1. ELECTRON CONFIGURATION
Where are the electrons of an atom located?
Why are electrons and their arrangement so important?
How do we show the arrangement or configuration of electrons in an atom or ion?

2. The Bohr Model
Nucleus
Electron
Orbit
Energy Levels 2

3. } Bohr’s Model Fifth Fourth Increasing energy Third Second First
Further away from the nucleus means more energy.
There is no “in between” energy
Fourth
Third
Increasing energy
Second
First
Nucleus 3

4. Find the Number of Electrons
Copper atom (Cu)
Neon atom (Ne)
Sodium ion (Na+1)
Fluoride ion (F-1)

5. 3-D Orientation of Electrons
Electrons are distributed among the atomic orbitals of the various energy levels. 
An “orbital” can be defined as the region or electron cloud where electrons are most likely to be found.
Electron probability cloud
Electron density cloud
The electron configuration of an element or ion shows this arrangement or distribution of electrons.

6. WRITING ELECTRON CONFIGURATION
The format consists of a series of numbers, letters and superscripts as shown below:
1s2
Here we see the electron configuration for the element helium. 
The “2” represents the total number of electrons in that sublevel
The “1” represents the principal energy level.
The “s” represents the type of sublevel

7. SHAPE OF ORBITALS in EACH SUBLEVEL
p holds 6 e-
s holds 2 e-
d holds 10 e-
f holds 14 e-

8. EACH ORBITAL CAN HOLD 2 ELECTRONS

9. Electron Capacity of the First Four Principle Energy Levels  
Electron Capacity of the First Four Principle Energy Levels
Principal energy level
(n)
Type of sublevel
Number of orbitals per type
Number of orbitals per level
(n2)
Maximum number of electrons (2n2) 1 s 2 4 8 p 3 9 18 d 5 16 32 f 7

10. WRITING ELECTRON CONFIGURATION
The electron configuration of hydrogen (1 electron). It must occupy the lowest possible energy levels. Aufbau principle. 1 1s
The electron configuration of lithium. 2 1 1s 2s
The electron configuration of beryllium?

11. ELEMENTS 1 - 5
The two electrons in He represent the complete filling of the first electronic shell. The up or down arrow each represents 1 electron.

12. FILLING UP “P” SUBLEVELS
Electrons repel each other, by occupying different orbitals the electrons remain as far as possible from one another.
They do not pair up in an orbital unless they have to. Hund’s Rule

13. Visualizing it together

14. OVERLAP
Since the 4s level has lower energy than 3d, 4s must fill up before 3d.

15. SUBLEVEL FILLING ORDER

16. ELECTRON CONFIGURATION AND THE PERIODIC TABLE

17. THE LONG VERSION OF THE PERIODIC TABLE

18. ORBITAL DIAGRAM Draw the orbital diagram for: 1s2 2s22p6 3s1
1s2 2s22p s2 3p4

19. PRACTICE Write the electron configuration and the orbital diagram for:
Fluorine:
Magnesium:
Silicon:

20. MORE PRACTICE
Write the electron configuration and the orbital diagram for:
Argon:
Potassium:
Bromine:

21. ADDITIONAL QUESTIONS
Answer the following questions based on the given electron configuration of an atom.
1s2 2s2 2p6 3s2 3p4
How many electrons are in this atom?
What element is represented by this electron configuration?
What is the outermost energy level?
How many principal energy levels are filled?
How many valance electrons are in this atom?

22. KEY TO STABILITY
Outermost s and p sublevels are completely filled - the most stable. (Noble Gases)
Outermost sublevels are completely filled. (i.e. 4s2 more stable than 4s23d1 or 3d10 is more stable than 3d9)
Outermost sublevels are half-filled. (i.e. 3d5 is more stable than 3d4 or 3d6 and 2p3 is more stable than 2p2 or 2p4)

23. Chromium and Copper Cr Cu [Ar] 3d5 4s1 - Por qué?
[Ar] 3d10 4s1 - Pourquoi? Wei shen me?

24. QUANTUM NUMBERS
Each electron in an atom is described by four different quantum numbers. The first three (n, l, ml) specify the particular orbital of interest, and the fourth (ms) specifies how many electrons can occupy that orbital.
Schrodinger derived an equation that described the energy and position of the electrons in an atom

25. n l m l ms 1 < n 1, 2, 3, 4 … Name Symbol Orbital meaning
Range of values
Principal quantum number n
energy level or shell
(size)
1 < n
1, 2, 3, 4 …
Azimuthal quantum number l
sublevel or subshell
(shape)
(s = 0, p =1, d = 2, f = 3)
Zero or integer less than n
Magnetic quantum number
m l
energy shift
(orientation)
(px, py or pz)
- l < ml < l
Spin projection quantum number ms
spin of the electron
(-1/2 = counter-clockwise, +1/2 = clockwise)
+1/2 or –1/2

26. Equating Electron Configuration with Quantum Numbers
1, 0, 0, +½
1, 0, 0, -½

27. 2nd Principal Energy Level
2, 0, 0, +½ or - ½
2, 1, 0, +½ or - ½
2, 1, 1, +½ or - ½
2, 1, -1, +½ or - ½
When filling orbitals of the same energy level i.e. 2px, 2py, and 2pz , electrons will fill each orbital before pairing up in an orbital – Hund’s Rule

28. 3rd Principal Energy Level
No two electrons in an atom can possess the same set of four quantum numbers.
– Pauli Exclusion Principle

29. EACH ORBITAL CAN HOLD 2 ELECTRONS

30. Diamagnetic vs Paramagnetic Elements
Diamagnetic elements have all of their electrons spin paired.
These materials are weakly magnetized in the direction opposite to the applied field.
Examples of diamagnetic material are bismuth, copper
Paramagnetic elements do not have all of their electrons spin paired and are strongly affected by magnetic fields.
These materials have atoms that have permanent magnetic moments.
These moments interact weakly with each other and randomly orient in different directions