Assembling the Modern Periodic Table

The periodic table is an elegant demonstration of properties of elements. You can determine the electron configuration of any atom, simply from its place. You can compare electronegativity, ionization energy, atomic radius, chemical reactivity, and more.

If I gave you all of the elements on cards and told you to recreate the periodic table, you probably wouldn’t have much trouble. You would order them by increasing atomic number and create a new row when you hit a noble gas. If you know about atomic numbers and electron shells, recreating the periodic table is simple. However, the periodic table predates knowledge of atomic numbers and subatomic particles (yes, including electrons). It even predates knowledge of the noble gases.

So how did Russian chemist Dmitri Ivanovich Mendeleev and the other creators of the periodic table (arguably six of them) bring order to the elements? How did they create a tool that would ultimately house 118 elements when they knew only 62 of them? And why does Mendeleev get all the credit?

Mass + reactions = periodic table

The modern periodic table didn’t spring fully formed from the genius of Mendeleev; it was shaped by key discoveries about the elements. One such discovery was that of atomic masses. Here is where we will begin our journey to periodicity. The modern definition of atomic mass (the weighted average of the atomic masses of all isotopes of an element) was meaningless 150 years ago. Chemists didn’t know about isotopes. In fact, many chemists held the view that atoms were the smallest units of matter possible. It would be 30 years after Mendeleev’s periodic table that scientists found out atoms were composed of smaller bits and pieces. The idea of isotopes wasn’t introduced until 1913, and neutrons weren’t discovered until 1932. 

So how did chemists of the 19th century define atomic mass? In 1803, English scientist John Dalton published an article in which he assigned hydrogen a weight of 1, and then used compounds of hydrogen to determine the relative weights of the other elements. For example, to determine the atomic mass of oxygen, he used the fact that 1 gram of hydrogen reacts with 8 grams of oxygen to make water. He then could use this ratio of 8:1 to determine the weight of oxygen compared with that of hydrogen. 

The problem with this method is that Dalton didn’t know the numbers of oxygen atoms and hydrogen atoms in a water molecule; he assumed water had a molecular formula of HO, leading to an incorrect relative atomic mass of 8 for oxygen.

To add even more confusion, in 1805 Prussian scientist Alexander von Humboldt and French scientist Joseph Louis Gay-Lussac determined that two volumes of gaseous hydrogen always combined with one volume of gaseous oxygen to form two volumes 

of water vapor. The pair found numerous other simple ratios, resulting in Gay-Lussac suggesting that equal volumes of gases have equal numbers of particles, what we now refer to as Avogadro’s law. The problem with this hypothesis was that for it to be true, somehow the gaseous oxygen had to be splitting in half. Many chemists, including Dalton, considered this possibility absurd: how could an atom—at the time believed to be the smallest unit of matter—split during the course of a chemical reaction? 

The mystery was solved in 1811 by Italian scientist Amedeo Carlo Avogadro, who argued that gaseous oxygen is composed not of atoms of oxygen but of molecules of oxygen: O₂. Unfortunately, although he was an accomplished scientist, Avogadro was not an accomplished writer, and his hypothesis was not accepted for another 50 years. 

The facts in two dimensions

The next scientist of mention on our road to periodicity is German chemist Meyer. Meyer’s epiphany occurred at the Karlsruhe conference when he learned of the work on atomic masses by Cannizzaro. He wrote that when he read Cannizzaro’s article, “the scales fell from my eyes and my doubts disappeared and were replaced by a feeling of quiet certainty.”

Meyer’s breakthrough was presented in his textbook Modern Theories of Chemistry and Their Significance for Chemical Statics in 1872. In his table, Meyer organized the elements according to their atomic masses and valences, the latter of which had been discovered in the 1850s.

Meyer accounted for two important features that are usually attributed only to Mendeleev: he reversed the order of tellurium and iodine, and he left gaps. Without atomic numbers, the placement of tellurium (atomic number of 52) and iodine (atomic number of 53) in the periodic table can be confusing. In order of increasing atomic mass, iodine, with a weight of 126.9 amu, should come before tellurium, with a mass of 127.6 amu, except that such an ordering doesn’t make sense when you consider their properties. Iodine is chemically more like chlorine and bromine, whereas tellurium is chemically more like selenium and sulfur. In constructing his table, Meyer decided that properties should override masses, and he put tellurium before iodine. 

The second distinguishing characteristic of Meyer’s table is that he left gaps in it. Other scientists of the day tried to eliminate gaps in their tables, often by forcing elements into illusionary categories, but Meyer simply left blank spots in his. While he didn’t go so far as to predict the properties of then-undiscovered elements, he left a gap between silicon and tin, for example, that would later be filled by germanium.

Interestingly, Meyer regarded periodicity and the similarities among elements in groups as evidence that elements were composed of smaller, more fundamental particles, an idea that Mendeleev himself never accepted.

Werthig is valence. The valency of an element was originally a measure of its combining power with other atoms when it forms chemical compounds or molecules. The concept of valence developed in the second half of the 19th century and helped successfully explain the molecular structure of inorganic and organic compounds.

*"J" in the table above is for iodine.