Isotopes are important when we talk about the atomic mass of an element. This idea is key to understanding the periodic table.

Now, what is atomic mass? It’s not just a whole number. Instead, it’s an average that takes into account all the isotopes of an element found in nature.

Let’s take carbon as an example. Carbon usually has two stable isotopes: carbon-12 and carbon-13. The atomic mass of carbon is about 12.01 atomic mass units (amu). This means that carbon-12, which is more common, has a bigger impact on the average atomic mass. On the other hand, carbon-13 is less common, so it has a smaller effect.

Isotopes mainly differ by the number of neutrons they have. Neutrons are particles in the atom that help add to the total mass. When we calculate the average atomic mass, we use this formula:

Atomic mass = (sum of (mass of isotope × abundance)) ÷ (total abundance)

Here, "abundance" is shown as a decimal. For example, if something is 99% abundant, we write it as 0.99.

So, when we talk about isotopes, we see that atomic mass isn’t just about counting protons and neutrons. It's about how these different isotopes work together. Understanding isotopes is very important for knowing how to find and understand atomic mass on the periodic table.