Lecture 21: Periodic Trends
• Reading: Zumdahl 12.14-12.16
• Outline
– Periodic Trends in the Periodic Table
• Ionization Energy, Electron Affinity, and Radii
– A Case Example
Periodic Trends
• The valence electron structure of atoms can be used
to explain various properties of atoms.
• In general, properties correlate down a group
(column in the table) of elements.
• A warning: such discussions are by nature very
generalized…exceptions do occur.
• General trends do have general reasons (usually in
terms of the valence electrons and the nuclear
charge).
• Often these forces make tradeoffs which can make
trends complex.
Electron Ionization
• If we put in enough energy, we can remove an
electron from an atom.
Z+Z
(Z-1)Energy
+Z
e-
• The electron is completely “removed” from the
atom (potential energy = 0).
• No neutral atom willingly gives up electrons
• However, for example, metals like to give up
electrons to non-metals
Na + 12 Cl2 → ( Na + )( Cl − )
Ionization
• To measure the ionization energy by itself,
generally use photons, with energy measured in
eV (1 eV = 1.6 x 10-19 J).
• The greater the propensity for an atom to “hold
on” to its electrons, the higher the ionization
potential will be.
• Koopmans’ Theorem: The ionization energy of an
electron is equal to the energy of the orbital from
where the electron came. (Not completely true
because there is some electron “rearrangement”)
Sequential Ionization
• One can perform multiple ionizations:
Al(g)
Al+(g) + e-
I1 = 580 kJ/mol
first
Al+(g)
Al2+(g) + e-
I2 = 1815 kJ/mol
second
Al2+(g)
Al3+(g) + e-
I3 = 2740 kJ/mol
third
Al3+(g)
Al4+(g) + e-
I4 = 11,600 kJ/mol
fourth
First Ionization Potential (I.P.)
Column 8
Column 1
First Ionization Potentials
• Increases as one goes from
left to right.
• Reason: increased Z+
• Decrease as one goes down
a group.
• Reason: increased distance
from nucleus
Ionization of Core Electrons
• Removal of valence versus core electrons
Na(g)
Na+(g) + e-
(removing “valence” electron)
[Ne]3s1
[Ne]
Na+(g)
Na2+(g) + e-
[Ne]
I1 = 495 kJ/mol
I2 = 4560 kJ/mol
1s22s22p5 (removing “core” electron)
• Takes significantly more energy to remove a core
electron….tendency for core configurations to be
energetically stable. C.F. first ionization energies.
Electron Affinity
• Electron Affinity: the energy change associated
with the addition of an electron to a gaseous atom.
eZ-
(Z+1)-
+Z
+Z
Energy
Electron Affinity
• We will stick with our thermodynamic definition,
with energy released being a negative quantity.
Wow!
Energetically favorable to take an electron if you fill out a shell (or go
to an inert gas electron configuration).
Periodic Trends: Electron Affinity
• Elements that have high electron affinity:
•Group 7 (the halogens) and Group 6 (O and S specifically).
•The more non-metal or higher the electronegativity of an atom
the greater (more negative) is the E.A.
Electron Affinity (E.A.)
• Some elements will not form ions:
N?
• Orbital configurations can explain both observations.
Periodic Trends: Electron Affinity
• Why is EA so great for the halogens?
F(g) + e1s22s22p5
F-(g)
1s22s22p6
EA = -327.8 kJ/mol
[Ne]
• Why is EA so poor for nitrogen?
N(g) + e1s22s22p3
N-(g)
EA > 0 (unstable)
1s22s22p4
(e- must go into occupied orbital),
loose half filling.
Electron Affinity
• How do these arguments do for O?
O(g) + e1s22s22p4
O-(g)
1s22s22p5
EA = -140 kJ/mol
Bigger Z+ overcomes
e- repulsion.
• What about the second EA for O?
O-(g) + e1s22s22p5
O2-(g)
EA > 0 (unstable)
1s22s22p6
[Ne] configuration, but electron
repulsion is just too great.
Atomic Radii
• Atomic Radii are defined as the covalent radii,
and are obtained by taking 1/2 the distance of a bond:
r = atomic radius
For metals, one can just take half the nearest neighbor
distance in the pure metal.
Atomic Radii of neutral atoms
• Decrease to right due
to increase in Z+
• Increase down column
due to population of
orbitals of greater n.
Summary and Looking
Ahead
We can partition the
periodic table into general
types of elements.
• Metals: tend to give up e• non-Metals: tend to gain e• Metalloids: can do either