Yesterday, the University of Calgary announced that there will be a significant revision to the periodic table, a phrasing that implies a radical upheaval of our understanding of basic matter. The reality is quite a bit less dramatic—so tame, in fact, that the publication that announced the revision was released on Sunday without causing any disturbance. But the announcement provides a good opportunity to give everyone a refresher on the whole concept of atomic weight.

A good periodic table (like this one) will typically have two numbers associated with each element. The first is the atomic number, which is the number of protons in the nucleus, and thus the number of electrons present when an atom isn't ionized. Since there are no fractional electrons, these numbers are integers. These electrons dictate the element's chemical properties, so they tell us something about its behavior.

The second is the atomic weight. This tells us, in grams, how much a mole (6.02 x 1023) of atoms will weigh, and is thus proportional to the weight of a single atom. These numbers are not integers, and that's not because atoms come in fractions. Instead, atoms have different isotopes that contain different numbers of neutrons in the nucleus. A pure isotope will have a well-defined weight, but in the natural world, most elements appear as a mixture of isotopes. A mass of hydrogen, for example, will mostly contain atoms with a single proton and no neutrons, but will have a few with one or two neutrons mixed in. As a result, when you have over 1023 atoms of hydrogen around, these heavier isotopes ensure that it weighs a touch more than just a gram.

So, it's possible for the atomic weight to change simply because we have a better idea of the typical isotope ratios found in the natural world, and that change wouldn't reflect any new knowledge of the atom's internal structure.

The changes announced this week, however, are a mix of new knowledge and a bookkeeping decision. The new knowledge is a better understanding of how isotope mixtures can differ based on the history of the material. This is easiest to understand in terms of carbon. Geological processes generally don't care what isotope of carbon gets incorporated into rocks or dissolved into the oceans. But biological processes are heavily biased towards the lighter 12C isotope. So, if you're looking at a bacteria-rich sediment, the atomic weight of carbon will be very different than if you looked at other sources. Carbon's not alone; other elements also have isotope ratios that differ based on their natural histories.

The changes to the atomic weights reflect this difference. The International Union of Pure and Applied Chemistry has decided to provide atomic weights as a range that reflects the differences found in different sources. The first set of ranges have been made available for 10 elements, including common ones like hydrogen, carbon, and oxygen, and they're pretty narrow (two decimal places in or less), so this shouldn't change that much in practical terms.

Next year is the International Year of Chemistry, though, so more announcements of this sort might be afoot.

Pure and Applied Chemistry, 2010. DOI: 10.1351/PAC-REP-10-09-14 (About DOIs).