An international governing body has adopted a new definition of atomic mass (aka atomic weight) changing from specific values to intervals of masses to resolve 15 years of debate on one of the most fundamental of scientific concepts. In a list that only singer-comedian Tom Lehrer could love, hydrogen, lithium, boron, carbon, nitrogen, oxygen, silicon, sulfur, chlorine and thallium now all have new mass definitions.
"Back in high school, you opened your chemistry book and saw a table of standard atomic weights. Your teacher probably told you these were constants of nature. It turns out that that is not true," says Tyler Coplen of the U.S. Geological Survey, co-author of the document that redefines the masses of the aforementioned elements. Swapping the ease of a single number for the apparently clumsier interval of numbers has been controversial among some chemists, but additional precision in measurements in recent years has shown the change is necessary for accuracy. After an arduous 18-month approval process, the new definitions became formal with their publication December 12 in Pure and Applied Chemistry.
Some elements have more than one stable (nonradioactive) isotope—variants of the same substance, but with different numbers of neutrons in their atomic nuclei that alter the mass. (The element's identity is determined by the number of protons.) In those cases the atomic mass listed on the periodic table has traditionally been defined as an average depending on how common each isotope is in nature. But that average is not the same every place on Earth, or in every situation. Just as radiocarbon dating can place a substance in time, isotopic analysis can also pinpoint its location.
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So, now instead of carbon listed as being 12.0107 atomic mass units with a measurement uncertainty of about 0.0008, it has an official atomic weight of [12.0096; 12.0116], where the brackets and semicolon indicate an interval of atomic weights. The interval doesn't reflect an uncertainty in measurement precision but rather a real variation of atomic weight from substance to substance. Only 10 elements will have these new intervals, because the others have at most only one stable isotope or because upper and lower bounds have not been quantified.*
Measurement techniques today are so sensitive that samples of food, water or pollutants can often be traced to their sources based on isotopic abundance. Scientists regularly use these techniques in oceanography, geology, ecology and forensic science. "When I was watching CSI: New York the other night, I saw they were going to analyze a pond water sample for the oxygen isotopes. They were essentially determining the atomic weight and using the fact that the atomic weight varies from place to place," Coplen says.
Although scientists might find this change useful, what about chemistry students who are struggling enough without having to deal with an interval of masses instead of just a number? Coplen sees it as an opportunity for education. "It will enable teachers to introduce students to things called stable isotopes, which is great because we're made up of stable isotopes, and that ratio varies depending on what we eat."
Beginning chemistry students who haven't yet heard about isotopes might have a more difficult time with the concept, and Coplen says he simply doesn't know what the best solution is, but that the problem is now in the hands of the International Union of Pure and Applied Chemistry's chemical education group.
The reeducation goes well beyond students. Coplen says, "One of the people in IUPAC said to me that he was talking with many other chemists and they did not know that atomic weights actually varied. Maybe for lots of people this is not something they've thought about, and it's going to catch them by surprise. In that case I think the commission has done something very useful."
Chemist Matt Hartings of American University in Washington, D.C., says that most chemists who don't think regularly about isotopic abundances or teach general chemistry at college would probably consider the atomic weights as constants. He said, however, most working chemists when pushed to consider the issue would understand the variation exists.
IUPAC also sets the international standards about elements for trade and commerce and so has defined a set of values for those uses when the specific atomic weight is not known. Those values are similar to the previous standards but are typically less precise to take into consideration the variation in atomic masses. For example, carbon now has a suggested atomic weight of 12.011.
*Editor's Note (12/23/10): This sentence was edited after posting. It originally stated that the other elements only have one stable isotope. The text was also edited throughout to correct statements that identified an element's atomic weight as atomic mass.