Transcript Periodic Trends

Periodic Trends
Elemental Properties and Patterns
The Periodic Law
• Dmitri Mendeleev - first to
publish an organized
periodic table of known
elements.
The Periodic Law
• Mendeleev even predicted the properties of
undiscovered elements.
https://www.youtube.com/watch?v=fPnwBITSmgU
The Periodic Law
• Says that:
“When arranged by increasing atomic
number, the chemical elements display a
regular and repeating pattern of chemical
and physical properties.”
So…
• Atoms with similar properties appear in
groups or families (vertical columns) on the
periodic table.
• They are similar because they all have the
same number of valence (outer shell)
electrons, which determines their chemical
behavior.
Valence Electrons
• For Groups 1, 2, 13, 14, 15, 16, 17, 18: The
digit farthest to the right is the number of
valence electrons.
Example:
• Groups 3-12?
• Many have 1-2, but figuring it out is more
complicated. 
A Different Type of Grouping
Besides the 4 blocks of the table, there is
another way of classifying element:
• Metals
• Nonmetals
• Metalloids
Metals, Nonmetals, Metalloids
Metals, Nonmetals, Metalloids
• Most elements that
border the stair
case are metalloids.
(“metal-like”)
Have properties of
both metals and
nonmetals!
Metals
• Lustrous (shiny),
malleable, ductile, and
are good conductors of
heat and electricity.
• Mostly hard – not
brittle
• They are mostly solids
at room temp.
Nonmetals
• The opposite!
• They are dull, brittle,
nonconductors
(insulators).
• Some are solid, but
many are gases, and a
few are liquid.
Metalloids
• Have properties both metals
& nonmetals!
• Shiny but brittle.
• Semiconductors.
The Octet Rule
• The “goal” of most atoms (except H, Li and Be)
is to have an octet (group of 8 electrons) in their
valence energy level.
• Metals generally give electrons, nonmetals take
them from other atoms.
Remember…
• Atoms that have gained or lost electrons are
called ions.
Ions
• When an atom gains an electron, it becomes
negatively charged and is called an (anion.)
• When an atom loses an electron, it becomes
positively charged (cation). Think of the “t”
like a “+” sign.
Cation Formation
Effective nuclear
charge on remaining
electrons increases.
Na atom
1 valence electron
11p+
Valence elost in ion
formation
Result: a smaller
sodium cation, Na+
Remaining e- are
pulled in closer to
the nucleus. Ionic
size decreases.
Anion Formation
Chlorine
atom with 7
valence e17p+
One e- is added
to the outer
shell.
Effective nuclear charge is
reduced and the e- cloud
expands.
A chloride ion is
produced. It is
larger than the
original atom.
Periodic Trends
• There are several important atomic
characteristics that show predictable trends
that you should know.
• The first and most important is atomic
radius.
Atomic Radius
• Is the distance from the center of the
nucleus to the “edge” of the electron cloud.
• Since that is difficult to define, scientists use
covalent radius, half the distance between the
nuclei of 2 bonded atoms.
• Atomic radii are usually measured in
picometers (pm) or angstroms (Å). An
angstrom is 1 x 10-10 m.
Covalent Radius
• Two Br atoms bonded together are 2.86
angstroms apart. So, the radius of each
atom is 1.43 Å.
2.86 Å
1.43 Å
1.43 Å
Atomic Radius
The trend for atomic radius in a GROUP is to
go from smaller at the top to larger at the
bottom.
• Why?
• With each step down the family, we add energy
levels to the electron cloud, making the atoms
larger.
Atomic Radius
• As you move ACROSS a PERIOD, atoms
are smaller because of an increased
attraction between nucleus and electron
cloud.
• The increased attraction pulls the cloud in,
making atoms smaller as we move from left to
right.
Effective Nuclear Charge
• What keeps electrons from simply flying off
into space?
• Effective nuclear charge is the pull that an
electron “feels” from the nucleus.
• The closer an electron is to the nucleus, the
more pull it feels.
• As effective nuclear charge increases, the
electron cloud is pulled in tighter.
Ionization
• If an electron is given enough energy to
overcome its attraction to the nucleus, it can
leave the atom completely.
• The atom has been “ionized” or charged.
• Number of protons ≠ number of electrons.
Ionization Energy
• Energy required to remove an electron from an
atom is ionization energy.
(measured in kilojoules, kJ)
• The larger the atom is, the easier its electrons
are to remove.
• Ionization energy and atomic radius are
inversely proportional.
• As you move down, it decreases!
• As you move across, it increases!
Electronegativity
• Electronegativity is a measure of an atom’s
attraction for another atom’s electrons.
• Metals are usually electron givers and have
low electronegativity.
• Nonmetals are electron takers and have high
electronegativity.
• What about the noble gases? Make a
prediction to your table partners.
Ionic Radius
• Cations are always smaller than the original
atom.
• Conversely, anions are always larger than
the original atom.
Review Video for Atomic Trends
• https://www.youtube.com/watch?v=0tP6bV
89log
Link to bonding animations
• http://bcs.whfreeman.com/thelifewire/conte
nt/chp02/02020.html
• http://www.youtube.com/watch?v=cZy8tGF
V8QE&list=TL6TJ_jx1X7SKnBP20kAt90e
SdTpgQOR8H