1. Electrons in Atoms Bohr Orbits vs. Electron-cloud Orbitals
Electron Configurations and Orbital Diagrams
Periodic Patterns
1s2 2s2 2p3

2. Bohr Model of the Atom Electrons are NOT randomly arranged
Energy of electrons is quantized
Vocabulary: quantized

3. The Bohr Diagram
The simplicity of the Bohr Model of the atom makes it useful for representing atoms of different elements.
Bohr Diagrams place electrons on concentric rings surrounding the nucleus.

4. Bohr Diagrams Show the important features of an atom number of protons
number of electrons
number of neutrons
arrangement of electrons

5. Number of Electrons Allowed
The Bohr Diagram
The following table shows how many electrons can occupy a given energy level in a Bohr Diagram.
Shell Number
Energy Level
Number of Electrons Allowed 1 2 8 3 18 4 32
Vocabulary: Shell number

6. The Bohr Diagram
The Bohr Diagram for the first 5 elements:

7. Learning Check
Draw the Bohr Diagram for Phosphorous (P)
remaining e-

8. Valence Electrons
Valence electrons: the electrons in the outermost occupied shell (orbital)
Valence electrons are in the highest energy shell (orbital)
Valence electrons are the electrons that will participate in a chemical reaction
The number of valence electrons largely influences the properties of an element
Vocabulary: valence electron

9. Valence Electrons
The valence electrons in our Bohr diagrams are the electrons on the outermost ring of the diagram, in this example the 4 electrons on the third ring.
14 P
14 N

10. Learning Check
How many valence electrons does Cl have?

11. Learning Check
Cl has 7 valence electrons. (Count the e- on the outermost ring)

12. Ground State Lowest energy state of an electron
Normal resting state most stable
The ground state is what you are familiar with:
Vocabulary: ground state, excited state

13. Excited States
Electrons have multiple possible excited states. Each one can produce a different line in the emission spectrum.

14. Excited States Excited states are NOT stable.
Excited states are higher energy.

15. Ground State vs. Excited State
An electron in the excited state will ALWAYS return to the ground state.
Examples: firecrackers and glowing burner

16. Emission Spectra
While investigating atoms and their properties, scientists discovered that atoms of different elements emit light in distinct patterns.
These patterns are called Line Emission Spectra, and are unique for each element, based on its electronic structure.

17. Each element has a "fingerprint" in terms of its line emission spectrum, as illustrated by the examples below.
Line spectrum for hydrogen.
Line spectrum for helium.
Line spectrum for neon.

18. Bohr Model vs. Electron Cloud
Both models agree that the arrangement of the electrons matters
Both models have quantized energy levels for the electrons
The Bohr model is simpler to understand
Why do we need the Electron cloud model?

19. Particle Behavior (Expectation before the Bohr model)
Electrons are subatomic particles of atoms.
Electrons are matter.
Matter is governed by certain laws of physics.
You can predict the trajectory of a particle. It does not bend around corners.
Matter has phases (solid, liquid, gas) with different amounts of kinetic energy.

20. Emission of Light (Caused change in thinking)
Electrons emit light when returning from an excited state to the ground state.
Emission of light indicates that electrons have wavelike properties.
Different colors of light correspond to different wavelengths and energies.
Line spectra exist for the different elements

21. Waves
Electromagnetic radiation propagates through space as a wave moving at the speed of light.
c = 
c = speed of light, a constant (3.00 x 108 m/s)
 = frequency, in units of hertz (hz, sec-1)
 = wavelength, in meters

22. E = h Long Wavelength = Low Frequency Low ENERGY Short Wavelength =
High Frequency
High ENERGY
E = h

23. The Electromagnetic Spectrum

24. Wave Behavior Light is a wave. Waves are NOT matter.
Light CAN bend around corners.
Light has different properties from matter.
demonstration

25. of light with definite wavelengths.
Electron transitions from excited state to ground state involve jumps of quantized amounts of energy.
This produces bands
of light with definite wavelengths.
What happens if the atom gets less than the needed amount of energy to reach the n = 2 level?

26. Proof of the Wave Nature of Electrons: Emission of Light?
Photons of light are emitted as an electron drops back to its ground state after being excited.

27. Flame Tests
Atoms become excited when they are heated in a flame. Their electrons move from their ground state to higher energy levels. As they return to their ground state they emit photons of very specific energy. The energy of the emitted photons corresponds to particular wavelengths of light, producing different colors of light.

28. This picture illustrates the distinctive colors produced by burning particular elements.

30. The Dilemma of the Atom
The Bohr Model predicted the observed line spectrum for H.
However, the Bohr Model does NOT correctly predict the line spectrum of any atom besides H.
Why did the Bohr Model fail for atoms with more than one electron?

31. Wave-Particle Duality
JJ Thomson won the Nobel prize (1906) for describing the electron as a particle.
His son, George Thomson won the Nobel prize (1937) for describing the wave-like nature of the electron.
The electron is a particle!
The electron is an energy wave!

32. Wave-Particle Duality
Electrons have BOTH particle and wave behavior.
This is known as the wave-particle duality.
All matter has both wave and particle properties, but we can usually only observe one or the other.
Only the Electron Cloud model of the atom uses the dual nature of the electron.

33. Quantum Mechanics Energy Levels
Quantum mechanics uses a principal quantum number, “n”. It represents the energy level just like in the Bohr diagram.
n = 1 describes the first energy level
n = 2 describes the second energy level
Each energy level is represented on the periodic table, and approximately correlates to a row.
Red n = 1
Orange n = 2
Yellow n = 3
Green n = 4
Blue n = 5
Indigo n = 6
Violet n = 7

34. Sublevels = Specific Atomic Orbitals
Energy levels have sublevels which describe the specific types of atomic orbitals for that level.
n = 1 has 1 subshell (the “s” orbital)
n = 2 has 2 subshells (“s” and “p” orbitals)
n = 3 has 3 subshells (“s”, “p” and “d”)
n = 4 has 4 subshells (“s”, “p”, “d” and “f”)
Vocabulary: subshell or sublevel

35. Energy Level and Orbital Size
Representation of how the energy level, n (shell), affects the size of an orbital.

36. Shells, Subshells, and Orbitals
Shell = energy level n = 1, n = 2, etc. Subshell = type of orbital s, p, d, f (ns, np) Orbital = exact orbital 2px

37. Shells, Subshells, and Orbitals
Shells define the energy level Subshells define the sublevels within a shell Orbitals exist for each sublevel
Shells
Orbitals
Subshells
Orbitals

38. Subshells of the energy levels
n = 1 1 subshell (s) red elements
n = 2 2 subshells (s, p) orange elements
n = 3 3 subshells (s, p, d) yellow elements 1
Red n = 1
Orange n = 2
Yellow n = 3
Green n = 4
Blue n = 5
Indigo n = 6
Violet n = 7 2 3 4 3 4

39. QM Atomic Orbitals
s 1 spherical orbital per energy level p 3 dumbbell shaped orbitals per energy level d 5 multi-lobed orbitals per energy level f 7 multi-lobed orbitals per energy level Each type of orbital corresponds to a block on the periodic table. Each atomic orbital can hold 2 e- (or fewer).

40. QM Atomic Orbitals

41. QM Atomic Orbitals
The quantum mechanical orbitals are identifiable by their shape. You should recognize each of the following orbitals.

42. Orbitals s p d
In the s block of the periodic table, electrons fill s orbitals.
In the p block, the s orbitals are full. New electrons begin filling the p orbitals.
In the d block, the s and p orbitals are full. New electrons are added to the d orbitals.
What about the f block?

44. Periodic Table & Electron Configuration

45. Orbital Diagrams and Electron Configurations
Orbital diagrams are a simple way of representing the complete electron arrangement in an atom. It contains more information than an electron configuration. An electron configuration is a simple way to give the basic electron arrangement in an atom, and emphasizes the valence electrons.
1s2 2s2 2p3

46. Capacity of the Levels Energy Level Sub-levels Total Orbitals
Total Electrons
Total Electrons per Energy 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

47. Rules for Electron Configurations
The Aufbau Principle states that each electron occupies the lowest energy orbital available.
Note the axis label: E

48. Orbital Filling Sequence
Note that 4d orbital is higher E than 5s orbital
Note the axis label: E
Shells (energy levels)
ALWAYS the lowest energy orbital

49. Aufbau Principle
Electrons fill orbitals starting at the lowest available energy level before filling higher levels 
1s is ALWAYS first

50. Aufbau Diagram
Start with 1s and follow the arrows.

51. Aufbau Diagram
Energy Levels: 1, 2, 3, 4… Sublevels: s, p, d, f Orbitals: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p…

52. Learning Check
Use the Aufbau Principle to order the orbitals occupied by electrons for Co.

53. Learning Check Co has 27 electrons Every orbital can hold 2 e-
Orbitals are filled beginning with the lowest energy (1s)
Using the Aufbau Diagram:
1s, 2s, 2p, 3s, 3p, 4s, 3d
27/2 = 13.5 orbitals
= 15 orbitals
Co will end in the 3d orbitals

54. Learning Check
Co:

55. 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

56. Pauli Exclusion Principle
The Pauli Exclusion Principle states that a maximum of two electrons can occupy a single orbital, but only if the paired electrons have opposite spins.
What orbitals are shown?
How many electrons are in each orbital? What is the valence energy level for Ne?

57. Pauli Exclusion Principle
An orbital can hold 0, 1, or 2 electrons
Full orbital contains 2 electrons
Half full orbital contains 1 electron
If there are two electrons paired in the orbital, they must have opposite spins.

58. Pauli Exclusion Principle
When we draw electrons, we use up and down arrows. If electrons are paired in a box, one arrow is up and the second must be down.
Incorrect; the paired electrons must have opposite spin, and unpaired electrons should have the same spin
Correct; the paired electrons have opposite spin, and the unpaired electrons have the same spin 

59. Learning Check
1. Write the correct orbital diagram for Si. 2. What is wrong with the following orbital diagram? Correct the orbital diagram and identify the element. 3p 3s 2p 2s 1s

60. Learning Check
1. Write the correct orbital diagram for Si. 2. What is wrong with the following orbital diagram? Correct the orbital diagram and identify the element.
1s s p s p 3p 3s Na 2p 2s 1s

61. No more than 2 Electrons in Any Orbital…ever!
The Pauli Exclusion Principle states that an atomic orbital may have up to 2 electrons and then it is full.
The spins have to be paired.
We usually represent this with an up arrow and a down arrow.

62. Hund’s Rule
Hund’s Rule states that single electrons with the same spin must occupy each degenerate orbital before electrons with opposite spins pair to occupy the same energy level orbitals.
Each 3p orbital has the same energy. They are degenerate.
Vocabulary: degenerate orbital

63. Hund’s Rule
When filling sublevels, electrons are placed individually in degenerate orbitals before they are paired up.
Think of the charge of an electron. Like charges repel. Thus, when a degenerate orbital is available, the electrons prefer to be unpaired.
degenerate orbitals

64. Hund’s Rule
So when working with the p sublevel, electrons fill like this....up, up, up...down, down, down 
atom
orbital box diagram B 1s 2s 2p C N O F

65. Learning Check Which of the following options is correct?
Apply Hund’s Rule (and what else?)

66. Hund’s Rule http://intro. chem. okstate
Hund’s Rule requires that you half fill degenerate orbitals before pairing. Begin pairing the electrons only when all orbitals are halfway full.
Degenerate orbitals = orbitals that have the same energy
The 3 p orbitals on each level are degenerate 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.

67. Summary: Rules for Filling Orbitals
1. Lower-energy orbitals fill first (Aufbau Principle). 2. An orbital can hold only 2 electrons with opposite spins (Pauli Exclusion Principle). 3. If degenerate orbitals are available, electrons remain unpaired with like spin in each orbital until all are half-full (Hund's Rule).

68. Chapter Vocabulary Week 3 (new) Week 2 Week 1
Aufbau Principle Pauli Exclusion Principle Hund’s Rule Degenerate Shell Subshell (sublevel) Orbital Ground state, Excited state Quantized Valence electron Shell number
Week 3 (new)
Week 2
Week 1

69. Electron Configurations and Orbital Diagrams
Orbital diagrams and electron configurations are closely related. The orbital diagram for N is The electron configuration for N is 1s22s22p3 Note that the number of electrons in the degenerative boxes adds up to the number in the orbital subshell of the electron configuration.

70. 4d 5s 4p 3d 4s s
1 2 p 1 3p 2 d 3 3s 4 5 2p 6 7 2s H f
1s1
That’s right: it goes in the 1s sublevel. And its e- config is 1s1. Notice in the table above where H is – in the area designated as 1s. So where does the next electron go? 1s +

71. 4d 5s 4p 3d 4s s
1 2 p 1 3p 2 d 3 3s 4 5 2p 6 7 2s He f
1s2
If you were thinking it went in the 2s, then you forgot that each orbital can hold up to two electrons. Note how He is right here in the area designated as 1s2 and so its e- config is 1s2. 1s +

72. 4d 5s 4p 3d 4s s
1 2 p 1 3p 2 d 3 3s 4 5 2p 6 7 2s Li f
1s2 2s1
Now that the 1s is filled, the next electron goes in the next sublevel – the 2s. Again note how Li is in 2s1. 1s +
Its full e- config is 1s2 2s1. What is the e- config of Be?

73. 4d 5s 4p 3d 4s s
1 2 p 1 3p 2 d 3 3s 4 5 2p 6 7 2s B f
1s2 2s2 2p1
Is this what you were thinking? Notice how B is in the 2p1 spot. 1s +
So its full e- config is 1s2 2s2 2p1. What’s next?

74. 4d 5s 4p 3d 4s s
1 2 p 1 3p 2 d 3 3s 4 5 2p 6 7 2s C f
1s2 2s2 2p2
Is this what you were thinking? Notice how C is in the 2p2 spot. So its e- config is 1s2 2s2 2p2 1s +
Notice also how we fill a sublevel…Hund’s Rule

75. 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 the numbers in the electron configuration add up to the atomic number for that element. Ex: for Ne (Z=10), = 10

76. Periodic Pattern
Look again at the electron configurations of hydrogen, lithium, sodium and potassium.
Do you see any similarities?
Since H and Li and Na and K are all in Group 1A, they all have the same ending. (s1)
Element
Configuration
H Z=1
1s1
Li Z=3
1s22s1
Na Z=11
1s22s22p63s1
K Z=19
1s22s22p63s23p64s1

77. Periodic Pattern
This similar electron configuration for members of a group (a column) causes them to behave the same chemically.
It’s for that reason the periodic table is arranged as it is.
Each group will have the same ending configuration, in this case the electron configuration always ends in s1.

78. Periodic Table & Electron Configuration

79. Noble Gas Electron Configurations
Simplified method of writing an electron configuration abbreviated form 1. Use noble gas symbol from last full shell 2. Place noble gas symbol in square brackets Represents the core (inner) electrons 3. Write the remaining valence electrons after the bracketed noble gas symbol

80. Noble Gas Electron Configurations
Ca electron configuration: 1s22s22p63s23p64s2 Ca abbreviated electron configuration (Noble gas electron configuration): [Ar]4s2 [Ar] = the 18 core electrons (1s22s22p63s23p6) 4s2 represents the 2 valence electrons (in n = 4 shell)

81. Learning Check
Write the abbreviated electron configuration of Al.

82. Learning Check
Abbreviated electron configuration of Al: Al is in the n = 3 shell and has 13 electrons Last full shell is n = 2 shell Noble gas element for n = 2 shell is Ne Ne has 10 electrons; we need to add 3 electrons [Ne]3s23p1

83. Learning Check
Write the abbreviated electron configuration of Sn. Sn is atomic number 50

84. Learning Check
Abbreviated electron configuration of Sn: Sn has 50 electrons, and is in the n = 5 shell Use the noble gas symbol from n = 4 Kr [Kr]5s24d105p2

85. Lewis Dot Diagrams Use only the valence electrons
Gives fundamental information about chemical behavior of the element
Simple representation that shows periodic trends

86. Lewis Dot Diagrams Steps: Write the chemical symbol for element
Count the valence electrons
Use one dot for each valence electron
Dots should be placed following the basic rules governing orbital diagrams: do not pair if not needed and place as far away as possible (electrons repel each other)

87. number of valence electrons
Lewis Dot Diagrams
Examples:
Column IA
Column IIA
Column IIIA
Column IVA
Column VA
Column VIA
Column VIIA
Column VIIIA
Element Li Be B C N O F Ne
Lewis Structure
number of valence electrons 1 2 3 4 5 6 7 8

88. Learning Check
Write the Lewis dot diagrams for Si, S, and Ar. Si = ? Valence electrons S = ? Valence electrons Ar = ? Valence electrons

89. Learning Check
Write the Lewis dot diagrams for Si, S, and Ar. 4 valence 6 valence 8 valence electrons electrons electrons