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Trends on the Periodic
Table
Periodic
Trends
 The
arrangement of the periodic table
reveals trends or general tendancies in the
physical and chemical properties of the
elements.
 Periodic Law states that the physical and
chemical properties of the elements are
periodic functions of their atomic numbers.
 Periodic means REPEATING.
What is a Trend?
A
trend is a predictable change in a
particular direction.
 For example, as you move down Group 1,
reactivity increases for each element.
 Knowing these various trends allows you
to make logical predictions about the
elements chemical and physical
properties.
Periodic Trends in Atomic
Radii
 The
exact size of an atom is difficult to
determine.
 The size is determined by the volume
occupied by the electrons surrounding the
nucleus.
 We model atoms as clouds, but in reality
there is no clear cut boundaries for the
edges of atoms.
Radii of Atoms

Are usually determined for atoms that are
chemically bonded or close together in a solid
state.
 Bond radius is half the distance from the
center of one like atom to the center of
another like atom bonded chemically
together.
 Bond radius is used to determine the size of
an atom.
Bond Radius
Equation
Another Way to Determine
Size of an Atom
 Van
der Waals is half the distance
between the nuclei in adjacent nonbonded molecules. Sometimes this radius
is used to denote the size of atoms, but is
available only for the main group
elements.
Van der Waals Forces
The Problem with Bond
Radius is…
 That
bonds between different atoms
create different size radii.
 For
example, bonds between Tin atoms in
metallic Tin are different than Tin Chloride,
SnCl4. But even with this limitation, it is still
a useful way to measure the size of atoms.
Atomic Radii Trends
 Atomic
 There
Radii increases down a group.
is a general trend toward LARGER
RADII as you proceed down a group.
Atomic Radii on the Periodic
Table
Reason for Larger Radii

1. The addition of one more principle
energy level from one period to the next.
2. The phenomenon of shielding.
Shielding-is the reduction of the attractive
force between a positively charged nucleus
and its outermost electrons due to the
cancellation of some of the positive charge by
the negative charge of the inner electrons.
Shielding
 Electrons
in the inner energy levels
“shield” the outer electrons from the
attractive force of the positive nucleus.
Hence, these inner electrons reduce the
positive force of the nucleus so the outer
electrons can move even further out from
the nucleus.
Shielding
Conclusion
 So
the outermost electrons (valence
electrons) are not subject to a full charge
of the nucleus, they, therefore are not held
as close to the nucleus.
Nuclear Charge

As you move down a group, nuclear charge
(positive charge) increases because more
protons are present.
 Also, at the same time, more shielding is present
do to the presence of more electrons.
 This creates a canceling effect regardless of
energ level.
Example
 Sodium,
Na and Cesium, Cs have about
the same net nuclear charge even though
Cesium has more protons in the nucleus.
Cesium also has more electrons to create
a greater shielding effect.
Atomic Radii Decrease
across a Period

Reason for Smaller Radii
 1. It is caused by increasing positive
charge of the nucleus-increased nuclear
charge. You gain one more proton than
the element before.
 2. You have added electrons however they
are going in the same energy level. So
therefore, shielding remains constant.
There is no cancellation of the positive
charge like there is when you move down
a group.
Conclusion
 So
therefore, more protons are added
without a canceling effect from the
electrons. Since no more energy levels
are being added to distance the electrons
from the positive nucleus, the electrons in
the outer energy level get pulled in closer.
This Trend Levels Off
 This
trend gets less pronounced when
there are many electrons between the
nucleus and the outermost energy level.
When many electrons are pulled in closer
to the nucleus, the electrons get too close
to each other and start to repel each other.
 See
Period Six
Conclusion
 Therefore,
there is a point where the
electrons will come no closer to each other
and they will level off in size.
Ionization Energy & Electron
Affinity
 Ionization
Energy is the energy you use
to remove an outer electron (valence
electron) from an atom.
 Equation:
 A + IE -> A+ + e Atom + ionization energy yields cation +
electron
Ionization Energy
Trends in Ionization Energy

Ionization Energy increases across a Period.
 Reason:
 1. More protons are added to the nucleus
as you move across a period, but they
are added to the same energy level.
 2. The nuclear charge (positive charge)
increases holding the electrons closer
because shielding is constant.
Conclusion
 Therefore,
you need MORE IONIZATION
ENERGY to pull electrons off the atom!
Ionization energy Decreases
down a Group

Reason:
 1. The number of energy levels increase
between the nucleus and outermost
electrons.
 2. The outer shell electrons are further
from the nucleus.
 3. The net nuclear charge (positive charge) is
the same all the way down the group due to
increased shielding.
Conclusion
 Therefore,
less energy is required to
remove an electron. The attractive forces
between the nucleus and the outermost
electrons decrease and the energy
necessary to remove them also
decreases.
Periodic Trends in Electron
Affinity
 Electron
Affinity is the energy emitted
(given off) upon the addition of an electron
to an atom or group of atoms in the gas
phase.
 Because inner electrons do not shield the
positive nucleus one hundred percent, an
approaching electron may experience a
net pull and be drawn into a vacant orbital
in an atom.
Electron Affinity
Equation
+ e- -> A- + energy
 Atom + electron yields anion + energy
given off
A
Electron Affinity Increases
going across a Period
 Reason:
 1.
Shielding is constant because the
energy level is the same.
 2. Nuclear charge (positive charge)
increases due to the addition of one
more proton.
Conclusion
 Therefore,
the atom’s attraction for extra
electrons increases left to right.
Electron Affinity Decreases
down a Group
 Reason:
 1.
Both shielding and nuclear charge
increase, but each cancel the other
making net nuclear charge constant.
 2. This constant nuclear charge allows
electrons to move further out from the
nucleus.
Conclusion
 The
atom’s attraction for extra electrons
decreases!
Electronegativity of an Atom
 Is
the tendancy of an atom to attract
electrons to itself when it is chemically
combined with another element.
Electronegativity is
expressed…
 In
terms of a relative scale with arbitrarily
selected standard units.
 Fluorine is the most electronegative
element and is assigned 4.0 as the largest
unit.
 Francium is the least electronegative and
is assigned a value of 0.7.
Trends in Electronegativity
 Electronegativity
increases across a
period.
 Reason:
1. Nuclear charge increases as you move
across a period. This is due to the fact
that shielding remains constant.
Electronegativity Chart
Conclusion
 Therefore,
the atom attracts electrons to
itself the greater the nuclear charge.
Electronegativity Decreases
down a Group

Reason:
1. Because increased shielding is present,
nuclear charge remains relatively constant
down a group.
Conclusions
 Therefore,
atoms further away from the
nucleus are less likely to be attracted by
the nuclear charge and the atom’s
attraction for more electrons decreases!
 Special Note: Chemists determine
electronegativity values by measuring the
polarities of the bonds between various
atoms.