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 1/12 of the mass of a ¹²C atom:

1 u=112 the mass of ¹²C 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 ¹H is 1.008 u, the mass of an atom of ¹⁶O is 15.995 u, and the mass of an atom of ³²S 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% ¹⁰B and 80.1% ¹¹B. 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% ¹²C and about 1% ¹³C, 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.