Molecules of what substances have a p bond. Chemical bond and structure of molecules. Covalent polar chemical bond

C 2s 2 2p 2 C + 1e \u003d C -

O 2s 2 2p 4 O -1e \u003d O +

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 an 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 that can accept this pair of electrons.

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

Like N 2 - CO, it has a high dissociation energy (1069 kJ), is poorly soluble in water, and is 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 an increase in temperature or irradiation, it exhibits the properties of a reducing agent.

Receipt:

in industry

CO 2 + C " 2CO

2C + O 2 ® 2CO

in the laboratory: H 2 SO 4, t

HCOOH ® CO + H 2 O;

H2SO4t

H 2 C 2 O 4 ® CO + CO 2 + H 2 O.

CO reacts 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 the oxides of many heavy metals (Fe, Co, Pb, etc.).

CO + Cl 2 = COCl 2 (phosgene)

CO + NH 3 ® HCN + H 2 O H-CºN

CO + H 2 O " CO 2 + H 2

CO + S ® COS

Of greatest interest are metal carbonyls (used to obtain pure metals). Chemical bonding by the donor-acceptor mechanism, there is p-overlapping by the dative mechanism.

5CO + Fe ® (iron pentacarbonyl)

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

→ 4CO + Ni (nickel carbonyl).

Like CO, metal carbonyls are toxic.

Chemical bond in the CO 2 molecule

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

O ═ S ═ O

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

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; equilibria are established:

K 1 \u003d 4 × 10 -7 K 2 \u003d 4.8 × 10 -11 is a weak, unstable, oxygen-containing, dibasic acid. Hydrocarbonates 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. Acid salts of carbonic acid are obtained by passing an excess of CO 2 into an aqueous solution of carbonate:

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

Na 2 CO 3 + HNO 3 ® NaHCO 3 + NaNO 3

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

Salts are hydrolyzed according to the equation:

I stage

Due to complete hydrolysis, carbonates Gr 3+ , Al 3+ , Ti 4+ , ​​Zr 4+ and others 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 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.

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

With metal sulfides, CS 2 forms salts of thiocarbonic acid - thiocarbonates.

This reaction is similar to the process

Thiocarbonates- yellow crystalline solids. Under the action of acids on them, free thiocarbonic acid is released.

It is more stable than H 2 CO 3 and at low temperature it is released from solution in the form of a yellow oily liquid, easily decomposing into:

Compounds of carbon with nitrogen (CN) 2 or C 2 N 2 - dicyan, highly flammable colorless gas. Pure dry cyanide 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, like halogens, it disproportionates:

(C N) 2 + 2NaOH = 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 poisonous). Get:

Hydrogen cyanide is obtained in industry by catalytic reactions.

2CH 4 + 3O 2 + 2NH 3 ® 2HCN + 6H 2 O.

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

Cyanide handling requires strict precautions. In agriculture, they are used to combat especially dangerous insect pests.

Cyanides get:

Compounds of carbon with a negative oxidation state:

1) covalent (SiC carborundum) ;

2) ionic-covalent;

3) metal carbides.

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

methanides(CH 4 stands out)

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

acetylenides(secured C 2 H 2)

H 2 C 2 + AgNO 3 ® Ag 2 C 2 + HNO 3

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

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 easily oxidized by oxygen to form the more stable oxide SiO 2 .

2SiO + O 2 t ® 2SiO 2

SiO2- silica, has several crystalline modifications. Low-temperature - quartz, has piezoelectric properties. Natural varieties of quartz: rock crystal, topaz, amethyst. Varieties of silica - chalcedony, opal, agate, sand.

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

Combinations of tetrahedra can be connected into chains, ribbons, nets and frameworks.

Important natural silicates 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 finely crystalline state - glass-ceramics - materials of increased strength. In nature, aluminosilicates are common - frame orthosilicates, some of the Si atoms are replaced by Al, for example, Na 12 [(Si,Al)O 4 ] 12.

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

Hexafluorosilicic acid (similar 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 as 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 existing only in solution)

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

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

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

Silicates - 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". Widely used as a silicate adhesive and as a wood preservative.

Fusion of Na 2 CO 3 , CaCO 3 and SiO 2 produces glass, which is a supercooled mutual solution of salts of polysilicic acids.

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

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

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

One of the most important issues in chemistry is the question of the chemical bond, which requires an explanation of the causes and identification of patterns of formation of bonds between atoms, ions, molecules based on the theory of the structure of the atom and the Periodic Law of D.I. Mendeleev, as well as the characteristics of these bonds through the interpretation of the physical and chemical properties of substances.

The formation of molecules, molecular ions, ions, crystalline, amorphous and other substances from atoms is accompanied by a decrease in energy compared to non-interacting atoms. In this case, the minimum energy corresponds to a certain arrangement of atoms relative to each other, which corresponds to a significant redistribution of the electron density. The forces that hold atoms in new formations have received the generalized name "chemical bond". The most important types of chemical bonds are: ionic, covalent, metallic, hydrogen, intermolecular.

When characterizing a chemical bond, such concepts as "valency", "oxidation state" and "bond multiplicity" are usually used.
Valence- the ability of an atom of a chemical element to form bonds with other atoms. For ionic compounds, the number of donated or received electrons is taken as the value of valency. For covalent compounds, valence is equal to the number of socialized electron pairs.

Depending on the method of redistribution of electrons, they emit covalent bonds, ionic and metal . According to the presence or absence of polarization, covalent bonds are divided into: polar between atoms of different elements, and non-polar between atoms of the same element. According to the method of formation, covalent bonds are divided into ordinary , donor-acceptor and dative.

According to the electronic theory of valence, a chemical bond arises due to the redistribution of electrons in valence orbitals, resulting in a stable electronic configuration of a noble gas (octet) due to the formation of ions (W. Kossel) or the formation of common electron pairs (G. Lewis). Quantum-mechanical theories (the theory of valence bonds and the method of molecular orbitals) are based on the concept of the wave function ψ, which describes the state of electrons in a molecule, based on approximate solutions of the Schrödinger equation. For the first time, such an approximate calculation was carried out by W. Heitler and F. London for the hydrogen molecule.

The energy of a system consisting of two hydrogen atoms a - the spins are parallel; b - spins are antiparallel; E is the energy of the system, r 0 is the internuclear distance in the molecule

As a result, equations were obtained that make it possible to find the dependence of the potential energy of the system E, consisting of two hydrogen atoms, on the distance r between the nuclei of these atoms. It turned out that the results of the calculation depend on whether the spins of the interacting electrons are the same or opposite in sign. With the same direction of spins (curve a), the approach of atoms leads to a continuous increase in the energy of the system. In this case, the approach of atoms requires energy expenditure, so that such a process is energetically unfavorable and no chemical bond between atoms occurs.

With oppositely directed spins (curve b), the approach of atoms to a certain distance r0 accompanied by a decrease in the energy of the system. At r = r0 the system has the lowest potential energy, i.e. is in the most stable state; further approach of the atoms again leads to an increase in energy. But this also means that in the case of oppositely directed spins of atomic electrons, an H 2 molecule is formed - a stable system of two hydrogen atoms located at a certain distance from each other.

The chemical bond is characterized energy and long . The measure of bond strength is the energy expended to break the bond, or the gain in energy during the formation of a compound from individual atoms ( E St.). Energy of chemical bonds - is the energy required to break chemical bonds. In this case, atoms, radicals, ions or excited molecules are formed from the molecule.

For example:

H 2 H + H, E sv \u003d 432 kJ / mol,

H 2 O H + OH E sv \u003d 461 kJ / mol,

NaCl (tv) Na + (g) + Cl - (g) E sv \u003d 788.3 kJ / mol,

C 2 H 6 ?H 3 + ?H 3, E St = 356 kJ/mol.

The bond energy, as can be seen, depends on the products that are obtained as a result of its rupture. On the basis of such data, the concept of ordinary (single), double, triple, and, in general, multiple bonds has been introduced.

Link length(nm, ?)- the distance between the nuclei of neighboring atoms in the molecule. It can be determined experimentally by modern physical methods (electronographic, radiographic, infrared introscopy, etc.). The bond length is approximately equal to the sum of the radii of neighboring atoms d A - B = r A + r B .

Like the radii of atoms, internuclear distances naturally change in series, subgroups of the Periodic system. For example, in the series HF - HCl - HBr - HI, the distance d H-G increases (1.0; 1.27; 1.41 and 1.62 ? , respectively). The distance between the same atoms in different compounds (with the same multiplicity) are close. So, ordinary C-C bonds in any compounds are d C-C from 1.54 to 1.58?. The higher the multiplicity of the bond, the shorter its length:

d C - C \u003d 1.54, d C \u003d C \u003d 1.34 and d C ≡ C \u003d 1.2?

the greater the bond energy, the shorter its length.

In compounds containing more than two atoms, an important characteristic is the bond angle formed by chemical bonds in the molecule and reflecting its geometry. They depend on the nature of the atoms (their electronic structure) and the nature of the chemical bond (covalent, ionic, hydrogen, metallic, ordinary, multiple). Bond angles are currently determined very accurately by the same methods as bond lengths.

For example, it has been shown that AB 2 molecules can be linear (CO 2) or angular (H 2 O), AB 3 - triangular (BF 3) and pyramidal (NH 3), AB 4 - tetrahedral (CH 4), or square (PtCl 4) -, or pyramidal (SbCl 4) -, AB 5 - trigonal-bipyramidal (PCl 5), or tetragonal-pyramidal (BrF 5), AB 6 - octahedral (AlF 6) 3-, etc. Valence angles naturally change with a change in the serial number in the periodic table. For example, the H-E-H angle for H 2 O, H 2 S, H 2 Se decreases (104.5; 92 and 90 0, respectively).

The polarity of a molecule is determined by the difference in the electronegativity of the atoms forming a two-center bond, the geometry of the molecule, and also by the presence of unshared electron pairs, since part of the electron density in the molecule can be localized not in the direction of the bonds. The polarity of a bond is expressed through its ionic component, that is, through the displacement of an electron pair to a more electronegative atom. The polarity of a molecule is expressed in terms of its dipole moment, which is equal to the vector sum of all dipole moments of the bonds of the molecule.

A dipole is a system of two equal but opposite charges located at a unit distance from each other. Dipole moment is measured in coulomb meters (C?m) or debyes (D); 1D \u003d 0.333? 10 -29 Cm.

Knowing the magnitude of the dipole moment, one can draw a conclusion about the nature of the chemical bond (ionic, covalent polar or nonpolar) and the geometric shape of the molecule. You can focus on the value of the differences in the electronegativity of the elements that make up the binary molecule: if? ? 1.7, then the bond in this compound is covalently polar, but what if? ? 1,7 - ionic.

The bond between atoms with the same electronegativity, for example, H 2 , Cl 2 , or similar values ​​of electronegativity - CH 4 does not have even a small contribution associated with charge separation. Such bonds and molecules are called covalent; they are non-polar, in them the centers of gravity of the charges coincide. A covalent bond is the most general type of chemical bond that occurs due to the socialization of an electron pair through an exchange mechanism.

To form a simple covalent bond, each of the atoms provides one electron: A.|.B. When donor-acceptor bond one atom - donor - provides two electrons, and another atom - acceptor - allocates a vacant electron orbital for this: A : | B. A classic example of a non-polar covalent bond (the difference in electronegativity is zero) is observed in homonuclear molecules: H-H, F-F, O + O = O 2. When a heteroatomic covalent bond is formed, the electron pair is shifted to a more electronegative atom, which makes such a bond polar (HCl, H 2 O): S + O 2 = O=S=O.

Except polarizability covalent bond has the property satiety - the ability of an atom to form as many covalent bonds as it has energetically available atomic orbitals. Electronic orbitals (except s-orbitals) have a spatial orientation . Therefore, the covalent bond, which is the result of the overlapping of electron clouds of interacting atoms, is located in a certain direction with respect to these atoms.

If the overlap of electron clouds occurs in the direction of a straight line connecting the nuclei of interacting atoms (i.e., along the bond axis), then σ -bond (sigma bond). During the interaction of p-electron clouds directed perpendicular to the bond axis, two overlapping regions are formed, located on both sides of this axis. Such a covalent bond is called a π-bond (pi-bond). The π-bond can arise not only due to p-electrons, but also due to the overlap of d- and p-electron clouds or d-clouds. Delta (δ) - bonds are due to the overlap of all four lobes d - electron clouds located in parallel planes.

Possible types of chemical orbital overlap

Based on the symmetry conditions, it can be shown that the electrons of s-orbitals can participate only in σ - bonding, p-electrons - already in σ - and π - bonding, and d - electrons - both in σ - and π -, and in δ - binding. For f-orbitals, the types of symmetry are even more diverse.

In most molecules, the bonds are intermediate in nature (including NaCl); such bonds and molecules are called polar (or polar covalent), in which the “centers of gravity” of charges do not coincide. A covalent bond is the most common type of bond; it is realized in most known substances. There are few compounds with a non-polar covalent bond and a bond close to a purely ionic one.

If the interacting atoms differ in electronegativity, then the electron density shifts to a more electronegative one and the atoms, in the limit, turn into charged ions. In this case, between the atoms is formed ionic connection. For example, the bond in the NaCl molecule can be approximately represented as the Coulomb interaction of Na + and Cl - ions.

An ionic bond is a special case of a covalent bond, when the resulting electron pair belongs entirely to a more electronegative atom that becomes an anion. The basis for separating this bond into a separate type is the fact that compounds with such a bond can be described in the electrostatic approximation, considering the ionic bond due to the attraction of positive and negative ions. Interaction of ions of opposite sign does not depend on direction, and the Coulomb forces are not have the saturation property. Therefore, each ion in an ionic compound attracts such a number of ions of the opposite sign that an ionic-type crystal lattice is formed. There are no molecules in an ionic crystal. Each ion is surrounded by a certain number of ions of a different sign (coordination number of the ion). Ion pairs can exist in the gaseous state as polar molecules.

In the gaseous state, NaCl has a dipole moment of ~3?10 -29 C?m, which corresponds to a shift of 0.8 electron charge per bond length of 0.236 nm from Na to Cl, i.e. Na 0.8+ Cl 0.8-. Metal atoms usually donate electrons, while acquiring the electronic configuration of an atom of the previous inert gas. atoms d- and f-elements exhibiting variable valence may have other stable electronic configurations. Atoms of non-metals often complete their outer electron layer. If a more electronegative element is present in the compound, the non-metal can donate electrons until it reaches a stable oxidation state (for example, for Cl it is +1, +3, +5, +7). When a metal atom forms a bond with a non-metal atom, the former donates electrons and the latter accepts. In the case of the interaction of a typical metal with a typical non-metal, between their atoms is formed ionic bond : 2Na + Cl 2 \u003d 2NaCl.

At present, two methods are mainly used to study chemical bonds: 1) valence bonds; 2) molecular orbitals.

Within the framework of the first method, individual atoms that interact are considered based on the principle of completeness of the electron shell (the octet rule). A covalent bond from the point of view of the method of valence bonds is formed due to the socialization of an electron pair. The simple method of valence bonds is the most understandable, convenient and illustrative for a chemist. The disadvantage of the valence bond method is that some experimental data cannot be explained within its framework.

The valence bond method (MVS) is otherwise called the theory of localized electron pairs, since the method is based on the assumption that a chemical bond between two atoms is carried out using one or more electron pairs that are localized predominantly between them. In the MVS connection is always two-electron and must be bicentric. The number of elementary chemical bonds that an atom or ion is able to form is equal to its valency; valence electrons take part in the formation of a chemical bond. The wave function that describes the state of the electrons that form a bond is called a localized orbital (LO).

covalent bond. The structure of the water molecule

Task 61.
What chemical bond is called a covalent bond? How can one explain the direction of a covalent bond? How does the method of valence bonds (BC) explain the structure of the water molecule?
Decision:
Communication carried out due to the formation of electron pairs, equally belonging to both atoms is called covalent non-polar. Covalent bonds are oriented in space in a certain way, that is, they have a direction. The reason that molecules can have a linear planar or some other structure is that atoms use different orbitals and different numbers of them to form bonds. Molecules that have a dipole moment are not linear, while molecules that do not have a dipole moment are linear.

The water molecule H 2 O has a dipole moment, which means that it has a nonlinear structure. One oxygen atom and two hydrogen atoms participate in the formation of bonds between oxygen and hydrogen atoms. Oxygen is the neutral atom in the water molecule, and it has four electron pairs, two lone pairs and two shared ones, which are formed by one s-electron and one p-electron of oxygen. Such a molecule has a tetrahedral structure in the center of the tetrahedron there is an oxygen atom, and at the corners of the tetrahedron there are two hydrogen atoms and two lone electron pairs of oxygen. In such a molecule, the angle between bonds should be equal to 109.5 0 . If the water molecule were flat, then the angle HOH should be 90 0 . But X-ray diffraction analysis of water molecules shows that the HOH angle is 104.5 0 . This explains that the water molecule does not have a linear shape, but has the shape of a distorted tetrahedron. This is explained by the fact that the oxygen atom undergoes sp 3 hybridization, when one s-orbital and three p-orbitals of the oxygen atom hybridize, forming four equivalent sp 3 hybrid orbitals. Of the four sp 3 hybrid orbitals, two are occupied by the s orbitals of the hydrogen atom. The difference between the values ​​of the bond angle and the tetrahedral angle is explained by the fact that the repulsion between lone electron pairs is greater than between bonding ones.

Polar covalent bond

Task 62.
What covalent bond is called polar? What is the quantitative measure of the polarity of a covalent bond? Based on the electronegativity values ​​of the atoms of the corresponding elements, determine which of the bonds: HCl, ICl, BrF is the most polar.
Decision:
A covalent bond formed by different atoms is called a polar bond. For example, H - Cl; the center of gravity of a negative charge (associated with electrons) does not coincide with the center of gravity of a positive charge (associated with the charge of the atomic nucleus). The electron density of common electrons is shifted to one of the atoms, which has a higher electronegativity value, to a greater extent. In H:Cl, the shared electron pair is biased towards the most electronegative chlorine atom. The polarity of the bond is quantified by the dipole moment (), which is the product of the dipole length (l) - the distance between two equal and opposite charges +g and -g by the absolute value of the charge: = lg. The dipole moments HCI, HBr, HI are equal to 1.04, respectively; 0.79; 0.38 D. Dipole moments of molecules are usually measured in debyes (D)*: 1D = 3.33 . 10 -30 C . m.

The dipole moment is a vector quantity and is directed along the dipole axis from a negative charge to a positive one. The bond dipole moment provides valuable information about the behavior of the molecule as a whole. Along with the dipole moment, a characteristic called the electronegativity of the element (EO) is used to assess the degree of polarity of the bond. EO is the ability of an atom to attract the valence electrons of other atoms to itself. The values ​​of EO elements are given in special scales (tables).

The EO values ​​of hydrogen, chlorine, bromine, iodine, fluorine, respectively, are: 2.1; 3.0; 2.8; 2.5; 4.0. Based on the values ​​of the EO elements in the compounds

the most polar bond in the BrF molecule, since the difference in electronegativity between fluorine and bromine is the largest - 1.2 (4.0 - 2.8 = 1.2) than that of HCl and ICl.

Donor-acceptor bond

Task 63.
What method of covalent bond formation is called donor-acceptor? What chemical bonds are present in NH 4+ and BF 4- ions? Specify the donor and acceptor.
Decision:

A donor-acceptor bond is a covalent bond in which only one of the atoms participating in the bond provides a shared pair of electrons. In this case, one of the atoms is a donor - a supplier of an electron pair, and the other is an acceptor - a supplier of a free quantum orbital.

The ammonium cation NH 4+ is formed by the donor-acceptor mechanism:

It has the shape of a regular tetrahedron:

In the ammonium ion, each hydrogen atom is bonded to the nitrogen atom by a common electron pair, one of which is realized by the donor-acceptor mechanism. It is important to note that the H - N bonds formed by various mechanisms do not have any differences, i.e. they are all equivalent. The donor is a nitrogen atom, and the acceptor is a hydrogen atom.

The BF 4- ion is formed from BF 3 and the F- ion. This ion is formed due to the fact that the unshared electron pair of the F- ion is “embedded” in the valence shell of the boron atom of the covalently bound BF 3 molecule:

In the BF 4 ion, the fluorine ion is the donor, and the boron atom of the BF 3 molecule is the acceptor.

The donor-acceptor bond in the structural formulas is depicted by an arrow that is directed from the donor to the acceptor.

Valence bond method (BC)

Task 64.
How does the method of valence bonds (BC) explain the linear structure of the BeCl 2 molecule and the tetrahedral CH 4?
Decision

a) Representations of the method of valence bonds make it possible to explain the geometry of many molecules. Thus, the BeCl2 molecule consists of one beryllium atom and two chlorine atoms. An excited beryllium atom has one s-electron and one p-electron. When BeCl 2 is formed, two covalent bonds appear. One of them should be an s - p bond formed due to the overlap of the s-cloud of the beryllium atom and the p-cloud of the chlorine atom, the other (p - p bond) due to the overlap of the p-cloud of the beryllium atom and the p-cloud of the chlorine atom.

p - p bond and s - p can be located at an angle relative to each other, i.e. the BeCl 2 molecule must be angular, but it is precisely established that the BeCl 2 molecule has a linear structure, and both bonds are equal in energy and length. The concept of hybridization of atomic orbitals is used to explain the geometry of the BeCl 2 molecule. The essence of the concept of atomic orbitals is that atomic orbitals can be geometrically modified and mixed with each other in such a way as to provide the greatest overlap with the orbits of other atoms and, therefore, the greatest gain in energy. This is achieved if, instead of orbitals having different shapes and energies, hybrid orbitals of the same shape and energy appear, which are linear combinations of the original atomic orbitals. So in the Be atom, the s-orbital and p-orbital interact, their energies are aligned and two identical sp-hybrid orbitals are formed. The two generated sp-hybrid electron clouds have the same energy and an asymmetric shape, which provides more overlap with p-electron clouds of the chlorine atom than overlap with pure unhybridized s- and p-clouds. Two hybrid sp-clouds are located relative to each other and the atomic nucleus at an angle of 180 0:

Rice. 1. Triatomic molecule BeCl 2

As a result of this arrangement of hybrid clouds, the BeCl 2 molecule has a linear structure.

b) The CH 4 molecule consists of one carbon atom and four hydrogen atoms, between which there are four covalent bonds. An excited carbon atom has four unpaired electrons, one in the s orbital and three in the p orbitals:

Filling the external energy level of the carbon atom in the ground state:

Filling the external energy level of the carbon atom in an excited state:

Of the four bonds in the CH 4 molecule, there should be one s - s and three s - p bonds formed due to the overlapping of the orbitals of the carbon atom with the s-orbital of hydrogen atoms. As a result of this overlap, an s - s bond should be formed, different from the three s - p bonds of length and energy, and located at an angle of about 125 0 to any of them. However, it is precisely established that the CH 4 molecule has the shape of a tetrahedron with an angle between bonds of 109.5 0, and all bonds are equivalent in length and energy. The tetrahedral structure of the CH 4 molecule can be explained by sp 3 hybridization. The carbon atom contains four sp 3 hybrid orbitals, resulting from a linear combination of an s orbital and three p orbitals. Four sp3-hybrid orbitals are located relative to each other at an angle of 109.5 0 . They are directed to the vertices of the tetrahedron, in the center of which is the nucleus of the carbon atom (Fig. 2.).

Rice. 2. Scheme of the structure of the CH 4 molecule;
Methane, there are no non-bonding electron pairs.

Thus, four equivalent chemical bonds are formed in the CH4 molecule due to the overlapping of sp3-hybrid orbitals of the carbon atom with s-orbitals of carbon atoms.

Formation of sigma bond and pi bond

Task 65.
Which covalent bond is called a -bond and which -bond? Explain the structure of the nitrogen molecule as an example.
Decision:
A bond formed by overlapping along a line connecting two atoms is called -bond (any simple bond) or “If the overlap of atomic orbitals occurs on the internuclear axis, then a sigma bond is formed (-connection). A sigma bond is formed by overlapping two s-orbitals (s-s bond), one s- and one p-orbital (s-p bond), two p-orbitals (p-p bond), one s- and one d -orbital (s - d bond), one p- and one d-orbital (p - d bond).

Options for overlapping atomic orbitals leading to the formation

162774 0

Each atom has a certain number of electrons.

Entering into chemical reactions, atoms donate, acquire, or socialize electrons, reaching the most stable electronic configuration. The configuration with the lowest energy is the most stable (as in noble gas atoms). This pattern is called the "octet rule" (Fig. 1).

Rice. one.

This rule applies to all connection types. Electronic bonds between atoms allow them to form stable structures, from the simplest crystals to complex biomolecules that eventually form living systems. They differ from crystals in their continuous metabolism. However, many chemical reactions proceed according to the mechanisms electronic transfer, which play an important role in the energy processes in the body.

A chemical bond is a force that holds together two or more atoms, ions, molecules, or any combination of them..

The nature of the chemical bond is universal: it is an electrostatic force of attraction between negatively charged electrons and positively charged nuclei, determined by the configuration of the electrons in the outer shell of atoms. The ability of an atom to form chemical bonds is called valence, or oxidation state. The concept of valence electrons- electrons that form chemical bonds, that is, those located in the most high-energy orbitals. Accordingly, the outer shell of an atom containing these orbitals is called valence shell. At present, it is not enough to indicate the presence of a chemical bond, but it is necessary to clarify its type: ionic, covalent, dipole-dipole, metallic.

The first type of connection isionic connection

According to Lewis and Kossel's electronic theory of valency, atoms can achieve a stable electronic configuration in two ways: first, by losing electrons, becoming cations, secondly, acquiring them, turning into anions. As a result of electron transfer, due to the electrostatic force of attraction between ions with charges of the opposite sign, a chemical bond is formed, called Kossel " electrovalent(now called ionic).

In this case, anions and cations form a stable electronic configuration with a filled outer electron shell. Typical ionic bonds are formed from cations of T and II groups of the periodic system and anions of non-metallic elements of groups VI and VII (16 and 17 subgroups - respectively, chalcogens and halogens). The bonds in ionic compounds are unsaturated and non-directional, so they retain the possibility of electrostatic interaction with other ions. On fig. 2 and 3 show examples of ionic bonds corresponding to the Kossel electron transfer model.

Rice. 2.

Rice. 3. Ionic bond in the sodium chloride (NaCl) molecule

Here it is appropriate to recall some of the properties that explain the behavior of substances in nature, in particular, to consider the concept of acids and grounds.

Aqueous solutions of all these substances are electrolytes. They change color in different ways. indicators. The mechanism of action of indicators was discovered by F.V. Ostwald. He showed that the indicators are weak acids or bases, the color of which in the undissociated and dissociated states is different.

Bases can neutralize acids. Not all bases are soluble in water (for example, some organic compounds that do not contain -OH groups are insoluble, in particular, triethylamine N (C 2 H 5) 3); soluble bases are called alkalis.

Aqueous solutions of acids enter into characteristic reactions:

a) with metal oxides - with the formation of salt and water;

b) with metals - with the formation of salt and hydrogen;

c) with carbonates - with the formation of salt, CO 2 and H 2 O.

The properties of acids and bases are described by several theories. In accordance with the theory of S.A. Arrhenius, an acid is a substance that dissociates to form ions H+ , while the base forms ions IS HE- . This theory does not take into account the existence of organic bases that do not have hydroxyl groups.

In line with proton Bronsted and Lowry's theory, an acid is a substance containing molecules or ions that donate protons ( donors protons), and the base is a substance consisting of molecules or ions that accept protons ( acceptors protons). Note that in aqueous solutions, hydrogen ions exist in a hydrated form, that is, in the form of hydronium ions H3O+ . This theory describes reactions not only with water and hydroxide ions, but also carried out in the absence of a solvent or with a non-aqueous solvent.

For example, in the reaction between ammonia NH 3 (weak base) and hydrogen chloride in the gas phase, solid ammonium chloride is formed, and in an equilibrium mixture of two substances there are always 4 particles, two of which are acids, and the other two are bases:

This equilibrium mixture consists of two conjugated pairs of acids and bases:

1)NH 4+ and NH 3

2) HCl and Cl ‑

Here, in each conjugated pair, the acid and base differ by one proton. Every acid has a conjugate base. A strong acid has a weak conjugate base, and a weak acid has a strong conjugate base.

The Bronsted-Lowry theory makes it possible to explain the unique role of water for the life of the biosphere. Water, depending on the substance interacting with it, can exhibit the properties of either an acid or a base. For example, in reactions with aqueous solutions of acetic acid, water is a base, and with aqueous solutions of ammonia, it is an acid.

1) CH 3 COOH + H 2 O ↔ H 3 O + + CH 3 SOO- . Here the acetic acid molecule donates a proton to the water molecule;

2) NH3 + H 2 O ↔ NH4 + + IS HE- . Here the ammonia molecule accepts a proton from the water molecule.

Thus, water can form two conjugated pairs:

1) H 2 O(acid) and IS HE- (conjugate base)

2) H 3 O+ (acid) and H 2 O(conjugate base).

In the first case, water donates a proton, and in the second, it accepts it.

Such a property is called amphiprotonity. Substances that can react as both acids and bases are called amphoteric. Such substances are often found in nature. For example, amino acids can form salts with both acids and bases. Therefore, peptides readily form coordination compounds with the metal ions present.

Thus, the characteristic property of an ionic bond is the complete displacement of a bunch of binding electrons to one of the nuclei. This means that there is a region between the ions where the electron density is almost zero.

The second type of connection iscovalent connection

Atoms can form stable electronic configurations by sharing electrons.

Such a bond is formed when a pair of electrons is shared one at a time. from each atom. In this case, the socialized bond electrons are distributed equally among the atoms. An example of a covalent bond is homonuclear diatomic H molecules 2 , N 2 , F 2. Allotropes have the same type of bond. O 2 and ozone O 3 and for a polyatomic molecule S 8 and also heteronuclear molecules hydrogen chloride HCl, carbon dioxide CO 2, methane CH 4, ethanol With 2 H 5 IS HE, sulfur hexafluoride SF 6, acetylene With 2 H 2. All these molecules have the same common electrons, and their bonds are saturated and directed in the same way (Fig. 4).

For biologists, it is important that the covalent radii of atoms in double and triple bonds are reduced compared to a single bond.

Rice. 4. Covalent bond in the Cl 2 molecule.

Ionic and covalent types of bonds are two limiting cases of many existing types of chemical bonds, and in practice most of the bonds are intermediate.

Compounds of two elements located at opposite ends of the same or different periods of the Mendeleev system predominantly form ionic bonds. As the elements approach each other within a period, the ionic nature of their compounds decreases, while the covalent character increases. For example, the halides and oxides of the elements on the left side of the periodic table form predominantly ionic bonds ( NaCl, AgBr, BaSO 4 , CaCO 3 , KNO 3 , CaO, NaOH), and the same compounds of the elements on the right side of the table are covalent ( H 2 O, CO 2, NH 3, NO 2, CH 4, phenol C6H5OH, glucose C 6 H 12 O 6, ethanol C 2 H 5 OH).

The covalent bond, in turn, has another modification.

In polyatomic ions and in complex biological molecules, both electrons can only come from one atom. It is called donor electron pair. An atom that socializes this pair of electrons with a donor is called acceptor electron pair. This type of covalent bond is called coordination (donor-acceptor, ordative) communication(Fig. 5). This type of bond is most important for biology and medicine, since the chemistry of the most important d-elements for metabolism is largely described by coordination bonds.

Pic. 5.

As a rule, in a complex compound, a metal atom acts as an electron pair acceptor; on the contrary, in ionic and covalent bonds, the metal atom is an electron donor.

The essence of the covalent bond and its variety - the coordination bond - can be clarified with the help of another theory of acids and bases, proposed by GN. Lewis. He somewhat expanded the semantic concept of the terms "acid" and "base" according to the Bronsted-Lowry theory. The Lewis theory explains the nature of the formation of complex ions and the participation of substances in nucleophilic substitution reactions, that is, in the formation of CS.

According to Lewis, an acid is a substance capable of forming a covalent bond by accepting an electron pair from a base. A Lewis base is a substance that has a lone pair of electrons, which, by donating electrons, forms a covalent bond with Lewis acid.

That is, the Lewis theory expands the range of acid-base reactions also to reactions in which protons do not participate at all. Moreover, the proton itself, according to this theory, is also an acid, since it is able to accept an electron pair.

Therefore, according to this theory, cations are Lewis acids and anions are Lewis bases. The following reactions are examples:

It was noted above that the subdivision of substances into ionic and covalent ones is relative, since there is no complete transfer of an electron from metal atoms to acceptor atoms in covalent molecules. In compounds with an ionic bond, each ion is in the electric field of ions of the opposite sign, so they are mutually polarized, and their shells are deformed.

Polarizability determined by the electronic structure, charge and size of the ion; it is higher for anions than for cations. The highest polarizability among cations is for cations of larger charge and smaller size, for example, for Hg 2+ , Cd 2+ , Pb 2+ , Al 3+ , Tl 3+. Has a strong polarizing effect H+ . Since the effect of ion polarization is two-sided, it significantly changes the properties of the compounds they form.

The third type of connection -dipole-dipole connection

In addition to the listed types of communication, there are also dipole-dipole intermolecular interactions, also known as van der Waals .

The strength of these interactions depends on the nature of the molecules.

There are three types of interactions: permanent dipole - permanent dipole ( dipole-dipole attraction); permanent dipole - induced dipole ( induction attraction); instantaneous dipole - induced dipole ( dispersion attraction, or London forces; rice. 6).

Rice. 6.

Only molecules with polar covalent bonds have a dipole-dipole moment ( HCl, NH 3, SO 2, H 2 O, C 6 H 5 Cl), and the bond strength is 1-2 debye(1D \u003d 3.338 × 10 -30 coulomb meters - C × m).

In biochemistry, another type of bond is distinguished - hydrogen connection, which is a limiting case dipole-dipole attraction. This bond is formed by the attraction between a hydrogen atom and a small electronegative atom, most often oxygen, fluorine and nitrogen. With large atoms that have a similar electronegativity (for example, with chlorine and sulfur), the hydrogen bond is much weaker. The hydrogen atom is distinguished by one essential feature: when the binding electrons are pulled away, its nucleus - the proton - is exposed and ceases to be screened by electrons.

Therefore, the atom turns into a large dipole.

A hydrogen bond, unlike a van der Waals bond, is formed not only during intermolecular interactions, but also within one molecule - intramolecular hydrogen bond. Hydrogen bonds play an important role in biochemistry, for example, for stabilizing the structure of proteins in the form of an α-helix, or for the formation of a DNA double helix (Fig. 7).

Fig.7.

Hydrogen and van der Waals bonds are much weaker than ionic, covalent, and coordination bonds. The energy of intermolecular bonds is indicated in Table. one.

Table 1. Energy of intermolecular forces

Note: The degree of intermolecular interactions reflect the enthalpy of melting and evaporation (boiling). Ionic compounds require much more energy to separate ions than to separate molecules. The melting enthalpies of ionic compounds are much higher than those of molecular compounds.

The fourth type of connection -metallic bond

Finally, there is another type of intermolecular bonds - metal: connection of positive ions of the lattice of metals with free electrons. This type of connection does not occur in biological objects.

From a brief review of the types of bonds, one detail emerges: an important parameter of an atom or ion of a metal - an electron donor, as well as an atom - an electron acceptor is its the size.

Without going into details, we note that the covalent radii of atoms, the ionic radii of metals, and the van der Waals radii of interacting molecules increase as their atomic number in the groups of the periodic system increases. In this case, the values ​​of the ion radii are the smallest, and the van der Waals radii are the largest. As a rule, when moving down the group, the radii of all elements increase, both covalent and van der Waals.

The most important for biologists and physicians are coordination(donor-acceptor) bonds considered by coordination chemistry.

Medical bioinorganics. G.K. Barashkov

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

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

Chemical bonding is inherently one: it is of electrostatic origin. But in various chemical compounds, the chemical bond is of various types; The most important types of chemical bonds are covalent (non-polar, polar), ionic, and metallic. Varieties of these types of bonds are donor-acceptor, hydrogen, etc. A metallic bond arises between metal atoms.

A chemical bond carried out by the formation of a common, or shared, pair or several pairs of electrons is called covalent. In the formation of one common pair of electrons, each atom contributes one electron, i.e. participates "in equal shares" (Lewis, 1916). Below are schemes for the formation of chemical bonds in H2, F2, NH3, and CH4 molecules. Electrons belonging to different atoms are designated by different symbols.

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

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

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

Fundamentals of the AIM

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

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

When writing structural formulas, the electron pairs that cause the bond are often depicted as dashes (instead of dots representing socialized electrons).

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

Valence is the property of an atom of a chemical element to attach or replace a certain number of atoms of another element. From this point of view, the valency of an atom is easiest to determine by the number of hydrogen atoms that form chemical bonds with it, or by the number of hydrogen atoms that are replaced by an atom of this element.

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

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

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

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

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

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

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

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

A chemical bond carried out by one electron pair is called a single bond. A single bond is always a σ-bond. Orbitals of type s can only form σ bonds.

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

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

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

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

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

Between different atoms, a pure covalent bond can manifest itself if the electronegativity (EO) of some molecules is electrosymmetric, i.e. The "centers of gravity" of the positive charges of the nuclei and the negative charges of the electrons coincide at one point, therefore they are called non-polar.

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

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

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

Dipole moments are measured in debyes D (D \u003d 10-18 el. st. units × cm, since the elementary charge is 4.810-10 el. st. units, and the length of the dipole is on average equal to the distance between two atomic nuclei, i.e. 10-8 cm) or coulometers (C × m) (1 D = 3.33 10-30 C × m) (electron charge 1.6 10-19 C multiplied by the distance between charges, for example, 0.1 nm, then Mp = 1.6 10-19 × 1 × 10-10 = 1.6 10-29 C m). Permanent dipole moments of molecules have values ​​from zero to 10 D.

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

At a very large difference in EC, the atoms have a clear unilateral polarization: the electron cloud of the bond shifts as much as possible towards the atom with the highest EC, the atoms pass into oppositely charged ions, and an ionic molecule arises:

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

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

The polarity of a bond can be predicted using the relative EO of atoms. The greater the difference in the relative EO of atoms, the stronger the polarity is expressed: DEO = 0 - non-polar covalent bond; DEO \u003d 0 - 2 - polar covalent bond; DEO \u003d 2 - ionic bond. It is more correct to speak about the degree of ionicity of a bond, since bonds are not 100% ionic. Even in the CsF compound, the bond is only 89% ionic.

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

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

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

Mechanisms of chemical bond formation

In the method of valence bonds, exchange and donor-acceptor mechanisms of formation of a chemical bond are distinguished.

exchange mechanism. The exchange mechanism for the formation of a chemical bond includes cases where one electron is involved in the formation of an electron pair from each atom.

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

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

From the quantum cells N and H, it can be seen that the nitrogen atom has 3

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

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

Donor-acceptor mechanism

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

NH3 + BF3 = NH3BF3 + 171.4 kJ/mol.

Consider the mechanism of this reaction:

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

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

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

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

Orientation of Bonds and Hybridization of Atomic Orbitals

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

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

When an ionic bond is formed, the electric field of the ion has spherical symmetry and therefore the ionic bond does not have directionality and saturation.

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

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

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

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

Hybridization of atomic orbitals

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

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

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

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

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

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

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

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

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

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

The hybridization of atomic orbitals (AO) leads to a shift of the electron cloud in the direction of bond formation with other atoms. As a result, the overlapping regions of hybrid orbitals turn out to be larger than for pure orbitals, and the bond strength increases.

Polarizability and polarizing effect of ions and molecules

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

The polarizability of anions, as a rule, is much higher than that of cations.

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

For ions of electronic analogs, the polarizability increases with an increase in the number of electronic layers, for example: or .

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

The electric field can be created both by a charged electrode and by an ion. Thus, the ion itself can exert a polarizing effect (polarization) on other ions or molecules. The polarizing effect of an ion increases with an increase in its charge and a decrease in its radius.

The polarizing effect of anions is, as a rule, much less than that of cations. This is due to the large size of anions compared to cations.

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

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

hydrogen bond

The hydrogen bond is a special type of chemical bond. It is known that hydrogen compounds with strongly electronegative non-metals, such as F, O, N, have abnormally high boiling points. If in the series Н2Тe – H2Se – H2S the boiling point naturally decreases, then in the transition from H2S to Н2О there is a sharp jump to an increase in this temperature. The same picture is observed in the series of hydrohalic acids. This indicates the presence of a specific interaction between H2O molecules and HF molecules. Such an interaction should hinder the separation of molecules from each other, i.e. reduce their volatility, and, consequently, increase the boiling point of the corresponding substances. Due to the large difference in ER, the H–F, H–O, and H–N chemical bonds are strongly polarized. Therefore, the hydrogen atom has a positive effective charge (δ+), and the F, O and N atoms have an excess of electron density, and they are negatively charged (d-). Due to Coulomb attraction, a positively charged hydrogen atom of one molecule interacts with an electronegative atom of another molecule. Due to this, the molecules are attracted to each other (bold dots indicate hydrogen bonds).

A hydrogen bond is such a bond that is formed by a hydrogen atom that is part of one of the two connected particles (molecules or ions). The energy of a hydrogen bond (21–29 kJ/mol or 5–7 kcal/mol) is approximately 10 times less than the energy of an ordinary chemical bond. Nevertheless, the hydrogen bond causes the existence of dimeric molecules (Н2О)2, (HF)2 and formic acid in pairs.

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

Intermolecular interactions

Previously, the bonds that cause the formation of molecules from atoms were considered. However, interactions also exist between molecules. It is the cause of the condensation of gases and their transformation into liquid and solid bodies. The first formulation of the forces of intermolecular interaction was given in 1871 by Van der Waals. Therefore, they are called van der Waals forces. The forces of intermolecular interaction can be divided into orientation, induction and dispersion.

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

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

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

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

A polar molecule can increase the polarity of an adjacent molecule. In other words, under the influence of the dipole of one molecule, the dipole of another molecule can increase, and a nonpolar molecule can become polar:

b

The dipole moment that appears as a result of polarization by another molecule or ion is called the induced dipole moment, and the phenomenon itself is called induction. Thus, the orientational interaction must always be superimposed by the inductive interaction of molecules.

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

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

Thus, the smallest component among the forces acting between molecules is the dispersion interaction. Between molecules with low polarity or without polarity (CH4, H2, HI), the acting forces are mainly dispersive. The greater the intrinsic dipole moment of the molecules, the greater the orientation forces of interaction between them.

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

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

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

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

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

The method of molecular orbitals (MO) is currently considered to be the best method for the quantum mechanical interpretation of a chemical bond. However, it is much more complicated than the VS method and is not as clear as the latter.

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

MO types. When an electron of one atom, when approaching, falls into the sphere of action of another atom, the nature of the motion, and, consequently, the wave function of the electron, changes. In the resulting molecule, the wave functions, or orbitals of the electrons, are unknown. There are several ways to determine the type of MO from known ARs. Most often, MO is obtained by a linear combination of atomic orbitals (LCAO). The Pauli principle, Hund's rule, the principle of least energy are also valid for the MO method.

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

In the simplest graphical form, MO, like LCAO, can be obtained by adding or subtracting wave functions. Figure 2.2 shows the formation of binding and loosening MO from the original AO.

AOs can form MOs if the energies of the corresponding AOs are close in magnitude and the AOs have the same symmetry about the bond axis.

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

When the wave functions are added, then in the overlapping region the density of the electron cloud, proportional to ψ2, becomes greater, an excess negative charge is created between the nuclei of atoms, and the nuclei of atoms are attracted to it. MO, obtained by adding the wave functions of hydrogen atoms, is called binding.

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

Since the hydrogen s-orbitals form only a σ-bond, the resulting MOs are designated σcv and σr. MO formed by 1s-atomic orbitals are denoted σcv1s and σр1s.

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

MOs are also formed from 2p orbitals of the same symmetry: bonding and antibonding σ orbitals from 2p orbitals located along the x axis. They are designated σcv2r and σr2r. The binding and antibonding p orbitals are formed from the 2pz orbitals. They are designated respectively πsv2rz, πp2pz. Similarly, πsv2py and πр2py orbitals are formed.

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

Let's consider an ion.

It can be seen from the diagram that the only electron is located along σcv - MO. A stable compound is formed with a binding energy of 255 kJ/mol, a bond length of 0.106 nm. The molecular ion is paramagnetic. If we accept that the bond multiplicity, as in the VS method, is determined by the number of electron pairs, then the bond multiplicity in is equal to ½. The process of education can be written as follows:

This notation means that there is one electron per MO formed from 1s AO.

The ordinary hydrogen molecule already contains two electrons with opposite spins in the σcv1s orbital: . The bond energy in H2 is greater than in - 435 kJ/mol, and the bond length (0.074 nm) is shorter. The H2 molecule has a single bond, the molecule is diamagnetic.

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

The molecular ion (+He+ ® He+2[(sc1s)2(sp1s)1]) already has one electron in the σ 1s orbital. The bond energy in - 238 kJ / mol (compared to H2 is lowered), and the bond length (0.108 us) is increased. The bond multiplicity is ½ (the bond multiplicity is equal to half the difference in the number of electrons in the bonding and loosening orbitals).

Li2	(Be)2	B2	N2	O2	(Ne)2	CO	NO
p2px	–	–	–	–	–	­¯	–	–
πp2py, πp2pz	– –	– –	– –	– –	­,­	­¯,­¯	– –	­,–
σcv2px	–	–	–	­¯	­¯	­¯	­¯	­¯
πcv2py, πcv2pz	– –	– –	­,­	­¯,­¯	­¯,­¯	­¯,­¯	­¯,­¯	­¯,­¯
σр2s	–	­¯	­¯	­¯	­¯	­¯	­¯	­¯
σcv2s	­¯	­¯	­¯	­¯	­¯	­¯	­¯	­¯
Communication multiplicity	1	0	1	3	2	0	3	2½
Bond energy, kJ/mol	105	–	288	941	566	–	1070	677

A hypothetical He2 molecule would have two electrons in the σc1s orbital and two electrons in the σp1s orbital. Since one electron in the antibonding orbital cancels the binding action of the electron in the bonding orbital, the He2 molecule cannot exist. The VS method leads to the same conclusion.

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

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

Physical properties of molecules and MMO

The existence of bonding and loosening MOs is confirmed by the physical properties of the molecules. The MO method makes it possible to foresee that if, during the formation of a molecule from atoms, the electrons in the molecule fall into bonding orbitals, then the ionization potentials of the molecules must be greater than the ionization potentials of atoms, and if the electrons fall into loosening orbitals, then vice versa.

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

Diatomic molecules with different atoms

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

The AO energies of the oxygen atom lie below the energies of the corresponding carbon orbitals (1080 kJ/mol), they are located closer to the nucleus. The 10 electrons present in the initial atoms on the outer layers fill the bonding sb2s- and loosening sp2s-orbitals and the binding - and psb2py,z-orbitals. The CO molecule turns out to be isoelectronic with the N2 molecule. The binding energy of atoms in a CO molecule (1105 kJ/mol) is even greater than in a nitrogen molecule (940 kJ/mol). The C–O bond length is 0.113 nm.

NO molecule

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

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

Bibliography

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

Shimanovich I.E., Pavlovich M.L., Tikavy V.F., Malashko P.M. General chemistry in formulas, definitions, schemes. - Minsk: Universitetskaya, 1996. - S. 51-77.

Vorobyov V.K., Eliseev S.Yu., Vrublevsky A.V. Practical and independent work in chemistry. - Minsk: UE "Donarit", 2005. - S. 21-30.

Electronegativity is a conditional value that characterizes the ability of an atom in a chemical compound to attract electrons to itself.