The relative atomic mass of an element is determined by calculating the weighted average of the masses of its isotopes, adjusted according to their natural abundances.

The relative atomic mass, denoted as $A_r$, is not merely the mass of a single atom of the element. Instead, it represents a weighted average of the masses of all the atoms in a naturally occurring sample of that element, factoring in the proportions of its various isotopes. This is important because most elements are found in nature as a mixture of different isotopes. An isotope is defined as a variant of a chemical element that has a different number of neutrons.

To compute the relative atomic mass, you must know both the masses of the isotopes and their relative abundances. The relative abundance refers to the percentage of each isotope present in a naturally occurring sample of the element. The mass of an isotope is measured on a scale where the carbon-12 atom has a defined mass of exactly $12$ atomic mass units (amu).

The calculation of the relative atomic mass involves multiplying the mass of each isotope by its relative abundance (expressed as a decimal) and then summing these values. For example, consider an element with two isotopes: one isotope has a mass of $20$ amu and a relative abundance of $90\%$ (or $0.9$ when converted to a decimal), while the other isotope has a mass of $22$ amu with a relative abundance of $10\%$ (or $0.1$). The calculation for the relative atomic mass would be:

This method of calculating relative atomic mass acknowledges that not all atoms of an element are identical due to the presence of isotopes. Consequently, it provides a more precise measure of atomic mass than simply considering the mass of a single atom. This understanding is crucial for analyzing the properties of elements and predicting their chemical reactions.