This guest post is written by Norman Holden, a Brookhaven scientist in the National Nuclear Data Center and a member of the International Union of Pure and Applied Chemistry (IUPAC). After receiving his Ph.D. in nuclear physics from the Catholic University of America, he spent a decade at the GE Knolls Atomic Power Lab before joining Brookhaven in 1974. He is the chair of an IUPAC subgroup that is producing a periodic table meant to show high school and college students the importance of isotopes in everyday life.

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Norman Holden

When the clock struck midnight on Saturday morning, we rang in a year dedicated to celebrating the wonders and achievements of chemistry.

The wide-ranging events held during the International Year of Chemistry 2011 coincide with the 100th anniversary of Marie Curie’s Nobel Prize and the 100th anniversary of the International Association of Chemical Societies (a forerunner of IUPAC). The worldwide celebration will demonstrate the achievements of a well-established field that continues to foster new discoveries every day. There’s no better evidence of chemistry’s vitality than a new change to one of its most familiar faces – the periodic table.

Recently, the way 10 elements are represented on the periodic table underwent a makeover. The aesthetic change reflects a shift in the way chemists think about atomic weight, an elemental property that was traditionally considered a constant of nature.

In the 19th century, atomic weight values were thought to be invariable numbers, like the speed of light. Scientists ordered the known elements according to their atomic weights and noted a repetition (or periodicity) in their chemical properties as you moved from low to high weights. This was represented by the periodic table of chemical elements.

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An example of a traditional periodic table (Click to enlarge. Credit: NIST)

Things got more complex at the start of the 20th century, when radioactive elements were discovered. Scientists concluded that these newly discovered materials were chemical elements with different radioactive properties and different atomic weights – meso-thorium (Ra-228) and radium (Ra-226), for example. But because the radioactive elements maintained the same chemical properties, they were placed in the same position in the periodic table. The word “isotope” (Greek for “in the same place”) was coined for the radioactive species.

Isotopes are made of protons and neutrons. The number of protons in an isotope determines the chemical element – hydrogen has one proton, while gold has 79. The neutrons and protons taken together determine the atomic mass of the isotope.

The way chemists viewed atomic weights was profoundly affected when the rare gas neon was found to be made up of two stable isotopes, 20Ne and 22Ne. The discovery of stable isotopes meant that an element’s atomic weight could be determined from the product of the mass and the relative percentage of each stable isotope in an element. This formula is complicated by the fact that the relative amount of each isotope can depend upon the source of an element’s sample.

Because small variations in an element’s atomic weight can greatly impact trade and commerce, an international committee was established to obtain agreement on the best value for each atomic weight value. Formed at the end of the 19th century, this committee is now part of IUPAC and is known as the Commission on Isotopic Abundances and Atomic Weights (CIAAW).

CIAWW members didn’t have an easy job.

In 1908, for instance, the atomic weight of “common” lead (from a non-radioactive source material) was measured to be 207.2. In 1914, a measurement of lead from thorium silicate had a value of 208.4, and a low value of 206.4 was measured for lead in a uranium sample. Similar variations were found for oxygen and carbon.

By mid-century, scientists discovered that variations in the number of stable isotopes in sulfur impacted the internationally accepted value of its atomic weight. To indicate the span of values that might apply to sulfur from different natural sources, the value ± 0.003 was attached to its atomic weight. Natural variations were added for six elements and experimental uncertainties were added for five others.

The CIAAW added uncertainties for all atomic weight values by 1969. The commission noted that isotopic mixtures, rather than a single stable isotope, “represent the normal and not the exceptional state of an element.” (Note, though, that 21 elements do have only one stable isotope and their atomic weights are constants of nature whose values are known to better than one part in a million.)

To determine the atomic weight of an element with a span of values, the CIAAW traditionally deemed the median value as the standard atomic weight and assigned an uncertainty to encompass all of the published values. But this presentation created confusion.

The uncertainty value of the standard atomic weight is often misinterpreted as a measurement uncertainty, causing people to wrongly question why atomic weight values cannot be determined more accurately. In addition, a standard atomic weight value is expected by readers to reflect a Gaussian distribution (or a bell-shaped curve). It does not reflect bimodal distributions of elements such as boron and sulfur, which have two main separated values and no samples in between. For these and other reasons (presented in this Chemistry International article), a change in presentation was considered necessary.

As a result, the atomic weights of 10 elements – hydrogen, lithium, boron, carbon, nitrogen, oxygen, silicon, sulfur, chlorine, and thallium – are now shown on the periodic table as intervals with upper and lower bounds rather than as values and uncertainties. Each element is also now associated with a pie chart, which shows the number and abundances of each isotope.

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Illustrations for elements in IUPAC’s new isotopic periodic table for the educational community with isotopic abundances shown as pie diagrams. a.) Element (chlorine) whose standard atomic weight is not a constant of nature and is an interval. b.) Element (mercury) whose standard atomic weight is not a constant of nature and is not an interval. c.) Element (arsenic) whose standard atomic weight is a constant of nature because it has a single stable isotope. d.) Element (americium) that has no stable isotopes and thus no standard atomic weight.

An IUPAC task group is now preparing a special periodic table to introduce teachers and students to the concept of isotopes, their impact upon atomic weights, and their importance in everyday life, ranging from applications like treating cancer to detecting sports doping. This new table is expected to be available later this year.

As we officially let go of the 19th-century idea of constant atomic weights, it seems like a more than appropriate time to launch the International Year of Chemistry and 21st-century research.

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