1. Chapter 5 Electrons in Atoms p. 126

2. Ernest Rutherford’s Model
Discovered dense + nucleus
e-s move like planets around sun
Mostly empty space
Didn’t explain chemical properties of elements

3. “Why don’t e-s fall into nucleus”?
The Bohr Model
Neils Bohr ( )
Danish physicist
Student of Rutherford
“Why don’t e-s fall into nucleus”?
Why don’t electrons fall into nucleus?
Niels Bohr

4. The Bohr Model
I pictured the electrons found in specific circular paths around the nucleus, and can jump from one level to another.
Furthermore, each level has a fixed amount of energy different from other levels
Niels Bohr

5. fixed energy e- have called Energy levels
Bohr’s model
fixed energy e- have called Energy levels
Like rungs of ladder
e- can’t exist btwn energy levels
energy levels not evenly spaced
High levels closer (less energy needed to jump)
Bohr’s model of the atom 5:17

6. The Quantum Mechanical Model
e-’s don’t move like big objects
Rutherford & Bohr model
Energy - “quantized” (in chunks)
exact energy needed to move e- 1 energy level called a quantum
energy never “in btwn”
quantum leap in energy must exist
Erwin Schrodinger (1926) mathematically described energy & position of e- in atom
Schrodinger
Quantum Leap TV intro

7. The Quantum Mechanical Model
energy levels for e-
Orbits not circular
Based on probability of finding e- certain distance from nucleus
electron cloud

8. Atomic Orbitals
Principal Quantum # (n) - energy level of e- (1, 2, 3 etc.)
atomic orbitals - regions of space w/ high probability of finding e- (not a true “orbit”)
within each energy level
Sublevels like rooms in a hotel
s, p, d, and f
Different shapes
How many e- in level 2? level 3?
Max # of e- that fit in energy level is:
2n2

9. s and p orbitals 1:20

10. atomic orbitals review (14:28)
d orbitals 3:40
atomic orbitals review (14:28)

11. atomic orbitals 2 6 10 14 First possible energy level
# of orbitals (regions of space)
Maximum electrons s
spherical 2 1
1st p
dumbell 3 6
2nd d
clover leaf 10 5
3rd f
complicated 7 14
4th

12. Summary of Principal Energy Levels, Sublevels, and Orbitals
Number of sublevels
Type of sublevel
Max # of electrons
Electron configuration
n = 1 1
1s (1 orbital) 2
1s2
n = 2
2s (1 orbital
2p (3 orbitals) 8
2s2
2p6
n = 3 3
3s (1 orbital)
3p (3 orbitals)
3d (5 orbitals) 18
3s2
3p6
3d10
n = 4 4
4s (1 orbital)
4p (3 orbitals)
4d (5 orbitals)
4f (7 orbitals) 32
4s2
4p6
4d10
4f14

13. ORDER OF ELECTRON SUBSHELL FILLING: NOT “IN ORDER”
energy levels overlap
Lowest energy fill first
1s2
2s2
2p6
3s2
3p6
3d10
Increasing energy
4s2
4p6
4d10
4f14
5s2
5p6
5d10
5f14
6s2
6p6
6d10
7s2
7p6
1s2
2s2
2p6
3s2
3p6
4s2
3d10
4p6
5s2
4d10
5p6
6s2
4f14
5d10
6p6
7s2
5f14
6d10
7p6

14. ELECTRON CONFIGURATION
group #
# valence e-
s: 1 or 2
p: 1-6
d: 1-10
f: 1-14
Total e- = Atomic #
1s1
Principal energy level
row #
1-7
7 rows
sublevel
s, p, d, or f
4 sublevels

15. Sublevels d and f are “special”
group # = # valence e- 1 2 3 4 5 6 7 d 3d
period # = # e- energy levels 4d 5d 6d f 4f 5f

16. Section 5.2 Electron Arrangement in Atoms p. 133
Increasing energy 1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f
aufbau diagram - page 133
Aufbau - German for “building up”

17. Electron Configurations…
….3 rules explain how e-’s fill their orbitals:
Aufbau principle – e-’s enter lowest energy level first.
Pauli Exclusion Principle - 2 e-’s max/orbital (hotel room) - different spins

18. Pauli Exclusion Principle
No 2 electrons in an atom can have the same four quantum numbers.
To show different direction of spin, a pair in the same orbital is written as:
Wolfgang Pauli

19. Quantum Numbers Each e- has unique set of 4 quantum #’s describing it
1) Principal quantum #
2) Angular momentum quantum #
3) Magnetic quantum #
4) Spin quantum #

20. Electron Configurations
Hund’s Rule- When e-’s occupy orbitals of same energy, they won’t pair up until they must
write e- configuration for Phosphorus
all 15 e-’s must be accounted for

21. Increasing energy 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p
The first 2 e-’s go into the 1s orbital
Notice opposite direction of spins

22. Increasing energy 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p 3s
The next e-’s go in 2s orbital

23. Increasing energy The next e-’s go in 2p orbital 7p 6d 5f 7s 6p 5d 6s

24. Increasing energy The next e-’s go in 3s orbital 7p 6d 5f 7s 6p 5d 6s

25. Increasing energy 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p 3s 2p
The last 3 e-’s go in 3p orbitals
They each go into separate shapes (Hund’s)
3 unpaired e-’s
= 1s22s22p63s23p3
Orbital notation

26. An internet program about electron configurations is:
I electron config (song) 3:24

27. Adding e-’s changes energy of orbital
Filling orbitals
Lowest  higher energy
Adding e-’s changes energy of orbital
Full orbitals best situation
half filled orbitals next best
more stable
Changes filling order

28. Write the electron configurations for these elements:
Titanium - 22 electrons
1s22s22p63s23p64s23d2
Vanadium - 23 electrons
1s22s22p63s23p64s23d3
Chromium - 24 electrons
1s22s22p63s23p64s23d4 (expected)
But this is not what happens!!

29. 1s22s22p63s23p64s13d5 Why? 2 half filled orbitals
Chromium is actually:
1s22s22p63s23p64s13d Why?
2 half filled orbitals
Half full slightly lower in energy
Same applies to copper

30. Copper’s e- configuration
Copper has 29 e-s so expect: 1s22s22p63s23p63d94s2
actual configuration is:
1s22s22p63s23p63d104s1
1 more full orbital & 1 half filled
Exceptions d4 d9

31. Irregular configurations of Chromium and Copper
Chromium steals a 4s e- to make its 3d sublevel HALF FULL
Copper steals a 4s electron to FILL its 3d sublevel

32. Section 5.3 Physics and the Quantum Mechanical Model p. 138
Light
Section 5.3 Physics and the Quantum Mechanical Model p. 138
Study of light led to quantum mechanical model
Light is electromagnetic radiation
EM radiation: gamma rays, x-rays, radio waves, microwaves
Speed of light = x 108 m/s
“c” - celeritas (Latin for speed)
All EM radiation travels same in vacuum

33. - Page 139
“R O Y G B I V”
Frequency Increases
Wavelength Longer

34. Parts of a wave
Crest
Wavelength
Amplitude
Trough

35. Electromagnetic radiation propagates through space as a wave moving at the speed of light.
Equation:
c =
c = is a constant (2.998 x 108 m/s)
 (lambda) = wavelength, in meters
 (nu) = frequency, in units of hertz (hz or sec-1)

36. Wavelength and Frequency
inversely related
one gets bigger, other smaller
Different frequencies = different colors
wide range of frequencies (spectrum)

37. - Page 140
Use Equation: c =

38. Low Energy
High Energy
Radiowaves
Microwaves
Infrared .
Ultra-violet
X-Rays
GammaRays
Low Frequency
High Frequency
Long Wavelength
Short Wavelength
Visible Light

39. Long  =
Low Frequency =
Low ENERGY
Short  =
High Frequency =
High ENERGY

40. White light all colors of visible spectrum
Atomic Spectra
White light all colors of visible spectrum
prism separates it according to λ

41. If the light is not white
heating gas with electricity will emit colors
this light thru prism is different

42. Atomic Spectrum elements emit own characteristic colors
composition of stars determined thru spectral analysis

43. atomic emission spectrum
Unique to each element, like fingerprints!
ID’s elements

44. Light is a Particle? Energy is quantized Light is energy…..
light must be quantized
photons smallest pieces of light
Photoelectric effect –
Matter emits e- when it absorbs energy
Albert Einstein Nobel Prize in chem
Energy & frequency: directly related

45. E = Energy, in units of Joules (kg·m2/s2)
energy (E ) of electromagnetic radiation directly proportional to frequency () of radiation.
Planck-Einstein Equation: E = h
E = Energy, in units of Joules (kg·m2/s2)
(Joule…metric unit of energy)
h = Planck’s constant (6.626 x J·s)
(reflecting sizes of energy quanta)
 = frequency, units of hertz (hz, sec-1)

46. c =  E = h The Math in Chapter 5 Put these on your 3 x 5 notecard!
There are 2 equations:
c = 
E = h
Put these on your 3 x 5 notecard!

47. What is the frequency of red light with a wavelength of 4.2 x 10-5 m?
Examples
What is the wavelength of blue light with a frequency of 8.3 x 1015 hz?
What is the frequency of red light with a wavelength of 4.2 x 10-5 m?
What is the energy of a photon of each of the above?

48. Explanation of atomic spectra
electron configurations written in lowest energy. energy level, and where electron starts from, called it’s ground state - lowest energy level.

49. Changing the energy
Let’s look at a hydrogen atom, with only one electron, and in the first energy level.

50. Changing the energy
Heat, electricity, or light can move e-’ up to different energy levels. The electron is now said to be “excited”

51. Changing the energy
As electron falls back to ground state, it gives energy back as light

52. Experiment #6, page 49-

53. Changing the energy may fall down in specific steps
Each step has different energy

54. { { { Balmer series (visible) Paschen series Lyman series (UV)
(infrared) { { {

55. further they fall, more energy released = higher frequency
Ultraviolet
Visible
Infrared
further they fall, more energy released = higher frequency
orbitals also have different energies inside energy levels
All electrons can move around.

56. = h/mv (from Louis de Broglie)
What is light?
Light is a particle - it comes in chunks.
Light is a wave - we can measure its wavelength and it behaves as a wave
combine E=mc2 , c=, E = 1/2 mv2 and E = h, then we can get:
= h/mv (from Louis de Broglie)
Calculates wavelength of a particle.
called de Broglie’s equation
He said particles exhibit properties of waves

57. Wave-Particle Duality
J.J. Thomson won the Nobel prize for describing the electron as a particle.
His son, George Thomson won the Nobel prize for describing the wave-like nature of the electron.
The electron is a particle!
The electron is an energy wave!

58. Confused? You’ve Got Company!
“No familiar conceptions can be woven around the electron; something unknown is doing we don’t know what.”
Physicist Sir Arthur Eddington
The Nature of the Physical World
1934

59. The physics of the very small
Quantum mechanics explains how very small particles behave
Quantum mechanics is an explanation for subatomic particles and atoms as waves
Classical mechanics describes the motions of bodies much larger than atoms

60. Heisenberg Uncertainty Principle
impossible to know exact location and velocity of particle better we know one, less we know other Measuring changes properties. True in quantum mechanics, but not classical mechanics

61. Heisenberg Uncertainty Principle
“One cannot simultaneously determine both the position and momentum of an electron.”
You can find out where the electron is, but not where it is going.
OR…
You can find out where the electron is going, but not where it is!
Werner Heisenberg

62. It is more obvious with the very small objects
To measure where e-, we use light But light energy (photon) moves e- due to small mass And hitting e- changes frequency of light

63. After Before Photon wavelength changes Photon Moving Electron
Electron velocity changes
Fig. 5.16, p. 145