2.5 Atomic Masses

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 unitOne-twelfth the mass of a ¹²C atom. (abbreviated u, although amu is also used) is defined as 1/12 of the mass of a ¹²C atom:

$$\text{1 u}=\frac{1}{12}{\text{ the mass of }}^{\text{12}}\text{C atom}$$

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 massA weighted average of the masses of all the element’s naturally occurring isotopes. (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.