1. Sections 7.1 – 7.3 Electron Spin, Orbital Energies and Electron Configurations

2. Atomic Electronic Structure
In these sections…
Electron Spin and Magnetism
Energies of Orbitals
Electron Configurations of Atoms

3. Electron Spin: 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

4. 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

5. 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 only two electrons, with opposite spins because ms can only be +1/2 or -1/2

6. Pauli Exclusion Principle
What’s allowed?

7. Orbital Energies Single Electron Atoms Multi-electron Atoms Why?
With a single electron, energy depends only on how far from the nucleus.
With multiple electrons, e-e- repulsions also play a role and differ depending on orbital shape.

8. Single Electron Atoms Multi-electron Atoms
For most atoms:
Energy increases as n increases: < 2 < 3 < 4 …
Energy increases as subshells progress: s < p < d < f

9. Atomic Electron Configurations
An atom has lots of electrons and lots of orbitals.
Which orbitals do the electrons occupy?

10. Atomic Electron Configurations
An atom has lots of electrons and lots of orbitals.
Which orbitals do the electrons occupy?
Electrons fill the lowest energy orbitals first.
Electron Configuration:
a listing of how many electrons
occupy each orbital.

11. 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.
An arrow represents an electron.

12. 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

13. 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

14. Electron Configurations and the Periodic Table: Periodic Blocks

15. Electron Configurations and the Periodic Table II

16. Periodic Table and the Order of Filling
In what order are
subshells filled?

17. Hund’s Rule: Subshells are filled to give the maximum number of unpaired electrons

18. Using Periodic Blocks: C

19. Using Periodic Blocks: Cl

20. Noble Gas Notation: Mg

21. Noble Gas Notation: Mg

22. Diamagnetic vs. Paramagnetic Elements

23. d-Block Elements: Fe

24. Two Anomalies:
Cr and Cu