Rarely chemical substances 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 internal energy molecules of this substance are 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 of different elements 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 formation of the fourth covalent N-H connections 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, 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 bonding 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:

М 0 — ne − = M n + ,

where M 0 is a neutral metal atom, and M n + 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 anomalous heat melting water. Besides water, also durable hydrogen bonds are formed in substances such as hydrogen fluoride, ammonia, oxygen-containing acids, phenols, alcohols, amines.

When two atoms of the same nonmetal element interact, a covalent chemical bond is formed between them using shared electron pairs. This covalent bond is called non-polar because the shared electron pairs are shared by both atoms equally and neither of them will have an excess or deficiency of the negative charge carried by the electrons.

However, if a covalent bond is formed between atoms of different non-metal elements, then the picture will be somewhat different. Consider, for example, the formation of the hydrogen chloride molecule HC1 from hydrogen and chlorine atoms.

1. The hydrogen atom has one electron in its only level, and is missing one more electron before its completion. The chlorine atom has seven electrons in its outer shell and is also one electron short of completion.

2. Hydrogen and chlorine atoms combine their unpaired electrons and form one common electron pair, i.e., a covalent bond occurs:

Structural formula of the hydrogen chloride molecule H-C1.

3. Since a covalent bond is formed between atoms of different non-metal elements, the common electron pair will no longer belong to the interacting atoms equally. In order to qualitatively determine which of these atoms the shared electron pair will belong to a greater extent, the concept of electronegativity is used.

EO can be characterized as a measure of the non-metallicity of chemical elements. In order of decreasing EO chemical elements are arranged in the following row:

The most electronegative element in D.I. Mendeleev’s table is fluorine. This is, so to speak, the “gold medalist” of electronegativity. " Silver medalist"is oxygen, and "bronze" is nitrogen.

The value of the EO of an element depends on its position in D.I. Mendeleev’s table: in each period it usually increases with increasing serial number element, and in each subgroup it decreases.

Using a number of EOs, it is possible to determine where the common electron pairs are displaced. They are always shifted towards the atoms of the element with greater EO. For example, in the hydrogen chloride molecule HC1, the common electron pair is shifted to the chlorine atom, since its EO is greater than that of hydrogen. As a result, partial charges are formed on the atoms , two poles appear in the molecule - positive and negative. Therefore, such a covalent bond is called polar.

The displacement of shared electron pairs in the case of a polar covalent bond is sometimes indicated by arrows, and the partial charge is Greek letterδ (“delta”): .

In formulas of compounds, the chemical sign of the less electronegative element is written first. Since a covalent polar bond is a type of covalent bond, the reasoning algorithm for its schematic representation is the same as for a covalent nonpolar bond (see § 11), only in this case one more step will be added - the fourth: from a series of EOs we will determine the more electronegative element and reflect the polarity of the bond in the structural formula with an arrow and the designation of partial charges.

For example, consider an algorithm for a schematic representation of bond formation for the compound OF 2 - oxygen fluoride.

1. Oxygen is an element of the main subgroup of group VI (VIA group) of D.I. Mendeleev’s Periodic Table. Its atoms have six electrons in their outer electron layer. There will be unpaired electrons: 8-6 = 2.

Fluorine is an element of the main subgroup of group VII (group VIIA) of D. I. Mendeleev’s periodic table. Its atoms contain seven electrons in their outer electron layer. One electron is unpaired.

2. Let’s write down the signs of chemical elements with the designation of outer electrons:

3. Let’s write down the electronic and structural formulas of the resulting molecules:

4. Based on a number of EOs, we determine that the common electron pairs will be shifted from oxygen to fluorine, as to a more electronegative element, i.e., the bond will be polar covalent: .

Water molecules are formed in a similar way:

In reality, the water molecule has not a linear, but an angular shape (∠HOH = 104°27"). The structure of a water molecule can be depicted different ways(Fig. 40).

Rice. 40.
Various models water molecules

A hydrogen atom forms only one covalent bond with other atoms. Therefore, they say that hydrogen is monovalent. The oxygen atom is connected to other atoms by two chemical bonds - it is divalent. When molecules are formed, atoms are combined in such a way that all their valences are used. It is clear that divalent oxygen must combine with two monovalent hydrogen atoms. If we denote valency with a dash, then the formation scheme of a water molecule can be represented as follows:

Similarly, trivalent nitrogen combines with three monovalent hydrogen atoms to form an ammonia molecule

Formulas in which the valencies of elements are indicated by dashes, as you know, are called structural.

The structural formula of methane CH 4 - a compound of tetravalent carbon with hydrogen - will be as follows:

How do they combine into a molecule? carbon dioxide C0 2 atoms of tetravalent carbon and divalent oxygen? Obviously, this method can only be reflected by the following structural formula:

Is valence a constant value? It turns out that this statement is true for hydrogen and oxygen, but not for nitrogen and carbon, since these elements can exhibit other valency values. For example, nitrogen can be mono-, di-, tri-, or tetravalent. Its compounds with oxygen will have different compositions. Therefore, a distinction is made between:

• elements with constant valency(for example, monovalent: H, F; divalent: O, Be; trivalent: B, A1);
• elements with variable valence (for example, S exhibits valences II, IV, VI; C1 - valences I, III, V and VII).

Let's learn how to derive formulas for two-element compounds based on valency.

To derive the formula for the compound of phosphorus with oxygen, in which phosphorus is pentavalent, the procedure is as follows:

Similarly, we derive the formula for the compound of nitrogen with oxygen, in which nitrogen is tetravalent.

Index 1 is not written in formulas.

Knowing the valence of chemical elements is necessary in order to correctly write down the formula of a substance. However, the opposite is also true: using the formula of a substance, you can determine the valency of one of the elements if the valency of the other is known. For example, let’s determine the valence of sulfur in a compound whose formula is SO 3:

Laboratory experiment No. 4
Making models of molecules of binary compounds

Using ball-and-stick kits, assemble models of the molecules of the following substances:

• option 1 - hydrogen chloride HC1, carbon tetrachloride CC1 4;
• option 2 - sulfur dioxide SO 2, aluminum chloride AlCl 3.

Key words and phrases

1. Covalent nonpolar and covalent polar chemical bonds.
2. Electronegativity.
3. Partial charge.
4. Valence.
5. Drawing up formulas of covalent compounds by valence.
6. Determination of valency using formulas.

Work with computer

1. Refer to the electronic application. Study the lesson material and complete the assigned tasks.
2. Find email addresses on the Internet that can serve as additional sources that reveal the content of keywords and phrases in the paragraph. Offer your help to the teacher in preparing a new lesson - send a message by keywords and phrases in the next paragraph.

Questions and tasks

1. Hydrogen and phosphorus atoms have almost the same EO values. What is the type of chemical bond in the phosphine molecule PH 3?
2. Determine the type of chemical bond and write down the scheme of its formation for substances with the formulas: a) S 2, K 2 O and H 2 S; b) N 2, Li 3 N and C1 3 N.
3. In which of the molecules - hydrogen chloride HC1 or hydrogen fluoride HF - is the covalent chemical bond more polar?
4. In the following sentences, fill in the missing words and expressions: “A covalent chemical bond is formed due to.... According to the number of shared electron pairs, it is.... According to EO, a covalent bond is divided into... and...”.
5. Determine the valences of elements in compounds with the formulas: PbS, PbO 2, FeS 2, Fe 2 S 3, SF 6.
6. Write down the formulas of chlorides - compounds of elements with monovalent chlorine: iron (III), copper (I), copper (II), manganese (IV), phosphorus (V).

Rice. 2.1. The formation of molecules from atoms is accompanied by redistribution of electrons of valence orbitals and leads to gain in energy, since the energy of molecules turns out to be less than the energy of non-interacting atoms. The figure shows a diagram of the formation of a nonpolar covalent chemical bond between hydrogen atoms.

§2 Chemical bond

Under normal conditions, the molecular state is more stable than the atomic state (Fig. 2.1). The formation of molecules from atoms is accompanied by a redistribution of electrons in valence orbitals and leads to a gain in energy, since the energy of molecules is less than the energy of non-interacting atoms(Appendix 3). The forces that hold atoms in molecules are collectively called chemical bond.

The chemical bond between atoms is carried out by valence electrons and is electrical in nature . There are four main types of chemical bonds: covalent,ionic,metal And hydrogen.

1 Covalent bond

A chemical bond carried out by electron pairs is called atomic or covalent . Compounds with covalent bonds are called atomic or covalent .

When a covalent bond occurs, an overlap of electron clouds of interacting atoms occurs, accompanied by the release of energy (Fig. 2.1). In this case, a cloud with an increased density of negative charge appears between the positively charged atomic nuclei. Due to the action of Coulomb forces of attraction between unlike charges, an increase in the density of the negative charge favors the bringing together of nuclei.

A covalent bond is formed by unpaired electrons in the outer shells of atoms . In this case, electrons with opposite spins form electron pair(Fig. 2.2), common to interacting atoms. If one covalent bond (one common electron pair) has arisen between atoms, then it is called single, double, double, etc.

Energy is a measure of the strength of a chemical bond. E sv spent on breaking the bond (gain in energy when forming a compound from individual atoms). This energy is usually measured per 1 mole. substances and are expressed in kilojoules per mole (kJ∙mol –1). The energy of a single covalent bond lies in the range of 200–2000 kJmol –1.

Rice. 2.2. Covalent bond is the most common type of chemical bond that arises due to the sharing of an electron pair through an exchange mechanism (A), when each of the interacting atoms supplies one electron, or through a donor-acceptor mechanism (b), when an electron pair is transferred for common use by one atom (donor) to another atom (acceptor).

A covalent bond has the properties saturation and focus . The saturation of a covalent bond is understood as the ability of atoms to form a limited number of bonds with their neighbors, determined by the number of their unpaired valence electrons. The directionality of a covalent bond reflects the fact that the forces holding atoms near each other are directed along the straight line connecting the atomic nuclei. Besides, covalent bond can be polar or non-polar .

When non-polar In a covalent bond, the electron cloud formed by a common pair of electrons is distributed in space symmetrically relative to the nuclei of both atoms. A non-polar covalent bond forms between atoms simple substances, for example, between identical gas atoms forming diatomic molecules (O 2, H 2, N 2, Cl 2, etc.).

When polar In a covalent bond, the electron cloud of the bond is shifted toward one of the atoms. The formation of polar covalent bonds between atoms is characteristic of complex substances. An example would be volatile molecules organic compounds: HCl, H 2 O, NH 3, etc.

The degree of displacement of the total electron cloud towards one of the atoms during the formation of a covalent bond (degree of bond polarity ) determined mainly by the charge of atomic nuclei and the radius of interacting atoms .

The greater the charge of an atomic nucleus, the more strongly it attracts a cloud of electrons. At the same time, the larger the radius of the atom, the weaker the outer electrons are held near the atomic nucleus. The combined effect of these two factors is expressed in the different ability of different atoms to “pull” the cloud of covalent bonds towards themselves.

The ability of an atom in a molecule to attract electrons to itself is called electronegativity. . Thus, electronegativity characterizes the ability of an atom to polarize a covalent bond: the greater the electronegativity of an atom, the more strongly the electron cloud of the covalent bond is shifted towards it .

A number of methods have been proposed to quantify electronegativity. In this case, the clearest physical meaning has the method proposed by the American chemist Robert S. Mulliken, who determined electronegativity  of an atom as half the sum of its energy E e electron affinity and energy E i ionization of atom:

. (2.1)

Ionization energy An atom is the energy that must be expended to “tear off” an electron from it and remove it to an infinite distance. Ionization energy is determined by photoionization of atoms or by bombarding atoms with electrons accelerated in an electric field. The smallest value of photon or electron energy that becomes sufficient to ionize atoms is called their ionization energy E i. This energy is usually expressed in electron volts (eV): 1 eV = 1.610 –19 J.

Atoms are most willing to give up outer electrons metals, which contain a small number of unpaired electrons (1, 2 or 3) on the outer shell. These atoms have the lowest ionization energy. Thus, the magnitude of the ionization energy can serve as a measure of the greater or lesser “metallicity” of an element: the lower the ionization energy, the more pronounced the metalproperties element.

In the same subgroup of the periodic system of elements of D.I. Mendeleev, with an increase in the atomic number of an element, its ionization energy decreases (Table 2.1), which is associated with an increase in the atomic radius (Table 1.2), and, consequently, with a weakening of the bond of external electrons with a core. For elements of the same period, ionization energy increases with increasing atomic number. This is due to a decrease in atomic radius and an increase in nuclear charge.

Energy E e, which is released when an electron is added to a free atom, is called electron affinity(also expressed in eV). The release (rather than absorption) of energy when a charged electron attaches to some neutral atoms is explained by the fact that the most stable atoms in nature are those with filled outer shells. Therefore, for those atoms in which these shells are “a little unfilled” (i.e., 1, 2 or 3 electrons are missing before filling), it is energetically favorable to attach electrons to themselves, turning into negatively charged ions 1. Such atoms include, for example, halogen atoms (Table 2.1) - elements of the seventh group (main subgroup) of D.I. Mendeleev’s periodic system. The electron affinity of metal atoms is usually zero or negative, i.e. It is energetically unfavorable for them to attach additional electrons; additional energy is required to keep them inside the atoms. The electron affinity of nonmetal atoms is always positive and the greater, the closer to the noble (inert) gas the nonmetal is located in periodic table. This indicates an increase non-metallic properties as we approach the end of the period.

From all that has been said, it is clear that the electronegativity (2.1) of atoms increases in the direction from left to right for elements of each period and decreases in the direction from top to bottom for elements of the same group of the Mendeleev periodic system. It is not difficult, however, to understand that to characterize the degree of polarity of a covalent bond between atoms, it is not the absolute value of electronegativity that is important, but the ratio of the electronegativities of the atoms forming the bond. That's why in practice they use relative electronegativity values(Table 2.1), taking the electronegativity of lithium as unity.

To characterize the polarity of a covalent chemical bond, the difference in the relative electronegativity of atoms is used. Typically, the bond between atoms A and B is considered purely covalent if |  A – B|0.5.

The idea of ​​forming a chemical bond using a pair of electrons belonging to both connecting atoms was expressed in 1916 by the American physical chemist J. Lewis.

Covalent bonds exist between atoms in both molecules and crystals. It occurs both between identical atoms (for example, in H2, Cl2, O2 molecules, in a diamond crystal) and between different atoms (for example, in H2O and NH3 molecules, in SiC crystals). Almost all bonds in molecules of organic compounds are covalent (C-C, C-H, C-N, etc.).

There are two mechanisms for the formation of covalent bonds:

1) exchange;

2) donor-acceptor.

Exchange mechanism of covalent bond formationlies in the fact that each of the connecting atoms provides one unpaired electron for the formation of a common electron pair (bond). The electrons of interacting atoms must have opposite spins.

Let us consider, for example, the formation of a covalent bond in a hydrogen molecule. When hydrogen atoms come closer, their electron clouds penetrate into each other, which is called overlapping of electron clouds (Fig. 3.2), the electron density between the nuclei increases. The nuclei attract each other. As a result, the energy of the system decreases. When atoms come very close together, the repulsion of nuclei increases. Therefore, there is an optimal distance between the nuclei (bond length l), at which the system has minimum energy. In this state, energy is released, called the binding energy E St.

Rice. 3.2. Diagram of electron cloud overlap during the formation of a hydrogen molecule

Schematically, the formation of a hydrogen molecule from atoms can be represented as follows (a dot means an electron, a line means a pair of electrons):

N + N→N: N or N + N→N - N.

IN general view for AB molecules of other substances:

A + B = A: B.

Donor-acceptor mechanism of covalent bond formationlies in the fact that one particle - the donor - represents an electron pair to form a bond, and the second - the acceptor - represents a free orbital:

A: + B = A: B.

donor acceptor

Let's consider the mechanisms of formation of chemical bonds in the ammonia molecule and ammonium ion.

1. Education

The nitrogen atom has two paired and three unpaired electrons at its outer energy level:

The hydrogen atom in the s sublevel has one unpaired electron.

In the ammonia molecule, the unpaired 2p electrons of the nitrogen atom form three electron pairs with the electrons of 3 hydrogen atoms:

.

In the NH 3 molecule, 3 covalent bonds are formed according to the exchange mechanism.

2. Formation of a complex ion - ammonium ion.

NH 3 + HCl = NH 4 Cl or NH 3 + H + = NH 4 +

The nitrogen atom remains with a lone pair of electrons, i.e. two electrons with antiparallel spins in one atomic orbital. The atomic orbital of the hydrogen ion contains no electrons (vacant orbital). When an ammonia molecule and a hydrogen ion approach each other, an interaction occurs between the lone pair of electrons of the nitrogen atom and the vacant orbital of the hydrogen ion. The lone pair of electrons becomes common to the nitrogen and hydrogen atoms, and a chemical bond occurs according to the donor-acceptor mechanism. The nitrogen atom of the ammonia molecule is the donor, and the hydrogen ion is the acceptor:

.

It should be noted that in the NH 4 + ion all four bonds are equivalent and indistinguishable; therefore, in the ion the charge is delocalized (dispersed) throughout the complex.

The considered examples show that the ability of an atom to form covalent bonds is determined not only by one-electron, but also by 2-electron clouds or the presence of free orbitals.

According to the donor-acceptor mechanism, bonds are formed in complex compounds: - ; 2+ ; 2- etc.

A covalent bond has the following properties:

- saturation;

- directionality;

- polarity and polarizability.

Data on ionization energy (IE), PEI and the composition of stable molecules - their actual values ​​and comparisons - both of free atoms and of atoms bound into molecules, allow us to understand how atoms form molecules through the mechanism of covalent bonding.

COVALENT BOND- (from the Latin “co” together and “vales” having force) (homeopolar bond), a chemical bond between two atoms that arises when the electrons belonging to these atoms are shared. Atoms in the molecules of simple gases are connected by covalent bonds. A bond in which there is one shared pair of electrons is called a single bond; There are also double and triple bonds.

Let's look at a few examples to see how we can use our rules to determine the number of covalent chemical bonds an atom can form if we know the number of electrons in a given atom's outer shell and the charge on its nucleus. The charge of the nucleus and the number of electrons in the outer shell are determined experimentally and are included in the table of elements.

Calculation of the possible number of covalent bonds

For example, let's count the number of covalent bonds that can form sodium ( Na), aluminum (Al), phosphorus (P), and chlorine ( Cl). Sodium ( Na) and aluminum ( Al) have, respectively, 1 and 3 electrons in the outer shell, and, according to the first rule (for the mechanism of covalent bond formation, one electron in the outer shell is used), they can form: sodium (Na)- 1 and aluminum ( Al)- 3 covalent bonds. After bond formation, the number of electrons in the outer shells of sodium ( Na) and aluminum ( Al) equal to 2 and 6, respectively; i.e., less maximum quantity(8) for these atoms. Phosphorus ( P) and chlorine ( Cl) have, respectively, 5 and 7 electrons on the outer shell and, according to the second of the above-mentioned laws, they could form 5 and 7 covalent bonds. In accordance with the fourth law, the formation of a covalent bond, the number of electrons on the outer shell of these atoms increases by 1. According to the sixth law, when a covalent bond is formed, the number of electrons on the outer shell of the bonded atoms cannot be more than 8. That is, phosphorus ( P) can only form 3 bonds (8-5 = 3), while chlorine ( Cl) can form only one (8-7 = 1).

Example: Based on the analysis, we discovered that a certain substance consists of sodium atoms (Na) and chlorine ( Cl). Knowing the regularities of the mechanism of formation of covalent bonds, we can say that sodium ( Na) can form only 1 covalent bond. Thus, we can assume that each sodium atom ( Na) bonded to the chlorine atom ( Cl) through a covalent bond in this substance, and that this substance is composed of molecules of an atom NaCl. The structural formula for this molecule: Na-Cl. Here the dash (-) denotes a covalent bond. The electronic formula of this molecule can be shown as follows:
. .
Na:Cl:
. .
In accordance with the electronic formula, on the outer shell of the sodium atom ( Na) V NaCl there are 2 electrons, and on the outer shell of the chlorine atom ( Cl) there are 8 electrons. In this formula, electrons (dots) between sodium atoms ( Na) And chlorine (Cl) are bonding electrons. Since the PEI of chlorine ( Cl) is equal to 13 eV, and for sodium (Na) it is equal to 5.14 eV, the bonding pair of electrons is much closer to the atom Cl than to an atom Na. If the ionization energies of the atoms forming the molecule are very different, then the bond formed will be polar covalent bond.

Let's consider another case. Based on the analysis, we discovered that a certain substance consists of aluminum atoms ( Al) and chlorine atoms ( Cl). In aluminum ( Al) there are 3 electrons in the outer shell; thus, it can form 3 covalent chemical bonds while chlorine (Cl), as in the previous case, can form only 1 bond. This substance is presented as AlCl3, and its electronic formula can be illustrated as follows:

Figure 3.1. Electronic formulaAlCl 3

whose formula of structure is:
Cl - Al - Cl
Cl

This electronic formula shows that AlCl3 on the outer shell of chlorine atoms ( Cl) there are 8 electrons, while the outer shell of the aluminum atom ( Al) there are 6 of them. According to the mechanism of formation of a covalent bond, both bonding electrons (one from each atom) go to the outer shells of the bonded atoms.

Multiple covalent bonds

Atoms that have more than one electron in their outer shell can form not one, but several covalent bonds with each other. Such connections are called multiple (more often multiples) connections. Examples of such bonds are the bonds of nitrogen molecules ( N= N) and oxygen ( O=O).

The bond formed when single atoms join together is called homoatomic covalent bond, e If the atoms are different, then the bond is called heteroatomic covalent bond[Greek prefixes "homo" and "hetero" respectively mean same and different].

Let's imagine what a molecule with paired atoms actually looks like. The simplest molecule with paired atoms is the hydrogen molecule.