1. Chapter 5
Electrons in Atoms

2. Rutherford’s Model Discovered the nucleus Small dense and positive
Electrons moved around in Electron cloud

3. Bohr’s Model Why don’t the electrons fall into the nucleus?
Move like planets around the sun.
In circular orbits at different levels.
Energy separates one level from another.

4. Bohr’s Model
Nucleus
Electron
Orbit
Energy Levels

5. Bohr’s Model
Nucleus
Electron
Orbit
Energy Levels

6. } Bohr’s Model Further away from the nucleus means more energy. Fifth
There is no “in between” energy
Energy Levels
Fifth
Fourth
Third
Increasing energy
Second
First
Nucleus

7. The Quantum Mechanical Model
Energy is quantized. It comes in chunks.
Quanta - the amount of energy needed to move from one energy level to another.
Quantum leap in energy.
Schrödinger derived an equation that described the energy and position of the electrons in an atom
Treated electrons as waves

8. The Quantum Mechanical Model
a mathematical solution
It is not like anything you can see.

9. The Quantum Mechanical Model
Does have energy levels for electrons.
Orbits are not circular.
It can only tell us the probability of finding an electron a certain distance from the nucleus.

10. The Quantum Mechanical Model
The electron is found inside a blurry “electron cloud”
An area where there is a chance of finding an electron.

11. Atomic Orbital’s and Quantum Numbers
Principal Quantum Number (n) = indicates the main energy level occupied by the electron.
Positive integers 1,2,3,…
Within each energy level the complex math of Schrödinger's equation describes several shapes.
These are called atomic orbitals
Regions where there is a high probability of finding an electron.
The total number of orbital’s that exist in a main energy level is equal to n2

12. Angular Momentum Quantum Number
(l) indicates the shape of the orbital.
Except at E1, orbitals of different shapes (sublevels) exist for a given value of n.
The number of orbital shapes possible is equal to n.
l = 0, 1, 2, … n-1 (all positive integers)

13. Shapes of Orbitals n = 1 l = 0 one orbital s
n = 2 l = 0, two orbitals s, p
n = 3 I = 0, 1, 2 three orbitals s, p, d

14. S orbitals One s orbital for every energy level Spherical shaped
Each s orbital can hold 2 electrons
Called the 1s, 2s, 3s, etc.. orbitals.

15. P orbitals Start at the second energy level 3 different directions
3 different shapes (dumbell)
Each can hold 2 electrons

16. P Orbitals

17. D orbitals Start at the third energy level 5 different shapes
Each can hold 2 electrons

18. F orbitals Start at the fourth energy level
Have seven different shapes
2 electrons per shape

19. F orbitals

20. Summary # of shapes Max electrons Starts at energy level s 1 2 1 p 3 6d 5 10 3 7 14 4 f

21. Magnetic Quantum Numberml
the orientation of the orbital in 3-D space. (x, y, z)
the values of ml range from –l to +l
Ex: n=1 l=0 ml=0
n=2 l=0, 1 ml= -1,0, 1
n= l= ? ml= ?

22. Spin Quantum Number ms
electrons are not stationary particles, they spin
they can only spin in two directions, clockwise and counterclockwise (designations we have assigned them)
the values of ms are +1/2 or – 1/2

23. The Address of an Electron
No two electrons have the same 4 quantum numbers.
what I know from the quantum numbers of an electron:
1, 0, 0, +1/2
first principal energy level, s orbital, (x,y,z) axis, spinning clockwise
3, 1, -1, -1/2
Third principal energy level, p orbital, x axis, spinning counterclockwise

24. Aufbau Principle German for Building Up
When the we build an atom with its various electrons we start with the lowest energy level and build up.

25. The easy way to remember filling order
2s 2p theoretical
3s 3p 3d beyond f are
4s 4p 4d 4f all orbitals
5s 5p 5d 5f 5g
6s 6p 6d 6f 6g 6h
7s 7p 7d 7f 7g 7h 7i

26. Increasing energy 7s 2p 3p 4p 5p 6p 7p 3d 4d 5d 6d 6s 5s 4f 5f 4s 3s

27. Increasing energy 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p 3s 2p

28. Electron Configurations
The way electrons are arranged in atoms.
Aufbau principle- electrons enter the lowest energy first.
This causes difficulties because of the overlap of orbitals of different energies.
Pauli Exclusion Principle- at most 2 electrons per orbital - different spins

29. Electron Configuration
Hund’s Rule- When electrons occupy orbitals of equal energy they don’t pair up until they have to .
Let’s determine the electron configuration for Phosphorus
Need to account for 15 electrons

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

31. Increasing energy 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p
The next electrons go into the 2s orbital
only 11 more

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

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

34. Increasing energy 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s
The last three electrons go into the 3p orbitals.
They each go into separate shapes
3 unpaired electrons
1s22s22p63s23p3

35. The easy way to remember
2s 2p
3s 3p 3d
4s 4p 4d 4f
5s 5p 5d 5f
6s 6p 6d 6f
7s 7p 7d 7f
1s2
2 electrons

36. Fill from the bottom up following the arrows
2s 2p
3s 3p 3d
4s 4p 4d 4f
5s 5p 5d 5f
6s 6p 6d 6f
7s 7p 7d 7f
1s2 2s2
4 electrons

37. Fill from the bottom up following the arrows
2s 2p
3s 3p 3d
4s 4p 4d 4f
5s 5p 5d 5f
6s 6p 6d 6f
7s 7p 7d 7f
1s2 2s2 2p6 3s2
12 electrons

38. Fill from the bottom up following the arrows
2s 2p
3s 3p 3d
4s 4p 4d 4f
5s 5p 5d 5f
6s 6p 6d 6f
7s 7p 7d 7f
1s2 2s2 2p6 3s2 3p6 4s2
20 electrons

39. Fill from the bottom up following the arrows
2s 2p
3s 3p 3d
4s 4p 4d 4f
5s 5p 5d 5f
6s 6p 6d 6f
7s 7p 7d 7f
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2
38 electrons

40. Fill from the bottom up following the arrows
2s 2p
3s 3p 3d
4s 4p 4d 4f
5s 5p 5d 5f
6s 6p 6d 6f
7s 7p 7d 7f
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2
56 electrons

41. Fill from the bottom up following the arrows
2s 2p
3s 3p 3d
4s 4p 4d 4f
5s 5p 5d 5f
6s 6p 6d 6f
7s 7p 7d 7f
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d10 6p6 7s2
88 electrons

42. Fill from the bottom up following the arrows
2s 2p
3s 3p 3d
4s 4p 4d 4f
5s 5p 5d 5f
6s 6p 6d 6f
7s 7p 7d 7f
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d10 6p6 7s2 5f14 6d10 7p6
118 electrons

43. Rewrite when done
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d10 6p6 7s2 5f14 6d10 7p6
Group the energy levels together
1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 4d10 4f14 5s2 5p6 5d105f146s2 6p6 6d10 7s2 7p6

44. Exceptions to Electron Configuration

45. Orbitals fill in order Lowest energy to higher energy.
Adding electrons can change the energy of the orbital.
Filled and half-filled orbitals have a lower energy.
Makes them more stable.
Changes the filling order of d orbitals

46. Write these electron configurations
Titanium - 22 electrons
1s22s22p63s23p63d24s2
Vanadium - 23 electrons 1s22s22p63s23p63d34s2
Chromium - 24 electrons
1s22s22p63s23p63d44s2 is expected
But this is wrong!!

47. Chromium is actually 1s22s22p63s23p63d54s1 Why?
This gives us two half filled orbitals.

48. Chromium is actually 1s22s22p63s23p63d54s1 Why?
This gives us two half filled orbitals.

49. Chromium is actually 1s22s22p63s23p63d54s1 Why?
This gives us two half filled orbitals.
Slightly lower in energy.
The same principle applies to copper.

50. Copper’s electron configuration
Copper has 29 electrons so we expect
1s22s22p63s23p63d94s2
But the actual configuration is
1s22s22p63s23p63d104s1
This gives one filled orbital and one half filled orbital.
Remember these exceptions
d4s2  d5 s1
d9s2  d10s1

51. In each energy level
The number of electrons that can fit in each energy level is calculated with
Max e- = 2n2 where n is energy level
1st
2nd
3rd

52. Light
The study of light led to the development of the quantum mechanical model.
Light is a kind of electromagnetic radiation.
Electromagnetic radiation includes many kinds of waves
All move at 3.00 x 108 m/s ( c)

53. Parts of a wave
Crest
Wavelength
Amplitude
Origin
Trough

54. Parts of Wave Origin - the base line of the energy.
Crest - high point on a wave
Trough - Low point on a wave
Amplitude - distance from origin to crest
Wavelength - distance from crest to crest
Wavelength - is abbreviated l -Greek letter lambda.

55. Frequency
The number of waves that pass a given point per second.
Units are cycles/sec or hertz (Hz)
Abbreviated n - the Greek letter nu
c = ln

56. Frequency and wavelength
Are inversely related
As one goes up the other goes down.
Different frequencies of light is different colors of light.
There is a wide variety of frequencies
The whole range is called a spectrum

57. Spectrum Low energy High energy Radiowaves Microwaves Infrared .
Ultra-violet
X-Rays
GammaRays
Low Frequency
High Frequency
Long Wavelength
Short Wavelength
Visible Light

58. Light is a Particle Energy is quantized. Light is energy
Light must be quantized
These smallest pieces of light are called photons.
Energy and frequency are directly related.

59. Energy and frequency E = h x n E is the energy of the photon
n is the frequency
h is Planck’s constant
h = x Joules sec.

60. The Math in Chapter 5 Only 2 equations c = ln E = hn
c is always x 108 m/s
h is always x J s

61. Examples
What is the frequency of red light with a wavelength of 4.2 x 10-5 cm?
What is the wavelength of KFI, which broadcasts at with a frequency of 640 kHz?
What is the energy of a photon of each of the above?

62. How color tells us about atoms
Atomic Spectrum
How color tells us about atoms

63. Prism
White light is made up of all the colors of the visible spectrum.
Passing it through a prism separates it.

65. If the light is not white
By heating a gas or with electricity we can get it to give off colors.
Passing this light through a prism does something different.

66. Atomic Spectrum Each element gives off its own characteristic colors.
Can be used to identify the atom.
How we know what stars are made of.

67. These are called line spectra
unique to each element.
These are emission spectra
Mirror images are absorption spectra
Light with black missing

68. An explanation of Atomic Spectra

69. Where the electron starts
When we write electron configurations we are writing the lowest energy.
The energy level an electron starts from is called its ground state.

70. Changing the energy
Let’s look at a hydrogen atom

71. Changing the energy
Heat or electricity or light can move the electron up energy levels

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

73. Changing the energy May fall down in steps
Each with a different energy

74. The Bohr Ring Atom
n = 4
n = 3
n = 2
n = 1

75. { { {

76. Ultraviolet
Visible
Infrared
Further they fall, more energy, higher frequency.
This is simplified
the orbitals also have different energies inside energy levels
All the electrons can move around.

77. What is light? Light is a particle - it comes in chunks.
Light is a wave- we can measure its wave length and it behaves as a wave
If we combine E=mc2 , c=ln, E = 1/2 mv2 and E = hn
We can get l = h/mv (de Broglie’s equation)
The wavelength of a particle.

78. Matter is a Wave Does not apply to large objects
Things bigger than an atom
A baseball has a wavelength of about m when moving 30 m/s
An electron at the same speed has a wavelength of 10-3 cm
Big enough to measure.

79. Diffraction
When light passes through, or reflects off, a series of thinly spaced lines, it creates a rainbow effect
because the waves interfere with each other.

80. A wave moves toward a slit.

81. A wave moves toward a slit.

82. A wave moves toward a slit.

83. A wave moves toward a slit.

84. A wave moves toward a slit.

87. Comes out as a curve

88. Comes out as a curve

89. Comes out as a curve

90. with two holes

91. with two holes

92. with two holes

93. with two holes

94. with two holes

95. Two Curves
with two holes

96. Two Curves
with two holes

97. Two Curves
with two holes
Interfere with each other

98. Two Curves
with two holes
Interfere with each other
crests add up

99. Several waves

100. Several waves

101. Several waves

102. Several waves

103. Several waves

104. Several waves

105. Several waves

106. Several waves

107. Several waves

108. Several waves

109. Several waves
Several Curves

110. Several waves
Several Curves

111. Several waves
Several Curves

112. Several waves
Several Curves

113. Several waves
Several waves
Several Curves
Interference Pattern

114. Diffraction Light shows interference patterns Light is a wave
What will an electron do when going through two slits?
Go through one slit or the other and make two spots
Go through both and make a interference pattern

115. Electron as Particle
Electron “gun”

116. Electron as wave
Electron “gun”

117. Which did it do? It made the diffraction pattern
The electron is a wave
Led to Schrödingers equation

118. The physics of the very small
Quantum mechanics explains how the very small behaves.
Quantum mechanics is based on probability because

119. Heisenberg Uncertainty Principle
It is impossible to know exactly the speed and position of a particle.
The better we know one, the less we know the other.
The act of measuring changes the properties.

120. More obvious with the very small
To measure where a electron is, we use light.
But the light moves the electron
And hitting the electron changes the frequency of the light.

121. After Before Photon changes wavelength Photon
Electron changes velocity
Moving Electron