1. The Periodic Table
Chapter 6

2. Objectives
Draw the periodic table and label the electron blocks and areas of non-metals, metals, and metalloids.
Relate the Lewis dot structure to its place in the periodic table.
Explain periodic trends as one moves along periods and down groups in the periodic table

3. Chapter 6 Vocabulary Chapter 6.1-6.2 Periodic Law Group Period
Representative Element
Transition Element
Metal
Alkali Metal
Alkaline Earth Metal
Transition Metal
Inner Transition Metal
Lanthanide Series
Actinide Series
Nonmetal
Halogen
Noble Gas
Metalloid

4. Early Periodic Law
Dmitri Mendeleev noticed in his table that there were repetitions of physical and chemical properties when the elements were arranged by atomic mass.

5. Mendeleev’s Periodic Table
The numbers above represent atomic masses.

6. Mendeleev’s Periodic table
Properties of Germanium (Ge)
Actual (1886)
Property
Predicted (1869)
Atomic Mass
72 u
Color
Dark gray
Density
5.5 g/mL
Melting Point
High
Density of Oxide
4.7 g/mL
Oxide solubility in HCl
Slightly dissolved by HCl
Formula of chloride
EsCl4

7. Modern Periodic Law
Periodic Law states that chemical and physical properties repeat in regular cyclic patterns when they are arranged by increasing atomic number.
Starts with metals at left and goes to non-metal (noble gas) on right
Properties change in orderly progression across a period.
Do Li, Na, K metals in water demo here if time.

8. Henry Mosley Rearranged by Atomic Number

9. Periodic Table Columns, Groups or Families Noble Gases Halogens
Alkali Metals
Alkaline Earth Metals
Representative Elements
Transition Elements
Periods
Inner Transition Elements
Metals
Metalloids
Nonmetals

10. Elemental Properties
What are some of the elemental properties that make the periodic table, well, periodic?
Classification by metals, nonmetals and metalloids
Metals - shiny ductile, malleable solids, good conductors of heat and electricity
Nonmetals - dull, brittle solids; or gas, poor conductors of heat and electricity
Metalloids - have chemical and physical properties of both metals and nonmetals

11. 6.2 - Fundamental Organizing Principle for Periodic Table
Representative Elements (Sometimes called A Group)
Group # = number of valence electrons
Means similar Lewis dot structure and similar properties.
s-block elements have 1-2 electrons in s-orbital
p-block elements have 1-6 electrons in p-orbitals
Noble gases have filled valence shells
Energy level of valence electrons is at energy level given by period (row) number

12. Fundamental Organizing Principle for Periodic Table
Transition Elements (Sometimes called B Group)
d-block elements have 1-10 electrons in d- orbitals
Columns 3-12 in periodic table
Energy level of valence electrons at n and partially filled n-1 d orbitals (example: 4s and 3d)
f-block (Lanthanides and Actinides) have electrons in f-orbitals

13. Practice Fill in the missing info for the following elements:
Identify the element fitting the description.
a) Group 2 (2A) element in 4th period:
b) Noble gas in 5th period:
c) Group 12 (2B) element in 4th period:
d) Group 16 (6A) element in 2nd period:
Configuration
Group
Period
Block
[Ne]3s2
[He]2s1
[Kr]5s24d105p5
7 (7B) 4

14. Ch PERIODIC TRENDS

15. Chapter 6.3 Vocabulary Effective Nuclear Charge (Z*) – Not in book!
Shielding (Not in book)
Ion
Ionization Energy
Octet Rule
Metallic Character (Not in book)
Electronegativity

16. General Periodic Trends
Atomic and ionic size
Ionization energy
Electronegativity
Metallic Character
Higher effective nuclear charge
Electrons held more tightly
Larger orbitals.
Electrons held less
tightly.
Can show a lot of these trends from the ACS periodic table.

17. Effective Nuclear Charge, Z*
Z* is the nuclear charge experienced by the outermost electrons. (Note: not in book!)
Z* increases across a period owing to shielding by inner electrons.
Shielding is blocking by inner electrons.
For a period (row), the number of shielding electrons remain the same, but the number of protons in the nucleus increases.
Example: All elements in the second period have the same underlying [He] noble gas configuration. However, the number of protons increase from left to right.

18. Effective Nuclear Charge, Z*
So we can estimate as
Z* = [ Z - (no. inner electrons) ] or
Z* = Z – S (inner electrons)
Z is total number of electrons
S is the number of electrons blocking the valence shell electrons, the underlying noble gas electrons.
Charge felt by 2s e- in Li Z* = = 1
Be Z* = = 2
B Z* = = 3 and so on!

19. Orbital Energies
Orbital energies “drop” as Z* increases

20. Atomic Size - Radius
Atomic size is a periodic trend influenced by electron configuration.
For metals, atomic radius is half the distance between adjacent nuclei in a crystal of the element.

21. Atomic Size - Radius
For other elements, the atomic radius is half the distance between nuclei of identical atoms that are bonded together.

22. Atomic Size - radii

23. Atomic Size
Size (radius) goes UP on going down a group. See previous slide.
Because electrons are further from the nucleus, there is less attraction.
Size (radius) goes DOWN on going across a period.

24. Atomic Size Increase in Z*
Size (radius) decreases across a period owing to increase in Z*. Each added electron feels a greater and greater positive charge. Note: Electrons in the same energy level don’t shield each other too much.
Large
Small
Increase in Z*

25. Trends in Atomic Size

26. Ionic Size The radius of an atom when it has become an ion.
An ion is an atom or bonded group of atoms that has an overall positive or negative charge.
An atom acquires a positive charge by losing electrons or negative charge by gaining electrons!!

27. Ion Configurations
To form positive ions from elements remove 1 or more e- from subshell of highest n [or highest (n + l)].
Al: [Ne] 3s2 3p e-  Al3+: [Ne] 3s0 3p0 1s 2s 3s 3p 2p

28. Atoms tend to gain, lose, or share electrons to get
Octet Rule
Atoms tend to gain, lose, or share electrons to get
8 valence electrons
(except small atoms up to Boron)

29. Practice
Write the electron configuration and orbital box diagram for Mg when it is an ion. Hints: What is its noble gas configuration? What will they do to get an octet?
Write the electron configuration and orbital box diagram for O when it is an ion.

30. Ion Sizes Forming a positive ion.Li +
, 78 pm
2e and 3 p
Forming a positive ion.
Li,152 pm
3e and 3p
Positive ions are SMALLER than the atoms from which they come.
The electron/proton attraction has gone UP and so size DECREASES.
Electron Configuration as ion is: [He] 2s0

31. Ion Sizes Forming a negative ion. F - , 133 pm 10 e and 9 p F, 71 pm
Negative ions are LARGER than the atoms from which they come.
The electron/proton attraction has gone DOWN and so size INCREASES.
Trends in ion sizes are the same as atom sizes.
Electron configuration as ion: 1s22s22p6 (just like neon.)

32. Trends in Ion Sizes
See Figure 6-14

33. Reaction Properties
Why do metals lose electrons in their reactions? Why does Mg form Mg2+ ions and not Mg3+? Why do nonmetals take on electrons?

34. Ionization Energy
IE = energy required to remove an electron from an atom in the gas phase.
Mg (g) kJ  Mg+ (g) + e-

35. Ionization Energy Mg+ (g) + 1451 kJ  Mg2+ (g) + e-
IE = energy required to remove an electron from an atom in the gas phase.
Mg (g) kJ  Mg+ (g) + e-
Mg+ (g) kJ  Mg2+ (g) + e-
Mg+ has 12 protons and only 11 electrons. Therefore, IE for Mg+ > Mg.

36. 3rd: Mg2+ (g) + 7733 kJ  Mg3+ (g) + e-
Ionization Energy
1st: Mg (g) kJ  Mg+ (g) + e- 2nd: Mg+ (g) kJ  Mg2+ (g) + e-
3rd: Mg2+ (g) kJ  Mg3+ (g) + e-
Energy cost is very high to dip into a shell of lower n.

37. Trends in Ionization Energy

38. Trends in Ionization Energy

39. As Z* increases, orbital energies “drop” and IE increases.

40. Subsequent Ionization Energies

41. Trends in Ionization Energy
IE increases across a period because Z* increases.
Metals lose electrons more easily than nonmetals.
Nonmetals lose electrons with difficulty.
High ionization energy: atoms want to hold on to electrons; likely to form negative ion
Low ionization energy: atom gives up electron easily; likely to form positive ion

42. Trends in Ionization Energy
IE decreases down a group
Because size increases.
Ability to lose electrons generally increases down the periodic table.
See reactions of Li, Na, K

43. Practice
Which element in each pair has the larger 1st ionization energy?
Na or Al
Ar or Xe
Ba or Mg

44. Periodic Trend in the Reactivity of Alkali Metals with Water
Lithium
Sodium
Potassium

45. Metallic character *Note: ‘metallic character’ not in book.
An element with metallic character is one that loses electrons easily.
Metallic character:
is more prevalent in metals on left side of periodic table
is less for nonmetals on right side of periodic table that do not lose electrons easily

46. Metallic character

47. Electronegativity
Relative ability of an element to attract electrons in a chemical bond.
Ionization energy reflects ability of atom to attract electrons in an isolated atom
Generally, the higher the ionization energy of an atom, the more electronegative the atom will be in a molecule
There are many electro negativity scales – we’ll use the one by Linus Pauling (values dimensionless)
Will be used to determine things like polarity of a chemical bond.
Give class an idea of derivation—see AP Chem notes. I believe that it is based on a value of hydrogen of 2.1

49. Electronegativity Trends
Decreases down a group
Why? Due to greater atomic radius
Increases across a period
Why? Increased positive charge in nucleus (Greater Z*)
Same trend as for ionization energy. Surprised?

50. Periodic Table Trend Summary
Moving Left  Right (periods)
Z* Increases
Atomic & ionic Radius Decrease
Ionization Energy Increases
Electronegativity Increases
Metallic Character Decreases
Moving Top  Bottom (groups)
Z* is roughly constant, but val e- distance increases
Atomic & Ionic Radius Increase
Ionization Energy Decreases
Electronegativity Decreases
Metallic Character Increases

51. Periodic Trends practice For each of the following properties, indicate whether fluorine or bromine has a larger value.
Electronegativity
Ionic Radius
Atomic Radius
Ionization Energy
Metallic character