1. ELECTROMAGNETIC
SPECTRUM
Chapter 5

2. Section 5.1 Physics and the Quantum Mechanical Model
OBJECTIVES:
Describe the relationship between the wavelength and frequency of light.

3. Section 5.1 Physics and the Quantum Mechanical Model
OBJECTIVES:
Identify the source of atomic emission spectra.

4. Section 5.1 Physics and the Quantum Mechanical Model
OBJECTIVES:
Explain how the frequencies of emitted light are related to changes in electron energies.

5. Section 5.1 Physics and the Quantum Mechanical Model
OBJECTIVES:
Distinguish between quantum mechanics and classical mechanics.

6. 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 types: gamma rays, x-rays, radio waves…
Speed of light = x 108 m/s, and is abbreviated “c”
All electromagnetic radiation travels at this same rate when measured in a vacuum

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

8. Parts of a wave

9. Parts of a wave
Crest
Wavelength
Amplitude
Origin
Trough

10. 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)

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

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

13. Long Wavelength = Low Frequency Low ENERGY Short Wavelength =
Wavelength Table
Long
Wavelength =
Low Frequency
Low ENERGY
Short
Wavelength =
High Frequency
High ENERGY

14. Water and sound waves transfer energy from one place to another- they require a medium through which to travel. They are mechanical waves. .

15. Nature of Electromagnetic Waves
They are Transverse waves without a medium. (They can travel through empty space)
They travel as vibrations in electrical and magnetic fields.
Have some magnetic and some electrical properties to them.
Speed of electromagnetic waves = 300,000,000 meters/second (Takes light 8 minutes to move from the sun to earth {150 million miles} at this speed.)

16. When an electric field changes, so does the magnetic field
When an electric field changes, so does the magnetic field. The changing magnetic field causes the electric field to change. When one field vibrates—so does the other.
RESULT-An electromagnetic wave.

17. Waves or Particles
Electromagnetic radiation has properties of waves but also can be thought of as a stream of particles. Example: Light
Light as a wave: Light behaves as a transverse wave which we can filter using polarized lenses.
Light as particles (photons)
When directed at a substance light can knock electrons off of a substance (Photoelectric effect)

18. Measuring the electromagnetic spectrum
You actually know more about it than you may think! The electromagnetic (EM) spectrum is just a name that scientists give a bunch of types of radiation when they want to talk about them as a group.
Radiation is energy that travels and spreads out as it goes-- visible light that comes from a lamp in your house or radio waves that come from a radio station are two types of electromagnetic radiation.

19. Measuring the electromagnetic spectrum
Other examples of EM radiation are microwaves, infrared and ultraviolet light, X-rays and gamma-rays.
Hotter, more energetic objects and events create higher energy radiation than cool objects.
Only extremely hot objects or particles moving at very high velocities can create high-energy radiation like X-rays and gamma-rays.

20. B. Waves of the Electromagnetic Spectrum
Electromagnetic Spectrum—name for the range of electromagnetic waves when placed in order of increasing frequency
Click here (Animation—Size of EMwaves)
GAMMA RAYS
ULTRAVIOLET RAYS
RADIO WAVES
INFRARED RAYS
X-RAYS
MICROWAVES
VISIBLE LIGHT

21. Here are the different types of radiation in the EM spectrum, in order from lowest energy to highest:

22. RADIO WAVES
A. Have the longest wavelengths and lowest frequencies of all the electromagnetic waves.
B. A radio picks up radio waves through an antenna and converts it to sound waves.
(Garage doors and wireless networks)
C. Each radio station in an area broadcasts at a different frequency. # on radio dial tells frequency.
D. MRI (MAGNETIC RESONACE IMAGING)
Uses Short wave radio waves with a magnet to create an image

23. MRI of the Brain

25. Radio
Radio: yes, this is the same kind of energy that radio stations emit into the air for your boom box to capture and turn into your favorite tunes. But radio waves are also emitted by other things ... such as stars and gases in space. You may not be able to dance to what these objects emit, but you can use it to learn what they are made of.

26. AM=Amplitude modulation—waves bounce off ionosphere can pick up stations from different cities.
(535kHz-1605kHz= vibrate at 535 to 1605 thousand times/second) +

27. FM=Frequency modulation—waves travel in a straight line & through the ionosphere--lose reception when you travel out of range.
(88MHz-108MHz = vibrate at 88million to 108million times/second) +

28. Bands of Radio/TV/Microwaves

29. MICROWAVES
Microwaves—have the shortest wavelengths and the highest frequency of the radio waves.
Used in microwave ovens.
Waves transfer energy to the water in the food causing them to vibrate which in turn transfers energy in the form of heat to the food.
Used by cell phones and pagers.
RADAR (Radio Detection and Ranging)
Used to find the speed of an object by sending out radio waves and measuring the time it takes them to return.

31. MICROWAVES
Microwaves: they will cook your popcorn in just a few minutes! In space, microwaves are used by astronomers to learn about the structure of nearby galaxies, including our own Milky Way!

32. INFRARED RAYS Infrared= below red
Shorter wavelength and higher frequency than microwaves.
You can feel the longest ones as warmth on your skin
Heat lamps give off infrared waves.
Warm objects give off more heat energy than cool objects.
Thermogram—a picture that shows regions of different temperatures in the body. Temperatures are calculated by the amount of infrared radiation given off. Therefore people give off infrared rays.

38. INFRARED RAYS
Infrared: we often think of this as being the same thing as 'heat', because it makes our skin feel warm. In space, IR light maps the dust between stars.

39. VISIBLE LIGHT
Shorter wavelength and higher frequency than infrared rays.
Electromagnetic waves we can see.
Longest wavelength= red light
Shortest wavelength= violet (purple) light
When light enters a new medium it bends (refracts). Each wavelength bends a different amount allowing white light to separate into it’s various colors ROYGBIV.

41. VISIBLE LIGHT
Visible: yes, this is the part that our eyes see. Visible radiation is emitted by everything from fireflies to light bulbs to stars ... also by fast-moving particles hitting other particles.

42. ULTRAVIOLET RAYS
Shorter wavelength and higher frequency than visible light
Carry more energy than visible light
Used to kill bacteria. (Sterilization of equipment)
Causes your skin to produce vitamin D (good for teeth and bones)
Used to treat jaundice ( in some new born babies.
Too much can cause skin cancer.
Use sun block to protect against (UV rays)

46. .

47. ULTRAVIOLET RAYS
Ultraviolet: we know that the Sun is a source of ultraviolet (or UV) radiation, because it is the UV rays that cause our skin to burn! Stars and other "hot" objects in space emit UV radiation.

48. X- RAYS Shorter wavelength and higher frequency than UV-rays
Carry a great amount of energy
Can penetrate most matter.
Bones and teeth absorb x-rays. (The light part of an x-ray image indicates a place where the x-ray was absorbed)
Too much exposure can cause cancer
(lead vest at dentist protects organs from unnecessary exposure)
Used by engineers to check for tiny cracks in structures.
The rays pass through the cracks and the cracks appear dark on film.

50. Type:
JPG

54. GAMMA RAYS Shorter wavelength and higher frequency than X-rays
Carry the greatest amount of energy and penetrate the most.
Used in radiation treatment to kill cancer cells.
Can be very harmful if not used correctly.

57. GAMMA RAYS
Gamma-rays: radioactive materials (some natural and others made by man in things like nuclear power plants) can emit gamma-rays.
Big particle accelerators that scientists use to help them understand what matter is made of can sometimes generate gamma-rays.
But the biggest gamma-ray generator of all is the Universe! It makes gamma radiation in all kinds of ways.

59. Brief Summary
A. All electromagnetic waves travel at the same speed. (300,000,000 meters/second in a vacuum.
B. They all have different wavelength and different frequencies.
Long wavelength-lowest frequency
Short wavelength highest frequency
The higher the frequency the higher the energy.

61. Section 5.2 Models of the Atom
OBJECTIVES:
Identify the inadequacies in the Rutherford atomic model.

62. Section 5.2 Models of the Atom
OBJECTIVES:
Identify the new proposal in the Bohr model of the atom.

63. Section 5.2 Models of the Atom
OBJECTIVES:
Describe the energies and positions of electrons according to the quantum mechanical model.

64. Section 5.2 Models of the Atom
OBJECTIVES:
Describe how the shapes of orbitals related to different sublevels differ.

65. Ernest Rutherford’s Model
Discovered dense positive piece at the center of the atom- “nucleus”
Electrons would surround and move around it, like planets around the sun
Atom is mostly empty space
It did not explain the chemical properties of the elements – a better description of the electron behavior was needed

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

67. 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

68. Bohr’s model Energy level of an electron
analogous to the rungs of a ladder
The electron cannot exist between energy levels, just like you can’t stand between rungs on a ladder
A quantum of energy is the amount of energy required to move an electron from one energy level to another

69. The Quantum Mechanical Model
Energy is quantized - It comes in chunks.
A quantum 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.
In 1926, Erwin Schrodinger derived an equation that described the energy and position of the electrons in an atom

70. Schrodinger’s Wave Equation
Equation for the probability of a single electron being found along a single axis (x-axis)
Erwin Schrodinger
Erwin Schrodinger

71. 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.

72. 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.

73. The Quantum Mechanical Model
The atom is found inside a blurry “electron cloud”
An area where there is a chance of finding an electron.
Think of fan blades

74. Atomic Orbitals
Principal Quantum Number (n) = the energy level of the electron: 1, 2, 3, etc.
Within each energy level, the complex math of Schrodinger’s equation describes several shapes.
These are called atomic orbitals - regions where there is a high probability of finding an electron.
Sublevels- like theater seats arranged in sections: letters s, p, d, and f

76. Principal Quantum Number
Generally symbolized by “n”, it denotes the shell (energy level) in which the electron is located.
Maximum number of electrons that can fit in an energy level:
2n2

77. Bohr Models of the Atom

78. Summary 2 6 10 14 # of shapes Max electrons Starts at energy level s 13 6 2 10 d 5 3 7 14 4 f

79. By Energy Level First Energy Level Has only s orbital only 2 electrons
Second Energy Level
Has s and p orbitals available
2 in s, 6 in p
2s22p6
8 total electrons

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

81. 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.

82. Section 4.3 Electron Arrangement in Atoms
OBJECTIVES:
Describe how to write the electron configuration for an atom.

83. Section 4.3 Electron Arrangement in Atoms
OBJECTIVES:
Explain why the actual electron configurations for some elements differ from those predicted by the aufbau principle.

84. Using the Periodic Table to Determine Electronic Arrangement
An electron configuration organizes an atom’s electrons according to certain rules. This configuration identifies the energy in an atom.
Electrons can only be in certain levels according to the following pattern:
1s,2s,2p,3s,3p,4s,3d,4p,5s,4d,5p,6s,4f, 5d,6p,7s,5f,6d,7p

85. BOHR MODELS OF THE ATOM Electrons by energy Level: (2n2)
1st level 2 e max
2nd level 8 e max
3rd level 18 e max
4th level 32 e max
5th level 50 e max
6th level 72 e max…
7th level 98 e max

86. You will learn three ways to write electron configurations. -Arrows
You will learn three ways to write electron configurations. -Arrows -Numbers -Noble Gas Configuration

87. Arrow Method

88. The Arrow Method
The Next couple of pages will teach you how to complete the arrow method for electron configuration.
Remember, you must fill-up on letter before moving on to a different type of letter.
Each space or dash can hold a maximum of two arrows.
Each space or dash (orbital) prefers to be single unless it has to be paired up.

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

90. Electron Configurations…
…are the way electrons are arranged in various orbitals around the nuclei of atoms. Three rules tell us how:
Aufbau principle - electrons enter the lowest energy first.
This causes difficulties because of the overlap of orbitals of different energies – follow the diagram!
Pauli Exclusion Principle - at most 2 electrons per orbital - different spins

91. 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

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

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

94. Increasing energy 7p 6d 5f 7s 6p 5d 6s 4f 5p 4d 5s 4p 3d 4s 3p 3s 2p
The first two electrons go into the 1s orbital
Notice the opposite direction of the spins
only 13 more to go...

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

96. 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...

97. 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...

98. 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

99. Orbitals fill in an order
Lowest energy to higher energy.
Adding electrons can change the energy of the orbital. Full orbitals are the absolute best situation.
However, half filled orbitals have a lower energy, and are next best
Makes them more stable.
Changes the filling order

100. Lets Try These:
Calcium
Fluorine
Titanium
Copper
Bromine

101. Numerical
Configuration

102. Numerical Electron Configuration
Similar to the arrow method, but instead of arrows, we use numerical exponents.
Same rules apply.

103. Examples: Hydrogen Chlorine Boron Copper Silicon Calcium Argon Xenon
Silver
Zirconium
1s1
1s2 2s2 2p6 3s2 3p5
1s2 2s2 2p1
1s2 2s2 2p6 3s2 3p6 4s2 3d9
1s2 2s2 2p6 3s2 3p2
1s2 2s2 2p6 3s2 3p6 4s2
1s2 2s2 2p6 3s2 3p6
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d9
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d2

104. Noble Gas
Configuration

105. Noble Gas Electron Configuration
(This is the shortest method)
**Color Periodic Table with s, p, d, f’s**
You must indicate the last noble gas element and then finish the configuration.
Calcium [Ar] 4s2

106. Examples: [He] 2s2 2p5 Fluorine [Ne] 3s2 Magnesium [Ne] 3s2 3p1
[Ar] 4s2 3d5
[Kr] 5s2 4d10 5p2
[Kr] 5s2 4d10 5p5
[Xe] 6s2 4f14 5d8
[Xe] 6s1
[Ne] 3s2 3p3
[Ar] 4s2 3d8
Fluorine
Magnesium
Aluminum
Manganese
Tin
Iodine
Platinum
Cesium
Phosphorus
Nickel

107. Do These and Turn in: Se Y Pd Os Au Ra Dy Am Ca Mn B Ga Mo Te Rb
Extra Credit (Long Method) V In Eu Cr

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

109. 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.

110. Copper’s electron configuration
Copper has 29 electrons so we expect: 1s22s22p63s23p64s23d9
But the actual configuration is:
1s22s22p63s23p64s13d10
This change gives one more filled orbital and one that is half filled.
Remember these exceptions: d4, d9

111. 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

112. Line and Continuous
Spectra

113. Atomic Spectra
White light is made up of all the colors of the visible spectrum.
Passing it through a prism separates it.

114. 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.

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

116. These are called the atomic emission spectrum
Unique to each element, like fingerprints!
Very useful for identifying elements

117. Light is a Particle? Energy is quantized. Light is a form of energy.
Therefore, light must be quantized
These smallest pieces of light are called photons.
Energy & frequency: directly related.

118. Explanation of atomic spectra
When we write electron configurations, we are writing the lowest energy.
The energy level, and where the electron starts from, is called it’s ground state - the lowest energy level.

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

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

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

122. Changing the energy They may fall down in specific steps
Each step has a different energy

123. { { {

124. This is a simplified explanation!
Ultraviolet
Visible
Infrared
The further they fall, more energy is released and the higher the frequency.
This is a simplified explanation!
The orbitals also have different energies inside energy levels
All the electrons can move around.