This document discusses the periodic table and chemical periodicity. It describes how Dmitri Mendeleev published one of the first periodic tables in 1869, organizing the known elements based on their chemical properties. It also notes that Julius Lothar Meyer published a similar table that year based on physical properties. The document outlines some of Mendeleev's predictions about undiscovered elements and how the periodic table was later updated based on Henry Moseley's work showing atomic number is the fundamental property.

1. CHEMICAL PERIODICITY
• PERIODIC TABLE
• MENDELEEV
• PREDICTION OF PROPERTIES
• Chapters 8.4; 9.5; 14.1-14.10
14-1

2. Goals & Objectives
• See the following learning objectives on pages
322-323, 356, 563.
• Understand the Concepts:
• 8.9-12; 9.13-15; 14.-8, 10, 14, 20, 23.
• Master these Skills: 8.6.
14-2

3. CHEMICAL PERIODICITY
• THE PERIODIC TABLE
– 1869 Russian chemist Dmitri Mendeleev published a periodic
table based on chemical properties of the elements known at
the time
– 1869 German chemist Julius Lothar Meyer published a similar
table based on physical properties
– Greater credit was given to Mendeleev because he was able to
predict the properties of several undiscovered elements
14-3

4. Mendeleev’s Periodic Table
Elements were grouped in order of
atomic mass; if the next known element
did not fit, Mendeleev left a space.
14-4

5. Pictures of Mendeleev &
Meyer
14-5

6. Mendeleev’s Predictions
14-6

7. The Periodic Law
• The properties of the elements are
periodic functions of their atomic
number
• This change is based on the work of
British chemist Henry G. J. Moseley,
who showed that the order of the
periodic table is not based on atomic
mass but rather on atomic number.
14-7

8. Henry Moseley picture
• During the first World
War, Moseley was
killed while serving
as a pilot in the
British army at the
age of 26. Thus a
brilliant career was
lost to science
14-8

9. Terminology
•
•
•
•
•
•
•
•
14-9
Group(family) of element - vertical column
Period(row) of elements - horizontal row
Group 1 metals - alkali metals
Group 2 metals - alkaline earth metals
Group 1 + Group 2 - representative metals
Group 16 - chalcogens
Group 17 - halogens
Group 18 - rare, inert or noble gases

10. Classification of the
Elements
14-10

11. Classification of the
Elements
Nonmetals
Transitions metals
14-11
Post trans.
Metals

12. This is a link to the Element
Song.
The Elements" (1959) is a song
by musical humorist Tom Lehrer,
which recites the names of all
the chemical elements known at
the time of writing, up to number
102, nobelium. It can be found
on his albums Tom Lehrer in
concert, More Songs by Tom
Lehrer and An Evening Wasted
with Tom Lehrer. The song is
sung to the tune of the Major
General's Song from The
Pirates of Penzance by Gilbert
and Sullivan.
14-12

13. This is a link to the New
Element Song. The New
Periodic Table Song
(2013) is a song by ASCAP
Science, which recites the
names of all the chemical
elements (in order) including
the new elements:
Flerovium with the symbol "Fl"
and atomic number 114 and
Livermorium with the symbol
“Lv” and atomic number 116.
14-13

14. Chapter 14
Periodic Patterns in the Main-Group Elements
14-14

15. Hydrogen
• Hydrogen has a very simple structure:
– the nucleus has a single positive charge, and has 1 electron.
• Hydrogen is the most abundant element in the universe.
• Hydrogen exists as a diatomic gas, H2.
– H2 is colorless and odorless with very low melting and boiling
points.
• H is abundant in combination with oxygen as H2O.
14-15

16. Hydrogen and Group 1
• Like the Group (1) elements, H has a half-filled valence
level.
• H is similar to the other Group 1 elements in terms of
– ionization energy,
– electron affinity,
– electronegativity, and
– bond energies.
14-16

17. Hydrogen and the Halogens
• Like the halogens or Group (17), hydrogen
– exists as a diatomic molecule and
– needs only 1 electron to fill its valence shell.
• Unlike the halogens
– H has a much lower electronegativity than any halogen,
– H lacks the three valence e- pairs that halogens have, and
– halide ions (X-) are common and stable, but the hydride ion (H-)
is rare and reactive.
14-17

18. Table 14.1
14-18
Trends in Atomic, Physical, and Chemical
Properties of the Period 2 Elements.

19. Table 14.1
14-19
Trends in Atomic, Physical, and Chemical
Properties of the Period 2 Elements.

20. Table 14.1
14-20
Trends in Atomic, Physical, and Chemical
Properties of the Period 2 Elements.

21. Table 14.1
14-21
Trends in Atomic, Physical, and Chemical
Properties of the Period 2 Elements.

22. Table 14.1
Trends in Atomic, Physical, and Chemical
Properties of the Period 2 Elements.
Trends in atomic radius, ionization energy, and electronegativity
across Period 2.
14-22

23. Group (1): The Alkali Metals
Family Portrait
KEY ATOMIC PROPERTIES, PHYSICAL PROPERTIES, AND REACTIONS
Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
14-23

24. Properties of the Alkali Metals
• Alkali metals are the largest elements in their respective
periods and their valence electron configuration is ns1.
– The valence e- is relatively far from the nucleus, resulting in
weak metallic bonding.
• Alkali metals are unusually soft for metals. They can be
cut easily with a knife.
• Alkali metals have lower melting and boiling points than
any other group of metals.
• Alkali metals have lower densities than most metals.
14-24

25. Periodic Properties
•
•
•
•
•
14-25
Electron configurations
Li [He]2s1
Na [Ne]3s1
K [Ar]4s1
Rb [Kr]5s1

26. Lithium floating in oil
floating on water. Alkali
metals have low densities.
14-26
Potassium reacting with
water. Alkali metals are
very reactive.

27. Demonstration of the Reactivity
of Alkali Metals
• The following YouTube video shows the
reactivity of the alkali metals.
• Brainiac Alkali Metals
14-27

28. Group (2): The Alkaline Earth Metals
• The oxides of Group (2) elements form basic solutions
and melt at extremely high temperatures.
• Group (2) elements have higher ionization energies than
Group (1) elements
– due to their higher effective nuclear charge and smaller size.
• Group (2) elements are strong reducing agents.
14-28

29. Group (2): The Alkaline Earth Metals
Family Portrait
KEY ATOMIC PROPERTIES, PHYSICAL PROPERTIES, AND REACTIONS
Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
14-29

30. Group (13): The Boron Family
Family Portrait
KEY ATOMIC PROPERTIES, PHYSICAL PROPERTIES, AND REACTIONS
Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
14-30

31. Influence of Transition Elements on
Group (13)
The larger 13 elements have smaller atomic radii and
larger ionization energies than electronegativities than
expected.
These properties influence the physical and chemical
behavior of these elements.
14-31

32. Group (14): The Carbon Family
Family Portrait
KEY ATOMIC PROPERTIES, PHYSICAL PROPERTIES, AND REACTIONS
Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
14-32

33. Allotropes
Allotropes are different crystalline or molecular forms of
the same element.
One allotrope of a particular element is usually more stable than
another at a particular temperature and pressure.
Carbon has several allotropes, including graphite,
diamond, and fullerenes.
Tin exhibits two allotropes; white β-tin and gray α-tin.
14-33

34. Figure 14.9
14-34
Phase diagram of carbon.

35. Figure 14.10 Crystalline buckminsterfullerene and a buckyball (A)
and nanotubes (B).
Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
Crystals of buckminsterfullerene (C60)
Nanotubes
14-35

36. Carbon in Organic Chemistry
The large number and wide variety of organic compounds
is due to the ability of C to bond to itself, and to form
multiple bonds.
Catenation is the process whereby carbon bonds to itself
to form stable chains, branches, and rings.
Since C is small, the C-C bond is short enough to allow
effective side-to-side overlap of p orbitals. C readily forms
double and triple bonds.
14-36

37. Figure 14.12 Three of the several million known organic
compounds of carbon.
Lysine, one of 20 amino
acids that occur in proteins
14-37

38. Inorganic Compounds of Carbon
Carbon bond with oxygen to form carbonates. Metal
carbonates such as CaCO3 are abundant in minerals.
Carbon forms two common gaseous oxides, CO and CO2,
which are molecular. Other Group (14) elements form
network-covalent or ionic oxides.
Carbon halides have major uses as solvents and in
structural plastics.
14-38

39. Figure 14.13 Freon-12 (CCl2F2), a chlorofluorocarbon.
Chlorofluorocarbons (CFC’s or Freons) are chemically and thermally
stable, nontoxic, and nonflammable. They are excellent cleaners,
refrigerants, and propellants, but they decompose extremely slowly
near the Earth’s surface. They readily enter the stratosphere, where
UV radiation causes them to release free Cl atoms that damage the
ozone layer.
14-39

40. Group (15) Elements
• Nitrogen is a diatomic gas (N2) with a very low boiling
point, due to its very weak intermolecular forces.
• Phosphorus exists most commonly as tetrahedral P4
molecules. It has stronger dispersion forces than N2.
• Arsenic exists as extended sheets of As atoms
covalently bonded together. The covalent network
structure gives it a high melting point.
• Antimony also has a covalent network structure.
• Bismuth has metallic bonding. Its melting point is lower
than that of As or Sb.
14-40

41. Group (15): The Nitrogen Family
Family Portrait
KEY ATOMIC PROPERTIES, PHYSICAL PROPERTIES, AND REACTIONS
Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
14-41

42. Figure 14.16
Two allotropes of phosphorous.
White phosphorous (P4)
Strained bonds in P4
Red phosphorous
14-42

43. Periodicvideos.com
• This is the periodic video for Phosphorus.
• Phosphorus Video
14-43

44. Patterns of Behavior in Group (15)
• N gains 3 electrons to form the anion N3-, but only in
compounds with active metals.
• The higher elements in the group are metallic and lose
electrons to form cations.
• Oxides change from acidic to amphoteric to basic as you
move down the group.
• All Group 5A(15) elements form gaseous hydrides with
the formula EH3.
– All except NH3 are extremely reactive and toxic.
14-44

45. Oxides of Nitrogen
• Nitrogen forms six stable oxides. Hf for all six oxides is
positive because of the great strength of the NΞN bond.
• NO is produced by the oxidation of ammonia:
– 4NH3(g) + 5O2(g) → 4NO(g) +6H2O(g)
– This is the first step in the production of nitric acid.
• NO is converted to 2 other oxides by heating:
3NO(g)
Δ
N2O(g) + NO2(g)
– This type of redox reaction is called disproportionation.
• NO2 is a component of photochemical smog.
14-45

46. Table 14.3
14-46
Structures and Properties of the Nitrogen Oxides

47. Group (16) Elements
• Oxygen, like nitrogen, occurs as a low-boiling diatomic
gas, O2.
• Sulfur, like phosphorus, occurs as a polyatomic
molecular solid.
• Selenium, like arsenic, commonly occurs as a gray
metalloid.
• Tellurium, like antimony, displays network covalent
bonding.
• Polonium, like bismuth, has a metallic crystal structure.
14-47

48. Group (16): The Oxygen Family
Family Portrait
KEY ATOMIC PROPERTIES, PHYSICAL PROPERTIES, AND REACTIONS
Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
14-48

49. Allotropes in the Oxygen Family
Oxygen has two allotropes:
- O2, which is essential to life, and
- O3 or ozone, which is poisonous.
Sulfur has more than 10 different forms, due to the ability
of S to catenate. S–S bond lengths and bond angles may
vary greatly.
Selenium has several allotropes, some consisting of
crown-shaped Se8 molecules.
14-49

50. Figure 14.20
The cyclo-S8 molecule.
top view
side view
At room temperature, the sulfur molecule is a crown-shaped ring of
eight atoms. The most stable S allotrope is orthorhombic α-S8,
which consists of cyclo-S8.
14-50

51. Hydrides of the Oxygen Family
• Oxygen forms two hydrides:
– water (H2O) and hydrogen peroxide (H2O2).
– H2O2 contains oxygen in a -1 oxidation state.
• The hydrides of the other 16 elements are foul-smelling,
poisonous gases.
– H2S forms naturally in swamps from the breakdown of organic
matter and is as toxic as HCN.
• H2O and H2O2 can form H bonds, and therefore have
higher melting and boiling points than other H2E
compounds.
• Hydride bond angles decrease and bond lengths
increase down the group.
14-51

52. Halides of the Oxygen Family
Except for O, the Group 16 elements form a wide range of
halides.
Their structure and reactivity patterns depend on the sizes of the
central atom and the surrounding halogens.
As the central atom becomes larger, the halides become
more stable.
14-52

53. Highlights of Sulfur Chemistry
• Sulfur forms two important oxides:
– SO2 has S in its +4 oxidation state. It is a colorless, choking gas
that forms when S, H2S or a metal sulfide burns in air.
– SO3 has S in the +6 oxidation state.
• Sulfur forms two important oxoacids.
– Sulfurous acid (H2SO3) is a weak acid with two acidic protons.
– Sulfuric acid (H2SO4) is a strong acid, and is an important
industrial chemical. It is an excellent dehydrating agent.
14-53

54. Figure 14.21 The dehydration of sugar by sulfuric acid.
14-54

55. 14-55

56. Group (17): The Halogens
Family Portrait
KEY ATOMIC PROPERTIES, PHYSICAL PROPERTIES, AND REACTIONS
Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
14-56

57. Reactivity of the Halogens
A halogen atom needs only one electron to fill its valence
shell. Halogens are therefore very reactive elements.
The halogens display a wide range of electronegativities, but
all are electronegative enough to behave as nonmetals.
A halogen will either
- gain one electron to form a halide anion or
- share an electron pair with a nonmetal atom.
The reactivity of the halogens decreases down the group,
reflecting the decrease in electronegativity.
14-57

58. Group (18): The Noble Gases
Family Portrait
KEY ATOMIC PROPERTIES, PHYSICAL PROPERTIES, AND REACTIONS
Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
14-58

59. Noble Gases
• Noble gases have a full valence shell.
• The noble gases are the smallest elements in their
respective periods, with the highest ionization energies.
• Atomic size increases down the group and IE decreases.
• Noble gases have very low melting and boiling points.
• Only Kr, Xe, and Rn are known to form compounds.
– Xe is the most reactive noble gas and exhibits all even oxidation
states from +2 to +8.
14-59

60. Figure 14.26
14-60
Crystals of xenon tetrafluoride (Xe(F4).

61. Chapter 8
Electron Configuration and Chemical Periodicity
14-61

62. Periodic Properties
• Electronic configurations
s
p
d
f
14-62
f

63. Electron Configurations
14-63

64. Electron Configuration
• Give the symbol for the element
having the outermost electron
configuration 3s2
• Give the symbol for the element
having the outermost electron
configuration 4p4
14-64

65. 14-65

66. Trends in Atomic Size
Atomic size increases as the principal quantum number n
increases.
As n increases, the probability that the outer electrons will be further
from the nucleus increases.
Atomic size decreases as the effective nuclear charge Zeff
increases.
As Zeff increases, the outer electrons are pulled closer to the nucleus.
For main group elements:
atomic size increases down a group in the periodic table
and decreases across a period.
14-66

67. Figure 8.13
Atomic radii of the maingroup and transition
elements.
14-67

68. Figure 8.14
14-68
Periodicity of atomic radius.

69. Sample Problem 8.3
PROBLEM:
Ranking Elements by Atomic Size
Using only the periodic table (not Figure 8.15), rank each
set of main-group elements in order of decreasing atomic
size:
(a) Ca, Mg, Sr
(b) K, Ga, Ca
(c) Br, Rb, Kr
(d) Sr, Ca, Rb
PLAN: Locate each element on the periodic table. Main-group
elements increase in size down a group and decrease in size
across the period.
14-69

70. Sample Problem 8.3
SOLUTION:
(a) Sr > Ca > Mg
Ca, Mg, and Sr are in Group 2A. Size increases down the group.
(b) K > Ca > Ga
K, Ga, and Ca are all in Period 4. Size decreases across the period.
(c) Rb > Br > Kr
Rb is the largest because it has one more energy level than the other
elements. Kr is smaller than Br because Kr is further to the right in the
same period.
(d) Rb > Sr > Ca
Ca is the smallest because it has one fewer energy level. Sr is smaller
than Rb because it is smaller to the right in the same period.
14-70

71. Atomic Radii
size decreases
14-71

72. Atomic Radii
• Give the symbol for the larger of the
two atoms As or N
• Give the symbol for the atom having
the smaller radius of Na or P
14-72

73. Ionization Energy (IE)
• The energy in kilojoules required for
the complete removal of one mole of
electrons from one mole of gaseous
atoms or ions
• Pulling an electron away from the
nucleus requires energy to overcome
the attraction of the nucleus.
14-73

74. Trends in Ionization Energy
Ionization energy (IE) is the energy required for the
complete removal of 1 mol of electrons from 1 mol of
gaseous atoms or ions.
Atoms with a low IE tend to form cations.
Atoms with a high IE tend to form anions (except the
noble gases).
Ionization energy tends to decrease down a group and
increase across a period.
14-74

75. Figure 8.15
14-75
Periodicity of first ionization energy (IE1).

76. Figure 8.16
14-76
First ionization energies of the main-group elements.

77. Ionization Energy
IE increases
14-77

78. Ionization Energy
14-78

79. Ionization Energy
14-79

80. Sample Problem 8.4
Ranking Elements by First Ionization
Energy
PROBLEM: Using the periodic table only, rank the elements in each of
the following sets in order of decreasing IE1:
(a) Kr, He, Ar
(c) K, Ca, Rb
(b) Sb, Te, Sn
(d) I, Xe, Cs
PLAN: Find each element on the periodic table. IE1 generally
decreases down a group and increases across a period.
SOLUTION:
(a) He > Ar > Kr
Kr, He, and Ar are in Group 8A. IE1 decreases down the group.
14-80

81. Sample Problem 8.4
SOLUTION:
(b) Te > Sb > Sn
Sb, Te, and Sn are in Period 5. IE1 increases across a period.
(c) Ca > K > Rb
K has a higher IE1 than Rb because K is higher up in Group 1A. Ca
has a higher IE1 than K because Ca is further to the right in Period 4.
(d) Xe > I > Cs
Xe has a higher IE1 than I because Xe is further to the right in the
same period. Cs has a lower IE1 than I because it is further to the left
in a higher period.
14-81

82. Ionization Energy
• Give the symbol for the element in
the same group as Sn having the
largest IE.
• Give the symbol for the element in
the same period as Si having the
smallest IE.
14-82

83. Electron Affinity (EA)
• Electron Affinity is the energy change
in kilojoules accompanying the
addition of one mole of electrons to
one mole of gaseous atoms or ions.
• In most cases, energy is released
when the first electron is added
because it is attracted to the atoms
nuclear charge.
14-83

84. Trends in Electron Affinity
Electron Affinity (EA) is the energy change that occurs
when 1 mol of electrons is added to 1 mol of gaseous
atoms or ions.
Atoms with a low EA tend to form cations.
Atoms with a high EA tend to form anions.
The trends in electron affinity are not as regular as those
for atomic size or IE.
14-84

85. Figure 8.18
14-85
Electron affinities of the main-group elements
(in kJ/mol).

86. Behavior Patterns for IE and EA
Reactive nonmetals have high IEs and highly negative
EAs.
These elements attract electrons strongly and tend to form negative
ions in ionic compounds.
Reactive metals have low IEs and slightly negative EAs.
These elements lose electrons easily and tend to form positive ions in
ionic compounds.
Noble gases have very high IEs and slightly positive EAs.
These elements tend to neither lose nor gain electtons.
14-86

87. Electron Affinity
14-87

88. Electron Affinity
EA increases
14-88

89. Electron Affinity
• Give the symbol for the
representative metal in the same
period as As having the larger
electron affinity.
• Give the symbol for the metalloid in
the same family as N having the
smaller EA.
14-89

90. Electronegativity (EN)
• Electronegativity is the relative ability of a
bonded atom to attract the shared electron pair
• It was developed over 50 years ago by
American Chemist Linus Pauling, who is the
only person to win the Nobel Prize in two
different categories – Chemistry in 1954 &
Peace in 1963. He died in 1994 at age 93.
14-90

91. Pictures of Linus Pauling
14-91

92. Trends in Electronegativity
The most electronegative element is fluorine.
In general electronegativity decreases down a group as
atomic size increases.
In general electronegativity increases across a period
as atomic size decreases.
Nonmetals are more electronegative than metals.
14-92

93. Figure 9.22
14-93
Electronegativity and atomic size.

94. Electronegativity
EN increases
14-94

95. Electronegativity
• Give the symbol for the nonmetal in
the same period as Al having the
greatest EN.
• Give the symbol for the element
having the greatest electronegativity.
14-95

96. Ionic Size vs. Atomic Size
Cations are smaller than their parent atoms while
anions are larger.
Ionic radius increases down a group as n increases.
Cation size decreases as charge increases.
An isoelectronic series is a series of ions that have
the same electron configuration. Within the series, ion
size decreases with increasing nuclear charge.
3- > 2- > 1- > 1+ > 2+ > 3+
14-96

97. Ionic Radii
• Cations are always smaller than their
corresponding neutral atoms
14-97

98. Ionic Radii
• Anions are always larger than their
corresponding neutral atoms
14-98

99. Ionic Radii
• Give the symbol for the element with
the largest ionic radius between Cl
and Cl-1.
• Give the symbol for the element with
the smallest ionic radius between Na
and Na+1.
14-99

100. Binary Compounds


A. Salt of metal cation with only one
charge and nonmetal anion
MgCl2



14-
metal + nonmetal--metal in Group 2
metal + stem of nonmetal + ide
magnesium chloride

101. Binary Compounds

Charges on metals







14-
Group 1
+1
Group 2
+2
Aluminum
+3
Zinc (Zn)
+2
Silver(Ag)
+1
Cadmium(Cd) +2
See also Table 2.3 on page 63.

102. Binary Compounds






14-
lithium sulfide
barium chloride
aluminum bromide
AgCl
CdO
Al2O3

103. Binary Compounds

B. Salt of metal cation with more than
one possible charge and nonmetal anion
All metals not covered in A
Name must include charge on metal
14-

104. Binary Compounds


IUPAC System
FeCl3, FeCl2


metal + stem of nonmetal+ide
charge on metal in Roman numerals


14-
iron(III) chloride, iron(II) chloride
See also Table 2.4 on page 65.

105. Binary Compounds






14-
NiO
CrCl3
SnCl4
Copper (II) bromide
Iron (III) oxide
Cobalt (III) nitride

106. Binary Compounds



14-
Older system
“ic “ suffix designates the higher of two
possible charges
“ous” suffix designates the lower of two
possible charges

107. Binary Compounds







14-
Memorize these four ions for the older
system
Fe2+ - ferrous
Fe3+ - ferric
Cu+1 - cuprousCu2+ - cupric
Sn2+ - stannous
Sn4+ - stannic
Hg22+ - mercurous
Hg2+ - mercuric
See also Table 2.4 on page 65.

108. Binary Compounds







14-
ferrous chloride
ferric sulfide
cuprous iodide
cupric bromide
mercuric chloride
mercurous chloride
stannous fluoride