3.1 Development of the Periodic Table and the Modern Periodic Law

By 1860 about sixty elements were known, and new ones were being found every few years with the help of the spectroscope. Each had its own atomic mass, density, melting point, valence and set of compounds. Learning them one at a time was hopeless; what chemists needed was an organising principle — a way of grouping elements so that knowing one member of a family tells you a great deal about the others, and so that the gaps in the pattern point to elements still waiting to be discovered.

This section follows that search. It begins with families of three (triads) and repeating eights (octaves), reaches its turning point in Mendeleev's table of 1869, which was bold enough to leave gaps and to correct accepted atomic masses, and ends with the discovery that the true ordering number is not the atomic mass but the atomic number $Z$. That gives the modern periodic law and the long form of the periodic table used today. Finally, you will learn how the newest, heaviest elements are named before they receive permanent names.

Early Attempts: Triads, Octaves and Atomic Volumes

The first clue came in 1829 from the German chemist Johann Döbereiner. He noticed that some elements with very similar properties come in groups of three, and that the middle member has an atomic mass close to the average of the other two.

In 1865 the English chemist John Newlands arranged the elements then known in order of increasing atomic mass and noticed a rhythm.

Beyond calcium the octaves break down: to keep the rhythm Newlands had to put dissimilar elements together (iron, a metal, fell in the same row as oxygen and sulphur) and to squeeze two elements into one place. The noble gases were not yet known; when neon and argon were found, the "eighth" element became the ninth, so the rule could not survive. Newlands' work was ridiculed at first, but in 1887 the Royal Society of London honoured him for it.

Around 1869 the German chemist Lothar Meyer plotted a physical property — the atomic volume, i.e. molar mass divided by density — against atomic mass. The curve rose and fell in a repeating way: the alkali metals sat at the peaks, and elements with similar properties occupied similar positions on the curve. Meyer had shown clearly that properties are periodic, repeating at regular intervals, and he produced a table very like Mendeleev's. Mendeleev published first, and went much further.

Mendeleev's Periodic Table

The Russian chemist Dmitri Mendeleev (1834–1907) arranged the $63$ elements known in 1869 in order of increasing atomic mass, in rows and columns, so that elements with similar properties — especially similar formulae of their oxides and hydrides — fell in the same vertical column (group).

What made Mendeleev's table so successful was that he did not follow atomic mass blindly. He trusted chemical similarity more than the measured masses, which were often unreliable at the time. This led him to three bold steps:

1. He left gaps for undiscovered elements and predicted their properties from those of their neighbours. He called them eka-boron, eka-aluminium and eka-silicon (Sanskrit eka = one: "one place beyond" boron, aluminium, silicon). Scandium (1879), gallium (1875) and germanium (1886) were later found with properties remarkably close to his predictions.
2. He corrected atomic masses. Where an element did not fit, he argued that its valence, and hence its atomic mass, had been wrongly assigned — beryllium and indium are famous examples (Example 4).
3. He reversed a few pairs when similarity demanded it: tellurium ($127.6$) was placed before iodine ($126.9$), so that iodine fell with the halogens and tellurium with sulphur and selenium.

How could anyone predict the properties of an element nobody had seen? Mendeleev's method was essentially interpolation: a property of the missing element was estimated from the same property of its neighbours above and below in its group and to the left and right in its row. Interactive 1 lets you repeat his reasoning with modern data.

The Modern Periodic Law and the Long Form of the Table

The explanation came from physics. In 1913 the English physicist Henry Moseley measured the frequencies $\nu$ of the characteristic X-rays emitted by different elements. A plot of $\sqrt{\nu}$ against atomic mass was irregular, but a plot of $\sqrt{\nu}$ against the position number of the element in the table was a perfectly straight line. That position number is the atomic number $Z$, the charge on the nucleus — the number of protons (Section 2.1). Atomic number, not atomic mass, is the fundamental property of an element; and since $Z$ also equals the number of electrons in the neutral atom, it fixes the electronic configuration, which decides chemical behaviour.

Ordering by $Z$ removes every anomalous pair at a stroke: $\mathrm{Ar}$ ($Z=18$) comes before $\mathrm{K}$ ($Z=19$), $\mathrm{Co}$ ($27$) before $\mathrm{Ni}$ ($28$), $\mathrm{Te}$ ($52$) before $\mathrm{I}$ ($53$). All isotopes of an element share one $Z$ and therefore one place. The table built on this law, and recommended by IUPAC, is the long form of the periodic table.

Naming Elements with Atomic Numbers above 100

The heaviest elements are made a few atoms at a time in particle accelerators, and their discovery takes years to confirm. Traditionally the discoverers proposed a name, which led to disputes when two laboratories claimed the same element. Since 1978 IUPAC has used a temporary systematic name and three-letter symbol built directly from the atomic number, used until a permanent name is approved.

Exercises