Introduction

In the early nineteenth century, about 30 known elements were sorted into three broad physical-property classes: metals, nonmetals and metalloids. As more elements were discovered, chemists looked for a numerical basis for classification, and two notable early attempts based on atomic mass were Johann Dobereiner's triads (groups of three chemically similar elements whose middle atomic mass was roughly the average of the other two) and John Newlands' law of octaves (arranging elements by atomic mass and noticing that every eighth element had similar properties, echoing musical octaves). In 1869, the Russian chemist Dmitri Mendeleev arranged the 63 elements then known in increasing order of atomic mass, and used the serial position of each element in this list as its 'atomic number'. He folded this ordered list wherever properties started repeating, producing a table of vertical groups (elements with similar properties) and horizontal series, later called periods, where properties varied gradually from left to right. Mendeleev deliberately left gaps at certain positions to preserve this periodicity, predicting that undiscovered elements would fill them — and when several such elements (like gallium and germanium) were indeed discovered later with closely-matching predicted properties, the scientific community accepted his table and his periodic law: 'the physical and chemical properties of elements are a periodic function of their atomic masses.' Even the noble gases, unknown to and unpredicted by Mendeleev, could later be accommodated by adding one new group. The picture changed after atomic structure was understood: Henry Moseley's 1913 X-ray spectroscopy studies showed that the frequency of X-rays emitted by an element correlates with its atomic number Z (the number of protons) rather than its atomic mass, establishing Z as the more fundamental property. This led to a revision of Mendeleev's law into the MODERN PERIODIC LAW: 'The physical and chemical properties of elements are a periodic function of their atomic numbers.' The periodic table was correspondingly revised into the final, modern long form used today, which the rest of this chapter explores.