What is characteristic of a covalent bond. Types of chemical bonds: ionic, covalent, metallic

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

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

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

Formation of a covalent bond according to the Lewis theory

The mechanism of formation of a covalent bond

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

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

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

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

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

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

The donor-acceptor bond is otherwise called semipolar, since a partial effective positive charge δ+ arises on the donor atom (due to the fact that its undivided pair of electrons deviated from it), and on the acceptor atom a partial effective negative charge δ- (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 addition to a molecule whose central atom has an unshared electron pair, for example, to NH 3, also leads to the formation of a complex ion NH 4 +, but with a positive charge:

Valence bond method

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

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

Hybridization of atomic orbitals

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

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

The main provisions of the theory of hybridization:

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

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

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

General scheme of sp hybridization

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

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

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

Scheme of sp 3 hybridization of atomic orbitals

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

General scheme of sp 3d hybridization

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

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

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

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

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

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

Characteristics of a covalent bond

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

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

Properties of compounds with a covalent bond

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

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

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

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

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

Crystals of diamond, silicon, quartz

Non-molecular (atomic) crystals, consisting of atoms of one or two elements interconnected by chemical bonds, belong to refractory substances. High melting temperatures are due to the need to spend 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.).

covalent bond

Characteristics of a chemical bond. Hybridization.

LECTURE №3. Chemical bond and structure of molecules. Valence.

Only a few chemical elements in natural conditions are in a monatomic state (for example, inert gases). Free atoms of other elements form more complex systems - molecules with more stable electronic configurations. This phenomenon is called the formation of a chemical bond.

chemical bond - this is the interaction of two or more atoms, as a result of which a chemically stable two- or polyatomic system is formed. The formation of a chemical bond is accompanied by a decrease in the total energy of the system.

The theory of chemical bonding is based on ideas about electronic interactions. The most stable (strong) groupings of electrons are the completed outer electron layers of atoms of inert gases (two-electron for helium and eight-electron for other noble gases). The incomplete outer electron layers of all other elements are unstable, and when such atoms are combined with other atoms, their electronic shells are rearranged. A chemical bond is formed by valence electrons, but is carried out in different ways.

Valence are called electrons that participate in the formation of chemical bonds, mainly these are electrons of the last or penultimate energy level.

There are several types of chemical bonds: ionic, metallic, covalent, and hydrogen.

The simplest example of a covalent bond is the formation of a hydrogen molecule. Hydrogen atoms have an electron shell of one unpaired s-electron, i.e. one electron is missing to complete the level. When hydrogen atoms approach each other up to a certain distance, electrons with antiparallel spins interact with the formation general electron pair. A common electron pair is formed as a result of partial overlapping of s-orbitals, and in this case, the greatest density is created in the region of overlapping orbitals.

The bonding of atoms using shared electron pairs is called covalent.

A molecule with a covalent bond can be written in the form of two formulas: electronic (an electron is indicated by a dot) and structural (a shared electron pair is indicated by a bar).

1. Link length is the distance between the nuclei of atoms. Expressed in nm. A chemical bond is stronger the shorter its length. However, the measure of bond strength is its energy.

2. Bond energy - this is the amount of energy that is released during the formation of a chemical bond and, therefore, this is the work that must be spent on breaking the bond. Expressed in kJ/mol. The bond energy increases with decreasing bond length.

3. Under satiety understand the ability of atoms to form a limited number of covalent bonds. For example, a hydrogen atom, having one unpaired electron, can form one bond, and an excited carbon atom can form no more than four bonds. Due to the saturation of the bonds, the molecules have a certain composition. However, even with saturated covalent bonds, more complex molecules can be formed according to the donor-acceptor mechanism.

4. multiplicity determined by the number of common electron pairs between atoms, i.e. the number of chemical bonds. In the considered hydrogen molecule, as well as in the molecules of fluorine and chlorine, the bond between atoms is carried out due to one electron pair, such a bond is called single. In an oxygen molecule double, and in the nitrogen molecule - triple.

Moreover, a covalent bond can be of two types:

1) If electron clouds overlap in the direction of a straight line that connects the nuclei of atoms (i.e. along communication axes ), such a covalent bond is called sigma bond . Covalent sigma bonds are formed by overlapping orbitals: s-s (hydrogen molecule), s-p (hydrogen chloride) and p-p (chlorine molecule).

2) If p-orbitals directed perpendicular to the bond axis overlap, two areas of overlap are formed on both sides of the bond axis and such a bond is called pi bond .

Despite the fact that the energy of a pi bond is less than sigma, the total energy of a double, and even more so a triple bond, is higher than a single one.

5. Polarity bonds are determined by the location of a common electron pair, if it is distributed in space symmetrically with respect to the nuclei of both atoms, then such a covalent bond is called non-polar . An example is diatomic molecules consisting of atoms of the same element, i.e. simple substances.

In the case polar covalent bond , the molecule is formed by atoms of different elements and the electron cloud of the bond, in this case, is shifted to the atom with a higher relative electronegativity. For example, during the formation of the HCl molecule, the common electron pair is shifted to the chlorine atom, since it has a greater EO.

EO- this is the ability of the atoms of elements to attract common electron pairs to themselves. An atom more than an EO element takes on an effective negative charge d-, and the second atom takes on an effective positive charge d+. As a result, there is dipole. The measure of bond polarity is electric dipole moment .

6. Orientation covalent bond determines the spatial structure of molecules, i.e. their geometric shape. Directionality is quantified valence angle is the angle between chemical bonds. Covalent bonds formed by multivalent atoms always have a spatial orientation.

In addition to the characteristics common to any chemical bond (energy, length), a covalent bond has additional features: multiplicity, saturation, directivity, conjugation, polarity and polarizability.

multiplicity

One, two or three covalent bonds can form between the connected atoms.

The multiplicity (or order) of a covalent bond is characterized by the number of common electron pairs between the connected atoms.

A pair of electrons between atoms is represented by a connecting line - valentine stroke.

In the presence of one electron pair between the connected atoms, they speak of a simple (ordinary, or single) covalent bond.

For example, in molecules H 2 , F 2 , HF, H 2 O, NH 3 , CH 4 , CH 3 CH 3 or complex ions OH - , + , 2- , 2+ all bonds between atoms are ordinary and are σ-bonds.

If the connected atoms have two or three common electron pairs, there is a double or triple covalent bond, respectively, between them, while one bond is necessarily a σ-bond, the rest are π-bonds.

Examples are molecules or polyatomic ions, where there are multiple (double or triple) bonds between atoms: N≡N (nitrogen), H 2 C=CH 2 (ethylene), H 2 C=O (formaldehyde), HC≡CH ( acetylene), O=N-O - , C≡N - (cyanide - ion).

With an increase in the multiplicity of a covalent bond, its length decreases and strength increases:

However, an increase in the energy of a covalent bond, as can be seen from the given values, is not proportional to an increase in its multiplicity, which indicates a difference in the energies of the σ- and π-bonds, and E σ > E π . This is due to the fact that the efficiency of overlapping atomic orbitals in the formation of a σ-molecular orbital is higher than in the formation of a π-molecular orbital.

Saturability

Each atom is able to form a certain number of covalent bonds, thanks to which the molecules have a certain composition: H 2, H 2 O, PCl 5, CH 4.

The number of possible covalent bonds formed by a given atom depends on the number of unpaired electrons at the external energy level of the atom in the ground and excited states during the exchange mechanism, and also on the number of free orbitals at the external levels in the donor-acceptor state.

When determining the number of covalent bonds that an atom of a given element can form by the exchange mechanism, it should be taken into account that when an atom passes into an excited state, the number of its unpaired electrons can increase as a result of the depairing of some electron pairs and the transition of electrons to higher energy sublevels. If the energy expended on the excitation of the atom is small, then it can be compensated by the energy of the formed chemical bond, and the excited state of the atom is stabilized.

A small expenditure of energy is accompanied by the transitions of electrons to higher energy sublevels within the level. The transitions of electrons from the energy sublevels of one level to the sublevels of another level require a lot of energy, so the excited states of the atoms of the elements of the first three periods of the Periodic system of chemical elements of D. I. Mendeleev, resulting from such transitions, cannot be stabilized by chemical bonds.

Let us determine the valencies 1 of the atoms of the elements of the first and second periods of the periodic system of chemical elements in the ground and excited states.

The hydrogen atom has one electron, so its valency is always I.

In a helium atom, two electrons occupy l s- orbital. The depairing and transition of one of these electrons to a higher energy level requires a lot of energy, so the helium atom is chemically inert.

The valencies of lithium atoms Li, nitrogen N, oxygen O, fluorine F and neon Ne are equal to the number of unpaired electrons in the ground state, since the depairing of electron pairs of atoms of these elements is possible only when an electron passes to a higher energy level:

From the above schemes of electronic formulas, it can be seen that the valency of the lithium atom is I, nitrogen - III, oxygen - II, fluorine - I, neon - 0. In the atoms of beryllium Be, boron B and carbon C, electron pairs can be depaired and electrons can transfer from 2 s- sublevel to vacant orbitals 2 R- sublevel.

The transition to a higher energy sublevel within the level does not require a large expenditure of energy, and it can be compensated by the formation of a chemical bond. And therefore such transitions are carried out under the conditions of ordinary chemical reactions. Therefore, the valences II, III, and IV, which are inherent in the Be, B, and C atoms in the excited state, respectively, are more characteristic than the valences I and II, respectively, of the B and C atoms, which are determined by the number of unpaired R- electrons in their ground state:

Starting from the third period, the atoms R- elements upon excitation of external electrons s- and R- sublevels can move to vacant d- sublevel, which causes an increase in the number of possible chemical bonds. This explains the ability of phosphorus atoms P to form five chemical bonds (PCl 5), sulfur atoms S - four (SO 2) or six (SO 3), and chlorine atoms Cl - three, five and even seven (the so-called expansion of the octet occurs ):

In most cases, when a bond is formed, the electrons of the bonded atoms are shared. This type of chemical bond is called a covalent bond (the prefix "co-" in Latin means compatibility, "valence" - having force). The binding electrons are predominantly located in the space between the bonded atoms. Due to the attraction of the nuclei of atoms to these electrons, a chemical bond is formed. Thus, a covalent bond is a chemical bond that occurs due to an increase in electron density in the region between chemically bonded atoms.

The first theory of the covalent bond belongs to the American physical chemist G.-N. Lewis. In 1916, he suggested that the bonds between two atoms are carried out by a pair of electrons, with an eight-electron shell usually forming around each atom (the octet rule).

One of the essential properties of a covalent bond is its saturation. With a limited number of outer electrons in the regions between the nuclei, a limited number of electron pairs are formed near each atom (and, consequently, the number of chemical bonds). It is this number that is closely related to the concept of the valency of an atom in a molecule (valency is the total number of covalent bonds formed by an atom). Another important property of a covalent bond is its orientation in space. This is manifested in approximately the same geometric structure of chemical particles with similar composition. A feature of the covalent bond is also its polarizability.

To describe a covalent bond, two methods are mainly used, based on different approximations in solving the Schrödinger equation: the method of molecular orbitals and the method of valence bonds. At present, the method of molecular orbitals is used almost exclusively in theoretical chemistry. However, the method of valence bonds, despite the great complexity of calculations, gives a more visual representation of the formation and structure of chemical particles.

Covalent bond parameters

The set of atoms that form a chemical particle differs significantly from the set of free atoms. The formation of a chemical bond leads, in particular, to a change in the atomic radii and their energy. There is also a redistribution of the electron density: the probability of finding electrons in the space between the bound atoms increases.

Chemical bond length

When a chemical bond is formed, atoms always approach each other - the distance between them is less than the sum of the radii of isolated atoms:

r(A−B) r(A) + r(B)

The radius of a hydrogen atom is 53 pm, that of a fluorine atom is 71 pm, and the distance between the nuclei of atoms in an HF molecule is 92 pm:

The internuclear distance between chemically bonded atoms is called the chemical bond length.

In many cases, the bond length between atoms in a molecule of a substance can be predicted by knowing the distances between these atoms in other chemicals. The bond length between carbon atoms in diamond is 154 pm, between halogen atoms in a chlorine molecule - 199 pm. The half-sum of distances between carbon and chlorine atoms calculated from these data is 177 pm, which coincides with the experimentally measured bond length in the CCl 4 molecule. At the same time, this is not always the case. For example, the distance between hydrogen and bromine atoms in diatomic molecules is 74 and 228 pm, respectively. The arithmetic mean of these numbers is 151 pm, but the actual distance between atoms in a hydrogen bromide molecule is 141 pm, that is, noticeably less.

The distance between atoms decreases significantly with the formation of multiple bonds. The higher the bond multiplicity, the shorter the interatomic distance.

Lengths of some simple and multiple bonds

Valence angles

The direction of covalent bonds is characterized by valence angles - the angles between the lines connecting the bonded atoms. The graphic formula of a chemical particle does not carry information about bond angles. For example, in the SO 4 2− sulfate ion, the bond angles between sulfur–oxygen bonds are 109.5 o , and in the tetrachloropalladate ion 2− 90 o . The combination of bond lengths and bond angles in a chemical particle determines its spatial structure. To determine bond angles, experimental methods are used to study the structure of chemical compounds. Valence angles can be estimated theoretically based on the electronic structure of a chemical particle.

Covalent Bond Energy

A chemical compound is formed from individual atoms only if it is energetically favorable. If the attractive forces prevail over the repulsive forces, the potential energy of the interacting atoms decreases, otherwise it increases. At some distance (equal to the bond length r 0) this energy is minimal.

Thus, when a chemical bond is formed, energy is released, and when it is broken, energy is absorbed. Energy E 0 , necessary to separate the atoms and remove them from each other at a distance at which they do not interact, is called binding energy. For diatomic molecules, the binding energy is defined as the energy of dissociation of a molecule into atoms. It can be measured experimentally.

In a hydrogen molecule, the binding energy is numerically equal to the energy that is released during the formation of an H 2 molecule from H atoms:

H + H \u003d H 2 + 432 kJ

The same energy must be expended to break the H-H bond:

H 2 \u003d H + H - 432 kJ

For polyatomic molecules, this value is conditional and corresponds to the energy of such a process in which a given chemical bond disappears, while all the others remain unchanged. If there are several identical bonds (for example, for a water molecule containing two oxygen-hydrogen bonds), their energy can be calculated using Hess' law. The values ​​of the energy of the decomposition of water into simple substances, as well as the energies of the dissociation of hydrogen and oxygen into atoms are known:

2H 2 O \u003d 2H 2 + O 2; 484 kJ/mol

H 2 \u003d 2H; 432 kJ/mol

O 2 \u003d 2O; 494 kJ/mol

Given that two water molecules contain 4 bonds, the oxygen-hydrogen bond energy is:

E(О−Н) \u003d (2. 432 + 494 + 484) / 4 \u003d 460.5 kJ / mol

In molecules of composition AB n the successive detachment of B atoms is accompanied by certain (not always identical) energy expenditures. For example, the energy values ​​(kJ/mol) of successive elimination of hydrogen atoms from a methane molecule differ significantly:

	427		368		519		335
CH 4	→	CH 3	→	CH 2	→	CH	→	With

In this case, the A–B bond energy is defined as the average value of the energy expended at all stages:

CH 4 \u003d C + 4H; 1649 kJ/mol

E(С−Н) = 1649 / 4 = 412 kJ/mol

The higher the energy of a chemical bond, the stronger the bond.. The bond is considered strong or strong if its energy exceeds 500 kJ/mol (for example, 942 kJ/mol for N 2), weak - if its energy is less than 100 kJ/mol (for example, 69 kJ/mol for NO 2). If during the interaction of atoms an energy of less than 15 kJ/mol is released, then it is considered that a chemical bond is not formed, but an intermolecular interaction is observed (for example, 2 kJ/mol for Xe 2). Bond strength usually decreases with increasing bond length.

A single bond is always weaker than multiple bonds - double and triple - between the same atoms.

Energies of some simple and multiple bonds

Polarity of a covalent bond

The polarity of a chemical bond depends on the difference in the electronegativity of the bonding atoms.

Electronegativity is a conditional value that characterizes the ability of an atom in a molecule to attract electrons. If, in a diatomic molecule A–B, the bonding electrons are attracted to the B atom more strongly than to the A atom, then the B atom is considered to be more electronegative.

The electronegativity scale was used by L. Pauling for quantitative characteristics of the ability of atoms to polarize covalent bonds. For a quantitative description of electronegativity, in addition to thermochemical data, data on the geometry of molecules (Sanderson method) or spectral characteristics (Gordy method) are also used. The Allred and Rochov scale is also widely used, in which the effective nuclear charge and the atomic covalent radius are used in the calculation. The method proposed by the American physical chemist R. Mulliken (1896-1986) has the clearest physical meaning. He defined the electronegativity of an atom as half the sum of its electron affinity and ionization potential. Electronegativity values ​​based on the Mulliken method and extended to a wide range of various objects are called absolute.

Fluorine has the highest electronegativity value. The least electronegative element is cesium. The higher the difference between the electronegativity of two atoms, the more polar is the chemical bond between them.

Depending on how the redistribution of electron density occurs during the formation of a chemical bond, several types of it are distinguished. The limiting case of chemical bond polarization is the complete transition of an electron from one atom to another. In this case, two ions are formed, between which an ionic bond occurs. In order for two atoms to form an ionic bond, their electronegativity must be very different. If the electronegativities of the atoms are equal (when molecules are formed from identical atoms), the bond is called non-polar covalent. Most often found polar covalent bond - it is formed between any atoms that have different values ​​of electronegativity.

Quantifying polarity("ionic") bonds can serve as the effective charges of atoms. The effective charge of an atom characterizes the difference between the number of electrons belonging to a given atom in a chemical compound and the number of electrons in a free atom. An atom of a more electronegative element attracts electrons more strongly. Therefore, the electrons are closer to him, and he receives some negative charge, which is called effective, and his partner has the same positive charge. If the electrons that form a bond between atoms belong equally to them, the effective charges are zero. In ionic compounds, the effective charges must coincide with the charges of the ions. And for all other particles they have intermediate values.

The best method for estimating the charges of atoms in a molecule is to solve the wave equation. However, this is possible only in the presence of a small number of atoms. Qualitatively, the charge distribution can be estimated using the electronegativity scale. Various experimental methods are also used. For diatomic molecules, the polarity of the bond can be characterized and the effective charges of the atoms can be determined based on the measurement of the dipole moment:

μ = q r,

where q is the charge of the dipole pole, which is equal to the effective charge for a diatomic molecule, r− internuclear distance.

The bond dipole moment is a vector quantity. It is directed from the positively charged part of the molecule to its negative part. Based on the measurement of the dipole moment, it was found that in the HCl molecule, the hydrogen atom has a positive charge of +0.2 fractions of the electron charge, and the chlorine atom has a negative charge of −0.2. Hence, the H–Cl bond is ionic by 20%. And the Na–Cl bond is 90% ionic.

1. Spatial Orientation. If electron clouds overlap in the direction of a straight line that connects the nuclei of atoms, such a connection is called s-bond(s–s overlap H 2 , р–рCl 2 , s–рHC1).

When the p-orbitals directed perpendicular to the bond axis overlap, two overlapping regions are formed on both sides of the bond axis. Such a covalent bond is called a p-bond. For example, in a nitrogen molecule, the atoms are linked by one s-bond and two p-bonds.

The direction of the bond determines the spatial structure of the molecules, i.e., their shape and is characterized by the presence of a strictly defined angle between the bonds. For example, the angle between s-bonds in a water molecule is 104.5°.

2. Communication polarity is determined by the asymmetry in the distribution of the common electron cloud along the bond axis.

If common electron pairs are located symmetrically with respect to both nuclei, then such a covalent bond is called non-polar.

If common electron pairs are shifted to one of the atoms (they are located asymmetrically with respect to the nuclei of various atoms), then such a covalent bond is called polar.

In the case when the electron pair is shifted towards a more electronegative atom, the centers (+) and (-) of the charges do not coincide, and a system (electric dipole) arises of two equal in magnitude, but opposite in sign charges, the distance between which ( l) is called the dipole length. The measure of polarity of molecules is dipole electric moment m, equal to the product of the absolute value of the electron charge
(q = 1.6 × 10 –19 C) per dipole length l:

m = q× l.

The unit of m is debye D, 1 D = 3.33×10 -30 C×m.

Exercise. The dipole length of the HCl molecule is 2.2×10 –9 cm. Calculate the electric moment of the dipole.

2.2×10 -9 cm = 2.2×10 -11 m

m = 1.6 × 10 -19 × 2.2 × 10 -11 = 3.52 × 10 -30 C × m = 3.52 × 10 -30 / 3.33 × 10 -30 = 1.06 D.

3. multiplicity covalent bond is determined by the number of shared electron pairs that link atoms. The bond between two atoms using one pair of electrons is called simple(bonds H - C1, C - H, H - O, etc.). Communication using two electron pairs is called double(ethylene H 2 C \u003d CH 2) , using three electron pairs triple(nitrogen N N, acetylene H - C C - H).

4.Link length is the equilibrium distance between the nuclei of atoms. The bond length is expressed in nanometers (nm). 1 nm = 10 -9 m. The shorter the bond length, the stronger the chemical bond.

5. Bond energy is equal to the work that must be expended to break the connection. Express the binding energy in kilojoules per mol (kJ/mol). The bond energy increases with a decrease in the bond length and with an increase in the bond multiplicity. The process of bond formation proceeds with the release of energy (exothermic process), and the process of breaking the bond - with the absorption of energy (endothermic process).

Hybridization

Hybridization– alignment of orbitals in shape and energy.

Sp hybridization

Consider the example of beryllium hydride BeH 2 . The electronic structure of the Be atom in the normal state is 1s 2 2s 2 . A beryllium atom can interact with hydrogen atoms only in an excited state (s ® p-transition).

Be - 1s 2 2s 1 2p 1

The two bonds formed must be different in energy, since the occurrence of one is associated with the overlap of two s-orbitals, the second
swarm - s- and p-orbitals. Then the hydrogen atoms in the molecule must also be chemically unequal: one is more mobile and reactive than the other. Experimentally, this is not the case - both hydrogen atoms are energetically equivalent. To explain this phenomenon, J.K. Slater and L. Pauling suggested that "when interpreting and calculating the angles between bonds and the length of bonds, it is advisable to replace bonds that are close in energy with an equal number of energetically equivalent bonds." The connections that arise in this way are hybrid.

Thus, one s- and one p-orbital of the beryllium atom are replaced by two energetically equivalent sp-orbitals located at an angle of 180° to each other, i.e. the molecule has a linear structure.

sp 2 hybridization

Consider the example of a boron hydride ВН 3 molecule. The electronic structure of the boron atom in the normal state is B - 1s 2 2s 2 2p 1 . It can form only one covalent bond. Three covalent bonds for the boron atom are characteristic only in the excited state B* – 1s 2 2s 1 2p 2

One bond formed by the overlap of two s-orbitals of the B and H atoms does not differ energetically from the other two, formed by the overlap of the s- and p-orbitals. Three sp 2 -hybrid orbitals located at an angle of 120 about to each other, the molecule has a flat structure. A similar picture is typical for any four-atomic molecules formed by three sp 2 hybrid bonds, for example, for boron chloride (BCl 3).

sp 3 hybridization

Consider the example of methane CH 4 . In the normal state, a carbon atom with the electronic structure 1s 2 2s 2 2p 2 can give only two covalent bonds. In an excited state, it is capable of being tetravalent with an electronic structure of 1s 2 2s 1 2p 3.

One s- and three p-orbitals of the carbon atom become hybrid, and four sp 3 -hybrid, energetically equivalent orbitals are formed. The methane molecule acquires a tetrahedral structure. In the center of the tetrahedron, all vertices of which are geometrically equivalent, there is a carbon atom, and hydrogen atoms at its vertices. The angle between the bonds is 109 about 28¢.

The forces of interaction between molecules are called van der Waals or intermolecular. This interaction is due to electrostatic attraction between individual molecules and is characterized by the following features:

It acts at relatively large distances, significantly exceeding the size of the molecules themselves;

It is characterized by low energy, therefore, it significantly weakens with increasing temperature;

It is non-saturable, i.e., the interaction of this molecule with the second does not exclude a similar effect in relation to the third, fourth, etc.

With an increase in relative molar masses, the forces of intermolecular interaction increase and, as a result, the melting and boiling points of substances increase.

Exercise. Calculate the difference in the electronegativity of atoms ΔEO for the O–H and O–Mg bonds in the Mg(OH) 2 compound and determine which of these bonds is more polar. EO(H) = 2.1 eV, EO(O) = 3.5 eV, EO(Mg) = 1.2 eV.

Decision:

ΔEO(O–H) = 3.5 – 2.1 = 1.4; ΔEO(O–Mg) = 3.5 – 1.2 = 2.3.

Thus, the Mg–O bond is more polar.

When compounds are formed from elements that are very different in electronegativity (typical metals and typical non-metals), the common electron pairs are completely shifted to the more electronegative atom. For example, during the combustion of sodium in chlorine, the unpaired 3s electron of the sodium atom pairs with the 3p electron of the chlorine atom. The shared electron pair is completely shifted to the chlorine atom (Δχ(Cl) = 2.83 eV, Δχ(Cl) = 0.93 eV). In order for an ionic bond to occur, it is necessary:

1. The presence of an atom with a pronounced tendency to give up an electron with the formation of a positively charged ion (cation), i.e. with low EI. The ionization potential is the energy that must be expended to remove the 1st electron from the outer orbit. The lower the ionization potential, the easier the atom loses electrons, the more pronounced the metallic properties of the element. The ionization potential increases within a period from left to right, decreases from top to bottom.

The process of donating electrons is called oxidation.

2. The presence of an atom with a pronounced tendency to attach an electron with the formation of negatively charged ions (anions), i.e. with a large SE. The process of adding electrons is called reduction.

Cl + e ® Cl –

Typical ionic compounds are formed by the combination of metal atoms of the main subgroups of groups I and II with atoms of non-metals of the main subgroup of group VII (NaCl, KF, CaCl 2).

There is no sharp boundary between ionic and covalent bonds. In the gas phase, substances are characterized by a purely covalent polar bond, but the same substances in the solid state are characterized by an ionic bond.