The name of the chemical bond. Chemical bond: types, reasons, main characteristics

С 2s 2 2p 2 С + 1е \u003d С -

О 2s 2 2p 4 О -1е \u003d О +

Another explanation for the formation of a triple bond in the CO molecule is possible.

An unexcited carbon atom has 2 unpaired electrons, which can form 2 common electron pairs with 2 unpaired electrons of the oxygen atom (by the exchange mechanism). However, the 2 paired p-electrons present in the oxygen atom can form a triple chemical bond, since there is one unfilled cell in the carbon atom, which can accept this pair of electrons.

The triple bond is formed by the donor-acceptor mechanism, the direction of the arrow is from the oxygen donor to the acceptor - carbon.

Like N 2 - CO has a high dissociation energy (1069 kJ), is poorly soluble in water, and chemically inert. CO is a colorless and odorless gas, indifferent, non-salt-forming, does not interact with acidic alkalis and water under normal conditions. Poisonous because interacts with iron, which is part of hemoglobin. With increasing temperature or irradiation, it exhibits the properties of a reducing agent.

Receiving:

in industry

CO 2 + C «2CO

2C + O 2 ® 2CO

in the laboratory: H 2 SO 4, t

HCOOH® CO + H 2 O;

H 2 SO 4 t

H 2 C 2 O 4 ® CO + CO 2 + H 2 O.

CO enters into the reaction only at high temperatures.

The CO molecule has a high affinity for oxygen, it burns to form CO 2:

CO + 1 / 2O 2 \u003d CO 2 + 282 kJ / mol.

Due to its high affinity for oxygen, CO is used as a reducing agent for oxides of many heavy metals (Fe, Co, Pb, etc.).

CO + Cl 2 \u003d COCl 2 (phosgene)

CO + NH 3 ® HCN + H 2 O H - C º N

CO + H 2 O «CO 2 + H 2

CO + S ® COS

Of greatest interest are metal carbonyls (used to obtain pure metals). The chemical bond is by the donor-acceptor mechanism; p-overlapping by the dative mechanism takes place.

5CO + Fe® (iron pentacarbonyl)

All carbonyls are diamagnetic substances, characterized by low strength, when heated, carbonyls decompose

→ 4CO + Ni (nickel carbonyl).

Like СО, metal carbonyls are toxic.

Chemical bond in the CO 2 molecule

In the CO 2 molecule sp-hybridization of a carbon atom. Two sp-hybrid orbitals form 2 s-bonds with oxygen atoms, and the remaining unhybridized p-orbitals of carbon give p-bonds with two p-orbitals of oxygen atoms, which are located in planes perpendicular to each other.

O ═ S ═ O

Under a pressure of 60 atm. and at room temperature, CO 2 condenses into a colorless liquid. With strong cooling, liquid CO 2 solidifies into a white, snow-like mass that sublimes at P \u003d 1 atm and t \u003d 195K (-78 °). The compressed solid mass is called dry ice, CO 2 does not support combustion. It burns only substances with a higher affinity for oxygen than carbon: for example,

2Mg + CO 2 ® 2MgO + C.

CO 2 reacts with NH 3:

CO 2 + 2NH 3 \u003d CO (NH 2) 2 + H 2 O

(carbamide, urea)

2CO 2 + 2Na 2 O 2 ® 2Na 2 CO 3 + O 2

Urea is decomposed by water:

CO (NH 2) 2 + 2H 2 O ® (NH 4) 2 CO 3 → 2NH 3 + CO 2

Cellulose is a carbohydrate that consists of b-glucose residues. It is synthesized in plants according to the following scheme

chlorophyll

6CO 2 + 6H 2 O ® C 6 H 12 O 6 + 6O 2 photosynthesis of glucose

CO 2 is obtained in technology:

2NaHCO 3 ® Na 2 CO 3 + H 2 O + CO 2

from coke C + O 2 ® CO 2

In the laboratory (in the Kipp apparatus):

Carbonic acid and its salts

Dissolving in water, carbon dioxide partially interacts with it, forming carbonic acid H 2 CO 3; in this case, equilibria are established:

K 1 \u003d 4 × 10 -7 K 2 \u003d 4.8 × 10 -11 - weak, unstable, oxygen-containing, dibasic acid. Bicarbonates are soluble in H 2 O. Carbonates are insoluble in water, except for alkali metal carbonates, Li 2 CO 3 and (NH 4) 2 CO 3. Acidic salts of carbonic acid are obtained by passing excess CO 2 into an aqueous carbonate solution:

or the gradual (dropwise) addition of a strong acid to an excess of an aqueous carbonate solution:

Na 2 CO 3 + HNO 3 ® NaHCO 3 + NaNO 3

When interacting with alkalis or heating (calcining), acidic salts turn into medium ones:

Salts are hydrolyzed according to the equation:

Stage I

Due to complete hydrolysis, carbonates Gr 3+, Al 3+, Ti 4+, Zr 4+, etc. cannot be isolated from aqueous solutions.

Salts are of practical importance - Na 2 CO 3 (soda), CaCO 3 (chalk, marble, limestone), K 2 CO 3 (potash), NaHCO 3 (baking soda), Ca (HCO 3) 2 and Mg (HCO 3) 2 determine the carbonate hardness of the water.

Carbon disulfide (CS 2)

When heated (750-1000 ° C), carbon reacts with sulfur, forming carbon disulfide, organic solvent (colorless volatile liquid, reactive substance), flammable and volatile.

CS 2 vapors are poisonous, used for fumigation (fumigation) of granaries against insect pests, in veterinary medicine it is used to treat ascariasis of horses. In technology - a solvent for resins, fats, iodine.

With metal sulfides, CS 2 forms thiocarbonate salts - thiocarbonates.

This reaction is similar to the process

Thiocarbonates - yellow crystalline substances. When acids act on them, free thiocarbonic acid is released.

It is more stable than H 2 CO 3 and at low temperatures it is released from the solution in the form of a yellow oily liquid, which readily decomposes into:

Compounds of carbon with nitrogen (СN) 2 or С 2 N 2 - dicyan, highly flammable colorless gas. Pure dry cyanogen is obtained by heating mercuric chloride with mercury (II) cyanide.

HgCl 2 + Hg (CN) 2 ® Hg 2 Cl 2 + (C N) 2

Other ways to get:

4HCN g + O 2 2 (CN) 2 + 2H 2 O

2HCN g + Cl 2 (CN) 2 + 2HCl

Dicyan is similar in properties to halogens in the molecular form X 2. So in an alkaline environment, it, like halogens, disproportionates:

(C N) 2 + 2NaOH \u003d NaCN + NaOCN

Hydrogen cyanide - HCN (), a covalent compound, a gas that dissolves in water to form hydrocyanic acid (a colorless liquid and its salts are extremely toxic). Receive:

Hydrogen cyanide is produced industrially by catalytic reactions.

2CH 4 + 3O 2 + 2NH 3 ® 2HCN + 6H 2 O.

Hydrocyanic acid salts - cyanides, are subject to strong hydrolysis. CN - - ion isoelectronic to the CO molecule, is included as a ligand in a large number of complexes of d-elements.

Handling cyanide requires strict precautions. In agriculture, they are used to combat especially dangerous insects - pests.

Cyanides get:

Carbon compounds with a negative oxidation state:

1) covalent (SiC carborundum) ;

2) ionic-covalent;

3) metal carbides.

Ionic-valent decompose by water with the release of gas, depending on which gas is released, they are divided into:

methanides (CH 4 is allocated)

Al 4 C 3 + 12H 2 O ® 4Al (OH) 3 + 3CH 4

acetylenides(C 2 H 2 is allocated)

H 2 C 2 + AgNO 3 ® Ag 2 C 2 + HNO 3

Metal carbides are compounds of stoichiometric composition formed by elements of 4, 7, 8 groups by introducing Me atoms into the crystal lattice of carbon.

Silicon chemistry

The difference between the chemistry of silicon and carbon is due to the large size of its atom and the possibility of using 3d orbitals. Because of this, the Si - O - Si, Si - F bonds are stronger than those of carbon.

For silicon, oxides of the composition SiO and SiO 2 are known. Silicon monoxide exists only in the gas phase at high temperatures in an inert atmosphere; it is readily oxidized with oxygen to form the more stable oxide SiO 2.

2SiO + О 2 t ® 2SiO 2

SiO 2- silica, has several crystalline modifications. Low-temperature - quartz, has piezoelectric properties. Natural varieties of quartz: rock crystal, topaz, amethyst. The varieties of silica are chalcedony, opal, agate, sand.

A wide variety of silicates (more precisely, oxosilicates) are known. In their structure, they have a common pattern: they all consist of SiO 4 4 tetrahedra - which are connected to each other through an oxygen atom.

Combinations of tetrahedrons can be connected in chains, ribbons, meshes, and wireframes.

Important natural silicates are 3MgO × H 2 O × 4SiO 2 talc, 3MgO × 2H 2 O × 2SiO 2 asbestos.

Like SiO 2, silicates are characterized by an (amorphous) glassy state. With controlled crystallization, it is possible to obtain a fine-crystalline state - sitalls - materials of increased strength. In nature, aluminosilicates are widespread - frame orthosilicates, some of the Si atoms are replaced by Al, for example Na 12 [(Si, Al) O 4] 12.

The strongest halide SiF 4 decomposes only under the action of an electric discharge.

Hexafluorosilicic acid (close in strength to H 2 SO 4).

(SiS 2) n - polymeric substance, decomposed by water:

Silicic acids.

The corresponding SiO 2 silicic acids do not have a specific composition, they are usually written in the form xH 2 O ySiO 2 - polymer compounds

Known:

H 2 SiO 3 (H 2 O × SiO 2) - metasilicon (does not really exist)

H 4 SiO 4 (2H 2 O × SiO 2) - orthosilicon (the simplest actually exists only in solution)

H 2 Si 2 O 5 (H 2 O × 2SiO 2) - dimetysilicon.

Silicic acids are poorly soluble substances; for H 4 SiO 4, a colloidal state is characteristic, as an acid is weaker than carbonic (Si is less metallic than C).

In aqueous solutions, orthosilicic acid condenses, resulting in the formation of polysilicic acids.

Silicates are salts of silicic acids, insoluble in water, except for alkali metal silicates.

Soluble silicates are hydrolyzed according to the equation

Jelly-like solutions of sodium salts of polysilicic acids are called "liquid glass". They are widely used as silicate glue and wood preservative.

By fusion of Na 2 CO 3, CaCO 3 and SiO 2 glass is obtained, which is a supercooled mutual solution of polysilicic acid salts.

6SiO 2 + Na 2 CO 3 + CaCO 3 ® Na 2 O × CaO × 6SiO 2 + 2CO 2 Silicate is written as a mixed oxide.

Silicates are most commonly used in construction. 1st place in the world for the production of silicate products - cement, 2nd - brick, 3rd - glass.

Building ceramics - facing tiles, ceramic pipes. For the manufacture of sanitary ware - glass, porcelain, earthenware, clay ceramics.

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 able to hold two positively charged ions in a single whole. In one-electron coupling, the Coulomb forces of repulsion 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 +:

A polar covalent bond occurs in heteronuclear diatomic molecules (Fig. 3). The bonding electron pair in a polar chemical bond is close to the atom with a higher first ionization potential.

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

Characterization of some polar substances

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

The distance between the centers of gravity of the positive and negative charges is called the dipole length. The polarity of a molecule, like the polarity of a bond, is estimated by the magnitude of the dipole moment μ, which is the product of the dipole length l by 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 concept of sigma and π-bonds, hybridization of atomic orbitals.

Pauling's hybridization for two S and two p electrons made it possible to explain the direction of chemical bonds, in particular, the tetrahedral configuration of methane. To explain the structure of ethylene, one p-electron has to be isolated from four equivalent Sp 3 - electrons of a carbon atom 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 Pauling's new theory, all bonding electrons became equal and equidistant from the line connecting the nuclei of the molecule. Pauling's bent chemical bond theory took into account the statistical interpretation of the M. Born wave function, the Coulomb electron correlation of electrons. A physical meaning has appeared - the nature of the chemical bond is completely determined by the electrical interaction of nuclei and electrons. The more bonding electrons, the shorter the internuclear distance and the stronger the chemical bond between carbon atoms.

Three-center chemical bond

Further development of the concept of chemical bonding was given by the American physicochemist 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 (borohydrides).

An electron pair in a three-center chemical bond becomes common for three atomic nuclei. In the simplest representative of the three-center chemical bond, the molecular hydrogen ion H 3 +, the electron pair holds three protons in a single whole (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, which was previously considered a standard monovalent element, turned out to be bound by the same bonds with two boron atoms and became formally a divalent element. The works of W. Lipscomb on deciphering the structure of boranes expanded the concept of chemical bonds. The Nobel Committee awarded the 1976 Chemistry Prize to William Nunn Lipscomb with the formulation "For studies of the structure of boranes (borohydrites), clarifying the problems of chemical bonds).

Multicenter chemical bond

Fig. 8 Ferrocene molecule

Fig. 9 Dibenzene chromium

Fig. 10 Uranocene

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

Chemical bond dynamics

The chemical bond is quite dynamic. Thus, a metal bond is transformed into a covalent bond during a phase transition during metal evaporation. The transition of a metal from a solid to a vapor state requires large amounts of energy.

In vapors, these metals consist essentially of homonuclear diatomic molecules and free atoms. Upon condensation of metal vapors, the covalent bond turns into a metallic one.

The evaporation of salts with a typical ionic bond, for example, alkali metal fluorides, leads to the destruction of the ionic bond and the formation of heteronuclear diatomic molecules with a polar covalent bond. In this case, the formation of dimeric molecules with bridging bonds takes place.

Characterization of the chemical bond in the molecules of alkali metal fluorides and their dimers.

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

Mechanism of transition from covalent to metallic bond

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

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

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

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

The formation of a covalent metal-metal bond is associated with deformation of the electron 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 the bond length in the hydrogen molecule) and the low energy of its breaking.

The specified ratio between r e and d determines the uneven distribution of electric charges in the molecule - the negative electric charge of the connecting electron pair is concentrated in the middle part of the molecule, and the positive electric charges of the two atomic cores are concentrated at the ends of the molecule.

The uneven distribution of electric charges creates conditions for the interaction of molecules due to orientational forces (van der Waals forces). Molecules of alkali metals tend to orient themselves in such a way that opposite electric charges appear in the neighborhood. As a result, attractive forces act between the molecules. Due to the presence of the latter, the molecules of alkali metals come closer and more or less firmly pull together. At the same time, some deformation of each of them occurs under the action of the 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, break off from the orbital radius of the atom and become free.

In this case, the bonding electron pair becomes common even for the system with six cations. The construction of the metal crystal lattice begins at the cluster stage. In the crystal lattice of alkali metals, the structure of a connecting link is clearly expressed, which has 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 constant of the translational lattice a w (Fig. 13).

The value of the translational 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 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 a geometrical place where electrons reside, providing the main property of a metal - to conduct an electric current.

When comparing the process of condensation of vapors of alkali metals with the process of condensation of gases, for example, hydrogen, a characteristic feature appears in the properties of the metal. So, if weak intermolecular interactions appear during the condensation of hydrogen, then during the condensation of metal vapors, processes characteristic of chemical reactions occur. The condensation of metal vapors itself goes through several stages and can be described by the following procession: 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 dimeric molecule can be viewed as an electric quadrupole (Fig. 15). At present, the main characteristics of dimers of alkali metal halides (chemical bond lengths and bond angles) are known.

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

E 2 X 2	X \u003d F		X \u003d Cl		X \u003d Br		X \u003d I
E 2 X 2	d EF, 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

In the process of condensation, the effect of orientational forces is enhanced, intermolecular interaction is accompanied by the formation of clusters, and then a solid. Alkali metal halides form crystals with a simple cubic and body-centered cubic lattice.

Type of crystal lattice and translation lattice constant for alkali metal halides.

In the process of crystallization, a further increase in the interatomic distance occurs, leading to the separation of an electron from the orbital radius of an alkali metal atom and the transfer of an electron to a halogen atom with the formation of the corresponding ions. The force fields of the ions are evenly distributed in all directions in space. In this regard, in crystals of alkali metals, the force field of each ion coordinates by no means 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 bound to six chlorine ions (in a sodium chloride crystal) and to eight chlorine ions (in a cesium chloride crystal. the interionic distances in crystals are equidistant.

Notes

Handbook of Inorganic Chemistry. Constants of inorganic substances. - M .: "Chemistry", 1987. - S. 124. - 320 p.
Lidin R.A., Andreeva L.L., Molochko V.A. Handbook of Inorganic Chemistry. Constants of inorganic substances. - M .: "Chemistry", 1987. - S. 132-136. - 320 p.
Gankin V.Yu., Gankin Yu.V. How a chemical bond is formed and chemical reactions take place. - M .: Publishing group "Border", 2007. - 320 p. - ISBN 978-5-94691296-9
B.V. Nekrasov General chemistry course. - M .: Goskhimizdat, 1962 .-- S. 88 .-- 976 p.
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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. - S. 63-86.
Chemical encyclopedic dictionary / Ch. ed. I.L. Knunyants. - M .: Sov. encyclopedia, 1983 .-- S. 607 .-- 792 p.
B.V. Nekrasov 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. - S. 155-161. - 320 p.
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Exchange and donor-acceptor mechanism of covalent bond formation

1) Exchange mechanism. Each atom gives one unpaired electron to a common electron pair.

2) Donor-acceptor mechanism. One atom (donor) provides an electron pair, and the other atom (acceptor) provides a free orbital for this pair.

Any interaction between atoms is possible only if there is a chemical bond. This bond is the reason for the formation of a stable polyatomic system - a molecular ion, a molecule, a crystal lattice. A strong chemical bond requires a lot of energy to break, which is why it is the basic value for measuring bond strength.

Conditions for the formation of a chemical bond

The formation of a chemical bond is always accompanied by the release of energy. This process occurs due to a decrease in the potential energy of the system of interacting particles - molecules, ions, atoms. The potential energy of the resulting system of interacting elements is always less than the energy of unbound outgoing particles. Thus, the basis for the occurrence of a chemical bond in the system is the decline in the potential energy of its elements.

The nature of the chemical interaction

Chemical bond is a consequence of the interaction of electromagnetic fields that arise around the electrons and nuclei of atoms of those substances that take part in the formation of a new molecule or crystal. After the discovery of the theory of atomic structure, the nature of this interaction became more accessible for study.

For the first time, the idea of \u200b\u200bthe electrical nature of chemical bonds arose from the English physicist G. Davy, who suggested that molecules are formed due to the electric attraction of oppositely charged particles. This idea interested the Swedish chemist and naturalist I.Ya. Bercellius, who developed the electrochemical theory of chemical bonding.

The first theory that explained the processes of chemical interaction of substances was imperfect, and over time it had to be abandoned.

Butlerov's theory

A more successful attempt to explain the nature of the chemical bond of substances was made by the Russian scientist A.M. Butlerov. This scientist based his theory on the following assumptions:

Atoms in a connected state are linked to each other in a specific order. A change in this order causes the formation of a new substance.
The atoms bond with each other according to the laws of valence.
The properties of a substance depend on the order in which atoms are joined in a substance molecule. A different order of arrangement causes a change in the chemical properties of a substance.
Atoms connected to each other have the greatest impact on each other.

Butlerov's theory explained the properties of chemicals not only by their composition, but also by the order of arrangement of atoms. Such an internal order of A.M. Butlerov called it "chemical structure".

The theory of the Russian scientist made it possible to put things in order in the classification of substances and made it possible to determine the structure of molecules by their chemical properties. The theory also gave an answer to the question: why molecules containing the same number of atoms have different chemical properties.

Prerequisites for the creation of chemical bond theories

In his theory of chemical structure, Butlerov did not touch on the question of what a chemical bond is. For this then there was too little data on the internal structure of matter. Only after the discovery of the planetary model of the atom, the American scientist Lewis began to develop a hypothesis that a chemical bond occurs through the formation of an electron pair that simultaneously belongs to two atoms. Subsequently, this idea became the foundation for the development of the theory of covalent bonds.

Covalent chemical bond

A stable chemical compound can be formed when the electron clouds of two neighboring atoms overlap. The result of this mutual intersection is an increasing electron density in the internuclear space. The nuclei of atoms, as you know, are positively charged, and therefore try to be attracted as close as possible to the negatively charged electron cloud. This attraction is much stronger than the repulsive forces between two positively charged nuclei, so this bond is stable.

For the first time, calculations of the chemical bond were carried out by chemists Geitler and London. They considered the bond between two hydrogen atoms. The simplest visual representation of it might look like this:

As you can see, the electron pair occupies a quantum place in both hydrogen atoms. This two-center arrangement of electrons is called "covalent chemical bond". The covalent bond is typical for molecules of simple substances and their compounds of non-metals. Substances created as a result of a covalent bond usually do not conduct electric current or are semiconductors.

Ionic bond

A chemical bond of the ionic type occurs when two oppositely charged ions are mutually attracted. Ions can be simple, consisting of one atom of matter. In compounds of this type, simple ions are most often positively charged atoms of metals of group 1, 2 that have lost their electron. The formation of negative ions is inherent in the atoms of typical non-metals and the bases of their acids. Therefore, among the typical ionic compounds there are many alkali metal halides, for example, CsF, NaCl, and others.

Unlike a covalent bond, an ion is not saturated: a different number of oppositely charged ions can attach to an ion or a group of ions. The number of attached particles is limited only by the linear dimensions of the interacting ions, as well as by the condition under which the forces of attraction of oppositely charged ions must be greater than the forces of repulsion of equally charged particles participating in the ionic type compound.

Hydrogen bond

Even before the creation of the theory of chemical structure, it was experimentally noticed that hydrogen compounds with various non-metals have somewhat unusual properties. For example, the boiling points of hydrogen fluoride and water are significantly higher than would be expected.

These and other features of hydrogen compounds can be explained by the ability of the H + atom to form another chemical bond. This type of compound is called "hydrogen bond". The reasons for the occurrence of hydrogen bonding lie in the properties of electrostatic forces. For example, in a molecule of hydrogen fluoride, the total electron cloud is so displaced towards fluorine that the space around the atom of this substance is saturated with a negative electric field. Around a hydrogen atom, devoid of its only electron, the field is much weaker and has a positive charge. As a result, an additional relationship arises between the positive fields of electron clouds H + and negative F -.

Chemical bond of metals

The atoms of all metals are located in space in a certain way. The order of the metal atoms is called the crystal lattice. In this case, the electrons of different atoms weakly interact with each other, forming a common electron cloud. This type of interaction between atoms and electrons is called "metal bond".

It is the free movement of electrons in metals that can explain the physical properties of metallic substances: electrical conductivity, thermal conductivity, strength, fusibility, and others.

Covalent chemical bond, its varieties and mechanisms of formation. Characterization of a covalent bond (polarity and bond energy). Ionic bond. Metallic bond. Hydrogen bond

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

A chemical bond is understood as such an interaction of atoms that binds them into molecules, ions, radicals, 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.

The ionic bond can be considered as the limiting case of the covalent polar bond.

The metallic bond combines the covalent interaction of atoms with the help of shared electrons and the electrostatic attraction between these electrons and metal ions.

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

For example, lithium fluoride $ LiF $ is referred to as ionic compounds. In fact, the bond in it is $ 80% $ ionic and $ 20% $ covalent. Therefore, it is more correct to speak about the degree of polarity (ionicity) of a chemical bond.

In the series of hydrogen halides $ HF — HCl — HBr — HI — HАt $, the degree of bond polarity decreases, because the difference in the values \u200b\u200bof electronegativity of halogen and hydrogen atoms decreases, and in hydrogen astate the bond becomes almost nonpolar $ (EO (H) \u003d 2.1; EO (At) \u003d 2.2) $.

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

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

Various types of links can transform one into another:

- during electrolytic dissociation of covalent compounds in water, the covalent polar bond transforms into an ionic one;

- upon evaporation of metals, the metal bond turns into a covalent non-polar one, 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 the electron-nuclear interaction of atoms, accompanied by the release of energy.

Methods for the formation of a covalent bond. Covalent bond characteristics: bond length and energy

A covalent chemical bond is a bond that occurs between atoms due to the formation of common electron pairs.

The mechanism for the formation of such a bond can be exchange and donor-acceptor.

I. Exchange mechanism acts when atoms form common 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 from $ 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 of $ p-p- $ orbitals):

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

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

The donor has an electron pair, the acceptor has a free 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 by the exchange mechanism, one - by the donor-acceptor mechanism.

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

The chemical bonds formed as a result of the overlap of electron orbitals along the bond line are called $ σ $ -links (sigma-links)... The sigma link 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 $ -links (pi-bonds).

By degree of bias common electron pairs to one of the atoms connected by them, a covalent bond can be polar and non-polar.

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

those. through a covalent non-polar bond, molecules of simple non-metallic substances are formed. A covalent chemical bond between atoms of elements whose electronegativities differ is called polar.

Covalent bond length and energy.

Characteristic covalent bond properties - its length and energy. Link length Is the distance between the nuclei of atoms. The shorter its length, the stronger the chemical bond. However, a measure of bond strength is bond energy, which is determined by the amount of energy required to break the 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 $ 0.074, 0.198 $, and $ 0.109 $ nm, respectively, and the binding energies are $ 436, 242 $, and $ 946 $ kJ / mol, respectively.

Jonah. Ionic bond

Let's imagine that two atoms "meet": a metal atom of group I and a non-metal atom of group VII. The metal atom has a single electron on the external energy level, and the non-metal atom just lacks just one electron for its external level to be complete.

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

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

The chemical bond that occurs between ions is called ionic.

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

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

This transformation of atoms into ions always occurs when the atoms of typical metals and typical non-metals interact.

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

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

Metal bond

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

Atoms of most metals on the outer level contain a small number of electrons - $ 1, 2, 3 $. These electrons are easily torn off, and the atoms are converted into positive ions. Detached electrons move from one ion to another, linking them into a single whole. Combining with ions, these electrons temporarily form atoms, then break off again and combine with another ion, etc. Consequently, in the bulk of the metal, atoms are continuously transformed into ions and vice versa.

The bond in metals between ions by means of 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.

The metallic bond has some similarities with the covalent bond, since it is based on the sharing of external electrons. However, with a covalent bond, the external unpaired electrons of only two neighboring atoms are socialized, while with a metal bond, all atoms take part in the socialization of these electrons. That is why crystals with a covalent bond are fragile, and crystals with a metal bond are usually ductile, electrically conductive and have a metallic luster.

The metallic bond is characteristic both for pure metals and for mixtures of various metals - alloys in solid and liquid states.

Hydrogen bond

The chemical bond between positively polarized hydrogen atoms of one molecule (or part of it) and negatively polarized atoms of strongly electronegative elements having lone electron pairs ($ F, O, N $ and less often $ S $ and $ Cl $), another molecule (or its parts) are called hydrogen.

The mechanism of hydrogen bonding is partly electrostatic and partly donor-acceptor.

Examples of intermolecular hydrogen bonds:

In the presence of such a bond, 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. Crystal lattice type. The 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. A substance under given conditions 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 its constituent particles - molecules, atoms or ions. By the type of bond, substances of molecular and non-molecular structure are distinguished.

Substances consisting of molecules are called molecular substances... The bonds between 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 composed 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.

Consider the physical properties of alkali metals. The relatively low bond strength between atoms causes low mechanical strength: alkali metals are soft, easily cut with a knife.

The large sizes of atoms lead to a low density 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 an increase in the ordinal number of the element, because the sizes of the atoms increase and the bonds weaken.

To substances non-molecular structures include ionic compounds. Most metal compounds with non-metals have such a 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

A substance, as you know, can exist in three states of aggregation: gaseous, liquid and solid.

Solids: amorphous and crystalline.

Consider how the features of chemical bonds affect the properties of solids. Solids are divided into crystallineand amorphous.

Amorphous substances do not have a clear melting point - when heated, they gradually soften and turn into a fluid state. In the amorphous state, for example, are plasticine and various resins.

Crystalline substances are characterized by the correct arrangement of those 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 where the crystal particles are located are called lattice points.

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

Ionic crystal lattices.

Ionic called crystal lattices, in the nodes of which there are ions. They are formed by substances with an ionic bond, which can be associated with both simple ions $ Na ^ (+), Cl ^ (-) $, and complex ions $ SO_4 ^ (2−), OH ^ - $. Consequently, salts, some oxides and hydroxides of metals have ionic crystal lattices. For example, a sodium chloride crystal is composed 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 a 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, atoms are linked together by very strong covalent bonds. An example of substances with this type of crystal lattice is diamond - one of the allotropic modifications of carbon.

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

Molecular crystal lattices.

Molecular are called crystal lattices, in the nodes of which molecules are located. Chemical bonds in these molecules can be both polar ($ HCl, H_2O $) and non-polar ($ N_2, O_2 $). Despite the fact that the atoms inside the molecules are bound by very strong covalent bonds, weak forces of intermolecular 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 a metallic bond have metallic crystal lattices. At the sites of such lattices are atoms and ions (either atoms or ions, into which metal atoms are easily transformed, donating their outer electrons "for general use"). This internal structure of metals determines their characteristic physical properties: ductility, plasticity, electrical and thermal conductivity, characteristic metallic luster.