1. Section 1 Electron structure and the periodic table
Periodic Table II
Section 1
Electron structure and the periodic table

2. Recall that the Periodic table is arranged according to properties?
From
reactive
metals
sodium
Mercury
To less
reactive metals

3. Recall that the Periodic table is arranged according to properties?
helium
…and
inert (noble)
gases
non-metals
silicon To
semi-metals
carbon

4. The Periodic table is also arranged according to electron arrangement8
Valence
electrons
1, 2, 3
Valence
electrons
4 to 7
Valence electrons

5. So the properties must depend on electron arrangement!
The Periodic table is also arranged according to electron arrangement
So the properties must depend on electron arrangement!

6. Periods (horizontal) 1, 2, 3 etc.
Periods have elements with same number of occupied energy levels (shells) OR
Have the same outer (valence) shell

7. Periods (horizontal) 1, 2, 3 etc.
Periods have elements with same number of occupied energy levels (shells) OR Have the same outer (valence) shell
Example:
Elements in period 3
have electrons in
three energy shells OR
the 3rd shell
is the valence level.
Al e- config. = 2 – 8 – 3

8. Periods (horizontal) 1, 2, 3 etc.
Periods have elements with same number of occupied energy levels (shells) OR Have the same outer (valence) shell
Within a period
electrons vary from 1 to 8 Valence electrons
Also,
the properties change as the atomic number increases

9. Groups (vertical) 1,2,…18 8 7 Valence electrons 6 1 2 5 4 3 1 or 2
Groups have elements with the same number of valence electrons 8 7
Valence electrons 6 1 2 5 3 4
1 or 2
Elements with the same number of valence electrons make for similar properties in groups

10. Crash course chemistry
Click here for
Crash course chemistry
on Periodic Table

11. Test Your Understanding
What aspect of an atoms structure allows you to predict its properties? How does the arrangement of elements on the table support your conclusion?
Why do the transition elements have such similar but unique properties?
What change in the atomic structure of the elements in a period accompanies the change in properties we observe?

12. HDYK?

13. Periodic Trends > Size of atoms > Strength of atoms
Section 2
Periodic Trends
> Size of atoms
> Strength of atoms

14. 2A) Periodic Trends: Size of atoms
Atoms get larger down the table
How can you explain these trends?
(how do the structure of atoms differ?
But get smaller left to right

15. Firstly, how are atoms measured?
Shooting X-rays at a sample reveals where the nuclei are located, but not the electrons.
Atomic nuclei

16. Firstly, how are atoms measured?
If you assume that the outer shell of each atom reaches half way between two atoms.
Then ½ the distance between the two nuclei is each atom’s Radius.
D/2
radius
radius

17. Secondly, what determines the radius?
The size of an atom depends on two factors: #1
The number of electron shells.
More shells take up more space
Atoms get bigger!

18. Periodic Trends: Size of atoms
As atomic number increases going down a group each atom has an additional energy level, further from the nucleus.
Radius goes UP since the valence shell is now further from the nucleus.

19. The pull from the protons in the nucleus
Secondly,
The size of an atom depends on two factors: #2
The pull from the protons in the nucleus

20. Secondly, what determines the radius?
The size of an atom depends on two factors:
A stronger nuclear charge will pull electrons closer.
so atoms are smaller.

21. As Atomic number increases across a period, atoms are gaining more protons in their nuclei (greater nuclear charge)
Radius decreases - since the extra positive charge from the protons pulls the electrons closer to the nucleus
11+
12+
13+
14+
15+
16+
17+
Large
Small

22. Summary: Atomic radii Click here for smaller Khan academy tutorial
on atomic radius
smaller
Overall Trend
larger

23. Test your learning How is the size of an atom expressed? Why?
Which atom is larger; H or He? Why?
Which is larger; and argon atom or a potassium atom? Why?
Describe the overall trend in atomic radius for elements on the periodic table.

25. Introduction to chemical reactions Atomic “strength”
Section 2B
Introduction to chemical reactions
Atomic “strength”

26. Why do atoms react with one another?
Recall that the group 18 Noble gases have 8 valence electrons?
Recall that the group 18 Noble gases are very stable?
Atoms with 8 valence electrons are very stable since they require a lot of energy to remove any of their electrons.

27. Why do atoms react with one another?
Nonmetal atoms like oxygen can become stable by gaining some additional electrons.
The new ion that is formed is no longer neutral

28. Why do atoms react with one another?
Ions like this now called “oxide “ are negative since they have two more electrons than protons
Negative ions like this are called an-ions

29. Why do atoms react with one another?
Its new symbol is O2-
Its new electron configuration is now 2-8 like neon
Negative ions like this are called an-ions

30. Why do atoms react with one another?
Metallic atoms like sodium can become stable by losing some electrons.

31. Why do atoms react with one another?
Sodium’s electron configuration was 2-8-1
Now becomes 2-8
Stable like neon.

32. Why do atoms react with one another?
The symbol for the sodium ion is Na1+
since it lose an electron it now has less electrons than protons and carries an excess positive charge.
Positive ions are called
Cat-ions

33. Forming an ion: So atoms lose or gain electrons to form ions Why?
O 2- = 8 P e-
= -2 charge
Na+ = 11 P e-
= + 1 charge
But, why are some atoms losers,
and other atoms gainers?

34. Atomic Strength: Ionization Energy
Energy is required to remove an electron from an atom.
This is called it ionization energy (energy to ion-ize)
Mg kJ ---> Mg+ + e-
Energy
Notice energy is added to Mg to remove one electron.
This is called the FIRST ionization energy because we removed only one of the outer electrons
ion and electron

35. Atomic Strength: Ionization Energy
Energy is required to remove an electron from an atom.
This is called it ionization energy (energy to ion-ize)
Mg kJ ---> Mg2+ + e-
It takes even more energy to remove a second electron.
This is the SECOND ionization energy.

36. Atomic Strength: Ionization Energy
Energy is required to remove an electron from an atom.
This is called it ionization energy (energy to ion-ize)
Metals like magnesium will lose electrons since their ionization energies are relatively low.

37. Atomic Strength: Ionization Energy
Energy is required to remove an electron from an atom.
This is called it ionization energy (energy to ion-ize)
In other words, metals are “weak”

38. Atomic Strength: Ionization Energy
Energy is required to remove an electron from an atom.
This is called it ionization energy (energy to ion-ize)
F kJ ---> F+ + e-
Removing an electron from a nonmetal requires too much energy

39. Atomic Strength: Ionization Energy
Energy is required to remove an electron from an atom.
This is called it ionization energy (energy to ion-ize)
F kJ ---> F+ + e-
Nonmetals like fluorine are too strong to lose electrons

40. Atomic Strength: Ionization Energy
Energy is required to remove an electron from an atom.
This is called it ionization energy (energy to ion-ize)
F kJ ---> F+ + e-
In fact, nonmetals gain electrons from metals and become stable.

41. Atomic Strength: Ionization Energy
Energy is required to remove an electron from an atom.
This is called it ionization energy (energy to ion-ize)
F + e-  F kJ
In fact energy is released when F gains and electron

42. Atomic Strength: Ionization Energy
Energy is required to remove an electron from an atom.
This is called it ionization energy (energy to ion-ize)
F + e-  F kJ
This is called electron affinity

43. Trends in Ionization Energy
Ionization energy increases across a period because the positive nuclear charge increases (more protons).
losers
11+
12+
13+
14+
15+
16+
17+
Period 3
496 kJ kJ kJ
As the nuclear charge gets stronger, they exert a greater pull on the outer electrons

44. Trends in Ionization Energy
Ionization energy increases across a period because the positive nuclear charge increases (more protons).
Ionization energy increases
gainers
losers
Metals (left side) are weak - lose electrons more easily: losers
Nonmetals (right side ) strong - don’t lose electrons: gainers

45. Going down a group
Atoms get larger: this means there is a greater distance from the nucleus to valence electrons.
So they get weaker.
and Ionization energy decreases.
decreases
Ionization energy

46. Atoms with lots of inner shells. Ionization energy decreases
Shielding Effect
Atoms with lots of inner shells.
Ionization energy
decreases
*Each additional energy level adds electrons that repel the valence electrons

47. Notice the Overall Trend
Like radius, ionization energy
follows the diagonal trend.

48. Non-metal atoms are strong And small.
Notice the Overall Trend
Metal atoms are weak
And large

49. Periodicity: a trend that occurs at regular intervals
Click here for Khan academy Tutorial
Ionization energy trends
Notice how ionization energy increases,
then decreases with each new period.

50. What does diagram this illustrate?

51. Learning check
What is ionization energy? What does it indicate about an element?
Which will have a higher ionization energy H or He? Why?
Which will have a larger ionization energy Ar or K? Why?
Describe the overall trend in ionization energy on the periodic table.

53. Other important trends
Section 3
Other important trends

54. Ion Sizes Metal ions are SMALLER than the atoms from which they come.
Forming a cation.
Metal ions are SMALLER than the atoms from which they come.
Loss of entire energy level so size DECREASES.

55. Forming an anion
Nonmetal ions are LARGER than the atoms from which they come.
Additional electrons repel each other more so size INCREASES (“swells up”)

56. Think about it Which of the following is the smallest:
O2- ion, Ne atom, or Mg2+ ion?
Why?
Hint: Each has 10 electrons.
But: How many protons?
+10 +8
+12
oxide
Neon
magnesium
8 protons
10 protons
12 protons

57. Think about it Which of the following is the smallest:
O2- ion, Ne atom, or Mg2+ ion? +8
+10
+12
Largest    middle    smallest
oxide
Neon
magnesium
8 protons
10 protons
12 protons
As nuclear charge increases, the size will decrease

58. Test your understanding
What is the electron configuration for a magnesium atom?
In terms of electrons, what will likely happen to magnesium during a chemical reaction?
What will be its new electron configuration?
Draw the Bohr diagram for the magnesium atom and the magnesium cation.
Which will be larger? Why?

59. Atomic # 12 = Mg
9 An atom of an element has a total of 12 electrons. An ion of the same element has a total of 10 electrons. Which statement describes the charge and radius of the ion?
(1) The ion is positively charged and its radius is
smaller than the radius of the atom.
(2) The ion is positively charged and its radius is
larger than the radius of the atom.
(3) The ion is negatively charged and its radius
is smaller than the radius of the atom.
(4) The ion is negatively charged and its radius
is larger than the radius of the atom.
12 P e-

60. Loser metal

61. Chemical strength: Electronegativity
measure of the ability of an atom in a molecule to attract electrons to itself.
Attraction for a pair of electrons
shared in a chemical bond
 Chemical bond 

62. Electronegativity: same as trends in Ionization energy
Why no values for group 18?
Increases: stronger nucleus = stronger force
Decreases:
More
Energy
Levels =
Weaker
force
decreases
increases

64. Why? HDYK? Which has a higher: 1st ionization energy? Mg or Ca ?
Al or S ?
Cs or Ba ?
Circle one in each pair
Why?
Which is more electronegative?
F or Cl ?
Na or K ?
Sn or I ?
HDYK?

65. Ionization Energy

66. Introduction to Chemical Reactivity
Atoms lose or gain electrons to become stable like group 18 elements
(8 valence e- ‘s)
Ex: Ne 2 - 8
What’s the magic number?

67. What elements sit beyond group 18?
Metals lose
To become
Like
Noble gases
Nonmetals gain
To become like
Noble gases
F: 2-7, becomes 2-8
Na: 2-8-1, becomes 2-8

68. Metals – group 1,2 and transitional, etc.
Chemically weak, tend to lose outer electrons to stronger atoms
Ex: Mg loses 2e-
becomes like Ne
Becomes a positive ION Mg2+
ION = charged “atom”
CATION – a positively charged ion

69. gain 1 e-
Nonmetals – groups 15, 16, 17
Chemically strong tend to gain electrons from weaker atoms
Ex: Fluorine atom F gains 1 e- becomes like Ne
Becomes a negative fluoride ION: F1-
ANION - a negative ion

71. 13 When a lithium atom forms an Li+ ion,
the lithium atom
(1) gains a proton
(2) gains an electron
(3) loses a proton
(4) loses an electron
37 What is the total number of electrons in a
Cu+ ion?
(1) 28 (3) 30
(2) 29 (4) 36

72. F gains e-

73. Activity / Reactivity More Active metals F More Active nonmetals Fr
Metals –”losers” must lose electrons during reactions
weakest (lowest electronegativity) are most active (francium)
Nonmetals – gainers must gain electrons during reactions
strongest (highest electronegativity) are most active (fluorine)
More
Active
metals F
More
Active
nonmetals Fr

74. Metallic / nonmetallic “character”
More
(loser)
metallic
character
More
(gainer)
Nonmetallic
character

75. 33 As the elements in Group 17 on the Periodic Table are considered from top to bottom, what happens to the atomic radius and the metallic character of each successive element?
(1) The atomic radius and the metallic character both increase.
(2) The atomic radius increases and the metallic character decreases.
(3) The atomic radius decreases and the metallic character increases.
(4) The atomic radius and the metallic character both decrease.

76. Hint: standard temperature = 00C
For questions 73 – 76 Refer to the table below and your knowledge of chemistry
Hint: standard temperature = 00C

77. 181 98 64

78. 1) How many elements are solids at STP?
2) Which element is a noble gas?
3) Letter Z below corresponds to a different element. Elements G,Q, L and Z are in the same group on the periodic table as shown to the right.
4) Based on the trend in melting points for G, Q and L, estimate the melting point of Element Z

80. Can you explain this trend in ionization energy?
+19
1) K kJ ---> K+ + e-
Energy
2) Na kJ ---> Na+ + e-
Energy
+11
3) Mg kJ ---> Mg+ + e-
Energy
4) Mg kJ ---> Mg+2 + e-
Energy
+13
Hint: compare the structures on the right 