Atomic Weight Versus Atomic Mass | Chemical & Engineering News
Volume 89 Issue 4 | p. 4 | Letters
Issue Date: January 24, 2011

Atomic Weight Versus Atomic Mass

Department: Letters

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Thank you for bringing to my attention the recent International Union of Pure & Applied Chemistry report “Atomic weights of the elements 2009” (C&EN, Dec. 20, 2010, page 36).” A link to the IUPAC report can be found at http://bit.ly/fmpcuq.

Unfortunately, there are some errors in the article that are also propagated in several general chemistry textbooks. In referring to the atomic weight of boron, the author of the article says, “For example, boron’s atomic mass until now has been given as 10.811 amu. The new tables will list the value as [10.806; 10.821] to reflect the element’s true atomic weight, which can vary over that range depending on the material’s source.”

First, 10.811 should be referred to as the atomic weight of boron, not the atomic mass, as the author correctly states in the article’s title and when discussing the range of atomic weights for boron. Second, listing the atomic weight of boron as 10.811 amu is incorrect. Atomic weights are, by convention, dimensionless, and the term has been used that way since 1808. A good discussion of the arguments for and against the term atomic weight can be found at http://bit.ly/fbeFPd. Third, using the unit “amu” (standing for atomic mass unit) is incorrect, as the correct unit since 1964 is “u” (standing for unified atomic mass unit).

The discovery of the isotopes of oxygen in 1929 led to two different atomic weight tables. Chemists used a table based on the abundance-weighted sum of the atomic masses of the three naturally occurring isotopes of oxygen. Physicists used a table based on only oxygen-16. This situation became untenable when chemists and physicists tried to communicate with each other. The reference for atomic weights was changed to carbon-12 in 1964 and the new symbol “u” replaced “amu.” For more information, see http://bit.ly/cJt6nn.

My reference above to textbook errors is clearly shown in the latest (10th) edition of one of the most widely used general chemistry texts. Since the text is otherwise an exceptional one, and it is not my intent to point fingers but to correct errors, I have not named the text or author. Not all general chemistry texts make these errors, but enough do so that the errors are perpetuated. It might interest readers to know that some scientifically oriented websites also make these errors.

The text I am referring to uses the unit “atomic mass unit” and the abbreviation “amu.” What is properly called the atomic weight is referred to as the “average atomic mass,” and is given the unit “amu.” Thus, the atomic weight of copper is listed as 63.55 amu. In other places in the text, the author refers to the “average atomic mass” as simply the “atomic mass.”

The IUPAC report defines atomic weight as follows: “The atomic weight of an element in a substance is the abundance-weighted sum of the atomic masses of its isotopes.” (Not stated in the definition, but implied by the context, is that only naturally occurring isotopes are included in the sum.) How is a student supposed to keep the distinction of the mass of an isotope of an element clear from the atomic weight of an element using this textbook’s nomenclature?

To clarify things, I suggest that the mass of an atom of a single isotope be called the “isotopic mass” with the unit “u.” The isotopic masses of the two stable isotopes of boron, boron-10 and boron-11, would be written as 10.012937 u and 11.009305 u, respectively. The atomic weight of boron would be written, as the IUPAC paper recommends in Table 6, if one number is needed for say, exams, as 10.81. Thus there will be a clear distinction between the mass of an individual atom and the “abundance-weighted sum,” which is the atomic weight.

Students should be able to distinguish the two with no confusion. “A Table of Isotopic Masses and Natural Abundances” can be found on Wikipedia at en.wikipedia.org/wiki/Isotopes_of_oxygen. Note that the natural abundances are what can vary in different samples and are what limit the number of significant figures in an atomic weight value and give rise to the range of values for certain elements. The exceptions are elements that have only one naturally occurring isotope, such as fluorine, where the numerical value of the isotopic mass of fluorine-19 would be equal to the atomic weight.

The problem with elements that are radioactive is discussed in great detail in the IUPAC report and will not be covered here.

It seems to me that if the international scientific community (as represented by IUPAC) has accepted the dimensionless term atomic weight and the unit of isotopic mass as “u,” then it behooves all textbook authors, scientists, and science writers to use the same term and unit. We must all communicate in the same language.

Harvey F. Carroll
Lake Forest Park, Wash.

 
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