1. Electrons in Atoms

2. Basic Properties of Waves

3. Amplitude
Amplitude is the maximum distance the particles of the medium carrying the wave move away from their rest positions.

4. Amplitude
Greater amplitude corresponds with greater energy.

5. Wavelength
Wavelength (λ) is the distance between two corresponding parts of a wave.

6. Frequency
Frequency equals the number of complete waves that pass a given point in a certain amount of time.

7. Frequency and Wavelength
Frequency and wavelength are inversely related, which means that as one goes up the other goes down.

8. Frequency and Wavelength
Different wavelengths (and frequencies) of light correspond to different colors of light.

9. The Quantum Concept
In 1900, the German physicist Max Planck began searching for an explanation as he studied the light emitted from heated objects.

11. The Quantum Concept
Matter can gain or lose energy only in small, specific amounts called quanta.
That is, a quantum is the minimum amount of energy that can be gained or lost by an atom.

12. The Quantum Concept 12

13. While a beam of light has many wavelike characteristics, it also can be thought of as a stream of tiny particles, or bundles of energy, called photons.
Thus, a photon is a particle of electromagnetic radiation with no mass that carries a quantum of energy.

14. Energy and Frequency
Planck showed that the energy of radiation increases as the radiation’s frequency increases.
Energy and frequency are directly proportional.

15. Wave-Particle Duality of Electrons15

16. STOP HERE

17. The Bohr Model of the Atom
Niels Bohr produced a model of the hydrogen atom based on experimental observations.

18. The Bohr Model of the Atom
This model indicated that
An electron circles the nucleus only in fixed energy ranges called orbits. 18

19. The Bohr Model of the Atom
An electron can neither gain nor lose energy inside this orbit, but could move up or down to another orbit. 19

20. The Bohr Model of the Atom
The lowest energy orbit is closest to the nucleus. 20

21. Modern View The atom has two regions and is 3- dimensional.
The nucleus is at the center and contains the protons and neutrons.

22. Modern View
The electron cloud is the region where you might find an electron (the region of high probability) and most of the volume of an atom.

23. Wave-Particle Duality of Electrons
Recall that electrons have wave-like properties, although they are normally thought of as particles. 23

24. Particle Properties of Electrons
There is always an integer number of electrons orbiting the nucleus.
Electrons jump between orbitals (orbits) in a particle-like fashion. 24

25. Wave Properties of Electrons
The electrons do not orbit the nucleus in the sense of a planet orbiting the sun, but instead exist as standing waves. 25

26. Wave Properties of Electrons
The electrons are never in a single point location, although there is an area of high probability (the electron cloud) of interacting with the electron at a single point. 26

27. The Quantum Mechanical Model
The lowest allowable energy state of an atom is called its ground state.

28. The Quantum Mechanical Model
When an electron gains an amount of energy equivalent to the energy difference, it is said to be in an excited state.
The electron moves from its ground state to a higher energy level. 28

29. The Quantum Mechanical Model29

30. The Quantum Mechanical Model
When the electron moves to a lower energy level, it releases an amount of energy equal to the energy difference in these levels as electromagnetic radiation (emissions spectrum).

31. The Quantum Mechanical Model
This electromagnetic radiation is given off as photons. 31

33. The spectrum of light released from excited atoms of an element is called the emission spectrum of that element.

34. These are called line spectra.
Each is unique to an element.

36. The further electrons fall, the greater the energy.
Ultraviolet
Visible
Infrared
The further electrons fall, the greater the energy.
This results in a higher frequency because energy and frequency are directly proportional.

37. Question
1. Use the Chemistry Reference Tables to answer the following:
An electron falls from energy level 5 to energy level 3. What is the wavelength of the light emitted?
(1282 nm)

38. Question
An electron falls from energy level 6 to energy level 2. What is the wavelength of the light emitted?
(410 nm)

39. Question
An electron falls from energy level 3 to energy level 1. What type of electromagnetic radiation is emitted (infrared, visible or ultraviolet)?
(UV)

40. Question
An electron falls from energy level 4 to energy level 2. What type of electromagnetic radiation is emitted (infrared, visible or ultraviolet)?
(visible)

41. Question
An electron falls from energy level 5 to energy level 2. What color of visible light is emitted?
(blue)

42. Question
An electron falls from energy level 3 to energy level 2. What color of visible light is emitted?
(red)

43. The Quantum Mechanical Model
The space around the nucleus of an atom where the atom’s electrons are found is called the electron cloud.

44. A three-dimensional region around the nucleus called an atomic orbital describes the electron’s probable location.

45. Energy Levels
In general, electrons reside in principal energy levels.

46. Energy Levels
As the energy level number increases, the orbital becomes larger, the electron spends more time farther from the nucleus, and the atom’s energy level increases.

47. Sublevels Principal energy levels contain energy sublevels.
Principal energy level 1 consists of a single sublevel, principal energy level 2 consists of two sublevels, principal energy level 3 consists of three sublevels, and so on.

48. Sublevels Sublevels are labeled s, p, d, or f.
The s sublevel can hold 2 electrons, the p sublevel can hold 6 electrons, the d sublevel can hold 10 electrons, and the f sublevel can hold 14 electrons.

49. s block
p block
d block
f block

50. Orbitals Sublevels contain orbitals.
Each orbital may contain at most two electrons.

51. s orbitals One s orbital for every energy level The s orbital
is spherical
shaped.
Called the 1s, 2s, 3s, etc… orbitals

52. p orbitals Start at the second energy level
p orbitals reside along 3 different directions
p orbitals have 3 different dumbbell shapes

53. d orbitals d orbitals start at the third energy level
d orbitals have 5 different shapes

54. f orbitals f orbitals start at the fourth energy level
f orbitals have seven different shapes

55. f orbitals

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

57. s block
p block
d block
f block

58. s- block s2 s1 s2 Really have to include helium.
Helium has the properties of the noble gases.

59. Transition Metals - d block

60. The p-block p1 p2 p3 p4 p6 p5

61. f - block inner transition elements f1 f5 f2 f3 f4 f6 f7 f8 f9 f10 f11

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

63. f orbitals start filling at 4f1 2 3 4 5 6 7 4f 5f
f orbitals start filling at 4f

64. 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.

66. Electron Configurations
Electron configurations represent the way electrons are arranged in atoms.
Electrons enter the lowest energy level first.

67. Electron Configurations
2. Determine the electron configuration for phosphorus (P).
The atomic number of phosphorus is 15, so we need to account for 15 electrons.

68. Electron Configurations
2p6
3s2
3p3
12 electrons
2 electrons
15 electrons
4 electrons
10 electrons
The 3p sublevel can hold 6 electrons, but only 3 are needed to get the 15 electrons.

69. Electron Configurations
3. Determine the electron configuration for chromium (Cr).
The atomic number of chromium is 24, so we need to account for 24 electrons.

70. Electron Configurations
2p6
3s2
3p6
4s2
3d4
12 electrons
24 electrons
20 electrons
18 electrons
2 electrons
4 electrons
10 electrons
The 3d orbital does not need all 10 electrons to get the 24 electrons. 70

71. Question
4. Write the electron configuration for aluminum (Al).
(1s2 2s2 2p6 3s2 3p1)

72. Question
5. Write the electron configuration for neon (Ne).
(1s2 2s2 2p6)

73. Question
6. Write the electron configuration for calcium (Ca).
(1s2 2s2 2p6 3s2 3p6 4s2)

74. Question
7. Write the electron configuration for iron (Fe).
(1s2 2s2 2p6 3s2 3p6 4s2 3d6)

75. Question
8. Write the electron configuration for bromine (Br).
(1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p5)

76. Electron Configuration and the Bohr Model76

77. Electron Configuration and the Bohr Model
Electron configuration was first conceived of under the Bohr model of the atom. 77

78. Electron Configuration and the Bohr Model
Consider the electron configuration and Bohr model image for silicon (Si).
1s2 2s2 2p6 3s2 3p2 78

79. Electron Configuration and the Bohr Model
1s2 2s2 2p6 3s2 3p2
There are 2 electrons in the first energy level (orbit). 79

80. Electron Configuration and the Bohr Model
1s2 2s2 2p6 3s2 3p2
There are 8 electrons in the second energy level (orbit). 80

81. Electron Configuration and the Bohr Model
1s2 2s2 2p6 3s2 3p2
There are 4 electrons in the third energy level (orbit). 81

82. STOP HERE 82

83. Orbital Diagrams
Remember sublevels contain orbitals, and each orbital may contain at most two electrons.
The 2 electrons in an orbital must have opposite “spins.” 83

84. Orbital Diagrams
In addition, when electrons occupy orbitals of equal energy, they don’t pair up with an electron of opposite spin until they have to.

85. Orbital Diagrams Ne

86. Example The electron configuration for oxygen is: 1s2 2s2 2p4
9. Draw the orbital diagram for oxygen (O).
The electron configuration for oxygen is: 1s2 2s2 2p4
The orbital diagram is:
1s 2s 2p 86

87. Problem The electron configuration for silicon is: 1s2 2s2 2p6 3s2 3p2
10. Draw the orbital diagram for silicon (Si).
The electron configuration for silicon is: 1s2 2s2 2p6 3s2 3p2
The orbital diagram is:
1s 2s 2p 3s 3p 87

88. Problem The electron configuration for is: 1s2 2s2 2p6 3s2 3p6 4s1
11. Draw the orbital diagram for potassium (K).
The electron configuration for is:
1s2 2s2 2p6 3s2 3p6 4s1
The orbital diagram is:
1s 2s 2p 3s 3p 4s 88

89. Using Electron Configurations for Element Identification
To identify an element with a given electron configuration, add the superscript numbers together and find the element with that atomic number. 89

90. Example
12. Identify the element with the following electron configuration: 1s2 2s2 2p6 3s1

91. Example
1s2 2s2 2p6 3s1
= 11
Element 11 is sodium (Na).

92. Question
13. Identify the element with the following electron configuration: 1s2 2s2 2p6 3s2 3p4
(sulfur - S)

93. Question
14. Identify the element with the following electron configuration: 1s2 2s2 2p6 3s2 3p6 4s2 3d9
(copper - Cu)

94. Question (germanium - Ge)
15. Identify the element with the following electron configuration: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p2
(germanium - Ge)

95. Electron Configuration Using a Noble Gas Abbreviation
In order to write this type of configuration, find the noble gas (from Group 8A) that comes before the element in question.

96. Electron Configuration Using a Noble Gas Abbreviation
Put the symbol for the noble gas in brackets and then write the part of the configuration that follows to reach the desired element.

97. Example
16. Write the electron configuration using a noble gas abbreviation for magnesium (Mg).
Neon is the noble gas that proceeds magnesium.

98. Example Put neon’s symbol in brackets. [Ne]
Now use the periodic table to determine the rest of the configuration.

99. The noble gas electron configuration for magnesium is: [Ne] 3s2
s block
p block 1 2
d block 3 4 5 6 7
f block
The noble gas electron configuration for magnesium is: [Ne] 3s2
Neon is in light blue; magnesium is in bright yellow.
The additional configuration is 3s2.

100. Example
17. Write the electron configuration using a noble gas abbreviation for nickel (Ni).
Argon is the noble gas that proceeds nickel.

101. Example Put argon’s symbol in brackets. [Ar]
Now use the periodic table to determine the rest of the configuration.

102. The noble gas electron configuration for nickel is: [Ar] 4s2 3d8
s block
p block 1 2
d block 3 4 5 6 7
f block
The noble gas electron configuration for nickel is: [Ar] 4s2 3d8
Argon is in light blue; nickel is in bright yellow.
The additional configuration is 4s2 3d8. (Remember you subtract 1 from the d sublevel row number.)

103. Question
18. Write the electron configuration using a noble gas abbreviation for fluorine (F).
([He] 2s2 2p5)

104. Question
19. Write the electron configuration using a noble gas abbreviation for silicon (Si).
([Ne] 3s2 3p2)

105. Question
20. Write the electron configuration using a noble gas abbreviation for zirconium (Zr).
([Kr] 5s2 4d2)

106. Valence Electrons
The electrons in the outermost energy level are called valence electrons.
You can also use the periodic table as a tool to predict the number of valence electrons in any atom in Groups 1, 2, 13, 14, 15, 16, 17, and 18.

107. Valence Electrons
All atoms in Group 1, like hydrogen, have one valence electron. Likewise, atoms in Group 2 have two valence electrons.

108. Valence Electrons All atoms in Group 13 have three valence electrons.
All atoms in Group 14 have four valence electrons.
All atoms in Group 15 have five valence electrons.

109. Valence Electrons All atoms in Group 16 have six valence electrons.
All atoms in Group 17 have seven valence electrons.
All atoms in Group 18 have eight valence electrons, except helium which only has two.

110. Valence Electrons
All atoms in Groups 3 – 11 (sublevels d and f ) have 2 valence electrons.

111. Question
21. How many valence electrons does each of the following elements have?
carbon (C)
(4)
b) bromine (Br)
(7)

112. Question
c) iron (Fe)
(2)
d) potassium (K)
(1)
e) aluminum (Al)
(3)

113. Question
22. How many valence electrons does each have based on the following electron configuration?
[Kr] 5s2 4d2
(2)
b) [He] 2s2 2p5
(7)
113

114. Question (2) (3) (2) c) 1s2 2s2 2p6 3s2 3p6 4s2 3d9
d) 1s2 2s2 2p6 3s2 3p1
(3)
e) 1s2 2s2 2p6 3s2 3p6 4s2
(2)
114

115. Lewis Dot Diagrams
Because valence electrons are so important to the behavior of an atom, it is useful to represent them with symbols.

116. Lewis Dot Diagrams
A Lewis dot diagram illustrates valence electrons as dots (or other small symbols) around the chemical symbol of an element.

117. Each dot represents one valence electron.
In the dot diagram, the element’s symbol represents the core of the atom—the nucleus plus all the inner electrons.

118. Electron (Lewis) Dot Diagrams
Write the symbol. X
Put one dot for each valence electron.
Don’t pair electrons until you have to.

119. Question
23. Write a Lewis dot diagram for chlorine.

120. Question
24. Write a Lewis dot diagram for calcium.

121. Question
25. Write a Lewis dot diagram for potassium.