1. Periodicity
Chapter 8, Sections 5-7

2. Periodic Repeats periodically Forms a pattern, like…
Day/Night
Cycle of the moon
TENDS to go up or go down

3. Oxidation and Reduction
Here’s a periodic property you’ve already seen…
Oxidation – The loss of electrons
Which elements tend to lose electrons?
Reduction – The gain of electrons
Which elements tend to gain electrons?
metals
So metals oxidize the best and are the best reducing agents.
nonmetals
So nonmetals reduce the best and are the best
oxidizing agents.

4. Atomic Radius Commonly known as covalent radius
This is 1/2 the distance between two nuclei of the same elements that are covalently bonded.

5. Atomic Radius
Notice, what happens to the atomic size going down a group?
Why? (What’s occurring within the atoms’ structures?)
There are more electrons.
There are more protons.
There are more energy levels.
Higher energy levels are further from the nucleus, so the atom gets larger.

6. Atomic Radius
Notice, what happens to the atomic size going across a period?
Why? (What’s occurring within the atoms’ structures?)
There are more electrons.
There are more protons.
The outer energy level DOES NOT CHANGE!!!
The outer energy level does not force the atom to get larger.
The increased attraction between e– and p causes the atom to get smaller.

7. Atomic Radius Br < Se < Te Example:
Refer to a periodic table and arrange the following in order of increasing atomic radius: Br, Se, Te.
Se and Br will be similar in size because they are in the same energy level. However, Se will be larger than Br because it has fewer electrons and protons attracting each other. Se Br Te
Te has to be the largest since it
is in the highest energy level.
Br < Se < Te

8. Ionic Radius v. Atomic Radius
What happens to the atom’s size when it turns into an ion?
If a positive ion is formed, what happens to the electrons?
In addition, the number of electrons is now less than the number of protons. So the nuclear attraction is stronger.
Often, when losing electrons, the outer energy level is lost as well.

9. Ionic Radius v. Atomic Radius
When a negative ion is formed, what happens to the electrons?
When there are more electrons than the protons, the nucleus can not attract the electrons as well. Because its attractions are weaker, the atom gets larger.
When an atom gains an electron, the number of electrons is now higher than the number of protons in the nucleus.

10. Ionic Radius v. Atomic Radius

11. Ionic Radius v. Atomic Radius
In summary…
If the number of electrons becomes lower than the number of protons, the nuclear attraction becomes stronger = the atom gets smaller
If the number of electrons is greater than the number of protons, the nucleus can not attract the electrons as well so the nuclear attractions are weaker = the atom gets larger.

12. Ionization Energy
When an atom becomes positively charged, it absorbs energy.
Breaking attractions = endothermic
This energy is called an ionization energy (IE).
The first ionization energy is the amount of energy to remove the first electron from an atom.
Energy + M  M+ + e– IE

13. Ionization Energy
What happens to the ionization energy going down a group?
Why?
What’s going on with the size of the atom?

14. Ionization Energy
What happens to the ionization energy going across a period?
Why?
What’s going on with the size of the atom?

15. Ionization Energy Energy + M  M+ + e– The first ionization energy…
becomes smaller as the atomic radius gets larger, i.e. going down a group.
This is because there are fewer attractions between the nucleus and the outermost electrons so less energy is required to remove the electron.

16. Ionization Energy Energy + M  M+ + e– The first ionization energy…
becomes larger as the atomic radius gets smaller, i.e. going across a period.
This is because the attractions between the nucleus and the outermost electrons are stronger so more energy is required to remove the electron.

17. Ionization Energy
Energy + M  M+ + e–
In summary, the greater the attraction for the electron, the more endothermic the ionization energy

18. Ionization Energy There are also successive ionization energies.
Electrons that are removed after having already taking off electrons are create successive ionization energy.
The ionization energy (IE) number will be indicated with a subscript (IEi).
IE1+ M  M+ + e–
IE2+ M+  M e–
IE3+ M2+  M e–

19. Ionization Energy For each element, where are the most
IE1
IE2
IE3
IE4
IE5
IE6
IE7 Na
498
4560
6910
9540
13400
16600
20100 Mg
736
1445
7730
10600
13600
18000
21700 Al
577
1815
2740
11600
15000
18310
23290 Si
787
1575
3220
4350
16100
1/900
23800 P
1063
1890
2905
4950
6270
21200
25400 S
1000
2260
3375
4565
6950
8490
27000 Cl
1255
2295
3850
5160
6560
9360
11000 Ar
1519
2665
3945
5770
7230
8780
12000
For each element, where are the most
distinct jumps in energy?

20. Ionization Energy Element IE1 IE2 IE3 IE4 IE5 IE6 IE7 Na 498 4560 6910
9540
13400
16600
20100

21. Ionization Energy Element IE1 IE2 IE3 IE4 IE5 IE6 IE7 Na 498 4560 6910
9540
13400
16600
20100
Why does it take almost nine times the amount of energy as the first ionization energy?

22. Ionization Energy IE1 + IE2 = 2181 kJ Element IE1 IE2 IE3 IE4 IE5 IE6Mg
736
1445
7730
10600
13600
18000
21700
In order to make a magnesium +2 ion, 2 electrons must be lost…
IE1 + IE2 = 2181 kJ

23. Ionization Energy IE3 = 7730kJ Element IE1 IE2 IE3 IE4 IE5 IE6 IE7 Mg
736
1445
7730
10600
13600
18000
21700
Why does it take so much energy to take off a 3rd electron?
IE3 = 7730kJ

24. Ionization Energy Sb < As < Br Example:
Refer to a periodic table and arrange the following in order of increasing ionization energy:
As, Br, Sb.
As and Br will be similar in size (and IE) because they are in the same energy level. However, As will be larger in size than Br so it will have a lower IE than Br. As Br Sb
Sb has to have the smallest ionization energy since its outer energy level is the furthest away.
Sb < As < Br

25. Electron Affinity
When an atom becomes negatively charged (gains an electron, it releases energy.
Forming attractions = exothermic
This energy is called electron affinity (EA).
e– + X  X– + Energy EA

26. Electron Affinity

27. Electron Affinity
Notice that the alkaline earth metals would need to add a subshell to hold another electron. Creating a higher energy subshell would be an endothermic process so gaining an electron won’t occur.

28. Electron Affinity
Notice that the noble gases would also need to add a subshell to hold another electron.
Creating a higher energy subshell would be an endothermic process so gaining an electron won’t occur.

29. Electron Affinity
The halogens, on the other hand, can use the added electron to complete the subshell.
This is a highly exothermic process so gaining an electron is very likely.

30. Electron Affinity
In general, what is the trend for electron affinities headed across the periods?

31. Electron Affinity
Going down a group, why does the electron affinity magnitude become smaller?

32. Electron Affinity
e– + X  X– + Energy
In summary, the greater the attraction for the electron, the more exothermic the electron affinity.

33. Electronegativity
Electronegativity is the measure of the tendency for an atom to attract an electron.
The measure of electronegativity occurs on a scale.
Not likely to attract an electron
4.00
Very likely to attract an electron

34. Electronegativity

35. Metallic Character
Metallic character includes all of the properties of metals.
Conductivity of electricity
Conductivity of heat
Luster
Ductility
Malleability
Reactivity with water
Reactivity with acids

36. Metallic Character
The properties of metals are created by their bonds…
metallic bonds which are produced when the electron clouds of the atoms fuse together to make an electron sea.

37. Metallic Character
What is the trend for metallic character?

38. Explanations Going down the periodic table…
As the principle quantum number (energy level) increases, the nuclear attractions to the outermost electrons…
decreases

39. Explanations Going down the periodic table…
As the principle quantum number (energy level) increases, the nuclear attractions to the outermost electrons…
decreases
As more energy levels fall in between the nucleus and the outermost electrons they shield (hinder) the nuclear attractions to those electrons…
the shielding effect

40. Explanations Going across the periodic table…
The number of electrons and protons increases while the energy level stays the same…
This increases the attractions to the nucleus

41. Explanations
Comparing one subshell to another subshell in the same energy level…
A full subshell will shield another subshell from nuclear attractions, making the nuclear attractions weaker.
A higher energy subshell is further from the nucleus, so the nuclear attractions are weaker.

42. Explanations
Comparing paired v. unpaired electrons of the same subshell…
UNLIKE comparing one subshell to another subshell, the amount of shielding remains the same.
So what happens when two electrons share the same space?

43. Explanations Comparing a charged atom to a neutral atom…
A neutral atom has the same number of e– as p.
A positive ion has                e– than p.
fewer
This causes the nuclear attractions to be significantly greater.

44. Explanations Comparing a charged atom to a neutral atom…
A neutral atom has the same number of e– as p.
A negative ion has                e– than p.
more
This causes the nuclear attractions to be significantly weaker.

45. Explanations Going down a group Going across a period Charged atoms
Principle quantum number Or
Shielding effect
Create weaker attractions
Going across a period
Same energy level but greater attractions
between p and e–.
Two different subshells
Shielding effect from inner subshell creating weaker attractions
Same subshell
Unpaired e–’s v. Paired e–’s
Paired e–’s repel
Charged atoms
Negative ion
More e–’s than creating weaker attractions
Positive ion
Fewer e–’s than p creating stronger attractions