1. 12/2/14 Today I will discuss the development of the periodic table
Warm Up – Name two ways we have used the periodic table.

2. CP Chemistry Chapters 5/6
The Periodic Table
CP Chemistry
Chapters 5/6

3. Periodic Table To organize the elements!
Why did scientists develop the periodic table?
To organize the elements!
The periodic table worked so well that scientists were actually able to predict elements that had not been discovered yet!

4. Forerunners 1700’s – only about 30 elements known
1800’s – about 60 elements known
Scientists needed a good way to classify the elements

5. Periodic Table
Using either the book, or your “smart” devices, you have 20 minutes to answer the questions regarding the scientists of the periodic table.
Ready, set, go!

6. J.W. Dobereiner – early 1800’s
Triads – groups of three elements with similar chemical properties
Li, Na, K
Ca, Sr, Ba
Cl, Br, I
The middle element has approximately the average mass and average density of the other two!

7. J.A.R. Newlands
Law of Octaves – when elements are arranged by increasing atomic mass, property patterns repeated every eight elements.
1st element like the ninth
2nd element like the tenth, etc.
No one believed him because of the relationship between chemistry and music!

8. Dmitri Mendeleev - 1869 The Father of the Periodic Table!
Placed elements on index cards and arranged them based on repeating patterns of properties
Increasing Atomic Mass
Greatest accomplishment was leaving blank spaces for elements not yet discovered!

9. Dmitri Mendeleev
The Father of the Periodic Table!

10. Henry Moseley – early 1900’s
Recognized that the patterns are based on atomic number, not atomic mass.
Fixed a few issues with Mendeleev’s table.
Ar (At# 18, AtW 39.95)
K (At# 19, AtW 39.01)
Periodic Law – properties of elements are periodic functions of their atomic number.
Cycles are observed as the Atomic Number increases

11. Homework
5-1 R&R (Front only)

12. 12/3/14 Objective – identify the organization of the periodic table
Warm up – What did Newlands contribute to the understanding of the elements?

13. Periodic Table Activity
Using your smart devices or the book, complete the periodic table activity

14. 12/4/14 Objective – identify the organization of the periodic table
Warm up – Give 2 examples of alkaline earth metals.

15. Reading the Periodic Table
Modern Periodic Table - Arrangement of elements in order of their atomic number so that elements with similar properties fall into the same column.
Groups or Families – vertical columns
Periods – Horizontal rows of elements

16. “s” n= 1 2 3 4 5 6 7
“p”
“d”
“f”

17. “s” & “p” = main-group elements
“d” = transition metals
“f” = inner transition metals

19. Periodic Blocks
What do the blocks of the periodic table tell us?

20. Group Numbers 1 18 1A 8A 2 13 14 15 16 17 2A 3A 4A 5A 6A 7A 3 4 5 6 79 10 11 12 3B 4B 5B 6B 7B 8B 1B 2B

21. Family Names Group 1 – Alkali metals Group 17 – Halogens
Group 2 – Alkaline-earth metals
Group 17 – Halogens
Group 18 – Noble gases
Groups 13-16 1 2 3 4 5 6 7 8 9 10 11 12 18 13 14 15 16 17

22. Lanthanides Actinides Groups 3-12 - Transition Metals
Below the Table - Inner Transition Metals
Lanthanides
Actinides 1 2 3 4 5 6 7 8 9 10 11 12 18 13 14 15 16 17

23. Family Properties
Alkali metals – very malleable, soft metals, very reactive
Alkaline Earth metals – malleable, harder metals, reactive
Halogens – gaseous (F, Cl, Br) or solid (I, At) nonmetals, often toxic, very reactive
Nobel gases – inert gases

24. =metaloid (semi-metal)
Metals/Nonmetals
Non-Metals
Metals

25. Properties of Metals/Nonmetals
Luster or shine
Good conductors of heat and electricity
Typically solid at room temp (except Hg)
Malleable – can be hammered into sheets
Ductile – can be drawn into wire

26. Properties of Metals/Nonmetals
Typically dull finish
Poor conductors of heat and electricity
Variety of forms at room temp (solid, liquid, gas)
Neither malleable nor ductile
Metalloids (semi-metals)
Properties of both metals and non-metals
Useful in computers

27. Families The electrons in their outermost shells are the same!
Why do elements in the same family have similar properties?
Let’s look at the electron configurations
H – 1s1
Li – 1s22s1
Na – 1s22s22p63s1
K – 1s22s22p63s23p64s1
Rb - 1s22s22p63s23p64s23d104p65s1
The electrons in their outermost shells are the same!

28. Families Why is that important?
Valence electrons – electrons in the outermost shell responsible for an atom’s chemical behavior and properties.

29. Homework
5-2 R&R

30. 12/5/14
Today I will determine the periodic trends for atomic radius and ionic radius Warm Up – How many valence electrons do the halogens have? Hint: Look at the electron configuration for F & Cl.

31. Periodic Trends
Periodic Trends- properties that change in predictable ways as you move through the periodic table.

32. Atomic Radius Atomic Radius –
The distance from the center of an atom’s nucleus to its outermost electron
(half the distance between two neighboring nuclei)

33. Atomic Radius Atoms get larger going down a column
New energy levels are added
Atoms get smaller moving across a period
???? – Aren’t we adding electrons?
We are also adding protons which pull the electrons in tighter!

34. Ionic Radius
Ionic Radius – Atomic radius of an atom that has gained or lost an electron
Where do ions come from?

35. Ionic Radius The electrons in their outermost shells are the same!
Remember Valence Electrons:
Let’s look at the electron configurations
H – 1s1
Li – 1s22s1
Na – 1s22s22p63s1
K – 1s22s22p63s23p64s1
Rb - 1s22s22p63s23p64s23d104p65s1
The electrons in their outermost shells are the same!

36. Valence Electrons
Elements of the same family also have the same number of valence electrons (outermost shell) 1 8 2 3 4 5 6 7

37. Ionic Charge – common charge
Octet Rule – Atoms will lose or gain electrons to have a full set of valence electrons
All atoms want to have 8 electrons in their outermost shells (except only 2 in the first shell).
When they do they have the same configuration as the closest noble gas!
Noble gas configurations are very stable!

38. Ionic Charge – common charge
Elements of the same family tend to form the same ions! +1 +2 +3 ±4 -3 -2 -1

39. Ionic Radius
How does gaining or losing an electron affect the atomic radius?
Cations – positive ions that have lost electrons
Smaller than the neutral counterpart because they have less electrons to pull in, so they can be pulled tighter!
Anions – negative ions that have gained electrons
Larger than the neutral counter part because they have more electrons to spread out!

40. Ionic Radius
Cations (smaller) tend to occur to the left of the periodic table
Anions (larger) tend to occur to the right of the periodic table

41. Homework
Atomic/Ionic Radius WS

42. 12/8/14
Today I will name and explain trends that occur on the periodic table.
Warm Up – Which atom in each pair is larger?
a) Al or Si
b) Ge or Sn
c) Ba or Pb
d) Mg or its most common ion
e) S or its most common ion

43. Valence Electrons
Elements of the same family also have the same number of valence electrons (outermost shell) 1 8 2 3 4 5 6 7 2

44. First Ionization Energy
Ionization Energy – the energy needed to remove an electron
The more tightly bound the electrons, the harder they are to remove.
Greater as you move across
More protons provide more pull
Smaller as you move down
Added energy levels because they are further from the protons’ pull
Electron shielding – Inner electrons shield outer ones from the pull of the nucleus

45. Successive Ionization Energy
Successive Ionization Energy – the energy needed to remove another electron after the first
Greater because of the greater number of protons than electrons
Much greater jump as you move to closer energy levels

46. Valence Electrons
Mg – Why would the third ionization be much larger than the second?

47. Electronegativity
Electronegativity – the ability of an atom to attract an electron
All atoms want to have 8 electrons in their outermost shells (except only 2 in the first shell).

48. Electronegativity Which elements want to gain electrons? 1 8 2 3 4 5 67 2

49. Electronegativity Higher as you move across
elements are more likely to gain an electron due to valence number. (until noble gases!)
Electronegativity decreases down a group
easier for the nucleus to pull in an electron that is closer (in a lower energy level).

50. Electron Affinity
Electron Affinity - the energy change that occurs when an electron is acquired by a neutral atom.
Related to electronegativity because the more electronegative an atom, the more energy it will aquire
Follows the same trend as electronegativity

51. Trend Review Atomic Radius Decreases Ionization Energy Increases
Electronegativity/Electron Affinity Increases
Atomic Radius Increases
Ionization Energy Decreases
Electronegativity/Electron Affinity Decreases

52. Trend Review Cations on the left are generally smaller
Anions on the right are generally larger

53. Special Cases/Exceptions
Hydrogen – hydrogen resides in group 1, but it’s not a metal! It is a very different element.
First energy level –
We said that all elements want 8 electrons in the outermost shell, but there is an exception.
The first energy level, only wants two electrons! (Hydrogen and Helium)
This is why helium (with only 2 valence) is a noble gas!

54. Homework Ionization energy, Electronegativity & Electron Affinity WS
Graphing Project starting tomorrow

55. 12/9/14 Today I will graph and analyze periodic trends
Warm Up – Explain why the first ionization energy trend is what it is…

56. 12/10/14 Today I will graph and analyze periodic trends
Warm Up – Explain the jump in ionization energies between a nobel gas and the following alkali metal.

57. 12/11/14 Today I will graph and analyze periodic trends
Warm Up – Sketch the periodic table and give an overview of each of the trends

58. 12/12/14 Today I will write noble gas configurations
Warm Up – Write the electron configuration for Ge.

59. Noble Gas Configuration
Let’s look at the electron configurations again
H – 1s1
Li – 1s22s1
Na – 1s22s22p63s1
K – 1s22s22p63s23p64s1
Rb - 1s22s22p63s23p64s23d104p65s1
Everything up to the last numbers is always the same!
Is there a shortcut we can take?

60. Noble Gas Configuration
1. Place the name of the noble gas before the element in brackets, and
2. continue writing the electron configuration.
Na – 1s22s22p63s1
[Ne]3s1
Ge - 1s22s22p63s23p64s23d104p2
[Ar] 4s23d104p2

61. Writing in Short-hand notation
Write the short hand notation for the following: V Rb I Hg U W
[Ar] 4s2 3d3
[Kr] 5s1
[Kr] 5s2 4d10 5p5
[Xe] 6s2 4f14 5d10
[Rn] 7s2 5f4
[Xe] 6s2 4f14 5d4

62. Homework
Finish the graphing project – due Monday!

63. 12/15/14 Today I will review periodic trends
Warm Up – Why do metals tend to form positive ions while non-metals tend to form negative ions?

64. Periodic Trends Review
1. Periodic Law states that elements are a periodic function of their atomic #. This means that you see repeating patterns when elements are lined up by atomic #. This is important because it’s the basis of the modern periodic table!

65. Periodic Trends WS

66. 12/16/14 Today I will review chapter 5/6
Warm Up – Write the noble gas configurations for Se, Nb, and Sm.

67. 12/17/14 Today I will review for the chapter 5/6 exam
Warm Up – List and describe the 4 people responsible for the development of the modern periodic table.