1. Chemistry Chapter 5
The Periodic Law

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. Giguere Periodic Table

10. Orbital filling table

11. Periodic Table with Group Names

12. 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

13. Properties of Metals
Metals are good conductors of heat and electricity
Metals are malleable
Metals are ductile
Metals have high tensile strength
Metals have luster

14. Examples of Metals
Potassium, K reacts with water and must be stored in kerosene
Copper, Cu, is a relatively soft metal, and a very good electrical conductor.
Zinc, Zn, is more stable than potassium
Mercury, Hg, is the only metal that exists as a liquid at room temperature

15. Properties of Nonmetals
Carbon, the graphite in “pencil lead” is a great example of a nonmetallic element.
Nonmetals are poor conductors of heat and
electricity
Nonmetals tend to be brittle
Many nonmetals are gases at room temperature

16. Examples of Nonmetals
Microspheres of phosphorus, P, a reactive nonmetal
Sulfur, S, was once known as “brimstone”
Graphite is not the only pure form of carbon, C. Diamond is also carbon; the color comes from impurities caught within the crystal structure

17. Properties of Metalloids
Metalloids straddle the border between metals and nonmetals on the periodic table.
They have properties of both metals and nonmetals.
Metalloids are more brittle than metals, less brittle than most nonmetallic solids
Metalloids are semiconductors of electricity
Some metalloids possess metallic luster

18. Silicon, Si – A Metalloid
Silicon has metallic luster
Silicon is brittle like a nonmetal
Silicon is a semiconductor of electricity
Other metalloids include:
Boron, B
Germanium, Ge
Arsenic, As
Antimony, Sb
Tellurium, Te

19. 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

20. Table of Atomic Radii

21. 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

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

23. Table of 1st Ionization Energies

24. Another Way to Look at Ionization Energy

25. 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 period
Electrons farther from the nucleus
experience less nuclear attraction
Some irregularities due to repulsive
forces in the relatively small p orbitals

26. Table of Electron Affinities

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

28. Summation of Periodic Trends

29. Table of Ion Sizes

30. 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

31. Periodic Table of Electronegativities