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Periodicity
And the Periodic Table
The Big Idea
• Periodic trends in the properties
of atoms allow us to predict
physical and chemical
properties
Development of the
Periodic Table
• The periodic table evolved over
time as scientists discovered
more useful ways to compare
and organize the elements
The Periodic Table
• Mendeleev (1869) – first periodic
table
• Based on increasing atomic
mass and repeating properties
of elements
• Had spaces for “missing”
elements he predicted
Henry Moseley
• Discovered in
1914 that
elements’
properties more
closely followed
their atomic
number
• Modern Periodic
Table based on
this discovery
Periodic Law
• Physical and chemical
properties of elements are
periodic functions of atomic
numbers
Classification of the
Elements
• Elements are organized into
different blocks in the periodic
table according to their electron
configurations
Organizing the Elements
by Electron Configuration
• Electrons in the highest principal energy
level are called valence electrons.
• All group 1 elements have one valence
electron.
Organizing the Elements
by Electron Configuration
• The energy level of an element’s valence
electrons indicates the period on the
periodic table in which it is found.
• The number of valence electrons for
elements in groups 13-18 is ten less than
their group number.
Organizing the Elements
by Electron Configuration
What is periodicity?
• Properties of the elements change in
a predictable way as you move
through the periodic table
• These properties include
–
–
–
–
–
Atomic radius
Octet Rule
Ionic radius
Ionization energy
Electronegativity
How are the elements
organized?
• Periodic Table: a complete
chart of all elements in the
universe
• Arranged according to physical
and chemical properties
• Each box on the table contains
the atomic number, atomic
mass, and chemical symbol
How are the elements
organized?
• Groups: also known as families;
are the columns
• Have similar properties
• Some have specific names:
–
–
–
–
Family 1: Alkali Metals
Family 2: Alkaline Earth Metals
Families 3 – 12: Transition Metals
Family 13: Boron Family
How are the elements
organized?
–
–
–
–
–
Family
Family
Family
Family
Family
14:
15:
16:
17:
18:
Carbon Family
Nitrogen Family
Oxygen Family
Halogen Family
Noble Gases
• Periods: the rows on the
periodic table
– Do not have similar properties
The Modern Periodic
Table
Metals
• To the left of the stair-step
line
• 88 elements
• Tend to lose electrons
• Most reactive in the s block
• Includes alkali metals and
alkaline earth metals
Properties of Metals
•
•
•
•
•
•
Shiny luster
Good conductors of heat
Good conductors of electricity
Usually solids at room temp
Malleable
Ductile
gold
lead
nickel
copper
Nonmetals
•
•
•
•
On right side of stair-step line
Tend to gain electrons
Most reactive group is halogens
Least reactive is Noble gases
Properties of Nonmetals
•
•
•
•
•
Dull luster
Poor conductors of heat
Poor conductors of electricity
Brittle
Many are gases at room temp
Carbon (graphite)
bromine
sulfur
Metalloids
• On either side of stair-step line
• Have properties of metals and
nonmetals
• Includes all elements that touch
the line except Al and Po
• Many are used in transistors
antimony
germanium
Alkali Metals
• Group 1 (except H)
• All have only 1 valence electron
• Most reactive metals; never
found in pure state in nature
• Soft, shiny, have relatively low
melting points
Alkaline Earth Metals
• Group 2
• All have 2 valence electrons
• Are the second most reactive metals;
never found naturally in pure state
• Harder, denser, 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 13
• Have 3 valence electrons
• Boron is a metalloid, while all of
the others are metals
Carbon Family
• Group 14
• All have four valence electrons
• Carbon is a nonmetal; Si and Ge
are metalloids; Sn and Pb are
metals
Nitrogen Family
• Group 15
• All have 5 valence electrons (in
s and p sublevels)
• N and P are nonmetals; As and
Sb are metalloids; Bi is a metal
Oxygen Family
• Group 16
• All have 6 valence electrons (in
s and p sublevels)
• All are nonmetals except Te,
which is a metalloid, and Po,
which is a metal.
Halogens
• Means “salt former”
• Group 17
• All have 7 valence electrons (in
s and p sublevels)
• Most reactive nonmetals
• All are nonmetals except At,
which is a metalloid
Noble Gases
• Group 18
• Complete, stable electron
configuration (no valence
electrons)
• Most unreactive elements
Rare Earth Elements
• Found in 2 rows at bottom of Periodic
Table
• Also known as the inner transition
metals
• Lanthanide series: starts with La
• Actinide series: starts with Ac
• Little variation in properties
• Actinides are radioactive; only first
three and Pu are found in nature
Periodic Trends
• Trends among elements in the
periodic table include their size
and their ability to lose or
attract electrons
Atomic Radius
• For elements that occur as molecules, the
atomic radius is half the distance between
nuclei of identical atoms.
Atomic Radius
Atomic Radius
• Atomic radius generally increases as you
move down a group.
• The outermost orbital size increases down a
group, making the atom larger.
Ionic Radius
• An ion is an atom or bonded group of
atoms with a positive or negative charge.
• When atoms lose electrons and form
positively charged ions, they always become
smaller for two reasons:
1. The loss of a valence electron can leave an empty
outer orbital resulting in a small radius.
2. Electrostatic repulsion decreases allowing the
electrons to be pulled closer to the radius.
Ionic Radius
• When atoms gain electrons, they can
become larger, because the addition of an
electron increases electrostatic repulsion.
Ionic Radius
Ionic Radius
• The ionic radii of
positive ions
generally decrease
from left to right.
• The ionic radii of
negative ions
generally decrease
from left to right,
beginning with group
15 or 16.
Ionization Energy
• The energy needed to remove one of
its electrons
• Decreases as you move down a
group
• Increases as you move across a
period
• Successive ionization energies
increase for every electron removed
Ionization Energy
Octet Rule (Rule of 8)
• Atoms tend to gain, share, or
lose in order to acquire a full set
of valence electrons (in most
cases, this is 8)
Electronegativity
• Reflects an atom’s ability to attract
electrons in a chemical bond
• Related to its ionization energy and
electron affinity
• Increases as you move across a
period
• Increases as you move up a group
Electronegativity