It is characteristic of a covalent bond. Covalent bond characteristics

Covalent bond is the most common type chemical bond, carried out when interacting with the same or close values \u200b\u200bof electronegativity.

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 work 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 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 bonded 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 nucleus. The emergence of a common electronic pair of bonds can proceed through different mechanisms, more often through exchange, and sometimes through donor-acceptor ones.

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

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

according to the donor-acceptor mechanism, a two-electron bond arises from the interaction of various particles. One of them is a donor AND: has an unseparated 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 pair for bonding (an undivided 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 for the 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 unseparated pair of electrons deviated from it), and on the acceptor atom - a partial effective negative charge δ- (due to the fact that that the unseparated electron pair of the donor is shifted towards it).

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

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

Valence bond method

The first quantum mechanical theory of covalent bond was created by Geitler 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 provisions of which can be summarized as follows:

each pair of atoms in a molecule is contained together using one or more common electron pairs, with the electron orbitals of the interacting atoms overlapping;
bond strength depends on the degree of overlapping of electron orbitals;
the condition for the formation of a covalent bond is the anti-directionality of the electron spins; this gives rise to a generalized electron orbital 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 a process of mixing and alignment of orbitals in shape and energy, during which there is a redistribution of electron densities of orbitals close in energy, as a result of which they become equivalent.

The main provisions of the theory of hybridization:

During hybridization, the initial shape and orbitals mutually change, while new, hybridized orbitals are formed, but with the same energy and the same shape, resembling an irregular figure eight.
The number of hybridized orbitals is equal to the number of output orbitals participating in hybridization.
Orbitals with similar energies (s- and p-orbitals of the outer energy level and d-orbital of the outer or preliminary levels) can participate in hybridization.
Hybridized orbitals are more elongated in the direction of the formation of chemical bonds and therefore provide better overlap with the orbitals of the neighboring atom, as a result of which it becomes more durable than that formed due to the electrons of individual non-hybrid orbitals.
Due to the formation of stronger bonds and a more symmetric 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.
Hybridized orbitals should be oriented in space in such a way as to ensure maximum mutual distance from each other; in this case, the repulsive energy is the smallest.
The type of hybridization is determined by the type and number of output orbitals and changes the size of the bond angle, as well as the spatial configuration of molecules.

The shape of the hybridized orbitals and bond 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

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

General sp-hybridization scheme

The bonds, which are formed with the participation of electrons of the 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 the 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 typical for molecules of other elements (0 \u003d C \u003d 0, HC≡CH), in which bonds are formed by sp-hybridized atoms.

Scheme of sp 2 -hybridization of atomic orbitals and planar triangular shape of the molecule, which is caused by 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 BF 3, BCl 3, AlF 3 and other molecules, bonds are formed due to sp 2 -hybridized orbitals of the central atom.

Scheme of sp 3 -hybridization of atomic orbitals

Placing 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, Cl 4, C 2 H 6 and other alkanes. Examples of compounds of other elements with a tetrahedral structure due to sp 3 -hybridization of 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 are oriented with 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 the 3p orbitals of five chlorine atoms.

sp - Hybridization. The combination of one s-i one p-orbitals gives rise to two sp-hybridized orbitals located symmetrically at an angle of 180 0.
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.
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 `.
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.
Other types of hybridization. In the case of sp 3 d 2 -hybridization, six sp 3 d 2 -hybridized orbitals are directed to the vertices of the octahedron. The orientation of 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.

The method of hybridization of atomic orbitals explains the geometric structure of a large number of molecules, however, according to experimental data, molecules with slightly different 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 (~ 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 bond angle deviates from the tetrahedral: 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 unseparated electron pair at nitrogen and oxygen atoms. The two-electron orbital, which contains an unseparated pair of electrons, repels the one-electron valence orbitals due to the increased density, which leads to a decrease in the valence angle. At the nitrogen atom in the NH 3 molecule, of the four sp 3 -hybridized orbitals, three one-electron orbitals form bonds with three H atoms, and the fourth orbital contains an unseparated pair of electrons.

An unbound electron pair, which 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 unseparated 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

At the oxygen atom in the H2O 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, while the two two-electron pairs remain unseparated, that is, they belong only to the H atom. This increases the asymmetry of the electron density distribution around the O atom and decreases 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 the molecules.

Covalent bond characteristics

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

1. Valence angle Is the angle between adjacent bond axes (that is, conventional 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 unseparated 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 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, the total number of bonds that 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 on the outer level can form more bonds.

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

The directionality 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 electron orbitals have different shapes and different orientations in space, their mutual overlap can be realized in different 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 overlapping region 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 communication axis. Overlapping of valence orbitals can occur on both sides of the bond axis - the so-called "lateral" overlap, which is most often carried out when π-bonds are formed.

Pi-bond (π-bond) is an overlap of electron orbitals, in which the maximum electron density is concentrated on both sides of the line connecting the nuclei of the 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 of the formation of π-bonds between conditional A and B atoms with lateral overlap of electron orbitals

4. Multiplicity.This characteristic is determined by the number of common electron pairs that connect the atoms. The multiplicity of covalent bonds can be single (simple), double and triple. The 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 atoms are connected by a single bond (H-H), in the oxygen molecule O 2 - by a double bond (B \u003d 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, a single bond (C-C) is carried out between the C atoms, in ethylene C 2 H 4 - a double bond (C \u003d C) in acetylene C 2 H 2 - triple (C ≡ C) (C≡C).

The multiplicity of a 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 bond energy.

The multiplicity of the bond between carbon atoms: a) single σ-bond in ethane Н3С-СН3; b) double σ + π-bond in ethylene Н2С \u003d СН2; c) triple σ + π + π-bond in acetylene HC≡CH

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

Polarity is a property of a covalent bond, which is determined by the region 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 a bond in which a common electron cloud is located symmetrically relative 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 usually manifests itself in homonuclear molecules (Н 2, Cl 2, N 2, etc.) or, less often, in compounds formed by atoms of elements with close electronegativity values, for example, SiC carborundum. Polar (or heteropolar) is a bond in which the common electron cloud is asymmetric and displaced towards one of the atoms.

Molecules with a polar bond are called polar, or heteronuclear. In molecules with a polar bond, the generalized electron pair is displaced towards the atom with greater electronegativity. As a result, a certain partial negative charge (δ-) arises on this atom, which is called effective, while an atom with a lower electronegativity has a partial positive charge of the same magnitude, but opposite in sign (δ +). For example, it was experimentally found that the effective charge on the hydrogen atom in the hydrogen chloride molecule HCl is δH \u003d + 0.17, and on the chlorine atom δCl \u003d -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 ascending order of electronegativity, the most common chemical elements are arranged 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 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 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 dipole length by the value of the electron charge: μ \u003d el.

In SI units, the dipole moment is measured in [Cm × m] (Coulomb meters), but more often the off-system unit [D] (Debye) is used: 1D \u003d 3.33 · 10 -30 Cm. The value of the dipole moments of covalent molecules changes in within 0-4 D, and ionic - 4-11D. The longer the dipole is, 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 electrons forming a bond under the action of an external electric field, including force field another particle.

The polarizability of a molecule depends on the electron mobility, 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 the electron clouds to deform. Under the action of an external field, non-polar molecules become polar, and polar ones 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

In contrast to constants, 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 breaking, in which the transition of the bonding 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 in relation 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 less than molecular.

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), 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 easily sublimate (S 8, P 4, I 2, sugar C 12 H 22 O 11, "dry ice" CO 2).

Low melting, sublimation and boiling points molecular substances are explained by very weak forces of intermolecular interaction in crystals. That is why molecular crystals are not characterized by great strength, hardness and electrical conductivity (ice or sugar). Moreover, substances with polar molecules have higher melting and boiling points than with non-polar ones. 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, which has non-polar molecules, does not dissolve in polar water, but dissolves in non-polar CCl 4 and low-polarity alcohol.

Nonmolecular (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 with 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 the diamond with Si atoms, without violating the skeleton structure of the crystal, then we get a silicon carbide crystal - silicon carbide SiC - very solid matterused as an abrasive. And if an O atom is inserted in the crystal lattice of silicon between every two Si atoms, then the crystal structure of quartz SiO 2 is formed - also a very solid substance, a kind 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", therefore their structural formulas can be depicted not completely, but only in the form a separate fragment, eg:

Crystals of diamond, silicon, quartz

Nonmolecular (atomic) crystals consisting of atoms of one or two elements interconnected by chemical bonds are referred to as refractory substances. High melting temperatures are due to the need to expend a large amount of energy to break strong chemical bonds during 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 sublimes at 3700 o C.

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

The 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 forces of interaction between electric charges. Electrons and atomic nuclei have such charges.

The electrons located at the external electronic levels (valence electrons) being the farthest from the nucleus interact the weakest with it, and therefore are able to break away from the nucleus. They are responsible for binding atoms to each other.

Types of interactions in chemistry

The types of chemical bonds can be represented in the following table:

Ionic bond characteristic

Chemical interaction that is formed due to ion attractionhaving different charges is called ionic. This happens if the atoms bonded have a significant difference in electronegativity (that is, the ability to attract electrons) and the electron pair goes to a more electronegative element. The result of such a transition of electrons from one atom to another is the formation of charged particles - ions. Attraction arises between them.

The lowest electronegativity indices have typical metals, and the largest are typical non-metals. Ions are thus formed by interactions between typical metals and typical non-metals.

Metal atoms become positively charged ions (cations), donating electrons to external electronic levels, and non-metals take electrons, thus turning into negatively charged ions (anions).

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

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

The arrangement of ions is such that around each 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 with the formation of the corresponding ions:

Na 0 - 1 e \u003d Na + (cation)

Cl 0 + 1 e \u003d Cl - (anion)

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

During the formation of interaction between atoms in barium sulfide, the following processes occur:

Ba 0 - 2 e \u003d Ba 2+

S 0 + 2 e \u003d S 2-

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

Metallic chemical bond

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

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

Examples include:

Mg - 2e<-> Mg 2+

Cs - e<-> Cs +

Ca - 2e<-> Ca 2+

Fe - 3e<-> Fe 3+

Covalent: polar and non-polar

The most common type of chemical interaction is the covalent bond. The values \u200b\u200bof the electronegativity of the elements that interact do not differ sharply, in this regard, only a shift of the common electron pair occurs to a more electronegative atom.

Covalent interaction can be formed by the exchange mechanism or by the donor-acceptor mechanism.

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

Covalent ones are divided by multiplicity into:

simple or single;
double;
triple.

Doubles provide the socialization of two pairs of electrons at once, and triples - three.

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

non-polar;
polar.

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

The interaction of atoms close in electronegativity is called a non-polar bond. The common pair of electrons in such a molecule is not attracted to any of the atoms, but belongs equally to both.

The interaction of elements differing in electronegativity leads to the formation of polar bonds. In this type of interaction, common electron pairs are attracted by a more electronegative element, but do not completely transfer to it (that is, no ions are formed). As a result of such a shift in the electron density, partial charges appear on the atoms: on the more electronegative one - a negative charge, and on the less - positive.

Properties and characteristics of covalence

Main characteristics of a covalent bond:

The length is determined by the distance between the nuclei of the interacting atoms.
Polarity is determined by the displacement of the electron cloud towards one of the atoms.
Directionality is the property to form space-oriented bonds and, accordingly, molecules that have 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 when exposed to an external electric field.
The energy required to break a bond, which determines its strength.

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

H + H → H-H molecule has a single non-polar connection,

O: +: O → O \u003d O molecule has double non-polar,

Ṅ: + Ṅ: → N≡N molecule has triple non-polar.

Molecules of carbon dioxide (CO2) and carbon monoxide (CO) gas, hydrogen sulfide (H2S), hydrochloric acid (HCL), water (H2O), methane (CH4), sulfur oxide (SO2) and many others can be cited as examples of the covalent bond of chemical elements. ...

In a CO2 molecule, the relationship between carbon and oxygen atoms is covalent polar, since the more electronegative hydrogen attracts the electron density to itself. Oxygen has two unpaired electrons at the outer level, and carbon can provide four valence electrons to form interactions. As a result, double bonds are formed and the molecule looks like this: O \u003d C \u003d O.

In order to determine the type of bond in a particular molecule, it is enough to consider the atoms that make it up. Simple substances - metals form metallic, metals with non-metals - ionic, simple substances non-metals are covalent non-polar, and molecules consisting of different non-metals are formed through a covalent polar bond.

Important quantitative characteristics of the covalent bond are bond energy, her length and dipole moment.

Communication energy - the energy released during its formation, or required to separate two bound atoms. The bond energy characterizes its strength.

Link length Is the distance between the centers of bound atoms. The shorter the length, the stronger the chemical bond.

Dipole moment couplings (μ) is a vector quantity that characterizes the polarity of the coupling (measured in debyes D or pendant-meters: 1 D \u003d 3.4 · 10 -30 C · m).

The vector length is equal to the product of the link length l effective charge q , which atoms acquire when the electron density is shifted: | μ | = l · q The vector of the dipole moment is directed from positive to negative. With the vector addition of the dipole moments of all bonds, the dipole moment of the molecule is obtained.
The characteristics of links are affected by their multiplicity:

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

The term covalent bond was first introduced by the laureate Nobel Prize Irving Langmuir in 1919. The term referred to a chemical bond due to the joint possession of electrons, as opposed to a metallic bond in which electrons were free, or an ionic bond in which one of the atoms donated an electron and became a cation, and another atom took an electron and became an anion.

Later (1927) F. London and W. Heitler, using the example of a hydrogen molecule, gave the first description of a covalent bond from the point of view of quantum mechanics.

Taking into account the statistical interpretation of the Born wave function, the probability density of finding bonding electrons is concentrated in the space between the nuclei of the molecule (Fig. 1). In the theory of repulsion of electron pairs, the geometric dimensions of these pairs are considered. So, for the elements of each period, there is a certain average radius of the electron pair (Å):

0.6 for elements up to neon; 0.75 for elements up to argon; 0.75 for elements up to krypton and 0.8 for elements up to xenon.

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

The directionality 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 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 electronegativities of atoms. According to this feature, covalent bonds are subdivided 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 is shifted towards one of the atoms, thereby forming an asymmetry in the distribution of the electric charge in the molecule, giving rise to the 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 another reacting particle. The polarizability is determined by the electron mobility. The polarity and polarizability of covalent bonds determines the reactivity of molecules in relation to polar reagents.

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

The molecular hydrogen ion H 2 + contains two protons and one electron. A single electron in 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 \u003d 0.53 Å and is the center of symmetry of the molecular hydrogen ion H 2 +.

9- question) Ways of forming a covalent bond. Give examples.

Methods for the formation of a covalent bond

There are two main ways of forming a covalent bond *.

1) An electron pair forming a bond can be formed due to unpaired electrons present in unexcited atoms.

However, the number of covalent bonds can be greater than the number of unpaired electrons. For example, in the unexcited state (also called the ground state), a carbon atom has two unpaired electrons, but it is characterized by compounds in which it forms four covalent bonds. This becomes possible as a result of the excitation of the atom. In this case, one of the s-electrons goes over to the p-sublevel:

An increase in the number of created covalent bonds is accompanied by the release of more energy than is spent on the excitation of the atom. Since the valence of an atom depends on the number of unpaired electrons, excitation increases the valence. At nitrogen, oxygen, fluorine atoms, the number of unpaired electrons does not increase, because there are no free orbitals * within the second level, and the transfer of electrons to the third quantum level requires much more energy than that which would be released during the formation of additional bonds. In this way, upon excitation of an atom, transitions of electrons to free orbitals are possible only within one energy level.

Elements of the 3rd period - phosphorus, sulfur, chlorine - can exhibit a valency equal to the group number. This is achieved by excitation of atoms with the transition of 3s and 3p electrons to the vacant orbitals of the 3d sublevel:

P * 1s 2 2s 2 2p 6 3s 1 3p 3 3d 1 (valency 5)

S * 1s 2 2s 2 2p 6 3s 1 3p 3 3d 2 (valency 6)

Cl * 1s 2 2s 2 2p 6 3s 1 3p 3 3d 3 (valency 7)

In the above electronic formulas * of excited atoms, sublevels * containing only unpaired electrons are underlined. Using the chlorine atom as an example, it is easy to show that the valence can be variable:

Unlike chlorine, the valence of the F atom is constant and equal to 1, because at the valence (second) energy level, there are no d-sublevel orbitals and other vacant orbitals.

2) Covalent bonds can be formed due to paired electrons available on the outer electron layer of the atom. In this case, the second atom must have a free orbital on the outer layer. For example, the formation of an ammonium ion from an ammonia molecule and a hydrogen ion can be represented by the diagram:

An atom that provides its electron pair for the formation of a covalent bond * is called a donor, and an atom that provides an empty orbital is called an acceptor. A covalent bond formed in this way is called a donor-acceptor bond. In the ammonium cation, this bond is absolutely identical in its properties to the other three covalent bonds formed by the first method, therefore the term “donor-acceptor” does not mean any special type of bond, but only the method of its formation.

10-question) Acid-base interaction - neutralization reactions. Acid and basic salts. Give examples.

NaOH + HCl \u003d NaCl + H2O - neutralization reaction
NaOH + H2SO4 \u003d NaHSO4 + H2O - the formation of an acidic salt of sodium hydrogen sulfate, acidic salts can form other basic acids, for example H3PO4 can form 2 acidic salts of NaH2PO4. Na2HPO4. -acid salts are a product of incomplete substitution of hydrogen cations in acid.
Al (OH) 3 + 3HCl \u003d AlCl3 + 3H2O - medium salt
Al (OH) 3 + 2HCl \u003d Cl2 + 2H2O - aluminum hydroxychloride - basic salt
Al (OH) 3 + HCl \u003d Cl + H2O - aluminum dihydroxochloride
The basic salt is the product of incomplete substitution of the hydroxyl groups of the base with anions of the acid residue.

Acid and base theory - a set of fundamental physicochemical concepts that describe the nature and properties of acids and bases. They all introduce definitions of acids and bases - two classes of substances that react with each other. The task of the theory is to predict the products of the reaction between an acid and a base and the possibility of its occurrence, for which the quantitative characteristics of the strength of the acid and base are used. The differences between the theories lie in the definitions of acids and bases, the characteristics of their strength and, as a consequence, in the rules for predicting the reaction products between them. They all have their own area of \u200b\u200bapplicability, which areas partially overlap.

Acid-base interactions are extremely common in nature and are widely used in scientific and industrial practice... Theoretical concepts of acids and bases are important in the formation of all conceptual systems of chemistry and have a multifaceted influence on the development of many theoretical concepts in all major chemical disciplines.

Based modern theory acids and bases, such branches of chemical sciences as the chemistry of aqueous and non-aqueous solutions of electrolytes, pH-metry in non-aqueous media, homo- and heterogeneous acid-base catalysis, the theory of acidity functions, and many others have been developed.

11- question) Ionic bond, its properties, give examples.

Unlike a covalent bond, an ionic bond is not saturable.
Strength of ionic bonds.
Substances with ionic bonds in molecules tend to have higher boiling and melting points.

Ionic bond - a very strong chemical bond formed between atoms with a large difference (\u003e 1.5 on the Pauling scale) of electronegativities, in which the total electron pair is completely transferred to an atom with a greater electronegativity. This is the attraction of ions as oppositely charged bodies. An example is the CsF compound, in which the "degree of ionicity" is 97%. Let us consider the method of formation using the example of sodium chloride NaCl. The electronic configuration of sodium and chlorine atoms can be represented: 11 Na 1s2 2s2 2p 6 3s1; 17 Cl 1s2 2s2 2p6 Зs2 3р5. These are atoms with incomplete energy levels. Obviously, for their completion, it is easier for a sodium atom to donate one electron than to attach seven, and it is easier for a chlorine atom to attach one electron than to donate seven. In chemical interaction, the sodium atom completely donates one electron, and the chlorine atom accepts it. It can be schematically written as: Na. - l е -\u003e Na + sodium ion, stable eight-electron 1s2 2s2 2p6 shell due to the second energy level. : Cl + 1e -\u003e .Cl - chlorine ion, stable eight electron shell. Forces of electrostatic attraction arise between the Na + and Cl- ions, as a result of which a compound is formed. Ionic bond is an extreme case of polarization of a covalent polar bond. Formed between typical metal and non-metal. In this case, the electrons of 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 of electronegativities (EO\u003e 1.7 according to Pauling), then the total electron paracomplete goes to the atom with a higher EO. This results in the formation of a compound of oppositely charged ions:

An electrostatic attraction arises between the formed ions, which is called ionic bond. Rather, this look is convenient. In fact, the pure ionic bond between atoms is not realized anywhere or almost nowhere; usually, in fact, the bond is partially ionic and partially covalent. 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 unsaturation. That is why crystals formed due to ionic bonding tend to different densest packing of the corresponding ions.

Characteristic such compounds are good solubility in polar solvents (water, acids, etc.). This is due to the charging of the 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, "pull" the substance molecule apart and surround them, preventing them from reuniting. The result is ions surrounded by solvent dipoles.

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

examples: (MgS, K2CO3), bases (LiOH, Ca (OH) 2), basic oxides (BaO, Na2O)
lattice type - metal

12) Exchange reactions in solutions. Give examples.

In practically irreversible reactions the equilibrium is strongly shifted towards the formation of reaction products.

There are often processes in which weak electrolytes or poorly soluble compounds are included in the number of initial and final products of the reaction. For instance,

HCN (p) + CH 3 COO - (p) ↔ CH 3 COOH (p) + CN - (p) (1), ΔG˚ \u003d 43 kJ

NH 4 OH (p) + H + (p) ↔ H 2 O (l) + NH 4 + (p) (2) ΔG˚ \u003d -84 kJ

weak electrolytes are present on both the left and right sides of the equations.

In these cases, the equilibrium of the reversible process is shifted towards the formation of a substance with a lower Kdissotz.

In reaction (1) the equilibrium is shifted to the left K HCN \u003d 4.9 · 10 -10< K CH 3 COOH = 1,8 · 10 -5 , в реакции (2) – сильно сдвинуто вправо (K H 2 O =1,8 · 10 -16 < K NH 4 OH = 1,8 · 10 -5).

Examples of processes in the reaction equation of which sparingly soluble substances enter left and right, can serve:

AgCl (k) ↓ + NaI (p) ↔ AgI ↓ (k) + NaCl (p) (1) ΔG˚ \u003d - 54 kJ

BaCO 3 ↓ (k) + Na 2 SO 4 (p) ↔ BaSO 4 ↓ (k) + Na 2 CO 3 (p) (2) ΔG˚≈ 0

Equilibrium shifts towards the formation of a less soluble compound. In reaction (1), the equilibrium is shifted to the right, because PRAgI \u003d 1.1 · 10 -16< ПРAgCl =1,8·

10 -10. In reaction (2), the equilibrium is only slightly shifted towards BaSO 4

(PR BaCO 3 \u003d 4.9 · 10 -9\u003e PR BaSO 4 \u003d 1.08 · 10 -10).

There are processes in the equations of which, on the one hand of equality, there is a poorly soluble compound, and on the other hand, a weak electrolyte. So, the balance in the system

AgCN (k) ↓ + H + (p) ↔ HCN (p) + Ag + (p) ΔG˚ \u003d - 46kJ

is significantly shifted to the right, since the СN - ion binds more strongly to the molecule of a very weak electrolyte HCN than to the molecule of the poorly soluble substance AgCN. Therefore, the AgCN precipitate dissolves upon the addition of nitric acid.

Why can atoms connect to each other and form molecules? What is the reason for the possible existence of substances, which include atoms of completely different chemical elements? These are global issues affecting the fundamental concepts of modern physical and chemical science. You can answer them if you have an idea of electronic structure atoms and knowing the characteristics of the covalent bond, which is the basic basis for most classes of compounds. The purpose of our article is to get acquainted with the mechanisms of formation of various types of chemical bonds and compounds containing them in their molecules.

Electronic structure of the atom

Electrically neutral particles of matter, which are its structural elements, have a structure, a mirror-reflecting device Solar system... As the planets revolve around the central star - the Sun, so the electrons in the atom move around the positively charged nucleus. For the characterization of the covalent bond, the electrons located at the last energy level and most distant from the nucleus will be significant. Since their connection with the center of their own atom is minimal, they can easily be attracted by the nuclei of other atoms. This is very important for the occurrence of interatomic interactions leading to the formation of molecules. Why exactly the molecular form is the main type of existence of matter on our planet? Let's figure it out.

The main property of atoms

The ability of electrically neutral particles to interact, leading to an energy gain, is their most important feature. Indeed, under normal conditions, the molecular state of matter is more stable than atomic. The main provisions of modern atomic-molecular doctrine explain both the principles of the formation of molecules and the characteristics of the covalent bond. Recall that there can be from 1 to 8 electrons per atom; in the latter case, the layer will be complete, which means it will be very stable. Atoms have such a structure on the outer level noble gases: argon, krypton, xenon - inert elements that complete each period in the system of D. I. Mendeleev. The exception here will be helium, which has not 8, but only 2 electrons at the last level. The reason is simple: in the first period there are only two elements, the atoms of which have a single electronic layer. All other chemical elements have from 1 to 7 electrons on the last, unfinished layer. In the process of interacting with each other, the atoms will tend to fill with electrons to the octet and restore the configuration of the atom of the inert element. This state can be achieved in two ways: the loss of one's own or the acceptance of foreign negatively charged particles. These forms of interaction explain how to determine which bond - ionic or covalent - will arise between the reacting atoms.

Mechanisms for the formation of a stable electronic configuration

Imagine that two simple substances enter into the reaction of a compound: metallic sodium and gaseous chlorine. Formed a substance of the class of salts - sodium chloride. It has an ionic type of chemical bond. Why and how did it arise? Let us turn again to the atomic structure of the initial substances. Sodium has only one electron in the last layer, weakly bound to the nucleus due to large radius atom. The ionization energy for all alkali metals, which includes sodium, is low. Therefore, the electron of the outer level leaves the energy level, is attracted by the nucleus of the chlorine atom and remains in its space. This creates a precedent for the transition of the Cl atom to the form of a negatively charged ion. Now we are no longer dealing with electrically neutral particles, but with charged sodium cations and chlorine anions. In accordance with the laws of physics, forces of electrostatic attraction arise between them, and the compound forms an ionic crystal lattice. The mechanism of the formation of an ionic type of chemical bond that we have considered will help to more clearly clarify the specificity and main characteristics of a covalent bond.

Common electronic pairs

If an ionic bond arises between the atoms of elements that are very different in electronegativity, i.e., metals and non-metals, then the covalent type appears when the atoms of both the same and different non-metallic elements interact. In the first case, it is customary to talk about the non-polar, and in the other, about the polar form of the covalent bond. The mechanism of their formation is common: each of the atoms partially donates electrons for common use, which are combined in pairs. But the spatial arrangement of electron pairs relative to atomic nuclei will be unequal. On this basis, the types of covalent bonds are distinguished - non-polar and polar. Most often in chemical compounds, consisting of atoms of non-metallic elements, there are pairs consisting of electrons with opposite spins, i.e., rotating around their nuclei in opposite directions. Since the movement of negatively charged particles in space leads to the formation of electron clouds, which ultimately ends with their mutual overlap. What are the consequences of this process for atoms and what does it lead to?

Physical properties of a covalent bond

It turns out that a two-electron cloud with a high density appears between the centers of two interacting atoms. The electrostatic forces of attraction between the negatively charged cloud itself and the nuclei of atoms increase. A portion of energy is released and the distance between atomic centers decreases. For example, at the beginning of the formation of the H2 molecule, the distance between the nuclei of hydrogen atoms is 1.06 A, after the clouds overlap and the formation of a common electron pair is 0.74 A. Examples of a covalent bond formed according to the above mechanism can be found both among simple and among complex inorganic substances. Her main distinctive feature - the presence of common electronic pairs. As a result, after the appearance of a covalent bond between atoms, for example, hydrogen, each of them acquires the electronic configuration of inert helium, and the resulting molecule has a stable structure.

Spatial shape of the molecule

Another very important physical property of a covalent bond is directionality. It depends on the spatial configuration of the substance molecule. For example, when two electrons overlap with a spherical cloud, the form of the molecule is linear (hydrogen chloride or hydrogen bromide). The shape of the water molecules, in which the s- and p-clouds hybridize, is angular, and the very strong particles of gaseous nitrogen look like a pyramid.

The structure of simple substances - non-metals

Having found out what kind of bond is called covalent, what signs it has, now is the time to deal with its varieties. If atoms of the same non-metal - chlorine, nitrogen, oxygen, bromine, etc. - interact with each other, the corresponding simple substances are formed. Their common electron pairs are located at the same distance from the centers of atoms, without shifting. For compounds with a non-polar type of covalent bond, the following features are inherent: low boiling and melting points, insolubility in water, dielectric properties. Next, we will find out which substances are characterized by a covalent bond, in which there is a displacement of common electron pairs.

Electronegativity and its effect on the type of chemical bond

The property of a certain element to attract electrons to itself from the atom of another element in chemistry is called electronegativity. The scale of this parameter, proposed by L. Pauling, can be found in all textbooks on inorganic and general chemistry. Its highest value - 4.1 eV - is possessed by fluorine, the lowest - by other active non-metals, and the lowest indicator is characteristic of alkali metals. If elements differing in their electronegativity react with each other, then inevitably one, more active, will attract negatively charged particles of an atom of a more passive element to its nucleus. Thus, the physical properties of a covalent bond directly depend on the ability of the elements to donate electrons for general use. The resulting common pairs are no longer located symmetrically with respect to the nuclei, but are shifted towards the more active element.

Features of compounds with polar connection

Substances in the molecules of which the joint electron pairs are asymmetric relative to the atomic nuclei include hydrogen halides, acids, chalcogen compounds with hydrogen, and acid oxides. These are sulfate and nitrate acids, oxides of sulfur and phosphorus, hydrogen sulfide, etc. For example, a hydrogen chloride molecule contains one common electron pair formed by unpaired electrons of hydrogen and chlorine. It is shifted closer to the center of the Cl atom, which is the more electronegative element. All substances with a polar bond in aqueous solutions dissociate into ions and conduct an electric current. The compounds we have given also have higher melting and boiling points in comparison with simple non-metallic substances.

Methods for breaking chemical bonds

IN organic chemistry saturated hydrocarbons with halogens proceed by a radical mechanism. A mixture of methane and chlorine, exposed to light and at ambient temperature, reacts in such a way that chlorine molecules begin to split into particles carrying unpaired electrons. In other words, the destruction of the common electron pair and the formation of very active radicals -Cl are observed. They are able to affect methane molecules in such a way that they break the covalent bond between carbon and hydrogen atoms. An active particle -H is formed, and the free valence of the carbon atom takes on the chlorine radical, and chloromethane becomes the first reaction product. This mechanism of molecular splitting is called homolytic. If the common pair of electrons is completely transferred into the possession of one of the atoms, then they speak of a heterolytic mechanism characteristic of reactions taking place in aqueous solutions. In this case, polar water molecules will increase the rate of destruction of the chemical bonds of the compound being dissolved.

Double and triple bonds

The overwhelming majority of organic substances and some inorganic compounds contain in their molecules not one, but several common electron pairs. The multiplicity of the covalent bond reduces the distance between atoms and increases the stability of the compounds. It is customary to speak of them as chemically resistant. For example, in a nitrogen molecule there are three pairs of electrons, they are indicated in the structural formula by three dashes and determine its strength. A simple substance, nitrogen is chemically inert and can react with other compounds, for example, with hydrogen, oxygen or metals, only when heated or at elevated pressure, as well as in the presence of catalysts.

Double and triple bonds are inherent in such classes organic compounds, as unsaturated diene hydrocarbons, as well as substances of the ethylene or acetylene series. Multiple connections determine the main chemical properties: addition and polymerization reactions occurring in the places of their rupture.

In our article, we gave general characteristics covalent bond and considered its main types.

It is extremely rare that chemicals are composed of separate, unconnected atoms of chemical elements. Only a small number of gases called noble gases have such a structure under normal conditions: helium, neon, argon, krypton, xenon and radon. More often than not, chemical substances do not consist of scattered atoms, but of their combinations into various groups. Such associations of atoms can number several units, hundreds, thousands, or even more atoms. The force that keeps these atoms in the composition of such groups is called chemical bond.

In other words, we can say that a chemical bond is an interaction that provides a bond between individual atoms into more complex structures (molecules, ions, radicals, crystals, etc.).

The reason for the formation of a chemical bond is that the energy of more complex structures is less than the total energy of the individual atoms that form it.

So, in particular, if an XY molecule is formed during the interaction of atoms X and Y, this means that the internal energy of the molecules of this substance is lower than the internal energy of the individual atoms from which it was formed:

E (XY)< E(X) + E(Y)

For this reason, when chemical bonds form between individual atoms, energy is released.

The formation of chemical bonds is attended by the electrons of the outer electron layer with the lowest binding energy with the nucleus, called valence... For example, in boron, these are electrons of 2 energy levels - 2 electrons for 2 s-orbitals and 1 by 2 p-orbitals:

When a chemical bond is formed, each atom seeks to obtain an electronic configuration of atoms of noble gases, i.e. so that there are 8 electrons in its outer electron layer (2 for the elements of the first period). This phenomenon is called the octet rule.

Achievement of the electronic configuration of a noble gas by atoms is possible if initially single atoms make part of their valence electrons common to other atoms. In this case, common electron pairs are formed.

Depending on the degree of electron socialization, covalent, ionic and metallic bonds can be distinguished.

Covalent bond

A covalent bond occurs most often between the atoms of nonmetal elements. If the atoms of non-metals that form a covalent bond belong to different chemical elements, such a bond is called covalent polar. The reason for this name lies in the fact that the atoms of different elements also have a different ability to attract a common electron pair. Obviously, this leads to a shift of the common electron pair towards one of the atoms, as a result of which a partial negative charge is formed on it. In turn, a partial positive charge is formed on the other atom. For example, in a hydrogen chloride molecule, an electron pair is shifted from a hydrogen atom to a chlorine atom:

Examples of substances with a covalent polar bond:

СCl 4, H 2 S, CO 2, NH 3, SiO 2, etc.

A covalent non-polar bond is formed between the atoms of non-metals of the same chemical element. Since the atoms are identical, their ability to pull off shared electrons is the same. In this regard, no displacement of the electron pair is observed:

The above mechanism for the formation of a covalent bond, when both atoms provide electrons for the formation of common electron pairs, is called exchange.

There is also a donor-acceptor mechanism.

When a covalent bond is formed by the donor-acceptor mechanism, a common electron pair is formed due to the filled orbital of one atom (with two electrons) and the empty orbital of another atom. An atom providing a lone electron pair is called a donor, and an atom with a free orbital is called an acceptor. Atoms with paired electrons act as donors of electron pairs, for example, N, O, P, S.

For example, according to the donor-acceptor mechanism, the formation of the fourth covalent communication N-H in the ammonium cation NH 4 +:

In addition to polarity, covalent bonds are also characterized by energy. Bond energy is the minimum energy required to break a bond between atoms.

The binding energy decreases with increasing radii of the bonded atoms. Since, as we know, atomic radii increase downward along subgroups, one can, for example, conclude that the strength of the halogen-hydrogen bond increases in the series:

HI< HBr < HCl < HF

Also, the bond energy depends on its multiplicity - the greater the bond multiplicity, the more its energy. The bond multiplicity refers to the number of common electron pairs between two atoms.

Ionic bond

The ionic bond can be considered as the limiting case of the covalent polar bond. If in a covalent-polar bond the total electron pair is partially displaced to one of the pair of atoms, then in the ionic it is almost completely "given" to one of the atoms. The atom that donated the electron (s) acquires a positive charge and becomes cation, and the atom, which took the electrons from it, acquires a negative charge and becomes anion.

Thus, an ionic bond is a bond formed due to the electrostatic attraction of cations to anions.

The formation of this type of bond is characteristic of the interaction of atoms of typical metals and typical non-metals.

For example, potassium fluoride. The potassium cation is obtained as a result of the abstraction of one electron from the neutral atom, and the fluorine ion is formed when one electron is attached to the fluorine atom:

A force of electrostatic attraction arises between the resulting ions, as a result of which an ionic compound is formed.

During the formation of a chemical bond, the electrons from the sodium atom passed to the chlorine atom and oppositely charged ions were formed, which have a complete external energy level.

It was found that the electrons from the metal atom are not completely detached, but only shifted towards the chlorine atom, as in a covalent bond.

Most binary compounds that contain metal atoms are ionic. For example, oxides, halides, sulfides, nitrides.

The ionic bond also occurs between simple cations and simple anions (F -, Cl -, S 2-), as well as between simple cations and complex anions (NO 3 -, SO 4 2-, PO 4 3-, OH -). Therefore, the ionic compounds include salts and bases (Na 2 SO 4, Cu (NO 3) 2, (NH 4) 2 SO 4), Ca (OH) 2, NaOH)

Metal bond

This type of bond is formed in metals.

The atoms of all metals on the outer electron layer have electrons that have a low binding energy with the atomic nucleus. For most metals, the process of loss of external electrons is energetically favorable.

In view of such a weak interaction with the nucleus, these electrons in metals are very mobile, and the following process continuously occurs in each metal crystal:

М 0 - ne - \u003d M n +,

where M 0 is a neutral metal atom, and M n + a cation of the same metal. The figure below shows an illustration of the ongoing processes.

That is, electrons "carry" along the metal crystal, detaching from one metal atom, forming a cation from it, joining another cation, forming a neutral atom. This phenomenon was called "electronic wind", and the set of free electrons in a crystal of a non-metal atom was called "electron gas". This type of interaction between metal atoms was called a metal bond.

Hydrogen bond

If a hydrogen atom in any substance is associated with an element with high electronegativity (nitrogen, oxygen or fluorine), such a substance is characterized by such a phenomenon as a hydrogen bond.

Since a hydrogen atom is bonded to an electronegative atom, a partial positive charge is formed on the hydrogen atom and a partial negative charge on the electronegative element. In this regard, electrostatic attraction becomes possible between the partially positively charged hydrogen atom of one molecule and the electronegative atom of another. For example, a hydrogen bond is observed for water molecules:

It is the hydrogen bond that explains the abnormally high melting point of water. In addition to water, strong hydrogen bonds are also formed in substances such as hydrogen fluoride, ammonia, oxygen-containing acids, phenols, alcohols, and amines.