Section 4.2—Atomic Structure
What are atoms?
Atom - smallest piece of matter that
has the chemical properties of the
element.
What’s in an atom?
An atom is made of three sub-atomic particles
Particle
Location
Mass
Charge
Proton
Nucleus
1 amu =
1.6710-27 kg
+1
Neutron
Nucleus
1 amu =
1.6710-27 kg
0
Electron
Outside the
nucleus
0.00055 amu
9.1010-31
kg
-1
1 amu (“atomic mass unit”) = 1.66  10-27 kg
What gives an atom its identity?
What makes an atom “carbon” as opposed
to “oxygen”?
Every atom has a different number of
protons.
The number of protons determines the
identity of the atom
The atomic number shows the number of
protons.
Atomic number = protons
The Nucleus & Mass
Since the nucleus has protons & neutrons,
and the mass of each one is 1 amu…
The mass of the nucleus (in amu’s) is the
number of protons + neutrons
Since electrons have relatively no mass
(0.054% of one proton or neutron), we
don’t need to worry about them when
determining mass of an atom
Mass # = protons + neutrons
Charges
Protons have a positive charge
Electrons have a negative charge
Neutrons have no charge
Overall charge = protons + (-1)×electrons
Charge = protons - electrons
How do we show information
about an element?
Element symbols
Element Symbol
1 or 2 letters, found on
the periodic table
Mass number
# protons + # neutrons
Charge
A
C
X
Z #
# protons - # electrons
(assumed to be “0” if
blank)
Atomic number
# of protons
Number
How many atoms do you
have?
Example: Element symbols
Element Symbol
O = Oxygen
Charge
-2
Mass number
16
16
-2
O
8
Atomic number
8
Number
Assumed to be “1” if blank
Let’s Practice
Example:
Fill in the
missing
values
Symbol
Name
Atomic
#
Magnesium-25
Mass
#
Charge Proton Neutron Electron
+2
82
126
82
Let’s Practice
Example:
Fill in the
missing
values
Symbol
Mg 2
Pb
25
12
208
82
Name
Remember:
Atomic number is the identity
Atomic number = protons
Charge = proton - electrons
Mass # = protons + neutrons
Atomic
#
Mass
#
12
25
Magnesium-25
Lead-208
82
208
Charge Proton Neutron Electron
12
+2
0
82
13
126
10
82
Isotopes
What are isotopes?
Isotopes - n. Atoms of the same
element with a different number of
neutrons
Some isotopes are radioactive—but not
all…many are quite stable!
Isotopes Example
Mass # = 2 amu
Mass # = 1 amu
Hydrogen-1
Hydrogen-2
 If they have different number of neutrons, and
neutrons have a mass of 1 amu…
 Then isotopes of the same element will have
different masses!
 But because their protons are the same, they
are the same element!
Identifying Isotopes
Isotopes can be differentiated by their different
mass numbers in the element symbol
12
C
Carbon-12
13
C
Carbon-13
Or by the mass number following their name.
Mass Number versus Atomic Mass
Mass Number
Average Atomic Mass
# of protons + # of neutrons
Average of actual masses
Always a whole number
Not a whole number
For one specific isotope
only
Weighted average of all
isotopes
Is not found on the periodic
table
Is found on the periodic table
Calculating Average Atomic Mass
Average atomic mass is a weighted average (it
takes into account how often each isotope occurs).
Actual mass
(not mass
number)
“Sum of”
Average
=
atomic
mass
(
Abundance of
isotope
 Mass of )
isotope
What fraction of the time is that
isotope present?
Example of Finding Avg Atomic Mass
Example:
Find the atomic mass of
chlorine if
Chlorine-35 has a mass of
34.969 amu
and Chlorine-37 has a
mass of 36.966 amu and is
present 24.22% of the time.
This chart summarizes
the information in the
problem:
Remember that percents add up to
100.
So they said the second isotope is
present 24.22% of the time.
This means that the first isotope is
present 100-24.22 = 75.78% of the
time
Isotope
Mass
Percent
Decimal
1
34.969 amu 75.78
0.7578
2
36.966 amu 24.22
0.2422
Avg Mass  0.7578  34.969amu  0.2422  36.966
= 35.45 amu
(this is what’s on the periodic table for Cl!)