Even though atoms are very tiny pieces of matter, they have mass. Their masses are so small, however, that chemists often use a unit other than grams to express them—the atomic mass unit.

The atomic mass unit: (abbreviated u, although amu is also used) is defined as [latex]\frac{1}{12}[/latex] of the mass of a C-12 atom:

1 u=[latex]\frac{1}{12}[/latex] the mass of C-12 atom

It is equal to 1.661 × 10⁻²⁴ g.

Masses of other atoms are expressed with respect to the atomic mass unit. For example, the mass of an atom of H-1 is 1.008 u, the mass of an atom of O-16 is 15.995 u, and the mass of an atom of S-32 is 31.97 u. Note, however, that these masses are for particular isotopes of each element. Because most elements exist in nature as a mixture of isotopes, any sample of an element will actually be a mixture of atoms having slightly different masses (because neutrons have a significant effect on an atom’s mass). How, then, do we describe the mass of a given element? By calculating an average of an element’s atomic masses, weighted by the natural abundance of each isotope, we obtain a weighted average mass called the atomic mass (also commonly referred to as the atomic weight) of an element.

For example, boron exists as a mixture that is 19.9% B-10 and 80.1% B-11. The atomic mass of boron would be calculated as (0.199 × 10.0 u) + (0.801 × 11.0 u) = 10.8 u. Similar average atomic masses can be calculated for other elements. Carbon exists on Earth as about 99% C-12 and about 1% C-13, so the weighted average mass of carbon atoms is 12.01 u.

The table in Chapter 21 “Appendix: Periodic Table of the Elements” also lists the atomic masses of the elements.