isotopes

(ÿ -sŏ-tohps) Forms of an element in which the nuclei contain the same number of protons but different numbers of neutrons. For example, there are three isotopes of hydrogen: ‘normal’ hydrogen has a single proton, deuterium has one proton and one neutron, and tritium, which is a radioactive isotope, has one proton and two neutrons.

Isotopes

varieties of a chemical element occupying the same place in the Mendeleev periodic system of elements but having different atomic masses.

The chemical properties of atoms, that is, their assignment to a particular element, depend on the number of electrons and their arrangement within the electron shell of the atom.The location of a chemical element in the periodic system of elements is determined by its atomic number Z, which is equal to the number of electrons in the shell of the atom or, what is equivalent, to the number of protons within the atomic nucleus. The nucleus also contains neutrons, the mass of each of which is approximately equal to that of the proton. The number of neutrons N in the nucleus with a given Z may differ only within certain limits. For example, the helium nucleus (Z = 2) may contain 1,2, 4, or 6 neutrons. The total number of protons Z and of neutrons N in the nucleus (combined under the general term of nucleons) determines the mass of the nucleus and virtually the entire mass of the atom. This number A = Z + N is called the mass number of the atom. The ratio of protons to neutrons in the nucleus determines the stability or lack of stability of the nucleus and the type of its radioactive decay, its spin, its magnetic dipole moment, its electric quadrupole moment, and other properties. Thus, atoms with the same Z but with different numbers of neutrons N have identical chemical properties but different masses and properties of the nucleus. These varieties of atoms are also called isotopes. The term “nuclide” is used to denote any variety of atom, regardless of whether or not it belongs to the same element.

The mass number of the isotope is indicated at the upper left of the chemical symbol of the element. The isotopes of helium are designated by ³He, ⁴He, ⁶He, and ⁸He. In more detailed notation and the lower index indicates the number of protons Z, the upper left index indicates the number of neutrons N, and the upper right index indicates the mass number. In designating the isotope without using the symbol of the element, the mass number A is given after the name of the element, for example, helium-3 and helium-4.

The masses of the atoms M, expressed in atomic mass units, differ very slightly from integral values. For this reason, the difference M — A is always a proper fraction, whose absolute value is less than 1/2. Thus, the mass number A is the nearest integer to the atomic mass M. Knowledge of the atomic mass permits the determination of the total binding energy ε of all nucleons in the nucleus. This energy is expressed by the relation ε = ΔMc², where c is the velocity of light in a vacuum and εM is the difference between the total mass of all nucleons in the nucleus in the free state and the mass of the nucleus, which is equal to the mass of the neutral atom less the mass of all the electrons.

The first proof that substances having identical chemical properties can have different physical properties was obtained from the studies of the radioactive transformations of atoms of heavy elements. It was found in 1906–07 that the product of the radioactive decay of uranium, which is ionium, and the product of the radioactive decay of thorium, which is radiothorium, have the same chemical properties as thorium but differ from the latter by their atomic masses and by the characteristics of their radioactive decay. Moreover, it was found later that all three elements (ionium, radiothorium, and thorium) have identical optical and X-ray spectra. Such substances, which had identical chemical properties but different atomic masses and some physical properties, began to be called isotopes, as proposed by the British scientist F. Soddy.

After the discovery of the isotopes of heavy radioactive elements, the search for the isotopes of stable elements was begun. In 1913 the British physicist J. Thomson discovered an isotope of neon. The parabola method developed by him permitted the determination of the relation between the mass of an ion and its charge from the deflection in mutually parallel electric and magnetic fields of a thin beam of positive ions produced by a high-voltage electric discharge. Simultaneously with the ²⁰Ne atoms, Thomson observed a small admixture of heavier atoms. However, convincing proofs that the second component (heavier atoms) consisted of isotopes of neon were not obtained. Only the use of the first mass spectrograph, constructed in 1919 by the British physicist F. Aston, provided a reliable proof of the existence of the two isotopes ²⁰Ne and ²²Ne, which are found in nature in the abundances of 91 percent and 9 percent, respectively. The isotope ²¹Ne with the abundance of 0.26 percent was discovered subsequently, as were the isotopes of chlorine, mercury, and other elements. By about 1940 isotopic analysis was achieved for all elements existing on the earth. These efforts resulted in the identification of practically all stable and long-lived radioisotopes of the natural elements.

In 1934, I. Curie and F. Joliot artificially produced radioactive isotopes of nitrogen (³N), silicon (²⁸Si), and phosphorus (³⁰P), which are absent in nature. These experiments demonstrated the possibility of synthesizing new radioactive nuclides. In subsequent years, many isotopes of known elements were synthesized, and about 20 new elements were obtained with the help of nuclear reactions induced by neutrons and accelerated charged particles. A total of 276 stable isotopes of 81 naturally occurring elements and 1,500 radioisotopes of 105 naturally occurring and synthesized elements are known.

Analysis of the ratio of neutrons to protons for various isotopes of the same element indicates that the nuclei of stable isotopes, as well as of radioisotopes, which are stable with respect to beta decay, contain at least one neutron for every proton. Only two nuclides, ¹H and ³He, are an exception to this rule. The ratio of neutrons to protons in the nucleus increases in going to the heavier nuclei and reaches 1.6 for uranium and the transuranium elements.

Elements with odd Z have no more than two stable isotopes. As a rule, the number of neutrons N in these nuclei is even and, hence, the mass number A is odd. Most elements with an even Z have several stable isotopes, of which no more than two have odd A values. Tin has the greatest number of isotopes (ten), xenon has nine, and cadmium and tellurium each have eight. Many elements have seven isotopes.

These large variations in the number of stable isotopes of various elements are due to the complex dependence of the binding energy on the number of protons and neutrons in the nucleus. A change in the number of neutrons N with a given number of protons Z leads to changes in the binding energy and its stability with respect to various types of decay. The addition of neutrons makes the nucleus unstable with respect to the emission of an electron with the transformation of one neutron in the nucleus into a proton. For this reason, the neutron-rich isotopes of all elements are β^--active. Conversely, neutron deficiency permits the nucleus either to capture an electron from the electron shell or to emit a positron. In this case, one proton is transformed into a neutron and the ratio of protons to neutrons optimum for the stability of the nucleus is restored. The neutron-deficient isotopes of all elements undergo either electron capture or positron decay. Heavy nuclei also undergo alpha decay and spontaneous nuclear fission. The production of neutron-rich isotopes of elements is possible by several methods. One method is the capture of neutrons by the nuclei of stable isotopes; another involves the fission of heavy nuclei by neutrons or charged particles, as a result of which two neutron-rich nuclei are formed from a heavy nucleus with a high relative content of neutrons. Neutron-rich isotopes of light elements are effected in the reactions of multinucleon exchange during the interaction of accelerated heavy ions with matter. The synthesis of neutron-deficient isotopes is accomplished in nuclear reactions by the action of accelerated charged light particles or heavy ions.

All the stable isotopes on the earth formed as a result of nuclear processes that occurred in the distant past, and their abundance depends on the properties of the nuclei and on the initial conditions under which the processes occurred. The isotopic composition of the natural elements on the earth is, as a rule, constant. This can be explained by the fact that it is not subject to appreciable changes during the chemical and physical processes occurring on the earth. However, small fluctuations in the relative abundance of the isotopes are nevertheless observed for the light elements, in which the difference between the atomic masses of the isotopes is relatively large. These fluctuations are due to changes in the isotopic composition of the elements (fractionation of isotopes), which take place as the result of diffusion or of changes in the state of aggregation of matter and during some chemical reactions and other processes taking place continuously in the atmosphere and in the earth’s crust. Changes in the isotopic composition of the elements which are characterized by intense migrations in the biosphere (H, C, N, O, S) are also related to the activity of living organisms.

In the case of nuclides formed as a result of radioactive decay —for example, for the isotopes of lead—the different content of isotopes in various samples is due to different initial contents of their parent elements (U or Th) and the different geologic ages of the samples.

The common origin of the bodies in the solar system leads to the assumption that the isotopic composition of the elements in the earth samples is characteristic of the entire solar system (with some variations). Meteors and the deep strata of the earth’s crust exhibit about the same ratio ¹⁶O/¹⁸O. Astrophysical research has detected deviations in the isotopic composition of elements in stellar material and the interstellar medium from that on the earth. For example, the ¹²C/¹³C ratio for the carbon R-type stars ranges from 4–5 to the terrestrial value.

Natural elements mixed with their radioisotopes make it possible, by using radiation detectors, to trace various chemical and physical processes in which a given element participates. This method has been widely used in biology, chemistry, medicine, and technology. Stable isotopes are added in some cases, and their presence is subsequently detected by mass-spectrographic methods.

The separation of certain isotopes from their natural or artificially obtained mixtures and the enrichment of these mixtures by certain isotopes constitute important problems in this field.

REFERENCES

Aston, F.W. Mass-spektry i izotopy. Moscow, 1948. (Translated from English.)
Kravtsov, V.A. Massy atomov i energii sviazi iader. Moscow, 1965.
Lederer, C. M., J.M. Hollander, and I. Perlman. Table of Isotopes, 6thed. New York, 1967.

N. I. Tarantin

单词	isotopes
释义	isotopes isotopes 1. Atoms of the same element (all chemically identical) having the same atomic number but containing different numbers of neutrons, giving a different mass number.2. Atoms of an element with an identical number of protons but differing numbers of neutrons. Translations isotopi isotopes isotopes (ÿ -sŏ-tohps) Forms of an element in which the nuclei contain the same number of protons but different numbers of neutrons. For example, there are three isotopes of hydrogen: ‘normal’ hydrogen has a single proton, deuterium has one proton and one neutron, and tritium, which is a radioactive isotope, has one proton and two neutrons. Isotopes varieties of a chemical element occupying the same place in the Mendeleev periodic system of elements but having different atomic masses. The chemical properties of atoms, that is, their assignment to a particular element, depend on the number of electrons and their arrangement within the electron shell of the atom.The location of a chemical element in the periodic system of elements is determined by its atomic number Z, which is equal to the number of electrons in the shell of the atom or, what is equivalent, to the number of protons within the atomic nucleus. The nucleus also contains neutrons, the mass of each of which is approximately equal to that of the proton. The number of neutrons N in the nucleus with a given Z may differ only within certain limits. For example, the helium nucleus (Z = 2) may contain 1,2, 4, or 6 neutrons. The total number of protons Z and of neutrons N in the nucleus (combined under the general term of nucleons) determines the mass of the nucleus and virtually the entire mass of the atom. This number A = Z + N is called the mass number of the atom. The ratio of protons to neutrons in the nucleus determines the stability or lack of stability of the nucleus and the type of its radioactive decay, its spin, its magnetic dipole moment, its electric quadrupole moment, and other properties. Thus, atoms with the same Z but with different numbers of neutrons N have identical chemical properties but different masses and properties of the nucleus. These varieties of atoms are also called isotopes. The term “nuclide” is used to denote any variety of atom, regardless of whether or not it belongs to the same element. The mass number of the isotope is indicated at the upper left of the chemical symbol of the element. The isotopes of helium are designated by ³He, ⁴He, ⁶He, and ⁸He. In more detailed notation and the lower index indicates the number of protons Z, the upper left index indicates the number of neutrons N, and the upper right index indicates the mass number. In designating the isotope without using the symbol of the element, the mass number A is given after the name of the element, for example, helium-3 and helium-4. The masses of the atoms M, expressed in atomic mass units, differ very slightly from integral values. For this reason, the difference M — A is always a proper fraction, whose absolute value is less than 1/2. Thus, the mass number A is the nearest integer to the atomic mass M. Knowledge of the atomic mass permits the determination of the total binding energy ε of all nucleons in the nucleus. This energy is expressed by the relation ε = ΔMc², where c is the velocity of light in a vacuum and εM is the difference between the total mass of all nucleons in the nucleus in the free state and the mass of the nucleus, which is equal to the mass of the neutral atom less the mass of all the electrons. The first proof that substances having identical chemical properties can have different physical properties was obtained from the studies of the radioactive transformations of atoms of heavy elements. It was found in 1906–07 that the product of the radioactive decay of uranium, which is ionium, and the product of the radioactive decay of thorium, which is radiothorium, have the same chemical properties as thorium but differ from the latter by their atomic masses and by the characteristics of their radioactive decay. Moreover, it was found later that all three elements (ionium, radiothorium, and thorium) have identical optical and X-ray spectra. Such substances, which had identical chemical properties but different atomic masses and some physical properties, began to be called isotopes, as proposed by the British scientist F. Soddy. After the discovery of the isotopes of heavy radioactive elements, the search for the isotopes of stable elements was begun. In 1913 the British physicist J. Thomson discovered an isotope of neon. The parabola method developed by him permitted the determination of the relation between the mass of an ion and its charge from the deflection in mutually parallel electric and magnetic fields of a thin beam of positive ions produced by a high-voltage electric discharge. Simultaneously with the ²⁰Ne atoms, Thomson observed a small admixture of heavier atoms. However, convincing proofs that the second component (heavier atoms) consisted of isotopes of neon were not obtained. Only the use of the first mass spectrograph, constructed in 1919 by the British physicist F. Aston, provided a reliable proof of the existence of the two isotopes ²⁰Ne and ²²Ne, which are found in nature in the abundances of 91 percent and 9 percent, respectively. The isotope ²¹Ne with the abundance of 0.26 percent was discovered subsequently, as were the isotopes of chlorine, mercury, and other elements. By about 1940 isotopic analysis was achieved for all elements existing on the earth. These efforts resulted in the identification of practically all stable and long-lived radioisotopes of the natural elements. In 1934, I. Curie and F. Joliot artificially produced radioactive isotopes of nitrogen (³N), silicon (²⁸Si), and phosphorus (³⁰P), which are absent in nature. These experiments demonstrated the possibility of synthesizing new radioactive nuclides. In subsequent years, many isotopes of known elements were synthesized, and about 20 new elements were obtained with the help of nuclear reactions induced by neutrons and accelerated charged particles. A total of 276 stable isotopes of 81 naturally occurring elements and 1,500 radioisotopes of 105 naturally occurring and synthesized elements are known. Analysis of the ratio of neutrons to protons for various isotopes of the same element indicates that the nuclei of stable isotopes, as well as of radioisotopes, which are stable with respect to beta decay, contain at least one neutron for every proton. Only two nuclides, ¹H and ³He, are an exception to this rule. The ratio of neutrons to protons in the nucleus increases in going to the heavier nuclei and reaches 1.6 for uranium and the transuranium elements. Elements with odd Z have no more than two stable isotopes. As a rule, the number of neutrons N in these nuclei is even and, hence, the mass number A is odd. Most elements with an even Z have several stable isotopes, of which no more than two have odd A values. Tin has the greatest number of isotopes (ten), xenon has nine, and cadmium and tellurium each have eight. Many elements have seven isotopes. These large variations in the number of stable isotopes of various elements are due to the complex dependence of the binding energy on the number of protons and neutrons in the nucleus. A change in the number of neutrons N with a given number of protons Z leads to changes in the binding energy and its stability with respect to various types of decay. The addition of neutrons makes the nucleus unstable with respect to the emission of an electron with the transformation of one neutron in the nucleus into a proton. For this reason, the neutron-rich isotopes of all elements are β^--active. Conversely, neutron deficiency permits the nucleus either to capture an electron from the electron shell or to emit a positron. In this case, one proton is transformed into a neutron and the ratio of protons to neutrons optimum for the stability of the nucleus is restored. The neutron-deficient isotopes of all elements undergo either electron capture or positron decay. Heavy nuclei also undergo alpha decay and spontaneous nuclear fission. The production of neutron-rich isotopes of elements is possible by several methods. One method is the capture of neutrons by the nuclei of stable isotopes; another involves the fission of heavy nuclei by neutrons or charged particles, as a result of which two neutron-rich nuclei are formed from a heavy nucleus with a high relative content of neutrons. Neutron-rich isotopes of light elements are effected in the reactions of multinucleon exchange during the interaction of accelerated heavy ions with matter. The synthesis of neutron-deficient isotopes is accomplished in nuclear reactions by the action of accelerated charged light particles or heavy ions. All the stable isotopes on the earth formed as a result of nuclear processes that occurred in the distant past, and their abundance depends on the properties of the nuclei and on the initial conditions under which the processes occurred. The isotopic composition of the natural elements on the earth is, as a rule, constant. This can be explained by the fact that it is not subject to appreciable changes during the chemical and physical processes occurring on the earth. However, small fluctuations in the relative abundance of the isotopes are nevertheless observed for the light elements, in which the difference between the atomic masses of the isotopes is relatively large. These fluctuations are due to changes in the isotopic composition of the elements (fractionation of isotopes), which take place as the result of diffusion or of changes in the state of aggregation of matter and during some chemical reactions and other processes taking place continuously in the atmosphere and in the earth’s crust. Changes in the isotopic composition of the elements which are characterized by intense migrations in the biosphere (H, C, N, O, S) are also related to the activity of living organisms. In the case of nuclides formed as a result of radioactive decay —for example, for the isotopes of lead—the different content of isotopes in various samples is due to different initial contents of their parent elements (U or Th) and the different geologic ages of the samples. The common origin of the bodies in the solar system leads to the assumption that the isotopic composition of the elements in the earth samples is characteristic of the entire solar system (with some variations). Meteors and the deep strata of the earth’s crust exhibit about the same ratio ¹⁶O/¹⁸O. Astrophysical research has detected deviations in the isotopic composition of elements in stellar material and the interstellar medium from that on the earth. For example, the ¹²C/¹³C ratio for the carbon R-type stars ranges from 4–5 to the terrestrial value. Natural elements mixed with their radioisotopes make it possible, by using radiation detectors, to trace various chemical and physical processes in which a given element participates. This method has been widely used in biology, chemistry, medicine, and technology. Stable isotopes are added in some cases, and their presence is subsequently detected by mass-spectrographic methods. The separation of certain isotopes from their natural or artificially obtained mixtures and the enrichment of these mixtures by certain isotopes constitute important problems in this field. REFERENCES Aston, F.W. Mass-spektry i izotopy. Moscow, 1948. (Translated from English.) Kravtsov, V.A. Massy atomov i energii sviazi iader. Moscow, 1965. Lederer, C. M., J.M. Hollander, and I. Perlman. Table of Isotopes, 6thed. New York, 1967. N. I. Tarantin MedicalSeeisotope
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