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Topic: Degenerate orbitals  (Read 22397 times)

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Offline zeoblade

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Degenerate orbitals
« on: March 01, 2010, 04:56:36 AM »
Why and how are orbitals said to be degenerate? I'm reading Housecroft's Inorganic Chemistry and I don't know how I've missed it but I have a feeling this is of some importance later on.

My understanding is that the electron's probability of being in a position degenerates with increasing distance from the nucleus, is this what degenerate orbitals are?

Then comes orbitals that are not degenerate such He, now I am a little confused. Would someone be kind to lead me with their rationale to an understanding of degenerate and non-degenerate orbitals?

Online Borek

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Re: Degenerate orbitals
« Reply #1 on: March 01, 2010, 06:22:33 AM »
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Offline zeoblade

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Re: Degenerate orbitals
« Reply #2 on: March 01, 2010, 05:50:50 PM »
Thanks Borek, now my understanding is electrons in the same n quantum number are degenerate.

So does this mean 2s2 electrons are degenerate because there are 2 electrons with opposing spin?

So that 2p4 have 2px as degenerate because 2 electrons fill the x-axis but 2py and 2pz are only half filled and therefore not degenerate?

Or does it mean when there is more than one electron in the n=2 quantum, the whole shell is degenerate except for 2s1?

Offline tamim83

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Re: Degenerate orbitals
« Reply #3 on: March 02, 2010, 08:36:57 AM »
Actually, all electrons in the H atom that have the same n are degenerate.  So in hydrogen 2s and 2p electrons have the same energy.  For other atoms all electrons with the same n and l are degenerate.  So in oxygen for example the 2s electrons are lower in energy than the 2p electrons.  This is due to screening; the electrons shield or screen one another from the full charge of the nucleus.  s electrons spend more time near the nucleus on average than p electrons.  So not only do s electrons experience less shielding but they do the most shielding as well. 

Hope that helps some. 

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