Laboratory staff. Cost: n.s. Gruner S.V., senior researcher Prishchenko A.A., associate professor Livantsova L.I., engineer Reutova T.O., researcher Novikova O.P., associate professor Livantsov M.V., senior researcher Demyanov P.I. Sitting: n.s. Meleshonkova N.N., engineer Shuvalova E.A., associate professor Gopius E.D., prof. Petrosyan V.S., researcher Kochetova E.K., researcher Averochkina I.A.
From the history of the laboratory
In 1957, academician Reutov O.A. created at Moscow State University the Laboratory of Theoretical Problems of Organic Chemistry, which, thanks to the results of kinetic, stereochemical and isotopic studies of the mechanisms of reactions of nucleophilic and electrophilic substitution at the carbon atom carried out by Reutov O.A. and his first doctors of science ( Beletskaya I.P., Bundel Yu.G., Sokolov V.I.) has received wide recognition in the scientific world.
The development of new methods in the laboratory (NMR spectroscopy, electrochemistry) made it possible to obtain unique data on the electronic and spatial structure of various organic and organoelement compounds and to study their behavior in solutions and the solid phase. New generation of doctors of science ( Butin K.P., Kurts A.L., Petrosyan V.S.) continued to grow the authority of the school of academician O.A. Reutov. In 1988, he transferred the leadership of the Laboratory to Professor V.S. Petrosyan. and since then it has been called the Laboratory of Physical Organic Chemistry. The research carried out in subsequent years was widely recognized and received many awards. Laboratory graduates (academicians Beletskaya I.P., Bubnov Yu.N., Egorov M.P., Professor Sokolov V.I., Bakhmutov V.I., Tretyakova N.Yu. head laboratories in our country and abroad. The research currently being carried out in the laboratory is highly appreciated by the Russian and international scientific community.
Main scientific directions of the laboratory
- physical organic chemistry
- chemistry of organoelement compounds
- environmental chemistry and toxicology
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In the scientific group, senior researcher Prishchenko A.A.(assoc. prof. Livantsov M.V., assistant professor Livantsova L.I., research scientist Novikova O.P., research scientist Meleshonkova N.N.) conduct research on new types of organic phosphorus compounds, study their structure and reactivity, as well as complexing and biological activity. Functionalized hydroxy- and aminomethyl derivatives of mono- and diphosphorus-containing acids - promising organophosphorus biomimetics of natural pyrophosphates and amino acids - are widely used as effective ligands and biologically active substances with various properties. The Laboratory of Physical Organic Chemistry has developed convenient methods for the synthesis of new types of these substances using highly reactive synthons - trimethylsilyl esters of trivalent phosphorus acids and functionalized carbonyl compounds, including aromatic, heterocyclic and unsaturated fragments. The resulting compounds are of interest for the production of diphosphorus-containing peptides, as well as effective polydentate ligands, promising antioxidants and cytoprotectors with multiple mechanisms of antioxidant action. The work was supported by several grants from the Russian Foundation for Basic Research. |
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The main direction of work of the senior researcher Demyanova P.I. consists of a theoretical study of the nature of intramolecular, primarily non-covalent, interactions between pairs of atoms in a molecule (complex, crystal), determining whether these interactions are bonding (stabilizing) or repulsive (destabilizing). The need for such work is dictated by the fact that many foreign and domestic researchers, based on the formal interpretation of the results of a topological analysis of the electron density distribution within the framework of the quantum theory of atoms in molecules (QTAVM) created by Bader, persistently assert the existence, for example, of bonding interactions between like-charged ions or the absence of intramolecular hydrogen bonds in ethylene glycol and other 1,2-diols (including sugars) and many other organic molecules. Another direction of theoretical calculations is aimed at obtaining information about the energy and nature of metal-metal interactions in solutions and crystals of Cu(I)-, Ag(I)- and Au(I)- organic compounds and complexes of these metals with organic and inorganic ligands. This information will help shed light on metallophilic interactions, the existence of which is still questioned. |
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N.s. Gruner S.V., having worked for many years in the chemistry of organic derivatives of silicon, germanium and tin, in recent years, with graduate students, he has obtained a large series of hypercoordination tin compounds that have interesting structural features and exhibit unusual reactivity. |
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Assistant professor Gopius E.D.– curator of teaching organic chemistry at the Faculty of Biology, deputy head of the laboratory. Scientific research focuses on the chemistry of carbocations. |
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Engineer Shuvalova E.A. deals with organic chemistry and toxicology of aquatic ecosystems. He pays a lot of attention to organizing the work of the Open Ecological University of Moscow State University, created in 1987 by prof. Petrosyan V.S. |
Is a long life of a free particle in non-free matter possible?
Possible. It is a stable free radical. They were discovered within the framework of physical organic chemistry, which deals with general issues of organic chemistry, paying special attention to the study of the mechanisms of organic reactions, as well as the quantitative relationship between the chemical structure of organic compounds, their properties and reactivity.
One of the achievements is the discovery and practical use of stable free radicals (a type of molecule or atom capable of independent existence and having one or two unpaired electrons).
The stability of free radicals is facilitated by so-called steric hindrances (steric hindrances), when the atom on which the unpaired electron is localized is reliably shielded from other reagents by nearby bulky substituents. It’s like an individual person in a crowd - it seems like he’s free to go wherever he wants - but no, your environment is holding you back!
Methods for the synthesis and technology for the production of stable nitroxyl radicals of the imidazoline series and their precursors have been developed, which are used as spin labels, probes and traps in scientific research and industry. Radicals are determined by electron paramagnetic resonance (EPR) in concentrations of 10 (to the power of -10) mole percent.
The technical and economic advantages of the developed nitroxyl radicals of the imidazoline series are determined by their uniqueness (there are no natural analogues), which allows them to be used as indicators of the movement of reservoir fluids (oil production), hidden marks, etc. Stable free radicals are used in various fields of science and technology:
In medicine and biology, stable nitroxyl radicals are usually used as spin labels. All molecules of spin labels, despite the diversity of their chemical structure, as a rule, contain the same paramagnetic fragment - a chemically stable nitroxyl radical (>N-O*). An unpaired electron is localized on this radical, serving as a source of the ESR signal. The specific choice of spin labels is determined by the research problem. For example, in order to monitor conformational rearrangements of proteins using spin tags, tag molecules are usually “sewn” to certain regions of the protein. In this case, the spin label must contain a special reaction group that can form a covalent chemical bond with the amino acid residues of the protein molecule.
To study the properties of artificial and biological membranes, lipid-soluble spin labels are usually used that can be incorporated into the lipid layer of the membrane:
- as spin pH probes for measuring the pH value in cellular organelles;
- when studying the processes of ion transfer through membranes;
- to determine the localization of drugs or other drugs in organs or tissues.
In analytical chemistry and geophysics:
- to create chelating agents capable of binding to metals;
- as indicators with a low detection threshold (for example, instead of tritium or organic dyes in oil exploration and production, instead of fluoroaromatic acids in the analysis of groundwater movement).
In other industries, for adding as hidden marks to fuels, alcohols, etc.
The Institute of Organic and Physical Chemistry of the Kazakh Scientific Center of the Russian Academy of Sciences (Kazan) operates. You can read: Gammet L. Fundamentals of physical organic chemistry. Speeds, equilibria and reaction mechanisms. M.: Mir, 1972.
The Department of Physical and Organic Chemistry of USPTU was created as a result of the merger of the departments of physical and organic chemistry in 1983. The history of the Department of Physical Chemistry should be traced back to 1947, when for the first time 3rd year students of the Ufa branch of the Moscow Petroleum Institute were not sent to continue their studies in Moscow. To organize laboratory work in those years, a glass-blowing workshop was created at the department, which greatly accelerated the introduction of various glass chemical equipment into the laboratory workshop. Over the course of several years, work was carried out on physical and colloidal chemistry. Much credit for this goes to the senior teacher Lyubov Nisovne Pirkis. The chief engineer of the Dubitel plant, D. M. Rapoport, was invited to give lectures on physical chemistry, who then joined the staff of the institute, working there until his retirement. The Department of Physical Chemistry received its official name in 1951, when Professor Boris Vasilievich Klimenok, who headed the department for 32 years until its reorganization in 1983. At the time of the creation of the department, educational and scientific work was carried out by 8 employees. Under the leadership of Professor B.V. Klimenok, the department began scientific research on the dewaxing of diesel fuels with urea. Comprehensive research made it possible to develop a dewaxing process using water-ethanol solutions of urea and build a pilot plant at the Ufa Oil Refinery. Pilot industrial results proved the real possibility of the process, however, large losses of alcohol and urea made the process ineffective. The department independently switched to studying the dewaxing of diesel fuels with a water-urea suspension.
Beginning with 1954 year, graduate students are involved in scientific work. Engineer A. T. Zemlyansky put a lot of effort into the development of an aqueous version of urea dewaxing and the practical implementation of the process. He was a man with “golden hands” and an excellent engineering mind. He independently carried out work on a lathe, performed welding and metalwork. It can be considered that a large pilot plant for dewaxing petroleum fractions with urea, built in the mechanical workshop of the institute, is the brainchild of A. T. Zemlyansky. Along with the pilot plant, a semi-industrial plant was built at the National Oil Refinery Plant. However, for a number of reasons, this installation ceased to exist, and the development of the process continued mainly in the pilot installation of the institute. The possibility of new process options was explored. Dissertations on this topic were defended by Z. V. Basyrov, F. A. Chegodaev, A. M. Syrkin, A. A. Krasnov, T. S. Kamkina, P. O. Ivanov, V. G. Abdula, R. M Abzalov. The expansion of the department occurred in 1968, when it became involved in research on the production of carbon materials with certain properties based on petroleum products and the use of waste sulfuric acid from oil refineries. 7 candidate dissertations were defended on this topic. It should be noted the work of A. A. Senko, who impeccably conducted both the laboratory workshop and all reporting at the department. Currently, this work is continued by engineer G. G. Medvedeva. She is responsible not only for the full provision of laboratory workshops, which is not easy in our time, but also for the management of departmental affairs and the organization of public events.
The Department of Organic Chemistry also dates back to 1947. Over the years, the department was reorganized, merging with other departments.
IN 1974 year, the department again became independent, it was headed by Professor R. B. Valitov. Under his leadership, scientific research in the field of petrochemical and heavy organic synthesis began to develop. Associate professors M. G. Safarov, D. M. Torikov, R. V. Faskhutdinova, A. I. Naimushin made a great contribution to the methodological work of the department. Associate Professor M. G. Safarov, together with Academician V. I. Isagulyants, created a school of chemistry of 1,3-dioxanes. Candidate's and doctoral dissertations in this area were defended by D. L. Rakhmankulov, U. B. Imashevm, E. A. Kantor, S. S. Zlotsky and others.
It is impossible not to note the great contribution to the organization of a laboratory workshop in organic chemistry by one of the oldest laboratory assistants at the university, T.V. Grigorieva, heads of laboratories R.T. Zinatullin and I.N. Gilmanov, senior laboratory assistants E.E. Sukhareva, who have been working at the department for many years, R. A. Bikimbetov, G. D. Vorshev.
IN 1981 each year a professor is elected head of the department Ural Bulatovich Imashev. After the merger of the departments of physical and organic chemistry, Professor U. B. Imashev continued to lead the department. With his arrival, research work on promising technologies for processing heavy oil raw materials intensified. The development and expansion of this direction led to the creation of a scientific school, the results of which were new, unparalleled technologies for processing oil residues in a wave field. Original methods for the synthesis of various functionally substituted mono-, di- and trithiocarbonates have been developed, starting from carbon sulphide and disulfide. Their properties and transformations in various chemical reactions have been studied. On organosulfur topics, 3 candidate's dissertations were defended at the department (R. F. Mavlyutov, N. T. Chanyshev, R. F. Ishteev). Associate Professor S.M. made a certain contribution to the scientific potential of the department. Kalashnikov on the synthesis of acetals and sulfur oxide derivatives. The scientific potential of the department has grown noticeably with the arrival of Professor E.M. Kuramshin. He developed a kinetic direction in the study of the chemistry of acetals. Professor E.M. Kuramshin leads work in the field of oxidative transformations of organic compounds, as well as the kinetics and mechanism of chemical reactions. Over the past five years, the department has completed research worth 1,800 thousand rubles.
The flowcharts developed by Associate Professor A.I. Naimushin were introduced into laboratory classes in organic chemistry. Professor E.M. Kuramshin published methodological recommendations for laboratory work in physical and colloidal chemistry using a flowchart, and a problem book in physical chemistry. Professor E.M. Kurashin, associate professors O.P. Zhurkin and O.B. Zvorygina published a laboratory workshop on physical chemistry.
Professor U. B. Imashev published the textbook “Fundamentals of Organic Chemistry” - M.: Kolos, 2011. - P.-464 p. and "Laboratory workshop on organic chemistry". - Ufa: USNTU, 2009.-236p. Together with Associate Professor S. M. Kalashnikov, Associate Professor N. T. Chanyshev and Professor E. A. Udalova, he published a textbook “Problems and exercises in organic chemistry” Ufa: USNTU, 2011. - 236 p. In co-authorship with Associate Professor O.P. Zhurkin, a textbook “Physical and chemical methods of analysis of organic compounds” was published. Ufa: - USNTU, 2009. - 211 p.
Over the past five years, teachers and staff of the department have published 12 textbooks and teaching aids, 7 monographs, 135 scientific articles, received 19 patents, 3 grants.
Since Lavoisier's time, chemists have been able to predict in which direction particular fast ionic reactions of relatively small molecules will go, and have been able to modify these reactions for practical use. Studying complex molecules was much more difficult. Slow reactions of organic compounds were also much more difficult to analyze. Often reactions could take several paths, and the chemist was allowed to direct the reaction along the desired path through his skill as an experimenter and intuition, and not through a deep understanding of the process.
With the advent of the electronic model of the atom, organic chemists were able to take a fresh look at their field of research. At the end of the 20s of the XX century. English chemist Christopher Ingold (1893-1970) and a number of other chemists tried to approach organic reactions from the standpoint of the theory of atomic structure, explaining the interaction of molecules by the transition of electrons. In organic chemistry, methods of physical chemistry began to be intensively used. An important discipline has become physical organic chemistry .
However, attempts to interpret organic reactions only as a result of the movement of electrons have not led to much success.
During the first quarter of the 20th century, since the discovery of the electron, it was considered proven that the electron was a very small, hard ball. However, in 1923, French physicist Louis Victor de Broglie (b. 1892) presented a theoretical justification that electrons (and all other particles) have wave properties. By the end of the 20s of the XX century. this hypothesis was confirmed experimentally.
Pauling (who was the first to suggest that molecules of proteins and nucleic acids have a spiral shape, see Chapter 10) in the early 30s developed methods that made it possible to take into account the wave nature of electrons when considering organic reactions.
He suggested that the socialization of a pair of electrons (according to Lewis and Langmuir) can be interpreted as the interaction of waves or the overlap of electron clouds. The chemical bond, depicted as a feature in Kekule’s structural theory, corresponds in the new concepts to the region of maximum overlap of electron clouds. It turned out that the overlap of electron clouds sometimes occurs not only in a single direction, represented by a valence bond in the structural formula. In other words, the true structure of a molecule cannot be represented even approximately by any single structural formula. It can, however, be considered as intermediate between several hypothetical structures, as a “resonant hybrid” of these structures. It is important to note that the energy of such a real molecule is lower than would be expected based on any single resonant "classical" structure. Such molecules are said to be “stabilized by resonance,” although resonance in this case, of course, is not a real physical phenomenon, but a convenient theoretical concept to explain the stability and properties of some molecules.
Resonance theory has proven particularly useful in understanding the structure of benzene, which has puzzled chemists since the time of Kekule (see Chapter 7). The formula for benzene was usually depicted as a hexagon with alternating single and double bonds. However, benzene is almost completely devoid of the properties characteristic of compounds with double bonds.
But for benzene, you can write a second, completely equivalent Kekulé formula, in which the simple and double bonds are swapped compared to the first formula. The actual benzene molecule is described as a resonant hybrid of two Kekulé structures; the electrons responsible for the formation of double bonds are delocalized, “spread” around the ring, so that all bonds between carbon atoms in benzene are equivalent and are intermediate between classical single and double bonds. This is precisely the reason for the increased stability and peculiarities of the chemical behavior of benzene.
In addition to the structure of benzene, ideas about the wave properties of electrons helped explain other issues. Since the four electrons located on the outer shell of a carbon atom are not entirely equivalent in energy, one could assume that the bonds formed between the carbon atom and its neighboring atoms differ somewhat depending on which of the electrons are involved in the formation of one or another communications.
However, it turned out that four electrons, like waves, interact with each other and form four “middle” bonds, which are completely equivalent and directed towards the vertices of the tetrahedron, as in the van’t Hoff-Le Bel tetrahedral atom.
At the same time, resonance helped explain the structure of a group of unusual compounds that chemists first encountered at the beginning of the 20th century. In 1900, the American chemist Moses Gomberg (1866-1947) tried to obtain hexaphenylethane, a compound in the molecule of which two carbon atoms are connected to six benzene rings (three for each carbon atom).
Instead of this compound, Gomberg received a colored solution of some very reactive compound. For a number of reasons, Gomberg believed that he had received triphenylmethyl- a “half-molecule” consisting of a carbon atom and three benzene rings, in which the fourth bond of the carbon atom is unsaturated.
This compound resembled one of those radicals whose concept was introduced in the 19th century. to explain the structure of organic compounds (see Chapter 6). However, unlike the radicals of the old theory, the molecule discovered by Gomberg existed in an isolated form, and not as a fragment of another compound, so it was called free radical .
With the development of electronic concepts of chemical bonding, it became clear that in free radicals, for example in triphenylmethyl, an unsaturated bond (in terms of Kekule’s theory) in the framework of new concepts corresponds to an unpaired electron. Typically, such molecules with an unpaired electron are extremely reactive and quickly transform into other substances.
However, if the molecule is flat and symmetrical (like a triphenylmethyl molecule), then the unpaired electron can be “smeared” throughout the molecule, which will lead to stabilization of the radical.
When the study of organic reactions was approached from the perspective of the theory of electronic structure, it became obvious that reactions often include the stage of formation of free radicals. Such free radicals, as a rule, not stabilized by resonance, exist only for a short time and are always formed with difficulty. It is because of the difficulty of forming free radical intermediates that most organic reactions proceed so slowly.
In the second quarter of the 20th century. Organic chemists began to penetrate ever deeper into the essence of organic reactions, and having studied the mechanism of reactions, comprehending the very essence of the process, they were able to synthesize molecules whose complexity amazed chemists of earlier generations.
However, the concepts of resonance theory are applicable not only in organic chemistry. Based on old ideas, it is impossible, in particular, to clearly explain the structure of borohydride molecules. The boron atom has too few valence electrons to form the required number of bonds. If we assume that the electrons are appropriately “smeared,” then we can propose an acceptable molecular structure.
Although since the discovery of inert gases it was believed that they did not enter into any reactions, in 1932 Pauling suggested that the atoms of these gases should form bonds.
Initially, this assumption of Pauling went unnoticed, but in 1962, as a result of the reaction of the inert gas xenon with fluorine, xenon fluoride. Soon after it, a number of other compounds of xenon with fluorine and oxygen, as well as compounds of radon and krypton, were obtained.
Half life
The study of the structure of the atom led to a new understanding of the problem, but at the same time, scientists faced a number of new questions.
In 1900, Crookes (see Chapter 12) discovered that freshly prepared compounds of pure uranium have only very little radioactivity and that the radioactivity of these compounds increases with time. By 1902, Rutherford and his collaborator, the English chemist Frederick Soddy (1877-1956), proposed that with the emission of an alpha particle the nature of the uranium atom changes and that the new atom produced gives off stronger radiation than uranium itself (thus, the observation was taken into account here Crookes). This second atom in turn also splits, forming another atom. Indeed, the uranium atom gives rise to a whole series of radioactive elements - radioactive series, including radium and polonium (see section "Ordinal Number") and ending with lead, which is not radioactive. It is for this reason that radium, polonium and other rare radioactive elements can be found in uranium minerals. The second radioactive series also begins with uranium, while the third radioactive series begins with thorium.
It is appropriate to ask why radioactive elements, constantly decaying, still continue to exist? In 1904, this issue was resolved by Rutherford. By studying the rate of radioactive decay, he showed that after a certain period, different for different elements, half of a given amount of a given radioactive element decays. This period, characteristic of each individual type of radioactive substance, Rutherford called half-life(Fig. 22).
Rice. 22. The half-life of radon is determined by measuring the amount of substance remaining at regular intervals. The resulting dependence is a “decaying” exponential curve y=e-ah .
The half-life of radium, for example, is just under 1600 years. Over the course of geological epochs, any amount of radium in the earth's crust would, of course, have disappeared long ago if it had not been constantly replenished by the decay of uranium. The same can be said about other uranium decay products, including those whose half-lives are measured in fractions of a second.
The half-life of uranium itself is 4,500,000,000 years. This is a huge period of time, and over the entire history of the Earth, only a part of the original uranium reserves could decay. Thorium decays even more slowly, with a half-life of 14,000,000,000 years.
Such huge periods of time can only be determined by counting the number of alpha particles emitted by a given mass of uranium (or thorium). Rutherford counted alpha particles by detecting small flashes that occurred when alpha particles collided with a zinc sulfide screen (i.e., using the so-called scintillation counter).
Each new alpha particle meant that another uranium atom had decayed, so Rutherford could determine how many atoms were decaying per second. From the mass of uranium he used, Rutherford determined the total number of uranium atoms. Having such data, it was no longer difficult to calculate the time required for the decay of half the available amount of uranium. As it turned out, we are talking about billions of years.
The decay of uranium is such a constant and characteristic process that it can be used to determine the age of the Earth. In 1907, the American chemist Bertram Borden Boltwood (1870-1927) suggested that such determinations could be based on the lead content of uranium minerals. If we assume that all the lead in the minerals came from the decay of uranium, then it is easy to calculate how long it took. Using this method, it was possible to determine that the age of the solid crust is at least four billion years.
Meanwhile, Soddy continued to describe the changes in the atom caused by the release of subatomic particles. If an atom loses an alpha particle (+2 charge), the total charge on its nucleus is reduced by two and the element moves two spaces to the left on the periodic table.
If an atom loses a beta particle (an electron with a charge of -1), the nucleus gains an additional positive charge and the element moves one space to the right in the periodic table. If an atom emits gamma rays (uncharged), the energy content changes, but the composition of the particles is not affected, so it remains the same element.
Guided by these rules, chemists were able to thoroughly study many radioactive series.
Isotopes
With the discovery of radioactive elements, scientists faced a serious problem: what to do with the various decay products of uranium and thorium? They were discovered in dozens, and in the periodic table there were only a maximum of nine places left (from polonium with serial number 84 to uranium with serial number 92) into which they could be placed.
Thus, a uranium atom (serial number 92) emits an alpha particle. The atomic number of the new element, according to Soddy's rule, is 90. This means that a uranium atom must form a thorium atom. However, the half-life of ordinary thorium is measured at 14 billion years, while the half-life of thorium derived from uranium is only 24 days.
Differences are observed even when obtaining non-radioactive elements. For example, Richards (a specialist in atomic masses, see Chapter 5) was able to show in 1913 that the atomic mass of lead obtained from the decay of uranium is somewhat different from the atomic mass of ordinary lead.
Soddy was determined enough to suggest that more than one type of atom could correspond to the same place on the periodic table. Place number 90 can be occupied by different varieties of thorium, place number 82 by different varieties of lead, etc. Soddy called these varieties of atoms occupying the same place in the table, isotopes(from the Greek tópos - place).
Isotopes occupying the same place in the table must have the same atomic number and, therefore, the same number of protons in the nucleus and the same number of electrons in the shells. Isotopes of an element must have the same chemical properties, since these properties depend on the number and location of electrons in the atoms.
But how, in this case, can we explain the difference in radioactive properties and atomic masses?
In the last century, Prout put forward his famous hypothesis (see Chapter 5), according to which all atoms are composed of hydrogen, so that all elements must have integer atomic masses. However, as it turned out, most atomic masses are non-integer, and this fact seemed to refute the hypothesis.
But, according to new ideas about the structure of the atom, an atom has a nucleus consisting of protons (and neutrons). Protons and neutrons are approximately equal in mass, and therefore the masses of all atoms must be multiples of the mass of a hydrogen atom (consisting of one proton). Prout's hypothesis was revived, but doubts arose again about what atomic masses should be.
In 1912, J. J. Thomson (who, as we said above, discovered the electron) exposed beams of positively charged neon ions to a magnetic field. The magnetic field caused the ions to deflect, causing them to fall onto the photographic plate. If all the ions were the same in mass, then they would all be deflected by the magnetic field at the same angle, and a discolored spot would appear on the photographic film. However, as a result of this experiment, Thomson obtained two spots, one of which was about ten times darker than the other. Thomson's collaborator Francis William Aston (1877-1945), who later improved this device, confirmed the correctness of the data obtained. Similar results were obtained for other elements. This device, which made it possible to separate chemically similar ions into beams of ions with different masses, was called mass spectrograph .
The amount of deflection of equally charged ions in a magnetic field depends on the mass of these ions; ions with higher mass are deflected less, and vice versa. Thus, the experiments of Thomson and Aston showed that there are two types of neon atoms. For one type of atom mass number equal to 20, for the other - 22. As a result of determining the relative blackness of the spots, it was found that the content of neon-20 is 10 times greater than neon-22. Later, the presence of small amounts of neon-21 was also discovered. If, when calculating the atomic mass of neon, we proceed from these data, it turns out that it is equal to approximately 20.2.
In other words, the mass of individual atoms is a whole number multiple of the mass of a hydrogen atom, but the atomic mass of an individual element is the average of the atomic masses of its constituent atoms, and so it may not be a whole number.
The average atomic mass of an element with a large number of isotopes may in some cases be greater than the average atomic mass of an element with a higher atomic number. For example, tellurium, whose atomic number is 52, has seven isotopes. Of these, the two heaviest isotopes, tellurium-126 and tellurium-128, are the most abundant. Consequently, the atomic mass of tellurium approaches 127.6. The atomic number of iodine is 53, i.e. one more than that of tellurium. But iodine has only one isotope, iodine-127, and therefore its atomic mass is 127. When Mendeleev placed iodine behind tellurium in his periodic table and thereby violated the order dictated by atomic mass, he, without knowing it, followed the charges of the nuclei, i.e. i.e. the physical essence of the periodic law.
Let's give another similar example. Potassium (serial number 19) has three isotopes - potassium-39, potassium-40 and potassium-41, but the lightest isotope is the most common - potassium-39. As a result, the atomic mass of potassium is 39.1. The atomic number of argon is 18, and it also has three isotopes - argon-36, argon-38 and argon-40, but the heaviest isotope - argon-40 - is the most common. As a result, the atomic mass of argon is approximately 40.
Using a mass spectrograph, you can measure the masses of individual isotopes and determine the content of these isotopes. Having obtained such data, it is possible to calculate the average atomic mass of the element. The accuracy of this method of determining atomic mass is much higher than that of chemical methods.
Different isotopes of a given element have the same nuclear charges but different mass numbers. Consequently, the nuclei of different isotopes contain the same number of protons, but a different number of neutrons. Neon-20, neon-21 and neon-22 each have 10 protons in the nucleus, the serial number of all these isotopes is 10, and the electrons are distributed among the shells as follows: 2, 8. However, the neon-20 nucleus contains 10 protons plus 10 neutrons, in the neon-21 nucleus has 10 protons plus 11 neutrons, and the neon-22 nucleus has 10 protons plus 12 neutrons.
Most elements (but not all) contain isotopes. In 1935, the American physicist Arthur Geoffrey Dempster (1886-1950) established, for example, that natural uranium, whose atomic mass (238.07) is very close to an integer, is a mixture of two isotopes. One of the isotopes is contained in a predominant amount (99.3%). The nuclei of this isotope consist of 92 protons and 146 neutrons, i.e., the total mass number is 238. This is uranium-238. The content of another isotope, uranium-235, is only 0.7%; there are three fewer neutrons in the nucleus of this isotope.
Since radioactive properties depend on the structure of the atomic nucleus, and not on the electronic environment, isotopes of the same element can have similar chemical properties and completely different radioactivity. While the half-life of uranium-238 is 4,500,000,000 years, the half-life of uranium-235 is only 700,000,000 years. Both of these elements are the first elements of two separate radioactive series.
There were theoretical premises that suggested that hydrogen, the simplest of elements, could also have a pair of isotopes. The nuclei of ordinary hydrogen atoms consist of one proton, i.e. ordinary hydrogen is hydrogen-1. In 1931, American chemist Harold Clayton Urey (1893-1980) proposed that the heavier isotope of hydrogen, if it existed, should boil at a higher temperature, evaporate more slowly, and accumulate in a residue.
In an attempt to detect this heavier isotope of hydrogen, Yuri began to slowly evaporate four liters of liquid hydrogen. And in the last cubic centimeter of hydrogen, Urey actually found unmistakable signs of the presence of hydrogen-2, an isotope whose nucleus contains one proton and one neutron. Hydrogen-2 was named deuterium .
Oxygen was no exception. In 1929, the American chemist Williams Francis Gioc (b. 1895) managed to show that oxygen has three isotopes. Oxygen-16 is the most abundant, accounting for about 99.8% of all atoms. There are 8 protons and 8 neutrons in the oxygen-16 nucleus. The nucleus of oxygen-18, the second most abundant isotope, has 8 protons and 10 neutrons; the nucleus of oxygen-17, which is found only in trace amounts, has 8 protons and 9 neutrons.
This created a problem. Since the time of Berzelius, the atomic masses of elements have been calculated under the assumption that the atomic mass of oxygen is 16.0000 (see Chapter 5). But the atomic mass of oxygen could only be the calculated average atomic mass of the three isotopes, and the ratio of oxygen isotopes could vary greatly from sample to sample.
Physicists began to determine atomic masses based on the atomic mass of oxygen-16, equal to 16.0000. As a result, a number of values were obtained ( physical atomic mass), which by a very small constant value exceeded the values that were used and which were gradually refined throughout the 19th century. ( chemical atomic weights).
In 1961, international organizations of both chemists and physicists agreed to adopt the atomic mass of carbon-12 as the standard, setting it to be exactly 12.0000. The atomic masses of the elements calculated using the new standard are almost exactly the same as the old chemical atomic weights, and, in addition, the new standard is associated with only one isotope, and not a galaxy of isotopes.
Chapter 14 Nuclear Reactions
New transformations
Once it became apparent that the atom was composed of smaller particles that rearranged randomly during radioactive transformations, the next step seemed almost preordained.
Man has learned to rearrange molecules at his discretion using ordinary chemical reactions. Why not try rearranging the nuclei of atoms using nuclear reactions? Protons and neutrons are much more tightly bound than the atoms in a molecule, and the usual methods used to carry out ordinary chemical reactions will naturally not lead to success. But you can try to develop new methods.
The first step in this direction was taken by Rutherford; he bombarded various gases with alpha particles and discovered that every time an alpha particle struck the nucleus of an atom, it disrupted its structure (Fig. 23).
In 1919, Rutherford was already able to show that alpha particles could knock out protons from nitrogen nuclei and combine with what remained of the nucleus. The most common isotope of nitrogen is nitrogen-14, which has 7 protons and 7 neutrons in its nucleus. If you knock out a proton from this nucleus and add 2 protons and 2 neutrons of an alpha particle, you will get a nucleus with 8 protons and 9 neutrons, i.e. an oxygen-17 nucleus. The alpha particle can be thought of as helium-4 and the proton as hydrogen-1. Thus, Rutherford was the first to successfully carry out an artificial nuclear reaction:
Nitrogen-14 + helium-4 → oxygen-17 + hydrogen-1
By transforming one element into another, he accomplished transmutation. So, in the 20th century. the most cherished dream of the alchemists came true.
Over the next five years, Rutherford conducted a series of other nuclear reactions using alpha particles. However, its capabilities were limited, since radioactive elements produced alpha particles only with average energy. Particles with much higher energies were needed.
Rice. 23. Scheme of Rutherford's experiment. The emitted alpha particles are deflected as they pass through the gold foil; the magnitude of the deviation is recorded when the particles collide with the fluorescent screen.
Physicists began to create devices designed to accelerate charged particles in an electric field. By forcing particles to move with acceleration, their energy could be increased. English physicist John Douglas Cockroft (1897-1967), together with his collaborator Irish physicist Ernest Thomas Sinton Walton (born 1903), were the first to develop the idea of an accelerator that made it possible to produce particles with energy sufficient to carry out a nuclear reaction. In 1929, such an accelerator was built. Three years later, the same physicists bombarded lithium atoms with accelerated protons and obtained alpha particles. This nuclear reaction can be written as follows:
Hydrogen-1 + lithium-7 → helium-4 + helium-4
In the Cockcroft-Walton accelerator and a number of other similar accelerators, particles moved along a straight path. It was possible to obtain high-energy particles in such an accelerator only if the particle path was sufficiently long, so accelerators of this type were extremely bulky. In 1930, the American physicist Ernest Orlando Lawrence (1901-1958) proposed an accelerator in which particles moved in a slightly diverging spiral. This one is relatively small cyclotron could produce particles with extremely high energy.
Lawrence's first very small cyclotron is the forerunner of today's gigantic installations, half a kilometer in circumference, that are used in the search for answers to the most complex questions related to the structure of matter.
In 1930, the English physicist Paul Adrien Morris Dirac (born in 1902) theoretically substantiated the assumption that both protons and electrons should have their own antiparticles . Antielectron must have the mass of an electron, but must be positively charged, antiproton must have the mass of a proton, but be negatively charged.
The antielectron was discovered in 1932 by American physicist Carl David Anderson (born 1905) during his research into cosmic rays. When cosmic rays collide with atomic nuclei in the atmosphere, they create particles that are deflected in the magnetic field at the same angle as electrons, but in the opposite direction. Anderson called particles of this kind positrons .
The antiproton could not be discovered for another quarter of a century. Since the mass of the antiproton is 1836 times greater than the mass of the antielectron, the formation of an antiproton requires 1836 times more energy, and therefore, until the 50s of the 20th century. this transformation was impossible. In 1955, American physicists Emilio Segre (born in 1905) and Owen Chamberlain (born in 1920) managed to obtain and detect an antiproton using powerful accelerators.
It was found that there may be such peculiar atoms in which negatively charged nuclei containing antiprotons are surrounded by positively charged positrons. Naturally, what is antimatter cannot exist for a long time either on Earth, or, probably, even within our Galaxy, since when matter comes into contact with antimatter, they annihilate (destroy), releasing a huge amount of energy. And yet, astronomers wonder whether galaxies built from antimatter could exist? If this is possible, then it will be very difficult to detect such Galaxies.
PrefaceFrom the author's prefaces to the first and second editions
Chapter 1. Introduction
Chapter 2. Thermodynamics of reactions in solutions
2.1. Chemical Potential
2.2. Change in Gibbs energy during a chemical reaction
2.3. Equilibrium condition for a chemical reaction
2.4. Approximation of a dilute solution
2.5. Standard potential
2.6. Equilibrium law in the dilute solution approximation
2.7. Activity factor
2.8. Exact equilibrium law
2.9. Sign of the derivative ((?)/ds)
2.10. Activity and standard activity
2.11. Equilibrium constants in various media
2.12. Number and selection of components
2.13. Formal and true values of chemical potential
2.14. Function q°
2.15. Gibbs-Duhem equation
2.16. Molecular weights in solution
2.17. Gibbs-Duhem equation for multicomponent systems
2.18. Electrolyte solutions
2.19. Temperature coefficient of chemical potential
2.20. Molar enthalpy and molar entropy as functions of concentration
2.21. Dependence of the equilibrium constant on temperature
2.22. Experimental determination of enthalpy and standard entropy of a reaction
2.23. The problem of the accuracy of determining (?) and (?)
2.24. Formal values of standard enthalpy, entropy and heat capacity
2.25. Effect of pressure
Literature
Chapter 3. Elements of statistical thermodynamics
3.1. The Need for Statistical Thermodynamics
3.2. Statistical thermodynamics of an ideal gas
3.3. The nature of statistical sums
3.4. Thermodynamic functions of a simple harmonic oscillator
3.5. Typical numerical examples
3.6. Calculation of statistical sums
3.7. Statistical mechanics and chemical equilibrium in the gas phase
3.8. Enthalpy change and potential energy change
3.9. Statistical mechanics of dilute solution
3.10. Barclay-Butler rule
3.11. Some features of aqueous solutions
3.12. Changes in standard entropy for reactions in solutions
3.13. Relationship to structure and reactivity
3.14. Statistical sum q°
Literature
Chapter 4. Interpretation of kinetic data
4.1. Law of Mass Action
4.2. Systematics and terminology
4.3. Relationship between rate constants and equilibriums
4.4. Experimental study of kinetics
4.5. Linear forms of equations for kinetically simple irreversible reactions
4.6. Accuracy issues
4.7. Impact of systematic deviations from linearity
4.8. Reliability of reaction order determination
4.9. Stirred flow reactors
4.10. Kinetically simple reversible reactions
4.11. Kinetically complex reactions
4.12. Experimental recognition of systems involving only first-order reactions
4.13. Integral forms of equations for first order reaction systems
4.14. Interpretation of kinetic equations containing two exponential terms
4.15. Bodenstein method (method of stationary concentrations)
4.16. The problem of accuracy of equations with a large number of parameters
4.17. Michaelis-Menten equation
4.18. Systems with pre-equilibrium
4.19. The influence of associative equilibrium
4.20. Recognizing Response Difficulty
4.21. Using a stirred flow reactor to study multi-step reactions
4.22. Flow tube reactor
4.23. Study of relative reactivity by the method of competing reactions
4.24. Very fast diffusion-limited reactions
4.25. Validity of the law of mass action for reactions in solutions
4.26. The principle of independent reactions
4.27. Catalytic and zero-order reactions
4.28. "Dark Times" and "Rebirth"
Literature
Chapter 5. Transition State Theory
5.1. General theory of transition state
5.2. Change in potential energy of a system during a reaction
5.3. Reaction coordinate
5.4. Lethal transition state theory
5.5. Theoretical foundations of the detailed transition state theory
5.6. Validity of a detailed transition state theory
5.7. Extension of the theory to other types of reactions
5.8. Extension of the theory to nonideal solutions
5.9. Transition state composition
5.10. Limiting transition state or limiting transition states
5.11. Parallel transition states
5.12. Is the formation of intermediate compounds necessary?
5.13. Progression from reactants to transition state
5.14. Questions about the mechanism that make and don't make sense
5.15. Progression from transition state to reaction products
5.16. Curtich's principle
5.17. Reaction of stilbene with bromine in methanol
5.18. The problem of the mechanism of reactions involving ions
5.19. Structure effects in systems with moving equilibrium
5.20. The problem of stoichiometric inclusion of solvents in the transition state
5.21. Entropy of activation
5.22 Structural isotope effects
5.23. Isotopic effects of solvents
5.24. Effect of pressure on reaction rates
5.25. Systems with mobile equilibria in initial and final substances
5.26. Catalysis
5.27. Progression from the transition state to the reaction products in the case of a catalytic process
5.28. Enzymatic and heterogeneous catalysis
5.29. Microscopic reversibility or detailed equilibrium
Literature
Chapter 6. Some substitution reactions
6.1. Second-order nucleophilic substitution at a saturated carbon atom
6.2. Racemization and isotope exchange
6.3. Solvolysis reactions
6.4. Carbonium ion mechanism
6.5. Salt effects in highly dissociating solvents
6.6. Role of the solvent
6.7. Terminology issues
6.8. Limited ideas about the formation of free carbonium ions
6.9. Ion pair hypothesis
6.10. Effect of Azide Additives
6.11. Participation of hydroxyl-free solvents in the formation of a transition complex
6.12. Allyl rearrangement
6.13. Solvolysis accompanied by racemization of the reagent
6.14. Additional Data on the Effects of Azide Supplementation
6.15. Special salt effect
6.16. Double Configuration Reversal
6.17. Bromonium ions as intermediates
6.18. Phenonium ion problem
6.19. Regroupings
6.20. Electrophilic aromatic substitution
6.21. Azo coupling reaction
6.22. Possible forms of reagents
6.23. Transition state of azo coupling reaction
6.24. Catalytic action of pyridine
6.25. Isotopic effects of hydrogen in the azo coupling reaction
6.26. Spatial influences in azo coupling reactions
6.27. Entropy of azocoupling
Literature
Chapter 7. Salt effects
7.1. Brønsted equation
7.2. Debye-Hückel limit law
7.3. Salt effects in azo coupling reaction
7.4. Applicability of the limit law to asymmetric ions
7.5. Applicability of the limit law to multiply charged ions
7.6. Effect of moderate salt concentrations on activity coefficients
7.7. The influence of moderate salt concentrations on the rate of reactions between ions
7.8. Salt effects as confirmation of the transition state theory
7.9. Olson-Simonson rule
7.10. Suppression of salt effects when studying reaction order
7.11. Physical meaning of the Debye-Hückel theory
7.12. Modification of the theory for the case of weak interactions
7.13. Bjerrum's approach
7.14. Ion pairs in chemical reactions
7.15. Ion pair formation and the Olson-Simonson rule
7.16. About the concept of ion pairs
7.17. On weak association of ions
7.18. Effect of salts and non-electrolytes in aqueous solutions
7.19. The influence of salts on reactions between ions and neutral molecules in aqueous solutions
7.20. The effect of salts on the reactions between ions and neutral molecules in liquid sulfur dioxide
7.21. Effect of salts on solvolysis reactions in aqueous solutions
7.22. The effect of salts on non-electrolytes in aqueous solutions
7.23. Historical information
Literature
Chapter 8. Effect of solvent and reactivity
8.1. The basic principle
8.2. Electrostatic contribution to the chemical potential of the electrolyte
8.3. Formation of ion pairs; case of Ideal electrostatic interaction
8.4. Association of ions in quaternary ammonium salts
8.5. Ionic triplets and salt polymers
8.6. Hydrogen bond
8.7. Hydrogen bonded ion pairs
8.8. Interaction of anions with hydrogen bond donors
8.9. Interaction of cations with hydrogen bond acceptors
8.10. Other types of interactions of ions with neutral molecules
8.11. Overview of Ion-Solvent Interactions
8.12. The role of dielectric constant
8.13. Organometallic compounds of elements of the first group
8.14. Effect of anion solvation on the rate of interaction of an anion with a neutral molecule
8.15. The influence of specific solvation of cations on the rates of reactions between anions and neutral molecules
8.16. Chelate solvation of cations
8.17. Kinetics of reactions involving ionic aggregates
8.18. Inertness of solid reagents
8.19. Effect of solvent on the reactivity of ambident anions
8.20. Influence of the nature of the solvent on the reactions of cations with anions
8.21. Dispersion forces
8.22. Cohesive Energy Density
8.23. Regular and ideal solutions
8.24. Transfer of a solute from one regular solution to another
8.25. Impact of strong interactions
8.26. Isokinetic dependence upon changing solvent
8.27. Solvolysis in water-alcohol mixtures
8.28. Other influences of the solvent on the rate of solvolysis
8.29. Charge transfer complexes
8.30. Complexation due to multipole or dipole interaction
Literature
Chapter 9. Quantitative study of acids and bases
9.1. Electrometric determination of pH in aqueous solutions
9.2. Electrometric determination of acidity in non-aqueous solutions
9.3. Measuring acidity using indicators
9.4. Acidity function Ho
9.5. Historical information
9.6. Acidity functions H"", Hi and Na
9.7. Acidity functions Hk and H"k
9.8. Other data on acidity functions
9.9. Three Variable Hypothesis
9.10. Solvation variable
9.11. The uncertainty of the concept of acidity
9.12. Standard activities in mixtures of strong acids with water
9.13. About the influence of water activity
9.14. Acidity functions for highly concentrated acid solutions
9.15. Functions of acidity and rate of acid-catalyzed reactions
9.16. Nitration of aromatic compounds in mixtures of sulfuric acid and water
9.17. Systems described by two acidity functions
9.18 Uncertainty of the concept of foundation force
9.19 Determination of pK°° values
9.20. Superposition of environmental effects in spectrophotometric determination of basicity
9.21. Methods for correcting for environmental effects
9.22. Other ways to determine the strength of weak bases
9.23. Temperature coefficients
9.24. Acid-base reactions in poorly dissociating solvents
9.25. Behavior of indicators in nitromethane and sulfolane
9.26. Acid-base reactions in dimethyl sulfoxide
9.27. Functions of acidity and base strength in other highly basic environments
9.28 Reaction rates in strongly basic media
9.29. Basicity of alkali metal alcoholates
Literature
Chapter 10. Reaction rates involving acids and bases
10.1. General acid and general base catalysis
10.2. Bronsted catalytic equation
10.3. Statistical adjustments
10.4. Spatial effects
10.5. Bronsted catalytic equation and transition state theory
10.6. Catalysis by solvent molecules, lyonium and lyate ions
10.7. Distinguishing between general acid and general base catalysis
10.8. Proton transfer reaction rates
10.9. Theory of electron displacements
10.10. Acetaldehyde hydration
10.11. Mechanism of base-catalyzed hydration
10.12. Mechanism of acid-catalyzed hydration
10.13. Catalysis in the enolisacin reaction
10.14. Cyclic transition states
10.15. Hydrolysis of esters
10.16. Koiids of the "bell" and "inverted bell" type
10.17. Formation of oximes and other similar reactions
10.18. Hydrolysis of phenyliminolactone
10.19. Lewis acid catalysis
Literature
Chapter 11. Quantitative relationships between structure and reactivity
11.1. Abstract theory
11.2. The principle of linear dependence of free energies
11.3. Relationship between rate constants and equilibriums
11.4. Linear relationships between changes in free energies
11.5. Gettler correlations
11.6. Strength of acids and rate of hydrolysis of their esters
11.7. Phenodes and carboxylic acids
11.8. Electronic interpretation
11.9. Hammett's equation
11.10. Spatial effects
11.11. Direct resonant interactions
11.12. Normal substituent constants
11.13. (?)Constants
11.14. (?)Constants
11.15. Insufficient two values (?)
11.16. Patterns of direct resonant interaction
11.17. Yukawa-Tsuno equation
11.18. The problem of random and systematic deviations
11.19 Attenuation of electronic influences in the side chain
11.20. Cumulative influence of several substituents
11.21. Spatial obstacles to resonance
11.22. The nature of spatial effects
11.23. Additivity of spatial effects
11.24. Further restrictions on the applicability of the principle of linear dependence of free energies
11.25. Induction and resonance effects
11.26. Taft's equation
11.27. Historical information
11.28. Correlation of reactivity of aliphatic and aromatic compounds
11.29. Exner analysis
11.30. "Flexible" and "rigid" molecules
11.31. Effect of solvents on (?) and (?)
Literature
Chapter 12. The influence of structure on changes in enthalpy and entropy
12.1. Isokinetic dependence
12.2. Isoentropic reaction series
12.3. Isenthalpic reaction series
12.4. Isokinetic temperature
12.5. Reliability of establishing isokinetic dependence
12.6. Limits of the existence of isokinetic dependence
12.7. Inconsistency of enthalpy and entropy influences
12.8. Structural influences and statistical thermodynamics
12.9. Assumptions
12.10. Consequences
12.11. Confirmations
12.12. Other approaches
Literature