The atomic masses that you see in the periodic table are not whole numbers primarily because they represent the weighted average of all the isotopes of an element found in nature.
Most elements exist as a mixture of isotopes, which are atoms of the same element that have different numbers of neutrons. For example, carbon has isotopes like carbon-12 and carbon-14. Each isotope has a different mass due to the varying number of neutrons. The atomic mass of an element takes into account the relative abundance of each isotope found in nature.
To calculate the atomic mass, scientists use the formula:
Atomic Mass = (mass of isotope 1 × relative abundance of isotope 1) + (mass of isotope 2 × relative abundance of isotope 2) + …
This averaging process is why atomic masses often end up being fractional or not whole numbers. For example, chlorine has two main isotopes: chlorine-35 and chlorine-37. The average of these isotopes, weighted by their natural abundance, results in a value that is approximately 35.5.
In summary, the atomic masses are not whole numbers because they reflect the natural distribution of isotopes and their masses, providing a more accurate representation of the element’s characteristics in the periodic table.