Multiple (double and triple) bonds

In many molecules, atoms are connected by double and triple bonds:

The possibility of the formation of multiple bonds is due to the geometric characteristics of atomic orbitals. The hydrogen atom forms its only chemical bond with the participation of the valence 5-orbital, which has a spherical shape. The rest of the atoms, including even the atoms of the elements of the 5-block, have valence p-orbitals, which have a spatial orientation along the coordinate axes.

In the hydrogen molecule, the chemical bond is carried out by an electron pair, the cloud of which is concentrated between atomic nuclei. Bonds of this type are called st-bonds (a - read "sigma"). They are formed by mutual overlapping of both 5- and ir-orbitals (Fig. 6.3).

Rice. 63

For one more pair of electrons, there is no room between the atoms. How then are double and even triple bonds formed? It is possible to overlap electron clouds oriented perpendicular to the axis passing through the centers of atoms (Fig. 6.4). If the axis of the molecule is aligned with the coordinate x y then the orbitals are oriented perpendicular to it plf and r 2 . Pairwise overlap RU and p 2 orbitals of two atoms gives chemical bonds, the electron density of which is concentrated symmetrically on both sides of the axis of the molecule. They are called l-bonds.

If the atoms have RU and/or p 2 orbitals have unpaired electrons, then one or two n-bonds are formed. This explains the possibility of the existence of double (a + z) and triple (a + z + z) bonds. The simplest molecule with a double bond between atoms is the hydrocarbon molecule ethylene C 2 H 4 . On fig. Figure 6.5 shows the n-bond cloud in this molecule, and the st-bonds are indicated schematically by dashes. The ethylene molecule consists of six atoms. It probably occurs to readers that a double bond between atoms is depicted in a simpler diatomic oxygen molecule (0=0). In fact, the electronic structure of the oxygen molecule is more complex, and its structure could only be explained on the basis of the molecular orbital method (see below). An example of the simplest molecule with a triple bond is nitrogen. On fig. 6.6 presents n-bonds in this molecule, the dots show the unshared electron pairs of nitrogen.

Rice. 6.4.

Rice. 6.5.

Rice. 6.6.

When n-bonds are formed, the strength of molecules increases. Let's take some examples for comparison.

Considering the above examples, we can draw the following conclusions:

• - bond strength (energy) increases with increasing bond multiplicity;
• - Using the example of hydrogen, fluorine and ethane, one can also be convinced that the strength of a covalent bond is determined not only by the multiplicity, but also by the nature of the atoms between which this bond arose.

It is well known in organic chemistry that molecules with multiple bonds are more reactive than so-called saturated molecules. The reason for this becomes clear when considering the shape of electron clouds. The electron clouds of a-bonds are concentrated between the nuclei of atoms and, as it were, are screened (protected) by them from the influence of other molecules. In case of i-connection electron clouds are not shielded by the nuclei of atoms and are more easily displaced when the reacting molecules approach each other. This facilitates the subsequent rearrangement and transformation of molecules. An exception among all molecules is the nitrogen molecule, which is characterized by both very high strength and extremely low reactivity. Therefore, nitrogen will be the main component of the atmosphere.

Topics USE codifier: Covalent chemical bond, its varieties and mechanisms of formation. Characteristics of a covalent bond (polarity and bond energy). Ionic bond. Metal connection. hydrogen bond

Intramolecular chemical bonds

Let us first consider the bonds that arise between particles within molecules. Such connections are called intramolecular.

chemical bond between atoms chemical elements has an electrostatic nature and is formed due to interactions of external (valence) electrons, in more or less degree held by positively charged nuclei bonded atoms.

The key concept here is ELECTRONEGNATIVITY. It is she who determines the type chemical bond between atoms and the properties of this bond.

is the ability of an atom to attract (hold) external(valence) electrons. Electronegativity is determined by the degree of attraction of external electrons to the nucleus and depends mainly on the radius of the atom and the charge of the nucleus.

Electronegativity is difficult to determine unambiguously. L. Pauling compiled a table of relative electronegativity (based on the bond energies of diatomic molecules). The most electronegative element is fluorine with meaning 4 .

It is important to note that in different sources you can find different scales and tables of electronegativity values. This should not be frightened, since the formation of a chemical bond plays a role atoms, and it is approximately the same in any system.

If one of the atoms in the chemical bond A:B attracts electrons more strongly, then the electron pair is shifted towards it. The more electronegativity difference atoms, the more the electron pair is displaced.

If the electronegativity values ​​of the interacting atoms are equal or approximately equal: EO(A)≈EO(V), then the shared electron pair is not displaced to any of the atoms: A: B. Such a connection is called covalent non-polar.

If the electronegativity of the interacting atoms differ, but not much (the difference in electronegativity is approximately from 0.4 to 2: 0,4<ΔЭО<2 ), then the electron pair is shifted to one of the atoms. Such a connection is called covalent polar .

If the electronegativity of the interacting atoms differ significantly (the difference in electronegativity is greater than 2: ΔEO>2), then one of the electrons almost completely passes to another atom, with the formation ions. Such a connection is called ionic.

The main types of chemical bonds are − covalent, ionic and metallic connections. Let's consider them in more detail.

covalent chemical bond

covalent bond – it's a chemical bond formed by formation of a common electron pair A:B . In this case, two atoms overlap atomic orbitals. A covalent bond is formed by the interaction of atoms with a small difference in electronegativity (as a rule, between two non-metals) or atoms of one element.

Basic properties of covalent bonds

• orientation,
• saturability,
• polarity,
• polarizability.

These bond properties affect the chemical and physical properties of substances.

Direction of communication characterizes the chemical structure and form of substances. The angles between two bonds are called bond angles. For example, in a water molecule, the H-O-H bond angle is 104.45 o, so the water molecule is polar, and in the methane molecule, the H-C-H bond angle is 108 o 28 ′.

Saturability is the ability of atoms to form a limited number of covalent chemical bonds. The number of bonds that an atom can form is called.

Polarity bonds arise due to the uneven distribution of electron density between two atoms with different electronegativity. Covalent bonds are divided into polar and non-polar.

Polarizability connections are the ability of bond electrons to be displaced by an external electric field(in particular, the electric field of another particle). The polarizability depends on the electron mobility. The farther the electron is from the nucleus, the more mobile it is, and, accordingly, the molecule is more polarizable.

Covalent non-polar chemical bond

There are 2 types of covalent bonding - POLAR and NON-POLAR .

Example . Consider the structure of the hydrogen molecule H 2 . Each hydrogen atom carries 1 unpaired electron in its outer energy level. To display an atom, we use the Lewis structure - this is a diagram of the structure of the external energy level of an atom, when electrons are denoted by dots. Lewis point structure models are a good help when working with elements of the second period.

H. + . H=H:H

Thus, the hydrogen molecule has one common electron pair and one H–H chemical bond. This electron pair is not displaced to any of the hydrogen atoms, because the electronegativity of hydrogen atoms is the same. Such a connection is called covalent non-polar .

Covalent non-polar (symmetrical) bond - this is a covalent bond formed by atoms with equal electronegativity (as a rule, the same non-metals) and, therefore, with a uniform distribution of electron density between the nuclei of atoms.

The dipole moment of nonpolar bonds is 0.

Examples: H 2 (H-H), O 2 (O=O), S 8 .

Covalent polar chemical bond

covalent polar bond is a covalent bond that occurs between atoms with different electronegativity (usually, different non-metals) and is characterized displacement common electron pair to a more electronegative atom (polarization).

The electron density is shifted to a more electronegative atom - therefore, a partial negative charge (δ-) appears on it, and a partial positive charge appears on a less electronegative atom (δ+, delta +).

The greater the difference in the electronegativity of atoms, the higher polarity connections and even more dipole moment . Between neighboring molecules and charges opposite in sign, additional attractive forces act, which increases strength connections.

Bond polarity affects the physical and chemical properties of compounds. The reaction mechanisms and even the reactivity of neighboring bonds depend on the polarity of the bond. The polarity of a bond often determines polarity of the molecule and thus directly affects such physical properties as boiling point and melting point, solubility in polar solvents.

Examples: HCl, CO 2 , NH 3 .

Mechanisms for the formation of a covalent bond

A covalent chemical bond can occur by 2 mechanisms:

1. exchange mechanism the formation of a covalent chemical bond is when each particle provides one unpaired electron for the formation of a common electron pair:

BUT . + . B= A:B

2. The formation of a covalent bond is such a mechanism in which one of the particles provides an unshared electron pair, and the other particle provides a vacant orbital for this electron pair:

BUT: + B= A:B

In this case, one of the atoms provides an unshared electron pair ( donor), and the other atom provides a vacant orbital for this pair ( acceptor). As a result of the formation of a bond, both electron energy decreases, i.e. this is beneficial for the atoms.

A covalent bond formed by the donor-acceptor mechanism, is not different by properties from other covalent bonds formed by the exchange mechanism. The formation of a covalent bond by the donor-acceptor mechanism is typical for atoms either with a large number of electrons in the external energy level (electron donors), or vice versa, with a very small number of electrons (electron acceptors). The valence possibilities of atoms are considered in more detail in the corresponding.

A covalent bond is formed by the donor-acceptor mechanism:

- in a molecule carbon monoxide CO(the bond in the molecule is triple, 2 bonds are formed by the exchange mechanism, one by the donor-acceptor mechanism): C≡O;

- in ammonium ion NH 4 +, in ions organic amines, for example, in the methylammonium ion CH 3 -NH 2 + ;

- in complex compounds, a chemical bond between the central atom and groups of ligands, for example, in sodium tetrahydroxoaluminate Na the bond between aluminum and hydroxide ions;

- in nitric acid and its salts- nitrates: HNO 3 , NaNO 3 , in some other nitrogen compounds;

- in a molecule ozone O 3 .

Main characteristics of a covalent bond

A covalent bond, as a rule, is formed between the atoms of non-metals. The main characteristics of a covalent bond are length, energy, multiplicity and directivity.

Chemical bond multiplicity

Chemical bond multiplicity - this is the number of shared electron pairs between two atoms in a compound. The multiplicity of the bond can be quite easily determined from the value of the atoms that form the molecule.

For example , in the hydrogen molecule H 2 the bond multiplicity is 1, because each hydrogen has only 1 unpaired electron in the outer energy level, therefore, one common electron pair is formed.

In the oxygen molecule O 2, the bond multiplicity is 2, because each atom has 2 unpaired electrons in its outer energy level: O=O.

In the nitrogen molecule N 2, the bond multiplicity is 3, because between each atom there are 3 unpaired electrons in the outer energy level, and the atoms form 3 common electron pairs N≡N.

Covalent bond length

Chemical bond length is the distance between the centers of the nuclei of atoms that form a bond. It is determined by experimental physical methods. The bond length can be estimated approximately, according to the additivity rule, according to which the bond length in the AB molecule is approximately equal to half the sum of the bond lengths in the A 2 and B 2 molecules:

The length of a chemical bond can be roughly estimated along the radii of atoms, forming a bond, or by the multiplicity of communication if the radii of the atoms are not very different.

With an increase in the radii of the atoms forming a bond, the bond length will increase.

For example

With an increase in the multiplicity of bonds between atoms (whose atomic radii do not differ, or differ slightly), the bond length will decrease.

For example . In the series: C–C, C=C, C≡C, the bond length decreases.

Bond energy

A measure of the strength of a chemical bond is the bond energy. Bond energy is determined by the energy required to break the bond and remove the atoms that form this bond to an infinite distance from each other.

The covalent bond is very durable. Its energy ranges from several tens to several hundreds of kJ/mol. The greater the bond energy, the greater the bond strength, and vice versa.

The strength of a chemical bond depends on the bond length, bond polarity, and bond multiplicity. The longer the chemical bond, the easier it is to break, and the lower the bond energy, the lower its strength. The shorter the chemical bond, the stronger it is, and the greater the bond energy.

For example, in the series of compounds HF, HCl, HBr from left to right the strength of the chemical bond decreases, because the length of the bond increases.

Ionic chemical bond

Ionic bond is a chemical bond based on electrostatic attraction of ions.

ions are formed in the process of accepting or giving away electrons by atoms. For example, the atoms of all metals weakly hold the electrons of the outer energy level. Therefore, metal atoms are characterized restorative properties the ability to donate electrons.

Example. The sodium atom contains 1 electron at the 3rd energy level. Easily giving it away, the sodium atom forms a much more stable Na + ion, with the electron configuration of the noble neon gas Ne. The sodium ion contains 11 protons and only 10 electrons, so the total charge of the ion is -10+11 = +1:

+11Na) 2 ) 8 ) 1 - 1e = +11 Na +) 2 ) 8

Example. The chlorine atom has 7 electrons in its outer energy level. To acquire the configuration of a stable inert argon atom Ar, chlorine needs to attach 1 electron. After the attachment of an electron, a stable chlorine ion is formed, consisting of electrons. The total charge of the ion is -1:

+17Cl) 2 ) 8 ) 7 + 1e = +17 Cl — ) 2 ) 8 ) 8

Note:

• The properties of ions are different from the properties of atoms!
• Stable ions can form not only atoms, but also groups of atoms. For example: ammonium ion NH 4 +, sulfate ion SO 4 2-, etc. Chemical bonds formed by such ions are also considered ionic;
• Ionic bonds are usually formed between metals and nonmetals(groups of non-metals);

The resulting ions are attracted due to electrical attraction: Na + Cl -, Na 2 + SO 4 2-.

Let us visually generalize difference between covalent and ionic bond types:

metal chemical bond

metal connection is the relationship that is formed relatively free electrons between metal ions forming a crystal lattice.

The atoms of metals on the outer energy level usually have one to three electrons. The radii of metal atoms, as a rule, are large - therefore, metal atoms, unlike non-metals, quite easily donate outer electrons, i.e. are strong reducing agents

Intermolecular interactions

Separately, it is worth considering the interactions that occur between individual molecules in a substance - intermolecular interactions . Intermolecular interactions are a type of interaction between neutral atoms in which new covalent bonds do not appear. The forces of interaction between molecules were discovered by van der Waals in 1869 and named after him. Van dar Waals forces. Van der Waals forces are divided into orientation, induction and dispersion . The energy of intermolecular interactions is much less than the energy of a chemical bond.

Orientation forces of attraction arise between polar molecules (dipole-dipole interaction). These forces arise between polar molecules. Inductive interactions is the interaction between a polar molecule and a non-polar one. A non-polar molecule is polarized due to the action of a polar one, which generates an additional electrostatic attraction.

A special type of intermolecular interaction is hydrogen bonds. - these are intermolecular (or intramolecular) chemical bonds that arise between molecules in which there are strongly polar covalent bonds - H-F, H-O or H-N. If there are such bonds in the molecule, then between the molecules there will be additional forces of attraction .

Mechanism of education The hydrogen bond is partly electrostatic and partly donor-acceptor. In this case, an atom of a strongly electronegative element (F, O, N) acts as an electron pair donor, and hydrogen atoms connected to these atoms act as an acceptor. Hydrogen bonds are characterized orientation in space and saturation .

The hydrogen bond can be denoted by dots: H ··· O. The greater the electronegativity of an atom connected to hydrogen, and the smaller its size, the stronger the hydrogen bond. It is primarily characteristic of compounds fluorine with hydrogen , as well as to oxygen with hydrogen , less nitrogen with hydrogen .

Hydrogen bonds occur between the following substances:

— hydrogen fluoride HF(gas, solution of hydrogen fluoride in water - hydrofluoric acid), water H 2 O (steam, ice, liquid water):

— solution of ammonia and organic amines- between ammonia and water molecules;

— organic compounds in which O-H or N-H bonds: alcohols, carboxylic acids, amines, amino acids, phenols, aniline and its derivatives, proteins, solutions of carbohydrates - monosaccharides and disaccharides.

The hydrogen bond affects the physical and chemical properties of substances. Thus, the additional attraction between molecules makes it difficult for substances to boil. Substances with hydrogen bonds exhibit an abnormal increase in the boiling point.

For example As a rule, with an increase in molecular weight, an increase in the boiling point of substances is observed. However, in a number of substances H 2 O-H 2 S-H 2 Se-H 2 Te we do not observe a linear change in boiling points.

Namely, at boiling point of water is abnormally high - not less than -61 o C, as the straight line shows us, but much more, +100 o C. This anomaly is explained by the presence of hydrogen bonds between water molecules. Therefore, under normal conditions (0-20 o C), water is liquid by phase state.

170762 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 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 + + 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 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 FROM 2 H 5 HE, sulfur hexafluoride SF 6, acetylene FROM 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. four. 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 transition 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

In the considered examples of the formation of a chemical bond, an electron pair took part. Such a connection is called single. Sometimes it is called ordinary, i.e. ordinary. This type of connection is usually denoted by a single line connecting the symbols of interacting atoms.

Overlapping electron clouds in a straight line connecting two nuclei leads to sigma bonds(o-bond). A single bond is in most cases an a-bond.

The bond formed by the overlapping of the side regions of p-electron clouds is called pi-bond(i-bond). Double and triple bonds are formed with the participation of two and three electron pairs, respectively. A double bond is one a-bond and one i-bond, a triple bond is one a-bond and two i-bonds.

Let us discuss the formation of bonds in the molecules of ethane C 2 H 6 , ethylene C 2 H 4 , acetylene C 2 H 2 and benzene C 6 H b.

The angles between bonds in a molecule ethane FROM. ; H (. almost exactly equal to each other (Fig. 1.18, a) and do not differ from the angles between C-H bonds in the methane molecule. Therefore, it can be assumed that the outer electron shells of carbon atoms are in a state of $p 3 hybridization. The C 2 H 6 molecule is diamagnetic and does not have an electric dipole moment. The C-C bond energy is -335 kJ/mol. All bonds in the C 9 H 6 molecule are a-bonds.

In a molecule ethylene C 2 H 4 bond angles are approximately 120° each. From this we can conclude that the $ p 2 hybridization of the outer electron orbitals of the carbon atom (Fig. 1.18, b). The C-H bonds lie in the same plane at angles of about 120°. Each carbon atom has one non-hybrid p-orbital containing

Rice. 1.18. Models of ethane molecules ( a ), ethylene (b) and acetylene (c)

holding one electron. These orbitals are located perpendicular to the plane of the figure.

The bond energy between carbon atoms in an ethylene C 2 H 4 molecule is -592 kJ/mol. If the carbon atoms were linked by the same bond as in the ethane molecule, then the binding energies in these molecules would be close.

However, the binding energy between carbon atoms in ethane is 335 kJ/mol, which is almost two times less than in ethylene. Such a significant difference in the binding energies between carbon atoms in ethylene and ethane molecules is explained by the possible interaction of non-hybrid p-orbitals, which in Fig. 1.18 , b depicted with wavy lines. The connection formed in this way is called the I-connection.

In the C 2 H 4 ethylene molecule, four CH bonds, as in the CH 4 methane molecule, are a-bonds, and the bond between carbon atoms is an a-bond and a p-bond, i.e. double bond, and the formula of ethylene is written as H 2 C=CH 2.

The acetylene C 2 H 2 molecule is linear (Fig. 1.18, in ), which speaks in favor of sp hybridization. The bond energy between carbon atoms is -811 kJ/mol, which suggests the existence of one a-bond and two n-bonds, i.e. it's a triple bond. The formula of acetylene is written as HC=CH.

One of the difficult questions of chemistry is to establish the nature of the bonds between carbon atoms in the so-called aromatic compounds , in particular, in the C 6 H benzene molecule (.. The benzene molecule is flat, the angles between the bonds of carbon atoms are equal in

Rice. 1.19.

a - formula model: 6 - ^-orbitals of carbon atoms and a-bonds between carbon atoms and carbon and hydrogen atoms; in- p-inhabitants and l-connections between

carbon atoms

120°, which allows us to assume the ^-hybridization of the outer orbitals of carbon atoms. Typically, the benzene molecule is depicted as shown in rice. 1.19, a.

It would seem that in benzene the bond between carbon atoms should be longer than the C=C double bond, as it is stronger. However, the study of the structure of the benzene molecule shows that all distances between carbon atoms in the benzene ring are the same.

This feature of the molecule is best explained by the fact that the non-hybrid p-orbitals of all carbon atoms are overlapped by "side" parts (Fig. 1.19, b) therefore, all internuclear distances between carbon atoms are equal. On fig. 1.19 in shows a-bonds between carbon atoms formed by overlapping sp2- hybrid orbitals.

Bond energy between atoms carbon in the benzene molecule C 6 H 6 is -505 kJ / mol, and this suggests that these bonds are intermediate between single and double bonds. Note that the electrons of the p-orbitals in the benzene molecule move along a closed hexagon, and they delocalized(does not refer to any specific place).

The forces that bind atoms to each other are of the same electrical nature. But due to the difference in the mechanism of formation and manifestation of these forces, chemical bonds can be of different types.

Distinguish three major typevalence chemical bond: covalent, ionic and metallic.

In addition to them, of great importance and distribution are: hydrogen connection that may be valence and non-valent, and non-valent chemical bond - m intermolecular ( or van der Waalsow), forming relatively small associates of molecules and huge molecular ensembles - super- and supramolecular nanostructures.

covalent chemical bond (atomic, homeopolar) -

this is chemical bond carried out general for interacting atoms one-threepairs of electrons .

This connection is two-electron and two-center(binds 2 atomic nuclei).

In this case, the covalent bond is most common and most common type valence chemical bond in binary compounds - between a) atoms of non-metals and b) atoms of amphoteric metals and non-metals.

Examples: H-H (in the hydrogen molecule H 2); four S-O bonds (in SO 4 2- ion); three Al-H bonds (in the AlH 3 molecule); Fe-S (in the FeS molecule), etc.

Peculiarities covalent bond - orientation and saturability.

Orientation - the most important property of a covalent bond, from

which depends on the structure (configuration, geometry) of molecules and chemical compounds. The spatial orientation of the covalent bond determines the chemical and crystal-chemical structure of the substance. covalent bond always directed in the direction of maximum overlap of atomic orbitals of valence electrons interacting atoms, with the formation of a common electron cloud and the strongest chemical bond. Orientation expressed in the form of angles between the directions of bonding of atoms in molecules of different substances and crystals of solids.

Saturability is a property, which distinguishes the covalent bond from all other types of particle interaction, manifested in the ability of atoms to form a limited number of covalent bonds, since each pair of binding electrons is formed only valence electrons with oppositely oriented spins, the number of which in an atom is limited valency, 1 - 8. In this case, it is forbidden to use the same atomic orbital twice to form a covalent bond (Pauli principle).

Valence - this is the ability of an atom to attach or replace a certain number of other atoms with the formation of valence chemical bonds.

According to the spin theory covalent bond valence determined the number of unpaired electrons in an atom in the ground or excited state .

Thus, for different elements ability to form a certain number of covalent bonds limited to receiving the maximum number of unpaired electrons in the excited state of their atoms.

Excited state of an atom - this is the state of an atom with additional energy received by it from the outside, causing steaming antiparallel electrons occupying one atomic orbital, i.e. the transition of one of these electrons from a paired state to a free (vacant) orbital the same or close energy level.

For example, scheme filling s-, r-AO and valence (AT) at the calcium atom Sa mostly and excited state the following:

It should be noted that the atoms with saturated valence bonds can form additional covalent bonds by a donor-acceptor or other mechanism (as, for example, in complex compounds).

covalent bond may bepolar andnon-polar .

covalent bond non-polar , e if socialized valence electrons evenly distributed between the nuclei of interacting atoms, the region of overlapping atomic orbitals (electron clouds) is attracted by both nuclei with the same force and therefore the maximum the total electron density is not biased towards either of them.

This type of covalent bond occurs when two identical element atoms. Covalent bond between identical atoms also called atomic or homeopolar .

Polar connection arises during the interaction of two atoms of different chemical elements, if one of the atoms due to a larger value electronegativity attracts valence electrons more strongly, and then the total electron density is more or less shifted towards this atom.

With a polar bond, the probability of finding an electron at the nucleus of one of the atoms is higher than that of the other.

Qualitative characteristic of the polar communications -

difference of relative electronegativity (|‌‌‌‌‌‌‌‌‌∆OEE |)‌‌‌ related atoms : the larger it is, the more polar the covalent bond is.

Quantitative characteristics of the polar communications, those. a measure of the polarity of a bond and a complex molecule - dipole electric moment μ St. , equal to workeffective charge δ per dipole length l d : μ St. = δ ∙ l d . unit of measurement μ St.- Debye. 1Debye = 3,3.10 -30 C/m.

electric dipole - this is an electrically neutral system of two electric charges equal and opposite in sign + δ and - δ .

Dipole moment (electric moment of the dipole μ St. ) – vector quantity . It is generally accepted that vector direction from (+) to (-) matches with the direction of displacement of the total electron density region(total electron cloud) polarized atoms.

General dipole moment of a complex polyatomic molecule depends on the number and spatial orientation of polar bonds in it. Thus, the determination of dipole moments makes it possible to judge not only the nature of bonds in molecules, but also their location in space, i.e. about the spatial configuration of the molecule.

With an increase in the difference of electronegativity | ‌‌‌‌‌‌‌‌‌∆OEE|‌‌‌ atoms forming a bond, the electric moment of the dipole increases.

It should be noted that the determination of the bond dipole moment is a complex and not always solvable problem (bond interaction, unknown direction μ St. etc.).

Quantum-mechanical methods for describing a covalent bond – explain the mechanism of formation of a covalent bond.

Conducted by W. Geytler and F. London, German. scientists (1927), the calculation of the energy balance of the formation of a covalent bond in the hydrogen molecule H 2 made it possible to make conclusion: the nature of the covalent bond, like any other type of chemical bond, lies inelectrical interaction occurring under the conditions of a quantum mechanical microsystem.

To describe the mechanism of formation of a covalent chemical bond, use two approximate quantum mechanical methods :

valence bonds and molecular orbitals – not exclusive, but mutually complementary.

2.1. Valence bond method (MVS orlocalized electron pairs ), proposed by W. Geytler and F. London in 1927, is based on the following provisions :

1) a chemical bond between two atoms arises as a result of the partial overlap of atomic orbitals with the formation of a common electron density of a joint pair of electrons with opposite spins, higher than in other regions of space around each nucleus;

2) covalent a bond is formed only when electrons with antiparallel spins interact, i.e. with opposite spin quantum numbers m S = + 1/2 ;

3) characteristics of a covalent bond (energy, length, polarity, etc.) are determined view connections (σ –, π –, δ –), degree of overlapping AO(the larger it is, the stronger the chemical bond, i.e. the higher the bond energy and the shorter the length), electronegativity interacting atoms;

4) a covalent bond can be formed by MVS two ways (two mechanisms) , fundamentally different, but having the same result – socialization of a pair of valence electrons by both interacting atoms: a) exchange, due to the overlap of one-electron atomic orbitals with opposite electron spins, when each atom contributes one electron per bond to overlap – the bond can be either polar or non-polar, b) donor-acceptor, due to the two-electron AO of one atom and the free (vacant) orbital of the other, on to whom one atom (donor) provides for bonding a pair of electrons in the orbital in a paired state, and the other atom (acceptor) provides a free orbital. This gives rise to polar bond.

2.2. Complex (coordination) compounds, many molecular ions that are complex,(ammonium, boron tetrahydride, etc.) are formed in the presence of a donor-acceptor bond - in other words, a coordination bond.

For example, in the reaction of formation of an ammonium ion NH 3 + H + = NH 4 + ammonia molecule NH 3 is an electron pair donor, and the H + proton is an acceptor.

In the reaction ВН 3 + Н - = ВН 4 - the hydride ion Н - plays the role of an electron pair donor, and the boron hydride molecule ВН 3, in which there is a vacant AO, plays the role of an acceptor.

The multiplicity of the chemical bond. Connections σ -, π – , δ –.

The maximum overlap of AO of different types (with the establishment of the strongest chemical bonds) is achieved with their specific orientation in space, due to the different shape of their energy surface.

The type of AO and the direction of their overlap determine σ -, π – , δ - connections:

σ (sigma) – connection – it's always aboutdinar (simple) bond arising from partial overlap one pair s -, p x -, d - JSCalong the axis , connecting the core interacting atoms.

Single bonds always are σ - connections.

Multiple bonds – π (pi) - (also δ (delta )–connections),double or triple covalent bonds carried out respectivelytwo orthree couples electrons when their atomic orbitals overlap.

π (pi) - connection carried out by overlapping R y -, p z - and d - JSC on both sides of the axis connecting the nuclei atoms, in mutually perpendicular planes ;

δ (delta )- connection occurs when overlapping two d orbitals located in parallel planes .

The most durable of σ -, π – , δ – connections is σ– bond , but π - connections based on σ – bond, form even stronger multiple bonds: double and triple.

Any double bond comprises one σ – and one π – connections, triple - from oneσ – and twoπ – connections.