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Chemistry Fix

To the Editors:

I enjoyed Sandra J. Ackerman’s article “A Different Kind of CSI: Crime and Stable Isotopes” in the January–February Science Observer section. In mentioning a topic that she says has “tripped up generations of chemistry students,” she also gets tripped up. The sum of the number of protons and neutrons in the nucleus of an atom is not called the atomic mass as she says. It is called the mass number. What a mass spectrometer measures is the atomic mass, which is the mass of the whole atom compared to the mass of an atom of carbon-12. That mass is exactly 12 u, where “u” stands for the unified atomic mass unit. For example, the mass number of carbon-13 is 13, but the atomic mass is 13.00335 u. (This number includes the weight of the electrons, protons and neutrons, and takes into account the binding energy released when the nucleus is formed.) Many textbooks and web sites incorrectly use the old abbreviation “amu” when they are referring to “u.”

The history behind the change from “amu” to “u” is interesting. 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: oxygen-16, oxygen-17 and oxygen-18. Physicists used a table based on only the atomic mass of oxygen-16. This situation became untenable when chemists and physicists tried to communicate with each other. So they decided to unify the atomic weight table in the early 1960s, and the result is the single table that we now use. The reference for atomic masses was changed to carbon-12 and the new symbol “u” replaced “amu.” The term dalton (Da) is sometimes used instead of “u,” especially for heavy molecules such as proteins and polymers.

In discussing the isotopes of hydrogen, Ackerman implies that there are only two, the stable isotopes hydrogen-1 (protium) and hydrogen-2 (deuterium). But there is a radioactive isotope, hydrogen-3 (tritium), which is made in trace amounts in the atmosphere (by cosmic rays colliding with the atoms of the atmosphere), as is the radioactive carbon-14 isotope she previously mentions. Mass spectrometers can be used to measure carbon-14, but I am not aware that they are used to measure naturally occurring tritium, because it is present in only minute amounts with the natural steady-state global mass being about 7.3 kilograms. About five times this amount remains from past atmospheric nuclear weapons tests, but this level will slowly decrease due to the radioactive decay of tritium. Of course, large amounts of tritium have been made in nuclear reactors, and concentrations (say in biological fluids) are typically measured using a liquid scintillation counter to detect the beta particles (negative electrons) that are emitted in the decay of tritium.

Harvey F. Carroll
Professor Emeritus of Physical Sciences
City University of New York
Lake Forest Park, WA

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