Types of chemical bonds. Single bond Single covalent bond in a molecule

Fig.1. Orbital radii of elements (r a) and single-electron length chemical bond(d)

The simplest one-electron chemical bond is created by a single valence electron. It turns out that one electron is capable of holding two positively charged ions together. In a one-electron bond, the Coulomb repulsive forces of positively charged particles are compensated by the Coulomb forces of attraction of these particles to a negatively charged electron. The valence electron becomes common to the two nuclei of the molecule.

Examples of such chemical compounds are molecular ions: H 2 +, Li 2 +, Na 2 +, K 2 +, Rb 2 +, Cs 2 +:

Polar covalent bonds occur in heteronuclear diatomic molecules (Fig. 3). The bonding electron pair in a polar chemical bond is closer to the atom with more high first ionization potential.

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

Characteristics of some polar substances

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

The distance between the centers of gravity of positive and negative charges is called the dipole length. The polarity of a molecule, as well as the polarity of a bond, is assessed by the value of the dipole moment μ, which is the product of the dipole length l and the value of the electronic charge:

Multiple covalent bonds

Multiple covalent bonds are represented by unsaturated organic compounds containing double and triple chemical bonds. To describe the nature of unsaturated compounds, L. Pauling introduces the concepts of sigma and π bonds, hybridization of atomic orbitals.

Pauling hybridization for two S and two p electrons made it possible to explain the directionality of chemical bonds, in particular the tetrahedral configuration of methane. To explain the structure of ethylene, from four equivalent Sp 3 electrons of the carbon atom, one p-electron has to be isolated to form an additional bond, called a π bond. In this case, the three remaining Sp 2 hybrid orbitals are located in the plane at an angle of 120° and form basic bonds, for example, a planar ethylene molecule (Fig. 5).

IN new theory Pauling, all bonding electrons became equal and equidistant from the line connecting the nuclei of the molecule. Pauling's theory of the bent chemical bond took into account the statistical interpretation of the M. Born wave function and the Coulomb electron correlation of electrons. Appeared physical meaning- the nature of the chemical bond is completely determined electrical interaction nuclei and electrons. The more bonding electrons, the smaller the internuclear distance and the stronger the chemical bond between carbon atoms.

Three-center chemical bond

Further development of ideas about chemical bonds was given by the American physical chemist W. Lipscomb, who developed the theory of two-electron three-center bonds and a topological theory that makes it possible to predict the structure of some more boron hydrides (hydrogen hydrides).

An electron pair in a three-center chemical bond becomes common to three atomic nuclei. In the simplest representative of a three-center chemical bond - the molecular hydrogen ion H 3 +, an electron pair holds three protons in a single whole (Fig. 6).

Fig. 7. Diboran

The existence of boranes with their two-electron three-center bonds with “bridging” hydrogen atoms violated the canonical doctrine of valence. The hydrogen atom, previously considered a standard monovalent element, turned out to be connected by identical bonds to two boron atoms and formally became a divalent element. W. Lipscomb's work on deciphering the structure of boranes expanded the understanding of chemical bonds. Nobel Committee awarded William Nunn Lipscomb the Chemistry Prize for 1976 with the wording "For his studies of the structure of boranes (borohydrites), clarifying the problems of chemical bonds."

Multisite chemical bond

Fig. 8. Ferrocene molecule

Fig. 9. Dibenzene chromium

Fig. 10. Uranocene

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

Dynamics of chemical bonding

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

In pairs, these metals consist practically of homonuclear diatomic molecules and free atoms. When metal vapor condenses, a covalent bond is converted into a metal bond.

Evaporation of salts with typical ionic bonds, such as fluorides alkali metals, leads to the destruction of ionic bonds and the formation of heteronuclear diatomic molecules with polar covalent bond. In this case, the formation of dimeric molecules with bridged bonds occurs.

Characteristics of chemical bonds in molecules of alkali metal fluorides and their dimers.

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

Mechanism of transition of covalent to metallic bond

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

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

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

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

The formation of a metal-to-metal covalent bond is associated with deformation electron shells interacting atoms - valence electrons create a bonding electron pair, the electron density of which is concentrated in the space between the atomic nuclei of the resulting molecule. A characteristic feature of homonuclear diatomic molecules of alkali metals is the long length of the covalent bond (3.6-5.8 times longer than the bond length in the hydrogen molecule) and the low energy of its rupture.

The indicated relationship between r e and d determines the uneven distribution of electric charges in the molecule - the negative electric charge of the bonding electron pair is concentrated in the middle part of the molecule, and the positive ones are concentrated at the ends of the molecule electric charges two atomic skeletons.

The uneven distribution of electric charges creates conditions for the interaction of molecules due to orientation forces (van der Waals forces). Molecules of alkali metals tend to orient themselves in such a way that opposite electric charges appear in their proximity. As a result, attractive forces act between molecules. Thanks to the presence of the latter, the molecules of alkali metals come closer and are more or less firmly pulled together. At the same time, some deformation of each of them occurs under the influence of closer poles of neighboring molecules (Fig. 12).

In fact, the binding electrons of the original diatomic molecule, falling into the electric field of the four positively charged atomic cores of alkali metal molecules, are torn away from the orbital radius of the atom and become free.

In this case, the bonding electron pair becomes common for a system with six cations. The construction of the metal crystal lattice begins at the cluster stage. IN crystal lattice alkali metals, the structure of the connecting link is clearly expressed, having the shape of a distorted flattened octahedron - a square bipyramid, the height of which and the edges of the basis are equal to the value of the translation lattice constant a w (Fig. 13).

The value of the translation lattice constant a w of an alkali metal crystal significantly exceeds the length of the covalent bond of an alkali metal molecule, therefore it is generally accepted that the electrons in the metal are in a free state:

The mathematical construction associated with the properties of free electrons in a metal is usually identified with the "Fermi surface", which should be considered as locus, where electrons reside, providing the main property of a metal - to conduct electric current.

When comparing the process of condensation of alkali metal vapors with the process of condensation of gases, for example, hydrogen, characteristic feature in the properties of the metal. Thus, if during the condensation of hydrogen weak intermolecular interactions appear, then during the condensation of metal vapor processes occur that are characteristic of chemical reactions. The condensation of metal vapor itself occurs in several stages and can be described by the following process: free atom → diatomic molecule with a covalent bond → metal cluster → compact metal with a metal bond.

The interaction of alkali metal halide molecules is accompanied by their dimerization. A dimer molecule can be considered an electric quadrupole (Fig. 15). Currently, the main characteristics of dimers of alkali metal halides are known (chemical bond lengths and bond angles between bonds).

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

E 2 X 2	X=F		X=Cl		X=Br		X=I
E 2 X 2	dEF, Å		d ECl, Å		d EBr, Å		d EI, Å
Li 2 X 2	1,75	105	2,23	108	2,35	110	2,54	116
Na 2 X 2	2,08	95	2,54	105	2,69	108	2,91	111
K 2 X 2	2,35	88	2,86	98	3,02	101	3,26	104
Cs 2 X 2	2,56	79	3,11	91	3,29	94	3,54	94

During the condensation process, the effect of orientation forces increases, intermolecular interaction is accompanied by the formation of clusters, and then a solid substance. Alkali metal halides form crystals with simple cubic and body-centered cubic lattices.

Crystal lattice type and translation lattice constant for alkali metal halides.

During the crystallization process, a further increase in the interatomic distance occurs, leading to the removal of an electron from the orbital radius of the alkali metal atom and the transfer of an electron to the halogen atom with the formation of the corresponding ions. The force fields of ions are evenly distributed in all directions in space. In this regard, in alkali metal crystals, the force field of each ion is coordinated by more than one ion with the opposite sign, as is customary to qualitatively represent the ionic bond (Na + Cl -).

In crystals of ionic compounds, the concept of simple two-ion molecules such as Na + Cl - and Cs + Cl - loses its meaning, since the alkali metal ion is associated with six chlorine ions (in a sodium chloride crystal) and with eight chlorine ions (in a cesium chloride crystal. However, all interionic distances in crystals are equidistant.

Notes

Handbook of Inorganic Chemistry. Constants of inorganic substances. - M.: “Chemistry”, 1987. - P. 124. - 320 p.
Lidin R.A., Andreeva L.L., Molochko V.A. Handbook of Inorganic Chemistry. Constants of inorganic substances. - M.: “Chemistry”, 1987. - P. 132-136. - 320 s.
Gankin V.Yu., Gankin Yu.V. How a chemical bond is formed and chemical reactions occur. - M.: publishing group "Granitsa", 2007. - 320 p. - ISBN 978-5-94691296-9
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Pauling L. The nature of chemical bonding / edited by Y.K. Syrkin. - per. from English M.E. Dyatkina. - M.-L.: Goskhimizdat, 1947. - 440 p.
Theoretical organic chemistry / ed. R.H. Freidlina. - per. from English Yu.G.Bundela. - M.: Publishing house. foreign literature, 1963. - 365 p.
Lemenovsky D.A., Levitsky M.M. Russian Chemical Journal (journal of the Russian Chemical Society named after D.I. Mendeleev). - 2000. - T. XLIV, issue 6. - P. 63-86.
Chemical encyclopedic dictionary / ch. ed. I.L. Knunyants. - M.: Sov. encyclopedia, 1983. - P. 607. - 792 p.
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Lidin R.A., Andreeva L.L., Molochko V.A. Handbook of Inorganic Chemistry. Constants of inorganic substances. - M.: “Chemistry”, 1987. - P. 155-161. - 320 s.
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Chemist's Handbook. - 2nd ed., revised. and additional - L.-M.: State Scientific and Technical Institute of Chemical Literature, 1962. - T. 1. - P. 402-513. - 1072 pp.
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In addition to them, the following are of great importance and distribution: hydrogen connection that could be valence And nonvalent, And nonvalent chemical bond - m intermolecular ( or van der Waals), forming relatively small molecular associates and huge molecular ensembles - super- and supramolecular nanostructures.

Covalent chemical bond (atomic, homeopolar) –

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

This connection is two-electron And two-center(links 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 the SO 4 2- ion); three Al-H bonds (in the AlH 3 molecule); Fe-S (in the FeS molecule), etc.

Peculiarities covalent bond - its focus And saturability.

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

which determines the structure (configuration, geometry) of molecules and chemical compounds. The spatial direction of the covalent bond determines the chemical and crystal chemical structure of the substance. Covalent bond always directed towards maximum overlap of atomic orbitals of valence electrons interacting atoms, with the formation of a common electron cloud and the strongest chemical bond. Focus 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 a covalent bond from all other types of particle interactions, manifested in the ability of atoms to form a limited number of covalent bonds, since each pair of bonding electrons is formed only valence electrons with oppositely oriented spins, the number of which in an atom is limited valency, 1 – 8. This prohibits the use of the same atomic orbital twice to form a covalent bond (Pauli principle).

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

According to spin theory covalent bond valence determined the number of unpaired electrons an atom has in its 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 the atom with additional energy received 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, schemes filling s-, r-AO And valence (IN) at the calcium atom Sa mostly And excited state the following:

It should be noted that 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 shared valence electrons evenly distributed between the nuclei of interacting atoms, the region of overlap of 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 any of them.

This type of covalent bond occurs when two identical atoms of the element. 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 the valence electrons more strongly, and then the total electron density is more or less shifted towards that atom.

In a polar bond, the probability of finding an electron in the nucleus of one of the atoms is higher than in the other.

Qualitative characteristics of polar communications –

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

Quantitative characteristics of polar communications, those. measure of bond polarity and complex molecule - electric dipole moment μ St. , equal 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 – is an electrically neutral system of two equal and opposite electric charges + δ And - δ .

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

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

With increasing electronegativity difference | ‌‌‌‌‌‌‌‌‌∆OEO|‌‌‌ atoms forming a bond, the electric dipole moment increases.

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

Quantum mechanical methods for describing covalent bonds – explain mechanism of covalent bond formation.

Conducted by W. Heitler and F. London, German. scientists (1927), calculation of the energy balance of the formation of a covalent bond in the hydrogen molecule H2 made it possible to make conclusion: nature of covalent bond, like any other type of chemical bond, iselectrical 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. Heitler and F. London in 1927, is based on the following provisions :

1) a chemical bond between two atoms results from the partial overlap of atomic orbitals to form 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 AO overlap(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 along the MBC can be formed in two ways (two mechanisms) , fundamentally different, but having the same result – sharing 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, By to whom one atom (donor) provides a pair of electrons in the orbital in a paired state for bonding, and the other atom (acceptor) provides a free orbital. In this case, there arises polar connection.

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 - otherwise, a coordination bond.

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

In the reaction BH 3 + H – = BH 4 – the role of electron pair donor is played by the hydride ion H –, and the acceptor is the boron hydride molecule BH 3, in which there is a vacant AO.

Multiplicity of chemical bond. Connections σ -, π – , δ –.

The maximum overlap of AOs of different types (with the establishment of the strongest chemical bonds) is achieved when they have a certain 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 Odinar (simple) connection , which occurs when there is partial overlap one pair s -, p x -, d - JSCalong the axis , connecting the nuclei interacting atoms.

Single bonds Always are σ – connections.

Multiple connections – π (pi) - (Also δ (delta )–connections),double or triples covalent bonds carried out accordinglytwo orthree pairs electrons when their atomic orbitals overlap.

π (pi) - connection carried out when overlapping r y -, p z - And d - JSC By both sides of the axis connecting the nuclei atoms, in mutually perpendicular planes ;

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

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

Any double bond consists of one σ – And one π – connections, triple - from oneσ – And twoπ – connections.

Atoms of most elements do not exist separately, as they can interact with each other. This interaction produces more complex particles.

The nature of a chemical bond is the action of electrostatic forces, which are the forces of interaction between electric charges. Electrons and atomic nuclei have such charges.

Electrons located on the outer electronic levels (valence electrons), being farthest from the nucleus, interact with it weakest, and therefore are able to break away from the nucleus. They are responsible for bonding atoms to each other.

Types of interactions in chemistry

Types of chemical bonds can be presented in the following table:

Characteristics of ionic bonding

Chemical reaction that occurs due to ion attraction having different charges is called ionic. This happens if the atoms being bonded have a significant difference in electronegativity (that is, the ability to attract electrons) and the electron pair goes to the more electronegative element. The result of this transfer of electrons from one atom to another is the formation of charged particles - ions. An attraction arises between them.

They have the lowest electronegativity indices typical metals, and the largest are typical non-metals. Ions are thus formed by the interaction between typical metals and typical nonmetals.

Metal atoms become positively charged ions (cations), donating electrons to their outer electron levels, and nonmetals accept electrons, thus turning into negatively charged ions (anions).

Atoms move into a more stable energy state, completing their electronic configurations.

The ionic bond is non-directional and non-saturable, since the electrostatic interaction occurs in all directions; accordingly, the ion can attract ions of the opposite sign in all directions.

The arrangement of the ions is such that around each there is a certain number of oppositely charged ions. The concept of "molecule" for ionic compounds doesn't make sense.

Examples of education

The formation of a bond in sodium chloride (nacl) is due to the transfer of an electron from the Na atom to the Cl atom to form the corresponding ions:

Na 0 - 1 e = Na + (cation)

Cl 0 + 1 e = Cl - (anion)

In sodium chloride, there are six chloride anions around the sodium cations, and six sodium ions around each chloride ion.

When interaction is formed between atoms in barium sulfide, the following processes occur:

Ba 0 - 2 e = Ba 2+

S 0 + 2 e = S 2-

Ba donates its two electrons to sulfur, resulting in the formation of sulfur anions S 2- and barium cations Ba 2+.

Metal chemical bond

The number of electrons in the outer energy levels of metals is small; they are easily separated from the nucleus. As a result of this detachment, metal ions and free electrons are formed. These electrons are called "electron gas". Electrons move freely throughout the volume of the metal and are constantly bound and separated from atoms.

The structure of the metal substance is as follows: the crystal lattice is the skeleton of the substance, and between its nodes electrons can move freely.

The following examples can be given:

Mg - 2е<->Mg 2+

Cs-e<->Cs+

Ca - 2e<->Ca2+

Fe-3e<->Fe 3+

Covalent: polar and non-polar

The most common type of chemical interaction is a covalent bond. The electronegativity values of the elements that interact do not differ sharply; therefore, only a shift of the common electron pair to a more electronegative atom occurs.

Covalent interactions can be formed by an exchange mechanism or a donor-acceptor mechanism.

The exchange mechanism is realized if each of the atoms has unpaired electrons on the outer electronic levels and the overlap of atomic orbitals leads to the appearance of a pair of electrons that already belongs to both atoms. When one of the atoms has a pair of electrons on the outer electronic level, and the other has a free orbital, then when the atomic orbitals overlap, the electron pair is shared and interacts according to the donor-acceptor mechanism.

Covalent ones are divided by multiplicity into:

simple or single;
double;
triples.

Double ones ensure the sharing of two pairs of electrons at once, and triple ones - three.

According to the distribution of electron density (polarity) between bonded atoms, a covalent bond is divided into:

non-polar;
polar.

A nonpolar bond is formed by identical atoms, and a polar bond is formed by different electronegativity.

The interaction of atoms with similar electronegativity is called a nonpolar bond. The common pair of electrons in such a molecule is not attracted to either atom, but belongs equally to both.

The interaction of elements differing in electronegativity leads to the formation of polar bonds. In this type of interaction, shared electron pairs are attracted to the more electronegative element, but are not completely transferred to it (that is, the formation of ions does not occur). As a result of this shift in electron density, partial charges appear on the atoms: the more electronegative one has a negative charge, and the less electronegative one has a positive charge.

Properties and characteristics of covalency

Main characteristics of a covalent bond:

The length is determined by the distance between the nuclei of interacting atoms.
Polarity is determined by the displacement of the electron cloud towards one of the atoms.
Directionality is the property of forming bonds oriented in space and, accordingly, molecules having certain geometric shapes.
Saturation is determined by the ability to form a limited number of bonds.
Polarizability is determined by the ability to change polarity under the influence of an external electric field.
The energy required to break a bond determines its strength.

An example of a covalent nonpolar interaction can be the molecules of hydrogen (H2), chlorine (Cl2), oxygen (O2), nitrogen (N2) and many others.

H· + ·H → H-H molecule has a single non-polar bond,

O: + :O → O=O molecule has a double nonpolar,

Ṅ: + Ṅ: → N≡N the molecule is triple nonpolar.

Examples of covalent bonds of chemical elements include molecules of carbon dioxide (CO2) and carbon monoxide (CO), hydrogen sulfide (H2S), hydrochloric acid (HCL), water (H2O), methane (CH4), sulfur oxide (SO2) and many others .

In the CO2 molecule, the relationship between carbon and oxygen atoms is covalent polar, since the more electronegative hydrogen attracts electron density. Oxygen has two unpaired electrons in its outer shell, while carbon can provide four valence electrons to form the interaction. As a result, double bonds are formed and the molecule looks like this: O=C=O.

In order to determine the type of bond in a particular molecule, it is enough to consider its constituent atoms. Simple metal substances form a metallic bond, metals with nonmetals form an ionic bond, simple nonmetal substances form a covalent nonpolar bond, and molecules consisting of different nonmetals form through a polar covalent bond.

170762 0

Each atom has a certain number of electrons.

When entering into chemical reactions, atoms donate, gain, or share electrons, achieving the most stable electronic configuration. The configuration with the lowest energy (as in noble gas atoms) turns out to be the most stable. This pattern is called the “octet rule” (Fig. 1).

Rice. 1.

This rule applies to everyone types of connections. Electronic connections between atoms allow them to form stable structures, from the simplest crystals to complex biomolecules that ultimately form living systems. They differ from crystals in their continuous metabolism. At the same time, many chemical reactions proceed according to the mechanisms electronic transfer, which play a critical role in energy processes in the body.

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

The nature of a 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 of 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, located in the highest energy orbitals. Accordingly, the outer shell of the atom containing these orbitals is called valence shell. Currently, 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 valence theory, 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 opposite signs, a chemical bond is formed, called by 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 T and II groups of the periodic table and anions of non-metallic elements of groups VI and VII (16 and 17 subgroups, respectively, chalcogens And halogens). The bonds of ionic compounds are unsaturated and non-directional, so they retain the possibility of electrostatic interaction with other ions. In Fig. Figures 2 and 3 show examples of ionic bonds corresponding to the Kossel model of electron transfer.

Rice. 2.

Rice. 3. Ionic bond in a molecule of table salt (NaCl)

Here it is appropriate to recall some properties that explain the behavior of substances in nature, in particular, consider the idea of acids And reasons.

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

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

For example, in the reaction between ammonia N.H. 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 conjugate pairs of acids and bases:

1)N.H. 4+ and N.H. 3

2) HCl And Cl ‑

Here, in each conjugate 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 Brønsted-Lowry theory helps 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 in reactions with aqueous solutions of ammonia, it is an acid.

1) CH 3 COOH + H2O ↔ H3O + + CH 3 COO- . Here, an acetic acid molecule donates a proton to a water molecule;

2) NH 3 + H2O ↔ NH 4 + + HE- . Here, an ammonia molecule accepts a proton from a water molecule.

Thus, water can form two conjugate pairs:

1) H2O(acid) and HE- (conjugate base)

2) H 3 O+ (acid) and H2O(conjugate base).

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

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

Thus, a characteristic property of an ionic bond is the complete movement of the bonding electrons to one of the nuclei. This means that between the ions there is a region 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 everyone atom. In this case, the shared bond electrons are distributed equally between the atoms. Examples of covalent bonds include homonuclear diatomic molecules H 2 , N 2 , F 2. The same type of connection is found in allotropes 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 N 5 HE, sulfur hexafluoride SF 6, acetylene WITH 2 N 2. All these molecules share the same electrons, and their bonds are saturated and directed in the same way (Fig. 4).

It is important for biologists that double and triple bonds have reduced covalent atomic radii compared to a single bond.

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

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

Compounds of two elements located at opposite ends of the same or different periods of the periodic system predominantly form ionic bonds. As elements move closer together within a period, the ionic nature of their compounds decreases, and the covalent character increases. For example, the halides and oxides of 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 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 one more modification.

In polyatomic ions and in complex biological molecules, both electrons can only come from one atom. It's called donor electron pair. An atom that shares 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 d-elements most important for metabolism is largely described by coordination bonds.

Fig. 5.

As a rule, in a complex compound the metal atom acts as an acceptor of an electron pair; 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 Brønsted-Lowry theory. Lewis's 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 electron pair, which, by donating electrons, forms a covalent bond with Lewis acid.

That is, Lewis's 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 capable of accepting an electron pair.

Therefore, according to this theory, the cations are Lewis acids and the anions are Lewis bases. An example would be the following reactions:

It was noted above that the division of substances into ionic and covalent is relative, since complete electron transfer from metal atoms to acceptor atoms does not occur in covalent molecules. In compounds with ionic bonds, 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; for anions it is higher than for cations. The highest polarizability among cations is for cations of higher charge and smaller size, for example, Hg 2+, Cd 2+, Pb 2+, Al 3+, Tl 3+. Has a strong polarizing effect N+ . Since the influence of ion polarization is two-way, it significantly changes the properties of the compounds they form.

The third type of connection isdipole-dipole connection

In addition to the listed types of communication, there are also dipole-dipole intermolecular interactions, also called 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 ( dispersive 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 Debaya(1D = 3.338 × 10‑30 coulomb meters - C × m).

In biochemistry, there is another type of connection - hydrogen connection that 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 similar electronegativity (such as chlorine and sulfur), the hydrogen bond is much weaker. The hydrogen atom is distinguished by one significant feature: when the bonding electrons are pulled away, its nucleus - the proton - is exposed and is no longer shielded 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, to stabilize the structure of proteins in the form of an a-helix, or for the formation of a double helix of DNA (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. 1.

Table 1. Energy of intermolecular forces

Note: The degree of intermolecular interactions is reflected by the enthalpy of melting and evaporation (boiling). Ionic compounds require significantly more energy to separate ions than to separate molecules. The enthalpy of melting of ionic compounds is much higher than that of molecular compounds.

The fourth type of connection ismetal connection

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

From a brief review of bond types, one detail becomes clear: an important parameter of a metal atom or ion - an electron donor, as well as an atom - an electron acceptor, is its 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 increases in groups of the periodic system. 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.

Of greatest importance 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. This connection is called single. Sometimes it is called ordinary, i.e. ordinary. This type of bond is usually denoted by a single line connecting the symbols of interacting atoms.

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

The connection formed by the overlap of the side regions of p-electron clouds is called pi bonds(I am connection). 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 r-bond, a triple bond is one a-bond and two r-bonds.

Let's 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.

Angles between bonds in a molecule ethane WITH. ; N (. almost exactly equal to each other (Fig. 1.18, A) and do not differ from the angles between C-H bonds in a 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 has no 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 With 2 H 4 the angles between the bonds are approximately 120°. From this we can conclude that the $p 2 hybridization of the outer electron orbitals of the carbon atom (Fig. 1.18, b). 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)

pressing one electron at a time. These orbitals are located perpendicular to the plane of the drawing.

The bond energy between carbon atoms in the ethylene molecule C 2 H 4 is equal to -592 kJ/mol. If the carbon atoms were connected 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 is shown in Fig. 1.18 , b depicted with wavy lines. A connection formed in this way is called an I-connection.

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

Acetylene molecule C 2 H 2 linear (Fig. 1.18, V ), 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 l-bonds, i.e. This is 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 benzene molecule C 6 H (.. The benzene molecule is flat, the angles between the bonds of the 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; V- p-inhabited and l-connections between

carbon atoms

120°, which suggests B-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, studying the structure of the benzene molecule shows that all the distances between the 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. In Fig. 1.19, V showing a-bonds between carbon atoms formed by overlapping sp 2 - hybrid orbitals.

Bond energy between atoms carbon in a benzene molecule C 6 H 6 is equal to -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(do not refer to any specific place).