The main objects of bio.chemistry.

Objects of study bioorganic chemistry are proteins and peptides, nucleic acids, carbohydrates, lipids, biopolymers, alkaloids, terpenoids, vitamins, antibiotics, hormones, toxins, as well as synthetic regulators of biological processes: drugs, pesticides, etc.

Isomerism of organic compounds, its types. Characteristics of types of isomerism, examples.

There are two types of isomerism: structural and spatial (i.e., stereoisomerism). Structural isomers differ from each other in the order of bonds of atoms in a molecule, stereoisomers - in the arrangement of atoms in space with the same order of bonds between them.

The following types of structural isomerism are distinguished: carbon skeleton isomerism, position isomerism, isomerism of various classes of organic compounds (interclass isomerism).

The isomerism of the carbon skeleton is due to the different bond order between the carbon atoms that form the skeleton of the molecule. For example: the molecular formula C4H10 corresponds to two hydrocarbons: n-butane and isobutane. Three isomers are possible for the C5H12 hydrocarbon: pentane, iso-pentane, and neopentane. C4H10 corresponds to two hydrocarbons: n-butane and isobutane. Three isomers are possible for the C5H12 hydrocarbon: pentane, iso-pentane, and neopentane.

Position isomerism is due to the different position of the multiple bond, substituent, functional group with the same carbon skeleton of the molecule

Interclass isomerism is the isomerism of substances belonging to different classes of organic compounds.

Modern classification and nomenclature of organic compounds.

Currently, the systematic nomenclature is widely used - IUPAC - the international unified chemical nomenclature. IUPAC rules are based on several systems:

1) radical-functional (the name is based on the name of the functional group),

2) connecting (names are made up of several equal parts),

3) substitution (the basis of the name is a hydrocarbon fragment).

covalent bonds. Pi and sigma bonds.

covalent bond is the main type of bond in organic compounds.

This is a bond formed by the overlap of a pair of valence electron clouds.

A pi bond is a covalent bond formed by overlapping p atomic orbitals.

A sigma bond is a covalent bond formed when s-atomic orbitals overlap.

If both s- and p-bonds are formed between atoms in a molecule, then a multiple (double or triple) bond is formed.

6. Modern ideas about the structure of organic compounds. The concept of "chemical structure", "configuration", "conformation", their definition. The role of structure in the manifestation of biological activity.

In 1861 A.M. Butlerov proposed a theory of the chemical structure of organic compounds, which underlies modern ideas about the structure of org. compounds, which consists of the following main provisions:

1. In the molecules of substances there is a strict sequence of chemical binding of atoms, which is called the chemical structure.

2. The chemical properties of a substance are determined by the nature of the elementary constituents, their quantity and chemical structure.

3. If substances with the same composition and molecular weight have a different structure, then the phenomenon of isomerism occurs.

4. Since only some parts of the molecule change in specific reactions, the study of the structure of the product helps to determine the structure of the original molecule.

5. The chemical nature (reactivity) of individual atoms in a molecule varies depending on the environment, i.e. on what atoms of other elements they are connected to.

The concept of "chemical structure" includes the idea of ​​a certain order of connection of atoms in a molecule and of their chemical interaction, which changes the properties of atoms.

14. Main characteristics of the covalent bond. Bond length and energy. saturation and direction. Communication multiplicity. Sigma and pi bonds.

- A chemical bond carried out by shared electronic pairs is called atomic or covalent. Each covalent chemical bond has certain qualitative or quantitative characteristics. These include:

Link length

Bond energy

Saturability

Direction of communication

Communication polarity

Communication multiplicity

- Link length is the distance between the nuclei of bound atoms. It depends on the size of the atoms and on the degree of overlap of their electron shells. The bond length is determined by the bond order: the higher the bond order, the shorter its length.

Bond energy is the energy released during the formation of a molecule from single atoms. It is usually expressed in J/mol (or cal/mol). The bond energy is determined by the bond order: the greater the bond order, the greater its energy. The bond energy is a measure of its strength. Its value is determined by the work required to break the bond, or the gain in energy during the formation of matter from individual atoms. A system that contains less energy is more stable. For diatomic molecules, the bond energy is equal to the dissociation energy, taken with the opposite sign. If more than 2 different atoms are connected in a molecule, then the average binding energy does not coincide with the value of the dissociation energy of the molecule. The bond energies in molecules consisting of identical atoms decrease in groups from top to bottom. Bond energies increase over the period.

- Saturability- shows how many bonds a given atom can form with others due to common electron pairs. It is equal to the number of common electron pairs by which this atom is connected to others. The saturation of a covalent bond is the ability of an atom to participate in the formation of a limited number of covalent bonds.

Orientation is a certain mutual arrangement of binding electron clouds. It leads to a certain arrangement in space of the nuclei of chemically bonded atoms. The spatial orientation of a covalent bond is characterized by the angles between the formed bonds, which are called valence angles.

- Communication multiplicity. It is determined by the number of electron pairs involved in the bond between atoms. If the bond is formed by more than one pair of electrons, then it is called a multiple. As the bond multiplicity increases, the energy increases and the bond length decreases. In molecules with multiple bonds, there is no rotation around the axis.

- Sigma - and pi bonds. The chemical bond is due to the overlap of electron clouds. If this overlap occurs along a line connecting the nuclei of atoms, then such a bond is called a sigma bond. It can be formed by s-s electrons, p-p electrons, s-p electrons. A chemical bond carried out by one electron pair is called a single bond. Single bonds are always sigma bonds. Orbitals of type s form only sigma bonds. But a large number of compounds are known in which there are double and even triple bonds. One of them is a sigma bond and the others are called pi bonds. When such bonds are formed, the overlap of electron clouds occurs in two regions of space symmetrical to the internuclear axis.

15. Hybridization of atomic orbitals on the example of molecules: methane, aluminum chloride, beryllium chloride. Valence angle and geometry of the molecule. Method of molecular orbitals (MO LCAO). Energy diagrams of homo- and hetero-nuclear molecules (N2, Cl2, NH3, Be2).

- Hybridization. The new set of mixed orbitals is called hybrid orbitals, and the mixing technique itself is called hybridization of atomic orbitals.

The mixing of one s- and one p-orbital, as in BeCl2, is called sp-hybridization. In principle, hybridization of the s-orbital is possible not only with one, but also with two, three or a non-integer number of p-orbitals, as well as hybridization with the participation of d-orbitals.

Consider a linear BeCl2 molecule. A beryllium atom in the valence state is capable of forming two bonds due to one s- and one p-electron. It is obvious that in this case two bonds of different lengths with chlorine atoms should be obtained, since the radial distribution of these electrons is different. The real BeCl2 molecule is symmetrical and linear; in it, two Be-Cl bonds are exactly the same. This means that they are provided with electrons of the same state, i.e. here, a beryllium atom in the valence state no longer has one s- and one p-electron, but two electrons located in orbitals formed by a “mixture” of s- and p-atomic orbitals. The methane molecule will have sp3 hybridization, and the aluminum chloride molecule will have sp2 hybridization.

Conditions for the stability of hybridization:

1) Compared to the original orbital atoms, the hybrid orbitals should overlap more tightly.

2) Atomic orbitals that are close in energy level take part in hybridization, therefore, stable hybrid orbitals should be formed on the left side of the periodic system.

Hybridization	Molecule shape	Valence angle
	Linear
	Triangle
	Tetrahedron

- Valence angle and geometry of the molecule. In each case, hybrid orbitals have a certain orientation, which contributes to the formation of molecules with certain angles between bonds, valence angles. Each type of hybridization corresponds to a specific bond angle and a specific shape of the molecule:

- MO LCAO. Molecular orbitals can be thought of as a linear combination of atomic orbitals. Molecular orbitals must have a certain symmetry. When filling atomic orbitals with electrons, the following rules must be taken into account:

1. If the atomic orbital is some function that is a solution of the Schrödinger Equation and describes the state of an electron in an atom, the MO method is also a solution of the Schrödinger equation, but for an electron in a molecule.

2. A molecular orbital is found by adding or subtracting atomic orbitals.

3. Molecular orbitals and their number are equal to the sum of the atomic orbitals of the reacting atoms.

If the solution for molecular orbitals is obtained by adding the functions of atomic orbitals, then the energy of the molecular orbitals will be lower than the energy of the original atomic orbitals. And such an orbital is called bonding orbital.

In the case of function subtraction, the molecular orbital has a large energy, and it is called loosening.

There are sigma and pi orbitals. They are filled according to Hund's rule.

The number of bonds (bond order) is equal to the difference between the total number of electrons in the bonding orbital and the number of electrons in the antibonding orbital, divided by 2.

The MO method uses energy diagrams:

16. Polarization of communication. The dipole moment of the bond. Characteristics of interacting atoms: ionization potential, electron affinity, electronegativity. The degree of ionicity of the bond.

- Dipole moment- physical quantity characterizing the electrical properties of a system of charged particles. In the case of a dipole (two particles with opposite charges), the electric dipole moment is equal to the product of the positive charge of the dipole and the distance between the charges and is directed from a negative charge to a positive one. The dipole moment of a chemical bond is due to the displacement of the electron cloud towards one of the atoms. A bond is said to be polar if the corresponding dipole moment differs substantially from zero. Cases are possible when individual bonds in a molecule are polar, and the total dipole moment of the molecule is zero; such molecules are called non-polar (eg CO 2 and CCl 4 molecules). If the dipole moment of the molecule is nonzero, the molecule is said to be polar. For example, the H 2 O molecule. The order of magnitude of the dipole moment of the molecule is determined by the product of the electron charge (1.6.10 -19 C) and the length of the chemical bond (of the order of 10 -10 m).

The chemical nature of an element is determined by the ability of its atom to lose and gain electrons. This ability can be quantified by the ionization energy of an atom and its electron affinity.

- Ionization energy atom is the amount of energy required to detach an electron from an unexcited atom. It is expressed in kilojoules per mole. For multielectron atoms, the ionization energies E1, E2, E3, ..., En correspond to the detachment of the first, second, etc. electrons. In this case, always E1

- Affinity of an atom for an electron- the energy effect of the attachment of an electron to a neutral atom with its transformation into a negative ion. The affinity of an atom for an electron is expressed in kJ/mol. The electron affinity is numerically equal, but opposite in sign, to the ionization energy of a negatively charged ion and depends on the electron configuration of the atom. Group 7 p-elements have the highest electron affinity. Atoms with the s2 (Be, Mg, Ca) and s2p6 (Ne, Ar, Kr) configurations or half-filled p-sublayers (N, P, As) do not exhibit electron affinity.

- Electronegativity is the average characteristic of the ability of an atom in a compound to attract an electron. In this case, the difference in the states of atoms in various compounds is neglected. Unlike the ionization potential and electron affinity, ER is not a strictly defined physical quantity, but a useful conditional characteristic. The most electronegative element is fluorine. EO depends on the ionization energy and electron affinity. According to one definition, the ER of an atom can be expressed as half the sum of its ionization energy and electron affinity. An element cannot be assigned a constant EC. It depends on many factors, in particular, on the valence state of the element, the type of compound in which it enters, etc.

17. Polarizing ability and polarizing action. Explanation of some physical properties of substances in terms of this theory.

- The theory of polarization considers all substances to be purely ionic. In the absence of an external field, all ions have a spherical shape. When the ions approach each other, the field of the cation affects the field of the anion, and they are deformed. Ion polarization is the displacement of the outer electron cloud of ions relative to their nucleus.

Polarization consists of two processes:

ion polarizability

polarizing action on another ion

The polarizability of an ion is a measure of the ability of the electron cloud of an ion to deform under the influence of an external electric field.

Regularities of polarizability of ions:

Anions are more polarized than cations. Excess electron density leads to a large diffuseness, friability of the electron cloud.

The polarizability of isoelectronic ions increases with decreasing positive charges and increasing negative charges. Isoelectronic ions have the same configuration.

In multiply charged cations, the nuclear charge far exceeds the number of electrons. This densifies the electron shell, it stabilizes, so such ions are less susceptible to deformation. The polarizability of cations decreases on passing from ions with an outer electron shell filled with 18 electrons to an unfilled one, and further to noble gas ions. This is due to the fact that for electrons of the same period, the d-electron shell is more diffuse compared to the s- and p-electron shells, because d-electrons spend more time near the nucleus. Therefore, d-electrons interact more strongly with the surrounding anions.

The polarizability of ions - analogues increases with an increase in the number of electron layers. Polarizability is most difficult for small and multiply charged cations, with the electron shell of noble gases. Such cations are called rigid. The most easily polarized are bulk multiply charged anions and low charged bulk cations. These are soft ions.

- Polarizing action. Depends on the charges, size and structure of the outer electron layer.

1. The polarizing effect of a cation increases with an increase in its charge and a decrease in radius. The maximum polarizing effect is characteristic of Catons with small radii and large charges; therefore, they form compounds of the covalent type. The greater the charge, the greater the polarizing bond.

2. The polarizing effect of cations increases with the transition of ions with an s-electron cloud to an incomplete one and to a 18-electron one. The greater the polarizing effect of the cation, the greater the contribution of the covalent bond.

- Application of polarization theory to explain physical properties:

The greater the polarizability of an anion (the polarizing effect of a cation), the more likely it is to form a covalent bond. Therefore, the boiling and melting points for compounds with a covalent bond will be lower than for compounds with an ionic bond. The greater the ionicity of the bond, the higher the melting and boiling points.

The deformation of the electron shell affects the ability to reflect or absorb light waves. From here, from the position of the theory of polarization, one can explain the color of the compounds: white - everything reflects; black - absorbs; transparent - passes. This is due: if the shell is deformed, then the quantum levels of electrons approach each other, reducing the energy barrier, so a small energy is required for excitation. Because absorption is associated with the excitation of electrons, i.e. with their transition to high-lying levels, then in the presence of high polarization, already visible light can excite external electrons and the substance will turn out to be colored. The higher the charge of the anion, the lower the color intensity. The polarizing effect affects the reactivity of compounds; therefore, for many compounds, salts of oxygen-containing acids are more stable than the salts themselves. The largest polarizing effect of the d-elements. The greater the charge, the greater the polarizing effect.

18. Ionic bond as a limiting case of covalent polar bond. Properties of substances with different types of bonds.

The nature of the ionic bond can be explained by the electrostatic interaction of ions. The ability of elements to form simple ions is due to the structure of their atoms. Cations most easily form elements with low ionization energy, alkali and alkaline earth metals. Anions are most easily formed by group 7 p-elements due to their high electron affinity.

The electric charges of ions cause their attraction and repulsion. Ions can be thought of as charged spheres whose force fields are uniformly distributed in all directions in space. Therefore, each ion can attract ions of the opposite sign to itself in any direction. An ionic bond, unlike a covalent bond, is non-directional.

The interaction of ions of the opposite sign with each other cannot lead to complete mutual compensation of their force fields. Because of this, they retain the ability to attract ions in other directions as well. Therefore, unlike a covalent bond, an ionic bond is characterized by unsaturation.

19.Metal connection. Similarities and differences with ionic and covalent bonds

A metallic bond is one in which the electrons of each individual atom belong to all the atoms in contact. The energy difference between the "molecular" orbitals in such a bond is small, so electrons can easily move from one "molecular" orbital to another and, therefore, move in the bulk of the metal.

Metals differ from other substances in high electrical conductivity and thermal conductivity. Under normal conditions, they are crystalline substances (with the exception of mercury) with high coordination numbers of atoms. In a metal, the number of electrons is much less than the number of orbitals, so electrons can move from one orbital to another. Metal atoms are characterized by high ionization energy - valence electrons are weakly retained in the atom, i.e. move easily in the crystal. The ability to move electrons through the crystal determines the electrical conductivity of metals.

Thus, unlike covalent and ionic compounds, in metals a large number of electrons simultaneously bind a large number of atomic nuclei, and the electrons themselves can move in the metal. In other words, a strongly delocalized chemical bond takes place in metals. The metallic bond has a certain similarity with the covalent bond, since it is based on the socialization of valence electrons. However, valence electrons of only two interacting atoms participate in the formation of a covalent bond, while in the formation of a metallic bond, all atoms participate in the socialization of electrons. That is why the metallic bond does not have a spatial orientation and saturation, which largely determines the specific properties of metals. The energy of a metallic bond is 3-4 times less than the energy of a covalent bond.

20. Hydrogen bond. Intermolecular and intramolecular. The mechanism of education. Features of the physical properties of substances with a hydrogen bond. Examples.

- The hydrogen bond is a special type of chemical bond. It is characteristic of hydrogen compounds with the most electronegative elements (fluorine, oxygen, nitrogen, and to a lesser extent chlorine and sulfur).

The hydrogen bond is very common and plays an important role in the association of molecules, in the processes of crystallization, dissolution, the formation of crystalline hydrates, etc. For example, in the solid, liquid and even gaseous state, hydrogen fluoride molecules are connected in a zigzag chain, which is due precisely to hydrogen bonding.

Its peculiarity is that the hydrogen atom, which is part of one molecule, forms a second, weaker bond with an atom in another molecule, as a result of which both molecules are combined into a complex. A characteristic feature of such a complex is the so-called hydrogen bridge – A – H...B–. The distance between atoms in a bridge is greater than between atoms in a molecule. Initially, the hydrogen bond was treated as an electrostatic interaction. At present, it has been concluded that the donor-acceptor interaction plays an important role in hydrogen bonding. A hydrogen bond is formed not only between molecules of different substances, but also in the molecules of the same substance H2O, HF, NH3, etc. This also explains the difference in the properties of these substances compared to related compounds. Hydrogen bonding is known within molecules, especially in organic compounds. Its formation is facilitated by the presence of an acceptor group A-H and a donor group B-R in the molecule. In an A-H molecule, A is the most electronegative element. Hydrogen bonding in polymers, such as peptides, results in a helical structure. DNA has similar structures - deoxyribonucleic acid - the keeper of the code of heredity. Hydrogen bonds are not strong. They easily form and break at ordinary temperature, which is very important in biological processes. It is known that hydrogen compounds with highly electronegative nonmetals have abnormally high boiling points.

Intermolecular interaction. The forces of attraction between saturated atoms and molecules are extremely weak compared to ionic and covalent bonds. Substances in which molecules are held together by extremely weak forces are more often gases at 20 degrees, and in many cases their boiling points are very low. The existence of such weak forces was discovered by van der Waals. The existence of such forces in the system can be explained by:

1. The presence of a permanent dipole in a molecule. In this case, as a result of a simple electrostatic attraction of dipoles, weak interaction forces arise - dipole-dipole (H2O, HCl, CO)

2. The dipole moment is very small, but when interacting with water, an induced dipole can form, which occurs as a result of polymerization of molecules by dipoles of surrounding molecules. This effect can be superimposed on the dipole-dipole interaction and increase the attraction.

3. Dispersion forces. These forces act between any atoms and molecules, regardless of their structure. This concept was introduced by London. For symmetrical atoms, the only acting forces are the London forces.

21. Aggregate states of matter: solid, liquid, gaseous. Crystalline and amorphous states. Crystal lattices.

- Under normal conditions, atoms, ions and molecules do not exist individually. It always constitutes only parts of a higher organization of a substance that practically participates in chemical transformations - the so-called state of aggregation. Depending on external conditions, all substances can be in different states of aggregation - in gas, liquid, solid. The transition from one state of aggregation to another is not accompanied by a change in the stoichiometric composition of the substance, but is necessarily associated with a greater or lesser change in its structure.

Solid state is a state in which matter has its own volume and its own shape. In solids, the forces of interaction between particles are very large. Almost all substances exist in the form of several solid bodies. The reactivity and other properties of these bodies are, as a rule, different. The ideal solid state corresponds to a hypothetical ideal crystal.

liquid state A state in which matter has its own volume but no shape of its own. The liquid has a certain structure. In terms of structure, the liquid state is intermediate between the solid state with a strictly defined periodic structure and the gas, in which there is no structure. Hence, a liquid is characterized, on the one hand, by the presence of a certain volume, and on the other, by the absence of a certain form. The continuous movement of particles in a liquid determines a strongly pronounced self-diffusion and its fluidity. The structure and physical properties of a liquid depend on the chemical identity of its constituent particles.

gaseous state. A characteristic feature of the gas state is that the molecules (atoms) of the gas are not held together, but move freely in the volume. The forces of intermolecular interaction appear when molecules come close to each other. Weak intermolecular interaction causes a low density of gases and their main characteristic properties - the desire for infinite expansion and the ability to exert pressure on the walls of vessels that prevent this desire. Due to the weak intermolecular interaction at low pressure and high temperatures, all typical gases behave approximately the same way, but even at ordinary temperatures and pressure, the individuality of gases begins to appear. The state of a gas is characterized by its temperature, pressure and volume. The gas is considered to be at N.O. if its temperature is 0 degrees and the pressure is 1 * 10 Pa.

- Crystal state. Among solids, the main one is the crystalline state, characterized by a certain orientation of particles (atoms, ions, molecules) relative to each other. This also determines the external form of the substance in the form of crystals. Single crystals - single crystals exist in nature, but they can be obtained artificially. But most often, crystalline bodies are polycrystalline formations - these are intergrowths of a large number of small crystals. A characteristic feature of crystalline bodies, which follows from their structure, is anisotropy. It manifests itself in the fact that the mechanical, electrical and other properties of crystals depend on the direction of the external action of forces on the crystal. Particles in crystals perform thermal vibrations near the equilibrium position or near the nodes of the crystal lattice.

amorphous state. The amorphous state is similar to the liquid state. It is characterized by incomplete ordering of the mutual arrangement of particles. The bonds between structural units are not equivalent, therefore, amorphous bodies do not have a specific melting point - in the process of heating, they gradually soften and melt. For example, the temperature range of melting processes for silicate glasses is 200 degrees. In amorphous bodies, the nature of the arrangement of atoms practically does not change when heated. Only the mobility of atoms changes - their vibrations increase.

- Crystal lattices:

Crystal lattices can be ionic, atomic (covalent or metallic) and molecular.

The ionic lattice consists of ions of opposite sign, alternating at the nodes.

In atomic lattices, atoms are linked by covalent or metallic bonds. Example: diamond (atomic-covalent lattice), metals and their alloys (atomic-metal lattice). The nodes of the molecular crystal lattice are formed by molecules. In crystals, molecules are bound by intermolecular interactions.

Differences in the type of chemical bond in crystals determine significant differences in the type of physical and chemical properties of a substance with all types of crystal lattice. For example, substances with an atomic-covalent lattice are characterized by high hardness, and those with an atomic-metal lattice are characterized by high plasticity. Substances with an ionic lattice have a high melting point and are not volatile. Substances with a molecular lattice (intermolecular forces are weak) are fusible, volatile, their hardness is not high.

22. Complex compounds. Definition. Compound.

Complex compounds are molecular compounds, the combination of components of which leads to the formation of complex ions capable of free existence, both in a crystal and in solution. Complex ions are the result of the interaction between the central atom (complexing agent) and the surrounding ligands. Ligands are both ions and neutral molecules. Most often, the complexing agent is a metal, which, together with ligands, forms an inner sphere. There is an outer sphere. The inner and outer spheres are interconnected by ionic bonding.

There are two types of covalent bonds: sigma and pi bonds. A sigma bond is a single covalent bond formed when the AO overlaps along a straight line (axis) connecting the nuclei of two bonded atoms with a maximum of overlap on this straight line. a sigma bond can arise when any (s-, p-hybrid) AO overlap. In organogens (carbon, nitrogen, oxygen, sulfur), hybrid orbitals can take part in the formation of sigma bonds, providing more efficient overlap. In addition to the axial overlap, another type of overlap is possible - the lateral overlap of p-AO, leading to the formation of a pi bond. A pi-bond is a bond formed by lateral overlap of unhybridized p-AO with a maximum of overlap on both sides of the straight line connecting the nuclei of atoms. Frequently occurring in organic compounds, multiple bonds are a combination of sigma and pi bonds; double - one sigma and one pi, triple - one sigma and two pi bonds.

Bond energy is the energy released when a bond is formed or required to separate two bonded atoms. It serves as a measure of bond strength: the greater the energy, the stronger the bond.

The bond length is the distance between the centers of the bonded atoms. A double bond is shorter than a single bond, and a triple bond is shorter than a double bond. The bonds between carbon atoms in different states of hybridization are characterized by a general pattern: with an increase in the proportion of the s-orbital in the hybrid orbital, the bond length decreases. For example, in the series of compounds propane CH3-CH2-CH3, propene CH3-CH=CH2, propyne CH3-C-=CH, the CH3-C bond length is 0.154, 0.150, and 0.146 nm, respectively.

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 in cases where the number of unpaired electrons in the ground state of an atom is less than the number of bonds formed. It is postulated that different atomic orbitals with similar energy levels interact with each other, forming hybrid orbitals with the same shape and energy. Hybrid orbitals, due to the greater overlap, form stronger bonds than unhybridized orbitals.

The type of hybridization determines the orientation of hybrid AOs in space and, consequently, the geometry of molecules. Depending on the number of orbitals entered into hybridization, the carbon atom can be in one of three states of hybridization. sp3 hybridization. As a result of sp3 hybridization, the carbon atom from the ground state 1s2-2s2-2p2, due to the transfer of an electron from the 2s- to the 2p-orbital, passes into the excited state 1s2-2s1-2p3. Mixing of four outer AOs of an excited carbon atom (one 2s and three 2p orbitals) gives rise to four equivalent sp hybrid orbitals. They have the shape of a volume eight, one of the blades of which is much larger than the other. Due to mutual repulsion, the sp3-hybrid AOs are directed in space to the vertices of the tetrahedron and the angles between them are equal to 109.5° (the most favorable arrangement). Each hybrid orbital in an atom is filled with one electron. The carbon atom in the state of sp3 hybridization has the electronic configuration 1s2(2sp3)4.

Such a state of hybridization is characteristic of carbon atoms in saturated hydrocarbons (alkanes) and, accordingly, in the alkyl radicals of their derivatives. sp2 hybridization. As a result of sp2 hybridization, due to the mixing of one 2s- and two 2p-AO of the excited carbon atom, three equivalent sp2-hybrid orbitals are formed, located in the same plane at an angle of 120'. The unhybridized 2p-AO is in the perpendicular plane. The carbon atom in the state of sp2 hybridization has the electronic configuration 1s2-(2sp2)3-2p1. Such a carbon atom is characteristic of unsaturated hydrocarbons (alkenes), as well as some functional groups, such as carbonyl, carboxyl, and others. As a result of sp hybridization due to mixing of one 2s and one 2p orbitals of the excited carbon atom, two equivalent sp hybrid AOs are formed, located linearly at an angle of 180°. The two 2p-AOs remaining unhybridized are located in mutually perpendicular planes. The carbon atom in the state of sp hybridization has the electronic configuration 1s2-(2sp)2-2p2. Such an atom is found in compounds having a triple bond, for example, in alkynes, nitriles. Atoms of other elements can also be in a hybridized state. For example, the nitrogen atom in the ammonium ion NH4+ and, accordingly, the alkylammonium RNН3+ is in the state of sp3 hybridization; in pyrrole and pyridine - sp2 hybridization; in nitriles - sp-hybridization.



Consists of one sigma and one pi-bond, triple - of one sigma- and two orthogonal pi-bonds.

The concept of sigma and pi bonds was developed by Linus Pauling in the 30s of the last century.

L. Pauling's concept of sigma and pi bonds became an integral part of the theory of valence bonds. At present, animated images hybridizations of atomic orbitals have been developed.

However, L. Pauling himself was not satisfied with the description of sigma and pi bonds. At a symposium on theoretical organic chemistry dedicated to the memory of F. A. Kekule (London, September 1958), he abandoned the σ, π description, proposed and substantiated the theory of a curved chemical bond. The new theory clearly took into account the physical meaning of the covalent chemical bond.

Encyclopedic YouTube

1 / 3

Pi-bonds and hybridized sp2 orbitals

The structure of the carbon atom. Sigma - and pi-bonds. Hybridization. Part 1

Chemistry. Covalent chemical bond in organic compounds. Foxford Online Learning Center

Subtitles

In the last video, we talked about the sigma bond. Let me draw 2 nuclei and orbitals. Here's the sp3 hybrid orbital of this atom, most of it here. And here, too, sp3-hybrid orbital. Here is a small part, here is a large part. A sigma bond is formed where the orbitals overlap. How can another type of connection be formed here? This will require some explanation. This is the sigma bond. It is formed when 2 orbitals overlap on the axis connecting the nuclei of atoms. Another type of bond can be formed by two p-orbitals. I will draw the nuclei of 2 atoms and one p-orbital each. Here are the cores. Now I will draw the orbitals. The P orbital is like a dumbbell. I'll draw them a little closer to each other. Here is a p-orbital in the shape of a dumbbell. This is one of the p-orbitals of the atom. I'll draw more of her. Here is one of the p-orbitals. Like this. And this atom also has a p-orbital parallel to the previous one. Let's say it's like this. Like this. Should have corrected it. And these orbitals overlap. That's it. 2 p-orbitals are parallel to each other. Here hybrid sp3-orbitals are directed at each other. And these are parallel. So p-orbitals are parallel to each other. They overlap here, up and down. This is a P-bond. I will sign. This is a 1 P bond. It is written with a single Greek small letter "P". Well, or so: "P-connection." And this - P bond is formed due to the overlap of p-orbitals. Sigma bonds are regular single bonds, and P bonds are added to them to form double and triple bonds. For a better understanding, consider the ethylene molecule. Its molecule is arranged like this. 2 carbons linked by a double bond, plus 2 hydrogens each. To better understand bond formation, we need to draw orbitals around carbon atoms. So that's it... First I'll draw the sp2 hybrid orbitals. I'll explain what's going on. In the case of methane, 1 carbon atom is bonded to 4 hydrogen atoms, thus forming a three-dimensional tetrahedral structure, like this. This atom is pointed at us. This atom lies in the plane of the page. This atom lies behind the plane of the page, And this one sticks up. This is methane. The carbon atom forms sp3 hybrid orbitals, each of which forms a single sigma bond with one hydrogen atom. Now let's write the electronic configuration of the carbon atom in the methane molecule. Let's start with 1s2. Next should be 2s2 and 2p2, but in fact everything is more interesting. See. There are 2 electrons on the 1s orbital, and instead of 2s and 2p orbitals with 4 electrons in total, they will have sp3 hybrid orbitals: here is one, here is the second, here is the third sp3 hybrid orbital and the fourth. An isolated carbon atom has a 2s orbital and 3 2p orbitals along the x-axis, along the y-axis, and along the z-axis. In the last video, we saw that they mix to form bonds in the methane molecule and the electrons are distributed like this. There are 2 carbon atoms in the ethylene molecule, and at the end it is clear that this is an alkene with a double bond. In this situation, the electronic configuration of carbon looks different. Here's the 1s orbital, and it's still full. It has 2 electrons. And for the electrons of the second shell, I will take a different color. So what's on the second shell? There are no s- and p-orbitals here, because these 4 electrons must be made unpaired in order to form bonds. Each carbon atom forms 4 bonds with 4 electrons. 1,2,3,4. But now the s-orbital hybridizes not with 3 p-orbitals, but with 2 of them. Here is the 2sp2 orbital. S-orbital mixes with 2 p-orbitals. 1s and 2p. And one p-orbital remains the same. And this remaining p-orbital is responsible for the formation of the P-bond. The presence of a P-bond leads to a new phenomenon. The phenomenon of lack of rotation around the axis of communication. Now you will understand. I will draw both carbon atoms in a volume. Now you will understand everything. I'll take a different color for this. Here is a carbon atom. Here is its core. I'll mark it with the letter C, it's carbon. First comes the 1s orbital, this little sphere. Then there are hybrid 2sp2 orbitals. They lie in the same plane, forming a triangle, well, or "pacific". I will show it in scale. This orbital points here. This one is directed there. They have a second, small part, but I won't draw it, because it's easier. They are similar to p-orbitals, but one of the parts is much larger than the second. And the last one is here. It looks a bit like a Mercedes badge if you draw a circle here. This is the left carbon atom. It has 2 hydrogen atoms. Here is 1 atom. There he is, right here. With one electron per 1s orbital. Here is the second hydrogen atom. This atom will be here. And now the right carbon atom. Now we draw it. I'll draw the carbon atoms close together. This is the carbon atom. Here is its 1s orbital. It has the same electronic configuration. 1s orbital around and the same hybrid orbitals. Of all the orbitals of the second shell, I drew these 3. I have not drawn the P-orbital yet. But I will. I'll draw the connections first. The first will be this bond formed by the sp2-hybrid orbital. I'll paint with the same color. This bond is formed by the sp2-hybrid orbital. And this is a sigma bond. The orbitals overlap on the bond axis. Everything is simple here. And there are 2 hydrogen atoms: one bond here, the second bond here. This orbital is slightly larger because it is closer. And this hydrogen atom is here. And that's also sigma bonds, if you notice. The S orbital overlaps with sp2, the overlap lies on the axis connecting the nuclei of both atoms. One sigma bond, the other. Here's another hydrogen atom, also linked by a sigma bond. All bonds in the figure are sigma bonds. In vain I sign them. I will mark them with small Greek letters "sigma". And here too. So this link, this link, this link, this link, this link is a sigma link. And what about the remaining p-orbital of these atoms? They do not lie in the plane of the Mercedes sign, they stick up and down. I'll take a new color for these orbitals. For example, purple. Here is the p-orbital. It is necessary to draw it more, very large. In general, the p-orbital is not that big, but I draw it like this. And this p-orbital is located, for example, along the z-axis, and the rest of the orbitals lie in the xy plane. The z-axis is up and down. The lower parts should also overlap. I'll draw more of them. Like this and like this. These are p orbitals and they overlap. This is how this connection is formed. This is the second component of the double bond. And here it is necessary to explain something. It's a P-bond, and that too. It's all the same P-bond. j The second part of the double bond. What's next? By itself, it is weak, but in combination with a sigma bond, it brings atoms closer together than an ordinary sigma bond. Therefore, a double bond is shorter than a single sigma bond. Now the fun begins. If there was one sigma bond, both groups of atoms could rotate around the axis of the bond. For rotation around the bond axis, a single bond is suitable. But these orbitals are parallel to each other and overlap, and this P-bond does not allow rotation. If one of these groups of atoms rotates, the other rotates with it. The P-bond is part of a double bond, and double bonds are rigid. And these 2 hydrogen atoms cannot rotate separately from the other 2. Their location relative to each other is constant. That's what's happening. I hope you now understand the difference between sigma and p bonds. For a better understanding, let's take the example of acetylene. It is similar to ethylene, but it has a triple bond. One hydrogen atom on each side. Obviously, these bonds are sigma bonds formed by sp orbitals. The 2s orbital hybridizes with one of the p orbitals, the resulting sp hybrid orbitals form sigma bonds, here they are. The remaining 2 bonds are P-bonds. Imagine another p-orbital pointing at us, and here another one, their second halves are directed away from us, and they overlap, and here one hydrogen atom. Maybe I should make a video about it. I hope I didn't confuse you too much.

Ideas about the mechanism of formation of a chemical bond, using the example of a hydrogen molecule, also apply to other molecules. The theory of chemical bonding, created on this basis, was called the method of valence bonds. (MVS).

Basic provisions:

1) a covalent bond is formed as a result of the overlap of two electron clouds with oppositely directed spins, and the formed common electron cloud belongs to two atoms;

2) the covalent bond is the stronger, the more the interacting electron clouds overlap. The degree of overlap of electron clouds depends on their size and density;

3) the formation of a molecule is accompanied by compression of electron clouds and a decrease in the size of the molecule compared to the size of atoms;

4) s- and p-electrons of the outer energy level and d-electrons of the pre-external energy level take part in the bond formation.

In the chlorine molecule, each of its atom has a completed external level of eight electrons s 2 p 6, and two of them (an electron pair) equally belong to both atoms. The overlap of electron clouds during the formation of a molecule is shown in the figure.

Scheme of the formation of a chemical bond in the molecules of chlorine Cl 2 (a) and hydrogen chloride HCl (b)

A chemical bond for which the line connecting the atomic nuclei is the symmetry axis of the bonding electron cloud is called sigma (σ)-bond. It occurs when the "frontal" overlapping of atomic orbitals. Bonds with overlapping s-s-orbitals in the H 2 molecule; p-p orbitals in the Cl 2 molecule and s-p orbitals in the HCl molecule are sigma bonds. Possible "lateral" overlapping of atomic orbitals. When overlapping p-electron clouds oriented perpendicular to the bond axis, i.e. along the y- and z-axes, two areas of overlap are formed, located on both sides of this axis.

This covalent bond is called pi(p)-bond. The overlap of electron clouds during the formation of a π bond is less. In addition, the areas of overlap lie farther from the nuclei than in the formation of a σ-bond. Due to these reasons, the π-bond is less strong than the σ-bond. Therefore, the energy of a double bond is less than twice the energy of a single bond, which is always a σ bond. In addition, the σ-bond has axial, cylindrical symmetry and is a body of revolution around the line connecting the atomic nuclei. The π-bond, on the contrary, does not have cylindrical symmetry.

A single bond is always a pure or hybrid σ bond. A double bond consists of one σ- and one π-bonds located perpendicular to each other. The σ-bond is stronger than the π-bond. In compounds with multiple bonds, there is always one σ-bond and one or two π-bonds.