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PERIODIC TRENDS
Unit 3 – Periodic Table
Lesson Essential Question:
What patterns exist on the
periodic table?
The Basics
The Basics
Metals, Nonmetals, Metalloids



There is a zig-zag or
staircase line that
divides the table.
Metals are on the left
of the line, in blue.
Nonmetals are on the
right of the line, in
orange.
METALS, Nonmetals,



Elements that border
the stair case, shown in
purple are the
metalloids or semimetals.
There is one important
exception.
Aluminum is more
metallic than not.
METALS, Nonmetals,



Metals are lustrous (shiny),
malleable, ductile, and are
good conductors of heat
and electricity.
They are mostly solids at
room temp.
Get oxidized (lose
electrons), giving a positive
oxidation state




Nonmetals are the
opposite.
They are dull, brittle,
nonconductors (insulators).
Some are solid, but many
are gases, and Bromine is a
liquid.
Get reduced (gain
electrons), giving a negative
oxidation state
METALS, Nonmetals,




Metalloids, aka semi-metals are just that.
They have characteristics of both metals and
nonmetals.
They are shiny but brittle.
And they are semiconductors.
INCREASES
METALLIC TREND


Metallic characteristics are greatest to the left of the
periodic table due to fact that the attraction between
valence electron and the nucleus is weaker, enabling an
easier loss of electrons.
Metallic characteristics increases as you move down a
group because the atomic size is increasing so valence
electrons are farther from the nucleus and easier to
remove.
ATOMIC RADIUS


Radius is the distance from the
center of the nucleus to the
“edge” of the electron cloud.
Atomic radii are usually
measured in picometers (pm) or
angstroms (Å). An angstrom is
1 x 10-10 m.
ATOMIC RADIUS

BROMINE = Br2
Since a cloud’s edge is difficult to
define, scientists use define covalent
radius, or half the distance between
the nuclei of 2 bonded atoms.
2.86 Å
1.43 Å
1.43 Å
ATOMIC RADII TRENDS
WHY?
DOWN A FAMILY OR GROUP

INCREASES
As you go down a
family the number of
energy levels
increases making the
radius larger.
ATOMIC RADII TRENDS
WHY?
ACROSS A PERIOD
DECREASES

As you go across a
period the number of
protons increases,
(nuclear charge)
pulling the electrons in
tighter making the
radius smaller.
IONS - remember
Metals

Lose electrons (oxidize)
becoming positive
(cations).
Calcium (Ca)
Loses 2 electrons becoming
Ca+2 and [Ar] Noble gas
Configuration.
(Octet Rule)
Nonmetals

Gain electrons (reduce)
becoming negative
(anions).
Chlorine (Cl)
Gains one e- becoming Cl-1
and [Ar] Noble gas
configuration.
(Octet Rule)
IONIC RADII TRENDS
WHY?
DOWN A FAMILY OR GROUP

INCREASES
As you go down a
family the number of
electron shells
increases making the
radius larger.
IONIC RADII TRENDS
WHY?
ACROSS A PERIOD
DECREASES then INCREASE


For the metals the
nuclear charge is greater
than then number of
electrons pulling them in
tighter making the radius
smaller.
For the nonmetals the
radius gets larger
because the ion has
gained electrons.
Shielding Effect
As more electrons are added to atoms, the
inner layers of electrons shield the outer
electrons from the nucleus.
The effective nuclear charge on those
outer electrons is less, and so the outer
electrons are less tightly held
Example of Shielding Effect
Ionization Energy
The energy required to remove an
electron from an atom.
(measured in kilojoules, kJ)
Trends in Ionization Energy
Why?
• Closer to nucleus
(more +)
• Electrons less
likely to be
removed
• Requires more
energy to form
ion
• Less shielding
INCREASES
IONIZATION TREND
IONIZATION ENERGY
The larger the atom is, the easier its electrons
are to remove. (Why?)
Ionization energy and atomic radius are
inversely proportional.
Ionization energy is always endothermic, that is
energy is added to the atom to remove the
electron.
IONIZATION TREND
NCREASES
Why?
• Elements in alkali
metals have 1
valence electron so
what to remove that
electron, they
therefore take the
least amount of
energy to remove an
electron
INCREASES
Electronegativity
is a measure of the tendency of an
atom to attract a bonding pair of
electrons.
Why?
• Closer to
nucleus
(more +) so
electrons are
more
attracted
INCREASES
Electronegativity
http://www.thecatalyst.org/electabl.html
Electronegativity
Why?
• Elements in halogens
only need 1 more
electron to have a
full valence shell so
are MOST likely to
attract electrons. As
you move to left
elements are more
likely to LOSE
electrons.
INCREASES
http://www.thecatalyst.org/electabl.html
In Summary….
Electronegativity
Electronegativity