1. Chapter 5 Electrons in Atoms

2. I Can………… Describe the Bohr Planetary Model of an atom
Describe light and how atoms emit light ( atomic emission spectra)
Describe ground state vs. excited state
Describe the quantum mechanics model of an atom – 3 parts

3. Niels Bohr Nobel Prize 1922 (1885 – 1962 )
Developed the model of the hydrogen atom at age 27
Nobel Prize 1922

4. Planetary Model of the Atom
Bohr’s Model
Planetary Model of the Atom

5. Bohr’s Planetary Model
TED- central mystery of quantum mechanics
Nucleus Electron
Orbit
Energy Levels

6. Bohr’s Model
Electrons move in circular orbits around the nucleus at fixed energy levels.
Electrons are never between energy levels or energy shells
An electron must have just the right amount of energy to jump from one level to another.
A quantum of energy is just the right amount of energy needed for an electron to jump levels.

7. As the orbits get further from the nucleus they have more energy...
Electrons in the outer orbits have more energy than those close to the nucleus.

8. Each different quantum jump corresponds to a different color of light energy that is released .

9. A ground state electron is at the lowest energy level possible.

10. Energy added in the form of heat, electricity, or light can move the electron up to higher energy levels which is called an excited state electron

11. As the electron falls back down to ground state it gives the absorbed energy back as a photon of light

12. As the electron falls back down to ground state it gives the absorbed energy back as a photon of light

13. Each photon of light corresponds to the energy difference between the two different energy levels.

14. 6 5 4 3 2 1

15. 6 5 4
656 nm 3 2 1

16. 6 5 4
656 nm 3 2 1

17. 6 5 4
656 nm 3
486 nm 2 1

18. 6 5 4
656 nm 3
486 nm 2 1

19. 6 5 4
434 nm
656 nm 3
486 nm 2 1

20. 6 5 4
434 nm
656 nm 3
486 nm 2 1

21. 6 5 4
434 nm
656 nm 3
410 nm
486 nm 2 1

22. The atom is not stable when the electrons are in higher energy levels .

23. The electron falls back down to its original energy level to stabilize the atom.

24. Quantum Mechanical Model of the Atom This is the modern description of an atom developed from mathematical solutions to the Schrödinger equation. 3 Parts of Quantum Mechanical Model of the Atom principal energy levels sublevels- s,p,d,f orbitals

25. Erwin Schrodinger Electron Structure: (1887 – 1961)

26. Schrodinger’s Equation
dd2ψ
+ V ψ = E ψ
8 π2mm ddxx22
These equations describe areas in atoms where electrons can exist, the electron clouds which come in many shapes and sizes

27. Heisenberg’s Uncertainty Principle
Since electrons are moving so fast, we can never know their exact position.

28. Heisenberg’s Uncertainty Principle
Due to this idea, electrons positions are described as probabilities.
Where is the most likely position of the electron?

29. Electron Clouds
If we think of fan blades in motion, we see a blurry area.
Electrons in motion also create these blurry areas, hence the name “clouds”.

30. Electron Clouds
Electrons in the “cloud” are defined by a specific location (different levels):
Principal energy level
Sublevel
Orbital

31. Principal Energy Levels
Represented by “n”
Indicates distance of the energy level from the nucleus
Have integer (whole number) values (n=1, n=2 etc.)
There are 7 principal (main) energy levels – Each period (row) on the periodic table represents an energy level

32. n = 2
n = 4 +
n = 3
n= 1

33. Sublevels, and Orbitals
Each of the 7 main energy levels (periods) is divided into a certain number of sublevels, labeled:
s, p, d, or f
The s sublevel is lowest in energy and the f sublevel is the highest
Each sublevel is divided into a certain number of orbitals.

34. The uncertain location of electrons.
What are Orbitals?
TED- 5 min
The uncertain location of electrons.
90% probability of finding an electron
Orbitals are a probability map of the electron within a sublevel- s,p,d,f

35. .l' X
2p. b ::
.,. d. d
d,. ,, : y y
.'f."

36. s sublevels and orbitals
Start at the first principal energy level/shell, n=1
S sub levels contain only 1 orbital
Each orbital can hold 2 electrons
1 orbital x 2 electrons = total of 2 electrons
Spherical in shape

37. S sublevel orbitals Every principal energy level has an s orbital
The orbitals get larger as the energy levels get higher

38. p sublevel & orbitals
p sublevels start at the second principal energy level , n=2
In each p sublevel there are 3 different p orbitals
Each orbital can hold 2 electrons
3 orbital x 2 electrons = total of 6 electrons
Shaped like dumb bells

39. d sublevel and orbitals
d sub levels start at the third principal energy level, n=3
In each d sublevel there are 5 different d orbitals
Each orbital can hold 2 electrons
For a total of 10 electrons

40. f sublevels and orbitals
Start at the fourth principal energy level, n=4
In each f sublevel there are seven different f orbitals
2 electrons per orbital (14 total electrons)

41. Energy levels & sublevels
Energy levels & sublevels
n = 1 1s
n = 2 2s 2p
n = 3 3s 3p 3d
n = 4 4s 4p 4d 4f

42. Summary
# of orbitals
# of electrons in sublevel
sublevel

43. Summary n = 1 s 1 2 n = 2 s,p 1,3 8 n = 3 s,p,d 1,3,5 18 n = 4 s,p,d,f
Principal energy
level
sublevels
# orbitals
total # of electrons
n = 1  s  1  2
n = 2
 s,p
 1,3  8
n = 3
 s,p,d
 1,3,5
 18
n = 4
 s,p,d,f
 1,3,5,7
 32

44. Writing Electron Configurations

45. Electron Configurations
We need a method for diagramming the locations of the electrons in their levels, sublevels, and orbitals.
There are rules that must be followed when filling the orbitals
The rules keep the electrons in the most stable arrangement

46. Filling Patterns
Electrons always occupy the lowest energy orbitals first. Therefore, electrons begin filling at energy level 1, in the s orbital.
This is written 1s1. This means one electron is in the s orbital of the first principal energy level.

47. LABEL THIS!
# e­ 1
sublevel
energy level = n

48. What does this tell us? 3

49. Using the Periodic Table to Write Electron Configurations
We can use our Periodic Table to help us write our electron configurations because the Periodic Table is arranged based on where electrons are located in the atom!
Electron configuration for d-block
7 min
Electron configuration for f-block
6 min

51. Hydrogen
Helium
Lithium
Note that the sum of the superscripts equals the total number of electrons in the atom.

52. Practice Write the electron configuration for Fluorine
Write the electron configuration for Phosphorus
Write the electron configuration for Iron

53. What does 2p4 mean?
There are 4 electrons in the p sublevel of the 2nd energy level!
Does 2p4 tell you how many electrons are in each of the three p orbitals?
We need more detail in our configuration!

54. We need two methods for drawing the electron arrangements
Basic Electron Configuration (as we just did)
Orbital Diagram (or notation) that shows how the electrons are placed in the individual orbitals.
Usually both are combined.

55. Example: Lithium (atomic number =3)
Electron Configuration: 1s2 2s1
Orbital diagram: ↑↓ ↑
Each arrow represents an electron.
Only two electrons can be placed in an orbital. If two electrons are present in the same orbital (on the same line), draw one up arrow, and one down arrow.

56. Pauli Exclusion Principle
Two electrons in the same orbital will have opposite spins.
Two electrons in one orbital will repel each other.
To prevent this, we say the electrons have opposite spins.
The arrows represent the opposing spins.

57. Pauli Exclusion Principle
There at most there are 2 electrons per orbital with opposite spins indicated by the up and down arrows
1s2
Summarize the Pauli Exclusion Principle on your worksheet now.

58. Hund’s Rule School Bus Rule - don’t sit next to a “Hun”
One electron in each orbital until each has one, then pair up.
When filling the p, d, or f sublevels, spread out the electrons into the orbitals as much as possible.
Ex) 2p3 is shown:   
instead of: ↑↓ ↑  __

59. 4
px py pz

60. Hund’s Rule
When one set of orbitals are filled, move to the next sublevel
Ex/ Oxygen (atomic number = 8)
1s s p4
__ __ __ __ __

61. Hund’s Rule 1s2 ↑↓ 2s2 ↑↓ 2p3 ↑ ↑ ↑ incorrect: 1s2 ↑↓ 2s2 ↑↓ 2p3 ↑↓ ↑
Example: Nitrogen (atomic number = 7)
correct:
1s2 ↑↓ 2s2 ↑↓ 2p3 ↑ ↑ ↑
incorrect:
1s2 ↑↓ 2s2 ↑↓ 2p3 ↑↓ ↑
What’s wrong?
Summarize the Hund’s Rule on your worksheet now.

62. Aufbau Principle Filling Chart
According to the Aufbau principle, electrons occupy the orbitals of lowest energy first. In the Aufbau diagram, each box represents an atomic orbital.
Increasing energy 6s 5s 4s 3s 2s 1s 6p 5p 5d 4p 4d 4f 3p 3d 2p

63. Aufbau filling chart sublevels fill from base of arrow to point3d 4d 5d 6d 7d 4f 5f
The order is 1s,2s,2p,3s,3p,4s,3d,4p,5s,4d,5p,6s,4f,5d,6p,7s,5f,6d,7p,7d

64. NOBLE GAS SHORT­ CUT METHOD
All electron configurations can be written in terms of a noble gas
Sodium - 1s2 2s2 2p6 3s 1
Can be written as: [Ne] 3s 1
Find the element then get the noble gas from the period ABOVE it

65. NOBLE GAS METHOD
Long form Noble Gas form Mg Al Ge Sb

66. Magnesium
• 1s2 2s2 2p6 3s2
• [Ne] 3s2
Aluminum
• 1s2 2s2 2p6 3s2 3p1
• [Ne] 3s2 3p1
Germanium
• 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p2
• [Ar] 4s2 3d10 4p2
Antimony
• 1s2 2s2 2p6 3s2 4s2 3d10 4p6 5s2 4d10 5p3
[Kr] 5s2 4d10 5p3