Atomic or covalent chemical bond. covalent bonds

For the first time about such a concept as covalent bond chemical scientists started talking after the discovery of Gilbert Newton Lewis, who described it as the socialization of two electrons. Later studies made it possible to describe the very principle of covalent bonding. Word covalent can be considered within the framework of chemistry as the ability of an atom to form bonds with other atoms.

Let's explain with an example:

There are two atoms with slight differences in electronegativity (C and CL, C and H). As a rule, these are which are as close as possible to the structure of the electron shell of noble gases.

When these conditions are met, the nuclei of these atoms are attracted to the electron pair common to them. In this case, the electron clouds do not simply overlap each other, as in the case of a covalent bond, which ensures a reliable connection of two atoms due to the fact that the electron density is redistributed and the energy of the system changes, which is caused by the "drawing" of one atom of the electron cloud of another into the internuclear space. The more extensive the mutual overlap of electron clouds, the stronger the connection is considered.

From here, covalent bond- this is a formation that has arisen by the mutual socialization of two electrons belonging to two atoms.

As a rule, substances with a molecular crystal lattice are formed through a covalent bond. Characteristics are melting and boiling at low temperatures, poor solubility in water and low electrical conductivity. From this we can conclude: the basis of the structure of such elements as germanium, silicon, chlorine, hydrogen is a covalent bond.

Properties characteristic of this type of connection:

Saturability. This property is usually understood as the maximum number of bonds that they can establish specific atoms. This number is determined by the total number of those orbitals in the atom that can participate in the formation of chemical bonds. The valency of an atom, on the other hand, can be determined by the number of orbitals already used for this purpose.
Orientation. All atoms tend to form the strongest possible bonds. The greatest strength is achieved in the case of the coincidence of the spatial orientation of the electron clouds of two atoms, since they overlap each other. In addition, it is precisely such a property of a covalent bond as directionality that affects the spatial arrangement of molecules, that is, is responsible for their "geometric shape".
Polarizability. This position is based on the idea that there are two types of covalent bonds:

polar or asymmetrical. A bond of this type can only be formed by atoms of different types, i.e. those whose electronegativity differs significantly, or in cases where the shared electron pair is not symmetrically separated.
arises between atoms, the electronegativity of which is almost equal, and the distribution of electron density is uniform.

In addition, there are certain quantitative:

Bond energy. This parameter characterizes the polar bond in terms of its strength. Energy is understood as the amount of heat that was necessary to break the bond of two atoms, as well as the amount of heat that was released when they were combined.
Under bond length and in molecular chemistry, the length of a straight line between the nuclei of two atoms is understood. This parameter also characterizes the bond strength.
Dipole moment- a value that characterizes the polarity of the valence bond.

This article tells about what a covalent non-polar bond is. Its properties are described, the types of atoms that form it. The place of the covalent bond among other types of atomic compounds is shown.

Physics or chemistry?

There is such a phenomenon in society: one part of a homogeneous group considers the other less intelligent, more clumsy. For example, the British laugh at the Irish, the musicians who play the strings - at the cellists, the inhabitants of Russia - at the representatives of the Chukchi ethnic group. Unfortunately, science is no exception: physicists regard chemists as second-rate scientists. However, they do it in vain: it is sometimes very difficult to separate where physics is and where chemistry is. Such an example can be the methods of connecting atoms in a substance (for example, a covalent non-polar bond): the structure of an atom is unambiguously physics, the production of iron sulfide from iron and sulfur with properties that are different from both Fe and S is exactly chemistry, but here’s how from two different atoms, a homogeneous combination is obtained - neither one nor the other. This is something in between, but traditionally the science of bonds is studied as a branch of chemistry.

Electronic levels

The number and arrangement of electrons in an atom is determined by four quantum numbers: principal, orbital, magnetic, and spin. So, according to the combination of all these numbers, there are only two s-electrons in the first orbital, two s-electrons and six p-electrons in the second, and so on. As the charge of the nucleus increases, the number of electrons also increases, filling more and more new levels. The chemical properties of a substance are determined by how many and which electrons are in the shell of their atoms. A covalent bond, polar and non-polar, is formed if there is one free electron each in the outer orbitals of two atoms.

Formation of a covalent bond

To begin with, it should be noted that it is incorrect to say “orbit” and “position” in relation to electrons in the electron shell of atoms. According to the Heisenberg principle, it is impossible to determine the exact location of an elementary particle. In this case, it would be more correct to speak of an electron cloud, as if "smeared" around the nucleus at a specific distance. So, if two atoms (sometimes the same, sometimes different chemical elements) have one free electron each, they can combine them into a common orbital. Thus, both electrons belong to two atoms at once. In this way, for example, a covalent non-polar bond is formed.

Properties of covalent bonds

There are four properties of a covalent bond: directionality, saturation, polarity, polarizability. Will vary depending on their quality. Chemical properties of the resulting substance: saturation shows how many bonds this atom is able to create, directionality shows the angle between bonds, polarizability is set by the density shift towards one of the bond participants. Polarity, on the other hand, is associated with such a concept as electronegativity, and indicates how a covalent non-polar bond differs from a polar one. In general terms, the electronegativity of an atom is the ability to attract (or repel) the electrons of neighbors in stable molecules. For example, the most electronegative chemical elements are oxygen, nitrogen, fluorine, chlorine. If the electronegativity of two different atoms is the same, a covalent non-polar bond appears. Most often this happens if two atoms of the same chemical substance are combined into a molecule, for example H 2, N 2, Cl 2. But this is not necessarily the case: in PH 3 molecules, the covalent bond is also nonpolar.

Water, crystal, plasma

In nature, there are several types of bonds: hydrogen, metallic, covalent (polar, non-polar), ionic. The connection is given by the structure of the unfilled electron shell and determines both the structure and properties of the substance. As the name implies, a metallic bond is inherent only in crystals of certain chemical substances. It is the type of bond between metal atoms that determines their ability to conduct electricity. In fact, modern civilization is built on this property. Water, the most important substance for humans, is the result of the covalent bonding of one oxygen atom and two hydrogen atoms. The angle between these two junctions determines the unique properties of water. Many substances, in addition to water, have useful properties only because their atoms are connected by a covalent bond (polar and non-polar). Ionic bonding most often exists in crystals. The most revealing are the useful properties of lasers. Now they are different: with a working fluid in the form of a gas, liquid, even an organic dye. But the solid-state laser still has the optimal ratio of power, size and cost. However, a covalent nonpolar chemical bond, like other types of interaction of atoms in molecules, is inherent in substances in three ways. states of aggregation: solid, liquid, gaseous. For the fourth state of aggregation of matter, plasma, it is meaningless to speak of a connection. In fact, it is a highly ionized heated gas. However, in the plasma state there can be molecules of substances that are solid under normal conditions - metals, halogens, etc. It is noteworthy that this aggregate state of matter occupies the largest volume of the Universe: stars, nebulae, even interstellar space are a mixture different types plasma. smallest particles, which are able to break through the solar panels of communication satellites and disable the GPS system, are dusty low-temperature plasma. Thus, the world familiar to people, in which it is important to know the type chemical bond substances, represents a very small part of the universe around us.

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

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

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

Formation of a covalent bond according to the Lewis theory

The mechanism of formation of a covalent bond

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

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

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

According to the donor-acceptor mechanism, a two-electron bond arises during the interaction of various particles. One of them is a donor 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 bond (an unshared pair of electrons) is called a donor, and a particle with a free orbital that accepts this electron pair is called an acceptor.

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

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

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

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

Valence bond method

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

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

Hybridization of atomic orbitals

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

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

The main provisions of the theory of hybridization:

During hybridization, the initial shape and orbitals change mutually, while new, hybridized orbitals are formed, but with the same energy and the same shape, resembling an irregular figure eight.
The number of hybridized orbitals is equal to the number of output orbitals involved in hybridization.
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.
Hybridized orbitals are more elongated in the direction of formation of chemical bonds and therefore provide better overlap with the orbitals of the neighboring atom, as a result, it becomes stronger than the individual non-hybrid orbitals formed due to electrons.
Due to the formation of stronger bonds and a more symmetrical distribution of electron density in the molecule, an energy gain is obtained, which more than compensates for the energy consumption required for the hybridization process.
Hybridized orbitals must be oriented in space in such a way as to ensure maximum mutual separation from each other; in this case, the repulsion energy is the smallest.
The type of hybridization is determined by the type and number of exit orbitals and changes the size of the bond angle, as well as the spatial configuration of the molecules.

The form of hybridized orbitals and valence angles (geometric angles between the axes of symmetry of the orbitals) depending on the type of hybridization: a) sp-hybridization; b) sp 2 hybridization; c) sp 3 hybridization

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

General scheme of sp hybridization

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

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

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

Scheme of sp 3 hybridization of atomic orbitals

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

General scheme of sp 3d hybridization

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

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

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

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

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

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

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

Characteristics of a covalent bond

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

Properties of compounds with a covalent bond

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

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

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

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

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

Crystals of diamond, silicon, quartz

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

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

It's no secret that chemistry is a rather complex and diverse science. Many different reactions, reagents, chemicals and other complex and incomprehensible terms - they all interact with each other. But the main thing is that we deal with chemistry every day, no matter if we listen to the teacher in the lesson and learn new material or we brew tea, which in general is also a chemical process.

In contact with

Classmates

It can be concluded that chemistry is a must, to understand it and to know how our world or some of its separate parts works is interesting, and, moreover, useful.

Now we have to deal with such a term as a covalent bond, which, by the way, can be both polar and non-polar. By the way, the very word "covalent" is formed from the Latin "co" - together and "vales" - having power.

Term occurrences

Let's start with the fact that The term "covalent" was first introduced in 1919 by Irving Langmuir - Nobel Prize Laureate. The concept of "covalent" implies a chemical bond in which both atoms share electrons, which is called co-ownership. Thus, it differs, for example, from a metallic one, in which electrons are free, or from an ionic one, where one gives electrons to another. It should be noted that it is formed between non-metals.

Based on the foregoing, we can draw a small conclusion about what this process is. It arises between atoms due to the formation of common electron pairs, and these pairs arise on the outer and pre-outer sublevels of electrons.

Examples, substances with a polar:

Types of covalent bond

Two types are also distinguished - these are polar, and, accordingly, non-polar bonds. We will analyze the features of each of them separately.

Covalent polar - education

What is the term "polar"?

It usually happens that two atoms have different electronegativity, therefore, common electrons do not belong to them equally, but they are always closer to one than to the other. For example, a molecule of hydrogen chloride, in which the electrons of the covalent bond are located closer to the chlorine atom, since its electronegativity is higher than that of hydrogen. However, in reality, the difference in electron attraction is small enough for complete transfer of an electron from hydrogen to chlorine.

As a result, at polarity, the electron density shifts to a more electronegative one, and a partial negative charge arises on it. In turn, the nucleus, whose electronegativity is lower, has, accordingly, a partial positive charge.

We conclude: polar arises between various non-metals, which differ in the value of electronegativity, and electrons are located closer to the nucleus with greater electronegativity.

Electronegativity - the ability of some atoms to attract the electrons of others, thereby forming a chemical reaction.

Examples of covalent polar, substances with a covalent polar bond:

The formula of a substance with a covalent polar bond

Covalent non-polar, difference between polar and non-polar

And finally, non-polar, we will soon find out what it is.

The main difference between non-polar and polar is symmetry. If, in the case of a polar bond, the electrons were located closer to one atom, then with a non-polar bond, the electrons are arranged symmetrically, that is, equally with respect to both.

It is noteworthy that non-polar arises between non-metal atoms of one chemical element.

Eg, substances with non-polar covalent bonds:

Also, a set of electrons is often called simply an electron cloud, based on this we conclude that the electron cloud of communication, which forms a common pair of electrons, is distributed in space symmetrically, or evenly with respect to the nuclei of both.

Examples of a covalent non-polar bond and a scheme for the formation of a covalent non-polar bond

But it is also useful to know how to distinguish between covalent polar and non-polar.

covalent non-polar are always atoms of the same substance. H2. CL2.

This article has come to an end, now we know what this chemical process is, we know how to determine it and its varieties, we know the formulas for the formation of substances, and in general a little more about our complex world, success in chemistry and the formation of new formulas.

A chemical bond is the interaction of particles (ions or atoms), which is carried out in the process of exchanging electrons located at the last electronic level. There are several types of such a bond: covalent (it is divided into non-polar and polar) and ionic. In this article, we will dwell in more detail on the first type of chemical bonds - covalent. And to be more precise, in its polar form.

A covalent polar bond is a chemical bond between the valence electron clouds of neighboring atoms. The prefix "ko-" - means in this case "together", and the basis of "valence" is translated as strength or ability. Those two electrons that bond with each other are called an electron pair.

Story

The term was first used in a scientific context by Nobel Prize-winning chemist Irving Lenngryum. It happened in 1919. In his work, the scientist explained that the bond in which electrons common to two atoms are observed differs from metallic or ionic. So, it requires a separate name.

Later, already in 1927, F. London and W. Heitler, taking as an example the hydrogen molecule as the chemically and physically simplest model, described a covalent bond. They got down to business from the other end, and substantiated their observations using quantum mechanics.

The essence of the reaction

The process of converting atomic hydrogen into molecular hydrogen is a typical chemical reaction, the qualitative feature of which is a large release of heat when two electrons combine. It looks something like this: two helium atoms are approaching each other, having one electron in their orbit. Then these two clouds approach each other and form a new one, similar to a helium shell, in which two electrons already rotate.

Completed electron shells are more stable than incomplete ones, so their energy is significantly lower than that of two separate atoms. During the formation of a molecule, excess heat is dissipated in the environment.

Classification

In chemistry, there are two types of covalent bonds:

A non-polar covalent bond formed between two atoms of the same non-metallic element, such as oxygen, hydrogen, nitrogen, carbon.
A covalent polar bond occurs between atoms of different non-metals. A good example is the hydrogen chloride molecule. When atoms of two elements combine with each other, the unpaired electron from hydrogen partially passes to the last electronic level of the chlorine atom. Thus, a positive charge is formed on the hydrogen atom, and a negative charge on the chlorine atom.

Donor-acceptor bond is also a type of covalent bond. It consists in the fact that one atom from a pair provides both electrons, becoming a donor, and the atom accepting them, respectively, is considered an acceptor. When a bond is formed between atoms, the charge of the donor increases by one, and the charge of the acceptor decreases.

Semipolar bond - e It can be considered a subspecies of donor-acceptor. Only in this case, atoms unite, one of which has a complete electron orbital (halogens, phosphorus, nitrogen), and the second has two unpaired electrons (oxygen). Communication is formed in two stages:

first, one electron is removed from the lone pair and joined to the unpaired ones;
the union of the remaining unpaired electrodes, that is, a covalent polar bond is formed.

Properties

A polar covalent bond has its own physical and chemical properties, such as directionality, saturation, polarity, and polarizability. They determine the characteristics of the resulting molecules.

The direction of the bond depends on the future molecular structure of the resulting substance, namely on geometric shape, which is formed by two atoms upon attachment.

Saturation shows how many covalent bonds one atom of a substance can form. This number is limited by the number of outer atomic orbitals.

The polarity of the molecule arises because the electron cloud, formed from two different electrons, is uneven along its entire circumference. This is due to the difference in negative charge in each of them. It is this property that determines whether a bond is polar or non-polar. When two atoms of the same element combine, the electron cloud is symmetrical, which means that the bond is covalent non-polar. And if atoms of different elements combine, then an asymmetric electron cloud is formed, the so-called dipole moment of the molecule.

Polarizability reflects how actively the electrons in a molecule are displaced under the action of external physical or chemical agents, such as electrical or magnetic field, other particles.

The last two properties of the resulting molecule determine its ability to react with other polar reagents.

Sigma bond and pi bond

The formation of these bonds depends on the distribution density of electrons in the electron cloud during the formation of the molecule.

The sigma bond is characterized by the presence of a dense accumulation of electrons along the axis connecting the nuclei of atoms, that is, in the horizontal plane.

The pi bond is characterized by the compaction of electron clouds at the point of their intersection, that is, above and below the nucleus of an atom.

Visualizing Relationships in a Formula Entry

Let's take the chlorine atom as an example. Its outer electronic level contains seven electrons. In the formula, they are arranged in three pairs and one unpaired electron around the designation of the element in the form of dots.

If the chlorine molecule is written in the same way, it will be seen that two unpaired electrons have formed a pair common to two atoms, it is called shared. In addition, each of them received eight electrons.

Octet-Doublet Rule

The chemist Lewis, who proposed how a polar covalent bond is formed, was the first of his colleagues to formulate a rule explaining the stability of atoms when they are combined into molecules. Its essence lies in the fact that chemical bonds between atoms are formed when a sufficient number of electrons are socialized to obtain an electronic configuration that repeats similar to the atoms of noble elements.

That is, when molecules are formed, for their stabilization it is necessary that all atoms have a complete external electronic level. For example, hydrogen atoms, uniting into a molecule, repeat the electron shell of helium, chlorine atoms, acquire similarity at the electronic level with the argon atom.

Link length

A covalent polar bond, among other things, is characterized by a certain distance between the nuclei of the atoms that form the molecule. They are located at such a distance from each other at which the energy of the molecule is minimal. In order to achieve this, it is necessary that the electron clouds of atoms overlap each other as much as possible. There is a directly proportional pattern between the size of the atoms and the long bond. The larger the atom, the longer the bond between the nuclei.

A variant is possible when an atom forms not one, but several covalent polar bonds. Then the so-called valence angles are formed between the nuclei. They can be from ninety to one hundred and eighty degrees. They determine the geometric formula of the molecule.