The atomic mass of an element is not always a whole number because it takes into account the existence of different isotopes of that element. Isotopes are atoms of the same element that have the same number of protons but different numbers of neutrons.
The atomic mass listed on the periodic table is a weighted average of the masses of all the naturally occurring isotopes of that element. The average is calculated by considering the abundance (relative proportion) of each isotope and its respective mass.
Since isotopes have different numbers of neutrons, their masses will be slightly different. The atomic mass of an element is expressed in atomic mass units (amu) and can be a decimal value due to this isotopic variation. The decimal part represents the weighted contribution of the different isotopes.
For example, carbon has three naturally occurring isotopes: carbon-12, carbon-13, and carbon-14. Carbon-12 is the most abundant, followed by carbon-13 and carbon-14. The atomic mass of carbon is listed as approximately 12.01 amu, which reflects the contribution of these isotopes to the average atomic mass.
So, the decimal nature of atomic mass is a result of the presence of different isotopes and their relative abundances in the naturally occurring samples of an element.