1. Recap of Last Class Recall:
Werner Heisenberg formulated the Uncertanity Principle that states it is impossible for us to know an electron’s exact position (where it is) and momentum (where it is going)
As a result, we cannot identify specific orbits that electrons travel in
We can only identify regions of space within an atom where an electron is most likely to be found
ORBITALS!
Schrodinger’s complex math equation allows us to:
Calculate the shape of the electron cloud
Probability of finding the electron at distinct locations within those clouds

2. How do the Orbitals Fill Up with Electrons?
An Introduction to Electron Configurations

3. Assigning an Electron’s Address Explore
Complete the activity “Welcome to Atomos Apartments!” on page 208

4. Predicting Electron Locations
We use electron configurations
The way electrons are arranged in atoms
There are rules to follow!
Aufbau principle
Electrons are added one at a time to the lowest energy orbitals available until all the electrons of the atom have been accounted for
“aufbau”
German for ‘build up or construct’

5. aufbau chart 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f

6. Predicting Electron Locations
Pauli’s Exclusion Principle
An orbital can hold only two electrons

7. Predicting Electron Locations
Hund’s Rule
“Electrons must fill a sub-level such that each orbital has a spin up electron before they are paired with spin down electrons”
A bus analogy:
If you enter a bus and don’t know anyone on it, you will pick a seat that is completely empty rather than one that already has a person in it

8. Orbital Diagrams and Electron Configurations
Electrons fill in order from lowest to highest energy
The Pauli exclusion principle holds. An orbital can hold only two electrons
Two electrons in the same orbital must have opposite signs (spins)
You must know how many electrons can be held by each orbital
2 for s
6 for p
10 for d
14 for f
Hund’s rule applies. The lowest energy configuration for an atom is the one having the maximum number of unpaired electrons for a set of orbitals
By convention, all unpaired electrons are represented as having parallel spins with the spin “up”

9. Electron Configuration Practice
Just a thought…
How do you determine the number of electrons in an element?
Examples:
Oxygen
Magnesium
Argon
Scandium

10. Orbital Notation

11. Short-Hand Notation
Use the Noble Gas symbol to abbreviate or shorten the electron configuration
Krypton
Rubidium
Zirconium

12. How Can We “Locate” an Electron?
Use Quantum Numbers!

13. Quantum Numbers
Each electron has a specific ‘address’ in the space around a nucleus
An electrons ‘address’ is given as a set of four quantum numbers
Each quantum number provides specific information on the electrons location

14. Electron Configuration
state
town
house number
street

15. Quantum Numbers state (energy level) - quantum number n
town (sub-level) - quantum number l
street (orbital) - quantum number ml
house number (electron spin) - quantum number ms

16. Principal Quantum Number (n)
Same as Bohr’s n
Integral values: 1, 2, 3, ….
Indicates probable distance from the nucleus
Higher numbers = greater distance from nucleus
Greater distance = less tightly bound = higher energy

17. Angular Momentum Quantum Number (l)
Integral values from 0 to n - 1 for each principal quantum number n
Indicates the shape of the atomic orbitals
Table 7.1 Angular momentum quantum numbers and corresponding atomic orbital numbers
Value of l 1 2 3 4
Letter used s p d f g

18. Magnetic Quantum Number (ml)
Integral values from l to -l, including zero
Relates to the orientation of the orbital in space relative to the other orbitals
3-D orientation of each orbital

19. Magnetic Quantum Number

20. Electron Spin Quantum Number (ms)
An orbital can hold only two electrons, and they must have opposite spins
Spin can have two values, +1/2 and -1/2
Pauli Exclusion Principle (Wolfgang Pauli)
"In a given atom no two electrons can have the same set of four quantum numbers"

22. Closer!
Complete the Closer on Page 206

23. Homework
Begin homework on page 209 – FRONT AND BACK!