Quantum Numbers

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. Hydrogen • 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? 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 (ml) • The magnetic quantum number (ml) indicates the direction of the electron orbit • Explains orientation of sublevels • ml can be represented by -l to +l i.e. n= 2 l can = 0, 1 If l = 1 ml = -1, 0, 1

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

9. The Four Quantum Numbers

10. So instead of this…

11. Electron distribution looks like this!