1. QUANTUM NUMBERS

2. Remember… The Bohr atomic theory incorporated Plank’s theory of quanta of energy Bohr’s atomic spectra theory failed to explain the atomic spectra for elements with multiple electrons Scientists found that instead of the atomic spectra being made up of 1 line, they could be made up of multiple lines

3. If each line on the spectrum represented an electron dropping back to its ground state and releasing energy, where did the additional lines (electrons/energy levels) come from for elements like iron? Hydrogen Iron

4. Principal Quantum Number (n) Bohr labelled the energy levels (a.k.a. orbits, shells) using the letter n. Principal quantum number: n n designates the main energy level of an electron in an atom i.e. n = 1, 2, 3, 4, etc…

5. Secondary Quantum Number (l) l : secondary quantum number Michelson theorized that within an energy level there were different orbits/paths that an electron could take – a subshell

6. Secondary Quantum Number ctd. The number of subshells in an energy level is equal to that energy level’s value i.e. if n=2, then l=0, 1 (2 subshells in total) Ie n=3, then l = 0, 1, 2 l = 0  (n-1)

7. Magnetic Quantum Number (m l ) The magnetic quantum number (m l ) indicates the direction of the electron orbit Explains orientation of sublevels m l can be represented by -l to +l i.e. n= 2 l can = 0, 1 If l = 1 m l = -1, 0, 1

8. Spin Quantum Number (m s ) The spin quantum number indicates the direction that the electron is spinning Within each subshell, two electrons spin in opposite directions m s = -1/2 or +1/2

9. The Four Quantum Numbers Principal Quantum Number (n) Secondary Quantum Number (l) Magnetic Quantum Number (m l ) Spin Quantum Number (m s ) 1 00+1/2, -1/2 2 0 -1, 0, +1+1/2, -1/2 1 3 0 -2, -1, 0, +1, +2+1/2, -1/2 1 2 4 0 -3, -2, -1, 0, +1, +2, +3 +1/2, -1/2 1 2 3

10. So instead of this…

11. Electron distribution looks like this!