Covalent bonding is the most common type of chemical bonding, carried out by interactions with the same or similar electronegativity values.

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

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

Formation of a covalent bond according to Lewis theory

Mechanism of covalent bond formation

The main feature 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 formation of a common electron bond pair can occur through different mechanisms, most often through exchange, and sometimes through donor-acceptor mechanisms.

According to the principle of the exchange mechanism of covalent bond formation, each of the interacting atoms supplies the same number of electrons with antiparallel spins to form the bond. Eg:

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

According to the donor-acceptor mechanism, a two-electron bond occurs when different particles interact. One of them is a donor A: has an unshared pair of electrons (that is, one that belongs to only one atom), and the other is an acceptor IN— has a vacant orbital.

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

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

A 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 unshared pair of electrons has deviated from it), and a partial effective negative charge δ- appears on the acceptor atom (due to , that there is a shift in its direction of the unshared 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 in the diagram as an empty quantum cell), for example BH 3, a complex complex ion BH 4 is formed — with a negative charge (N — + VN 3 ⟶⟶ [VN 4 ] -) :

The electron pair acceptor is a hydrogen ion, or simply a H + proton. Its addition 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 bonding was created by Heitler and London (in 1927) to describe the hydrogen molecule, and was later applied by Pauling to polyatomic molecules. This theory is called valence bond method, the main provisions of which can be briefly summarized as follows:

• each pair of atoms in a molecule is held together by one or more shared pairs of electrons, 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 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 from 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 turn out to be equivalent. To explain this phenomenon, the concept of “hybridization” was introduced.

Hybridization is the process of mixing and alignment of orbitals in shape and energy, during which the electron densities of orbitals close in energy are redistributed, as a result of which they become equivalent.

Basic provisions of the theory of hybridization:

1. During hybridization, the initial shape and orbitals mutually change, and new, hybridized orbitals are formed, but with the same energy and the same shape, reminiscent of 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 a neighboring atom, as a result, it becomes stronger than that formed by the electrons of individual non-hybrid orbitals.
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 compensates with a margin for the energy consumption required for the hybridization process.
6. Hybridized orbitals must be oriented in space in such a way as to ensure mutual maximum distance from each other; in this case the repulsion energy is minimal.
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 shape of hybridized orbitals and bond angles (geometric angles between the symmetry axes of orbitals) depending on the type of hybridization: a) sp-hybridization; b) sp 2 hybridization; c) sp 3 hybridization

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

General scheme of sp hybridization

The bonds that are formed with the participation of electrons from 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), the atoms of which in the valence state have unpaired s- and p-electrons. The linear form is also characteristic of 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 the 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, the atoms of which in the 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. Thus, in molecules BF 3, BCl 3, AlF 3 and other 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 molecules to have a tetrahedral shape. 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 the following 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 often found in nonmetal 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 are oriented with their elongated ends towards the corners of a mental trigonal bipyramid. This determines the shape of the PCl 5 molecule, which is formed by the overlap of five s 1 p 3 d-hybridized orbitals with the 3p-orbitals of five chlorine atoms.

1. sp - Hybridization. When one s-i and one p-orbital are combined, 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 shape 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 the 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 the 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 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 large quantity molecules, however, according to experimental data, molecules with slightly different bond angles are more often observed. For example, in the molecules CH 4, NH 3 and H 2 O, the central atoms are in the sp 3 hybridized state, so one would expect that the bond angles in them are tetrahedral (~ 109.5 0). It has been experimentally established that the bond angle in the CH 4 molecule is actually 109.5 0. However, in the NH 3 and H 2 O molecules, the value of the bond angle deviates from the tetrahedral one: it is equal to 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 unshared electron pair on 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. For 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 unbonded electron pair that occupies one of the sp 3 -hybridized orbitals directed towards the vertices of the tetrahedron, repelling the 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 a decrease in the bond angle from 109.5 0 to 107 0 as a result of the action of an unshared electron pair of the N atom is observed in the NCl 3 molecule.

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

The oxygen atom in the H 2 O molecule has two one-electron and two two-electron orbitals per four sp 3 -hybridized orbitals. One-electron hybridized orbitals participate in the formation of two bonds with two H atoms, and two two-electron pairs remain unshared, 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 unbonded electron pairs of the central atom and their placement in hybridized orbitals affects the geometric configuration of the molecules.

Characteristics of a covalent bond

A covalent bond has a set of specific properties that define it specific features, or characteristics. These, in addition to the already discussed characteristics of “bond energy” and “bond length,” include: bond angle, saturation, directionality, polarity, and the like.

1. Bond 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 magnitude of the bond angle depends on the nature of the orbitals, the type of hybridization of the central atom, and 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, by 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, by the donor-acceptor mechanism. However total The bonds an atom can form are limited.

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

Thus, of the second period, which have four orbitals at 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 larger number of orbitals at the outer level can form more bonds.

3. Focus. 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 directionality of the covalent bond.

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

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

A sigma bond (σ bond) is an overlap of electron orbitals such that the maximum electron density is concentrated along an imaginary line connecting the 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 overlap of electron orbitals; it is always single, that is, it is formed by only one electron pair.

The variety of forms of spatial orientation of “pure” orbitals and hybridized orbitals does not always allow for the possibility of overlapping orbitals on the bond axis. 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.

A pi bond (π bond) is an overlap of electron orbitals in which the maximum electron density is concentrated on either side of the line connecting the atomic nuclei (i.e., 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 with lateral overlap of electronic orbitals

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

The bond multiplicity affects the energy: as the multiplicity increases, its strength increases. Increasing the multiplicity leads to a decrease in the internuclear distance (bond length) and an increase in 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 nonpolar covalent bonds are distinguished. A nonpolar bond is a bond in which the common electron cloud is located symmetrically relative to the nuclei of the connected atoms and belongs equally to both atoms.

Molecules with this type of bond are called nonpolar or homonuclear (that is, those that contain atoms of the same element). A nonpolar bond usually manifests itself in homonuclear molecules (H 2 , Cl 2 , N 2 , etc.) or, less commonly, in compounds formed by atoms of elements with similar electronegativity values, for example, carborundum SiC. Polar (or heteropolar) is a bond in which the overall electron cloud is asymmetrical and is shifted towards one of the atoms.

Molecules with polar bonds are called polar, or heteronuclear. In molecules with a polar bond, the generalized electron pair is shifted towards the atom with higher electronegativity. As a result, a certain partial negative charge (δ-) appears on this atom, which is called effective, and an atom with lower electronegativity has a partial positive charge (δ+) of the same magnitude but opposite in sign. For example, it has been experimentally established that the effective charge on the hydrogen atom in the hydrogen chloride HCl molecule 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 the positive charges of nuclei and the negative charges of electrons do not coincide.

A dipole is a system that is a combination 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 dipole length and is designated 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 dipole length and the electron charge: μ=el.

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

The shared electron cloud in a molecule can be displaced under the influence of 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 influence 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 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 influence 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 non-polar molecule under the influence of the force field of a polar particle - dipole

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

The polarity and polarizability of covalent bonds determines the reactivity of molecules towards polar reagents.

Properties of compounds with covalent bonds

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

Under normal conditions, molecular compounds can be in various states of aggregation: in the form of gases (CO 2, NH 3, CH 4, Cl 2, O 2, NH 3), highly volatile liquids (Br 2, H 2 O, C 2 H 5 OH ) or solid crystalline substances, most of which, even with very slight heating, can quickly melt and easily sublimate (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 very weak forces intermolecular interaction in crystals. That is why molecular crystals are not characterized by great strength, hardness and electrical conductivity (ice or sugar). In this case, substances with polar molecules have higher melting and boiling points than those with non-polar ones. Some of them are soluble in or other polar solvents. On the contrary, substances with non-polar molecules dissolve better in non-polar solvents (benzene, carbon tetrachloride). Thus, iodine, whose molecules are non-polar, does not dissolve in polar water, but dissolves in non-polar CCl 4 and low-polar 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 diamond crystal lattice is a regular three-dimensional framework in which each sp 3 -hybridized carbon atom is connected to four neighboring atoms with σ bonds. In fact, the entire diamond crystal is one huge and very strong molecule. Silicon crystals, which are widely used in radio electronics and electronic engineering, have a similar structure. If you replace half of the C atoms in diamond with Si atoms without disturbing the framework structure of the crystal, you will get a crystal of carborundum - silicon carbide SiC - a very hard substance used as an abrasive material. And if in the crystal lattice of silicon an O atom is inserted between every two Si atoms, then the crystal structure of quartz SiO 2 is formed - also a very hard substance, a variety of which is also used as an abrasive material.

Crystals of diamond, silicon, quartz and similar structures are atomic crystals, they are huge “supermolecules”, therefore their structural formulas can not be depicted in full, but only in the form 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, are classified as refractory substances. High temperatures melting is due to the need to expend a large amount of energy to break strong chemical bonds when melting atomic crystals, and not weak intermolecular interactions, as in the case of molecular substances. For the same reason, many atomic crystals do not melt when heated, but decompose or immediately go 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 is inherently conductive, and semiconductors - silicon, germanium, etc.).

Atoms of most elements do not exist separately, as they can interact with each other. This interaction produces more complex particles.

The nature of a chemical bond is the action of electrostatic forces, which are the forces of interaction between electric charges. Electrons and atomic nuclei have such charges.

Electrons located on the outer electronic levels (valence electrons), being farthest from the nucleus, interact with it weakest, and therefore are able to break away from the nucleus. They are responsible for bonding atoms to each other.

Types of interactions in chemistry

Types of chemical bonds can be presented in the following table:

Characteristics of ionic bonding

Chemical reaction that occurs due to ion attraction having different charges is called ionic. This happens if the atoms being bonded have a significant difference in electronegativity (that is, the ability to attract electrons) and the electron pair goes to the more electronegative element. The result of this transfer of electrons from one atom to another is the formation of charged particles - ions. An attraction arises between them.

They have the lowest electronegativity indices typical metals, and the largest are typical non-metals. Ions are thus formed by the interaction between typical metals and typical nonmetals.

Metal atoms become positively charged ions (cations), donating electrons to their outer electron levels, and nonmetals accept electrons, thus turning into negatively charged ions (anions).

Atoms move into a more stable energy state, completing their electronic configurations.

The ionic bond is non-directional and non-saturable, since electrostatic interaction occurs in all directions, accordingly the ion can attract ions opposite sign in all directions.

The arrangement of the ions is such that around each there is a certain number of oppositely charged ions. The concept of "molecule" for ionic compounds doesn't make sense.

Examples of education

The formation of a bond in sodium chloride (nacl) is due to the transfer of an electron from the Na atom to the Cl atom to form the corresponding ions:

Na 0 - 1 e = Na + (cation)

Cl 0 + 1 e = Cl - (anion)

In sodium chloride, there are six chloride anions around the sodium cations, and six sodium ions around each chloride ion.

When interaction is formed between atoms in barium sulfide, the following processes occur:

Ba 0 - 2 e = Ba 2+

S 0 + 2 e = S 2-

Ba donates its two electrons to sulfur, resulting in the formation of sulfur anions S 2- and barium cations Ba 2+.

Metal chemical bond

The number of electrons in the outer energy levels of metals is small; they are easily separated from the nucleus. As a result of this detachment, metal ions and free electrons are formed. These electrons are called "electron gas". Electrons move freely throughout the volume of the metal and are constantly bound and separated from atoms.

The structure of the metal substance is as follows: crystal cell is the skeleton of matter, and between its nodes electrons can move freely.

The following examples can be given:

Mg - 2е<->Mg 2+

Cs-e<->Cs+

Ca - 2e<->Ca2+

Fe-3e<->Fe 3+

Covalent: polar and non-polar

The most common type of chemical interaction is a covalent bond. The electronegativity values ​​of the elements that interact do not differ sharply; therefore, only a shift of the common electron pair to a more electronegative atom occurs.

Covalent interactions can be formed by an exchange mechanism or a donor-acceptor mechanism.

The exchange mechanism is realized if each of the atoms has unpaired electrons on the outer electronic levels and the overlap of atomic orbitals leads to the appearance of a pair of electrons that already belongs to both atoms. When one of the atoms has a pair of electrons on the outer electronic level, and the other has a free orbital, then when the atomic orbitals overlap, the electron pair is shared and interacts according to the donor-acceptor mechanism.

Covalent ones are divided by multiplicity into:

• simple or single;
• double;
• triples.

Double ones ensure the sharing of two pairs of electrons at once, and triple ones - three.

According to the distribution of electron density (polarity) between bonded atoms, a covalent bond is divided into:

• non-polar;
• polar.

A nonpolar bond is formed by identical atoms, and a polar bond is formed by different electronegativity.

The interaction of atoms with similar electronegativity is called a nonpolar bond. The common pair of electrons in such a molecule is not attracted to either atom, but belongs equally to both.

The interaction of elements differing in electronegativity leads to the formation of polar bonds. In this type of interaction, shared electron pairs are attracted to the more electronegative element, but are not completely transferred to it (that is, the formation of ions does not occur). As a result of this shift in electron density, partial charges appear on the atoms: the more electronegative one has a negative charge, and the less electronegative one has a positive charge.

Properties and characteristics of covalency

Main characteristics of a covalent bond:

• The length is determined by the distance between the nuclei of interacting atoms.
• Polarity is determined by the displacement of the electron cloud towards one of the atoms.
• Directionality is the property of forming bonds oriented in space and, accordingly, molecules having certain geometric shapes.
• Saturation is determined by the ability to form a limited number of bonds.
• Polarizability is determined by the ability to change polarity under the influence of an external electric field.
• The energy required to break a bond determines its strength.

An example of a covalent non-polar interaction can be the molecules of hydrogen (H2), chlorine (Cl2), oxygen (O2), nitrogen (N2) and many others.

H· + ·H → H-H molecule has a single non-polar bond,

O: + :O → O=O molecule has a double nonpolar,

Ṅ: + Ṅ: → N≡N the molecule is triple nonpolar.

As examples of covalent bonding chemical elements we can cite molecules of carbon dioxide (CO2) and carbon monoxide (CO), hydrogen sulfide (H2S), of hydrochloric acid(HCL), water (H2O), methane (CH4), sulfur oxide (SO2) and many others.

In the CO2 molecule, the relationship between carbon and oxygen atoms is covalent polar, since the more electronegative hydrogen attracts electron density. Oxygen has two unpaired electrons in its outer shell, while carbon can provide four valence electrons to form the interaction. As a result, double bonds are formed and the molecule looks like this: O=C=O.

In order to determine the type of bond in a particular molecule, it is enough to consider its constituent atoms. Simple metal substances form a metallic bond, metals with nonmetals form an ionic bond, simple nonmetal substances form a covalent nonpolar bond, and molecules consisting of different nonmetals form through a polar covalent bond.

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

Now let's consider distinctive features 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, covalent bonds are characterized saturability– formation of a limited number of connections.

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 direction.

If a covalent bond in a molecule is formed by identical atoms or atoms with equal electronegativity, then such a bond has no polarity, that is, the electron density is distributed symmetrically. It's called non-polar covalent bond ( H2, Cl2, O2 ). Bonds can be single, double, or triple.

If the electronegativity of atoms differs, then when they combine, the electron density is distributed unevenly between the atoms and is formed covalent polar bond(HCl, H 2 O, CO), the multiplicity of which can also be different. When this type of bond is formed, the more electronegative atom acquires a partial negative charge, and the atom with less 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 connection is more pronounced, the greater the dipole moment. The molecules will be non-polar if the dipole moment is zero.

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

Categories ,

Data on ionization energy (IE), PEI and the composition of stable molecules - their actual values ​​and comparisons - both of free atoms and of atoms bound into molecules, allow us to understand how atoms form molecules through the mechanism of covalent bonding.

COVALENT BOND- (from the Latin “co” together and “vales” having force) (homeopolar bond), a chemical bond between two atoms that arises when the electrons belonging to these atoms are shared. Atoms in the molecules of simple gases are connected by covalent bonds. A bond in which there is one shared pair of electrons is called a single bond; There are also double and triple bonds.

Let's look at a few examples to see how we can use our rules to determine the number of covalent chemical bonds an atom can form if we know the number of electrons in a given atom's outer shell and the charge on its nucleus. The charge of the nucleus and the number of electrons in the outer shell are determined experimentally and are included in the table of elements.

Calculation of the possible number of covalent bonds

For example, let's count the number of covalent bonds that can form sodium ( Na), aluminum (Al), phosphorus (P), and chlorine ( Cl). Sodium ( Na) and aluminum ( Al) have, respectively, 1 and 3 electrons in the outer shell, and, according to the first rule (for the mechanism of covalent bond formation, one electron in the outer shell is used), they can form: sodium (Na)- 1 and aluminum ( Al)- 3 covalent bonds. After bond formation, the number of electrons in the outer shells of sodium ( Na) and aluminum ( Al) equal to 2 and 6, respectively; i.e., less maximum quantity(8) for these atoms. Phosphorus ( P) and chlorine ( Cl) have, respectively, 5 and 7 electrons on the outer shell and, according to the second of the above-mentioned laws, they could form 5 and 7 covalent bonds. In accordance with the fourth law, the formation of a covalent bond, the number of electrons on the outer shell of these atoms increases by 1. According to the sixth law, when a covalent bond is formed, the number of electrons on the outer shell of the bonded atoms cannot be more than 8. That is, phosphorus ( P) can only form 3 bonds (8-5 = 3), while chlorine ( Cl) can form only one (8-7 = 1).

Example: Based on the analysis, we discovered that a certain substance consists of sodium atoms (Na) and chlorine ( Cl). Knowing the regularities of the mechanism of formation of covalent bonds, we can say that sodium ( Na) can form only 1 covalent bond. Thus, we can assume that each sodium atom ( Na) bonded to the chlorine atom ( Cl) through a covalent bond in this substance, and that this substance is composed of molecules of an atom NaCl. The structural formula for this molecule: Na-Cl. Here the dash (-) denotes a covalent bond. The electronic formula of this molecule can be shown as follows:
. .
Na:Cl:
. .
In accordance with the electronic formula, on the outer shell of the sodium atom ( Na) V NaCl there are 2 electrons, and on the outer shell of the chlorine atom ( Cl) there are 8 electrons. In this formula, electrons (dots) between sodium atoms ( Na) And chlorine (Cl) are bonding electrons. Since the PEI of chlorine ( Cl) is equal to 13 eV, and for sodium (Na) it is equal to 5.14 eV, the bonding pair of electrons is much closer to the atom Cl than to an atom Na. If the ionization energies of the atoms forming the molecule are very different, then the bond formed will be polar covalent bond.

Let's consider another case. Based on the analysis, we discovered that a certain substance consists of aluminum atoms ( Al) and chlorine atoms ( Cl). In aluminum ( Al) there are 3 electrons in the outer shell; thus, it can form 3 covalent chemical bonds while chlorine (Cl), as in the previous case, can form only 1 bond. This substance is presented as AlCl3, and its electronic formula can be illustrated as follows:

Figure 3.1. Electronic formulaAlCl 3

whose formula of structure is:
Cl - Al - Cl
Cl

This electronic formula shows that AlCl3 on the outer shell of chlorine atoms ( Cl) there are 8 electrons, while the outer shell of the aluminum atom ( Al) there are 6 of them. According to the mechanism of formation of a covalent bond, both bonding electrons (one from each atom) go to the outer shells of the bonded atoms.

Multiple covalent bonds

Atoms that have more than one electron in their outer shell can form not one, but several covalent bonds with each other. Such connections are called multiple (more often multiples) connections. Examples of such bonds are the bonds of nitrogen molecules ( N= N) and oxygen ( O=O).

The bond formed when single atoms join together is called homoatomic covalent bond, e If the atoms are different, then the bond is called heteroatomic covalent bond[Greek prefixes "homo" and "hetero" respectively mean same and different].

Let's imagine what a molecule with paired atoms actually looks like. The simplest molecule with paired atoms is the hydrogen molecule.

Covalent bond(atomic bond, homeopolar bond) - a chemical bond formed by the overlap (socialization) of paravalent electron clouds. The electronic clouds (electrons) that provide communication are called shared electron pair.

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

The direction of the connection is determined by the molecular structure of the substance and the geometric shape of its molecule. The angles between two bonds are called bond angles.

Saturability is 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 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 relative 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 electric charge in the molecule, generating a dipole moment of the molecule).

The polarizability of a bond is expressed in the displacement of the 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 determines the reactivity of molecules towards polar reagents.

Education Communications

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 will be nothing more than the bond energy).

Filling of atomic (along the edges) and molecular (in the center) orbitals in the H 2 molecule with electrons. The vertical axis corresponds to the energy level, electrons are indicated by arrows reflecting their spins.

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 (MO): linking MO And anti-binding (loosening) MO. The shared electrons are located on the lower energy bonding MO.

Types of covalent bond

There are three types of covalent chemical bonds, differing in the mechanism of formation:

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

· If the atoms forming a simple covalent bond are the same, then the true charges of the atoms in the molecule are also the same, since the atoms forming the bond equally own a shared electron pair. This connection is called non-polar covalent bond. Simple substances have such a connection, for example: O 2, N 2, Cl 2. But not only nonmetals of the same type can form a covalent nonpolar bond. Non-metal elements whose electronegativity is of equal importance can also form a covalent nonpolar bond, for example, in the PH 3 molecule the bond is covalent nonpolar, since the EO of hydrogen is equal to the EO of phosphorus.

· If the atoms are different, then the degree of possession of a shared pair of electrons is determined by the difference in the electronegativity of the atoms. An atom with greater electronegativity attracts a pair of bonding electrons more strongly toward itself, and its true charge becomes negative. An atom with lower electronegativity acquires, accordingly, a positive charge of the same magnitude. If a compound is formed between two different non-metals, then such a compound is called covalent polar bond.

2. Donor-acceptor bond. To form this type of covalent bond, both electrons are provided by one of the atoms - donor. The second of the atoms involved in the formation of a bond is called acceptor. In the resulting molecule, the formal charge of the donor increases by one, and the formal charge of the acceptor decreases by one.

3. Semipolar connection. It can be considered as a polar donor-acceptor bond. This type of covalent bond is formed between an atom with a lone pair of electrons (nitrogen, phosphorus, sulfur, halogens, etc.) and an atom with two unpaired electrons (oxygen, sulfur). The formation of a semipolar bond occurs in two stages:

1. Transfer of one electron from an atom with a lone pair of electrons to an atom with two unpaired electrons. As a result, an atom with a lone pair of electrons turns into a radical cation (a positively charged particle with an unpaired electron), and an atom with two unpaired electrons turns into a radical anion (a negatively charged particle with an unpaired electron).

2. Sharing of unpaired electrons (as in the case of a simple covalent bond).

When a semipolar bond is formed, an atom with a lone pair of electrons increases its formal charge by one, and an atom with two unpaired electrons decreases its formal charge by one.

σ bond and π bond

Sigma (σ)-, pi (π)-bonds are an approximate description of the types of covalent bonds in molecules of various compounds; the σ-bond is characterized by the fact that the density of the electron cloud is maximum along the axis connecting the nuclei of atoms. When a -bond is formed, the so-called lateral overlap of electron clouds occurs, and the density of the electron cloud is maximum “above” and “below” the σ-bond plane. For example, let's take ethylene, acetylene and benzene.

In the ethylene molecule C 2 H 4 there is a double bond CH 2 = CH 2, its electronic formula: H:C::C:H. The nuclei of all ethylene atoms are located in the same plane. The three electron clouds of each carbon atom form three covalent bonds with other atoms in the same plane (with angles between them of approximately 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 the carbon atoms. The first, stronger covalent bond between carbon atoms is called a σ bond; the second, less strong covalent bond is called an -bond.

In a linear acetylene molecule

N-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 σ bonds between the same carbon atoms. Two -bonds are located above the sphere of action of the σ-bond in two mutually perpendicular planes.

All six carbon atoms of the cyclic benzene molecule C 6 H 6 lie in the same plane. There are σ bonds between carbon atoms in the plane of the ring; Each carbon atom has the same bonds with hydrogen atoms. Carbon atoms spend three electrons to make these bonds. Clouds of fourth valence electrons of carbon atoms, shaped like figures of eight, are located perpendicular to the plane of the benzene molecule. Each such cloud overlaps equally with the electron clouds of neighboring carbon atoms. In a benzene molecule, not three separate -bonds are formed, but a single -electronic system of six electrons, common to all carbon atoms. The bonds between the carbon atoms in the benzene molecule are exactly the same.

Examples of substances with covalent bonds

A simple covalent bond connects atoms in the molecules of simple gases (H 2, Cl 2, etc.) and compounds (H 2 O, NH 3, CH 4, CO 2, HCl, etc.). Compounds with a donor-acceptor bond - ammonium NH 4 +, tetrafluoroborate anion BF 4 - etc. Compounds with a semipolar bond - nitrous oxide N 2 O, O - -PCl 3 +.

Crystals with covalent bonds are dielectrics or semiconductors. Typical examples of atomic crystals (atoms in which are interconnected by covalent (atomic) bonds are diamond, germanium and silicon.

The only one known person a substance with an example of a covalent bond between a metal and a carbon is cyanocobalamin, known as vitamin B12.

Ionic bond- a very strong chemical bond formed between atoms with a large difference (> 1.5 on the Pauling scale) of electronegativity, in which the common electron pair is completely transferred to an atom with greater electronegativity. This is the attraction of ions as oppositely charged bodies. An example is the compound CsF, in which the “degree of ionicity” is 97%. Let's consider the method of formation using sodium chloride NaCl as an example. The electronic configuration of sodium and chlorine atoms can be represented as: 11 Na 1s2 2s2 2p 6 3s1; 17 Cl 1s2 2s2 2p6 3s2 3р5. These are atoms with incomplete energy levels. Obviously, to complete them, it is easier for a sodium atom to give up one electron than to gain seven, and for a chlorine atom it is easier to gain one electron than to give up seven. During a chemical interaction, the sodium atom completely gives up one electron, and the chlorine atom accepts it. Schematically, this can be written as follows: Na. - l e -> Na+ sodium ion, stable eight-electron 1s2 2s2 2p6 shell due to the second energy level. :Cl + 1е --> .Cl - chlorine ion, stable eight electron shell. Electrostatic attraction forces arise between the Na+ and Cl- ions, resulting in the formation of a compound. Ionic bonding is an extreme case of polarization of a polar covalent bond. Formed between a typical metal and non-metal. In this case, the electrons from the metal are completely transferred to the non-metal. Ions are formed.

If a chemical bond is formed between atoms that have a very large difference in electronegativity (EO > 1.7 according to Pauling), then the common electron pair is completely transferred to the atom with a higher EO. The result of this is the formation of a compound of oppositely charged ions:

An electrostatic attraction occurs between the resulting ions, which is called ionic bonding. Or rather, this look is convenient. In practice ionic bond between atoms in its pure form is not realized anywhere or almost nowhere; usually, in fact, the bond is partly ionic and partly covalent in nature. At the same time, the bond of complex molecular ions can often be considered purely ionic. The most important differences between ionic bonds and other types of chemical bonds are non-directionality and non-saturation. That is why crystals formed due to ionic bonds gravitate towards various dense packings of the corresponding ions.

Characteristics Such compounds have good solubility in polar solvents (water, acids, etc.). This occurs due to the charged parts of the molecule. In this case, the dipoles of the solvent are attracted to the charged ends of the molecule, and, as a result of Brownian motion, they “tear” the molecule of the substance into pieces and surround them, preventing them from connecting again. The result is ions surrounded by solvent dipoles.

When such compounds are dissolved, energy is usually released, since the total energy of the formed solvent-ion bonds is greater than the energy of the anion-cation bond. Exceptions are many salts of nitric acid (nitrates), which absorb heat when dissolved (solutions cool). Last fact explained on the basis of laws that are considered in physical chemistry.