2.39

Moderators: Chem_Mod, Chem_Admin

104754311
Posts: 13
Joined: Fri Sep 29, 2017 7:06 am

2.39

Postby 104754311 » Tue Oct 24, 2017 1:46 am

Why is Carbon's electron configuration (as shown in the diagram) an excited state? I thought an excited state only happens when you go from a an s orbital to a p orbital for example.

Jingyi Li 2C
Posts: 56
Joined: Fri Sep 29, 2017 7:06 am

Re: 2.39

Postby Jingyi Li 2C » Tue Oct 24, 2017 10:24 am

An excited state can happen when you go from a an s-orbital to a p-orbital. But its definition is that an atom with electrons in energy states higher than predicted by the building-up principle is said to be in an excited state. In this case, carbon should have two electrons in different p-orbital with parallel spins. However, in the diagram, these two electrons have paired spin in the same orbital. Therefore, this arrangement has slightly higher energy than predicted by the building-up principle. So, the excited state definition can also be applied.

David Zhou 1L
Posts: 61
Joined: Fri Sep 29, 2017 7:04 am

Re: 2.39

Postby David Zhou 1L » Fri Nov 03, 2017 6:52 pm

Due to Hund's Rule, the electron configuration isn't the ground state. One of the electrons in the single filled 2p orbital should move to an empty p-orbital, with spin up or down, whichever the one remaining in the first orbital is oriented in; they should be parallel.

Excited doesn't have to mean jumping to different energy levels, there are many ways that electrons can be configured such that they aren't at the lowest energy level.

Jasmin Tran 1J
Posts: 54
Joined: Thu Jul 27, 2017 3:00 am

Re: 2.39

Postby Jasmin Tran 1J » Mon Nov 06, 2017 10:27 pm

For further explanation for the whole question, (a, Carbon) shows 2 electrons in one p orbital when there are empty p orbitals, so it is in an excited state because ground state electrons want to be in separate orbitals. For (b, Nitrogen), one electron has opposite spin compared to the other two, but there are only three electrons in the p orbital so they all should have parallel spin with one electron in each orbital. For (c, Beryllium), there is one electron in each of the 2s and 2p orbitals, but in ground state, they both should be in the 2s orbital because the s orbital is supposed to be filled before the 2p. For (d, Oxygen), this is the only ground state configuration because there are 2 electrons each in the 1s and 2s orbitals and 4 electrons in the 2p orbitals, three with parallel spin and one with opposite spin, which is normal.

Hope this helps to clarify a bit more!

Joanna Pham - 2D
Posts: 113
Joined: Fri Apr 06, 2018 11:04 am

Re: 2.39

Postby Joanna Pham - 2D » Thu May 03, 2018 10:41 am

Jasmin Tran 1J wrote:For further explanation for the whole question, (a, Carbon) shows 2 electrons in one p orbital when there are empty p orbitals, so it is in an excited state because ground state electrons want to be in separate orbitals. For (b, Nitrogen), one electron has opposite spin compared to the other two, but there are only three electrons in the p orbital so they all should have parallel spin with one electron in each orbital. For (c, Beryllium), there is one electron in each of the 2s and 2p orbitals, but in ground state, they both should be in the 2s orbital because the s orbital is supposed to be filled before the 2p. For (d, Oxygen), this is the only ground state configuration because there are 2 electrons each in the 1s and 2s orbitals and 4 electrons in the 2p orbitals, three with parallel spin and one with opposite spin, which is normal.

Hope this helps to clarify a bit more!



Can you please explain part b? Why is it that the electrons must have parallel spins?


Return to “Quantum Numbers and The H-Atom”

Who is online

Users browsing this forum: No registered users and 1 guest