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Lesson 2.03
Isotopes and
Weighted Averages
Introduction
• How can more than one form of atom
make up an element?
• What can be different and what is always
the same among the different atoms of a
given element?
Objectives
After completing this lesson, you will be able
to:
• Identify the structural differences between
the isotopes of elements.
• Determine the weighted average mass of
an element when given the appropriate
data.
What do you know?
• The number of protons in an atom’s
nucleus identifies the atom as a specific
element.
• Proton’s never change for a particular
element
EXAMPLE: All oxygen atoms have eight
protons. We know this because Oxygen’s
atomic number = 8
What is an Isotope?
• Isotopes are atoms of the
same element that have
different masses due to their
varying numbers of neutrons
Ways to represent an Isotope
Method One:
Add the mass number of the isotope, with a
hyphen, to the end of the element’s name.
EXAMPLES:
an isotope of a carbon (C) atom that has six protons, six
neutrons, and six electrons you can represent the isotope as
Carbon-12. The 12 is the mass number of the isotope (six
protons + six neutrons).
Uranium-235would be an isotope of uranium with a mass of
235 (92 protons + 143 neutrons)
Ways to represent an Isotope
Method Two:
The mass number is given as a superscript to the upper left
side of the element’s symbol, with the atomic number as a
subscript to the lower left side of the symbol.
EXAMPLE: If you have an isotope of a carbon (C) atom that
has six protons, eight neutrons, and six electrons you can
represent the isotope by writing C with 14 a superscript on
the upper left side and six as a subscript in the lower left side
of the symbol. The 14 is the mass number and the six is the
atomic number.
Mass Number
Atomic Number
Method Two Examples
EXAMPLE:
An isotope of hydrogen (H) that has a mass of two
would be written as H with two as a superscript on
the upper left side and one as a subscript in the
lower left side of the symbol. The two is the mass
number and the one is the atomic number.
Practice
• Refer to the periodic table and each
element’s atomic number to answer the
following questions:
1- How many protons, neutrons, and
electrons are in an atom of chlorine-37?
Chlorine has an atomic number of 17.
2-How many protons, neutrons, and
electrons are in an atom of Br with an atomic
number of 35 and a mass number of 80?
More Practice
3-How many protons, neutrons, and
electrons in an atom of beryllium-9?
Beryllium has an atomic number of 4.
4- How many protons, neutrons, and
electrons are in an atom of Ar with an atomic
number of 18 and a mass number of 38?
Check your Answers
1- How many protons, neutrons, and electrons are in an atom of
chlorine-37? Chlorine has an atomic number of 17.
Protons = 17 (atomic number)
Electron = 17 (atomic number)
Neutrons = 20 (neutrons = mass number (37)—atomic
number (17)
2- How many protons, neutrons, and electrons are in an atom of Br
with an atomic number of 35 and a mass number of 80?
Protons = 35 (atomic number)
Electron = 35 (atomic number)
Neutrons = 45 (neutrons = mass number (80)—atomic
number (35)
Check Your Answers
3- How many protons, neutrons, and electrons in an atom of
beryllium-9? Beryllium has an atomic number of 4.
Protons = 4 (atomic number)
Electron = 4 (atomic number)
Neutrons = 5 (neutrons = mass number (9)—atomic
number (4)
4- How many protons, neutrons, and electrons are in an atom of
Ar with an atomic number of 18 and a mass number of 38?
Protons = 18 (atomic number)
Electron = 18 (atomic number)
Neutrons = 20 (neutrons = mass number (38)—atomic
number (18)
Atomic Mass
• Have you noticed that the atomic masses
on the periodic table have decimal values
even though the mass numbers are whole
numbers?
Average Atomic Mass
• The average mass of an element calculated by averaging
the mass values for all of that element’s isotopes.
• Because pure elements are made up of mixtures of
isotopes, it is important to take into account the
percentage and mass of each isotope, especially when
some isotopes make up a greater percentage of the
mixture than other isotopes.
• the average atomic mass of each element is calculated
using a weighted average.
What is a weighted average?
• A weighted average accounts for the percent
abundance and mass of each isotope in an element.
• Percent abundance is a measure of how common or
rare that isotope is.
• The average atomic mass of an element can be
calculated by multiplying the mass of each isotope by
its relative abundance (the percentage represented in
decimal form) and adding the results
• Isotope A [percent abundance × mass] + Isotope B
[percent abundance × mass] + continue for each
isotope = average atomic mass
Example
Isotope
Mass
Percent Abundance
Oxygen- 16
16.0 u
99.72 %
Oxygen- 17
17.0 u
0.038%
Oxygen- 18
18.0 u
0.200%
Step 1: Convert the percentages to decimals by dividing by 100
Isotope
Mass
Percent Abundance (decimals)
Oxygen- 16
16.0 u
99.72 /100= .99762
Oxygen- 17
17.0 u
0.038/100= .00038
Oxygen- 18
18.0 u
0.200/100= .00200
Example continued
Step 2- multiply each isotopes mass by the
decimal of the abundance
Isotope
Mass
Decimal of the Abundance
Total
Oxygen- 16
16.0 u
.99762
15.96192
Oxygen- 17
17.0 u
.00038
0.00646
Oxygen- 18
18.0 u
.00200
0.036
Step 3- add the individual totals together
(15.96192 + 0.00646 + 0.036) = 16.0u
Practice 1
Copper has two naturally occurring isotopes. Copper-63,
with an atomic mass of 62.94 u, makes up 69.17 percent of
the sample, and Copper-65, with an atomic mass of 64.93 u,
makes up the other 30.83 percent.
Calculate the weighted average atomic mass of copper based
on the given information. Please give your answer in a total
of four significant figures.
Isotope
Mass
Percent Abundance
Copper-63
62.94 u
69.17%
Copper-65
64.93 u
30.83%
Practice 2
Three isotopes of argon (Ar) occur in nature. Argon-36
(35.968 u) makes up 0.337 percent of the sample, Argon-38
(37.963 u) makes up 0.063 percent of the sample, and Argon40 (39.962 u) makes up the other 99.6 percent of the
sample.
Calculate the weighted average atomic mass of argon, giving
your answer in a total of five significant figures.
Isotope
Mass
Percent Abundance
Argon-36
35.968 u
0.337 %
Argon-38
37.963 u
0.063%
Argon-40
39.963 u
99.6%
Answers to Practice
Question 1[.6917 × 62.94 u] + [.3083 × 64.93 u] =
63.55 u (average atomic mass)
Question 2[0.00337 × 35.968 u] + [0.00063 × 37.963 u] + [0.99600 ×
39.962 u] =
39.947 u (average atomic mass)
Review
Term
Isotope
Definition
atoms of the same element
that have different masses
due to their varying numbers
of neutrons
Example
Carbon-12, Carbon-14
Percent Abundance
A calculated percentage
based on how common a
specific isotope is in any
given sample of an element
Copper-63, with an atomic
mass of 62.94 u, makes up
69.17% of the sample and
Copper-65, with an atomic
mass of 64.93 u, makes up
the other 30.83%
Average Atomic Mass
average mass of an element
calculated by averaging the
weighted mass values for all
of that element’s isotopes
[.6917 × 62.94 u] + [.3083 ×
64.93 u] = 63.55 u (average
atomic mass)
Congrats!
You should now be able to:
• Identify the structural differences between
the isotopes of elements.
• Determine the weighted average mass of
an element when given the appropriate
data.
Please review this presentation again if you
need additional practice