1. 10/31/14
Today I will define the characteristics of a wave and compare the major regions of the electromagnetic spectrum.
Warm Up – What are the three types of energy we discussed in a previous chapter?

2. Chapter 4 – Electron Configuration

3. Radiant Energy
Light
Properties of both a particle and a wave!
We call this the dual nature of light

4. Nature of Light
Einstein proposed that: The particles or “packets” of light are called “PHOTONS”
Light travels in waves
We use four main terms to define the characteristics of these waves

5. Waves Wavelength (λ) units in meters (often nm)
distance between successive crests (high points) or troughs (low points)
units in meters (often nm)

6. Waves Amplitude (A) height of the wave (from origin to crest)
units in meters
Brightness of light

7. Waves Frequency () – how fast the wave is moving up and down
how many waves pass a fixed point in a given time
units in per seconds or Hertz (Hz)
Notice that the higher the frequency, the shorter the wavelength.

8. Wave Speed (c) – how fast light moves through space
All light (regardless of wavelength) moves through space at the speed of light!
c = 3.00 x 108 m/s

9. Electromagnetic Spectrum
Visible light is only one form of radiant energy.
7 Parts of the Electromagnetic Spectrum
Longest wavelength to Shortest:
Radio
Microwaves
Infrared
Visible Light
U.V. Light
X-Rays
Gamma-Rays

10. V I B G Y O R

11. Homework
4-1 Review & Reinforcement

12. 11/3/14
Today I will
determine the mathematical relationships in the electromagnetic spectrum
solve for wavelength and frequency
Warm up –
Draw a wave and label the origin, amplitude, wavelength, crest, and trough.
Now draw a wave with a larger frequency. Does it have a longer or shorter wavelength?

13. Wave Relationships
We said that wavelength and frequency are inversely related
When one goes up, the other goes down.

14. Wave Relationships Mathematical Relationship
Notice that speed of light is in m/s so wavelength must be in meters and frequency must be hertz!
Wavelength is often given in nm.
You must convert first! 1m = 1x109nm
So 415 nm = ? m

15. Wave Relationships
A helium-neon laser produces a red light whose wavelength is 633 nm. Calculate the frequency of the radiation.
 = 4.74 x 1014 Hz
Do you remember what to do when the variable is on the bottom???

16. Wave Relationships
What is the wavelength of U.V light that has a frequency of 4.50 x 1016 Hz?
Answer: 6.67 x m

17. Wave Relationships
What is the wavelength of light, that has a frequency of 6.00 x kHz?
Convert from kHz to Hz
Answer: 5.00 x m

18. Homework
Electromagnetic Radiation – Frequency & Wavelength Calculations

19. 11/5/14 Today I will Determine relationships between energy and waves
Solve for energy
Warm Up –
What is the frequency of light that has a wavelength of 245 nm?

20. Wave Relationships - energy
The higher the frequency, the more energy.

21. Wave Relationships - energy
Maxwell Plank
E=h
Where
E is energy
h is plank’s constant
 is frequency
Plank’s constant = x J·s
Energy and frequency are directly proportional
When one goes up, the other goes up!

22. Wave Relationships - energy
What is the energy of U.V. light with a frequency of 4.50 x Hz?
E = h 
E = (6.626 x 10 –34 Js)(4.50 x Hz)
E = 2.98 x 10 –17 J

23. Wave Relationships - energy
Determine the energy of light that has a wavelength of 450nm.
Do we know anything about wavelength and energy?
E=h

24. Wave Relationships - energy
Determine the energy of light that has a wavelength of 450nm.

25. Homework
Electromagnetic Radiation - Energy Calculations

26. 11/6/14
Today I will review the electromagnetic spectrum & calculations
Warm Up – How much energy is in a wave whose wavelength is 432 nm?

27. 11/7/14 Today I will explain the quantization of light
Warm Up – What is the relationship (in words) between:
Wavelength and frequency
Frequency and energy
Wavelength and energy

28. Quantization of Energy
Scientists were confused by the observation that the wavelength of the radiation changed with temperature

29. Quantization of Energy
Maxwell Planck – proposed that there are fixed amounts of energy that an object emits.
He called each of these pieces of energy a quantum.
Quantum means a fixed amount

30. Quantization of Energy
What is the difference between quantized and continuous?
Going up a ramp going up steps
If energy is in quanta, why aren’t we aware of it?
Quanta of energy are very, very small. Too small to notice.

31. Quantization of Energy
Not all light will have the same effects.
We are surrounded by radio waves every day and they do not harm us, but gamma rays are very dangerous.
Very energetic light is dangerous (UV, X-rays, gamma)
Low frequency light does not have enough energy to be dangerous.

32. Photoelectric Effect
Electrons can be ejected from the surface of a metal when light shines on it.
Not all lights will cause this to happen on all metals.
For example. Red light is not able to release electrons from sodium, but violet light is.

33. Photoelectric Effect
Einstein proposed the idea of photons of light. Some photons have enough energy to release the electrons of a given metal and some do not.
Compton
said that Light packets collide with the electrons to release them.
You can’t add the energy up as more light hits.

34. Photoelectric cells
One everyday use of the photoelectric effect is photoelectric cells.
Automatic lights
Automatic doors
When the light shines, electrons are emitted and a current is produced. When that light is broken (by darkness or a body crossing it), the current is shut off and the desired effect is triggered.

35. Homework
4-2 Review & Reinforcement

36. 11/10/14 Objective – to describe the Bohr model of the atom
Warm Up – Which light will have a greater chance of producing the photoelectric effect: radio waves or ultraviolet waves? Why?

37. Bright Line spectra
Scientists observed another phenomenon they could not explain.
When certain salts are heated, they produce colorful light

38. Bright Line spectra
These distinct colors are known as that element’s continuous spectrum
Lithium – red
Potassium – purplish pink
Sodium - yellow

39. Bright Line spectra Sometimes difficult to differentiate
When these lights are passed through a prism, they separate into distinct lines.
These combinations of lines are called line spectra
Bright Line spectra are unique to certain elements

40. Bohr Model
Neils Bohr – 1911
Wanted to explain line spectra
Electrons orbit the nucleus in defined energy levels
Gave each level a quantum number (n)

41. Bohr Model
When energy is absorbed, an electron can “jump” to a higher energy level
Energy absorbed

42. Bohr Model These electrons then “fall” back down and release energy.
This energy is released as light.

43. Bohr Model
Ground State- when an electron is on the energy level at which it normally resides
Excited State – when an electron is on a higher energy level than usual

44. Bohr Model Excited state Ground state
Energy absorbed
Ground state
Energy released as light
In this example, n=1 is the ground state and n=2 is the excited state

45. Bohr Model

47. Bohr Model
The amount of energy released is determined by how that atom’s electrons jump and fall.
The wavelengths (or colors) of light are determined by those energies!

48. Bohr Model
Bohr’s model and the calculations of energy and wavelength work very well for Hydrogen.
However, with larger atoms, the model works only approximately.
Ultimately, this model does not work, but it was a good starting point for scientists, especially in thinking of distinct energy levels!

49. LAB - Bohr Model
Line spectra are particularly useful with metallic elements
Electrons in metals are more loosely bound and better able to jump and fall – metallic bonding
Since we can’t burn metals, we will use metallic salts

50. Lighting a Bunsen Burner!

51. Lighting a Bunsen Burner

52. Lighting a Bunsen Burner
CLOSED = PERPENDICULAR

53. Lighting a Bunsen Burner
OPEN = PARALLEL

54. Homework
Read the lab & do the pre-lab questions on a separate sheet of paper to turn in before tomorrow’s lab.

55. 11/11/14
LAB - Today I will observe line spectra and describe where an electron might exist in an atom.

56. 11/12/14
Today I will discuss key ideas that lead us to the modern day theory of an atom.
Warm Up – Why do different salts burn different colors?

57. Matter Waves When light travels through space, it acts like a wave.
When light interacts with matter, it acts like a particle.
Could all matter exhibit this dual nature?

58. Matter Waves Louis deBroglier - 1924 Mathematical manipulation
Using Einstein and Plank’s work to relate wavelength and mass
All matter exhibits a wavelength-like energy
Matter Waves – wavelike behavior of particles

59. Matter Waves
Why don’t I notice the wave of a golf ball traveling through the air?
The wavelength is too small to notice.
In order to have an observable wavelength, the mass must be very very small (like an electron).
Electron microscopes utilize the wave nature of electrons

60. Heisenberg’s Uncertainty Principle
Werner Heisenberg – 1927
Heisenberg Uncertainty principle – the position and momentum of an object cannot be known simultaneously.
We can see things because light bounces off of them and hits our eyes.
When light hits something, it changes it!

61. Heisenberg’s Uncertainty Principle
Measuring something inherently changes it!
Obviously with large objects the effect is too small to notice, but with tiny objects such as electrons, this is huge!

62. Quantum Mechanics
Because we cannot measure an electron’s exact position, we cannot know where exactly it is in an atom.
Since we don’t know exactly, we talk about electrons in terms of probability.

63. Quantum Mechanics
Electron Cloud – a visual representation of the probability of the location of an electron.
Electron Density – the density of the electron cloud at a certain point.

64. Quantum Mechanics
Quantum-Mechanical Model – a new model of the atom that better explains its properties by treating the electron as a wave and quantizing its energy.

65. Homework
4-3 Review & Reinforcement

66. 11/13/14
Today I will determine the probable location of each electron in an atom
Warm Up – Why do we use an electron cloud to represent an electron?

67. Crazy Landlord 7 6 5 4 3 2 1 f d f p d f s p d f s p d s p d s p s p s

68. Quantum Mechanics
Quantum-Mechanical Model – a new model of the atom that better explains its properties by treating the electron as a wave and quantizing its energy.

69. Quantum Mechanics
Remember that different electrons have different energy levels.
Atomic Orbitals – a region around the nucleus where an electron with a given energy is likely to be found.

70. Electron Configurations
We talked about the rules of a strange landlord. Those rules are actually real scientific principles!
Aufbau Principle – Electrons are added one at a time to the lowest energy level available.
Pauli Exclusion Principle – An orbital can hold a maximum of 2 electrons. These electrons must have opposite spins (clockwise & counter-clockwise).
Hund’s Rule – Electrons fill orbitals so that a maximum number of unpaired electrons exist. (Ladies first!)

71. Quantum Mechanics
The orbitals replace Bohr’s paths, and we describe them using four parts
1. Principal Quantum Number (n)
Like Bohr’s original quantum numbers (n=1,2,3…)
Principle energy levels
Size of each orbital. Bigger as they get further away from the nucleus (distance from the nucleus.)
Has values of n =1 to 7, 1 is the closest; 7 is the farthest from the nucleus.
“Apartment Buildings”

72. Primary Quantum Number (n)
Quantum Mecahnics
Primary Quantum Number (n)
# of sublevels 1
1 (s) 2
2 (s,p) 3
3 (s,p,d) 4
4 (s,p,d,f) 5 6 7
2. Energy Sublevels (l)
“Floors”
s, p, d, and f
Each principle level can hold a different number of sublevels
Each type of sublevel has a different shape

73. Quantum Mechanics
Like electron clouds of all different shapes and sizes

74. Primary Quantum Number (n)
Quantum Mechanics
3. Each sublevel is broken up into orbitals
“Apartments”
s has 1 orbital
p has 3 orbitals
d has 5 orbitals
f has 7 orbitals
Primary Quantum Number (n)
# of sublevels
orbitals 1
1 (s)
1(s) 2
2 (s,p)
1(s) + 3(p) = 4 3
3 (s,p,d)
1(s) + 3(p) + 5(d) = 9 4
4 (s,p,d,f)
1(s) + 3(p) + 5(d) + 7(f) = 16 5 6 7

75. Quantum Mechanics 3. Each sublevel is broken up into orbitals
“Apartments”
s has 1 orbital
p has 3 orbitals
d has 5 orbitals
f has 7 orbitals

76. Quantum Mechanics Pauli Exclusion Principle
4. Each orbital can hold up to two electrons.
How can two negative charges occupy the same space?
They must have opposite spins. Clockwise and Counter-clockwise.
“Boy and Girl”
Pauli Exclusion Principle

77. Electron Configurations
How many orbitals on an “s” sublevel?
How many electrons on an “s” sublevel?
How many orbitals on a “p” sublevel?
How many electrons on a “p” sublevel?
How many orbitals on a “d” sublevel?
How many electrons on a “d” sublevel?
How many orbitals on an “f” sublevel?
How many electrons on an “f” sublevel? 1 2 3 6 5 10 7 14

78. Homework
4-4 Review & Reinforcement

79. 11/14/14 Today I will draw orbital diagrams
Warm Up – Explain the distribution of electrons in the n=3 energy level

80. Electron Filling Order
Each orbital gets labeled with the principle quantum number and the sublevel.
Remember:
Principle quantum numbers are the apartment buildings! n = 1-7
Sublevels are the floors  s, p, d, f

81. Electron Filling Order
Also Remember, the crazy landlord had a special way of filling the apartments!
Based on the Aufbau Principle of lowest energy first:
1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s 5f 6d 7p

82. Electron Filling Order
Aufbau Principle
1s 2s 3s 4s 5s 6s 7s 2p 3p 4p 5p 6p 7p 3d 4d 5d 6d 7d 4f 5f 6f 7f

83. Orbital Diagrams
Orbital Diagram – Representation of an atom using boxes and arrows to show the distribution of electrons in an atom

84. Orbital Diagrams 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s 5f 6d

85. Orbital Diagrams First Ask: How many electrons do you need?
Add electrons based on the filling order, paying attention to Hund’s rule!
Draw the orbital diagram for oxygen.
How many electrons?

86. Orbital Diagrams How many unpaired electrons? 1s 2s 2p 3s 3p 4s 3d 4p4f 5d 6p 7s 5f 6d 7p
How many unpaired electrons?

87. Orbital Diagrams
Draw the orbital diagram for Nickel. How many unpaired electrons does Nickel have?

88. Homework
Orbital Diagrams Worksheet

89. 11/17/14
Today I will write electron configurations
Warm Up – Draw the orbital diagram for the Tungsten atom

90. Tungsten (W) = 74 electrons2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d

91. Orbital Diagrams Pros Visual representation of electrons
Easily shows the number of unpaired electrons
Cons
Large, takes up space
Time Consuming

92. Electron Configuration
Electron Configuration – the distribution of electrons among the orbitals.
Shortened version of the orbital diagram.
How many electrons?
Write the principle quantum number and the orbital name in filling order.
Write the number of electrons in each orbital as a superscript

93. Electron Configuration
Sn –
Electron Configuration 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 2

94. Electron Configuration
How many electrons can an “s”sublevel hold? 2
How many electrons can a “p” sublevel hold? 6
How many electrons can a “d” sublevel hold? 10
How many electrons can an “f” sublevel hold? 14

95. Electron Configuration
Rules:
How many electrons?
Using the filling order, add the maximum number of electrons (as superscripts) to the orbitals until you reach the point where the maximum would be too many.
Partially fill the final orbital to the number of electrons needed.

96. Electron Configuration
Write the electron configuration for Gold. How many unpaired electrons?
Write the electron configuration for Strontium. How many unpaired electrons?

97. Electron Configuration - Ions
What is different about ions?
They have a different number of electrons.
When counting up the total number of electrons, make sure to take the charge into account.
The rest of the process is the same.

98. Electron Configuration - Ions
Draw the orbital diagram for Br-1
Write the electron configuration for Ba+2
Write the electron configuration for N-3 1s 2s 2p 3s 3p 4s 3d 4p

99. Homework
4-5 Practice Problems

100. 11/18/14
Today I will draw orbital diagrams and write electron configurations for exceptions.
Write the electron configuration for Tellurium.

101. Exceptions to Aufbau Chromium’s column Cr – 1s22s22p63s23p64s23d4 ??

102. Exceptions to Aufbau Copper’s column Cu – 1s22s22p63s23p64s23d9 ??

103. Write the electron configurations for:
Se-2
Hg+2 Mo
Cs+1 Au

104. Se s22s22p63s23p64s23d104p6
Hg s22s22p63s23p64s23d104p65s24d105p66s24f145d8
Mo 1s22s22p63s23p64s23d104p65s14d5
Cs+1 1s22s22p63s23p64s23d104p65s24d105p6
Au 1s22s22p63s23p64s23d104p65s24d105p66s14f145d10

105. Homework
Exceptions Worksheet

106. 11/19/14
Objective – to write electron configurations based on the periodic table.
Warm Up – Write the electron configuration for Hf

108. n
“p” 1 2
“d” 3 4 5 6 7
“f”

109. 1 2
d = n-1 3 4 5 6 7
f = n-2

110. d = n-1 f = n-2 Write the electron configuration for Os 1s2 2s2 2p63 4 5 6 7
f = n-2

111. d = n-1 f = n-2 Write the electron configuration for Fr+1 1s2 2s2 2p63 4 5 6 7
f = n-2

112. Write the electron configuration for:
Pr (# 59)
Rf (#104)
Pb+4(#82)

113. Homework
Orbital Diagram and Electron Configuration Review

114. Today I will review chapter 4
11/20/14
Today I will review chapter 4
Warm Up
If light has an energy of 1.2 x J, what is its wavelength?
c = 3.00 x 108 m/s h=6.626 x J·s

115. Homework
Chapter 4 Review