In the space around the nuclei in comparison with the distribution of electron density in the neutral atoms forming this bond.

The so-called effective charges on atoms are used as a quantitative measure of bond polarity.

The effective charge is defined as the difference between the charge of electrons located in some region of space near the nucleus and the charge of the nucleus. However, this measure has only a conditional and approximate meaning, since it is impossible to single out unambiguously a region in a molecule that belongs exclusively to a single atom, and in the case of several bonds, to a specific bond.

The presence of an effective charge can be indicated by the symbols of the charges of atoms (for example, H + δ - Cl - δ, where δ is some fraction of the elementary charge).

Almost all chemical bonds, with the exception of bonds in diatomic homonuclear molecules, are polar to one degree or another. Covalent bonds are usually weakly polar. Ionic bonds are highly polar.

see also

Sources


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See what "Chemical bond polarity" is in other dictionaries:

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When a covalent bond is formed between dissimilar atoms, the bonding pair of electrons shifts towards the more electronegative atom. This leads to the polarization of molecules, so all diatomic molecules consisting of dissimilar elements turn out to be polar to some extent. In more complex molecules, the polarity also depends on the geometry of the molecule. For the appearance of polarity, it is necessary that the centers of distribution of positive and negative charges do not coincide.

In the CO 2 molecule, the carbon-oxygen bonds are polar, with a certain positive charge on the carbon atom, and the same negative charge on each of the oxygen atoms. Therefore, the center of positive charge is concentrated on the carbon atom. Since the oxygen atoms are located on the same straight line, but both sides of the carbon atom (linear molecule) are at equal distances, the positive charge is neutralized. Thus, despite the polarity of each bond in CO., the entire molecule as a whole is non-polar and the reason for this is

Rice. 434. Examples of the structure and polarity of a molecule is its linear structure. On the contrary, the S=C=0 molecule is polar, since the carbon-sulfur and carbon-oxygen bonds have different lengths and different polarities. On fig. 4.34 shows the structures and polarity of some molecules.

It follows from the above examples that if the atoms or groups of atoms attached to the central atom are the same or located symmetrically relative to it (linear, flat triangular, tetrahedral and other structures), then the molecule will be nonpolar. If unequal groups are attached to the central atom or if there is an asymmetric arrangement of groups, then the molecules are polar.

When considering polar bonds, the effective charge of atoms in a molecule is important. For example, in the HC1 molecule, the binding electron cloud is shifted towards the more electronegative chlorine atom, as a result of which the charge of the hydrogen nucleus is not compensated, and the electron density on the chlorine atom becomes excessive compared to the charge of its nucleus. Therefore, the hydrogen atom is positively polarized, and the chlorine atom is negatively polarized. The hydrogen atom has a positive charge, and the chlorine atom has a negative charge. This charge 8, called the effective charge, is usually established experimentally. So, for hydrogen 8 H \u003d +0.18, and for chlorine 5 C, \u003d -0.18 of the absolute electron charge, as a result, the bond in the HC1 molecule is 18% ionic (i.e., the degree of ionicity is 0.18 ).

Since the polarity of the bond depends on the degree of displacement of the bonding pair of electrons towards the more electronegative element, the following must be taken into account:

  • a) electronegativity (EO) is not a strict physical quantity that can be determined directly experimentally;
  • b) the value of electronegativity is not constant, but depends on the nature of the other atom with which this atom is bonded;
  • c) the same atom in a given chemical bond can sometimes function both as an electropositive and as an electronegative one.

Experimental evidence suggests that relative electronegativities (RERs) can be assigned to elements, the use of which makes it possible to judge the degree of polarity of the bond between atoms in a molecule (see also paragraphs 3.6 and 4.3).

In a molecule consisting of two atoms, the greater the polarity of the covalent bond, the higher the RER of one of them, therefore, with an increase in the RER of the second element, the degree of ionicity of the compound increases.

To characterize the reactivity of molecules, not only the nature of the electron density distribution is important, but also the possibility of its change under the influence of an external influence. The measure of this change is the polarizability of the bond, i.e. its ability to become polar or even more polar. Bond polarization occurs both under the influence of an external electric field and under the influence of another molecule that is a reaction partner. The result of these influences may be the polarization of the bond, accompanied by its complete break. In this case, the binding pair of electrons remains at the more electronegative atom, which leads to the formation of opposite ions. This type of bond breaking is called teterolytic. For example:

In the above example of an asymmetric bond cleavage, hydrogen is split off in the form of an H + -ion, and the binding pair of electrons remains with chlorine, so the latter is converted into an anion C1.

In addition to this type of bond rupture, a symmetrical bond rupture is also possible, when not ions are formed, but atoms and radicals. This type of bond breaking is called homolytic.

Electronegativity of atoms of elements. Relative electronegativity. Change in periods and groups of the Periodic system. The polarity of a chemical bond, the polarity of molecules and ions.

Electronegativity (e.o.) is the ability of an atom to displace electron pairs towards itself.
Meroy e.o. is the energy arithmetically equal to ½ the sum of the ionization energy I and the electron similarity energy E
E.O. = ½ (I+E)

Relative electronegativity. (OEO)

Fluorine, as the strongest e.o element, is assigned a value of 4.00 relative to which the other elements are considered.

Changes in periods and groups of the Periodic system.

Within periods, as the nuclear charge increases from left to right, electronegativity increases.

Least value is observed in alkali and alkaline earth metals.

Greatest- for halogens.

The higher the electronegativity, the stronger the non-metallic properties of the elements.

Electronegativity (χ) is a fundamental chemical property of an atom, a quantitative characteristic of the ability of an atom in a molecule to displace common electron pairs towards itself.

The modern concept of the electronegativity of atoms was introduced by the American chemist L. Pauling. L. Pauling used the concept of electronegativity to explain the fact that the energy of the A-B heteroatomic bond (A, B are symbols of any chemical elements) is generally greater than the geometric mean of the A-A and B-B homoatomic bonds.

The highest value of e.o. fluorine, and the lowest is cesium.

The theoretical definition of electronegativity was proposed by the American physicist R. Mulliken. Based on the obvious position that the ability of an atom in a molecule to attract an electronic charge to itself depends on the ionization energy of the atom and its electron affinity, R. Mulliken introduced the concept of the electronegativity of the atom A as the average value of the binding energy of the outer electrons during the ionization of valence states ( for example, from A− to A+) and on this basis proposed a very simple relation for the electronegativity of an atom:

where J1A and εA are the ionization energy of an atom and its electron affinity, respectively.
Strictly speaking, an element cannot be ascribed a permanent electronegativity. The electronegativity of an atom depends on many factors, in particular, on the valence state of the atom, the formal oxidation state, the coordination number, the nature of the ligands that make up the environment of the atom in the molecular system, and some others. Recently, more and more often, to characterize electronegativity, the so-called orbital electronegativity is used, which depends on the type of atomic orbital involved in the formation of a bond, and on its electron population, i.e., whether the atomic orbital is occupied by an unshared electron pair, singly populated by an unpaired electron, or is vacant. But, despite the well-known difficulties in interpreting and determining electronegativity, it always remains necessary for a qualitative description and prediction of the nature of bonds in a molecular system, including the bond energy, electronic charge distribution and degree of ionicity, force constant, etc. One of the most developed in the current approach is the Sanderson approach. This approach was based on the idea of ​​equalizing the electronegativity of atoms during the formation of a chemical bond between them. Numerous studies have found relationships between the Sanderson electronegativity and the most important physicochemical properties of inorganic compounds of the vast majority of the elements of the periodic table. A modification of Sanderson's method, based on the redistribution of electronegativity between the atoms of a molecule for organic compounds, also proved to be very fruitful.

2) The polarity of the chemical bond, the polarity of molecules and ions.

What is in the abstract and in the textbook - Polarity is associated with a dipole moment. It appears as a result of the displacement of a common electron pair to one of the atoms. Polarity also depends on the difference in the electronegativity of the atoms being bonded. two atoms, the more polar is the chemical bond between them. Depending on how the electron density is redistributed during the formation of a chemical bond, several types of it are distinguished. The limiting case of chemical bond polarization is a complete transition from one atom to another.

In this case, two ions are formed, between which an ionic bond arises. In order for two atoms to be able to create an ionic bond, it is necessary that their e.o. differed greatly. If e.o. are equal, then a non-polar covalent bond is formed. The most common polar covalent bond is formed between any atoms that have different e.o.

The effective charges of atoms can serve as a quantitative estimate of the polarity of a bond. The effective charge of an atom characterizes the difference between the number of electrons belonging to a given atom in a chemical compound and the number of electrons of a free atom. An atom of a more electronegative element attracts electrons more strongly, so the electrons are closer to it, and it receives some negative charge, which is called effective, and its partner has the same positive effective charge. If the electrons that form a bond between atoms belong to them equally, the effective charges are zero.

For diatomic molecules, it is possible to characterize the polarity of the bond and determine the effective charges of atoms based on measuring the dipole moment M = q * r where q is the charge of the dipole pole, which is equal to the effective charge for a diatomic molecule, r is the internuclear distance. The dipole moment of the bond is a vector quantity. It is directed from the positively charged part of the molecule to its negative part. The effective charge on the atom of an element does not coincide with the oxidation state.

The polarity of molecules largely determines the properties of substances. Polar molecules turn towards each other with oppositely charged poles, and mutual attraction arises between them. Therefore, substances formed by polar molecules have higher melting and boiling points than substances whose molecules are non-polar.

Liquids whose molecules are polar have a higher dissolving power. Moreover, the greater the polarity of the solvent molecules, the higher the solubility of polar or ionic compounds in it. This dependence is explained by the fact that the polar molecules of the solvent, due to the dipole-dipole or ion-dipole interaction with the solute, contribute to the decomposition of the solute into ions. For example, a solution of hydrogen chloride in water, whose molecules are polar, conducts electricity well. A solution of hydrogen chloride in benzene does not have an appreciable electrical conductivity. This indicates the absence of hydrogen chloride ionization in the benzene solution, since the benzene molecules are nonpolar.

Ions, like an electric field, have a polarizing effect on each other. When two ions meet, their mutual polarization occurs, i.e. displacement of the electrons of the outer layers relative to the nuclei. The mutual polarization of ions depends on the charges of the nucleus and ion, the radius of the ion, and other factors.

Within the groups of e.o. decreases.

The metallic properties of the elements increase.

Metallic elements at the external energy level contain 1,2,3 electrons and are characterized by a low value of ionization potentials and e.o. because metals show a pronounced tendency to donate electrons.
Non-metallic elements have a higher ionization energy.
As the outer shell of nonmetals is filled, the atomic radius decreases within the periods. On the outer shell, the number of electrons is 4,5,6,7,8.

The polarity of a chemical bond. Polarity of molecules and ions.

The polarity of a chemical bond is determined by the displacement of the bonds of an electron pair to one of the atoms.

A chemical bond arises due to the redistribution of electrons in valence orbitals, resulting in a stable electronic configuration of a noble gas, due to the formation of ions or the formation of common electron pairs.
A chemical bond is characterized by energy and length.
The measure of bond strength is the energy expended to break the bond.
For example. H - H = 435 kJmol-1

Electronegativity of atomic elements
Electronegativity is a chemical property of an atom, a quantitative characteristic of the ability of an atom in a molecule to attract electrons to itself from atoms of other elements.
Relative electronegativity

The first and most famous scale of relative electronegativity is the L. Pauling scale, obtained from thermochemical data and proposed in 1932. The electronegativity of the most electronegative element fluorine, (F) = 4.0, is arbitrarily taken as the reference point in this scale.

Elements of group VIII of the periodic system (noble gases) have zero electronegativity;
The conditional boundary between metals and non-metals is considered to be the value of relative electronegativity equal to 2.

The electronegativity of the elements of the periodic system, as a rule, increases sequentially from left to right in each period. Within each group, with a few exceptions, electronegativity consistently decreases from top to bottom. Electronegativity is used to characterize a chemical bond.
Bonds with a smaller difference in the electronegativity of atoms are referred to as polar covalent bonds. The smaller the difference in the electronegativity of the atoms forming a chemical bond, the lower the degree of ionicity of this bond. The zero difference in the electronegativity of atoms indicates the absence of an ionic character in the bond formed by them, i.e., its pure covalence.

Polarity of a chemical bond, polarity of molecules and ions
The polarity of chemical bonds, a characteristic of a chemical bond, showing the redistribution of electron density in space near the nuclei compared to the initial distribution of this density in the neutral atoms that form this bond.

Almost all chemical bonds, with the exception of bonds in diatomic homonuclear molecules, are polar to one degree or another. Usually covalent bonds are weakly polar, ionic bonds are strongly polar.

For example:
covalent non-polar: Cl2, O2, N2, H2,Br2

covalent polar: H2O, SO2, HCl, NH3, etc.

In homonuclear molecules (H 2 , F 2 , etc.), the electron pair that forms the bond equally belongs to each atom, so the centers of positive and negative charges in the molecule coincide. Such molecules are non-polar.

However, in heteronuclear molecules, the contribution to the coupling of the wave functions of different atoms is not the same. Near one of the atoms, an excess electron density appears, therefore, an excess negative charge, and near the other, a positive one. In this case, one speaks of a displacement of an electron pair from one atom to another, but this should not be understood literally, but only as an increase in the probability of finding an electron pair near one of the nuclei of the molecule.

To determine the direction of such a shift and a semiquantitative estimate of its magnitude, the concept of electronegativity is introduced.

There are several scales of electronegativity. However, the elements are arranged in the electronegativity series in the same order, so the differences are insignificant, and the electronegativity scales are quite comparable.

According to R. Mulliken, electronegativity is half the sum of ionization energies and electron affinity (see Section 2.10.3):

The valence electron pair is shifted to a more electronegative atom.

It is more convenient to use not absolute values ​​of electronegativity, but relative ones. The unit is the electronegativity of lithium 3 Li. The relative electronegativity of any element A is:

Heavy alkali metals have the lowest electronegativity. (X Fr = 0.7). The most electronegative element is fluorine (X F = 4.0). By periods, there is a general trend of increasing electronegativity, and by subgroups - its decrease (Table 3.4).

In the practical use of the data in this table (as well as data from other electronegativity scales), it should be borne in mind that in molecules consisting of three or more atoms, the value of electronegativity under the influence of neighboring atoms can change noticeably. Strictly speaking, a constant electronegativity cannot be attributed to an element at all. It depends on the valence state of the element, the type of compound, etc. Nevertheless, this concept is useful for a qualitative explanation of the properties of chemical bonds and compounds.

Table 3.4

Electronegativity of s- and p-elements according to Pauling

Period

Group

The polarity of the bond is determined by the displacement of the valence electron pair in diatomic molecules and is quantitatively characterized dipole moment, or dipole electric moment, molecules. It is equal to the product of the distance between the nuclei G in the molecule and the effective charge 5 corresponding to this distance:

Insofar as G considered to be a vector directed from positive to negative charge, the dipole moment is also a vector and has the same direction. The unit of dipole moment is the debye D (1D = 3.33 10 -30 C m).

The dipole moment of a complex molecule is defined as the vector sum of the dipole moments of all bonds. Therefore, if the AB I molecule is symmetrical with respect to the line of each bond, the total dipole moment of such a molecule, despite the polarity

the number of links A-B is equal to zero: D = ^ D; = 0. Examples are

live the previously considered symmetrical molecules, the bonds in which are formed by hybrid orbitals: BeF 2, BF 3, CH 4, SF 6, etc.

Molecules in which bonds are formed by non-hybrid orbitals or hybrid orbitals involving lone pairs of electrons are asymmetric with respect to bond lines. The dipole moments of such molecules are not equal to zero. Examples of such polar molecules: H 2 S, NH 3 , H 2 0, etc. In fig. 3.18 shows a graphical interpretation of the summation of polar bond vectors in a symmetric BeF 2 (fl) molecule and an asymmetric H 2 S molecule (b).


Rice. 3.18. Dipole moments of BeF 2 (a) and H 2 S (b) molecules

As already noted, the greater the difference in the electronegativity of the atoms forming the bond, the more the valence electron pair shifts, the more polar the bond and, consequently, the greater the effective charge b, which is illustrated in Table. 3.5.

Table 3.5

Changing the nature of the bond in a series of compounds of elements of the II period with fluorine

In a polar bond, two components can be conditionally distinguished: ionic, due to electrostatic attraction, and covalent, due to overlapping orbitals. As the difference in electronegativity increases OH the valence electron pair shifts more and more towards the fluorine atom, which acquires an increasingly negative effective charge. The contribution of the ionic component to the bond increases, while the proportion of the covalent component decreases. Quantitative changes turn into qualitative ones: in the UF molecule, the electron pair almost completely belongs to fluorine, and its effective charge approaches unity, i.e. to the charge of the electron. We can assume that two ions were formed: the Li + cation and the anion F~ and the bond is due only to their electrostatic attraction (the covalent component can be neglected). Such a connection is called ionic. It can be considered as extreme case of a covalent polar bond.

The electrostatic field does not have preferred directions. So ionic bond as opposed to covalent no directionality. An ion interacts with any number of ions of opposite charge. This is due to another distinctive property of the ionic bond - lack of saturation.

For ionic molecules, the binding energy can be calculated. If we consider ions as non-deformable balls with charges ±e, then the force of attraction between them, depending on the distance between the centers of the ions G can be expressed by the Coulomb equation:

The attraction energy is determined by the relation

When approaching, a repulsive force appears due to the interaction of electron shells. It is inversely proportional to the distance to the power P:

where AT is some constant. Exponent P is much greater than unity and for various configurations of ions lies in the range from 5 to 12. Taking into account that the force is the derivative of energy with respect to distance, from equation (3.6) we obtain:

With change G change F np and F qtt . At some distance g 0 these forces are equalized, which corresponds to the minimum of the resulting interaction energy U Q . After transformation, you can get

This equation is known as the Born equation.

Minimum on the dependence curve U=f(r) correspond to the equilibrium distance r 0 and the energy U Q . This is the binding energy between ions. Even P is unknown, then we can estimate the value of the binding energy by taking 1 /P equal to zero:


The error will not exceed 20%.

For ions with charges z l and z 2 equations (3.7) and (3.8) take the form:


Since the existence of a bond approaching a purely ionic one in molecules of this type is problematic, the last equations should be considered a very rough approximation.

At the same time, the problems of polarity and ionicity of the bond can be approached from the opposite position - from the point of view of ion polarization. It is assumed that there is a complete transfer of electrons, and the molecule consists of isolated ions. Then the electron clouds are displaced under the action of the electric field created by the ions, - polarization ions.

Polarization is a two-pronged process that combines polarizing effect ions from their polarizability. Polarizability is the ability of an electron cloud of an ion, molecule or atom to deform under the action of the electrostatic field of another ion. The strength of this field determines the polarizing effect of the ion. It follows from equation (3.10) that the greater the polarizing effect of an ion, the greater its charge and the smaller its radius. The radii of cations, as a rule, are much smaller than the radii of anions; therefore, in practice, it is more often necessary to encounter polarization of anions under the action of cations, and not vice versa. The polarizability of ions also depends on their charge and radius. Ions of large size and charge are more easily polarized. The polarizing effect of an ion is reduced to pulling the electron cloud of an ion of opposite charge towards itself. As a result, the ionicity of the bond decreases; the bond becomes polar covalent. Thus, the polarization of ions reduces the degree of ionicity of the bond and is opposite to the polarization of the bond in its effect.

Polarization of ions in a molecule, i.e. an increase in the proportion of a covalent bond in it increases the strength of its decay into ions. In a series of compounds of a given cation with anions of the same type, the degree of dissociation in solutions decreases with an increase in the polarizability of the anions. For example, in the series of lead halides PbCl 2 - PbBr 2 - Pb 2, the radius of the halide anions increases, their polarizability increases, and the decay into ions weakens, which is expressed in a decrease in solubility.

When comparing the properties of salts with the same anion and sufficiently large cations, one should take into account the polarization of the cations. For example, the radius of the Hg 2+ ion is greater than the radius of the Ca 2+ ion, so Hg 2+ polarizes more strongly than Ca 2+ . As a result, CaCl 2 is a strong electrolyte; dissociates completely in solution, and HgCl 2 - as a weak electrolyte, i.e. practically does not dissociate in solutions.

The polarization of ions in a molecule reduces its strength during decay into atoms or molecules. For example, in the series CaCl 2 - CaBr 2 - Ca1 2, the radius of halide ions increases, their polarization by the Ca 2+ ion increases, therefore, the temperature of thermal dissociation into calcium and halogen decreases: CaNa1 2 \u003d Ca + Na1 2.

If the ion is easily polarized, then its excitation requires little energy, which corresponds to the absorption of visible light quanta. This is the reason for the color of solutions of such compounds. An increase in polarizability leads to an increase in color, for example, in the series NiCl 2 - NiBr 2 - Nil 2 (increase in the polarizability of the anion) or in the series KC1 - CuCl 2 (increase in the polarizability of the cation).

The boundary between covalent polar and ionic bonds is very conditional. For molecules in the gaseous state, it is believed that with a difference in electronegativity AH > 2.5 bond is ionic. In solutions of polar solvents, as well as in the crystalline state, solvent molecules and neighboring particles at the sites of the crystal lattice, respectively, exert a strong influence. Therefore, the ionic nature of the bond manifests itself at a much smaller difference in electronegativity. In practice, we can assume that the bond between typical metals and nonmetals in solutions and crystals is ionic.

In molecules, the positive charges of the nuclei are compensated by the negative charges of the electrons. However, positive and negative charges can be spatially separated. Let us assume that the molecule consists of atoms of different elements (HC1, CO, etc.). In this case, the electrons are shifted to an atom with a higher electronegativity and the centers of gravity of positive and negative charges do not coincide, forming electric dipole- a system of two equal in magnitude and opposite in sign charges q, located at a distance l called dipole length. The length of the dipole is a vector quantity. Its direction is conditionally taken from the negative charge to the positive one. Such molecules are called polar molecules or dipoles.

The polarity of the molecule is greater, the greater the absolute value of the charge and the length of the dipole. The measure of polarity is the product q . l, called the electric moment of the dipole μ: μ = q. l.

Unit of measurement μ serves as Debye (D). 1 D \u003d 3.3. 10 -30 C. m.

In molecules consisting of two identical atoms, μ = 0. They are called nonpolar. If such a particle enters an electric field, then under the action of the field, polarization- displacement of the centers of gravity of positive and negative charges. An electric dipole moment arises in the particle, called induced dipole.

The dipole moment of a diatomic AB molecule can be identified with the dipole moment of the A-B bond in it. If the common electron pair is shifted to one of the atoms, then the electric moment of the bond dipole is not equal to zero. The relationship in this case is called polar covalent bond. If the electron pair is symmetrically located relative to the atoms, then the bond is called non-polar.

In a polyatomic molecule, a certain electric dipole moment can be assigned to each bond. Then the electric moment of the dipole of the molecule can be represented as the vector sum of the electric moments of the dipole of individual bonds. The existence or absence of a dipole moment in a molecule is related to its symmetry. Molecules that have a symmetrical structure are non-polar (μ = 0). These include diatomic molecules with identical atoms (H 2, C1 2, etc.), a benzene molecule, molecules with polar bonds BF 3, A1F 3, CO 2, BeC1 2, etc.

The electric moment of the dipole of a molecule is an important molecular parameter. Knowing the value of μ can indicate the geometric structure of the molecule. So, for example, the polarity of a water molecule indicates its angular structure, and the absence of a CO 2 dipole moment indicates its linearity.

Ionic bond

The limiting case of a covalent polar bond is an ionic bond. If the electronegativities of atoms differ very much (for example, atoms of alkali metals and halogens), then when they approach each other, the valence electrons of one atom completely transfer to the second atom. As a result of this transition, both atoms become ions and take on the electronic structure of the nearest noble gas. For example, when sodium and chlorine atoms interact, they turn into Na + and Cl - ions, between which there is an electrostatic attraction. The ionic bond can be described in terms of the VS and MO methods, but it is usually considered using the classical laws of electrostatics.

Molecules in which there is a pure ionic bond are found in the vapor state of matter. Ionic crystals are made up of endless rows of alternating positive and negative ions bound by electrostatic forces. When ionic crystals are dissolved or melted, positive and negative ions pass into the solution or melt.

It should be noted that ionic bonds are very strong, therefore, to destroy ionic crystals, it is necessary to expend a lot of energy. This explains the fact that ionic compounds have high melting points.

Unlike a covalent bond, an ionic bond does not have the properties of saturation and directionality. The reason for this is that the electric field created by the ions has spherical symmetry and acts equally on all ions. Therefore, the number of ions surrounding a given ion and their spatial arrangement are determined only by the magnitudes of the charges of the ions and their sizes.

Considering the ionic bond, it must be borne in mind that during the electrostatic interaction between ions, their deformation occurs, called polarization. On fig. 2.1, a two interacting electrostatically neutral ions are shown, which retain a perfectly spherical shape. On fig. 2.1, b the polarization of ions is shown, which leads to a decrease in the effective distance between the centers of positive and negative charges. The greater the polarization of the ions, the lower the degree of ionicity of the bond, i.e., the greater the covalent nature of the bond between them. In crystals, the polarization turns out to be low, since the ions are symmetrically surrounded by ions of the opposite sign and the ion is subjected to the same action in all directions.