1. Electron Configurations

2. Quantum Mechanics
Better than any previous model, quantum mechanics does explain how the atom behaves.
Quantum mechanics treats electrons not as particles, but more as waves (like light waves) which can gain or lose energy.
But they can’t gain or lose just any amount of energy. They gain or lose a “quantum” of energy(Packets of E)
A quantum is just an amount of energy that the electron needs to gain (or lose) to move to the next energy level.
In this case it is losing the energy and dropping a level.

3. Atomic Orbitals
Much like the Bohr model, the energy levels in quantum mechanics describe locations where you are likely to find an electron.
Remember that orbitals are area around the nucleus where electrons are found.
Quantum mechanics calculates the probabilities where you are “likely” to find electrons.
Neil Bohr

4. Atomic Orbitals
Atomic orbital is where an electron is 90% more likely to be found
No more than 2 electrons can ever be in 1 orbital. The orbital just defines an “area” where you can find an electron.
What is the chance of finding an electron in the nucleus? Yes, of course, it’s zero. There aren’t any electrons in the nucleus.

5. Energy Levels
Quantum mechanics has a principal quantum number. It is represented by a little n. It represents the “energy level” similar to Bohr’s model.
n=1 describes the first energy level
n=2 describes the second energy level
Etc.
Each energy level represents a period or row on the periodic table.
Red n = 1
Orange n = 2
Yellow n = 3
Green n = 4
Blue n = 5
Indigo n = 6
Violet n = 7

6. Sub-levels = Specific Atomic Orbitals
Each energy level has 1 or more “sub-levels” which describe the specific “atomic orbitals” for that level.
n = 1 has 1 sub-level (the “s” orbital)
n = 2 has 2 sub-levels (“s” and “p”)
n = 3 has 3 sub-levels (“s”, “p” and “d”)
n = 4 has 4 sub-levels (“s”, “p”, “d” and “f”)
There are 4 types of atomic orbitals:
s, p, d and f
Each of these sub-levels represent the blocks on the periodic table.
Blue = s block
Yellow = p block
Red = d block
Green = f block

7. Shapes of Orbitals
There is 1 spherical shaped s orbital, so the s sublevel can hold 2 electrons.

8. Shapes of Orbitals
There are 3 dumbbell shaped p orbitals so p sublevel can hold 6 electrons

9. Shapes of Orbitals
There are 5 d orbitals so 10 electrons can fit in the d sublevel There are 7 f orbitals so 14 electrons can fit in the f sublevel)

10. Complete the chart in your notes as we discuss this.
Energy Level
Sub-levels
Total Orbitals
Total Electrons
Total Electrons per Level
n = 1 s
1 (1s orbital) 2
n = 2 p
1 (2s orbital)
3 (2p orbitals) 6 8
n = 3 d
1 (3s orbital)
3 (3p orbitals)
5 (3d orbitals) 10 18
n = 4 f
1 (4s orbital)
3 (4p orbitals)
5 (4d orbitals)
7 (4f orbitals) 14 32
Complete the chart in your notes as we discuss this.
The first level (n=1) has an s orbital. It has only 1. There are no other orbitals in the first energy level.
We call this orbital the 1s orbital.

11. Where are these Orbitals?2p 3p 3d 4p 4d 5p 5d 6p 6d 7p 4f 5f

12. Electron Configurations
What do I mean by “electron configuration?”
The electron configuration is the specific way in which the atomic orbitals are filled.
Think of it as being similar to your address. The electron configuration tells where all the electrons “live.”

13. Rules for Electon Configurations
In order to write an electron configuration, we need to know the RULES.
3 rules govern electron configurations.
Aufbau Principle
Pauli Exclusion Principle
Hund’s Rule
Using the orbital filling diagram at the right will help you figure out HOW to write them
Start with the 1s orbital. Fill each orbital completely and then go to the next one, until all of the electrons have been acounted for.

14. Fill Lower Energy Orbitals FIRST http://www. meta-synthesis
Each line represents an orbital.
1 (s), 3 (p), 5 (d), 7 (f)
The Aufbau Principle states that electrons enter the lowest energy orbitals first.
The lower the principal quantum number (n) the lower the energy.
Within an energy level, s orbitals are the lowest energy, followed by p, d and then f. F orbitals are the highest energy for that level.
High Energy
Low Energy

15. No more than 2 Electrons in Any Orbital…ever.
The next rule is the Pauli Exclusion Principal.
The Pauli Exclusion Principle states that an atomic orbital may have up to 2 electrons and then it is full.
The electrons in an orbital should be of opposite spins
We usually represent this with an up arrow and a down arrow.
Since there is only 1 s orbital per energy level, only 2 electrons fill that orbital.
Wolfgang Pauli, yet another German Nobel Prize winner

16. Hund’s Rule
Hunds Rule states that when you get to degenerate orbitals, you fill them all half way first, and then you start pairing up the electrons.
What are degenerate orbitals?
Degenerate means they have the same energy.
So, the 3 p orbitals on each level are degenerate, because they all have the same energy.
Similarly, the d and f orbitals are degenerate too.
Don’t pair up the 2p electrons until all 3 orbitals are half full.

17. NOW that we know the rules, we can try to write some electron configurations.
Remember to use your orbital filling guide to determine WHICH orbital comes next.

18. Refer to the following electron configuration and answer the questions below: 1s22s22p63s23p64s23d5
a. How many total electrons does this element have?
b. What element is this?
c. How many energy levels are represented?
d. How many sublevels are represented?

19. Try these!!!
Write the electron configuration for each element: Mg: Au:

20. Try these!!!
Mg 1s22s22p63s2 Au 1s22s22p63s23p64s23d104p65s24d105p66s2 4f145d9

21. Electron configuration for Orbital Box diagrams: Each arrow represents an electron Ex. Oxygen Atomic number = 8, so 8 electrons. Start with s, then fill up orbitals as you go.
1s s p4

22. Draw a box diagram for each of these elements:
Na :
P ;

23. Electron Configurations
Element
Configuration
H Z=1
1s1
He Z=2
1s2
Li Z=3
1s22s1
Be Z=4
1s22s2
B Z=5
1s22s22p1
C Z=6
1s22s22p2
N Z=7
1s22s22p3
O Z=8
1s22s22p4
F Z=9
1s22s22p5
Ne Z=10
1s22s22p6
(2p is now full)
Na Z=11
1s22s22p63s1
Cl Z=17
1s22s22p63s23p5
K Z=19
1s22s22p63s23p64s1
Sc Z=21
1s22s22p63s23p64s23d1
Fe Z=26
1s22s22p63s23p64s23d6
Br Z=35
1s22s22p63s23p64s23d104p5
Note that all the numbers in the electron configuration add up to the atomic number for that element. Ex: for Ne (Z=10), = 10

24. Shorthand electron configuration
You can use the noble gases to help make electron configuration easier to do.
First identify the noble gas (Group 8A) that is right BEFORE the element.
Write that noble gas symbol in [brackets]
Then write the rest of the configuration.
Ex. Na: 1s22s22p63s1 or [Ne]3s1

25. examples: shorthand electron configuration
1.Mg
[Ne]3s2
2. Al
[Ne]3s2 3p1
3. Sn
[Kr] 5s24d105p2
4. Br
[Ar] 4s23d104p5
5. Po
[Xe]6s24f145d106p4

26. One last thing. Look at the previous slide and look at just hydrogen, lithium, sodium and potassium.
Notice their electron configurations. Do you see any similarities?
Since H and Li and Na and K are all in Group 1A, they all have a similar ending. (s1)

27. Electron Configurations
Element
Configuration
H Z=1
1s1
Li Z=3
1s22s1
Na Z=11
1s22s22p63s1
K Z=19
1s22s22p63s23p64s1
This similar configuration causes them to behave the same chemically.
It’s for that reason they are in the same family or group on the periodic table.
Each group will have the same ending configuration, in this case something that ends in s1.