1. Trends & the Periodic Table

3. Trends
more than 20 properties change in predictable way based location of elements on Periodic Table
some properties:
Density
melting point/boiling point
atomic radius
ionization energy
electronegativity

4. Going down column 1: Fr 7 Cs 6 Rb 5 K 4
2-8-1 Na 3
2-1 Li 2 1 H
Configuration
Element
Period
increasing # energy levels as go down

5. Increasing number of energy levels

6. Atomic Radius
Atomic radius: defined as ½ distance between neighboring nuclei in molecule or crystal
Affected by
1. # of energy
levels
2. Proton Pulling
Power

8. Increasing Atomic Radius
Increasing number of energy levels
Increasing Atomic Radius

9. Cs has more energy levels, so it’s bigger
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Li: Group 1 Period Cs: Group 1 Period 6

10. As we go across, elements gain electrons, but they are getting smaller!
2-8 Ne
VIIIA or 18
2-7 F
VIIA or 17
2-6 O
VIA or 16
2-5 N
VA or 15
2-4 C
IVA or 14
2-3 B
IIIA or 13
2-2 Be
IIA or 2
2-1 Li
IA or 1
Configuration
Element
Family

11. Decreasing Atomic Radius
Increasing number of energy levels
Increasing Atomic Radius

12. previous | index | next
                                                                                                                                                                                                                  

13. Why does this happen..
Moving from left to right, more protons are added (the atomic number increases)
The nucleus has greater “proton pulling power”
Remember the nucleus is + and the electrons are - so they get pulled towards the nucleus
The more protons your have, the more Proton Pulling Power

14. Why does this happen..
Moving from left to right, more electrons are added too, but NOT TO ADDITIONAL ENERGY LEVELS.
This means there is not extra space between nucleus and valence electrons.

15. Why does this happen..
The greater proton pulling power causes the electrons to be pulled closer to the nucleus and the radius decreases.

16. previous | index | next
                                                                                                                                                                                                                  
as go across row size tends to decrease a bit because of greater PPP “proton pulling power”

17. SAMPLE PROBLEM 8.3
Ranking Elements by Atomic Size
PROBLEM:
Using only the periodic table (not Figure 8.15)m rank each set of main group elements in order of decreasing atomic size:
(a) Ca, Mg, Sr
(b) K, Ga, Ca
(c) Br, Rb, Kr
(d) Sr, Ca, Rb
PLAN:
Elements in the same group decrease in size as you go up; elements decrease in size as you go across a period.
SOLUTION:
(a) Sr > Ca > Mg
These elements are in Group 2A(2).
(b) K > Ca > Ga
These elements are in Period 4.
(c) Rb > Br > Kr
Rb has a higher energy level and is far to the left. Br is to the left of Kr.
(d) Rb > Sr > Ca
Ca is one energy level smaller than Rb and Sr. Rb is to the left of Sr.

18. Increased Space between nucleus and electrons
Decreasing Atomic Radius
Increasing number of energy levels
Increasing Atomic Radius
Increased Space between nucleus and electrons

19. Francium's one valance electron barely feels the proton pull from the nucleus. It is very far away from the nucleus and has 86 other electrons between it and the nucleus. No wonder it will lose its one electron the easiest. No wonder it’s the most reactive metal

20. Ionization Energy
= amount energy required to remove a valence electron from an atom in gas phase
1st ionization energy = energy required to remove the first e- [farthest from nucleus]

21. Cs valence electron lot farther away from nucleus than Li
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Cs valence electron lot farther away from nucleus than Li
electrostatic attraction much weaker so easier to steal electron away from Cs
Li has a higher Ionization Energy then Cs

22. First ionization energies of the main-group elements
Trends in the Periodic Table

23. Increased Ionization Energy (harder to remove an electron)
Decreasing Atomic Radius
Increased Electron Shielding
Increasing number of energy levels
Increasing Atomic Radius
Decreased Ionization Energy
(easier to remove an electron)

24. SAMPLE PROBLEM 8.4
Ranking Elements by First Ionization Energy
PROBLEM:
Using the periodic table only, rank the elements in each of the following sets in order of decreasing IE1:
(a) Kr, He, Ar
(b) Sb, Te, Sn
(c) K, Ca, Rb
(d) I, Xe, Cs
PLAN:
IE increases as you proceed up in a group; IE increases as you go across a period.
SOLUTION:
(a) He > Ar > Kr
Group 8A(18) - IE decreases down a group.
(b) Te > Sb > Sn
Period 5 elements - IE increases across a period.
(c) Ca > K > Rb
Ca is to the right of K; Rb is below K.
(d) Xe > I > Cs
I is to the left of Xe; Cs is further to the left and down one period.

25. Electronegativity ability of atom to attract electrons in bond
noble gases tend not to form bonds [octet rule], so don’t have electronegativity values
Unit = Pauling
Fluorine: most electronegative element
= 4.0 Paulings

26. Decreased Electronegativity
Increased Electronegativity
Increased Ionization Energy (harder to remove an electron)
Decreasing Atomic Radius
Increased Electron Shielding
Decreased Ionization Energy (easier to remove an electron)
Increasing number of energy levels
Increasing Atomic Radius
Decreased Electronegativity

30. Reactivity of Metals
judge reactivity of metals by how easily give up electrons
Why do they lose electrons? Sodium Atom: 1s22s22p63s1

31. Reactivity of Metals Sodium Cation: 1s22s22p6
judge reactivity of metals by how easily give up electrons
Sodium Cation: 1s22s22p6

32. Most reactive metal = Fr
Increased Electronegativity
Increased Ionization Energy (harder to remove an electron)
Decreasing Atomic Radius
More metallic
Increased Electron Shielding
Decreased Ionization Energy (easier to remove an electron)
Increasing number of energy levels
Decreased Electronegativity
Increasing Atomic Radius
Most reactive metal = Fr
(the most metallic)

33. Reactivity of Non-metals
judge reactivity of non-metals by how easily gain electrons
Why do they gain? Chlorine Atom: 1s22s22p63s23p5

34. Reactivity of Non-metals
judge reactivity of non-metals by how easily gain electrons
Why do they gain? Chlorine Anion: 1s22s22p63s23p6

35. Most reactive metal = Fr
Increased Electronegativity
Increased Ionization Energy (harder to remove an electron)
Decreasing Atomic Radius
Most Reactive Nonmetal
= F
More metallic
Increased Electron Shielding
Decreased Ionization Energy (easier to remove an electron)
Increasing number of energy levels
Decreased Electronegativity
Increasing Atomic Radius
Most reactive metal = Fr
(the most metallic)
Nonreactive
BACK

36. How do you know if an atom gains or loses electrons?
Atoms form ions to get a valence of 8 (or 2 for H)
Metals tend to have 1, 2, or 3 valence electrons
It’s easier to lose them
Nonmetals tend to have 5, 6, or 7 valence electrons
It’s easier to add some
Noble gases already have 8 so they don’t form ions very easily.
The easier an atom will lose an electron, the MORE METALLIC it is.

37. Trends in metallic behavior

38. Positive ions (cations)
Formed by loss of electrons
Cations always smaller than neutral atom
Strong proton pull since more protons than electrons. 2e 8e 8e 8e 8e 2e 2e Ca Ca
Ca+2

39. Negative ions or (anions)
Formed by gain of electrons
Anions always larger than neutral atom
Electrons repel each other and create space between each other.
Weaker proton pull
since less protons
than electrons

40. Ionic vs. atomic radius

41. Ranking Ions by Size
PROBLEM:
Rank each set of ions in order of decreasing size, and explain your ranking:
(a) Ca2+, Sr2+, Mg2+
(b) K+, S2-, Cl -
(c) Au+, Au3+
PLAN:
Compare positions in the periodic table, formation of positive and negative ions and changes in size due to gain or loss of electrons.
SOLUTION:
(a) Sr2+ > Ca2+ > Mg2+
These are members of the same Group (2A/2) and therefore decrease in size going up the group.
The ions are isoelectronic; S2- has the smallest proton pull and therefore is the largest while K+ is a cation with a large proton pull is the smallest.
(b) S2- > Cl - > K+
(c) Au+ > Au3+
The higher the + charge, the smaller the ion.