1. Chapter 13: Electrons in the Atom
College Prep Chemistry
Orbital Interactive

2. Evolution of Atomic Models
1. John Dalton
– indestructible mass
- no subatomic particles

3. Evolution of Atomic Models
2. J.J. Thomson – plum-pudding model
Discovered electrons
electrons stuck into a lump of positively charged material
“mint chocolate chip ice cream model”

4. Evolution of Atomic Models
3. Ernest Rutherford
nucleus of the atom is positively charged
Evidence – gold foil experiment

5. Evolution of Atomic Models
4. Niels Bohr
electrons travel in definite orbitals around the nucleus
Energy level = the region around the nucleus where the electron is likely to be moving
Fixed energy levels analogous to the rungs of a ladder
The lowest rung of the ladder corresponds to the lowest energy level. A person can climb up or down a ladder by going from rung to rung. Similarly, an electron can jump from one energy level to another. A person on the ladder cannot stand between the rungs. Similarly, an electron cannot exist between energy levels. To move from one rung to another, a person climbing the ladder must move just the right distance. To move from one energy level to another, an electron must gain or lose just the right amount of energy.

6. Evolution of Atomic Models
Quantum Mechanical Model
Describes the probability of finding an electron in a region of space around the nucleus
It is impossible to know the exact position and momentum of an electron at the same time
Based on probability rather than certainty
Instead of traveling in defined orbits as Bohr proposed, electrons actually travel in diffuse clouds around the nucleus
When we say orbital, this is the image we should picture in our minds. Electrons have fuzzy positions but definite energy levels.

7. Principal Energy Levels (n)
Principal (main) energy levels are assigned numbers according to their energy
n=1, 2, 3, 4 …
Generally, energy increases with increasing n
Distance of the electron from the nucleus increases with increasing n

8. Sublevels
For each principal energy level, there are one or more sublevels
s, p, d, f
# of sublevels = the principal energy level (n)
For example,
1st principal energy level has 1 sublevel (s)
2nd principal energy level has 2 sublevel (s,p)
3rd principal energy level has 3 sublevels (s,p,d)

9. Sublevels and Orbitals
Each sublevel has a specific shape
Each sublevel houses a specific # of electron orbitals (“where the electrons live”)
Each orbital can contain 2 electrons

10. Sublevels and Orbitals
s-sublevel:
Spherical in shape
One orbital
Contains 2e- maximum (2e- per each s orbital)

11. Sublevels and Orbitals
p-sublevel:
dumbbell in shape
3 orbitals
Holds 6 e- maximum (2e- per each p orbital)

12. Sublevels and Orbitals
d-sublevel:
Double dumbbell in shape
5 orbitals
Holds 10 e- maximum (2e- per each d orbital)

13. Sublevels and Orbitals
f-sublevel:
complex in shape
7 types of f-orbital
Holds 14 e- maximum in each orbital (2e- per each f orbital)

14. # of Orbitals in each Sublevel
Atomic Orbital Chart
Energy Level
# of
Sublevels
 Sublevels
Total # of
Electrons
# of Orbitals in each Sublevel 1 3 5 7
n=1 1 s 2 p 2 s
n=2 8
n=3 3 p d 18 s
n=4 4 s p d f 32

15. Electron Arrangement in Atoms
Electron Configuration: the ways in which electrons are arranged around the nuclei of atoms
Gives information about principal energy levels, sublevels, orbitals

16. Three Rules determine electron configurations
The Aufbau Principal
Hund’s Rule
Pauli Exclusion Principle

17. Rules for Electron Configurations
Aufbau principle:
Electrons enter orbitals of the lowest energy first
The s sublevel is always the lowest in energy

18. Rules for Electron Configurations
Aufbau principle: Draw your own Aufbau filling diagram
7s 7p 7d 7f
6s 6p 6d 6f
5s 5p 5d 5f
4s 4p 4d 4f (the #4 spelled out has an f)
3s 3p 3d
2s 2p 1s
7s 7p 7d
6s 6p 6d
5s 5p 5d
4s 4p 4d
3s 3p 3d (we see in 3 dimensions!)
2s 2p 1s 7s 6s 5s 4s 3s 2s
1s (the #7 spelled out has an s)
7s 7p
6s 6p
5s 5p
4s 4p
3s 3p
2s 2p (two “peas” in a pod) 1s

19. Rules for Electron Configurations
Aufbau principle: Draw your own Aufbau filling diagram
7s 7p 7d 7f
6s 6p 6d 6f
5s 5p 5d 5f
4s 4p 4d 4f
3s 3p 3d
2s 2p 1s

20. ↑ ↑ ↑_ 2. Hund’s Rule: (“hogs don’t like each other”)
Rules for Electron Configurations
2. Hund’s Rule: (“hogs don’t like each other”)
Every orbital in a sublevel is singly occupied before any orbital is doubly occupied
All of the electrons in singly occupied orbitals have the same spin
↑ ↑ ↑_

21. 3. Pauli Exclusion Principle:
Rules for Electron Configurations
3. Pauli Exclusion Principle:
an atomic orbital may have a maximum of two electrons
Two electrons that occupy the same orbital must have opposite spins
designated with 

22. Orbital Notation Examples
Li ↑↓
1s 2s
B ↑↓ ↑↓ ↑
1s 2s 2p
C ↑↓ ↑↓ ↑ ↑
Hund’s Rule

23. Orbital Notation practice
Elements #1-20
Use the Aufbau Diagram Provided
Remember all 3 rules!

25. Electron Configurations
A shorthand way of identifying the location of electrons
C ↑↓ ↑↓ ↑ ↑
1s 2s 2p
C 1s2 2s2 2p2

26. Electron Configurations

27. Rearranging the e- configurations
Scientists rearrange the configurations so that all similar energy levels stay together
#21 = Sc (fill with Aufbau first!)
1s22s22p63s23p64s23d1
Then Switch!
1s22s22p63s23p63d14s2

28. Noble Gas electron configurations
Noble gas configurations – look on the periodic table!
#11 = Na
The noble gas that precedes Na is Ne 1s22s22p6
So instead of 1s22s22p63s1 use [Ne]3s1

29. Orbital blocks on the Periodic Table

30. Exceptions to electron configurations
#24 = Cr
1s22s22p63s23p64s23d4
↑↓ ↑ ↑ ↑ ↑ __
Takes less energy to half fill all orbitals
1s22s22p63s23p64s13d5
↑_ ↑ ↑ ↑ ↑ ↑_

31. Exceptions to the electron configurations
Also true for the rest of column 6 & 11
#29 = Cu
Following the rules
1s22s22p63s23p64s23d9
↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑_
Actual configuration
1s22s22p63s23p64s13d10
↑_ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓

32. Physics and the Quantum Mechanical Model… Electrons and Light
Section 13.3
We are going to backtrack a bit today to delve further into the development of the quantum mechanical model. The model actually was born from the study of light.
Physics and the Quantum Mechanical Model…
Electrons and Light

33. Back to Bohr Bohr’s model is based on atomic emission spectra
Atoms only give off light of certain colors (wavelengths)

34. Wave model of light
Light is a wave with a frequency, speed and wavelength
Emission of light is related to the behavior of electrons in an atom

35. Parts of a Wave (1 of 2) Origin – center line
Wavelength (l) – distance from crest to crest
Amplitude (A) – distance from origin to crest
amplitude
origin

36. Parts of a Wave (2 of 2) Frequency (f) - # of waves per second (s-1)
SI Unit (Hertz)
1 Hz = 1 s-1
inversely related to wavelength

37. Electromagnetic Radiation (EMR)
a form of energy that exhibits wave-like behavior as it travels through space
Includes radio waves, microwaves, infrared waves, Visible light, ultraviolet waves, X-rays and gamma rays
All waves travel at the speed of light
3.0 X 108 m/s

38. Visible Light A prism separates sunlight into a spectrum of colors
Sunlight consists of a continuous range of colors
Each color has a specific wavelength and f
Red light = lowest f, longest wavelength

39. Electromagnetic Spectrum

40. c = l x f Wave model of light Speed m/s Wavelength m Frequency
Hz (hertz) 1/s
c = Speed of light = 3.00 x 108 m/s

41. Calculating Frequency
Determine the frequency of light with a wavelength of 500 nm?
c = l x f
l = 500 nm =5 x 10-7 m
c = 3.00 x 108 m/s
3.00 x 108 m/s = (5 x 10-7 m) f
f = 6.00 x 1014 Hz

42. Calculating Wavelength
What is the wavelength of the yellow light emitted by a sodium lamp if the frequency of the radiation is 5.10 x 1014 s-1?

43. Sample Problem Answer Given: f = 5.10 x 1014 s-1 c = 3.00 x 108 m/s
Unknown: l
Parent Equation: c = f x l
Answer: x 10-7 m = 588 nm

44. Your turn #1
What is the wavelength of radiation with a frequency of is 1.50 x 1013 s-1?
2.00 x m
Active Inspire

45. Your turn #2
What is the frequency of radiation with a wavelength of 5.00 x m?
6.00 x s -1
Active Inspire

46. Physics and the Quantum Mechanical Model
Max Planck ( )
Discovered a direct relationship between frequency and energy
Higher frequency = higher energy
E = h x f
h = Planck’s Constant = 6.63 x J·s

47. Sample Problem
What is the energy of a photon with a frequency of 5.00 x 1015 s-1?

48. Sample Problem Answer Given: f = 5.00 x 1015 s-1 h = 6.63 x 10-34 J-s
Unknown: E = ?
Parent Equation: E = h x f
Answer: x J

49. Your turn #1
What is the energy of radiation with a frequency of is 5.50 x 1014 Hz?
Active Inspire

52. Photoelectric Effect Photon - electron in the form of light
Ground state – normal position of electron
Excited state – electron jumps to higher energy level when energy is applied
Photoelectric effect – release of energy in the form of light when electron falls back to ground state
Color of light emitted depends on the energy level of the electrons
Blue light has a higher energy than red light

53. Photoelectric Effect

54. Atomic Emission Spectrum
Elements emit light when electrocuted in gaseous form
The light is passed through a prism and an atomic emission spectrum is obtained
Atomic emission spectra are NOT continuous like sunlight
Each line represents one distinct wavelength and frequency
Every element has a unique emission spectra

55. Atomic Emission Spectrum

56. Spectrums
When a narrow beam of emitted light is shined through a prism, it is separated into colors of the visible SPECTRUM

57. Types of Spectrums Visible light Spectrum: Atomic emission spectrum
Continuous range
Atomic emission spectrum
Light emitted by an element through a prism
Every element has a distinct emission spectrum
Discontinuous
Each line = one frequency

58. Different Elements Have Different Spectrums

59. Video of Spectra via Spectroscopes

60. More practice problems with energy, frequency, and wavelength
Last page of your chapter 13 packet

61. Chapter 13 Review Problems

62. Evolution of Atomic Models
5. Quantum Mechanical Model (1 of 2)
Quantum of energy = amount of energy required to move an electron from one energy level to the next higher level
Energy levels are not equally spaced
Levels get closer together the further from the nucleus
The further away from the nucleus, the less energy is required for an electron to escape ?