.6. It should be easy to extrapolate to the highest lying MO of this type - antibonding all the way around. These all had the "original" (as in, if resonance didn't occur) quantum numbers $n=2,l=1,m=\text$ orbitals, so this orbital has antibonding character for two of the carbon-carbon pairs. This seems to disobey Pauli's exclusion principle. ![]() I've always been a bit uncomfortable with the concept of more than two electrons in a single orbital-like region(probability-wise) which occurs in resonance.
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