What can be concluded about the relative atomic mass of an element with stable isotopes?

The relative atomic mass of an element is calculated as the weighted average of the masses of its naturally occurring isotopes. Each isotope contributes to the overall atomic mass based on its abundance and its mass. Because isotopes of an element can vary in mass while coexisting at specific proportions in nature, this average does not simply reflect whole number values unless all isotopes have whole number masses and identical abundances.

This concept ties nicely into the understanding of stable isotopes, as they are isotopes that do not undergo radioactive decay. Therefore, the atomic mass reflects the average of these stable forms, weighted by how much of each isotope exists in a particular sample of the element. The result is often a decimal number, indicative of the mixture of isotopes, rather than a whole number that would denote a simplistic representation of a single mass.

The other conclusions related to the relative atomic mass do not accurately encapsulate its calculation principle. For instance, while the mass of heavier isotopes may influence the average, it is not an absolute rule that the relative atomic mass will be higher for them since the average considers all isotopes in the sample. Similarly, while the atomic mass will indeed never exceed the highest isotope mass, this does not fully describe the calculation method, which