1. Electron Configuration and the Periodic Table
Mallard Creek Chemistry - Rines

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

3. Electromagnetic Radiation
Mathematical Relations
C = speed of light = 3.0 x 108 m/s
λ (lamda) = wavelength (m)
f= frequency (Hz or s-1)
This is how we know what color light is emitted!
C = λ  f

4. 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

5. Bohr Model Nucleus: Neutrons and Protons Orbits: Electrons
We know both specific energy and location of each electron
Electrons orbit the nucleus in certain fixed energy levels (or shells)
Energy Levels
Nucleus

6. Bohr Model 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

7. Bohr’s Model The Photoelectric Effect Atomic Emission Spectra
Light releases electrons
Not all colors work
Atomic Emission Spectra
Hydrogen gas emitted specific bands of light
Bohr’s calculated energies matched the IR, visible, and UV lines for the H atom 6 5 4 3 2 1

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

9. Electromagnetic Radiation
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
Based on the frequency and wavelength of the photon
Photoelectric Effect

10. Bohr’s Model Excited electrons
Energy added to atom – electrons “jump” up energy levels
When the atom relaxes - electron “falls” to lower energy levels and emits photon
Bohr Model of hydrogen
Reference Sheets!!!!!

11. Electromagnetic Radiation
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

12. Electromagnetic Radiation
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

13. Electromagnetic Radiation
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
E = hf

14. Other Scientists Contributions
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. Other Scientists Contributions
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. Quantum Mechanical Model
Nucleus: Neutrons and protons
Orbitals: region in space surrounding the nucleus where there is a 95% probability of finding an electron.
We know either energy or location of each electron.

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

18. Quantum Numbers 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

19. Quantum Numbers l – Orbital Quantum Number l = 0,1,2,3,4,….(n-1)
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

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

21. Sublevels are named for their shape
Quantum Numbers
Sublevels are named for their shape s
l = 0
Spherical in shape p
l = 1
Dumbbell in shape d
l = 2 f
l = 3 f d s p

22. Quantum Numbers 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 l +1 orbital's
“Street” you live on
m – Magnetic Quantum Number

23. Can only be one s orbital
Quantum Numbers
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

24. Orbital Designations 3d 3 2 -2,-1,0,+1,+2 5 10 3p 1 -1,0,+1 6 3s 2p 2sM
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

25. Orbital Rules 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

26. Reflection
How is the Bohr model different from the earlier models of the atom?
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.

27. Does Your Head Hurt Yet??
Quantum…. What….. ?

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

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

30. Hund’s Rule 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. Periodic Table & Electron Configuration

32. Periodic Table & Electron Configuration
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

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

34. Orbital Energy Diagram and Electron Configuration
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

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

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

37. Orbital Notation ___ __ __ __ __ __ __ __ __ 1s 2s 2p 3s 3p
Write the orbital notation for S
S (atomic number 16)
___ __ __ __ __ __ __ __ __
1s 2s p s p
1s22s22p63s23p4
How many unpaired electrons does sulfur have?
2 unpaired electrons!

38. Orbital Notation ___ ____ ____ ____ ____ ____ ____ ____ ____
Write the orbital notation for S
S (atomic number 16)
___ ____ ____ ____ ____ ____ ____ ____ ____
1s 2s p s p
How many unpaired electrons does sulfur have?

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

40. Valence Electrons Valence Electrons As (atomic number 33)
The electrons in the outermost energy level.
s and p electrons in last shell

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

42. Noble Gas Configuration
Example - Germanium
X X X X X X X X X X X X X
[Ar]
4s2
3d10
4p2

43. Electron Configuration
Let’s Practice
P (atomic number 15)
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

44. Electron Configuration
Noble Gas
Configuration
[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

45. Stability
Full energy level
Full sublevel
Half full sublevel

46. Exceptions 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

47. Atoms take electron configuration of the closest noble gas
Stability
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

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

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

50. Try Some P-3 (atomic number 15) Ca+2 (atomic number 20)
Zn+2 (atomic number 30)

51. X Lewis Structures 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

52. Try Some 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 • • • •

53. Try Some
Write the Electron Configuration & Lewis structures for:
Oxygen (O)
Iron (Fe)
Bromine (Br) O Fe Br

54. What Do I Need to Know? How the periodic table is arranged
Be able to identify subcategories of the periodic table
How the elements within a group are similar
How the elements within a period are similar
Be able to compare and contrast the electronegativities, ionization energies, and radii of metals and non-metals

55. Periodic Table What He Did Some Problems
Dmitri Mendeleev – Father of the Periodic Table
What He Did
Put elements in rows by increasing atomic weight
Put elements in columns by similar properties
Some Problems
He left blank spaces for what he said were undiscovered elements (he was right!)
He broke the pattern of increasing atomic weight to keep similar reacting elements together

56. Arranged by Atomic # Columns = Groups
Mosley
Arranged by Atomic # Columns = Groups
Rows = Periods

57. Periodic Table Organization
Metalloids
Metals
Non-Metals

58. Periodic Table Organization
Representative Elements
Transition Metals
Inner Transition Metals

59. Metals and Non-metals Metals Non-metals Shiny Malleable
Ductile (pulled into wires)
Conduct heat and electricity
Low specific heat
High melting points
Solids
Lose electrons
Dull
Brittle
Poor conductors
Low melting/boiling points
Varied properties
Varied phases

60. Atomic Radius Atomic Radius = ½ the distance between adjacent nuclei
Increases towards Francium

61. Ionic Radius K Cl- K+ Cl Cations Anions Positive Ion Metals
Lose electrons
Radius gets smaller!
Negative Ion
Non-metals
Gain electrons
Radius gets larger! K
Cl- K+ Cl

62. Ionization Energy Energy required to remove an electron from an atom
Why are there peaks in this trend?

63. Noble gases have the highest first Ionization Energy

64. Electronegativity Increases
Pull of electrons in a covalent bond
“Attraction” of atoms towards an electron
Fluorine is “the man”
Electronegativity Increases

65. Periodic Trends Nuclear Charge increases Atomic radius decreases
Ionization energy increases
Electronegativity increases
Ionization energy decreases
Electronegativity decreases
Orbital Size increases
Atomic radius increases

66. What Do I Need To Know? 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