Chapter 7: Periodic Properties of the Elements
Electronegativity
Ionization Energy & Electron Affinity
Ionization Energy & Electron Affinity
Electronegativity
There are a number of atomic characteristics that either increase
or decrease along the periodic table.
Atomic Radius:
As you go down a
group, the atomic
radius increases with
increasing energy
level.
Atomic radius
decreases as you
go left to right along
a period because the
greater nuclear
charge pulls the
electrons in closer to
the nucleus.
s-block
p-block
Why ionization energy decreases and atomic/ionic radius
increases as go down a group:
Shielding effect: The inner electron shells insulate the
valence electrons from some of the electrical attraction with
the positive charge of the nucleus.
+
valence enucleus
core electrons
Why ionization energy increases and atomic/ionic radius
decreases as go across a period:
Increasing Effective Nuclear charge(Zeff): Electrons in the
outermost energy levels do not effectively screen each other
from an increasingly positive nucleus.
nucleus
-
-
-
-
+
+
+
+
+
+
+
Zeff = Z - S
Li: Zeff = 3 – 2 = 1
N: Zeff = 7 – 2 = 5
A few definitions:
van der Waals radius: The nonbonding radius of atoms
Covalent radius:
Isoelectronic:
The radius of atoms covalently bonded to
another atom; is smaller than the van der
Waals radius
Different ions that have the same number of
electrons
Ex: O2-, F-, Na+, Mg2+, Al3+ all have 10 electrons
Ionization energy: The amount of energy required to
remove an electron from an atom or ion.
Electron affinity:
The amount of energy released when an
electron is added to an atom or ion.
There is a spike in ionization energy whenever a noble gas
electron configuration is disrupted.
Q: What is the valence electron configuration for an atom that
has the following ionization energies:
1st: 734 kJ/mol, 2nd: 1850 kJ/mol, 3rd: 16,432 kJ/mol
A: ns2
Large spike in IE indicates noble gas core is disrupted
Electron Affinity: the ability of an atom to gain an electron. This
is closely related to ionization energy, and increases going left to
right, and decreases going down.
Electronegativity: Ability of an atom to attract electrons
when in a molecule.
Electronegativity increases going left to right, and decreases
going down.
Electron configurations of ions
•If electrons are added to make an anion, they fill the
lowest energy levels first. (Auf bau principle)
•If electrons are removed to make a cation, the are taken
from the highest energy levels first.
Atom
Li
1s2 2s1
Ion
1s2 2s0
Li+
Fe [Ar] 4s2 3d6 [Ar] 4s0 3d6
Fe2+
[Ar] 4s0 3d5
Fe3+
Fe3+:
4s
3d
Notice that the Fe3+ ion
has the maximum
multiplicity possible.
Hence, its greater
stability wrt Fe2+
Electron configuration exceptions and their ions
Atom
Cr
Cu
Au
[Ar] 4s1 3d5
[Ar] 4s1 3d10
[Xe] 6s1 4f14 5d10
Ion
[Ar] 4s0 3d5
Cr+
[Ar] 4s0 3d3
Cr3+
[Ar] 4s0 3d0
Cr6+
[Ar] 4s0 3d10
Cu+
[Ar] 4s0 3d9
Cu2+
[Xe] 6s0 4f14 5d10
Au+
[Xe] 6s0 4f14 5d8
Au3+
Characteristic Properties of Metals and Nonmetals
Metals
Metalloids
• Have a shiny luster
and are usually silvery
in color
Notable exceptions: Cu Au
Nonmetals
• Do not have luster;
various colors
• Solids are usually brittle
• Solids are malleable
• Poor conductors of heat
and ductile
Hg is only liquid metal at RT and electricity
• Good conductors of heat
and electricity
• Metal oxides are ionic solids
and form basic solutions
• Form cations in solution
• Most nonmetal oxides are
molecular solids that form
acidic solutions.
• Tend to form anions or
oxyanions in solution
Elements with Color
Pale yellow
gas
Light green
gas
Dark orange
liquid/gas
Dark violet
crystals/gas
Properties of SOLID ionic compounds (M + NM)
• Poor conductors of electricity and heat.
• Generally high melting (more than 150°C).
• Crystalline, hard and brittle.
NaCl crystal lattice structure
Properties of ionic compounds (cont.)
• Molten ionic compounds form liquids that are an electrical
conductors.
• Ionic solids that are water soluble, dissolve to form
solutions that are electrical conductors.
• The solubility of ionic compounds depends upon the lattice
energy. The greater the lattice energy, the lower the
solubility.
+
-
Acid-Base Behavior of Oxides
• Most metal oxides form basic solutions
• Most nonmetal oxides form acidic solutions
• The acidity of the solution increases with oxidation
number of the central atom
pH: SO3(aq) < SO2(aq)
Hint: an easy way to evaluate the acidity is that pH 
as the # oxygen atoms 
• Amphoteric substances can act as either an acid or a base
Group Trends
Alkali metals: low IE  highly reactive, soft silvery metals
with a low density and low melting point.
While lithium reacts with oxygen to form lithium oxide
4 Li + O2  2 Li2O
the other alkali metals form peroxides. (peroxide = O22-)
2 Na + O2  Na2O2
Potassium, rubidium and cesium react with oxygen to form
superoxides (superoxide = O2)
K + O2  KO2
Alkaline earth metals
• Harder, more dense and have a higher melting point than
Group 1 metals.
• 1st IE is low, but not as low as Group 1 because disrupting
pseudo-noble gas configuration (s2)
• Increasing reactivity with increasing atomic number due
to increase in nuclear shielding.
• Beryllium will not react with water, but the other
alkaline earth metals will to form the metal hydroxide
and hydrogen gas.
Ca + 2 H2O  Ca(OH)2 + H2