1. Chapter 6 “The Periodic Table”
The Elements by Tom Lehrer

2. Organizing the Elements
used properties of elements to sort into groups.
1829 J. W. Dobereiner arranged elements into triads – groups of 3 w/ similar properties
One element in each triad
had properties intermediate
of the other two elements
Cl, Br, and I look different,
but similar chemically

3. Mendeleev’s Periodic Table
mid-1800s, about 70 elements known
Dmitri Mendeleev – Russian chemist & teacher
Arranged elements by
increasing atomic mass

4. blanks for undiscovered elements
Mendeleev
blanks for undiscovered elements
When discovered, his predictions accurate
Problems w/ order
Co to Ni
Ar to K
Te to I

5. A better arrangement
1913, Henry Moseley – British physicist, arranged elements according to increasing atomic number

6. The Elements by Tom Lehrer

7. Periodic Law
When elements arranged in order of increasing atomic #, periodic repetition of phys & chem props
Horizontal rows = periods
7 periods
Vertical column = group (or family)
Similar phys & chem prop.
ID’ed by # & letter (IA, IIA)

8. Areas of periodic table
3 classes of elements:
1) Metals: electrical conductors, have luster, ductile, malleable
2) Nonmetals: generally brittle and non-lustrous, poor conductors of heat and electricity
Some gases (O, N, Cl)
some brittle solids (B, S)
fuming red liquid (Br)

9. 3) Metalloids: border the line-2 sides
Properties are intermediate between metals and nonmetals

10. Section 6.2 Classifying the Elements
OBJECTIVES:
Describe the information in a periodic table.

11. Section 6.2 Classifying the Elements
OBJECTIVES:
Classify elements based on electron configuration.

12. Section 6.2 Classifying the Elements
OBJECTIVES:
Distinguish representative elements and transition metals.

14. Groups of elements - family names
Group IA – alkali metals
Forms “base” (or alkali) when reacting w/ H2O (not just dissolved!)
Group 2A – alkaline earth metals
Also form bases with H2O; don’t dissolve well, hence “earth metals”
Group 7A – halogens
“salt-forming”

15. Electron Configurations in Groups
Elements sorted based on e- configurations:
Noble gases
Representative elements
Transition metals
Inner transition metals
Let’s now take a closer look at these.

16. Electron Configurations in Groups
Noble gases in Group 8A (also called Group 18)
very stable = don’t react
e- configuration w/ outer s & p sublevels full

17. Electron Configurations in Groups
Representative Elements Groups 1A - 7A
wide range of properties
“Representative” of all elements
s & p sublevels of highest energy level NOT filled
Group # equals # of e- in highest energy level

18. Electron Configurations in Groups
Transition metals in “B” columns
outer s sublevel full
Start filling “d” sublevel
“Transition” btwn metals & nonmetals

19. Electron Configurations in Groups
Inner Transition Metals below main body of PT, in 2 horizontal rows
outer s sublevel full
Start filling “f” sublevel
Once called “rare-earth” elements
not true b/c some abundant

20. Elements 1A-7A groups called representative elements
outer s or p filling 1A 8A 2A 3A 4A 5A 6A 7A

21. The group B called transition elements
These are called the inner transition elements, and they belong here

22. Group 1A called alkali metals (but NOT H)
Group 2A called alkaline earth metals H

23. Group 8A are noble gases
Group 7A called halogens

24. Let’s take a quick break……
Periodic table rap

25. 1s1 1s22s1 1s22s22p63s1 1s22s22p63s23p64s1 1s22s22p63s23p64s23d104p65s1 1s22s22p63s23p64s23d104p65s24d10 5p66s1 1s22s22p63s23p64s23d104p65s24d105p66s24f145d106p67s1 H 1 Li 3 Na 11 K 19 Rb 37 Cs 55 Fr 87
Do you notice any similarity in these configurations of the alkali metals?

26. 1s2 1s22s22p6 1s22s22p63s23p6 1s22s22p63s23p64s23d104p6 1s22s22p63s23p64s23d104p65s24d105p6 1s22s22p63s23p64s23d104p65s24d10 5p66s24f145d106p6 He
Do you notice any similarity in the configurations of the noble gases? 2 Ne 10 Ar 18 Kr 36 Xe 54 Rn 86

27. Elements in the s - blocks
Alkali metals end in s1
Alkaline earth metals end in s2
should include He, but…
He has properties of noble gases
has a full outer level of e-’s
group 8A.

28. Transition Metals - d block
Note the change in configuration.
s1 d5
s1 d10 d1 d2 d3 d5 d6 d7 d8
d10

29. The P-block p1 p2 p3 p4 p6 p5

30. F - block f1 f5 f2 f3 f4 f6 f7 f8 f9 f10 f11 f12 f14 f13
Called “inner transition elements” f1 f5 f2 f3 f4 f6 f7 f8 f9
f10
f11
f12
f14
f13

31. Each row (or period) is energy level for s & p orbitals.1 2 3 4 5 6 7
Period Number
Each row (or period) is energy level for s & p orbitals.

32. 3d 4d 5d “d” orbitals fill up in levels 1 less than period #
first d is 3d found in period 4. 1 2 3 4 5 6 7 3d 4d 5d

33. f orbitals start filling at 4f….2 less than period #1 2 3 4 5 6 7 4f 5f
f orbitals start filling at 4f….2 less than period #

34. Demo p. 165

35. Section 6.3 Periodic Trends
OBJECTIVES:
Describe trends among the elements for atomic size.

36. Section 6.3 Periodic Trends
OBJECTIVES:
Explain how ions form.

37. Section 6.3 Periodic Trends
OBJECTIVES:
Describe periodic trends for first ionization energy, ionic size, and electronegativity.

38. } Trends in Atomic Size Radius
Measure Atomic Radius - half distance btwn 2 nuclei of diatomic molecule (i.e. O2)
Units of picometers (10-12 m… 1 trillionth)

39. ALL Periodic Table Trends
Influenced by 3 factors:
1. Energy Level
Higher energy levels further away from nucleus.
2. Charge on nucleus (# protons)
More charge pulls electrons in closer. (+ and – attract each other)
3. Shielding effect

40. What do they influence?
Energy levels & Shielding have effect on GROUP (  )
Nuclear charge has effect on PERIOD (  )

41. #1. Atomic Size - Group trendsH
Going down a group, each atom has another energy level (floor)
atoms get bigger Li Na K Rb

42. #1. Atomic Size - Period Trends
left to right across period:
size gets smaller
e-’s occupy same energy level
more nuclear charge
Outer e-’s pulled closer
Here is an animation to explain the trend. Na Mg Al Si P S Cl Ar

43. Rb K
Period 2 Na Li
Atomic Radius (pm) Kr Ar Ne H 3 10
Atomic Number

44. Trends of Atomic Radius

45. ion is atom (or group of atoms) w/ + or - charge
Ions
Some compounds composed of “ions”
ion is atom (or group of atoms) w/ + or - charge
Atoms are neutral because the number of protons = electrons
+ & - ions formed when e- transferred (lost or gained) btwn atoms

46. Metals LOSE electrons, from outer energy level
Ions
Metals LOSE electrons, from outer energy level
Sodium loses 1 e-
more p+ (11) than e- (10)
+ charge particle formed…“cation”
Na+ called “sodium ion”

47. Nonmetals GAIN one or more electrons
Ions
Nonmetals GAIN one or more electrons
Cl gains 1 e-
p+ (17) & e- (18), so charge of -1
Cl1- called “chloride ion”
anions

48. #2. Trends in Ionization Energy
Ionization energy - energy required to completely remove e- (from gaseous atom)
energy required to remove only 1st e-called first ionization energy.

49. Ionization Energy
second ionization energy is E required to remove 2nd e-
Always greater than first IE.
third greater than 1st or 2nd IE.

50. Symbol First Second Third
Table 6.1, p. 173
Symbol First Second Third
HHeLiBeBCNO F Ne

51. Symbol First Second Third
HHeLiBeBCNO F Ne
Why did these values increase so much?

52. What factors determine IE
greater nuclear charge = greater IE
Greater distance from nucleus decreases IE
Filled & half-filled orbitals have lower energy
Easier to achieve (lower IE)
Shielding effect

53. Shielding
e-’s in outer energy level “looks through” all other energy levels to see nucleus

54. Ionization Energy - Group trends
going down group
first IE decreases b/c...
e- further away from nucleus attraction
more shielding

55. Ionization Energy - Period trends
Atoms in same period:
same energy level
Same shielding
Increasing nuclear charge
So IE generally increases left - right
Exceptions…full & 1/2 full orbitals

56. Both have same shielding (e- in 1st level)He
He has greater IE than H.
Both have same shielding (e- in 1st level)
He = greater nuclear charge H
First Ionization energy
Atomic number

57. These outweigh greater nuclear charge
Li lower IE than H
more shielding
further away
These outweigh greater nuclear charge H
First Ionization energy Li
Atomic number

58. greater nuclear chargeHe
Be higher IE than Li
same shielding
greater nuclear charge
First Ionization energy H Be Li
Atomic number

59. greater nuclear charge He
B has lower IE than Be
same shielding
greater nuclear charge
By removing an electron we make s orbital half-filled
First Ionization energy H Be B Li
Atomic number

60. First Ionization energyHe
First Ionization energy H C Be B Li
Atomic number

61. First Ionization energyHe N
First Ionization energy H C Be B Li
Atomic number

62. He
Oxygen breaks the pattern, because removing an electron leaves it with a 1/2 filled p orbital N
First Ionization energy H C O Be B Li
Atomic number

63. First Ionization energyHe F N
First Ionization energy H C O Be B Li
Atomic number

64. Ne has a lower IE than He Both are full, Ne has more shielding
Greater distance F N
First Ionization energy H C O Be B Li
Atomic number

65. Na has a lower IE than Li Both are s1 Na has more shieldingHe Ne
Na has a lower IE than Li
Both are s1
Na has more shielding
Greater distance F N
First Ionization energy H C O Be B Li Na
Atomic number

66. First Ionization energy
Atomic number

67. Trends in Ionization Energy (IE)

68. Driving Forces Full Energy Levels require high E to remove e-
Noble Gases = full orbitals
Atoms want noble gas configuration

69. 2nd Ionization Energy
For elements w/ filled or ½ filled orbital by removing 2 e-, 2nd IE lower than expected.
True for s2
Alkaline earth metals form 2+ ions.

70. 3rd IE
Using the same logic s2p1 atoms have an low 3rd IE.
Atoms in the aluminum family form 3+ ions.
2nd IE and 3rd IE are always higher than 1st IE!!!

71. Trends in Ionic Size: Cations
Cations form by losing electrons.
metals
Cations are smaller than the atom they came from –
they lose electrons
they lose an entire energy level.
Cations of representative elements have noble gas configuration before them.

72. Trends in Ionic size: Anions
Anions gain electrons
Anions bigger than the atom they came from –
same energy level
greater area the nuclear charge needs to cover
Nonmetals

73. Configuration of Ions
Ions always have noble gas configurations (full outer level)
Na atom is: 1s22s22p63s1
Forms a 1+ sodium ion: 1s22s22p6
Same as Ne

74. Configuration of Ions
Non-metals form ions by gaining electrons to achieve noble gas configuration.
They end up with the configuration of the noble gas after them.

75. Ion Group trends Each step down a group is adding an energy level
Li1+
Na1+
Each step down a group is adding an energy level
Ions get bigger going down, b/c of extra energy level
K1+
Rb1+
Cs1+

76. Ion Period Trends Across period
nuclear charge increases
Ions get smaller.
energy level changes between anions and cations.
N3-
O2-
F1-
B3+
Li1+
Be2+
C4+

77. Size of Isoelectronic ions
Iso- means “the same”
Isoelectronic ions have the same # of electrons
Al3+ Mg2+ Na1+ Ne F1- O2- and N3-
all have 10 electrons
all have the same configuration: 1s22s22p6 (which is the noble gas: neon)

78. Size of Isoelectronic ions?
Positive ions that have more protons would be smaller (more protons would pull the same # of electrons in closer)
N3-
O2-
F1- Ne
Na1+
Al3+ 10 9 8 7 13 12 11
Mg2+

79. #3. Trends in Electronegativity
Electronegativity is tendency for atom to attract e-’s when atom in a compound
Sharing e-, but how equally do they share it?
Element with big electronegativity means it pulls e- towards itself strongly!

80. Electronegativity Group Trend
Further down a group, farther e- is away from nucleus, plus the more e-’s an atom has
more willing to share
Low electronegativity

81. Electronegativity Period Trend
Metals let e-’s go easily
low electronegativity
Nonmetals want more electrons
take them away from others
High electronegativity.

82. Trends in Electronegativity

83. Chemistry Song "Elemental Funkiness" - Mark Rosengarten