CHAPTER 2. CHEMICAL BOND AND MUTUAL INFLUENCE OF ATOMS IN ORGANIC COMPOUNDS
The chemical properties of organic compounds are determined by the type of 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. The electronic structure of the carbon atom
The part of the atomic space in which the probability of finding an electron is maximum is called the 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 on its outer electronic level, only two unpaired electrons are in the ground state 1s 2 2s 2 2p 2 (Fig. 2.1, a and Appendix 2-1). In these cases, it is postulated that different atomic orbitals, close in energy, can mix with each other, forming hybrid orbitals of the same shape and energy.
Hybrid orbitals, due to the greater overlap, form stronger bonds compared to non-hybridized orbitals.
Depending on the number of hybridized orbitals, a carbon atom can be in one of three states
Rice. 2.1.The distribution of electrons in orbitals at the carbon atom in the ground (a), excited (b) and hybridized states (c - sp 3 , 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 molecules, i.e., their spatial structure.
The spatial structure of molecules is the mutual arrangement of atoms and atomic groups in space.
sp 3-Hybridization.When mixing four external AOs of an excited carbon atom (see Fig. 2.1, b) - one 2s- and three 2p-orbitals - 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). Such a 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 to 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 denoted by a regular line. A bold line or a bold wedge denotes a connection that extends forward from the plane of the drawing and is directed towards the observer; a dotted line or a hatched wedge (..........) - a connection that goes away from the observer beyond the plane of the drawing
Rice. 2.2.Types of hybridization of the carbon atom. The dot in the center is the nucleus of the atom (small fractions of hybrid orbitals are omitted to simplify the figure; unhybridized p-AOs are shown in color)
zha (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-AO of the excited carbon atom, three equivalent sp 2-hybrid orbitals and remains unhybridized 2p-AO. The carbon atom is in the state sp 2-hybridization has an electronic configuration 1s 2 2(sp 2) 3 2p 1 (see Fig. 2.1, d). This state of hybridization of the carbon atom is typical for unsaturated hydrocarbons (alkenes), as well as for some functional groups, such as carbonyl and carboxyl.
sp 2 - Hybrid 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 triangular configuration. The carbon atoms bound 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, while two p AOs remain unhybridized. The carbon atom in the sp hybridization state has the 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. 2.1e). This state of hybridization of the carbon atom occurs in compounds having a triple bond, for example, in alkynes, nitriles.
sp-hybrid orbitals are located at an angle of 180?, and two unhybridized AOs are in mutually perpendicular planes (see Fig. 2.2, c). The carbon atom in the sp hybridization state has line configuration, for example, in an acetylene molecule, all four atoms are on the same straight line (see Fig. 2.3, in).
Atoms of other organogen elements can also be in a hybridized state.
2.2. Chemical bonds of carbon atom
Chemical bonds in organic compounds are mainly represented by covalent bonds.
A covalent bond is a chemical bond formed as a result of the socialization of the electrons of the bonded atoms.
These shared electrons occupy molecular orbitals (MOs). As a rule, MO is a multicenter orbital and the electrons filling it are delocalized (dispersed). Thus, MO, like AO, can be vacant, filled with one electron or two electrons with opposite spins*.
2.2.1. σ- andπ -Communications
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 the overlap maximum on this straight line.
The σ-bond arises when any AO overlaps, including hybrid ones. 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 C-H σ-bonds by overlapping the hybrid sp 3 -AO of carbon and the 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 overlapping of AO (small fractions of hybrid orbitals are omitted, color shows sp 3 -AO carbon, black - s-AO hydrogen)
In addition to the axial overlap, another type of overlap is possible - the lateral overlap of the p-AO, leading to the formation of a π bond (Fig. 2.5).
p-atomic orbitals
Rice. 2.5.π-bond formation in ethylene by lateral overlap r-AO
A π-bond is a bond formed by lateral overlap of unhybridized p-AOs with a maximum of 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 in terms of characteristics such as energy, length, polarity, and polarizability.
Bond energyis the energy released during the formation of a bond or required to separate two bonded atoms. It serves as a measure of bond strength: the greater 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). The bonds between carbon atoms in different states of hybridization have a common pattern -
Table 2.1.Main characteristics of covalent bonds
with an increase in the fraction of the s-orbital in the hybrid orbital, 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 CH 3 bond length -C, respectively, is equal to 0.154; 0.150 and 0.146 nm.
Communication polarity due to the uneven distribution (polarization) of the 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 the polarity of the bond is the difference in the electronegativity of the bonded atoms.
Electronegativity characterizes the ability of an atom in a molecule to hold valence electrons. With an increase in the electronegativity of an atom, the degree of displacement of the bond electrons in its direction increases.
Based on the bond energies, the American chemist L. Pauling (1901-1994) proposed a quantitative characteristic of the relative electronegativity of atoms (Pauling scale). In this scale (row), typical organogenic 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 effect of substituents. For example, the electronegativity of a carbon atom in the state of sp 2 - or sp-hybridization is higher than in the state of sp 3 -hybridization, 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 length of the hybrid orbital gradually decreases (especially in the direction that provides the greatest overlap during the formation of the σ-bond), which means that in the same sequence, the electron density maximum is located 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. With a difference of up to 0.4, they speak of a weakly polar, more than 0.5 - of a strongly polar covalent bond, and more than 2.0 - of an ionic bond. Polar covalent bonds are prone to heterolytic cleavage
(see 3.1.1).
Communication polarizability is expressed in the displacement of bond electrons under the influence of an external electric field, including another reacting particle. Polarizability is determined by the electron mobility. Electrons are more mobile the farther they are from the nuclei of atoms. In terms of polarizability, the π-bond significantly exceeds the σ-bond, since the maximum electron density of the π-bond is located farther from the bonded nuclei. Polarizability largely determines the reactivity of molecules with respect to 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 the English. non-bonding) are nitrogen, oxygen, halogen atoms.
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 a donor-acceptor interaction.
A covalent bond formed by a pair of electrons from one of the bond partners is called a donor-acceptor bond.
The formed donor-acceptor bond differs only in the way of formation; its properties are the same as other covalent bonds. The donor atom acquires a positive charge.
Donor-acceptor bonds are characteristic of complex compounds.
2.2.3. Hydrogen bonds
A hydrogen atom bound to a strongly electronegative element (nitrogen, oxygen, fluorine, etc.) is able to interact 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 kind of donor-
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 the electrostatic interaction.
Intermolecular hydrogen bonds cause 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. For example, the boiling point of ethanol C 2H5 OH (78.3 ? C) is significantly higher than that of dimethyl ether CH 3 OCH 3 (-24 ? C) of the same molecular weight, which is not associated due to hydrogen bonds.
Hydrogen bonds can also be intramolecular. Such a bond in the anion of salicylic acid leads to an increase in its acidity.
Hydrogen bonds play an important role in the formation of the spatial structure of macromolecular compounds - proteins, polysaccharides, nucleic acids.
2.3. Related systems
A covalent bond can be localized or delocalized. A bond is called localized, the electrons of which are actually divided between the two nuclei of the bonded atoms. If the bond electrons are shared by more than two nuclei, then one speaks of a delocalized bond.
A delocalized bond is a covalent bond whose molecular orbital spans more than two atoms.
Delocalized bonds in most cases are π-bonds. They are characteristic of coupled systems. In these systems, a special kind of mutual influence of atoms occurs - conjugation.
Conjugation (mesomeria, from the Greek. mesos- medium) 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 participating in conjugation can belong either to two or more π-bonds, or to 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). Carbon and hydrogen atoms and, consequently, 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. 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. Overlapping occurs not only between the p-AO of the C-1 and C-2, C-3 and C-4 atoms, but also between the p-AO of the C-2 and C-3 atoms, resulting in the formation of a single π spanning four carbon atoms -system, i.e., a delocalized covalent bond arises (see Fig. 2.6, b).
Rice. 2.6.Atomic orbital model of the 1,3-butadiene molecule
This is reflected in the change in bond lengths in the molecule. The bond length C-1-C-2, as well as C-3-C-4 in butadiene-1,3 is somewhat 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 the alignment of bond lengths.
Hydrocarbons with a large number of conjugated double bonds are common in the plant kingdom. 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 member of the chain of conjugation of three sp 2 -hybridized carbon atoms and an oxygen atom. Each of these atoms contributes one p-electron to the single π-system.
pn-pairing.This type of conjugation is most often manifested in compounds containing the structural fragment -CH=CH-X, where X is a heteroatom having an unshared 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 the orbital of an 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 exists 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 particles, 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) suggests 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 resonant structures, which differ from each other only in the distribution of electrons. In conjugated systems, the main contribution to the resonant hybrid is made by structures with different π-electron density distributions (the two-sided arrow connecting these structures is a special symbol of the resonance theory).
Limit (boundary) structures do not really exist. However, they "contribute" to some extent to the real distribution of electron density in a molecule (particle), which is represented as a resonant hybrid obtained by superimposition (superposition) of limiting structures.
In ρ,π-conjugated systems with a carbon chain, conjugation can occur if there is a carbon atom with an unhybridized p-orbital next to the π-bond. Such systems can be intermediate particles - carbanions, carbocations, free radicals, for example, allyl structures. Free radical allyl fragments play an important role in the processes of lipid peroxidation.
In the allyl anion CH 2 \u003d CH-CH 2 sp 2 -hybridized carbon atom C-3 delivers to the common conjugated
Rice. 2.7.Electron density map of the COONa group in penicillin
two electron system, in the allyl radical CH 2=CH-CH 2+ - one, and in the allyl carbocation CH 2=CH-CH 2+ does not supply any. As a result, when the p-AO overlaps three sp 2 -hybridized carbon atoms, a delocalized three-center bond is formed, containing four (in the carbanion), three (in the 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 has an unpaired electron, and in the allyl anion it has a negative charge. In fact, in such conjugated systems, there is a delocalization (dispersal) of the electron density, which leads to the alignment of bonds and charges. The C-1 and C-3 atoms are equivalent in these systems. 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 with the C-2 atom.
Thus, conjugation leads to a significant difference in the electron density distribution in real structures compared to structures represented 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 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
enter into substitution reactions, not addition, resistance to oxidizing agents and temperature.
Typical representatives of aromatic systems are arenes and their derivatives. Features 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 R-AO. As a result of this overlap, a single delocalized π-system arises, in which the highest electron density is located above and below the σ-skeleton plane and covers all 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 a 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 planar cyclic system must contain (4n + 2) π-electrons, where n= 1, 2, 3, etc. (Hückel's rule, 1931). Taking into account these data, it is possible to concretize the concept of "aromaticity".
A compound is aromatic if it has a planar ring and a conjugatedπ -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 that are common to more than
Rice. 2.8.Atomic orbital model of the benzene molecule (hydrogen atoms omitted; see text for explanation)
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 the degree of overlapping of the orbitals increases and delocalization (dispersal) occurs. R-electrons. In this regard, conjugated and aromatic systems have increased thermodynamic stability. They contain a smaller amount of internal energy and in the ground state occupy a lower energy level compared to non-conjugated systems. The difference between these levels can be used to 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. With an increase in the length of the conjugated chain, 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 the nearest 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 transmission of the influence of atoms is called the inductive effect (/-effect).
Inductive effect - 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 is attenuated after three or four bonds in the circuit. Its action 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 estimated by comparing it with the hydrogen atom, the inductive effect of which is taken as zero. Graphically, the result of the /-effect is depicted by an arrow coinciding with the position of the valence line and pointing towards the more electronegative atom.
/in\stronger than a hydrogen atom, exhibitsnegativeinductive effect (-/-effect).
Such substituents generally lower the electron density of the system, they are called electron-withdrawing. These include most of the functional groups: OH, NH 2, COOH, NO2 and cationic groups, such as -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 main role in the transfer of electronic influence is played by π-electrons of delocalized covalent bonds. The effect that manifests itself in a shift in the electron density of a delocalized (conjugated) π-system is called the mesomeric (M-effect), or conjugation effect.
Mesomeric effect - the transfer of the electronic influence of substituents along the conjugated system.
In this case, the substitute is itself a member of the conjugated system. It can introduce into the conjugation system either a π-bond (carbonyl, carboxyl groups, etc.), or an unshared pair of electrons of a heteroatom (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).
Substituents that include atoms with a lone pair of electrons (for example, an amino group in an aniline molecule) or a whole negative charge have an M-Effect. These substitutes are capable
to the transfer of a pair of electrons to a common conjugated system, that is, they are electron donor.
A substituent that lowers the electron density in a conjugated system exhibitsnegativemesomeric effect (-M- effect).
The M-effect in the conjugated system is possessed by oxygen or nitrogen atoms bound by a double bond to a carbon atom, as shown in the example of acrylic acid and benzaldehyde. Such groupings are electron-withdrawing.
The displacement of the electron density is indicated by a curved arrow, the beginning of which shows which p- or π-electrons are being displaced, and the end is the bond or atom to which they are displaced. The mesomeric effect, in contrast to the inductive effect, is transmitted over 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 action of the inductive and mesomeric effects (Table 2.2).
Table 2.2.Electronic effects of some substituents
The electronic effects of substituents make it possible to give a qualitative estimate of the electron density distribution in a nonreacting molecule and to predict its properties.
Continuation. For the beginning, see № 15, 16/2004
Lesson 5
atomic orbitals of carbon
A covalent chemical bond is formed using common bonding electron pairs of the type:
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 by one in the orbital cell.
atomic orbital is a function that describes the density of the electron cloud at each point in space around the nucleus of an atom. An electron cloud is a region of space in which an electron can be found with a high probability.
To harmonize the electronic structure of the carbon atom and the valence of this element, the concepts of 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 pass to free R-orbital. Then four unpaired electrons appear in the carbon atom:
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, 2r y and 2 pz according to the axes along which these orbitals are located.
When chemical bonds are formed, the electron orbitals acquire the same shape. So, in saturated hydrocarbons, one s-orbital and three R-orbitals of a carbon atom to form four identical (hybrid) sp 3-orbitals:
This is - 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. At the same time, the axes of four sp 3-hybrid orbitals turn out to be directed to the vertices of the tetrahedron (regular triangular pyramid).
Accordingly, the angles between these orbitals are tetrahedral, equal to 109°28".
The tops 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()-bond. For example, in a C 2 H 6 ethane molecule, 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 hybrid 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-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-hybridization of carbon 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 a bond. (The other bonding orbitals of the atom are directed in opposite directions.) The second bond is formed as a result of the overlap of non-hybrid R-orbitals on both sides of the line connecting the nuclei of atoms.
Covalent bond formed by lateral overlap R-orbitals of neighboring carbon atoms is called pi()-bond.
Education
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Due to less overlap of orbitals, the -bond is less strong than the -bond.
sp-Hybridization is a mixing (alignment in form and energy) of one s- and one
R-orbitals with the formation of two hybrid sp-orbitals. sp- 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 perpendicular to each other.
directions - connections. On the image sp-orbitals are shown along the axis y, and the unhybridized two
R-orbitals - along the axes X and z.
The triple carbon-carbon bond CC consists of a -bond that occurs when 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 bonds of carbon
| 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 is connection |
|
| 2 | sp | Two - connections and two connections |
H-CC-H |
Exercises.
1. What 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 are the 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 in the second energy level of a C atom and how do they differ?
6. What is the difference between hybrid orbitals and the original orbitals from which they were formed?
7. What types of hybridization are known for the carbon atom and what are they?
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) the type of hybridization; b) types of its chemical bonds; c) bond angles.
Answers to exercises for topic 1
Lesson 5
1. Electrons that are one per 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 participating in the formation of one chemical bond are called common electron pair. Usually, before the formation of a chemical bond, one of the electrons of this pair belonged to one atom, and the other electron belonged to another atom:
3. The electronic state of the atom, in which the order of filling of electronic orbitals is observed: 1 s 2 , 2s 2 , 2p 2 , 3s 2 , 3p 2 , 4s 2 , 3d 2 , 4p 2 etc. are called main state. AT 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:
Whereas in the ground state there were only two valence unpaired electrons, in the excited state there are four such electrons.
5.
An atomic orbital is a function that describes the density of an electron cloud at each point in space around the nucleus of a given atom. There are four orbitals on the second energy level of the carbon atom - 2 s, 2p x, 2r y,
2pz. These orbitals are:
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 denoted p x, r y,
pz.
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 initial s-
and p-orbitals.
7. Three types of hybridization are known for the carbon atom: sp 3 , sp 2 and sp (see the text of lesson 5).
9.
-bond - a covalent bond formed by frontal overlapping of orbitals along a line connecting the centers of atoms.
-bond - a covalent bond formed by lateral overlap R-orbitals on either side of the line connecting the centers of atoms.
- Bonds are shown by the second and third lines between the connected atoms.
In the ground state, the carbon atom C (1s 2 2s 2 2p 2) has two unpaired electrons, due to which only two common electron pairs can form. However, in most of its compounds, carbon is tetravalent. This is due to 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. able to form four covalent bonds and take part in the formation of four common electron pairs:
6 C 1 s 2 2s 2 2 p 2 6 C * 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 between them (Table 7).
Table 7
Types of hybridization and geometry of molecules
Simple (single) s - communication is carried out when sp 3-hybridization, in which all four hybrid orbitals are equivalent and have an orientation in space at an angle of 109 about 29 ’ to each other and are oriented towards the vertices of a regular tetrahedron.
Rice. 19. The formation of a methane molecule CH 4
If hybrid orbitals of carbon overlap with spherical 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. 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 the C-H angles in the methane molecule.
Rice. 21. Ethane molecule C 2 H 6
Therefore, the carbon atoms are in the state sp 3-hybridization.
Hybridization of electron orbitals of carbon atoms can be incomplete, i.e. it can involve two sp 2 hybridization) or one ( sp-hybridization) of three R- orbitals. In this case, between the carbon atoms are formed multiples(double or triple) connections. Hydrocarbons with multiple bonds are called unsaturated or unsaturated. A double bond (C = C) is formed when sp 2- hybridization. In this case, each of the carbon atoms 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 about to each other, and non-hybrid 2 R-orbital is perpendicular to this plane. Two carbon atoms are connected to each other, forming 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 an ethylene molecule C 2 H 4 (Fig. 22), the simplest representative of unsaturated hydrocarbons.
Rice. 22. The formation of an ethylene molecule C 2 H 4
The overlap of electronic orbitals in the case of p-bonds is less and the zones with increased electron density lie farther from the nuclei of atoms, so 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 degrees relative to each other, and the remaining two non-hybrid R-orbitals are located in two mutually perpendicular planes. The formation of a triple bond takes place in the acetylene C 2 H 2 molecule.
A special type of bond arises 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 cycle (benzene ring), while each carbon atom is 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 state of sp 2 hybridization 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 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 are combined into 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 overlapping, 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 there is a delocalization of electrons and the corresponding bonds.
169375 0
Each atom has a certain number of electrons.
Entering into chemical reactions, atoms donate, acquire, or socialize electrons, reaching the most stable electronic configuration. The configuration with the lowest energy is the most stable (as in noble gas atoms). This pattern is called the "octet rule" (Fig. 1).
Rice. one.
This rule applies to all connection types. Electronic bonds between atoms allow them to form stable structures, from the simplest crystals to complex biomolecules that eventually form living systems. They differ from crystals in their continuous metabolism. However, many chemical reactions proceed according to the mechanisms electronic transfer, which play an important role in the energy processes in the body.
A chemical bond is a force that holds together two or more atoms, ions, molecules, or any combination of them..
The nature of the chemical bond is universal: it is an electrostatic force of attraction between negatively charged electrons and positively charged nuclei, determined by the configuration of the electrons in the outer shell of atoms. The ability of an atom to form chemical bonds is called valency, or oxidation state. The concept of valence electrons- electrons that form chemical bonds, that is, those located in the most high-energy orbitals. Accordingly, the outer shell of an atom containing these orbitals is called valence shell. At present, it is not enough to indicate the presence of a chemical bond, but it is necessary to clarify its type: ionic, covalent, dipole-dipole, metallic.
The first type of connection isionic connection
According to Lewis and Kossel's electronic theory of valency, atoms can achieve a stable electronic configuration in two ways: first, by losing electrons, becoming cations, secondly, acquiring them, turning into anions. As a result of electron transfer, due to the electrostatic force of attraction between ions with charges of the opposite sign, a chemical bond is formed, called Kossel " electrovalent(now called ionic).
In this case, anions and cations form a stable electronic configuration with a filled outer electron shell. Typical ionic bonds are formed from cations of T and II groups of the periodic system and anions of non-metallic elements of groups VI and VII (16 and 17 subgroups - respectively, chalcogens and halogens). The bonds in ionic compounds are unsaturated and non-directional, so they retain the possibility of electrostatic interaction with other ions. On fig. 2 and 3 show examples of ionic bonds corresponding to the Kossel electron transfer model.
Rice. 2.
Rice. 3. Ionic bond in the sodium chloride (NaCl) molecule
Here it is appropriate to recall some of the properties that explain the behavior of substances in nature, in particular, to consider the concept of acids and grounds.
Aqueous solutions of all these substances are electrolytes. They change color in different ways. indicators. The mechanism of action of indicators was discovered by F.V. Ostwald. He showed that the indicators are weak acids or bases, the color of which in the undissociated and dissociated states is different.
Bases can neutralize acids. Not all bases are soluble in water (for example, some organic compounds that do not contain -OH groups are insoluble, in particular, triethylamine N (C 2 H 5) 3); soluble bases are called alkalis.
Aqueous solutions of acids enter into characteristic reactions:
a) with metal oxides - with the formation of salt and water;
b) with metals - with the formation of salt and hydrogen;
c) with carbonates - with the formation of salt, CO 2 and H 2 O.
The properties of acids and bases are described by several theories. In accordance with the theory of S.A. Arrhenius, an acid is a substance that dissociates to form ions H+ , while the base forms ions IS HE- . This theory does not take into account the existence of organic bases that do not have hydroxyl groups.
In line with proton Bronsted and Lowry's theory, an acid is a substance containing molecules or ions that donate protons ( donors protons), and the base is a substance consisting of molecules or ions that accept protons ( acceptors protons). Note that in aqueous solutions, hydrogen ions exist in a hydrated form, that is, in the form of hydronium ions H3O+ . This theory describes reactions not only with water and hydroxide ions, but also carried out in the absence of a solvent or with a non-aqueous solvent.
For example, in the reaction between ammonia NH 3 (weak base) and hydrogen chloride in the gas phase, solid ammonium chloride is formed, and in an equilibrium mixture of two substances there are always 4 particles, two of which are acids, and the other two are bases:
This equilibrium mixture consists of two conjugated pairs of acids and bases:
1)NH 4+ and NH 3
2) HCl and Cl ‑
Here, in each conjugated pair, the acid and base differ by one proton. Every acid has a conjugate base. A strong acid has a weak conjugate base, and a weak acid has a strong conjugate base.
The Bronsted-Lowry theory makes it possible to explain the unique role of water for the life of the biosphere. Water, depending on the substance interacting with it, can exhibit the properties of either an acid or a base. For example, in reactions with aqueous solutions of acetic acid, water is a base, and with aqueous solutions of ammonia, it is an acid.
1) CH 3 COOH + H 2 O ↔ H 3 O + + CH 3 SOO- . Here the acetic acid molecule donates a proton to the water molecule;
2) NH3 + H 2 O ↔ NH4 + + IS HE- . Here the ammonia molecule accepts a proton from the water molecule.
Thus, water can form two conjugated pairs:
1) H 2 O(acid) and IS HE- (conjugate base)
2) H 3 O+ (acid) and H 2 O(conjugate base).
In the first case, water donates a proton, and in the second, it accepts it.
Such a property is called amphiprotonity. Substances that can react as both acids and bases are called amphoteric. Such substances are often found in nature. For example, amino acids can form salts with both acids and bases. Therefore, peptides readily form coordination compounds with the metal ions present.
Thus, the characteristic property of an ionic bond is the complete displacement of a bunch of binding electrons to one of the nuclei. This means that there is a region between the ions where the electron density is almost zero.
The second type of connection iscovalent connection
Atoms can form stable electronic configurations by sharing electrons.
Such a bond is formed when a pair of electrons is shared one at a time. from each atom. In this case, the socialized bond electrons are distributed equally among the atoms. An example of a covalent bond is homonuclear diatomic H molecules 2 , N 2 , F 2. Allotropes have the same type of bond. O 2 and ozone O 3 and for a polyatomic molecule S 8 and also heteronuclear molecules hydrogen chloride Hcl, carbon dioxide CO 2, methane CH 4, ethanol With 2 H 5 IS HE, sulfur hexafluoride SF 6, acetylene With 2 H 2. All these molecules have the same common electrons, and their bonds are saturated and directed in the same way (Fig. 4).
For biologists, it is important that the covalent radii of atoms in double and triple bonds are reduced compared to a single bond.
Rice. 4. Covalent bond in the Cl 2 molecule.
Ionic and covalent types of bonds are two limiting cases of many existing types of chemical bonds, and in practice most of the bonds are intermediate.
Compounds of two elements located at opposite ends of the same or different periods of the Mendeleev system predominantly form ionic bonds. As the elements approach each other within a period, the ionic nature of their compounds decreases, while the covalent character increases. For example, the halides and oxides of the elements on the left side of the periodic table form predominantly ionic bonds ( NaCl, AgBr, BaSO 4 , CaCO 3 , KNO 3 , CaO, NaOH), and the same compounds of the elements on the right side of the table are covalent ( H 2 O, CO 2, NH 3, NO 2, CH 4, phenol C6H5OH, glucose C 6 H 12 O 6, ethanol C 2 H 5 OH).
The covalent bond, in turn, has another modification.
In polyatomic ions and in complex biological molecules, both electrons can only come from one atom. It is called donor electron pair. An atom that socializes this pair of electrons with a donor is called acceptor electron pair. This type of covalent bond is called coordination (donor-acceptor, ordative) communication(Fig. 5). This type of bond is most important for biology and medicine, since the chemistry of the most important d-elements for metabolism is largely described by coordination bonds.
Pic. 5.
As a rule, in a complex compound, a metal atom acts as an electron pair acceptor; on the contrary, in ionic and covalent bonds, the metal atom is an electron donor.
The essence of the covalent bond and its variety - the coordination bond - can be clarified with the help of another theory of acids and bases, proposed by GN. Lewis. He somewhat expanded the semantic concept of the terms "acid" and "base" according to the Bronsted-Lowry theory. The Lewis theory explains the nature of the formation of complex ions and the participation of substances in nucleophilic substitution reactions, that is, in the formation of CS.
According to Lewis, an acid is a substance capable of forming a covalent bond by accepting an electron pair from a base. A Lewis base is a substance that has a lone pair of electrons, which, by donating electrons, forms a covalent bond with Lewis acid.
That is, the Lewis theory expands the range of acid-base reactions also to reactions in which protons do not participate at all. Moreover, the proton itself, according to this theory, is also an acid, since it is able to accept an electron pair.
Therefore, according to this theory, cations are Lewis acids and anions are Lewis bases. The following reactions are examples:
It was noted above that the subdivision of substances into ionic and covalent is relative, since there is no complete transition of an electron from metal atoms to acceptor atoms in covalent molecules. In compounds with an ionic bond, each ion is in the electric field of ions of the opposite sign, so they are mutually polarized, and their shells are deformed.
Polarizability determined by the electronic structure, charge and size of the ion; it is higher for anions than for cations. The highest polarizability among cations is for cations of larger charge and smaller size, for example, for Hg 2+ , Cd 2+ , Pb 2+ , Al 3+ , Tl 3+. Has a strong polarizing effect H+ . Since the effect of ion polarization is two-sided, it significantly changes the properties of the compounds they form.
The third type of connection -dipole-dipole connection
In addition to the listed types of communication, there are also dipole-dipole intermolecular interactions, also known as van der Waals .
The strength of these interactions depends on the nature of the molecules.
There are three types of interactions: permanent dipole - permanent dipole ( dipole-dipole attraction); permanent dipole - induced dipole ( induction attraction); instantaneous dipole - induced dipole ( dispersion attraction, or London forces; rice. 6).
Rice. 6.
Only molecules with polar covalent bonds have a dipole-dipole moment ( HCl, NH 3, SO 2, H 2 O, C 6 H 5 Cl), and the bond strength is 1-2 debye(1D \u003d 3.338 × 10 -30 coulomb meters - C × m).
In biochemistry, another type of bond is distinguished - hydrogen connection, which is a limiting case dipole-dipole attraction. This bond is formed by the attraction between a hydrogen atom and a small electronegative atom, most often oxygen, fluorine and nitrogen. With large atoms that have a similar electronegativity (for example, with chlorine and sulfur), the hydrogen bond is much weaker. The hydrogen atom is distinguished by one essential feature: when the binding electrons are pulled away, its nucleus - the proton - is exposed and ceases to be screened by electrons.
Therefore, the atom turns into a large dipole.
A hydrogen bond, unlike a van der Waals bond, is formed not only during intermolecular interactions, but also within one molecule - intramolecular hydrogen bond. Hydrogen bonds play an important role in biochemistry, for example, for stabilizing the structure of proteins in the form of an α-helix, or for the formation of a DNA double helix (Fig. 7).
Fig.7.
Hydrogen and van der Waals bonds are much weaker than ionic, covalent, and coordination bonds. The energy of intermolecular bonds is indicated in Table. one.
Table 1. Energy of intermolecular forces
Note: The degree of intermolecular interactions reflect the enthalpy of melting and evaporation (boiling). Ionic compounds require much more energy to separate ions than to separate molecules. The melting enthalpies of ionic compounds are much higher than those of molecular compounds.
The fourth type of connection -metallic bond
Finally, there is another type of intermolecular bonds - metal: connection of positive ions of the lattice of metals with free electrons. This type of connection does not occur in biological objects.
From a brief review of the types of bonds, one detail emerges: an important parameter of an atom or ion of a metal - an electron donor, as well as an atom - an electron acceptor is its the size.
Without going into details, we note that the covalent radii of atoms, the ionic radii of metals, and the van der Waals radii of interacting molecules increase as their atomic number in the groups of the periodic system increases. In this case, the values of the ion radii are the smallest, and the van der Waals radii are the largest. As a rule, when moving down the group, the radii of all elements increase, both covalent and van der Waals.
The most important for biologists and physicians are coordination(donor-acceptor) bonds considered by coordination chemistry.
Medical bioinorganics. G.K. Barashkov
