1. The Periodic Table Chapter 5
Video clip

2. What Does Periodic Mean?
To answer:
Turn to the picture on pg. 133, Fig. 1
What things occur periodically?
Moon phases, magazine publications
Keep these in mind as we learn about the periodic table

3. History
Russian scientist Dmitri Mendeleev taught chemistry in terms of properties.
Mid molar masses of elements were known.
Wrote down the elements in order of increasing mass.
Found a pattern of repeating properties.

4. The Table
Compare this table with our modern table

5. The Table

6. Mendeleev’s Table
Grouped elements in columns by similar properties in order of increasing atomic mass.
Found some inconsistencies - felt that the properties were more important than the mass, so switched order
Example: Iodine after Tellurium
Found some gaps.
Must be undiscovered elements.
Predicted their properties before they were found.

7. Interestingly… Mendeleev never won a Nobel Prize.
He was nominated shortly before his death, but lost to Henri Moissan, who discovered Fluorine in 1906

8. Periodic Law
Why could most of the elements be arranged in the order of increasing mass, but a few could not?
Henry Moseley- discovered that atomic number should be the basis for the table, not mass
Example: I = 53, Te = 52
Physical and chemical properties of the elements are periodic functions of their atomic numbers

9. The Modern Table Elements are still grouped by properties.
Similar properties are in the same column.
Order is in increasing atomic number.
Added a column of elements Mendeleev didn’t know about.
The noble gases weren’t found because they didn’t react with anything.
So a new group was formed

10. Horizontal rows are called periods
There are 7 periods

11. Vertical columns are called groups.
Elements are placed in columns by similar properties.
Also called families

12. The elements in the A groups are called the representative elements

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

14. Group 1A are the alkali metals
Group 2A are the alkaline earth metals

15. Group 7A is called the Halogens
Group 8A are the noble gases

16. Review Rewrite the periodic law in your own words
What group are the Alkali Metals in?
Halogens?

17. Why do families have the same properties?
The part of the atom another atom “sees” is the electron cloud.
More specifically the outside (valence) orbitals.
A “family’s” orbitals fill up in a regular pattern.
The outside orbital electron configuration repeats.
The properties of atoms repeat.

18. H 1 Li 3 Na 11 K 19 Rb 37 Cs 55 Fr 87
1s1
1s22s1
1s22s22p63s1
1s22s22p63s23p64s1
1s22s22p63s23p64s23d104p65s1
1s22s22p63s23p64s23d104p65s24d10 5p66s1
1s22s22p63s23p64s23d104p65s24d105p66s24f145d106p67s1

19. 1s2
1s22s22p6
1s22s22p63s23p6
1s22s22p63s23p64s23d104p6
1s22s22p63s23p64s23d104p65s24d105p6
1s22s22p63s23p64s23d104p65s24d10 5p66s24f145d106p6 He 2 Ne 10 Ar 18 Kr 36 Xe 54 Rn 86

20. S- block: Reactive Metals
Alkali metals (column 1) all end in s1
Alkaline earth metals (column 2) all end in s2
really have to include He but it fits better later.
He has the properties of the noble gases.
Demo- Ca and Mg in test tube (pg. 142)

21. Transition Metals -d block

22. The P-block: Main Group Elements

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

24. 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.

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

26. 1 2 3 4 5 6 7 4f 5f
f orbitals start filling at 4f

28. Review
Create an acronym for remembering the position of the s, p and d blocks of the table

29. The s-Block Elements Quantum Formula: ns1,2 Group 1 Group 2
1 s electron
2 s electrons
Alkali metals
Alkaline-earth metals
Extremely reactive
A little less reactive
silvery
vary
soft
Stronger, harder
Not found in nature
Melt at lower temps
Melt at higher temps
Quantum Formula: ns1,2

30. Example problems Quantum Formula: ns1,2 Sample Problem A on pg. 143
Without looking at the periodic table, identify the group, period, and block that [Xe] 6s2 is located
Answer: group 2, sixth period, s block
Write the electron config. For the Group 1 element in the third period. Will it be more reactive or less?
Answer: Grp 1, third period= 1s22s22p63s1
Must be more reactive, because it’s in
group 1

31. The d-Block d block Quantum Formula: ns0-2 (n-1)d1-10
Groups 3-12
5 orbitals, 10 e’ total
Transition metals
Extremely reactive
Conduct electricity
High luster
Less reactive
Not always the same outer e’configuration
d block

32. Example Problem Sample Problem B on pg. 146
For d-block problems, identify group using this formula: d + s
Without using the periodic table, identify the period, block and group of [Kr] 4d55s1
Answer: 5 period, d-block, group 6
Molybedenum
Quantum Formula: ns0-2(n-1)d1-10

33. p- Block Group number – 10 for electrons Main group elements
Groups 13-18
Group number – 10 for electrons
Main group elements
Properties vary
Nonmetals, metalloids, halogens
Not always the same outer e’configuration
Quantum Formula: ns2np1-6

34. Example Sample Problem C on pg. 148
Write the outer electron configuration for the Group 14 element in the second period, name it and identify it as a metal, nonmetal or metalloid
Group number is higher than 12- so it’s the P block, second period makes n=2, group number -10 = 14-10= 4
So, you have 2 left over for the p’s
2s22p2
Quantum Formula: ns2np1-6

35. Halogens (p-block) Metalloids (p-block) Group 17 Gases, mostly
Most reactive of nonmetals
Seven valence electrons
Metalloids (p-block)
Groups btw metals and nonmetals
solids
Electrical conductivity

36. Metals (p-block) Problem Examples
Harder and denser than alkaline-earth metals, but softer than d-block
Found only in compounds
Problem Examples
Quantum Formula: ns2np1-6
Without the table, write the outer electron config. For Group 14 in second period. Name it and classify it as a metal, nonmetal or metalloid
Answer: 14= p block, 14-10= 4 so 2s22p2, carbon= nonmetal

37. f- Block Sixth and seventh periods La-Hf = Cerium-Lutetium
Ac-Rf = Thorium-Lawrencium
Mostly lab made (sythetic)

38. 1 2 3 4 5 6 7 4f 5f
f orbitals start filling at 4f

39. More Example Problems
Name the block and group for each and identify as metal, nonmetal or metalloid
[Xe]4f145d96s1
Answer: d-block, group 10, Pt, metal (period 6)
[Ne]3s23p6
Answer: p-block, group 18, Ar, nonmetal (period 3)
Group #
Group Config
Block
Comments
1, 2
ns1,2 s
One or two electrons in ns
3-12
ns0-1(n-1)d1-10 d
Sum of electrons in ns and (n-1)d equals group number
13-18
ns2np1-6 p
Number of electrons in np sublevel equals group number +/- 10

40. Metals Nonmetals Metalloids Halogens Nobel gases Alkali Alkali-Earth
Transition

41. Driving Force of Atoms Full Energy Levels are very stable
Noble Gases have full orbitals.
Atoms behave in ways to achieve noble gas configuration.

42. Atomic Size First problem where do you start measuring.
The electron cloud doesn’t have a definite edge.
They get around this by measuring more than 1 atom at a time.

43. Atomic Size }
Radius
Atomic Radius = half the distance between two nuclei of a diatomic molecule.

44. What’s A Trend? Name some fashion trends Color trends?
Behavior trends?

45. Trends in Atomic Size Influenced by two factors. Energy Level
Higher energy level is further away.
Charge on nucleus
More charge pulls electrons in closer.

46. Periodic Trends Na Mg Al Si P S Cl Ar
As you go across a period the radius gets smaller.
They have the same energy level, though.
More nuclear charge.
Outermost electrons are closer. Na Mg Al Si P S Cl Ar

47. Group trends H Li Na K Rb As we go down a group
Each atom has another energy level,
So the atoms get bigger. Li Na K Rb

48. Ionic Size Cations form by losing electrons. Groups 1-3
Form positive ions
Groups 1-3
Cations are smaller than the atom they come from.
Metals form cations

49. Ionic size Anions form by gaining electrons.
Form negative ions
Anions are bigger than the atom they come from.
Nonmetals form anions.

50. Ionization Energy- Pg. 153
The amount of energy required to remove an electron from an atom (only deals with losing an e’)
Removing one electron makes a +1 ion.
The energy required is called the first ionization energy.
Measured in kilojoules per mole
If an atom has a low IE, it will release an electron easier than one with a high IE (making them more reactive)
So, which group would have the highest IE?
Nobel gases

51. Ionization Energy
The second ionization energy is the energy required to remove the second electron.
Always greater than first IE.
The third IE is the energy required to remove a third electron.
Greater than 1st of 2nd IE.

52. Symbol First Second Third
HHeLiBeBCNO F Ne

53. Symbol First Second Third
HHeLiBeBCNO F Ne

54. Group trends As you go down a group first IE decreases because
The electron is further away.
More shielding occurs
What’s shielding?

55. Shielding
The electron on the outside energy level has to look through all the other energy levels to see the nucleus
Inner shell electrons “shield” nuclear charge from outer shell electrons

56. Period trends
All the atoms in the same period have the same energy level.
Same shielding.
Increasing nuclear charge
So IE generally increases from left to right
A higher charge will more strongly attract electrons, holding them “hostage”

57. First Ionization energyHe
He has a greater IE than H.
same shielding
greater nuclear charge H
First Ionization energy
Atomic number

58. outweighs greater nuclear charge First Ionization energyHe
Li has lower IE than H
more shielding
further away
outweighs greater nuclear charge H
First Ionization energy Li
Atomic number

59. greater nuclear charge First Ionization energyHe
Be has higher IE than Li
same shielding
greater nuclear charge
First Ionization energy H Be Li
Atomic number

60. 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

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

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

63. First Ionization energyHe
Breaks the pattern because removing an electron gets to 1/2 filled p orbital N
First Ionization energy H C O Be B Li
Atomic number

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

65. First Ionization energyHe Ne
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

66. 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

67. First Ionization energy
Atomic number

68. Electron Affinity- Pg 157
The energy associated with adding an electron to an atom (only deals with gaining e’)
Easiest to add to group 7A.
A highly negative number = a high EA, that means the atom will gain electrons easily
Trends:
EA increases from left to right because atoms become smaller, with greater nuclear charge.
EA decrease as we go down a group.
Also measured in kJ/mol

69. Electronegativity- Pg. 161
The tendency for an atom to attract electrons to itself when it is chemically combined with another element.
How fair it shares the electron with the atom its bonding with
Big electronegativity means it pulls the electron toward it- it’s a bully!
Atoms with large negative electron affinity have larger electronegativity.
Flouine is the boss!

70. Valence Electron
So, Chlorine isn’t sharing the electron fairly with sodium, because it has such a large electronegativity Na Cl
Electronegativity= 3.0
Electronegativity= 0.9

71. Valence Electron
So, Chlorine isn’t sharing the electron fairly with sodium, because it has such a large electronegativity Na Cl
Electronegativity= 3.0
Electronegativity= 0.9

72. Group Trend
The further down a group, the farther the electron is away from the nucleus because there are more energy levels
Therefore, atoms are more willing to share these electrons.
Low electronegativity.

73. Periodic Trend Cations let their electrons go easily
Low electronegativity
Anions want more electrons- try to steal them
High electronegativity

74. Ionization energy, electronegativity
Electron affinity INCREASE

75. Atomic size increases, shielding constant
Ionic size increases