Periodic Table
Chapter 6
Periodic Table

Many different versions of the Periodic Table
exist

All try to arrange the known elements into an
organized table
Alternate Periodic Tables
Alternate Periodic Tables
Alternate Periodic Tables
Alternate Periodic Tables
Alternate Periodic Tables
Alternate Periodic Tables
Alternate Periodic Tables
Elements known since Ancient times
Elements Discovered in 1600’s
Elements Discovered in 1700’s
Elements Discovered 1800-1810
Elements Discovered 1810-1863
Elements Discovered 1875-1899
Elements Discovered 1900-1940
Elements Discovered 1944-1961
Elements Discovered 1966-1996
Elements Discovered since1999
History

1869 - Russian chemist and teacher, Dmitri
Mendeleev proposed a table for organizing
elements

Mendeleev arranged the elements in a table
based on increasing atomic mass.
History

Mendeleev placed elements next to each
other with similar chemical properties

He would leave elements out of order based
on atomic mass if they lined up better based
on chemical properties
History

Mendeleev left spaces for elements not yet
discovered
–

He predicted properties of elements that would fit
in those spots
He predicted very closely the properties of
Ge, Ga, Sc, and 5 others
History

1913 - British physicist, Henry Moseley,
determined the atomic numbers for the
elements

The modern periodic table is arranged in
order of increasing atomic number.
Periodic Table
Arrangement

Columns are called Groups
–
Numbered 1-18

Rows are called Periods

Elements in the same group have similar
properties
Group Names

Group 1 - Alkali Metals

Group 2 - Alkaline earth metals

Group 17 - Halogens

Group 18 - Inert or Noble gases.
Group Names

Groups 3-11 – Transition Metals

Bottom 2 rows – Inner Transition
Phases at STP

Most elements are solids at STP

Hg and Br are liquids at STP

H, N, O, F, Cl and Noble Gases are all gases
at STP
Periodic Law

Periodic Law – When elements are arranged
in order of increasing atomic number, there is
a periodic repetition of their physical and
chemical properties.
Valence Electrons

Electrons in outermost occupied energy level

Valence Electrons are responsible for most
chemical properties
–
Elements in the same group have similar
properties because they have the same number
of valence electrons
Classifying Elements

Elements are classified into 3 groups based
on their properties:

Metals – Left and Middle

Nonmetals – Right

Metalloids - Staircase
Metals




Good conductors of heat and electrical
current
High luster or sheen
Many are ductile, meaning they can be
drawn into wires
Most are malleable, meaning they can be
hammered into thin sheets
Metals

Metallic Character increases as you move
towards the lower left

Most Metallic Element is Francium, Fr
Nonmetals

Most are gases at room temperature, some
are solids, and one is liquid

Most are poor conductors

Most solids are brittle
Nonmetals

Non-Metallic Character increases as you
move towards upper right

Most nonmetallic element is Fluorine, F
Metalloids

B, Si, Ge, As, Sb, Te

Have properties of both metals and
nonmetals, based on conditions

Exceptions:
–
–
Al and Po are metals
At is a nonmetal
Group Characteristics

Alkali Metals (Group 1)
–
–
–
–
–
H, Li, Na, K, Rb, Cs, Fr
All have 1 valence electron, tend to form +1 ions
Most reactive metals
Not found in nature by themselves, always
combined with someone else
Have properties of metals but are softer and less
dense
Group Characteristics (cont)

Alkaline Earth Metals (Group 2)
–
–
–
–
–
Be, Mg, Ca, Sr, Ba, Ra
All have 2 valence electrons, tend to form +2 ions
Harder and more dense than alkali metals, but
also have higher melting and boiling points
Highly reactive, but not as much as alkali metals
Not found by themselves in nature
Group Characteristics (cont)

Halogens (Group 17)
–
–
–
–
–
–
F, Cl, Br, I, At
All have 7 valence electrons, tend to form -1 ions
Strongly non-metallic
Most active nonmetals
Have low melting and boiling points
Combine readily with metals to form salts
Group Characteristics (cont)

Noble Gases (Group 18)
–
–
–
–
He, Ne, Ar, Kr, Xe, Rn
Colorless gases that are extremely non-reactive
Full valence shell, non-reactive
All are found in small amounts in our atmosphere
Group Characteristics (cont)

Transition Metals (Groups 3-11)
–
–
–
–
–
–
Most are excellent heat and electrical conductors
Most have high melting points and are hard,
except Hg
Less active than group 1 and 2 metals
Many combine with Oxygen to form oxides
(Chemical property)
Many have more than one oxidation number
Form compounds that are colorful
Reminder


STP
Standard Temperature and Pressure
–

1 atm, 0°C
Reference Point for most measurements
Diatomics

Eight elements are diatomic molecules when
alone in nature (exist as two atoms bonded
together)

H2, N2, O2, F2, Cl2, Br2, I2, At2
Diatomics

Hydrogen and the Magic 7
Coloring

Color in the
specific groups
with your own
color choices
Coloring


Color in the
different
classifications
with your own
color choices
Metals,
Nonmetals,
Metalloids
Orbital Blocks
s - block
p - block
d - block
f - block
Periodic Trends

How a property changes either across a
period or down a group
–
–
–
–
–
–
Atomic Number
Atomic Mass
Atomic Radius
Ionic Radius
Ionization Energy
Electronegativity
Trends

Atomic number increases across a period.
–

Increasing number of protons
Atomic number increases down a group
–
Increasing number of protons
Trends

Atomic mass generally increases across a
period.
–

Increasing protons, neutrons, and electrons.
Atomic mass increases down a group.
–
Increasing protons, neutrons, and electrons.
Radius

Atomic Radius – measure of the size of the
atom
–

Half the distance between two nuclei
Ionic Radius – measure of the size of an ion
Trends

Atomic Radius decreases across a period
–

More protons to pull on the electrons
Atomic Radius increases down a group
–
Increasing electrons into more energy levels
(more shells)
Ions

Atom, or group of atoms, that has gained or
lost electrons

Cation – positive ion
Anion – negative ion

Ions

When an atom loses an electron, it becomes
positively charged
–
–
The radius becomes smaller
Metals tend to lose electrons
Ions

When an atom gains an electron, it becomes
negatively charged
–
–
The radius becomes larger
Nonmetals tend to gain electrons
Trends

Ionic Radius decreases for positive ions
across a period
–

More protons to pull on the electrons
Ionic Radius decreases for negative ions
across a period
–
More protons to pull on the electrons
Ionic Radius
+1
+2
+3
+4
-3
-2
-1
Trends

Ionic Radius increases down a group
–
Increasing electrons into more energy levels
(more shells)
Ionization Energy (IE)

Amount of energy required to remove an electron
from an atom
–
Ca  Ca+ + e-
590kJ/mol

First ionization energy is removing the first electron

Second Ionization energy is removing the second
electron after having the first removed
–
Ca+  Ca2+ + e-
1145kJ/mol
IE Trends

Ionization energy tends to increase across a
period
–

More protons are able to hold on tighter to
electrons
Ionization energy tends to decrease down a
group
–
Electrons are farther away from the protons (more
shells)
Electronegativity (EN)

Ability of an atom to attract an electron from
another atom when in a compound.
–
–
–
Noble gases are usually omitted since they don’t
form compounds
Fluorine, F, is the most electronegative element
with a value of 4.0
Francium, Fr, is the least electronegative element
with a value of 0.7
EN Trends

Electronegativity tends to increase across a
period
–

More protons are able to attract electrons better
Electronegativity tends to decrease down a
group
–
Electrons are farther away from the protons (more
shells)
Trends Summary
Property
Atomic Number
Atomic Mass
Atomic Radius
Ionic Radius
Ionization Energy
Electronegativity
Period (LR)
Group (TB)
Reactivity


Elements that are more reactive tend to
either gain or lose electrons very easily
Elements that lose electrons easily have low
IE and low EN
–

Lower left, Fr
Elements that gain electrons easily have high
IE and high EN
–
Upper right, F