Periodic Properties of
the Elements
The Periodic Table
The modern periodic table was developed in
1872 by Dmitri Mendeleev (1834-1907). A
similar table was also developed independently
by Julius Meyer (1830-1895).
The table groups elements with similar
properties (both physical and chemical) in
vertical columns. As a result, certain properties
recur periodically.
The Periodic Table
Mendeleev left empty spaces in his table for
elements that hadn’t yet been discovered. Based
on the principle of recurring properties, he was
able to predict the density, atomic mass, melting
or boiling points and formulas of compounds
for several “missing” elements.
The Periodic Table
The Periodic Table
metal/non-metal
line
The Periodic Table
The periodic table is based on observations
of chemical and physical behavior of the
elements. It was developed before the discovery
of subatomic particles or knowledge of the
structure of atoms.
The basis of the periodic table can be
explained by quantum theory and the electronic
structure of atoms.
Quantum Numbers
In addition to n, the principal quantum
number, there are three additional quantum
numbers which describe the type of orbital ( l ) ,
the spatial orientation of the orbital (ml ) , and
the spin of the electron (ms ) .
Quantum Numbers
The magnetic quantum number (ml )
specifies the spatial orientation of the orbital.
An example is to distinguish between the px, py
or pz orbitals.
Electron Spin
Each orbital, regardless of type, can contain
zero, one or two electrons. If two electrons
occupy the same orbital, they must spin in
opposite directions.
The spin is quantized, and can be expressed
using quantum numbers, or simply specifying
the spin as up or down or clockwise and
counter-clockwise.
The Pauli Exclusion Principle
Quantum mechanics dictates that no two
electrons in an atom can have the same four
quantum numbers. Another way of stating the
Pauli Exclusion Principle is that if electrons occupy
the same orbital, they must have opposite spins.
Multi-electron Atoms
Orbitals of any
type can be empty, or
have 1 or two
electrons.
Experimental data
indicate that if two
electrons are in the
same orbital, they will
spin in opposite
directions.
Energy Levels
In any atom or ion
with only 1 electron,
the principal quantum
number, n,
determines the energy
of the electron. For
n=2, the 2s and 2p
orbitals all have the
same energy.
Energy Levels
Likewise, the 3s,
3p and 3d orbitals are
all degenerate, with
the same energy.
Energy Levels
In a multi-electron atom, there is interaction
between electrons. As a result of this
interaction, the various subshells of a principal
quantum level will vary in energy.
Energy Levels
Energy Levels
Energy Levels
The energy
diagram for the first
three quantum levels
shows the splitting
of energies.
Energy Levels
For a given value
of n, the energies of
the subshells is as
follows:
ns<np<nd<nf
Energy Levels
The subshells
have different
energies due to the
penetrating ability for
each type of orbital.
Electrons in a 2s
orbital can get nearer
to the nucleus than
those in a 2p orbital.
Energy Levels
The electrons in
the 3s orbital (top
diagram) have higher
probability to be
found near the
nucleus, and thus
greater penetrating
ability than those in
3p or 3d orbitals.
Multi-electron Atoms
Electron configurations are a way of noting
which subshells of an atom contain electrons.
Although much of the periodic table was
developed before the concept of electron
configurations, it turns out that the position of
an element on the periodic table is directly
related to its electron configuration.
Multi-electron Atoms
Electron Configurations



Write the complete electron configurations for
nitrogen and zinc.
How many unpaired electrons does each atom
have?
What is the short hand notation for each
element.
Hund’s Rule
When electrons occupy degenerate orbitals,
they occupy separate orbitals with parallel spins.
This is the lowest energy, or ground state,
configuration.
Multi-electron Atoms
The electron configurations for Cr and Cu differ
from that expected based on their positions in
the periodic table.
Multi-electron Atoms
Electron configurations also get less
predictable for the elements near the bottom of
the periodic table.
With many quantum levels (n) occupied, the energy
levels overlap and the lowest energy arrangement
becomes more difficult to predict.
Periodic Trends
Many of the properties of atoms show clear
trends in going across a period (from left to
right) or down a group.
In going across a period, each atom gains a
proton in the nucleus as well as a valence
electron.
Periodic Trends
The increase of positive charge in the
nucleus isn’t completely cancelled out by the
addition of the electron.
Electrons added to the valence shell don’t
shield each other very much. As a result, in
going across a period, the effective nuclear charge
(Zeff) increases.
Effective Nuclear Charge
The effective nuclear charge (Zeff) equals the
atomic number (Z) minus the shielding factor
(σ).
Zeff= Z-σ
Effective Nuclear Charge
Zeff= Z-σ
Effective Nuclear Charge
Electrons in the
valence shell are
partially shielded
from the nucleus by
core electrons.
Effective Nuclear Charge
Electrons in p or d
orbitals don’t get too
close to the nucleus, so
they are less shielding
than electrons in s
orbitals. As a result,
effective nuclear charge
increases across a period.
Periodic Trends
Periodic Trends
In going down a group or family, a full
quantum level of electrons, along with an equal
number of protons, is added.
As n increases, the valence electrons are, on
average, farther from the nucleus, and
experience less nuclear pull due to the shielding
by the “core” electrons. As a result, Zeff
decreases slightly going down a group.
Trends- Atomic Radii
Atomic radii are obtained in a variety of ways:
1. For metallic elements, the radius is half the
internuclear distance in the crystal, which is
obtained from X-ray data.
2. For diatomic molecules, the radius is half the
bond length.
3. For other elements, estimates of the radii are
made.
Trends- Atomic Radii
Atomic radii follow trends directly related to
the effective nuclear charge. As Zeff increases
across a period, the electrons are pulled closer to
the nucleus, and atomic radii decrease.
As Zeff decreases down a group, the valence
electrons experience less nuclear attraction, and
the radius increases.
TrendsAtomic Radii
Atomic size roughly
halves across a
period, and doubles
going down a group.
Electron Configurations of Ions
The atoms of the main group elements
(groups IA-VIIA) will form ions by losing or
gaining electrons. The resulting ion will have
the same electron configuration as a noble gas
(group VIIIA). These configurations are usually
very stable.
Electron Configurations of Ions

Atoms or ions with the same electron
configuration (or number of electrons) are called
isoelectronic.
For example, Na+, Mg2+, Ne, F-, and O2- are
isoelectronic. The size will decrease with
increasing positive charge.
O2- > F- >Ne> Na+> Mg2+
Electron Configurations of Ions
When atoms lose electrons, the electrons are
always removed from the highest quantum level
first.
For the first row of transition metals, this
means that the electrons in 4s subshell are lost
before the 3d subshell.
Fe: [Ar]4s23d6
Fe2+: [Ar] 3d6 or [Ar]4s03d6
Common Ionic Charges
The charges of ions of elements in groups
1A-7A (the main groups) are usually predictable.
Group 1A metals form +1 ions, group 2A
metals form +2 ions, etc.
The non-metals of group 5A have a -3
charge, those of group 6A have a -2 charge, and
the halogens form ions with a -1 charge.
Typical Ionic Charges
Trends – Ionic Size
Cations are always smaller than the neutral
atom. The loss of one or more electrons
significantly increases Zeff, resulting in the
valence electrons being pulled closer to the
nucleus.
Ionic Size - Cations
Within a group, assuming the same ionic charge,
the size of the ion increases going down the
group, due to more core electrons shielding the
nucleus as n increases.
Trends – Ionic Size
Across period,
the cations get
more positive,
and as a result,
considerably
smaller.
Trends – Ionic Size
Anions are always larger in size than the
neutral atom. The addition of one or more
electrons results in greater electronelectron repulsion, which causes the
valence electrons to “spread out” a bit.
Size of Anions
Anions are always
larger than the neutral
atom.
Size of Anions
Within a group, assuming the same
ionic charge, the size of the ion increases
going down the group, due to more core
electrons shielding the nucleus as n
increases.
Trends – Ionic Size
Trends – Ionization Energy
Ionization energy is the energy required to remove
an electron from a mole of gaseous atoms or
ions.
X(g) + energy  X+(g) + eElements can lose more than one electron,
so there are 1st, 2nd, 3rd, etc., ionization energies.
Ionization Energy
It always requires energy to remove an
electron from a neutral atom.
As more electrons are removed and the ion
becomes positively charged, it requires
increasingly greater energy to remove electrons.
Trends – Ionization Energy
Ionization energy is a measure of how tightly
the electrons in the highest occupied orbitals are
held by the nucleus. As a result, it is directly
related to the effective nuclear charge.
Ionization energy increases going across a
period, and decreases going down a group.
Trends – Ionization Energy
Trends – Ionization Energy
Ionization Energy
Trends – Electron Affinity

Electron Affinity is the energy change when an
electron is added to a mole of gaseous atoms.
X(g) + e-  X-(g)
∆E = electron affinity
A negative value for the electron affinity
indicates that the process releases energy, and
that the anion is easily formed.
Trends – Electron Affinity
There is less of a predictable trend in
electron affinities. In going across a period
(ignoring the noble gases), the electron affinity
should become more negative. Although this is
observed, there are many inconsistencies.
Trends – Electron Affinity
Trends – Electron Affinity
Trends- Electron Affinity
In going down a group, the electron affinity
should become less negative. Although this
trend is observed, there is only a slight change in
electron affinities within a group. There may
also be inconsistencies in the general trend.
Trends – Electron Affinity
Trends- Electron Affinity
The electron affinity
of fluorine is less
negative than
expected. This may
be due to additional
electron-electron
repulsion when an
electron is added to
such a small atom.
Metallic Character
Metallic Character
Across a period, metallic behavior decreases.
Non-metals are often crumbly solids, liquids or
gases at room temperature.
Metallic Character
Metallic behavior
increases going down a
group.
Group IA – the Alkali Metals
In discussing the chemistry, preparation and
properties of the group IA elements, it is
important to remember that hydrogen is not a
group IA metal. It’s properties and reactivity
would place it within group 7A (diatomic nonmetals), rather than group IA.
Group 1A Metals
The group 1A metals are soft shiny metals
with fairly low densities (Li, Na and K are less
dense than water) and low melting points.
Sodium melts at 98oC, and cesium melts at 29oC.
The softness, low density and low melting
points are the result of weaker metallic bonding
due to only one valence electron in this group.
Group 1A Metals - Production
Due to the high reactivity with oxygen and
water, all of the metals are found in nature in
ionic form (M1+).
The pure metal must be produced in an
oxygen and water-free environment. Typically,
an electrical current is passed through the
melted chloride salt. The metal and the chlorine
gas are collected separately.
Reactivity Trends
The chemical behavior of the group IA
metals illustrates periodic trends. As the valence
electron occupies a higher quantum level, it
experiences less nuclear attraction, and is more
easily removed.
Group 1A Metals + Water
The reaction with water forms hydrogen gas
and the aqueous metal hydroxide. The reaction
is so vigorous, that the hydrogen may ignite.
2 M(s) + 2 H2O(l)  H2(g) + 2 MOH(aq)
Metallic Character
The group IA metals react with water to
produce hydrogen and the metal hydroxide.
Metallic behavior increases going down a group.