1. Chapter 5
Electrons in Atoms

2. Greek Idea Democritus and Leucippus
Matter is made up of indivisible particles
Dalton - one type of atom for each element

3. Thomson’s Model Discovered electrons Atoms were made of positive stuff
Negative electron floating around
“Plum-Pudding” model

4. Rutherford’s Model
Discovered dense positive piece at the center of the atom
Nucleus
Electrons moved around
Mostly empty space

5. Bohr’s Model Why don’t the electrons fall into the nucleus?
Move like planets around the sun.
In circular orbits at different levels.
Amounts of energy separate one level from another.

6. Bohr’s Model
Nucleus
Electron
Orbit
Energy Levels

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

8. 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)

9. Parts of a Wave
Crest
Wavelength
Amplitude
Origin
Trough

10. Parts of Wave Origin - the base line of the energy.
Crest - highest 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.

11. 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

12. Wavelength and Frequency

13. 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
Movie Flame Test

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

16. Aurora Borealis
Energy entering the earth’s atmosphere causes gas atom electrons to become excited. When they fall to the ground state they give of photons of light.
Movie Aurora

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

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

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

20. 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.
Movie Emission Spectrum

21. Atomic Emission Spectrum

22. These are called discontinuous spectra
Or line spectra
unique to each element.
These are emission spectra
The light is emitted given off.

24. 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.

25. Energy and Frequency E = hn E is the energy of the photon
n is the frequency
h is Planck’s constant
h = x Joules sec.
joule is the metric unit of Energy

26. The Math in Chapter 5
Only 2 equations
c = ln
E = h n
Plug and chug.

27. 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?
Given h = x Joules sec and the Speed of light equals 3.00 x 108 m/s ( c)

28. An Explanation of Atomic Spectra

29. Where the Electron Starts
When we write electron configurations we are writing the lowest energy.
The energy level where an electron starts from is called its ground state.
Movie Energy Levels

30. Changing the Energy
Let’s look at a hydrogen atom

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

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

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

34. { { {

35. 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.

36. What Makes These Glow?

37. 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= n l, E = 1/2 mv2 and E = hn
We can get De Broglie’s equation l = h/mv
The wavelength of a particle.

38. Matter is a Wave Does not apply to large objects
Things bigger than an atom
A baseball has a wavelength of about
1x10-32 m when moving 30 m/s
An electron at the same speed has a wavelength of 1x10-3 cm
Big enough to measure.

39. The Physics of the Very Small
Quantum mechanics explains how the very small behaves.
Classic physics is what you get when you add up the effects of millions of packages.
Quantum mechanics is based on probability because

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

41. Measuring an Electron To measure where an electron is, we use light
But the light moves the electron
And hitting the electron changes the frequency of the light

42. After Before Photon Changes Wavelength Photon
Electron Changes Velocity
Moving Electron

43. The Quantum Mechanical Model
Remember energy is quantized and comes in chunks.
A quanta is the amount of energy needed to move from one energy level to another.
Since the energy of an atom is never “in between” there must be a quantum leap in energy.
Schrodinger derived an equation that described the energy and position of the electrons in an atom.

44. Atoms Are Never “In Between” Levels

45. The Quantum Mechanical Model
Things that are very small behave differently from things big enough to see.
The quantum mechanical model is a mathematical solution
It is not like anything you can see.

46. The Quantum Mechanical Model
Has 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.

47. The Quantum Mechanical Model
The atom is found inside a blurry “electron cloud”
An area where there is a chance of finding an electron.
Can be divided into smaller regions in the electron cloud.

48. Atomic Orbitals
Principle Quantum Number (n) = the energy level of the electron
Within each energy level the complex math of Schrodinger’s equation describes several geometric shapes.
The group shape is called the sublevel s,p,d,f
Individual shapes are the electron orbitals
The orbitals represent regions where there is a high probability of finding an electron.

49. Sublevels and Energy Levels

50. s Sublevel 1 s orbital is found on every energy level Spherical shaped
Each s orbital can hold 2 electrons
Called the 1s, 2s, 3s, etc.. orbitals. The 1,2 and 3 designate the energy levels.

51. 6s Sublevel

52. S- block s1 s2 Alkali metals all end in s1
Alkaline earth metals all end in s2
really should include He, but it fits better later.
He has the properties of the noble gases.

53. s, p, d, & f blocks

54. p Sublevel Start at the second energy level 3 different directions
3 different shapes
Each can hold 2 electrons

55. 6p Sublevel

56. s & p orbitals for neon
Glencoe ‘Chemistry Matter and Change’ 2002, page 137

57. The P-block p1 p2 p3 p4 p6 p5

58. Each row (or period) is the energy level for s and p orbitals.1 2 3 4 5 6 7
Each row (or period) is the energy level for s and p orbitals.

59. d Sublevel Start at the third energy level 5 different shapes
Each can hold 2 electrons

60. 6d Sublevel

61. Transition Metals -d block

62. d orbitals fill up after previous energy level, so first d is 3d even though it’s in row 4.1 2 3 4 5 6 7 3d

63. f Sublevel Start at the fourth energy level
Have seven different shapes
2 electrons per shape

64. f Sublevel

65. 6f Sublevel

66. F - block f1 f5 f2 f3 f4 f6 f7 f8 f9 f10 f11 f12 f14 f13
inner transition elements f1 f5 f2 f3 f4 f6 f7 f8 f9
f10
f11
f12
f14
f13

67. 1 2 3 4 5 6 7 4f 5f
f orbitals start filling at 4f

68. s, p, d, & f blocks

69. 6g Sublevel Sublevel g starts on the 5th level. 9 shapes
18 electrons max

70. Summary # of shapes Max electrons Starts at energy level SL s 1 2 1 p3 6 2 d 5 10 3 7 14 4 f

71. By Energy Level First Energy Level only s orbital only 2 electrons 1s2
Second Energy Level
s and p orbitals are available
2 in s, 6 in p
2s22p6
8 total electrons

72. By Energy Level Third energy level s, p, and d orbitals
2 in s, 6 in p, and 10 in d
3s23p63d10
18 total electrons
Fourth energy level
s,p,d, and f orbitals
2 in s, 6 in p, 10 in d, and 14 in f
4s24p64d104f14
32 total electrons

73. By Energy Level
Any more than the fourth and not all the orbitals will fill up.
You simply run out of electrons
The orbitals do not fill up in a neat order.
The energy levels overlap
Lowest energy fill first.

74. Writing Electron Configurations the Easy Way
Yes there is a shorthand

75. Electron Configurations repeat
The shape of the periodic table is a representation of this repetition.
When we get to the end of the column the outermost energy level is full.
This is the basis for our shorthand.

76. The Shorthand Method
Write symbol of the noble gas before the element, in [ ].
Then, the rest of the electrons.
Aluminum’s full configuration:
1s22s22p63s23p1
previous noble gas Ne is: 1s22s22p6
so, Al is: [Ne] 3s23p1

77. The Shorthand Again Sn- 50 electrons The noble gas before it is Kr
Takes care of 36
Next 5s2
Then 4d10
Finally 5p2
[ Kr ]
5s2
4d10
5p2

78. Increasing energy 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p
Aufbau Diagram shows the energy of each sublevel and the order they fill with electrons
Each box represents an atomic orbital

79. 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

80. 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

81. 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 spins
Same as Helium
1s2
only 13 more

82. 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
‘Be’ on periodic table
1s22s2
only 11 more

83. Increasing energy 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p
The next electrons go into the 2p orbitals
Neon on periodic table
1s22s22p6
only 5 more

84. 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
‘Mg’ on periodic table
Only 3 more
1s22s22p63s2

85. 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 orbitals
3 unpaired electrons
1s22s22p63s23p3

86. 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

87. 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

88. 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

89. 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

90. 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

91. 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

92. 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

93. 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

94. Exceptions to the Rule in Electron Configuration

95. Orbitals Fill Order Lowest energy to higher energy
Adding electrons can change the energy of the orbital
Half filled orbitals have a lower energy
Makes them more stable
Sometimes changes the filling order

96. Write These Electron Configurations
Titanium - 22 electrons
1s22s22p63s23p64s23d2
Vanadium - 23 electrons 1s22s22p63s23p64s23d3
Chromium - 24 electrons
1s22s22p63s23p64s23d4 is expected
But this is wrong!!

97. Chromium is Actually 1s22s22p63s23p64s13d5 Why?
This gives us two half filled orbitals
Slightly lower in energy
Holds true for other group 6 elements Cr, Mo, W, and probably Sg
The same principal applies to copper Cu and group 11 elements

98. Transition Metals -d block

99. Cu’s Electron Configuration
Copper has 29 electrons so we expect
1s22s22p63s23p64s23d9
But the actual configuration is
1s22s22p63s23p64s13d10
This gives one filled orbital and one half filled orbital.
Same is true for ‘Silver’ Ag and ‘Gold’ Au and probably ‘Unununium’ Uuu

100. Lewis/Electron Dot Diagram