How do the natural abundances of isotopes affect an element’s atomic weight?

The atomic weight of an element is determined by the natural abundances of its isotopes. Because most elements exist in nature as a mixture of different isotopes—atoms with the same number of protons but varying numbers of neutrons—the atomic weight represents the weighted average of the atomic masses of these naturally occurring isotopes.12

To calculate this value, the mass of each individual isotope is multiplied by its relative abundance (the percentage of that isotope found in a natural sample), and these products are then summed. For example, while helium-3 and helium-4 have distinct atomic masses, the atomic weight of helium is approximately 4.002602 atomic mass units (AMU) because the helium-4 isotope is vastly more abundant than helium-3, which accounts for less than 1 percent of the total.1

Consequently, the atomic weight is not the mass of a single atom but a statistical average reflecting the specific distribution of isotopes found in nature.32