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Electron configuration 3d orbitals

  1. Aug 4, 2012 #1
    Hello.

    I'm having a problem understanding the 3d orbitals when I'm doing electron configurations.
    When you fill up electrons for let's say nickel, Ni, you get the configuration:
    [Ar]3d8 4s2,
    if I'm not mistaken.

    But what I have read, at least I think I have, the 3d orbital gets really stable whenever it is possible to half fill it, or fully fill it (5 or 10 electrons).
    So why is it that the configuration for nickel isn't:
    [Ar]3d10 ?

    I'm currently not sure about when it is favourable to fill up the 3s orbital instead of making a 5 electron or 10 electron d-orbital.
    I mean, if I got a, lets say Ni-1 ion, I would be able to make a 3d orbital with 10 electrons also, just as I would imagine it without the -1, and then, in my mind, get a configuration that is:
    [Ar]3d10 4s1

    But is that more or less correct than:
    [Ar]3d9 4s2


    So, I'm kinda confused actually to when you fill up 3d-orbitals with 5 or 10 electrons instead of filling up the 4s-orbital first.

    Anyone who could maybe enlighten me ? :)


    Regards
     
  2. jcsd
  3. Aug 5, 2012 #2
    As long as i know,

    Ni -- > [Ar]3d8 4s2
    Ni- -- > [Ar]3d10 4s1

    when you have Ni, you cannot fit the electrons for a full or half filled structure. but you have scope for 3d10 4s1 in case of nickle with one extra electron.

    higher secondary school level chemistry books contain much information about such issues.
     
  4. Aug 5, 2012 #3
    It is soothing to remember that the further you go down the periodic table, the more exceptions to rules one will find, and you should not be perturbed by that too greatly.

    If you read the Wiki article on "en.wikipedia.org/wiki/Aufbau_principle" [Broken], you will see that, naively, one will fill the 4s subshell before the 3d subshell. You will also note that there are exceptions to this rule. If you go on to examine the article on "en.wikipedia.org/wiki/Nickel" [Broken], you will note a section on its electron configuration and some debate regarding its nature.

    In short, unless you know it's an exception of some sort, you can usually feel safe filling the s orbital of level i before the d orbitals of i -1.
     
    Last edited by a moderator: May 6, 2017
  5. Aug 6, 2012 #4
    Great, thank you...

    That was pretty much what I was hoping to hear :)

    But one last thing. So it is incorrect to compare the configuration of Co+2 to the electron configuration of Manganese ?
    They do have the same amount of electrons, but if I'm not mistaken Manganese has the electron configuration:
    [Ar] 4s2 3d5,

    where Co+2 apparently got (According to a textbook I have):
    [Ar] 3d7

    Why isn't it the same then ?


    Thanks in advance :)
     
    Last edited by a moderator: May 6, 2017
  6. Aug 6, 2012 #5
    The previously linked-to Wiki article on the Aufbau principle notes the following:

    Bolded for emphasis. Co2+ is a cation, after all, and not a neutral atom in its ground state.
     
  7. Aug 7, 2012 #6
    What I do know is that the shells do not necessarily be full ( duplet or octet ); they ( the orbiting electrons ) can move to the next shell by quantum jump.

    I don't know much details about it though.
     
    Last edited: Aug 7, 2012
  8. Aug 8, 2012 #7
    Does anyone else wonder if this stuff (aufbau etc) is more trouble than it's worth? It gets taught as some ironclad fact and then people get confused by many complicated exceptions (transition metals, ions of those transition metals etc). When this stuff is "used" in practice isn't it usually a quantum chemistry calculation with a check to see that it matches the experimental spectra and them move on? My point is, does anyone IN PRACTICE use aufbau and think deeply about its exceptions and if not, why foist it on the young people?
     
  9. Aug 9, 2012 #8
    It's the basis of qualitative predictions in chemistry! It's used in organic chemistry everyday (and for those elements there no exceptions). It's also used more as a starting point for quantum chemistry calculations more than a final goal.
     
  10. Aug 9, 2012 #9
    It makes sense to use in cases like O-chem where it works, and it being taught as a special case that applies to that portion of the table won't cause any confusion. But it's so all-over-the-place for the transition metals and their ions and there is so much effort put into "it's d before s and then s before d for the next one but then it's all d and no s and then..." that it spends more time than it saves. As for it being a "starting point for calculations, that's certainly not important enough of a consideration to confuse gen chem students if it's even really true. If aufbau had never occurred to to anyone and we were sitting down to do a QC calculation without any idea of confusing "rules" for which orbital is lower than which other, we'd be just fine. The variational principal really has our back on this one. We'd wait a short bit of time and we'd have the right answer and we'd get on with out lives.
     
  11. Aug 9, 2012 #10
    I taught (as a TA) gen chem. to first year engineering and physics students and we did not insist on exceptions (which are few) too much. In fact most of the course focused on s and p elements. I understand your frustration though. The thing is these approximations are very limited beyond s and p elements; beyond that you're just memorizing properties of certain elements and not really rationalizing anything (although I've met people who disagree with me). I have to admit I never quite understood all the BS rationalization of properties of transition metals (although I shun away from trying to understand them too hard).
     
  12. Aug 11, 2012 #11
    These are the ground states. They do not show a way of transition among themselves.
    Similarly for a positive nickel ion:
    Ni+ -- > [Ar]3d9
     
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