Menu Close

# Isotopes and Atomic Mass

## Isotopes

Isotopes differ from their original element in that they have the same number of protons in their atomic nucleus but a different number of neutrons. This difference in neutron number results in isotopes having the same atomic number but different atomic masses, which affects their physical and chemical properties. To identify isotopes, their atomic number and atomic mass are used, such as carbon-12, carbon-13, and carbon-14. Due to variations in their atomic mass and nuclear stability, isotopes can exhibit distinct physical and chemical properties compared to the original element and to each other. Additionally, many isotopes are radioactive and undergo radioactive decay, emitting radiation until they become more stable forms.

## Atomic Mass

To account for the infinitesimal size of protons, neutrons, and electrons, a new unit of measurement was introduced: the atomic mass unit (amu). This unit is approximately one-twelfth of the mass of a carbon-12 atom. A carbon-12 atom consists of 6 protons and 6 neutrons, and therefore, it is reasonable to approximate the mass of a proton and a neutron as 1 amu each.

Upon examining the Periodic Table of Elements, one will notice that the atomic mass of carbon is typically recorded as 12.01 amu. This value represents an average mass of different isotopes of carbon, taking into account their relative abundance in nature. Therefore, the recorded mass refers to the average mass of a carbon isotope.

## Calculating Atomic Mass

To calculate atomic mass, you use the following formula:

Mass = (mass1Ă—abundance1)+(mass2Ă—abundance2)+…

Mass would be the mass of a specific isotope in amu. Abundance is the relative abundance of a specific isotope as a decimal. The abundance values should sum to 1.