The relative atomic masses of elements are not always whole numbers. In fact, most of them are not. The relative atomic mass of an element is a weighted average of the masses of its naturally occurring isotopes, taking into account the abundance of each isotope.
Isotopes are atoms of the same element that have different numbers of neutrons in their nuclei. Since the number of neutrons can vary, the mass of an atom can also vary. For example, carbon has two stable isotopes: carbon-12 and carbon-13, with atomic masses of approximately 12 and 13 atomic mass units (u), respectively. The relative atomic mass of carbon is a weighted average of these isotopes, with carbon-12 being more abundant, so the value is closer to 12.
However, there are some elements that have only one stable isotope, such as helium-4 and nitrogen-14, whose atomic masses are both close to whole numbers. In these cases, the relative atomic mass is equal to the atomic mass of the isotope since there is no need for a weighted average.
It's worth noting that the atomic masses listed on the periodic table are not always whole numbers due to the inclusion of other isotopes with lower abundance. These values are typically given with several decimal places to reflect the average atomic mass accurately.