The Periodic Table of the Elements had a great influence on the subsequent development of chemistry. Not only was it the first natural classification of the chemical elements, which showed that they form a coherent system and are in close connection with each other, but it was also a powerful tool for further research.
At the time when Mendeleev compiled his table on the basis of the periodic law he had discovered, many elements were still unknown. So, element 4 of the scandium period was unknown. In terms of atomic mass, Ca was followed by Ti, but Ti could not be placed immediately after Ca, because it would fall into group 3, but according to the properties of Ti it should be assigned to group 4. Therefore, Mendeleev skipped one cell. On the same basis, two free cells were left between Zn and As in period 4. There were also empty seats in other rows. Mendeleev was not only convinced that there must be still unknown elements that fill these places, but also predicted in advance the properties of such elements, based on their position among other elements of the periodic system. These elements were given the names ekabor (since its properties were supposed to resemble boron), ekaaluminum, ekasilicium ..

Over the next 15 years, Mendeleev's predictions were brilliantly confirmed; all three expected elements were opened. First, the French chemist Lecoq de Boisbaudran discovered gallium, which has all the properties of ekaaluminum. Following that, in Sweden, L.F. Nilson discovered scandium, and finally, a few more years later, in Germany, K.A.
The discovery of Ga, Sc, Ge was the greatest triumph of the periodic law. The periodic system was also of great importance in establishing the valency and atomic masses of certain elements. Similarly, the periodic system gave impetus to the correction of the atomic masses of some elements. For example, Cs used to be assigned an atomic mass of 123.4. Mendeleev, arranging the elements in a table, found that, according to its properties, Cs should be in the main subgroup of the first group under Rb and therefore will have an atomic mass of about 130. Modern definitions show that the atomic mass of Cs is 132.9054 ..
And at present the periodic law remains the guiding star of chemistry. It was on its basis that transuranium elements were artificially created. One of them, element No. 101, first obtained in 1955, was named mendelevium in honor of the great Russian scientist.
The subsequent development of science made it possible, relying on the periodic law, to understand the structure of matter much more deeply, than was possible during Mendeleev's lifetime.
.Brilliant confirmation was found by the prophetic words of Mendeleev: "The periodic law is not threatened with destruction, but only a superstructure and development are promised."

The prerequisite for the discovery of the Periodic Law was the decision of the international congress of chemists in the city of Karlsruhe in 1860, when the atomic and molecular doctrine was finally established, the first unified definitions of the concepts of a molecule and atom, as well as atomic weight, which we now call relative atomic mass, were made.

D. I. Mendeleev in his discovery relied on clearly formulated starting points:

A common invariable property of the atoms of all chemical elements is their atomic mass;

The properties of elements depend on their atomic masses;

The form of this dependence is periodic.

The prerequisites discussed above can be called objective, that is, independent of the personality of the scientist, since they were due to the historical development of chemistry as a science.

III Periodic law and the Periodic system of chemical elements.

Mendeleev's discovery of the Periodic Law.

The first version of the Periodic Table of the Elements was published by D. I. Mendeleev in 1869 - long before the structure of the atom was studied. At this time, Mendeleev taught chemistry at St. Petersburg University. Preparing for lectures, collecting material for his textbook "Fundamentals of Chemistry", D. I. Mendeleev thought about how to systematize the material in such a way that information about the chemical properties of elements did not look like a set of disparate facts.

The reference point in this work for D. I. Mendeleev was the atomic masses (atomic weights) of the elements. After the World Congress of Chemists in 1860, in which D. I. Mendeleev also participated, the problem of the correct determination of atomic weights was constantly at the center of attention of many leading chemists of the world, including D. I. Mendeleev.Arranging the elements in ascending order of their atomic weights, D. I. Mendeleev discovered a fundamental law of nature, which is now known as the Periodic Law:

The properties of elements change periodically according to their atomic weight.

The above formulation does not in the least contradict the modern one, in which the concept of "atomic weight" is replaced by the concept of "nuclear charge". The nucleus is made up of protons and neutrons. The number of protons and neutrons in the nuclei of most elements is approximately the same, so the atomic weight increases in much the same way as the number of protons in the nucleus increases (nuclear charge Z).

The fundamental novelty of the Periodic Law was as follows:

1. A connection was established between elements NOT SIMILAR in their properties. This relationship lies in the fact that the properties of the elements change smoothly and approximately equally with an increase in their atomic weight, and then these changes are PERIODICALLY REPEATED.

2. In those cases where it seemed that some link was missing in the sequence of changes in the properties of elements, the Periodic Table provided for GAPS that had to be filled with yet undiscovered elements.

In all previous attempts to determine the relationship between elements, other researchers sought to create a complete picture in which there was no place for elements not yet discovered. On the contrary, D. I. Mendeleev considered the most important part of his Periodic Table to be those of its cells that were still empty. This made it possible to predict the existence of still unknown elements.

It is admirable that D. I. Mendeleev made his discovery at a time when the atomic weights of many elements were determined very approximately, and only 63 elements were known - that is, a little more than half of those known to us today.

A deep knowledge of the chemical properties of various elements allowed Mendeleev not only to point to yet undiscovered elements, but also to accurately predict their properties! D. I. Mendeleev accurately predicted the properties of the element, which he called "eka-silicon". After 16 years, this element was indeed discovered by the German chemist Winkler and named germanium.

Comparison of the properties predicted by D. I. Mendeleev for the yet undiscovered element "eka-silicon" with the properties of the element germanium (Ge). In the modern Periodic Table, germanium occupies the place of "eka-silicon".

Property

Predicted by D. I. Mendeleev for "eka-silicon" in 1870

Determined for germanium Ge, discovered in 1886

Color, appearance

brown

light brown

Atomic weight

72,59

Density (g/cm3)

5,5

5,35

Oxide Formula

XO2

GeO2

Chloride Formula

XCl4

GeCl4

Density of chloride (g/cm3)

1,9

1,84

In the same way, the properties of "eka-aluminum" (the element gallium Ga, discovered in 1875) and "eka-boron" (the element scandium Sc discovered in 1879) predicted by D. I. Mendeleev were brilliantly confirmed.

After that, it became clear to scientists all over the world that the Periodic Table of D. I. Mendeleev not only systematizes the elements, but is a graphic expression of the fundamental law of nature - the Periodic Law.

Structure of the Periodic System.

Based on the Periodic Law of D.I. Mendeleev created the Periodic Table of Chemical Elements, which consisted of 7 periods and 8 groups (short period version of the table). At present, the long-period version of the Periodic Table is more often used (7 periods, 8 groups, the elements - lanthanides and actinides are shown separately).

Periods are horizontal rows of the table, they are divided into small and large. In small periods there are 2 elements (1st period) or 8 elements (2nd, 3rd periods), in large periods - 18 elements (4th, 5th periods) or 32 elements (6th, 7th period). Each period begins with a typical metal and ends with a non-metal (halogen) and a noble gas.

Groups are vertical sequences of elements, they are numbered with Roman numerals from I to VIII and Russian letters A and B. The short-period version of the Periodic System included subgroups of elements (main and secondary).

A subgroup is a collection of elements that are unconditional chemical analogues; often the elements of a subgroup have the highest oxidation state corresponding to the group number.

In A-groups, the chemical properties of elements can vary over a wide range from non-metallic to metallic (for example, in the main subgroup of group V, nitrogen is a non-metal, and bismuth is a metal).

In the Periodic system, typical metals are located in group IA (Li-Fr), IIA (Mg-Ra) and IIIA (In, Tl). Non-metals are located in groups VIIA (F-Al), VIA (O-Te), VA (N-As), IVA (C, Si) and IIIA (B). Some elements of the A-groups (beryllium Be, aluminum Al, germanium Ge, antimony Sb, polonium Po and others), as well as many elements of the B-groups, exhibit both metallic and non-metallic properties (amphoteric phenomenon).

For some groups, group names are used: IA (Li-Fr) - alkali metals, IIA (Ca-Ra) - alkaline earth metals, VIA (O-Po) - chalcogens, VIIA (F-At) - halogens, VIIIA (He-Rn ) are noble gases. The form of the Periodic system, which was proposed by D.I. Mendeleev, was called short-period or classical. Currently, another form of the Periodic System is used more - long-period.

Periodic law D.I. Mendeleev and the Periodic Table of Chemical Elements became the basis of modern chemistry. Relative atomic masses are given according to the International Table of 1983. For elements 104-108, the mass numbers of the longest-lived isotopes are given in square brackets. The names and symbols of the elements given in parentheses are not generally accepted.

IV Periodic law and the structure of the atom.

Basic information about the structure of atoms.

At the end of the 19th - beginning of the 20th century, physicists proved that the atom is a complex particle and consists of simpler (elementary) particles. Were discovered:

cathode rays (English physicist J. J. Thomson, 1897), whose particles are called electrons e− (carry a unit negative charge);

natural radioactivity of elements (French scientists - radiochemists A. Becquerel and M. Sklodowska-Curie, physicist Pierre Curie, 1896) and the existence of α-particles (helium nuclei 4He2+);

the presence of a positively charged nucleus in the center of the atom (English physicist and radiochemist E. Rutherford, 1911);

artificial conversion of one element into another, for example, nitrogen into oxygen (E. Rutherford, 1919). From the nucleus of an atom of one element (nitrogen - in Rutherford's experiment), upon collision with an α-particle, the nucleus of an atom of another element (oxygen) was formed and a new particle carrying a unit positive charge and called a proton (p +, nucleus 1H)

the presence in the nucleus of an atom of electrically neutral particles - neutrons n0 (English physicist J. Chadwick, 1932).

As a result of the studies, it was found that in the atom of each element (except 1H) there are protons, neutrons and electrons, and protons and neutrons are concentrated in the nucleus of the atom, and electrons - on its periphery (in the electron shell).

The number of protons in the nucleus is equal to the number of electrons in the shell of an atom and corresponds to the serial number of this element in the Periodic system.

The electron shell of an atom is a complex system. It is divided into subshells with different energies (energy levels); levels, in turn, are subdivided into sublevels, and sublevels include atomic orbitals, which can differ in shape and size (denoted by the letters s, p, d, f, etc.).

So, the main characteristic of an atom is not the atomic mass, but the magnitude of the positive charge of the nucleus. This is a more general and precise description of the atom, and hence of the element. All the properties of the element and its position in the periodic system depend on the magnitude of the positive charge of the atomic nucleus. Thus, the serial number of a chemical element numerically coincides with the charge of the nucleus of its atom. The periodic system of elements is a graphic representation of the periodic law and reflects the structure of the atoms of the elements.

The theory of the structure of the atom explains the periodic change in the properties of elements. An increase in the positive charge of atomic nuclei from 1 to 110 leads to a periodic repetition of elements of the structure of the external energy level in atoms. And since the properties of the elements mainly depend on the number of electrons at the outer level, they also repeat periodically. This is the physical meaning of the periodic law.

Each period in the periodic system begins with elements whose atoms at the outer level have one s-electron (incomplete outer levels) and therefore exhibit similar properties - they easily give up valence electrons, which determines their metallic character. These are alkali metals - Li, Na, K, Rb, Cs.

The period ends with elements whose atoms at the outer level contain 2 (s2) electrons (in the first period) or 8 (s2p6) electrons (in all subsequent ones), that is, they have a completed outer level. These are the noble gases He, Ne, Ar, Kr, Xe, which have inert properties.

In 1869, D. I. Mendeleev, based on an analysis of the properties of simple substances and compounds, formulated the Periodic Law: "The properties of simple bodies and compounds of elements are in a periodic dependence on the magnitude of the atomic masses of the elements." On the basis of the periodic law, the periodic system of elements was compiled. In it, elements with similar properties were combined into vertical columns of the group. In some cases, when placing elements in the Periodic system, it was necessary to violate the sequence of increasing atomic masses in order to observe the periodicity of the repetition of properties. For example, tellurium and iodine, as well as argon and potassium, had to be "swapped". The reason is that Mendeleev proposed the periodic law at a time when nothing was known about the structure of the atom. After the planetary model of the atom was proposed in the 20th century, the periodic law is formulated as follows:

"The properties of chemical elements and compounds are in a periodic dependence on the charges of atomic nuclei."

The charge of the nucleus is equal to the number of the element in the periodic system and the number of electrons in the electron shell of the atom. This formulation explained the "violations" of the Periodic Law. In the Periodic system, the period number is equal to the number of electronic levels in the atom, the group number for elements of the main subgroups is equal to the number of electrons in the outer level.

Scientific significance of the periodic law. The periodic law made it possible to systematize the properties of chemical elements and their compounds. When compiling the periodic system, Mendeleev predicted the existence of many yet undiscovered elements, leaving free cells for them, and predicted many properties of undiscovered elements, which facilitated their discovery. The first of these followed four years later.

But not only in the discovery of a new great merit of Mendeleev.

Mendeleev discovered a new law of nature. Instead of disparate, unrelated substances, a single harmonious system arose before science, uniting all the elements of the Universe into a single whole, atoms began to be considered as:

1. organically interconnected by a common pattern,

2. detecting the transition of quantitative changes in atomic weight to qualitative changes in their chemical. personalities,

3. indicating that the opposite of metallic. and non-metallic properties of atoms is not absolute, as previously thought, but only relative.

24. The emergence of structural theories in the development of organic chemistry. Atomic-molecular theory as a theoretical basis for structural theories.

Organic chemistry. Throughout the 18th century in the question of the chemical relationships between organisms and substances, scientists were guided by the doctrine of vitalism - a doctrine that considered life as a special phenomenon, subject not to the laws of the universe, but to the influence of special vital forces. This view was inherited by many scientists of the 19th century, although its foundations were shaken as early as 1777, when Lavoisier suggested that respiration is a process analogous to combustion.

In 1828, the German chemist Friedrich Wöhler (1800–1882), heating ammonium cyanate (this compound was unconditionally considered an inorganic substance), obtained urea, a waste product of humans and animals. In 1845, Adolf Kolbe, a student of Wöhler, synthesized acetic acid from the starting elements carbon, hydrogen, and oxygen. In the 1850s, the French chemist Pierre Berthelot began systematic work on the synthesis of organic compounds and obtained methyl and ethyl alcohols, methane, benzene, and acetylene. A systematic study of natural organic compounds has shown that they all contain one or more carbon atoms and very many contain hydrogen atoms. Type theory. The discovery and isolation of a huge number of complex carbon-containing compounds sharply raised the question of the composition of their molecules and led to the need to revise the existing classification system. By the 1840s, chemists realized that Berzelius' dualistic ideas only applied to inorganic salts. In 1853 an attempt was made to classify all organic compounds by type. A generalized "theory of types" was proposed by the French chemist Charles Frederic Gerard, who believed that the association of various groups of atoms is determined not by the electric charge of these groups, but by their specific chemical properties.

Structural chemistry. In 1857, Kekule, proceeding from the theory of valency (by valency was understood as the number of hydrogen atoms that combine with one atom of a given element), suggested that carbon is tetravalent and therefore can combine with four other atoms, forming long chains - straight or branched. Therefore, organic molecules began to be depicted not as combinations of radicals, but as structural formulas - atoms and bonds between them.

In 1874 a Danish chemist Jacob van't Hoff and the French chemist Joseph Achille Le Bel (1847–1930) extended this idea to the arrangement of atoms in space. They believed that molecules are not flat, but three-dimensional structures. This concept made it possible to explain many well-known phenomena, such as spatial isomerism, the existence of molecules of the same composition but with different properties. The data fit very well. Louis Pasteur about the isomers of tartaric acid.

6. Periodic law and periodic system d.I. Mendeleev Structure of the periodic system (period, group, subgroup). The meaning of the periodic law and the periodic system.

Periodic D.I. law Mendeleev:Properties of simple bodies, as well as shapes and properties of compoundselements are in a periodic dependence onthe values ​​of the atomic weights of the elements. (The properties of the elements are in a periodic dependence on the charge of the atoms of their nuclei).

Periodic system of elements. Series of elements within which properties change sequentially, such as a series of eight elements from lithium to neon or from sodium to argon, Mendeleev called periods. If we write these two periods one below the other so that sodium is under lithium, and argon is under neon, then we get the following arrangement of elements:

With this arrangement, elements that are similar in their properties and have the same valency, for example, lithium and sodium, beryllium and magnesium, etc., fall into the vertical columns.

Dividing all the elements into periods and arranging one period under another so that elements similar in properties and type of compounds formed fall under each other, Mendeleev compiled a table, which he called the periodic system of elements by groups and series.

The value of the periodic systemWe. The Periodic Table of the Elements had a great influence on the subsequent development of chemistry. Not only was it the first natural classification of the chemical elements, which showed that they form a coherent system and are in close connection with each other, but it was also a powerful tool for further research.

7. Periodic change in the properties of chemical elements. Atomic and ionic radii. Ionization energy. Affinity for an electron. Electronegativity.

The dependence of atomic radii on the charge of the atomic nucleus Z has a periodic character. Within one period, with an increase in Z, there is a tendency to a decrease in the size of the atom, which is especially clearly observed in short periods.

With the beginning of the construction of a new electron layer, more distant from the nucleus, i.e., during the transition to the next period, the atomic radii increase (compare, for example, the radii of fluorine and sodium atoms). As a result, within the subgroup, as the charge of the nucleus increases, the sizes of atoms increase.

The loss of electron atoms leads to a decrease in its effective size, and the addition of excess electrons leads to an increase. Therefore, the radius of a positively charged ion (cation) is always less, and the radius of a negatively charged non (anion) is always greater than the radius of the corresponding electrically neutral atom.

Within one subgroup, the radii of ions of the same charge increase with increasing nuclear charge. This pattern is explained by an increase in the number of electron layers and a growing distance of outer electrons from the nucleus.

The most characteristic chemical property of metals is the ability of their atoms to easily give up external electrons and turn into positively charged ions, while non-metals, on the contrary, are characterized by the ability to attach electrons to form negative ions. To detach an electron from an atom with the transformation of the latter into a positive ion, it is necessary to expend some energy, called the ionization energy.

The ionization energy can be determined by bombarding atoms with electrons accelerated in an electric field. The smallest field voltage at which the electron velocity becomes sufficient for the ionization of atoms is called the ionization potential of atoms of a given element and is expressed in volts. With the expenditure of sufficient energy, two, three or more electrons can be torn off from an atom. Therefore, they speak of the first ionization potential (the energy of detachment from the atom of the first electron). The second ionization potential (the energy of detachment of the second electron)

As noted above, atoms can not only donate, but also add electrons. The energy released when an electron is attached to a free atom is called the affinity of the atom for the electron. Electron affinity, like ionization energy, is usually expressed in electronvolts. So, the electron affinity of a hydrogen atom is 0.75 eV, oxygen - 1.47 eV, fluorine - 3.52 eV.

The electron affinity of metal atoms is usually close to zero or negative; from this it follows that for the atoms of most metals, the addition of electrons is energetically unfavorable. The electron affinity of the atoms of non-metals is always positive and the greater, the closer to the noble gas the non-metal is located in the periodic system; this indicates an increase in non-metallic properties as the end of the period is approached.

The possibility of scientific prediction of unknown elements became a reality only after the discovery of the periodic law and the periodic system of elements. D. I. Mendeleev predicted the existence of 11 new elements: ekabor, ekasilicon, ekaaluminum, etc. The "coordinates" of the element in the periodic system (serial number, group and period) made it possible to roughly predict the atomic mass, as well as the most important properties of the predicted element. The accuracy of these predictions increased especially when the predicted element was surrounded by known and sufficiently studied elements.

Thanks to this, in 1875 in France, L. de Boisbaudran discovered gallium (ekaaluminum); in 1879 L. Nilson (Sweden) discovered scandium (ekabor); in 1886 in Germany K. Winkler discovered germanium (ecasilicon).

With regard to the undiscovered elements of the ninth and tenth rows, the statements of D. I. Mendeleev were more cautious, because their properties were studied extremely poorly. So, after bismuth, on which the sixth period ended, two dashes were left. One corresponded to an analogue of tellurium, the other belonged to an unknown heavy halogen. In the seventh period, only two elements were known - thorium and uranium. D. I. Mendeleev left several cells with dashes, which should have belonged to the elements of the first, second and third groups, preceding thorium. An empty cage was also left between thorium and uranium. Five empty places were left for uranium, i.e. almost 100 years later, transuranium elements were foreseen.

To confirm the accuracy of D. I. Mendeleev’s predictions regarding the elements of the ninth and tenth series, we can give an example with polonium (serial number 84). Predicting the properties of the element with the atomic number 84, D. I. Mendeleev designated it as an analogue of tellurium and called it ditellurium. For this element, he assumed an atomic mass of 212 and the ability to form an oxide of the EO e type. This element should have a density of 9.3 g/cm 3 and be a low-melting, crystalline and non-volatile gray metal. Polonium, which was obtained in its pure form only in 1946, is a soft, fusible, silver-colored metal with a density of 9.3 g/cm 3 . Its properties are similar to those of tellurium.

The periodic law of D. I. Mendeleev, being one of the most important laws of nature, is of exceptional importance. Reflecting the natural relationship that exists between the elements, the stages of development of matter from simple to complex, this law laid the foundation for modern chemistry. With his discovery, chemistry ceased to be a descriptive science.

The periodic law and the system of elements of D. I. Mendeleev are one of the reliable methods for understanding the world. Since the elements are united by a common property or structure, this indicates the patterns of interconnection and interdependence of phenomena.

All elements together constitute one line of continuous development from the simplest hydrogen to the 118th element. Such a pattern was first noticed by D. I. Mendeleev, who managed to predict the existence of new elements, thereby showing the continuity of the development of matter.

By comparing the properties of elements and their compounds within groups, one can easily detect the manifestation of the law on the transition of quantitative changes into qualitative ones. So, within any period there is a transition from a typical metal to a typical non-metal (halogen), however, the transition from a halogen to the first element of the next period (an alkali metal) is accompanied by the appearance of properties that are sharply opposite to this halogen. The discovery of D. I. Mendeleev laid an accurate and reliable foundation for the theory of the structure of the atom, having a huge impact on the development of all modern knowledge about the nature of matter.

The work of D. I. Mendeleev on the creation of the periodic system marked the beginning of a scientifically based method for the purposeful search for new chemical elements. Numerous advances in modern nuclear physics can serve as examples. Over the past half century, elements with serial numbers 102-118 have been synthesized. The study of their properties, as well as obtaining, would be impossible without knowledge of the patterns of the relationship between chemical elements.

The evidence for such a statement is results research on the synthesis of elements 114, 116, 118 .

The isotope of the 114th element was obtained by the interaction of plutonium with the 48Ca isotope, and the 116th isotope by the interaction of curium with the 48Ca isotope:

The stability of the resulting isotopes is so high that they do not spontaneously fission, but undergo alpha decay, i.e. fission of the nucleus with the simultaneous emission of alpha particles.

The obtained experimental data fully confirm the theoretical calculations: as successive alpha decays, the nuclei of the 112th and 110th elements are formed, after which spontaneous fission begins:

Comparing the properties of the elements, we are convinced that they are interconnected by a common structural features. Thus, by comparing the structure of the outer and pre-outer electron shells, it is possible to predict with high accuracy all types of compounds characteristic of a given element. Such a clear relationship is very well illustrated by the example of the 104th element - rutherfordium. Chemists predicted that if this element is an analogue of hafnium (72 Hf), then its tetrachloride properties should be approximately the same as HfCl 4 . Experimental chemical studies confirmed not only the forecast of chemists, but also the discovery of a new superheavy element 1(M Rf. The same analogy can be traced in the properties - Os (Z = 76) and Ds (Z = 110) - both elements form volatile oxides of the R0 4 type. All this speaks of manifestation of the law of interconnection and interdependence of phenomena.

Comparison of the properties of elements both within groups and periods, and their comparison with the structure of the atom, indicate the law transition from quantity to quality. The transition of quantitative changes into qualitative ones is possible only throughnegation of negation. Within periods, with an increase in the charge of the nucleus, there is a transition from an alkali metal to a noble gas. The next period begins again with an alkali metal - an element that completely negates the properties of the noble gas that preceded it (for example, He and Li; Ne and Na; Ar and Kg, etc.).

In each period, the charge of the nucleus of the next element increases by one compared to the previous one. This process is observed from hydrogen to the 118th element and indicates the continuity of the development of matter.

Finally, the combination of opposite charges (proton and electron) in an atom, the manifestation of metallic and non-metallic properties, the existence of amphoteric oxides and hydroxides is a manifestation of the law unity and struggle of opposites.

It should also be noted that the discovery of the periodic law was the beginning of fundamental research concerning the properties of matter.

In the words of Niels Bohr, the periodic system is "a guiding star for research in the field of chemistry, physics, mineralogy, and technology."

• Elements 112, 114, 116, 118 were obtained at the Joint Institute for Nuclear Research (Dubna, Russia). Elements 113 and 115 were obtained jointly by Russian and American physicists. The material was kindly provided by Yu. Ts. Oganesyan, Academician of the Russian Academy of Sciences.