1. Unit 2 The Periodic Table

2. The Periodic Table The most important tool in chemistry
Used to understand and predict the properties of elements

3. Dmitri Mendeleev Russian chemist Proposed the first Periodic Table
1871

4. The First Periodic Table
Arranged according to increasing atomic mass
"The properties of the elements are a periodic function of their atomic masses" – Dmitri Mendeleev

5. The First Periodic Table
Columns were organized so that elements with similar properties were in the same column.

6. The First Periodic Table

7. Henry Moseley British physicist 1914
Similar arrangement as that of the Modern Periodic Table

8. Moseley Arranged elements according to increasing atomic number
Rearrangement cleared up inconsistencies and contradictions of Mendeleev’s Periodic Table

9. 1930’s Periodic Table

10. Modern Periodic Table

11. Group or family A vertical column of elements in the periodic table
18 groups
Elements in the same group have similar chemical and physical properties

12. Groups 18 1

13. A horizontal row of elements in the periodic table
7 periods

14. Period 1 7

15. Classes of Elements
Metals
Nonmetals
Metalloids

16. Properties of Metals Lustrous (shiny) Malleable (pound into sheets)
Ductile (drawn into wires)

17. Properties of Metals Good Conductors (heat / electricity)
Solids at room temperature
Exception: Mercury (liquid)

18. Properties of Nonmetals
Not lustrous (not shiny)
Poor conductors (heat and electricity)
At room temperature, most are brittle solids or gases. One element, bromine, is a liquid.

19. Properties of Metalloids
Properties of both metals and nonmetals
Metalloids are semiconductors.

20. Element Classification
Elements classified into 4 categories based on their electron configurations
Noble Gases
Representative Elements (A)
Transition Metals (B)
Rare Earth Elements
(La & Ac rows)

21. The Noble Gases

22. The Noble Gases Elements in which outermost electron levels are filled
Sometimes referred to as Group 0, 8A or 18
Sometimes called inert gases

23. The Representative Element

24. The Representative Elements
Elements in which the outermost electron levels are only partially filled
Sometimes called Group A elements

25. The Representative Elements
Some groups have special names
Group 1A (1): Alkali Metals
Group 2A (2): Alkaline Earth Metals
Group 6A (16): Chalcogens
Group 7A (17): Halogens

26. Alkali Metals *except H

27. Alkali Metals Group 1A (1) of the Periodic Table
Elements contain 1 valence electron in the outermost s sublevel.

28. Electron Placement

29. Alkali Metals
Link

30. Alkaline Earth Metals

31. Alkaline Earth Metals Group 2A (2) of the Periodic Table
Elements contain 2 valence electrons in the outermost s sublevel.

32. Halogens

33. Halogens Group 7A (17) on the Periodic Table
Elements contain 5 valence electrons in the p sublevel and 2 in the s sublevel.
Total of 7 valence electrons.

34. The Transition Metals

35. The Transition Metals
Metallic elements in which the outer most s sublevel is filled and nearby d sublevel contains electrons.
Part of the Group B elements

36. Electron Placement

37. Rare Earth Elements

38. Rare Earth Elements
Sometimes called Inner Transition Metals or Lanthanides and Actinides
Metallic elements in which the outermost s sublevel is filled and nearby f sublevel contains electrons

39. The Inner Transition Metals
Part of the Group B elements
Moved to the bottom of the Periodic Table to save space

42. Objective
Review the atomic structure of the atom.

43. Quantum Mechanical Model
Erwin Schrodinger – 1926
Mathematical solution to Schrodinger’s Equation

44. Quantum Mechanical Model
Nucleus:
Protons
Neutrons
Electron Cloud
Area where
there is a 90%
chance electrons
can be found

45. Subatomic Particles 1+ 1- nucleus cloud Subatomic Particle Relative
Charge
Mass
Location
Proton 1+
1 amu
nucleus
neutron
electron 1-
0 amu
cloud

46. Objective
Define atomic radius and account for the trend in atomic radii for elements within a group and period in the periodic table.

47. Atomic Radius
One-half of the distance between the nuclei in a molecule consisting of identical atoms.

48. Atomic Radius
Atomic
Radius

49. Atomic Radius Cs Rb K Na Li Xe Kr Ar Ne He

50. Atomic Radius

51. Atomic Radius

53. Atomic Radius (Period)
As you go across a period you add more electrons to the SAME energy level
As you go across the period you are also adding more protons to the nucleus

54. Atomic Radius (Period)
Additional protons in the nucleus create a higher “effective nuclear charge.”

55. Atomic Radius (Period)
The higher “effective nuclear charge” means that there is a stronger force pulling the electrons toward the nucleus (like a magnet).

56. Atomic Radius Trend (Period)
Atomic radius decreases (it gets smaller) as you go from left to right on the periodic table.

57. Atomic Radius (Period 2)

58. Atomic Radius (Group)
The number of energy levels increases as you go down a group.

59. Atomic Radius (Group)
Each additional energy level is further from the nucleus.

60. Atomic Radius (Group)
As you go down a group, the number of protons in the nucleus also increases and so does the “effective nuclear charge.”

61. Atomic Radius (Group)
Electrons in the inner energy levels “shield” those electrons on the outer energy levels. (shielding effect)

62. Atomic Radius Trend (Group)
Atomic radius increases (it gets larger) as you go down a group.

63. Atomic Radius (Group 1)

64. Circle the Element with the Larger Atomic Radius
P or S? P
B or Al? Al
N, O, P or S?

65. Circle the Element with the Larger Atomic Radius
Rb or Cs? Cs
Na or K? K
Li, Be, Na, or Mg? Na

66. Atomic Radius

67. Objective
Define ionization energy and account for the trend in ionization energy within a group and period in the periodic table.

68. Objective
Distinguish and account for the differences between first, second, and third ionization energies.

69. Ion and Ionization
Ionization: Any process that results in the formation of an ion.
Ion: An atom or group of atoms that has a positive or negative charge.

70. Ionization Energy
The energy required to remove an electron from an atom.

71. Ionization Energy
Removing an electron results in the formation of a cation (positively charged ion).
“cat-eye-on”

72. Ionization Energy
+1 ion + Na
Na+
electron
Na(g)  Na+(g) + e-

73. Energy required to remove the first outermost electron.
1st Ionization Energy
Energy required to remove the first outermost electron.

74. First Ionization Energies

75. First Ionization Energies (Group)
As you go down a group electrons are added to energy levels that are further from the nucleus

76. First Ionization Energies (Group)
Outer energy level electrons are also “shielded” from the full affect of the effective nuclear charge by the inner energy level electrons.

77. First Ionization Energies (Group)
As you go down a group the first ionization energy decreases (gets smaller).
Less energy is required to remove electrons from an element as you go down a group.

78. First Ionization Energies (Period)
The effective nuclear charge increases as you go across the period (left to right).
Electrons are held tighter by the nucleus and require more energy to be removed.

79. First Ionization Energies (Period)
First Ionization Energy increases (gets bigger) as you go left to right across a period.

80. 1st Ionization Energy

81. 2nd Ionization Energy
Energy required to remove a second electron from the outermost electrons.

82. Third Ionization Energy
Energy required to remove a third electron from the outermost electrons.

83. Ionization Energies
Once an electron is removed from an atom, the nucleus holds onto the other electrons tighter.
It requires more energy to remove additional electrons.

84. 1st IE < 2nd IE < 3rd IE
Ionization Energies
1st IE < 2nd IE < 3rd IE

85. Circle the Element with the Larger Ionization Energy
P or S? S
B or Al? B
N, O, P or S? O

86. Circle the Element with the Larger Ionization Energy
Rb or Cs? Rb
Na or K? Na
Li, Be, Na, or Mg? Be

87. Objective
Define electronegativity and state the group and period trend in electronegativity.

88. Electronegativity
The tendency for an atom to attract electrons to itself when it is chemically combined with another element.

89. Electronegativity Pair of Shared Electrons
Tug of war for shared electrons
The electrons will be closer to the
more electronegative atom.
Pair of Shared
Electrons

90. Electronegativity
Fluorine is the most electronegative element – it is assigned a value of 4.0
All other elements are compared to fluorine on this relative scale.

91. Electronegativity

92. Electronegativity
Noble gases are often omitted because they tend not to form compounds.

93. Electronegativity F Cl Br I At Li Na K Rb Cs Fr

94. Electronegativity

95. Electronegativity (Group)
As you go down a group electrons are added to energy levels that are further from the nucleus

96. Electronegativity (Group)
Outer energy level electrons are also “shielded” from the full affect of the effective nuclear charge by the inner energy level electrons.

97. Electronegativity (Group)
As you go down a group the electronegativity decreases (gets smaller)
The nucleus has less of an pull or attraction on the electrons shared in a chemical bond.

98. Electronegativity Group 1

99. Electronegativity (Period)
The effective nuclear charge increases as you go across the period (left to right)
Electrons are held tighter and pulled closer to the nucleus.

100. Electronegativity (Period)
Electronegativity increases (gets bigger) as you go left to right across a period.

101. Electronegativity Period 2

102. Circle the Element with the Larger Electronegativity
P or S? S
B or Al? B
N, O, P or S? O

103. Circle the Element with the Larger Electronegativity
Rb or Cs? Rb
Na or K? Na
Li, Be, Na, or Mg? Be

104. Fr Metallic Properties Increase down a group.
Decrease across a period (L R).
Most metallic element- Fr

105. Objective
Write the electron configurations of an element using the periodic table.

106. Electron
The electron is the most important subatomic particle in determining physical and chemical properties.

107. Periodic Table
The periodic table is arranged in a way that elements with the same number of “valence electrons” are in the same group.

108. Valence Electrons
Electrons in the highest occupied energy level of an atom.

109. Blocks The periodic table can be divided into sections or blocks
s-block
p-block
d-block
f-block

110. The Periodic Table

111. Blocks
s-block
Groups 1 – 2
p-block
Groups 13 – 18

112. Blocks d-block f-block Groups 3 – 12 Transition Metals
Lanthanide and Actinide series
Inner Transition Metals

113. Electron Configurations
Using the periodic table to write electron configurations for elements.

114. Electron Configurations
Read the periodic table as you would a book
Left to Right
Top to Bottom

115. Electron Configurations
Write down all filled sublevels and stop at the element of interest

116. Shorthand Notation and Electron Configuration
Use shorthand notation by writing the symbol of the previous noble gas and continue writing the electron configuration from there.

117. Electron ConfigurationsP
1st and 2nd energy levels are full  1s2 2s2 2p6
3s is also full  3s2
3p is not, 3 squares in  3p3
1s2 2s2 2p6 3s2 3p3

118. Electron Configurations
Use the periodic table to write electron configurations for the following elements.

119. Electron ConfigurationP
Shorthand Notation – start with previous noble gas [Ne]
Next period is 3 – 3 s is full and 3 blocks into 3p  3s2 3p3
[Ne] 3s2 3p3

120. Electron ConfigurationMn
Look at the periodic table
all of 1s, 1p, 2s, 2p, 3s and 3 p are full  takes you up to Ar
4s is full and then we are 5 squares into 4d

121. Electron ConfigurationMn
[Ar] 4s2 3d5

122. Atomic Structure Electron Configuration
Review
Atomic Structure
Electron Configuration

123. Lithium (Li) atomic number: ____ number of electrons: ___
group number: ___
period number: ___
group name (if any): ________________

124. Lithium (Li)
electron configuration:

125. Argon (Ar) atomic number: ____ number of electrons: ___
group number: ___
period number: ___
group name (if any): ________________

126. Argon (Ar)
electron configuration:

127. Titanium atomic number: ____ number of electrons: ___
group number: ___
period number: ___
group name (if any): ________________

128. Titanium
electron configuration: