1. Electrons in Atoms Joke of the day:
Why did Erwin Schrödinger, Neils Bohr and Werner Heisenberg work in very small garages?

2. Dalton’s Atomic Model

3. Plum Pudding Model (Thomson)

4. The modern view of the atom was developed by Ernest Rutherford (1871-1937).
Screen 2.9

6. Niels Bohr’s Model (1913)
Electrons orbit the nucleus in circular paths of fixed energy (energy levels).

7. Neils Bohr and the Atom
                                                                           
Proposed electrons are found in concentric orbits around the nucleus call energy levels
Originated the idea of quantum energy

8. Atomic Line Spectra and Niels Bohr
Bohr’s theory was a great accomplishment.
Rec’d Nobel Prize, 1922
Problems with theory —
theory only successful for H.
introduced quantum idea artificially.
So, we go on to QUANTUM or WAVE MECHANICS
Niels Bohr
( )

9. Electron Motion Around Atom Shown as a de Broglie Wave

10. Electrons in Atoms Joke of the day:
Why did Erwin Schrödinger, Neils Bohr and Werner Heisenberg work in very small garages?

11. Electrons in Atoms Joke of the day:
Why did Erwin Schrödinger, Neils Bohr and Werner Heisenberg work in very small garages?
They were quantum mechanics.

12. Quantum Numbers
There are four quantum numbers that describe the properties of an electron and the “orbital” that it occupies within an atom.
Apple

13. What are orbitals?
The main division of energy for an electron within an atom is called the energy level or shell.
The energy level (shell) is subdivided into distinct areas called sublevels (subshells).
The sublevels are divided into orbitals. The orbitals are where the electrons are found.
Each orbital can hold no more than a pair of electrons.

14. Four Quantum Numbers describe the following:
Energy level (distance from the nucleus)
Also Called a “Shell”
Sublevel (shape of the orbital)
Also Called a “Sub-Shell”
Orbital (3-dimensional positioning)
Direction of the electron’s spin within the orbital

15. The First Quantum Number The Principal Quantum Number, “Shell”
Abbreviated as “n”
n = 1,2,3...
Represents the distance of an electron from the nucleus & the main energy level of the electron.

16. What else does the Principal Quantum Number (n) tell us?
n = number of sublevels in the energy level.

17. What else does the Principal Quantum Number (n) tell us?
n2 = number of orbitals in the energy level.

18. What else does the Principal Quantum Number (n) tell us?
2n2 = maximum number of electrons possible in the energy level.

19. Think! In the n=3 shell How many sub-shells are there?
How many orbitals?

20. Think! In the n=3 shell How many sub-shells are there? 3
How many orbitals? n2=9

21. Second Quantum Number L
Sub-level

22. Types of sublevels The four common sublevels in atoms.
The sublevels are designated s, p, d, f.
When l=0, “s”
l=1, “p”
l=2, “d”
l=3, “f”

23. Types of sublevels The four common sublevels in atoms.
The sublevels are designated s, p, d, f.

24. s Orbitals
All s orbitals are spherical in shape.

25. 1s Orbital

26. 2s Orbital

27. 3s Orbital

28. “s” sublevel and “p” sublevel
s orbitals have a spherical shape
p orbitals have a dumbbell shape a

29. Think!
For the 4d sub-shell
What is the n-value?
What is the l-value?

30. Think! For the 4d sub-shell What is the n-value? 4
What is the l-value? 2

31. The Third Quantum Number
The third quantum number identifies the orbital that the electron is in.
An orbital is an area within a sublevel that can hold up to two electrons.
Value of the 3rd quantum number tells us the orientation of the orbital in that sub-shell

32. P (n=2, l=1) orbital shape

33. d orbital shapes
(n=3, l=2)

34. Electron Spin 4th Quantum Number
Opposite spins produce opposite magnetic fields.
+1/2
-1/2

35. Orbitals Pauli Exclusion Principle
No more than 2 e- assigned to an orbital
Orbitals grouped in s, p, d (and f) sublevels
s sublevels
p sublevels
d sublevels

36. ATOMIC ELECTRON CONFIGURATIONS AND PERIODICITY
To play the movies and simulations included, view the presentation in Slide Show Mode.

37. Arrangement of Electrons in Atoms
Electrons in atoms are arranged as
Principle Energy Levels (n)
SUBLEVELS (l)
ORBITALS (ml)

38. Principle Energy Levels

39. Assigning Electrons to Atoms
Electrons generally assigned to orbitals of successively higher energy. Aufbau Principle

40. Sublevel Filling of Electrons
Electrons fill orbitals from the bottom up: Aufbau Principle

41. 1 3 5 2 6 10 s orbitals d orbitals p orbitals s orbitals p orbitals
No.
orbs. 1 3 5
No. e- 2 6 10

42. Summary of 4 Quantum Numbers
n = Energy level (distance from the nucleus)
-Also Called a “Shell”
l= Sublevel (shape of the orbital)
-Also Called a “Sub-Shell”
m=Orbital (3-dimensional positioning)
Spin = Direction of the electron’s spin within the orbital

43. End for today!

44. Writing Atomic Electron Configurations
Two ways of writing configs. One is called the spdf notation. 1 s
value of n
value of l
no. of
electrons
spdf notation
for H, atomic number = 1

45. Writing Atomic Electron Configurations
Two ways of writing configs. Other is called the orbital box notation.

46. Electron Filling Order Figure 8.5

47. See “Toolbox” for Electron Configuration tool.

48. Electron Configurations and the Periodic Table
Figure 8.7

49. Lithium
Group 1A
Atomic number = 3
1s22s1 ---> 3 total electrons

50. Beryllium
Group 2A
Atomic number = 4
1s22s2 ---> 4 total electrons

51. Boron
Group 3A
Atomic number = 5
1s2 2s2 2p1 --->
5 total electrons

52. Carbon Group 4A Atomic number = 6 1s2 2s2 2p2 --->
6 total electrons
Here we see for the first time HUND’S RULE. When placing electrons in a set of orbitals having the same energy, we place them singly as long as possible.

53. Nitrogen Group 5A Atomic number = 7 1s2 2s2 2p3 --->
7 total electrons

54. Oxygen Group 6A Atomic number = 8 1s2 2s2 2p4 --->
8 total electrons

55. Fluorine Group 7A Atomic number = 9 1s2 2s2 2p5 --->
9 total electrons

56. Neon Group 8A Atomic number = 10 1s2 2s2 2p6 --->
10 total electrons
Note that we have reached the end of the 2nd period, and the 2nd shell is full!

57. Sodium All Group 1A elements have [core]ns1 configurations. Group 1A
Atomic number = 11
1s2 2s2 2p6 3s1 or
“neon core” + 3s1
[Ne] 3s1 (uses rare gas notation)
Note that we have begun a new period.
All Group 1A elements have [core]ns1 configurations.

58. Electron Configurations of p-Block Elements

59. Aluminum Group 3A Atomic number = 13 1s2 2s2 2p6 3s2 3p1 [Ne] 3s2 3p1
All Group 3A elements have [core] ns2 np1 configurations where n is the period number.

60. Phosphorus Group 5A Atomic number = 15 1s2 2s2 2p6 3s2 3p3
[Ne] 3s2 3p3
All Group 5A elements have [core ] ns2 np3 configurations where n is the period number.

61. Calcium Group 2A Atomic number = 20 1s2 2s2 2p6 3s2 3p6 4s2 [Ar] 4s2
All Group 2A elements have [core]ns2 configurations where n is the period number.

62. Electron Configurations and the Periodic Table

63. Transition Metals Table 8.4
All 4th period elements have the configuration [argon] nsx (n - 1)dy and so are “d-block” elements.
Chromium
Iron
Copper

64. Transition Element Configurations
3d orbitals used for Sc-Zn (Table 8.4)

65. Lanthanides and Actinides
All these elements have the configuration [core] nsx (n - 1)dy (n - 2)fz and so are “f-block” elements.
Cerium
[Xe] 6s2 5d1 4f1
Uranium
[Rn] 7s2 6d1 5f3

66. Lanthanide Element Configurations
4f orbitals used for Ce - Lu and 5f for Th - Lr (Table 8.2)

67. Arrangement of Electrons in Atoms
Each orbital can be assigned no more than 2 electrons! Pauli Exclusion Principle
This is tied to the existence of a 4th quantum number, the electron spin quantum number, ms.

68. Electron Spin Quantum Number, ms
Can be proved experimentally that electron
has a spin. Hunds Rule

69. Electron Spin Quantum Number
Diamagnetic: NOT attracted to a magnetic field
Paramagnetic: substance is attracted to a magnetic field. Substance has unpaired electrons.

70. Assigning Electrons to Subshells
In H atom all subshells of same n have same energy.
In many-electron atom:
a) subshells increase in energy as value of n + l increases.
b) for subshells of same n + l, subshell with lower n is lower in energy.

71. 3px Orbital
2px Orbital

72. d Orbitals
s orbitals have no planar node (l = 0) and so are spherical.
p orbitals have l = 1, and have 1 planar node,
and so are “dumbbell” shaped.
This means d orbitals (with l = 2) have
2 planar nodes
See Figure 7.16

73. 3dxy Orbital

74. 3dxz Orbital

75. 3dyz Orbital

76. 3dx2- y2 Orbital

77. 3dz2 Orbital

78. Orbital Filling: The Aufbau Principle & Hund’s Rule
Aufbau Principle: Lower energy orbitals fill first.
Hund’s Rule:
Degenerate orbitals (those of the same energy) are filled with electrons until all are half filled before pairing up of electrons can occur.
Pauli exclusion principle:
Individual orbitals only hold two electrons, and each should have different spin.
“s” orbitals can hold 2 electrons
“p” orbitals hold up to 6 electrons
“d” orbitals can hold up to 10 electrons

79. Orbital Filling: The Aufbau Principle & Hund’s Rule
Aufbau Principle: Lower energy orbitals fill first.
Hund’s Rule:
Degenerate orbitals (those of the same energy) are filled with electrons until all are half filled before pairing up of electrons can occur.
Pauli exclusion principle:
Individual orbitals only hold two electrons, and each should have different spin.
“s” orbitals can hold 2 electrons
“p” orbitals hold up to 6 electrons
“d” orbitals can hold up to 10 electrons

80. Pauli Exclusion Principle
No two electrons in the same atom can have the same set of 4 quantum numbers.
That is, each electron has a unique address.