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Emission spectra and energy levels in atoms

  1. Mar 11, 2013 #1

    trollcast

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    We're just been learning about emission and absorption spectra and how these relate to the energy levels of electrons in an atom but its brought up 2 questions.

    1. In chemistry our notes say that the energy change between levels gets smaller as you move out from the nucleus, however in physics we were given an unnamed example where n=4 to n=3 was smaller than n=5 to n=4?

    2. I know it says that each emission spectra is unique to that element but are there any similarities between groups or periods of elements?


    Thanks
     
  2. jcsd
  3. Mar 11, 2013 #2

    sophiecentaur

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    Are you sure about this? Where is the diagram? Was it sketched or in a text book?
     
  4. Mar 11, 2013 #3

    trollcast

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    Just found my notes and its one of the exam style question examples: (Sorry about the typed diagram but its the best I can do:

    The lowest energy levels of a mercury atom are shown in the diagram below. The diagram is not to scale. (Energies in Jx10^-18)


    .................................................................. 0eV
    ------------------------------------------------------------------ -0.26 n = 4


    ------------------------------------------------------------------ -0.59 n = 3

    ------------------------------------------------------------------ -0.88 n= 2





    ------------------------------------------------------------------ -2.13 n=1 (ground state)

    I got the numbers a bit mixed up in the OP.

    Does that help?
     
  5. Mar 11, 2013 #4
    You can't assume that the "distance from the nucleus" always increases with n.


    Distance really means "average distance." The electrons are not really located at definite distances from the nucleus. In the "classical physics approximation," the electrons are in elliptical orbits around the nucleus. In the "quantum mechanics theory", the position of the electrons are smeared out due to the "uncertainty principle."

    The position of an electron in an energy level has both an "average", a "standard deviation", and many other statistical moments. Even though the energy of a level may be well determined, the position is not. There are also several types of averaging.

    The "average distance" from the nucleus increases with n for the hydrogen atom. However, this will not always be true for other atoms. I suspect that the chemistry textbook is talking only about elements in column I of the periodic table.

    For a neutral atom, the "energy" of a valence-electron tends to increase with "n+l". So if the "average distance" increases with "n+l", then the textbooks statement would be correct.

    Note n=1,2,3,.. and l=0, 1, 2, 3, ... So one can have an decreasing "average distance"

    I think that law which you are describing is far too ambiguous to be useful. Especially in chemistry. The distance from the electron to the nucleus is not something that can easily be measured in either chemistry or physics. I don't think there is any chemistry technique that can measure the distance from electron to nucleus.

    I recommend shelving that statement for a while. It may be more important in a different context, which may come later.

    In column III, there are two boxes that each contain 15 elements. The elements in each box is called a series. One box is for the "lanthanide series" and one box is for the "actinide series". Some of the emission lines span parts of each series. Some elements in each series resemble each other in chemical and spectroscopic terms.

    I have done a lot of work with the elements "terbium" and "europium". There is one emission line at a wavelength of 620 nm which belongs to both terbium and europium. There are other "coincidental" emission lines within the lanthanide series.

    The reason each series has its own properties has to do with the filling up of the f-orbitals with electrons. Again, this concept may be something that is not of immediate concern.
     
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