1. Unit 5 The Periodic Table
The how and why

2. Newlands -1865
Arranged known elements according to properties & order of increasing atomic mass
Law of Octaves – pattern of chemical & physical properties repeated every 8 elements

4. Mendeleev - 1869 Created 1st periodic table (63 elements)
Ordered by increasing atomic mass
Predicted pattern of missing elements
Started new rows and lined up columns to organize elements with similar properties
Rearranged elements so similar properties would line up correctly

6. The Modern Table
Moseley- determined the atomic number for each known element.
Elements are still grouped by properties
Similar properties are in the same column
Ordered by increasing atomic number
Added a column of elements Mendeleev didn’t know about – noble gases

8. Periodic Law
When elements are arranged in order of increasing atomic number, elements with similar properties appear at regular intervals

9. Horizontal rows are called periods
There are 7 periods

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

11. Other Systems IA IIA IIIB IVB VB VIB VIIB VIIIB IIIA IVA VA VIA VIIA
VIIIA IB
IIB 1 2 13 14 15 16 17 18 3 4 5 6 7 8 9 10 11 12 1A 2A 3A 4A 5A 6A 7A 8A 3B 4B 5B 6B 7B 8B 1B 2B

12. The elements in the s & p blocks are called the representative elements

13. Transition metals
D block elements

14. These are called the inner transition elements and they belong here

15. Three Classes of Elements
Metals
Nonmetals
Metalloids

16. Metals

17. Metals Ductile – drawn into wires Malleable – hammered into sheets
All solid at room temperature (except Hg- Mercury)
Conductors of heat and electricity

18. Group 1 are the alkali metals
VERY reactive because one valence e-
Found as compounds in nature
Not including H!

19. Group 2 are the alkaline earth metals
Still highly reactive but not as much so
as alkali metals (2 valence e-)

20. Transition Metals The weird ones…
May lose different #s of valence electrons depending on the element with which it reacts
Less reactive than alkali or alkaline earth metals

21. Inner Transition Metals
1st row = lanthanides
Shiny metals similar in reactivity to alkaline earth metals
2nd row = actinides
Unstable nuclei – all radioactive

22. Non-metals

23. Non-metals Most are gases, some solid, and 1 liquid (Br)
More variation than metals

24. Group 7 is called the Halogens
Most reactive non-metals – 7 valence
React frequently with alkali metals

25. Group 8 are the noble gases
Low reactivity, very stable, inert

26. Metalloids or Semimetals

27. Metalloids Border the staircase between metals and nonmetals
Properties – similar to metals and nonmetals
Some used in semiconductors (silicon in electronics)

28. Identifying the patterns
Part 2 Periodic trends
Identifying the patterns

29. What we will investigate
Atomic size
how big the atoms are
Ionization energy
How much energy to remove an electron
Electronegativity
The attraction for the electron in a compound

30. What we will look for Periodic trends
How those things vary as you go across a period
Group trends
How those things vary as you go down a group
Why?
Explain why these variations exist

31. Atomic Size Where do you start measuring?
The electron cloud doesn’t have a definite edge.
Scientists focused first on diatomic elements -- measured more than 1 atom at a time

32. Atomic Size }
Radius
Atomic Radius = half the distance between two nuclei of molecule

33. Atomic Size - Periodic Trends
The positive nucleus pulls on electrons
Periodic trend
As you move across a period, elements have more protons
The charge on the nucleus gets bigger
The outermost electrons of each element are in the same energy level
So there is more pull on the outermost electrons as you move across

34. Periodic Trends Na Mg Al Si P S Cl Ar
As you go across a period, the radius gets smaller.
Same outermost energy level
More nuclear charge
Pulls outermost electrons closer Na Mg Al Si P S Cl Ar

35. Atomic Size – Group Trends
The positive nucleus pulls on electrons
Group Trend
As you go down a group, you add energy levels
Outermost electrons not as attracted by the nucleus

36. Shielding
Increasing numbers of electrons between the nucleus and the valence electrons tends to decrease the force between the nucleus & the valence electrons + 36

37. Shielding
The electron on the outside energy level has to look through all the other energy levels to see the nucleus +

38. Shielding
The electron on the outside energy level has to look through all the other energy levels to see the nucleus
A second electron has the same shielding
In the same energy level (period) shielding is the same +

39. Shielding As the energy levels changes the shielding changes
Moving down the group
More energy levels
More shielding
Outer electron less attracted +
Three shields
Two shields
No shielding
One shield

40. Group trends H Li Na K Rb As we go down a group
Each atom has another energy level
More shielding
The atoms get bigger Li Na K Rb

41. Rb K
Overall Na Li
Atomic Radius (nm) Kr Ar Ne H 10
Atomic Number

42. Atomic size

44. IONIZATION ENERGY

45. It’s all about stability
Alkali metals are more stable if they lose an electron
Example
Sodium ([Ne] 3s1)
Getting rid of the 3s1 electron makes sodium more stable and creates a sodium ion (Na1+)

46. Ionization Energy
The amount of energy required to completely remove an electron from a neutral atom.
The energy required for the 1st electron is called the first ionization energy

47. Ionization Energy
The 2nd ionization energy is the energy required to remove the second electron
Always greater than 1st IE
The 3rd IE is the energy required to remove a third electron
Greater than 1st or 2nd IE

48. Symbol First Second Third
HHeLiBeBCNO F Ne

49. Group trends As you go down a group first IE decreases
Valence e- farther from nucleus
More shielding

50. Periodic trends First IE increases from left to right across a period
Increased nuclear charge from added proton
Electron shielding not an issue b/c valence are all in same energy level
Exceptions at full and 1/2 full orbitals
Lower IE b/c offer stability to atom

51. How to remember? HI LO LO

52. 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 F N
First Ionization energy H C O Be B Li Na
Atomic number

53. First Ionization energy
Atomic number

54. Electronegativity

55. Electronegativity
There’s an electron tug of war between atoms in a compound
The tendency for an atom to attract electrons to itself when it is chemically combined with another element
How “greedy”
Large electronegativity means the atom pulls the electron towards itself

56. Group Trend As you move down a group More shielding
Less attraction for electrons
Lower electronegativity

57. Periodic Trend As you move across a period from left to right,
Nuclear charge increases
Greater electronegativity

58. How to remember? HI LO LO

59. All 3 trends HI LO LO