1. Electron Configuration

2. Principal energy levels
Electron Clouds
Electron cloud
The electron cloud is made of energy levels
Principal energy levels
Subshells
Energy levels are composed of subshells
Orbitals
Subshells have orbitals.

3. Subshell versus Orbital
Subshell – A set of orbitals with equal energy
Orbital – Area of high probability of the electron being located.
Each orbital can hold 2 electrons

4. Number of equal energy orbitals Total number of electrons possible
Types of Subshells
Subshell
Begins in energy level
Number of equal energy orbitals
Total number of electrons possible s 1 1 2 p 2 3 6
Energy increases d 3 5 10 f 4 7 14

5. What are electron configurations?
They show the grouping and position of electrons in an atom.
Electron configurations use boxes for orbitals and arrows for electrons.

6. Energy and Subshells 6p 5d 4f 6s 5p 4d 5s 4p 3d 4s 3p 3s 2p
Subshells are filled from the lowest energy level to increasing energy levels. 2s
Energy 1s

7. Aufbau Principle
The first of 3 rules that govern electron configurations
Aufbau Principle: Electrons fill subshells (and orbitals) so that the total energy of atom is the minimum 1
What does this mean?
Electrons must fill the lowest available subshells and orbitals before moving on to the next higher energy subshell/orbital.

8. Hund’s Rule
Hund’s Rule: Place electrons in unoccupied orbitals of the same energy level before doubling up. 2
How does this work?
If you need to add 3 electrons to a p subshell, add 1 to each before beginning to double up.

9. Pauli Exclusion Principle
Pauli Exclusion Principle: Two electrons that occupy the same orbital must have different spins. 3
“Spin” describes the angular momentum of the electron
“Spin” is designated with an up or down arrow.
How does this work?
If you need to add 4 electrons to a p subshell, you’ll need to double up. When you double up, make them opposite spins.

10. Determining the Number of Electrons
Charge = # of protons – # of electrons
Atomic number = # of protons
Example:
How many electrons does Br-1 have?

11. Determining the Number of Electrons
Charge = # of protons – # of electrons
Atomic number = # of protons
Example:
How many electrons does Br-1 have?
Charge = -1
Atomic number for Br = 35 = # of protons
-1 = 35 - electrons
Electrons = 36

12. Writing Electron Configurations
Aufbau Principle: Electrons fill subshells (and orbitals) so that the total energy of atom is the minimum 1 2
Hund’s Rule: Place electrons in unoccupied orbitals of the same energy level before doubling up.
Pauli Exclusion Principle: Two electrons that occupy the same orbital must have different spins. 3
Example:
Write the boxes & arrows configuration for Cl

13. Writing Electron Configurations
Aufbau Principle: Electrons fill subshells (and orbitals) so that the total energy of atom is the minimum 1 2
Hund’s Rule: Place electrons in unoccupied orbitals of the same energy level before doubling up.
Pauli Exclusion Principle: Two electrons that occupy the same orbital must have different spins. 3
Example:
Write the boxes & arrows configuration for Cl
No charge written  Charge is 0
Atomic number for Cl = 17 = # of protons
0 = 17 - electrons
Electrons = 17 4 13 12 11 14 15 8 17 16 10 9 1 2 5 3 6 7 1s 2s 2p 3s 3p

14. Sub-Energy Levels

15. Electron Configuration PT

16. Arrow-Orbital Diagrams
Energy 3d 4s 3p 3s 2p 2s 1s

17. Electron Configuration Symbols
# of e- in sub-energy level
5f 3
Sub-Energy Level
Energy Level

18. Bohr Models vs. e- ConfigsK K:
1s2
2s2
2p6
3s2
3p6
4s1

19. Spectroscopic Notation

20. Spectroscopic Notation
Shorthand way of showing electron configurations
The number of electrons in a subshell are shown as a superscript after the subshell designation 1s 2s 2p 3s 3p
1s2 2s2 2p6 3s2 3p5

21. Writing Spectroscopic Notation1
Determine the number of electrons to place 2
Follow Aufbau Principle for filling order
Fill in subshells until they reach their max (s = 2, p = 6, d = 10, f = 14). 3
The total of all the superscripts is equal to the number of electrons. 4
Example:
Write spectroscopic notation for S
No charge written  Charge is 0
Atomic number for S = 16 = # of protons
0 = 16 - electrons 1s 2s 2p 3s 3p 2 2 6 2 4
Electrons = 16 2 + 2 + 6 + 2 + 4 = 16

22. Noble Gas Configuration

23. Noble Gases & Noble Gas Notation
Noble Gas – Group 8 of the Periodic Table. They contain full valence shells.
Noble Gas Notation – Noble gas is used to represent the core (inner) electrons and only the valence shell is shown. Br
Spectroscopic 1s 2s 2p 3s 3p 2 6 4s 3d 10 4p 5
Noble gas
[Ar] 4s 2 3d 10 4p 5
The “[Ar]” represents the core electrons and only the valence electrons are shown

24. Which Noble Gas Do You Choose?
How do you know which noble gas to use to symbolize the core electrons?
Think: Price is Right.
How do you win on the Price is Right?
By getting as close as possible without going over.
Choose the noble gas that’s closest without going over!
Noble Gas
# of electrons He 2 Ne 10 Ar 18 Kr 36 Xe 54

25. Noble Gas Notation Example1
Determine the number of electrons to place 2
Determine which noble gas to use
Start where the noble gas left off and write spectroscopic notation for the valence electrons 3
Example:
Write noble gas notation for As

26. Noble Gas Notation Example1
Determine the number of electrons to place 2
Determine which noble gas to use
Start where the noble gas left off and write spectroscopic notation for the valence electrons 3
No charge written  Charge is 0
Example:
Write noble gas notation for As
Atomic number for As = 33 = # of protons
0 = 33 - electrons
[Ar] 4s 3d 4p 2 10 3
Electrons = 33 18 + 2 + 10 + 3 = 33
Closest noble gas: Ar (18)
Ar is full up through 3p