Lesson Plan for Atomic Orbitals

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Lesson Plan for Atomic Orbitals Powered By Docstoc
					Atomic Orbitals Lesson
What do an onion, test tubes full of colored dyes, and a light bulb have to do with chemistry?

Part 1 – The Onion

Bohr Model1:
    proposed in 1913
    electrons orbit the nucleus at a fixed distance with a definite energy

Heisenburg Uncertainty Principle2:

      proposed in 1920
      impossible to know the exact location and energy of an electron
      can only know where an electron is probably found in an atom
      visualized as a “cloud of electrons”

Energy level
      Energy sub-level

Energy Level (Shell)3:
    layer where electrons are found (specific distance from the nucleus)
    identified by numbers beginning with 1 (closest to the nucleus)
    energy of electrons increases as the shell number increases (farther away from the

Sub-Level (Sub-Shell)4:
    energy level within a shell at which all electrons have the same energy
    identified by a letter (s, p, d, f) and a number (shell number)
    total number of sub-levels for an element = the energy shell number

    region within a sub-level where an electron may be found
    never contains more than 2 electrons
        Orbitals & Electron Capacity of the First Four Principle Energy Levels
                                        Number of           Number of            Maximum
                 Type of sublevel       orbitals per        orbitals per         number of
energy level (n)
                                           type               level(n2)        electrons (2n2)
        1                  s                  1                   1                   2
                           s                  1
        2                                                         4                   8
                           p                  3
                           s                  1
        3                  p                  3                   9                   18
                           d                  5
                           s                  1
                           p                  3
        4                                                        16                   32
                           d                  5
                           f                  7

The sub-levels are filled with electrons by starting at the beginning of each arrow and following
it to the end. In other words, the order for filling in the sublevels becomes: 1s, 2s, 2p, 3s, 3p, 4s,
3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d,7p.

Electron Configuration:
    Form of notation that tells us where the electrons are located
    Example: Helium = 1s2
                        1 = tells us that the electrons are in the first energy level
                             which is the closest to the nucleus
                        s = tells us that the electrons are in the s sub-shell
                        2 = tells us that the orbital is full

Part 2 – The Colored Dyes

   1. Energy added to the atom (heat, friction, etc.)
          In this demonstration we’re adding energy in the form of ultra-violet light.
   2. Electron absorbs this energy and jumps to a higher energy level.
   3. Electron falls back to original energy level.
   4. As electron falls, it releases the extra energy as a photon (light).
                  In this demonstration, we see that the liquid is glowing -- fluorescing (light is in
                   the ultra-violet range).

Part 3 – The Light Bulb

This is the same concept as explained with the colored dyes demonstration, except that energy
added is heat and the light is emitted in the visible spectrum. For this reason we can see the light
produced by a light bulb.

Argon, Ar 5
    gas inside the light bulb
    1s22s2 2p63s2 3p6
    or [Ne] 3s2 3p6
Tungsten, W 5
    filament inside the light bulb
    1s22s2 2p63s2 3p63d10 4s2 4p64d10 5s2 5p64f145d46s2
    or [Xe] 4f14 5d4 6s2

    Science Insights. Scott Foresman-Addison Wesley.1999. pg 159-160, 409, 468.

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