Why are the atomic masses for most elements not whole numbers?

Question

My title is my question. Does anyone have an idea why?

Is this correct?:

5. The atomic masses for most elements are not whole numbers because it is all the relative weights added together, and to get that figure the relative abundance is multiplied by the mass of each isotope, which produces fairly small numbers.

Answer 1

that sounds basically correct

Answer 2

yes. that's pretty much it.

There are different isotopes for each atom.

Carbon 12, the most common for example has an isotope Carbon 14 which is used in dating things that used to be living.

But it's rather rare in nature. So Carbon is 12.01 i think.

Answer 3

An element consists of several isotopes, each with the same atomic number or number of protons, but with different numbers of neutrons. Usually one isotope is common and the others consist of a small percentage of the total.

Listings of atomic weights or unified atomic mass units usually average the isotopes of an element. Thus, even if the old integer system for atomic weights held, the number still wouldn't be simple. For example, the atomic weight of Carbon-12 is 12 amu. But Carbon consists of C-12, C-13, and C-14 isotopes. Thus the average atomic weight of Carbon is 12.01 amu.

Answer 4

That is essentially correct. Also, the contribution of the mass of electrons must be taken into account; even though they mass about 2500 times less than protons and neutrons, they now provide a measurable contribution to atomic mass within current technology.