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all the atoms of a given atomic number—that is, having the same number of protons in their nuclei. The quantity of electrons in an atom generally equals the number of protons in the nucleus.
In addition to protons, the atomic nucleus contains neutrons, whose number may vary in atoms of the same element. Atoms having the same atomic number but a different mass number (the sum of protons and neutrons in the nucleus) are called isotopes. Many chemical elements in nature have two or more isotopes; scientists have identified 276 stable isotopes of 81 natural chemical elements, and they have isolated some 1,500 radioactive ones. On the earth, atoms of a given element generally contain constant proportions of isotopes, thus giving the element a virtually constant atomic weight, one of the element’s most important characteristics.
As of 1978,107 chemical elements were known. They are predominantly nonradioactive. All simple and compound substances are formed by these few elements. The element in its free state is called a simple substance. Some chemical elements in the free state exist in two or more allotropic forms, which differ in their chemical and physical properties; for example, carbon may take the forms of graphite and diamond. Nearly 400 simple substances are known. The concepts of simple substance and element are sometimes confused, since, in the vast majority of cases, the two are identical and have the same name. “Nevertheless, a conceptual distinction should always be maintained,” wrote D. I. Mendeleev in 1869 (Soch., vol. 13, 1949, p. 490). A chemical compound consists of atoms of two or more different elements that are chemically bonded together. Scientists have identified more than 100,000 inorganic compounds and more than 3 million organic compounds.
The chemical symbols used to designate the elements consist of the first letter of the element’s Latin name or a combination of the first and one of the subsequent letters. In chemical formulas and equations, such a symbol represents both the name and the atomic weight of the element. The study of chemical elements constitutes chemistry—in particular, inorganic chemistry.
Historical survey. In the prescientific period, the doctrine of Empedocles was widely accepted; Empedocles held that all matter was composed of four primary “elements”—fire, air, water, and earth. This theory, as elaborated by Aristotle, was fully accepted in alchemy. In the eighth and ninth centuries, alchemists added to the theory the idea that sulfur and mercury were components of all metals; sulfur was thought to be the source (principle) of combustibility, and mercury was the source of metallic properties. A third principle arose in the 16th century: salt was responsible for making substances nonvolatile and fire-resistant.
In 1661, R. Boyle rejected the theory of the four elements and three principles by offering the first scientific definition of chemical elements as pure and simple substances that may be combined to form compounds. In the 18th century, the hypothesis formulated by J. J. Becher and G. E. Stahl was generally recognized. According to this hypothesis, natural substances consist of water, earth, and phlogiston, the last of these being the source of combustibility. At the end of the 18th century, the hypothesis was refuted by A. L. Lavoisier, who essentially repeated Boyle’s formulation, defining chemical elements as substances that cannot be broken down into simpler substances and that combine to form compounds. Lavoisier, however, carried the concept further by offering in 1789 the first list of actual chemical elements in the history of science. This list included all the nonmetals known at the time (oxygen, nitrogen, hydrogen, sulfur, phosphorus, and carbon) and all the metals then known (silver, arsenic, bismuth, cobalt, calcium, tin, iron, manganese, mercury, molybdenum, nickel, gold, platinum, lead, tungsten, and zinc), as well as “radicals” (containing chlorine, fluorine, and boron; seeRADICAL THEORY) and “earths” (lime [CaO], magnesium oxide [MgO], barium oxide [BaO], aluminum oxide [Al2O3], and silica [SiO2]). Lavoisier conjectured that the earths were compound substances, but since this assertion had not yet been experimentally proved, he classified them as chemical elements. In recognition of contemporary views, he also included the weightless “fluids” light and caloric in his list of chemical elements. Lavoisier considered the bases sodium hydroxide and potassium hydroxide to be compound substances, even though it had not been shown that they could be broken down into simpler substances. (Not until 1807 did H. Davy do so by electrolysis.)
The research of J. Dalton led to the more refined concept that an element was an atomic species with a specific mass. In 1803, Dalton compiled the first table of relative atomic weights. Assigning a weight of 1 atomic mass unit to the hydrogen atom, he expressed the atomic weights of oxygen, nitrogen, carbon, sulfur, and phosphorus in terms of that unit. He thereby laid the foundation for the general recognition of atomic weight as a major characteristic of an element. Like Lavoisier, Dalton considered chemical elements to be substances that could not be broken down into simpler substances.
Thereafter, chemistry developed rapidly; in particular, a large number of chemical elements were discovered. Lavoisier had listed only 25 elements (including the radicals, but excluding the fluids and earths), but by 1869, when D. I. Mendeleev discovered the periodic law, 63 elements were known. Mendeleev’s periodic table established the interrelationships and proper classification of the elements and led to the prediction of the existence and properties of elements that were then unknown.
The discovery of radioactivity at the end of the 19th century undermined the belief, held for more than 100 years, that the atom was indivisible. Taking the phenomenon of radioactivity into consideration, scientists continued to debate the nature of the chemical elements until virtually the mid-20th century. The debate was at last resolved by the modern theory of atomic structure, which includes the strictly objective definition of chemical elements that was given above.
Abundances in nature. The abundance of chemical elements in the universe are determined by nucleosynthesis within stars. The sun, terrestrial planets, and meteorites in our solar system appear to contain almost exactly the same chemical elements, although in different proportions. The nuclei of the elements are formed as a result of nuclear reactions in stellar interiors. Thus, in different stages of their evolution, stars and stellar systems have differing chemical compositions (see COSMOGONY). A total of 99.9 percent of the matter in the universe consists of hydrogen and helium. Cosmochemistry deals with the abundances and distribution of chemical elements in the universe, the motion and combination of atoms in the formation of cosmic matter, and the chemical composition of celestial bodies; geochemistry is the most developed branch of cosmochemistry.
Of the 107 chemical elements, only 89 are found in nature. The remainder—namely, technetium (element 43), promethium (element 61), astatine (element 85), francium (element 87), and the transuranium elements—have been obtained artificially by means of nuclear reactions; tiny amounts of technetium, promethium, neptunium, and francium are formed upon the spontaneous fission of uranium and are present in uranium ores. The ten most common elements in the earth’s crust have atomic numbers ranging from 8 to 26. These ten elements account for 99.92 percent of the mass of the crust. Table 1 shows the relative proportions in which they are found.
|Table 1. Ten most common elements in the earth’s crust|
|Element||Atomic weight||Percentage (by weight) of the earth’s crust|
Classification and properties. The periodic table of elements that was compiled by D. I. Mendeleev provides the most advanced natural classification of the chemical elements; it shows their interrelationships and indicates the variation in properties that results from differences in atomic number. On the basis of their properties, chemical elements are divided into metals and nonmetals; the distinction between the two is shown in the periodic table.
The tendency to lose outer electrons and thus form cations is the most characteristic property exhibited by metals in chemical reactions. The tendency to pick up electrons and thus form anions is the most characteristic feature of nonmetals in chemical reactions; in other words, the nonmetals exhibit high electronegativity. A distinction is also made between the nontransition and transition elements. The former are those groups of elements in which successive filling of the s and p orbitals occurs; the latter are those in which the d and f orbitals are filled.
At room temperature, two chemical elements (mercury and bromine) exist in the liquid state, and eleven elements (hydrogen, nitrogen, oxygen, fluorine, chlorine, helium, neon, argon, krypton, xenon, and radon) exist in the gaseous state; the remainder are solids. The melting points of the solid elements vary widely, ranging from approximately 30°C (28.5°C for cesium and 29.8°C for gallium) to 3000°C or more (2996°C for tantalum, 3410°C for tungsten, and 3800° ± 200°C for graphite at a pressure of 125 kilobars).
More detailed information on the properties, production, and uses of the chemical elements is given in the articles on the individual elements and groups of elements.
REFERENCESKedrov, B. M. Evoliutsiia poniatiia elementa v khimii. Moscow, 1956.
Seaborg, G. T., and E. G. Velens. Elementy Vselennoi. Moscow, 1962. (Translated from English.)
Seaborg, G. Iskusstvennye transuranovye elementy. Moscow, 1965. (Translated from English.)
Figurovskii, N. A. Otkrytie khimkheskikh elementov i proiskhozhdenie ikh nazvanii. Moscow, 1970.
Populiarnaia biblioteka khimicheskhikh elementov. Moscow, 1971–73.
Nekrasov, B. V. Osnovy obshchei khimii, 3rd ed., [vols.] 1–2. Moscow, 1973.
Pauling, L. Obshchaia khimiia. Moscow, 1974. (Translated from English.)
Giua, M. Istoriia khimii, 2nd ed. Moscow, 1975. (Translated from Italian.)
Weeks, M. E. Discovery of the Elements, 6th ed. Easton, Pa., 1956.
S. A. POGODIN