1. Electrons exhibit a magnetic field We think of them as spinning They can spin only two ways: think of it as left or right Spin quantum number: ms can be +1/2 or -1/2

2. Magnetic Properties come from additive effects of electron spins
Magnetic Properties come from additive effects of electron spins. Diamagnetic: all electrons are paired Paramagnetic: 1 or more unpaired electrons Ferromagnetic (real magnets): unpaired electrons all lined up in the same direction

3. Pauli Exclusion Principle
No two electrons in an atom can have the same 4 quantum numbers
n, ℓ, mℓ define an orbital
Therefore: an orbital can hold two electrons, with opposite spins because ms can only be +1/2 or -1/2

4. Orbital Energies
Why?
For a single electron, it only depends on how far from the nucleus
For many electrons, e-e repulsions also play a role and
differ depending on orbital shape

5. Orbital Energies For most atoms:
Energy increases as n increases: 1 < 2 < 3 < 4 …
Energy increases as subshells go from s < p < d < f

6. Electron Configurations
General Rule: electrons fill lowest energy orbitals first
Sodium, Na as an example
Na has 11 electrons.
Fill 2 electrons per orbital till you run out
A box represents an orbital.
A arrow represents an electron.

7. Electron Configurations: Three Notation Types1.
2. spdf (or spectroscopic) notation:
List subshells and how many
electrons they contain:
1s22s22p63s1
3. Noble gas notation: short
[Ne]3s1
Where [Ne] = 1s22s22p6

8. Electron Configurations and the Periodic Table
Examples using Electron Configuration Simulation
Periodic Blocks
Hund’s Rule (using the p block)
n value increases as you move down table
Anomalies: Cr and Cu

9. Electron Configurations and the Periodic Table I

10. Electron Configurations and the Periodic Table II

11. Electron Configurations and the Periodic Table III
What would the periodic table look
like if the rules were different?
For example, what if electrons could
only have a spin of +1/2 (and not -1/2)?
Sketch it.

12. Notes:
There’s no known reason electrons have spin, or have only two of them. The other stuff about orbitals is theoretically derived from Schrod. Equn., but the whole spin thing is just something we see. Can’t explain it- just know it’s true. Like gravity or Coulombic attractions.
The reason why different subshells have different energies: for example: The energy of the 2s subshell has to do with how well the 2s electrons are attacted to the nucleus minus how much they are repelled by the 1s electrons.
Same thing for the 2p electrons. Difference is, the 1s electrons repel the 2p electrons more than the 2s electrons, so the 2p electrons are less stable, and higher energy.
Same reasoning happens when you go to higher subshells (e.g. d > p)
Why does the 4s subshell come before the 3d subshell? The reason above about d being higher in energy plus the fact that as you go up in n value, the orbital energies all get closer together. So, 2 is much higher than 1; 3 is less higher than 2; 4 is not much higher than 3, etc. This comes from the En = -constant/n2
The reason for Hund’s Rule: there is less e-e repulsion if electrons are in different orbitals because they are in different
Places. That’s why they go to different orbitals in a subshell first. I don’t know why they go with the same spin.
Have them fill in the blanks for a set of elements as you use the simulation. Be sure to include a Hund’s rule one: B, C, N, or something like it.
Answer to hard question: the pt looks the same, but is half as wide for each block because each orbital can only hold a single electron.

13. Predicting Electron Configurations

14. Predicting Electron Configurations

15. Predicting Electron Configurations

16. Predicting Electron Configurations

21. Electron Configurations of Cations

22. Electron Configurations of Anions

23. Transition Metal Cations: Lose s electrons first

24. Diamagnetic vs. Paramagnetic Elements