Review – Periodic Table
• The modern periodic table is not arranged
by increasing atomic mass, but rather
increasing atomic number
• Periodic Law: States that the properties of
elements are periodic functions of their atomic
numbers
• The placement of an element on the Periodic
Table gives an indication of the chemical and
physical properties of that element
Periods

Periods are horizontal rows on the periodic table

Period 1 contains H and He

The number at the beginning of the period
indicates the energy level where valence
electrons are located for atoms in that period
Groups or Families

Vertical Columns

Elements in a group have the same number of valence
electrons.

Elements in a group have similar properties and react
similarly because they have the same number of valence
electrons (this is known as Periodic Law)
Common Group Names
Group 1: Alkali Metals
Group 2: Alkaline Earth Metals
Groups 3-12: Transition Elements (metals)
Group 17: Halogen Group
Group 18: Noble Gases
Properties of Elements
Physical Properties of Metals
Metallic luster (shine)
Malleable (can be made into sheets)
Ductile (can be made into wires)
Metals are good conductors of heat and
electricity
• Most metals are solids at room temperature or
STP which stands for standard temperature and
pressure (exception is mercury which is liquid at
room temperature).
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Properties of Elements
Physical Properties of Non-metals,
Metalloids, and Noble Gases
• Non-metals: Lack metallic luster and are brittle
in the solid phase; they are poor conductors of
heat and electricity; non-metals are usually gases,
molecular solids or network solids (exception is
bromine which is a liquid at room temperature).
• Metalloids: Have some properties of both metals
and nonmetals.
• Noble Gases: Are all gases. Have 8 valence
electrons (helium is the exception with 2).
• SOLIDS: Most of the elements
of the PT
• LIQUIDS: Hg & Br
• GASES: 11 TOTAL
(always reactive)H2, N2, O2, F2,
Cl2,
(typically non-reactive) Noble
Gases
Periodic Trends (Table S)
Ionization Energy (IE): The energy
required to lose an electron.
• The lower the IE, the easier it is to lose an
electron
• Remember, metals want to lose electrons,
and non-metals want to gain
Ionization energy down a group
As you go down a group ionization energy
decreases because the number of PELS increase,
and so the valence electrons are further away
from the positive nucleus; there is less attraction
which results in less energy needed to remove
electrons
 Look at the ionization energies on Table S
Li
520
Na
496
K
419
Rb
403

Ionization Energy across a Period

As you go from left to right ionization energy
increases because as you go across a period, the
number of valence electrons increase by 1, and
therefore the less likely an element wants to lose
an electron (on the contrary, they want to gain
them).

Look at ionization energies on Table S
Li (520)  Ne (2081)
Periodic Trends (Table S)
Electronegativity (EN): The
affinity to gain an electron.
• The larger the EN the more the atom
attracts electrons
• Remember, metals want to lose electrons,
and non-metals want to gain
Electronegativity energy down a group
• As you go down a group EN decreases because as
you go down a group, the number of electron
orbitals increase, and therefore the outermost
shell of electrons is further from the positively
charged nucleus making the attraction for
electrons decrease.
• Look at the electronegativities on Table S
Li
1.0
Na
.9
K
.8
Rb
.8
Electronegativity across a Period
• As you go from left to right EN increases
because as you go across a period, the
number of valence electrons increase by 1,
and therefore the more likely an element
wants to gain an electron (metal  nonmetal)
• Look at ionization energies on Table S
Li (1.0)  F (4.0)
Periodic Trends (Table S)
Atomic Radius (AR): Is the distance from the
nucleus to the outer valence electrons.
radius increases.
Ionic Radius
• RULE: Positive Ion
LOSING AN ELECTRON MEANS THAT THE IONIC RADIUS
IS LESS THAN THE ATOMIC RADIUS
(LOSE = LESS)
AR > IR
• RULE: Negative Ion
GAINING AN ELECTRON MEANS THAT THE IONIC
RADIUS IS GREATER THAN THE ATOMIC RADIUS
(GAIN = GREATER)
AR < IR