1. Sunday, September 30 at 4:10pm

2. Some stuff you might have forgotten (already)

3. Some practice  A mercury atom is initially in its lowest possible (or ground state) energy level. The atom absorbs a photon with a wavelength of 185nm, and then emits a photon with a frequency of 6.88x10 14 Hz. At the end of this series of transitions, the atom will still be in an energy level above ground state.  Q: Draw an energy level diagram for this process, and find the energy of this resulting excited state, assuming we assign E = 0 to the ground state.

4. Photoelectric Practice  What is the kinetic energy of an electron removed from iron using light with a wavelength of 101 nm? The binding energy of an electron in iron is 7.5x10 –19 J.  Bonus: what is the speed of this electron?

5. A brief look at

6. Bohr’s major contribution was the suggestion of stable orbits that electrons occupy. His model is the most common depiction of atoms, but sadly is also incorrect.

7. Now to the real show!

8. What the heck are these things?  The principle quantum number indicates the shell  The second indexes energy differences between orbitals in the same shell (subshells)—also, gives letter designations  The magnetic quantum number gives insight to the orientation of the orbital

9. Some ways to remember…  Principal quantum number—periods on the periodic table  Second quantum number—region of the periodic table (s,p,d, or f)  Each orbital actually has a 4 th quantum number, the “spin” – we’ll learn this tomorrow

10. Practice  An orbital has quantum numbers of n = 4, l = 2, and m l = -1. Which type of orbital is this?  How many orbitals are there in an s sub- shell? p sub-shell, d, f?

11. More practice  Which of the following represent valid sets of quantum numbers? For a set that is invalid, explain briefly why it is not correct. n = 3, l = 3, m l = 0 n = 2, l = 1, m l = 0 n = 6, l = 5, m l = -1 n = 4, l = 3, m l = -4

12. Don’t forget!