Periodic Trends
Periodic Trends
• What is a trend?
• A trend is the general direction in which
something tends to move.
Periodic Trends
• The elements on the
Periodic Table of
Elements show many
trends in their
physical and chemical
properties.
• Across the rows
(periods or series)
• Down the columns
(groups or families)
Atomic Radius
• ½ the distance between the nuclei of 2 like atoms in a
diatomic molecule
• The atoms of the 8 main groups are shown here.
Atomic Radius - Groups
• Atomic radius increases as
you move down a group
• Why?
• More electrons in more
Principal Energy Levels
• Atomic size increases
Atomic Radius - Periods
• Atomic radius decreases as you move across a
period
• Why?
• (-) electrons increase, but so do (+) protons !!!
• Increased (+) nuclear charge pulls the (-)
electrons closer to the nucleus
• Atomic size decreases
Atomic Radius - Periods
• The size trend in periods is less pronounced
than in groups because of the electron shielding
effect.
Shielding Effect
• Reduction in effective nuclear
charge on an electron that is
caused by the repulsive forces
of other electrons between it
and the nucleus
• In an atom with one electron,
that electron experiences the
full charge of the positive
nucleus. However, in an atom
with many electrons, the outer
electrons are simultaneously
attracted to the positive
nucleus and repelled by the
negatively charged electrons.
Atomic Radius - Graph
Atomic Radius
• Click on source to see a short video
• Source:
http://cwx.prenhall.com/petrucci/medialib/
media_portfolio/text_images/046_AtomicR
adii.MOV
Atomic Radius
Ionic Radius
• What are ions?
• Ions are charged atoms,
either + or • Cations are positive ions
• Cations form when atoms
lose electrons
• Anions are negative ions
• Anions form when atoms
gain electrons
Ionic Radius
Ionic Radius – Cations Group
• Cations are smaller than
their parent atoms.
• Why?
• By losing their valence
electrons, they lose their
entire valence shell
• Cations are formed by the
metals on the left side of the
Periodic Table
Ionic Radius – Cations Groups
• Ionic size increases as you
move down a group for the
same reason atomic size
increases
• Number of principal energy
levels increases
Ionic Radius – Anions Groups
• Anions are larger than their
parent ions
• Why?
• When extra (-) electrons are
added, extra (+) protons are
NOT added to the nucleus
• Effective nuclear attraction is
less for the increases number
of electrons
Ionic Radius – Anions Group
• Ionic size increases as you
move down a group for the
same reason atomic size
increases
• Number of principal energy
levels increases
• Cations are formed by
nonmetals on the right side of
the Periodic Table
Ionic Radius - Periods
• Just like their parent atoms…
• Cations get smaller as you move from left to
right
• Anions get smaller as you move from left to right
• Increased (+) nuclear charge pulls the (-)
electrons closer to the nucleus
Ionic Radius
Ionization Energy
• Energy is needed to remove an electron from an
atom
• The energy needed to overcome the attraction
of the nuclear charge and remove an electron
(from a gaseous atom) is called the Ionization
Energy
1st Ionization Energy
• The energy needed to remove the 1st electron
from an atom is the 1st Ionization Energy
Factors Affecting
Ionization Energy
• Atomic Radius
• Smaller atoms hang on
to valence electrons
more tightly, and so
have higher ionization
energy
Factors Affecting
Ionization Energy
• Charge
• The higher the positive
charge becomes, the
harder it is to pull away
additional electrons
• Second ionization
energy is always higher
than the first
Factors Affecting
Ionization Energy
• Orbital Type
• It's easier to remove
electrons from p orbitals
than from s orbitals,
which are “deeper”
Factors Affecting
Ionization Energy
• Electron Pairing
• Within a subshell,
paired electrons are
easier to remove than
unpaired ones
• Reason: repulsion
between electrons in
the same orbital is
higher than repulsion
between electrons in
different orbitals
Factors Affecting
Ionization Energy
• Electron Pairing – Example
• On the basis of gross periodic
trends, one might expect O to
have a higher ionization
energy than N. However, the
ionization energy of N is 1402
kJ/mol and the ionization
energy of O is only 1314
kJ/mol.
• Taking away an electron from
O is much easier, because the
O contains a paired electron in
its valence shell which is
repelled by its partner.
1st Ionization Energy
2nd, 3rd, 4th Ionization Energies, etc.
• Subsequent electrons require more energy
to remove than the first electron
• How much more energy is needed
depends on what energy levels and
orbitals the electrons are in
2nd, 3rd, 4th Ionization Energies, etc.
•
Source: http://chemed.chem.purdue.edu/genchem/topicreview/bp/ch7/ie_ea.html
Ionization Energy
• Click on source to see a short video.
• Source:
http://cwx.prenhall.com/petrucci/medialib/
media_portfolio/text_images/047_Ionizatio
nEner.MOV
Electronegativity
• Ability of an atom to attract electrons toward
itself in a chemical bond
Electronegativity
• The difference between the electronegativities of
two atoms will determine what kind of bond they
form
• Linus Pauling used an element's ionization energy
and electron affinity to predict how it will behave in a
bond.
• The more energy it takes to pull off the outer
electron of an atom, the less likely it is to allow
another atom to take those electrons. The more
energy the atom releases when it gains an electron,
the more likely it is to take electrons from another
atom in bonding. These two energies were used to
compute a numerical score.
Electronegativity - Periods
• Electronegativity increases going left to right across the
periodic table.
• Fluorine's high nuclear charge coupled with its small size
make it hold onto bonding electrons more tightly than
any other element. Lithium has a lower nuclear charge
and is actually larger than fluorine. Its valence electron is
not tightly held and it tends to surrender it in chemical
bonds.
Li
Be
B
C
N
O
F
1.0
1.5
2.0
2.5
3.0
3.5
4.0
Electronegativity - Groups
• Electronegativity decreases
going down a group
• The bonding electrons are
increasingly distant from the
attraction of the nucleus
H
2.1
Li
1.0
Na
0.9
K
0.8
Rb
0.8
Cs
0.7
Fr
0.7
Electronegativity
Electronegativity
Electron Affinity
• The electron affinity is a measure of the energy
absorbed when an electron is added to a neutral
atom to form a negative ion.
• Most elements have a negative electron affinity.
This means they do not require energy to gain
an electron; instead, they release energy.
• Atoms more attracted to extra electrons have a
more negative electron affinity.
• The more negative the value, the more stable
the ion is.
Electron Affinity
• Click on source to see a short video.
• Source:
http://cwx.prenhall.com/petrucci/medialib/
media_portfolio/text_images/049_ElectrAff
inity.MOV
Electron Affinity
• Electron affinity is essentially the opposite of the
ionization energy.
Electron Affinity - Trends
• Click on source to see a short video.
• Source:
http://cwx.prenhall.com/petrucci/medialib/
media_portfolio/text_images/050_PeriodEl
ectron.MOV
Image Sources
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http://www.Chem4kids.com
http://images.encarta.msn.com
http://antoine.frostburg.edu
http://www.webelements.com
http://cwx.prenhall.com
http://www.800mainstreet.com
http://intro.chem.okstate.edu/1215
Credits
• PowerPoint: Adela J. Dziekanowski