Trends & the Periodic Table
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
When you’re
done it will
look like this so
leave room for
writing!
Going down column 1:
Period
Element
Configuration
1
H
1
2
Li
2-1
3
Na
2-8-1
4
K
2-8-8-1
5
Rb
2-8-18-8-1
6
Cs
2-8-18-18-8-1
7
Fr
2-8-18-32-18-8-1
increasing # energy levels as go down
Increasing number
of energy levels
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
Increasing Atomic
Radius
Increasing number of energy levels
Cs has more energy levels, so it’s bigger
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Li: Group 1 Period 2
Cs: Group 1 Period 6
As we go across, elements gain electrons, but they
are getting smaller!
Family
IA or 1
IIA or 2
IIIA or 13
IVA or 14
VA or 15
VIA or 16
VIIA or 17
VIIIA or 18
Element
Li
Be
B
C
N
O
F
Ne
Configuration
2-1
2-2
2-3
2-4
2-5
2-6
2-7
2-8
Increasing number of energy levels
Increasing Atomic Radius
Decreasing
Atomic
Radius
previous | index | next
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
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as go across row size tends to decrease a bit
because of greater PPP “proton pulling power”
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-)
previous | index | next
+7
+1
Calculate “effective nuclear charge”
• # protons minus # inner electrons
What the inner electrons do….
They Shield the charge felt by the valance electrons.
previous | index | next
H and He:
only elements
whose valence
electrons feel
full nuclear
charge (pull)
NOTHING
TO
SHIELD
THEM
Increasing number of energy levels
Increasing Atomic Radius
Decreasing Atomic Radius
Increased Electron
Shielding
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
Positive ions (cations)
• Formed by loss of electrons
• Cations always smaller than parent
atom
2e
8e
8e
2e
Ca
Ca
8e
8e
2e
Ca+2
Negative ions or (anions)
• Formed by gain of electrons
• Anions always larger than parent
atom
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
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)
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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
•THEREFORE, Li has a higher Ionization energy then Cs
Increased Ionization Energy (harder to remove an electron)
Increasing number of energy levels
Increasing Atomic Radius
Increased Electron Shielding
Decreasing Atomic Radius
Decreased Ionization Energy
(easier to remove an electron)
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
Increasing number of energy levels
Increasing Atomic Radius
Increased Electron Shielding
Decreased Ionization Energy (easier to remove an electron)
Increased Electronegativity
Increased Ionization Energy (harder to remove an electron)
Decreasing Atomic Radius
Decreased
Electronegativity
Reactivity of Metals
• judge reactivity of metals by how easily
give up electrons (they’re losers)
Increasing number of energy levels
Increasing Atomic Radius
Increased Electron Shielding
Decreased Ionization Energy (easier to remove an electron)
Decreased Electronegativity
Increased Electronegativity
Increased Ionization Energy (harder to remove an electron)
Decreasing Atomic Radius
More metallic
Most
reactive
metal = Fr
(the most
metallic)
Reactivity of Non-metals
• judge reactivity of non-metals by how
easily gain electrons (they are
winners)
Increased Ionization Energy (harder to remove an electron)
Increasing number of energy levels
Increasing Atomic Radius
Decreasing Atomic Radius
Increased Electron Shielding
Decreased Ionization Energy (easier to remove an electron)
Decreased Electronegativity
Increased Electronegativity
Most Reactive
Nonmetal
=F
More metallic
Most
reactive
metal = Fr
(the most
metallic)
Nonreactiv
e
BACK
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