- Introduction. Subject and content of the course: physical organic chemistry. Methods of physical organic chemistry. The role of diffusion and collision frequency in the kinetics of chemical reactions. Quantitative assessment of the role of diffusion in the kinetics of chemical reactions. A chemical reaction is the transformation of a diffusion pair of reacting particles. Equilibrium concentration of diffusion pairs. The number of collisions between particles included in a diffusion pair and the reaction rate constant.
- Interaction between particles in solution. Electrostatic interactions. Pair interactions. Interaction of two ions. Interaction of an ion with a dipole. Orientation interaction. Interaction of an ion and a permanent dipole with an induced dipole. Electrostatic interactions of particles with solvent. Free energy of an ion in an electrolyte medium. Dispersion interactions. Effect of solvent on pairwise dispersion interactions. Energy of dispersion interaction of a dissolved particle with a medium. Hydrogen bond.
- Solvation of ions and molecules in solutions. Theoretical calculations of ion solvation energy. Born equation. Further development of Born's theory. Solvation in non-aqueous solutions. Features of association in non-aqueous solutions. Consideration of the association process and the formation of ion pairs from the perspective of electrostatic theory. Solvation and coordination with solvent.
- Transition state theory and its application to liquid-phase reactions. Basic concepts of the theory of absolute reaction rates. Energy surface of a chemical reaction. Derivation of the basic equation of the activated complex theory. Application of transition state theory equations to reactions in the liquid phase.
- Solvation of the activated complex. Evaluation of the thermodynamic characteristics of solvation of the activated complex based on experimental data. Activation process in the liquid phase and preliminary reorganization of the solvation shell. The theory of the activated complex and taking into account the formation of diffusion pairs.
- Quantitative consideration of the influence of the medium on the rate of liquid-phase reaction. The simplest electrostatic models of the reaction of two ions. Primary salt effect.
- Electrostatic models that take into account the nature of charge distribution in reacting particles. Kirkwood model. Laidler and Landskroner model. Hiromi method. Application of electrostatic theories when considering kinetic and activation parameters of reactions.
- Correlation relationships and the influence of the environment on the kinetics of reactions. Quantitative accounting of environmental influences using semi-empirical correlation equations. Winstein–Grundald equation. Empirical solvent polarities Z and E T . The influence of the solvent on the dependence of the reactivity of compounds on their structure. Hammett's equation. The principle of linearity of free energies.
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) in 1913 was able to show that the atomic mass of lead obtained as a result of 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 (born in 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 was left 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.
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. |
This is the first time a book of this kind has been published in Russian; it is intended for organic chemists - researchers, teachers, graduate students and senior students of chemical universities.
PREFACE
The author of this book, Professor L. Hammett, was born in 1894 in Wilmington (USA). He graduated from Harvard University, where he received a Bachelor of Science degree in 1916 and then worked for a year in Zurich with Staudinger. In 1923, he defended his dissertation for the degree of Doctor of Philosophy at Columbia University. Until 1961 he taught at this university, from 1951 to 1957 heading the department of chemistry. Currently, L. Hammett is a retired professor emeritus.
L. Hammett was one of the pioneers of a new branch of science that arose in the 20s and 30s of our century - physical organic chemistry. Three fundamental discoveries are associated with his name: the creation of the acidity function, the establishment of a connection between the rate of reactions catalyzed by acids and the acidity function, as well as the introduction into chemistry of correlation equations like \gk - and thereby the principle of linearity of free energies. It is now clear that even one of these discoveries would be enough to leave its mark on science. Naturally, Professor Hammett was awarded numerous scientific prizes and medals: Nichols (1957), Norris (1960, 1966), Priestley (1961), Gibbs (1961), Lewis (1967), Chandler (1968), National Science Medal (1968) . He is a member of the National Academy of Sciences (USA) and an honorary member of the Chemical Society (London).
The book “Fundamentals of Physical Organic Chemistry” by L. Hammett, which is brought to the attention of readers, occupies an outstanding place in the world chemical literature. Its first edition, published in 1940, was far ahead of its time and, to use Hammett’s own words about another book, became “the bible of thinking organic chemists.” The reason for this lies, firstly, in the depth of many of the original ideas presented in it, which determined entire areas of scientific research for decades. Secondly, the book contained a large number of logical predictions, which were later brilliantly confirmed. Thirdly, for many years this book was the only one in world literature that outlined the problems of a new branch of science - physical organic chemistry. The first edition of the book was not translated into Russian, but is probably known to almost all Soviet chemists from numerous references to it, and to many from the English original. This translation is made from the second, completely revised and expanded edition, published in the USA in 1970.
Unlike the first edition, this book is not a monograph, since the author did not strive for either an exhaustive coverage of all world literature or coverage of all issues. The book is a beautifully written textbook designed to provide a thorough understanding of the most important issues in physical organic chemistry. There is no need to write about the principles of the book’s construction and its specific content, since they are fully reflected in the author’s preface and detailed table of contents. It should only be noted that it differs significantly from existing books of the same purpose in its more general and more rigorous “physical” approach to problems and their quantitative interpretation.
The reader of this book has a rare opportunity to obtain information, so to speak, first-hand: L. Hammett was a contemporary and an active participant in the development of physical organic chemistry as a science. The “lyrical digressions” in the book, permeated with good humor, allow you to feel the atmosphere in which this or that discovery was made, and to feel the originality of the author as a scientist and a person. At the same time, Soviet chemists cannot help but feel sympathy for L. Hammett, who, in a dispute with his invisible opponents, claims that the goal of theoretical chemistry is the ability to foresee and control chemical processes, or, in another place, expresses the conviction that especially valuable results are obtained not those researchers who are engaged in clarifying what is already known, but those who enter into dispute with established views, even if they seem to be immutable laws. However, it is noted jokingly. Hammet, at the same time, you should clearly understand what you can argue about and what you cannot argue about, so as not to waste your time and other people’s funds.
When working on the Russian edition of this book, we tried to convey its contents as accurately as possible, made notes where necessary and compiled small lists of additional literature.
There is no doubt that L. Hammet's excellent book will deservedly enjoy wide popularity in the circles of Soviet chemists, employees of universities, research institutes and enterprises.
In conclusion, here is an excerpt from L. Hammett’s letter to the editor: “I am glad that my book is published in Russian, one of the main languages of the world. I strongly recommend that young chemists, including my own grandson, study Russian in order to read Russian scientific literature. Although my knowledge of the Russian language, unfortunately, does not extend beyond the alphabet, it can be assumed that the translation of my book into Russian will serve as an incentive for me to further study it.”
L. Efros Y. Kaminsky
FROM THE AUTHOR'S PREFACE TO THE FIRST AND SECOND EDITIONS
One of the general trends in the development of science is a temporary weakening of attention to phenomena located at the intersection of various fields of science. Sooner or later this shortcoming becomes too obvious, and then a new branch of science appears. Something similar happened in the twenties and thirties at the intersection of physical and organic chemistry: a set of facts, generalizations and theories appeared, which it would be correct to call physical organic chemistry. This name implies the study of organic chemistry phenomena using quantitative and mathematical methods.
One of the main directions in which physical organic chemistry developed was the study by quantitative methods of reaction mechanisms, as well as the influence of structure and environment on reactivity. No other direction produces results that have such direct practical value for the main task of chemistry - the control of chemical processes.
Sometimes fellow physicists mockingly called this kind of work “the study of soap making.” But soap plays an important role in the civilization of mankind, and I am not at all sure that we know more about the basics of soap making, which, as they say, is the hydrolysis of esters, than about the structure of the atomic nucleus.
No matter how rapid the development of physical organic chemistry has been in the thirty years since the publication of the first edition of this book, the current situation is still very far from the gloomy future depicted by Kachalsky: “Whether we like it or not, the ultimate goal of every science is to to become trivial, to become a well-regulated apparatus for solving textbook exercises or for practical application in the construction of machines.” This goal is still a long way off if, as was the case a few years ago, we are able to be surprised by the discovery that reactions involving bases can proceed 1013 times faster in dimethyl sulfexide than in methanol. And the time is still far off when predicting the catalyst for a given reaction will become an exercise for students.
Yet much of what was conjectural thirty years ago has become certain; rough approximations gave way to more accurate ones; Physical organic chemistry itself has grown significantly, as has the volume of knowledge of researchers working in this field. Apparently, the time has come for a radical revision of the topics covered in the first edition of the book.
Like the first edition, the second edition of this book deals with rates, equilibria, and reaction mechanisms. To be more precise, the consideration will be limited to the range of heterolytic reactions in solutions. Radical reactions, as well as the theory of molecular orbitals, are not considered, since they have become the subjects of separate monographs.
Even with these limitations, the size of the book would exceed all reasonable limits if I had tried, as I did in the first edition, to make it encyclopedic and discuss all reactions for which mechanistic information is available, instead of selecting examples to illustrate ways of investigating the mechanism .
However, I hope that the basic principles that can be used in the study of heterolytic reactions in solutions are presented quite thoroughly, deeply and thoughtfully. I also hope that this book will be of benefit to both physical chemists and organic chemists; Therefore, I have strived to make the presentation of the material accessible to students with a solid (albeit elementary) knowledge of both physical and organic chemistry. I can only apologize if there are passages in the book that seem trivial to persons already working in any of these areas.
The theories and principles discussed in the book are my own only to a very small extent. To justify the errors that undoubtedly exist in the indications of the true authors, I can only refer to the difficulty of determining priority in ideas.
I am deeply indebted to E. M. Arnett, M. M. Davis, G. L. Hering, D. E. Kimball, R. W. Taft, and G. Zollinger, each of whom read portions of the manuscript and made valuable comments. I thank my graduate students who taught me more than I taught them, and most of all, three great teachers: E. P. Kohler, G. Staudinger, and J. M. Nelson, for they sowed the seeds from which this book grew.
L. Hammett
The main task of a chemist, as I imagine it, is the ability to anticipate and control the course of reactions. In this case, as with any other attempt by man to master the laws of nature, two approaches can be used. One is to create general theories from which consequences concerning the particular properties of matter are deduced. The second, relying on empirical generalizations, builds particular and approximate theories that can explain observed phenomena or suggest an interesting direction for experimental research. Because of the nature of our science, we chemists are forced to follow mainly the second path. As I once noted, “chemists had arrived at effective working principles long before Schrödinger’s equation became the embodiment of the theoretical key to all problems of chemistry. Even today, the amount of information a chemist can obtain directly from this equation represents only a small part of what he knows."
Some chemists seem ashamed of this and envy scientists working in fields where, in Dirac's words, “the beauty of equations is more important than their agreement with experiments.” I feel a sense of pride in science, which has achieved so much through the inventive use of every means, be it rough and clumsy or refined and elegant. I quote myself again: “I hope that nothing I have said will lead you to think that I am neglecting the theory or belittling its importance. But I think that the respect with which we treat theory should not obscure, as sometimes happens, the fact that science is equally indebted to empirical generalizations. Let us recall, for example, what enormous theoretical consequences resulted from the discovery by a Swiss schoolteacher of a quantitative relationship between the frequencies of lines in the spectrum of hydrogen - a connection that looked strange and unexpected.
I think that sometimes we forget the big difference between exact and approximate theories. Let me dwell on our attitude towards the latter. If, for example, my colleague Breslow, on the basis of molecular orbital theory, predicts the stability and aromaticity of such an exotic structure as the cyclopropenyl cation, then it is obvious that, despite its approximate nature, molecular orbital theory is a powerful tool for the discovery of unexpected phenomena. But if molecular orbital theory predicts the impossibility of some new phenomenon or relationship, its conclusions should be considered only somewhat (but not completely) discouraging. If someone starts looking for the effect predicted by this kind of theory, then most likely time and money will not be wasted. But if someone starts looking for an effect that such a theory considers impossible, then there will be little chance of a favorable outcome. Fortunately, among scientists, as well as among betting players, there are people who prefer to bet against the odds, while the vast majority always bet on the favorite. I think that in science we should strongly support people who decide to take risks in the face of such unequal odds.
This does not mean that we should support fools and ignoramuses who ignore the facts that kill them, people who want to spend their time and, as a rule, other people's money in search of an effect that is not consistent, for example, with the conclusions made by Willard Gibbs. Gibbs proceeded from strictly proven generalizations - the first and second laws of thermodynamics,
used precise mathematical apparatus and created a theory that is the best example of an exact theory known to me. It is futile to fight against such a theory.
However, the general laws discovered by Gibbs are abstract, and their transformation into concrete relationships, including such prosaic quantities as the concentration of dissolved substances, requires either exact empirical equations of state or theories that are inevitably approximate. Therefore, caution must be exercised, especially in cases where the prestige of a famous scientist is mixed with the approximate nature of the theory. Thus, in Nernst’s old age there was a period when a rare daredevil dared to publish any conclusions that were not consistent with the approximate partial equations of state, which allowed Nernst to accomplish many useful things during his youth. Those who tried faced the wrath of Jupiter, who usually destroyed the offender (until the time came for G. N. Lewis).
The moral of all this is: have great respect for accurate theory, but be really sure that a theory that tells you not to do what you would like to do is actually an accurate theory, and not just the favorite creation of an established authority.”
KOHETS FRAGMEHTA