The fundamental basis of the chemical bond was theory of chemical structure of A. M. Butlerov(1861), according to which the properties of compounds depend on the nature and number of their constituent particles and chemical structure. This theory has been confirmed not only for organic, but also for inorganic substances, so it should be considered the fundamental theory of chemistry.

The concept of a chemical bond

Most simple substances and all complex substances (compounds) consist of atoms interacting with each other in a certain way. In other words, a chemical bond is established between the atoms.

chemical bond- an electronic phenomenon, which consists in the fact that at least one electron, which was in the force field of its nucleus, finds itself in the force field of another nucleus or several nuclei at the same time. When a chemical bond is formed, energy is always released, i.e. the energy of the resulting particle must be less than the total energy of the initial particles.

Chemical bond means different kinds interactions that cause the stable existence of two- and polyatomic compounds: molecules, ions, crystalline and other substances.

The main features of a chemical bond include:

• reduction in the total energy of a two- or polyatomic system compared to the total energy of isolated particles from which this system is formed;
• redistribution of the electron density in the region of a chemical bond compared to a simple superposition of the electron densities of unbonded atoms brought together by a bond distance.

By its nature, a chemical bond is an interaction between positively charged nuclei and negatively charged electrons, as well as electrons with each other.

The transition of an electron from one atom to another, resulting in the formation of oppositely charged ions with stable electronic configurations, between which an electrostatic attraction is established, is the simplest model of ionic bonding:

X ⇒ X++e- Y+e- ⇒ Y- X+Y-

Chemical bond parameters

The chemical bond is carried out s- and p- electrons of external and d- electrons in the outer layer. This connection is characterized by the following parameters:

1. Bond length - the internuclear distance between two chemically bonded atoms.
2. Valence angle - the angle between imaginary lines passing through the centers of chemically bonded atoms.
3. Bond energy - the amount of energy expended on breaking it in the gaseous state.
4. The multiplicity of bonds - the number of electron pairs through which a chemical bond between atoms is carried out.

If we bring two protons together, then repulsive forces will arise between them, and there is no need to talk about obtaining a stable system. Let's place one electron in their field. Two cases may arise here.

The first, when the electron is between the protons (left), and the second, when it is located behind one of them (right).

In both cases, attractive forces arise. In the first case, the components of these forces (projections) on the axis passing through the proton centers are directed in opposite directions with repulsive forces (left) and can compensate them. This creates an energetically stable system. In the second case, the components of the attractive forces are directed in different directions (on the right) and it is difficult to talk about balancing the repulsive forces between the protons. It follows that for the occurrence of a chemical bond with the formation of a molecule or an ion, the electrons must be predominantly in the internuclear space. This area is called binding, because When electrons are present, a chemical bond is formed. The area behind the nuclei is called loosening, because when electrons enter it, no chemical bond is formed.

Applying similar reasoning to the hydrogen molecule, we can conclude that the appearance of a second electron in the binding region stabilizes the system even more. Therefore, at least one electron pair is required to form a stable chemical bond. The electron spins in this case must be antiparallel, i.e. directed in different directions. The formation of a chemical bond must be accompanied by a decrease in the total energy of the system.

Consider the change in the potential energy of the system on the example of the approach of two hydrogen atoms. When the atoms are at a very large distance from each other, they do not interact and the energy of such a system is close to zero. As they approach, attractive forces arise between the electron of one atom and the nucleus of another, and vice versa. These forces increase inversely with the square of the distance between the atoms. The energy of the system is reduced. As the atoms approach each other, the repulsive force between their nuclei and electrons begins to play the role. The increase in repulsive forces is inversely proportional to the sixth power of the distance. The potential energy curve passes through a minimum, and then sharply goes up.

The distance corresponding to the position of the minimum on the curve is the equilibrium internuclear distance and determines the length of the chemical bond. Since the atoms in a molecule are involved in oscillatory motion relative to the equilibrium position, the distance between them is constantly changing, i.e. atoms are not rigidly bonded to each other. The equilibrium distance corresponds at a given temperature to some average value. As the temperature rises, the oscillation amplitude increases. At some sufficiently high temperature, atoms can fly apart to an infinitely large distance from each other, which will correspond to the breaking of a chemical bond. The depth of the minimum along the energy axis determines the chemical bond energy, and the value of this energy, taken with the opposite sign, will be equal to the dissociation energy of a given diatomic particle. If hydrogen atoms approach each other, the electrons of which have parallel spins, only repulsive forces arise between the atoms, and the potential energy of such a system will increase.

The amount of energy released when a chemical bond is formed is called chemical bond energy E St. It has a unit of measure [kJ/mol]. For polyatomic compounds with bonds of the same type, its average value is taken as the bond energy, calculated by dividing the energy of formation of a compound from atoms by the number of bonds. For example, the binding energy in methane is determined by dividing the energy of formation of the molecule CH 4 from hydrogen and carbon atoms into four (1640: 4 = 410 kJ / mol). The greater the energy of a chemical bond, the more stable the molecules. For example, a molecule HF more stable than a molecule HBr.

An important characteristic of a chemical bond is its length. l sv, equal to the distance between the nuclei in the compound. It depends on the size of the electron shells and the degree of their overlap. There is a definite correlation between bond length and bond energy: with a decrease in the bond length, the binding energy usually increases and, accordingly, the stability of the molecules. For example, in the series of hydrogen halides from HF before HI The bond length increases and its energy decreases.

Energies and lengths of some chemical bonds

Connection	E St, kJ/mol	l sv, nm	Connection	E St, kJ/mol	l sv, nm	Connection	E St, kJ/mol	l sv, nm	Connection	E St, kJ/mol	l sv, nm
	536	0,092		348	0,154		432	0,128		614	0,134
	360	0,142		495	0,121		299	0,162		839	0,120
	436	0,074		1040	0,113		380	0,134		940	0,110

Types of chemical bond

involved in the formation of a chemical bond s-, R- and d-electrons having different geometrical configuration of electron clouds and different signs of wave functions in space. For the occurrence of a chemical bond, it is necessary to overlap the parts of the electron shells with the same sign of the wave function. Otherwise, no chemical bond is formed. This statement can be easily explained by the example of the superposition of two sinusoids, which, in the first approximation, can be identified with wave functions.

In the case of superposition of two sinusoids with different signs in the same area (on the left), their total component will be equal to zero - there is no connection. In the opposite case, the oscillation amplitudes are added and a new sinusoid is formed - a chemical bond has formed (on the right).

Depending on the symmetry of the electron clouds, as a result of the overlap of which a chemical bond is formed, the total electron cloud will have a different symmetry, according to which they fall into three types: σ -, π - and δ -connections.

σ bond is carried out when clouds overlap along the line connecting the centers of atoms, while the maximum electron density is achieved in the internuclear space and has cylindrical symmetry relative to the line connecting the centers of atoms. In education σ -bonds, due to their spherical symmetry, always take part s-electrons. They form σ - bond as a result of overlap with the following electrons of another atom: s-, p x-, d x 2 -y 2-electrons. With electrons in other orbitals, for example, RU or p x, the occurrence of a chemical bond is impossible, since overlapping occurs in areas where the electron density has opposite signs. Education Opportunity σ - communications s-electrons are not exhausted, it can be formed in case of overlapping of other electron clouds, such as two p x or p x and d x 2 -y 2.

π bonds arise when electron clouds overlap above and below the line connecting the centers of atoms. The total electron clouds are also symmetrical about this axis, but they do not have cylindrical symmetry, as in the case σ -connections. Due to its spatial location π -bond is formed by electrons on such pairs of orbitals as p y -p y,p z -p z,p y -d xy.

δ bond form only d-electrons due to the overlapping of all four of its petals of electron clouds located in parallel planes. This is possible when dxy-dxy, dxz-dxz, dyz-dyz-electrons.

There is another approach to the classification of a chemical bond based on the nature of the electron density distribution between atoms in a molecule, i.e. a chemical bond is considered from the point of view of belonging of an electron pair to one or another atom. Three cases are possible:

1. An electron pair links two identical atoms in a molecule. In this case, it belongs equally to both of them. In a molecule there is no separation of the centers of gravity of positive and negative charges. They coincide, and such a relationship is called covalent non-polar.
2. If an electron pair binds two different atoms, then it shifts towards a more electronegative atom. The centers of gravity of the positive and negative charges are separated, the bond becomes polar and is called covalent polar bond.
3. The third case is connected with the complete transfer of the electron pair into the possession of one of the atoms. This happens during the interaction of two atoms that differ sharply in electronegativity, i.e. the ability to hold an electron pair in its electric field. In this case, the atom that donated electrons becomes a positively charged ion, and the atom that accepted them becomes negative. In this case, the relationship is called ionic.

A chemical bond arises due to the interaction of electric fields created by electrons and nuclei of atoms, i.e. the chemical bond is electrical in nature.

Under chemical bond understand the result of the interaction of 2 or more atoms leading to the formation of a stable polyatomic system. The condition for the formation of a chemical bond is a decrease in the energy of the interacting atoms, i.e. the molecular state of matter is energetically more favorable than the atomic state. When a chemical bond is formed, atoms tend to obtain a complete electron shell.

There are: covalent, ionic, metallic, hydrogen and intermolecular.

covalent bond- the most general type of chemical bond that arises due to the socialization of an electron pair through exchange mechanism -, when each of the interacting atoms supplies one electron, or donor-acceptor mechanism, if an electron pair is transferred for common use by one atom (donor - N, O, Cl, F) to another atom (acceptor - atoms of d-elements).

Chemical bond characteristics.

1 - multiplicity of bonds - only 1 sigma bond is possible between 2 atoms, but along with it, there can be pi and delta bonds between the same atoms, which leads to the formation of multiple bonds. The multiplicity is determined by the number of common electron pairs.

2 - bond length - the internuclear distance in the molecule, the greater the multiplicity, the smaller its length.

3 - bond strength - this is the amount of energy required to break it

4 - saturation of the covalent bond is manifested in the fact that one atomic orbital can take part in the formation of only one c.s. This property determines the stoichiometry of molecular compounds.

5 - directivity of the c.s. Depending on the shape and direction of electron clouds in space, when they overlap, compounds with linear and angular molecular shapes can be formed.

Ionic bond – formed between atoms that are very different in electronegativity. These are compounds of the main subgroups of groups 1 and 2 with elements of the main subgroups of groups 6 and 7. Ionic is a chemical bond, which is carried out as a result of mutual electrostatic attraction of oppositely charged ions.

The mechanism of formation of ionic bonds: a) the formation of ions of interacting atoms; b) the formation of a molecule due to the attraction of ions.

Non-directionality and unsaturation of the ionic bond

The force fields of the ions are evenly distributed in all directions, so each ion can attract ions of the opposite sign in any direction. This is the non-directionality of the ionic bond. The interaction of 2 ions of the opposite sign does not lead to complete mutual compensation of their force fields. Therefore, they retain the ability to attract ions in other directions as well, i.e. an ionic bond is characterized by unsaturation. Therefore, each ion in an ionic compound attracts such a number of ions of the opposite sign that an ionic-type crystal lattice is formed. There are no molecules in an ionic crystal. Each ion is surrounded by a certain number of ions of a different sign (coordination number of the ion).

metal connection- chem. Communication in metals. Metals have an excess of valence orbitals and a lack of electrons. When atoms approach each other, their valence orbitals overlap, due to which electrons move freely from one orbital to another, and a connection is made between all metal atoms. The bond that is carried out by relatively free electrons between metal ions in a crystal lattice is called a metallic bond. The connection is strongly delocalized and does not have directionality and saturation, because valence electrons are evenly distributed throughout the crystal. The presence of free electrons determines the existence of common properties of metals: opacity, metallic luster, high electrical and thermal conductivity, malleability and plasticity.

hydrogen bond– bond between the H atom and a strongly negative element (F, Cl, N, O, S). Hydrogen bonds can be intra- and intermolecular. BC is weaker than a covalent bond. The emergence of VS is explained by the action of electrostatic forces. The H atom has a small radius and, when a single electron H is displaced or donated, it acquires a strong positive charge, which affects the electronegativity.

Covalent bond - it is a bond between two atoms through the formation of a common electron pair.

Covalent non-polar bond– this bond between atoms with equal

electronegativity. For example: H 2, O 2, N 2, Cl 2, etc. The dipole moment of such bonds is zero.

covalent polar bond – this bond is between atoms with different electronegativity. The electron cloud overlap zone shifts towards the more electronegative atom.

For example, H–Cl (H b+ →Cl b–).

A covalent bond has the following properties:

- saturation - the ability of an atom to form the number of chemical bonds corresponding to its valency;

- directionality - the overlap of electron clouds occurs in the direction that provides the maximum overlap density.

Ionic bond– it is a bond between oppositely charged ions. It can be considered as an extreme case of a covalent polar bond. Such a bond occurs when there is a large difference in the electronegativity of atoms,

forming a chemical bond. For example, in the NaF molecule, the difference

electronegativity is 4.0 – 0.93 \u003d 3.07, which leads to an almost complete transition of an electron from sodium to fluorine:

The interaction of ions of the opposite sign does not depend on the direction, and the Coulomb forces do not have the saturation property. Because of this, the ionic bond does not have directionality and saturation.

metal connection– is the bonding of positively charged metal ions with free electrons.

Most metals have a number of properties that have general character and different from the properties of other substances. These properties are relatively high temperatures melting, the ability to reflect light, high thermal and electrical conductivity. This is a consequence of the formation of a special type of bond between metal atoms - a metallic bond.

In metal atoms, valence electrons are weakly bound to their nuclei and can easily be detached from them. As a result, positively charged metal ions and "free" electrons appear in the crystal lattice of the metal, the electrostatic interaction of which provides a chemical bond.

hydrogen bond– is a bond through a hydrogen atom bonded to a highly electronegative element.

A hydrogen atom bound to a highly electronegative element (fluorine, oxygen, nitrogen, etc.) gives up almost completely an electron from the valence orbital. The resulting free orbital can interact with the lone pair of electrons of another electronegative atom, resulting in a hydrogen bond. Using the example of water and acetic acid molecules, the hydrogen bond is shown by dashed lines:

This bond is much weaker than other chemical bonds (the energy of its formation is 10÷40 kJ/mol). Hydrogen bonds can occur both between different molecules and within a molecule.

Exclusively important role the hydrogen bond plays in such inorganic substances as water, hydrofluoric acid, ammonia, etc., as well as in biological macromolecules.

Any interaction between atoms is possible only in the presence of a chemical bond. Such a connection is the reason for the formation of a stable polyatomic system - a molecular ion, a molecule, a crystal lattice. A strong chemical bond requires a lot of energy to break, which is why it is the base value for measuring bond strength.

Conditions for the formation of a chemical bond

The formation of a chemical bond is always accompanied by the release of energy. This process occurs due to a decrease in the potential energy of a system of interacting particles - molecules, ions, atoms. The potential energy of the resulting system of interacting elements is always less than the energy of unbound outgoing particles. Thus, the basis for the occurrence of a chemical bond in the system is the decline in the potential energy of its elements.

The nature of the chemical interaction

A chemical bond is a consequence of the interaction of electromagnetic fields that arise around the electrons and nuclei of atoms of those substances that take part in the formation of a new molecule or crystal. After the discovery of the theory of the structure of the atom, the nature of this interaction became more accessible for study.

For the first time, the idea of ​​​​the electrical nature of a chemical bond arose from the English physicist G. Davy, who suggested that molecules are formed due to the electrical attraction of oppositely charged particles. This idea interested the Swedish chemist and naturalist I.Ya. Berzellius, who developed the electrochemical theory of the formation of a chemical bond.

The first theory, which explained the processes of chemical interaction of substances, was imperfect, and over time it had to be abandoned.

Butlerov's theory

A more successful attempt to explain the nature of the chemical bond of substances was made by the Russian scientist A.M. Butlerov. This scientist based his theory on the following assumptions:

• Atoms in the connected state are connected to each other in a certain order. A change in this order causes the formation of a new substance.
• Atoms bind to each other according to the laws of valency.
• The properties of a substance depend on the order of connection of atoms in a molecule of a substance. A different arrangement causes a change in the chemical properties of the substance.
• Atoms bonded together have the strongest influence on each other.

Butlerov's theory explained the properties chemical substances not only by their composition, but also by the arrangement of atoms. Such an internal order of A.M. Butlerov called "chemical structure".

The theory of the Russian scientist made it possible to put things in order in the classification of substances and made it possible to determine the structure of molecules by their chemical properties. The theory also gave an answer to the question: why molecules containing the same number of atoms have different chemical properties.

Prerequisites for the creation of chemical bond theories

In his theory of the chemical structure, Butlerov did not touch on the question of what a chemical bond is. There was too little data for this at the time. internal structure substances. Only after the discovery of the planetary model of the atom, the American scientist Lewis began to develop a hypothesis that a chemical bond arises through the formation of an electron pair, which simultaneously belongs to two atoms. Subsequently, this idea became the foundation for the development of the theory of covalent bonds.

covalent chemical bond

sustainable chemical compound can be formed when the electron clouds of two neighboring atoms overlap. The result of such mutual crossing is an increasing electron density in the internuclear space. The nuclei of atoms, as you know, are positively charged, and therefore they try to be attracted as close as possible to the negatively charged electron cloud. This attraction is much stronger than the repulsive forces between two positively charged nuclei, so this bond is stable.

The first chemical bond calculations were performed by the chemists Heitler and London. They considered the bond between two hydrogen atoms. The simplest visual representation of it might look like this:

As can be seen, the electron pair occupies a quantum place in both hydrogen atoms. This two-center arrangement of electrons is called a "covalent chemical bond". A covalent bond is typical for molecules of simple substances and their compounds of non-metals. Substances created as a result of a covalent bond usually do not conduct electricity or are semiconductors.

Ionic bond

An ionic type chemical bond occurs when two oppositely charged ions are attracted electrically. Ions can be simple, consisting of one atom of a substance. In compounds of this type, simple ions are most often positively charged atoms of metals of the 1.2 group that have lost their electron. The formation of negative ions is inherent in the atoms of typical non-metals and bases of their acids. Therefore, among the typical ionic compounds, there are many alkali metal halides, such as CsF, NaCl, and others.

Unlike a covalent bond, an ion does not have saturation: a different number of oppositely charged ions can join an ion or group of ions. The number of attached particles is limited only by the linear dimensions of the interacting ions, as well as the condition under which the attractive forces of oppositely charged ions must be greater than the repulsive forces of identically charged particles participating in an ionic type connection.

hydrogen bond

Even before the creation of the theory of chemical structure, it was empirically observed that hydrogen compounds with various non-metals have several unusual properties. For example, the boiling points of hydrogen fluoride and water are much higher than might be expected.

These and other features of hydrogen compounds can be explained by the ability of the H + atom to form another chemical bond. This type of connection is called a "hydrogen bond". The causes of hydrogen bonding lie in the properties of electrostatic forces. For example, in a hydrogen fluoride molecule, the general electron cloud is so shifted towards fluorine that the space around the atom of this substance is saturated with negative electric field. Around the hydrogen atom, deprived of its only electron, the field is much weaker and has a positive charge. As a result, there is an additional relationship between the positive fields of electron clouds H + and negative F - .

Chemical bonding of metals

The atoms of all metals are located in space in a certain way. The arrangement of metal atoms is called the crystal lattice. In this case, the electrons of different atoms weakly interact with each other, forming a common electron cloud. This type of interaction between atoms and electrons is called a "metal bond".

It is the free movement of electrons in metals that can explain the physical properties of metallic substances: electrical conductivity, thermal conductivity, strength, fusibility, and others.

chemical bond - this is the interaction of two atoms, carried out in the process of redistribution of electrons of valence orbitals, resulting in a stable eight- or two-electron configuration of the nearest noble gas (octet or doublet) due to the formation of ions (W. Kossel) or the formation of common electron pairs (G. Lewis ). As a result, the total energy of the system decreases.

3.1 Main characteristics of the chemical bond

3.1.1 Bond energy is the energy required to break a chemical bond in all molecules that make up one mole of a substance, or the energy gain in the formation of a compound from individual atoms ( E St.). The greater the energy of a chemical bond, the stronger the bond itself, the more stable the molecule.

Bond energy is usually measured in kilojoules per mole. kJ/mol.

kJ/mol . For polyatomic compounds with bonds of the same type, its average value is taken as the bond energy, calculated by dividing the energy of formation of a compound from atoms by the number of bonds. Thus, 432.1 kJ/∙mol is spent on breaking the H–H bond, and 1648 kJ/∙mol is spent on the atomization of methane CH 4, in this case E C–H = 1648: 4 = 412 kJ.

Ionic and covalent bonds are the strongest. , whose energies range from tens to hundreds of kJ/mol. The metallic bond, as a rule, is somewhat weaker than ionic and covalent bonds, but the values ​​of the binding energies in metals are close to the values ​​of the energies of ionic and covalent bonds. The energy of a hydrogen bond is very small and usually amounts to 20-40 kJ/mol, while the energy of covalent bonds can reach several hundred kilojoules per mole, kJ/mol.

3.1.2 Link length l St. . When a chemical bond is formed, overlap electron clouds of two atoms and the distance between the nuclei of atoms becomes less than the sum of the distances from the nuclei to the outer zones of the highest electron density in atoms.

Link length is equal to the distance between the nuclei of the interacting atoms in the compound. It is measured in nanometers, nm, or angstroms, A (1A = 10 -8 cm). It depends on the size of the electron shells and the degree of their overlap. There is a definite correlation between bond length and bond energy: with a decrease in the bond length, the binding energy usually increases and, accordingly, the resistance of molecules to decay or the action of other substances.

3.1.3 Communication polarity characterized by an ionic component, that is, a shift of an electron pair to a more electronegative atom, resulting in the formation of a dipole. Dipole - a system of two equal, but opposite in sign charges, located at a unit distance from each other. The polarity of a bond can be expressed in terms of its dipole moment μ , equal to the product of the elementary charge and the length of the dipole μ= el. Dipole moment is measured in coulombs per meter, C⋅m, or debyes, D.

1D = 0.333∙10 -29 C∙m. It is a vector quantity and is directed along the dipole axis from a negative charge to a positive one.

Molecule polarity in general, it is determined by the difference in the electronegativity of the atoms that form a two-center bond, the geometry of the molecule, as well as the presence of unshared electron pairs, since part of the electron density in the molecule can be localized not in the direction of the bonds. She expresses herself through her dipole moment, which is equal to the vector sum of all dipole moments of the bonds of the molecule .

It is necessary to distinguish between the dipole moments (polarity) of the bond and the molecule as a whole. For example, for a linear CO 2 molecule, μ \u003d 0 (although each of the bonds is polar, and the molecule as a whole is non-polar, since the O \u003d C \u003d O molecule is linear, and the dipole moments of the C \u003d O bonds compensate each other), but for H 2 O μ ≠ 0. The presence of a dipole moment in a water molecule means that it is non-linear, i.e. O-N connections located at an angle not equal to 180°.

3.1.4 Spatial structure of molecules - this is the shape and location in space of electron clouds, taking into account the nature of the chemical bond.

In compounds containing more than two atoms, an important characteristic is the bond angle formed by chemical bonds in the molecule and reflecting its geometry.

3.1.5 Link order (link multiplicity) is the number of shared shared pairs between two bonded atoms. The higher the bond order, the stronger the bonds between the atoms and the shorter the bond itself. The order of communication above three does not occur. For example, the bond order in H 2 , O 2 and N 2 molecules is 1, 2 and 3, respectively, since the bond in these cases is formed due to the overlap of one, two and three pairs of electron clouds.

3.1.6 Bond saturation - the ability of an atom to give a certain number of chemical bonds. Some types of chemical interaction do not have saturation, that is, particles can form a different number of bonds with their neighbors. This property is inherent in ionic bonds.

3.1.7 How electron clouds overlap. According to the method of overlapping electronic clouds, communication is divided into σ - communication and π – communication (Figure 4).

Figure 4 - Scheme of σ - and π - bonds

σ - the connection is formed due to the overlapping of electron clouds along the line connecting the centers of interacting atoms. It can be carried out, for example, between two s-clouds, between two p-clouds, between s- and p-clouds, or between s- and d-clouds. π - the connection is formed due to the overlap of electron clouds on both sides of the line connecting the centers of interacting atoms (due to the lateral overlap of electron clouds). It is formed mainly by overlapping p-orbitals. σ - bond is stronger than π - connection, since it provides a more complete overlap and therefore more energy is required to break it.

Theories explaining the chemical bond

Two theories are currently in use: valence bond method (MVS) and molecular orbital method (MMO) .

3.2.1 Valence bond method otherwise called the theory of localized electron pairs, since the method is based on the assumption that during the formation of a molecule, atoms retain their atomic orbitals, but an increased electron density (common electron pair) is formed, which belongs to both atoms. In contrast to MMO, in which the simplest chemical bond can be both two- and multicenter, in MVS it is always two-electron and necessarily two-center.

Note that, according to the Pauli principle, electrons must have oppositely directed spins, that is, in the MVS all spins are paired, and all molecules must be diamagnetic (since the magnetic properties are determined by the presence of free electrons). Consequently, the MVS fundamentally cannot explain the magnetic properties of molecules.

3.2.2 Molecular orbital method proceeds from the fact that each molecular orbital is represented as an algebraic sum (linear combination) of atomic orbitals. That is, when a molecule is formed, atomic orbitals as such disappear, and new molecular orbitals appear instead. Moreover, the number of molecular orbitals is equal to the sum of the initial atomic orbitals, but some of the molecular orbitals are lower in energy (bonding MO), and some are higher in energy (loosening MO) than the initial atomic orbitals.

For example, in a hydrogen molecule, only 1s atomic orbitals of two hydrogen atoms can participate in the formation of an MO, which give two MOs. Since the nuclei in the interacting hydrogen atoms are the same, the contribution of the atomic orbitals will be the same. And since interaction in a hydrogen molecule is possible only along the axis of the molecule, each of the MOs can be redesignated as σb and σ* and named, respectively, bonding (σb) and loosening (σ*) molecular orbitals.

The transition of two electrons to MO σ st contributes to a decrease in the energy of the system; this energy gain is equal to the binding energy between atoms in the H–H hydrogen molecule. That is, the population of the bonding MO σb with electrons stabilizes the system, and the population of loosening MOs destabilizes it.

According to the MO method order (multiplicity) of communication ndetermined by the half-difference of the number of binding N St. and loosening N once electrons

The greater the bond multiplicity, the stronger the bond in the molecule. At zero multiplicity of bonds, a molecule is not formed.

Let us consider some cases of the structure of molecules according to MMO.

The H 2 molecule is formed from two H atoms, the atomic valence band of which is represented by one electron at the 1s sublevel. What is necessary for the electrons to have opposite spins. Let's depict the structure of the H 2 molecule in the following energy diagram (Figure 5). When filling out this diagram, one should take into account the principles of filling electron orbitals (Pauli's principle, Hund's principle, minimum energy principle).

Since the orbitals with lower energy are filled first, the electrons will go to the bonding orbital. And since, according to the Pauli principle, only two of them can fit there, there are no more electrons left for the loosening orbital. The multiplicity of the connection in this case will be equal to

n \u003d (2-0) / 2 \u003d 1, that is, the connection is carried out by one pair of electrons.

As for the molecular ion H 2 +, it is formed from the atom H and the proton H +, which does not have a valence electron (Figure 6).

The multiplicity of the bond in this case will be equal to n = (1-0)/2 = 1/2, that is, the bond is carried out by one electron. The MVS cannot explain such a phenomenon.

The molecular ion H 2 - will have a similar structure, which is formed from the H atom and the atomic ion H -, which has one extra valence electron (two electrons go to the bonding orbital, and one to the loosening orbital). The multiplicity of the bond in this case will be equal to n = (2-1)/2 = 1/2, that is, the bond is also carried out by one electron. From the foregoing, it follows that the H 2 molecule will be stronger than molecular ions, since it is characterized by a higher bond multiplicity; and the molecular ions H 2 + and H 2 - are absolutely identical in strength.

MMO can be used to explain the absence of noble gas molecules. Let us consider this using the example of the He 2 molecule, which should have been formed from two He atoms with AID 1s 2 (Figure 7).

In this case, the bond multiplicity will be equal to n = (2-2)/2 = 0, that is, the bond is not carried out, since there is no common electron pair and there is no gain in energy. And under such conditions, the molecule is not formed.

Hydrogen and helium are followed by elements that have a more complex structure of the electron shell, therefore, the molecules formed by these substances will have the corresponding structure.

Consider this on the example of the O 2 molecule (Figure 8). It is formed from two oxygen atoms with ABZ 2s 2 2р 4 . Since s-orbitals have less energy than p-, they will be located lower on the energy diagram. Notice again that the valence electrons of different atoms have antiparallel spins. Since each of the atoms provides one s-orbital and three p-orbitals, the total number of molecular orbitals will be eight. From two atomic s-orbitals, two molecular orbitals are formed: σ s - bonding and σ s - loosening (since the overlap of electron clouds occurs along the line connecting the centers of atoms).

The six p orbitals form three bonding and three antibonding orbitals. Due to the fact that one pair of electron clouds overlaps along the line connecting the centers of atoms, σр– bonding and σр– loosening molecular orbitals are formed between them. Between the remaining two pairs of p-orbitals, lateral overlap will be observed, therefore, two p-bonding orbitals, identical in energy, and two p-loosening orbitals, also identical in energy, are formed. On the energy diagram, σ p - the bonding orbital is located below the p- binding ones, since more energy is released during the formation of the σ- bond, therefore the bond itself will have a lower energy reserve. The filling of the diagram with electrons (12 electrons for both atoms) occurs as the energy of molecular orbitals increases (from bottom to top). On loosening orbitals, two electrons are in different cells, since in this case the total spin of the system will be maximum. The multiplicity of the connection in this case will be equal to n \u003d (8-4) / 2 \u003d 2, that is, the connection is also carried out by two pairs of electrons.

The structure of the molecular ion O 2 + or O 2 2+ will differ in the absence of one or two electrons, respectively, in loosening orbitals. In this case, the multiplicity of communication will increase ( n = (8-3) / 2 = 2.5 and n \u003d (8-2) / 2 \u003d 3) and the strength of the ions will increase compared to the strength of the O 2 molecule. That is, the removal of an electron from a bonding orbitals reduces the binding energy in a molecular ion, and the removal of an electron from antibonding orbital leads to an increase in the binding energy in the molecular ion in comparison with the molecule.

MMO makes it possible to explain the existence of electron-deficient compounds (B 2 H 6 , NO) and noble gas compounds; magnetism and color of matter; stability of molecular ions and atoms in comparison with molecules.