1. Development of the Modern Periodic Table
DAY 1

2. Why is the Periodic Table important to me?
The periodic table is the most useful tool to a chemist.
You apply information from it on every test.
It organizes lots of information about all the known elements.

3. Pre-Periodic Table Chemistry …
…was a mess!!!
There were at least 55 elements discovered
No organization of elements.
Varying properties amongst the elements, but some similar trends
Difficult to find information.

4. History of the Periodic Table

5. 1. During the nineteenth century, how did the chemists began to categorize the elements on the periodic table? What was the end result of this method?
The chemists began to categorize the elements according to similarities in their physical and chemical properties.
The end result of this method was our modern periodic table we use today 

6. Physical Properties? Chemical Properties?

7. Physical properties:
Physical properties are characteristics that can be observed or measured without changing the chemical formula of the substance
Mass and volume are physical properties. So are color, shape, flexibility, physical state (solid, liquid, gas) and conductivity.

8. Summary: Physical Properties
Properties that can be observed or measured without changing what you have
Examples: mass, viscosity, density, conductivity, color, size, melting point

9. Matter also has chemical properties:
Chemical property: the ability of something to react with or change into a different substance
For example, magnesium is
flammable. When it burns,
it turns into magnesium oxide.
Flammability is a chemical
property.

10. Chemical properties Examples: Metals can react with acids
Chlorine can react with sodium
Copper can react with oxygen
Iron can rust
Neon doesn’t react with anything!

11. WHO: Dmitri Mendeleev Modern Periodic Table 1834 - 1907 WHEN: WHAT:
EXAMPLE:
Dmitri Mendeleev
1869
Modern Periodic Table
Organized periodic table by increasing atomic mass

12. Mendeleev’s Periodic Table
Mendeleev looked at all known elements and began to arrange them
Studied
Number of valence electrons
Similar chemical and physical properties
Atomic weight

13. Dmitri Mendeleev: Father of the Table
HOW HIS WORKED…
Put elements in rows by increasing atomic weight.
Put elements in columns by the way they reacted.
stated that if the atomic weight of an element caused it to be placed in the wrong group, then the weight must be wrong. (He corrected the atomic masses of Be, In, and U)
was so confident in his table that he used it to predict the physical properties of three elements that were yet unknown.

14. Elements known at this time
Mendeleev arranged the elements in order of increasing atomic mass.
He left vacant spaces where unknown elements should fit.

15. FUN FACT
Prior to his development of the Periodic Table, he was known more for his wild hair and beard, which he had trimmed just once a year.

16. Periodic Table Geography

17. Periodic Law
When elements are arranged in order of increasing atomic number, there is a periodic pattern in their physical and chemical properties.

18. The horizontal rows of the periodic table are called PERIODS.

19. The vertical columns are called GROUPS, or FAMILIES.
The elements in the same group have similar physical and chemical properties!

20. Nonmetals to the right of the zig zag line
Metals to the left of the zig zag line
Metalloids border the zig zag line
(B to At, NOT Al or Po)

21. Periodic Table Geography Fill out the following on your Periodic Table!

23. Metals vs. Nonmetals
Outline the metals red.

24. metals

25. Metals vs. Nonmetals Outline the metals red.
Outline the nonmetals blue.

26. metals
nonmetals

27. Metals vs. Nonmetals Metalloids Outline the metals red.
Outline the nonmetals blue.
Outline the metalloids green
Metalloids

28. metals
nonmetals
metalloids

29. CUT!

30. Development of the Modern Periodic Table
DAY 2

31. Characteristics of Elements
Metals
Shiny when smooth and clean
Solid at room temp.
Good conductors of heat and electricity
Malleable
Ductile
Nonmetals
Generally gases or brittle, dull-looking solids
Poor conductors of heat and electricity
Metalloids
– Have physical and chemical properties of both metals and nonmetals!

32. Metals and nonmetals can be split into 5 subgroups
Alkai Metals
Alkaline Earth Metals
Transition Metals
Halogens
Noble Gases

33. _s1
Alkali Metals

34. ALKALI METALS
Group1; 1 valence electron (one electron in outer orbital), representative elements, very reactive metals
Have to be stored in a solution, Do not occur individually in nature
malleable, ductile, good conductors of heat and electricity. (properties of metals)
softer than most other metals, can cut with a knife
can explode if they are exposed to water

35. _s2
Alkaline Earth Metals

36. Group 2; representative elements, two valence electrons
ALKLINE EARTH METALS
Group 2; representative elements, two valence electrons
Not quite as reactive because has to lose two electrons in outer orbital to combine, still very reactive
Properties that of METALS
Not found free in nature

37. d - block
Transition Metals

38. TRANSITION METALS
Groups 3-12; Can’t easily tell how many valence electrons
ductile and malleable; conduct electricity and heat
iron, cobalt, and nickel, are the only elements known to produce a magnetic field.

39. InnerTransition Metals
These elements are also called the rare-earth elements.
f block
InnerTransition Metals

40. INNER TRANSITION ELEMENTS
“F-Block”
4f = lanthanide and 5f = actinide
Many are man-made; not necessarily rare
Used in metallurgy, ceramics, glass making, dyes, computers, televisions and other electrical components
They tend to be black or dark brown in color with reddish, yellowish, or more commonly brownish streaks

41. 2p1 = Boron Group, 3 valence electrons
2p2 = Carbon group, 4 valence electrons
2p3 = Nitrogen group, 5 valence electrons
2p4 = Oxygen group, 6 valence electrons

42. _ p5
Halogens

43. HALOGENS Group 17; 7 valence electrons, representative elements
“Halogen" means "salt-former" and compounds containing halogens are called "salts“.
Exist in all three states of matter:
Solid- Iodine, Astatine
Liquid- Bromine
Gas- Fluorine, Chlorine
Extremely reactive nonmetals

44. _ p6
Noble Gases

45. NOBLE GASES
Do not form compounds easily because outer shell of electrons is full and does not lose or gain electrons (does not need to bond to be in most stable form)
He has only 2 valence electrons, but has a full valence shell. All other noble gases have 8 valence electrons.
He: used with O2 for deep sea dives and hot air balloons; Ar and Ne in lights; Kr/Xe in photographic flashes and strobe lamps

46. Cold Call Questions
Name an element that would have properties similar to calcium.
What is the name of the group that bromine is in?
How many valence electrons does a strontium atom have?

47. CUT!

48. Classification of the Elements
DAY 3

49. Lewis Dot Structures
Lewis dot structures show us how many electrons are in the valence shell.
Ex: B
3 valence electrons
Ex: H
1 valence electron

50. Draw a lewis dot structure for the following elements:
Nitrogen Magnesium Chlorine

51. Draw a lewis dot structure for the following elements:
Sulfur Lithium Bromine

52. Bromine Why does Bromine have only 7 valence electrons?
Remember energy levels and sublevels?
How many electrons are in bromine’s outermost level?

53. An atom with MORE than 4 valence electrons
Octet Rule: atoms ability to gain, lose or share electrons to acquire a full set of 8 valence electrons
An atom with MORE than 4 valence electrons
Greater ability to attract an electron, most likely GAIN an electron.
An atom with LESS than 4 valence electrons
Lower ability to attract an electron, most likely DONATE an electron.
An atom WITH 4 valence electrons, it has relatively the same ability to either donate or attract electrons.

54. **REMEMBER**
If an atom GAINS an electron (which holds a negative charge) it becomes NEGATIVE, this is called an anion (NEGATIVE).
If an atom LOSES an electron (which holds a negative charge) it becomes POSITIVE, this is called a cation (POSITIVE ION).

55. CUT!

56. Periodic Trends
DAY 4

57. Periodic Trends A periodic trend is a trend that repeats
Periodic Trends A periodic trend is a trend that repeats. Which of these actions is periodic? A. School starts in August. B. A hurricane causes massive power outages and shuts down the schools for two weeks.

58. All the periodic trends can be understood in terms of three basic rules: 1. Electrons are attracted to the protons in the nucleus. A. The closer an electron is to the nucleus, the more strongly it is attracted. B. The more protons in the nucleus, the more strongly an electron is attracted.

59. 2. Electrons are repelled by other electrons in an atom
2. Electrons are repelled by other electrons in an atom. If there are other electrons between a valence electron and the nucleus, the valence electron will be less attracted to the nucleus. That’s called shielding. Definition: Shielding – the effect of inner electrons on the attraction between the valence electrons and the nucleus

60. 3. Completed shells are very stable
3. Completed shells are very stable. Atoms prefer to add or subtract valence electrons to get 8 electrons in the outer energy level. Exception: Helium is stable with 2 electrons in its outer (and only) energy level.

61. Trend Number One: Atomic radius

62. Trend Number One: Atomic radius Moving from left to right across a period, atomic radius decreases. Why? As you move across a row, the effective nuclear charge increases. Effective nuclear charge is the charge “felt” by the valence electrons after you have taken into account the number of shielding electrons that surround the nucleus. Which is larger, an atom of sulfur or an atom of argon?

63. Trend: Atomic radius Moving down a group, atomic radius increases. Why
Trend: Atomic radius Moving down a group, atomic radius increases. Why? As you move down a group, there are more energy levels occupied by electrons. Which is larger, an atom of oxygen or an atom of sulfur?

65. Trend: Atomic radius Metallic ions are smaller than atoms. Why
Trend: Atomic radius Metallic ions are smaller than atoms. Why? Metals lose electrons to form ions, and an entire energy level disappears. Which is larger, a sodium atom or a sodium ion?

66. Trend: Atomic radius Non-metals form negatively charged ions
Trend: Atomic radius Non-metals form negatively charged ions. Negatively charged ions are larger than atoms. Why? Adding electrons to an energy level increases the repulsion between the electrons, so they spread out more. Which is larger, a chlorine atom or a chlorine ion?

67. Practice Questions
Use the periodic table to choose the largest atom or ion: a. Al, Ga, Ge b. N, F, P c. F atom, F ion, Cl atom, Cl ion

68. Trend Number Two: Ionization Energy But, wait
Trend Number Two: Ionization Energy But, wait! What is ionization energy? Ionization energy is the amount of energy required to remove an electron from an atom in the gaseous state.

69. Trend Number Two: Ionization Energy Moving across a period, ionization energy increases. Why? As you move across a period, the effective nuclear charge increases, so the electrons are more attracted to the nucleus. Which has a higher ionization energy, oxygen or fluorine?

70. Trend: Ionization Energy Moving down a group, ionization energy decreases. Why? As you move down a group, the valence electrons are farther from the nucleus and held more loosely. It takes less energy to remove a “loose” electron.

71. Trend 3: Electronegativity Electonegativity refers to how strongly the nucleus of an atom attracts the electrons from other atoms it is bonding with.

72. Trend 3: Electronegativity Moving from left to right across a period, electronegativity increases. Why? 2 reasons: 1. The effective nuclear charge increases as you move across a period. 2. The closer the nucleus is to the outside of the atom, the more likely it is to attract an electron that is in a chemical bond. Atoms get smaller going left to right, so the nucleus is closer to the outside of the atom.

73. Electronegativity
Moving down a group, electronegativity decreases. Why? As you move down a group, the atoms get larger and larger, so the nucleus is farther away from the bonding electrons. Which element is the most electronegative? Fluorine