Just as the weight listed on your driver’s license doesn’t necessarily reflect your actual poundage, the official atomic weights of most chemical elements are actually more like ballpark estimates than precise constants. In acknowledgment of this natural variation, the official weights of 10 chemical elements will no longer be expressed as single numbers, but as ranges. The adjustments, published online December 12 in Pure and Applied Chemistry, are the first in an overhaul of the atomic weight of almost every element on the periodic table.
Instead of being described by a single fuzzy number, the atomic weights of oxygen, hydrogen, lithium, boron, carbon, nitrogen, silicon, sulfur, chlorine and thallium will now be expressed as intervals. The change, long overdue, explicitly acknowledges the fact that most of the 118 elements come in multiple forms of varying weight.
Most elements have a preferred, energetically stable form that dominates in nature. For example, oxygen, the most abundant element in the Earth’s crust, is most comfortable having eight neutrons and eight protons in its nucleus (the latter of which defines it as oxygen). But oxygen can gain an extra neutron or two, changing the element’s weight (electrons are also variable but so light that their weight isn’t taken into account). These heavier versions, or isotopes, have been presented as existing in constant quantities no matter the source. For example, it’s commonly said that more than 99 percent of oxygen is the normal eight-neutron isotope — called oxygen-16 — while the heavier versions exist in fractions of one percent.
But those proportions aren’t set in stone, and the new adjustment to the official weights acknowledges that, says Tyler Coplen, head of the U.S. Geological Survey’s Reston Stable Isotope Laboratory in Virginia.
For example, ratios of the three oxygen isotopes will differ depending on whether the oxygen is in air, groundwater, fruit juice or bone. This variation is what makes isotopes such a powerful scientific tool: the relative ratios of the different carbon isotopes can tell scientists if ivory came from an elephant that ate shrubby savanna plants or woody jungle trees. Similarly, testosterone supplements are plant-derived and have a different isotopic carbon signature than testosterone made by the body (to Tour de France cyclist Floyd Landis’ chagrin).
“Isotope studies extend from studies of previous climates to dating artifacts to weapons programs and biomedical applications,” says James Adelstein, a professor at Harvard Medical School and coeditor of a National Research Council report on isotopes in medicine and the life sciences.
Previously, a given element’s official atomic weight was actually an average of this variation. But as the number of discovered isotopes grew — there are more than 2,000, but only 118 elements — weights kept needing adjustment. These numerical tweaks implied that the numbers couldn’t be pinned down with precision, when in fact such measurements are more precise than ever, says Coplen, who headed the international task force charged with surveying various isotope abundances in nature so that the numbers could be revised.
“It should have been done a decade ago,” Coplen says.
Now that it has completed the initial round, the International Union of Pure and Applied Chemistry’s commission in charge of atomic weights will reassess the rest of the elements in the coming years. Gold, fluorine, aluminum and sodium, each of which exists in only one stable version, will be left alone.
Credit (from left to right): Scewing/Wikimedia Commons; Kuebi/Wikimedia Commons; Mav/Wikimedia Commons
Back Story – Chemistry by committee
Atomic weight refers to the averaged mass of the atoms of a chemical element using a scale based on a standard atomic nucleus. Currently the standard is the nucleus of a carbon atom containing six protons and six neutrons (carbon-12). An atomic weight of 1 corresponds to an average mass equal to one-twelfth the mass of the carbon-12 nucleus. Previous scales were based on hydrogen or oxygen:
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1803 John Dalton (left), an English schoolteacher, compiles the first table of atomic weights for various elements, using units where the atomic weight of hydrogen was equal to 1.
1810 Swedish chemist Jöns Jacob Berzelius (center) begins work on developing an atomic weight scale based on oxygen equal to 100. But most chemists continue to use the scale based on hydrogen equal to 1.
1860 At a conference held in Karlsruhe, Germany, chemists discuss the need for an improved and consistent atomic weight scale.
1869 Listing the elements in order of atomic weight, Russian chemist Dmitri Mendeleyev (right) creates a table in which elements with similar properties fall in the same row (columns in later versions of the table). Mendeleyev’s chart becomes known as the periodic table of the elements. A similar table was created at about the same time by the German chemist Lothar Meyer.
1898 A German committee of chemists recommends basing atomic weights on a scale fixing oxygen equal to 16.
1906 The International Committee on Atomic Weights adopts the oxygen equals 16 scale.
1929 Discovery of heavy forms of oxygen (O-17 and O-18) creates a discrepancy between chemists’ atomic weight for oxygen and physicists’ atomic mass for the oxygen-16 isotope.
1961 Following approval by the International Union of Pure and Applied Physics and the International Union of Pure and Applied Chemistry, atomic weights are based on the carbon-12 nucleus equals 12 scale.
