Isotopes

Atoms of the same element that have different mass numbers are called isotopes. Atoms of isotopes of the same element have the same number of protons (Z) and differ from each other in the number of neutrons (N).

Isotopes of various elements do not have their own names, but repeat the name of the element; in this case, the atomic mass of a given isotope - its only difference from other isotopes of the same element - is reflected using a superscript in the chemical formula of the element: for example, for uranium isotopes - 235 U, 238 U. The only exception to the rules of isotope nomenclature is element No. 1 - hydrogen. All three currently known isotopes of hydrogen have not only their own special chemical symbols, but also their own name: 1 H - protium, 2 D - deuterium, 3 T - tritium; in this case, the protium nucleus is simply one proton, the deuterium nucleus contains one proton and one neutron, the tritium nucleus contains one proton and two neutrons. The names of hydrogen isotopes have historically developed this way because the relative difference in the masses of hydrogen isotopes caused by the addition of one neutron is the maximum among all chemical elements.

All isotopes can be divided into stable (stable), that is, not subject to spontaneous decay of atomic nuclei into parts (decay in this case is called radioactive), and unstable (unstable) - radioactive, that is, subject to radioactive decay. Most elements widespread in nature consist of a mixture of two or more stable isotopes: for example, 16 O, 12 C. Of all the elements, tin has the largest number of stable isotopes (10 isotopes), and, for example, aluminum exists in nature in the form of only one stable isotope - the rest of its known isotopes are unstable. The nuclei of unstable isotopes decay spontaneously, releasing b particles and c particles (electrons) until a stable isotope of another element is formed: for example, the decay of 238 U (radioactive uranium) ends with the formation of 206 Pb (a stable isotope of lead). When studying isotopes, it was found that they do not differ in chemical properties, which, as we know, are determined by the charge of their nuclei and do not depend on the mass of the nuclei.

Electronic shells

The electron shell of an atom is a region of space where electrons are likely to be located, characterized by the same value of the principal quantum number n and, as a consequence, located at close energy levels. Each electron shell can have a certain maximum number of electrons.

Starting from the value of the main quantum number n = 1, the energy levels (layers) are designated K, L, M and N. They are divided into sublevels (sublayers) that differ from each other in the binding energy with the nucleus. The number of sublevels is equal to the value of the main quantum number, but does not exceed four: the 1st level has one sublevel, the 2nd - two, the 3rd - three, the 4th - four sublevels. Sublevels, in turn, consist of orbitals. It is customary to denote sublevels with Latin letters, s is the first sublevel of each energy level closest to the nucleus; it consists of one s-orbital, p - the second sublevel, consists of three p-orbitals; d is the third sublevel, it consists of five d-orbitals; f is the fourth sublevel, contains seven f orbitals. Thus, for each value of n there are n 2 orbitals. Each orbital can contain no more than two electrons - the Pauli principle. If there is one electron in an orbital, then it is called unpaired; if there are two, then these are paired electrons. The Pauli principle explains the formula N=2n 2. If the first level K(n=1) contains 1 2 = 1 orbital, and each orbital has 2 electrons, then the maximum number of electrons will be 2*1 2 =2; L (n = 2) =8; M (n = 3) =18; N (n = 4) =32.

When studying the properties of radioactive elements, it was discovered that the same chemical element can contain atoms with different nuclear masses. At the same time, they have the same nuclear charge, that is, these are not impurities of foreign substances, but the same substance.

What are isotopes and why do they exist?

In Mendeleev's periodic table, both this element and atoms of a substance with different nuclear masses occupy one cell. Based on the above, such varieties of the same substance were given the name “isotopes” (from the Greek isos - identical and topos - place). So, isotopes- these are varieties of a given chemical element, differing in the mass of atomic nuclei.

According to the accepted neutron-proton model of the nucleus, it was possible to explain the existence of isotopes as follows: the nuclei of some atoms of a substance contain different numbers of neutrons, but the same number of protons. In fact, the nuclear charge of isotopes of one element is the same, therefore, the number of protons in the nucleus is the same. Nuclei differ in mass; accordingly, they contain different numbers of neutrons.

Stable and unstable isotopes

Isotopes can be stable or unstable. To date, about 270 stable isotopes and more than 2000 unstable ones are known. Stable isotopes- These are varieties of chemical elements that can exist independently for a long time.

Most of unstable isotopes was obtained artificially. Unstable isotopes are radioactive, their nuclei are subject to the process of radioactive decay, that is, spontaneous transformation into other nuclei, accompanied by the emission of particles and/or radiation. Almost all radioactive artificial isotopes have very short half-lives, measured in seconds or even fractions of seconds.

How many isotopes can a nucleus contain?

The nucleus cannot contain an arbitrary number of neutrons. Accordingly, the number of isotopes is limited. Even number of protons elements, the number of stable isotopes can reach ten. For example, tin has 10 isotopes, xenon has 9, mercury has 7, and so on.

Those elements the number of protons is odd, can have only two stable isotopes. Some elements have only one stable isotope. These are substances such as gold, aluminum, phosphorus, sodium, manganese and others. Such variations in the number of stable isotopes of different elements are associated with the complex dependence of the number of protons and neutrons on the binding energy of the nucleus.

Almost all substances in nature exist in the form of a mixture of isotopes. The number of isotopes in a substance depends on the type of substance, atomic mass and the number of stable isotopes of a given chemical element.

Isotopes- varieties of atoms (and nuclei) of a chemical element that have the same atomic (ordinal) number, but at the same time different mass numbers.

The term isotope is formed from the Greek roots isos (ἴσος "equal") and topos (τόπος "place"), meaning "same place"; Thus, the meaning of the name is that different isotopes of the same element occupy the same position in the periodic table.

Three natural isotopes of hydrogen. The fact that each isotope has one proton has variants of hydrogen: the identity of the isotope is determined by the number of neutrons. From left to right, the isotopes are protium (1H) with zero neutrons, deuterium (2H) with one neutron, and tritium (3H) with two neutrons.

The number of protons in the nucleus of an atom is called the atomic number and is equal to the number of electrons in a neutral (non-ionized) atom. Each atomic number identifies a specific element, but not an isotope; An atom of a given element can have a wide range in the number of neutrons. The number of nucleons (both protons and neutrons) in the nucleus is the mass number of the atom, and each isotope of a given element has a different mass number.

For example, carbon-12, carbon-13, and carbon-14 are three isotopes of elemental carbon with mass numbers 12, 13, and 14, respectively. The atomic number of carbon is 6, which means that each carbon atom has 6 protons, so the neutron numbers of these isotopes are 6, 7 and 8 respectively.

Nuklides And isotopes

Nuclide refers to a nucleus, not an atom. Identical nuclei belong to the same nuclide, for example, each nucleus of the nuclide carbon-13 consists of 6 protons and 7 neutrons. The nuclide concept (relating to individual nuclear species) emphasizes nuclear properties over chemical properties, while the isotope concept (grouping all the atoms of each element) emphasizes chemical reaction over nuclear reaction. The neutron number has a large influence on the properties of nuclei, but its effect on chemical properties is negligible for most elements. Even in the case of the lightest elements, where the ratio of neutrons to atomic number varies most between isotopes, it usually has only a minor effect, although it does matter in some cases (for hydrogen, the lightest element, the isotope effect is large to have a large effect for biology). Because isotope is an older term, it is better known than nuclide and is still sometimes used in contexts where nuclide may be more appropriate, such as nuclear technology and nuclear medicine.

Designations

An isotope or nuclide is identified by the name of the specific element (this indicates the atomic number), followed by a hyphen and mass number (for example, helium-3, helium-4, carbon-12, carbon-14, uranium-235, and uranium-239). When a chemical symbol is used, e.g. "C" for carbon, standard notation (now known as "AZE-notation" because A is the mass number, Z is the atomic number, and E is for the element) - indicate the mass number (number of nucleons) with a superscript at the top left of chemical symbol and indicate the atomic number with a subscript in the lower left corner). Because the atomic number is given by the symbol of the element, usually only the mass number is given in a superscript and no atomic index is given. The letter m is sometimes added after the mass number to indicate a nuclear isomer, a metastable or energetically excited nuclear state (as opposed to the lowest energy ground state), for example, 180m 73Ta (tantalum-180m).

Radioactive, primary and stable isotopes

Some isotopes are radioactive and are therefore called radioisotopes or radionuclides, while others have never been observed to decay radioactively and are called stable isotopes or stable nuclides. For example, 14 C is the radioactive form of carbon, while 12 C and 13 C are stable isotopes. There are approximately 339 naturally occurring nuclides on Earth, of which 286 are primordial nuclides, meaning they have existed since the formation of the Solar System.

The original nuclides include 32 nuclides with very long half-lives (over 100 million years) and 254 that are formally considered "stable nuclides" because they were not observed to decay. In most cases, for obvious reasons, if an element has stable isotopes then those isotopes dominate the elemental abundance found on Earth and in the Solar System. However, in the case of three elements (tellurium, indium and rhenium), the most common isotope found in nature is actually one (or two) extremely long-lived radioisotope(s) of the element, despite the fact that these elements have one or more stable isotopes.

The theory predicts that many apparently "stable" isotopes/nuclides are radioactive, with extremely long half-lives (ignoring the possibility of proton decay, which would make all nuclides eventually unstable). Of the 254 nuclides that have never been observed, only 90 of them (all of the first 40 elements) are theoretically stable to all known forms of decay. Element 41 (niobium) is theoretically unstable by spontaneous fission, but this has never been discovered. Many other stable nuclides are in theory energetically susceptible to other known decay forms, such as alpha decay or double beta decay, but the decay products have not yet been observed, and so these isotopes are considered to be "observationally stable". The predicted half-lives for these nuclides often greatly exceed the estimated age of the Universe, and in fact there are also 27 known radionuclides with half-lives longer than the age of the Universe.

Radioactive nuclides created artificially, currently there are 3,339 known nuclides. These include 905 nuclides that are either stable or have half-lives greater than 60 minutes.

Properties of isotopes

Chemical and molecular properties

A neutral atom has the same number of electrons as protons. Thus, different isotopes of a given element have the same number of electrons and have similar electronic structures. Since the chemical behavior of an atom is largely determined by its electronic structure, different isotopes exhibit nearly identical chemical behavior.

The exception to this is the kinetic isotope effect: due to their large masses, heavier isotopes tend to react somewhat more slowly than lighter isotopes of the same element. This is most pronounced for protium (1 H), deuterium (2 H), and tritium (3 H), since deuterium has twice the mass of protium and tritium has three times the mass of protium. These differences in mass also affect the behavior of their respective chemical bonds, changing the center of gravity (reduced mass) of atomic systems. However, for heavier elements the relative mass differences between isotopes are much smaller, so mass difference effects in chemistry are usually negligible. (Heavy elements also have relatively more neutrons than lighter elements, so the ratio of nuclear mass to total electron mass is somewhat larger).

Likewise, two molecules that differ only in the isotopes of their atoms (isotopologues) have the same electronic structure and hence almost indistinguishable physical and chemical properties (again, with the primary exceptions being deuterium and tritium). The vibrational modes of a molecule are determined by its shape and the masses of its constituent atoms; Therefore, different isotopologues have different sets of vibrational modes. Because vibrational modes allow a molecule to absorb photons of appropriate energies, isotopologues have different optical properties in the infrared.

Nuclear properties and stability

Isotopic half-lives. The graph for stable isotopes deviates from the Z = N line as the element number Z increases

Atomic nuclei consist of protons and neutrons bound together by a residual strong force. Because protons are positively charged, they repel each other. Neutrons, which are electrically neutral, stabilize the nucleus in two ways. Their contact pushes the protons apart slightly, reducing the electrostatic repulsion between the protons, and they exert an attractive nuclear force on each other and on the protons. For this reason, one or more neutrons are required for two or more protons to bind to a nucleus. As the number of protons increases, so does the ratio of neutrons to protons required to provide a stable nucleus (see graph on the right). For example, although the neutron:proton ratio of 3 2 He is 1:2, the neutron:proton ratio is 238 92 U
More than 3:2. A number of lighter elements have stable nuclides with a 1:1 ratio (Z = N). Nuclide 40 20 Ca (calcium-40) is the observationally heaviest stable nuclide with the same number of neutrons and protons; (Theoretically, the heaviest stable one is sulfur-32). All stable nuclides heavier than calcium-40 contain more neutrons than protons.

Number of isotopes per element

Of the 81 elements with stable isotopes, the highest number of stable isotopes observed for any element is ten (for the element tin). No element has nine stable isotopes. Xenon is the only element with eight stable isotopes. Four elements have seven stable isotopes, eight of which have six stable isotopes, ten have five stable isotopes, nine have four stable isotopes, five have three stable isotopes, 16 have two stable isotopes, and 26 elements have only one (of which 19 are so-called mononuclide elements, having a single primordial stable isotope that dominates and fixes the atomic weight of the natural element with high accuracy; 3 radioactive mononuclide elements are also present). There are a total of 254 nuclides that have not been observed to decay. For the 80 elements that have one or more stable isotopes, the average number of stable isotopes is 254/80 = 3.2 isotopes per element.

Even and odd numbers of nucleons

Protons: The neutron ratio is not the only factor affecting nuclear stability. It also depends on the parity or oddness of its atomic number Z, the number of neutrons N, hence their sum of mass number A. Odd both Z and N tend to lower the nuclear binding energy, creating odd nuclei that are generally less stable. This significant difference in nuclear binding energy between neighboring nuclei, especially odd isobars, has important consequences: unstable isotopes with suboptimal numbers of neutrons or protons decay by beta decay (including positron decay), electron capture, or other exotic means such as spontaneous fission and decay clusters.

Most stable nuclides are an even number of protons and an even number of neutrons, where the Z, N and A numbers are all even. Odd stable nuclides are divided (approximately evenly) into odd ones.

Atomic number

The 148 even proton, even neutron (NE) nuclides account for ~58% of all stable nuclides. There are also 22 primordial long-lived even nuclides. As a result, each of the 41 even-numbered elements from 2 to 82 has at least one stable isotope, and most of these elements have multiple primary isotopes. Half of these even-numbered elements have six or more stable isotopes. The extreme stability of helium-4, due to the double compound of two protons and two neutrons, prevents any nuclides containing five or eight nucleons from existing long enough to serve as platforms for the accumulation of heavier elements through nuclear fusion.

These 53 stable nuclides have an even number of protons and an odd number of neutrons. They are a minority compared to the even isotopes, which are approximately 3 times more abundant. Among the 41 even-Z elements that have a stable nuclide, only two elements (argon and cerium) do not have even-odd stable nuclides. One element (tin) has three. There are 24 elements that have one even-odd nuclide and 13 that have two odd-even nuclides.

Because of their odd neutron numbers, odd-even nuclides tend to have large neutron capture cross sections due to the energy that arises from neutron coupling effects. These stable nuclides may be unusually abundant in nature, mainly because to form and enter primordial abundance they must escape neutron capture to form yet other stable even-odd isotopes during the s process and r neutron capture process during nucleosynthesis.

Odd atomic number

The 48 stable odd-proton and even-neutron nuclides, stabilized by their even number of paired neutrons, form the majority of stable isotopes of odd elements; Very few odd-proton-odd neutron nuclides make up the others. There are 41 odd elements from Z = 1 to 81, of which 39 have stable isotopes (the elements technetium (43 Tc) and promethium (61 Pm) have no stable isotopes). Of these 39 odd Z elements, 30 elements (including hydrogen-1, where 0 neutrons are even) have one stable even-odd isotope, and nine elements: chlorine (17 Cl), potassium (19K), copper (29 Cu), gallium ( 31 Ga), Bromine (35 Br), silver (47 Ag), antimony (51 Sb), iridium (77 Ir) and thallium (81 Tl) each have two odd-even stable isotopes. This gives 30 + 2 (9) = 48 stable even-even isotopes.

Only five stable nuclides contain both an odd number of protons and an odd number of neutrons. The first four "odd-odd" nuclides occur in low molecular weight nuclides for which changing a proton to a neutron or vice versa will result in a very lopsided proton-neutron ratio.

The only completely "stable", odd-odd nuclide is 180m 73 Ta, which is considered the rarest of the 254 stable isotopes and is the only primordial nuclear isomer that has not yet been observed to decay, despite experimental attempts.

Odd number of neutrons

Actinides with an odd number of neutrons tend to fission (with thermal neutrons), while those with an even neutron number generally do not, although they do fission with fast neutrons. All observationally stable odd-odd nuclides have non-zero integer spin. This is because a single unpaired neutron and an unpaired proton have a greater nuclear force attraction towards each other if their spins are aligned (producing a total spin of at least 1 unit) rather than aligned.

Occurrence in nature

Elements are made up of one or more naturally occurring isotopes. Unstable (radioactive) isotopes are either primary or postprimary. The primordial isotopes were the product of stellar nucleosynthesis or another type of nucleosynthesis such as cosmic ray fission, and have persisted down to the present day because their decay rates are so low (e.g., uranium-238 and potassium-40). Post-natural isotopes were created by cosmic ray bombardment as cosmogenic nuclides (eg tritium, carbon-14) or the decay of a radioactive primordial isotope into the daughter of a radioactive radiogenic nuclide (eg uranium to radium). Several isotopes are naturally synthesized as nucleogenic nuclides by other natural nuclear reactions, such as when neutrons from natural nuclear fission are absorbed by another atom.

As discussed above, only 80 elements have stable isotopes, and 26 of them have only one stable isotope. Thus, about two-thirds of the stable elements occur naturally on Earth in several stable isotopes, with the largest number of stable isotopes for an element being ten, for tin (50Sn). There are about 94 elements on Earth (up to and including plutonium), although some are only found in very small quantities, such as plutonium-244. Scientists believe that elements that occur naturally on Earth (some only as radioisotopes) occur as 339 isotopes (nuclides) in total. Only 254 of these natural isotopes are stable in the sense that they have not been observed to date. Another 35 primordial nuclides (for a total of 289 primordial nuclides) are radioactive with known half-lives, but have half-lives of more than 80 million years, allowing them to exist since the beginning of the Solar System.

All known stable isotopes occur naturally on Earth; Other naturally occurring isotopes are radioactive, but because of their relatively long half-lives or other means of continuous natural production. These include the cosmogenic nuclides mentioned above, nucleogenic nuclides, and any radiogenic isotopes resulting from the ongoing decay of a primary radioactive isotope such as radon and radium from uranium.

Another ~3000 radioactive isotopes not found in nature have been created in nuclear reactors and particle accelerators. Many short-lived isotopes not found naturally on Earth have also been observed by spectroscopic analysis, naturally produced in stars or supernovae. An example is aluminum-26, which is not naturally found on Earth but 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. Before the discovery of isotopes, empirically determined, non-integrated atomic mass values ​​confused scientists. For example, a sample of chlorine contains 75.8% chlorine-35 and 24.2% chlorine-37, giving an average atomic mass of 35.5 atomic mass units.

According to the generally accepted theory of cosmology, only isotopes of hydrogen and helium, traces of some isotopes of lithium and beryllium, and possibly some boron, were created in the Big Bang, and all other isotopes were synthesized later, in stars and supernovae, and in interactions between energetic particles , such as cosmic rays, and previously obtained isotopes. The corresponding isotopic abundances of isotopes on Earth are determined by the quantities produced by these processes, their propagation through the galaxy, and the decay rate of the isotopes, which are unstable. After the initial solar system merger, isotopes were redistributed according to mass and the isotopic composition of elements varies slightly from planet to planet. This sometimes allows one to trace the origin of meteorites.

Atomic mass of isotopes

The atomic mass (mr) of an isotope is determined primarily by its mass number (i.e., the number of nucleons in its nucleus). Small corrections are due to the binding energy of the nucleus, the small difference in mass between the proton and neutron, and the mass of the electrons associated with the atom.

Mass number - dimensionless quantity. Atomic mass, on the other hand, is measured using an atomic mass unit based on the mass of a carbon-12 atom. It is denoted by the symbols "u" (for the unified atomic mass unit) or "Da" (for the dalton).

The atomic masses of an element's natural isotopes determine the atomic mass of the element. When an element contains N isotopes, the following expression applies for the average atomic mass:

Where m 1, m 2, ..., mN are the atomic masses of each individual isotope, and x 1, ..., xN are the relative abundance of these isotopes.

Application of isotopes

There are several applications that take advantage of the properties of different isotopes of a given element. Isotopic separation is an important technological problem, especially with heavy elements such as uranium or plutonium. Lighter elements such as lithium, carbon, nitrogen and oxygen are usually separated by gaseous diffusion of their compounds such as CO and NO. The separation of hydrogen and deuterium is unusual because it is based on chemical rather than physical properties, such as in the Girdler sulfide process. Uranium isotopes were separated by volume by gas diffusion, gas centrifugation, laser ionization separation, and (in the Manhattan Project) mass spectrometry-type production.

Use of chemical and biological properties

• Isotope analysis is the determination of the isotope signature, the relative abundance of isotopes of a given element in a particular sample. For nutrients in particular, significant variations in the C, N, and O isotopes can occur. The analysis of such variations has a wide range of applications, such as detecting adulteration in food products or the geographic origin of products using isoscapes. The identification of some meteorites that originated on Mars is based in part on the isotopic signature of the trace gases they contain.
• Isotopic substitution can be used to determine the mechanism of a chemical reaction through the kinetic isotope effect.
• Another common application is isotope labeling, the use of unusual isotopes as indicators or markers in chemical reactions. Usually the atoms of a given element are indistinguishable from each other. However, by using isotopes of different masses, even different non-radioactive stable isotopes can be distinguished using mass spectrometry or infrared spectroscopy. For example, in “stable isotope labeling of 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 radioisotope tagging).
• Isotopes are commonly used to determine the concentration of various elements or substances using the isotope dilution method, in which known quantities of isotopically substituted compounds are mixed with samples and the isotopic signatures of the resulting mixtures are determined using mass spectrometry.

Using Nuclear Properties

• A similar method to radioisotope tagging is radiometric dating: using the known half-life of an unstable element, the time that has passed since the existence of a known concentration of the isotope can be calculated. The most widely known example is radiocarbon dating, which is used to determine the age of carbonaceous materials.
• Some forms of spectroscopy rely on the unique nuclear properties of specific isotopes, both radioactive and stable. For example, nuclear magnetic resonance (NMR) spectroscopy can only be used for isotopes with non-zero nuclear spin. The most common isotopes used in NMR spectroscopy are 1 H, 2 D, 15 N, 13 C and 31 P.
• Mössbauer spectroscopy also relies on nuclear transitions of specific isotopes, such as 57Fe.

A certain element that has the same but different. They have nuclei with the same number and diversity. number, have the same structure of electron shells and occupy the same place in the periodicity. chemical system elements. The term "isotopes" was proposed in 1910 by F. Soddy to designate chemically indistinguishable varieties that differ in their physical properties. (primarily radioactive) Saints. Stable isotopes were first discovered in 1913 by J. Thomson using the so-called he developed. the method of parabolas - the prototype of the modern one. . He found that Ne has at least 2 varieties with a wt. parts 20 and 22. The names and symbols of isotopes are usually the names and symbols of the corresponding chemicals. elements; point to the top left of the symbol. For example, to indicate natural isotopes use the notation 35 Cl and 37 Cl; sometimes the element is also indicated at the bottom left, i.e. write 35 17 Cl and 37 17 Cl. Only isotopes of the lightest element, hydrogen, with wt. parts 1, 2 and 3 have special. names and symbols: (1 1 H), (D, or 2 1 H) and (T, or 3 1 H), respectively. Due to the large difference in masses, the behavior of these isotopes differs significantly (see,). Stable isotopes occur in all even and most odd elements with[ 83. The number of stable isotopes of elements with even numbers may be equals 10 (e.g. y); Odd-numbered elements have no more than two stable isotopes. Known approx. 280 stable and more than 2000 radioactive isotopes of 116 natural and artificially obtained elements. For each element, the content of individual isotopes in nature. the mixture undergoes small fluctuations, which can often be neglected. More means. fluctuations in the isotopic composition are observed for meteorites and other celestial bodies. The constancy of the isotopic composition leads to the constancy of the elements found on Earth, which is the average value of the mass of a given element, found taking into account the abundance of isotopes in nature. Fluctuations in the isotopic composition of light elements are associated, as a rule, with changes in the isotopic composition during decomposition. processes occurring in nature (, etc.). For the heavy element Pb, variations in the isotopic composition of different samples are explained by different factors. content in, and other sources and - the ancestors of natural sciences. . Differences in the properties of isotopes of a given element are called. . Important practical The task is to obtain from nature. mixtures of individual isotopes -

The content of the article

ISOTOPES– varieties of the same chemical element that are similar in their physicochemical properties, but have different atomic masses. The name "isotopes" was proposed in 1912 by the English radiochemist Frederick Soddy, who formed it from two Greek words: isos - identical and topos - place. Isotopes occupy the same place in the cell of Mendeleev's periodic table of elements.

An atom of any chemical element consists of a positively charged nucleus and a cloud of negatively charged electrons surrounding it. The position of a chemical element in the periodic table of Mendeleev (its serial number) is determined by the charge of the nucleus of its atoms. Isotopes are therefore called varieties of the same chemical element, the atoms of which have the same nuclear charge (and, therefore, practically the same electron shells), but differ in nuclear mass values. According to the figurative expression of F. Soddy, the atoms of isotopes are the same “outside”, but different “inside”.

The neutron was discovered in 1932 – a particle that has no charge, with a mass close to the mass of the nucleus of a hydrogen atom - a proton , and created proton-neutron model of the nucleus. As a result in science, the final modern definition of the concept of isotopes has been established: isotopes are substances whose atomic nuclei consist of the same number of protons and differ only in the number of neutrons in the nucleus . Each isotope is usually denoted by a set of symbols, where X is the symbol of the chemical element, Z is the charge of the atomic nucleus (the number of protons), A is the mass number of the isotope (the total number of nucleons - protons and neutrons in the nucleus, A = Z + N). Since the charge of the nucleus appears to be uniquely associated with the symbol of the chemical element, simply the notation A X is often used for abbreviation.

Of all the isotopes known to us, only hydrogen isotopes have their own names. Thus, the isotopes 2 H and 3 H are called deuterium and tritium and are designated D and T, respectively (the isotope 1 H is sometimes called protium).

Occurs in nature as stable isotopes , and unstable - radioactive, the nuclei of atoms of which are subject to spontaneous transformation into other nuclei with the emission of various particles (or processes of so-called radioactive decay). About 270 stable isotopes are now known, and stable isotopes are found only in elements with atomic number Z Ј 83. The number of unstable isotopes exceeds 2000, the vast majority of them were obtained artificially as a result of various nuclear reactions. The number of radioactive isotopes of many elements is very large and can exceed two dozen. The number of stable isotopes is significantly smaller. Some chemical elements consist of only one stable isotope (beryllium, fluorine, sodium, aluminum, phosphorus, manganese, gold and a number of other elements). The largest number of stable isotopes - 10 - was found in tin, for example in iron there are 4, and in mercury - 7.

Discovery of isotopes, historical background.

In 1808, the English naturalist John Dalton first introduced the definition of a chemical element as a substance consisting of atoms of the same type. In 1869, the chemist D.I. Mendeleev discovered the periodic law of chemical elements. One of the difficulties in substantiating the concept of an element as a substance occupying a certain place in a cell of the periodic table was the experimentally observed non-integer atomic weights of elements. In 1866, the English physicist and chemist Sir William Crookes put forward the hypothesis that each natural chemical element is a certain mixture of substances that are identical in their properties, but have different atomic masses, but at that time such an assumption did not yet have experimental confirmation and therefore did not last long noticed.

An important step towards the discovery of isotopes was the discovery of the phenomenon of radioactivity and the hypothesis of radioactive decay formulated by Ernst Rutherford and Frederick Soddy: radioactivity is nothing more than the decay of an atom into a charged particle and an atom of another element, different in its chemical properties from the original one. As a result, the idea of ​​radioactive series or radioactive families arose , at the beginning of which there is the first parent element, which is radioactive, and at the end - the last stable element. Analysis of the chains of transformations showed that during their course, the same radioactive elements, differing only in atomic masses, can appear in one cell of the periodic table. In fact, this meant the introduction of the concept of isotopes.

Independent confirmation of the existence of stable isotopes of chemical elements was then obtained in the experiments of J. J. Thomson and Aston in 1912–1920 with beams of positively charged particles (or so-called channel beams ) emanating from the discharge tube.

In 1919, Aston designed an instrument called a mass spectrograph. (or mass spectrometer) . The ion source still used a discharge tube, but Aston found a way in which successive deflection of a beam of particles in electric and magnetic fields led to the focusing of particles with the same charge-to-mass ratio (regardless of their speed) at the same point on the screen. Along with Aston, a mass spectrometer of a slightly different design was created in the same years by the American Dempster. As a result of the subsequent use and improvement of mass spectrometers through the efforts of many researchers, by 1935 an almost complete table of the isotopic compositions of all chemical elements known by that time had been compiled.

Methods for isotope separation.

To study the properties of isotopes and especially for their use for scientific and applied purposes, it is necessary to obtain them in more or less noticeable quantities. In conventional mass spectrometers, almost complete separation of isotopes is achieved, but their quantity is negligibly small. Therefore, the efforts of scientists and engineers were aimed at searching for other possible methods for separating isotopes. First of all, physicochemical methods of separation were mastered, based on differences in such properties of isotopes of the same element as evaporation rates, equilibrium constants, rates of chemical reactions, etc. The most effective among them were the methods of rectification and isotope exchange, which are widely used in the industrial production of isotopes of light elements: hydrogen, lithium, boron, carbon, oxygen and nitrogen.

Another group of methods consists of the so-called molecular kinetic methods: gas diffusion, thermal diffusion, mass diffusion (diffusion in a vapor flow), centrifugation. Gas diffusion methods, based on different rates of diffusion of isotopic components in highly dispersed porous media, were used during the Second World War to organize the industrial production of uranium isotope separation in the United States as part of the so-called Manhattan Project to create the atomic bomb. To obtain the required quantities of uranium enriched to 90% with the light isotope 235 U, the main “combustible” component of the atomic bomb, plants were built, occupying an area of ​​about four thousand hectares. More than 2 billion dollars were allocated for the creation of an atomic center with plants for the production of enriched uranium. After the war, plants for the production of enriched uranium for military purposes, also based on the diffusion method of separation, were developed and built in the USSR. In recent years, this method has given way to the more efficient and less expensive method of centrifugation. In this method, the effect of separating an isotope mixture is achieved due to the different effects of centrifugal forces on the components of the isotope mixture filling the centrifuge rotor, which is a thin-walled cylinder limited at the top and bottom, rotating at a very high speed in a vacuum chamber. Hundreds of thousands of centrifuges connected in cascades, the rotor of each of which makes more than a thousand revolutions per second, are currently used in modern separation plants both in Russia and in other developed countries of the world. Centrifuges are used not only to produce the enriched uranium needed to power the nuclear reactors of nuclear power plants, but also to produce isotopes of about thirty chemical elements in the middle part of the periodic table. Electromagnetic separation units with powerful ion sources are also used to separate various isotopes; in recent years, laser separation methods have also become widespread.

Application of isotopes.

Various isotopes of chemical elements are widely used in scientific research, in various fields of industry and agriculture, in nuclear energy, modern biology and medicine, in environmental studies and other fields. In scientific research (for example, in chemical analysis), as a rule, small quantities of rare isotopes of various elements are required, calculated in grams and even milligrams per year. At the same time, for a number of isotopes widely used in nuclear energy, medicine and other industries, the need for their production can amount to many kilograms and even tons. Thus, due to the use of heavy water D 2 O in nuclear reactors, its global production by the early 1990s of the last century was about 5000 tons per year. The hydrogen isotope deuterium, which is part of heavy water, the concentration of which in the natural mixture of hydrogen is only 0.015%, along with tritium, will in the future, according to scientists, become the main component of the fuel of power thermonuclear reactors operating on the basis of nuclear fusion reactions. In this case, the need for the production of hydrogen isotopes will be enormous.

In scientific research, stable and radioactive isotopes are widely used as isotopic indicators (tags) in the study of a wide variety of processes occurring in nature.

In agriculture, isotopes (“labeled” atoms) are used, for example, to study the processes of photosynthesis, the digestibility of fertilizers and to determine the efficiency of plants’ use of nitrogen, phosphorus, potassium, trace elements and other substances.

Isotope technologies are widely used in medicine. Thus, in the USA, according to statistics, more than 36 thousand medical procedures are performed per day and about 100 million laboratory tests using isotopes. The most common procedures involve computed tomography. The carbon isotope C13, enriched to 99% (natural content about 1%), is actively used in the so-called “diagnostic breathing control”. The essence of the test is very simple. The enriched isotope is introduced into the patient's food and, after participating in the metabolic process in various organs of the body, is released in the form of carbon dioxide CO 2 exhaled by the patient, which is collected and analyzed using a spectrometer. The differences in the rates of processes associated with the release of different amounts of carbon dioxide, labeled with the C 13 isotope, make it possible to judge the condition of the patient’s various organs. In the US, the number of patients who will undergo this test is estimated at 5 million per year. Now laser separation methods are used to produce highly enriched C13 isotope on an industrial scale.

Vladimir Zhdanov