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). The 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 the stable isotopes of the 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.

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 scientist 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 system. 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

Even ancient philosophers suggested that matter is built from atoms. However, scientists began to realize that the “building blocks” of the universe themselves consist of tiny particles only at the turn of the 19th and 20th centuries. Experiments proving this produced a real revolution in science at one time. It is the quantitative ratio of its constituent parts that distinguishes one chemical element from another. Each of them is assigned its place in according to the serial number. But there are varieties of atoms that occupy the same cells in the table, despite differences in mass and properties. Why this is so and what isotopes are in chemistry will be discussed further.

Atom and its particles

Studying the structure of matter through bombardment with alpha particles, E. Rutherford proved in 1910 that the main space of the atom is filled with void. And only in the center is the core. Negative electrons move around it in orbitals, making up the shell of this system. This is how a planetary model of the “building blocks” of matter was created.

What are isotopes? Remember from your chemistry course that the nucleus also has a complex structure. It consists of positive protons and neutrons that have no charge. The number of the former determines the qualitative characteristics of the chemical element. It is the number of protons that distinguishes substances from each other, giving their nuclei a certain charge. And on this basis they are assigned a serial number in the periodic table. But the number of neutrons in the same chemical element differentiates them into isotopes. The definition in chemistry of this concept can therefore be given as follows. These are varieties of atoms that differ in the composition of the nucleus, have the same charge and atomic numbers, but have different mass numbers due to differences in the number of neutrons.

Designations

While studying chemistry in the 9th grade and isotopes, students will learn about the accepted conventions. The letter Z indicates the charge of the nucleus. This figure coincides with the number of protons and is therefore their indicator. The sum of these elements with neutrons marked with N is A - mass number. A family of isotopes of one substance is usually designated by the symbol of that chemical element, which in the periodic table is assigned a serial number that coincides with the number of protons in it. The left superscript added to the indicated icon corresponds to the mass number. For example, 238 U. The charge of an element (in this case, uranium, marked with the serial number 92) is indicated by a similar index below.

Knowing these data, you can easily calculate the number of neutrons in a given isotope. It is equal to the mass number minus the serial number: 238 - 92 = 146. The number of neutrons could be less, but this would not make this chemical element cease to remain uranium. It should be noted that most often in other, simpler substances the number of protons and neutrons is approximately the same. Such information helps to understand what an isotope is in chemistry.

Nucleons

It is the number of protons that gives a certain element its individuality, and the number of neutrons does not affect it in any way. But the atomic mass is made up of these two specified elements, which have the common name “nucleons,” representing their sum. However, this indicator does not depend on those forming the negatively charged shell of the atom. Why? All you have to do is compare.

The fraction of proton mass in an atom is large and amounts to approximately 1 a. e.m. or 1.672 621 898(21) 10 -27 kg. The neutron is close to the performance of this particle (1.674 927 471(21)·10 -27 kg). But the mass of an electron is thousands of times smaller, is considered insignificant and is not taken into account. That is why, knowing the superscript of an element in chemistry, the composition of the isotope nucleus is not difficult to find out.

Isotopes of hydrogen

Isotopes of some elements are so well known and widespread in nature that they have received their own names. The most striking and simplest example of this is hydrogen. It is naturally found in its most common form, protium. This element has a mass number of 1, and its nucleus consists of one proton.

So what are hydrogen isotopes in chemistry? As is known, the atoms of this substance have the first number in the periodic table and, accordingly, are endowed with a charge number of 1 in nature. But the number of neutrons in the nucleus of an atom is different. Deuterium, being heavy hydrogen, in addition to the proton, has another particle in its nucleus, that is, a neutron. As a result, this substance exhibits its own physical properties, unlike protium, having its own weight, melting and boiling points.

Tritium

Tritium is the most complex of all. This is superheavy hydrogen. According to the definition of isotopes in chemistry, it has a charge number of 1, but a mass number of 3. It is often called a triton because in addition to one proton, it has two neutrons in its nucleus, that is, it consists of three elements. The name of this element, discovered in 1934 by Rutherford, Oliphant and Harteck, was proposed even before its discovery.

This is an unstable substance exhibiting radioactive properties. Its core has the ability to split into a beta particle and an electron antineutrino. The decay energy of this substance is not very high and amounts to 18.59 keV. Therefore, such radiation is not too dangerous for humans. Ordinary clothing and surgical gloves can protect against it. And this radioactive element obtained from food is quickly eliminated from the body.

Isotopes of uranium

Much more dangerous are the various types of uranium, of which science currently knows 26. Therefore, when talking about what isotopes are in chemistry, it is impossible not to mention this element. Despite the variety of types of uranium, only three isotopes occur in nature. These include 234 U, 235 U, 238 U. The first of them, having suitable properties, is actively used as fuel in nuclear reactors. And the latter is for the production of plutonium-239, which itself, in turn, is irreplaceable as a valuable fuel.

Each of the radioactive elements is characterized by its own This is the length of time during which the substance is split in a ratio of ½. That is, as a result of this process, the amount of the remaining part of the substance is halved. This period of time is huge for uranium. For example, for isotope-234 it is estimated at 270 thousand years, but for the other two specified varieties it is much more significant. Uranium-238 has a record half-life, lasting billions of years.

Nuclides

Not every type of atom, characterized by its own and strictly defined number of protons and electrons, is so stable as to exist for at least a long period sufficient for its study. Those that are relatively stable are called nuclides. Stable formations of this kind do not undergo radioactive decay. Unstable ones are called radionuclides and, in turn, are also divided into short-lived and long-lived. As you know from 11th grade chemistry lessons about the structure of isotope atoms, osmium and platinum have the largest number of radionuclides. Cobalt and gold have one stable nuclide each, and tin has the largest number of stable nuclides.

Calculating the atomic number of an isotope

Now we will try to summarize the information described earlier. Having understood what isotopes are in chemistry, it’s time to figure out how to use the knowledge gained. Let's look at this with a specific example. Suppose it is known that a certain chemical element has a mass number of 181. Moreover, the shell of an atom of this substance contains 73 electrons. How can you use the periodic table to find out the name of a given element, as well as the number of protons and neutrons in its nucleus?

Let's start solving the problem. You can determine the name of a substance by knowing its serial number, which corresponds to the number of protons. Since the number of positive and negative charges in an atom are equal, it is 73. This means it is tantalum. Moreover, the total number of nucleons in total is 181, which means that the protons of this element are 181 - 73 = 108. Quite simple.

Isotopes of gallium

The element gallium has atomic number 71. In nature, this substance has two isotopes - 69 Ga and 71 Ga. How to determine the percentage of gallium species?

Solving problems on isotopes in chemistry almost always involves information that can be obtained from the periodic table. This time you should do the same. Let us determine the average atomic mass from the indicated source. It is equal to 69.72. Having designated by x and y the quantitative ratio of the first and second isotope, we take their sum equal to 1. This means that this will be written in the form of an equation: x + y = 1. It follows that 69x + 71y = 69.72. Expressing y in terms of x and substituting the first equation into the second, we find that x = 0.64 and y = 0.36. This means that 69 Ga is found in nature 64%, and the percentage of 71 Ga is 34%.

Isotopic transformations

Radioactive fission of isotopes with their transformation into other elements is divided into three main types. The first of these is alpha decay. It occurs with the emission of a particle representing the nucleus of a helium atom. That is, this is a formation consisting of a combination of pairs of neutrons and protons. Since the amount of the latter determines the charge number and number of the atom of a substance in the periodic table, as a result of this process there is a qualitative transformation of one element into another, and in the table it shifts to the left by two cells. In this case, the mass number of the element decreases by 4 units. We know this from the structure of isotope atoms.

When the nucleus of an atom loses a beta particle, essentially an electron, its composition changes. One of the neutrons transforms into a proton. This means that the qualitative characteristics of the substance change again, and the element shifts in the table one cell to the right, without practically losing weight. Typically, such a transformation is associated with electromagnetic gamma radiation.

Radium isotope transformation

The above information and knowledge from grade 11 chemistry about isotopes again help solve practical problems. For example, the following: 226 Ra during decay turns into a chemical element of group IV, with a mass number of 206. How many alpha and beta particles should it lose?

Taking into account the changes in the mass and the group of the daughter element, using the periodic table, it is easy to determine that the isotope formed during splitting will be lead with a charge of 82 and a mass number of 206. And taking into account the charge number of this element and the original radium, it should be assumed that its nucleus has lost five alpha -particles and four beta particles.

Use of radioactive isotopes

Everyone is well aware of the harm radioactive radiation can cause to living organisms. However, the properties of radioactive isotopes can be useful for humans. They are successfully used in many industries. With their help, it is possible to detect leaks in engineering and construction structures, underground pipelines and oil pipelines, storage tanks, and heat exchangers in power plants.

These properties are also actively used in scientific experiments. For example, the tsetse fly is a carrier of many serious diseases for humans, livestock and domestic animals. In order to prevent this, males of these insects are sterilized using weak radioactive radiation. Isotopes are also indispensable in studying the mechanisms of certain chemical reactions, because atoms of these elements can be used to label water and other substances.

Tagged isotopes are also often used in biological research. For example, this is how it was established how phosphorus affects the soil, growth and development of cultivated plants. The properties of isotopes are also successfully used in medicine, which has made it possible to treat cancerous tumors and other serious diseases and determine the age of biological organisms.

It has been established that every chemical element found in nature is a mixture of isotopes (hence they have fractional atomic masses). To understand how isotopes differ from one another, it is necessary to consider in detail the structure of the atom. An atom forms a nucleus and an electron cloud. The mass of an atom is influenced by electrons moving at stunning speeds through orbitals in the electron cloud, neutrons and protons that make up the nucleus.

What are isotopes

Isotopes is a type of atom of a chemical element. There are always equal numbers of electrons and protons in any atom. Since they have opposite charges (electrons are negative, and protons are positive), the atom is always neutral (this elementary particle does not carry a charge, it is zero). When an electron is lost or captured, an atom loses neutrality, becoming either a negative or a positive ion.
Neutrons have no charge, but their number in the atomic nucleus of the same element can vary. This does not in any way affect the neutrality of the atom, but it does affect its mass and properties. For example, any isotope of a hydrogen atom contains one electron and one proton. But the number of neutrons is different. Protium has only 1 neutron, deuterium has 2 neutrons, and tritium has 3 neutrons. These three isotopes differ markedly from each other in properties.

Comparison of isotopes

How are isotopes different? They have different numbers of neutrons, different masses and different properties. Isotopes have identical structures of electron shells. This means that they are quite similar in chemical properties. Therefore, they are given one place in the periodic table.
Stable and radioactive (unstable) isotopes have been found in nature. The nuclei of atoms of radioactive isotopes are capable of spontaneously transforming into other nuclei. During the process of radioactive decay, they emit various particles.
Most elements have over two dozen radioactive isotopes. In addition, radioactive isotopes are artificially synthesized for absolutely all elements. In a natural mixture of isotopes, their content varies slightly.
The existence of isotopes made it possible to understand why, in some cases, elements with lower atomic mass have a higher atomic number than elements with higher atomic mass. For example, in the argon-potassium pair, argon includes heavy isotopes, and potassium contains light isotopes. Therefore, the mass of argon is greater than that of potassium.

TheDifference.ru determined that the difference between isotopes is as follows:

They have different numbers of neutrons.
Isotopes have different atomic masses.
The value of the mass of ion atoms affects their total energy and properties.

Isotopes

ISOTOPES-s; pl.(unit isotope, -a; m.). [from Greek isos - equal and topos - place] Specialist. Varieties of the same chemical element, differing in the mass of atoms. Radioactive isotopes. Isotopes of uranium.

◁ Isotopic, oh, oh. I. indicator.

isotopes

History of research
The first experimental data on the existence of isotopes were obtained in 1906-10. when studying the properties of radioactive transformations of atoms of heavy elements. In 1906-07. It was discovered that the radioactive decay product of uranium, ionium, and the radioactive decay product of thorium, radiothorium, have the same chemical properties as thorium, but differ from the latter in atomic mass and radioactive decay characteristics. Moreover: all three elements have the same optical and x-ray spectra. At the suggestion of the English scientist F. Soddy (cm. SODDIE Frederick), such substances began to be called isotopes.
After isotopes were discovered in heavy radioactive elements, the search for isotopes in stable elements began. Independent confirmation of the existence of stable isotopes of chemical elements was obtained in the experiments of J. J. Thomson (cm. THOMSON Joseph John) and F. Aston (cm. ASTON Francis William). Thomson discovered stable isotopes of neon in 1913. Aston, who conducted research using an instrument he designed called a mass spectrograph (or mass spectrometer), using the mass spectrometry method (cm. MASS SPECTROMETRY), proved that many other stable chemical elements have isotopes. In 1919, he obtained evidence of the existence of two isotopes 20 Ne and 22 Ne, the relative abundance (abundance) of which in nature is approximately 91% and 9%. Subsequently, the isotope 21 Ne was discovered with an abundance of 0.26%, isotopes of chlorine, mercury and a number of other elements.
A mass spectrometer of a slightly different design was created in the same years by A. J. Dempster (cm. DEMPSTER Arthur Jeffrey). As a result of the subsequent use and improvement of mass spectrometers, an almost complete table of isotopic compositions was compiled through the efforts of many researchers. In 1932, a neutron was discovered - a particle without a charge, with a mass close to the mass of the nucleus of a hydrogen atom - a proton, and a proton-neutron model of the nucleus was created. As a result, science has established the final definition of the concept of isotopes: isotopes are substances whose atomic nuclei consist of the same number of protons and differ only in the number of neutrons in the nucleus. Around 1940, isotope analysis was carried out for all chemical elements known at that time.
During the study of radioactivity, about 40 natural radioactive substances were discovered. They were grouped into radioactive families, the ancestors of which are isotopes of thorium and uranium. Natural ones include all stable varieties of atoms (there are about 280 of them) and all naturally radioactive ones that are part of radioactive families (there are 46 of them). All other isotopes are obtained as a result of nuclear reactions.
For the first time in 1934 I. Curie (cm. JOLIO-CURIE Irene) and F. Joliot-Curie (cm. JOLIO-CURIE Frederic) artificially obtained radioactive isotopes of nitrogen (13 N), silicon (28 Si) and phosphorus (30 P), which are absent in nature. With these experiments they demonstrated the possibility of synthesizing new radioactive nuclides. Among the currently known artificial radioisotopes, more than 150 belong to transuranium elements (cm. TRANSURANE ELEMENTS), not found on Earth. Theoretically, it is assumed that the number of varieties of isotopes capable of existence can reach about 6000.

encyclopedic Dictionary. 2009 .

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