The periodic table, that ingeniously arranged display of atomic weights, numbers and elements that’s probably posted on a door way in the back of your lab, isn’t a static document. It’s changed a fair amount in its lifetime, largely as weights are refined and new elements added.

However in 2011, the International Union of Pure and Applied Chemistry (IUPAC), the U.S.-based organization in charge of all things periodic table, made a big change to a few elements (the paper making the announcement was published this year). Instead of the single number for atomic weights (you probably had to memorize these once), IUPAC ordered that a range of atomic weights be displayed for 12 elements (Hydrogen, lithium, boron, carbon, nitrogen, oxygen, magnesium, silicon, sulfur, chlorine, bromine, and thallium).

 Why do this?

Because atomic weights are not necessarily a constant. In fact, as early as 1951, IUPAC noted that sulfur atomic weights varied according to the element’s (natural) sources. In addition, oxygen has a different atomic weight in seawater than in air, and the slight differences in atomic weights of hydrogen led to the discovery of deuterium (which then led to the development of the atom bomb, among other things). And many others, like carbon, have naturally occurring isotopes with differing atomic weights. So, IUPAC decided that for at least these 12 elements, they would offer an “interval,” with a high and low value for atomic weight.

The IUPAC committee responsible for these things, called the Commission on Isotopic Abundance and Atomic Weights, also changed the standard atomic weights for germanium, indium, and mercury. And, they’ve approved the names and symbols of elements 114 (flerovium, symbolized by “Fl”) and 116 (livermorium, symbolized by “Lv”).

 Where did the table come from, anyway?

Most history texts credit the Russian chemist and inventor Dmitri Mendeleev, who introduced his version of the table in 1869, in an attempt to corral what we knew about the 56-or-so elements known at the time. However, as the Royal Society points out, many other scientists had been trying to make organizational sense of the elements. Nearly a century before Mendeleev’s presentation, Antoine Lavoisier organized the elements into metals, non-metals, gases, and “earths.” And, just a few years before Mendeleev’s table, British chemist John Newlands introduced his “Law of Octaves,” in which the element’s atomic weights were arranged in seven distinct periods. He was almost right; the eighth period would be included after the discovery of noble gases. Meanwhile, German scientists Lothar Meyer noticed patterns between atomic valences and weights and created his own table of 28 elements in 1968. Unfortunately, he didn’t get the table published until 1870, which allowed Mendeleev to steal most of Meyer’s thunder (an example of publish or perish, 19th-century style).

One sticky little problem with Mendeleev’s table persisted, however; arrangements by atomic number. A few elements, like iodine, had atomic weights that didn’t quite “fit” in order on the table like they should. British chemist Henry Moseley, using a new x-ray beam “gun,” found a way for the first time to accurately measure atomic numbers. Then, these “renegade” elements fell into place (tragically, Moseley enlisted as a soldier in World War I, and was killed in 1915).

 Why not atomic mass?

One term that probably won’t change anytime soon is the term itself; “atomic weight.” In 1992 (and the debate rages on today), IUPAC decided that they would continue with the term “weight,” since atomic weight is actually a relative measure and not a true measurement in grams, pounds or otherwise. But most important, chemists have always used the word “weight” when referring to elements, and were, in IUPAC’s view, always comfortable with the term!

We’ve come a somewhat long way since Mendeleev, but we’re virtually guaranteed more changes to the periodic table. At least you’ll have something new to post on your lab refrigerator!


Royal Society periodic table pages

Wieser, M.E., et al. (2013) Atomic weights of the elements 2011. Pure Appl. Chem  85(5), 1047–1078.

de Bièvre, P., and Peiser, H. (1992) ‘Atomic weight’: The name, its history, definition, and units. Pure Appl. Chem., 64, 1535-1543.