At the same time Mendeleyev, who was engaged in writing a chemistry textbook that later became world-famous, reached a conclusion similar to that of Meyer but based instead largely on the chemical properties of the elements. He also left blank spaces where elements were obviously missing, and he went well beyond Meyer in his published tables. He predicted in detail what the chemical and physical properties of the missing elements would be when they were found. Mendeleyev's predictions were soon confirmed and new atomic weight determinations corrected the values he had questioned, and the discovery of the actual elements gallium, scandium, and germanium showed that these had almost exactly the properties that he had anticipated.
From that time on, the periodic table assumed the basic form that, in spite of some modifications, it has retained ever since. Hydrogen, which was recognized as an anomalous element, was placed by itself at the beginning of the chart. The periodic table can be read either horizontally or vertically. If horizontally, the elements are arranged in a series not only by atomic weight, but also by ‘valence’. Originally there were seven vertical columns below each element in the first period, designated by the Roman numerals I to VII. Further down the table appeared the ‘transition elements’ (iron, cobalt, and nickel). Because the properties of these transition elements set them apart, another column, VIII, was designated for them.
When the table is read vertically, each column comprises a family of elements having similar properties. In this form the table was able to meet the needs of inorganic chemists for the organization of what had previously been a mass of uncoordinated facts. No theoretical reasons for its existence have ever been given.
The first serious challenge to the neat organization of the periodic table occurred in 1894 when argon was discovered. No similar unreactive gases were then known. To fit the substance into the table, Mendeleyev proposed adding a new vertical column to the table. Since the gas formed no compounds, it had no valence and so could take its place before the elements of column I. This, of course, implied that other similar inert gases must also exist. With the table as a guide, the other members of the ‘noble-gas’ family (helium, neon, krypton, and xenon) were quickly found.
At the end of the 19th and the beginning of the 20th century great progress was made in clarifying the structure of the atom itself. It was quickly realized that the atomic number represented the positive charge of the atomic nucleus, and that this charge increased by one as the elements advanced along the table. The atomic number was therefore a more fundamental value than the atomic weight. In fact, it was found that the atomic weight was not a unique value for each element. Studies of radioactive elements and their end products showed that elements from different locations sometimes had different atomic weights, yet showed identical chemical properties. These were named ‘isotopes’.
Gilbert Lewis and Irving Langmuir between 1916 and 1920 pointed out that the chemical properties of an element are determined by the number of electrons in the outermost shell. The noble gases do not normally form compounds, because their outermost shells are filled. Other elements tend to form the ideal number to attain this maximum, either by losing electrons or by gaining electrons.
