Isotopic composition of chemical elements. isotopes

Probably, there is no such person on earth who would not have heard about isotopes. But not everyone knows what it is. The phrase "radioactive isotopes" sounds especially frightening. These obscure chemical elements terrify humanity, but in fact they are not as scary as it might seem at first glance.

Definition

To understand the concept of radioactive elements, it is first necessary to say that isotopes are samples of the same chemical element, but with different masses. What does it mean? Questions will disappear if we first remember the structure of the atom. It consists of electrons, protons and neutrons. The number of the first two elementary particles in the nucleus of an atom is always constant, while neutrons having their own mass can occur in the same substance in different quantities. This circumstance gives rise to a variety of chemical elements with different physical properties.

Now we can give a scientific definition of the concept under study. So, isotopes are a cumulative set of chemical elements similar in properties, but having different masses and physical properties. According to more modern terminology, they are called a galaxy of nucleotides of a chemical element.

A bit of history

At the beginning of the last century, scientists discovered that the same chemical compound under different conditions can have different masses of electron nuclei. From a purely theoretical point of view, such elements could be considered new and they could begin to fill empty cells in the periodic table of D. Mendeleev. But there are only nine free cells in it, and scientists discovered dozens of new elements. In addition, mathematical calculations showed that the discovered compounds cannot be considered previously unknown, because their chemical properties fully corresponded to the characteristics of existing ones.

After lengthy discussions, it was decided to call these elements isotopes and place them in the same cell as those whose nuclei contain the same number of electrons with them. Scientists have been able to determine that isotopes are just some variations of chemical elements. However, the causes of their occurrence and the duration of life were studied for almost a century. Even at the beginning of the 21st century, it is impossible to assert that humanity knows absolutely everything about isotopes.

Persistent and non-persistent variations

Each chemical element has several isotopes. Due to the fact that there are free neutrons in their nuclei, they do not always enter into stable bonds with the rest of the atom. After some time, free particles leave the core, which changes its mass and physical properties. This is how other isotopes are formed, which eventually leads to the formation of a substance with an equal number of protons, neutrons and electrons.

Those substances that decay very quickly are called radioactive isotopes. They release a large number of neutrons into space, forming powerful ionizing gamma radiation, known for its strong penetrating ability, which negatively affects living organisms.

More stable isotopes are not radioactive, since the number of free neutrons they release is not capable of producing radiation and significantly affecting other atoms.

Quite a long time ago, scientists established one important pattern: each chemical element has its own isotopes, persistent or radioactive. Interestingly, many of them were obtained in the laboratory, and their presence in their natural form is small and not always recorded by instruments.

Distribution in nature

Under natural conditions, most often there are substances whose isotope mass is directly determined by its ordinal number in the D. Mendeleev table. For example, hydrogen, denoted by the symbol H, has serial number 1, and its mass is equal to one. Its isotopes, 2H and 3H, are extremely rare in nature.

Even the human body has a certain amount of radioactive isotopes. They get inside through food in the form of isotopes of carbon, which, in turn, is absorbed by plants from the soil or air and passes into the composition of organic matter during photosynthesis. Therefore, both humans, animals, and plants emit a certain radiation background. Only it is so low that it does not interfere with normal functioning and growth.

The sources that contribute to the formation of isotopes are the inner layers of the earth's core and radiation from outer space.

As you know, the temperature on the planet largely depends on its hot core. But only recently it became clear that the source of this heat is a complex thermonuclear reaction, in which radioactive isotopes participate.

Isotope decay

Since isotopes are unstable formations, it can be assumed that, over time, they always decay into more permanent nuclei of chemical elements. This statement is true, because scientists have not been able to detect a huge number of radioactive isotopes in nature. And most of those that were mined in laboratories lasted from a couple of minutes to several days, and then turned back into ordinary chemical elements.

But there are also isotopes in nature that are very resistant to decay. They can exist for billions of years. Such elements were formed in those distant times, when the earth was still being formed, and there was not even a solid crust on its surface.

Radioactive isotopes decay and are re-formed very quickly. Therefore, in order to facilitate the assessment of the stability of the isotope, scientists decided to consider the category of its half-life.

Half life

It may not be immediately clear to all readers what is meant by this concept. Let's define it. The half-life of an isotope is the time during which the conditional half of the substance taken ceases to exist.

This does not mean that the rest of the connection will be destroyed in the same amount of time. With regard to this half, it is necessary to consider a different category - the period of time during which its second part, that is, a quarter of the original amount of the substance, will disappear. And this consideration continues ad infinitum. It can be assumed that the time of complete decay of the initial amount of matter is simply impossible to calculate, since this process is practically endless.

However, scientists, knowing the half-life, can determine how much of the substance existed in the beginning. These data are successfully used in related sciences.

In the modern scientific world, the concept of complete decay is practically not used. For each isotope, it is customary to indicate its half-life, which varies from a few seconds to many billions of years. The lower the half-life, the more radiation comes from the substance and the higher its radioactivity.

Enrichment of minerals

In some branches of science and technology, the use of a relatively large amount of radioactive substances is considered mandatory. But at the same time, in natural conditions, there are very few such compounds.

It is known that isotopes are uncommon variants of chemical elements. Their number is measured by a few percent of the most resistant variety. That is why scientists need to carry out artificial enrichment of fossil materials.

Over the years of research, it was possible to find out that the decay of an isotope is accompanied by a chain reaction. The released neutrons of one substance begin to influence another. As a result of this, heavy nuclei break up into lighter ones and new chemical elements are obtained.

This phenomenon is called a chain reaction, as a result of which more stable, but less common isotopes can be obtained, which are later used in the national economy.

Application of decay energy

Scientists also found that during the decay of a radioactive isotope, a huge amount of free energy is released. Its quantity is usually measured by the Curie unit, equal to the fission time of 1 g of radon-222 in 1 second. The higher this indicator, the more energy is released.

This was the reason for the development of ways to use free energy. This is how nuclear reactors appeared, in which a radioactive isotope is placed. Most of the energy it gives off is collected and converted into electricity. Based on these reactors, nuclear power plants are created, which provide the cheapest electricity. Reduced versions of such reactors are put on self-propelled mechanisms. Considering the danger of accidents, most often such machines are submarines. In the event of a reactor failure, the number of victims on the submarine will be easier to minimize.

Another very scary option for using half-life energy is atomic bombs. During World War II, they were tested on humanity in the Japanese cities of Hiroshima and Nagasaki. The consequences were very sad. Therefore, the world has an agreement on the non-use of these dangerous weapons. At the same time, large states with a focus on militarization continue research in this industry today. In addition, many of them, secretly from the world community, are making atomic bombs, which are thousands of times more dangerous than those used in Japan.

Isotopes in medicine

For peaceful purposes, the decay of radioactive isotopes has learned to use in medicine. By directing radiation to the affected area of ​​the body, it is possible to stop the course of the disease or help the patient to recover completely.

But more often radioactive isotopes are used for diagnostics. The thing is that their movement and the nature of the cluster is easiest to fix by the radiation that they produce. So, a certain non-dangerous amount of a radioactive substance is introduced into the human body, and doctors use instruments to observe how and where it gets.

Thus, the diagnosis of the work of the brain, the nature of cancerous tumors, and the features of the work of the endocrine and external secretion glands are carried out.

Application in archeology

It is known that in living organisms there is always radioactive carbon-14, the half-life of which isotope is 5570 years. In addition, scientists know how much of this element is contained in the body until the moment of his death. This means that all cut trees emit the same amount of radiation. Over time, the intensity of radiation decreases.

This helps archaeologists determine how long ago the tree from which the galley or any other ship was built died, and therefore the very time of construction. This research method is called radioactive carbon analysis. Thanks to him, it is easier for scientists to establish the chronology of historical events.

A certain element that has the same but different . Possess nuclei with the same number and different. number , have the same structure of electron shells and occupy the same place in the periodic. chemical system. elements. The term "isotopes" was proposed in 1910 by F. Soddy to denote chemically indistinguishable varieties that differ in their physical. (primarily radioactive) St. you. Stable isotopes were first discovered in 1913 by J. Thomson with the help of the so-called. method of parabolas - the prototype of modern. . He found that Ne has at least 2 varieties with wt. hours 20 and 22. The names and symbols of isotopes are usually the names and symbols of the corresponding chems. elements; point to the top left of the symbol. For example, to designate nature. isotopes use the record 35 Cl and 37 C1; 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, wt. Parts 1, 2 and 3 have specials. 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 are found in all even and most odd elements with[ 83. The number of stable isotopes for elements with even numbers can be. equals 10 (e.g. y); elements with odd numbers have at most two stable isotopes. Known ca. 280 stable and more than 2000 radioactive isotopes in 116 natural and artificially obtained elements. For each element, the content of individual isotopes in nature. 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 a change in the isotopic composition during decomp. processes occurring in nature (, etc.). For the heavy element Pb, fluctuations in the isotopic composition of different samples are explained by decomp. content in, and other sources and - the founders of nature. . Differences St. in the isotopes of a given element called. . An important practical the task is to obtain from nature. mixtures of individual isotopes -

isotopes- varieties of atoms (and nuclei) of a chemical element that have the same atomic (ordinal) number, but 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 isotope identity 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 particular 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 a nucleus is the mass number of an 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.

Huclides And isotopes

The nuclide belongs to the nucleus, not to the atom. Identical nuclei belong to the same nuclide, for example, each carbon-13 nuclide nucleus is made up of 6 protons and 7 neutrons. The concept of nuclides (referring 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. The neutron number has a great influence on the properties of nuclei, but its influence on the chemical properties is negligible for most elements. Even in the case of the lightest elements, where the ratio of neutrons to atomic number varies the 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 greatly affect to biology). Since isotope is an older term, it is better known than nuclide and is still occasionally used in contexts where nuclide might be more appropriate, such as nuclear technology and nuclear medicine.

Notation

An isotope or nuclide is identified by the name of a particular element (this indicates the atom number), followed by a hyphen and a 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 for the element) is to 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). Since the atomic number is given by the symbol of the element, usually only the mass number in the superscript is given, and the atom index is not given. The letter m is sometimes appended after the mass number to indicate a nuclear isomer, a metastable or energetically excited nuclear state (as opposed to the lowest energy ground state), such as 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 a radioactive form of carbon, while 12 C and 13 C are stable isotopes. There are about 339 naturally occurring nuclides on Earth, of which 286 are primordial nuclides, meaning they have been around 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 have not been 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 abundant 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 (not considering 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 resistant to all known decay forms. 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 forms of decay, such as alpha decay or double beta decay, but decay products have not yet been observed, and thus 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, artificially created, currently 3339 nuclides are known. These include 905 nuclides that are either stable or have half-lives greater than 60 minutes.

Isotope Properties

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 a similar electronic structure. Since the chemical behavior of an atom is largely determined by its electronic structure, different isotopes exhibit almost identical chemical behavior.

An 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 by changing the center of gravity (reduced mass) of atomic systems. However, for heavier elements, the relative mass difference between isotopes is much smaller, so that the effects of mass difference 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.)

Similarly, 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 deuterium and tritium being the primary exceptions). 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 the appropriate energies, isotopologues have different optical properties in the infrared.

Nuclear properties and stability

Atomic nuclei are made up of protons and neutrons bound together by a residual strong force. Because the protons are positively charged, they repel each other. Neutrons, which are electrically neutral, stabilize the nucleus in two ways. Their contact pushes the protons back a little, 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 the nucleus. As the number of protons increases, so does the ratio of neutrons to protons needed to provide a stable nucleus (see graph on the right). For example, although the ratio neutron: proton 3 2 He is 1:2, the ratio neutron: proton 238 92 U
Over 3:2. A number of lighter elements have stable nuclides with a ratio of 1:1 (Z = N). The nuclide 40 20 Ca (calcium-40) is the observable heaviest stable nuclide with the same number of neutrons and protons; (Theoretically, the heaviest stable 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 largest number of stable isotopes observable 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 the so-called mononuclide elements, which have a single primordial stable isotope that dominates and fixes the atomic weight of the natural element with high precision, 3 radioactive mononuclide elements are also present). In total, there are 254 nuclides that have not been observed to decay. For 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 ratio of neutrons is not the only factor affecting nuclear stability. It depends also on the parity or oddness of its atomic number Z, the number of neutrons N, hence the sum of their mass number A. Odd both Z and N tends 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 Z, N, and A are all even. Odd stable nuclides are divided (approximately evenly) into odd ones.

atomic number

The 148 even proton, even neutron (EE) nuclides make up ~58% of all stable nuclides. There are also 22 primordial long-lived even nuclides. As a result, each of the 41 even elements from 2 to 82 has at least one stable isotope, and most of these elements have multiple primary isotopes. Half of these even elements have six or more stable isotopes. The extreme stability of helium-4, due to the binary bonding 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 about 3 times as numerous. 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 odd-even nuclide and 13 that have two odd-even nuclides.

Because of their odd neutron numbers, even-odd nuclides tend to have large neutron capture cross sections due to the energy that comes from neutron coupling effects. These stable nuclides may be unusually abundant in nature, mainly because in order to form and enter into primordial abundance, they must escape neutron capture in order to form yet other stable even-odd isotopes over the course of how s is the process and r is the 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 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 is even) have one stable odd-even 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. Thus, 30 + 2 (9) = 48 stable even-even isotopes are obtained.

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 from 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 tend not to, although they do fission into fast neutrons. All observationally stable odd-odd nuclides have a non-zero integer spin. This is because a single unpaired neutron and an unpaired proton have more nuclear force attraction to 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 post-example. The original isotopes were the product of stellar nucleosynthesis, or another type of nucleosynthesis such as cosmic ray splitting, and have persisted up to the present because their decay rate is so slow (eg uranium-238 and potassium-40). Post-natural isotopes have been 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 a few stable isotopes, with the highest number of stable isotopes for an element being ten, for tin (50Sn). About 94 elements exist on Earth (up to and including plutonium), although some are only found in very small amounts, 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 naturally occurring isotopes are stable in the sense that they have not been observed to date. An additional 35 primordial nuclides (a total of 289 primordial nuclides) are radioactive with known half-lives, but have half-lives in excess of 80 million years, allowing them to exist since the beginning of the solar system.

All known stable isotopes naturally occur on Earth; Other natural isotopes are radioactive, but because of their relatively long half-lives, or because of other continuous natural production methods. These include the cosmogenic nuclides mentioned above, nucleogenic nuclides, and any radiogenic isotopes resulting from the continued 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 created in stars or supernovae. An example is aluminum-26, which does not naturally occur on Earth, but is found in abundance on an astronomical scale.

The tabulated atomic masses of the elements are averages that explain the presence of multiple isotopes with different masses. Prior to the discovery of isotopes, empirically determined non-integrated values ​​for atomic mass 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 the 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, as well as in the interactions between energetic particles , such as cosmic rays, and previously obtained isotopes. The corresponding isotopic abundance of isotopes on Earth is due to the quantities produced by these processes, their propagation through the galaxy, and the rate of decay of the isotopes, which are unstable. After the initial merger of the solar system, isotopes were redistributed according to mass, and the isotopic composition of the elements varies slightly from planet to planet. This sometimes makes it possible to trace the origin of meteorites.

Atomic mass of isotopes

The atomic mass (mr) of an isotope is determined mainly 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 is a dimensionless quantity. Atomic mass, on the other hand, is measured using the unit of atomic mass, based on the mass of the 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 element's atomic mass. When an element contains N isotopes, the expression below applies to the average atomic mass:

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

Application of isotopes

There are several applications that exploit the properties of the various isotopes of a given element. Isotope separation is an important technological issue, 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 have been separated by volume by gaseous diffusion, gas centrifugation, laser ionization separation and (in the Manhattan Project) by type of mass spectrometry production.

Use of chemical and biological properties

• Isotope analysis is the determination of the isotopic signature, the relative abundance of the isotopes of a given element in a particular sample. For nutrients in particular, significant variations in C, N and O isotopes can occur. The analysis of such variations has a wide range of applications, such as the detection of adulteration in foods or the geographic origin of foods using isoscapes. The identification of some meteorites originating 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 isotopic labeling, the use of unusual isotopes as tracers 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 marking).
• Isotopes are commonly used to determine the concentration of various elements or substances using the isotopic dilution method, in which known amounts of isotopically substituted compounds are mixed with samples and the isotopic characteristics of the resulting mixtures are determined using mass spectrometry.

Using nuclear properties

• A method similar to radioisotope tagging is radiometric dating: using the known half-life of an unstable element, one can calculate the time elapsed since the existence of a known isotope concentration. The most widely known example is radiocarbon dating, which is used to determine the age of carbonaceous materials.
• Some forms of spectroscopy are based 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 the nuclear transitions of specific isotopes such as 57 Fe.

Studying the phenomenon of radioactivity, scientists in the first decade of the XX century. discovered a large number of radioactive substances - about 40. There were significantly more of them than free places in the periodic table of elements in the interval between bismuth and uranium. The nature of these substances has been controversial. Some researchers considered them to be independent chemical elements, but in this case the question of their placement in the periodic table turned out to be insoluble. Others generally denied them the right to be called elements in the classical sense. In 1902, the English physicist D. Martin called such substances radioelements. As they were studied, it turned out that some radioelements have exactly the same chemical properties, but differ in atomic masses. This circumstance contradicted the main provisions of the periodic law. The English scientist F. Soddy resolved the contradiction. In 1913, he called chemically similar radioelements isotopes (from the Greek words meaning "same" and "place"), i.e., occupying the same place in the periodic system. Radioelements turned out to be isotopes of natural radioactive elements. All of them are combined into three radioactive families, the ancestors of which are the isotopes of thorium and uranium.

Isotopes of oxygen. Isobars of potassium and argon (isobars are atoms of different elements with the same mass number).

Number of stable isotopes for even and odd elements.

It soon became clear that other stable chemical elements also have isotopes. The main merit in their discovery belongs to the English physicist F. Aston. He discovered stable isotopes in many elements.

From a modern point of view, isotopes are varieties of atoms of a chemical element: they have different atomic masses, but the same nuclear charge.

Their nuclei thus contain the same number of protons, but a different number of neutrons. For example, natural oxygen isotopes with Z = 8 contain 8, 9, and 10 neutrons in their nuclei, respectively. The sum of the numbers of protons and neutrons in the nucleus of an isotope is called the mass number A. Therefore, the mass numbers of the indicated oxygen isotopes are 16, 17 and 18. The following designation of isotopes is now accepted: the Z value is given at the bottom left of the element symbol, the A value is given at the top left. For example: 16 8 O, 17 8 O, 18 8 O.

After the discovery of the phenomenon of artificial radioactivity, about 1800 artificial radioactive isotopes were obtained using nuclear reactions for elements with Z from 1 to 110. The vast majority of artificial radioisotopes have very short half-lives, measured in seconds and fractions of seconds; only a few have a relatively long lifespan (for example, 10 Be - 2.7 10 6 years, 26 Al - 8 10 5 years, etc.).

Stable elements are present in nature with about 280 isotopes. However, some of them turned out to be slightly radioactive, with huge half-lives (for example, 40 K, 87 Rb, 138 La, l47 Sm, 176 Lu, 187 Re). The lifetime of these isotopes is so long that they can be considered stable.

There are still many problems in the world of stable isotopes. So, it is not clear why their number in different elements varies so much. About 25% of stable elements (Be, F, Na, Al, P, Sc, Mn, Co, As, Y, Nb, Rh, I, Cs, Pt, Tb, Ho, Tu, Ta, Au) are present in nature only one kind of atom. These are the so-called single elements. Interestingly, all of them (except Be) have odd Z values. In general, for odd elements, the number of stable isotopes does not exceed two. On the contrary, some elements with even Z consist of a large number of isotopes (for example, Xe has 9, Sn - 10 stable isotopes).

The set of stable isotopes of a given element is called a galaxy. Their content in the galaxy often fluctuates greatly. It is interesting to note that the abundance of isotopes with mass numbers that are multiples of four (12 C, 16 O, 20 Ca, etc.) is the highest, although there are exceptions to this rule.

The discovery of stable isotopes made it possible to solve the long-term mystery of atomic masses - their deviation from integers, due to the different percentages of stable isotopes of elements in the galaxy.

In nuclear physics, the concept of "isobars" is known. Isobars are called isotopes of different elements (i.e., with different Z values) that have the same mass numbers. The study of isobars contributed to the establishment of many important regularities in the behavior and properties of atomic nuclei. One of these regularities is expressed by the rule formulated by the Soviet chemist S. A. Shchukarev and the Yemenite physicist I. Mattauch. It says: if the two isobars differ in Z values ​​by 1, then one of them will necessarily be radioactive. A classic example of a pair of isobars is 40 18 Ar - 40 19 K. In it, the potassium isotope is radioactive. The Shchukarev-Mattauch rule made it possible to explain why the elements technetium (Z = 43) and promethium (Z = 61) have no stable isotopes. Since they have odd Z values, more than two stable isotopes could not be expected for them. But it turned out that the isobars of technetium and promethium, respectively, the isotopes of molybdenum (Z = 42) and ruthenium (Z = 44), neodymium (Z = 60) and samarium (Z = 62), are represented in nature by stable varieties of atoms in a wide range of mass numbers . Thus, physical laws impose a ban on the existence of stable isotopes of technetium and promethium. That is why these elements do not actually exist in nature and they had to be synthesized artificially.

Scientists have long been trying to develop a periodic system of isotopes. Of course, it is based on other principles than the basis of the periodic system of elements. But these attempts have not yet led to satisfactory results. True, physicists have proved that the sequence of filling proton and neutron shells in atomic nuclei is in principle similar to the construction of electron shells and subshells in atoms (see Atom).

The electron shells of the isotopes of a given element are built in exactly the same way. Therefore, their chemical and physical properties are almost identical. Only the isotopes of hydrogen (protium and deuterium) and their compounds show noticeable differences in properties. For example, heavy water (D 2 O) freezes at +3.8, boils at 101.4 ° C, has a density of 1.1059 g / cm 3, does not support the life of animal and plant organisms. During the electrolysis of water into hydrogen and oxygen, H 2 0 molecules are predominantly decomposed, while heavy water molecules remain in the electrolyzer.

The separation of isotopes of other elements is an extremely difficult task. Nevertheless, in many cases, isotopes of individual elements with a significantly changed content compared to the natural abundance are needed. For example, when solving the problem of atomic energy, it became necessary to separate the isotopes 235 U and 238 U. For this purpose, the mass spectrometry method was first applied, with the help of which the first kilograms of uranium-235 were obtained in 1944 in the USA. However, this method turned out to be too expensive and was replaced by the gaseous diffusion method, which used UF 6 . Now there are several methods for separating isotopes, but all of them are quite complex and expensive. Nevertheless, the problem of “separation of the inseparable” is being successfully solved.

A new scientific discipline appeared - the chemistry of isotopes. It studies the behavior of various isotopes of chemical elements in chemical reactions and isotope exchange processes. As a result of these processes, the isotopes of a given element are redistributed between the reacting substances. Here is the simplest example: H 2 0 + HD = HD0 + H 2 (a water molecule exchanges a protium atom for a deuterium atom). The geochemistry of isotopes is also developing. It investigates fluctuations in the isotopic composition of various elements in the earth's crust.

The most widely used are the so-called labeled atoms - artificial radioactive isotopes of stable elements or stable isotopes. With the help of isotope indicators - labeled atoms - they study the ways of movement of elements in inanimate and living nature, the nature of the distribution of substances and elements in various objects. Isotopes are used in nuclear technology: as materials for the construction of nuclear reactors; as a nuclear fuel (isotopes of thorium, uranium, plutonium); in thermonuclear fusion (deuterium, 6 Li, 3 He). Radioactive isotopes are also widely used as radiation sources.

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 the electrons moving at a staggering speed in orbits in the electron cloud, the neutrons and protons that make up the nucleus.

Definition

isotopes 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 equal to zero). When an electron is lost or captured, the atom loses its 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 be different. This does not affect the neutrality of the atom, but it does affect its mass and properties. For example, each isotope of a hydrogen atom has one electron and one proton each. And 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

They have a different number of neutrons, different masses and different properties. Isotopes have an identical structure of electron shells. This means that they are quite similar in chemical properties. Therefore, they are assigned one place in the periodic system.

Stable and radioactive (unstable) isotopes have been found in nature. The nuclei of atoms of radioactive isotopes are able to spontaneously transform into other nuclei. In 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 fluctuates slightly.

The existence of isotopes made it possible to understand why, in some cases, elements with a lower atomic mass have a higher serial number than elements with a larger atomic mass. For example, in an argon-potassium pair, argon includes heavy isotopes, and potassium includes light isotopes. Therefore, the mass of argon is greater than that of potassium.

Findings site

1. They have different numbers of neutrons.
2. Isotopes have different masses of atoms.
3. The value of the mass of atoms of ions affects their total energy and properties.