Chapter 6:
Periodic Table and Trends
Dmitri Mendeleev
• First Periodic Table
• Based on increasing
Atomic mass and repeating
properties of elements
• Had spaces for “missing”
elements that he predicted
Henry G. J. Moseley
• Discovered that
elements’ properties
were more closely
associated with atomic
number
• Modern periodic table is
based on this discovery
Periodic Law
• When elements are arranged in order of
increasing atomic number, their physical and
chemical properties show a periodic pattern
Reading the Periodic Table
Inner or
Groups
Transition
Nitrogen
Alkaline
Oxygen
Carbon
Boron
Noble
Alkali
Halogens
Periods
Metals
Family
Gases
Transition
Earth
Families
Family
Metals
Metals
Metals
Metals
• Left of the stairstep line
• Majority of the
elements
• Tend to LOSE
electrons
• Most reactive in
the s-block
Properties of Metals
•
•
•
•
•
•
Shiny luster
Good conductors of heat
Good conductors of electricity
Most are solids at room temperature
Malleable
Ductile
gold
lead
copper
nickel
Nonmetals
• Right of Stair-Step line
• Tend to GAIN electrons
• Most reactive group is
halogens
• Least reactive group is
Noble Gases
Properties of Nonmetals
•
•
•
•
•
Dull Luster
Poor conductors of heat
Poor conductors of electricity
Brittle
Many are gases at room temperature
CARBON
BROMINE
SULFUR
ARGON
Metalloids
• Along Stair-step line
• Have properties of
metals AND nonmetals
• Many are used in
transistors, found in
electronics
Silicon
Antimony
Boron
Alkali Metals
• Group 1 (Except H)
• All have only 1 valence electron
• Most reactive metals; never found in
elemental form in nature
• Soft and shiny
• Relatively low melting points
Alkaline Earth Metals
• Group 2
• All have 2 valence electrons
• Second most reactive metals; never
found in pure state in nature
• Harder, denser, and stronger than
alkali metals
• Have higher melting points than
alkali metals
Transition Metals
• Groups 3-12
• All have 1 or 2
valence electrons
(in s sublevels)
• Do not fit into any
other group or
family
• Have many
irregularities in
their electron
configurations
Boron Family
•
•
•
•
Group 3A
Have 3 valence electrons
Boron is a metalloid
All others are metals
Carbon Family
•
•
•
•
•
Group 4A
All have 4 valence electrons
Carbon is a nonmetal
Si and Ge are metalloids
Sn and Pb are metals
Nitrogen Family
• Group 5A
• All have 5 valence electrons (s and p
sublevels)
• N and P are nonmetals
• As and Sb are metalloids
• Bi is a metal
Oxygen Family
• Group 6A
• All have 6 valence electrons
• Oxygen, Sulfur, and Selenium are
nonmetals
• Tellurium and Polonium are
metalloids
Halogens
•
•
•
•
•
Means “salt former”
Group 7A
All have 7 valence electrons
Most reactive nonmetals
All are nonmetals
Noble Gases
• Group 8A
• 8 Valence electrons makes a full
electron shell: s2 p6
• Complete, stable electron
configuration (Complete outer energy
level)
• Least reactive of all elements
Rare Earth Elements
(Inner Transition metals)
•
•
•
•
•
Found in 2 rows at bottom of periodic table
Lanthanide series follows La
Actinide series follows Ac
Little variation in properties
Actinides are radioactive; only first three and
Pu are found in nature
Summary
• Groups: Up and Down
• Periods: Across
• Main Group Elements are in groups 1-2, 1318
• Elements along the stair step line are
metalloids
• Elements to the left of the stair step line are
metals
• Elements to the right of the stair step line are
nonmetals
Octet Rule
“Noble Gas Envy”
• Atoms tend to gain, lose, or share electrons
in order to acquire a full set of valence
electrons (typically 8)
Periodicity
• Properties of the elements change in a
predictable way as you move through the
periodic table
• These properties include
• Atomic Radius
• Ionization energy
• Electronegativity
Atomic Radius
• Distance from nucleus to outermost valence
electrons
Atomic Radius
• Increases down groups
• Decreases from left to right
Ionization Energy
• The energy needed to remove 1 of an atom’s
electrons
• Decreases as you move down a group
• Increases from left to right, across a period
• Successive ionization energies increase for
every electron removed
1st ionization energy
Electronegativity
• Reflects an atom’s ability to attract electrons
in a chemical bond
• Related to its ionization energy and electron
affinity
• Increases from left to right, across a period
• Decreases from top to bottom, down a group
Shielding
• Shielding electrons are electrons located
between the nucleus and the valence
electrons
• For example:
• Chlorine has the following electron configuration:
1s2 2s2 2p6 3s2 3p5
• The shielding electrons would be 1s2 2s2 2p6
• The valence electrons would be 3s2 3p5
• Then we have 10 shielding electrons and 7 valence
electrons, right?
Shielding
• You try one
• Try Sodium (Na)… Remember, first you have to know
the electron configuration
• Did you get:
•
•
•
•
Electron configuration: 1s2 2s2 2p6 3s1
Shielding electrons: 1s2 2s2 2p6
Valence electrons: 3s1
So we have 10 shielding electrons and 1 valence
electron, right?
Zeff
• Effective nuclear charge (Zeff) is the charge
felt by the valence electrons after you’ve
taken into account the number of shielding
electrons that surround the nucleus.
• Huh?
• Let’s put it in an equation
• Zeff =# of protons - # of shielding electrons
• So, calculate the effective nuclear charge for
the all the elements in period 3
• Now, calculate the effective nuclear charge
for all the elements in group 2
• What pattern do you see arising?
• What is the correlation between Zeff and
atomic radius? (Remember opposite
charges attract)
• The greater the Zeff the smaller the atomic
radius
• I’m still lost…..
• Greater effective nuclear charge means that
the valence electrons are feeling a greater
pull toward the nucleus, making the atom
smaller in size
In summary…
• Effective nuclear charge can be used to
predict trends in atomic radius
• Increases from left to right and decreases from top to
bottom
• Zeff = Z - σ
• Effective nuclear charge is dependent upon
electron shielding
• Electronegativity increases from left to right
and decreases from top to bottom