Chapter 6: Periodic Trends
Notes goals
• Be able to use the periodic table to predict
chemical trends
Periodic Trends
•
•
•
•
Effective Nuclear Charge
Atomic Radius
Ionic Radius
Ionization Energy
Effective Nuclear Charge
• Valence electrons = electrons in outer level
• Core electrons = electrons in levels
underneath outer level
• Valence electrons are:
– Attracted to protons in nucleus
– Repelled by core electrons which are in the way
(shielding)
Effective Nuclear Charge
• Effective Nuclear Charge = the net
attraction of the valence electrons to the
nucleus
– Greater effective nuclear charge means valence
electrons are closer to the nucleus
Effective Nuclear Charge (Zeff )
Felt
by
Shielding electrons (S)
Effective Nuclear Charge (Zeff)
Atomic Number (Z = #p+)
Zeff = Z Attractive
Charge felt
by valence
electrons
Atomic
Number
(#p+)
S
Inner Core of
Shielding e= screening
constant
{Eff.Nuclear.Charge*}
Effective Nuclear Charge (Zeff )
depends on both the fact that
valence electrons are both:
• attracted to the nucleus
• repelled by the other (shielding)
electrons.
Effective Nuclear Charge (Zeff.) Increasing
+1
Shielding
electrons
(0e-)
(2e-)
(10e-)
+8
+2
+3 +4 +5 +6 +7
Effective Nuclear Charge Trend
• Effective Nuclear Charge
Increases DOWN and to the RIGHT
1
2
3
4
5
6
7
6.3
Atomic Radius
• The atomic radius is one half of the distance
between the nuclei of two atoms of the same
element when the atoms are joined.
Atomic Radius
• Why larger going down?
– Higher energy levels have larger orbitals
• Why smaller to the right?
– Increased nuclear charge without additional
shielding pulls electrons in tighter
Atomic Radius
Atomic Radius
Atomic Radius Trend
• Atomic Radius
Increases to the LEFT and DOWN
1
2
3
4
5
6
7
Ionic Radius
• Cations (positive ions)
– Lose electron
– Become smaller than
the neutral atom
• Anions (negative ions)
– Gain electron
– Become larger than the
neutral atom
• Trend is the same as
the atomic radius
Ionic Radius
Ionic Radius Trend
• Ionic Radius
Increases to the LEFT and DOWN within
the cations and the anions
Cations (+)
1
2
3
4
5
6
7
Anions (-)
Ionization Energy
• Ionization Energy = Energy required to
remove an electron from a neutral atom
• Smaller atoms have electrons closer to the
nucleus so they are held tighter.
– This means a higher ionization energy is
required to take an electron away.
• There are a few exceptions where where s,
p, or d orbitals are filled or half-filled which
creates a more stable structure
Ionization Energy
Ionization Energy Trend
• First Ionization Energy
Increases UP and to the RIGHT
1
2
3
4
5
6
7