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Electron Configuration Exceptions

  1. Sep 17, 2012 #1
    This is the electron configuration for Ag (Silver) found on the wikipedia page:
    1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 4d10 5s1

    I used the Aufbau Principle 1jlz7.png and got this instead:
    1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 5s2 4d9

    1. I can see that Ag is an exception to the rule, but are there any generalizations I can make so that I don't write down the wrong configurations just because of an exception?

    2. Also, is it traditional to write these in order of increasing quantum number (just like it's written above), because the Aufbau Principle is meant for the orbitals to be filled in increasing energy?

    For example, Br is written, on wikipedia, as 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p5, but written using Aufbau Principle, it's actually
    1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p5. Which one is more correct?
     
    Last edited: Sep 17, 2012
  2. jcsd
  3. Sep 17, 2012 #2

    AGNuke

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    1.) In case of configuration like (n-1)d4 ns2 OR (n-1)d9 ns2, to attain a more stable configuration, the electron from s orbital jumps to the previous d orbital. So, the new configuration becomes (n-1)d5 ns1 and (n-1)d10 ns1 respectively.

    This phenomenon is called pseudo-inert effect, and according to this, half filled or fully filled orbitals cluster (like 2p, 3d orbitals) are more stable, and if this is attainable by one electron jump, most atoms are ready to do so. Example are Chromium Group (Group 6) and Copper Series (Group 11).

    2) It is OK if they clubbed the orbitals on the basis of their respective Principle Quantum number or by Aufbau's Principle, as long as it do not look absurd. But conventionally, people fill the orbitals with increasing Principle Quantum number, rather than by the diagram. Diagram is useful for allotting electrons, but writing configuration that way may seem confusing.
     
  4. Sep 17, 2012 #3
    One more thing: Why is filling or half-filling a p or d orbital more stable than filling or half-filling an s orbital?
     
  5. Sep 17, 2012 #4

    AGNuke

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    Probably owing to the fact that s orbital is a single orbital, one can half-fill it and second can fully. Moreover, p and d orbitals are clusters of 3 and 5 orbitals respectively, so they are stable if each orbital possess equal amount of electron.
     
  6. Sep 18, 2012 #5
    This question (original post) is uncovering a lot of grey areas and textbook misinformation and misunderstandings in the area of atomic structure. I will address several of these:

    First, the convention of writing electron configurations is always done principal quantum number first, regardless of relative energies.

    Thus the electron configuration of neodymium is written either

    Nd: 1s22s22p63s23p63d104s24p64d104f45s25p66s2

    or Nd: {Xe}4f46s2

    The electronic structure of silver is written

    Ag: {Kr}4d105s1, or its longer-winded equivalent.

    Second, "Aufbau" is not the name of a German physicist. It is actually a German word, that does not have an exact English equivalent. It describes the process of building up something piece by piece on the foundation of the simpler structure that has been arrived at by placement of the previous piece. Like the process of building up a brick wall by placing bricks in the structure one at a time. The Aufbau principle (never "Aufbau's principle") refers to obtaining the electron configuration for nitrogen by looking at the previous element C: 1s22s22p2, and saying "nitrogen needs an extra electron. Where can we put it?" to arrive at N: 1s22s22p3.

    Third, it is not the case that the periodic table groups elements with similar valence electron configuration. The classic example is in group 10, which contains
    Ni: {Ar} 3d84s2
    Pd: {Kr} 4d10
    Pt: {Xe} 5d96s1

    Fourth, it is not the case that an electron configuration with a half-filled d or f subshell has a lower energy than one that is one electron away from such a state. There is a second factor involved: Spin-Orbit coupling. Russell Saunders coupling is a model for this factor. It shows that electron configurations like those above can split into several different spectroscopic terms (electronic energy levels), with quite widely separated energies. The "electron configuration" that is quoted for each element in a periodic table is the one associated with the spectroscopic term that happens to be associated with the ground electronic state. The lowest energy electron configuration averaged over all of the spectroscopic terms associated with it can often come out quite differently from the actual electron configuration associated with the lowest energy spectroscopic term.
     
    Last edited: Sep 18, 2012
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