1. Periodic Law:
A periodic repetition of chemical and physical properties of elements when the elements are arranged by increasing atomic number.
For example:
Na behaves much like Li
Na is the eighth element after Li
Mg behaves much like Be
Mg is the eighth element after Be
P behaves much like N
P is the eighth element after N
S behaves much like O
S is the eighth element after O
Cl behaves much like F
Cl is the eighth element after F
Ar behaves much like Ne
Ar is the eighth element after Ne
Russian Chemist Dmitri Mendeleev is usually given credit for understanding the implications of this repetition in properties and he arranged elements based on similarities of properties.
Henry Moseley first arranged elements in order of increasing Atomic Number.

2. Mendeleev's Early Periodic Table, Published in 1872
Mendeleev is from Russia, but the chart is in German. Why?

3. Periods: Horizontal Rows across the periodic table (Side to Side)
Groups or Families: Vertical Columns in the periodic table (Up and Down)
Periods: Horizontal Rows across the periodic table (Side to Side)
Representative Elements: Groups 1-2 and (the S and P blocks)
These groups generally display a wide range of properties which are said to be representative of most elements.
Transition Elements: Groups 3-12 (the D block)
These elements gradually change (or transition) their properties and are more similar to each other than to many elements in the representative group.

4. Metals: elements that typically have the following properties
good conductors of heat and electricity
malleable-can be beaten into a flat sheet
ductile-can be drawn into a wire
usually shiny solids when clean and smooth
Metals are on the left side of the periodic table.
Non-Metals: elements that are not like metals
poor conductors of heat and electricity
brittle-break into pieces when hit or are gases
dull looking solids if not a gas
Non-metals are on the right side of the periodic table.
Metalloids: elements that have properties of both metals and non-metals (these are also called semiconductors)
B, Si, Ge, As, Sb, Te

5. Non-metals Metals Metaloids1 2 3 4 5 6 7
Non-metals
Metals
Metaloids
Notice that Hydrogen is considered to be a non-metal even though it is placed with the group 1 metals. Hydrogen does form positive ions like the metals do.

6. Specially Named Groups or Families
Alkali Metals:
Group 1 metals, all have s1 electron configuration
Very reactive metals-form ions with positive one charge (+1)
Alkaline Earth Metals:
Group 2 metals, all have s2 electron configuration
Reactive metals-form ions with positive two charge (+2)
Halogens:
Group 7 nonmetals, all have s2, p5 electron configuration
Very reactive nonmetals-form ions with negative one charge (-1) when reacting with metals
Noble Gases:
Group nonmetals, all have s2, p6 electron configuration (except He)
Nonreactive elements-almost no know compounds containing these elements. Remember that a full energy level is not reactive.

7. Lanthanides Actinides Alkali Metals Halogens Transition Metals
Color coding the groups on a periodic table may help you to remember where the named groups and series are. 1 2 3 4 5 6 7
Lanthanides
Series
Actinides
Alkali Metals
Halogens
Transition Metals
Groups
Alkaline Earth Metals
Nobel Gases

8. Periodic Table Labeling Activity (Using your book, notes, or collaboration with others)
1 Obtain a “fresh” periodic table
2 Obtain a set of colored pencils
Shade in appropriate areas of the periodic table (using different colors) that correspond to the
following sections of the periodic table:
Alkali Metals; Alkaline Earth Metals; Halogen; Noble Gases
Transition Metals; Metalloids
Create a color-coded key that clearly indicates which color corresponds to which section of the
periodic table.
Bonus Work: Review/Tie-in with Chapter 4 material (Begin in class, finish at home)
Using another fresh periodic table, create a similar color-coded key that shows the location of atoms
that have each of the following designations:
“s” block; “p” block; “d” block; “f” block
Within each block label each column as 1, 2, 3 ….. (for example which column would be s1, p3, d6,
or f2).
Choose an element within each block and write out an electron configuration for that element (use
the noble gas shortcut method).
Choose an element from each s or p column and write a dot diagram for that element.
For each element for which you drew a dot diagram, identify two other elements that would have the
same number of dots in their dot diagrams (do not draw the dot diagrams, just list the element
symbols)

9. 1 2 3 4 5 6 7
Gases
Liquids
Solids
Diatomic Elements

10. F Na F -1 Na+1 9 protons, 9 electrons Zero charge
Ion Formation +9 F
+11 Na
9 protons, 9 electrons
Zero charge
11 protons, 11 electrons
Zero charge +9
+11
F -1
Na+1
11 protons, 10 electrons
+1 charge
9 protons, 10 electrons
-1 charge

11. Periodic Trends:
Many properties of the elements change in a regular way as you move either up and down or left and right on the periodic table.
Atomic Radius: half the distance between the nuclei in a molecule consisting of identical atoms
Ionic Radius: radius of an ion (an atom that has become charged)
Ionization Energy: the energy required to remove an electron from a gaseous atom
Electronegativity: a measure of the ability of an atom to attract electrons that are in a chemical bond

12. Atomic Radius: half the distance between the nuclei in a molecule consisting of identical atoms
Since atoms do not actually have an edge, we define the radius based on one half the distance between two nuclei in a compound.
As you move from left to right across a period, what happens to the radius of the elements?
As you move from top to bottom in a group, what happens to the radius of the elements?

13. Period Trend
Increasing Radii
Group
Trend
Increasing Radii
Increasing Radii
Overall Trend
Atomic Radii in picometers (pm)

14. Increasing Atomic Radius Arrows1 2 3 4 5 6 7
Increasing Atomic Radius Arrows
Summary Arrow for Increasing Atomic Radius
You need to remember that atoms get larger when you move from the right to the left or from top to bottom on the periodic table!
This trend is related to several other trends that follow.

15. Trends in Ionization Energies (kJ/mol) for the Representative Elements The trend can be hard to see in this form.

16. The same data plotted in a different way. Wow! What a complicated graph!
We need to make sense of this information!

17. If we look at data from a single period (across), we get this graph.
In general, ionization energy increases as you move from left to right across each period.

18. If we look at data from a single group (vertical), we get this graph.
Ionization energy decreases as you move from top to bottom in a group or family.
Or, Ionization energy increases as you move from bottom to top in a group or family.

19. Increasing Ionization Energy Arrows1 2 3 4 5 6 7
Increasing Ionization Energy Arrows
Summary Arrow for Increasing Ionization Energy
This is exactly opposite the trend for atomic size!!! Said another way, small atoms hold onto their outer shell electrons more tightly than larger atoms do.

20. Electronegativity

21. Increasing Electronegativity Arrows1 2 3 4 5 6 7 F
Increasing Electronegativity Arrows
Summary Arrow for Increasing Electronegativity
Note: Nobel gases do not form compounds, so they have no electronegativity values.
This trend is the same as ionization energy which makes sense because atoms with high electronegativity tend to hold onto their electrons firmly.

22. Ionic Radii in picometers
Increasing Radii
Increasing Radii
Positive Ions
Negative Ions
Increasing Radii
Ionic Radii in picometers
Notice that the negative ions are larger than all but the very largest positive ion.

23. Shielding: Inner shell electrons “blocking” the pull of the nucleus for the outer shell electrons.
The stronger the shielding, the larger the atom becomes. The outer shell electrons do not feel the pull of the nucleus as much if they are strongly shielded.
The weaker the shielding, the smaller the atom becomes. The outer shell electrons feel the pull of the nucleus more strongly if they are weakly shielded.
Increasing the positive charge in the nucleus increases the pull of the nucleus for the all electrons.
If the number of outer shell electrons increases while the number of inner shell electrons remains constant, the shielding is decreasing which causes the atoms to be smaller. This is what happens as we move from left to right across a period.

24. Li Be B C N O F Ne Strongest
Shielding
Look at a series of atoms to see how shielding works?
+4 in Nucleus Pulling on 2 inner shell and 2 outer shell electrons-
+5 in Nucleus Pulling on 2 inner shell and 3 outer shell electrons-
+7 in Nucleus Pulling on 2 inner shell and 5 outer shell electrons-
+8 in Nucleus Pulling on 2 inner shell and 6 outer shell electrons-
+6 in Nucleus Pulling on 2 inner shell and 4 outer shell electrons-
+3 in Nucleus Pulling on 2 inner shell and one outer shell electron-
+9 in Nucleus Pulling on 2 inner shell and 7 outer shell electrons-
like 2 blockers and 5 receivers
like 2 blockers and 4 receivers
like 2 blockers and 3 receivers
like 2 blockers and 2 receivers
like 2 blockers and 6 receivers
like 2 blockers and 7 receivers
like 2 blockers and one receiver +3 +4
+10 in Nucleus Pulling on 2 inner shell and 8 outer shell electrons-
like 2 blockers and 8 receivers
More receivers with the same number of blockers is weaker shielding which means that the electrons will feel more pull from the nucleus-smaller atom. Li Be
Weakest
Shielding +5 +6 +7 +8 +9
+10 B C N O F Ne

25. When a shell becomes full, the next electron must go into a higher energy level. Higher energy means farther from the nucleus, so the new shells are larger than the previous shells.
Going down a group, each element has the same type of electron in its outer shell, but the electron is in the next higher energy level.
Therefore, going down a group makes the atoms larger even though there are more protons in the nuclei as you go down a group.
Larger numbers of electrons in inner shells also increase shielding as you move down a group-increased shielding means larger atoms.

26. Smallest Atom Ne Li Ar Na Largest Atom Increasing Shielding+3
+10 Ne Li
2 shells present-2 inner shell electrons and 1 outer
2 shells present-2 inner shell
electrons and 8 outer
Increasing Shielding
+18
+11 Ar Na
3 shells present-10 inner shell electrons and 1 outer
3 shells present-10 inner shell electrons and 8 outer
Largest Atom
Increasing Shielding

27. A positive ion will be smaller than the atom that created it because there will be more protons than electrons in the ion. The electrons will be pulled more strongly by the nucleus.
A negative ion will be larger than the atom that created because there will be more electrons than protons in the ion.
The greater the charge of the ion, the more pronounced the size change will be. For example, a +3 ion will be smaller than a +2 or +1 ion in the same period and a -3 ion will be larger than a -2 or -1 ion in the same period.

28. F Na F -1 Na+1 9 protons, 9 electrons Zero charge+9 F
+11 Na
9 protons, 9 electrons
Zero charge
11 protons, 11 electrons
Zero charge +9
+11
F -1
Na+1
11 protons, 10 electrons
+1 charge
9 protons, 10 electrons
-1 charge

29. Ionization energy depends upon how close the electron is to the nucleus-the closer the electron, the higher the ionization energy.
Smaller atoms have higher ionization energies!
Larger atoms have lower ionization energies!
Ionization energy will have a trend opposite to atomic radius.
Increasing Atomic Radius
Increasing Ionization Energy

30. 1st Na+ 1s2, 2s2, 2p6 4th Si+4 1s2, 2s2, 2p6 2nd Mg+2 1s2, 2s2, 2p6
Note: a very large increase in ionization energy occurs when ionization would be removing what was previously an “inner shell” electron.
1st
Na+ 1s2, 2s2, 2p6
4th
Si+4 1s2, 2s2, 2p6
2nd
Mg+2 1s2, 2s2, 2p6
5th
P+5 1s2, 2s2, 2p6
3rd
Al+3 1s2, 2s2, 2p6
6th
S+6 1s2, 2s2, 2p6

31. An atom that attracts its outer shell electrons strongly will have a high electronegativity. Within a period, atoms are smaller to the right because the nucleus attracts the outer shell electrons strongly because of increased nuclear charge and decreased shielding.
Therefore electronegativity will have a trend opposite to atomic radius.
Increasing Atomic Radius
Increasing Electronegativity

32. Increasing Ionization Energy Arrows1 2 3 4 5 6 7
Increasing Ionization Energy Arrows
Summary Arrow for Increasing Ionization Energy

33. Increasing Electronegativity Arrows1 2 3 4 5 6 7
Increasing Electronegativity Arrows
Summary Arrow for Increasing Electronegativity
Note: Nobel gases do not form compounds, so they have no electronegativity values.

34. Key Skills for chapter 5:
1) Based on an element’s position on the periodic table, you need to be able to
identify: the name of the family or group, the period, the general classification
(metal, non-metal, or metalloid), and what larger group it belongs to
(representative, transition, inner transition, etc…).
2) Based on the identity of an element, you also need to be able to identify
whether it is a solid, liquid, or gas, and whether it is a diatomic element or not.
3) When given two or more elements you need to be able to correctly arrange
them in order of increasing: atomic size (radii), electronegativity, and ionization
energy.
4) When given a representative element, you need to be able to predict what type
of ion the element will form. Type includes both the sign of the charge (+ or -)
and the amount of charge (1, 2, or 3 etc…).
5) When given two or more ions you need to be able to correctly arrange them in
order of increasing size.

35. Practice 1 2 3 4 5 6 7

36. Which of the following atoms should have the smallest atomic radius?
O element 8
Be element 4
C element 6
Li element 3

37. Which of the following atoms should have the largest atomic radius?
Sr element 38
Be element 4
Ca element 20
Ba element 56

38. Which of the following atoms should have the highest electronegativity?
Cl element 17
Na element 11
Si element 14
Ar element 18

39. Which of the following atoms should have the lowest electronegativity?
S element 16
Te element 52
O element 8
Po element 84

40. Which of the following atoms should have the lowest ionization energy?
Li element 3
K element 19
Na element 11
Cs element 55

41. Which of the following atoms should have the highest ionization energy?
Rb element 37
Zn element 30
S element 16
V element 23

42. Which of the following atoms is an alkali metal?
Rb element 37
Zn element 30
Mg element 12
Al element 13

43. Which of the following atoms is a halogen?
Br element 35
Si element 14
Ca element 20
Ne element 10

44. Which of the following atoms is a noble gas?
N element 7
H element 1
Kr element 36
Be element 4

45. Which of the following atoms is an alkaline earth metal?
Ni element 28
K element 19
Cl element 17
Ca element 20

46. Which of the following atoms is metalloid?
Al element 13
K element 19
P element 15
Ge element 32

47. Which of the following atoms would form a +1 ion when it becomes an ion?
Al element 13
K element 19
F element 9
Ne element 10

48. Which of the following atoms would form a -1 ion when it becomes an ion?
O element 8
Li element 3
Br element 35
Mg element 12

49. Which of the following atoms would be least likely to form an ion?
S element 16
He element 2
Cs element 55
Cl element 17

50. Which of the following ions would be the smallest?
S element 16
Na element 11
Mg element 12
Cl element 17

51. Which of the following ions would be the largest?
S element 16
Na element 11
Mg element 12
Cl element 17

52. Which of the following ions would have the same electron configuration as argon (element 18)?
Br element 35
K element 19
Mg element 12
Ar element 18

53. Which of the following atoms is a transition metal?
Sr element 38
Cu element 29
Xe element 54
Li element 3

54. Which of the following atoms is correctly called a diatomic element?
S element 16
Be element 4
He element 2
N element 7