Atoms in a chemical element that have different numbers of neutrons than protons and electrons are called isotopes. The atoms in a particular element have an identical number of protons and electrons but can have varying numbers of neutrons.

Hydrogen is a common element on earth. Hydrogen’s atomic number is 1 — its nucleus contains 1 proton. The hydrogen atom also has 1 electron. Because it has the same number of protons as electrons, the hydrogen atom is neutral (the positive and negative charges cancel each other out).

Most of the hydrogen atoms on earth contain no neutrons. You can use symbolization to represent hydrogen atoms that don’t contain neutrons, as shown in part (a) of the following diagram.

However, approximately one hydrogen atom out of 6,000 contains a neutron in its nucleus. These atoms are still hydrogen, because they have one proton and one electron; they simply have a neutron that most hydrogen atoms lack. So these atoms are called isotopes.

The isotopes of hydrogen.
The isotopes of hydrogen.

Part (b) of the diagram shows an isotope of hydrogen, called deuterium. It’s still hydrogen, because it contains only one proton, but it’s different from the hydrogen in part (a), because it also has one neutron. Because it contains one proton and one neutron, its mass number is two.

There’s even an isotope of hydrogen containing two neutrons. This one’s called tritium, and it’s represented in part (c) of the diagram. Tritium doesn’t occur naturally on earth, but it can easily be created.

Notice that the diagram also shows an alternative way of representing isotopes: Write the element symbol, a dash, and then the mass number.

Now you may be wondering, “If I’m doing a calculation involving the atomic mass of hydrogen, which isotope do I use?” Well, you use an average of all the naturally occurring isotopes of hydrogen. But not a simple average.

You have to take into consideration that there’s a lot more H-1 than H-2, and you don’t even consider H-3, because it’s not naturally occurring. You use a weighted average, which takes into consideration the abundances of the naturally occurring isotopes.