Isotopic abundance significantly influences the relative atomic mass of an element by contributing to the average mass of its atoms.

To elaborate, isotopes are variants of the same element that possess the same number of protons but differ in their neutron count. This variation in neutrons leads to differences in atomic mass. For instance, carbon has two common isotopes: carbon-12, which contains 6 neutrons, and carbon-14, which contains 8 neutrons.

Isotopic abundance refers to the proportion of each isotope of an element found in nature, expressed as a percentage. For example, approximately $99\%$ of naturally occurring carbon is in the form of carbon-12, while only about $1\%$ exists as carbon-14. This indicates that the isotopic abundance of carbon-12 is substantially greater than that of carbon-14.

The relative atomic mass of an element is calculated as the weighted average of the atomic masses of its isotopes, considering their respective isotopic abundances. This means that the calculation is not a simple arithmetic average; instead, each isotope’s mass is multiplied by its abundance, and these values are summed to yield the final result.

Consequently, if an element has one isotope that is significantly more abundant than the others, that isotope will exert a greater influence on the element’s relative atomic mass. For carbon, the relative atomic mass is approximately $12$, primarily because carbon-12 is so much more prevalent compared to carbon-14.

In summary, the isotopic abundance of each isotope of an element is vital in determining the element’s relative atomic mass. This is why the relative atomic masses of most elements are not whole numbers; rather, they are expressed as decimals.