Chemical Forums
Chemistry Forums for Students => Undergraduate General Chemistry Forum => Topic started by: zeoblade 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?
-
http://en.wikipedia.org/wiki/Degenerate_energy_level
-
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?
-
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.