Transcript Periodic Patterns

PERIODIC PATTERNS
Unit 3 – Periodic Table
Lesson Essential Question:
What patterns exist on the
periodic table?
INCREASES
METALLIC TREND
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 becoming
positive.
Calcium (Ca)
Loses 2 electrons becoming
Ca+2 and [Ar] Noble gas
Configuration.
(Octet Rule)
Nonmetals

Gain electrons becoming
negative.
Chlorine (Cl)
Gains one e- becoming Cl-1
and [Ar] Noble gas
configuration.
(Octet Rule)
IONS – How can I remember?
Metals
Nonmetals
This is Cat-ion - CATION
This is Ann ion - ANION
He is a “plussy” cat!
She is unhappy
and negative.
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.
At the nonmetals the
radius gets larger
because the ion has
gained electrons.
METALLIC ATOM AND ION COMPARISON
NONMETALLIC ATOM AND ION COMPARISON
Why do the Noble
Gases not have an
ionic Radius?
ATOM AND ION COMPARISON
Why does
Hydrogen not have
an ionic Radius?
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)
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