1. The Periodic Table Periodicity Unit IV Ch. 6

2. Pictionary Words
Period Atomic radius Group Transition metal Metal Reactivity Metalloid Ionization energy Electronegativity Non-metal

3. Warm Up 1. What do you KNOW about the periodic table?
2. What WOULD you like to know about the periodic table?
3. What did you like to LEARN about the periodic table?

4. Warm Up What is the law of octaves?
Identify the groups that correspond to the following family names: alkali metals, alkali earth metals, chalcogens, halogens, noble gases and transition metals.
State the periodic law.
Who is credited with organizing the modern periodic table.

5. Warm Up-02/19/13
Identify the family in which the following appear: K, Cl, Pr, Ba, F, Mn, Th and W.
List 5 transition elements.
The alkili metals are also known as what?
The noble gas family is also known as what?

6. Warm Up-09/24/13
Place the following elements in order from largest to smallest: Si, Na, Ar, Al and S.
Place the following elements in order from largest to smallest: Be, Ba, Ca, Mg, and Sr.
Repeat questions 4 and 5 for ionic radius excluding argon.
Place the following in order from greatest to lowest ionization energies: Kr, Ne, Xe, He and Ar.

7. Expectations:
1.) Each group will have created a neat, colorful, and easy to read periodic table.
2.) Atomic Number, Symbol and Name must be present.
3.) Ionization/Electronegativity/Radii value must be written in the place of the atomic mass
4.) 3-d dimensions must be to scale

8. Atomic Radius- p. 914- 916 scale: 1 mm= 1pm.
Atomic Radius- p scale: 1 mm= 1pm.
Ionization Energy- p scale: 1 mm= 10 kJ/mol.
Electronegativity- p. 169 scale: 10 cm = 1 Paulings.
Ionic Radii- handout scale: 1 cm = .1 Angstroms.

9. Warm Up
If 1 mm= 1pm, how long would you cut a straw to represent fluorine’s atomic radius?
If 1 mm= 1pm, how long would you cut a straw to represent germanium’s atomic radius?
If 10 cm= 1 pauling, how long would you cut a straw to represent silver’s electronegativity?
If 10 cm= 1 pauling, how long would you cut a straw to represent gold’s electronegativity?
If 1 mm= 10 kJ/mol., how long would you cut a straw to represent neon’s ionization energy?
If 1 mm= 10 kJ/mol., how long would you cut a straw to represent potassium’s ionization energy?

10. Warm Up
If 1 mm= 1pm, how long would you cut a straw to represent gallium's atomic radius?
If 1 mm= 1pm, how long would you cut a straw to represent oxygen’s atomic radius?
If 10 cm= 1 pauling, how long would you cut a straw to represent manganese's electronegativity?
If 1 mm= 1 pauling, how long would you cut a straw to represent calcium’s electronegativity?
If 1 mm= 10 kJ/mol., how long would you cut a straw to represent kripton’s ionization energy?

11. Johann Dobereiner, 1817 Triads
Groups of three elements having similar physical and chemical properties.
These three elements are in the same Group or Family

12. John Newlands, 1863 Law of Octaves
Arranged elements in order of their atomic masses.
Noticed that their properties repeated every 8th element
Law of Octaves
The same properties repeat every eighth element

13. Dmitri Mendeleev, 1869 The Periodic Law
Believed that similar properties occurred after periods that could vary in length
Properties of the elements repeat in an orderly way. Such a pattern is “periodic”
The Periodic Law
“The properties of the elements are a periodic function of their atomic masses”

14. Henry Moseley, 1913 Periodic Law
Using X-Rays, he showed that the nucleus had a positive charge. Thus, the Periodic Law was revised:
Periodic Law
Properties of the elements are a periodic function of their atomic numbers

15. Electron Configuration and The Periodic Table
Electron Configuration determines a chemical's reactivity.
s-Block Elements (Groups 1 and 2; or Groups I A and II A)
Outermost electrons are added to an s-orbital
Group 1: s1 - Alkali Metals
Group 2: s2 - Alkaline Earth Metals

16. Electron Configuration and The Periodic Table
p-Block Elements (Groups 13-18; Groups III A through VIII A)
Outermost electrons are added to a p-orbital
Group 13: p1
Group 14: p2
Group 15: p3
Group 16: p4 - Chalcogens
Group 17: p5 - Halogens
Group 18: p6 - Noble Gases (Inert Gases)

17. Electron Configuration and The Periodic Table
d-Block Elements (Groups 3-12; Groups I B through VIII B)
Outermost electrons are added to a d-orbital
Known as the transition metals
All metals

18. Electron Configuration and The Periodic Table
Reading the electron configuration directly off the Periodic Table.
s-Block
p-Block
Main Group Elements (s- & p-blocks)
d-Block (Transition Elements)
f-Block (Lanthanides and Actinides)

19. Warm Up
What are the three largest divisions within the periodic table?
What are the characteristics of metals?
What are the characteristics of non-metals?
What are the characteristics of metalloids?
What is the rule of thumb for metals?
What is the rule of thumb for non-metals?
State the octet rule?

20. Categories of Elements
Metals - hard and shiny; conduct electricity
Nonmetals - gases or brittle solids; good insulators
Metalloids - properties of both metals and nonmetals
Rule of Thumb:
Metals have 1-3 electrons in outer level
Nonmetals have >5 electrons in the outer level
Metals: on left side of the periodic table
Nonmetals: on the right side of table

21. The Octet Rule
Eight electrons in the outer level render an atom essentially unreactive
Rule of Thumb:
An atom having a filled or half filled sublevel is slightly more stable (less reactive) than an atom without a filled or half-filled sublevel.

22. Relative Atomic Stability
Decreasing order of stability
Full outer shell (s2 and p6)
Full sublevel (s2)
Half-filled sublevel
No special arrangement

23. Electron Sublevel Structure
If the last electron for an atom is in a full or half-full sublevel, then the atom is inherently more stable.

24. Periodic Properties Atomic Radii Ionic Radii First Ionization Energy
Electronegativity
Electron Affinity
Oxidation Numbers

25. Warm Up
What happens to atomic radius as you go down the periodic table?
What is ionization energy?
As atomic number increases in a period, what happens to ionization energy?
As atomic number increases in a group, what happens to ionization energy?
List three factors that affect ionization energy?
Define electronegativity?
What happens to electronegativity as you move left to right across the periodic table?
What happens to electronegativity as you move top to bottom across the periodic table?

26. Atomic Radius
As the principal quantum number (n) increases, the size of the electron cloud increases. That is, the atomic size increases as you go down the table.
The reason for this is that you are adding energy levels as you go down the table (1, 2, 3,...).
The positive charge of the nucleus increase as you go from left to right across the table. This increase in nuclear charge increases the pull on the electron cloud by the nucleus - pulling the the electron cloud in tighter to the nucleus. Thus, the atoms decrease in size.
Summary: The atomic radii increase from top to bottom, and from right to left.

27. Ionic Radii Metallic Ions Formed by the loss of electrons
Smaller than the atoms from which they were formed
Nonmetallic Ions
Formed by the gain of electrons
Larger than the atoms from which they were formed

28. First Ionization Energy
The energy required to remove the most loosely-held electron from a neutral atom.
Increases as atomic number increases in any period.
Decreases as the atomic number increases in any group.
Affected by:
Shielding Effect
Radius of the atom
Nuclear Charge
Electron Sublevel Structure

29. Factors Affecting Ionization Energy
Nuclear Charge - ionization energy is proportional to the nuclear charge
Shielding Effect - Ionization energy is inversely proportional to the shielding effect
Radius - Ionization energy is inversely proportional to the distance between the nucleus and the outer electrons
Sublevel - an electron from a full or half-full sublevel requires additional energy to be removed
Metals - Low ionization energy
Nonmetals - High ionization energy

30. Electronegativity
Electronegativity indicates the ability of an element’s atom to attract electrons in a chemical bond.

31. Electronegativity
Influenced by the same factors which affect ionization energy and electron affinity
Size
Shielding effect
Nuclear charge
The Trends (in the Periodic Table) are the same
increases from left to right
increases from bottom to top

32. Electronegativity
The most active metals have the lowest electronegativity
The most active nonmetals have the highest electronegativity

33. Nuclear Charge
As the positive charge of the nucleus increases, it becomes increasingly harder to remove an electron from an outer shell
Tends to raise the First Ionization Energy

34. Shielding Effect
The effect when “inner” electrons block the attraction of the nucleus for the outer electrons
Tends to lower the First Ionization Energy

35. Radius
Ionization energy is inversely proportional to the distance between the nucleus and the outer electrons
Tends to lower the First Ionization Energy

36. Sublevel
An electron from a full or half-full sublevel requires additional energy to be removed
If a sublevel is full or half-filled, then this tends to raise the First Ionization Energy

37. Electron Affinity
The attraction an isolated atom has for an additional electron
Shows the same trend as First Ionization Energy (increases from left to right, and decreases from the top down).

38. Oxidation Numbers Group 1 lose 1 electron +1
Group 2 lose 2 electrons +2
Group multiple gain/loss
Group 13 lose 3 electrons +3
Group 14 lose/gain 2,4 electrons ±4, +2
Group 15 gain 3 electrons -3
Group 16 gain 2 electrons -2
Group 17 gain 1 electron -1
Group 18 stable

39. LEO says GER

40. Warm Up What is the oxidation number for group 1?
Name the following groups: 1, 2, 16,17 and 18.

41. Web Activity