PPT on Periodic trends

1. Trends & the Periodic Table

2. Trends
• more than 20 properties change in predictable
way based location of elements on PT
• some properties: - anyone know where we can find these numbers?!
– Density
– melting point/boiling point
– atomic radius
– ionization energy
– electronegativity

3. When you’re
done it will
look like this so
leave room for
writing!

4. 2-8-18-32-18-8-1
Fr
7
2-8-18-18-8-1
Cs
6
2-8-18-8-1
Rb
5
2-8-8-1
K
4
2-8-1
Na
3
2-1
Li
2
1
H
1
Configuration
Element
Period
Going down column 1:
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
number
of
energy
levels
Increasing Atomic
Radius

9. previous | index | next
Li: Group 1 Period 2 Cs: Group 1 Period 6
Cs has more energy levels, so it’s bigger

10. 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
As we go across, elements gain electrons, but they
are getting smaller!

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

12. previous | index | next

13. Why does this happen..
• As you go from left to right, you again
more protons (the atomic number
increases)
• You have 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. as go across row size tends to decrease a bit
because of greater PPP “proton pulling power”
previous | index | next

15. We can “measure” the Proton
Pulling Power by determining
the Effective nuclear charge
• It is the charge actually felt by valence electrons
• The equation
Nuclear charge - # inner shell electrons
(doesn’t include valance e-)

16. previous | index | next
Calculate “effective nuclear charge”
• # protons minus # inner electrons
+7 +1

17. What the inner electrons do….
They Shield the charge felt by the valance electrons.

18. previous | index | next
H and He:
only elements
whose valence
electrons feel
full nuclear
charge (pull)
NOTHING
TO
SHIELD
THEM

19. Increasing
number
of
energy
levels
Increasing
Atomic
Radius
Decreasing Atomic Radius
Increased Electron
Shielding

20. Look at all the shielding Francium's one valance
electron has. It barely feels the proton pull from the
nucleus. No wonder it will lose it’s one electron the
easiest. No wonder it’s the most reactive metal

21. Ionization Energy
• = amount energy required to remove a
valence electron from an atom in gas
phase
• 1st ionization energy = energy required to
remove the most loosely held valence
electron (e- farthest from nucleus)

22. •Cs valence electron lot farther away from nucleus than Li
•electrostatic attraction much weaker so easier to steal
electron away from Cs
•THEREFORE, Li has a higher Ionization energy then Cs
previous | index | next

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

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

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

27. Reactivity of Metals
• judge reactivity of metals by how easily
give up electrons (they’re losers)

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

29. Reactivity of Non-metals
• judge reactivity of non-metals by how
easily gain electrons (they are
winners)

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

31. How do you know if an atom gains
or loses electrons?
• Think back to the Lewis structures of ions
• 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

32. Positive ions (cations)
• Formed by loss of electrons
• Cations always smaller than parent
atom
Ca
2e
8e
8e
2e
Ca+2
2e
8e
8e
Ca

33. Negative ions or (anions)
• Formed by gain of electrons
• Anions always larger than parent
atom

34. Allotropes
• Different forms of element in same phase
– different structures and properties
• O2 and O3 - both gas phase
–O2 (oxygen) - necessary for life
–O3 (ozone) - toxic to life
• Graphite, diamond:
–both carbon in solid form