CURRENT STATE OF THE THEORY OF THE MUTUAL INFLUENCE OF ATOMS IN MOLECULES

B. V. TRONOV

The theory of the mutual influence of atoms as a component part of the theory of chemical structure

The great Russian organic chemist A. M. Butlerov, the founder of the theory of chemical structure, already in his first work [5, 68—81] gives a definition of chemical structure and goes on to say:

“The chemical nature of a complex particle is determined by the nature of its elementary constituent parts, by their number, and by their chemical structure.” Detailing this proposition, he raises the question that the properties of each atom in a molecule are determined not only by the nature of this atom and of the atoms directly connected with it, but also depend on more distant atoms:

“… the further development of the view expressed here will show how inadequate the stated rule is, and what mutual influence two atoms, located within one and the same chemical particle but not acting chemically directly upon one another, may exert.”

These ideas were expressed by Butlerov as early as 1861 (a report at the congress of German natural scientists and physicians in Speyer and an article in Zeitschrift für Chemie und Pharmacie) and in 1862 (the article “On the Chemical Structure of Substances,” published in the Scientific Notes of Kazan University).

All of A. M. Butlerov’s subsequent scientific activity was devoted chiefly to strengthening and developing the theory of chemical structure, and in many of his works we find new indications of the mutual influence of all atoms in a molecule, an explanation of this idea with concrete examples, and certain rules concerning just how atoms influence one another in particular organic compounds. In the article “On Tertiary Isobutyl Alcohol (trimethylated methyl alcohol)” [5, 119—137] there is an indication of the difference in the properties of hydrogen bound to oxygen or to carbon, as well as in the properties of hydroxyl in ethyl alcohol and acetic acid and of chlorine in ethyl chloride and acetyl chloride. In the large article “The Present Significance of the Theory of Chemical Structure” [5, 410—434] the doctrine of the mutual influence of atoms is developed in greatest detail. A. M. Butlerov points out that the hydrogen atoms in water, hydrogen chloride, and methane:

$$ \mathrm{H_2O,\ HCl,\ CH_4} $$

are sharply different in their chemical character. In the methane molecule the hydrogens are completely incapable of direct replacement by any metal; the hydrogen of water is replaced under the action of sodium; for the liberation of hydrogen from hydrogen chloride, less energetic metals, for example zinc or iron, are sufficient. On the other hand, in methane the hydrogen atoms are rather easily replaced under the action of chlorine.

These examples illustrate the dependence of the properties of hydrogen on the atom with which it is directly bonded. Then, analyzing the structure of the molecule of methylene chloride, $\mathrm{CH_2Cl_2}$, Butlerov notes that here too, although the hydrogen and chlorine atoms are not directly bonded to one another, there undoubtedly exists a dependence between them—a dependence of another category, which also manifests itself quite clearly in chemical properties. Thus, the compound of carbon with chlorine alone, $\mathrm{CCl_4}$, does not react with solutions of silver salts; but if part of the affinity of the carbon atom is saturated by oxygen, then a substance is obtained that reacts vigorously and rapidly with solutions of silver salts.

Butlerov draws the conclusion: “Thus, it is wholly unfounded to reproach the doctrine of chemical structure with the fact that, in considering the order of bonding between the elementary atoms in a particle, it leads to an ignoring of the mutual influence of atoms which, as is commonly expressed, are not directly connected with one another.”

Thus, the founder of the theory of chemical structure himself viewed his theory more broadly than did certain Western European scholars. He did not believe that the theory of chemical structure had a purely formal character, but saw in structural formulas a deep inner meaning; he tried to understand the influence on the properties of each atom of the composition and structure of the entire molecule as a whole.

In the work “On the derivatives of trimethylcarbinol” [5, 180—207], dating to 1867, Butlerov arrives—although on only one example—at an important particular regularity concerning the mutual influence of atoms in molecules. He noted that “in the chlorination of trimethylformene (i.e., trimethylmethane), the action of chlorine is undergone by that hydrogen which is connected with the carbon bonded to the other three carbon units, i.e., the hydrogen bonded to the least hydrogenized carbon”:

\[ (\mathrm{CH_3})_3\mathrm{CH} + \mathrm{Cl_2} \rightarrow (\mathrm{CH_3})_3\mathrm{CCl} + \mathrm{HCl}. \]

The doctrine of the mutual influence of atoms received particular development in the works of V. V. Markovnikov, the closest pupil of A. M. Butlerov. In his doctoral dissertation “Materials on the Question of the Mutual Influence of Atoms in Chemical Compounds” [10], Markovnikov wrote:

“The question of the influence of elementary atoms on the direction of the chemical reactions of a complex body belongs among the most vital questions of contemporary chemistry and, as has already been noted, is attracting more and more of the attention of chemists. It naturally had to arise as soon as the doctrine of chemical structure was assimilated by the majority, and it is the direct continuation of that doctrine.”

There follows this characterization of the interrelation of these theories:

“Defining their mutual relations in a few words, one may say that the theory of structure is the external mechanism, the action of which is directed and regulated by the theory of the mutual influence of atoms. This is that inner cause which gives meaning to all manifestations of the mechanism.”

Defining the concrete tasks of chemists in the development of the doctrine of the mutual influence of atoms, Markovnikov pointed out that the first stage here should be the study of the most striking manifestation of mutual influence in examples of the most characteristic reactions.

In the above-mentioned work, Markovnikov himself considered the reactions of substitution of hydrogen in saturated hydrocarbons and oxygen-containing compounds, reactions of addition to unsaturated hydrocarbons, and the formation of ethylene hydrocarbons by the elimination of certain atoms from complex molecules.

Concerning the substitution of hydrogen by halogens, V. V. Markovnikov gives the following rules:

“In unsaturated hydrocarbons the hydrogens belonging to the carbon atoms that have expended most of their affinity in bonding with other carbon atoms are more readily replaced than the others.” Further, on the influence of oxygen: “The hydrogen pairs that are under the nearest influence of oxygen are always subject to replacement.”

Everyone knows Markovnikov’s rule, which speaks of the direction of addition reactions in unsaturated hydrocarbons: hydrogen attaches to the more hydrogenated carbon atom, while the halogen, acid residue, or hydroxyl attaches to the less hydrogenated one. As for elimination reactions, Markovnikov showed first of all that the elimination of atoms always proceeds from different carbon atoms, and not from one, as some scientists had thought. This served as an important argument in favor of the doctrine of double and triple bonds—a doctrine finally introduced into science by A. M. Butlerov in 1870 [5, 231—253] on the basis of the results of the study of reactions of isocrotyl bromide

\[ (\mathrm{CH}_3)_2\mathrm{C}=\mathrm{CHBr}. \]

Relevant to the problem of the mutual influence of atoms in this part of Markovnikov’s work is the proposition, which turned out to be incorrect, that hydrogen leaves the most hydrogenated carbon atom; for example, 2-iodobutane, on elimination of hydrogen iodide, would supposedly give butene-1:

\[ \mathrm{CH}_3-\mathrm{CH}_2-\mathrm{CHI}-\mathrm{CH}_3 \to \mathrm{CH}_3-\mathrm{CH}_2-\mathrm{CH}=\mathrm{CH}_2. \]

This proposition was derived from erroneous literature data. A. M. Zaitsev [6] proved that in reality hydrogen is split off from the less hydrogenated link of the carbon chain, i.e., elimination reactions proceed as should have been expected in view of the regularity of substitution processes formulated by Markovnikov himself. The fewer hydrogens are bound to a given carbon atom, the more readily these hydrogens leave.

Russian chemists of the late nineteenth century, pupils and followers of A. M. Butlerov, derived a number of other rules concerning the direction of reactions in organic chemistry.

An outstanding achievement of Western European scientists at the end of the nineteenth century in the study of the problem of the mutual influence of atoms was the substitution rules concerning reactions of aromatic compounds. It was established that the direction of substitution reactions in homologues and derivatives of benzene depends in a definite way on the nature of the first substituent already present on the benzene ring. Some substituents, the so-called substituents of the first kind—for example, halogens, hydrocarbon radicals, and the groups OH, NH₂—orient chiefly to the ortho and para positions; others—substituents of the second kind, for example the groups NO₂, SO₂OH, COOH, CN—direct to the meta position. This regularity was noted by Hübner in 1875 and developed in greater detail by Nölting in 1876.

In general, by the end of the nineteenth century organic chemistry had already accumulated a considerable amount of material on the question of the mutual influence of atoms. From individual, sometimes large groups of facts, definite rules were derived, which were of great benefit in the practical work of chemists. These rules made it possible to orient oneself better in methods for obtaining organic substances, to predict to some extent in advance the properties of new compounds not yet obtained, and to judge which compounds should prove practically useful.

A shortcoming of the concrete regularities known at that time in this field was that an empirical rule has force only for the range of phenomena from the study of which it was derived. For example, Markovnikov’s rule says that in addition reactions of unsaturated hydrocarbons the hydrogen of the adding reagent attaches to the more

to the hydrogenated carbon atom, while halogen, other acid residues, and hydroxyl attach to the less hydrogenated one. But this rule generalized the results of experiments carried out by no means with all unsaturated compounds and only under definite conditions. How a reaction would proceed, for example, with an unsaturated carboxylic acid or with unsaturated hydrocarbons, but under other conditions, and whether the conditions could be changed so that it would proceed in another direction—about this, guided only by Markovnikov’s rule, we could say nothing. Subsequently it became known, again from direct experiments, that in unsaturated acids addition does not follow Markovnikov’s rule:

\[ \mathrm{CH_2{=}CH{-}COOH + HBr \longrightarrow CH_2Br{-}CH_2{-}COOH.} \]

It turned out that unsaturated hydrocarbons, in the presence of peroxides, also react with acids in the other direction:

\[ \mathrm{CH_3{-}CH{=}CH_2 + HBr\ (+\text{peroxide}) \longrightarrow CH_3{-}CH_2{-}CH_2Br.} \]

In the chemistry of aromatic compounds, the substitution rules stated that, when chlorobenzene is chlorinated under the usual conditions for carrying out this reaction, chiefly ortho- and para-dichlorobenzenes are formed; but whether meta-dichlorobenzene could be prepared in some way, and how the chlorination conditions would have to be changed for this, again nothing could be said. Further experiments showed that the formation of the meta derivative here is favored by high temperature.

The predictive possibilities of the theory of chemical structure, together with the empirical rules adjoining it, were quite considerable; but one could hope to increase these possibilities still further if a theoretical foundation were given to the empirically found rules. To provide a theoretical foundation for the rules determining the direction of reactions, speaking of the dependence of the course of reactions on the composition and structure of the reacting substances and on various external factors, it was necessary to elucidate the origin of the force of chemical affinity, to find the connection of this force with other forces of nature. And this could be done only after the structure of atoms had been clarified.

Electronic and Wave Theory of the Structure of Atoms and Molecules

Development of Electronic Ideas about the Structure of Atoms and Molecules

Already by the beginning of the nineteenth century, a connection between chemical phenomena and electricity had been established. In 1800 Nicholson and Carlisle decomposed water by the action of an electric current. Soon the electrolysis of alkalis and salts was carried out. These facts led to the appearance of the electrochemical theories of Davy, Berzelius, and M. G. Pavlov, which stated that chemical bonds arise because of the attraction of unlike electric charges. Then the idea of the electrical origin of chemical forces lay dormant for a long time and again began to develop only at the end of the nineteenth and the beginning of the twentieth century, on the basis of new discoveries in physics and chemistry.

It was proved that atoms of chemical elements do indeed consist of particles bearing positive and negative electric charges. Particles of negative electricity—electrons—were found and studied after being separated from chemical atoms. As for the positive charge, at the very beginning of the twentieth century J. J. Thomson proposed that the entire space of the atom (linear dimensions of the order of \(10^{-8}\) cm) is occupied by a positive electric field; in this positive field are located electrons very small in volume (linear dimensions of the order of \(10^{-13}\) cm).

In 1911 Rutherford showed that the positive charge and almost the entire mass of the atom are concentrated in its central part, in the atomic nucleus, whose linear dimensions are \(10^{-13}\)—\(10^{-11}\) cm. Electrons must move around the nucleus in closed orbits. In 1913 Bohr refined the nuclear model of the atom, assuming that for an electron in an atom only a definite discontinuous series of orbits is possible. This explained, above all, the line spectra of the elements. The chemical properties of the elements also received an explanation, at least in general outline. It was proved that the periodicity in the change of properties with increasing atomic weight is due precisely to the complex structure of the atom’s electron shell. In chemical reactions only the electrons of the outer shell, the so-called valence electrons, directly participate; their number ranges from 1 to 8. After the outer shell has been filled to 8 electrons (an octet), the next electrons form a new shell.

According to Bohr’s first theory, as somewhat modified by Kossel, the electrons of each shell of the atom move in a common circular orbit. Later, elliptical electron orbits were introduced into the atomic model; then it was accepted that the planes of the orbits of individual electrons are situated at certain angles to one another. For example, the four valence electrons of the carbon atom move in separate elliptical orbits whose planes are parallel to the planes of a tetrahedron described around the atomic nucleus.

In atomic physics the proposition was adopted concerning four quantum numbers that determine the motion of an electron; moreover, in an atom there cannot be two electrons for which all four quantum numbers coincide (the Pauli principle). In a visual representation, the first quantum number determines the radius of a circular orbit or the major semi-axis of an elliptical orbit; the second quantum number determines the eccentricity of the orbit; the third number is connected with the angles between the planes of electron orbits; and the fourth expresses the direction of rotation of the electron about its own axis (“spin”).

The first fully developed electronic theory that explained the occurrence of chemical bonds between atoms and the formation of molecules was the ionic-electronic theory, based on Thomson’s model of the atom. This theory was developed in greatest detail and widely applied in chemistry by our scientists A. M. Berkenheim in Moscow [4] and Ya. I. Mikhailenko in Tomsk [11].

According to the ionic-electronic theory, every bond between two atoms has an electrostatic nature and arises owing to the transfer of an electron from one atom to another. The general scheme of bond formation is:

\[ \pm + \pm \longrightarrow + \mp . \]

In abbreviated notation we have, for example, for two hydrogen atoms:
\( \mathrm{H} + \mathrm{H} \longrightarrow \mathrm{H}\cdot\mathrm{H}' \), i.e., the atoms are converted into positive and negative ions, which attract one another by electrostatic force.

For polyvalent elements the ionic-electronic theory admitted the possibility of the appearance of different charges in the parts of an atom corresponding to separate units of affinity. For example, according to this theory methyl chloride has the structure \( \mathrm{CH_3Cl} \), i.e., the carbon atom, in its three parts turned toward the hydrogens, has excess negative charges (owing to the transfer of electrons from the hydrogen atoms); on the fourth side, carbon, on the contrary, has lost its electron, giving it to chlorine as the more electronegative atom.

The ionic-electronic theory explained many facts of organic chemistry well. For example, the regularities formulated by Markovnikov ...

in the course of substitution reactions in saturated hydrocarbons and addition reactions in hydrocarbons of the ethylene series become understandable if one writes:

\[ \mathrm{CH_3 - CH_2 - CH_3 + Cl_2 \longrightarrow CH_3 - CHCl - CH_3 + HCl} \]

\[ \mathrm{CH_3 - CH = CH_2 + HCl \longrightarrow CH_3 - CHCl - CH_3}. \]

However, in more complex cases the conclusions of this theory sometimes contradict the facts. A major fundamental shortcoming of the ionic-electronic theory was that it ascribed an electrostatic origin to all bonds and, moreover, assumed that when a bond is formed between any atoms, including completely identical ones, one atom necessarily gives an electron to the other. Yet experiment gave no indication that, for example, the bond between hydrogen atoms in the molecule \(\mathrm{H_2}\) had an ionic character.

In 1914 Thomson pointed out the possibility of the emergence of a bond of another type—also electrostatic, but neutral. If an atom consists of opposite electric charges, then two atoms can be arranged in such a way that the positive charge of the first is closer to the negative charge of the second, and the negative charge of the first closer to the positive charge of the second.

\[ \pm \mp . \]

Since the force of electrostatic interaction is greater the closer the charged bodies are to one another, in such a system attraction will predominate over repulsion, and the atoms will combine into a molecule.

This theory was developed in detail, particularly in its application to organic chemistry, by Shishkovskii, but it did not become widely accepted in science, since by that time more advanced theories had appeared.

In 1913 Bohr drew attention to the fact that not only electrostatic forces but also electromagnetic forces may participate in the formation of chemical bonds. Suppose that two atoms, for example hydrogen atoms, approach each other in such a way that the electron orbits are parallel and the electrons move along them in the same direction. Then electromagnetic attraction will arise between the atoms, as between parallel conductors with the same direction of current. As a result, the atoms must combine into a single system in which the positive nuclei are at some distance from each other, while the electrons move along a common orbit whose plane is located between the atomic nuclei and is perpendicular to the straight line joining the nuclei. Thus was constructed the first Bohr model of the hydrogen molecule:

The octet theory of Lewis and Langmuir became very widespread among chemists; according to it, the electrons in atoms are for the most part located at the corners of cubes. In the atoms of inert gases, except helium, the outer shell contains a complete octet of electrons, i.e., all corners of the cube are occupied. Hence their inability to react. The atom of an alkali metal has ...

in its outer shell one electron and very readily gives it up. A halogen atom has seven valence electrons, i.e., it lacks one electron for the formation of a stable cubic octet system. Such atoms should therefore react readily with one another.

Let us write the reaction between sodium and fluorine in the abbreviated electronic formulas now commonly used in electronic chemistry. The sign of a chemical element denotes not the atom, but the positive part of it, which remains if all the valence electrons are removed from the atom. Electrons are denoted by dots. Sodium and fluorine react as follows:

\[ Na\cdot + \cdot \ddot{F} : \longrightarrow Na^{+} + :\ddot{F}:^{-} \]

Thus an ionic, or electrovalent, bond arises. But atoms can also combine in another way, remaining neutral. For example, when two fluorine atoms combine, each of them completes its valence shell to an octet by attracting the electron of the other atom and at the same time retaining all its own electrons. A system of two cubes with a common edge is obtained. Denoting the electrons of one atom by circles and those of the other by crosses, we have:

\[ \begin{array}{c} \begin{array}{c} \circ\ \circ\\ \circ\,\ddot{F}\,\circ\\ \circ\ \circ \end{array} + \begin{array}{c} \times\times\\ \times\,F\,\times\\ \times\times \end{array} \longrightarrow \begin{array}{c} \circ\ \circ\\ \circ\,F\,\circ\\ \circ\ \circ \end{array} \begin{array}{c} \times\times\\ F\,\times\\ \times\times \end{array} \quad \text{or} \quad :\ddot{F}:\ddot{F}: \end{array} \]

in the usual electronic formulas.

When a double bond is formed, four electrons are shared, and the valence cubes of the atoms acquire a common face

\[ \begin{array}{c} \circ\\[-2mm] \circ\,O\,\circ\\[-2mm] \circ \end{array} : + \begin{array}{c} \times\\[-2mm] \times\,O\,\times\\[-2mm] \times \end{array} \longrightarrow \begin{array}{c} \circ\\[-2mm] \circ\,O\,\circ\\[-2mm] \circ \end{array} :: \begin{array}{c} \times\\[-2mm] O\,\times\\[-2mm] \times \end{array}. \]

What is valuable in Lewis’s theory is the clear indication that a simple (covalent) bond arises through the sharing of two electrons, a double bond is formed by four electrons, and a triple bond by six electrons. In Bohr’s initial ideas on this question there was nothing definite. Thus, in the water molecule each hydrogen–oxygen bond was represented as having arisen through three electrons. A shortcoming of Lewis’s theory was that it essentially ignored the motion of electrons. Lewis only said, and rather vaguely, that the electrons oscillate about points lying at the corners of the cube.

In 1921–22 Bohr developed a second theory of atomic structure. He accepted, in accordance with the facts of physics, that electrons move in separate orbits.

When two atoms combine with one another by sharing a pair of electrons, these electrons too must move in different ways. The following opinions were expressed concerning the form of electron orbits in a covalent bond:

1) the electrons move in separate circular orbits inclined to one another at a certain angle (I);

2) the electrons move along ellipses or an ellipse, in the focus of which the atomic nuclei are located (II);

3) the electrons move in a figure-eight, which also embraces both nuclei (III).

![schematic drawings of electron orbits labeled I, II, III]

The last scheme is the most probable. In this case the orbit may be common to two electrons, but the latter have opposite “spins.”

In general, as a result of the work of many scientists, the following ideas about the electronic types of chemical bonding have been adopted in chemistry.

A simple valence bond may be electrovalent (ionic, heteropolar) or covalent (neutral, homeopolar). In the formation of an electrovalent bond, one atom gives an electron to another. A covalent bond arises through two electrons, one from each atom, the electrons beginning to move in an orbit that embraces both atomic nuclei. A covalent bond may be nonpolar, for example in the molecule \(H_2\), with completely identical atomic nuclei, or polar, for example in the molecule HCl, where the electrons are attracted considerably more strongly by the chlorine nucleus, which leads to an asymmetric shape of the orbit.

A complex or coordination bond arises through attraction between an unshared pair of electrons of some atom of one molecule and an atomic nucleus of another molecule insufficiently surrounded by electrons. The principal feature of this bond is that the electrons forming it belong to only one atom; the character of the bond that is formed, however, may be different. In ammonium salts we have:

\[ \begin{array}{c} \mathrm{H}\\[-2pt] {:}\\[-2pt] \mathrm{H:N:}+\mathrm{H:Cl:} \end{array} \;\longrightarrow\; \begin{array}{c} \mathrm{H}\\[-2pt] {:}\\[-2pt] \mathrm{H:N:H:Cl:}\\[-2pt] {:}\\[-2pt] \mathrm{H} \end{array} \;\longrightarrow\; \left[ \begin{array}{c} \mathrm{H}\\[-2pt] {:}\\[-2pt] \mathrm{H:N:H}\\[-2pt] {:}\\[-2pt] \mathrm{H} \end{array} \right]^+ \left[ \begin{array}{c} \mathrm{\ddot{Cl}:} \end{array} \right]^- . \]

Molecules of ammonia and hydrogen chloride combined owing to the attraction between the unshared pair of nitrogen electrons and the proton of hydrogen chloride, insufficiently surrounded by electrons. Then this proton completely left the orbit of the electrons that bound it to chlorine, and the electrons of the unshared pair of nitrogen began, in their motion, to include the new proton. As a result there arose an \(N\)–\(H\) bond completely identical with the other three nitrogen–hydrogen bonds. This bond may be called a covalent complex bond.

In other cases the formation of a complex bond is limited to electrostatic attraction, without the electrons embracing a new atomic nucleus.

For example, in the complex of nitrobenzene with aniline

           :O:   H
            :    :
C6H5 : N : N : C6H5
            :    :
           :O:   H

the unshared pair of electrons of the amine nitrogen is attracted to the nitrogen of the nitro group, but, apparently, the electrons of this pair in their motion embrace only the nitrogen nucleus of the aniline molecule. Such a bond between two atoms may be called an electrovalent complex bond.

Bonds differing in their electronic character also occur in complex molecules or ions that are seemingly completely analogous in composition. Thus, by physical methods of investigation it was proved that the complex ion \([Fe(CN)_6]^{\prime\prime\prime}\) has a covalent structure, whereas the ion \([FeF_6]^{\prime\prime\prime}\) is constructed electrovalently.

Four electrons participate in the formation of a double bond. This bond, too, may have a different character. In the ethylene molecule the double bond, as the results of physical investigations show (for example, the dipole moment is zero), is covalent, i.e., all its electrons move around both atomic nuclei:

   H : C :: C : H
       H    H

In some compounds with a double bond the orbits of two electrons embrace both atomic nuclei, while the other pair of electrons moves only around one nucleus. Such a bond is called semipolar. It is assumed, for example, in the nitro group between nitrogen and one of the oxygen atoms (see above). In structural formulas a semipolar bond is often denoted by one line and by the signs \(+\) and \(-\) on the atoms:

\[ \mathrm{C_6H_5—N^{+}—O^{-}} \]
\[ \hspace{2.8em}\Vert \]
\[ \hspace{2.8em}\mathrm{O} \]

The electronic structure of molecules with a triple bond may be still more varied. One may propose, for example, the following structures for acetylene:

\[ \mathrm{H : C ::: C : H} \qquad \mathrm{H : \dot{C} :: C : H} \qquad \mathrm{H : \dot{C} : C : H} \qquad \mathrm{H : \dot{C} : C : H,} \]

In addition, acetylene may also pass into the tautomeric form \(\mathrm{C=CH_2}\) with divalent carbon. This diversity in electronic structures may be the cause of certain contradictions in the data on the chemical properties of ethylenic and acetylenic hydrocarbons. According to the strain theory, a triple bond should break more readily than a double bond; in actual fact, it sometimes undergoes addition reactions more readily, sometimes with more difficulty. One may suppose that under certain conditions, for example under the action of certain reagents, acetylene passes into a form more reactive than the active form with a semipolar bond probable for ethylene, while under other conditions a less reactive form is fixed in it.

Wave theory of the structure of atoms and molecules

In 1924 de Broglie expressed the idea that all material particles—electrons, protons, etc.—are not only particles but have a wave nature, representing concentrations of certain waves. The wave nature.

electrons, and then of atomic nuclei, was soon proved experimentally. It was shown that they produce the phenomenon of diffraction, which is characteristic only of waves.

Physicists assert \([8, 355]\) that condensations of waves, or wave packets, cannot be stable and must quickly spread out. In this connection some speak of “probability waves,” of the fact that diffraction phenomena, characteristic of waves, are produced as if only by beams of electrons. However, diffraction is observed, as L. Biberman, N. Sushkin, and B. Fabrikant \([17]\) have shown, also for electron beams of very low intensity, where individual electrons can be followed. As for the fact that electrons are stable, whereas wave packets, according to the requirements of classical mechanics, cannot be stable, one may recall the introduction into science of Bohr’s postulates. It was known that a system consisting of a positive nucleus and electrons moving around it, according to the laws of electrodynamics, must radiate energy and therefore cannot be stable. On the other hand, it had been proved that the atom must consist of a nucleus and electrons moving around it and at the same time be stable.

Explaining this contradiction, Bohr violated the laws of electrodynamics by adopting as a postulate that in certain orbits an electron can move without loss of energy. As later became clear, this is explained by the fact that, in addition to the laws of electrodynamics, the electron obeys still other laws connected with its wave nature. But the waves themselves that constitute the electron and other particles are something distinctive and, perhaps, do not in every respect obey laws derived for other kinds of wave motion. In particular, for these waves it may be necessary to put forward the postulate that certain wave packets, satisfying definite conditions, can be stable. What these conditions are and what explains the stability of such wave packets—the answers to these questions must be provided by the further development of science; but in any case the wave nature of electrons is a fact, and their stability is also a fact. If these facts seem incompatible from the point of view of contemporary physics, then from the point of view of dialectical materialism there is nothing strange in this. It means only that some new physical laws must be discovered, corresponding to the new forms of matter and of its motion being discovered by science.

In atomic physics, wave theory first of all explained the quantum laws. Thus, the existence of stable electronic orbits, contradicting the old electrodynamics, was explained by the fact that on these orbits there is fitted an integral number of waves corresponding to the given electron. The detailed development of wave theory and the quantum mechanics associated with it was achieved in the works of Schrödinger, Heisenberg, Dirac, and a number of other scientists. In 1926 Schrödinger derived the basic equation used by physicists in their considerations of the structure of atoms:

\[ \frac{\partial^2 \psi}{\partial x^2} + \frac{\partial^2 \psi}{\partial y^2} + \frac{\partial^2 \psi}{\partial z^2} + \frac{8\pi^2 m}{h^2}(\varepsilon - V)\psi = 0 . \]

Here \(\psi\) is the amplitude of the electron waves, \(\varepsilon\) is the total energy of the electron, \(V\) is the potential energy, and \(h\) is Planck’s radiation constant.

Great disagreements arose over the question of the real physical meaning of this equation, and in particular of the basic quantity \(\psi\). Some physicists hold that \(\psi^2\) expresses the probability of finding an electron in a definite element of the volume of the atom. Others speak of the density of the “electron cloud” in this element of volume, believing that there are no individual electrons in the atom. However, considerations may be adduced in favor of the presence in the atom of electrons as, to a considerable degree, separate particles. D. D. Iva-

...nenko [7], in proving the impossibility of the existence of electrons in atomic nuclei, motivates this proposition by the fact that the binding energy of an electron inside the nucleus should be about 7 million electron-volts, whereas the electron’s own energy, calculated from the Lebedev—Einstein equation, \(\varepsilon = mc^2\), reaches only 0.5 million electron-volts. Consequently, an electron in the atomic nucleus must be, so to speak, crushed. But in the outer part of the atom, especially in its valence shell, the binding energy of the electrons is measured only in individual electron-volts or in tens of electron-volts. Hence there is no reason to think that an electron which in cathode rays undoubtedly flies as a separate particle, on entering an atom, molecule, or ion for some reason ceases to be a particle. The processes of acquisition or loss of electrons by atoms are energetically far too insignificant for so sharp a change.

Of course, one cannot regard the electron as a completely separate particle, completely independent of the surrounding matter, as a “foreign body floating in the ether,” as was thought earlier. An experimental solution of this question is provided at least by the experiments mentioned above on the diffraction of very weak electron beams [17]. The diffraction gratings used in these experiments had spacings of atomic size (of the order of \(10^{-8}\ \mathrm{cm}\)). Consequently, if electrons were only particles with linear dimensions of the order of \(10^{-13}\ \mathrm{cm}\), then most of them would have had to pass freely through such a grating without undergoing diffraction. Consequently, when speaking of the dimensions of the electron, one must bear in mind its dual nature. The electron represents a condensation of waves, and the action on these waves is manifested at least in the space of the atom. But the densest part of the condensation, possessing the electrical properties of the electron and already representing a qualitatively different form than the surrounding matter, apparently occupies a very small space, so that it is quite possible to speak of the place occupied by the electron in the atom and of the orbit along which it moves¹.

Consequently, Schrödinger’s equation speaks not of the density of an “electron cloud” in general (one may speak of an “electron cloud” only conditionally, and it would be better to abandon this expression altogether), but of the probability of finding the electron at a definite place in the atom or molecule.

Among physicists who recognize electrons as small particles, rather sharply delimited from the surrounding matter, there are also disagreements. Schrödinger’s equation has a statistical character; it speaks only of probability. Some scientists [27] believe that the laws of motion of electrons in atoms have, in principle, a statistical character, that no dynamical regularities are hidden behind the statistical laws of the intra-atomic world. It is said, for example, that electrons possess “free will.” One also cannot in any way agree with such assertions. The statistical character of the laws of atomic physics known to us by no means indicates the absence of dynamical regularities in intra-atomic motions. The motion of electrons in the atom is more complex than was thought earlier, and has not been sufficiently elucidated by us. The electrons themselves have proved to be complex systems, with whose structure and properties we are still little acquainted, but

From the point of view of dialectical materialism, this is precisely what was to be expected. V. I. Lenin, in his book Materialism and Empirio-Criticism, wrote as early as 1909 [1]: “The electron is as inexhaustible as the atom; nature is infinite, but it exists infinitely, and this sole categorical, sole unconditional recognition of its existence outside the consciousness and sensation of man is what distinguishes dialectical materialism from relativist agnosticism and idealism.”

In these words of Lenin we find an answer also to the arguments of those bourgeois physicists who, as it were, acknowledge the existence of electrons and the dynamic laws governing their motion, but assert that we cannot elucidate these laws. They advance the proposition of “fundamental unobservability,” i.e., they believe that if at present we do not possess experimental means of investigation that would allow us to trace the motion of individual electrons in an atom or molecule, then, consequently, we shall never have such means of investigation.

However, A. M. Butlerov developed his theory of chemical structure and drew concrete conclusions about the structure of organic compounds at a time when science knew of no means for the direct determination of the structure of molecules and did not even have direct proof of the very existence of atoms and molecules. Butlerov’s conclusions were based on indirect considerations, but these conclusions proved to be correct. They are fully confirmed by all the data of modern science.

Soviet chemists concerned with questions of the structure of atoms and molecules must not forget the statement of J. V. Stalin in the work On Dialectical and Historical Materialism [2]:

“In contrast to idealism, which disputes the possibility of knowing the world and its laws, which does not believe in the reliability of our knowledge, which does not recognize objective truth and holds that the world is full of ‘things-in-themselves’ which can never be known by science—Marxist philosophical materialism proceeds from the principle that our knowledge of the laws of nature, tested by experiment and practice, is reliable knowledge having the significance of objective truths, that there are no unknowable things in the world, but only things not yet known, which will be disclosed and known by the forces of science and practice.”

We must be confident in the possibility of knowing not only the statistical but also the dynamic laws of the intra-atomic world, and we have every right to speak of the motion of individual electrons in atoms and molecules and to explain by this motion various physical and chemical phenomena.

Application of Wave Theory and Quantum Mechanics in Organic Chemistry

Having developed a general conception of the various electronic types of chemical bond, electronic chemistry approached the quantitative calculation of chemical processes. Considerable success was achieved in applying electronic considerations to compounds with electrovalent bonds and to complex compounds. The values of the energy of formation of the electrovalent bond in individual molecules and the energies of the crystal lattice for many substances were calculated. The coordination numbers of elements in complex compounds received a satisfactory explanation. It turned out that the coordination number is determined primarily by the charge and radius of the complex-forming ion.

Electronic chemistry by itself could not cope with the covalent bond. Calculation of the energy of formation even of the simplest molecule, \(H_2\), gave completely unsatisfactory results if one proceeded only from

from the electrostatic interaction of atomic nuclei and electrons. Nor was taking electromagnetic forces into account of any help here.

Wave theory, together with quantum mechanics, brought new possibilities. An entire new branch of science—quantum chemistry—developed, and by the present time it has a number of achievements:

1. It has been clarified that, in the process of formation of a covalent bond, the wave interaction of atomic nuclei and electrons plays a significant role. Only by taking into account the quantum laws that are the expression of wave interaction was it possible to calculate the energy of formation of the hydrogen molecule and of certain other simplest molecules.

2. A doctrine of the valence states of atoms has been developed, which is already useful in solving specific problems in chemistry. For organic chemists the valence states of the nitrogen atom are of great importance. This element is in group 5 of the system, i.e., its highest valence would seemingly have to be equal to 5. However, the experience of chemists shows that nitrogen can in no way be combined with five identical atoms or groups of atoms: \(NH_5\), \(NF_5\), \(N(CH_3)_5\) do not exist, while ammonium salts do exist, for example, \(NH_4Cl\), in which four identical or analogous atoms or radicals are bonded to the nitrogen atom, while the fifth atom or atomic group constitutes an anion and in solution moves separately from the ammonium cation, and in the crystal occupies a special point of the crystal lattice.

According to quantum laws, nitrogen cannot be pentacovalent, but can be tetra-covalent positive [16, 62], which fully explains the indicated facts.

In the new way we must also conceive the structure of nitro compounds. In the nitro group there cannot be two covalent double bonds between nitrogen and oxygen. One of them must be semipolar; i.e., here nitrogen is also tetracovalent positive. This conclusion is also confirmed by experiment. It is indicated by the parachor of nitro compounds and by their considerable polarity, manifested, for example, in the dipole moment and in the capacity for complex formation [26].

In accordance with the requirements of quantum mechanics, none of the elements of the 2nd row of the periodic system (from lithium to neon) can have more than four covalent bonds. The conclusions drawn by chemists from experimental data are in full agreement with this.

New confirmation of the same proposition was given by our experiments on the action of esters of complex acids on benzene in the presence of catalysts [21]. It turned out that esters of nitric and sulfuric acids behave differently in this reaction. With nitrates, nitrobenzene is obtained; with sulfates—alkylbenzenes:

\[ C_6H_6 + C_2H_5ONO_2(+AlCl_3) \longrightarrow C_6H_5NO_2 + C_2H_5OH, \]

\[ 2C_6H_6 + (C_2H_5O)_2SO_2(+AlCl_3) \longrightarrow 2C_6H_5CH_2CH_3 + H_2SO_4. \]

Meanwhile, \(HNO_3\) and \(H_2SO_4\) are acids of almost identical strength, and one might have expected that their esters would react with benzene in one and the same direction.

The difference becomes understandable if one takes into account that nitrogen and sulfur are in different rows of the periodic system. According to quantum theory, nitrogen cannot have more than four covalent bonds, whereas sulfur can quite well be hexacovalent. This is also clearly indicated by the existence of \(SF_6\), which has the properties of a typical substance with covalent bonds. Hence the esters of nitric and sulfuric acids must have different electronic structures. The complexes of the esters with catalysts must also be different, which in turn affects the course of reactions with benzene:

\[ \mathrm{C_6H_6 + C_2H_5O{-}\overset{O}{\overset{\|}{N}}{}^{+}{-}O^{-}\cdots AlCl_3 \longrightarrow C_2H_5O{-}\overset{O}{\overset{\|}{N}}{}^{+}{-}O\cdots AlCl_3 \longrightarrow} \]
\[ \begin{array}{c} \\[-1.5em] \mathrm{\ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ C_6H_6} \end{array} \]

\[ \mathrm{ C_6H_5{-}\overset{O}{\overset{\|}{N}}{}^{+}{-} + C_2H_5OH } \]
\[ \mathrm{ \ \ \ \ \ \ \ \ \overset{\displaystyle O^{-}\cdots AlCl_3}{} } \]

\[ \mathrm{ C_6H_5H + C_2H_5{-}O{-}\overset{O}{\overset{\|}{S}}{-}OC_2H_5 \longrightarrow C_6H_5{-}C_2H_5 + HO{-}\overset{O}{\overset{\|}{S}}{-}OC_2H_5 \ etc. } \]
\[ \mathrm{ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ AlCl_3 \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ \ AlCl_3 } \]

1. Quantum mechanics makes it possible to calculate the magnitudes of the angles between bonds in multivalent atoms. The results obtained in this way correspond to what had earlier been found by chemists and later confirmed by physical measurements. Thus, for carbon, quantum considerations lead to a tetrahedral structure.

2. Quantum mechanics has led to a new understanding of double and triple bonds. The two pairs of electrons by means of which a double bond arises prove to differ in the character of their motion. One pair, the σ-electrons, moves, like the electrons of a simple bond, along a stable orbit (or orbits) in the sphere of attraction of the nuclei of both atoms. The second electron pair, the π-electrons, in a covalent double bond also moves in the field of attraction of both nuclei, but the orbits of these electrons have a different form and are unstable.

Owing to the easy changeability of the orbits of π-electrons, molecules with double bonds can be polarized under the influence of various factors. The physical index characterizing polarizability is refraction. The refraction of a simple carbon bond for the sodium D line is 1.21; the refraction of a double bond is 4.16 [19, 46].

When there is a large difference in the electrochemical character of the atoms, the molecule of a compound with a double bond is polar already in itself, and sometimes the π-electrons move only around one nucleus, i.e., the bond becomes semipolar (for example, in the nitro group). With a smaller difference, incomplete polarization arises, with only some asymmetry in the orbits. Unsymmetrical homologs of ethylene may serve as an example. This explains the definite direction of addition reactions (according to Markovnikov’s rule). Finally, even molecules of entirely symmetrical structure, such as the ethylene molecule, are in themselves completely nonpolar (according to physical data), but are readily polarized under the influence of external conditions, for example, by the action of one reagent or another. It is in the polar form that unsaturated compounds enter (or, at any rate, enter more readily into) reactions.

The π-electrons move very readily in compounds with conjugated systems of bonds, the extreme representative of which is benzene. The mobility of the electron orbits here reaches the point that all six π-electrons of the benzene ring have, in essence, merged into one system, in which there are no localized double bonds. The results of magnetic investigations indicate the presence in the benzene molecule of superconductivity, i.e., free motion of the π-electrons along the ring [19, 238]. The structure of benzene is apparently best expressed by formula I, in which the dots denote π-electrons. In reactions, under the influence of external effects, the arrangement of the π-electrons

changes, and the molecule can pass into forms of type II with covalent double bonds (the Kekulé formula) or, perhaps, of type III with semipolar bonds

As is shown by the properties of homologues and derivatives of benzene, polarization of the benzene ring also occurs under the influence of atoms and atomic groups entering into the molecule of the substance itself. This explains the substitution rules.

In a triple bond we have one pair of $\sigma$-electrons and two pairs of $\pi$-electrons. According to the laws of quantum mechanics, there should be no quadruple bonds. In this, too, conclusions long ago drawn by chemists are confirmed.

It is interesting that the magnitudes of the energies of formation of double and triple bonds are by no means two and three times greater than the energies of formation of single bonds. Thus, the average value of the energy of a single bond between carbon atoms is $62.77$ kcal/g-mol, of a double bond $101.2$ kcal/g-mol, and of a triple bond $128.2$ kcal/g-mol; i.e., energetically the second and third bonds are not equivalent to the first.

1. It will be shown below that wave theory and quantum mechanics contribute much to explaining the mutual influence of atoms.

Types of the Mutual Influence of Atoms

Effect of Electrical Interaction

In evaluating the influence of some atom or atomic group on other atoms entering into the composition of a molecule, we, following the suggestion of A. M. Butlerov and V. V. Markovnikov, compare the properties of the substance under study with the properties of a compound in which hydrogen is present in place of the atom or group of interest to us. Replacement of a hydrogen atom by another atom or atomic group leads to a change in the arrangement of the electrons, more precisely, of the electronic orbits. More often this is called a change in the distribution of electron density.

In its original form, electronic theory explained these changes exclusively by the action of electrostatic forces. At present, only one of the types of mutual influence of atoms, known as the inductive effect, is associated with electrostatic forces; another type of mutual influence of atoms is distinguished from it and is usually called the conjugation effect.

The effect of electrical interaction was first studied, above all, in carboxylic acids. Comparison of many data indicates that the ability of acids to undergo electrolytic dissociation is determined mainly by the electrostatic interaction of the atoms making up the acid molecule. Ostwald had already observed that the constant of electrolytic dissociation of an acid increases upon the introduction of electronegative atoms.

or groups. This regularity was explained from the electronic point of view [18, 35—43] by Falk in 1911, and later by Lewis in 1923.

If, in the acetic acid molecule, instead of a hydrogen of the methyl group, an atom of a strongly electronegative element is introduced—for example, chlorine—then the positive part of the chlorine atom attracts electrons more strongly than the hydrogen nucleus (proton) can do. As a result, the pair of electrons joining chlorine to carbon is displaced toward chlorine and away from carbon (in comparison with what was present in the acetic acid molecule). The nucleus of the carbon atom now becomes less well neutralized and attracts more strongly the electrons of the remaining bonds of this carbon, including the pair of electrons that joins it to the carbon of the carboxyl group; the nucleus of the carboxyl carbon must now attract and neutralize more strongly the electrons of the bond with the hydroxyl oxygen, and this oxygen attracts and neutralizes more strongly the electrons of its bond with hydrogen. Naturally, the hydrogen of the hydroxyl in chloroacetic acid must be detached more readily in the form of a cation, i.e., the acidic properties must be strengthened:

\[ \begin{array}{cc} \begin{array}{c} H\\ H:\ddot{C}:C:\ddot{O}:H\\ H\quad \ddot{O}: \end{array} & \begin{array}{c} H\\ :\ddot{Cl}:\ \ddot{C}:C:\ddot{O}:H\\ H\quad \ddot{O}: \end{array} \end{array} \]

In abbreviated form such a displacement of electrons is expressed by arrows:

\[ \mathrm{Cl}\leftarrow \mathrm{CH}_2 \leftarrow \mathrm{C}\leftarrow \mathrm{O}\leftarrow \mathrm{H} \]
\[ \phantom{\mathrm{Cl}\leftarrow \mathrm{CH}_2 \leftarrow}\Vert \]
\[ \phantom{\mathrm{Cl}\leftarrow \mathrm{CH}_2 \leftarrow}\mathrm{O} \]

The influence of halogens, and also of hydroxyl, when they replace hydrogen atoms in the acid radical is well illustrated by the following table.

Table 1

We see that halogens greatly increase the constant of electrolytic dissociation of acids of the limiting series. The most electronegative element, fluorine, acts most strongly; iodine acts most weakly. Hydroxyl has an even weaker effect, but in the same direction. An increase in the number of halogen atoms intensifies the acidity. The position of the halogen or hydroxyl also has an effect. As should be expected from the laws of electrostatics, this effect weakens as the distance between the atoms increases. \(\alpha\)-Chlorobutyric aci-

almost 100 times stronger than butyric acid, β-chlorobutyric acid only 6 times, and γ-chlorobutyric acid less than twice.

In the present case we have considered the propagation of the effect of electrical interaction along a chain of atoms, but the electric charges of different parts of a molecule can act upon one another also directly through space. This “field effect” [18, 45] is most directly observed in those cases where the order of bonding of the atoms in the compounds being compared is the same, while the distances between the atoms are different, i.e., in cis–trans geometrical isomerism. Of the two geometrically isomeric acids, maleic and fumaric,

\[ \begin{array}{cc} \mathrm{H{-}C{-}COOH} & \mathrm{H{-}C{-}COOH} \\ \| & \| \\ \mathrm{H{-}C{-}COOH} & \mathrm{HOOC{-}C{-}H} \\ \text{maleic acid} & \text{fumaric acid,} \end{array} \]

the first is considerably stronger. The carboxyl is an electronegative group that increases the acidic properties of the other carboxyl. The closer two such groups are to one another, the stronger the acidity (Table 2).

Table 2

However, this holds only for the first dissociation constant, corresponding to the splitting off of one proton. In maleic acid the first dissociation constant is 13 times greater than in fumaric acid. The second dissociation constant, on the contrary, is 60 times greater in fumaric acid than in maleic acid, since in the cis position the second proton is firmly held by the negative charge of the COO group located near it.

In the materials of the All-Union Conference on the Theory of Chemical Structure in Organic Chemistry, a number of further examples are given of a very noticeable manifestation of the effect of electrical interaction. It is known, for example, that in tertiary alcohols the reactivity of the hydroxyl hydrogen is greatly weakened in comparison with primary alcohols. The results of determining the electromotive force arising in the reaction of alcohols with sodium, and the rate of this reaction, showed that trimethylcarbinol is much less active with respect to metals than methanol [22]. At the same time, trifluoromethyl-dimethylcarbinol \(\mathrm{CF_3.C(CH_3)_2OH}\) decomposes even carbonic acid salts [19, 48].

In the presentation by G. P. Miklukhin [19, 293–5], the influence of the cations of acetic salts on the reactivity of the hydrogens of the methyl group was noted. It turns out that if sodium is introduced into the carboxyl, the methyl hydrogens become very little mobile. Evidently, the negative charge of the acetic acid anion prevents the splitting off of protons from the methyl group as well. If, however, a lead atom is introduced into the molecule, a weakly dissociating salt is obtained, in which the cation promotes the expulsion of protons. Lead acetate could even be condensed with benzaldehyde without the use of acetic anhydride.

The effect we have considered may be static, i.e., it may manifest itself in a nonreacting molecule and be reflected, for example, in the dipole moment. There may also be an electrostatic effect of another kind, when the molecule undergoes not polarization (or a change in the degree of polarization), noticeable from the corresponding physical properties of the substance, but only a capacity for polarization (polarizability). Such an influence of atoms or atomic groups is dynamic; it manifests itself at the moment of reaction under the action of the attacking reagent. For the chemist, dynamic influence is no less important than static influence.

Mutual influence of atoms not explained by electrical forces

In the examples discussed above, we dealt with the manifestation of the effect of electrical interaction, not obscured or almost not obscured by other influences. However, in the properties of many organic compounds, peculiarities are observed that indicate the presence of some other kind of interaction of atoms.

Thus, the electrolytic dissociation constants of isomeric nitrophenols do not obey the general rule of electrical interaction \[19, 48\]:

\[ \begin{array}{cccc} \mathrm{OH} & \mathrm{OH} & \mathrm{OH} & \mathrm{OH} \\ \text{phenol} & o\text{-}\mathrm{NO_2}\text{ phenol} & m\text{-}\mathrm{NO_2}\text{ phenol} & p\text{-}\mathrm{NO_2}\text{ phenol} \\ K_{25}=1.7\cdot10^{-10} & 0.75\cdot10^{-7} & 1.0\cdot10^{-8} & 0.96\cdot10^{-7} \end{array} \]

The nitro group in all cases increases the acidity of the phenolic hydroxyl, but this influence proves strongest in the para position, i.e., at the greatest distance from \(\mathrm{NO_2}\), and weakest in the meta position, i.e., at an intermediate distance.

An even more striking example of such an anomaly in properties is found in the oxybenzoic acids:

\[ \begin{array}{cccc} \mathrm{COOH} & \mathrm{COOH} & \mathrm{COOH} & \mathrm{COOH} \\ \text{benzoic acid} & o\text{-}\mathrm{OH}\text{ acid} & m\text{-}\mathrm{OH}\text{ acid} & p\text{-}\mathrm{OH}\text{ acid} \\ K_{25}\cdot10^{4}=0.652 & 10.6 & 0.87 & 0.285 \end{array} \]

The hydroxyl in the ortho position sharply increases acidity, in the meta position it increases it slightly, and in the para position it even weakens it. The same is observed for methoxybenzoic acids.

It is known that in ortho- and para-chloronitrobenzenes the halogen atoms are distinguished by great reactivity, whereas in meta-chloronitrobenzene chlorine is almost as inert as in chlorobenzene. In triphenylmethane dyes, color is associated with the presence of electron-donor substituents in the ortho and para positions, but not in the meta position. Many medicinal substances, for example sulfonamide preparations, belong to para-disubstituted compounds.

derivatives of benzene. Among the isomeric nitrotoluenes, it is precisely in para-nitrotoluene that the methyl hydrogens are distinguished by a special reactivity, although they are very far from the nitro group.

A comparison of the dipole moments of organic halides, the dipole moments of furan, tetrahydrofuran, divinyl and diethyl ethers, and a number of other compounds also leads to results that cannot be explained by electrical effects alone.

Thus, alongside the electrical interaction of electrons and atomic nuclei, there exists some other interaction as well. Since electrons and atomic nuclei are not only particles carrying definite electric charges, but also have a wave nature, one must think that they can act upon one another not only electrostatically but also through waves. The new type of mutual influence of atoms that we are considering is most probably explained precisely by their wave interaction.

This type of mutual influence of atoms manifests itself especially strongly in molecules with conjugated bond systems and is evidently connected with the easy variability of the orbitals of \(\pi\)-electrons. Many call such an influence of atoms the conjugation effect. The conjugation effect explains the peculiarities in the properties of aromatic compounds, for example the phenomena of orientation in the benzene ring, the anomalies noted above in the influence of substituents on the acidity of the carboxyl and phenolic hydroxyl groups, on the reactivity of halogens, etc. In compounds with an open chain, transfer of the activating action is also observed, sometimes to a great distance from the activating group itself (“vinylogy”). Thus the methyl hydrogens in acetic acid esters

\[ \begin{array}{c} \mathrm{CH_3{-}C{=}O}\\ \quad |\\ \mathrm{OR} \end{array} \]

have an increased reactivity owing to the influence of the carbonyl group. The hydrogens of the \(\mathrm{CH_3}\) group in the esters of crotonic and sorbic acids are distinguished by the same increased reactivity [19, 56]:

\[ \mathrm{CH_3{-}CH{=}CH{-}C{=}O} \qquad \mathrm{CH_3{-}CH{=}CH{-}CH{=}CH{-}C{=}O}. \]

\[ \begin{array}{c} \phantom{\mathrm{CH_3{-}CH{=}CH{-}}}|\\[-0.4em] \phantom{\mathrm{CH_3{-}CH{=}CH{-}}}\mathrm{OR} \end{array} \qquad \begin{array}{c} \phantom{\mathrm{CH_3{-}CH{=}CH{-}CH{=}CH{-}}}|\\[-0.4em] \phantom{\mathrm{CH_3{-}CH{=}CH{-}CH{=}CH{-}}}\mathrm{OR} \end{array} \]

In the regularities of its manifestation and propagation, the conjugation effect differs sharply from the effect of electrical interaction:

1. The directions of action of these effects sometimes coincide and sometimes are opposite, as is shown by the example of oxybenzoic and methoxybenzoic acids.

2. The conjugation effect weakens less with increasing distance.

3. The conjugation effect is in a peculiar dependence on the spatial structure of molecules. It acts only in planar (coplanar) systems of atoms. All atomic nuclei of a conjugated system must lie in one plane. In A. N. Terenin’s report at the conference on the theory of chemical structure, the following examples were given [19, 54].

\[ \begin{array}{cccc} \begin{array}{c} \mathrm{NO_2}\\ |\\ \text{benzene ring } I\\ |\\ \mathrm{Cl} \end{array} & \begin{array}{c} \mathrm{NO_2}\\ |\\ \text{benzene ring } II\\ |\\ \mathrm{Cl}\\ \text{with } \mathrm{CH_3}\text{ groups at both ortho positions to } \mathrm{Cl} \end{array} & \begin{array}{c} \mathrm{NO_2}\\ |\\ \text{benzene ring } III\\ |\\ \mathrm{Cl}\\ \text{with } \mathrm{CH_3}\text{ groups at both ortho positions to } \mathrm{NO_2} \end{array} & \begin{array}{c} \mathrm{NO_2}\\ |\\ \text{benzene ring } IV\\ |\\ \mathrm{Cl}\\ \text{with } \mathrm{CH_3}\text{ groups at all four ortho positions} \end{array} \end{array} \]

In compounds I and II chlorine is mobile owing to the influence of the nitro group, transmitted without hindrance along the benzene ring. If, however, methyls stand on both sides of the nitro group (compounds III and IV), then the nitro group loses coplanarity with the benzene ring and with chlorine. Therefore its influence on the halogen is destroyed, although the distance between \(\mathrm{NO_2}\) and \(\mathrm{Cl}\) remains the same.

The same has been observed for nitrophenols. p-Nitrophenol (I) has considerably stronger acidic properties than phenol \((-\lg K = 7.21)\); 2,6-dimethyl-4-nitrophenol (II) has almost the same acidity \((-\lg K = 7.16)\), but the dissociation constant of 3,5-dimethyl-4-nitrophenol (III) is considerably smaller \((-\lg K = 8.24)\); the acidity-increasing influence of the nitro group is preserved here (perhaps owing to the electrical effect), but has become much weaker.

\[ \begin{array}{ccc} \text{I: } p\text{-nitrophenol} & \text{II: } 2,6\text{-dimethyl-}4\text{-nitrophenol} & \text{III: } 3,5\text{-dimethyl-}4\text{-nitrophenol} \\[4pt] \mathrm{HO{-}C_6H_4{-}NO_2} & \mathrm{HO{-}C_6H_2(CH_3)_2{-}NO_2}\ (2,6\text{-dimethyl}) & \mathrm{HO{-}C_6H_2(CH_3)_2{-}NO_2}\ (3,5\text{-dimethyl}) \end{array} \]

The same cause explains the instability of cyclooctatetraene despite the great stability of benzene. The benzene ring is a planar hexagon with conjugated bonds and, in any case, without any great deviation of the carbon valence units from their normal positions.

\[ \text{cyclooctatetraene: an eight-membered ring with alternating } \mathrm{CH{=}CH} \text{ links} \]

In cyclooctatetraene one of two things is possible: either a planar ring, but with spatial strain owing to displacement of the valence units, or absence of displacement of the valence units, but with violation of the planar arrangement of the atoms; in that case the conjugation effect must disappear, i.e. the double bonds must retain the activity characteristic of them in ethylenic hydrocarbons.

As indicated in the report of A. N. Terenin, “The disturbance of conjugation when a system is taken out of coplanarity is a purely quantum effect, inexplicable from the standpoint of the classical conception of electrons” [19, 55]. And at the basis of quantum effects lies the wave nature of electrons.

1. It is not difficult to see that the mutual influence of atoms in conjugated systems has a periodic character. The properties of the links of the benzene ring change periodically: the links situated in the ortho and para positions to a substituent are similar to one another, but differ sharply from the meta links. Periodicity is observed (with period \(-\mathrm{CH{=}CH}-\)) in the transmission of the influence of a carbonyl group to a methyl group in the esters of acetic, crotonic, and sorbic acids. All this is quite natural for wave processes, to which periodicity is inherent by their very nature.

Less studied is the effect of wave interaction in chains of atoms with single bonds. From the standpoint of wave theory, attention must be paid to the fact that all electrons and atomic nuclei have a wave nature. Consequently, in principle one may expect the manifestation of wave interaction for all atoms and for all types of bond. If in systems with \(\pi\)-electrons arranged in a definite manner it is manifested more strongly, it still does not follow from this that in chains with single bonds, where there are

only σ-electrons, there can be no wave interaction at all, even if weaker. σ-bonds must enter into conjugation both with π-bonds and with one another. These phenomena of “superconjugation,” “hyperconjugation,” and “σ-conjugation” [13] have in recent years been studied in detail by A. N. Nesmeyanov and co-workers. The features in the properties of “quasi-complex” compounds, which include the products of the addition of halides of certain elements to unsaturated compounds, are explained by “superconjugation.” For example, adducts of acetylene with $\mathrm{HgCl_2}$, of ethylene with $\mathrm{Hg(OH)Cl}$, etc., in many reactions behave in accordance with the formulas of products of valence addition

\[ \mathrm{Cl{-}CH{=}CH{-}HgCl} \quad \text{or} \quad \mathrm{HO{-}CH_2{-}CH_2{-}HgCl}, \]

whereas in other cases they very readily give up $\mathrm{HgCl_2}$ or $\mathrm{Hg(OH)Cl}$, reacting like complexes. The results of an investigation of the spectra of combination scattering of light [12] decide the question in favor of the valence structure of the compounds indicated. Their ability readily to decompose with liberation of acetylene or ethylene is explained by the strong mutual influence of atoms through a $\sigma$–$\pi$ and even through a $\sigma$–$\sigma$ conjugated system.

An interesting example of a more complex system with a clear influence of such conjugation is dimethylvinylcarbinol acetate, which adds a molecule of mesitylmagnesium bromide in positions 1 and 6

\[ \begin{aligned} &\underset{6}{\mathrm{CH_2}} -\underset{5}{\mathrm{CH}} -\underset{\substack{4\\[-0.2em] |\,\mathrm{CH_3}\\[-0.2em] |\,\mathrm{CH_3}}}{\mathrm{C}} -\underset{3}{\mathrm{O}} -\underset{\substack{2\\[-0.2em] |\,\mathrm{CH_3}}}{\mathrm{C}} -\underset{1}{\mathrm{O}} + \mathrm{(CH_3)_3C_6H_2MgBr} \ \xrightarrow{\mathrm{H_2O}} \\[1.0em] &\mathrm{(CH_3)_3C_6H_2} -\underset{6}{\mathrm{CH_2}} -\underset{5}{\mathrm{CH}} {=}\underset{\substack{4\\[-0.2em] |\,\mathrm{CH_3}\\[-0.2em] |\,\mathrm{CH_3}}}{\mathrm{C}} +\mathrm{HO} -\underset{\substack{2\\[-0.2em] |\,\mathrm{CH_3}}}{\mathrm{C}} {=}\mathrm{O} \end{aligned} \]

Such a direction of the reaction is explained [16] by the fact that the simple bond between the 3rd and 4th atoms of the chain is capable of entering into conjugation with the double bonds $\mathrm{C{=}O}$ and $\mathrm{C{=}C}$ and, owing to this, is very readily cleaved.

Like ordinary conjugation, “superconjugation” depends on the spatial structure of molecules. Here the spatial arrangement of the electron orbits is also important. Thus, having determined the reactivity of hydrogen atoms in several cyclic ketones, A. N. Nesmeyanov, D. N. Kursanov, K. A. Pecherskaya, and Z. N. Parnes [14] considered it probable that in some ketones “conjugation is absent because of the mutual perpendicularity of the axis of the π-cloud of the carbonyl group and the σ-cloud of the $\mathrm{C{-}H}$ bond, for the same reasons for which conjugation of two π-bonds is excluded in the case of a mutually perpendicular arrangement of the axes of the corresponding clouds.” The same explanation is given for the disturbance of conjugation in molecules of cobalt and aluminum acetylacetonates [15].

At the conference on the theory of chemical structure in organic chemistry, the question of the manifestation in chains of atoms with simple bonds of yet another feature characteristic of wave interaction—namely periodicity in the mutual influence of atoms—was discussed, but was not resolved. The idea of an alternating effect, or alternating polarity in various forms, has been advanced in the works of a number of scientists since the very beginning of the 20th century (Flurscheim, Fry, Berkenheim, Mikhailenko, Lapworth, Kerman, Robinson, Stieglitz). Other scientists (Lucas, Ingold) deny the alternation of charges in saturated chains [18, 43].

Considering periodicity to be an inevitable feature of that type of mutual influence of atoms whose cause is wave interaction, I shall give examples confirming the presence of such periodicity even in fairly long chains of atoms with simple bonds. Conant, measuring the rate of reaction of organic chlorides with potassium iodide, drew attention to the alternation of increases and decreases in the reactivity of the halogen in the series:

\[ \mathrm{C_6H_5Cl,\ C_6H_5\cdot CH_2Cl,\ C_6H_5\cdot CH_2\cdot CH_2Cl,\ C_6H_5\cdot CH_2\cdot CH_2\cdot CH_2Cl.} \]

\[ \mathrm{C_6H_5COCl,\ C_6H_5CO\cdot CH_2Cl,\ C_6H_5CO\cdot CH_2\cdot CH_2Cl,\ C_6H_5CO\cdot CH_2\cdot CH_2\cdot CH_2Cl.} \]

In chlorobenzene the halogen atom, directly bonded to the carbon of the benzene ring, is held very firmly; in benzyl chloride the halogen, separated from the benzene ring by one group, is distinguished, on the contrary, by extraordinary activity; with two \(\mathrm{CH_2}\) groups the reactivity of the halogen again decreases, and with three \(\mathrm{CH_2}\) groups it increases once more, although not to the same extent as in \(\mathrm{C_6H_5CH_2Cl}\). An analogous change in the activity of chlorine is observed in the second of the series given here as well. Thus, the influence of the \(\mathrm{C_6H_5}\) and \(\mathrm{C_6H_5CO}\) groups on chlorine is transmitted along the chain, gradually weakening, but the direction of this influence changes periodically.

The same may also be noted for aromatic alcohols. Table 3 gives the values of the relative rate of reaction (for \(\mathrm{CH_3OH}\) the rate is taken as 100) of benzyl, phenylethyl, and phenylpropyl alcohols with p-nitrobenzoyl chloride according to the data of Norris and Ashdown, and also the values determined by us for the electromotive force arising in the reaction of these alcohols with sodium [23].

Table 3

Here a distinct periodicity is observed in the change in the reactivity of the hydroxyl hydrogen.

In the presentation by O. A. Reutov at a meeting of the Academy of Sciences of the USSR [19, 22], it was noted that the chlorine of \(\beta\)-chloropropionic acid departs as an anion much more readily than in \(\alpha\)-chloropropionic acid; that, in general, substituents in the \(\alpha\)- and \(\gamma\)-positions behave similarly, while compounds with the same substituents in the \(\beta\)-position differ more from the \(\alpha\)- and \(\gamma\)-derivatives. Consequently, the influence of the carboxyl group is also transmitted along the chain with simple bonds with a change in direction.

An interesting fact was noted in one of our works on determining the reactivity of organohalogen compounds [24]. We drew attention to the fact that, according to the literature data, halogen derivatives of cyclohexane differ sharply, by the low activity of their halogen, from the analogous derivatives of cyclopentane and cycloheptane. In our earlier works [22] a concept had been substantiated of two mechanisms of cleavage of the carbon–halogen bond—ionic and neutral (atomic-radical):

\[ \mathrm{R:X \longrightarrow R^+ + :X;\quad R:X \longrightarrow R\cdot + X.} \]

This explained the experimentally derived division of organohalogen compounds, and of the reagents acting on them, into groups according to type of activity. It turned out that tertiary halides react more readily with alkalis and alcoholates, while primary ones react more readily with amines and iodide salts. It was suggested that alkalis and alcoholates should be regarded as reagents with a predominantly ionic type of action, whereas in the case of amines and iodide salts the atomic-radical action predominates (only predominates!) [25].

From this point of view, the different behavior of cyclic halogen derivatives with different rings can be explained by the influence of the halogen, transmitted through the ring, on the bond of that same halogen atom with carbon. An electronegative halogen makes the carbon atom bound to it, as it were, more electropositive; the neighboring carbon becomes, as it were, more electronegative, and so on. When the number of ring members is even, this influence reaches the first carbon atom with the same sign; when the number of members is odd, with the opposite sign.

Consequently, in halocyclohexanes the electropositive character of the carbon atom directly bonded to the halogen is intensified. This should facilitate ionic cleavage of the bond and hinder neutral cleavage.

Such a conclusion was also drawn with respect to tertiary halides. Experimental verification showed that, indeed, the chloro, bromo, and iodo derivatives of cyclohexane, in terms of the type of reactivity of the halogens, approach tertiary organohalogen compounds; more readily than primary and secondary halides they react with alcoholates, and less readily with amines.

Reality proved to be more complex than the considerations presented here. It has now been established that in primary halogen derivatives ionic cleavage of the bond occurs more readily, whereas in tertiary derivatives neutral cleavage occurs more readily. However, the noted fact of periodicity in the change of properties among cyclic halides retains its significance.

In general, many facts nevertheless indicate that the alternating effect, which must be regarded as having a wave origin, is manifested also in chains with simple bonds, although at times this effect is imperceptible, since it may be masked by other influences.

The concepts of an effect having an electrostatic origin and of an effect whose cause is wave interaction allow us to penetrate more deeply into the mechanism of the mutual influence of atoms and explain many facts of organic chemistry—facts that sometimes had seemed strange.

Conclusion

Interrelations among the theory of chemical structure, the electron theory, and the wave theory in elucidating the laws governing the mutual influence of atoms in molecules

In discussing the problem of the mutual influence of atoms in molecules and explaining the laws of mutual influence from the standpoint of the electron and wave theories with quantum mechanics, one must not forget, however, that these theories pertain directly not to chemistry, but to physics; they deal with the properties and laws of motion not of atoms and molecules, but of smaller particles. Atoms and molecules do indeed consist of atomic nuclei and electrons; in chemical reactions changes occur in the motion of electrons, and the motion of electrons obeys the laws of quantum mechanics. However, it cannot be concluded from this that chemistry, as a separate science, with its own theories, laws, and methods of research, is no longer needed, since the laws of chemistry are wholly reducible to quantum mechanics.

An extreme expression of the tendency to reduce chemical processes to quantum mechanics was the “theory of electronic resonance,” recently exposed by Soviet scientists. As B. M. Kedrov showed [19, 309], such pseudo-scientific constructions, created by bourgeois scientists, have a definite political aim. In the final analysis they are directed toward the denial of objective laws both in nature and in social development.

The “theory of electronic resonance” also penetrated Soviet science. Some of our scientists accepted it uncritically and attempted to use it in their work. Facts of chemistry—especially those facts which, owing to the complexity of the phenomena occurring, were difficult to explain—were interpreted from the point of view of “resonance theory.” However, the “theory,” vicious in its very foundation, naturally proved to be practically fruitless as well. What seemingly was explained only by the “resonance theory” proved possible to explain without this “theory”; and in a number of cases there was even a direct contradiction with the facts. For example, the advocates of “resonance” came to the conclusion that the nitro group in the ortho and para positions should lower the reactivity of the halogen in aromatic compounds [20], whereas organic chemists observe precisely the opposite.

Even if one discards the completely incorrect conceptions of the “theory of electronic resonance” and applies from physical theories only what has really justified itself in atomic physics, calculations based solely on the use of physical quantities—such as dipole moment, refraction, bond lengths—all these calculations do not lead to good results in solving chemical questions. They do yield something, since in chemical processes there occur displacements of atomic nuclei and changes of electronic orbits; consequently, the laws of quantum physics must operate here. Therefore certain aspects of the mutual influence of atoms find an electron-quantum explanation. The direction of chemical reactions can often be predicted qualitatively, and sometimes even quantitatively with some approximation, on the basis of physical data. But the accuracy of these predictions is considerably less than in solving physical problems, for example problems of the structure of spectra. Electronic theory has existed for 50 years already, and quantum mechanics for about 30 years, and to this day they have not given chemistry very much. Meanwhile A. M. Butlerov’s theory of chemical structure, in a shorter period of time, made it possible to unite all the factual material of organic chemistry and to make a number of predictions very important for practice. On the basis of Butlerov’s theory many important branches of the chemical industry developed.

The main reason for such a difference in the practical successes of the theories under consideration cannot be regarded as the imperfection of the mathematical apparatus of quantum mechanics. The essential point here is that each form of matter has its own laws of motion. Electronic theory and quantum mechanics speak of the motion of electrons and atomic nuclei. Atoms and molecules are composed of atomic nuclei and electrons, but they already constitute a new, qualitatively different form of matter. Therefore chemical processes, taking place directly with atoms and molecules, must have their own specific laws. A. M. Butlerov already emphasized the peculiarity of the chemical form of motion. Therefore one must not think that on the basis of the mathematical conclusions of quantum mechanics it will be possible to predict precisely the course of chemical reactions. The laws of quantum mechanics are of significance in chemistry, but they are not, for us, leading ones. Engels [3] characterizes the relation between the various forms of motion as follows:

“Every motion includes within itself mechanical motion and the displacement of larger or smallest parts of matter; to know these mechanical motions is the first task of science, but only the first. This mechanical motion by no means exhausts motion in general. Motion

“in general is not simple displacement, a simple change of place; in supra-mechanical domains it is also a change of quality.”

Chemistry too is such a supra-mechanical domain. Hence chemists must, without being satisfied with what electron-quantum physics provides, elucidate for chemical phenomena the regularities proper precisely to them, and in practical use of chemical processes be guided chiefly by these laws. And for elucidating the specific regularities of the chemical form of the motion of matter, a direct study of chemical phenomena by chemical means is absolutely necessary, although physical means of investigation must not be neglected either.

The resolution of the Conference on the Theory of Chemical Structure [19, 374—378], expressing the opinion of the overwhelming majority of Soviet organic chemists, states that “the leading direction in organic chemistry is the synthetic direction and the investigations connected with it of the structure and properties of organic substances, including natural substances, by the methods of synthetic organic chemistry...”.

“The Conference notes that the most important task of theoretical organic chemistry is the creative development of Butlerov’s doctrine, the deepening and expansion of its basic propositions in the light of modern conceptions of the structure of matter. The development of theoretical organic chemistry, and first of all the development of the theory of chemical structure, must proceed in a resolute struggle against idealistic and mechanistic theories in chemistry, on the basis of the world outlook of dialectical materialism.”

The resolution of the Conference also notes: “It is necessary to concentrate the attention of Soviet scientists on deepening and expanding knowledge of the chemical structure of matter and on developing the theory of the mutual influence of atoms in the molecule.”

Thus, in creatively developing Butlerov’s doctrine, one of the most important tasks is the further elucidation of the regularities of the mutual influence of atoms. These regularities must, of course, receive not only a qualitative explanation, but also a quantitative expression.

B. A. Kazanskii and G. V. Bykov, in the article “A. M. Butlerov and the Theory of Chemical Structure,” write [5, 585]: “A. M. Butlerov firmly believed in the future successes of science: he predicted the emergence of stereochemistry as early as 1863; he was also convinced that we would know the nature of chemical energy and that ‘general laws’ of the mutual influence of atoms in the molecule would be created.”

It is the task of Soviet chemists to discover these general laws and to use them in the practice of creating the higher technology necessary for the construction of communism.

Conclusions

1. The proposition on the mutual influence of all atoms in the molecule constitutes an integral component of the theory of chemical structure created by A. M. Butlerov. The basic idea of the mutual influence of atoms was first formulated by Butlerov and developed in detail by his pupil V. V. Markovnikov and by other outstanding scholars of the end of the 19th century, chiefly Russian organic chemists.

2. Electronic theories connect the mutual influence of atoms with the electronic structure of atoms and molecules. Modern electronic theory distinguishes two principal mechanisms of the mutual influence of atoms:

a) the effect of the electrostatic interaction of electrons and atomic nuclei; b) the conjugation effect, which cannot be explained by electrical interaction and must be connected with the wave nature of electrons.

1. The effect of electrical interaction is transmitted along a chain of atoms or directly through space, gradually weakening with increasing distance, but retaining the direction of its action. A characteristic feature of the effect of wave interaction is periodicity in the direction of action. It is transmitted better along a chain of atoms with conjugated bonds, i.e., with alternating single and double bonds, but it can also be transmitted along a chain with single bonds.

2. The opinion must be recognized as incorrect that the mutual influence of atoms, like all chemical phenomena in general, can be completely explained on the basis of the propositions of quantum mechanics. The proponents of this opinion point out that atoms and molecules consist of atomic nuclei and electrons, and that the motion of both must proceed according to the laws of quantum mechanics. From the standpoint of dialectical materialism, chemistry deals with a special form of motion, qualitatively different from other forms. Chemical phenomena are connected with the mechanical displacement of atomic nuclei and electrons, but they are not reduced to this displacement. This, and not merely the imperfection of the mathematical apparatus of quantum mechanics, explains the failures of attempts at an exact quantitative calculation of chemical processes solely by physical indices.

3. The chief task of Soviet scientists in the further development of theoretical organic chemistry must be the creative development of A. M. Butlerov’s theory of chemical structure and, in particular, of Butlerov–Markovnikov’s teaching on the mutual influence of atoms in molecules. This work must be carried out on the basis of materialist dialectics, on the basis of the brilliant works of V. I. Lenin and I. V. Stalin, using the methods of chemical experiment, but also drawing upon the achievements of modern physics. Soviet organic chemists must elucidate the general laws of the mutual influence of atoms in the molecule and use these laws in the practice of building communist society.

LITERATURE

1. V. I. Lenin, Collected Works, 4th ed., vol. 14, p. 249.

2. I. V. Stalin, On Dialectical and Historical Materialism. Questions of Leninism, 11th ed., p. 543, 1939.

3. F. Engels, Anti-Dühring, 6th ed., p. 249, 1934.

4. A. M. Berkenheim, “Electron chemistry of organic compounds.” Journal of the Russian Physico-Chemical Society, Chemical Section (ZhRKhO), vol. 49, no. 1, p. 7.

5. A. M. Butlerov, Selected Works on Organic Chemistry. Publishing House of the Academy of Sciences of the USSR, 1951.

6. A. M. Zaitsev, “On the question of the order of addition and elimination of the elements of hydrogen iodide in organic compounds,” ZhRKhO 7, 289–293, 1876.

7. D. D. Ivanenko, Introduction to the Theory of Elementary Particles. Uspekhi fizicheskikh nauk 32, no. 2, pp. 149–81, 1947.

8. A. F. Ioffe, Fundamental Concepts of Modern Physics, 1949.

9. E. Larsen, “Advances in the chemistry of zirconium and hafnium.” Uspekhi khimii 21, 824, 1952.

10. V. V. Markovnikov, Materials on the Question of the Mutual Influence of Atoms in Chemical Compounds. Kazan, 1869.

11. Ya. I. Mikhailenko, Compounds of Carbon, part one (2nd edition, 1923), part two (1921).

12. A. N. Nesmeyanov, M. I. Batuev, A. E. Borisov, “Combination light-scattering spectra of chloro-vinyl derivatives of mercury and sulfur.” Izv. AN SSSR, Division of Chemical Sciences, no. 6, pp. 557–566, 1949.

13. A. N. Nesmeyanov, A. K. Kochetkov, R. Kh. Freidlina, “Interaction of β-chlorovinyl compounds of mercury with bases.” Izv. AN SSSR, Division of Chemical Sciences, no. 6, pp. 657–659, 1947.

14. A. N. Nesmeyanov, D. N. Kursanov, K. A. Pecherskaya, E. N. Parnes, “On the influence of structural factors in phenomena of conjugation. Mobility of hydrogen atoms in vinyl ketones of various structures.” Izv. AN SSSR, Division of Chemical Sciences, no. 6, pp. 592–597, 1949.

1. Nesmeyanov A. N., Kursanov D. N., Smolina T. A., Parnes Z. N. Mobility of hydrogen atoms in acetylacetone and in the acetylacetonates of cobalt and aluminum. Izv. AN SSSR, Otd. khim. nauk, No. 6, pp. 598—600, 1949.

2. Nesmeyanov A. N., Freidlina R. Kh., Kochetkov A. K. 1,6-addition of bromo esters of mesitylmagnesium to dimethylvinylcarbinol acetate. Izv. AN SSSR, Otd. khim. nauk, No. 6, pp. 623—630, 1949.

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13. Einstein A., Infeld L. The Evolution of Physics, pp. 251—255, 1948.

1. The considerations presented concern not only electrons, but also other “elementary particles,” and also atomic nuclei. For example, comparing the effective cross sections of various atomic nuclei when neutrons are captured, and the sizes of the nuclei determined by other methods, we arrive at different conclusions about the magnitude of these nuclei. For the nucleus of the hafnium atom the effective cross section is a hundred times greater than for the nucleus of the zirconium atom [9], which in no way agrees with the ratio of the sizes of the nuclei. Evidently, a neutron experiences some kind of action from atomic nuclei even without colliding with them mechanically. The neutron has no electric charge, and one may think that this action must be wave-like. ↩