(based on materials from the site http://chemel.ru/2008-05-24-19-19-34/2008-06-01-15-23-43/18-2008-05-29-22-08-32. html)
It is known that nonmetals interact with each other. Let's consider the mechanism of occurrence covalent bond using the example of the formation of a hydrogen molecule:
H+H=H 2 H= - 436 kJ/mol
Let's imagine that we have two separate isolated hydrogen atoms. The nucleus of each free hydrogen atom is surrounded by a spherical symmetric electron cloud formed by a 1s electron (see Fig. 1). When atoms come closer to a certain distance, partial overlap occurs electronic shells(orbitals) (Fig. 2).
As a result, a molecular two-electron cloud appears between the centers of both nuclei, which has a maximum electron density in the space between the nuclei; An increase in the negative charge density favors a strong increase in the attractive forces between the nuclei and the molecular cloud.
So, a covalent bond is formed as a result of the overlap of electron clouds of atoms, accompanied by the release of energy. If the distance between the nuclei of hydrogen atoms approaching before touching is 0.106 nm, then after the electron clouds overlap (formation of the H 2 molecule), this distance is 0.074 nm (Fig. 2).
Typically, the greatest overlap of electron clouds occurs along the line connecting the nuclei of two atoms.
The greater the overlap of electron orbitals, the stronger the chemical bond.
As a result of the formation of a chemical bond between two hydrogen atoms, each of them reaches electronic configuration noble gas atom.
Chemical bonds are usually depicted in different ways:
1) using electrons in the form of points placed at chemical sign element.
Then the formation of a hydrogen molecule can be shown by the diagram:
N + N N:N
2) using quantum cells (Hund cells), like placing two electrons with opposite spins in one molecular quantum cell:
The diagram on the left shows that the molecular energy level is lower than the original atomic levels, which means that the molecular state of matter is more stable than the atomic state.
3) often, especially in organic chemistry, a covalent bond is represented by a dash (prime)
(for example H-H), which symbolizes a pair of electrons.
The covalent bond in the chlorine molecule is also carried out using two shared electrons, or an electron pair: 
As you can see, each chlorine atom has three lone pairs and one unpaired electron.
The formation of a chemical bond occurs due to unpaired electrons every atom. Unpaired electrons bond into a common pair of electrons, also called a shared pair.
If one covalent bond (one common electron pair) has arisen between atoms, then it is called a single bond; if more, then multiple (two common electron pairs), triple (three common electron pairs).
A single bond is represented by one dash (prime), a double bond by two, and a triple bond by three. The dash between two atoms shows that they have a generalized pair of electrons, as a result of which chemical bond. With the help of such dashes, the sequence of connections of atoms in a molecule is depicted.
So, in a chlorine molecule, each of its atoms has a complete external level of eight electrons (s 2 p 6), and two of them (electron pair) belong equally to both atoms.
The bond in the O2 oxygen molecule is depicted somewhat differently. It has been experimentally established that oxygen is a paramagnetic substance (it is drawn into a magnetic field). Its molecule has two unpaired electrons. The structure of this molecule can be depicted as follows:
An unambiguous solution to depicting the electronic structure of the oxygen molecule has not yet been found. However, it cannot be depicted like this:
In the nitrogen molecule N2, the atoms have three common electron pairs:
It is obvious that a nitrogen molecule is stronger than an oxygen or chlorine molecule, which explains the significant inertness of nitrogen in chemical reactions.
A chemical bond carried out by electron pairs is called covalent.
This is a two-electron and two-center (holds two nuclei) bond.
Compounds with a covalent bond are called homeopolar, or atomic.
There are two types of covalent bonds: nonpolar and polar.
In the case of a nonpolar covalent bond, the electron cloud formed by a common pair of electrons, or the electron cloud of the bond, is distributed in space symmetrically relative to the nuclei of both atoms.
An example is diatomic molecules consisting of atoms of one element: H 2 Cl 2, O 2, N 2, F 2, etc., in which the electron pair belongs equally to both atoms.
In the case of a polar covalent bond, the electron cloud of the bond is shifted towards the atom with higher relative electronegativity.
An example would be volatile molecules inorganic compounds: HC1, H 2 O, H 2 S, NH 3, etc.
The formation of the HC1 molecule can be represented by the following diagram:
The electron pair is shifted towards the chlorine atom, since the relative electronegativity of the chlorine atom (2.83) is greater than that of the hydrogen atom (2.1).
A covalent bond is formed not only due to the overlap of one-electron clouds, it is an exchange mechanism for the formation of a covalent bond.
Another mechanism for the formation of a covalent bond is also possible - donor-acceptor. In this case, a chemical bond occurs due to the two-electron cloud of one atom and the free orbital of another atom. Let us consider as an example the mechanism of formation of the ammonium ion NH +4. In an ammonia molecule, the nitrogen atom has a lone pair of electrons (two electrons).
new cloud):
The hydrogen ion has a free (unfilled) 1s orbital, which can be denoted as follows: H+. When an ammonium ion is formed, the two-electron cloud of nitrogen becomes common to the nitrogen and hydrogen atoms, i.e. it turns into a molecular electron cloud. This means that a fourth covalent bond appears.
The process of formation of ammonium ion can be represented by the diagram:
The charge of the hydrogen ion becomes common (it is delocalized, i.e. dispersed between all atoms), and the two-electron cloud (lone electron pair) belonging to nitrogen becomes common with hydrogen. In diagrams, the cell image is often omitted.
The atom that provides a lone pair of electrons is called a donor, and the atom that accepts it (i.e., providing a vacant orbital) is called an acceptor.
The mechanism of formation of a covalent bond due to the two-electron cloud of one atom (donor) and the free orbital of another atom (acceptor) is called donor-acceptor. The covalent bond formed in this way is called a donor-acceptor, or coordination, bond.
However, this is not a special type of bond, but only a different mechanism (method) for the formation of a covalent bond. The properties of the fourth N-H bond in the ammonium ion are no different from the other bonds.
Metal connection
The atoms of most metals contain a small number of electrons at the outer energy level. Thus, 16 elements contain one electron each, 58 elements contain two electrons, 4 elements contain three electrons, and only Pd contains none. The atoms of the elements Ge, Sn and Pb have 4 electrons at the outer level, Sb and Bi - 5, Po - 6, but these elements are not characteristic metals.
The elements metals form simple substances- metals. Under normal conditions, these are crystalline substances (except mercury). In Fig. Figure 3 shows a diagram of the sodium crystal lattice.
As you can see, each sodium atom is surrounded by eight neighboring ones. Using sodium as an example, let us consider the nature of the chemical bond in metals.
The sodium atom, like other metals, has an excess of valence orbitals and a deficiency of electrons.
Thus, the valence electron (3s 1) can occupy one of nine free orbitals - 3s (one), 3p (three) and 3d (five).
When approaching atoms as a result of the formation changes in the crystal lattice, the valence orbitals of neighboring atoms overlap,
due to which electrons move freely from one orbital to another, communicating between all atoms of the metal crystal. This type of chemical bond is called metal bond.
A metallic bond is formed by elements whose atoms at the outer level have few valence electrons compared to the total number of outer orbitals that are energetically close, and the valence electrons, due to their low ionization energy, are weakly retained in the atom.
The chemical bond in metal crystals is highly delocalized, i.e. the electrons that carry out the communication are socialized (“electron gas”) and move throughout the entire piece of metal, which is generally electrically neutral.
Metallic bonding is characteristic of metals in solid and liquid states. This is a property of aggregates of atoms located in close proximity to each other. However, in the vapor state, metal atoms, like all substances, are connected to each other by covalent bonds. Metal pairs consist of individual molecules (monatomic and diatomic). The bond strength in a crystal is greater than in a metal molecule, and therefore the process of formation of a metal crystal occurs with the release of energy.
The metallic bond has some similarities with the covalent bond, since it is also based on the sharing of valence electrons. However, the electrons that perform a covalent bond are located close to the bonded atoms and are tightly bound to them. The electrons that carry out the metallic bond move freely throughout the crystal and belong to all its atoms. That is why crystals with a covalent bond are brittle, while those with a metal bond are ductile, i.e. they change shape when struck, are rolled into thin sheets, and drawn into wire.
Metallic bond explains physical properties metals
Hydrogen bond
A hydrogen bond is a kind of chemical bond. It can be intermolecular and intramolecular.
Intermolecular hydrogen bonding occurs between molecules that contain hydrogen and a strongly electronegative element - fluorine, oxygen, nitrogen, and less commonly chlorine and sulfur. Since in such a molecule the shared electron pair is strongly displaced from hydrogen to the atom of the electronegative element, and the positive charge of hydrogen is concentrated in a small volume, the proton interacts with the lone electron pair of another atom or ion, sharing it. As a result, a second, weaker bond is formed, called a hydrogen bond.
Previously, hydrogen bonding was reduced to electrostatic attraction between a proton and another polar group. But it should be more correct to consider that the donor-acceptor interaction also contributes to its formation. This connection is characterized by directionality in space and saturation.
Typically, a hydrogen bond is indicated by dots and this indicates that it is much weaker than a covalent bond (about 15-20 times). However, it is responsible for the association of molecules. For example, the formation of dimers (in the liquid state they are most stable) of water and acetic acid can be represented by diagrams:
As can be seen from these examples, two molecules of water, and in the case of acetic acid, two molecules of acid, are combined through hydrogen bonding to form a cyclic structure.
The presence of hydrogen bonds explains the higher boiling point of water (100° C) compared to hydrogen compounds of elements of the oxygen subgroup ( H2O, H2S, H2Te). In the case of water, additional energy must be expended to break hydrogen bonds.
Chemical bond.
Various substances have different structures. Of all substances known today, only inert gases exist in the form of free (isolated) atoms, which is due to their high stability electronic structures. All other substances (and more than 10 million of them are currently known) consist of bonded atoms.
Note: italics indicate those parts of the text that you do not need to learn or understand.
The formation of molecules from atoms leads to a gain in energy, since under normal conditions the molecular state is more stable than the atomic state.
An atom can have from one to eight electrons in its outer energy level. If the number of electrons in the outer level of an atom is the maximum that it can accommodate, then such a level is called completed. Completed levels are characterized by great strength. These are the outer levels of noble gas atoms: helium has two electrons at the outer level (s 2), the rest have eight electrons (ns 2 np 6). The outer levels of atoms of other elements are incomplete and in process chemical interaction they are completed.
A chemical bond is formed by valence electrons, but it occurs in different ways. There are three main types of chemical bonds: covalent, ionic and metallic.
Covalent bond
Let us consider the mechanism of the formation of a covalent bond using the example of the formation of a hydrogen molecule:
H + H = H 2; Q = 436 kJ
The nucleus of a free hydrogen atom is surrounded by a spherically symmetrical electron cloud formed by a 1 s electron. When atoms approach a certain distance, their electron clouds (orbitals) partially overlap.
As a result, a molecular two-electron cloud appears between the centers of both nuclei, which has a maximum electron density in the space between the nuclei; an increase in the negative charge density favors a strong increase in the forces of attraction between the nuclei and the molecular cloud.
So, a covalent bond is formed as a result of the overlap of electron clouds of atoms, accompanied by the release of energy. If the distance between the nuclei of hydrogen atoms approaching before touching is 0.106 nm, then after the electron clouds overlap (formation of the H2 molecule), this distance is 0.074 nm. The greatest overlap of electron clouds occurs along the line connecting the nuclei of two atoms (this occurs when a σ bond is formed). The greater the overlap of electron orbitals, the stronger the chemical bond. As a result of the formation of a chemical bond between two hydrogen atoms, each of them reaches the electronic configuration of an atom of the noble gas helium.
Chemical bonds are usually depicted in different ways:
1) using electrons in the form of dots placed at the chemical sign of the element. Then the formation of a hydrogen molecule can be shown by the diagram
H∙ + H∙ →H:H
2) often, especially in organic chemistry, a covalent bond is represented by a dash (prime) (for example, H-H), which symbolizes a shared pair of electrons.
The covalent bond in the chlorine molecule is also carried out using two shared electrons, or an electron pair:
Lone pair of electrons, there are 3 in an atom
← Lone pair of electrons,
There are 6 of them in a molecule.
unpaired electron shared or shared pair of electrons
As you can see, each chlorine atom has three lone pairs and one unpaired electron. The formation of a chemical bond occurs due to the unpaired electrons of each atom. Unpaired electrons bond into a shared pair of electrons, also called a shared pair.
If one covalent bond (one common electron pair) has arisen between atoms, then it is called a single bond; if more, then multiple double (two shared electron pairs), triple (three shared electron pairs).
A single bond is represented by one dash (prime), a double bond by two, and a triple bond by three. The dash between two atoms shows that they have a shared pair of electrons, as a result of which a chemical bond is formed. With the help of such dashes they depict structural formulas molecules.
So, in a chlorine molecule, each of its atoms has a complete outer level of eight electrons (s 2 p 6), and two of them (electron pair) belong equally to both atoms. The overlap of electronic orbitals during the formation of a molecule is shown in Fig: 
In the nitrogen molecule N2, the atoms have three common electron pairs:
:N· + ·N: → :N:::N:
Obviously, a nitrogen molecule is stronger than a hydrogen or chlorine molecule, which explains the significant inertness of nitrogen in chemical reactions.
A chemical bond carried out by electron pairs is called covalent.
Mechanisms of covalent bond formation.
A covalent bond is formed not only due to overlap single-electron clouds is an exchange mechanism for the formation of covalent bonds.
In an exchange mechanism, atoms share the same number of electrons.
Another mechanism of its formation is also possible - donor-acceptor. In this case, the chemical bond occurs due to unshared electron pair of one atom and free orbitals of another atom.
Let us consider as an example the mechanism of formation of the ammonium ion NH 4 +
When ammonia reacts with HCl, chemical reaction:
NH 3 + HCl = NH 4 Cl or in abbreviated ionic form: NH 3 + H + = NH 4 +
At the same time, in the ammonia molecule the nitrogen atom has unshared a couple of electrons (two-electron cloud):
Extremely rare chemicals consist of individual, unrelated atoms of chemical elements. Under normal conditions, only a small number of gases called noble gases have this structure: helium, neon, argon, krypton, xenon and radon. Most often, chemical substances do not consist of isolated atoms, but of their combinations into various groups. Such associations of atoms can number a few, hundreds, thousands, or even more atoms. The force that holds these atoms in such groups is called chemical bond.
In other words, we can say that a chemical bond is an interaction that provides the connection of individual atoms into more complex structures (molecules, ions, radicals, crystals, etc.).
The reason for the formation of a chemical bond is that the energy of more complex structures is less than the total energy of the individual atoms that form it.
So, in particular, if the interaction of atoms X and Y produces a molecule XY, this means that the internal energy of the molecules of this substance is lower than the internal energy of the individual atoms from which it was formed:
E(XY)< E(X) + E(Y)
For this reason, when chemical bonds are formed between individual atoms, energy is released.
Electrons of the outer electron layer with the lowest binding energy with the nucleus, called valence. For example, in boron these are electrons of the 2nd energy level - 2 electrons per 2 s- orbitals and 1 by 2 p-orbitals:
When a chemical bond is formed, each atom tends to obtain the electronic configuration of noble gas atoms, i.e. so that there are 8 electrons in its outer electron layer (2 for elements of the first period). This phenomenon is called the octet rule.
It is possible for atoms to achieve the electron configuration of a noble gas if initially single atoms share some of their valence electrons with other atoms. In this case, common electron pairs are formed.
Depending on the degree of electron sharing, covalent, ionic and metallic bonds can be distinguished.
Covalent bond
Covalent bonds most often occur between atoms of nonmetal elements. If the nonmetal atoms forming a covalent bond belong to different chemical elements, such a bond is called a polar covalent bond. The reason for this name lies in the fact that atoms different elements They also have different abilities to attract a common electron pair. Obviously, this leads to a displacement of the common electron pair towards one of the atoms, as a result of which a partial negative charge is formed on it. In turn, a partial positive charge is formed on the other atom. For example, in a hydrogen chloride molecule the electron pair is shifted from the hydrogen atom to the chlorine atom:
Examples of substances with polar covalent bonds:
CCl 4, H 2 S, CO 2, NH 3, SiO 2, etc.
A covalent nonpolar bond is formed between nonmetal atoms of the same chemical element. Since the atoms are identical, their ability to attract shared electrons is also the same. In this regard, no displacement of the electron pair is observed:
The above mechanism for the formation of a covalent bond, when both atoms provide electrons to form common electron pairs, is called exchange.
There is also a donor-acceptor mechanism.
When a covalent bond is formed by the donor-acceptor mechanism, a shared electron pair is formed due to the filled orbital of one atom (with two electrons) and the empty orbital of another atom. An atom that provides a lone pair of electrons is called a donor, and an atom with a vacant orbital is called an acceptor. Atoms that have paired electrons, for example N, O, P, S, act as donors of electron pairs.
For example, according to the donor-acceptor mechanism, the fourth covalent N-H bond is formed in the ammonium cation NH 4 +:
In addition to polarity, covalent bonds are also characterized by energy. Bond energy is the minimum energy required to break a bond between atoms.
The binding energy decreases with increasing radii of bonded atoms. Since we know that atomic radii increase down the subgroups, we can, for example, conclude that the strength of the halogen-hydrogen bond increases in the series:
HI< HBr < HCl < HF
Also, the bond energy depends on its multiplicity - the greater the bond multiplicity, the greater its energy. Bond multiplicity refers to the number of shared electron pairs between two atoms.
Ionic bond
An ionic bond can be considered as an extreme case of a polar covalent bond. If in a covalent-polar bond the common electron pair is partially shifted to one of the pair of atoms, then in an ionic bond it is almost completely “given” to one of the atoms. The atom that donates electron(s) acquires a positive charge and becomes cation, and the atom that has taken electrons from it acquires a negative charge and becomes anion.
Thus, an ionic bond is a bond formed due to the electrostatic attraction of cations to anions.
The formation of this type of bond is typical during the interaction of atoms of typical metals and typical non-metals.
For example, potassium fluoride. The potassium cation is formed by the removal of one electron from a neutral atom, and the fluorine ion is formed by the addition of one electron to the fluorine atom:
An electrostatic attraction force arises between the resulting ions, resulting in the formation of an ionic compound.
When a chemical bond was formed, electrons from the sodium atom passed to the chlorine atom and oppositely charged ions were formed, which have a completed external energy level.
It has been established that electrons from the metal atom are not completely detached, but are only shifted towards the chlorine atom, as in a covalent bond.
Most binary compounds that contain metal atoms are ionic. For example, oxides, halides, sulfides, nitrides.
Ionic bond also occurs between simple cations and simple anions (F −, Cl −, S 2-), as well as between simple cations and complex anions (NO 3 −, SO 4 2-, PO 4 3-, OH −). Therefore, ionic compounds include salts and bases (Na 2 SO 4, Cu(NO 3) 2, (NH 4) 2 SO 4), Ca(OH) 2, NaOH).
Metal connection
This type of bond is formed in metals.
Atoms of all metals have electrons in their outer electron layer that have a low binding energy with the nucleus of the atom. For most metals, the process of losing outer electrons is energetically favorable.
Due to such a weak interaction with the nucleus, these electrons in metals are very mobile and the following process continuously occurs in each metal crystal:
M 0 - ne - = M n + , where M 0 is a neutral metal atom, and M n + is a cation of the same metal. The figure below provides an illustration of the processes taking place.
That is, electrons “rush” across a metal crystal, detaching from one metal atom, forming a cation from it, joining another cation, forming a neutral atom. This phenomenon was called “electron wind,” and the collection of free electrons in a crystal of a nonmetal atom was called “electron gas.” This type of interaction between metal atoms is called a metallic bond.
Hydrogen bond
If a hydrogen atom in a substance is bonded to an element with high electronegativity (nitrogen, oxygen, or fluorine), that substance is characterized by a phenomenon called hydrogen bonding.
Since a hydrogen atom is bonded to an electronegative atom, a partial positive charge is formed on the hydrogen atom, and a partial negative charge is formed on the atom of the electronegative element. In this regard, electrostatic attraction becomes possible between a partially positively charged hydrogen atom of one molecule and an electronegative atom of another. For example, hydrogen bonding is observed for water molecules:
It is the hydrogen bond that explains the abnormally high melting point of water. In addition to water, strong hydrogen bonds are also formed in substances such as hydrogen fluoride, ammonia, oxygenated acids, phenols, alcohols, amines.
A covalent bond, depending on how the shared electron pair occurs, can be formed by exchange or donor-acceptor mechanism.
Exchange mechanism The formation of a covalent bond is realized in cases where both an atomic orbital and an unpaired electron located in this orbital participate in the formation of a common electron pair from each atom.
For example, in a hydrogen molecule. Interacting hydrogen atoms containing single electrons with opposite spins in atomic s-orbitals form a common electron pair, the movement of which in the H2 molecule occurs within the boundaries of the σ-molecular orbital, which arises when two s-atomic orbitals merge:
In the ammonia molecule, the nitrogen atom, having three single electrons and one electron pair in four atomic orbitals of the outer energy level, forms three common electron pairs with the s-electrons of three hydrogen atoms. These electron pairs in the NH 3 molecule are located in three σ-molecular orbitals, each of which arises when the atomic orbital of a nitrogen atom merges with the s-orbital of a hydrogen atom:
Thus, in an ammonia molecule, the nitrogen atom forms three σ-bonds with hydrogen atoms and has unshared electron pair.
Donor-acceptor mechanism the formation of a covalent bond occurs in cases where one neutral atom or ion (donor) has an electron pair in the atomic orbital of the outer energy level, and the other ion or neutral atom (acceptor)- free (vacant) orbital. When atomic orbitals merge, a molecular orbital appears in which there is a common electron pair that previously belonged to the donor atom:
According to the donor-acceptor mechanism, for example, the formation of a covalent bond between an ammonia molecule and a hydrogen ion occurs with the appearance of ammonium + ion. In the ammonia molecule, the nitrogen atom in the outer layer has a free electron pair, which allows this molecule to act as a donor. The hydrogen ion (acceptor) has a free s-orbital. Due to the fusion of the atomic orbitals of the nitrogen atom and the hydrogen ion, a σ-molecular orbital arises, and the free pair of electrons of the nitrogen atom becomes common to the connecting atoms:
Or H + + NH 3 [ H NH 3 ] +
In the ammonium ion + the covalent N-H bond, formed by the donor-acceptor mechanism, is equal in energy and length to three other covalent bonds N-H bonds, formed by the exchange mechanism.
The boron atom forms the boron fluoride molecule BF 3 due to the overlap of electron orbitals occupied in the excited state by unpaired electrons, with electron orbitals fluorine In this case, the boron atom retains one vacant orbital, due to which a fourth chemical bond can be formed through the donor-acceptor mechanism.
A bond formed by a donor-acceptor mechanism is often called donor-acceptor, coordination or coordinated. However, this is not a special type of bond, but only a different mechanism for the formation of a covalent bond.
The donor-acceptor mechanism of covalent bond formation is characteristic of complex compounds: The acceptor role is usually performed by d-metal ions, which can usually provide two, four or six free atomic orbitals of the s-, p-, d-type, which significantly expands their ability to form covalent bonds.
For example, Ag + and Cu 2+ ions, respectively, provide two and four free atomic orbitals, and the donor of electron pairs can be, for example, two or four molecules of ammonia or cyanide ion:
Acceptor Donor 
In these cases, covalent bonds arise between the donors and the acceptor with the formation of complex cations (silver and copper ammonia) or an anion (copper cyanide).
There are two main ways (mechanisms) for the formation of a covalent bond.
1) Spinvalent (exchange) mechanism : The electron pair that forms the bond is formed by the unpaired electrons present in the unexcited atoms.
However, the number of covalent bonds may be greater than the number of unpaired electrons. For example, in the unexcited state (also called the ground state), the carbon atom has two unpaired electrons, but it is characteristic of compounds in which it forms four covalent bonds. This turns out to be possible as a result of excitation of the atom. In this case, one of the s-electrons moves to the p-sublevel:
An increase in the number of covalent bonds created is accompanied by the release of more energy than is expended on excitation of the atom. Since the valence of an atom depends on the number of unpaired electrons, excitation leads to an increase in valence. In nitrogen, oxygen, and fluorine atoms, the number of unpaired electrons does not increase, because There are no free orbitals within the second level, and the movement of electrons to the third quantum level requires significantly more energy than that that would be released during the formation of additional bonds. Thus, when an atom is excited, electron transitions to free orbitals are possible only within one energy level.
Elements of the 3rd period - phosphorus, sulfur, chlorine - can exhibit a valency equal to the group number. This is achieved by excitation of atoms with the transition of 3s and 3p electrons to vacant orbitals of the 3d sublevel:
P* 1s 2 2s 2 2p 6 3s 1 3p 3 3d 1 (valency 5)
S* 1s 2 2s 2 2p 6 3s 1 3p 3 3d 2 (valence 6)
Cl* 1s 2 2s 2 2p 6 3s 1 3p 3 3d 3 (valency 7)
In the above electronic formulas for excited atoms, sublevels containing only unpaired electrons are underlined. Using the example of a chlorine atom, it is easy to show that valence can be variable:
Unlike chlorine, the valence of the F atom is constant and equal to 1, because At the valence (second) energy level there are no d-sublevel orbitals and other vacant orbitals.
2) Donor-acceptor mechanism : Covalent bonds are formed due to the paired electrons present in the outer electron layer of the atom. In this case, the second atom must have a free orbital on the outer layer. For example, the formation of an ammonium ion from an ammonia molecule and a hydrogen ion can be represented by the diagram: (the representation of electrons with crosses and dots in the diagram below is very conditional, since in reality the electrons are indistinguishable):
An atom that provides its electron pair to form a covalent bond is called a donor, and an atom that provides an empty orbital is called an acceptor. A covalent bond formed in this way is called a donor-acceptor bond. In the ammonium cation, this bond is absolutely identical in its properties to the other three covalent bonds formed by the exchange method.
Hybridization of atomic orbitals
To explain the difference between the bond angles in the H 2 O (104.5) and NH 3 (107.3) molecules from 90, it should be taken into account that the stable state of the molecule corresponds to its geometric structure with the lowest potential energy. Therefore, during the formation of a molecule, the shape and relative arrangement of atomic electron clouds changes in comparison with their shape and arrangement in free atoms. As a result, a more complete overlap of orbitals is achieved during the formation of a chemical bond. This deformation of the electron clouds requires energy, but more complete overlap results in a stronger bond and an overall energy gain. This explains the emergence of hybrid orbitals.
The shape of the hybrid orbital can be determined mathematically by adding the wave functions of the original orbitals:
As a result of adding the wave functions of the s- and p-orbitals, taking into account their signs, it turns out that the density of the electron cloud (value 2) on one side of the nucleus is increased, and on the other it is decreased.
In general, the hybridization process includes the following stages: excitation of the atom, hybridization of the orbitals of the excited atom, formation of bonds with other atoms. The energy costs for the first two stages are compensated by the energy gain during the formation of stronger bonds with hybrid orbitals. The type of hybridization is determined by the type and number of orbitals involved.
Below are examples of various types of hybridization of s- and p-orbitals.
Hybridization of one s- and one p-orbital (sp-hybridization) occurs, for example, during the formation of beryllium hydride, beryllium halides, zinc, and cadmium-mercury. Atoms of these elements in the normal state have two paired s-electrons in the outer layer. As a result of excitation, one of the s-electrons goes into the p-state - two unpaired electrons appear, one of which is an s-electron and the other a p-electron. When a chemical bond is formed, these two different orbitals are converted into two identical hybrid orbitals. Total number of orbitals during hybridization does not change . Two sp-hybrid orbitals are directed at an angle of 180º to each other and form two bonds (Figure 2):
Figure 2 - Overlapping sp-orbitals of beryllium and p-orbitals of chlorine in the BeCl 2 molecule
Experimental determination of the structure of the molecules BeG 2, ZnG 2, CdG 2, HgG 2 (G-halogen) showed that these molecules are linear, and both metal bonds with halogen atoms have the same length.
Hybridization of one s and two p orbitals (sp 2 hybridization) occurs, for example, in the formation of boron compounds. An excited boron atom has three unpaired electrons - one s-electron and two p-electrons. Three equivalent sp 2 -hybrid orbitals are formed from three orbitals, located in the same plane at an angle of 120 to each other (Figure 3). Indeed, as experimental studies show, the molecules of such boron compounds as BG 3 (G-halogen), B(CH 3) 3 - trimethylboron, B(OH) 3 - boric acid, have a flat structure. Moreover, the three boron bonds in these molecules have the same length and are located at an angle of 120.
Figure 3– Overlap of sp 2 -orbitals of boron and p-orbitals of chlorine in the BCl 3 molecule
Hybridization of one s- and three p-orbitals (sp 3 -hybridization) is characteristic, for example, of carbon and its analogues - silicon and germanium. In this case, the four hybrid sp3 orbitals are located at an angle of 10928 to each other; they are directed towards the vertices of the tetrahedron (in molecules CH 4, CCl 4, SiH 4, GeBr 4, etc.). The bond angles of the H 2 O (104.5º) and NH 3 (107.3º) molecules do not exactly correspond to the relative positions of the “pure” p-orbitals (90º). This is due to some contribution of s-electrons to the formation of a chemical bond. Such a contribution is nothing other than hybridization. The valence electrons in these molecules occupy four orbitals, which are close to sp 3 hybrid. The slight difference between the bond angles and the tetrahedral 109º28" is explained, according to Gillespie's theory, by the fact that unshared hybrid orbitals occupy a larger volume in space.
In many molecules the central atom does not undergo hybridization. Thus, bond angles in molecules H 2 S, PH 3, etc. are close to 90, i.e. the formation of bonds occurs with the participation of “pure” p-orbitals located at right angles to each other. 
