1. Unit. 5 Electron Configuration

2. Pauli Exclusion Principle
No two electrons have the same quantum #’s
Maximum electrons in any orbital is two ()
Pauli Exclusion Principle

3. How are electrons arranged in an atom
The two natures of electromagnetic radiation: Particles vs. Waves
How to use the periodic table to list the configuration or orbital diagram
What quantum numbers are and how they are related to electron configuration.
How the periodic table is arranged
The basic periodic trends 3

4. Property of Waves Wave Nature of Electromagnetic Radiation Frequency
No. of waves per second
Wave Length
Distance between corresponding points in a wave
Amplitude
Size of the wave peak
Wave Nature of Electromagnetic Radiation

5. C = λ  f 5

6. Frequency is inversely proportional to Wavelength
If λ increases f decreases
If f increases λ decreases
Speed of the wave is always constant at 3.0 x 108 m/s

7. Wave nature could not explain all observations (Plank & Einstein)
Photoelectric Effect
When light strikes a metal electrons are ejected
Atomic Line Spectra
When elements are heated, they emit a unique set of frequencies of visible and non-visible light.
Particle Nature of Light

8. Photoelectric Effect – There is a minimum frequency to eject the electron

9. E = h f Photoelectric Effect
Only explained by “energy packets” of light called a quantum
Quantum - minimum amount of energy that can be gained or lost by an atom
Photons are massless particles of light of a certain quantum of energy
Photoelectric Effect
E = h f
H = Plank’s constant
As frequency increases
energy increases

10. Electrons in an atom add energy to go to an “excited state”.
Atomic Line Spectra
Electrons in an atom add energy to go to an “excited state”.
When they relax back to the ground state, they emit energy in specific energy quanta

11. 5 ______ 4 ______ 3 ______ 2 ______ 1 ______ 5 ______ 4 ______
These observations suggested that electrons must exist in defined energy levels
First, the electron absorbs energy and jumps from the ground state to an excited state
Next, the excited electron relaxes to a lower excited state or ground state
5 ______
4 ______
3 ______
2 ______
1 ______ hv
5 ______
4 ______
3 ______
2 ______
1 ______
5 ______
4 ______
3 ______
2 ______
1 ______ hv

12. Bohr’s Atomic Model of Hydrogen
Bohr - electrons exist in energy levels AND defined orbits around the nucleus.
Each orbit corresponds to a different energy level.
The further out the orbit, the higher the energy level

13. Transitions of the visible spectrum
There are three main series of transition
Paschen Series- IR
Drops into 3rd Level
Balmer Series - Visible
Drops into 2nd Level
Lyman Series - UV
Drops into 1st Level

14. De Broglie Heisenburg Modeled electrons as waves
Heisenberg Uncertainty Principle: states one cannot know the position and energy of an electron
Electrons exist in orbital’s of probability
Orbital - the area in space around the nucleus where there is a 90% probability of finding an electron

15. Schrödinger
Schrödinger Wave Equation - mathematical solution of an electron’s energy in an atom
quantum mechanical model of the atom – current model of the atom treating electrons as waves.

16. Wave Equation generates 4 variable solutions
n - size
l - shape
m - orientation
s – spin
Address of an electron
Quantum Numbers

17. n – Primary Quantum Number
Describes the size and energy of the orbital
n is any positive #
n = 1,2,3,4,….
Found on the periodic table
Like the “state” you live in

18. l – Orbital Quantum Number
Sub-level of energy
Describes the shape of the orbital
l = 0,1,2,3,4,….(n-1)
“City” you live in
l – Orbital Quantum Number
n = 3
l = 0,1,2
n = 2
l = 0,1
n = 1
l = 0

19. l – Orbital Quantum Number
# level = # sublevels
1st level – 1 sublevel
2nd level – 2 sublevels
4th level = 4 sublevels
l – Orbital Quantum Number

20. Sublevels are named for their shape
Spherical in shape p
l = 1
Dumbbell in shape d
l = 2 f
l = 3 f d s p

21. m – Magnetic Quantum Number
Describes the orientation of the orbital in space
Also denotes how many orbital's are in each sublevel
For each sublevel there are 2l +1 orbital's
m = 0, ±1, ±2, ±3, ±l
“Street” you live on
m – Magnetic Quantum Number

22. Can only be one s orbital
Look at Orbital's as Quantum Numbers
l = 1 m = -1, 0, +1
For each p sublevel there are 3 possible orientations, so three 3 orbital's
l = 0 m = 0
Can only be one s orbital

23. Orbital Designation n l M
2l+1
No. of Orbital
No. of Electron 3d 3 2
-2,-1,0,+1,+2 5 10 3p 1
-1,0,+1 6 3s 2p 2s 1s

24. n n2 2n2 4 s, p, d, f 16 32 3 s, p, d 9 18 2 s, p 8 1 s Energy Level
Possible sub-levels
Number of Sub-levels n
No. of Orbitals n2
No. of Electrons
2n2 4
s, p, d, f 16 32 3
s, p, d 9 18 2
s, p 8 1 s

25. How is the Bohr model different from the model of Dalton?
Who contributed to the modern model of the atom? How is it different from Bohr’s?
Why do atoms give unique atomic line spectra?
What are ground and excited states?
Is 2d possible? 4f ? 2s ? 6p? 1p?
How many total orbital's in the 2nd level? 4th level.

26. Quantum…. What….. ?

27. Aufbau Principal Lowest energy orbital available fills first
“Lazy Tenant Rule” 27

28. Pauli Exclusion Principle
No two electrons have the same quantum #’s
Maximum electrons in any orbital is two ()
Pauli Exclusion Principle

29. Hund’s Rule RIGHT WRONG
When filling degenerate orbital's, electrons will fill an empty orbital before pairing up with another electron.
Empty room rule
RIGHT
WRONG

31. Using the periodic table for the filling order of orbitals, by going in atomic number sequence until you use all the needed electrons in the element

32. An energy diagram for the first 3 main energy levels
Sub-level (l)
Increasing Energy
1 s ______
2 s ______
2 p ______ ______ ______
3 s ______
3 p ______ ______ ______
Level (n)
Orbitals (m)
An energy diagram for the first 3 main energy levels

33. An energy diagram for Neon
Increasing Energy
p ______ ______ ______
3 s ______
2 s ______
1 s ______
1s2
2s2
2px2
2py2
2pz2
1s2
2s2
2p6
Electron Configuration Notation
Electron Spin
An energy diagram for Neon

34. 1s22s22p4 electron configuration!
Orbital Notation shows each orbital
O (atomic number 8)
____ ____ ____ ____ ____ ____
1s s px 2py 2pz 3s
1s22s22p4 electron configuration!

35. ____ ____ ____ ____ ____ ____ ____ ____ ____ 1s 2s 2p 3s 3p
Write the orbital notation for S
S (atomic number 16)
____ ____ ____ ____ ____ ____ ____ ____ ____
1s s p s p
1s22s22p63s23p4
How many unpaired electrons does sulfur have?
2 unpaired electrons!

36. Bellringer What is the electron configuration for As ?
A. Use the long form
B. Use the noble gas notation
Log onto SFHS website – navigate to Staff / Curtis and select the Gizmo link

37. Valence Electrons As (atomic number 33)
1s22s22p63s23p64s23d104p3
The electrons in the outermost energy level.
s and p electrons in last shell
5 valence electrons 37

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

39. Example - Germanium
X X X X X X X X X X X X X
[Ar]
4s2
3d10
4p2

40. Let’s Practice P (atomic number 15) Ca (atomic number 20)
1s22s22p63s23p3
Ca (atomic number 20)
1s22s22p63s23p64s2
As (atomic number 33)
1s22s22p63s23p64s23d104p3
W (atomic number 74)
1s22s22p63s23p64s23d104p65s24d105p66s24f145d4
Noble Gas Configuration
[Ne] 3s23p3
[Ar] 4s2
[Ar] 4s23d104p3
[Xe] 6s24f145d4

41. Noble Gas Configuration Your Turn N (atomic number 7)
[He] 2s22p3
[Ne] 3s1
[Kr]5s24d105p3
[Ar] 4s23d4
Your Turn
N (atomic number 7)
1s22s22p3
Na (atomic number 11)
1s22s22p63s1
Sb (atomic number 51)
1s22s22p63s23p64s23d104p65s24d105p3
Cr (atomic number 24)
1s22s22p63s23p64s23d4

42. Full energy level
Full sublevel
Half full sublevel

43. Exceptions are explained, but not predicted!
Copper
Expect: [Ar] 4s2 3d9
Actual: [Ar] 4s1 3d10
Silver
Expect: [Kr] 5s2 4d9
Actual: [Kr] 5s1 4d10
Chromium
Expect: [Ar] 4s2 3d4
Actual: [Ar] 4s1 3d5
Molybdenum
Expect: [Kr] 5s2 4d4
Actual: [Kr] 5s1 4d5
Exceptions are explained, but not predicted!
Atoms are more stable with half full sublevel

44. Atoms take electron configuration of the closest noble gas
Atoms create stability by losing, gaining or sharing electrons to obtain a full octet
Isoelectronic with noble gases +1 +2 +3 +4 -3 -2 -1
Atoms take electron configuration of the closest noble gas

45. Na (atomic number 11) 1s22s22p63s1 1s22s22p6 = [Ne] 1 Valence electron
Metal = Loses Ne Na

46. Full Octet P-3 (atomic number 15) Ca+2 (atomic number 20)
1s22s22p63s23p6
Ca+2 (atomic number 20)
Zn+2 (atomic number 30)
1s22s22p63s23p63d10
Last valence electrons (s and p)
Full Octet

47. X 6 3 4 1 s electrons 7 2 5 8 p electrons
Shows valence electrons only!
s & p electrons
Write noble gas configuration for the element
Place valence electrons around element symbol in order X 6 3 4 1
s electrons
p electrons 7 2 5 8

48. O Fe Br Valence electrons Write the Lewis structures for: Oxygen (O)
[He] 2s2 2p4
Iron (Fe)
[Ar] 4s2 3d6
Bromine (Br)
[Ar] 4s2 3d10 4p5 • • • O • • •
Valence electrons Fe • • • • • Br • • • •