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In honour of the International Year of Chemistry, Dimensions takes a peek behind recent changes to the periodic table — that marvel of simplicity and scientific achievement that represents the building blocks of chemistry.
Remember the periodic table with its neat rows and columns of elements grouped into alkali metals, lanthanides, and noble gases? Maybe you had to memorize the order of the elements, or even their atomic weights?
Well, that periodic table you struggled to learn in high school has changed. This in itself is nothing new; the periodic table changes regularly, as newly created elements are added and atomic weights are updated.
What is atomic weight?
Atomic weight represents the relative average mass of all isotopes of a given element. For example, bromine, with its two almost equally abundant isotopes weighing 79 and 81 daltons (the unit of atomic mass), has a relative average mass of 79.904 daltons on Earth.
But in 2010, something unprecedented happened. The standard atomic weights of 10 elements morphed from a single number to a low and a high value, known as an interval.
The update is meant to better reflect how these elements vary in natural substances. For instance, the atomic weight of oxygen is slightly greater in air than in seawater.
It came as a surprise to many that atomic weights — long presented as constants — can actually vary in nature. “There are even chemists who are surprised that standard atomic weights are not constants of nature,” says Dr. Tyler Coplen of the U.S. Geological Survey.
Learn the periodic table in 60 seconds
To learn more about the periodic table, including why it's arranged the way it is, visit NRC's periodic table of the elements.
The new atomic weights will soon appear in chemistry textbooks worldwide. But where did this change come from? Who decided that the atomic weights of these elements should change from a single value to an interval, and why?
It turns out that the change, though based on science, came down to one very important element — the human element.
The world authority on chemistry standards is the International Union of Pure and Applied Chemistry (IUPAC), based in the United States.
IUPAC has committees of experts who, among other things, approve the names of the elements and who recommend terminology and symbols in chemistry, “making sure that chemists worldwide speak the same language,” says Dr. Coplen.
One committee within IUPAC is responsible for atomic weights. Called the Commission on Isotopic Abundances and Atomic Weights, this select group of 12 chemists from around the world meets every two years to review the latest developments in atomic weight measurement science.
The Commission has a history going back more than 100 years. Its members have included Nobel Prize-winning scientists such as Marie Curie, who discovered radium and polonium, and Frederick Soddy, who along with Ernest Rutherford helped to unlock the secrets of radioactivity.
An elemental change
The decision to represent the atomic weights of 10 elements as intervals was by no means an easy one.
Throughout the table’s history, the atomic weights had always been presented as a fixed number, thus giving the impression to generations of students and teachers that they were constants of nature.
But the Commission wanted to make clear to the public that atomic weights are variable. The question was how?
“That was the most intense discussion that I had ever seen since I joined the Commission in 1985,” says Dr. Coplen, who co-authored the formal “Atomic Weights of the Elements 2009” report. “Some people wanted to maintain what we had but explain it more clearly. There are all sorts of ways you could go.”
Although the use of intervals would cement the fact that atomic weights of many elements vary significantly in nature, there were concerns over how educators would work with the new interval values.
For example, how would students calculate the molecular weights of substances such as carbon dioxide without a defined standard atomic weight for either carbon or oxygen?
“The debate was very intense and charged,” says Dr. Juris Meija, a chemist at NRC and also a member of the Commission. “You can imagine 12 people sitting in a room, and you have six people saying this is a great idea and six people saying it’s a terrible idea. So where do you go from there?”
In the end, the fate of the periodic table was decided by a close vote. “You can think of it as a five to four Supreme Court decision,” says Dr. Meija.
Why so much attention to atomic weight? Does it really matter whether the atomic weight of hydrogen is, say, 1.00777 or 1.00756?
Probably not to most people, but in the 1920s, that discrepancy led to the discovery of deuterium — the isotope in heavy water. That discovery made possible the Manhattan project and the nuclear bomb, which soon altered the outcome of World War II. “That’s all from the fourth decimal place of hydrogen,” says Dr. Meija.
Atomic weight has many practical uses, from detecting doping in athletes (synthetic testosterone betrays itself by the lower atomic weight of its carbon) to determining the age of the Earth (the precise atomic weight of lead and uranium changes over time due to radioactive decay). The element strontium is used to verify the provenance of food products such as wine: scientists can tell where the grapes were grown by the atomic weight of strontium in the wine.
Atomic weight also tells the story of human impact on the environment. After the 1986 Chernobyl disaster, radioactive cesium was deposited in southern France. Researchers now can tell whether a bottle of wine from that region dates from before or after 1986 by the presence or absence of heavy cesium isotopes.
A notorious past
Controversy is nothing new in the history of the periodic table. Element 118 (ununoctium) is rather infamous, having been dropped from the table in 2002 when it was revealed that a scientist had fabricated data supporting its discovery. It has since been reinstated with more reliable references.
Hydrogen is another controversial member: due to its unique chemical properties, no one can agree on its correct position. Some tables place it atop the noble gases, others above the alkali metals. In a kind of compromise, some tables plunk it above all the elements, “like the mother of all,” says NRC’s Dr. Juris Meija.
The naming of elements is especially contentious. In 1994, when the discoverers of element 106 proposed that it be named after American chemist Glenn Seaborg, the name was initially rejected on the basis that he was still alive. This led to years of controversy, during which Dr. Seaborg wryly commented that IUPAC had refused to accept the name, “because I'm still alive and they can prove it, they say."
In the end, “seaborgium” was confirmed in 1997, two years before Dr. Seaborg died, thus allowing him to enjoy his own immortality in elemental form.