1. Chemistry Chapter 6
The Periodic Law

2. Mendeleev’s Periodic Table
Dmitri Mendeleev

3. Mendeleev – organized periodic table
Vertical columns in atomic mass order
Made some exceptions to place elements in rows with similar properties (Te and I)
Horizontal rows have similar chemical properties
Gaps for “yet to be discovered” elements
Left questions: why didn’t some elements fit in order of increasing mass? Why did some elements exhibit periodic behavior?

4. Moseley
Discovered that periodic table was in atomic number order, not atomic mass order
Explained the Te-I anomaly

5. Periodic Law
Physical and chemical properties of the elements are periodic functions of their atomic numbers

6. Modern Periodic Table
Discovery of noble gases yields new family (Group 18 – aka inert gases)
Lanthanides (#58 - #71)
Actinides (#90 – #103)

7. Periods and Blocks of the Periodic Table
Periods – horizontal rows
Period number corresponds to the highest principal quantum number
Groups/Families – vertical columns; these elements share similar chemical properties (they have the same number of valence electrons)
Blocks – periodic table can be broken into blocks corresponding to s, p, d, f sublevels

8. Orbital filling table

9. Blocks and Groups – s block
Group1 – “The alkali metals”
One s electron in outer shell
Soft, silvery metals of low density and low melting points
Highly reactive, never found pure in nature

10. Blocks and Groups – s block
Group 2 – “Alkaline Earth Metals”
2 s electrons in outer shell
Denser, harder, stronger, less reactive than Group 1
Too reactive to be found pure in nature

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
React with halogens to form salts

13. Blocks and Groups – d block
Metals with typical metallic properties
Referred to as transition metals
Group number = sum of outermost s and d electrons

14. 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

15. 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

16. Blocks and Groups – p block
Properties vary greatly – metals, metalloids, and nonmetals
Group 17 – halogens are most reactive of non metals
Group 18 – noble gases are NOT reactive

17. 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

18. 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

19. 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

20. 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

21. Blocks and Groups – f block
Lanthanides – shiny metals similar to group 2
Actindes – all are radioactive; plutonium – lawrencium are man-made

22. Ions: atom or group of atoms that have a charge due to loss or gain of electrons monatomic ion – only one element polyatomic ion – more than one element
Cation: A positive ion
Mg2+, NH4+
Anion: A negative ion
Cl-, SO42-

23. Predicting Ionic Charges
Group 1:
Lose 1 electron to form 1+ ions H+
Li+
Na+ K+

24. Predicting Ionic Charges
Group 2:
Loses 2 electrons to form 2+ ions
Be2+
Mg2+
Ca2+
Sr2+
Ba2+

25. Predicting Ionic Charges
Loses 3
electrons to form
3+ ions
Group 13:
B3+
Al3+
Ga3+

26. Predicting Ionic Charges
Lose 4
electrons or gain
4 electrons?
Group 14:
Neither! Group 14 elements rarely form ions, if they do, they can do both!

27. Predicting Ionic Charges
Nitride
Gains 3
electrons to form
3- ions
Group 15:
P3-
Phosphide
As3-
Arsenide

28. Predicting Ionic Charges
Oxide
Gains 2
electrons to form
2- ions
Group 16:
S2-
Sulfide
Se2-
Selenide

29. Predicting Ionic ChargesF-
Fluoride
Br-
Bromide
Group 17:
Gains 1
electron to form
1- ions
Cl-
Chloride I-
Iodide

30. Predicting Ionic Charges
Stable Noble gases do not form ions!
Group 18:

31. Predicting Ionic Charges
Many transition elements
have more than one possible oxidation state.
Groups :
Iron(II) = Fe2+
Iron(III) = Fe3+

32. Predicting Ion Charges
Some transition elements
have only one possible oxidation state.
Groups :
Zinc = Zn2+
Silver = Ag+ Cadmium = Cd2+

33. Write formulas for: Rubidium ion Gallium ion Sulfide Bromide
Cadmium ion

34. 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 increase in protons (energy level stays same)
Radius increases down a group
Addition of energy levels results in valence
electrons farther from nucleus - decrease in
nuclear shielding

35. Table of Atomic Radii

36. 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 are harder to remove because they
are closer to the nucleus, so IE is high
Tends to decrease down a group
Outer electrons are farther from the
nucleus and easier to remove, IE is low

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

38. Table of 1st Ionization Energies

39. Another Way to Look at Ionization Energy

40. Ionic Radii Cations Anions Positively charged ions
Smaller than the corresponding
atom due to loss of energy level
Anions
Negatively charged ions
Larger than the corresponding
atom; more e- means less attraction for nucleus

41. Table of Ion Sizes

42. Electronegativity A measure of the ability of an atom in a chemical
compound to attract electrons
Electronegativities tend to increase across
a period
* more nuclear charge, more power to attract electrons
Electronegativities tend to decrease down a group or remain the same
* additional energy levels result in less attraction to the nucleus
* Noble gases do not have EN values b/c they do not form bonds!

43. Periodic Table of Electronegativities