Spectroscopy and the Atomic Number

Dec 01, 2013
Volume 28, Issue 12

The year 2013 represents the centennial of a seminal discovery about our universe. Key data that led to the discovery were obtained from — you guessed it — spectroscopy.

Table I: Dalton’s list of relative masses (1)
Ever since the enunciation of the modern atomic theory in 1803, the element has taken center stage in our understanding of matter. However, at the very beginning, elements were characterized by their atomic mass (1). John Dalton, the originator of the modern atomic theory, correctly assumed that hydrogen was the lightest element and so assigned it a relative mass of one. Masses of other substances were determined as a ratio to that of hydrogen (see Table I). Given the rather early state of the chemical sciences in the early 19th century, Dalton made several errors in his mass determinations, nor did he recognize that elemental hydrogen was a diatomic molecule. Nonetheless, the concept was sound, and it continues today with much more precise definitions of mass and instrumental techniques.

Figure 1: De Chancourtois’s organization of the known elements (as of 1862), based on atomic masses. (This media file is in the public domain in the United States [PD-US, PD-1923].)
As the discovery of elements occurred through the 1800s, scientists noted similarities in the properties of the elements and began to group them together according to those similarities. According to Scerri (2), the first person to publish a chart in an attempt to systematize elements by properties was Alexandre-Emile Béguyer de Chancourtois, a French geologist. Figure 1 shows de Chancourtois's "periodic table" (my label, not his) (3). This figure, meant to be displayed on the surface of a tube so that the lines plot out a spiral, used more accurate measures of atomic mass that came out in 1858, right before the famed Karlsruhe Conference on chemistry in 1860. The work by de Chancourtois was not spotted by chemists because he published mostly in the field of geology. Over the next several years, contributions to element organization were made by Newlands, Odling, and Hinrichs (4).

In 1862, German chemist Julius Lothar Meyer developed a table that not only had gaps to accommodate potentially new elements, but also included at least one deviation from a strict ordering by atomic mass — tellurium (with a reported mass of 128.3 relative to hydrogen) was listed before iodine (mass of 126.8). Published as part of a chemistry text in 1864 (5), Lothar Meyer apparently felt that chemical properties trumped atomic masses, although an argument could be (and was) made that the measurements of the atomic masses were inaccurate and that atomic masses were still the fundamental organizational principle. Although Lothar Meyer was preparing a second edition of his book to be published in 1868, which would include even more accurate mass data for even more elements, the book was never published and his periodic table did not become well known until later — after he lost a historical priority dispute with Mendeleev.

Figure 2: Mendeleev’s 1871 version of his periodic table. Note the gaps for masses 44, 68, 72, and 100. (This media file is in the public domain in the United States [PD-US].)
In 1869, Russian chemist Dmitri Mendeleev published his version of a periodic table (6), complete with gaps and predictions for the chemical and physical properties of the elements that would fit in those gaps. An 1871 version of the table is shown in Figure 2 (7). Note a bit of revisionist science — in Figure 2, tellurium is given an atomic mass of 125 so that it appears before iodine (whose symbol is J). Mendeleev firmly believed that the atomic mass was controlling, so much so that he argued that the atomic masses of tellurium and iodine were incorrect. Because of Mendeleev's activism, his outlets of publication, and the success of his predictions in filling the gaps, he is widely (and perhaps incorrectly) recognized as the father of the modern periodic table.

There things stood.