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Chemistry - Isotopes

The atoms of a particular element have identical numbers of protons and electrons, but can have varying numbers of neutrons. If atoms have different number of neutrons, then the atoms are called isotopes, so all the isotopes of an atom have the same atomic number but have different mass numbers.


The image above shows symbolization chemists use to respond a specific isotope of an element.

In the image above, the X represents the symbol of the element found on the periodic table; Z represents the atomic number (number of protons in the nucleus), and A represents the mass number (the sum of the protons and neutrons in a particular isotope).

If you subtract the atomic number from the mass number (A – Z) you get the number of neutrons in a particular isotope.

Scientists attach the mass number to an element’s name to differentiate between isotopes, for example, helium with a mass number of three is called helium-3, and helium with a mass number of four is called helium-4.

The percentage of each isotope found in nature is known as the isotope’s isotopic abundance. For example, the isotopic abundance of helium-3 is very small. 0.00014%, while the abundance of helium-4 is 99.99986%. This means, only about one of every 1 million helium atoms is helium-3 (the rest are helium-4).

Elements also have unstable isotopes, which are prone to breaking down or decaying that their isotopes of an element. When atoms decay, the number of protons in the nucleus changes. This decay changes one element into another, because the number of protons in the nucleus of an atom determines what element the atom belongs to.

Different isotopes decay at different rates – one way to measure the rate of decay of an isotope is to find its half-life. The half-life of an isotope is the time that passes until half of a sample of an isotope has decayed.

Various isotopes of a given element have nearly identical chemical properties and many similar physical properties. They differ in their mass – the mass of helium-3 is 3.016amu (atomic mass units), while the mass of helium-4 is 4.003amu.

Scientists do not usually specify the atomic weight of an element, in terms of one isotope or another. They express atomic weight as an average of all the naturally occurring isotopes of the element, taking into account the isotopic abundance.

For example:

The element copper has two naturally occurring isotopes:

Copper-63, with a mass of 62.930 amu, and an isotope abundance of 69.2%.

Copper-65, with a mass of 64.928 amu, and an isotopic abundance of 30.8%.

The average mass of naturally occurring copper atoms is equal to the sum of the atomic mass for each isotope multiplied by its isotopic abundance:

(62.930amu X 0.692) + (64.928amu X 0.308) = 63.545amu.

The atomic weight of copper is therefore, 63.545g.

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