isotope

One of two or more atoms having the same atomic number but different mass numbers.

[ISO– + Greek topos, place (so called because the isotopes of a chemical element occupy the same position in the periodic table of elements).]

Concept

Isotopes are atoms of the same element that have different masses due to differences in the number of neutrons they contain. Many isotopes are stable, meaning that they are not subject to radioactive decay, but many more are radioactive. The latter, also known as radioisotopes, play a significant role in modern life. Carbon-14, for instance, is used for estimating the age of objects within a relatively recent span of time—up to about 5,000 years—whereas geologists and other scientists use uranium-238 to date minerals of an age on a scale with that of the Earth. Concerns over nuclear power and nuclear weapons testing in the atmosphere have heightened awareness of the dangers posed by certain kinds of radioactive isotopes, which can indeed be hazardous to human life. However, the reality is that people are subjected to considerably more radiation from non-nuclear sources.

How It Works

Atoms and Elements

The elements are substances that cannot be broken down into other matter by chemical means, and an atom is the fundamental particle in an element. As of 2001, there were 112 known elements, 88 of which occur in nature; the rest were created in laboratories. Due to their high levels of radioactivity, they exist only for extremely short periods of time. Whatever the number of elements—and obviously that number will increase over time, as new elements are synthesized—the same number of basic atomic structures exists in the universe.

What distinguishes one element from another is the number of protons, subatomic particles with a positive electric charge, in the nucleus, or center, of the atom. The number of protons, whatever it may be, is unique to an element. Thus if an atom has one proton, it is an atom of hydrogen, because hydrogen has an atomic number of 1, as shown on the periodic table of elements. If an atom has 109 protons, on the other hand, it is meitnerium. (Meitnerium, synthesized at a German laboratory in 1982, is the last element on the periodic table to have been assigned a name as of 2001.)

The Nucleus and Electrons

Together with protons in the nucleus are neutrons, which exert no charge. The discovery of these particles, integral to the formation of isotopes, is discussed below. The nucleus, with a diameter about 1/10,000 that of the atom itself, makes up only a tiny portion of the atom's volume, but the vast majority of its mass. Thus a change in the mass of a nucleus, as occurs when an isotope is formed, is reflected by a noticeable change in the mass of the atom itself.

Far from the nucleus (in relative terms, of course), at the perimeter of the atom, are the electrons, which have a negative electric charge. Whereas the protons and neutrons have about the same mass, the mass of an electron is less than 0.06% of either a proton or neutron. Nonetheless, electrons play a highly significant role in chemical reactions and chemical bonding. Just as isotopes are the result of changes in the number of neutrons, ions—atoms that are either positive or negative in electric charge—are the result of changes in the number of electrons.

Unless it loses or gains an electron, thus becoming an ion, an atom is neutral in charge, and it maintains this electric-charge neutrality by having an equal number of protons and electrons. There is, however, no law of the universe stating that an atom must have the same number of neutrons as it does protons and electrons: some do, but this is far from universal, as we shall see.

Neutrons

The number of neutrons is variable within an element precisely because they exert no charge, and thus while their addition or removal changes the mass, it does not affect the electric charge of the atom. Therefore, whereas the importance of the proton and the electron is very clear to anyone who studies atomic behavior, neutrons, on the other hand, might seem at first glance as though they are only "along for the ride." Yet they are all-important to the formation of isotopes.

Not surprisingly, given their lack of electric charge, neutrons were the last of the three major subatomic particles to be discovered. English physicist J. J. Thomson (1856-1940) identified the electron in 1897, and another English physicist, Ernest Rutherford (1871-1937), discovered the proton in 1914. Rutherford's discovery overturned the old "plum pudding" model, whereby atoms were depicted as consisting of electrons floating in a positively charged cloud, rather like raisins in an English plum pudding. As Rutherford showed, the atom must have a nucleus—yet protons alone could not account for the mass of the nucleus.

There must be something else at the heart of the atom, and in 1932, yet another English physicist, James Chadwick (1891-1974), identified what it was. Working with radioactive material, he found that a certain type of subatomic particle could penetrate lead. All types of radiation known at the time were stopped by the lead, and therefore Chadwick reasoned that this particle must be neutral in charge. In 1932, he won the Nobel Prize in physics for his discovery of the neutron.

Neutrons and Nuclear Fusion

Neutrons played a critical role in the development of the atomic bomb during the 1940s. In nuclear fission, atoms of uranium are bombarded with neutrons. The result is that the uranium nucleus splits in half, releasing huge amounts of energy. As it does so, it emits several extra neutrons, which split more uranium nuclei, creating still more energy and setting off a chain reaction.

This explains the destructive power in an atomic bomb, as well as the constructive power—providing energy to homes and businesses—in a nuclear power plant. Whereas the chain reaction in an atomic bomb becomes an uncontrolled explosion, in a nuclear plant, the reaction is slowed and controlled. One of the means used to do this is by the application of "heavy water," which, as we shall see, is water made with a hydrogen isotope.

Isotopes: the Basics

Two atoms may have the same number of protons, and thus be of the same element, yet differ in their number of neutrons. Such atoms are called isotopes, atoms of the same element having different masses. The name comes from the Greek phrase isos topos, meaning "same place": because they have the same atomic number, isotopes of the same element occupy the same position on the periodic table.

Also called nuclides, isotopes are represented symbolically as follows: where S is the symbol of the element, a is the atomic number, and m is the mass number—the sum of protons and neutrons in the atom's nucleus. For the stable silver isotope designated as for instance, Ag is the element symbol; 47 its atomic number; and 93 the mass number. From this, it is easy to discern that this particular stable isotope has 46 neutrons in its nucleus.

Because the atomic number of any element is established, sometimes isotopes are represented simply with the mass number, thus: ⁹³Ag. They may also be designated with a subscript notation indicating the number of neutrons, so that this information can be obtained at a glance without having to do the arithmetic. For the silver isotope shown here, this is written as Isotopes can also be indicated by simple nomenclature: for instance, carbon-12 or carbon-13.

Stable and Unstable Isotopes

Radioactivity is a term describing a phenomenon whereby certain materials are subject to a form of decay brought about by the emission of high-energy particles or radiation. Forms of particles or energy emitted in radiation include alpha particles (positively charged helium nuclei); beta particles (either electrons or subatomic particles called positrons); or gamma rays, which occupy the highest energy level in the electromagnetic radiation emitted by the Sun. Radioactivity will be discussed below, but for the present, the principal concern is with radioactive properties as a distinguishing factor between the two varieties of isotope.

Isotopes are either stable or unstable. The unstable variety, known as radioisotopes, are subject to radioactive decay, but in this context, "decay" does not mean what it usually does. A radioisotope does not "rot"; it decays by turning into another isotope of the same element—or even into another element entirely. (For example, uranium-238 decays by emitting alpha particles, ultimately becoming lead-206.) A stable isotope, on the other hand, has already become what it is going to be, and will not experience further decay.

Most elements have between two and six stable isotopes. On the other hand, a few elements—for example, technetium—have no stable isotopes. Twenty elements, among them gold, fluorine, sodium, aluminum, and phosphorus, have only one stable isotope each. The element with the most stable isotopes is easy to remember because its name is almost the same as its number of stable isotopes: tin, with 10.

As for unstable isotopes, there are over 1,000, some of which exist in nature, but most of which have been created synthetically in laboratories. This number is not fixed; in any case, it is not necessarily important, because many of these highly radioactive isotopes last only for fractions of a second before decaying to form a stable isotope. Yet radioisotopes in general have so many uses, in comparison to stable isotopes, that they are often referred to simply as "isotopes."

Understanding Isotopes

Before proceeding with a discussion of isotopes and their uses, it is necessary to address a point raised earlier, when it was stated that some atoms do have the same numbers of neutrons and protons, but that this is far from universal. In fact, nuclear stability is in part a function of neutron-to-proton ratio.

Stable nuclei with low atomic numbers (up to about 20) have approximately the same number of neutrons and protons. For example, the most stable and abundant form of carbon is carbon-12, with six protons and six neutrons. Beyond atomic number 20 or so, however, the number of neutrons begins to grow: in other words, the lowest mass number is increasingly high in comparison to the atomic number.

For example, uranium has an atomic number of 92, but the lowest mass number for a uranium isotope is not 184, or 92 multiplied by two; it is 218. The ratio of neutrons to protons necessary for a stable isotope creeps upward along the periodic table: tin, with an atomic number of 50, has a stable isotope with a mass number of 120, indicating a 1.4 to 1 ratio of neutrons to protons. For mercury-200, the ratio is 1.5 to 1.

The higher the atomic number, by definition, the greater the number of protons in the nucleus. This means that more neutrons are required to "bind" the nucleus together. In fact, all nuclei with 84 protons or more (i.e., starting at polonium and moving along the periodic table) are radioactive, for the simple reason that it is increasingly difficult for the neutrons to with stand the strain of keeping so many protons in place.

One can predict the mode of radioactive decay by noting whether the nucleus is neutron-rich or neutron-poor. Neutron-rich nuclei undergo beta emission, which decreases the numbers of protons in the nucleus. Neutron-poor nuclei typically undergo positron emission or electron capture, the first of these being more prevalent among the lighter nuclei. Elements with atomic numbers of 84 or greater generally undergo alpha emission, which decreases the numbers of protons and neutrons by two each.

Real-Life Applications

Deuterium and Tritium

Only three isotopes are considered significant enough to have names of their own, as opposed to being named after a parent atom (for example, carbon-12, uranium-238). These are protium, deuterium, and tritium, all three isotopes of hydrogen. Protium, or ¹H, is hydrogen in its most basic form—one proton, no neutrons—and the name "protium" is only applied when necessary to distinguish it from the other two isotopes. Therefore we will focus primarily on the two others.

Deuterium, designated as ²H, is a stable isotope, whereas tritium—³H—is radioactive. Both, in fact, have chemical symbols (D and T respectively), just as though they were elements on the periodic table. What makes these two so special? They are, as it were, "the products of a good home"—in other words, their parent atom is the most basic and plentiful element in the universe. Indeed, the vast majority of the universe is hydrogen, along with helium, which is formed by the fusion of hydrogen atoms. If all atoms were numbers, then hydrogen would be 1; but of course, this is more than a metaphor, since its atomic number is indeed 1.

Ordinary hydrogen or protium, as noted, consists of a single proton and a single electron, the simplest possible atomic form possible. Its simplicity has made it a model for understanding the atom, and therefore when physicists discovered the existence of two hydrogen atoms that were just a bit more complex, they were intrigued.

Just as hydrogen represented the standard against which atoms could be measured, scientists reasoned, deuterium and tritium could offer valuable information regarding stable and unstable isotopes respectively. Furthermore, the pronounced tendency of hydrogen to bond with other substances—it almost never appears by itself on Earth—presented endless opportunities for study regarding hydrogen isotopes in association with other elements.

Isolation of Deuterium

Deuterium is sometimes called "heavy hydrogen," and its nucleus—with one proton and one neutron—is called a deuteron. It was first isolated in 1931 by American chemist Harold Clayton Urey (1893-1981), who was awarded the 1934 Nobel Prize in Chemistry for his discovery.

Serving at that time as a professor of chemistry at Columbia University in New York City, Urey started with the assumption that any hydrogen isotopes other than protium must exist in very minute quantities. This assumption, in turn, followed from an awareness that hydrogen's average atomic mass—measured in atomic mass units—was only slightly higher than 1. There must be, as Urey correctly reasoned, a very small quantity of "heavy hydrogen" on Earth.

To separate deuterium, Urey collected a relatively large sample of liquid hydrogen: 4.2 quarts (4 l). Then he allowed the liquid to evaporate very slowly, predicting that the more abundant protium would evaporate more quickly than the isotope whose existence he had hypothesized. After all but 0.034 oz (1 ml) of the sample had evaporated, he submitted the remainder to a form of analysis called spectroscopy, adding a burst of energy to the atoms and then analyzing the light spectrum they emitted for evidence of differing varieties of atom.

Characteristics and Uses of Deuterium

Deuterium, with an atomic mass of 2.014102 amu, is almost exactly twice as heavy as protium, which has an atomic mass of 1.007825. Its melting point, or the temperature at which it changes from a solid to a liquid −426°F (−254°C), is much higher than for protium, which melts at −434°F (−259°C). The same relationship holds for its boiling point, or the temperature at which it changes from a liquid to its normal state on Earth, as a gas: −417°F (−249°C), as compared to −423°F (−253°C) for protium. Deuterium is also much, much less plentiful than protium: protium represents 99.985% of all the hydrogen that occurs naturally, meaning that deuterium accounts for just 0.015%.

Often, deuterium is applied as a tracer, an atom or group of atoms whose participation in a chemical, physical, or biological reaction can be easily observed. Radioisotopes are most often used as tracers, precisely because of their radioactive emissions; deuterium, on the other hand, is effective due to its almost 2:1 mass ratio in comparison to protium. In addition, it bonds with other atoms in a fashion slightly different from that of protium, and this contrast makes its presence easier to trace.

Its higher boiling and melting points mean that when deuterium is combined with oxygen to form "heavy water" (D₂O), the water likewise has higher boiling and melting points than ordinary water. Heavy water is often used in nuclear fission reactors to slow down the fission process, or the splitting of atoms.

Deuterium in Nuclear Fusion

Deuterium is also applied in a type of nuclear reaction much more powerful that fission: fusion, or the joining of atomic nuclei. The Sun produces energy by fusion, a thermonuclear reaction that takes places at temperatures of many millions of degrees Celsius. In solar fusion, it appears that two protium nuclei join to form a single deuteron.

During the period shortly after World War II, physicists developed a means of duplicating the thermonuclear fusion process. The result was the hydrogen bomb—more properly called a fusion bomb—whose detonating device was a compound of lithium and deuterium called lithium deuteride. Vastly more powerful than the "atomic" (that is, fission) bombs dropped by the United States over Japan in 1945, the hydrogen bomb greatly increased the threat of worldwide nuclear annihilation in the postwar years.

Yet the power that could destroy the world also has the potential to provide safe, abundant fusion energy from power plants—a dream that as yet remains unrealized. Among the approaches being attempted by physicists studying nuclear fusion is a process in which two deuterons are fused. The result is a triton, the nucleus of tritium, along with a single proton. The triton and deuteron would then be fused to create a helium nucleus, with a resulting release of vast amounts of energy.

Tritium

Whereas deuterium has a single neutron, tritium—as its mass number of 3 indicates—has two. And just as deuterium has approximately twice the mass of protium, tritium has about three times the mass, 3.016 amu. As is expected, the thermal properties of tritium are different from those of protium. Again, the melting and boiling points are higher: thus tritium heavy water (T₂O) melts at 40°F (4.5°C), as compared with 32°F (0°C) for H₂O.

Because it is radioactive, tritium is often described in terms of half-life, the length of time it takes for a substance to diminish to one-half its initial amount. The half-life of tritium is 12.26 years. As it decays, its nucleus emits a low-energy beta particle, and this results in the creation of the helium-3 isotope. Due to the low energy levels involved, the radioactive decay of tritium poses little danger to humans.

Like deuterium, tritium is applied in nuclear fusion, though due to its scarcity, it is usually combined with deuterium. Furthermore, tritium decay requires that hydrogen bombs containing the radioisotope be recharged periodically. Also, like deuterium, tritium is an effective tracer. Sometimes it is released in small quantities into groundwater as a means of monitoring subterranean water flow. It is also used as a tracer in biochemical processes.

Separating Isotopes

As noted in the discussion of deuterium, tritium can only be separated from protium due to the differences in mass. The chemical properties of isotopes with the same parent element make them otherwise indistinguishable, and hence purely chemical means cannot be used to separate them.

Physicists working on the Manhattan Project, the U.S. effort to develop atomic weaponry during World War II, were faced with the need to separate ²³⁵U from ²³⁸U. Uranium-238 is far more abundant, but what they wanted was the uranium-235, highly fissionable and thus useful in the processes they were attempting.

Their solution was to allow a gaseous uranium compound to diffuse, or separate, the uranium through porous barriers. Because uranium-238 was heavier, it tended to move more slowly through the barriers, much like grains of rice getting caught in a sifter. Another means of separating isotopes is by mass spectrometry.

Radioactivity

One of the scientists working on the Manhattan Project was Italian physicist Enrico Fermi (1901-1954), who used radium and beryllium powder to construct a neutron source for making new radioactive materials. Fermi and his associates succeeded in producing radioisotopes of sodium, iron, copper, gold, and numerous other elements. As a result of Fermi's work, for which he won the 1938 Nobel Prize in Physics, scientists have been able to develop radioactive versions of virtually all elements.

Interestingly, the ideas of radioactivity, fission reactions, and fusion reactions collectively represent the realization of a goal sought by the medieval alchemists: the transformation of one element into another. The alchemists, forerunners of chemists, believed they could transform ordinary metals into gold by using various potions—an impossible dream. Yet as noted in the preceding paragraph, among the radioisotopes generated by Fermi's neutron source was gold. The "catch," of course, is that this gold was unstable; furthermore, the amount of energy and human mental effort required to generate it far outweighed the monetary value of the gold itself.

Radioactivity is, in the modern imagination, typically associated with fallout from nuclear war, or with hazards resulting from nuclear power—hazards that, as it turns out, have been greatly exaggerated. Nor is radioactivity always harmful to humans. For instance, with its applications in medicine—as a means of diagnosing and treating thyroid problems, or as a treatment for cancer patients—it can actually save lives.

Hazards Associated With Radioactivity

It is a good thing that radiation, even the harmful variety known as ionizing radiation, is not fatal in small doses, because every person on Earth is exposed to small quantities of radiation every year. About 82% of this comes from natural sources, and 18% from manmade sources. Of course, some people are at much greater risk of radiation exposure than others: coal miners are exposed to higher levels of the radon-222 isotope present underground, while cigarette smokers ingest much higher levels of radiation than ordinary people, due to the polonium-210, lead-210, and radon-222 isotopes present in the nitrogen fertilizers used to grow tobacco.

Nuclear weapons, as most people know, produce a great deal of radioactive pollution. However, atmospheric testing of nuclear armaments has long been banned, and though the isotopes released in such tests are expected to remain in the atmosphere for about a century, they do not constitute a significant health hazard to most Americans. (It should be noted that nations not inclined to abide by international protocols might still conduct atmospheric tests in defiance of the test bans.) Nuclear power plants, despite the great deal of attention they have received from the media and environmentalist groups, do not pose the hazard that has often been claimed: in fact, coal-and oil-burning power plants are responsible for far more radioactive pollution in the United States.

This is not to say that nuclear energy poses no dangers, as the disaster at Chernobyl in the former Soviet Union has shown. In April 1986, an accident at a nuclear reactor in what is now the Ukraine killed 31 workers immediately, and ultimately led to the deaths of some 10,000 people. The fact that the radiation was allowed to spread had much to do with the secretive tactics of the Communist government, which attempted to cover up the problem rather than evacuate the area.

Another danger associated with nuclear power plants is radioactive waste. Spent fuel rods and other waste products from these plants have to be dumped somewhere, but it cannot simply be buried in the ground because it will create a continuing health hazard through the water supply. No fully fail-safe storage system has been developed, and the problem of radioactive waste poses a continuing threat due to the extremely long half-lives of some of the isotopes involved.

Dating Techniques

In addition to their uses in applications related to nuclear energy, isotopes play a significant role in dating techniques. The latter may sound like a subject that has something to do with romance, but it does not: dating techniques involve the use of materials, including isotopes, to estimate the age of both organic and inorganic materials.

Uranium-238, for instance, has a half-life of 4.47 · 10⁹ years, which is nearly the age of Earth; in fact, uranium-dating techniques have been used to determine the planet's age, which is estimated at about 4.7 billion years. As noted elsewhere in this volume, potassium-argon dating, which involves the isotopes potassium-40 and argon-40, has been used to date volcanic layers in east Africa. Because the half-life of potassium-40 is 1.3 billion years, this method is useful for dating activities that are distant in the human scale of time, but fairly recent in geological terms.

Another dating technique is radiocarbon dating, used for estimating the age of things that were once alive. All living things contain carbon, both in the form of the stable isotope carbon-12 and the radioisotope carbon-14. While a plant or animal is living, there is a certain proportion between the amounts of these two isotopes in the organism's body, with carbon-12 being far more abundant. When the organism dies, however, it ceases to acquire new carbon, and the carbon-14 present in the body begins to decay into nitrogen-14. The amount of nitrogen-14 that has been formed is thus an indication of the amount of time that has passed since the organism was alive.

Because it has a half-life of 5,730 years, carbon-14 is useful for dating activities within the span of human history, though it is not without controversy. Some scientists contend, for instance, that samples may be contaminated by carbon from the surrounding soils, thus affecting ratios and leading to inaccurate dates.

Where to Learn More

"Carbon 14 Dating Calculator" (Web site). <http://www.museum.mq.edu.au/eegypt2/carbdate.html> (May 15, 2001).

Ebbing, Darrell D.; R. A. D. Wentworth; and James P. Birk. Introductory Chemistry. Boston: Houghton Mifflin, 1995.

"Exploring the Table of Isotopes" (Web site). <http://ie.lbl.gov/education/isotopes.htm> (May 15, 2001).

Goldstein, Natalie. The Nature of the Atom. New York: Rosen Publishing Group, 2001.

"The Isotopes" (Web site). <http://chemlab.pc.maricopa.edu/periodic/isotopes.html> (May 15, 2001).

"Isotopes" University of Colorado Department of Physics (Web site). <http://www.colorado.edu/physics/2000/isotopes/index.html> (May 15, 2001).

Milne, Lorus Johnson and Margery Milne. Understanding Radioactivity. Illustrated by Bill Hiscock. New York: Atheneum, 1989.

Smith, Norman F. Millions and Billions of Years Ago: Dating Our Earth and Its Life. New York: F. Watts, 1993.

"Stable Isotope Group." Martek Biosciences (Web site). <http://www.martekbio.com/frmain.htm> (May 15, 2001).

"Tracking with Isotopes" (Web site). <http://whyfiles.org/083isotope/2.html> (May 15, 2001).

One member of a (chemical-element) family of atomic species which has two or more nuclides with the same number of protons (Z) but a different number of neutrons (N). Because the atomic mass is determined by the sum of the number of protons and neutrons contained in the nucleus, isotopes differ in mass. Since they contain the same number of protons (and hence electrons), isotopes have the same chemical properties. However, the nuclear and atomic properties of isotopes can be different. The electronic energy levels of an atom depend upon the nuclear mass. Thus, corresponding atomic levels of isotopes are slightly shifted relative to each other. A nucleus can have a magnetic moment which can interact with the magnetic field generated by the electrons and lead to a splitting of the electronic levels. The number of resulting states of nearly the same energy depends upon the spin of the nucleus and the characteristics of the specific electronic level. See also Atomic structure and spectra; Hyperfine structure; Isotope shift.

Of the 12 elements onfirmed thus far, 81 have at least one stable isotope whereas the others exist only in the form of radioactive nuclides. Some radioactive nuclides (for example, ¹¹⁵In, ²³²Th, ²³⁵U, ²³⁸U) have survived from the time of formation of the elements. Several thousand radioactive nuclides produced through natural or artificial means have been identified. See also Radioisotope.

Of the 83 elements which occur naturally in significant quantities on Earth, 20 are found as a single isotope (mononuclidic), and the others as admixtures containing from 2 to 10 isotopes. Isotopic composition is mainly determined by mass spectroscopy. See also Mass spectroscope.

Nuclides with identical mass number (that is, A = N + Z) but differing in the number of protons in the nucleus are called isobars. Nuclides having different mass number but the same number of neutrons are called isotones. See also Isobar (nuclear physics); Isotone.

Isotopic abundance refers to the isotopic composition of an element found in its natural terrestrial state. The isotopic composition for most elements does not vary much from sample to sample. This is true even for samples of extraterrestrial origin such as meteorites and lunar materials brought back to Earth by space missions. However, there are a few exceptional cases for which variations of up to several percent have been observed. There are several phenomena that can account for such variations, the most likely being some type of nuclear process which changes the abundance of one isotope relative to the others. For some of the lighter elements, the processes of distillation or chemical exchange between different chemical compounds could be responsible for isotopic differences. See also Nuclear reaction; Radioactivity.

The areas in which separated (or enriched) isotopes are utilized have become fairly extensive, and a partial list includes nuclear research, nuclear power generation, nuclear weapons, nuclear medicine, and agricultural research. For many applications there is a need for separated radioactive isotopes. These are usually obtained through chemical separations of the desired element following production by means of a suitable nuclear reaction. Separated radioactive isotopes are used for a number of diagnostic studies in nuclear medicine, including the technique of positron tomography. See also Isotope separation; Nuclear medicine.

Studies of metabolism, drug utilization, and other reactions in living organisms can be done with stable isotopes such as ¹³C, ¹⁵N, ¹⁸O, and ²H. Molecular compounds are “spiked” with these isotopes, and the metabolized products are analyzed by using a mass spectrometer to measure the altered isotopic ratios. See also Isotope dilution techniques; Radioisotope (biology).

One member of a family of chemical elements that has the same chemical properties (the same atomic number) but differs in mass. Isotopes have the same number of protons and electrons, but a different number of neutrons. The mass is determined by the total number of nucleons (neutrons and protons). See allotrope.

Download Computer Desktop Encyclopedia to your iPhone/iTouch

Forms of elements with the same chemical properties, differing in atomic mass because of differing numbers of neutrons in the nucleus. Thus, hydrogen has three isotopes, of atomic masses 1, 2, and 3, generally written as ¹H, ²H (deuterium), and ³H (tritium). ¹H is the most abundant isotope of hydrogen; ²H is stable, while ³H is radioactive.

Radioactive isotopes are unstable, and decay to stable elements, emitting radiation in the process. This may be α-radiation, β-radiation (electrons), γ-radiation, or X-rays, depending on the isotope. The time taken for half the radioactivity to decay is the half-life of the isotope, and can vary from a fraction of a second, through several days to years (e.g. the half-life of ³H is 12½ years, that of ¹⁴C is 5200 years).

Stable isotopes can be detected only by their different atomic mass. Since they emit no radiation, they are safe for use in labelled compounds given to human beings. Examples of stable isotopes commonly used in nutrition research include ²H, ¹³C, ¹⁵N, and ¹⁸O.

One of several nuclides having the same number of protons in their nuclei, and hence having the same atomic number but differing in the number of neutrons, and therefore in the mass number. The isotopes of a particular element have virtually identical chemical properties.

One of two or more alternative forms of an element that have the same number of protons in their nucleus, but have different numbers of neutrons. An isotope of carbon is used in carbon dating.

For more information on isotope, visit Britannica.com.

Atoms of an element that have different numbers of neutrons are isotopes of the same element. Isotopes of an element have the same atomic number but different mass numbers. Common examples are the isotopes of carbon: ¹²C and ¹⁴C. ¹²C has 6 protons, 6 electrons, and 6 neutrons, while ¹⁴C has 6 protons, 6 electrons, and 8 neutrons. Some isotopes are physically stable, while others, known as radioisotopes, are unstable. Radioisotopes undergo radioactive decay, emitting both particles and energy. If the decay leads to a change in the number of protons, the atomic number changes, transforming the isotopes into a different element.

In physics, different forms of the same element, with nuclei that have the same number of protons but different numbers of neutrons. Isotopes are distinguished from each other by giving the combined number of protons and neutrons in the nucleus. For example, uranium 235 is the isotope of uranium that has 235 protons and neutrons in its nucleus rather than the more commonly occurring 238. All elements have isotopes.

A chemical element having the same atomic number as another (i.e. the same number of nuclear protons), but having a different atomic mass (i.e. a different number of nuclear neutrons).

• radioactive i. — one having an unstable nucleus and which emits characteristic radiation during its decay to a stable form. See also radioisotope.
• stable i. — one that does not transmute into another element with emission of corpuscular or electromagnetic radiations.

For other uses, see Isotope (disambiguation).

Isotopes (Greek isos = "equal", tópos = "site, place") are any of the different types of atoms (nuclides) of the same chemical element, each having a different atomic mass (mass number).^[1] Isotopes of an element have nuclei with the same number of protons (the same atomic number) but different numbers of neutrons. Therefore, isotopes of the same element have different mass numbers (number of nucleons).

A nuclide is any particular atomic nucleus with a specific atomic number Z and mass number A; it is equivalently an atomic nucleus with a specific number of protons and neutrons. Collectively, all the isotopes of all the elements form the set of nuclides. The distinction between the terms isotope and nuclide has somewhat blurred, and they are often used interchangeably. If they are to be distinguished in use, isotope is better used in its original sense, when referring to several different nuclides of the same chemical element. Nuclide is a later and more generic term, and is used when referencing to only one type of nucleus, and may also be used to refer to several types of nuclei of different elements. For example, it is better to say that an element such as fluorine consists of one stable nuclide rather than that it has one stable isotope, because the latter word is usually reserved to refer to more than one nuclide. On the other hand, carbon can be correctly said to have two stable isotopes, and fluorine to have several radioactive isotopes.

Isotopes and nuclides are specified by the name of the particular element, implicitly giving the atomic number, followed by a hyphen and the mass number (e.g. helium-3, carbon-12, carbon-13, iodine-131 and uranium-238). In symbolic form, the number of nucleons is denoted as a superscripted prefix to the chemical symbol (e.g. ³He, ¹²C, ¹³C, ¹³¹I and ²³⁸U).

About 339 nuclides occur naturally on Earth^[2], of which 256 (about 75%) are stable (or, to be careful, have never been observed to decay; this note is necessary because many "stable" isotopes are predicted to be radioactive with very long half-lives). Counting the radioactive nuclides not found in nature that have been created artificially, more than 3100 nuclides are currently known.^[3]

Contents 1 History of the term 2 Variation in properties between isotopes 2.1 Chemical and atomic properties 2.2 Nuclear properties and stability 3 Occurrence in nature 4 Atomic mass of isotopes 5 Applications of isotopes 5.1 Use of chemical and biological properties 5.2 Use of nuclear properties 6 See also 7 References 8 External links

History of the term

In the bottom right corner of JJ Thomson's photographic plate are markings for the two isotopes of neon: neon-20 and neon-22.

The term isotope was coined in 1913 by Margaret Todd, a Scottish doctor, during a conversation with Frederick Soddy (to whom she was distantly related by marriage).^[4] Soddy, a chemist at Glasgow University, explained that it appeared from his investigations as if several elements occupied each position in the periodic table. Hence Todd suggested the Greek term for "at the same place" as a suitable name. Soddy adopted the term and went on to win the Nobel Prize for Chemistry in 1921 for his work on radioactive substances.

Soddy's use of the word isotope was initially with regard to radioactive (unstable) atoms. However, in 1913, as part of his exploration into the composition of canal rays, J. J. Thomson channeled a stream of ionized neon through a magnetic and an electric field and measured its deflection by placing a photographic plate in its path. Thomson observed two patches of light on the photographic plate (see image on right), which suggested two different parabolas of deflection. This was the first observation of different stable isotopes for an element. Thomson eventually concluded that some of the atoms in the neon gas were of higher mass than the rest.

Variation in properties between isotopes

Chemical and atomic properties

A neutral atom has the same number of electrons as protons. Thus, different isotopes of a given element all have the same number of protons and electrons and the same electronic structure, and because the chemical behavior of an atom is largely determined by its electronic structure, different isotopes exhibit nearly identical chemical behavior. The main exception to this is the kinetic isotope effect: due to their larger masses, heavier isotopes tend to react somewhat more slowly than lighter isotopes of the same element. This is most pronounced for protium (¹H) vis-à-vis deuterium (²H), because deuterium has twice the mass of protium. The mass effect between deuterium and the relatively light protium also affects the behavior of their respective chemical bonds, by means of changing the center of gravity (reduced mass) of the atomic systems. However, for heavier elements, the absolute mass of nucleus relative to electrons is far more, and the relative mass difference between isotopes is much less, and thus the mass-difference effects on chemistry are usually negligible.

Isotope half lifes. Note that the darker more stable isotope region departs from the line of protons Z = neutrons N, as the element number Z becomes larger

Similarly, two molecules which differ only in the isotopic nature of their atoms (isotopologues) will have identical electronic structure and therefore almost indistinguishable physical and chemical properties (again with deuterium providing the primary exception to this rule). The vibrational modes of a molecule are determined by its shape and by the masses of its constituent atoms. Consequently, isotopologues will have different sets of vibrational modes. Since vibrational modes allow a molecule to absorb photons of corresponding energies, isotopologues have different optical properties in the infrared range.

Nuclear properties and stability

See also: Stable isotope and List of elements by nuclear stability

Atomic nuclei consist of protons and neutrons bound together by the nuclear force. Because protons are positively charged, they repel each other. Neutrons, which are electrically neutral, allow some separation between the positively charged protons, reducing the electrostatic repulsion. Neutrons also stabilize the nucleus because at short ranges they attract each other and protons equally by the nuclear force, and this extra binding force also offsets the electrical repulsion between protons. For this reason, one or more neutrons are necessary for two or more protons to be bound into a nucleus. As the number of protons increases, an increasing ratio of neutrons are needed to form a stable nucleus (see graph at right). For example, although the neutron:proton ratio of ³He is 1:2, the neutron:proton ratio of ²³⁸U is greater than 3:2.

Of the 80 elements with a stable isotope, the largest number of stable isotopes observed for any element is ten (for the element tin). Xenon is the only element which has nine stable isotopes. There is no element with exactly eight stable isotopes. See list of elements by nuclear stability for a complete list. Five elements have seven stable isotopes, eight have six stable isotopes, nine have five stable isotopes, nine have four stable isotopes, nine have three stable isotopes, 16 have two stable isotopes (counting Ta-180m as stable), and 26 elements have only a single stable isotope (of these, 19 are so-called mononuclidic elements, having a single primordial stable isotope which dominates and fixes the atomic weight of the natural element to high precision; 3 radioactive mononuclidic elements occur as well).^[5]. In total, there are 256 nuclides which have not been observed to decay (see List of elements by nuclear stability). For the 80 elements which have one or more stable isotopes, the average number of stable isotopes is 256/80 = 3.20 isotopes per element.

Other effects besides the bulk ratio of protons and neutrons affect nuclear stability. For example, the extreme stability of helium-4 due to a double pairing of 2 protons and 2 neutrons prevents any nuclides containing five nucleons from existing for long enough to serve as platforms for building up of heavier elements during fusion formation in stars (see triple alpha process). A similar pairing pattern shows in the fact that the 256 known stable nuclides contain only five that have both an odd number of protons and an odd number of neutrons (odd-odd nuclei): ²H, ⁶Li, ¹⁰B, ¹⁴N, ^180mTa (the last one was predicted to decay but this process was never observed). Also, four long-lived radioactive odd-odd nuclides (⁴⁰K, ⁵⁰V, ¹³⁸La, ¹⁷⁶Lu) occur naturally. Most odd-odd nuclides are highly unstable with respect to beta decay, because the decay products are even-even, and are therefore more strongly bound, due to nuclear pairing effects.

Although isotopes exhibit nearly identical electronic and chemical behavior, their nuclear behavior varies dramatically. Adding neutrons to isotopes can vary their nuclear spins and nuclear shapes, causing differences in neutron capture cross-sections and gamma spectroscopy and nuclear magnetic resonance properties.

Occurrence in nature

Elements are composed of one or more naturally occurring isotopes, which are normally stable. Some elements have unstable (radioactive) isotopes, either because their decay is so slow that a fraction still remains since they were created (examples: uranium, potassium), or because they are continually created through cosmic radiation (tritium, carbon-14) or by decay from an isotope in the first category (radium, radon).

As discussed above, only 80 elements have any stable isotopes, and 26 of these have only one stable isotope. Thus, about two-thirds of stable elements occur naturally on Earth in multiple stable isotopes, with the largest number of stable isotopes for an element being ten, for tin (element number 50). There are about 94 elements found naturally on Earth (up to plutonium, element 94, inclusive), though some are detected only in very tiny amounts, such as plutonium-244. Scientists[1] estimate that the elements which occur naturally on Earth (some only as radioisotopes) occur as 339 isotopes (nuclides) in total. Only 256 of these naturally-occurring isotopes are stable in the sense of never having been observed to decay as of the present time. All the known stable isotopes occur naturally on Earth); the other 85 naturally-occurring isotopes are radioactive, but occur on Earth due to their relatively long half-lives, or else due to other means of ongoing natural production. An additional ~ 2700 radioactive isotopes not found in nature have been created in nuclear reactors and in particle accelerators. Many short-lived isotopes not found naturally on Earth have also been observed by spectroscopic analysis, being naturally created in stars or supernovae. An example is aluminum-26, which is not naturally found on Earth, but which is found in abundance on an astronomical scale.

The tabulated atomic masses of elements are averages that account for the presence of multiple isotopes with different masses. A good example is chlorine, having the composition ³⁵Cl, 75.8%, and ³⁷Cl, 24.2%, giving an atomic mass of 35.5. Values like this confounded scientists before the discovery of isotopes, as most light element atomic masses are close to integer multiples of hydrogen.

According to generally accepted cosmology only isotopes of hydrogen and helium, and traces of some isotopes of lithium, beryllium and boron were created at the Big Bang, while all other isotopes were synthesized later, in stars and supernovae, and in interactions between energetic particles such as cosmic rays, and previously-produced isotopes. The most common isotope of hydrogen has no neutrons at all. (See nucleosynthesis for details of the various processes thought to be responsible for isotope production.) The respective abundances of isotopes on Earth result from the quantities formed by these processes, their spread through the galaxy, and the rates of decay for isotopes that are unstable. After the initial coalescence of the solar system, isotopes were redistributed according to mass, and the isotopic composition of elements varies slightly from planet to planet. This sometimes makes it possible to trace the origin of meteorites.

Atomic mass of isotopes

The atomic mass (M_r) of an isotope is determined mainly by its mass number (i.e. number of nucleons in its nucleus). Small corrections are due to the binding energy of the nucleus (see Mass defect), to slightly different masses of neutron and proton, and to the mass of electron shell of the atom. One should take into account that the mass number is an integer dimensionless quantity, whereas the atomic mass is a real number expressed in atomic mass units. But the difference between these quantities is less than 1% in any case.

The atomic masses of naturally occurring isotopes of an element determine the atomic weight of the element. When the element contains N isotopes, the equation below is applied for the atomic weight M:

where M₁, M₂, ..., M_N are the atomic masses of each individual isotope, and η₁,...,η_N are the relative abundances of these isotopes.

Applications of isotopes

Several applications exist that capitalize on properties of the various isotopes of a given element.

Use of chemical and biological properties

• Isotope analysis is the determination of isotopic signature, the relative abundances of isotopes of a given element in a particular sample. For biogenic substances in particular, significant variations of isotopes of C, N and O can occur. Analysis of such variations has a wide range of applications, such as the detection of adulteration of food products.^[6] The identification of certain meteorites as having originated on Mars is based in part upon the isotopic signature of trace gases contained in them.^[7]
• Another common application is isotopic labeling, the use of unusual isotopes as tracers or markers in chemical reactions. Normally, atoms of a given element are indistinguishable from each other. However, by using isotopes of different masses, they can be distinguished by mass spectrometry or infrared spectroscopy (see "Properties"). For example, in 'stable isotope labeling with amino acids in cell culture (SILAC)' stable isotopes are used to quantify proteins. If radioactive isotopes are used, they can be detected by the radiation they emit (this is called radioisotopic labeling).
• A technique similar to radioisotopic labelling is radiometric dating: using the known half-life of an unstable element, one can calculate the amount of time that has elapsed since a known level of isotope existed. The most widely known example is radiocarbon dating used to determine the age of carbonaceous materials.
• Isotopic substitution can be used to determine the mechanism of a reaction via the kinetic isotope effect.

Use of nuclear properties

• Several forms of spectroscopy rely on the unique nuclear properties of specific isotopes. For example, nuclear magnetic resonance (NMR) spectroscopy can be used only for isotopes with a nonzero nuclear spin. The most common isotopes used with NMR spectroscopy are ¹H, ²D,¹⁵N, ¹³C, and ³¹P.
• Mössbauer spectroscopy also relies on the nuclear transitions of specific isotopes, such as ⁵⁷Fe.
• Radionuclides also have important uses. Nuclear power and nuclear weapons development require relatively large quantities of specific isotopes. The process of isotope separation represents a significant technological challenge, but more so with heavy elements such as uranium or plutonium, than with lighter elements such as hydrogen, lithium, carbon, nitrogen, and oxygen. The lighter elements are commonly separated by gas diffusion of their compounds such as CO and NO. Uranium isotopes have been separated in bulk by gas diffusion, gas centrifugation, laser ionization separation, and (in the Manhattan Project) by a type of production mass spectroscopy.

See also

• Atom
• Table of nuclides
• Isotopomer
• List of particles
• Isotopes are nuclides having the same number of protons; compare:
  • Isotones are nuclides having the same number of neutrons.
  • Isobars are nuclides having the same mass number, i.e. sum of protons plus neutrons.
  • Nuclear isomers are different excited states of the same type of nucleus. A transition from one isomer to another is accompanied by emission or absorption of a gamma ray, or the process of internal conversion. (Not to be confused with chemical isomers.)
• Bainbridge mass spectrometer

References

1. ^ IUPAC http://goldbook.iupac.org/I03331.html
2. ^ "Radioactives Missing From The Earth". http://www.don-lindsay-archive.org/creation/isotope_list.html. 
3. ^ "NuDat 2 Description". http://www.nndc.bnl.gov/nudat2/help/index.jsp. 
4. ^ Budzikiewicz H, Grigsby RD (2006). "Mass spectrometry and isotopes: a century of research and discussion". Mass spectrometry reviews 25 (1): 146–57. doi:10.1002/mas.20061. PMID 16134128. 
5. ^ Sonzogni, Alejandro. "Interactive Chart of Nuclides". National Nuclear Data Center: Brook haven National Laboratory. http://www.nndc.bnl.gov/chart/. Retrieved on 2008-06-06. 
6. ^ E. Jamin et al. (2003). "Improved Detection of Added Water in Orange Juice by Simultaneous Determination of the Oxygen-18/Oxygen-16 Isotope Ratios of Water and Ethanol Derived from Sugars"". J. Agric. Food Chem. 51: 5202. doi:10.1021/jf030167 m. http://pubs.acs.org/cgi-bin/article.cgi/jafcau/2003/51/i18/pdf/jf030167 m.pdf. 
7. ^ A. H. Treiman, J. D. Gleason and D. D. Bogard (2000). ""The SNC meteorites are from Mars"". Planet. Space. Sci. 48: 1213. doi:10.1016/S0032-0633(00)00105-7. http://www.sciencedirect.com/science?_ob=ArticleURL&_udi=B6V6T-41WBDHD-8&_user=2400262&_coverDate=10%2F31%2F2000&_alid=678948366&_rdoc=3&_fmt=summary&_orig=search&_cdi=5823&_sort=r&_docanchor=&view=c&_ct=89&_acct=C000057185&_version=1&_urlVersion=0&_userid=2400262&md5=c5ae2aa8ea60dbd76c2870048730a299. 

External links

• Nucleonica Nuclear Science Portal
• Nucleonica Nuclear Science Wiki
• International Atomic Energy Agency
• Atomic weights of all isotopes
• Atomgewichte, Zerfallsenergien und Halbwertszeiten aller Isotope
• Chart of the Nuclides produced by the Knolls Atomic Power Laboratory $25
• Exploring the Table of the Isotopes at the LBNL
• Current isotope research and information
• Radioactive Isotopes by the CDC
• Interacive Chart of the nuclides, isotopes and Periodic Table

This entry is from Wikipedia, the leading user-contributed encyclopedia. It may not have been reviewed by professional editors (see full disclaimer)

Dansk (Danish)
n. - isotop

Nederlands (Dutch)
isotoop

Français (French)
n. - isotope

Deutsch (German)
n. - (Phys.) Isotop

Ελληνική (Greek)
n. - ισότοπο

Italiano (Italian)
isotopo

Português (Portuguese)
n. - isótopo (m) (Quím.)

Русский (Russian)
изотоп

Español (Spanish)
n. - isótopo

Svenska (Swedish)
n. - isotop (kem.), (om atomer med samma atomnummer men olika antal neutroner)

中文（简体）(Chinese (Simplified))
同位素

中文（繁體）(Chinese (Traditional))
n. - 同位素

한국어 (Korean)
n. - 동위원소, 동위체, 핵종

日本語 (Japanese)
n. - 同位元素, 同位体, アイソトープ, 核種

العربيه (Arabic)
‏(الاسم) النظير : واحد النظائر‏

עברית (Hebrew)
n. - ‮צורות שונות של אותו יסוד השונות במשקלן האטומי אך לא בתכונות הכימיות, איזוטופ‬

If you are unable to view some languages clearly, click here.

To select your translation preferences click here.