1. Electron Configuration

2. Why we study electrons
Knowing the position of electrons helps us with many topics:
Why chemical reactions occur
Why some atoms are more stable than others
Why some elements react with only certain atoms
We want to know how many electrons an atom has, and where they’re located.

3. The Electron Cloud
Bohr’s model of the atom only showed us that electrons have different energy levels
Our current “quantum mechanical” model of the atom goes even further and talks about the different orbital shapes inside these energy levels.
Quantum numbers link
The world of quantum mechanics 1 2 3 4 5

4. 4 QUANTUM NUMBERS n l m s N - Principal Energy Level
Higher values of n mean more energy for the electron and the corresponding radius of the electron cloud or orbital is further away from the nucleus. Values of n start at 1 and go up by integer amounts. The higher the value of n, the closer the corresponding energy levels are to each other.

5. QUANTUM NUMBERS n l m s L - Angular Quantum Number
In chemistry, there are names for each values of ℓ. The first value, ℓ = 0 called an s orbital. s orbitals are spherical, centered on the nucleus. The second, ℓ = 1 is called a p orbital. p orbitals are usually polar and form a teardrop petal shape with the point towards the nucleus. ℓ = 2 orbital is called a d orbital. These orbitals are similar to the p orbital shape, but with more 'petals' like a clover leaf. They can also have ring shapes around the base of the petals. The next orbital, ℓ=3 is called an f orbital. These orbitals tend to look similar to d orbitals, but with even more 'petals'.

6. QUANTUM NUMBERS n l m s M - Magnetic Quantum Number
The third quantum number is the magnetic quantum number, m. These numbers were first discovered in spectroscopy when the gaseous elements were exposed to a magnetic field. This relationship shows for every value of ℓ, a corresponding set of values of m ranging from -ℓ to ℓ is found. This number determines the orbital's orientation in space.
For example, p orbitals correspond to ℓ=1, can have m values of -1,0,1. This would represent three different orientations in space for the twin petals of the p orbital shape. They are usually defined to be px, py, pz to represent the axes they align with.

7. QUANTUM NUMBERS n l m s S - Spin Quantum Number
The fourth quantum number is the spin quantum number, s. There are only two values for s, +½ and -½. These are also referred to as 'spin up' and 'spin down'. This number is used to explain behavior of individual electrons as if they were spinning in a clockwise or counterclockwise. The important part to orbitals is the fact that each value of m has two electrons and needed a way to distinguish them from one another..

8. QUANTUM NUMBERS

9. QUANTUM NUMBERS SUMMARY
Principal
Angular Momentum
Magnetic
Spin
Symbol N l Ml Ms
Values
N-1
-l to l
What it describes
How far the electron is f rom nucleus
Describes the shape or type of orbital
Describes the electron’s position in the orbital
Describes the spin
N=1, 2, 3, 4, 5, 6, 7
l=0, 1, 2, 3
s, p, d, f
Ml=
s 1 position
p 3 positions -1, 0, 1
D 5 positions -2, -1, 0, 1, 2
F 7 positions -3, -2, -2, 0, 1, 2, 3
+1/2 counter-clockwise
-1/2 clockwise
Maximum # e-
2N2
2e for each position
S=2, p=6, d=10, f=14
Maximum # sub levels
Eg. e level 2
l=0 or l=1 s, p
Pauli Exclusion Principle:
NO 2 ELECTRONS in the same atom have the same 4 quantum numbers

10. Orbital – region where electrons of a specific energy are likely to be.
-up to 2 electrons per orbital

11. Summary of Principal Energy Levels, Sublevels, Orbitals
Principal E Level N
Number of Sublevels =N
Types of Sublevels
Orbitals
=N2
# Electrons in Energy Level
=2N2
N=1 1
1s (1 orbital) 2
N=2
2s (1 orbital), 2p (3 orbitals) 4 8
N=3 3
3s (1 orbital)
3p (3 orbitals)
3d(5 orbitals) 9 18
N=4
4s (1 orbital)
4p (3 orbitals)
4d(5 orbitals)
4f (7 orbitals) 16 32
N=5
N=6
N=7

12. Orbital Diagrams
The order in which will fill a sublevel all comes down to energy!
Electrons fill the lowest energy sublevels first
The periodic table helps to show this order!

13. C. Periodic Patterns s p d (n-1) f (n-2) 1 2 3 4 5 6 7 6 7
© 1998 by Harcourt Brace & Company

14. Help a brother out….
When working out electron configuration, it’s important to understand orbital diagrams.
Here are 3 rules that will help us understand and interpret orbital diagrams.

15. A. General Rules Electrons fill the lowest energy sublevel available.
Aufbau Principle
Electrons fill the lowest energy sublevel available.
You cannot skip any sublevels

16. A. General Rules Pauli Exclusion Principle
Each orbital can hold only TWO electrons with opposite spins.
Represents an electron rotating clockwise
Represents an electron rotating counter - clockwise

17. A. General Rules WRONG RIGHT Hund’s Rule
Within a sublevel, place one e- per orbital with the same spin before pairing them.
“Empty Seat Rule”
WRONG
RIGHT

18. A Few more things….
Each arrow you draw represents an electron. (remember the three rules!)
Each orbital holds 2 electrons and has different shapes
The S level has only 1 orbital= total of 2 e-
The P level has 3 orbitals = total of 6 e-
The D level has 5 orbitals = total of 10 e-
The F level has 7 orbitals = total of 14 e-

19. Practice - Fill in the orbital diagram for Magnesium
*Use your PT to help you out!

20. Electron Configuration
An atom’s electron configuration…
Describes the location of electrons within the atom
Identifies the shape of the electron clouds
- regions where the electrons are held.
Uses numbers and letters to describe electrons’ location and which electron cloud it is part of.
Essentially, summarizes the orbital diagram!

21. 1s2 An example The “s” tells you the electron’s cloud shape.
In this case it’s a spherically shaped cloud. It is called the s sublevel. Sublevel = Shape
1s2
The “2” simply tells you how many electrons are in this cloud.
The “1” tells you how far from the nucleus the electrons can go.
In this case 2 electrons are creating the cloud.
In this case, its in the 1st energy level, which is the closest level to the nucleus.
Note: Every orbital can only hold 2 electrons. They must have opposite spins.

22. 1s2 2s2 2p4 O B. Notation 8e- 1s 2s 2p Orbital Diagram
Electron Configuration
1s2 2s2 2p4

23. C. Periodic Patterns Shorthand Configuration
This is a shorter way to do e- configurations
Steps…
1) Write the symbol of the element your doing
2) Find the noble gas (group 18) that comes before the element you are doing.
3) To begin, put that noble gas in [brackets]
4) Starting with that noble gas, finish the electron configuration

24. S 16e- 1s2 2s2 2p6 3s2 3p4 S 16e- [Ne] 3s2 3p4 B. Notation
Longhand Configuration
S 16e-
1s2
2s2
2p6
3s2
3p4
Core Electrons
Valence Electrons
Shorthand Configuration
S 16e- [Ne] 3s2 3p4

25. C. Shorthand practice
Example - Germanium
[Ar]
4s2
3d10
4p2

26. Practice these in shorthand
Li N Mg C

27. Draw this! This will help It is a way to remember
the order in which electrons
fill their orbitals.
When you reach the far
Left side, you reach a
“wall” and must go back to
The right hand side.