Section 2.3 – The Periodic Table
and Atomic Theory
Text Pages 60-69
Periodic Table and Atomic Theory

The periodic table has a series of
patterns! For example metals appear on
the left and non-metals appear on the
right.
Periodic Table and Atomic Theory

These patterns are no accident, and they
result from changes in the structure of
atoms as you move across the periodic
table.

Atomic Structure (Recap)
Recall

The atomic # for an
element tells you the
number of protons (p+)
and electrons (e-) in an
atom of that element
Recall

Electrons are found outside
the nucleus of an atom.

Neils Bohr said that
electrons exist in energy
levels.
*Electrons do not actually move in
circles or ellipses. We have drawn it
this way because it is the easiest way
to represent it.

In this section we will learn about the way that
electrons are organized in energy levels and how
this influences the reactivity of different elements.
STRUCTURE OF AN ATOM:
ENERGY LEVELS
Energy Level: a region surrounding the
nucleus of an atom that may be occupied by
one or more electrons. They are also called
electron shells.
Important Points about Energy
Levels

The period # of an element tells you how
many energy levels its atoms have.
Ex: Carbon has ______ energy levels with
electrons in them.
Ex: Sulfur has ______ energy levels with
electrons in them.
Important Points about Energy
Levels

Energy levels can be numbered from 1 to 7.

Electrons in higher energy levels are
further from the nucleus.

Electrons in higher energy levels have more
energy.
The highest energy level in an atom that
has electrons in it is known as the valence
energy level. Any electrons in that level
are known as valence electrons.
Closer,
Less
Energy
Further,
More
energy

Not all energy levels can hold the same
number of electrons. The rules below
outline how many electrons each level
can hold below.
Energy Level
Max. Number of
Electrons
1st Level
2
2nd Level
8
3rd Level
8
4th Level
18

Electrons enter these energy levels
starting closest to the nucleus.

An energy level must be full before
electrons start to occupy a higher level!

Assembly analogy
Bohr-Rutherford Diagrams

Bohr-Rutherford Diagram: shows the
arrangement of the subatomic particles in
an atom. It shows the number of protons
and neutrons in the nucleus. It also shows
the number of electrons in each energy
level.
Drawing Bohr-Rutherford Diagrams

Step 1: Draw a circle to represent the
nucleus. Write the element symbol inside
at the top of the circle.
Drawing Bohr-Rutherford Diagrams

Step 2: Find atomic number  gives #
of electrons (e-) and # of protons (p+)
Drawing Bohr-Rutherford Diagrams

Step 3: Calculate # of neutrons (no)
using Mass # - Atomic #
Drawing Bohr-Rutherford Diagrams

Step 4: Label inside the nucleus with #
of p+ AND # of no (write these below
the element symbol!)
Drawing Bohr-Rutherford Diagrams

Step 5: Fill the energy levels using the
2, 8, 8, 18 energy level pattern for
electrons.
Each energy level looks like a dashed
horizontal line with # of e- written
inside it: --5e--
Worksheet #11

Complete diagrams for the first 18
elements.
Remember:
Valence Electrons and Reactivity

Do you notice any relationship between the maximum
# of electrons in an energy level and the # of elements
in each period?
What do you notice about their electron
structures?
 Valence energy levels?
 # of valence electrons?
Elements in the same family (group) have
the same number of valence electrons. This
explains why they have similar properties.

The number of energy levels in the atoms
of an element increases as you move
down a group.

Elements in the same period have
electrons in the same valence energy
level.
Valence electrons are very important
because the number of valence electrons
determines how reactive an element will
be.

If an element has a full valence energy
level, it will not be reactive. We say that it
is stable. This helps explain why the Noble
Gases are stable.

Example: Neon

Elements that do not have full valence
energy levels will react with other
elements in order to become stable!
These reactions involve gaining, losing, or
sharing electrons!

Examples: Lithium and Fluorine