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Chemical Periodicity
Mendeleev’s Periodic Table

In 1871, Dmitri
Mendeleev arranged
all of the then-known
elements by order of
increasing atomic
mass.
The Modern Periodic Table

In 1913, Henry Moseley
created a more accurate
positioning of the
elements in the modern
periodic table. He
arranged the elements in
order of increasing
atomic number.

He was killed in action at
Gallipoli in 1915 at the age of
28.
The Modern Periodic Table
The Periodic Table
Remember…
 Vertical columns are groups.
 Horizontal rows are periods.
and…
 Valence electrons – electrons in the
outermost (highest) principal energy level
of an atom
 Core electrons – inner electrons
Patterns in the electron
configurations of the elements
The noble gases:
 these are elements in which the
outermost s and p sublevels are
completely filled.
• He: 1s2
• Ne: 1s2 2s2 2p6
• Ar: 1s2 2s2 2p6 3s2 3p6
Patterns in the electron
configurations of the elements
The representative elements:
 these elements have the outermost s
and p sublevel only partially filled.
• Li: 1s2 2s1
• Na: 1s2 2s2 2p6 3s1
• K: 1s2 2s2 2p6 3s2 3p6 4s1
Patterns in the electron
configurations of the elements
The transition metals:
 In these metallic elements, the
outermost s and nearby d sublevel
contain electrons.
The inner transition metals:
 In these metallic elements, the
outermost s and nearby f sublevel
generally contain electrons.
Patterns in the electron
configurations of the elements
 Elements
with the same valence electron
arrangement show very similar chemical
behavior.
Electron Configurations and the
Periodic Table
Electron configurations for K through Kr
Practice writing electron
configurations from the periodic table.
Covalent atomic radius

It is not possible to directly measure the size of
an atom.
 Chemists determine the distance between the
nuclei in atoms of diatomic elements, then divide
this by 2 to find the radius of one atom.
I. Periodic Trend in
Atomic Size
Atomic radius tends to increase
down the group.
 Atomic radius tends to decrease
across the period.

Atomic radius
Atomic radius
Reason for the trends


Atomic size decreases across the period
because the effective nuclear charge increases
across the period.
Atomic size increases down the group because
the shielding effect increases down the group.
 (Shielding effect occurs as the core electrons
shield the attraction of the nucleus on the
outermost electrons.)
II. Periodic Trend in
Ionization Energy (IE)
 Ionization
Energy (IE) is the
energy required to remove an
electron from a gaseous atom.

Na(g)

+
Na
Ionization energy
(g) +
e

Ionization
Energy (IE) in
kilojoules/mole

First IE is the energy
required to remove
the first electron.
Second IE is the
energy required to
remove the second
electron (NOT both
first + second), etc.


Explain why there is
a large increase
between the first
and second IE for
Na.
 Explain why there is
a large increase
between the
second and third IE
for Mg.
 Predict where there
will be a large
increase in IE for
Al.
Ionization Energy (IE)



Ionization energy
tends to
decrease down a
group.
Ionization energy
tends to increase
across a period.
Explain why the trend is as
described (same reason
as for the size trend).
Explain why these elements do not follow
the trend...

Consider the electron
configurations as you
explain why the trend
is not always
consistent as you
move across the
period (such as
across period 2, from
Li to Ne).
 Why is the IE lower
than predicted by the
trend for B? …for O?
Ionization energy
IE in Metals

Metals tend to lose electrons to form
positive ions because their ionization
energies are low.
III. Periodic Trend in
Electron Affinity (EA)
 Electron
affinity (EA) is the energy
change that accompanies the addition
of an electron to a gaseous atom.

Br(g) +
e

Electron affinity
Br (g)
Electron Affinity

Electron Affinities for the Representative Elements (in kilojoules/mole)
Electron Affinity
Electron Affinity tends to decrease
down a group.
 Electron Affinity tends to increase
across a period.


Explain why the trend is as described
(same reason as for the size trend).
Electron affinity
IV. Periodic Trend in Ionic Size
 Cations
are smaller than their original
atoms. (Remember; atoms lose electrons
to form “+” cations.)
 Anions are larger than their original atoms.
(Atoms gain electrons to form “-” anions)
 Ionic
radius increases down the group (as
does atomic radius).
Ionic Size
V. Periodic Trend in Electronegativity
– the tendency for the
atom to attract electrons to itself in a
chemical bond
 Electronegativity


Increases from left to right across a
period
Decreases down a group
Pauli Electronegativity Values
Electronegativity
Electronegativity

Linus Pauling
established the
concept of
electronegativity and
developed a scale
that helped to predict
the nature of chemical
bonding.

In 1954, Pauling was
awarded the Nobel
Prize in Chemistry.
Summary of Periodic Trends
Review Group Names
 Remember:
 Group
1A = Alkali metals
 Group 2A = Alkaline earth metals
 Group 7 A = Halogens
 Group 0 = Noble Gases