Covalent, ionic, and metallic are the three main types of chemical bonds.

Let's get to know more about covalent chemical bond. Let's consider the mechanism of its occurrence. Let's take the formation of a hydrogen molecule as an example:

A spherically symmetric cloud formed by a 1s electron surrounds the nucleus of a free hydrogen atom. When atoms approach each other up to a certain distance, their orbitals partially overlap (see Fig.), as a result, a molecular two-electron cloud appears between the centers of both nuclei, which has a maximum electron density in the space between the nuclei. With an increase in the density of the negative charge, there is a strong increase in the forces of attraction between the molecular cloud and the nuclei.

So, we see that a covalent bond is formed by overlapping electron clouds of atoms, which is accompanied by the release of energy. If the distance between the nuclei of the atoms approaching to touch is 0.106 nm, then after the overlap of the electron clouds it will be 0.074 nm. The greater the overlap of electron orbitals, the stronger the chemical bond.

covalent called chemical bonding carried out by electron pairs. Compounds with a covalent bond are called homeopolar or atomic.

Exist two types of covalent bond: polar and non-polar.

With non-polar covalent bond formed by a common pair of electrons, the electron cloud is distributed symmetrically with respect to the nuclei of both atoms. An example can be diatomic molecules that consist of one element: Cl 2, N 2, H 2, F 2, O 2 and others, in which the electron pair belongs to both atoms equally.

At polar In a covalent bond, the electron cloud is displaced towards the atom with a higher relative electronegativity. For example, molecules of volatile inorganic compounds such as H 2 S, HCl, H 2 O and others.

The formation of the HCl molecule can be represented as follows:

Because the relative electronegativity of the chlorine atom (2.83) is greater than that of the hydrogen atom (2.1), the electron pair shifts towards the chlorine atom.

In addition to the exchange mechanism for the formation of a covalent bond - due to overlap, there is also donor-acceptor the mechanism of its formation. This is a mechanism in which the formation of a covalent bond occurs due to a two-electron cloud of one atom (donor) and a free orbital of another atom (acceptor). Let's look at an example of the mechanism for the formation of ammonium NH 4 +. In the ammonia molecule, the nitrogen atom has a two-electron cloud:

The hydrogen ion has a free 1s orbital, let's denote it as .

In the process of ammonium ion formation, the two-electron cloud of nitrogen becomes common for nitrogen and hydrogen atoms, which means it is converted into a molecular electron cloud. Therefore, a fourth covalent bond appears. The process of ammonium formation can be represented as follows:

The charge of the hydrogen ion is dispersed among all atoms, and the two-electron cloud that belongs to nitrogen becomes common with hydrogen.

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covalent chemical bond occurs between atoms with close or equal values ​​of electronegativity. Suppose that chlorine and hydrogen tend to take electrons and take on the structure of the nearest noble gas, then neither of them will give up an electron to the other. How are they connected anyway? It's simple - they will share with each other, a common electron pair is formed.

Now consider the distinctive features of a covalent bond.

Unlike ionic compounds, the molecules of covalent compounds are held together by "intermolecular forces", which are much weaker than chemical bonds. In this regard, the covalent bond is characteristic saturability– the formation of a limited number of bonds.

It is known that atomic orbitals are oriented in space in a certain way, therefore, when a bond is formed, the overlap of electron clouds occurs in a certain direction. Those. such a property of a covalent bond is realized as orientation.

If a covalent bond in a molecule is formed by identical atoms or atoms with equal electronegativity, then such a bond has no polarity, i.e. the electron density is distributed symmetrically. It's called non-polar covalent bond ( H 2 , Cl 2 , O 2 ). Bonds can be single, double or triple.

If the electronegativity of the atoms differ, then when they combine, the electron density is distributed unevenly between the atoms and forms covalent polar bond(HCl, H 2 O, CO), the multiplicity of which can also be different. When this type of bond is formed, a more electronegative atom acquires a partial negative charge, and an atom with a lower electronegativity acquires a partial positive charge (δ- and δ+). An electric dipole is formed, in which charges of opposite sign are located at a certain distance from each other. The dipole moment is used as a measure of bond polarity:

The polarity of the compound is all the more pronounced, the greater the dipole moment. Molecules will be non-polar if the dipole moment is zero.

In connection with the above features, it can be concluded that covalent compounds are volatile and have low melting and boiling points. Electric current cannot pass through these connections, hence they are poor conductors and good insulators. When heat is applied, many covalently bonded compounds ignite. For the most part, these are hydrocarbons, as well as oxides, sulfides, halides of non-metals and transition metals.

In which one of the atoms donated an electron and became a cation, and the other atom accepted an electron and became an anion.

The characteristic properties of a covalent bond - directionality, saturation, polarity, polarizability - determine the chemical and physical properties of compounds.

The direction of the bond is due to the molecular structure of the substance and the geometric shape of their molecule. The angles between two bonds are called bond angles.

Saturation - the ability of atoms to form a limited number of covalent bonds. The number of bonds formed by an atom is limited by the number of its outer atomic orbitals.

The polarity of the bond is due to the uneven distribution of the electron density due to differences in the electronegativity of the atoms. On this basis, covalent bonds are divided into non-polar and polar (non-polar - a diatomic molecule consists of identical atoms (H 2, Cl 2, N 2) and the electron clouds of each atom are distributed symmetrically with respect to these atoms; polar - a diatomic molecule consists of atoms of different chemical elements , and the general electron cloud shifts towards one of the atoms, thereby forming an asymmetry in the distribution of the electric charge in the molecule, generating a dipole moment of the molecule).

The polarizability of a bond is expressed in the displacement of bond electrons under the influence of an external electric field, including that of another reacting particle. Polarizability is determined by electron mobility. The polarity and polarizability of covalent bonds determine the reactivity of molecules with respect to polar reagents.

However, twice Nobel Prize winner L. Pauling pointed out that "in some molecules there are covalent bonds due to one or three electrons instead of a common pair." A single-electron chemical bond is realized in the molecular ion hydrogen H 2 + .

The molecular hydrogen ion H 2 + contains two protons and one electron. The single electron of the molecular system compensates for the electrostatic repulsion of two protons and keeps them at a distance of 1.06 Å (the length of the H 2 + chemical bond). The center of the electron density of the electron cloud of the molecular system is equidistant from both protons by the Bohr radius α 0 =0.53 A and is the center of symmetry of the molecular hydrogen ion H 2 + .

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  A covalent bond is formed by a pair of electrons shared between two atoms, and these electrons must occupy two stable orbitals, one from each atom.

  A + B → A: B

  As a result of socialization, electrons form a filled energy level. A bond is formed if their total energy at this level is less than in the initial state (and the difference in energy is nothing more than the bond energy).

  According to the theory of molecular orbitals, the overlap of two atomic orbitals leads in the simplest case to the formation of two molecular orbitals (MOs): binding MO and antibonding (loosening) MO. Shared electrons are located on a lower energy binding MO.

  Formation of a bond during the recombination of atoms

  However, the mechanism of interatomic interaction remained unknown for a long time. Only in 1930, F. London introduced the concept of dispersion attraction - the interaction between instantaneous and induced (induced) dipoles. At present, the attractive forces due to the interaction between fluctuating electric dipoles of atoms and molecules are called "London forces".

  The energy of such an interaction is directly proportional to the square of the electronic polarizability α and inversely proportional to the sixth power of the distance between two atoms or molecules.

  Bond formation by the donor-acceptor mechanism

  In addition to the homogeneous mechanism for the formation of a covalent bond described in the previous section, there is a heterogeneous mechanism - the interaction of oppositely charged ions - the proton H + and the negative hydrogen ion H -, called the hydride ion:

  H + + H - → H 2

  When the ions approach, the two-electron cloud (electron pair) of the hydride ion is attracted to the proton and eventually becomes common to both hydrogen nuclei, that is, it turns into a binding electron pair. The particle that supplies an electron pair is called a donor, and the particle that accepts this electron pair is called an acceptor. Such a mechanism for the formation of a covalent bond is called donor-acceptor.

  H + + H 2 O → H 3 O +

  A proton attacks the lone electron pair of a water molecule and forms a stable cation that exists in aqueous solutions of acids.

  Similarly, a proton is attached to an ammonia molecule with the formation of a complex ammonium cation:

  NH 3 + H + → NH 4 +

  In this way (according to the donor-acceptor mechanism for the formation of a covalent bond), a large class of onium compounds is obtained, which includes ammonium, oxonium, phosphonium, sulfonium and other compounds.

  A hydrogen molecule can act as an electron pair donor, which, upon contact with a proton, leads to the formation of a molecular hydrogen ion H 3 + :

  H 2 + H + → H 3 +

  The binding electron pair of the molecular hydrogen ion H 3 + belongs simultaneously to three protons.

  Types of covalent bond

  There are three types of covalent chemical bonds that differ in the mechanism of formation:

  1. Simple covalent bond. For its formation, each of the atoms provides one unpaired electron. When a simple covalent bond is formed, the formal charges of the atoms remain unchanged.

  • If the atoms that form a simple covalent bond are the same, then the true charges of the atoms in the molecule are also the same, since the atoms that form the bond equally own a shared electron pair. Such a connection is called non-polar covalent bond. Simple substances have such a connection, for example: 2, 2, 2. But not only non-metals of the same type can form a covalent non-polar bond. Non-metal elements whose electronegativity is of equal value can also form a covalent non-polar bond, for example, in the PH 3 molecule, the bond is covalent non-polar, since the EO of hydrogen is equal to the EO of phosphorus.
  • If the atoms are different, then the degree of ownership of a socialized pair of electrons is determined by the difference in the electronegativity of the atoms. An atom with greater electronegativity attracts a pair of bond electrons to itself more strongly, and its true charge becomes negative. An atom with less electronegativity acquires, respectively, the same positive charge. If a compound is formed between two different non-metals, then such a compound is called polar covalent bond.

  In the ethylene molecule C 2 H 4 there is a double bond CH 2 \u003d CH 2, its electronic formula: H: C:: C: H. The nuclei of all ethylene atoms are located in the same plane. Three electron clouds of each carbon atom form three covalent bonds with other atoms in the same plane (with angles between them of about 120°). The cloud of the fourth valence electron of the carbon atom is located above and below the plane of the molecule. Such electron clouds of both carbon atoms, partially overlapping above and below the plane of the molecule, form a second bond between carbon atoms. The first, stronger covalent bond between carbon atoms is called a σ-bond; the second, weaker covalent bond is called π (\displaystyle \pi )-communication.

  In a linear acetylene molecule

  H-S≡S-N (N: S::: S: N)

  there are σ-bonds between carbon and hydrogen atoms, one σ-bond between two carbon atoms and two π (\displaystyle \pi ) bonds between the same carbon atoms. Two π (\displaystyle \pi )-bonds are located above the sphere of action of the σ-bond in two mutually perpendicular planes.

  All six carbon atoms of the C 6 H 6 cyclic benzene molecule lie in the same plane. σ-bonds act between carbon atoms in the plane of the ring; the same bonds exist for each carbon atom with hydrogen atoms. Each carbon atom spends three electrons to make these bonds. Clouds of the fourth valence electrons of carbon atoms, having the shape of eights, are located perpendicular to the plane of the benzene molecule. Each such cloud overlaps equally with the electron clouds of neighboring carbon atoms. In the benzene molecule, not three separate π (\displaystyle \pi )-connections, but a single π (\displaystyle \pi ) dielectrics or semiconductors. Typical examples of atomic crystals (the atoms in which are interconnected by covalent (atomic) bonds) are

  A covalent bond is the most common type of chemical bond that occurs when interacting with the same or similar electronegativity values.

  A covalent bond is a bond between atoms using shared electron pairs.

  Since the discovery of the electron, many attempts have been made to develop an electronic theory of chemical bonding. The most successful were the works of Lewis (1916), who proposed to consider the formation of a bond as a consequence of the appearance of electron pairs common to two atoms. To do this, each atom provides the same number of electrons and tries to surround itself with an octet or doublet of electrons, characteristic of the external electronic configuration of inert gases. Graphically, the formation of covalent bonds due to unpaired electrons according to the Lewis method is depicted using dots indicating the outer electrons of the atom.

  Formation of a covalent bond according to the Lewis theory

  The mechanism of formation of a covalent bond

  The main sign of a covalent bond is the presence of a common electron pair belonging to both chemically connected atoms, since the presence of two electrons in the field of action of two nuclei is energetically more favorable than the presence of each electron in the field of its own nucleus. The emergence of a common electron pair of bonds can take place through different mechanisms, more often through exchange, and sometimes through donor-acceptor.

  According to the principle of the exchange mechanism for the formation of a covalent bond, each of the interacting atoms supplies the same number of electrons with antiparallel spins to the formation of a bond. For example:

  
  The general scheme for the formation of a covalent bond: a) by the exchange mechanism; b) according to the donor-acceptor mechanism

  According to the donor-acceptor mechanism, a two-electron bond arises during the interaction of various particles. One of them is a donor BUT: has an unshared pair of electrons (that is, one that belongs to only one atom), and the other is an acceptor AT has a vacant orbital.

  A particle that provides a two-electron bond (an unshared pair of electrons) is called a donor, and a particle with a free orbital that accepts this electron pair is called an acceptor.

  The mechanism of formation of a covalent bond due to a two-electron cloud of one atom and a vacant orbital of another is called the donor-acceptor mechanism.

  The donor-acceptor bond is otherwise called semipolar, since a partial effective positive charge δ+ arises on the donor atom (due to the fact that its undivided pair of electrons has deviated from it), and a partial effective negative charge δ- arises on the acceptor atom (due to the fact that that there is a shift in its direction of the undivided electron pair of the donor).

  An example of a simple electron pair donor is the H ion. — , which has an unshared electron pair. As a result of the addition of a negative hydride ion to a molecule whose central atom has a free orbital (indicated as an empty quantum cell in the diagram), for example, ВН 3 , a complex complex ion ВН 4 is formed — with a negative charge (N — + VN 3 ⟶⟶ [VN 4] -):

  

  The electron pair acceptor is a hydrogen ion, or simply a proton H +. Its attachment to a molecule whose central atom has an unshared electron pair, for example, to NH 3, also leads to the formation of a complex ion NH 4 +, but with a positive charge:

  

  Valence bond method

  First quantum mechanical theory of covalent bond was created by Heitler and London (in 1927) to describe the hydrogen molecule, and then was applied by Pauling to polyatomic molecules. This theory is called valence bond method, the main points of which can be summarized as follows:

  • each pair of atoms in a molecule is held together by one or more shared electron pairs, with the electron orbitals of the interacting atoms overlapping;
  • bond strength depends on the degree of overlap of electron orbitals;
  • the condition for the formation of a covalent bond is the antidirection of the electron spins; due to this, a generalized electron orbital arises with the highest electron density in the internuclear space, which ensures the attraction of positively charged nuclei to each other and is accompanied by a decrease in the total energy of the system.

  Hybridization of atomic orbitals

  Despite the fact that electrons of s-, p- or d-orbitals, which have different shapes and different orientations in space, participate in the formation of covalent bonds, in many compounds these bonds are equivalent. To explain this phenomenon, the concept of "hybridization" was introduced.

  Hybridization is the process of mixing and aligning orbitals in shape and energy, in which the electron densities of orbitals with similar energies are redistributed, as a result of which they become equivalent.

  The main provisions of the theory of hybridization:

  1. During hybridization, the initial shape and orbitals change mutually, while new, hybridized orbitals are formed, but with the same energy and the same shape, resembling an irregular figure eight.
  2. The number of hybridized orbitals is equal to the number of output orbitals involved in hybridization.
  3. Orbitals with similar energies (s- and p-orbitals of the outer energy level and d-orbitals of the outer or preliminary levels) can participate in hybridization.
  4. Hybridized orbitals are more elongated in the direction of formation of chemical bonds and therefore provide better overlap with the orbitals of the neighboring atom, as a result, it becomes stronger than the individual non-hybrid orbitals formed due to electrons.
  5. Due to the formation of stronger bonds and a more symmetrical distribution of electron density in the molecule, an energy gain is obtained, which more than compensates for the energy consumption required for the hybridization process.
  6. Hybridized orbitals must be oriented in space in such a way as to ensure maximum mutual separation from each other; in this case, the repulsion energy is the smallest.
  7. The type of hybridization is determined by the type and number of exit orbitals and changes the size of the bond angle, as well as the spatial configuration of the molecules.
  
  The form of hybridized orbitals and valence angles (geometric angles between the axes of symmetry of the orbitals) depending on the type of hybridization: a) sp-hybridization; b) sp 2 hybridization; c) sp 3 hybridization

  During the formation of molecules (or individual fragments of molecules), the following types of hybridization most often occur:

  
  General scheme of sp hybridization

  Bonds that are formed with the participation of electrons of sp-hybridized orbitals are also placed at an angle of 180 0, which leads to a linear shape of the molecule. This type of hybridization is observed in the halides of elements of the second group (Be, Zn, Cd, Hg), whose atoms in the valence state have unpaired s- and p-electrons. The linear form is also characteristic of the molecules of other elements (0=C=0,HC≡CH), in which bonds are formed by sp-hybridized atoms.

  
  Scheme of sp 2 hybridization of atomic orbitals and a flat triangular shape of the molecule, which is due to sp 2 hybridization of atomic orbitals

  This type of hybridization is most typical for molecules of p-elements of the third group, whose atoms in an excited state have an external electronic structure ns 1 np 2, where n is the number of the period in which the element is located. So, in the molecules of ВF 3 , BCl 3 , AlF 3 and in others bonds are formed due to sp 2 -hybridized orbitals of the central atom.

  
  Scheme of sp 3 hybridization of atomic orbitals

  Placing the hybridized orbitals of the central atom at an angle of 109 0 28` causes the tetrahedral shape of the molecules. This is very typical for saturated compounds of tetravalent carbon CH 4 , CCl 4 , C 2 H 6 and other alkanes. Examples of compounds of other elements with a tetrahedral structure due to sp 3 hybridization of the valence orbitals of the central atom are ions: BH 4 - , BF 4 - , PO 4 3- , SO 4 2- , FeCl 4 - .

  
  General scheme of sp 3d hybridization

  This type of hybridization is most commonly found in non-metal halides. An example is the structure of phosphorus chloride PCl 5 , during the formation of which the phosphorus atom (P ... 3s 2 3p 3) first goes into an excited state (P ... 3s 1 3p 3 3d 1), and then undergoes s 1 p 3 d-hybridization - five one-electron orbitals become equivalent and orient with their elongated ends to the corners of the mental trigonal bipyramid. This determines the shape of the PCl 5 molecule, which is formed when five s 1 p 3 d-hybridized orbitals overlap with 3p orbitals of five chlorine atoms.

  1. sp - Hybridization. When one s-i is combined with one p-orbitals, two sp-hybridized orbitals arise, located symmetrically at an angle of 180 0 .
  2. sp 2 - Hybridization. The combination of one s- and two p-orbitals leads to the formation of sp 2 -hybridized bonds located at an angle of 120 0, so the molecule takes the form of a regular triangle.
  3. sp 3 - Hybridization. The combination of four orbitals - one s- and three p leads to sp 3 - hybridization, in which four hybridized orbitals are symmetrically oriented in space to the four vertices of the tetrahedron, that is, at an angle of 109 0 28 `.
  4. sp 3 d - Hybridization. The combination of one s-, three p- and one d-orbitals gives sp 3 d-hybridization, which determines the spatial orientation of five sp 3 d-hybridized orbitals to the vertices of the trigonal bipyramid.
  5. Other types of hybridization. In the case of sp 3 d 2 hybridization, six sp 3 d 2 hybridized orbitals are directed towards the vertices of the octahedron. The orientation of the seven orbitals to the vertices of the pentagonal bipyramid corresponds to the sp 3 d 3 hybridization (or sometimes sp 3 d 2 f) of the valence orbitals of the central atom of the molecule or complex.

  Atomic Orbital Hybridization Method Explains Geometric Structure a large number molecules, however, according to experimental data, molecules with slightly different values ​​of bond angles are more often observed. For example, in CH 4, NH 3 and H 2 O molecules, the central atoms are in the sp 3 hybridized state, so one would expect that the bond angles in them are equal to tetrahedral ones (~ 109.5 0). It has been experimentally established that the bond angle in the CH 4 molecule is actually 109.5 0 . However, in NH 3 and H 2 O molecules, the value of the bond angle deviates from the tetrahedral one: it is 107.3 0 in the NH 3 molecule and 104.5 0 in the H 2 O molecule. Such deviations are explained by the presence of an undivided electron pair at the nitrogen and oxygen atoms. A two-electron orbital, which contains an unshared pair of electrons, due to its increased density, repels one-electron valence orbitals, which leads to a decrease in the bond angle. At the nitrogen atom in the NH 3 molecule, out of four sp 3 hybridized orbitals, three one-electron orbitals form bonds with three H atoms, and the fourth orbital contains an unshared pair of electrons.

  An unbound electron pair, which occupies one of the sp 3 -hybridized orbitals directed to the vertices of the tetrahedron, repels one-electron orbitals, causes an asymmetric distribution of the electron density surrounding the nitrogen atom, and as a result, compresses the bond angle to 107.3 0 . A similar picture of the decrease in the bond angle from 109.5 0 to 107 0 as a result of the action of the unshared electron pair of the N atom is also observed in the NCl 3 molecule.

  
  Deviation of the bond angle from the tetrahedral (109.5 0) in the molecule: a) NH3; b) NCl3

  At the oxygen atom in the H 2 O molecule, four sp 3 hybridized orbitals have two one-electron and two two-electron orbitals. One-electron hybridized orbitals participate in the formation of two bonds with two H atoms, and two two-electron pairs remain undivided, that is, belonging only to the H atom. This increases the asymmetry of the electron density distribution around the O atom and reduces the bond angle compared to the tetrahedral one to 104.5 0 .

  Consequently, the number of unbound electron pairs of the central atom and their placement in hybridized orbitals affects the geometric configuration of molecules.

  Characteristics of a covalent bond

  A covalent bond has a set of specific properties that define its specific features, or characteristics. These, in addition to the characteristics already considered "bond energy" and "bond length", include: bond angle, saturation, directivity, polarity, and the like.

  1. Valence angle- this is the angle between adjacent bond axes (that is, conditional lines drawn through the nuclei of chemically connected atoms in a molecule). The value of the bond angle depends on the nature of the orbitals, the type of hybridization of the central atom, the influence of unshared electron pairs that do not participate in the formation of bonds.

  2. Saturation. Atoms have the ability to form covalent bonds, which can be formed, firstly, according to the exchange mechanism due to the unpaired electrons of an unexcited atom and due to those unpaired electrons that arise as a result of its excitation, and secondly, according to the donor-acceptor mechanism. However, the total number of bonds an atom can form is limited.

  Saturation is the ability of an atom of an element to form a certain, limited number of covalent bonds with other atoms.

  So, the second period, which have four orbitals on the external energy level (one s- and three p-), form bonds, the number of which does not exceed four. Atoms of elements of other periods with a large number of orbitals at the outer level can form more bonds.

  3. Orientation. According to the method, the chemical bond between atoms is due to the overlap of orbitals, which, with the exception of s-orbitals, have a certain orientation in space, which leads to the direction of the covalent bond.

  The orientation of a covalent bond is such an arrangement of the electron density between atoms, which is determined by the spatial orientation of the valence orbitals and ensures their maximum overlap.

  Since electronic orbitals have different shapes and different orientations in space, their mutual overlap can be realized in various ways. Depending on this, σ-, π- and δ-bonds are distinguished.

  A sigma bond (σ bond) is an overlap of electron orbitals in which the maximum electron density is concentrated along an imaginary line connecting two nuclei.

  A sigma bond can be formed by two s electrons, one s and one p electron, two p electrons, or two d electrons. Such a σ-bond is characterized by the presence of one region of overlapping electron orbitals, it is always single, that is, it is formed by only one electron pair.

  

  A variety of forms of spatial orientation of "pure" orbitals and hybridized orbitals do not always allow the possibility of overlapping orbitals on the bond axis. The overlap of valence orbitals can occur on both sides of the bond axis - the so-called "lateral" overlap, which most often occurs during the formation of π bonds.

  Pi-bond (π-bond) is the overlap of electron orbitals, in which the maximum electron density is concentrated on both sides of the line connecting the nuclei of atoms (i.e., from the bond axis).

  A pi bond can be formed by the interaction of two parallel p orbitals, two d orbitals, or other combinations of orbitals whose axes do not coincide with the bond axis.

  
  Schemes for the formation of π-bonds between conditional A and B atoms in the lateral overlap of electron orbitals

  4. Multiplicity. This characteristic is determined by the number of common electron pairs that bind atoms. A covalent bond in multiplicity can be single (simple), double and triple. A bond between two atoms using one common electron pair is called a single bond (simple), two electron pairs - a double bond, three electron pairs - a triple bond. So, in the hydrogen molecule H 2, the atoms are connected by a single bond (H-H), in the oxygen molecule O 2 - double (B \u003d O), in the nitrogen molecule N 2 - triple (N≡N). Of particular importance is the multiplicity of bonds in organic compounds - hydrocarbons and their derivatives: in ethane C 2 H 6 a single bond (C-C) occurs between C atoms, in ethylene C 2 H 4 - double (C \u003d C) in acetylene C 2 H 2 - triple (C ≡ C)(C≡C).

  The multiplicity of the bond affects the energy: with an increase in the multiplicity, its strength increases. An increase in the multiplicity leads to a decrease in the internuclear distance (bond length) and an increase in the binding energy.

  
  Multiplicity of bonds between carbon atoms: a) single σ-bond in ethane H3C-CH3; b) double σ + π-bond in ethylene H2C = CH2; c) triple σ+π+π-bond in acetylene HC≡CH

  5. Polarity and polarizability. The electron density of a covalent bond can be located differently in the internuclear space.

  Polarity is a property of a covalent bond, which is determined by the location of the electron density in the internuclear space relative to the connected atoms.

  Depending on the location of the electron density in the internuclear space, polar and non-polar covalent bonds are distinguished. A non-polar bond is such a bond in which the common electron cloud is located symmetrically with respect to the nuclei of the connected atoms and equally belongs to both atoms.

  Molecules with this type of bond are called non-polar or homonuclear (that is, those that include atoms of one element). A non-polar bond appears as a rule in homonuclear molecules (H 2, Cl 2, N 2, etc.) or, more rarely, in compounds formed by atoms of elements with similar electronegativity values, for example, carborundum SiC. A polar (or heteropolar) bond is a bond in which the common electron cloud is asymmetric and shifted to one of the atoms.

  Molecules with a polar bond are called polar, or heteronuclear. In molecules with a polar bond, the generalized electron pair shifts towards the atom with a higher electronegativity. As a result, a certain partial negative charge (δ-), which is called effective, appears on this atom, and an atom with a lower electronegativity has a partial positive charge of the same magnitude, but opposite in sign (δ+). For example, it has been experimentally established that the effective charge on the hydrogen atom in the hydrogen chloride molecule HCl is δH=+0.17, and on the chlorine atom δCl=-0.17 of the absolute electron charge.

  To determine in which direction the electron density of a polar covalent bond will shift, it is necessary to compare the electrons of both atoms. In order of increasing electronegativity, the most common chemical elements are placed in the following sequence:

  Polar molecules are called dipoles - systems in which the centers of gravity of positive charges of nuclei and negative charges of electrons do not coincide.

  A dipole is a system that is a collection of two point electric charges, equal in magnitude and opposite in sign, located at some distance from each other.

  The distance between the centers of attraction is called the length of the dipole and is denoted by the letter l. The polarity of a molecule (or bond) is quantitatively characterized by the dipole moment μ, which in the case of a diatomic molecule is equal to the product of the length of the dipole and the value of the electron charge: μ=el.

  In SI units, the dipole moment is measured in [C × m] (Coulomb meters), but more often they use the off-system unit [D] (debye): 1D = 3.33 10 -30 C × m. The value of the dipole moments of covalent molecules varies in within 0-4 D, and ionic - 4-11D. The longer the dipole length, the more polar the molecule is.

  A joint electron cloud in a molecule can be displaced by an external electric field, including the field of another molecule or ion.

  Polarizability is a change in the polarity of a bond as a result of the displacement of the electrons forming the bond under the action of an external electric field, including the force field of another particle.

  The polarizability of a molecule depends on the mobility of electrons, which is the stronger, the greater the distance from the nuclei. In addition, polarizability depends on the direction of the electric field and on the ability of electron clouds to deform. Under the action of an external field, non-polar molecules become polar, and polar molecules become even more polar, that is, a dipole is induced in the molecules, which is called a reduced or induced dipole.

  
  Scheme of the formation of an induced (reduced) dipole from a nonpolar molecule under the action of the force field of a polar particle - a dipole

  Unlike permanent ones, induced dipoles arise only under the action of an external electric field. Polarization can cause not only the polarizability of the bond, but also its rupture, in which the transition of the binding electron pair to one of the atoms occurs and negatively and positively charged ions are formed.

  The polarity and polarizability of covalent bonds determine the reactivity of molecules with respect to polar reagents.

  Properties of compounds with a covalent bond

  Substances with covalent bonds are divided into two unequal groups: molecular and atomic (or non-molecular), which are much smaller than molecular ones.

  Molecular compounds under normal conditions can be in various states of aggregation: in the form of gases (CO 2, NH 3, CH 4, Cl 2, O 2, NH 3), volatile liquids (Br 2, H 2 O, C 2 H 5 OH ) or solid crystalline substances, most of which, even with very slight heating, are able to quickly melt and sublimate easily (S 8, P 4, I 2, sugar C 12 H 22 O 11, "dry ice" CO 2).

  The low melting, sublimation, and boiling points of molecular substances are explained by the very weak forces of intermolecular interaction in crystals. That is why molecular crystals are not characterized by high strength, hardness and electrical conductivity (ice or sugar). Moreover, substances with polar molecules have higher melting and boiling points than those with non-polar molecules. Some of them are soluble in or other polar solvents. And substances with non-polar molecules, on the contrary, dissolve better in non-polar solvents (benzene, carbon tetrachloride). So, iodine, whose molecules are non-polar, does not dissolve in polar water, but dissolves in non-polar CCl 4 and low-polarity alcohol.

  Non-molecular (atomic) substances with covalent bonds (diamond, graphite, silicon Si, quartz SiO 2 , carborundum SiC and others) form extremely strong crystals, with the exception of graphite, which has a layered structure. For example, the crystal lattice of diamond is a regular three-dimensional framework in which each sp 3 hybridized carbon atom is connected to four neighboring C atoms by σ bonds. In fact, the entire diamond crystal is one huge and very strong molecule. Silicon crystals Si, which is widely used in radio electronics and electronic engineering, have a similar structure. If we replace half of the C atoms in diamond with Si atoms without disturbing the frame structure of the crystal, we get a crystal of carborundum - silicon carbide SiC - a very hard substance used as an abrasive material. And if an O atom is inserted between each two Si atoms in the crystal lattice of silicon, then the crystal structure of quartz SiO 2 is formed - also a very solid substance, a variety of which is also used as an abrasive material.

  Crystals of diamond, silicon, quartz and similar in structure are atomic crystals, they are huge "supermolecules", so their structural formulas can not be depicted in full, but only as a separate fragment, for example:

  
  Crystals of diamond, silicon, quartz

  Non-molecular (atomic) crystals, consisting of atoms of one or two elements interconnected by chemical bonds, belong to refractory substances. High melting temperatures are due to the need to expend a large amount of energy to break strong chemical bonds during the melting of atomic crystals, and not weak intermolecular interaction, as in the case of molecular substances. For the same reason, many atomic crystals do not melt when heated, but decompose or immediately pass into a vapor state (sublimation), for example, graphite sublimates at 3700 o C.

  Non-molecular substances with covalent bonds are insoluble in water and other solvents, most of them do not conduct electric current (except for graphite, which has electrical conductivity, and semiconductors - silicon, germanium, etc.).