1. Chapter 5 “Electrons in Atoms”

2. Ernest Rutherford’s Model
Discovered dense + piece at nucleus
Electrons move around it, like planets around sun
Mostly empty space
Didn’t explain chemical properties of elements

3. Niels Bohr’s Model Why don’t electrons fall into nucleus?
Move like planets around sun.
specific circular orbits at different levels.
An amount of fixed energy separates one level from another.

4. The Bohr Model of the Atom
I pictured the electrons orbiting the nucleus much like planets orbiting the sun.
However, electrons are found in specific circular paths around the nucleus, and can jump from one level to another.
Niels Bohr

5. Bohr’s model Energy level measure of fixed energy e- can have
Like rungs of ladder
e- can’t exist between energy levels
can’t stand between rungs on ladder
Unlike ladder, energy levels not evenly spaced
Higher energy levels closer together – less energy needed for jump

6. The Quantum Mechanical Model
Energy is “quantized” - in chunks.
A quantum is exact energy needed to move e- one energy level to another.
Since energy of atom is never “in between” quantum leap in energy must exist.
Erwin Schrodinger (1926) derived equation described energy and position of e- in an atom
Schrodinger

7. The Quantum Mechanical Model
Very small things behave differently from big things
quantum mechanical model is a mathematical solution
not like anything you can see (like plum pudding!)

8. The Quantum Mechanical Model
Has energy levels for e-
Orbits not circular.
only tells us probability of finding e- a certain distance from nucleus.

9. The Quantum Mechanical Model
e- found inside blurry “electron cloud” (area where there’s chance of finding e-
Like fan blades or bicycle wheel

10. Atomic Orbitals
Principal Quantum Number (n) = the energy level of e- (1, 2, 3, etc.)
Within each energy level, Schrodinger’s equation describes several shapes.
called atomic orbitals - regions of space w/ high probability of finding e-
Sublevels like rooms in a hotel
letters s, p, d, and f

11. Principal Quantum Number (n)
denotes the energy level (shell) e- is located.
Maximum number of e- that can fit in energy level is:
2n2
How many e- in level 2? level 3?

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

15. By Energy Level Energy Level 2 (n=2) Has s and p sublevels available
2s (1 orbital)
w/ 2 e-
2p (3 orbitals)
w/ 6 e-
2s22p6
8 total e-
Energy Level 1 (n=1)
Has only s sublevel
1s (1 orbital)
w/ only 2 e-
2 total e- 2 2 6

16. By Energy Level Energy level 4 (n=4) Energy level 3 (n=3)
s, p, d, and f sublevels
4s (1 orbital)
w/ 2 e-
4p (3 orbitals)
w/ 6 e-
4d (5 orbitals)
w/ 10 e-
4f (7 orbitals)
w/ 14 e-
4s24p64d104f14
32 total e-
Energy level 3 (n=3)
s, p, and d sublevels
3s (1 orbital)
w/ 2 e-
3p (3 orbitals)
w/ 6 e-
3d (5 orbitals)
w/ 10 e-
3s23p63d10
18 total e- 2 2 6 6 10 14 10

17. By Energy Level Any more than the fourth and not all orbitals fill up.
You simply run out of e-’s
orbitals do not fill up in neat order.
energy levels overlap
Lowest energy fill first.
s and p orbitals 1:20
d orbitals 3:40

18. 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

19. Electron distribution in an Atom
Energy Level 3d 3 3p 3s
ENERGY 2 2p 2s 1 1s
NUCLEUS

20. Section 5.2 Electron Arrangement in Atoms
OBJECTIVES:
Describe how to write the electron configuration for an atom.

21. Section 5.2 Electron Arrangement in Atoms
OBJECTIVES:
Explain why the actual electron configurations for some elements differ from those predicted by the aufbau principle.

22. Increasing energy 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p 3s 2p
aufbau diagram - page 133
Aufbau is German for “building up”

23. Electron Configurations…
…the way electrons arranged in various orbitals around nuclei of atoms. Three rules tell us how:
Aufbau principle - electrons enter the lowest energy first.
causes difficulties b/c overlap of orbitals of different energies – follow diagram!
Pauli Exclusion Principle - 2 electrons max per orbital (hotel room) - different spins

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

25. Quantum Numbers
Each electron in an atom has a unique set of 4 quantum numbers which describe it.
1) Principal quantum number
2) Angular momentum quantum number
3) Magnetic quantum number
4) Spin quantum number

26. Electron Configurations
Hund’s Rule- When electrons occupy orbitals of equal energy, they don’t pair up until they have to.
write e- configuration for Phosphorus
We need to account for all 15 e-’s in phosphorus

27. Increasing energy 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p
The first two electrons go into the 1s orbital
Notice the opposite direction of the spins
only 13 more to go...

28. Increasing energy 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p 3s
The next electrons go into the 2s orbital

29. Increasing energy The next electrons go into the 2p orbital3d 4d 5d 7p 6d 4f 5f
The next electrons go into the 2p orbital
only 5 more...

30. Increasing energy The next electrons go into the 3s orbital2p 3p 4p 5p 6p 3d 4d 5d 7p 6d 4f 5f
The next electrons go into the 3s orbital
only 3 more...

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

32. Stairstep pattern

33. An internet program about electron configurations is:
(Just click on the above link)
I electron config 3:24

34. Orbitals fill in an order
Lowest energy  higher energy.
Adding e-’s changes energy of orbital. Full orbitals are best situation.
half filled orbitals  lower energy, next best
more stable.
Changes filling order

35. 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!!

36. Chromium is actually: 1s22s22p63s23p64s13d5 Why?
This gives us two half filled orbitals (the others are all still full)
Half full is slightly lower in energy.
The same principal applies to copper.

37. Copper’s electron configuration
Copper has 29 electrons so we expect: 1s22s22p63s23p63d94s2
But the actual configuration is:
1s22s22p63s23p63d104s1
This change gives one more filled orbital and one that is half filled.
Remember these exceptions: d4, d9

38. Irregular configurations of Cr and Cu
Chromium steals a 4s electron to make its 3d sublevel HALF FULL
Copper steals a 4s electron to FILL its 3d sublevel

39. Section 5.3 Physics and the Quantum Mechanical Model
OBJECTIVES:
Describe the relationship between the wavelength and frequency of light.

40. Section 5.3 Physics and the Quantum Mechanical Model
OBJECTIVES:
Identify the source of atomic emission spectra.

41. Section 5.3 Physics and the Quantum Mechanical Model
OBJECTIVES:
Explain how the frequencies of emitted light are related to changes in electron energies.

42. Section 5.3 Physics and the Quantum Mechanical Model
OBJECTIVES:
Distinguish between quantum mechanics and classical mechanics.

43. Light Study of light led to development of quantum mechanical model.
Light is electromagnetic radiation.
EM radiation: gamma rays, x-rays, radio waves, microwaves
Speed of light = x 108 m/s
abbreviated “c”
“celeritas“ Latin for speed
All EM radiation travels same in vacuum

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

45. Parts of a wave
Crest
Wavelength
Amplitude
Origin
Trough

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

47. Wavelength and Frequency
inversely related
one gets bigger, other smaller
Different frequencies of light different colors
wide variety of frequencies (spectrum)

48. - Page 140
Use Equation: c =

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

50. Long
Wavelength =
Low Frequency
Low ENERGY
Short
Wavelength =
High Frequency
High ENERGY

51. Atomic Spectra White light all colors of visible spectrum.
prism separates it.

52. If the light is not white
heating gas with electricity will emit colors.
Passing this light through prism different.

53. Atomic Spectrum Each element gives off own characteristic colors.
This is how we know composition of stars

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

55. 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.

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

57. 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!

58. 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?

59. 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.

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

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

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

63. Experiment #6, page 49-

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

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

66. 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.

67. = 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

68. 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!

69. 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

70. 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

71. 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

72. 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

73. 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

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

75. End of Chapter 5