1. Pre-IB/Pre-AP CHEMISTRY
Chapter 4 – Arrangement of Electrons in Atoms

2. Section 1 Objectives
Be able to define: electromagnetic radiation, electromagnetic spectrum, wavelength, amplitude, frequency, photoelectric effect, quantum(pl. quanta), photon, ground state, excited state, line emission spectrum, continuous spectrum, energy level.

3. Section 1 Objectives
Be able to explain the mathematical relationship between speed, wavelength, and frequency of a wave.
Be able to describe what is meant by the wave-particle duality of light.
Be able to discuss how the photoelectric effect and the line emission spectrum of hydrogen lead to the development of the atomic model.

4. Section 1 Objectives
Be able to describe the Bohr model of the atom.

6. Wave
A wave is a method of transferring energy. This transfer does not require matter as a medium.

7. Wave
Some waves travel through matter (sound, water waves, etc.).

8. Wave
Some waves do not require matter and can travel through empty space (light).

9. Wave Properties
Waves can be described by their wavelength, amplitude, and frequency.

10. Wavelength A crest is the highest point on a wave.
A trough is the lowest point on a wave.
Crest
Trough

11. Wavelength
Wavelength is simply the length of a wave. It is the distance between two crests or two troughs.
Wavelength is measured in m, mm, or nm.
Wavelength
Crest
Trough

12. Amplitude
Amplitude is simply the height of a wave. It is the distance between the crest and trough of a wave.
Amplitude is measured in units of distance.
Amplitude

13. Frequency
Frequency is the number of waves passing a given point in a given time.
Frequency describes the energyof a wave.

14. Frequency
Frequency describes the energy of a wave: the higher the frequency, the greater the energyof that wave.

15. Frequency
Frequency is measured in hertz or cycles per secondor vibrations per second or 1/sec or sec-1 - they all mean the same thing.

16. Frequency
As the wavelength increases, frequency decreases. This is called an inverse relationship.

17. Wave Properties
Wavelength and amplitude give waves their distinctive properties. For example, the loudness of a sound wave is its amplitude, the color of visible light is its wavelength.

18. Types of waves
Electromagnetic waves do not require a medium or matter in order to travel. Light is an example.

19. Light Light is an electromagnetic wave.
Visible light is a small part of the electromagnetic spectrum that humans are able to see.

20. Light
The electromagnetic spectrum consists of different kinds of light of different wavelengths.

21. EM Spectrum

22. EM Spectrum

23. EM Spectrum

24. EM Spectrum

25. EM Spectrum

26. EM Spectrum

27. EM Spectrum

28. Light Interactions
White light is light consisting of all colorsof visible light. These colors are visible in a rainbow or through a prism.

29. Velocity
The velocity of a wave is a product of its frequency and wavelength.
v= fl
v = velocity
f = frequency
l = wavelength

30. Velocity
The velocity of light through a vacuum(c) is about 3.0 x 108 m/sec. It is slightly slower through matter.

31. Photoelectric Effect
Photoelectric effect refers to the emission of electrons from a metal when light shines on the metal.

32. Photoelectric Effect
It was found that light of a certain frequency would cause electrons to be emitted by a particular metal. Light below that frequency had no effect.

33. Emission Spectra
If an object becomes hot enough it will begin to emit light.

34. Emission Spectra
Max Planck suggested that hot objects emit light in specific amounts called quanta (sing. quantum).

35. Emission Spectra Equantum= hf
Planck showed the relationship between a quantum of energy and the frequency of the radiation.
Equantum= hf
Equantum= energy of a quantum in joules
h = Planck’s constant
f = frequency

36. Wave-Particle Duality
Einstein later said that light had a dual nature – it behaved as both a particle and a wave.

38. Wave-Particle Duality
Each particle of light, Einstein said, carries a particular quantum of energy.

39. Wave-Particle Duality
Einstein called the “particles” of light photons which had zero mass and carried a quantum of energy. The energy is described as:
Ephoton= hf

40. Photoelectric Effect
Einstein explained photoelectric effect by saying in order for an electron to be ejected from a metal, the photon striking it must have enough energy to eject it.

41. Attraction
Different metals have stronger attraction for their electrons than other. Therefore, some must absorb more energy than others to emit electrons.

42. Ground State
The lowest energy state of an atom is called its ground state.

43. Excited State
When a current is passed through a gas at low pressure, the atoms become “excited.”

44. Excited State
Atoms in an excited state have a higher potential energy than their ground state.

45. Excited State
An “excited” atom will return to its ground state by releasing energy in the form of electromagnetic radiation.

46. Emission Spectra
Elements will emit radiation of certain frequencies. This reflects the energy states of its electrons and is called a bright-line or emission spectrum.

47. Emission Spectra
The emission spectrum of an element is like its “fingerprint”.
Sodium
Helium
Mercury

48. Energy Levels
Studying the emission spectrum of hydrogen lead Niels Bohr to the idea of energy levels.

49. Energy Levels
The spectrum Bohr and others observed was the result of excited electrons releasing photons as they returned to their ground states.

50. Energy Levels
The difference in the energy of photons was reflected in the different frequencies of light they observed.

51. Section 2 Objectives
Be able to define: diffraction, interference, Heisenberg Uncertainty Principle, Quantum Theory, quantum numbers, principal quantum number, angular momentum quantum number, magnetic quantum number, spin quantum number.
Be able to distinguish between the Bohr model and the quantum model of the atom.

52. Section 2 Objectives
Be able to explain how the Heisenberg Uncertainty Principle and the Schroedinger Wave Equation led to the idea of atomic orbitals.
Be able to list the four quantum numbers that describe each electron in an atom.

53. Section 2 Objectives
Be able to relate the number of sublevels corresponding to each of an atom’s main energy levels, the number of orbitals per sublevel, and the number of orbitals per main energy level.

55. Electrons as Waves
French scientist Louis De Broglie demonstrated that electrons had a dual nature also.

56. Electrons as Waves E = hf
De Broglie showed that electrons behaved as waves confined to the atom. The energy of those electrons could be found like that of waves:
E = hf

57. Electrons as Waves
Electron beams were shown to exhibit the wave properties of diffraction and interference.

58. Heisenberg Uncertainty
Werner Heisenberg tried to find the location and velocity of electrons in the atom.

59. Heisenberg Uncertainty
Heisenberg found that it is impossible to simultaneously determine the position and velocity of an electron or any other particle (The Heisenberg Uncertainty Principle).

60. Schrödinger Wave Equation
Erwin Schrödinger said that electrons had a dual nature(like light) and treated them as waves.

61. Quantum Theory
Schrödinger’s wave equation and Heisenberg’s Uncertainty Principle laid the foundation of modern quantum theory.

63. Quantum Theory
Quantum theory describes mathematically the wave properties of electrons and other very small particles.

64. Quantum Theory
According to the Bohr model we should be able to predict the location and velocity of an electron at any time.

65. Quantum Theory
Quantum theory disagrees with the Bohr model and says that electrons can be found in regions of high probability but cannot be pinpointed.

66. Orbitals
Quantum theory describes electrons as inhabiting a three-dimensional region around the nucleus that indicates their probable locations. These regions are called orbitals.

68. s and p orbitals

71. Orbitals
Scientists use quantum numbers to describe orbitals. These numbers describe the properties of the orbitals and the electrons which occupy them.

72. Atomic Orbitals and Quantum Numbers
Section 2 The Quantum Model of the Atom
Chapter 4
Atomic Orbitals and Quantum Numbers
Principal Quantum Number (n) - indicates the main energy level occupied by the electron.
(n) values are positive integers only.
- As (n) increases, the electron’s energy and its average distance from the nucleus increase
- Electrons are sometimes said to be in the same electron shell.

73. Atomic Orbitals and Quantum Numbers, continued
Section 2 The Quantum Model of the Atom
Chapter 4
Atomic Orbitals and Quantum Numbers, continued
Angular Momentum Quantum Number (l) – indicates the shape of the orbital
Orbitals (except for first main energy level) of different shapes (sublevels) exist for a given value of n. Values are allowed to be zero and all positive integers less than or equal to n – 1
l = (n – 1)
s – spherical orbitals
p – dumbbell shape
d – more complex .

74. Quantum Numbers and Orbitals
Section 2 The Quantum Model of the Atom
Chapter 4
Quantum Numbers and Orbitals
Angular Momentum Quantum Number (l)
1st energy level – only one sublevel possible [n=1 (s orbital)]
2nd energy level – two sublevels [n = 1 (s) & n=2 (p)]
3rd energy level – three sublevels [n=1 (s), n=2 (p) & n=3(d)]
4th energy level – four sublevels [n=1 (s), n=2 (p), n=3(d) & n=4 (f)]

75. Shapes of s, p, and d Orbitals
Section 2 The Quantum Model of the Atom
Chapter 4
Shapes of s, p, and d Orbitals

76. Atomic Orbitals and Quantum Numbers, continued
Section 2 The Quantum Model of the Atom
Chapter 4
Atomic Orbitals and Quantum Numbers, continued
Magnetic Quantum Number (m) – indicates the orientation of an orbital around the nucleus
m = (-l to +l)
- s is spherical and it is centered about the nucleus, it only has one possible orientation (m=0, s)
- p orbitals extend along the x, y, or z axis of the coordinate system – there are three p orbitals in each p sublevel (px – m=-1, py – m=0, and pz - m=+1)
- 5 d orbitals in each d sublevel (m=-2, m=-1, m=0, m=+1, m=+2)
7 f orbitals in each f sublevel

77. Chapter 4 Quantum Numbers Section 2 The Quantum Model of the Atom
n = energy level (Rows on PT)
l = shape (l = n-1) (Sublevel)
m = orientation (m= -l to +l)(Orbital) 1
l = 0 (s)
m = 0 2
m = -1, 0, +1
l= 1 (p) 3
m = -2, -1, 0, +1, +2
l = 2 (d) 4
m = -3, -2, -1, 0, +1, +2, +3
l = 3 (f)

78. Electrons in Energy Levels and Sublevels
Section 2 The Quantum Model of the Atom
Chapter 4
Electrons in Energy Levels and Sublevels
Principal
Energy
Level
(n=?)
Orbital Types
(sublevels available)
# of orbitals per level
Total # of orbitals per P.E.L.
(n2)
# of e’s per orbital type
Total # of e’s per P.E.L.
(2n2) 1 s 2 p 6 3 5 4 14

79. Pg. 110

80. Section 3 Objectives
Be able to define: Aufbau Principle, Pauli Exclusion Principle, Hund’s rule.
Be able to list the number of electrons it takes to fill each energy level.
Be able to describe the electron configurations for the atoms of any element using orbital notation, electron configuration notation, and, when appropriate, noble-gas notation.

82. Chapter 4 Preview Lesson Starter Objectives Electron Configurations
Section 3 Electron Configurations
Chapter 4
Preview
Lesson Starter
Objectives
Electron Configurations
Rules Governing Electron Configurations
Representing Electron Configurations
Elements of the Second Period
Elements of the Third Period
Elements of the Fourth Period
Elements of the Fifth Period

83. Section 3 Electron Configurations
Chapter 4
Objectives
List the total number of electrons needed to fully occupy each main energy level.
State the Aufbau principle, the Pauli exclusion principle, and Hund’s rule.
Describe the electron configurations for the atoms of any element using orbital notation, electron-configuration notation, and, when appropriate, noble-gas notation.

84. Chapter 4 Lesson Starter
Section 3 Electron Configurations
Chapter 4
Lesson Starter
The electron configuration of carbon is 1s22s22p2.
An electron configuration describes the arrangement of electrons in an atom.
The integers indicate the main energy level of each orbital occupied by electrons.
The letters indicate the shape of the occupied orbitals.
The superscripts identify the number of electrons in each sublevel.

85. Electron Configurations
Section 3 Electron Configurations
Chapter 4
Electron Configurations
An electron configuration describes the arrangement of electrons in an atom. The integers indicate the main energy level of each orbital occupied by electrons. The letter indicate the shape of the occupied orbitals. The superscripts identify the number of electrons in each sublevel. The arrangement of electrons in an atom is known as the atom’s electron configuration. The lowest-energy arrangement of the electrons for each element is called the element’s ground-state electron configuration

86. Relative Energies of Orbitals
Section 3 Electron Configurations
Chapter 4
Relative Energies of Orbitals

87. Electron Configuration
Section 3 Electron Configurations
Chapter 4
Electron Configuration
The bottom energy level is level 1 - it has the lowest energy.
Each box represents an orbital.
You can see that there is 1 orbital for an s subshell
There are 3 orbitals for a p subshell, 5 for a d, and 7 for an f subshell.
Each orbital can hold 2 electrons.
Therefore, the s subshell can hold 2 electrons; the p can hold 6; the d can hold 10; and the f can hold 14.
The first energy level can hold 2 electrons (1s = 2), the second energy level can hold 8 electrons (2s2p = 2 + 6), the third energy level can hold 18 electrons (3s3p3d = ), and the forth energy level can hold 32 (4s4p4d4f = ).

88. Rules Governing Electron Configurations
Section 3 Electron Configurations
Chapter 4
Rules Governing Electron Configurations
According to the Aufbau principle, an electron occupies the lowest-energy orbital that can receive it.
According to the Pauli exclusion principle, no two electrons in the same atom can have the same set of four quantum numbers.

89. Rules Governing Electron Configurations, continued
Section 3 Electron Configurations
Chapter 4
Rules Governing Electron Configurations, continued
According to Hund’s rule, orbitals of equal energy are each occupied by one electron before any orbital is occupied by a second electron, and all electrons in singly occupied orbitals must have the same spin state.

90. Chapter 4 Apartment Analogy:
Section 3 Electron Configurations
Chapter 4
Apartment Analogy:
Imagine you are the landlord of a very strange apartment building. Your job is to fill the apartments in the building in the most efficient way possible. You are required by the owner of the building to fill the rooms in a certain way. The rules you have to follow are as strange as the building because quantum mechanics is not like anything you might have expected. The rules are summarized in the table below.

91. Chapter 4 Apartment Analogy:
Section 3 Electron Configurations
Chapter 4
Apartment Analogy:
In the building the different floors are like the different energy levels (or shells) in an atom. The energy levels are numbered starting from one, just like the floors in an apartment. Each room corresponds to one orbital. The rooms have a capacity of two electrons (two people) each. In each room only a man and a woman may be paired together. In the strange world of quantum mechanics there are no same-gender roommates.

92. Chapter 4 Apartment Analogy: Section 3 Electron Configurations
Apartment Rules
Electron Rules
From the Bottom Up: Rooms must be filled from the ground floor up. Fill the one room on the first floor before starting to put new tenants on the second floor. Then fill the s room before the p rooms. At higher floors the order might change a bit.
Aufbau Principle: the electrons fill the available orbitals from lowest energy to highest energy. In the ground state all the electrons are in the lowest possible energy level.
Singles First: the owner of the building wants to have the tenants spread out as much as possible. For that reason singles are placed in rooms before couples. If couples must be placed into a room then all of the other rooms on that floor must already have a single in them.
Hund’s Rule: The electrons must be placed into the orbitals in such a way that no pairs are put together unless absolutely necessary. That is, single electrons must be placed into boxes first and then paired up if necessary.
Opposite Gender Only: When two people are placed in a room they must be of opposite genders. No men may room together and no women may room together. This is an arbitrary rule on the part of the owners: in a just world we wouldn’t have to follow it. But quantum mechanics has nothing to do with justice.
Pauli Exclusion Principle: Electrons come in two varieties based on the direction they are ‘spinning’. There is an Up spin and a Down spin. Up and Down spins are always paired together and Up-Up or Down-Down combinations are not allowed. No two electrons can ever be in the same place at the same time.

93. Chapter 4 Electron Filling: Section 3 Electron Configurations
The rule is to start at the beginning of each arrow, and then follow it all of the way to the end, filling in the sublevels that it passes through.  In other words, the order for filling in the sublevels becomes; 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d,7p.

94. Blocks on the Periodic Table:
Section 3 Electron Configurations
Chapter 4
Blocks on the Periodic Table:

95. Section 3 Electron Configurations
Chapter 4
Activity #1:
On the blank periodic table you were given, please color/outline the s, p, d and f blocks.
(See example)

96. Writing electron configurations of a neutral atom:
Section 3 Electron Configurations
Chapter 4
Writing electron configurations of a neutral atom:
In a neutral atom, the number of electrons equals the number of protons of the atom (ATOMIC NUMBER).
When the electrons fill the orbitals, they occupy the lowest energy orbitals that are available.

97. Writing electron configurations of a neutral atom:
Section 3 Electron Configurations
Chapter 4
Writing electron configurations of a neutral atom:
For example:
Hydrogen is atomic number 1 (has 1 proton)
The one electron that it has occupies the lowest orbital, 1s
Electron configuration = 1s1
Helium has 2 protons
Electron configuration = 1s2
Boron has 5 protons
Electron configuration = 1s22s22p1
Calcium has 20 electrons
Electron configuration of 1s2 2s2 2p6 3s2 3p6 4s2

98. Activity #2: (Groups of 2)
Section 3 Electron Configurations
Chapter 4
Activity #2: (Groups of 2)
On the periodic table you just broke into s, p, d and f blocks, write the orbital being completed for each element
For example:
Hydrogen’s block should say 1s1
Rubidium’s block should say 5s1

99. Section 3 Electron Configurations
Chapter 4

100. Chapter 4: Section 3 Continued
Section 3 Electron Configurations
Chapter 4
Chapter 4: Section 3 Continued
Objectives:
List the total number of electrons needed to fully occupy each main energy level.
State the Aufbau principle, the Pauli exclusion principle, and Hund’s rule.
Describe the electron configurations for the atoms of any element using orbital notation, electron-configuration notation, and, when appropriate, noble-gas notation.

101. Writing electron configurations of a neutral atom:
Section 3 Electron Configurations
Chapter 4
Writing electron configurations of a neutral atom:
REVIEW:
Hydrogen is atomic number 1 (has 1 proton)
The one electron that it has occupies the lowest orbital, 1s
Electron configuration = 1s1
Helium has 2 protons
Electron configuration = 1s2
Boron has 5 protons
Electron configuration = 1s22s22p1
Calcium has 20 electrons
Electron configuration of 1s2 2s2 2p6 3s2 3p6 4s2

102. Section 3 Electron Configurations
Chapter 4
Keep In Mind -
Aufbau principle - an electron occupies the lowest- energy orbital that can receive it.
Pauli exclusion principle - no two electrons in the same atom can have the same set of four quantum numbers.
Hund’s rule - orbitals of equal energy are each occupied by one electron before any orbital is occupied by a second electron, and all electrons in singly occupied orbitals must have the same spin state.

103. Representing Electron Configurations
Section 3 Electron Configurations
Chapter 4
Representing Electron Configurations
Orbital Notation
An unoccupied orbital is represented by a line, with the orbital’s name written underneath the line.
An orbital containing one electron is represented as:

104. Representing Electron Configurations
Section 3 Electron Configurations
Chapter 4
Representing Electron Configurations
Orbital Notation
An orbital containing two electrons is represented as:
The lines are labeled with the principle quantum number and sublevel letter. For example, the orbital notation for helium is written as follows: 1s He

105. Drawing Orbital Notation
Section 3 Electron Configurations
Chapter 4
Drawing Orbital Notation
Click below to watch the Visual Concept.
Visual Concept

106. Representing Electron Configurations, continued
Section 3 Electron Configurations
Chapter 4
Representing Electron Configurations, continued
Electron-Configuration Notation
Electron-configuration notation eliminates the lines and arrows of orbital notation.
Instead, the number of electrons in a sublevel is shown by adding a superscript to the sublevel designation.
The helium configuration is represented by 1s2.
The superscript indicates that there are two electrons in helium’s 1s orbital.

107. Writing Electron Configurations
Section 3 Electron Configurations
Chapter 4
Writing Electron Configurations

108. Representing Electron Configurations, continued
Section 3 Electron Configurations
Chapter 4
Representing Electron Configurations, continued
Sample Problem A
The electron configuration of boron is 1s22s22p1. How many electrons are present in an atom of boron? What is the atomic number for boron? Write the orbital notation for boron.

109. Section 3 Electron Configurations
Chapter 4
Representing Electron Configurations, continued Sample Problem A Solution
The number of electrons in a boron atom is equal to the sum of the superscripts in its electron-configuration notation: = 5 electrons. The number of protons equals the number of electrons in a neutral atom. So we know that boron has 5 protons and thus has an atomic number of 5. To write the orbital notation, first draw the lines representing orbitals. 1s 2s 2p

110. Section 3 Electron Configurations
Chapter 4
Representing Electron Configurations, continued Sample Problem A Solution, continued
Next, add arrows showing the electron locations. The first two electrons occupy n = 1 energy level and fill the 1s orbital. 1s 2s 2p

111. Section 3 Electron Configurations
Chapter 4
Representing Electron Configurations, continued Sample Problem A Solution, continued
The next three electrons occupy the n = 2 main energy level. Two of these occupy the lower-energy 2s orbital. The third occupies a higher-energy p orbital. 1s 2s 2p

112. Drawing Bohr’s Models:
Section 3 Electron Configurations
Chapter 4
Drawing Bohr’s Models:

113. Drawing Bohr’s Models:
Section 3 Electron Configurations
Chapter 4
Drawing Bohr’s Models:
Draw the nucleus.
Write the number of neutrons and the number of protons in the nucleus.
Draw the first energy level.
Draw the electrons in the energy levels according to the rules below. Make sure you draw the electrons in pairs.
Keep track of how many electrons are put in each level and the number of electrons left to use.

114. Rules for Energy Levels:
Section 3 Electron Configurations
Chapter 4
Rules for Energy Levels:
Level 1 (closest to the nucleus) can hold a maximum of 2e.
Level 2 can hold a max of 8e.
Level 3 can hold a max of 18e.
Level 4 can hold a max of 32e.
You must fill one level before going on to draw the next level! (Fill the lowest energy level 1st!)

115. Protons: 6 Neutrons: 6 Electrons: 6
Guided Practice
In order to draw Bohr models of these elements, you must first determine the number of protons, neutrons, and electrons. Once you have found this information, follow the directions to draw your model. 6 C
Carbon
12.011
Protons: 6 Neutrons: 6 Electrons: 6

116. Section 3 Electron Configurations
Chapter 4
Activity #1:
Complete the following activities for a Rubidium (Rb) atom.
Electron-configuration notation
Orbital Diagram
Bohr’s Model
I will be asking for volunteers to come work this on the board and EXPLAIN your answer!

117. Classwork for Today/Beginning of Period Tomorrow:
Section 3 Electron Configurations
Chapter 4
Classwork for Today/Beginning of Period Tomorrow:
Electron Configuration Resource/Practice Worksheet
Please complete as much of this as possible today…
I will give you a few minutes in class to finish this
tomorrow (due to Nine Weeks Exams). Once your time is up, I will call on people to write their answers on the board and EXPLAIN.

118. Section 3 Electron Configurations
Chapter 4
FINISHING CHAPTER 4

119. Elements of the Third Period
Section 3 Electron Configurations
Chapter 4
Elements of the Third Period
Noble-Gas Notation
The Group 18 elements (helium, neon, argon, krypton, xenon, and radon) are called the noble gases.
A noble-gas configuration refers to an outer main energy level occupied, in most cases, by eight electrons.

120. Orbital Notation for Three Noble Gases
Section 3 Electron Configurations
Chapter 4
Orbital Notation for Three Noble Gases

121. Chapter 4 Noble-Gas Notation Section 3 Electron Configurations
Click below to watch the Visual Concept.
Visual Concept

122. Chapter 4 Sample Problem B
Section 3 Electron Configurations
Chapter 4
Sample Problem B
a. Write both the complete electron-configuration notation and the noble-gas notation for iron, Fe.
b. How many electron-containing orbitals are in an atom of iron? How many of these orbitals are completely filled? How many unpaired electrons are there in an atom of iron? In which sublevel are the unpaired electrons located?

123. Sample Problem B Solution
Section 3 Electron Configurations
Chapter 4
Sample Problem B Solution
a. The complete electron-configuration notation of iron is 1s22s22p63s23p63d64s2. Iron’s noble-gas notation is [Ar]3d64s2.
b. An iron atom has 15 orbitals that contain electrons.
They consist of one 1s orbital, one 2s orbital, three 2p orbitals, one 3s orbital, three 3p orbitals, five 3d orbitals, and one 4s orbital.
Eleven of these orbitals are filled, and there are four unpaired electrons.
They are located in the 3d sublevel.
The notation 3d6 represents 3d

124. End of Chapter 4 Show