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THEORY OF CHEMICAL BONDING AND CHEMICAL STRUCTURE

SUMMARY 

In the course of the work conducted during these years, answers were given to questions, which gradually arose during the development of chemistry from alchemy to electronic chemistry. As far as the latter is concerned, back in the 1930s it was realized that the structures and transformations of chemical substances are defined by the change of the electrons' energies in the outer layer of the elements' atoms.

On the basis of the studies and conclusions made on chemical materials accumulated by Mendeleyev, Lewis, and Pauling, rules were formulated to better explain the structure of chemical compounds and their physical and chemical properties. The Table of Elements, the Lewis Rules, the Resonance Rules, and the Valence Shell Electron Pair Repulsion Rules (VSEPR) allowed foretelling the chemical and spatial structure of chemical compounds.

The discovery of a great number of chemical reactions (especially in organic chemistry) has allowed suggesting the synthesis method for chemical compounds. In the beginning of the 19th century the scientific basis for chemistry was formed: the atom-molecular theory of substance structure and the theory of chemical transformation.

In the framework of the structure theory, it was found that chemical compounds (molecules) consist of atoms, which are connected to chemical bonds. As far as the theory of organic compound structure is concerned, rules were formulated (first of all the four-valence carbon rule) which allowed foretelling the structure of organic compounds. The Lewis Rules and the Resonance Rules allowed foretelling the structure of both organic and inorganic compounds.

Most urgent was the general question concerning the physical nature of chemical bonding. Also urgent were the questions concerning the elucidation of the physical nature of the above mentioned rules (Periodic Law, Lewis Rules, Resonance Rules, etc.).

The following questions, concerning the theory, were unclear: Why is it that two electrons (one from each atom) and not one or three take part in the formation of the bond?  Why is it that atoms taking part in chemical bonding strive to build their shells to the level of those of the inert (noble) gas? Why is there an exception to this rule? Etc.

Such questions arose in the course of the development of traditional (classical) chemistry as a science. In the 1930s the traditional development of chemistry was exchanged for quantum chemistry, which, in the 1980s, proved to be a false path in the development of science. At the beginning of the 1980s we continued the traditional path directed at deepening the understanding of the main chemical phenomena (i.e., chemical bonding and chemical reactions).

The improvement of the knowledge of the physical nature of chemical bonding involved a shift from the conclusion that during chemical bond formation the system's energy decreases to the explanation why the system's energy decreases (qualitative explanation of the enthalpy factor) and by how much the system's energy actually decreases (quantitative evaluation of the enthalpy factor for homoatomic and heteroatomic molecules).

We have managed to realize why more energy is required for breaking a chemical bond as compared to the difference in electronic energies of atoms and molecules (qualitative evaluation of the entropy factor). We have also managed to quantitatively evaluate the influence of the entropy factor on the chemical bond-breaking energy. The evaluation of the enthalpy and entropy factors has allowed us to compile a system of three algebra equations with three unknowns (which we will further refer to as a system of equations), on the basis of a model which presupposes that covalent (homo- and heteroatomic) bonding is formed via electrons which rotate in one circle whose plane is perpendicular to the axis connecting the nuclei. Upon deducing the system's equations, it was supposed that electrons were particles with a definite mass, a charge, and an orbit speed.

The solution of the system of equations, with regard to a hydrogen molecule, has shown that the value of a hydrogen molecule's electronic energy differs from that received via the experiment by less than 4% which proves the correctness of the model. This system of equations has helped to deepen the understanding of the main questions raised in the course of the development of the theory of chemical substance structure.

The solution of this system of equations has allowed realizing the physical essence of chemical bonding and the main regularities observed when studying the latter, which are given in the form of the Lewis Rules and as additions to these rules, which explained the exclusions concerning the Lewis Rules (connections with surplus, the resonance rules, etc.) For example, the system of equations has allowed calculating the optimal number of bonding electrons and the radius of the circle in which the bonding electrons rotate.

This system has allowed elucidating the functional dependence of bonding energy in dual-atom molecules on the first ionization energy (FIE) of the bonding atoms and on the number of bonding electrons. It was found that when the FIEs of the bonding atoms are equal, a parabola describes the dependence of the bonding energy on the FIE. If the FIEs of the atoms differ, the bonding energy increases proportionally to the difference of the atoms' FIEs.

The solution of the system of equations has made it possible to calculate the dependence of the following on the FIE of the bonding atoms: bond lengths, bond polarity, radius of the circle where the bonding electrons rotate, and the radius of the circle where nonbonding electrons rotate.  All the calculated dependencies coincided with those received as a result of treating the available experimental data on the FIE of atoms, bonding energies, bonding lengths, and bond polarity.     

According to the model whose correctness has been proven by comparing the calculated and experimental data when one atom is bonded to several atoms, covalent chemical bonds should have a definite direction. The corners between the bonds should be defined by the repulsion between the circles in which the bonding electrons rotate.

The calculation of the radius of the circle where the nonbonding electrons rotate has shown that this radius is greater than the radius of the bonding electrons' circle. According to this model, the repulsion between the electronic pairs increases proportionally to the increase of the radius of the circle where the electrons rotate. That is, the interelectronic repulsion should, according to the calculation, increase in the row: nonbonding pairs —nonbonding pairs > nonbonding pairs — bonding pairs > bonding pairs — bonding pairs. Such a sequence is observed in the experiment summarized by the VSEPR rules.

On the background of a satisfactory coincidence in the calculated and experimental dependencies, for most of the elements there were great discrepancies between the calculated and experimental data for elements, which were called anomalous.

The reason for these discrepancies is explained as follows: The anomalously small bonding energy, as compared to the model calculation, had dual-atomic covalent bonds in a number of noble gases and elements with two electrons on the atoms' outer electronic shell. It was found that the anomalously small bonding energy in molecules formed out of anomalous elements is connected with the atom's anomalous properties of these elements, which is obvious in the fact that unlike normal elements, the anomalous ones had negative values of affinity towards the electron.

As a result of solving the system of equations it was found that during the formation of a covalent bond, both bonding electrons enter the outer (previously existing) electronic shells of the bonding atoms. That is, in the process of covalent bond formation, the number of electrons in the outer shells of the bonding atoms increases by one unit. The coincidence of the rows of anomalous atoms (with a negative affinity towards the electron) and the rows of atoms with an anomalously small covalent bonding energy in dual-atomic molecules formed of these atoms, has shown that the anomalously small bonding energy is due to the additional energy expenditure connected with the entrance of the bonding electron to the outer layer of the atom being bonded.

We realized the anomalous behavior of atoms: inter-electronic repulsion in the outer electronic layers. The mutual repulsion of electrons in one layer causes their distribution in other layers. The number of electrons on the inner layers is the same in all atoms; that is why, when the nuclear charge increases, the number of electrons in the atom's outermost layer changes periodically. According to electrostatics, the FIE of the atoms and their affinity towards electrons also change periodically. These three values thus define the physical and chemical properties of elements: valence, bonding energy, and reaction possibilities.

That is, the elucidation of the physical nature of the elements' anomalous behavior helps to understand the following:

   1) the physical nature of valence;

   2) the connection between the FIE, the affinity of atoms towards electrons, the number of electrons in the outermost layer with the physical and chemical properties of elements, and the cause of their periodic change;

   3) the physical nature of the Periodic Law.

We have deepened the understanding of the Lewis Rules.  

Instead of phrases like: "A molecule is composed of atoms that are bound together by sharing pairs of electrons using the atomic orbitals of the bound atoms..." and "The most important requirement for the formation of a stable compound is that the atoms achieve noble gas and electronic configuration....", the following wording has been suggested:          

"According to the solution of the system of equations, it has become clear that:  

   1) the minimal electronic energy (during the formation of two atomic molecules out of atoms) is achieved when two electrons (not one or three) rotate between connected nuclei; 

   2) during the formation of a covalent bond the number of electrons in the outermost shell of the atoms being bonded is increased by 1." 

The interelectronic repulsion (between the electrons of the outermost layer) limits the number of electrons that can enter the outermost layer of the atom when chemical bonding occurs. That is why the number of bonds that an atom can form with other atoms is limited. The number of electrons in the outermost layer, plus those added during bond formation, cannot exceed 8 — for the first 20 elements in the Mendeleyev Table of Elements. 

The improvement of the Resonance Rules included the elucidation of the physical nature of the phenomena described by these rules. It became known that the effects, described by the Resonance Rules, are conditioned by electron-nuclear isomerization.

We have managed to answer many special questions like:

Why are inert (noble) gases inert? Why is bonding energy in F2 smaller than in Cl2?  Why is F2  a more active oxidizer than Cl2? Why is mercury, being a metal, always in liquid form at room temperature? Why is it that during molecule formation the atoms construct their shells up to the shells of inert (noble) gases?