single covalent bond. Types of chemical bonds

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

The simplest one-electron chemical bond is created by a single valence electron. It turns out that one electron is able to hold two positively charged ions in a single whole. In 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 + :

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

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

Characterization of some polar substances

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

The distance between the centers of gravity of positive and negative charges is called the length of the dipole. The polarity of the molecule, as well as the polarity of the bond, is estimated by the value of the dipole moment μ, which is the product of the length of the dipole 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's hybridization for two S- and two p-electrons allowed the directionality of chemical bonds to be explained, in particular the tetrahedral configuration of methane. To explain the structure of ethylene, it is necessary to isolate one p-electron from four equivalent Sp 3 - electrons of the carbon atom to form an additional bond, called the π-bond. In this case, the three remaining Sp 2 -hybrid orbitals are located in the plane at an angle of 120° and form the main bonds, for example, a flat ethylene molecule (Fig. 5).

In Pauling's new theory, all binding electrons became equal and equidistant from the line connecting the nuclei of the molecule. Pauling's theory of a bent chemical bond took into account the statistical interpretation of the wave function by M. Born, the Coulomb electron correlation of electrons. A physical meaning appeared - the nature of the chemical bond is completely determined by the electrical interaction of 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 the chemical bond 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 (borohydrides).

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 "bridge" hydrogen atoms violated the canonical doctrine of valency. The hydrogen atom, previously considered a standard univalent element, turned out to be bound by identical bonds with two boron atoms and became formally a divalent element. The work of W. Lipscomb on deciphering the structure of boranes expanded the understanding of the chemical bond. The Nobel Committee awarded the William Nunn Lipscomb Prize in Chemistry in 1976 with the wording "For his investigations into the structure of boranes (borohydrites) which elucidate the problems of chemical bonds".

Multicenter chemical bond

Fig. 8. Ferrocene molecule

Fig. 9. Dibenzenechromium

Fig. 10. Uranocene

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

Chemical bond dynamics

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

In vapors, these metals consist practically of homonuclear diatomic molecules and free atoms. When metal vapor condenses, the covalent bond turns into a metal one.

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

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

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

The mechanism of the transition of a covalent to a metallic bond

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

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

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

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

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

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

The uneven distribution of electric charges creates conditions for the interaction of molecules due to orientational forces (van der Waals forces). Molecules of alkali metals tend to orient themselves in such a way that opposite electric charges appear in the neighborhood. As a result, attractive forces act between the molecules. Due to the presence of the latter, alkali metal molecules approach each other and are more or less firmly drawn together. At the same time, some deformation of each of them occurs under the action of closer located poles of neighboring molecules (Fig. 12).

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

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

The value of the translational lattice constant a w of an alkali metal crystal significantly exceeds the length of the covalent bond of an alkali metal molecule, therefore it is generally accepted that 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 a geometric place where electrons reside, providing the main property of the 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, a characteristic feature appears in the properties of the metal. So, if weak intermolecular interactions appear during the condensation of hydrogen, then during the condensation of metal vapors, processes characteristic of chemical reactions occur. The condensation of metal vapor itself proceeds in several stages and can be described by the following procession: a free atom → a diatomic molecule with a covalent bond → a metal cluster → a compact metal with a metal bond.

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

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	d EF , Å		d ECl , Å		d EBr , Å		d EI , Å
Li 2 X 2	1,75	105	2,23	108	2,35	110	2,54	116
Na 2 X 2	2,08	95	2,54	105	2,69	108	2,91	111
K2X2	2,35	88	2,86	98	3,02	101	3,26	104
Cs 2 X 2	2,56	79	3,11	91	3,29	94	3,54	94

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

Lattice type and translational lattice constant for alkali metal halides.

In the process of crystallization, a further increase in the interatomic distance occurs, leading to the removal of an electron from the orbital radius of an alkali metal atom and the transfer of an electron to a halogen atom with the formation of the corresponding ions. 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 coordinates by no means one ion with the opposite sign, as it 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 chloride ions (in a sodium chloride crystal) and eight chlorine ions (in a cesium chloride crystal. In this case, all interionic distances in crystals are equidistant.

Notes

Handbook of inorganic chemistry. Constants of inorganic substances. - M .: "Chemistry", 1987. - S. 124. - 320 p.
Lidin R.A., Andreeva L.L., Molochko V.A. Handbook of inorganic chemistry. Constants of inorganic substances. - M .: "Chemistry", 1987. - S. 132-136. - 320 s.
Gankin V.Yu., Gankin Yu.V. How chemical bonds are formed and how chemical reactions proceed. - M .: publishing group "Border", 2007. - 320 p. - ISBN 978-5-94691296-9
Nekrasov B.V. General chemistry course. - M .: Goshimizdat, 1962. - S. 88. - 976 p.
Pauling L. The nature of the chemical bond / edited by Ya.K. Syrkin. - per. from English. M.E. Dyatkina. - M.-L.: Goshimizdat, 1947. - 440 p.
Theoretical organic chemistry / ed. R.Kh. Freidlina. - per. from English. Yu.G. Bundel. - M .: Ed. 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. - S. 63-86.
Chemical Encyclopedic Dictionary / Ch. ed. I.L.Knunyants. - M .: Sov. Encyclopedia, 1983. - S. 607. - 792 p.
Nekrasov B.V. General chemistry course. - M .: Goshimizdat, 1962. - S. 679. - 976 p.
Lidin R.A., Andreeva L.L., Molochko V.A. Handbook of inorganic chemistry. Constants of inorganic substances. - M .: "Chemistry", 1987. - S. 155-161. - 320 s.
Gillespie R. Geometry of molecules / per. from English. E.Z. Zasorina and V.S. Mastryukov, ed. Yu.A. Pentina. - M .: "Mir", 1975. - S. 49. - 278 p.
Handbook of a chemist. - 2nd ed., revised. and additional - L.-M.: GNTI Chemical Literature, 1962. - T. 1. - S. 402-513. - 1072 p.
Lidin R.A., Andreeva L.L., Molochko V.A. Handbook of inorganic chemistry. Constants of inorganic substances .. - M .: "Chemistry", 1987. - S. 132-136. - 320 s.
Zieman J. Electrons in metals (introduction to the theory of Fermi surfaces). Advances in physical sciences .. - 1962. - T. 78, issue 2. - 291 p.

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 acceptor is the boron hydride molecule ВН 3, in which there is a vacant AO.

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.

Simple (single) bond Types of bonds in bioorganic compounds.

Parameter name	Meaning
Article subject:	Simple (single) bond Types of bonds in bioorganic compounds.
Rubric (thematic category)	Chemistry

covalent bond. Multiple connection. non-polar connection. polar connection.

valence electrons. Hybrid (hybridized) orbital. Link length

Keywords.

Characterization of chemical bonds in bioorganic compounds

AROMATICITY

LECTURE 1

CONNECTED SYSTEMS: ACYCLIC AND CYCLIC.

1. Characteristics of chemical bonds in bioorganic compounds. Hybridization of the orbitals of the carbon atom.

2. Classification of conjugate systems: acyclic and cyclic.

3 Types of conjugation: π, π and π, p

4. Criteria for the stability of conjugated systems - ʼʼ conjugation energyʼʼ

5. Acyclic (non-cyclic) conjugate systems, types of conjugation. The main representatives (alkadienes, unsaturated carboxylic acids, vitamin A, carotene, lycopene).

6. Cyclic adjoint systems. Aromatic criteria. Hückel's rule. The role of π-π-, π-ρ-conjugation in the formation of aromatic systems.

7. Carbocyclic aromatic compounds: (benzene, naphthalene, anthracene, phenanthrene, phenol, aniline, benzoic acid) - structure, formation of an aromatic system.

8. Heterocyclic aromatic compounds (pyridine, pyrimidine, pyrrole, purine, imidazole, furan, thiophene) - structure, features of the formation of an aromatic system. Hybridization of electronic orbitals of the nitrogen atom in the formation of five- and six-membered heteroaromatic compounds.

9. Medico-biological significance of natural compounds containing conjugated bond systems, and aromatic.

The initial level of knowledge for mastering the topic (school chemistry course):

Electronic configurations of elements (carbon, oxygen, nitrogen, hydrogen, sulfur, halogens), the concept of ʼʼorbitalʼʼ, hybridization of orbitals and spatial orientation of orbitals of elements of period 2., types of chemical bonds, features of the formation of covalent σ- and π-bonds, changes in the electronegativity of elements in a period and group, classification and principles of the nomenclature of organic compounds.

Organic molecules are formed through covalent bonds. Covalent bonds arise between two atomic nuclei due to a common (socialized) pair of electrons. This method refers to the exchange mechanism. Non-polar and polar bonds are formed.

Non-polar bonds are characterized by a symmetrical distribution of electron density between the two atoms that this bond connects.

Polar bonds are characterized by an asymmetric (non-uniform) distribution of electron density, it shifts towards a more electronegative atom.

Electronegativity series (composed downwards)

A) elements: F> O> N> C1> Br> I ~~ S> C> H

B) carbon atom: C (sp) > C (sp 2) > C (sp 3)

Covalent bonds are of two types: sigma (σ) and pi (π).

In organic molecules, sigma (σ) bonds are formed by electrons located on hybrid (hybridized) orbitals, the electron density is located between atoms on the conditional line of their binding.

π-bonds (pi-bonds) arise when two unhybridized p-orbitals overlap. Their main axes are parallel to each other and perpendicular to the σ-bond line. The combination of σ and π bonds is called a double (multiple) bond, it consists of two pairs of electrons. A triple bond consists of three pairs of electrons - one σ - and two π -bonds. (It is extremely rare in bioorganic compounds).

σ - Bonds are involved in the formation of the skeleton of the molecule, they are the main ones, and π -bonds can be considered as additional, but imparting special chemical properties to molecules.

1.2. Hybridization of the orbitals of the carbon atom 6 C

Electronic configuration of the unexcited state of the carbon atom

expressed by the distribution of electrons 1s 2 2s 2 2p 2.

At the same time, in bioorganic compounds, as well as in most inorganic substances, the carbon atom has a valence equal to four.

There is a transition of one of the 2s electrons to a free 2p orbital. Excited states of the carbon atom arise, creating the possibility of the formation of three hybrid states, denoted as С sp 3 , С sp 2 , С sp .

A hybrid orbital has characteristics different from the "pure" s, p, d orbitals and is a "mixture" of two or more types of unhybridized orbitals.

Hybrid orbitals are characteristic of atoms only in molecules.

The concept of hybridization was introduced in 1931 by L. Pauling, Nobel Prize winner.

Consider the arrangement of hybrid orbitals in space.

C sp 3 --- -- -- ---

In the excited state, 4 equivalent hybrid orbitals are formed. The location of the bonds corresponds to the direction of the central angles of a regular tetrahedron, the angle between any two bonds is equal to 109 0 28 , .

In alkanes and their derivatives (alcohols, haloalkanes, amines), all carbon, oxygen, and nitrogen atoms are in the same sp 3 hybrid state. A carbon atom forms four, a nitrogen atom three, an oxygen atom two covalent σ -connections. Around these bonds, the parts of the molecule can freely rotate relative to each other.

In the excited state sp 2, three equivalent hybrid orbitals arise, the electrons located on them form three σ -bonds that are located in the same plane, the angle between the bonds is 120 0 . Unhybridized 2p - orbitals of two neighboring atoms form π -connection. It is located perpendicular to the plane in which they are σ -connections. The interaction of p-electrons in this case is called ʼʼ lateral overlapʼʼ. A double bond does not allow free rotation of parts of the molecule around itself. The fixed position of the parts of the molecule is accompanied by the formation of two geometric planar isomeric forms, which are called: cis (cis) - and trans (trans) - isomers. (cis- lat- on one side, trans- lat- through).

π -connection

Atoms linked by a double bond are in a state of sp 2 hybridization and

present in alkenes, aromatic compounds, form a carbonyl group

>C=O, azomethine group (imino group) -CH= N-

With sp 2 - --- -- ---

The structural formula of an organic compound is depicted using Lewis structures (each pair of electrons between atoms is replaced by a dash)

C 2 H 6 CH 3 - CH 3 H H

1.3. Polarization of covalent bonds

A covalent polar bond is characterized by an uneven distribution of electron density. Two conditional images are used to indicate the direction of electron density shift.

Polar σ - bond. The electron density shift is indicated by an arrow along the communication line. The end of the arrow points towards the more electronegative atom. The appearance of partial positive and negative charges is indicated using the letter ʼʼ bʼʼ ʼʼ deltaʼʼ with the desired charge sign.

b + b- b+ b + b- b + b-

CH 3 -\u003e O<- Н СН 3 - >C1 CH 3 -\u003e NH 2

methanol chloromethane aminomethane (methylamine)

Polar π bond. The electron density shift is indicated by a semicircular (curved) arrow above the pi bond, also directed towards the more electronegative atom. ()

b + b- b + b-

H 2 C \u003d O CH 3 - C \u003d== O

methanal |

CH 3 propanone -2

1. Determine the type of hybridization of carbon, oxygen, nitrogen atoms in compounds A, B, C. Name the compounds using the IUPAC nomenclature rules.

A. CH 3 -CH 2 - CH 2 -OH B. CH 2 \u003d CH - CH 2 - CH \u003d O

B. CH 3 - N H - C 2 H 5

2. Make the designations characterizing the direction of polarization of all the indicated bonds in the compounds (A - D)

A. CH 3 - Br B. C 2 H 5 - O- H C. CH 3 -NH- C 2 H 5

G. C 2 H 5 - CH \u003d O

Simple (single) bond Types of bonds in bioorganic compounds. - concept and types. Classification and features of the category "Single (single) bond Types of bonds in bioorganic compounds." 2017, 2018.

covalent chemical bond occurs in molecules between atoms due to the formation of common electron pairs. The type of covalent bond can be understood as both the mechanism of its formation and the polarity of the bond. In general, covalent bonds can be classified as follows:

According to the mechanism of formation, a covalent bond can be formed by an exchange or donor-acceptor mechanism.
The polarity of a covalent bond can be non-polar or polar.
According to the multiplicity of the covalent bond, it can be single, double or triple.

This means that a covalent bond in a molecule has three characteristics. For example, in a molecule of hydrogen chloride (HCl), a covalent bond is formed by the exchange mechanism, it is polar and single. In the ammonium cation (NH 4 +), a covalent bond between ammonia (NH 3) and a hydrogen cation (H +) is formed according to the donor-acceptor mechanism, in addition, this bond is polar, is single. In the nitrogen molecule (N 2), the covalent bond is formed by the exchange mechanism, it is non-polar, it is triple.

At exchange mechanism the formation of a covalent bond, each atom has a free electron (or several electrons). Free electrons of different atoms form pairs in the form of a common electron cloud.

At donor-acceptor mechanism the formation of a covalent bond, one atom has a free electron pair, and the other has an empty orbital. The first (donor) gives a pair for common use with the second (acceptor). So in the ammonium cation, nitrogen has a lone pair, and the hydrogen ion has a free orbital.

Non-polar covalent bond formed between atoms of the same chemical element. So in the molecules of hydrogen (H 2), oxygen (O 2), etc., the bond is non-polar. This means that the common electron pair equally belongs to both atoms, since they have the same electronegativity.

Polar covalent bond formed between atoms of different chemical elements. A more electronegative atom displaces an electron pair towards itself. The greater the difference in the electronegativity of the atoms, the more the electrons will be displaced, and the bond will be more polar. So in CH 4, the shift of common electron pairs from hydrogen atoms to carbon atom is not so large, since carbon is not much more electronegative than hydrogen. However, in hydrogen fluoride, the HF bond is highly polar, since the difference in electronegativity between hydrogen and fluorine is significant.

Single covalent bond formed when atoms share the same electron pair double- if two triple- if three. An example of a single covalent bond can be hydrogen molecules (H 2), hydrogen chloride (HCl). An example of a double covalent bond is an oxygen molecule (O 2), where each oxygen atom has two unpaired electrons. An example of a triple covalent bond is a nitrogen molecule (N 2).

covalent bond. Multiple connection. non-polar connection. polar connection.

valence electrons. Hybrid (hybridized) orbital. Link length

Keywords.

Characterization of chemical bonds in bioorganic compounds

AROMATICITY

LECTURE 1

CONNECTED SYSTEMS: ACYCLIC AND CYCLIC.

1. Characterization of chemical bonds in bioorganic compounds. Hybridization of the orbitals of the carbon atom.

2. Classification of conjugate systems: acyclic and cyclic.

3 Types of conjugation: π, π and π, p

4. Criteria for the stability of conjugated systems - "conjugation energy"

5. Acyclic (non-cyclic) conjugate systems, types of conjugation. The main representatives (alkadienes, unsaturated carboxylic acids, vitamin A, carotene, lycopene).

6. Cyclic adjoint systems. Aromatic criteria. Hückel's rule. The role of π-π-, π-ρ-conjugation in the formation of aromatic systems.

7. Carbocyclic aromatic compounds: (benzene, naphthalene, anthracene, phenanthrene, phenol, aniline, benzoic acid) - structure, formation of an aromatic system.

8. Heterocyclic aromatic compounds (pyridine, pyrimidine, pyrrole, purine, imidazole, furan, thiophene) - structure, features of the formation of an aromatic system. Hybridization of electron orbitals of the nitrogen atom in the formation of five- and six-membered heteroaromatic compounds.

9. Medico-biological significance of natural compounds containing conjugated bond systems, and aromatic.

The initial level of knowledge for mastering the topic (school chemistry course):

Electronic configurations of elements (carbon, oxygen, nitrogen, hydrogen, sulfur, halogens), the concept of "orbital", hybridization of orbitals and spatial orientation of the orbitals of elements of the 2nd period., types of chemical bonds, features of the formation of covalent σ- and π-bonds, changes in the electronegativity of elements in period and group, classification and principles of nomenclature of organic compounds.

Non-polar bonds are characterized by a symmetrical distribution of electron density between the two atoms that this bond connects.

Polar bonds are characterized by an asymmetric (non-uniform) distribution of electron density; it shifts towards a more electronegative atom.

Electronegativity series (composed downwards)

A) elements: F> O> N> C1> Br> I ~~ S> C> H

B) carbon atom: C (sp) > C (sp 2) > C (sp 3)

Covalent bonds can be of two types: sigma (σ) and pi (π).

In organic molecules, sigma (σ) bonds are formed by electrons located on hybrid (hybridized) orbitals, the electron density is located between atoms on the conditional line of their binding.

1.2. Hybridization of the orbitals of the carbon atom 6 C

Electronic configuration of the unexcited state of the carbon atom

is expressed by the distribution of electrons 1s 2 2s 2 2p 2 .

However, in bioorganic compounds, as well as in most inorganic substances, the carbon atom has a valency of four.

A hybrid orbital has characteristics different from the "pure" s, p, d orbitals and is a "mixture" of two or more types of unhybridized orbitals.

Hybrid orbitals are characteristic of atoms only in molecules.

The concept of hybridization was introduced in 1931 by L. Pauling, Nobel Prize winner.

Consider the arrangement of hybrid orbitals in space.

C sp 3 --- -- -- ---

In the excited state sp 2, three equivalent hybrid orbitals arise, the electrons located on them form three σ -bonds that are located in the same plane, the angle between the bonds is 120 0 . Unhybridized 2p orbitals of two neighboring atoms form π -connection. It is located perpendicular to the plane in which they are σ -connections. The interaction of p-electrons in this case is called "lateral overlap". A double bond does not allow free rotation of parts of the molecule around itself. The fixed position of the parts of the molecule is accompanied by the formation of two geometric planar isomeric forms, which are called: cis (cis) - and trans (trans) - isomers. (cis- lat- on one side, trans- lat- through).

π -connection

Atoms linked by a double bond are in a state of sp 2 hybridization and

present in alkenes, aromatic compounds, form a carbonyl group

>C=O, azomethine group (imino group) -CH= N-

With sp 2 - --- -- ---

The structural formula of an organic compound is depicted using Lewis structures (each pair of electrons between atoms is replaced by a dash)

C 2 H 6 CH 3 - CH 3 H H

1.3. Polarization of covalent bonds

A covalent polar bond is characterized by an uneven distribution of electron density. Two conditional images are used to indicate the direction of electron density shift.

b + b- b+ b + b- b + b-

CH 3 -\u003e O<- Н СН 3 - >C1 CH 3 -\u003e NH 2

methanol chloromethane aminomethane (methylamine)

Polar π bond. The electron density shift is indicated by a semicircular (curved) arrow above the pi bond, also directed towards the more electronegative atom. ()

b + b- b + b-

H 2 C \u003d O CH 3 - C \u003d== O

methanal |

CH 3 propanone -2

1. Determine the type of hybridization of carbon, oxygen, nitrogen atoms in compounds A, B, C. Name the compounds using the IUPAC nomenclature rules.

A. CH 3 -CH 2 - CH 2 -OH B. CH 2 \u003d CH - CH 2 - CH \u003d O

B. CH 3 - N H - C 2 H 5

2. Make the designations characterizing the direction of polarization of all the indicated bonds in the compounds (A - D)

A. CH 3 - Br B. C 2 H 5 - O- H C. CH 3 -NH- C 2 H 5

Portal for the student. Self-training

single covalent bond. Types of chemical bonds

Multiple covalent bonds

Three-center chemical bond

Multicenter chemical bond

Chemical bond dynamics

The mechanism of the transition of a covalent to a metallic bond

Notes

see also

Portal for the student. Self-training

Multiple covalent bonds

Three-center chemical bond

Multicenter chemical bond

Chemical bond dynamics

The mechanism of the transition of a covalent to a metallic bond

Notes

see also

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