1. Part I - Introduction to the Periodic Table
Unit 3.2
Part I - Introduction to the Periodic Table

2. History of the Periodic Table

3. Until 1800’s – no clear system
The Search for a Periodic Table
Until 1800’s – no clear system
Elements grouped by similar properties or atomic mass
In 1829, J.W. Döbereiner classified some elements into groups of three, which he called triads.
1600’s – chemistry still tied to alchemy
Belief that urine could be changed to gold led to discovery of matches
1700’s – elan vital – force of life
Then mid 1700’s = Lavoisier
1800’s – no clear system used individual abbreviations, symbols. Formulas meant different things to different people
1850’s – Mendeleev – grouped by atomic mass or common properties (solids, gases)

4. Döbereiner’s Triads
The elements in a triad had similar chemical properties, and their physical properties varied in an orderly way according to their atomic masses.
Element
Atomic mass (g)
Density (g/mL)
Melting point (C)
Boiling point (C)
Chlorine
35.5
-101
-34
Bromine
79.9
3.12 -7 59
Iodine
127
4.93
114
185

5. Density increases with increasing atomic mass.
Döbereiner’s Triads
Element
Atomic mass (g)
Density (g/mL)
Melting point (C)
Boiling point (C)
Chlorine
35.5
-101
-34
Bromine
79.9
3.12 -7 59
Iodine
127
4.93
114
185
Density increases with increasing atomic mass.
The concept of triads suggested that the properties of an element are related to its atomic mass.

6. Which of the Dobereiner triads shown are still listed in the same column of the modern periodic table?
Triad 1
Triad 2
Triad 3 Li Mn S Na Cr Se K Fe Te
Triad 1 and triad 3

7. organized the elements according to increasing atomic mass.
Mendeleev’s Periodic Table
The Russian chemist, Dmitri Mendeleev, developed a periodic table of elements.
organized the elements according to increasing atomic mass.

8. Mendeleev’s Periodic Table
Mendeleev later developed an improved version of his table with the elements arranged in horizontal rows.

9. Mendeleev’s Periodic Table
Patterns of changing properties repeated for the elements across the horizontal rows.
Elements in vertical columns have similar properties.

10. Periodicity is the tendency to recur at regular intervals.
Mendeleev’s Periodic Table
properties of the elements repeat in an orderly way from row to row of the table.
This repeated pattern is an example of periodicity in the properties of elements.
Periodicity is the tendency to recur at regular intervals.

11. Mendeleev’s Periodic Table
In order to group elements with similar properties in the same columns, Mendeleev had to leave some blank spaces in his table.
He suggested that these spaces represented undiscovered elements.
***Mendeleev correctly predicted the properties of several undiscovered elements.
Why is this important?

12. Mendeleev’s Periodic Table

13. What are two factors that contributed to the acceptance of Mendeleev’s periodic law?
Grouping of elements with similar chemical properties
Ability to predict properties of undiscovered elements

15. The Modern Periodic Table
the basis for ordering the elements in the table is the atomic number, not atomic mass.
The atomic number of an element is equal to the number of protons in the nucleus.
Each row (except the first) begins with a metal and ends with a noble gas.

16. The Modern Periodic Table
In between, the properties of the elements change in an orderly progression from left to right.
This regular cycle illustrates periodicity in the properties of the elements.

17. The Modern Periodic Table
periodic law - physical and chemical properties of the elements repeat in a regular pattern when they are arranged in order of increasing atomic number

18. Use the periodic table to separate these 12 elements into 6 pairs fo elements having similair properties. Ca, K, Ga, P, Si, Rb, B, Sr, Sn, Cl, Bi, Br Ca K Ga P Si Cl Sr Rb B Bi Sn Br

19. Layout of the Periodic Table
Have them get out their periodic tables for labeling

20. Layout of the periodic table
A group, also called a family, consists of the elements in a vertical column.

21. Groups are numbered 1 – OR IA – VIIIA for main group elements and IB – VIIIB for transition elements
Make sure to point out the weird numbering with the transitions elements

22. As you move left to right across a period the number of valence electrons increases by one

23. Elements in the same group have same number of valence electrons and similar properties

24. A period consists of the elements in a horizontal row

25. Periods are numbered 1-7 and each new row begins a new energy level

26. The elements in the middle are called transition elements

27. The others are main group elements

28. Lithium is the first element in Group 1 and in Period 2
Lithium is the first element in Group 1 and in Period 2. Check this location on the periodic table.

29. 4 groups have commonly used names: alkali metals in Group 1 (IA)

30. the alkaline earth metals in Group 2 (IIA)

31. the halogens in Group 17 (VIIA) -from the Greek words for “salt former” , compounds that halogens form with metals are salt-like.

32. the noble gases in Group 18 (VIIIA) – full outer shell (8 valence electrons), generally unreactive

33. In the periodic table, two series of elements are placed below the main body of the table.
The elements in these two series are known as the inner transition elements.

34. The first series of inner transition elements is called the lanthanides because they follow element number 57, lanthanum.
Because of their natural abundance on Earth is less than 0.01 percent, the lanthanides are sometimes called the rare earth elements.

35. The second series of inner transition elements are the actinides
All of the actinides are radioactive, and all beyond uranium (92) are man made (synthetic).

36. Elemental Funkiness By Mark Rosengarten UF

37. Classification of elements

38. Elements are classified as metals, metalloids, or nonmetals on the basis of their physical and chemical properties.
The majority of the elements are metals (solids). They occupy the entire left side and center of the periodic table.

39. Physical States and Classes of the Elements
Nonmetals occupy the upper-right-hand corner. – green, yellow, orange

40. Physical States and Classes of the Elements
Metalloids are located along the staircase boundary between metals and nonmetals. - purple

41. Metals
Metals are elements that have luster, conduct heat and electricity, and usually bend without breaking.
All metals except mercury are solids at room temperature; in fact, most have extremely high melting points.
Click box to view movie clip.

42. Metals
With the exception of tin, lead, and bismuth, metals have one, two, or three valence electrons.
The periodic table shows that most of the metals (coded blue) are not main group elements.

43. Nonmetals
Most nonmetals don’t conduct electricity, are much poorer conductors of heat than metals, and are brittle when solid.
Their melting points tend to be lower than those of metals.
Many are gases at room temperature

44. With the exception of carbon, nonmetals have five, six, seven, or eight valence electrons.

45. Properties of Metals and Nonmetals
Copy the chart

46. Metalloids
Metalloids have some chemical and physical properties of metals and other properties of nonmetals. - purple
In the periodic table, the metalloids lie along the border between metals and nonmetals.

47. some metalloids are semiconductors
A semiconductor is an element that does not conduct electricity as well as a metal, but does conduct slightly better than a nonmetal.
Some metalloids such as silicon, germanium (Ge), and arsenic (As) are semiconductors.
Silicon’s semiconducting properties made the computer revolution possible.

48. Part II - Periodic Trends

49. Periodic Properties of the Elements
The electron structure of an atom determines many of its chemical and physical properties.
Understanding the relationship between electron configuration and position in the periodic table enables you to predict the properties of the elements and the outcome of many chemical reactions.

50. Atomic Radius
size of atom
© 1998 LOGAL

51. Atomic Size
size of an atom INCREASES in any group as you go DOWN the column because the valence electrons are in energy levels farther from the nucleus.

52. “shielding effect” – electrons in energy levels closer to the nucleus “shield” the valance electrons from the positive pull of the nucleus

53. The shielding effect
Increases down a group because electrons are being added to higher energy levels
There is no shielding effect as you go across a period because electrons are being added to the same principal energy level

54. size of an atom DECREASES in any period as you go to the RIGHT in any row because there is an increased nuclear (+) charge pulling e- in tighter.

55. Atomic Radius Why larger going down? Why smaller to the right?
Higher energy levels have larger orbitals
Shielding - core e- block the attraction between the nucleus and the valence e-
Why smaller to the right?
Increased nuclear charge without additional shielding pulls e- in tighter

56. Atomic Radii of Main Group Elements

57. Examples
Which atom has the larger radius?
Be or Ba
Ca or Br Ba Ca

58. For each of the following pairs, predict which atom is larger.
a. Mg, Sr Sr
b. Sr, Sn Sr
c. Ge, Sn Sn
d. Ge, Br Ge
e. Cr, W W

59. Octet Rule
reactivity of atoms is based on achieving a complete octet of valence electrons(8/8)
Everybody wants to be like a noble gas! Ne

60. Atoms achieve noble gas configuration by gaining or losing their valence electrons An ion is an atom or group of atoms that has a charge because of the loss or gain of electrons.

61. cation - An ion that has LOST an e- and now has a positive (+) charge anion – an ion that has GAINED an e- and now has a negative (-) charge

62. Common Ion Charges aka oxidation number1+ 1- 3+ 3- 2- 2+

63. Group IVA 4 can lose or gain Group VA 5 gains 3 Group VIA 6 gains 2
GROUP #: VALENCE # WHEN FORMING IONS:
OUT OF 8:
Group IA loses 1
Group IIA loses 2
Group IIIA loses 3
Group IVA 4 can lose or gain
Group VA gains 3
Group VIA gains 2
Group VIIA gains 1
Group VIIIA 8 does not form ions

64. Ionic Size
positive ions + (cations) acquire the configuration of the noble gas in the preceding period.
the outermost electrons of the ion are in a lower energy level than the valence electrons of the neutral atom.

65. The electrons that are not lost by the atom experience a greater attraction to the nucleus and pull together in a tighter bundle with a smaller radius.
all cations ions have smaller radii than their corresponding atoms.

66. anions acquire the electron configuration of the noble gas at the end of its period.
But the nuclear charge doesn’t increase with the number of electrons.

67. In the case of fluorine, a nuclear charge of 9+ must hold ten electrons in the F– ion; all the electrons are held less tightly
the radius of the anion is larger than the neutral atom.

68. Ionic Radius Ionic Radius Cations (+) lose e- smaller Anions (–)
gain e-
larger
© 2002 Prentice-Hall, Inc.

69. S or S2- Al or Al3+ S2- Al Examples
Which particle has the larger radius?
S or S2-
Al or Al3+
S2- Al

70. For each of the following pairs, predict which atom or ion is larger
a. Mg, Mg2+ Mg
b. S, S2–
S2–
c. Ca2+, Ba2+
Ba2+
d. Cl–, I– I–
e. Na+, Al3+
Na+

71. First Ionization Energy
ionization energy - the energy needed to REMOVE an electron from an atom, in kJ/mol
First Ionization Energy
Energy required to remove the 1st e- from a neutral atom.
© 1998 LOGAL

72. group trends
(first) ionization energy decreases from top to bottom along a group
reason: outermost electron is farther and farther from the nucleus in larger atoms, so it is more easily removed

73. periodic trends
(first) ionization energy increases from left to right in a period
reason: ―nuclear charge(+) increases; more attraction between electrons and protons

74. Successive Ionization Energies
Large jump in I.E. occurs when a CORE e- is removed.
Mg 1st I.E kJ
2nd I.E. 1,445 kJ
Core e- 3rd I.E. 7,730 kJ

75. Successive Ionization Energies
Large jump in I.E. occurs when a CORE e- is removed.
Al 1st I.E kJ
2nd I.E. 1,815 kJ
3rd I.E. 2,740 kJ
Core e- 4th I.E. 11,600 kJ

76. Examples
Which atom has the higher 1st I.E.?
N or Bi
Ba or Ne N Ne

77. For each of the following pairs, predict which atom has the higher first ionization energy.
a. Mg, Na Mg
b. S, O O
c. Ca, Ba Ca
d. Cl, I Cl
e. Na, Al Al
f. Se, Br Br

78. Periodic Trends in Electronegativity
electronegativity— tendency of an atom to attract electrons.
noble gases do not have electronegativity values
chemical bonds are determined by electronegativity differences between the bonding partners

79. electronegativity trends are not completely regular
fluorine = most electronegative element with a value of 4.0 (smallest anion formed)
cesium = least electronegative element (largest cation formed)

80. electronegativity decreases from top to bottom in a group

81. electronegativity increases from left to right in a period

82. RECAP

83. Atomic Radius
Atomic Radius
Increases to the LEFT and DOWN

84. First Ionization Energy
Increases UP and to the RIGHT

85. electronegativity
Increases UP and to the RIGHT

86. Electron affinity
Increases UP and to the RIGHT

87. Part 3: Electron Configuration

88. Electrons are arranged in orbits around the nucleus
Electrons in Atoms
Niels Bohr.
Electrons are arranged in orbits around the nucleus
The energy level of an electron is the region around the
nucleus where the
electron is likely to
be moving.

89. modern 3-D electron-cloud model - probability model
Heisenberg Uncertainty Principle — it is not possible to know both the exact position and velocity of an object simultaneously

90. Modern electron cloud model
orbitals are areas of high probability (~95%) of finding electrons

91. Electrons can change energy level, by absorbing energy
Electrons can change energy level, by absorbing energy. When an electron absorbs a quantum of energy, it moves up to a higher energy level.
When the electron falls from a higher energy level to a lower energy level, energy is released, and we see light

92. Energy levels have sublevels —divisions within an energy level
1) many similar energy states grouped together in a level
2) different shapes: spherical, dumbbell, cloverleaf

93. There are 4 sublevels
s, p, d, f
(s p d f stand for sharp, principal, diffuse, fundamental)
maximum number of e- in a principal
energy level = 2n 2
n = principal quantum number
= electron energy level or ― “shell” (period) number
n = 1, 2, 3, 4, 5, 6, 7

94. electron maximums in the sublevels
s can hold 2 e-
p can hold 6 e-
d can hold 10 e-
f can hold 14 e-

95. Electrons fill orbitals in a certain way
electron configuration - a specific electron arrangement in orbitals

96. Electron configuration: General Rules
Pauli Exclusion Principle
Each orbital can hold 2 electrons with opposite spins.

97. Aufbau Principle
Electrons fill the lowest energy orbitals first.
“Lazy Tenant Rule”

98. WRONG RIGHT Hund’s Rule
Within a sublevel, place one e- per orbital before pairing them.
“Empty Bus Seat Rule”
WRONG
RIGHT

99. Different sections of the periodic table correspond to the different sublevels
Groups IA & IIA = s block
Groups IIIA – VIIIA = p block
Transition = d block
Inner transition = f block

101. Diagonal rule -to help us remember the order in which energy level subshells fill - follow the arrows
1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 6f 7s 7p

102. 1 s2 2s2 2p4 O Example 1s2 2s2 2p4 8e- Orbital Diagram
Electron Configuration =
1 s2 2s2 2p4

103. 1s2 2s2 2p4
the sum of the superscripts = the atomic number of the element
superscripts are NOT exponents (nothing is being squared, etc.)

104. *** valence configurations will be
s OR s and p ***

105. Longhand Configuration 1s2 2s2 2p6 3s2 3p4 S 16e-
Condensed (Abbreviated) Electron Configurations
use the previous Noble Gas as the starting point in brackets, then finish the configuration
Longhand Configuration
1s2
2s2
2p6
3s2
3p4
S 16e-
Core Electrons
Valence Electrons
Shorthand Configuration
S 16e- [Ne] 3s2 3p4

106. Example: Indium #49
1) complete: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 5p1
2) condensed: [Kr] 5s2 5p1
3) valence 5s2 5p1

107. I Heart Electron Configuration
by Mark Rosengarten UF

108. THE END !