1. Chemistry Chapter 5&6
The Periodic Law Notes 5

2. Mendeleev’s Periodic Table
Dmitri Mendeleev

3. Modern Russian Table

4. Chinese Periodic Table

5. Stowe Periodic Table

6. A Spiral Periodic Table

7. Triangular Periodic Table

8. “Mayan” Periodic Table

9. Orbital filling table

10. Periodic Table with Group Names

11. The Properties of a Group: the Alkali Metals
Easily lose valence electron
(Reducing agents)
React violently with water
Large hydration energy
React with halogens to form salts

12. Sublevel Blocks of the Periodic Table
Figure 5-5 p. 129

13. s-block Group 1 Alkali Metals Group 2 Alkaline-earth metals
ns1 is highest level
Silvery appearance
Soft—cut with knife
Highly reactive—never found free in nature
Low melting points <100oC
Group 2
Alkaline-earth metals
ns2 is highest level
Harder & denser, w/ higher melting points than Group 1
Highly reactive—never found free in nature

14. Special exceptions to s-block
Hydrogen
Has ns1
Totally different properties from alkali metals
Helium
Has ns2
Highest level is completely full
Stable like noble gases

15. d-block d sublevel for preceding energy level is filling
d sublevel filling has some deviations—Group 11: Cu, Ag, Au
Outer s & d sublevels still have same # e-
Transition elements: d-block metals w/ typical metallic properties
Less reactive than Group 1 & 2
Exist free in nature
Good conductors of electricity
High luster (shiny)

16. p-block All elements of Groups 13-18 except Helium
Properties vary greatly
Nonmetals (right hand end)
All six metalloids
Brittle solids
Some properties of metals and nonmetals
Eight metals (left hand side and bottom of the block)
Harder and denser then s-block alkaline-earth metals
Softer and less dense than d-block metals
Stable in the presence of air
Group 17 Halogens
Most reative of the nonmetal
7 electrons in outer shell

17. f-block Lanthanides & Actinides
14 elements—seven 4f orbitals are filling
Lanthanides
Similar reactivity to Group 2
Shiny metals
Actinides
Only 1st four found in nature
All are radioactive

18. Determination of Atomic Radius:
Half of the distance between nucli in
covalently bonded diatomic molecule
"covalent atomic radii"
Periodic Trends in Atomic Radius
Radius decreases across a period
Increased effective nuclear charge due
to decreased shielding
Radius increases down a group
Addition of principal quantum levels

19. Table of Atomic Radii

20. Ionization Energy - the energy required to remove an electron from an atom
Increases for successive electrons taken from
the same atom
Tends to increase across a period
Electrons in the same quantum level do
not shield as effectively as electrons in
inner levels
    Irregularities at half filled and filled
sublevels due to extra repulsion of
electrons paired in orbitals, making them
easier to remove
Tends to decrease down a group
Outer electrons are farther from the
nucleus

21. Ionization of Magnesium
Mg kJ  Mg+ + e-
Mg kJ  Mg e-
Mg kJ  Mg e-

22. Table of 1st Ionization Energies

23. Another Way to Look at Ionization Energy

24. Electron Affinity - the energy change associated with the addition of an electron
Affinity tends to increase across a period
Affinity tends to decrease as you go down
in a Group or family
Electrons farther from the nucleus
experience less nuclear attraction
Some irregularities due to repulsive
forces in the relatively small p orbitals

25. Table of Electron Affinities

26. Ionic Radii Cations Anions Positively charged ions
Smaller than the corresponding
atom
Anions
Negatively charged ions
Larger than the corresponding
atom

27. Summation of Periodic Trends

28. Table of Ion Sizes

29. Electronegativity A measure of the ability of an atom in a chemical
compound to attract electrons
Electronegativities tend to increase across
a period
Electronegativities tend to decrease down a
group or remain the same

30. Periodic Table of Electronegativities