Before 1961 there were two definitions for atomic mass unit. The elements defined the atomic mass unit as ‘1/16th of the atomic weight of oxygen’. Apart from the fact that the term mass should have been used instead of weight, this definition does not specify which oxygen isotope, and implies the average mass of an oxygen atom in the mixture of 16O, 17O, and 18O existing in nature. The relative abundances of these three isotopes (approximately 99.762%, 0.038%, and 0.200%, repsectively) vary in different natural substances (water, air, minerals, etc.). As a result, the chemical definition was not very precise. Physicists, desirous of greater precision, defined their atomic mass unit as 1/16th of the mass of the oxygen isotope 16O. This made their atomic mass unit about 0.027% smaller than the chemical one. The symbol for both was amu (atomic mass unit). This was of course very confusing, and in 1961 the International Union of Pure and Applied Chemistry (IUPAC) decided to redefine the...
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Bibliography
Emiliani, C. (1990) Avogadro's number. Chem. Eng. News, 68, 3, 34.
Emiliani, C. (1991) Avogadro's number and mole: a royal confusion. J. Geol. Educ., 39, 31–3.
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Emiliani, C. (1998). Atomic mass unit, avogadro constant and mole. In: Geochemistry. Encyclopedia of Earth Science. Springer, Dordrecht. https://doi.org/10.1007/1-4020-4496-8_18
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