ISOTOPE

one of two or more species of atoms of a chemical element with the same atomic number and position in the periodic table and nearly identical chemical behaviour but with different atomic masses and physical properties. Every chemical element has one or more isotopes. An atom is first identified and labeled according to the number of protons in its nucleus. This atomic number is ordinarily given the symbol Z. The great importance of the atomic number derives from the observation that all atoms with the same atomic number have nearly, if not precisely, identical chemical properties. A large collection of atoms with the same atomic number constitutes a sample of an element. A bar of pure uranium, for instance, would consist entirely of atoms with atomic number 92. The periodic table of the elements assigns one place to every atomic number, and each of these places is labeled with the common name of the element, as, for example, calcium, radon, or uranium. Not all the atoms of an element need have the same number of neutrons in their nuclei. In fact, it is precisely the variation in the number of neutrons in the nuclei of atoms that gives rise to isotopes. Hydrogen is a case in point. It has the atomic number 1. Three nuclei with one proton are known that contain 0, 1, and 2 neutrons, respectively. The three share the place in the periodic table assigned to atomic number 1 and hence are called isotopes (from the Greek isos, meaning same, and topos, signifying place) of hydrogen. Many important properties of an isotope depend on its mass. The total number of neutrons and protons (symbol A), or mass number, of the nucleus gives approximately the mass measured on the so-called atomic-mass-unit (amu) scale. The numerical difference between the actual measured mass of an isotope and A is called the mass defect (symbol D; see Table). The specification of Z, A, and the chemical symbol (a one- or two-letter abbreviation of the element's name, say Sy) in the form A/ZSy identifies an isotope adequately for most purposes. Thus in the standard notation, 1/1H refers to the simplest isotope of hydrogen and 235/92U to an isotope of uranium widely used for nuclear power generation and nuclear weapons fabrication. (Authors who do not wish to use symbols sometimes write out the element name and mass numberhydrogen-1 and uranium-235 in the examples above.) The term nuclide is used to describe particular isotopes, notably in cases where the nuclear rather than the chemical properties of an atom are to be emphasized. The lexicon of isotopes includes three other frequently used terms: isotones for isotopes of different elements with the same number of neutrons; isobars for isotopes of different elements with the same mass number; and isomers for isotopes identical in all respects except for the total energy content of the nuclei. one of two or more species of atoms of a chemical element with the same atomic number and position in the periodic table and nearly identical chemical behaviour but with different atomic masses and physical properties. Before the early 1900s it was generally assumed that the mass of a standard number of atoms of any given element was a basic characteristic of the element. It was also thought that all the atoms of an element were the same and, in particular, had the same mass. The first evidence that two substances with the same chemical properties did not have to be physically identical came from the study of the radioactivity of the heavy elements. Between 1906 and 1907 several investigators showed that ionium (a decay product of uranium) and radiothorium (a decay product of thorium), when mixed with thorium, could not be separated from it by any chemical means. The two substances had radioactive properties quite different from those of thorium and could be shown to have atomic masses differing by several units from that of thorium. The term isotope was introduced in 1913 by the English chemist Frederick Soddy to cover such situations. Not long after the acceptance of these ideas as applied to the heavy elements came indications that isotopy might exist in the main group of naturally occurring stable elements. In 1919 Francis William Aston of England showed conclusively that neon consisted chiefly of two atomic species. This success was followed by the discovery that chlorine had two isotopes. It soon became clear that most elements consisted of a mixture of isotopes, each with an atomic mass close to an integer on the atomic mass scale. The majority of elements, as found in the Earth's crust and atmosphere, are now known to be mixtures of several isotopes. Such mixtures are in almost unvarying proportions. In its natural form, tin, for example, has 10 isotopes whose atomic masses range from values approximating 112 to 124, but any given sample of the element has an average value of 118.69. In effect, ordinary tin is nature's standardized blend of these 10 atomic species. In most cases, only the stable isotopes of elements can be found in nature. The unstable, or radioactive, forms decay (break down) spontaneously into entirely different elements at characteristic rates because their ratio of neutrons to protons is either too low or too high for stability. Isotopes of all the elements heavier than bismuth are radioactive. Some of these, such as uranium, do occur naturally because their isotopes have long half-lives. Additional reading F.W. Aston, Mass Spectra and Isotopes, 2nd ed. (1942), the history of the discovery of radioactive and stable isotopes; Gerhart Friedlander et al., Nuclear and Radiochemistry, 3rd ed. (1981); Keith J. Laidler, Chemical Kinetics, 3rd ed. (1987); Stelio Villani, Isotope Separation, trans. from Italian (1967); and James W. Truran, Nucleosynthesis, Annual Review of Nuclear and Particle Science, 34:5397 (1984). Gregory F. Herzog