Atomic Mass and Atomic Weight
Atoms are not very massive; a carbon atom weighs about 2 × 10−23 grams. Because atoms have so little mass, a unit much smaller than the gram is used. In the current system (adopted in 1960–61) the unit of atomic mass, called atomic mass unit (amu), is defined as exactly 1⁄12 the mass of an atom of carbon-12. The atomic weight of an element is the mean (weighted average) of the atomic masses of all the naturally occurring isotopes. Carbon has two principal naturally occurring isotopes, carbon-12 and carbon-13. Carbon-12, whose mass is defined as exactly 12 amu, constitutes 98.89% of naturally occurring carbon; carbon-13, whose mass is 13.00335 amu, constitutes 1.11%. (There are also small traces of the radioactive isotope carbon-14.) The atomic weight of the element is determined by multiplying the percent abundance of each isotope by the atomic mass of the isotope, adding these products, and dividing by 100. However, isotope abundance is often determined by the medium of the source, solid, liquid, or gas, and the average atomic weight may fluctuate. Thus, for carbon, [(98.89 × 12.000) + (1.11 × 13.00335)]/100 = 12.01115, which is the atomic weight of the element carbon in amu, but because the proportions of the isotopes vary depending on where the carbon is found, carbon's atomic weight is now expressed as an interval defined by the lower and upper bounds within which the atomic weight ranges: [12.0096; 12.0116]. Certain synthetic elements exist only momentarily in the form of a few short-lived isotopes; in such cases the concept of atomic weight cannot be applied.
Sections in this article:
The Columbia Electronic Encyclopedia, 6th ed. Copyright © 2024, Columbia University Press. All rights reserved.
See more Encyclopedia articles on: Chemistry: General
