Chemical bond: definition, types, classification and features of the definition. Types of chemical bonds The type of chemical bond formed between atoms is determined by

The smallest particle of a substance is a molecule formed as a result of the interaction of atoms between which chemical bonds or chemical bonds act. The doctrine of chemical bonding forms the basis theoretical chemistry. A chemical bond occurs when two (sometimes more) atoms interact. Bond formation occurs with the release of energy.

A chemical bond is an interaction that binds individual atoms into molecules, ions, and crystals.

The chemical bond is uniform in nature: it is of electrostatic origin. But in various chemical compounds the chemical bond is of different types; The most important types of chemical bonds are covalent (non-polar, polar), ionic, and metallic. Varieties of these types of bonds are donor-acceptor, hydrogen, etc. Between metal atoms there is metal connection.

A chemical bond carried out through the formation of a common, or shared, pair or several pairs of electrons is called covalent. Each atom contributes one electron to the formation of one common pair of electrons, i.e. participates “in equal share” (Lewis, 1916). Below are diagrams of the formation of chemical bonds in the molecules H2, F2, NH3 and CH4. Electrons belonging to different atoms are represented by different symbols.

As a result of the formation of chemical bonds, each of the atoms in the molecule has a stable two- and eight-electron configuration.

When a covalent bond occurs, the electron clouds of atoms overlap to form a molecular electron cloud, accompanied by a gain in energy. The molecular electron cloud is located between the centers of both nuclei and has an increased electron density compared to the density of the atomic electron cloud.

The implementation of a covalent bond is possible only in the case of antiparallel spins of unpaired electrons belonging to different atoms. With parallel electron spins, atoms do not attract, but repel: a covalent bond does not occur. The method of describing a chemical bond, the formation of which is associated with a common electron pair, is called the valence bond method (VBC).

Basic provisions of the MBC

A covalent chemical bond is formed by two electrons with opposite spins, and this electron pair belongs to two atoms.

The more the interacting electron clouds overlap, the stronger the covalent bond.

When writing structural formulas, electron pairs that determine the bond are often depicted with dashes (instead of dots representing shared electrons).

The energy characteristics of a chemical bond are important. When a chemical bond is formed, the total energy of the system (molecule) is less than the energy components(atoms), i.e. EAB<ЕА+ЕB.

Valency is the property of an atom of a chemical element to attach or replace a certain number of atoms of another element. From this point of view, the valency of an atom is most easily determined by the number of hydrogen atoms forming chemical bonds with it, or by the number of hydrogen atoms replaced by an atom of this element.

With the development of quantum mechanical concepts of the atom, valence began to be determined by the number of unpaired electrons involved in the formation of chemical bonds. In addition to unpaired electrons, the valence of an atom also depends on the number of empty and fully filled orbitals of the valence electron layer.

Binding energy is the energy released when a molecule is formed from atoms. Binding energy is usually expressed in kJ/mol (or kcal/mol). This is one of the most important characteristics of a chemical bond. The system that contains less energy is more stable. It is known, for example, that hydrogen atoms tend to unite into a molecule. This means that a system consisting of H2 molecules contains less energy than a system consisting of the same number of H atoms, but not combined into molecules.

Rice. 2.1 Dependence of the potential energy E of a system of two hydrogen atoms on the internuclear distance r: 1 - during the formation of a chemical bond; 2 – without her education.

Figure 2.1 shows an energy curve characteristic of interacting hydrogen atoms. The approach of atoms is accompanied by the release of energy, which will be greater the more the electron clouds overlap. However, under normal conditions, due to Coulomb repulsion, it is impossible to achieve the fusion of the nuclei of two atoms. This means that at some distance, instead of the attraction of atoms, their repulsion will occur. Thus, the distance between atoms r0, which corresponds to the minimum on the energy curve, will correspond to the length of the chemical bond (curve 1). If the electron spins of the interacting hydrogen atoms are the same, then their repulsion will occur (curve 2). The binding energy for different atoms varies within the range of 170–420 kJ/mol (40–100 kcal/mol).

The process of electron transition to a higher energy sublevel or level (i.e., the process of excitation or vaporization, which was discussed earlier) requires energy. When a chemical bond is formed, energy is released. In order for a chemical bond to be stable, it is necessary that the increase in atomic energy due to excitation be less than the energy of the chemical bond being formed. In other words, it is necessary that the energy expended on the excitation of atoms be compensated by the release of energy due to the formation of a bond.

A chemical bond, in addition to bond energy, is characterized by length, multiplicity and polarity. For a molecule consisting of more than two atoms, the angles between the bonds and the polarity of the molecule as a whole are significant.

The multiplicity of a bond is determined by the number of electron pairs connecting two atoms. Thus, in ethane H3C–CH3 the bond between carbon atoms is single, in ethylene H2C=CH2 it is double, in acetylene HCºCH it is triple. As the bond multiplicity increases, the bond energy increases: the C–C bond energy is 339 kJ/mol, C=C - 611 kJ/mol and CºC - 833 kJ/mol.

The chemical bond between atoms is caused by the overlap of electron clouds. If the overlap occurs along the line connecting the atomic nuclei, then such a bond is called a sigma bond (σ bond). It can be formed by two s-electrons, s- and p-electrons, two px-electrons, s and d electrons (for example):

A chemical bond carried out by one electron pair is called a single bond. A single bond is always a σ bond. Type s orbitals can only form σ bonds.

The bond between two atoms can be accomplished by more than one pair of electrons. This relationship is called multiple. An example of the formation of a multiple bond is the nitrogen molecule. In a nitrogen molecule, the px orbitals form one σ bond. When a bond is formed by pz orbitals, two regions arise
overlaps – above and below the x-axis:

Such a bond is called a pi bond (π bond). The formation of a π bond between two atoms occurs only when they are already connected by a σ bond. The second π bond in the nitrogen molecule is formed by the py orbitals of the atoms. When π bonds are formed, the electron clouds overlap less than in the case of σ bonds. As a result, π bonds are generally less strong than σ bonds formed by the same atomic orbitals.

p orbitals can form both σ and π bonds; in multiple bonds, one of them is necessarily a σ-bond: .

Thus, out of three bonds in a nitrogen molecule, one is a σ bond and two are π bonds.

The bond length is the distance between the nuclei of bonded atoms. The bond lengths in various compounds are tenths of a nanometer. As the multiplicity increases, the bond lengths decrease: the bond lengths N–N, N=N and NºN are equal to 0.145; 0.125 and 0.109 nm (10-9 m), and the lengths of the C-C, C=C and CºC bonds are, respectively, 0.154; 0.134 and 0.120 nm.

Between different atoms, a pure covalent bond can appear if the electronegativity (EO)1 of the atoms is the same. Such molecules are electrosymmetric, i.e. The “centers of gravity” of the positive charges of nuclei and the negative charges of electrons coincide at one point, which is why they are called non-polar.

If the connecting atoms have different EO, then the electron cloud located between them shifts from a symmetrical position closer to the atom with a higher EO:

The displacement of the electron cloud is called polarization. As a result of unilateral polarization, the centers of gravity of positive and negative charges in a molecule do not coincide at one point, and a certain distance (l) appears between them. Such molecules are called polar or dipoles, and the bond between the atoms in them is called polar.

A polar bond is a type of covalent bond that has undergone slight one-sided polarization. The distance between the "centers of gravity" of positive and negative charges in a molecule is called the dipole length. Naturally, the greater the polarization, the greater the length of the dipole and the greater the polarity of the molecules. To assess the polarity of molecules, they usually use the permanent dipole moment (Mp), which is the product of the value of the elementary electric charge (e) and the length of the dipole (l), i.e. .

Dipole moments are measured in debyes D (D = 10-18 electric units × cm, since the elementary charge is 4.810-10 electric units, and the dipole length is on average equal to the distance between two atomic nuclei, i.e. 10-8 cm) or coulometers (C×m) (1 D = 3.33·10-30 C×m) (electron charge 1.6·10-19 C multiplied by the distance between charges, for example, 0.1 nm, then Mr = 1.6 10-19 × 1 × 10-10 = 1.6 10-29 C m). The permanent dipole moments of molecules range from zero to 10 D.

For non-polar molecules l = 0 and Мр = 0, i.e. they do not have a dipole moment. For polar molecules, Мр > 0 and reaches values of 3.5 – 4.0 D.

With a very large difference in EO between atoms, there is a clear one-sided polarization: the electron cloud of the bond is shifted as much as possible towards the atom with the highest EO, the atoms turn into oppositely charged ions and an ionic molecule appears:

The covalent bond becomes ionic. The electrical asymmetry of the molecules increases, the length of the dipole increases, and the dipole moment increases to 10 D.

The total dipole moment of a complex molecule can be considered equal to the vector sum of the dipole moments of individual bonds. The dipole moment is usually considered to be directed from the positive end of the dipole to the negative.

Bond polarity can be predicted using the relative EO of atoms. How greater difference relative EO of atoms, the more pronounced the polarity: DEO = 0 – non-polar covalent bond; DEO = 0 – 2 – polar covalent bond; DEO = 2 – ionic bond. It is more correct to talk about the degree of ionicity of a bond, since bonds are not 100% ionic. Even in the CsF compound the bond is only 89% ionic.

A chemical bond that arises due to the transfer of electrons from atom to atom is called ionic, and the corresponding molecules of chemical compounds are called ionic. Ionic compounds in the solid state are characterized by an ionic crystal lattice. In a molten and dissolved state they conduct electricity, have high melting and boiling points and a significant dipole moment.

If we consider compounds of elements of any period with the same element, then as we move from the beginning to the end of the period, the predominantly ionic nature of the bond changes to covalent. For example, in the fluorides of the 2nd period LiF, BeF2, CF4, NF3, OF2, F2, the degree of ionicity of the bond from lithium fluoride gradually weakens and is replaced by a typically covalent bond in the fluorine molecule.

Thus, the nature of the chemical bond is the same: there is no fundamental difference in the mechanism of formation of polar covalent and ionic bonds. These types of bonds differ only in the degree of polarization of the electron cloud of the molecule. The resulting molecules differ in the lengths of the dipoles and the values of the permanent dipole moments. In chemistry, the dipole moment is very important. As a general rule, the larger the dipole moment, the higher the reactivity of the molecules.

Mechanisms of chemical bond formation

The valence bond method distinguishes between exchange and donor-acceptor mechanisms for the formation of a chemical bond.

Exchange mechanism. The exchange mechanism for the formation of a chemical bond includes cases when one electron from each atom participates in the formation of an electron pair.

In H2, Li2, Na2 molecules, bonds are formed due to the unpaired s-electrons of the atoms. In F2 and Cl2 molecules - due to unpaired p-electrons. In HF and HCl molecules, bonds are formed by s-electrons of hydrogen and p-electrons of halogens.

A feature of the formation of compounds by the exchange mechanism is saturation, which shows that the atom forms not any, but a limited number of bonds. Their number, in particular, depends on the number of unpaired valence electrons.

From the quantum cells N and H we can see that the nitrogen atom has 3

unpaired electrons, and the hydrogen atom has one. The principle of saturation indicates that the stable compound should be NH3 and not NH2, NH or NH4. However, there are molecules containing an odd number of electrons, for example, NO, NO2, ClO2. All of them are characterized by increased reactivity.

At certain stages chemical reactions Valently unsaturated groups can also be formed, which are called radicals, for example, H, NH2, O, CH3. The reactivity of radicals is very high and therefore their lifetime is usually short.

Donor-acceptor mechanism

It is known that valence-saturated compounds ammonia NH3 and boron trifluoride BF3 react with each other according to the reaction

NH3 + BF3 = NH3BF3 + 171.4 kJ/mol.

Let's consider the mechanism of this reaction:

It can be seen that out of four boron orbitals, three are occupied, and one remains vacant. In the ammonia molecule, all four nitrogen orbitals are occupied, three of them are occupied by the exchange mechanism with electrons of nitrogen and hydrogen, and one contains an electron pair, both electrons of which belong to nitrogen. Such an electron pair is called a lone electron pair. The formation of the H3N · BF3 compound occurs due to the fact that the lone electron pair of ammonia occupies the vacant orbital of boron fluoride. In this case, the potential energy of the system decreases and an equivalent amount of energy is released. Such a formation mechanism is called donor-acceptor; a donor is an atom that donates its electron pair to form a bond (in in this case nitrogen atom); and the atom that, by providing a vacant orbital, accepts an electron pair, is called an acceptor (in this case, a boron atom). A donor-acceptor bond is a type of covalent bond.

In the H3N · BF3 compound, nitrogen and boron are tetravalent. The nitrogen atom increases its valency from 3 to 4 as a result of using a lone pair of electrons to form an additional chemical bond. The boron atom increases its valency due to the presence of a free orbital at the valence electronic level. Thus, the valence of elements is determined not only by the number of unpaired electrons, but also by the presence of lone electron pairs and free orbitals at the valence electronic level.

A simpler case of the formation of a chemical bond by the donor-acceptor mechanism is the reaction of ammonia with a hydrogen ion:

. The role of electron pair acceptor is played by the empty orbital of the hydrogen ion. In the ammonium ion NH4+, the nitrogen atom is tetravalent.

Directionality of bonds and hybridization of atomic orbitals

An important characteristic of a molecule consisting of more than two atoms is its geometric configuration. It is determined by the mutual arrangement of atomic orbitals involved in the formation of chemical bonds.

Overlapping of electron clouds is possible only with a certain relative orientation of the electron clouds; in this case, the overlap region is located in a certain direction with respect to the interacting atoms.

When an ionic bond is formed electric field The ion has spherical symmetry and therefore the ionic bond is not directional and saturable.

k.ch. = 6 k.h. = 6

The angle between bonds in a water molecule is 104.5°. Its magnitude can be explained on the basis of quantum mechanical concepts. Electronic circuit oxygen atom 2s22p4. Two unpaired p-orbitals are located at an angle of 90° to each other - the maximum overlap of electron clouds of s-orbitals of hydrogen atoms with p-orbitals of an oxygen atom will occur if the bonds are located at an angle of 90°. In a water molecule, the O–H bond is polar. The effective positive charge on the hydrogen atom is δ+, on the oxygen atom - δ-. Therefore, the increase in the angle between bonds to 104.5° is explained by the repulsion of the effective positive charges of hydrogen atoms, as well as electron clouds.

The electronegativity of sulfur is significantly less than the EO of oxygen. Therefore, the polarity of the H–S bond in H2S is less than the polarity of the H–O bond in H2O, and the length of the H–S bond (0.133 nm) is greater than H–O (0.56 nm) and the angle between the bonds approaches a right angle. For H2S it is 92o, and for H2Se – 91o.

For the same reasons, the ammonia molecule has a pyramidal structure and the angle between the H–N–H valence bonds is greater than straight (107.3°). When moving from NH3 to PH3, AsH3 and SbH3, the angles between the bonds are 93.3°, respectively; 91.8o and 91.3o.

Hybridization of atomic orbitals

An excited beryllium atom has a configuration of 2s12p1, an excited boron atom has a configuration of 2s12p2, and an excited carbon atom has a configuration of 2s12p3. Therefore, we can assume that not the same, but different atomic orbitals can participate in the formation of chemical bonds. For example, in compounds such as BeCl2, BeCl3, CCl4 there should be bonds of unequal strength and direction, and σ-bonds from p-orbitals should be stronger than bonds from s-orbitals, because for p-orbitals there are more favorable conditions for overlapping. However, experience shows that in molecules containing central atoms with different valence orbitals (s, p, d), all bonds are equivalent. An explanation for this was given by Slater and Pauling. They concluded that different orbitals, not very different in energy, form a corresponding number of hybrid orbitals. Hybrid (mixed) orbitals are formed from different atomic orbitals. The number of hybrid orbitals is equal to the number of atomic orbitals involved in hybridization. Hybrid orbitals are identical in electron cloud shape and energy. Compared to atomic orbitals, they are more elongated in the direction of formation of chemical bonds and therefore provide better overlap of electron clouds.

Hybridization of atomic orbitals requires energy, so hybrid orbitals in an isolated atom are unstable and tend to turn into pure AOs. When chemical bonds are formed, the hybrid orbitals are stabilized. Due to the stronger bonds formed by the hybrid orbitals, more energy is released from the system and therefore the system becomes more stable.

sp-hybridization occurs, for example, during the formation of Be, Zn, Co and Hg (II) halides. In the valence state, all metal halides contain s and p-unpaired electrons at the appropriate energy level. When a molecule is formed, one s and one p orbital form two hybrid sp orbitals at an angle of 180°.

Experimental data show that Be, Zn, Cd and Hg(II) halides are all linear and both bonds are of the same length.

sp2 hybridization. As a result of the hybridization of one s-orbital and two p-orbitals, three hybrid sp2 orbitals are formed, located in the same plane at an angle of 120° to each other.

sp3 hybridization is characteristic of carbon compounds. As a result of the hybridization of one s-orbital and three p-orbitals, four hybrid sp3 orbitals are formed, directed towards the vertices of the tetrahedron with an angle between the orbitals of 109.5°.

Hybridization is manifested in the complete equivalence of the bonds of a carbon atom with other atoms in compounds, for example, in CH4, CCl4, C(CH3)4, etc.

Hybridization can involve not only s- and p-orbitals, but also d- and f-orbitals.

With sp3d2 hybridization, 6 equal clouds are formed. It is observed in such compounds as,.

Ideas about hybridization make it possible to understand such structural features of molecules that cannot be explained in any other way.

Hybridization of atomic orbitals (AO) leads to a displacement of the electron cloud in the direction of forming bonds with other atoms. As a result, the overlap areas of hybrid orbitals turn out to be larger than for pure orbitals and the bond strength increases.

Polarizability and polarizing effect of ions and molecules

In an electric field, an ion or molecule is deformed, i.e. in them there is a relative displacement of nuclei and electrons. This deformability of ions and molecules is called polarizability. Since the electrons of the outer layer are least tightly bound in the atom, they experience displacement first.

The polarizability of anions, as a rule, is significantly higher than the polarizability of cations.

With the same structure of electronic shells, the polarizability of the ion decreases as the positive charge increases, for example, in the series:

For ions of electronic analogues, polarizability increases with increasing number of electronic layers, for example: or.

The polarizability of molecules is determined by the polarizability of their constituent atoms, geometric configuration, number and multiplicity of bonds, etc. A conclusion about relative polarizability is possible only for similarly constructed molecules that differ in one atom. In this case, the difference in the polarizability of molecules can be judged by the difference in the polarizability of atoms.

An electric field can be created either by a charged electrode or by an ion. Thus, the ion itself can have a polarizing effect (polarization) on other ions or molecules. The polarizing effect of an ion increases with increasing its charge and decreasing radius.

The polarizing effect of anions is, as a rule, much less than the polarizing effect of cations. This is explained large sizes anions versus cations.

Molecules have a polarizing effect if they are polar; The greater the dipole moment of the molecule, the higher the polarizing effect.

The polarizing ability increases in the series, because the radii increase and the electric field created by the ion decreases.

Hydrogen bond

A hydrogen bond is special kind chemical bond. It is known that hydrogen compounds with highly electronegative nonmetals, such as F, O, N, have an anomalous high temperatures boiling. If in the series H2Te – H2Se – H2S the boiling point naturally decreases, then when moving from H2S to H2O there is a sharp jump to an increase in this temperature. The same picture is observed in the series of hydrohalic acids. This indicates the presence of a specific interaction between H2O molecules and HF molecules. Such interaction should make it difficult for molecules to separate from each other, i.e. reduce their volatility, and, consequently, increase the boiling point of the corresponding substances. Due to the large difference in EO, the chemical bonds H–F, H–O, H–N are highly polarized. Therefore, the hydrogen atom has a positive effective charge (δ+), and the F, O and N atoms have an excess of electron density and are negatively charged (d-). Due to Coulomb attraction, the positively charged hydrogen atom of one molecule interacts with the electronegative atom of another molecule. Due to this, the molecules are attracted to each other (thick dots indicate hydrogen bonds).

A hydrogen bond is a bond that is formed through a hydrogen atom that is part of one of two bonded particles (molecules or ions). The energy of a hydrogen bond (21–29 kJ/mol or 5–7 kcal/mol) is approximately 10 times less than the energy of a conventional chemical bond. Nevertheless, the hydrogen bond determines the existence of dimeric molecules (H2O)2, (HF)2 and formic acid in pairs.

In a series of combinations of atoms HF, HO, HN, HCl, HS, the energy of the hydrogen bond decreases. It also decreases with increasing temperature, so substances in the vapor state exhibit hydrogen bonding only to a small extent; it is characteristic of substances in liquid and solid states. Substances such as water, ice, liquid ammonia, organic acids, alcohols and phenols are associated into dimers, trimers and polymers. IN liquid state dimers are the most stable.

Intermolecular interactions

Previously, the bonds that determine the formation of molecules from atoms were considered. However, there is also interaction between molecules. It causes gases to condense and transform into liquids and solids. The first formulation of the forces of intermolecular interaction was given in 1871 by Van der Waals. Therefore, they are called van der Waals forces. Intermolecular interaction forces can be divided into orientational, inductive and dispersive.

Polar molecules, due to the electrostatic interaction of opposite ends of dipoles, are oriented in space so that the negative ends of the dipoles of some molecules are turned to positive

ends of dipoles of other molecules (orientational intermolecular interaction).

The energy of such interaction is determined by the electrostatic attraction of two dipoles. The larger the dipole, the stronger the intermolecular attraction (H2O, HCl).

The thermal movement of molecules prevents the mutual orientation of molecules, therefore, with increasing temperature, the orientation effect weakens. Inductive interaction is also observed in substances with polar molecules, but it is usually much weaker than orientational interaction.

A polar molecule can increase the polarity of a neighboring molecule. In other words, under the influence of the dipole of one molecule, the dipole of another molecule can increase, and a non-polar molecule can become polar:

The dipole moment resulting from polarization by another molecule or ion is called an induced dipole moment, and the phenomenon itself is called induction. Thus, the orientational interaction must always be superimposed on the inductive interaction of molecules.

In the case of non-polar molecules (for example, H2, N2 or noble gas atoms), there is no orientational and inductive interaction. However, hydrogen, nitrogen and noble gases are known to be burned. To explain these facts, London introduced the concept of dispersion forces of intermolecular interaction. These forces interact between any atoms and molecules, regardless of their structure. They are caused by instantaneous dipole moments occurring in concert across a large group of atoms:

At any given moment in time, the direction of the dipoles may be different. However, their coordinated occurrence provides weak interaction forces leading to the formation of liquid and solids. In particular, it causes the transition of noble gases to the liquid state at low temperatures.

Thus, the smallest component among the forces acting between molecules is the dispersion interaction. Between molecules with little or no polarity (CH4, H2, HI) active forces are mainly dispersive. The greater the intrinsic dipole moment of the molecules, the greater the orientational forces of interaction between them.

In a series of substances of the same type, dispersion interaction increases with increasing sizes of the atoms that make up the molecules of these substances. For example, in HCl, dispersion forces account for 81% of the total intermolecular interaction; for HBr this value is 95%, and for HI – 99.5%.

Description of chemical bonds in the molecular orbital (MO) method

The BC method is widely used by chemists. In this method, a large and complex molecule is viewed as consisting of individual two-center and two-electron bonds. It is accepted that the electrons responsible for the chemical bond are localized (located) between two atoms. The BC method can be successfully applied to most molecules. However, there are a number of molecules to which this method is not applicable or its conclusions are in conflict with experiment.

It has been established that in a number of cases the decisive role in the formation of a chemical bond is played not by electron pairs, but by individual electrons. The possibility of a chemical bond using one electron is indicated by the existence of an ion. When this ion is formed from a hydrogen atom and a hydrogen ion, 255 kJ (61 kcal) of energy is released. Thus, the chemical bond in the ion is quite strong.

If we try to describe the chemical bond in an oxygen molecule using the BC method, we will come to the conclusion that, firstly, it must be double (σ- and p-bonds), and secondly, in the oxygen molecule all electrons must be paired, i.e. .e. the O2 molecule must be diamagnetic. [In diamagnetic substances, the atoms do not have a permanent magnetic moment and the substance is pushed out of the magnetic field. A paramagnetic substance is one whose atoms or molecules have a magnetic moment, and it has the property of being drawn into a magnetic field]. Experimental data show that the energy of the bond in the oxygen molecule is indeed double, but the molecule is not diamagnetic, but paramagnetic. It has two unpaired electrons. The BC method is powerless to explain this fact.

The best method for quantum mechanical interpretation of chemical bonding is currently considered to be the molecular orbital (MO) method. However, it is much more complicated than the BC method and is not as visual as the latter.

The MO method considers all the electrons of a molecule to be in molecular orbitals. In a molecule, an electron is located at a certain MO, described by the corresponding wave function ψ.

Types of MO. When an electron of one atom, upon approaching, falls into the sphere of action of another atom, the nature of the movement, and therefore the wave function of the electron, changes. In the resulting molecule, the wave functions, or orbitals, of the electrons are unknown. There are several ways to determine the type of MO based on known AOs. Most often, MOs are obtained by linear combination of atomic orbitals (LCAO). The Pauli principle, Hund's rule, and the principle of least energy are also valid for the MO method.

Rice. 2.2 Formation of bonding and antibonding molecular orbitals from atomic orbitals.

In its simplest graphical form, MOs, like LCAO, can be obtained by adding or subtracting wave functions. Figure 2.2 shows the formation of binding and antibonding MOs from the initial AO.

AOs can form MOs if the energies of the corresponding AOs are close in value and the AOs have the same symmetry relative to the bond axis.

The wave functions, or orbitals, of hydrogen 1s can give two linear combinations - one when added, the other when subtracted (Fig. 2.2).

When the wave functions add up, in the overlap region the density of the electron cloud, proportional to ψ2, becomes greater, an excess negative charge is created between the atomic nuclei and the atomic nuclei are attracted to it. A MO obtained by adding the wave functions of hydrogen atoms is called a bonding MO.

If the wave functions are subtracted, then in the region between the atomic nuclei the density of the electron cloud becomes zero, the electron cloud is “pushed out” from the region located between the atoms. The resulting MO cannot bond atoms and is called antibonding.

Since the s-orbitals of hydrogen form only a σ bond, the resulting MOs are designated σcв and σр. MOs formed by 1s-atomic orbitals are designated σcв1s and σр1s.

At the bonding MO, the potential (and total) energy of electrons turns out to be less than at the AO, and at the antibonding MO, it is greater. In absolute value, the increase in the energy of electrons in antibonding orbitals is somewhat greater than the decrease in energy in bonding orbitals. An electron located in a bonding orbital ensures the connection between atoms, stabilizing the molecule, and an electron in an antibonding orbital destabilizes the molecule, i.e. the bond between atoms weakens. Erazr. > Esv.

MOs are also formed from 2p orbitals of the same symmetry: bonding and antibonding σ orbitals from 2p orbitals located along the x axis. They are designated σcв2р and σр2р. Bonding and antibonding p orbitals are formed from 2pz orbitals. They are designated πсв2рz, πp2pz, respectively. The πsv2py and πр2у orbitals are formed similarly.

Filling out MO. The filling of MOs with electrons occurs in the order of increasing orbital energy. If MOs have the same energy (πst or πp orbitals), then filling occurs according to Hund’s rule so that the spin moment of the molecule is greatest. Each MO, like an atomic one, can accommodate two electrons. As noted, the magnetic properties of atoms or molecules depend on the presence of unpaired electrons: if a molecule has unpaired electrons, then it is paramagnetic, if not, it is diamagnetic.

Consider the ion.

From the diagram it is clear that the only electron is located along σcв - MO. A stable compound is formed with a binding energy of 255 kJ/mol and a bond length of 0.106 nm. The molecular ion is paramagnetic. If we assume that the bond multiplicity, as in the BC method, is determined by the number of electron pairs, then the bond multiplicity in is equal to ½. The formation process can be written as follows:

This entry means that there is one electron in the σc MO formed from a 1s AO.

The ordinary hydrogen molecule already contains two electrons with opposite spins in the σcв1s orbital: . The bond energy in H2 is greater than in H2 - 435 kJ/mol, and the bond length (0.074 nm) is shorter. The H2 molecule contains a single bond and the molecule is diamagnetic.

Rice. 2.3. Energy diagram of AO and MO in a system of two hydrogen atoms.

The molecular ion (+He+ ® He+2[(sсв1s)2(sр1s)1]) already has one electron in the σdischarge 1s orbital. The bond energy is 238 kJ/mol (reduced compared to H2), and the bond length (0.108 nm) is increased. The bond multiplicity is ½ (the bond multiplicity is equal to half the difference in the number of electrons in the bonding and antibonding orbitals).

A hypothetical He2 molecule would have two electrons in the σcв1s orbital and two electrons in the σр1s orbital. Since one electron in the antibonding orbital destroys the bonding effect of the electron in the bonding orbital, the He2 molecule cannot exist. The BC method leads to the same conclusion.

The order in which MOs are filled with electrons during the formation of molecules by period II elements is shown below. In accordance with the diagrams, the B2 and O2 molecules are paramagnetic, and the Be2 molecule cannot exist.

The formation of molecules from atoms of elements of period II can be written as follows (K - internal electronic layers):

Physical properties of molecules and MMOs

The existence of binding and loosening MOs is confirmed physical properties molecules. The MO method allows us to predict that if, during the formation of a molecule from atoms, the electrons in the molecule fall into bonding orbitals, then the ionization potentials of the molecules should be greater than the ionization potentials of atoms, and if the electrons fall into antibonding orbitals, then vice versa.

Thus, the ionization potentials of hydrogen and nitrogen molecules (bonding orbitals) - 1485 and 1500 kJ/mol, respectively - are greater than the ionization potentials of hydrogen and nitrogen atoms - 1310 and 1390 kJ/mol, and the ionization potentials of oxygen and fluorine molecules (antibonding orbitals) - 1170 and 1523 kJ/mol are less than those of the corresponding atoms - 1310 and 1670 kJ/mol. When molecules are ionized, bond strength decreases if an electron is removed from a bonding orbital (H2 and N2), and increases if an electron is removed from an antibonding orbital (O2 and F2).

Diatomic molecules with different atoms

MOs for molecules with different atoms (NO, CO) are constructed similarly if the initial atoms do not differ very much in ionization potential values. For the CO molecule, for example, we have:

The AO energies of the oxygen atom are lower than the energies of the corresponding carbon orbitals (1080 kJ/mol); they are located closer to the nucleus. The 10 electrons present in the initial atoms on the outer layers fill the bonding scb2s and antibonding sp2s orbitals and the bonding and pscb2ry,z orbitals. The CO molecule turns out to be isoelectronic with the N2 molecule. The binding energy of atoms in a CO molecule (1105 kJ/mol) is even greater than in a nitrogen molecule (940 kJ/mol). The C–O bond length is 0.113 nm.

NO molecule

has one electron in the antibonding orbital. As a result, the binding energy of NO (680 kJ/mol) is lower than that of N2 or CO. Removing an electron from the NO molecule (ionization to form NO+) increases the binding energy of atoms to 1050–1080 kJ/mol.

Let us consider the formation of MO in the hydrogen fluoride molecule HF. Since the ionization potential of fluorine (17.4 eV or 1670 kJ/mol) is greater than that of hydrogen (13.6 eV or 1310 kJ/mol), the 2p orbitals of fluorine have lower energy than the 1s orbital of hydrogen. Due to the large difference in energy, the 1s orbital of the hydrogen atom and the 2s orbital of the fluorine atom do not interact. Thus, the 2s orbital of fluorine becomes without changing the energy of the MO in HF. Such orbitals are called non-bonding orbitals. The 2py and 2рz orbitals of fluorine also cannot interact with the 1s orbital of hydrogen due to the difference in symmetry relative to the bond axis. They also become non-binding MOs. The bonding and antibonding MOs are formed from the 1s orbital of hydrogen and the 2px orbital of fluorine. The hydrogen and fluorine atoms are connected by a two-electron bond with an energy of 560 kJ/mol.

Bibliography

Glinka N.L. General chemistry. – M.: Chemistry, 1978. – P. 111-153.

Shimanovich I.E., Pavlovich M.L., Tikavyy V.F., Malashko P.M. General chemistry in formulas, definitions, diagrams. – Mn.: Universitetskaya, 1996. – P. 51-77.

Vorobyov V.K., Eliseev S.Yu., Vrublevsky A.V. Practical and independent work in chemistry. – Mn.: UE “Donarit”, 2005. – P. 21-30.

Fig.1. Orbital radii of elements (r a) and length of one-electron chemical bond (d)

The simplest one-electron chemical bond is created by a single valence electron. It turns out that one electron is capable of holding two positively charged ions together. In a one-electron bond, the Coulomb repulsive forces of positively charged particles are compensated by the Coulomb forces of attraction of these particles to a negatively charged electron. The valence electron becomes common to the two nuclei of the molecule.

Examples of such chemical compounds are molecular ions: H 2 +, Li 2 +, Na 2 +, K 2 +, Rb 2 +, Cs 2 +:

Polar covalent bonds occur in heteronuclear diatomic molecules (Fig. 3). The bonding electron pair in a polar chemical bond is brought closer to the atom with a higher first ionization potential.

The distance d between atomic nuclei, which characterizes the spatial structure of polar molecules, can be approximately considered as the sum of the covalent radii of the corresponding atoms.

Characteristics of some polar substances

The shift of a bonding electron pair to one of the nuclei of a polar molecule leads to the appearance of an electric dipole (electrodynamics) (Fig. 4).

The distance between the centers of gravity of positive and negative charges is called the dipole length. The polarity of a molecule, as well as the polarity of a bond, is assessed by the value of the dipole moment μ, which is the product of the dipole length l and the value of the electronic charge:

Multiple covalent bonds

Multiple covalent bonds are represented by unsaturated organic compounds containing double and triple chemical bonds. To describe the nature of unsaturated compounds, L. Pauling introduces the concepts of sigma and π bonds, hybridization of atomic orbitals.

Pauling hybridization for two S and two p electrons made it possible to explain the directionality of chemical bonds, in particular the tetrahedral configuration of methane. To explain the structure of ethylene, from four equivalent Sp 3 electrons of the carbon atom, one p-electron has to be isolated to form an additional bond, called a π bond. In this case, the three remaining Sp 2 hybrid orbitals are located in the plane at an angle of 120° and form basic bonds, for example, a planar ethylene molecule (Fig. 5).

IN new theory Pauling, all bonding electrons became equal and equidistant from the line connecting the nuclei of the molecule. Pauling's theory of the bent chemical bond took into account the statistical interpretation of the M. Born wave function and the Coulomb electron correlation of electrons. A physical meaning has emerged - the nature of a chemical bond is completely determined by the electrical interaction of nuclei and electrons. The more bonding electrons, the smaller the internuclear distance and the stronger the chemical bond between carbon atoms.

Three-center chemical bond

Further development of ideas about chemical bonds was given by the American physical chemist W. Lipscomb, who developed the theory of two-electron three-center bonds and a topological theory that makes it possible to predict the structure of some more boron hydrides (hydrogen hydrides).

An electron pair in a three-center chemical bond becomes common to three atomic nuclei. In the simplest representative of a three-center chemical bond - the molecular hydrogen ion H 3 +, an electron pair holds three protons together (Fig. 6).

Fig. 7. Diboran

The existence of boranes with their two-electron three-center bonds with “bridging” hydrogen atoms violated the canonical doctrine of valence. The hydrogen atom, previously considered a standard monovalent element, turned out to be connected by identical bonds to two boron atoms and formally became a divalent element. W. Lipscomb's work on deciphering the structure of boranes expanded the understanding of chemical bonds. The Nobel Committee awarded William Nunn Lipscomb the Chemistry Prize for 1976 with the wording "For his studies of the structure of boranes (borohydrites), clarifying the problems of chemical bonds."

Multisite chemical bond

Fig. 8. Ferrocene molecule

Fig. 9. Dibenzene chromium

Fig. 10. Uranocene

All ten bonds (C-Fe) in the ferrocene molecule are equivalent, the value of the internuclear Fe-c distance is 2.04 Å. All carbon atoms in a ferrocene molecule are structurally and chemically equivalent, the length of each C-C connections 1.40 - 1.41 Å (for comparison, in benzene the C-C bond length is 1.39 Å). A 36-electron shell appears around the iron atom.

Dynamics of chemical bonding

The chemical bond is quite dynamic. Thus, a metal bond is transformed into a covalent bond in the process phase transition when metal evaporates. The transition of a metal from a solid to a vapor state requires costs large quantities energy.

In pairs, these metals consist practically of homonuclear diatomic molecules and free atoms. When metal vapor condenses, a covalent bond is converted into a metal bond.

Evaporation of salts with typical ionic bonds, such as fluorides alkali metals, leads to the destruction of ionic bonds and the formation of heteronuclear diatomic molecules with a polar covalent bond. In this case, the formation of dimeric molecules with bridged bonds occurs.

Characteristics of chemical bonds in molecules of alkali metal fluorides and their dimers.

During the condensation of vapors of alkali metal fluorides, the polar covalent bond is transformed into an ionic bond with the formation of the corresponding salt crystal lattice.

Mechanism of transition of covalent to metallic bond

Fig. 11. The relationship between the orbital radius of an electron pair r e and the length of a covalent chemical bond d

Fig. 12. Orientation of dipoles of diatomic molecules and the formation of a distorted octahedral fragment of a cluster during condensation of alkali metal vapors

Fig. 13. Body-centered cubic arrangement of nuclei in crystals of alkali metals and a connecting link

Dispersive attraction (London forces) determines interatomic interaction and the formation of homonuclear diatomic molecules from alkali metal atoms.

The formation of a metal-metal covalent bond is associated with deformation of the electronic shells of interacting atoms - valence electrons create a bonding electron pair, the electron density of which is concentrated in the space between the atomic nuclei of the resulting molecule. A characteristic feature of homonuclear diatomic molecules of alkali metals is the long length of the covalent bond (3.6-5.8 times longer than the bond length in the hydrogen molecule) and the low energy of its rupture.

The indicated relationship between r e and d determines the uneven distribution of electric charges in the molecule - the negative electric charge of the bonding electron pair is concentrated in the middle part of the molecule, and the positive ones are concentrated at the ends of the molecule electric charges two atomic skeletons.

The uneven distribution of electric charges creates conditions for the interaction of molecules due to orientation forces (van der Waals forces). Molecules of alkali metals tend to orient themselves in such a way that opposite electric charges appear in their proximity. As a result, attractive forces act between molecules. Thanks to the presence of the latter, the molecules of alkali metals come closer and are more or less firmly pulled together. At the same time, some deformation of each of them occurs under the influence of closer poles of neighboring molecules (Fig. 12).

In fact, the binding electrons of the original diatomic molecule, falling into the electric field of the four positively charged atomic cores of alkali metal molecules, are torn away from the orbital radius of the atom and become free.

In this case, the bonding electron pair becomes common for a system with six cations. The construction of the metal crystal lattice begins at the cluster stage. IN crystal lattice alkali metals, the structure of the connecting link is clearly expressed, having the shape of a distorted flattened octahedron - a square bipyramid, the height of which and the edges of the basis are equal to the value of the translation lattice constant a w (Fig. 13).

The value of the translation lattice constant a w of an alkali metal crystal significantly exceeds the length of the covalent bond of an alkali metal molecule, therefore it is generally accepted that the electrons in the metal are in a free state:

The mathematical construction associated with the properties of free electrons in a metal is usually identified with the “Fermi surface”, which should be considered as the geometric location where electrons reside, providing the main property of a metal - to conduct electric current.

When comparing the process of condensation of alkali metal vapors with the process of condensation of gases, for example, hydrogen, characteristic feature in the properties of the metal. Thus, if during the condensation of hydrogen weak intermolecular interactions appear, then during the condensation of metal vapor processes occur that are characteristic of chemical reactions. The condensation of metal vapor itself occurs in several stages and can be described by the following process: free atom → diatomic molecule with a covalent bond → metal cluster → compact metal with a metal bond.

The interaction of alkali metal halide molecules is accompanied by their dimerization. A dimer molecule can be considered an electric quadrupole (Fig. 15). Currently, the main characteristics of dimers of alkali metal halides are known (chemical bond lengths and bond angles between bonds).

Chemical bond length and bond angles in dimers of alkali metal halides (E 2 X 2) (gas phase).

E 2 X 2	X=F		X=Cl		X=Br		X=I
E 2 X 2	dEF, Å		d ECl, Å		d EBr , Å		d EI, Å
Li 2 X 2	1,75	105	2,23	108	2,35	110	2,54	116
Na 2 X 2	2,08	95	2,54	105	2,69	108	2,91	111
K 2 X 2	2,35	88	2,86	98	3,02	101	3,26	104
Cs 2 X 2	2,56	79	3,11	91	3,29	94	3,54	94

During the condensation process, the effect of orientation forces increases, intermolecular interaction is accompanied by the formation of clusters, and then a solid substance. Alkali metal halides form crystals with simple cubic and body-centered cubic lattices.

Crystal lattice type and translation lattice constant for alkali metal halides.

During the crystallization process, a further increase in the interatomic distance occurs, leading to the removal of an electron from the orbital radius of the alkali metal atom and the transfer of an electron to the halogen atom with the formation of the corresponding ions. The force fields of ions are evenly distributed in all directions in space. In this regard, in alkali metal crystals, the force field of each ion is coordinated by more than one ion with the opposite sign, as is customary to qualitatively represent the ionic bond (Na + Cl -).

In crystals of ionic compounds, the concept of simple two-ionic molecules such as Na + Cl - and Cs + Cl - loses its meaning, since the alkali metal ion is associated with six chlorine ions (in a sodium chloride crystal) and with eight chlorine ions (in a cesium chloride crystal. However, all interionic distances in crystals are equidistant.

Notes

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Lidin R.A., Andreeva L.L., Molochko V.A. Handbook of Inorganic Chemistry. Constants of inorganic substances. - M.: “Chemistry”, 1987. - P. 132-136. - 320 s.
Gankin V.Yu., Gankin Yu.V. How a chemical bond is formed and chemical reactions occur. - M.: publishing group "Granitsa", 2007. - 320 p. - ISBN 978-5-94691296-9
Nekrasov B.V. General chemistry course. - M.: Goskhimizdat, 1962. - P. 88. - 976 p.
Pauling L. The nature of chemical bonding / edited by Y.K. Syrkin. - per. from English M.E. Dyatkina. - M.-L.: Goskhimizdat, 1947. - 440 p.
Theoretical organic chemistry / ed. R.H. Freidlina. - per. from English Yu.G.Bundela. - M.: Publishing house. foreign literature, 1963. - 365 p.
Lemenovsky D.A., Levitsky M.M. Russian Chemical Journal (journal of the Russian Chemical Society named after D.I. Mendeleev). - 2000. - T. XLIV, issue 6. - P. 63-86.
Chemical encyclopedic dictionary / ch. ed. I.L. Knunyants. - M.: Sov. encyclopedia, 1983. - P. 607. - 792 p.
Nekrasov B.V. General chemistry course. - M.: Goskhimizdat, 1962. - P. 679. - 976 p.
Lidin R.A., Andreeva L.L., Molochko V.A. Handbook of Inorganic Chemistry. Constants of inorganic substances. - M.: “Chemistry”, 1987. - P. 155-161. - 320 s.
Gillespie R. Geometry of molecules / trans. from English E.Z. Zasorina and V.S. Mastryukova, ed. Yu.A Pentina. - M.: “Mir”, 1975. - P. 49. - 278 p.
Chemist's Handbook. - 2nd ed., revised. and additional - L.-M.: State Scientific and Technical Institute of Chemical Literature, 1962. - T. 1. - P. 402-513. - 1072 pp.
Lidin R.A., Andreeva L.L., Molochko V.A. Handbook of Inorganic Chemistry. Constants of inorganic substances.. - M.: “Chemistry”, 1987. - P. 132-136. - 320 s.
Ziman J. Electrons in metals (introduction to the theory of Fermi surfaces). Advances in physical sciences.. - 1962. - T. 78, issue 2. - 291 p.

Basic Concepts

Chemical bond called a set of interactions that lead to the bonding of atoms with the formation of stable particles of a more complex structure (molecules, ions, radicals), as well as aggregates (crystals, glasses, etc.). The nature of these interactions is electrical in nature, and they arise during the distribution of valence electrons in approaching atoms.

Valence accepted name the ability of an atom to form a certain number of bonds with other atoms. In ionic compounds, the number of electrons given up or gained is taken as the valence value. In covalent compounds it is equal to the number of shared electron pairs.

Under the degree of oxidation is understood as a conditional the charge that could be on an atom if all polar covalent bonds were ionic in nature.

The multiplicity of a connection is called the number of shared electron pairs between the atoms under consideration.

The bonds considered in various branches of chemistry can be divided into two types of chemical bonds: those that lead to the formation of new substances (intramolecular) , And those that occur between molecules (intermolecular).

Basic communication characteristics

Energy of communication is the energy required to break all existing bonds in a molecule. It is also the energy released during bond formation.

Link length is the distance between neighboring nuclei of atoms in a molecule at which the forces of attraction and repulsion are balanced.

These two characteristics of a chemical bond between atoms are a measure of its strength: the shorter the length and the greater the energy, the stronger the bond.

Bond angle it is customary to call the angle between the represented lines passing in the direction of communication through the nuclei of atoms.

Methods for describing connections

The most common two approaches to explaining chemical bonding, borrowed from quantum mechanics:

Molecular orbital method. He views the molecule as a collection of electrons and atomic nuclei, with each individual electron moving in the field of action of all other electrons and nuclei. The molecule has an orbital structure, and all its electrons are distributed in these orbits. This method is also called MO LCAO, which stands for “molecular orbital - linear combination

Valence bond method. Represents a molecule as a system of two central molecular orbitals. Moreover, each of them corresponds to one bond between two neighboring atoms in the molecule. The method is based on the following provisions:

The formation of a chemical bond is carried out by a pair of electrons having opposite spins, which are located between the two atoms in question. The electron pair formed belongs equally to the two atoms.
The number of bonds formed by one or another atom is equal to the number of unpaired electrons in the ground and excited states.
If electron pairs do not participate in the formation of a bond, then they are called lone pairs.

Electronegativity

The type of chemical bond in substances can be determined based on the difference in the electronegativity values of its constituent atoms. Under electronegativity understand the ability of atoms to attract shared electron pairs (electron cloud), which leads to bond polarization.

There are various ways to determine the electronegativity values of chemical elements. However, the most used is the scale based on thermodynamic data, which was proposed back in 1932 by L. Pauling.

The greater the difference in electronegativity of atoms, the more pronounced its ionicity. On the contrary, equal or similar electronegativity values indicate the covalent nature of the bond. In other words, it is possible to determine mathematically what chemical bond is observed in a particular molecule. To do this, you need to calculate ΔХ - the difference in electronegativity of atoms using the formula: ΔХ=|Х 1 -X 2 |.

If ΔХ>1.7, then the bond is ionic.
If 0.5≤ΔХ≤1.7, then the covalent bond is polar.
If ΔХ=0 or close to it, then the bond is classified as covalent nonpolar.

Ionic bond

An ionic bond is a bond that appears between ions or due to the complete withdrawal of a common electron pair by one of the atoms. In substances, this type of chemical bond is carried out by forces of electrostatic attraction.

Ions are charged particles formed from atoms by gaining or losing electrons. If an atom accepts electrons, it acquires a negative charge and becomes an anion. If an atom gives up valence electrons, it becomes a positively charged particle called a cation.

It is characteristic of compounds formed by the interaction of atoms of typical metals with atoms of typical non-metals. The main reason for this process is the desire of atoms to acquire stable electronic configurations. And for this, typical metals and non-metals need to give or accept only 1-2 electrons, which they do with ease.

The mechanism of formation of an ionic chemical bond in a molecule is traditionally considered using the example of the interaction of sodium and chlorine. Alkali metal atoms easily give up an electron, drawn by a halogen atom. As a result, the Na + cation and the Cl - anion are formed, which are held together by electrostatic attraction.

There is no ideal ionic bond. Even in such compounds, which are often classified as ionic, the final transfer of electrons from atom to atom does not occur. The formed electron pair still remains in common use. Therefore, they talk about the degree of ionicity of a covalent bond.

An ionic bond is characterized by two main properties related to each other:

non-directionality, i.e. the electric field around the ion has the shape of a sphere;
unsaturation, i.e., the number of oppositely charged ions that can be placed around any ion, is determined by their sizes.

Covalent chemical bond

A bond formed by overlapping electron clouds of nonmetal atoms, that is, carried out by a common electron pair, is called a covalent bond. The number of shared electron pairs determines the multiplicity of the bond. Thus, hydrogen atoms are connected by a single H··H bond, and oxygen atoms form an O::O double bond.

There are two mechanisms for its formation:

Exchange - each atom represents one electron to form a common pair: A· + ·B = A:B, while external atomic orbitals, on which one electron is located, participate in the bonding.
Donor-acceptor - to form a bond, one of the atoms (donor) provides a pair of electrons, and the second (acceptor) provides a free orbital for its placement: A + : B = A: B.

The ways in which electron clouds overlap during the formation of a covalent chemical bond are also different.

Direct. The region of cloud overlap lies on a straight imaginary line connecting the nuclei of the atoms in question. In this case, σ bonds are formed. The type of chemical bond that occurs in this case depends on the type of electron clouds that overlap: s-s, s-p, p-p, s-d or p-d σ bonds. In a particle (molecule or ion), only one σ bond is possible between two neighboring atoms.
Lateral. It is carried out on both sides of the line connecting the nuclei of atoms. This is how a π bond is formed, and its varieties are also possible: p-p, p-d, d-d. A π bond is never formed separately from a σ bond; it can occur in molecules containing multiple (double and triple) bonds.

Properties of covalent bonds

They determine the chemical and physical properties of compounds. The main properties of any chemical bond in substances are its directionality, polarity and polarizability, as well as saturation.

Focus connections are determined by the features of the molecular structure of substances and the geometric shape of their molecules. Its essence is that the best overlap of electron clouds is possible at a certain orientation in space. The options for the formation of σ- and π-bonds have already been discussed above.

Under saturation understand the ability of atoms to form a certain number of chemical bonds in a molecule. The number of covalent bonds for each atom is limited by the number of outer orbitals.

Polarity bond depends on the difference in the electronegativity values of the atoms. The uniformity of the distribution of electrons between the nuclei of atoms depends on it. According to this characteristic, a covalent bond can be polar or nonpolar.

If the common electron pair belongs equally to each of the atoms and is located at the same distance from their nuclei, then the covalent bond is non-polar.
If a common pair of electrons is displaced towards the nucleus of one of the atoms, then a covalent polar chemical bond is formed.

Polarizability is expressed by the displacement of bond electrons under the influence of an external electric field, which may belong to another particle, neighboring bonds in the same molecule, or come from external sources of electromagnetic fields. Thus, a covalent bond under their influence can change its polarity.

Hybridization of orbitals is understood as a change in their shapes during a chemical bond. This is necessary to achieve the most effective overlap. The following types of hybridization exist:

sp3. One s and three p orbitals form four “hybrid” orbitals of the same shape. Outwardly it resembles a tetrahedron with an angle between the axes of 109°.
sp2. One s- and two p-orbitals form a flat triangle with an angle between the axes of 120°.
sp. One s- and one p-orbital form two “hybrid” orbitals with an angle between their axes of 180°.

A special feature of the structure of metal atoms is their rather large radius and the presence of a small number of electrons in outer orbitals. As a result, in such chemical elements the bond between the nucleus and valence electrons is relatively weak and is easily broken.

Metal A bond is an interaction between metal atoms and ions that occurs with the help of delocalized electrons.

In metal particles, valence electrons can easily leave the outer orbitals, as well as occupy vacant positions on them. Thus, at different moments of time the same particle can be an atom and an ion. The electrons detached from them move freely throughout the entire volume of the crystal lattice and carry out a chemical bond.

This type of bond has similarities with ionic and covalent bonds. Just like ionic bonds, metallic bonds require ions to exist. But if cations and anions are needed to carry out electrostatic interaction in the first case, then in the second the role of negatively charged particles is played by electrons. When comparing a metallic bond with a covalent bond, both require shared electrons to form. However, unlike polar chemical bonds, they are not localized between two atoms, but belong to all metal particles in the crystal lattice.

Metallic bonding is responsible for the special properties of almost all metals:

plasticity is present due to the possibility of displacement of layers of atoms in a crystal lattice held by an electron gas;
metallic luster, which is observed due to the reflection of light rays from electrons (in the powder state there is no crystal lattice and, therefore, electrons moving through it);
electrical conductivity, which is carried out by a flow of charged particles, and in this case small electrons move freely among large metal ions;
thermal conductivity is observed due to the ability of electrons to transfer heat.

This type of chemical bond is sometimes called intermediate between covalent and intermolecular interactions. If a hydrogen atom has a bond with one of the highly electronegative elements (such as phosphorus, oxygen, chlorine, nitrogen), then it is capable of forming an additional bond, called a hydrogen bond.

It is much weaker than all the types of bonds discussed above (energy no more than 40 kJ/mol), but it cannot be neglected. This is why a hydrogen chemical bond appears as a dotted line in the diagram.

The occurrence of a hydrogen bond is possible due to the simultaneous donor-acceptor electrostatic interaction. A large difference in electronegativity values leads to the appearance of excess electron density on the O, N, F and other atoms, as well as to its deficiency on the hydrogen atom. In the event that there is no existing chemical bond between such atoms, when they are close enough, attractive forces are activated. In this case, the proton is the acceptor of the electron pair, and the second atom is the donor.

Hydrogen bonds can occur both between neighboring molecules, for example, water, carboxylic acids, alcohols, ammonia, and within a molecule, for example, salicylic acid.

The presence of hydrogen bonds between water molecules explains a number of its unique physical properties:

The values of its heat capacity, dielectric constant, boiling and melting points, in accordance with calculations, should be significantly less than real ones, which is explained by the connectivity of molecules and the need to expend energy on breaking intermolecular hydrogen bonds.
Unlike other substances, the volume of water increases as the temperature decreases. This occurs due to the fact that the molecules occupy a certain position in the crystal structure of ice and move away from each other by the length of the hydrogen bond.

This connection plays a special role for living organisms, since its presence in protein molecules determines their special structure, and therefore their properties. In addition, nucleic acids, making up the double helix of DNA, are also connected by hydrogen bonds.

Bonds in crystals

The vast majority of solids have a crystal lattice - a special relative arrangement of the particles that form them. In this case, three-dimensional periodicity is observed, and atoms, molecules or ions are located at the nodes, which are connected by imaginary lines. Depending on the nature of these particles and the connections between them, all crystalline structures are divided into atomic, molecular, ionic and metallic.

The nodes of the ionic crystal lattice contain cations and anions. Moreover, each of them is surrounded by a strictly defined number of ions with only the opposite charge. A typical example is sodium chloride (NaCl). They tend to have high melting points and hardness because they require a lot of energy to break down.

At the nodes of the molecular crystal lattice there are molecules of substances formed by covalent bonds (for example, I 2). They are connected to each other by a weak van der Waals interaction, and therefore such a structure is easy to destroy. Such compounds have low boiling and melting points.

The atomic crystal lattice is formed by atoms of chemical elements with high valency values. They are connected by strong covalent bonds, which means that the substances have high boiling and melting points and great hardness. An example is a diamond.

Thus, all types of bonds present in chemical substances have their own characteristics, which explain the subtleties of the interaction of particles in molecules and substances. The properties of the compounds depend on them. They determine all processes occurring in the environment.

Covalent chemical bond, its varieties and mechanisms of formation. Characteristics of covalent bonds (polarity and bond energy). Ionic bond. Metal connection. Hydrogen bond

The doctrine of chemical bonding forms the basis of all theoretical chemistry.

A chemical bond is understood as the interaction of atoms that binds them into molecules, ions, radicals, and crystals.

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

The division of chemical bonds into types is conditional, since they are all characterized by a certain unity.

An ionic bond can be considered as an extreme case of a polar covalent bond.

A metallic bond combines the covalent interaction of atoms using shared electrons and the electrostatic attraction between these electrons and metal ions.

Substances often lack limiting cases of chemical bonding (or pure chemical bonding).

For example, lithium fluoride $LiF$ is classified as an ionic compound. In fact, the bond in it is $80%$ ionic and $20%$ covalent. It is therefore more correct, obviously, to talk about the degree of polarity (ionicity) of a chemical bond.

In the series of hydrogen halides $HF—HCl—HBr—HI—HAt$ the degree of bond polarity decreases, because the difference in the electronegativity values of the halogen and hydrogen atoms decreases, and in astatine hydrogen the bond becomes almost nonpolar $(EO(H) = 2.1; EO(At) = 2.2)$.

Different types of bonds can be found in the same substances, for example:

in bases: between the oxygen and hydrogen atoms in hydroxo groups the bond is polar covalent, and between the metal and the hydroxo group it is ionic;
in salts of oxygen-containing acids: between the non-metal atom and the oxygen of the acidic residue - covalent polar, and between the metal and the acidic residue - ionic;
in ammonium, methylammonium salts, etc.: between nitrogen and hydrogen atoms - covalent polar, and between ammonium or methylammonium ions and the acid residue - ionic;
in metal peroxides (for example, $Na_2O_2$), the bond between oxygen atoms is covalent nonpolar, and between the metal and oxygen is ionic, etc.

Different types of connections can transform into one another:

— during electrolytic dissociation of covalent compounds in water, the covalent polar bond turns into an ionic bond;

- when metals evaporate, the metal bond turns into a nonpolar covalent bond, etc.

The reason for the unity of all types and types of chemical bonds is their identical chemical nature - electron-nuclear interaction. The formation of a chemical bond in any case is the result of electron-nuclear interaction of atoms, accompanied by the release of energy.

Methods for forming covalent bonds. Characteristics of a covalent bond: bond length and energy

A covalent chemical bond is a bond formed between atoms through the formation of shared electron pairs.

The mechanism of formation of such a bond can be exchange or donor-acceptor.

I. Exchange mechanism operates when atoms form shared electron pairs by combining unpaired electrons.

1) $H_2$ - hydrogen:

The bond arises due to the formation of a common electron pair by $s$-electrons of hydrogen atoms (overlapping $s$-orbitals):

2) $HCl$ - hydrogen chloride:

The bond arises due to the formation of a common electron pair of $s-$ and $p-$electrons (overlapping $s-p-$orbitals):

3) $Cl_2$: in a chlorine molecule, a covalent bond is formed due to unpaired $p-$electrons (overlapping $p-p-$orbitals):

4) $N_2$: in a nitrogen molecule three common electron pairs are formed between the atoms:

II. Donor-acceptor mechanism Let us consider the formation of a covalent bond using the example of the ammonium ion $NH_4^+$.

The donor has an electron pair, the acceptor has an empty orbital that this pair can occupy. In the ammonium ion, all four bonds with hydrogen atoms are covalent: three were formed due to the creation of common electron pairs by the nitrogen atom and hydrogen atoms according to the exchange mechanism, one - through the donor-acceptor mechanism.

Covalent bonds can be classified by the way the electron orbitals overlap, as well as by their displacement towards one of the bonded atoms.

Chemical bonds formed as a result of overlapping electron orbitals along a bond line are called $σ$ -bonds (sigma bonds). The sigma bond is very strong.

$p-$orbitals can overlap in two regions, forming a covalent bond due to lateral overlap:

Chemical bonds formed as a result of “lateral” overlap of electron orbitals outside the communication line, i.e. in two areas are called $π$ -bonds (pi-bonds).

By degree of displacement shared electron pairs to one of the atoms they bond, a covalent bond can be polar And non-polar.

A covalent chemical bond formed between atoms with the same electronegativity is called non-polar. Electron pairs are not shifted to any of the atoms, because atoms have the same EO - the property of attracting valence electrons from other atoms. For example:

those. molecules of simple non-metal substances are formed through covalent non-polar bonds. A covalent chemical bond between atoms of elements whose electronegativity differs is called polar.

Length and energy of covalent bonds.

Characteristic properties of covalent bond- its length and energy. Link length is the distance between the nuclei of atoms. The shorter the length of a chemical bond, the stronger it is. However, a measure of the strength of the connection is binding energy, which is determined by the amount of energy required to break a bond. It is usually measured in kJ/mol. Thus, according to experimental data, the bond lengths of $H_2, Cl_2$ and $N_2$ molecules are respectively $0.074, 0.198$ and $0.109$ nm, and the bond energies are respectively $436, 242$ and $946$ kJ/mol.

Ions. Ionic bond

Let's imagine that two atoms “meet”: an atom of a group I metal and a non-metal atom of group VII. A metal atom has a single electron at its outer energy level, while a non-metal atom just lacks one electron for its outer level to be complete.

The first atom will easily give the second its electron, which is far from the nucleus and weakly bound to it, and the second will provide it with a free place on its outer electronic level.

Then the atom, deprived of one of its negative charges, will become a positively charged particle, and the second will turn into a negatively charged particle due to the resulting electron. Such particles are called ions.

The chemical bond that occurs between ions is called ionic.

Let's consider the formation of this bond using the example of the well-known compound sodium chloride (table salt):

The process of converting atoms into ions is depicted in the diagram:

This transformation of atoms into ions always occurs during the interaction of atoms of typical metals and typical non-metals.

Let's consider the algorithm (sequence) of reasoning when recording the formation of an ionic bond, for example, between calcium and chlorine atoms:

Numbers showing the number of atoms or molecules are called coefficients, and numbers showing the number of atoms or ions in a molecule are called indexes.

Metal connection

Let's get acquainted with how atoms of metal elements interact with each other. Metals usually do not exist as isolated atoms, but in the form of a piece, ingot, or metal product. What holds metal atoms in a single volume?

The atoms of most metals contain a small number of electrons at the outer level - $1, 2, 3$. These electrons are easily stripped off and the atoms become positive ions. The detached electrons move from one ion to another, binding them into a single whole. Connecting with ions, these electrons temporarily form atoms, then break off again and combine with another ion, etc. Consequently, in the volume of the metal, atoms are continuously converted into ions and vice versa.

The bond in metals between ions through shared electrons is called metallic.

The figure schematically shows the structure of a sodium metal fragment.

In this case, a small number of shared electrons bind a large number of ions and atoms.

A metallic bond has some similarities with a covalent bond, since it is based on the sharing of external electrons. However, with a covalent bond, the outer unpaired electrons of only two neighboring atoms are shared, while with a metallic bond, all atoms take part in the sharing of these electrons. That is why crystals with a covalent bond are brittle, but with a metal bond, as a rule, they are ductile, electrically conductive and have a metallic luster.

Metallic bonding is characteristic of both pure metals and mixtures of various metals—alloys in solid and liquid states.

Hydrogen bond

A chemical bond between positively polarized hydrogen atoms of one molecule (or part thereof) and negatively polarized atoms of strongly electronegative elements having lone electron pairs ($F, O, N$ and less commonly $S$ and $Cl$) of another molecule (or its part) is called hydrogen.

The mechanism of hydrogen bond formation is partly electrostatic, partly donor-acceptor in nature.

Examples of intermolecular hydrogen bonding:

In the presence of such a connection, even low-molecular substances can, under normal conditions, be liquids (alcohol, water) or easily liquefied gases (ammonia, hydrogen fluoride).

Substances with hydrogen bonds have molecular crystal lattices.

Substances of molecular and non-molecular structure. Type of crystal lattice. Dependence of the properties of substances on their composition and structure

Molecular and non-molecular structure of substances

It is not individual atoms or molecules that enter into chemical interactions, but substances. Under given conditions, a substance can be in one of three states of aggregation: solid, liquid or gaseous. The properties of a substance also depend on the nature of the chemical bond between the particles that form it - molecules, atoms or ions. Based on the type of bond, substances of molecular and non-molecular structure are distinguished.

Substances made up of molecules are called molecular substances. The bonds between the molecules in such substances are very weak, much weaker than between the atoms inside the molecule, and even at relatively low temperatures they break - the substance turns into a liquid and then into a gas (sublimation of iodine). The melting and boiling points of substances consisting of molecules increase with increasing molecular weight.

Molecular substances include substances with an atomic structure ($C, Si, Li, Na, K, Cu, Fe, W$), among them there are metals and non-metals.

Let's consider the physical properties of alkali metals. The relatively low bond strength between atoms causes low mechanical strength: alkali metals are soft and can be easily cut with a knife.

Large atomic sizes lead to low densities of alkali metals: lithium, sodium and potassium are even lighter than water. In the group of alkali metals, the boiling and melting points decrease with increasing atomic number of the element, because Atom sizes increase and bonds weaken.

To substances non-molecular structures include ionic compounds. Most compounds of metals with nonmetals have this structure: all salts ($NaCl, K_2SO_4$), some hydrides ($LiH$) and oxides ($CaO, MgO, FeO$), bases ($NaOH, KOH$). Ionic (non-molecular) substances have high melting and boiling points.

Crystal lattices

Matter, as is known, can exist in three states of aggregation: gaseous, liquid and solid.

Solids: amorphous and crystalline.

Let us consider how the characteristics of chemical bonds influence the properties of solids. Solids are divided into crystalline And amorphous.

Amorphous substances do not have a clear melting point; when heated, they gradually soften and turn into a fluid state. For example, plasticine and various resins are in an amorphous state.

Crystalline substances are characterized by the correct arrangement of the particles of which they are composed: atoms, molecules and ions - at strictly defined points in space. When these points are connected by straight lines, a spatial framework is formed, called a crystal lattice. The points at which crystal particles are located are called lattice nodes.

Depending on the type of particles located at the nodes of the crystal lattice and the nature of the connection between them, four types of crystal lattices are distinguished: ionic, atomic, molecular And metal.

Ionic crystal lattices.

Ionic are called crystal lattices, in the nodes of which there are ions. They are formed by substances with ionic bonds, which can bind both simple ions $Na^(+), Cl^(-)$, and complex $SO_4^(2−), OH^-$. Consequently, salts and some oxides and hydroxides of metals have ionic crystal lattices. For example, a sodium chloride crystal consists of alternating positive $Na^+$ and negative $Cl^-$ ions, forming a cube-shaped lattice. The bonds between ions in such a crystal are very stable. Therefore, substances with an ionic lattice are characterized by relatively high hardness and strength, they are refractory and non-volatile.

Atomic crystal lattices.

Atomic are called crystal lattices, in the nodes of which there are individual atoms. In such lattices, the atoms are connected to each other by very strong covalent bonds. An example of substances with this type of crystal lattices is diamond, one of the allotropic modifications of carbon.

Most substances with an atomic crystal lattice have very high melting points (for example, for diamond it is above $3500°C), they are strong and hard, and practically insoluble.

Molecular crystal lattices.

Molecular called crystal lattices, in the nodes of which molecules are located. Chemical bonds in these molecules can be both polar ($HCl, H_2O$) and nonpolar ($N_2, O_2$). Despite the fact that the atoms inside the molecules are connected by very strong covalent bonds, weak intermolecular forces of attraction act between the molecules themselves. Therefore, substances with molecular crystal lattices have low hardness, low melting points, and are volatile. Most solid organic compounds have molecular crystal lattices (naphthalene, glucose, sugar).

Metal crystal lattices.

Substances with metallic bonds have metallic crystal lattices. At the sites of such lattices there are atoms and ions (either atoms or ions, into which metal atoms easily transform, giving up their outer electrons “for common use”). This internal structure of metals determines their characteristic physical properties: malleability, ductility, electrical and thermal conductivity, characteristic metallic luster.

All chemical compounds are formed through the formation of a chemical bond. And depending on the type of connecting particles, several types are distinguished. The most basic– these are covalent polar, covalent nonpolar, metallic and ionic. Today we will talk about ionic.

In contact with

What are ions

It is formed between two atoms - as a rule, provided that the difference in electronegativity between them is very large. The electronegativity of atoms and ions is assessed using the Paulling scale.

Therefore, in order to correctly consider the characteristics of compounds, the concept of ionicity was introduced. This characteristic allows you to determine what percentage of a particular bond is ionic.

The compound with the highest ionicity is cesium fluoride, in which it is approximately 97%. Ionic bonding is characteristic for substances formed by metal atoms located in the first and second groups of the D.I. table. Mendeleev, and atoms of non-metals located in the sixth and seventh groups of the same table.

Note! It is worth noting that there is no compound in which the relationship is exclusively ionic. For currently discovered elements, it is not possible to achieve such a large difference in electronegativity to obtain a 100% ionic compound. Therefore, the definition of an ionic bond is not entirely correct, since in reality compounds with partial ionic interaction are considered.

Why was this term introduced if such a phenomenon does not really exist? The fact is that this approach helped explain many of the nuances in the properties of salts, oxides and other substances. For example, why are they highly soluble in water, and why are they solutions are capable of conducting electric current. This cannot be explained from any other perspective.

Education mechanism

The formation of an ionic bond is possible only if two conditions are met: if the metal atom participating in the reaction is able to easily give up electrons located in the last energy level, and the non-metal atom is able to accept these electrons. Metal atoms by their nature are reducing agents, that is, they are capable of electron donation.

This is due to the fact that the last energy level in a metal can contain from one to three electrons, and the radius of the particle itself is quite large. Therefore, the force of interaction between the nucleus and electrons at the last level is so small that they can easily leave it. The situation with nonmetals is completely different. They have small radius, and the number of own electrons at the last level can be from three to seven.

And the interaction between them and the positive nucleus is quite strong, but any atom strives to complete the energy level, so non-metal atoms strive to obtain the missing electrons.

And when two atoms - a metal and a non-metal - meet, electrons transfer from the metal atom to the non-metal atom, and a chemical interaction is formed.

Connection diagram

The figure clearly shows how exactly the formation of an ionic bond occurs. Initially, there are neutrally charged sodium and chlorine atoms.

The first has one electron at the last energy level, the second seven. Next, an electron transfers from sodium to chlorine and the formation of two ions. Which combine with each other to form a substance. What is an ion? An ion is a charged particle in which the number of protons is not equal to the number of electrons.

Differences from covalent type

Due to its specificity, an ionic bond has no directionality. This is due to the fact that the electric field of the ion is a sphere, and it decreases or increases in one direction uniformly, obeying the same law.

Unlike covalent, which is formed due to the overlap of electron clouds.

The second difference is that covalent bond is saturated. What does it mean? The number of electronic clouds that can take part in interaction is limited.

And in the ionic one, due to the fact that the electric field has a spherical shape, it can connect with an unlimited number of ions. This means that we can say that it is not saturated.

It can also be characterized by several other properties:

Bond energy is a quantitative characteristic, and it depends on the amount of energy that must be expended to break it. It depends on two criteria - bond length and ion charge involved in its education. The stronger the bond, the shorter its length and the greater the charges of the ions that form it.
Length - this criterion has already been mentioned in the previous paragraph. It depends solely on the radius of the particles involved in the formation of the compound. The radius of atoms changes as follows: it decreases over the period with increasing atomic number and increases in the group.

Substances with ionic bonds

It is characteristic of a significant number of chemical compounds. This is a large part of all salts, including the well-known table salt. It occurs in all connections where there is a direct contact between metal and non-metal. Here are some examples of substances with ionic bonds:

sodium and potassium chlorides,
cesium fluoride,
magnesium oxide.

It can also manifest itself in complex compounds.

For example, magnesium sulfate.

Here is the formula of a substance with ionic and covalent bonds:

An ionic bond will form between oxygen and magnesium ions, but sulfur is connected to each other using a polar covalent bond.

From which we can conclude that ionic bonds are characteristic of complex chemical compounds.

What is an ionic bond in chemistry

Types of chemical bonds - ionic, covalent, metallic

Conclusion

Properties directly depend on the device crystal lattice. Therefore, all compounds with ionic bonds are highly soluble in water and other polar solvents, conduct and are dielectrics. At the same time, they are quite refractory and fragile. The properties of these substances are often used in the design of electrical devices.