The Periodic Table and
Periodic Law
Chemistry Unit 5
Why does the
Periodic Table look
the way it does?
Main Ideas

Periodic trends in the properties of atoms allow us to
predict physical and chemical properties.

The periodic table evolved over time as scientists
discovered more useful ways to compare and organize
elements.

Elements are organized into different blocks in the
periodic table according to their electron
configurations.

Trends among elements in the periodic table include
their size and their ability to lose or attract electrons
Development of the
Modern Periodic Table
Objectives:
 Trace
the development of the periodic table
 Identify
key features of the periodic table
Development of the
Periodic Table
In the 1700’s, Lavoisier compiled a list of all
the known elements of the time.
 33
elements
Development of the
Periodic Table
The 1800s brought large amounts of
information and scientists needed a way to
organize knowledge about elements.
 Advent
of electricity – break down
compounds
 Development
of the spectrometer –
identify newly isolated elements
Development of the
Periodic Table
The 1800s brought large amounts of
information and scientists needed a way to
organize knowledge about elements.
 Industrial
revolution – new chemistry
based ingredients and compounds.
 70
known elements by 1870
Development of the
Periodic Table
The 1800s brought large amounts of
information and scientists needed a way to
organize knowledge about elements.
 John
Newlands proposed an arrangement
where elements were ordered by
increasing atomic mass.
Law of Octaves
Newlands (1864)
noticed when the
elements were
arranged by
increasing atomic
mass, their properties
repeated every eighth
element.
Law of Octaves
Octaves was used due
to the musical
analogy, but was
widely dismissed.
Some elements didn’t
follow the pattern
The Periodic Table
 Meyer
and Mendeleev both demonstrated a
connection between atomic mass and
elemental properties.
The Periodic Table
 Mendeleev’s
Table – A Russian scientist –
gets the most credit because he published
first.
Arranged elements by increasing mass and
columns with similar properties.
 Predicted the existence and properties of
undiscovered elements.
 Still some inconsistencies.

The Periodic Table
 Moseley
discovered that each element had a
distinct number of protons.
 Once rearranged by increasing atomic
number, the table resulted in a clear
periodic pattern.
The Periodic Table
Periodic repetition of chemical and physical
properties of the elements when they are
arranged by increasing atomic number is
called periodic law.
The Modern Periodic
Table

The modern periodic table contains boxes which
contain the element's name, symbol, atomic
number, and atomic mass.
The Modern Periodic
Table

Rows of elements are called periods. (total of 7)

Columns of elements are called groups. (total of 18)

Elements in groups 1,2, and 13-18 possess a wide
variety of chemical and physical properties and are
called the representative elements.

Elements in groups 3-12 are known as the transition
elements .
Types of Elements
Elements are classified as metals, non-metals, and
metalloids.

Metals are made up of most of the representative
elements and all of the transition elements.
They are generally shiny when smooth and clean,
solid at room temperature, and good conductors of
heat and electricity.
 Most are Ductile and Malleable –
 Ductile – the ability to be drawn into wire.
 Malleable – the ability to be pounded into sheets

Types of Elements
Elements are classified as metals,
non-metals, and metalloids.

Alkali metals are all the elements in group 1,
except hydrogen, and are very reactive.

Alkaline earth metals are in group 2, and are also
highly reactive.
Alkali Metals

Alkali metals and water
Types of Elements
The transition elements (groups 3 - 12) are divided
into transition metals and inner transition
metals.

The two sets of inner transition metals are called
the lanthanide series and actinide series and are
located at the bottom of the periodic table.

Lanthanides are phosphors – elements that
emit light when struck by electrons.
The Modern Periodic
Table

Non-metals are elements that are generally gases
or brittle, dull-looking solids, and poor conductors
of heat and electricity.

Group 17 is composed of highly reactive
elements called halogens.

Group 18 gases are extremely unreactive and
commonly called noble gases.
The Modern Periodic
Table

Metalloids have physical and chemical properties
of both metals and non-metals, such as silicon
and germanium. They are found along the stair
step of the table starting with Boron
Questions
What is a row of elements on the periodic
table called?
A. octave
B. period
C. group
D. transition
Questions
What is silicon an example of ?
A. metal
B. non-metal
C. inner transition metal
D. metalloid
Practice Problems

CALM 5.1
Classification of the
Elements
Objectives:
 Explain
why elements in the same group
have similar properties.
 Identify
the four blocks of the periodic table
on their electron configuration.
Organizing the Elements
by Electron Configuration
Electron configuration determines the
chemical properties of an element.
 Recall
electrons in the highest principal
energy level are called valence electrons.
Organizing the Elements
by Electron Configuration
 All
group 1 elements have one valence
electron.
 All
group 2 elements have two valence
electrons.
Organizing the Elements
by Electron Configuration
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.

After the s-orbital is filled, valence electrons occupy
the p-orbital.
Organizing the Elements
by Electron Configuration
Organizing the Elements
by Electron Configuration
Organizing the Elements
by Electron Configuration

The d-block contains the transition metals and is the
largest block.


There are exceptions, but d-block elements usually
have filled outermost s orbital, and filled or partially
filled d orbital.
The five d orbitals can hold 10 electrons, so the d-block
spans ten groups on the periodic table.
Organizing the Elements
by Electron Configuration

The f-block contains the inner transition metals.


f-block elements have filled or partially filled outermost
s orbitals and filled or partially filled 4f and 5f orbitals.
The 7 f orbitals hold 14 electrons, and the inner
transition metals span 14 groups.
Practice Problems

CALM 5:2
Periodic Trends
Objectives:
 Compare
period and group trends of several
properties.
 Relate
period and group trends in atomic
radii to electron configuration
Atomic Radius
Atomic radius – is determined by the amount of
positive charge in the nucleus and the number of
valence electrons of an atom. It is usually measured
in picometers (10-12).

For metals, atomic radius is half the distance between
adjacent nuclei in a crystal of the element.

For diatomic nonmetals, the atomic radius is the
distance between nuclei of identical atoms.
Diatomic Nonmetals
 H2,
N2, O2, F2, Cl2, Br2
Atomic Radius
Organizing the Elements
by Electron Configuration
Atomic Radius
The periodic trend: decreases from left to right
(periods) and increases top to bottom (groups)
due to the increasing positive charge in the
nucleus.
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.

Valence electrons are not shielded from the
increasing nuclear charge because no additional
electrons come between the nucleus and the valence
electrons.
Ionic Radius
Ions – atom(s) that gain or lose one or more electrons to
form a net charge.
Ionic radius is the radius of a charged atom.

When atoms lose electrons and form positively charged
ions, they always become smaller.
 Lost electrons are usually valence electrons and could
leave the outer orbital empty and therefore smaller.
 Electrostatic repulsion between remaining electrons
decreases and pulls closer to nucleus.
Ionic Radius

When atoms gain electrons and forms a
negatively charged ion, they become larger.
 Increased electrostatic repulsion increases
distance of outer electrons.
Organizing the Elements
by Electron Configuration
Ionic Radius
Periodic Trend: radius of an ion decreases from left
to right (periods) until charge changes and then
the radii increases dramatically. After the change,
the radius continues to decrease. Ionic radii
increases top to bottom (groups) until change in
charge.
Ionic Radius
Ionization Energy
Ionization energy is the energy needed to remove
an electron from the positive charge of the
nucleus of a gaseous atom. (how strongly a
nucleus holds on to an electron.)
 First
ionization energy is the energy required to
remove the first electron.
 Removing
the second electron requires more
energy, and is called the second ionization
energy.
Ionization Energy
 Atoms
with large ionization energies have a
strong hold of its electrons and are less likely to
form positive ions.
 Atoms
with low ionization energies lose their
outer electrons easily and readily form positive
ions.
 The
ionization at which the large increase in
energy occurs is related to the number of valence
electrons.
Organizing the Elements
by Electron Configuration
Ionization Energy
Periodic Trend: First ionization energy increases from left
to right across a period. First ionization energy
decreases down a group because atomic size increases
and less energy is required to remove an electron
farther from the nucleus.
Ionization Energy

The octet rule states that atoms tend to gain, lose
or share electrons in order to acquire a full set of
eight valence electrons. The octet rule is useful for
predicting what types of ions an element is likely
to form.
Ionization Energy
Electronegativity
Electronegativity of an element indicates its
relative ability to attract electrons in a
chemical bond. Measured in Paulings:
numbers 4 and less.
Electronegativity
Periodic Trend: electronegativity decreases down a group
and increases left to right across a period.
Questions
The lowest ionization energy is the ____.
A. first
B. second
C. third
D. fourth
Questions
The ionic radius of a negative ion becomes
larger when:
A. moving up a group
B. moving right to left across period
C. moving down a group
D. the ion loses electrons
Practice Problems

CALM 5:3
Ion Formation
Objective:

Learn the common list of cations

Learn the common list of anions
Cations

Cations are atoms or groups of atoms
that have lost electrons.
Anions

Anions are atoms or groups of atoms
that have gained electrons.
Practice Problems

CALM 5:4
5:4 Accumulating Content

How does the electron configuration of an atom
relate to ion formation?
5:4 Accumulating Content

How do chemical and physical properties relate to
periodic trends? What other properties might
have periodic trends?
5:4 Accumulating Content

What does a group tell us about valence electrons
and ion formation?
Key Concepts

The elements were first organized by increasing
atomic mass, which led to inconsistencies. Later,
they were organized by increasing atomic number.

The periodic law states that when the elements are
arranged by increasing atomic number, there is a
periodic repetition of their chemical and physical
properties.

The periodic table organizes the elements into
periods (rows) and groups (columns); elements with
similar properties are in the same group.
Key Concepts

Elements are classified as either metals, nonmetals, or
metalloids.

The periodic table has four blocks (s, p, d, f).

Elements within a group have similar chemical properties.

The group number for elements in groups 1 and 2 equals the
element’s number of valence electrons.

The energy level of an atom’s valence electrons equals its
period number.
Key Concepts

Atomic and ionic radii decrease from left to right across
a period, and increase as you move down a group.

Ionization energies generally increase from left to right
across a period, and decrease as you move down a
group.

The octet rule states that atoms gain, lose, or share
electrons to acquire a full set of eight valence electrons.

Electronegativity generally increases from left to right
across a period, and decreases as you move down a
group.
Chapter Questions
The actinide series is part of the
A. s-block elements.
B. inner transition metals.
C. non-metals.
D. alkali metals.
Chapter Questions
In their elemental state, which group has a
complete octet of valence electrons?
A. alkali metals
B. alkaline earth metals
C. halogens
D. noble gases
Chapter Questions
Which block contains the transition metals?
A. s-block
B. p-block
C. d-block
D. f-block
Chapter Questions
An element with a full octet has how
many valence electrons?
A. two
B. six
C. eight
D. ten
Chapter Questions
How many groups of elements are
there?
A. 8
B. 16
C. 18
D. 4
Chapter Questions
Which group of elements are the least
reactive?
A. alkali metals
B. inner transition metals
C. halogens
D. noble gases
Chapter Questions
On the modern periodic table, alkaline earth
metals are found only in ____.
A. group 1
B. s-block
C. p-block
D. groups 13–18
Chapter Questions
Bromine is a member of the
A. noble gases.
B. inner transition metals.
C. earth metals.
D. halogens.
Chapter Questions
How many groups does the d-block span?
A. two
B. six
C. ten
D. fourteen
THE END
Chapter Questions
Chapter Questions
Chapter Questions
Chapter Questions
Chapter Questions