Why are atomic masses not whole numbers?

Atomic masses are not whole numbers because they are an average of the masses of different isotopes of an element.

Atomic masses, also known as relative atomic masses, are calculated by taking into account the masses of all the different isotopes of an element and their relative abundance. Isotopes are atoms of the same element that have different numbers of neutrons, and therefore different masses. For example, carbon has two stable isotopes, carbon-12 and carbon-13, which have different masses due to the different number of neutrons in their nuclei.

The atomic mass of an element is calculated by multiplying the mass of each isotope by its relative abundance (the percentage of that isotope found in nature), and then adding these values together. This results in a weighted average, which is why the atomic mass is not a whole number.

For instance, let's consider carbon. About 98.9% of carbon on Earth is carbon-12, which has a mass of exactly 12 atomic mass units, and about 1.1% is carbon-13, which has a mass of approximately 13 atomic mass units. The atomic mass of carbon is therefore calculated as (0.989 x 12) + (0.011 x 13), which equals approximately 12.01 atomic mass units.

This method of calculating atomic masses allows scientists to account for the natural variation in the isotopes of an element. It provides a more accurate representation of the mass of an element as it is found in nature, rather than just considering the mass of one isotope. This is why atomic masses are usually not whole numbers, but rather decimals.