The atomic mass of an element, expressed in atomic mass units (amu), represents the average mass of the atoms of that element. It takes into account the relative abundance of different isotopes of the element.
Atoms of the same element can have different numbers of neutrons in their nuclei, resulting in different isotopes. Each isotope has a slightly different mass due to the varying number of neutrons. For example, carbon-12 and carbon-14 are two isotopes of carbon with atomic masses of approximately 12 amu and 14 amu, respectively.
The atomic mass listed on the periodic table represents the weighted average of the masses of all the naturally occurring isotopes of that element, taking into account their relative abundances. It is calculated by multiplying the mass of each isotope by its abundance (as a decimal), and summing these values. The resulting value is the atomic mass.
For instance, if an element has an atomic mass of 40 amu, it means that the weighted average mass of all the naturally occurring isotopes of that element is approximately 40 atomic mass units. This value provides a convenient reference for comparing the masses of different elements.