ISOTOPES – PROBLEM SET
IPC – Mr. Coburn
Introduction
Atoms of a given element which have the same number of protons but different numbers of neutrons
are called isotopes. Thus, isotopes have the same position in the periodic table, the same chemical
properties and the same atomic charge.
The simplest example of an atom with different isotopes is hydrogen. The three isotopes of hydrogen
are shown below:
The increasing number of neutrons in the nucleus of the hydrogen atom adds mass to the atom and
thus each isotope of a given element has a different mass.
For the isotopes of hydrogen, 1H (or hydrogen-1), 2H (or hydrogen-2) and 3H (or hydrogen-3)
represent Protium (usually just referred to as hydrogen), Deuterium and Tritium, respectively. Most of
the light elements contain different proportions of at least two isotopes. Usually one isotope is the
predominantly abundant isotope. For example, the average abundance of 12C is 98.89%, while the
average abundance for 13C is 1.11%.
Calculating Atomic Mass
The calculation of the atomic mass of an element is performed using the relative abundance data
from the isotope of each atom.
Atomic Mass = [(% abundance of isotope)(mass of isotope)] + [(% abundance of isotope)(mass of isotope)] + [….]
100
For example:
The natural abundance for boron isotopes is 19.9% 10B (10.013 amu) and 80.1% 11B (11.009 amu).
Calculate the atomic mass of boron.
Average atomic mass =
[(19.9%)(10.013)] + [(80.1%)(11.009)]
100
=
10.811 amu
Part I – Answer the following questions
1. What is an isotope? _____________________________________________________________
________________________________________________________________________________
2. What does the number next to isotopes signify? _______________________________________
________________________________________________________________________________
3. How can you tell isotopes apart? ___________________________________________________
________________________________________________________________________________
Part II - Use the periodic table and the information provided to complete any missing
information (including isotope numbers) for the charts below.
Chromium-52 Chromium-54
# of protons
# of protons
# of neutrons
# of neutrons
# of electrons
# of electrons
Iron-
Iron-
# of protons
# of protons
# of neutrons
# of neutrons
29
31
# of electrons
# of electrons
# of protons
# of protons
Selenium-35
Selenium-36
Iodine-
Iodine-
73
75
146
147
79
# of neutrons
# of neutrons
117
# of electrons
118
# of electrons
92
Part III - Calculate the atomic mass for each element based on the information provided.
1. Calculate the average atomic mass of iron if its abundance in nature is 15% iron-55 and 85% iron-56.
2. What is the average atomic mass of silicon if 92.21 % of its atoms have a mass of 27.977 amu, 4.07 %
have a mass of 28.976 amu, and 3.09 % have a mass of 29.974 amu?
3. Calculate the average atomic mass for neon if its abundance in nature is 90.5% neon-20 (19.922 amu),
0.3% neon-21 (20.994 amu), and 9.2% neon-22 (21.991 amu).
4. Calculate the average atomic mass of silver if 13 out of 25 atoms are silver-107 and 12 out of 25 atoms are
silver-109.
5. Calculate the average atomic mass of chromium.
Isotope
Mass (amu)
Chromium – 50
49.946
Chromium – 52
51.941
Chromium – 53
52.941
Chromium – 54
53.939
Relative Abundance
0.043500
0.83800
0.095000
0.023500
6. Boron has three naturally occurring isotopes: boron-10, boron-11, and boron-12. If the average atomic
mass of boron is 10.811 amu, which isotope is the most abundant? How do you know?
7. Bromine has two isotopes. Bromine-79 occurs at a relative abundance of 50.69% and Bromine-81 has a
relative abundance of 49.31%. What is the atomic mass of bromine?
8. Determine the atomic mass of Chlorine with the following data:
i. Cl-35 with a relative abundance of 75.77%
ii. Cl-37 with a relative abundance of 24.23%