1. An investigation into electrons and their location and behavior within the atom Learning Targets:  Describe the process of excitation and emission of energy by an electron.

2.  The spectra that were shown through emission spectroscopy led Niels Bohr to question the structure of the atom.

3.  With white light, all of the colors of the visible spectrum are shown.

4.  Since that was NOT what the spectra of elements looked like, Bohr began to look at why only certain wavelengths of color appeared.

5. E = hc λ Energy h = 6.63 x 10 -34 Js wavelength in meters c = speed of light = 3x10 8 m/s This equation shows that larger wavelengths indicate lower amounts of energy and smaller wavelengths indicate higher amounts of energy... an inverse relationship. Bohr realized that the specific wavelengths revealed specific amounts of energy.

6. Specific amounts of energy!! That inferred that energy within the atom existed at specific amounts. Bohr called these orbits, or energy levels. An electron cannot be in-between energy levels, it can only be within an energy level. Therefore, energy is quantized.

7. Bohr realized that the spectra were being created as electrons moved between these energy levels:  If an electron absorbs energy, it may jump to a higher energy level.  When an electron is at a higher energy level we say that the electron is in its “excited” state.  When the electron releases energy in the form of radiation, we say that the electron has returned to its “ground” state.  The type of radiation that is emitted depends on the amount of energy released.

8. Nucleus 1 st Energy Level 3 rd Energy Level 2 nd Energy Level 4 th Energy Level Energy Coming In!

9. Nucleus 1 st Energy Level 3 rd Energy Level 2 nd Energy Level 4 th Energy Level Energy emitted (infrared) Energy emitted (red light) Energy emitted (ultraviolet light)

10. Nucleus 1 st Energy Level 3 rd Energy Level 2 nd Energy Level 4 th Energy Level Energy emitted (blue/green light) Energy emitted (ultraviolet light)

11. Nucleus 1 st Energy Level 3 rd Energy Level 2 nd Energy Level 4 th Energy Level Energy emitted (blue/green light) Energy emitted (red light)

12.  This is the full electromagnetic spectrum.

13.  Bohr saw Visible Light:  wavelength is in the range of 400 to 700 nanometers (4 x 10 -7 meters)  ROY G. BIV  White light is made of all the colors of light

14.  Gamma rays: cosmic radiation, very high energy  Ultraviolet rays (UV): solar radiation, high energy  Infrared rays (IR): thermal radiation, remote controls, low energy  Microwave rays: microwave oven, very low energy

15. 2 --> 1Ultraviolet 3 --> 1Ultraviolet 4 --> 1Ultraviolet 3 --> 2Visible Red 4 --> 2Visible Blue/Green 5 --> 2Visible Blue 4 --> 3Infrared Energy Level Change Spectra Emission

16.  In addition to knowing that there were energy levels in the atom, three scientists began to notice other things...  Heisenberg – impossible to know the exact position and exact speed of an electron at the same time  De Broglie – electrons have wave-like properties, as in they move in wave patterns  Schroedinger – developed probability of finding each electron in a given location

17.  Heisenberg  Bohr suggested that the electrons move in perfect circles around the nucleus.  Heisenberg showed that, instead, the electron moves in a three dimensional cloud of probability that is smeared out over the orbit – Heisenberg uncertainty principle

18.  DeBroglie  Bohr suggested that the electrons move in perfect circles around the nucleus.  DeBroglie showed that there were other shapes because the electrons moved like waves – wave-particle duality.

19.  Schrodinger  Schroedinger realized how to put the theories of Bohr, Heisenberg, and DeBroglie together by creating a mathematical equation to find the most likely location for each electron within an atom – wave equation. Watch this YouTube video.YouTube

20. Every electron within an atom has “coordinates”. Schrodinger gave these coordinates numerical values, known as quantum numbers. Each quantum number describes part of the coordinates that determine the energy and probable location of any electron for any atom.

21. Energy levels begin at the number 1. Each level is higher in energy than the next. The higher in energy, the farther away from the nucleus.

22.  Atoms are three dimensional.  Within the energy levels exist different shapes, or subshells.  The shapes are determined by how much energy is required to create them.

23. Did you notice that there were different positions of some of the subshells?  The different positions, or orientations, are called orbitals, not orbits.  The orbitals are determined by which subshell they are in and in which positions they are. The s orbital does NOT have a different position. The p orbital has THREE different orientations – x, y, and z.

24. Each electron can be spin up (+ 1 / 2 ) or spin down (- 1 / 2 ) No two electrons in the same orbital orientation can have the same spin. With only one spin up and one spin down, the maximum number of electrons that can fit into any given orbital orientation is two. This is called the Pauli Exclusion Principle.