1. Periodic Table and Trends
Chapter 6:
Periodic Table and Trends

2. Dmitri Mendeleev First Periodic Table
Based on increasing Atomic mass and repeating properties of elements
Had spaces for “missing” elements that he predicted

3. Henry G. J. Moseley
Discovered that elements’ properties were more closely associated with atomic number
Modern periodic table is based on this discovery

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

5. Reading the Periodic Table
Nitrogen Family
Oxygen Family
Groups or Families
Inner Transition Metals
Periods
Carbon Family
Boron Family
Alkali Metals
Alkaline Earth Metals
Halogens
Noble Gases
Transition Metals

6. Metals Left of the stair-step line Majority of the elements
Tend to LOSE electrons
Most reactive in the s-block

7. Properties of Metals Shiny luster Good conductors of heat
Good conductors of electricity
Most are solids at room temperature
Malleable
Ductile

8. gold
lead
copper
nickel

9. Nonmetals Right of Stair-Step line Tend to GAIN electrons
Most reactive group is halogens
Least reactive group is Noble Gases

10. Properties of Nonmetals
Dull Luster
Poor conductors of heat
Poor conductors of electricity
Brittle
Many are gases at room temperature

11. CARBON
BROMINE
SULFUR
ARGON

12. Metalloids Along Stair-step line
Have properties of metals AND nonmetals
Many are used in transistors, found in electronics

13. Silicon
Antimony
Boron

14. Alkali Metals Group 1 (Except H) All have only 1 valence electron
Most reactive metals; never found in elemental form in nature
Soft and shiny
Relatively low melting points

15. Alkaline Earth Metals Group 2 All have 2 valence electrons
Second most reactive metals; never found in pure state in nature
Harder, denser, and stronger than alkali metals
Have higher melting points than alkali metals

16. Transition Metals Groups 3-12
All have 1 or 2 valence electrons (in s sublevels)
Do not fit into any other group or family
Have many irregularities in their electron configurations

17. Boron Family Group 3A Have 3 valence electrons Boron is a metalloid
All others are metals

18. Carbon Family Group 4A All have 4 valence electrons
Carbon is a nonmetal
Si and Ge are metalloids
Sn and Pb are metals

19. Nitrogen Family Group 5A
All have 5 valence electrons (s and p sublevels)
N and P are nonmetals
As and Sb are metalloids
Bi is a metal

20. Oxygen Family Group 6A All have 6 valence electrons
Oxygen, Sulfur, and Selenium are nonmetals
Tellurium and Polonium are metalloids

21. Halogens Means “salt former” Group 7A All have 7 valence electrons
Most reactive nonmetals
All are nonmetals

22. Noble Gases
Group 8A
8 Valence electrons makes a full electron shell: s2 p6
Complete, stable electron configuration (Complete outer energy level)
Least reactive of all elements

23. Rare Earth Elements (Inner Transition metals)
Found in 2 rows at bottom of periodic table
Lanthanide series follows La
Actinide series follows Ac
Little variation in properties
Actinides are radioactive; only first three and Pu are found in nature

24. Summary Groups: Up and Down Periods: Across
Main Group Elements are in groups 1-2, 13-18
Elements along the stair step line are metalloids
Elements to the left of the stair step line are metals
Elements to the right of the stair step line are nonmetals

25. Octet Rule “Noble Gas Envy”
Atoms tend to gain, lose, or share electrons in order to acquire a full set of valence electrons (typically 8)

26. Periodicity
Properties of the elements change in a predictable way as you move through the periodic table
These properties include
Atomic Radius
Ionization energy
Electronegativity

27. Atomic Radius
Distance from nucleus to outermost valence electrons

28. Atomic Radius
Increases down groups
Decreases from left to right

29. Ionization Energy The energy needed to remove 1 of an atom’s electrons
Decreases as you move down a group
Increases from left to right, across a period
Successive ionization energies increase for every electron removed

30. 1st ionization energy

31. Electronegativity
Reflects an atom’s ability to attract electrons in a chemical bond
Related to its ionization energy and electron affinity
Increases from left to right, across a period
Decreases from top to bottom, down a group

33. Shielding
Shielding electrons are electrons located between the nucleus and the valence electrons
For example:
Chlorine has the following electron configuration:
1s2 2s2 2p6 3s2 3p5
The shielding electrons would be 1s2 2s2 2p6
The valence electrons would be 3s2 3p5
Then we have 10 shielding electrons and 7 valence electrons, right?

34. Shielding You try one Did you get:
Try Sodium (Na)… Remember, first you have to know the electron configuration
Did you get:
Electron configuration: 1s2 2s2 2p6 3s1
Shielding electrons: 1s2 2s2 2p6
Valence electrons: 3s1
So we have 10 shielding electrons and 1 valence electron, right?

35. Zeff
Effective nuclear charge (Zeff) is the charge felt by the valence electrons after you’ve taken into account the number of shielding electrons that surround the nucleus.
Huh?
Let’s put it in an equation

36. Zeff =# of protons - # of shielding electrons
So, calculate the effective nuclear charge for the all the elements in period 3
Now, calculate the effective nuclear charge for all the elements in group 2
What pattern do you see arising?

37. What is the correlation between Zeff and atomic radius
What is the correlation between Zeff and atomic radius? (Remember opposite charges attract)
The greater the Zeff the smaller the atomic radius
I’m still lost…..
Greater effective nuclear charge means that the valence electrons are feeling a greater pull toward the nucleus, making the atom smaller in size

38. In summary…
Effective nuclear charge can be used to predict trends in atomic radius
Increases from left to right and decreases from top to bottom
Zeff = Z - σ
Effective nuclear charge is dependent upon electron shielding
Electronegativity increases from left to right and decreases from top to bottom