1. Periodic Table

2. Antoine Lavoisier (1700’s)
Listed all known elements (33) at the time
4 groups: gases, metals, nonmetals, and earths

3. Dobereiner (early 1800’s)
Dobereiner arranged the elements into triads (groups of three elements) based on similarities in properties.

4. John Newlands (1864) Arranged elements by increasing atomic mass (70)
Noticed a repeating pattern of properties
Created the law of octaves (repeating patterns at every eighth element)

5. Newlands

6. Lothar Meyer (1869)
Identified and proved that there was a connection between atomic mass and the property of the element
Arranged the elements by increasing atomic mass (added the new ones)

7. Dmitri Mendeleev (1869)
Proved a connection between atomic mass and element properties
Arranged elements by increasing atomic mass
Predicted the existence and properties of elements yet to be discovered

8. Henry Moseley (1913) Discovered atomic number
Arranged elements by increasing atomic number
By doing this a pattern of properties was discovered and fixed previous problems

9. Periodic Law
When elements are arranged by increasing atomic number, there is a periodic repetition of physical and chemical properties.

10. Modern Periodic Table
Periods (rows)- contain a variety of elements ranging from metals to nonmetals to Noble gases. There are 7.
Groups or Family (columns)- contain elements that share similar properties. There are 18.
Periods 

11. Representative (Main) Elements
Marked by “A” on most groups.
Elements in the ‘s’ and ‘p’ block
Wide range of characteristics
This is Newlands octaves

12. Transition Elements (B)
Consists of only metals.
Found in the center of the period table.
Elements of the “d” block

13. Metals, Nonmetals, and Metalloids

14. Metals Make up most of the periodic table
Solid at room temperature (except Mercury)
Good conductors of heat and electricity
Ductile and malleable
Have Luster (shiny)

15. Nonmetals
Gases or solids at room temperature (except Br, it is a liquid)
Poor conductors of heat or electricity
Brittle
Dull

16. Metalloids Combination of characteristics of both metals and nonmetals
Silicon and Germanium both used in computer chips

17. Answer the following questions about iron:
In what period is iron found?
In what group is iron found?
How many protons does an atom of iron have?
Write the electron configuration for iron.
Write the dot notation for iron.
Is iron a representative or transition element?
Is iron a metal or nonmetal?
List three properties of iron.
Describe one way in which classification was used in biology.

18. Color Code the Periodic Table
Elements
Color
Hydrogen
Aqua
Group 1: Alkali Metals
Red
Group 2: Alkali Earth Metals
Yellow
Groups 3 -12: Transition Elements
Blue
Group 13: Boron Group
Grey
Group 14: Carbon Group
Peach
Group 15: Nitrogen Group
Brown
Group 16: Oxygen Group
Purple
Group 17: Halogens
Orange
Group 18: Noble Gases
Green
Rare Earth Elements: Lanthanide Series
Pink
Rare Earth Elements: Actinide Series
Magenta 18

20. S- Block Alkali Metals 1 valence e-. This makes them highly reactive
Exist only as compounds
Silvery white in color
Often bond with halogens
Used in salts and batteries
Forms ions with a 1+ charge.
Alkaline Earth Metals
2 valence e-. Makes them highly reactive
Ca and Mg are important components of living cells
Silvery in color
Used to make laptop casings
Forms ions with a 2+ charge.

21. S- Block

22. P- Block (Families 13-18) Boron Family (13) Carbon Family (14)
3 valence e-
Tends to give its valence e- away
Most are metals
Not as reactive as group one and two
Forms ions with a 3+ charge
Carbon Family (14)
4 valence e-
Can either give away its valence electrons or take additional electrons
Sn and Pb will form ions with 4+charges

23. P- Block Nitrogen Family (15) Oxygen Family (16)
5 valence e-, but will form 3- (it prefers to gain 3 e- rather than give away 5)
N and P are reactive and found in many molecular compounds
Oxygen Family (16)
6 valence e-
Forms 2- ions. It prefers to gain 2 e- rather than give away 6)
O and S are reactive and found in many compounds

24. P- Block Halogens (17) Noble Gases (18) 7 valence e-
Form 1- ions (gains 1 e-)
Highly reactive nonmetals
Will often bond with metals to make salts
Noble Gases (18)
8 valence e-, full p sublevel
Does not form ions
Inert gases (unreactive)
They do not bond with other elements because they do not need anymore e-

25. P- Block

26. D- Block (3-12) All transition metals Most are hard metals
All can exist as free elements in nature
Will form a variety of charged ions due to the fact that the s and d sublevels are close in energy amounts

27. D- Block (3-12)

28. F- Block (Period 6 and 7) Lanthanide Series Actinide Series
Shiny metals
Highly reactive
Fits in period 6
Actinide Series
Radioactive
The first 4 are naturally occurring the rest are lab created
Fits in period 7

29. F- Block

30. Identify each of the following elements described below:
1. Nonmetal of the second period and group 4A.
2. The noble gas in period 3.
This element has two more protons than phosphorus.
The only nonmetal in group 1A.
Metal in period 7 with two valence electrons.
The element whose electron configuration ends with 3p1.
The nonreactive element consisting of 4 energy levels.
The metalloid with three valence electrons.
The only noble gas that does not have 8 valence electrons.
10. The element in group 2A that has fewer energy levels than magnesium.

31. Periodic Trends
Patterns in the periodic table that can be determined by comparing a period or a group
Atomic Radius, Ionic Radius, Ionization energy, Electronegativity

32. Atomic Radius Size of the atom
increases
Size of the atom
Half the distance between two identical nuclei
increases

33. Atomic Radius Increases:
Top to Bottom within a group because as you move down a column, the number of energy levels is increasing.
Right to Left because as you move left across a period the atomic number decreases which results in a larger atomic radius

34. Ionization Energy Energy needed to remove the outermost electron.
increases
Energy needed to remove the outermost electron.
Follows this trend because small atoms have a high ionization energy
increases

35. Formation of Ions
A positive ion (called a cation) results when an atom loses electrons.
Metals have low I.E. and therefore, form positive ions.
Positive ions are smaller than the parent atom.
A negative ion (called an anion) results when an atom gains electrons.
Nonmetals have high I.E. and therefore, gain electrons.
Negative ions are larger than the parent atom.

36. Octet Rule
Every Element wants 8 valence e-

38. Ionic Radius The formation of an ion
Increases
The formation of an ion
When atoms lose e- they become smaller
When atoms gain e- they become larger
increases

39. Ionization Energy
Energy required to remove an e- from a gaseous atom. (Pulling an e- off)

40. Electronegativity
The ability of an atom to attract an electron while in a chemical bond