1. Electron Configurations
Section 4.3

2. Arrangement of Electrons
Electron configuration: the arrangement of electrons in an atom
Electrons in atoms assume arrangements that have the lowest possible energies
This makes them more stable
Ground-state electron configuration: the lowest energy arrangement of the electrons for each element

3. Rules for Electron Location
Aufbau principle: an electron occupies the lowest-energy orbital that can receive it
Pauli exclusion principle: no two electrons in the same atom can have the same set of four quantum numbers

4. Last Rule
Hund’s rule: orbitals of equal energy are each occupied by one electron before any orbital is occupied by a second electron and all electrons in single occupied orbitals must have the same spin state

5. Electron Configuration
Find the element on the periodic table
Start with the lowest energy level, fill the sublevels with electrons until you get to your element
The principle quantum number is written first, followed by the sublevel superscripted with the number of electrons in that sublevel

6. Use the Periodic Table to Find the Sublevels for Your Element

7. Blocks of Sublevels 1 2 3 4 5 6 7
3d1

8. Examples Oxygen has 8 electrons
The first two electrons will go in the 1s orbital, 1s2
The next two go in the 2s, 2s2
The next 4 go in the 2p orbitals for 2p4
Complete electron configuration:
1s22s22p4 (make sure you have 8 e-)
Do beginning practice worksheet here.

9. Another Way for Aufbau

10. Exception You would expect chromium’s electron configuration to be:
1s22s22p63s23p64s23d4
But it is more stable if the 3d sublevel is half full, so what really happens is:
1s22s22p63s23p64s13d5 as Cr steals a 4s electron and puts it in the 3d sublevel

11. Another Exception
You would expect copper’s electron configuration to be:
1s22s22p63s23p64s23d9
But it is more stable if the 3d sublevel is full, so what really happens is:
1s22s22p63s23p64s13d10 as Cu steals a 4s electron and puts it in the 3d sublevel
Do other worksheet here

12. Orbital Notation (Diagrams)
Sometimes we want to see exactly which orbital is occupied and we want to specify the spin
Only two electrons can be in any one orbital, and they must have opposite spin
Every orbital in a sublevel gets one electron before any get two, and they all have the same spin

13. Basic Directions Write the electron configuration
Write lines based on the orbitals
One line for s, 3 lines for p, 5 lines for d
Draw arrows to represent the electrons, filling according to Hund’s rule and the Pauli exclusion principle

14. Practice For oxygen: 1s22s22p4 ___ ___ ___ ___ ___ 1s 2s 2p
___ ___ ___ ___ ___
1s s p
8 electrons have been placed and Hund’s rule and the Pauli exclusion principle has been followed
Practice worksheet here

15. Shorthand or Noble Gas Notation
Noble gases: group 18 elements
Noble gas notation: AKA shorthand, when the preceding noble gas’s configuration is followed by the rest of the element’s configuration
Example: for vanadium write the preceding noble gas in brackets, [Ar] followed by 4s23d3 to be written as [Ar] 4s23d3
Practice more

16. Excited State vs. Ground State
An atom is in an excited state when the electron configuration shows electrons in higher than normal orbitals
Na would normally be: 1s22s22p63s1
An excited state would be 1s22s22p53s2