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Homework Help: Electron Configuration Calculations

  1. Jun 17, 2008 #1

    I've been trying to learn about electron configuration in atoms, and how it is calculated.

    I pretty much figured out a lot of it from doing my own research. From what I understand, each group in the periodic table represents a different electron shell, and the blocks S, P, d, and F represent subshells. Once it got up to group 4, I figured it needed to fill 3D before working on 4P, and I was correct. I thought the same principal would apply once I got to the 6th group, so I expected Lanthanum to be [Xe]6s^2 4f^1. But, apparently, an electron is put in 5D first so that it is [Xe]6S^2 5d^1, and then Cerium is [Xe]6s^2 4f^1 5d^1. This makes no sense to me.

    I admit, I am very new to learning about this. I am a physics and math major, and my first chemistry course starts next semester, and I was jus ttrying to get a bit of a head start. But I'd like to understand why it progresses in this fashion, when it's so unlike anything else before it.

    it seems to repeat in Actinium, so at least it is persistent. But I want to know, why.

    Thanks for any help. :)
  2. jcsd
  3. Jun 17, 2008 #2


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    Staff: Mentor

    Start with Aufbau principle - that gives a general idea. But then there are exclusions. In general electron enters the orbital that will give the lowest energy state. But electrons interact, and the more electrons are present, the more closely spaced energy levels of different states are, thus sometimes electrons fill orbitals in slightly different order that what we expect. At such situations we usually say something like "in this particular case configuration with 5 unpaired electrons on d and 1 on s has less energy then the configuration with 4 electrons on d and fully filled s" (that's for chromium). And that's true, it just doesn't have any predictable power, as it works only for chromium.

    In the end the only real answer to question "why" is "because that's the way it is" :wink:
    Last edited by a moderator: Aug 13, 2013
  4. Jun 17, 2008 #3


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    Actually, you missed an even "earlier" pattern-breaking. The 4s gets filled first ahead of the 3d. Look at the transition metals.

    When you add more and more electrons to the shell, you also start to increase the possibly of shielding of the nuclear potential. Since the p, d, and higher orbitals tend to extend even further from the nucleus than the s orbital, on average, they get shielded by the inner shell electrons even more. This can cause them to have a higher energy state than the s-orbital of the next principle quantum number state.

  5. Jun 17, 2008 #4
    Thank you for the replies.

    Borek, interesting. I just expected that when it comes to chemistry, there would not be exceptions to rules.

    I didn't even realize that about chromium. Thanks. :)

    ZapperZ, yeah I already encountered that throughout all of groups 4 and 5. That's why I expected group 6 to be the same, except applied to F as well.

    Then, what is the configuration of Yttrium? I have seen [Kr]5s^1 4d^1, but also [Kr]5s^2 4d^1. The former breaks the rule, but now that apparently isn't impossible, so, I don't know which it is.
  6. Jun 17, 2008 #5


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    Staff: Mentor

    There different kinds of rules. There are rules like - say - first law of thermodynamics, or mass and energy conservation. These are rock hard.

    But then there are rules like Aufbau principle, which states that "The orbitals of lower energy are filled in first with the electrons and only then the orbitals of high energy are filled." This is not a hard rule, more like rule of thumb - usually works, sometimes not, as sometimes to obtain lower energy of the whole system it is better to put electrons on different orbitals.
    Last edited by a moderator: Aug 13, 2013
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