PERIODIC CHANGE: A suggested presentation for the elements on the periodic table. The abundances of different isotopes are shown as pie charts. Where there is more than one stable isotope, such as for chlorine, the atomic weight is given as an interval. In the case of arsenic there is only one stable isotope, so that atomic weight retains its single value.
Image: Chemistry International
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.
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.
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16 Comments
Add Comment"they did not know that atomic weights actually varied"
Reply | Report Abuse | Link to thisI do sincerely hope that something was lost in the third-hand reporting about these chemists. Perhaps they were expressing surprise that isotopic variation was such a large effect, or that it was widely used for purposes other than absolute dating?
If 'chemists' actually did not know about isotopes, I would have to ask how they became chemists at all.
It seems to me that anyone who needs to know that the atomic mass of chlorine varies from 35.446 to 35.457 in different places, probably needs to know MORE than that. So I'm not sure how useful this is. And I think I'd rather see a number with a +/- than a range.
Reply | Report Abuse | Link to thisI like the pie charts, though. Very informative.
--Dave
@billsmith, I think the article meant that many people didn't know that atomic mass varied significantly from place to place.
This statement, "Only 10 elements will have these new intervals because the others have at most only one stable isotope." is simply wrong. Every element from carbon to lead with an even atomic numbers has at least two stable isotopes, as do helium and quite a few odd-numbered elements. These ten are simply the ones where the source really matters.
Reply | Report Abuse | Link to thisNext these crazy scientist will be telling us Pluto is not a planet...
Reply | Report Abuse | Link to thisThis is moronic. As was explained to me by a general chemistry teacher, there are isotopes and the current atomic weights in the periodic table are an average of the isotopes. Complicating chemistry calculations up by having multiple numbers makes no sense, especially when they are only including a range, which does not specify all of the isotopes! Why would you do this? This is a bastard compromise that will only serve to muddle up teaching the basic elements of chemistry. As slaven41 pointed out, anyone who cares, will be looking up the AMU of the isotope they want to use.
The IUPAC... also responsible for illogical naming conventions. Why would you name the alkene first, if both alkynes and alkenes are primary, yet alkynes take priority everywhere else in a molecule?
IUPAC - thank you for making chemistry less accessible.
Also wrong is the caption that says "In the case of mercury there is only one stable isotope, so that atomic mass retains its single value."
Reply | Report Abuse | Link to thisTo get even pickier at errors in the initial report is the comment, "different numbers of neutrons..." There seems to be evidence accruing that the blindly accepted idea of "neutrons and protons" making up nuclei is a handy "bookkeeping device" but is simply not the case in "Reality." The situation appears to be far more complex and more interesting, with a possibility that the basic units of "nuclei" may be "Deuterons" and "Tritons" rather than Protons and Neutrons....
Reply | Report Abuse | Link to thisQuite deep "beneath the Radar of the General Scientific Community," there seems to be brewing a definite "Revolution" in some science concepts which have not been revised in 70 to 100 years.
There's no reason to get alarmed. The new atomic weight ranges are useful, but only for specialists engaging in very high precision work. These new weights will not at all affect the work of the vast majority of practicing chemists, for whom using a weight of 12.011 for carbon provides more than enough precision (we often get away with simply using 12.!). The simpler, conventional weights are still available and you can still use them. They are listed on Table 6 of the article in Pure and Applied Chemistry.
Reply | Report Abuse | Link to thisI certainly hope that teachers of introductory general chemistry courses won't muddle their teaching by introducing such levels of detail prematurely.
I also got stuck on Mercury sample chart. The example shown could be correct if there is only one stable isotope with a mass of 200.59 but Mercury is most often found in the unstable isotopes shown in the pie chart. This seems unlikely, however.
Reply | Report Abuse | Link to thisI'm no chemist, but I thought this article was clearly written and very helpful to me, with the exception of the Mercury example. Perhaps the author or someone can clarify?
Thanks in advance.
Mercury has seven *stable*, naturally occurring isotopes. The average atomic weight of natural mercury on Earth is 200.59. Arsenic, on the other hand, does have only one stable isotope. Americium has no stable isotopes.
Reply | Report Abuse | Link to thisThanks for the excellent clarification!
Reply | Report Abuse | Link to thisI'm almost 75 and it's a long time since I've studied or used either high school chemistry or college (no! never say "chemistry"!) quantitative analysis--but. . . I do recall that our methods of calculation changed in college. In high school, atomic "weights" (as averages) were all the rage, as was valence. But in college (in the later 1950s) atomic weight and valence were more verboten than those 4-letter words in "f".
Reply | Report Abuse | Link to thisWe used atomic MASS and our calculations were by moles. Isotopes were quite well known, thanks, and sometimes entered into our calculations. So the only surprise is that it took so long to register this knowledge within the periodic table, and that the authors think it's such a big deal.
PS: weight or mass, whole integers or ranges, the periodic table, like the basic harmonic series, is still one of the most ELEGANTLY BEAUTIFUL symbolizations in science!
This is a very good comment. Thanks njcjuhkuusb. It would be interesting to have a copy of that article in Pure and Applied Chemistry.
Reply | Report Abuse | Link to thisThis isn't about making chemistry "accessible". It is about making it accurate. And quite honestly, if a student cannot understand that the quantity of a given isotope can vary then they clearly aren't interested enough or aren't smart enough to pass a quality chem class. What's the point in teaching and conducting science if we have to water down the concepts so much that five-year-olds could understand them?
Reply | Report Abuse | Link to thisI agree wholeheartedly with the sentiments expressed in your postscript - thanks!
Reply | Report Abuse | Link to thisHowever, IMO it is not the symbolic representation of atomic elements that is beautiful but the diversity of natural characteristics produced by the incremental composition of elements and our place in the universe that provides the abundant availability of these elements.
It's not the periodic table that clarifies the complexities of nature but the clarity of elementary nature that makes this table so understandable.
Everybody who has to explain the C14 method in his curriculum, has to explain first how atomic masses can vary. And in our chemistry courses we used most of the time rounded values for atomic masses: 35.5 for chlorine e.g.
Reply | Report Abuse | Link to thisInteresting article. I guess I have to start learning the new masses. The article could be improved with an image of periodic table. This may be useful: http://ygraph.com/chart/1576
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