Bonds between atoms, depicted in structural formulas in the form of dashes or sticks (rods), are formed by the interaction of external (valence) electrons of 2 atoms of one molecule. Based on the nature of this interaction, two main or extreme types of bonds between atoms are distinguished.

1st type. Ionic, or electrovalent, or salt bonding is most pronounced (in pure form) is presented in the case of interaction between an atom of a strong metal (for example, an alkali) and an atom of a strong non-metal (for example, a halogen). An alkali metal atom, losing a single external electron, becomes a positively charged particle, and a halogen atom, gaining one external electron, becomes negatively charged: Na + - CI -. The strength of this bond is due to the Coulomb forces of attraction between differently charged particles and the energy released during the formation of a new electron pair at the halogen atom. Examples are salts of organic and inorganic acids.

2nd extreme type. More characteristic of organic compounds is covalent (or atomic) connection brightly (in its pure form) presented in 2-atomic gas molecules: H 2, O 2, N 2, C1 2, etc. in which the bond is formed between completely identical atoms. As a result of the pairing of two electrons of two atoms with opposite spins, when they approach each other, a certain amount of energy is released (≈ 400 kJ/mol) and the new electron pair acquires a molecular orbit, occupying one cell in both atoms. Moreover highest density the electron cloud of this pair is located between the atoms (there is, as it were, an overlap of the atomic orbits of both electrons, Fig. a) or the formation of a molecular orbit - Fig. b).

Although image b) is closer to the truth, both images of covalent bonds are valid and apply. In a purely covalent bond, there is no deviation of the centers of the positive and negative charges of the molecule; they coincide - the molecule is non-polar.

In addition to these two extreme types of bonds (ionic and covalent), there are intermediate types: 3) polar, 4) semipolar, 5) coordination, found mainly in the so-called onium compounds (oxonium, ammonium, sulfonium).

IN polar connection an electron pair located simultaneously in the field of the centers of both atoms (their nuclei) is deviated towards a more electronegative atom, as for example, in the HC1 molecule the electron pair is more in the chlorine cell than in the hydrogen cell:

Because of this deviation in electron density, the centers of positive and negative charges of the molecule diverged. It has become polar, having a dipole moment (i.e. the product of a charge and half the distance between charges) that is not equal to zero.

Semi-polar connection most clearly represented in oxygen compounds of pentavalent nitrogen:

Image a) although allowed, is not valid, because the nitrogen atom has only two electronic levels (layers), where on the outer (2nd) layer there are only four cells and five pairs of electrons (five bonds) will not find a place there. In this regard, image b) is more correct, which shows the resonance of the extreme structures (I and III) and a transition to a more energetically favorable middle structure. Thus, due to the symmetrical distribution of electron density, the negative charge is divided in half between two equal oxygen atoms. But in the nitrogen atom, in fact, there is no whole positive charge, but there is a positive charge close to the whole, respectively, on the oxygen atoms (structure P) there are negative charges close to half the whole charge.

Coordination link is most stable in ammonium compounds, where the nitrogen atom becomes tetravalent, giving up one electron to the proton (and in an aqueous medium to the hydronium cation), acquiring a positive charge (or, in other words, adding a proton to the lone pair of nitrogen electrons): H 3 N: + H + → N 4 N +

ammonia proton ammonium cation

In the ammonium cation, the new bond formed modifies the nature of the three other N - H bonds previously present in nitrogen, which now become more elongated, and all four ammonium hydrogen atoms, mutually repelling, become in the most distant positions from each other, namely at the corners of a regular tetrahedron when the nitrogen atom is in the center of this tetrahedron:

The new bond formed (coordination) is no different from the modified N - H bonds previously present in the ammonia molecule. Here s 1 p 3 takes place - hybridization of the four electrons remaining on nitrogen (like methane carbon).

If the methane molecule is a relatively strong formation that has no charge, then the ammonium cation (which differs from it in structure only in that the nucleus of the central atom contains one more proton and one neutron) is less stable and can easily dissociate into ammonia and hydronium cation, overcoming a certain an energy barrier that determines the strength of ammonium compounds.

The concept of σ-(sigma) and π- (pi) connections

Covalent bonds that predominate in organic compounds generally determine chemical activity these substances. However, weak multiple bonds (double, triple) and the characteristics of bonds in functional groups are of great importance.

Carbon is the central element of the organic world; carbon skeletons (the backbones of molecules) determine their stability and their almost limitless diversity. It is therefore necessary to consider in somewhat more detail the electronic nature of its bonds.

A carbon atom has a nucleus with a charge of +6, and an electron shell: 1s 2, 2s 2, 2p 2, where the four outer electrons are valence electrons (2s 2, 2p 2). But in order for these electrons to form four bonds, the 2s 2 electrons that are in the volume of the ball in a paired form must go into an unpaired state. And the excited carbon atom must have a shell: 1s 2, 2s 1, 2p 3, where, along with the unpaired s-electron (spherical shape), there are three p-electrons (shaped like three-dimensional eights or dumbbells), located along the coordinate axes (x, y, z) three-dimensional space (Fig. 3). However, in order to form four bonds of equal value to each other, as in methane, one s-electron and three p-electrons must go into a modified hybridized (s 1 p 3 -hybridization) state, and all four outer electrons already have exactly the same direction (hybrid) shapes of clouds, and the hydrogen atoms of methane are in equal, most distant positions from each other:

which correspond to the vertices of a regular tetrahedron if a methane carbon atom is placed in its center (Fig. 4). The state of carbon when hybridization of valence electrons takes place in the ratio s 1 p 3 is called the first valence state of carbon, and the bonds of such an atom with other atoms are called b(sigma)-bonds(Fig. 5, 6).

Thus, σ bond is any single bond of carbon to another atom. And most common in organic molecules substances C-H and C-H σ bond have the following basic data (Fig. 6b, 6c).

C-N energy bond ~93-96 kcal/mol (~370-380 kJ/mol).

Communication length 1.1 A 0 (0.11 nm)

S-S energy bond ~84-86 kcal/mol (~340-360 kJ/mol)

Bond length 1.54 A 0 (0.154 nm)

Second valence state of carbon characteristic of ethylene and other compounds with a double bond. In ethylene carbon, hybridization of the valence electrons of the excited atom (2s 1, 2p 3) takes place in the ratio s 1 p 2, when one p-electron (out of three) does not participate in hybridization and remains in the p-form. And hybrid clouds of a directional (elongated) shape are located in a plane at an angle of 120° to each other (Fig. 7).

And the p-electrons of two carbons are paired in the p-form in addition to the σ bond between these carbons (Fig. 7). This additional (multiple) bond formed by the pairing of electrons in the p-form is called π (pi)- communication The energy released during its formation is less than the energy of the σ bond, because the energy of the C = C double bond is ~140 kcal/mol (~560-580 kJ/mol). If we subtract the energy of the C - C σ bond from here (~85 kcal/mol), then π -bond remains ~55 kcal/mol (140-85=55).

Third valence state of carbon characteristic of acetylene and other compounds with a triple bond. For acetylene carbon, out of the four valence electrons (2s 1, 2p 3) of the excited atom, one s- and p-electron each participates in hybridization (s 1 p 1 - hybridization). And the hybrid (elongated) two clouds are located on the same straight line, forming σ-connections (Fig. 8). That is, they occupy the most distant positions (z-coordinate) from the 2 electrons remaining in the p-form, located along the coordinate axes (x, y) of three-dimensional space, paired in the p-form to form two π - connections in mutually perpendicular planes(Fig. 8). The energy released during the formation of a triple bond is ~200 kcal/mol. If we subtract 85 kcal/mol from here - the energy of the σ bond, then ~115 kcal/m remains for two π bonds, i.e. ~57 kcal/mol for each π bond. Here are the main characteristics of single, double and triple bonds for comparison:

C - C bond length 1.54 A 0, bond formation energy ~85 kcal/mol

C = C bond length 1.34 A 0 , bond formation energy ~140 kcal/mol

C ≡ C bond length 1.21 A 0, bond formation energy ~200 kcal/mol Electrons of π bonds are more voluminous, weaker bound to atomic nuclei, more accessible to attack by the reagent, and are easily exposed to electric and magnetic fields, for example, incident light or attack by charged particles. That's why π -bonds, having a completely different nature compared to σ-bonds, are less stable and cause high chemical activity of unsaturated compounds compared to saturated (saturated) compounds.

CHAPTER 2. CHEMICAL BONDING AND MUTUAL INFLUENCE OF ATOMS IN ORGANIC COMPOUNDS

CHAPTER 2. CHEMICAL BONDING AND MUTUAL INFLUENCE OF ATOMS IN ORGANIC COMPOUNDS

The chemical properties of organic compounds are determined by the type chemical bonds, the nature of the bonded atoms and their mutual influence in the molecule. These factors, in turn, are determined by the electronic structure of atoms and the interaction of their atomic orbitals.

2.1. Electronic structure of the carbon atom

The part of the atomic space in which the probability of finding an electron is maximum is called an atomic orbital (AO).

In chemistry, the concept of hybrid orbitals of the carbon atom and other elements is widely used. The concept of hybridization as a way of describing the rearrangement of orbitals is necessary when the number of unpaired electrons in the ground state of an atom is less than the number of bonds formed. An example is the carbon atom, which in all compounds manifests itself as a tetravalent element, but in accordance with the rules for filling orbitals, its outer electronic level in the ground state 1s 2 2s 2 2p 2 contains only two unpaired electrons (Fig. 2.1, A and Appendix 2-1). In these cases, it is postulated that different atomic orbitals, similar in energy, can mix with each other, forming hybrid orbitals of the same shape and energy.

Hybridized orbitals, due to greater overlap, form stronger bonds compared to non-hybridized orbitals.

Depending on the number of orbitals that have entered into hybridization, a carbon atom can be in one of three states

Rice. 2.1.Distribution of electrons over orbitals of a carbon atom in the ground (a), excited (b) and hybridized states (c - sp3, g-sp2, d-sp)

hybridization (see Fig. 2.1, c-e). The type of hybridization determines the orientation of hybrid AOs in space and, consequently, the geometry of the molecules, i.e., their spatial structure.

The spatial structure of molecules is the relative arrangement of atoms and atomic groups in space.

sp 3-Hybridization.When four external AOs of an excited carbon atom (see Fig. 2.1, b) - one 2s and three 2p orbitals - are mixed, four equivalent sp 3 hybrid orbitals arise. They have the shape of a three-dimensional “eight”, one of the blades of which is much larger than the other.

Each hybrid orbital is filled with one electron. The carbon atom in the state of sp 3 hybridization has the electronic configuration 1s 2 2(sp 3) 4 (see Fig. 2.1, c). This state of hybridization is characteristic of carbon atoms in saturated hydrocarbons (alkanes) and, accordingly, in alkyl radicals.

Due to mutual repulsion, sp 3 -hybrid AOs are directed in space towards the vertices tetrahedron, and the angles between them are 109.5? (the most advantageous location; Fig. 2.2, a).

The spatial structure is depicted using stereochemical formulas. In these formulas, the sp 3 -hybridized carbon atom and its two bonds are placed in the plane of the drawing and graphically indicated by a regular line. A thick line or a thick wedge indicates a connection extending forward from the plane of the drawing and directed towards the observer; dotted line or shaded wedge (..........) - a connection extending from the observer beyond the plane of the drawing -

Rice. 2.2.Types of carbon atom hybridization. The point in the center is the atomic nucleus (small fractions of hybrid orbitals are omitted to simplify the figure; non-hybridized p-AOs are shown in color)

lady (Fig. 2.3, a). The carbon atom is in the state sp 3-hybridization has a tetrahedral configuration.

sp 2-Hybridization.When mixing one 2s- and two 2p-AOs of an excited carbon atom, three equivalent ones are formed sp 2-hybrid orbitals and remains unhybridized 2p-AO. The carbon atom is in the state sp 2-hybridization has the electronic configuration 1s 2 2(sp 2) 3 2p 1 (see Fig. 2.1, d). This state of carbon atom hybridization is typical for unsaturated hydrocarbons (alkenes), as well as for some functional groups, such as carbonyl and carboxyl.

sp 2 -Hybridized orbitals are located in the same plane at an angle of 120?, and the non-hybridized AO is in a perpendicular plane (see Fig. 2.2, b). The carbon atom is in the state sp 2-hybridization has trigonal configuration. Carbon atoms connected by a double bond are in the plane of the drawing, and their single bonds directed towards and away from the observer are designated as described above (see Fig. 2.3, b).

sp-Hybridization.When one 2s- and one 2p-orbitals of an excited carbon atom are mixed, two equivalent sp-hybrid AOs are formed, and two p-AOs remain unhybridized. The carbon atom in the sp-hybridized state has an electronic configuration

Rice. 2.3.Stereochemical formulas of methane (a), ethane (b) and acetylene (c)

1s 2 2(sp 2) 2 2p 2 (see Fig. 2.1, d). This state of hybridization of the carbon atom occurs in compounds that have a triple bond, for example, in alkynes and nitriles.

sp-Hybridized orbitals are located at an angle of 180°, and two non-hybridized AOs are located in mutually perpendicular planes (see Fig. 2.2, c). The carbon atom in the sp-hybridized state has linear configuration for example, in an acetylene molecule, all four atoms are on the same straight line (see Fig. 2.3, V).

Atoms of other organogenic elements may also be in a hybridized state.

2.2. Chemical bonds of a carbon atom

Chemical bonds in organic compounds are represented mainly by covalent bonds.

A covalent bond is a chemical bond formed as a result of the sharing of electrons between bonded atoms.

These shared electrons occupy molecular orbitals (MOs). As a rule, a MO is a multicenter orbital and the electrons filling it are delocalized (dispersed). Thus, a MO, like an AO, can be vacant, filled with one electron or two electrons with opposite spins*.

2.2.1. σ- Andπ -Connections

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

A σ-bond is a covalent bond formed when an AO overlaps along a straight line (axis) connecting the nuclei of two bonded atoms with a maximum overlap on this straight line.

The σ-bond occurs when any AO, including hybrid ones, overlap. Figure 2.4 shows the formation of a σ bond between carbon atoms as a result of the axial overlap of their hybrid sp 3 -AO and σ bonds C-H way overlap of hybrid sp 3 -AO of carbon and s-AO of hydrogen.

* For more details see: Popkov V.A., Puzakov S.A. General chemistry. - M.: GEOTAR-Media, 2007. - Chapter 1.

Rice. 2.4.Formation of σ bonds in ethane by axial overlap of AOs (small fractions of hybrid orbitals are omitted and shown in color sp 3 -AO carbon, black - s-AO hydrogen)

In addition to axial overlap, another type of overlap is possible - lateral overlap of p-AO, leading to the formation of a π bond (Fig. 2.5).

p-atomic orbitals

Rice. 2.5.Formation of π bond in ethylene by lateral overlap r-AO

A π-bond is a bond formed by the lateral overlap of unhybridized p-AOs with a maximum overlap on both sides of the straight line connecting the nuclei of atoms.

Multiple bonds found in organic compounds are a combination of σ- and π-bonds: double - one σ- and one π-, triple - one σ- and two π-bonds.

The properties of a covalent bond are expressed through characteristics such as energy, length, polarity and polarizability.

Communication energyis the energy released when a bond is formed or required to separate two bonded atoms. It serves as a measure of the strength of the bond: the higher the energy, the stronger the bond (Table 2.1).

Link lengthis the distance between the centers of bonded atoms. A double bond is shorter than a single bond, and a triple bond is shorter than a double bond (see Table 2.1). Bonds between carbon atoms in different states of hybridization have a common pattern -

Table 2.1.Basic characteristics of covalent bonds

As the fraction of the s orbital in the hybrid orbital increases, the bond length decreases. For example, in a series of compounds propane CH 3 CH 2 CH 3, propene CH 3 CH=CH 2, propyne CH 3 C=CH bond length CH 3 -C is correspondingly equal to 0.154; 0.150 and 0.146 nm.

Communication polarity due to uneven distribution (polarization) of electron density. The polarity of a molecule is quantified by the value of its dipole moment. From the dipole moments of a molecule, the dipole moments of individual bonds can be calculated (see Table 2.1). The larger the dipole moment, the more polar the bond. The reason for bond polarity is the difference in electronegativity of the bonded atoms.

Electronegativity characterizes the ability of an atom in a molecule to hold valence electrons. As the electronegativity of an atom increases, the degree of displacement of bond electrons in its direction increases.

Based on the values ​​of bond energy, the American chemist L. Pauling (1901-1994) proposed a quantitative characteristic of the relative electronegativity of atoms (Pauling scale). In this scale (series), typical organogen elements are arranged according to relative electronegativity (two metals are given for comparison) as follows:

Electronegativity is not an absolute constant of an element. It depends on the effective charge of the nucleus, the type of AO hybridization and the influence of substituents. For example, the electronegativity of a carbon atom in the sp 2 or sp hybridization state is higher than in the sp 3 hybridization state, which is associated with an increase in the proportion of the s orbital in the hybrid orbital. During the transition of atoms from sp 3 - to sp 2 - and further to sp-hybridized state, the extent of the hybrid orbital gradually decreases (especially in the direction that provides the greatest overlap during the formation of a σ bond), which means that in the same sequence the maximum electron density is located closer and closer to the nucleus of the corresponding atom.

In the case of a non-polar or practically non-polar covalent bond, the difference in the electronegativity of the bonded atoms is zero or close to zero. As the difference in electronegativity increases, the polarity of the bond increases. A difference of up to 0.4 is said to be weakly polar, more than 0.5 is a strongly polar covalent bond, and more than 2.0 is an ionic bond. Polar covalent bonds are prone to heterolytic cleavage

(see 3.1.1).

Bond polarizability is expressed in the displacement of bond electrons under the influence of external electric field, including another reacting particle. Polarizability is determined by electron mobility. Electrons are more mobile the further they are from the nuclei of atoms. In terms of polarizability, the π bond is significantly superior to the σ bond, since the maximum electron density of the π bond is located further from the bonded nuclei. Polarizability largely determines the reactivity of molecules towards polar reagents.

2.2.2. Donor-acceptor bonds

The overlap of two one-electron AOs is not the only way to form a covalent bond. A covalent bond can be formed by the interaction of a two-electron orbital of one atom (donor) with a vacant orbital of another atom (acceptor). Donors are compounds containing either orbitals with a lone pair of electrons or π-MO. Carriers of lone pairs of electrons (n-electrons, from English. non-bonding) are atoms of nitrogen, oxygen, halogens.

Lone pairs of electrons play an important role in the manifestation of the chemical properties of compounds. In particular, they are responsible for the ability of compounds to enter into donor-acceptor interactions.

A covalent bond formed by a pair of electrons from one of the bonding partners is called a donor-acceptor bond.

The resulting donor-acceptor bond differs only in the method of formation; its properties are identical to other covalent bonds. The donor atom thereby acquires a positive charge.

Donor-acceptor bonds are characteristic of complex compounds.

2.2.3. Hydrogen bonds

A hydrogen atom bonded to a strongly electronegative element (nitrogen, oxygen, fluorine, etc.) is capable of interacting with the lone pair of electrons of another sufficiently electronegative atom of the same or another molecule. As a result, a hydrogen bond arises, which is a type of donor bond.

acceptor bond. Graphically, a hydrogen bond is usually represented by three dots.

The hydrogen bond energy is low (10-40 kJ/mol) and is mainly determined by electrostatic interaction.

Intermolecular hydrogen bonds determine the association of organic compounds, such as alcohols.

Hydrogen bonds affect the physical (boiling and melting points, viscosity, spectral characteristics) and chemical (acid-base) properties of compounds. Thus, the boiling point of ethanol is C 2 H 5 OH (78.3°C) is significantly higher than dimethyl ether CH 3 OCH 3 (-24°C), which has the same molecular weight and is not associated through hydrogen bonds.

Hydrogen bonds can also be intramolecular. This bond in the salicylic acid anion leads to an increase in its acidity.

Hydrogen bonds play an important role in the formation of the spatial structure of high-molecular compounds - proteins, polysaccharides, nucleic acids.

2.3. Conjugate systems

A covalent bond can be localized or delocalized. A localized bond is one whose electrons are actually shared between the two nuclei of the bonded atoms. If the bonding electrons are shared between more than two nuclei, then they speak of a delocalized bond.

A delocalized bond is a covalent bond whose molecular orbital spans more than two atoms.

Delocalized bonds are in most cases π bonds. They are characteristic of coupled systems. In these systems, a special type of mutual influence of atoms occurs—conjugation.

Conjugation (mesomerism, from Greek. mesos- average) is the alignment of bonds and charges in a real molecule (particle) in comparison with an ideal, but non-existent structure.

The delocalized p-orbitals involved in conjugation can belong to either two or more π-bonds, or a π-bond and one atom with a p-orbital. In accordance with this, a distinction is made between π,π-conjugation and ρ,π-conjugation. The conjugation system can be open or closed and contain not only carbon atoms, but also heteroatoms.

2.3.1. Open circuit systems

π,π -Pairing. The simplest representative of π,π-conjugated systems with a carbon chain is butadiene-1,3 (Fig. 2.6, a). The carbon and hydrogen atoms and, therefore, all σ bonds in its molecule lie in the same plane, forming a flat σ skeleton. Carbon atoms are in a state of sp 2 hybridization. The unhybridized p-AOs of each carbon atom are located perpendicular to the plane of the σ-skeleton and parallel to each other, which is a necessary condition for their overlap. Overlap occurs not only between p-AO of atoms C-1 and C-2, C-3 and C-4, but also between p-AO of atoms C-2 and C-3, resulting in the formation of a single π covering four carbon atoms -system, i.e., a delocalized covalent bond appears (see Fig. 2.6, b).

Rice. 2.6.Atomic orbital model of the 1,3 butadiene molecule

This is reflected in changes in bond lengths in the molecule. The length of the C-1-C-2 as well as C-3-C-4 bonds in 1,3-butadiene is slightly increased, and the distance between C-2 and C-3 is shortened compared to conventional double and single bonds. In other words, the process of electron delocalization leads to equalization of bond lengths.

Hydrocarbons with a large number of conjugated double bonds are common in the plant world. These include, for example, carotenes, which determine the color of carrots, tomatoes, etc.

An open conjugation system can also include heteroatoms. An example of open π,π-conjugated systems with a heteroatom in the chainα,β-unsaturated carbonyl compounds can serve. For example, the aldehyde group in acrolein CH 2 =CH-CH=O is a participant in the conjugation chain of three sp 2 -hybridized carbon atoms and an oxygen atom. Each of these atoms contributes one p-electron to a single π-system.

pn-Pairing.This type of conjugation most often occurs in compounds containing the structural fragment -CH=CH-X, where X is a heteroatom having a lone pair of electrons (primarily O or N). These include, for example, vinyl ethers, in the molecules of which the double bond is conjugated with r-orbital of the oxygen atom. A delocalized three-center bond is formed by overlapping two p-AO sp 2 -hybridized carbon atoms and one r-AO of a heteroatom with a pair of n-electrons.

The formation of a similar delocalized three-center bond occurs in the carboxyl group. Here, the π-electrons of the C=O bond and the n-electrons of the oxygen atom of the OH group participate in conjugation. Conjugated systems with fully aligned bonds and charges include negatively charged species, such as the acetate ion.

The direction of electron density shift is indicated by a curved arrow.

There are other graphical ways to display pairing results. Thus, the structure of the acetate ion (I) assumes that the charge is evenly distributed over both oxygen atoms (as shown in Fig. 2.7, which is true).

Structures (II) and (III) are used in resonance theory. According to this theory, a real molecule or particle is described by a set of certain so-called resonance structures, which differ from each other only in the distribution of electrons. In conjugated systems, the main contribution to the resonance hybrid is made by structures with different distributions of π-electron density (the double-sided arrow connecting these structures is a special symbol of resonance theory).

Limit (boundary) structures do not really exist. However, to one degree or another, they “contribute” to the real distribution of electron density in a molecule (particle), which is represented as a resonant hybrid obtained by superposition of limiting structures.

In ρ,π-conjugated systems with a carbon chain, conjugation can occur if there is a carbon atom with a non-hybridized p-orbital next to the π bond. Such systems can be intermediate particles - carbanions, carbocations, free radicals, for example, with an allylic structure. Free radical allylic moieties play an important role in the processes of lipid peroxidation.

In the allyl anion CH 2 =CH-CH 2 sp 2 -hybridized carbon atom C-3 supplies to the common conjugate

Rice. 2.7.Electron density map of the COONa group in penicillin

system two electrons, in the allylic radical CH 2 =CH-CH 2+ - one, and in the allylic carbocation CH 2 =CH-CH 2+ does not supply any. As a result, when the p-AO of three sp 2 -hybridized carbon atoms overlaps, a delocalized three-center bond is formed containing four (in the carbanion), three (in free radical) and two (in the carbocation) electrons, respectively.

Formally, the C-3 atom in the allyl cation carries a positive charge, in the allyl radical it carries an unpaired electron, and in the allyl anion it carries a negative charge. In fact, in such conjugated systems there is delocalization (dispersal) of the electron density, which leads to the alignment of bonds and charges. The C-1 and C-3 atoms in these systems are equivalent. For example, in an allyl cation, each of them carries a positive charge+1/2 and is connected by a one-and-a-half bond to the C-2 atom.

Thus, conjugation results in a significant difference in the electron density distribution in real structures compared to the structures depicted by conventional structure formulas.

2.3.2. Closed-loop systems

Cyclic conjugated systems are of great interest as a group of compounds with increased thermodynamic stability compared to conjugated ones. open systems. These compounds also have other special properties, the totality of which is united by the general concept aromaticity. These include the ability of such formally unsaturated compounds

engage in substitution reactions rather than addition, resistance to oxidizing agents and temperature.

Typical representatives of aromatic systems are arenes and their derivatives. The peculiarities of the electronic structure of aromatic hydrocarbons are clearly manifested in the atomic orbital model of the benzene molecule. The benzene framework is formed by six sp 2 -hybridized carbon atoms. All σ bonds (C-C and C-H) lie in the same plane. Six unhybridized p-AOs are located perpendicular to the plane of the molecule and parallel to each other (Fig. 2.8, a). Each r-AO can equally overlap with two neighboring ones r-AO. As a result of such overlap, a single delocalized π-system arises, the highest electron density in which is located above and below the plane of the σ-skeleton and covers all the carbon atoms of the cycle (see Fig. 2.8, b). The π-Electron density is evenly distributed throughout the cyclic system, which is indicated by a circle or dotted line inside the cycle (see Fig. 2.8, c). All bonds between carbon atoms in the benzene ring have the same length (0.139 nm), intermediate between the lengths of single and double bonds.

Based on quantum mechanical calculations, it was established that for the formation of such stable molecules, a flat cyclic system must contain (4n + 2) π-electrons, where n= 1, 2, 3, etc. (Hückel’s rule, 1931). Taking these data into account, the concept of “aromaticity” can be specified.

A compound is aromatic if it has a planar ring and a conjugateπ -electronic system covering all atoms of the cycle and containing(4n+ 2) π-electrons.

Hückel's rule applies to any planar condensed systems in which there are no atoms shared by more than

Rice. 2.8.Atomic orbital model of the benzene molecule (hydrogen atoms omitted; explanation in text)

two cycles. Compounds with condensed benzene rings, such as naphthalene and others, meet the criteria for aromaticity.

Stability of coupled systems. The formation of a conjugated and especially aromatic system is an energetically favorable process, since this increases the degree of overlap of orbitals and delocalization (dispersal) occurs. r-electrons. In this regard, conjugated and aromatic systems have increased thermodynamic stability. They contain a smaller supply of internal energy and in the ground state occupy a lower energy level compared to non-conjugated systems. From the difference between these levels, one can quantify the thermodynamic stability of the conjugated compound, i.e., its conjugation energy(delocalization energy). For butadiene-1,3 it is small and amounts to about 15 kJ/mol. As the length of the conjugated chain increases, the conjugation energy and, accordingly, the thermodynamic stability of the compounds increase. The conjugation energy for benzene is much higher and amounts to 150 kJ/mol.

2.4. Electronic effects of substituents 2.4.1. Inductive effect

A polar σ bond in a molecule causes polarization of nearby σ bonds and leads to the appearance of partial charges on neighboring atoms*.

Substituents cause polarization not only of their own, but also of neighboring σ-bonds. This type of transfer of influence of atoms is called the inductive effect (/-effect).

The inductive effect is the transfer of the electronic influence of substituents as a result of the displacement of electrons of σ bonds.

Due to the weak polarizability of the σ bond, the inductive effect fades after three or four bonds in the circuit. Its effect is most pronounced in relation to the carbon atom adjacent to the one that has a substituent. The direction of the inductive effect of the substituent is qualitatively assessed by comparing it with the hydrogen atom, the inductive effect of which is taken to be zero. Graphically, the result of the /-effect is represented by an arrow coinciding with the position of the valence line and pointing towards the more electronegative atom.

/V\stronger than the hydrogen atom, exhibitsnegativeinductive effect (-/- effect).

Such substituents generally reduce the electron density of the system; they are called electron-withdrawing. These include most functional groups: OH, NH 2, COOH, NO 2 and cationic groups, for example -NH 3+.

A substituent that shifts the electron density compared to the hydrogen atomσ -bonds towards the carbon atom of the chain, exhibitspositiveinductive effect (+/- effect).

Such substituents increase the electron density in the chain (or ring) and are called electron donor. These include alkyl groups located at the sp 2 -hybridized carbon atom, and anionic centers in charged particles, for example -O -.

2.4.2. Mesomeric effect

In conjugated systems, the π-electrons of delocalized covalent bonds play the main role in the transmission of electronic influence. The effect manifested in a shift in the electron density of a delocalized (conjugated) π-system is called the mesomeric (M-effect), or conjugation effect.

The mesomeric effect is the transfer of the electronic influence of substituents through a conjugated system.

In this case, the deputy himself is a participant in the coupled system. It can introduce into the conjugation system either a π-bond (carbonyl, carboxyl groups, etc.), or a lone pair of heteroatom electrons (amino and hydroxy groups), or a vacant or one-electron-filled p-AO.

A substituent that increases the electron density in a conjugated system exhibitspositivemesomeric effect (+M- effect).

The M-effect is exhibited by substituents that include atoms with a lone pair of electrons (for example, an amino group in an aniline molecule) or an entire negative charge. These substituents are capable

to the transfer of a pair of electrons to a common conjugate system, i.e. they are electron donor.

A substituent that lowers the electron density in a conjugated system exhibitsnegativemesomeric effect (-M- effect).

The M-effect in a conjugated system is caused by oxygen or nitrogen atoms linked by a double bond to a carbon atom, as shown in the example of acrylic acid and benzaldehyde. Such groups are electron-withdrawing.

An electron density shift is indicated by a curved arrow, the beginning of which shows which p or π electrons are displaced, and the end of which shows the bond or atom to which they are displaced. The mesomeric effect, in contrast to the inductive effect, is transmitted through a system of conjugated bonds over a much greater distance.

When assessing the influence of substituents on the distribution of electron density in a molecule, it is necessary to take into account the resulting effect of inductive and mesomeric effects (Table 2.2).

Table 2.2.Electronic effects of some substituents

Electronic effects of substituents make it possible to qualitatively assess the distribution of electron density in a non-reacting molecule and predict its properties.

Continuation. See the beginning in № 15, 16/2004

Lesson 5. Hybridization
carbon atomic orbitals

A covalent chemical bond is formed using shared bonding electron pairs like:

Form a chemical bond, i.e. Only unpaired electrons can create a common electron pair with a “foreign” electron from another atom. When writing electronic formulas, unpaired electrons are located one at a time in an orbital cell.
Atomic orbital is a function that describes the density of the electron cloud at each point in space around the atomic nucleus. An electron cloud is a region of space in which an electron can be detected with a high probability.
To harmonize the electronic structure of the carbon atom and the valency of this element, concepts about the excitation of the carbon atom are used. In the normal (unexcited) state, the carbon atom has two unpaired 2 r 2 electrons. In an excited state (when energy is absorbed) one of 2 s 2 electrons can go to free r-orbital. Then four unpaired electrons appear in the carbon atom:

Let us recall that in the electronic formula of an atom (for example, for carbon 6 C – 1 s 2 2s 2 2p 2) large numbers in front of the letters - 1, 2 - indicate the number of the energy level. Letters s And r indicate the shape of the electron cloud (orbital), and the numbers to the right above the letters indicate the number of electrons in a given orbital. All s-spherical orbitals:

At the second energy level except 2 s-there are three orbitals 2 r-orbitals. These 2 r-orbitals have an ellipsoidal shape, similar to dumbbells, and are oriented in space at an angle of 90° to each other. 2 r-Orbitals denote 2 p x, 2p y and 2 p z in accordance with the axes along which these orbitals are located.

When chemical bonds are formed, the electron orbitals acquire the same shape. Thus, in saturated hydrocarbons one s-orbital and three r-orbitals of the carbon atom to form four identical (hybrid) sp 3-orbitals:

This - sp 3 -hybridization.
Hybridization– alignment (mixing) of atomic orbitals ( s And r) with the formation of new atomic orbitals called hybrid orbitals.

Hybrid orbitals have an asymmetric shape, elongated towards the attached atom. Electron clouds repel each other and are located in space as far as possible from each other. In this case, the axes of four sp 3-hybrid orbitals turn out to be directed towards the vertices of the tetrahedron (regular triangular pyramid).
Accordingly, the angles between these orbitals are tetrahedral, equal to 109°28".
The vertices of electron orbitals can overlap with the orbitals of other atoms. If electron clouds overlap along a line connecting the centers of atoms, then such a covalent bond is called sigma()-connection. For example, in the ethane molecule C 2 H 6, a chemical bond is formed between two carbon atoms by overlapping two hybrid orbitals. This is a connection. In addition, each of the carbon atoms with its three sp 3-orbitals overlap with s-orbitals of three hydrogen atoms, forming three -bonds.

In total, three valence states with different types of hybridization are possible for a carbon atom. Except sp 3-hybridization exists sp 2 - and sp-hybridization.
sp 2 -Hybridization- mixing one s- and two r-orbitals. As a result, three hybrids are formed sp 2 -orbitals. These sp 2-orbitals are located in the same plane (with axes X, at) and are directed to the vertices of the triangle with an angle between the orbitals of 120°. Unhybridized
r-the orbital is perpendicular to the plane of the three hybrid sp 2-orbitals (oriented along the axis z). Upper half r-orbitals are above the plane, the lower half is below the plane.
Type sp 2-carbon hybridization occurs in compounds with a double bond: C=C, C=O, C=N. Moreover, only one of the bonds between two atoms (for example, C=C) can be an - bond. (The other bonding orbitals of the atom are directed in opposite directions.) The second bond is formed as a result of overlapping non-hybrid r-orbitals on both sides of the line connecting the atomic nuclei.

Covalent bond formed by lateral overlap r-orbitals of neighboring carbon atoms is called pi()-connection.

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Due to less orbital overlap, the -bond is less strong than the -bond.
sp-Hybridization– this is mixing (alignment in shape and energy) of one s- and one
r-orbitals to form two hybrid sp-orbitals. sp-The orbitals are located on the same line (at an angle of 180°) and directed in opposite directions from the nucleus of the carbon atom. Two
r-orbitals remain unhybridized. They are placed mutually perpendicular
directions of connections. In the picture sp-orbitals are shown along the axis y, and the unhybridized two
r-orbitals – along the axes X And z.

A carbon-carbon triple bond CC consists of an -bond formed by overlapping
sp-hybrid orbitals, and two -bonds.
The relationship between such parameters of the carbon atom as the number of attached groups, the type of hybridization and the types of chemical bonds formed is shown in Table 4.

Table 4

Covalent carbon bonds

Number of groups related with carbon	Type hybridization	Types participating chemical bonds	Examples of compound formulas
4	sp 3	Four - connections
3	sp 2	Three - connections and one - connection
2	sp	Two - connections and two - connections	H–CC–H

Exercises.

1. Which electrons of atoms (for example, carbon or nitrogen) are called unpaired?

2. What does the concept of “shared electron pairs” mean in compounds with a covalent bond (for example, CH 4 or H 2 S )?

3. What electronic states of atoms (for example, C or N ) are called basic, and which are excited?

4. What do the numbers and letters mean in the electronic formula of an atom (for example, C or N )?

5. What is an atomic orbital? How many orbitals are there in the second energy level of the C atom? and how do they differ?

6. How are hybrid orbitals different from the original orbitals from which they were formed?

7. What types of hybridization are known for the carbon atom and what do they consist of?

8. Draw a picture of the spatial arrangement of orbitals for one of the electronic states of the carbon atom.

9. What chemical bonds are called and what? Specify-And-connections in connections:

10. For the carbon atoms of the compounds below, indicate: a) type of hybridization; b) types of its chemical bonds; c) bond angles.

Answers to exercises for topic 1

Lesson 5

1. Electrons that are located one at a time in an orbital are called unpaired electrons. For example, in the electron diffraction formula of an excited carbon atom there are four unpaired electrons, and the nitrogen atom has three:

2. Two electrons involved in the formation of one chemical bond are called shared electron pair. Typically, before a chemical bond is formed, one of the electrons in this pair belonged to one atom, and the other electron belonged to another atom:

3. The electronic state of an atom in which the order of filling the electron orbitals is observed: 1 s 2 , 2s 2 , 2p 2 , 3s 2 , 3p 2 , 4s 2 , 3d 2 , 4p 2, etc., are called underlying condition. IN excited state one of the valence electrons of the atom occupies a free orbital with a higher energy; such a transition is accompanied by the separation of paired electrons. Schematically it is written like this:

While in the ground state there were only two unpaired valence electrons, in the excited state there are four such electrons.

5. An atomic orbital is a function that describes the density of the electron cloud at each point in space around the nucleus of a given atom. At the second energy level of the carbon atom there are four orbitals - 2 s, 2p x, 2p y, 2p z. These orbitals differ:
a) the shape of the electron cloud ( s– ball, r– dumbbell);
b) r-orbitals have different orientations in space - along mutually perpendicular axes x, y And z, they are designated p x, p y, p z.

6. Hybrid orbitals differ from the original (non-hybrid) orbitals in shape and energy. For example, s-orbital – the shape of a sphere, r– symmetrical figure eight, sp-hybrid orbital – asymmetric figure eight.
Energy differences: E(s) < E(sp) < E(r). Thus, sp-orbital – an orbital averaged in shape and energy, obtained by mixing the original s- And p-orbitals.

7. For a carbon atom, three types of hybridization are known: sp 3 , sp 2 and sp (see text of lesson 5).

9. -bond - a covalent bond formed by head-on overlapping of orbitals along a line connecting the centers of atoms.
-bond – a covalent bond formed by lateral overlap r-orbitals on both sides of the line connecting the centers of the atoms.
-Bonds are shown by the second and third lines between connected atoms.

In the ground state, carbon atom C (1s 2 2s 2 2p 2) has two unpaired electrons, due to which only two common electron pairs can be formed. However, in most of its compounds, carbon is tetravalent. This is explained by the fact that the carbon atom, absorbing a small amount of energy, goes into an excited state in which it has 4 unpaired electrons, i.e. capable of forming four covalent bonds and take part in the formation of four common electron pairs:

6 С 1 s 2 2s 2 2 p 2 6 С * 1 s 2 2s 1 2 p 3

↓

1

↓

p

↓

p

s

s

The excitation energy is compensated by the formation of chemical bonds, which occurs with the release of energy.

Carbon atoms have the ability to form three types of hybridization of electron orbitals ( sp 3, sp 2, sp) and the formation of multiple (double and triple) bonds among themselves (Table 7).

Table 7

Types of hybridization and molecular geometry

Simple (single) s - communication is carried out at sp 3-hybridization, in which all four hybrid orbitals are equivalent and are directed in space at an angle of 109 o 29 'to each other and oriented to the vertices of a regular tetrahedron.

Rice. 19. Formation of a methane molecule CH 4

If hybrid carbon orbitals overlap with spherical ones s-orbitals of the hydrogen atom, then the simplest organic compound methane CH 4 is formed - a saturated hydrocarbon (Fig. 19).

Rice. 20. Tetrahedral arrangement of bonds in the methane molecule

Of great interest is the study of the bonds of carbon atoms with each other and with atoms of other elements. Let's consider the structure of the molecules of ethane, ethylene and acetylene.

The angles between all bonds in the ethane molecule are almost exactly equal to each other (Fig. 21) and do not differ from angles C-H in a methane molecule.

Rice. 21. Ethane molecule C 2 H 6

Therefore, the carbon atoms are in a state sp 3-hybridization.

The hybridization of the electronic orbitals of carbon atoms may be incomplete, i.e. two ( sp 2–hybridization) or one ( sp-hybridization) of three r- orbitals. In this case, between the carbon atoms there are formed multiples(double or triple) communications. Hydrocarbons with multiple bonds are called unsaturated or unsaturated. A double bond (C=C) is formed when sp 2– hybridization. In this case, each carbon atom has one of three r- orbitals are not involved in hybridization, resulting in the formation of three sp 2– hybrid orbitals located in the same plane at an angle of 120° to each other, and non-hybrid 2 r The -orbital is located perpendicular to this plane. Two carbon atoms bond together to form one s-bond due to overlapping hybrid orbitals and one p-bond due to overlapping r-orbitals. The interaction of free hybrid orbitals of carbon with 1s-orbitals of hydrogen atoms leads to the formation of the ethylene molecule C 2 H 4 (Fig. 22), the simplest representative of unsaturated hydrocarbons.

Rice. 22. Formation of an ethylene molecule C 2 H 4

The overlap of electron orbitals in the case of a p-bond is less and the zones with increased electron density lie further from the atomic nuclei, therefore this bond is less strong than the s-bond.

A triple bond is formed by one s-bond and two p-bonds. In this case, the electron orbitals are in a state of sp-hybridization, the formation of which occurs due to one s- and one r- orbitals (Fig. 23).

Rice. 23. Formation of an acetylene molecule C 2 H 2

Two hybrid orbitals are located at an angle of 180° relative to each other, and the remaining non-hybrid two r-orbitals are located in two mutually perpendicular planes. The formation of a triple bond takes place in the acetylene molecule C 2 H 2.

A special type of bond occurs during the formation of a benzene molecule (C 6 H 6), the simplest representative of aromatic hydrocarbons.

Benzene contains six carbon atoms linked together in a ring (benzene ring), with each carbon atom in a state of sp 2 hybridization (Fig. 24).

All carbon atoms included in the benzene molecule are located in the same plane. Each carbon atom in the sp 2 hybridization state has one more non-hybrid p-orbital with an unpaired electron, which forms a p-bond (Fig. 25).

The axis of such a p-orbital is located perpendicular to the plane of the benzene molecule.

Rice. 24. sp 2 - orbitals of the benzene molecule C 6 H 6

Rice. 25. - bonds in the benzene molecule C 6 H 6

All six non-hybrid p orbitals form a common bonding molecular p orbital, and all six electrons combine to form a p electron sextet.

The boundary surface of such an orbital is located above and below the plane of the carbon s - skeleton. As a result of circular overlap, a single delocalized p-system arises, covering all carbon atoms of the cycle. Benzene is schematically depicted as a hexagon with a ring inside, which indicates that delocalization of electrons and corresponding bonds takes place.