The basis of the periodic system

Other chemical and physical classifications

Certain methods of classifying elements on the basis of chemical properties are not strictly related to the groups in which the elements appear. Such classification schemes illustrate the fact that useful horizontal as well as vertical relationships exist in the periodic table. Thus, the transition elements, either as a whole or as three horizontal series, are often considered together when chemical properties are discussed. The transition elements in each horizontal series exhibit much less variation in atomic size than do the elements in other parts of the same periods, leading to a similarity in chemical and physical properties. The lanthanoid and actinoid elements exhibit an even greater similarity for the same reason. The metallic elements in Groups Ia and IIa are often classed together because they are markedly more reactive than the other metallic elements. At the other extreme, elements of the platinum group—including ruthenium, rhodium, palladium, osmium, iridium, and platinum—are chemically inert, as are silver and gold; these elements are collectively designated the noble metals because they do not readily enter into combination with other elements.

Of all the 118 known elements, 11 are gaseous, 3 are liquid, and the remainder are solids under ordinary conditions. With the exception of hydrogen and mercury, the gaseous and liquid elements occur in the right-hand part of the periodic table, the region associated with the nonmetallic elements.

The physical characteristics of the elements provide convenient means of identification. The melting points of the various elements range from −272 °C (for helium) to greater than 3,500 °C (for carbon in the form of diamond). Properties such as boiling points, electrical conductivity, and thermal conductivity also can be used for identification because they are unique for each element.

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chemical bonding: The periodic table

Perhaps the single most useful characteristic for identifying an element is its pattern of light absorption or emission, which is called a spectrum. An element exhibits its own characteristic spectrum whether it exists in the free state, in a mixture, or in chemical combination with other elements. Since the intensity of the spectrum is dependent on the amount of the element contained in the sample, the spectrum also can be used as a means for quantitative analysis of the elements. There are several chemical methods for estimating the percentage of an element present in a sample; these, however, require a detailed knowledge of the chemistry of the element in question (see analysis).

All naturally occurring elements with atomic numbers of 84 or greater are radioactive. In addition, several naturally occurring isotopes of the lighter elements are radioactive. The atomic nuclei of all radioactive elements are unstable and emit highly energetic particles. In the process, the number of protons in the nucleus changes, and the atom is transformed into one of a different element.

The half-life of a radioactive isotope is the time required for half of any amount of the isotope to disintegrate by radioactive decay. The common modes of decay of radioactive isotopes are loss of beta or alpha particles or the capture of an electron. The loss of a beta particle, or electron, from the nucleus increases the atomic number by one unit; the loss of an alpha particle, or helium nucleus (two protons and two neutrons), decreases the atomic number by two units; and the process of electron capture, in which an electron from an inner shell is drawn into the nucleus, corresponds to a decrease of atomic number by one unit.

Elements with atomic numbers greater than 92, the so-called transuranium elements, have been synthetically prepared and are all radioactive. Two radioactive nontransuranium elements—promethium and technetium—were also first produced artificially and, like the transuranium elements, exist in nature (if at all) only in trace amounts. Although the remaining elements generally are not considered to be radioactive, some do have radioactive isotopes that exist naturally in very small concentrations, and more than 1,000 radioactive isotopes of these elements have been prepared in the laboratory.

History of the periodic law

The early years of the 19th century witnessed a rapid development in analytical chemistry—the art of distinguishing different chemical substances—and the consequent building up of a vast body of knowledge of the chemical and physical properties of both elements and compounds. This rapid expansion of chemical knowledge soon necessitated classification, for on the classification of chemical knowledge are based not only the systematized literature of chemistry but also the laboratory arts by which chemistry is passed on as a living science from one generation of chemists to another.

Relationships were discerned more readily among the compounds than among the elements; it thus occurred that the classification of elements lagged many years behind that of compounds. In fact, no general agreement had been reached among chemists as to the classification of elements for nearly half a century after the systems of classification of compounds had become established in general use.

Johann Wolfgang Döbereiner in 1817 showed that the combining weight, meaning atomic weight, of strontium lies midway between those of calcium and barium, and some years later he showed that other such “triads” exist (chlorine, bromine, and iodine [halogens] and lithium, sodium, and potassium [alkali metals]). Jean-Baptiste-André Dumas, Leopold Gmelin, Ernst Lenssen, Max von Pettenkofer, and Josiah Parsons Cooke expanded Döbereiner’s suggestions between 1827 and 1858 by showing that similar relationships extended further than the triads of elements, fluorine being added to the halogens and magnesium to the alkaline-earth metals, whereas oxygen, sulfur, selenium, and tellurium were classed as one family and nitrogen, phosphorus, arsenic, antimony, and bismuth as another family of elements.

Attempts were later made to show that the atomic weights of the elements could be expressed by an arithmetic function, and in 1862 Alexandre-Émile Béguyer de Chancourtois proposed a classification of the elements based on the new values of atomic weights given by Stanislao Cannizzaro’s system of 1858. De Chancourtois plotted the atomic weights on the surface of a cylinder with a circumference of 16 units, corresponding to the approximate atomic weight of oxygen. The resulting helical curve brought closely related elements onto corresponding points above or below one another on the cylinder, and he suggested in consequence that “the properties of the elements are the properties of numbers,” a remarkable prediction in the light of modern knowledge.

Classification of the elements

In 1864 John Newlands proposed classifying the elements in the order of increasing atomic weights, the elements being assigned ordinal numbers from unity upward, and divided into seven groups having properties closely related to the first seven of the elements then known: hydrogen, lithium, beryllium, boron, carbon, nitrogen, and oxygen. This relationship was termed the law of octaves, by analogy with the seven intervals of the musical scale.

Then, on March 6, 1869, as a result of an extensive correlation of the properties and the atomic weights of the elements, with special attention to valency (that is, the number of single bonds the element can form), Dmitri Mendeleev presented the first version of his periodic table to the Russian Chemical Society and proposed the periodic law, by which “the elements arranged according to the magnitude of atomic weights show a periodic change of properties.” (Lothar Meyer had independently reached a similar conclusion, published after the appearance of Mendeleev’s paper.)