Isotopes/Average Atomic Mass Tutorial

Isotopes: All the atoms of an element have the same atomic number, but they can have different numbers of neutrons and different mass numbers. Isotopes of an element are atoms that have the same number of protons, but different numbers of neutrons. Examples of isotopes are the three different kinds of carbon atoms where all have 6 protons, but different numbers of neutrons specifically 8, 7 and 6 neutrons respectively.
Examples of the carbon isotopes:
  • Carbon-14 = 8 neutrons
  • Carbon-13 = 7 neutrons
  • Carbon-12 = 6 neutrons
Naming Isotopes: An isotope name contains the name of the element and the mass of the isotope.
Carbon-14 Carbon-13 Carbon-12
Greek "iso" means same and "topos" means place. This fits the idea that isotopes are in the same place in the periodic table, but have different masses. Periodic table entries provide the information shown here. The periodic table does not indicate isotope information.
Isotope abundances
The isotopes of an element do not occur with equal frequency. The abundance (how common it is) depends on the stability of the isotope. The isotopes contribute to the average atomic mass based on their abundance. The result is that the most abundant or common isotope dictates the “average” mass for the atoms of an element. The atomic mass in the periodic table are massed averages. This means the tabulated value doesn't match any actual atom, but is closer to the most common isotope.

Stable vs. Unstable Isotopes

While there may be several isotopes of the same element,some of isotopes are stable while others are unstable or radioactive. Radioactive isotopes emit nuclearradiation in the form of rapidly moving particles or high energy electromagnetic waves. The particles are emitted from the nucleus itself and their removal results in changing the atom from one isotope to another. This change may occur once or emission of particles may continue until the atom becomes a stable isotope.

The Downside of Isotopes

The downside to any application involving isotopes is how to safely dispose of the radioactive waste generated during processing. Any nuclear process involves the manufacture of nuclear waste whether they are low level (gloves, cotton balls, hospital gowns) or high level (spent nuclear fuel rods) waste products. Concern about the management of nuclear waste materials has caused much controversy and concern among government agencies, industrial, scientific and medical users and citizenry who have nuclear repository facilities in their areas. This becomes a real issue of risks versus benefits.

How to Calculate an Average Atomic Mass

To do these problems you need some information: the exact atomic mass for each naturally-occurring stable isotope and its percent abundance.

Example #1: Carbon

Mass number / Exact mass / Percent abundance
12 / 12.000000 / 98.90%
13 / 13.003355 / 1.10%

To calculate the average atomic mass, each exact atomic mass is multiplied by its percent abundance (expressed as a decimal). Then, add the results together and round off to an appropriate number of significant figures (hundredths place will be fine for now).

This is the solution for carbon:

Work  (12.000000) (0.9890) + (13.003355) (0.0110) = 12.01amu

Example #2: Nitrogen

Mass number / Exact mass / Percent abundance
14 / 14.003074 / 99.63%
15 / 15.000108 / 0.37%

This is the solution for nitrogen:

Work  (14.003074) (0.9963) + (15.000108) (0.0037) = 14.01amu

Try These On Your Own: (show your work)

Example #3: Chlorine / Example #4: Silicon
Mass number / Exact mass / Percent abundance / Mass number / Exact mass / Percent abundance
35 / 34.968852 / 75.77 / 28 / 27.976927 / 92.23
37 / 36.965903 / 24.23 / 29 / 28.976495 / 4.67
30 / 29.973770 / 3.10
The answer for chlorine: / The answer for silicon:

show your work here: