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Energy of an electron

  1. Feb 22, 2007 #1
    Is it possible to find the energy of an electron of a particular element in a particular orbit?
     
  2. jcsd
  3. Feb 22, 2007 #2
    This question takes me back to my chemistry days, yes I believe it is. I seem to remember also that with hydrogen oribitals it's fairly easy, however the more electrons you have the harder it is to work out an individual electrons energy and thus: the larger the electron number the harder it becomes to determine the exact energy of an electron. At a simple level though it should be possible to give values for all electrons in a certain arrangement, although of course at the quantum level the arrangements are hardly that clearly defined :/

    You'll have to forgive me though it's been years since I had to work out energy levels and oribitals etc, I'm sure someone else has something more detailed; all I can vaguely remember really is that an electron needs a precise amount of energy or quanta to jump to a higher state or orbital, which is both where the term quantum mechanics comes from and quantum leap. IIRC

    Here's a simple-more complicated overview, should explain it:-

    http://en.wikipedia.org/wiki/Atomic_orbital
     
    Last edited: Feb 22, 2007
  4. Feb 22, 2007 #3
    Thnx .But the webpage duznt give me clear equations or sumthing lik that to calculate the energy.Anywyz the help was appreciated
     
    Last edited: Feb 22, 2007
  5. Feb 22, 2007 #4

    cristo

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    The energy of the electron in the ground state of hydrogen is -13.6eV. Here's a link discussing the other energy levels of the hydrogen atom: http://hyperphysics.phy-astr.gsu.edu/hbase/hyde.html#c2.

    It get's a bit more complicated when we discuss elements other than hydrogen, so it's best to study hydrogen first.
     
  6. Feb 22, 2007 #5
    What about the energy need for an electron of hydrogen to jump to another orbital?Say when a photon hits it?
     
  7. Feb 22, 2007 #6

    Hootenanny

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    It is the difference between the energy levels of the atom. For example, an electron in hydrogen in the ground state (n=1) has an binding energy of -13.6eV. The binding energy at n=2 is -3.4eV, hence for an electron transition from n=1 to n=2 the electron must 'absorb' a photon of energy E = |-13.6 + 3.4 | = 10.2eV. This energy corresponds to the wavelength of a photon in the UV range. Equally, when an electron becomes 'de-excited' i.e. transitions from n=2 to n=1 the electron will emit a photon of the same corresponding energy and hence wavelength. The process becomes a little more complicated if you start looking at the hyperfine structure of hydrogen or other elements in general (as cristo said).
     
  8. Feb 22, 2007 #7
    The energy of an electron in any orbital of any element is the solution to its Schrodinger equation for that specific case. With more than one electron it's hard; with more than two no analytic solutions exist and these must be computed at great length. The difficulty arises because each electron in the system also has a potential that must be applied to the nuclear potential - each of which varies in an equally complex fashion as the wavefunction of the electron you're studying. Things are complicated further by relativistic corrections to the electron's kinetic energy which become more important the larger the atom and the higher the energy levels. The Schrodinger equation with the simplest set of relativistic corrections is known as the Dirac Equation.
     
  9. Feb 22, 2007 #8
    Looking at Hootenanny's reply,one question comes to my mind.If a photon hits an electron of the H atom with not enough energy to make it jump up an energy level,what happends to the photon?
     
  10. Feb 22, 2007 #9

    jtbell

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    Nothing. It doesn't even really "hit" the atom, it just goes right on past.
     
  11. Feb 22, 2007 #10
    Er..how doesnt it hit the atom?Suppose it hits the electron but not with enough energy to make it jump an energy level,then does the electron start vibrating?Thats what one of my friends suggested!
     
  12. Feb 23, 2007 #11

    jtbell

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    Don't think of photons and electrons as little tiny balls that can be aimed so as to guarantee that they "hit" each other. At the quantum-mechanical level, interactions are always probabilistic in nature. There is always a certain probability for interaction versus a corresponding probability for non-interaction. For atomic excitation, the probability for interaction decreases very rapidly towards zero, as the photon energy gets further away from the nominal difference between two energy levels.
     
  13. Feb 23, 2007 #12
    Thanks a lot for the explanation.This might be able to explain why hydrogen is colourless.As Hootenanny said,ultraviolet rays are emitted by an excited hydrogen electron.hence we cannot see it.Is that correct?

    Also i have another doubt regarding hoot's explanation.If an infrared photon strikes an electron why is heat produced?The energy of the photon is so low that as u said it shouldn't interact with the electron or should have very low probability of interacting
     
  14. Feb 23, 2007 #13

    Hootenanny

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    Actually, Hydrogen does have spectral lines in the visible range, take a transition from n=3 to n=2 for example; this correspond to an energy of about 1.9eV which results in a wavelength of about 656nm which is in visible [red] light. There are more transitions which occur at the violet end of the spectrum. The energy of IR radiation also 'happens' to correspond to the energies separating the quantum states of molecular vibrations (which results in temperature), its not only atoms and electrons that have quantum states...
     
    Last edited: Feb 23, 2007
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