Consists of one sigma and one pi bond, a triple bond consists 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 was included integral part into the theory of valence bonds. Animated images of atomic orbital hybridization have now 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. New theory clearly took into account physical meaning covalent chemical bond.
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Pi bonds and sp2 hybridized orbitals
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 sigma communications. Let me draw 2 nuclei and orbitals. This is the sp3 hybrid orbital of this atom, most of it is here. And here too there is an sp3 hybrid orbital. Here is a small part of it, here is a large part. Where the orbitals overlap, a sigma bond is formed. How can a different type of connection be formed here? To do this, you will have to explain something. This is the sigma connection. It is formed when two 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. Here are the kernels. 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 it. 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. It would be necessary to correct it. And these orbitals overlap. Just like that. The 2 p orbitals are parallel to each other. Here are the hybrid sp3 orbitals directed towards each other. And these are parallel. So the p orbitals are parallel to each other. They overlap here, above and below. This is a P-bond. I'll sign it. This is 1 P-connection. It is written with one Greek small letter "P". Or so: “P-connection”. And this P bond is formed due to the overlap of p-orbitals. Sigma bonds are ordinary 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 structured like this. 2 carbon atoms linked by a double bond, plus 2 hydrogen atoms each. To better understand bond formation, we need to diagram the orbitals around the carbon atoms. So... First I'll draw the sp2 hybrid orbitals. I'll explain what's happening. In the case of methane, 1 carbon atom is bonded to 4 hydrogen atoms, forming a three-dimensional tetrahedral structure, like this. This atom is directed towards 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 describe the electronic configuration of the carbon atom in the methane molecule. Let's start with 1s2. Next should go 2s2 and 2p2, but in fact everything is more interesting. Look. There are 2 electrons in the 1s orbital, and instead of 2s and 2p orbitals with 4 electrons, they will have sp3 hybrid orbitals in total: 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 it is an alkene with a double bond. In this situation, the electron 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 or p orbitals here because these 4 electrons must be made unpaired 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 a 2sp2 orbital. The S orbital mixes with 2 p orbitals. 1 s and 2 p. 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 connection axis. Now you will understand. I'll draw both carbon atoms in volume. Now you will understand everything. I'll take a different color for this. Here's a carbon atom. Here is its core. I'll label it C, which is carbon. First comes the 1s orbital, this little sphere. Then there are the hybrid 2sp2 orbitals. They lie in the same plane, forming a triangle, or “pacific”. I'll show it in full. This orbital is directed 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 other. And the last one is sent here. It looks a bit like the Mercedes logo if you draw a circle here. This is the left-handed carbon atom. It has 2 hydrogen atoms. Here is 1 atom. There he is, right here. With one electron in the 1s orbital. Here's the second one hydrogen atom. This atom will be here. And now the right carbon atom. Now let's draw it. I'll draw the carbon atoms close together. This carbon atom here. Here is its 1s orbital. He 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 do it. First I'll draw the connections. The first one will be this bond formed by the sp2 hybrid orbital. I'll paint it with the same color. This bond is formed by an sp2 hybrid orbital. And this is a sigma connection. 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 these are also sigma connections, if you noticed. The S orbital overlaps with sp2, the overlap lies on the axis connecting the nuclei of both atoms. One sigma connection, the second. Here's another hydrogen atom, also connected by a sigma bond. All bonds in the figure are sigma bonds. I shouldn't sign them. I will mark them with small Greek letters “sigma”. And here too. So this bond, this bond, this bond, this bond, this bond are sigma bonds. What about the remaining p-orbital of these atoms? They do not lie in the plane of the Mercedes sign, they stick out up and down. I'll take a new color for these orbitals. For example, purple. This is the p orbital. We need to draw it bigger, very big. In general, the p-orbital is not that large, but I draw it like this. And this p-orbital is located, for example, along the z axis, and the remaining orbitals lie in the xy plane. And the z axis is directed up and down. The bottom 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 we need to clarify something. It's a P-bond and that too. It's all the same P-connection. j Second part of the double bond. What's next? By itself it is weak, but in combination with the sigma bond it brings atoms closer together than a regular sigma bond. Therefore, a double bond is shorter than a single sigma bond. Now the fun begins. If there were one sigma bond, both groups of atoms could rotate around the bond axis. For rotation around the coupling axis, a single coupling is suitable. But these orbitals are parallel to each other and overlap, and this P-bond prevents 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 better understanding, let's look at the example of acetylene. It is similar to ethylene, but it has a triple bond. There is a hydrogen atom on each side. It is obvious that 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 directed towards us, and here is another one, their second halves are directed away from us, and they overlap, and here there is one hydrogen atom each. Maybe I should make a video about this. I hope I haven't confused you too much.
14. Basic characteristics of covalent bonds. Bond length and energy. Saturation and direction. Multiplicity of communication. Sigma and pi connections.
- A chemical bond carried out by shared electron pairs is called atomic or covalent. Each covalent chemical bond has certain qualitative or quantitative characteristics. These include:
Link length
Communication energy
Saturability
Communication direction
Communication polarity
Communication multiplicity
- Link length– the distance between the nuclei of bonded atoms. It depends on the size of the atoms and the degree of overlap of their electron shells. The length of a bond is determined by the order of the bond: the higher the order of the bond, the shorter its length.
Communication energy is the energy that is released when a molecule is formed from single atoms. It is usually expressed in J/mol (or cal/mol). Bond energy is determined by the bond order: the greater the bond order, the greater its energy. Bond energy is a measure of its strength. Its value is determined by the work required to break a bond, or the gain in energy when a substance is formed from individual atoms. The 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 combine in a molecule, then the average binding energy does not coincide with the dissociation energy of the molecule. Bond energies in molecules consisting of identical atoms decrease in groups from top to bottom. The bond energies increase over the period.
- Saturability– shows how many bonds a given atom can form with others due to shared electron pairs. It is equal to the number of common electron pairs with which a given atom is connected to others. Saturability covalent bond is the ability of an atom to participate in the formation of a limited number of covalent bonds.
Focus- this is a certain mutual arrangement of connecting 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 bonds formed, which are called bond angles.
- Multiplicity of communication. Determined by the number of electron pairs involved in the bond between atoms. If a bond is formed by more than one pair of electrons, then it is called multiple. As the bond multiplicity increases, the energy increases and the bond length decreases. In molecules with a multiple bond there is no rotation around an axis.
- Sigma and pi bonds. The chemical bond is caused by the overlap of electron clouds. If this overlap occurs along a line connecting the atomic nuclei, then the 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– these are always sigma connections. S-type orbitals form only sigma bonds. But a large number of compounds are known that have double and even triple bonds. One of them is sigma bond and the others are called pi bonds. When such bonds are formed, overlapping electron clouds occur in two regions of space symmetrical to the internuclear axis.
15. Hybridization of atomic orbitals using the example of molecules: methane, aluminum chloride, beryllium chloride. Bond angle and molecular geometry. Molecular orbital method (MO LCAO). Energy diagrams of homo- and hetero-nuclear molecules (N2, Cl2, N.H.3, 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 an 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 involving d-orbitals.
Let's consider the linear BeCl2 molecule. A beryllium atom in the valence state is capable of forming two bonds due to one s- and one p-electron. Obviously, this should result in two bonds with chlorine atoms of different lengths, since the radial distribution of these electrons is different. The real BeCl2 molecule is symmetrical and linear; its two Be-Cl bonds are exactly the same. This means that they are provided with electrons of the same state, i.e. here the beryllium atom in the valence state no longer has one s- and one p-electron, but two electrons located in orbitals formed by the “mixing” of s- and p-atomic orbitals. A methane molecule will have sp3 hybridization, and an aluminum chloride molecule will have sp2 hybridization.
Conditions for hybridization stability:
1) Compared to the original orbital atoms, the hybrid orbitals should overlap more closely.
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 table.
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Hybridization |
Molecule shape |
Bond angle | |
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Linear | |||
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Triangle | |||
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Tetrahedron | |||
- 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, it is necessary to take into account the following rules:
1. If an atomic orbital is a certain function that is a solution to the Schrödinger Equation and describes the state of an electron in an atom, the MO method is also a solution to 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 subtraction of functions, the molecular orbital has a higher energy, and it is called loosening.
There are sigma and pi orbitals. They are filled in 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. Dipole moment of connection. Characteristics of interacting atoms: ionization potential, electron affinity, electronegativity. The degree of ionicity of the bond.
- Dipole moment- a 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 the negative charge to the positive one. The dipole moment of a chemical bond is caused by the displacement of the electron cloud towards one of the atoms. A bond is called polar if the corresponding dipole moment differs significantly from zero. There are cases when individual bonds in a molecule are polar, and the total dipole moment of the molecule is zero; such molecules are called non-polar (for example, CO 2 and CCl 4 molecules). If the dipole moment of a molecule is non-zero, the molecule is called 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 (about 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.
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Ionization energy of an atom is the amount of energy required to remove an electron from an unexcited atom. It is expressed in kilojoules per mole. For multi-electron atoms, the ionization energies E1, E2, E3, ..., En correspond to the separation of the first, second, etc. electrons. In this case, always E1 -
Atom electron affinity– the energetic effect of adding an electron to a neutral atom, transforming it into a negative ion. The electron affinity of an atom is expressed in kJ/mol. Electron affinity is numerically equal but opposite in sign to the ionization energy of a negatively charged ion and depends on the electronic 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) configuration or half-filled with a p-sublayer (N, P, As) do not exhibit electron affinity. -
Electronegativity- an averaged 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 different compounds is neglected. Unlike ionization potential and electron affinity, EO is not a strictly defined physical quantity, but a useful conditional characteristic. The most electronegative element is fluorine. EO depends on ionization energy and electron affinity. According to one definition, the EO of an atom can be expressed as half the sum of its ionization energy and electron affinity. An element cannot be assigned a constant EO. It depends on many factors, in particular on the valence state of the element, the type of compound in which it is included, etc. 17. Polarizing ability and polarizing effect. Explanation of some physical properties of substances from the point of view of this theory. -
Polarization theory 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 effect on another ion The polarizability of an ion is a measure of the ability of the ion's electron cloud to deform under the influence of an external electric field. Regularities of ion polarizability: Anions are more polarized than cations. Excessive electron density leads to high diffuseness and looseness of the electron cloud. The polarizability of isoelectronic ions increases with a decrease in positive and increasing negative charges. Isoelectronic ions have the same configuration. In multiply charged cations, the nuclear charge is much greater than the number of electrons. This compacts the electron shell and stabilizes it, so such ions are less susceptible to deformation. The polarizability of cations decreases upon transition from ions with an outer electron shell filled with 18 electrons to an unfilled one, and then 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 surrounding anions. The polarizability of analogue ions increases with increasing number of electronic layers. Polarizability is most difficult for small-sized and multiply charged cations, with an electron shell of noble gases. Such cations are called hard. Bulk multicharged anions and low-charge bulk cations are most easily polarized. These are soft ions. -
Polarizing effect. Depends on the charges, size and structure of the outer electronic layer. 1. The polarizing effect of a cation increases with increasing its charge and decreasing radius. The maximum polarizing effect is characteristic of Katons with small radii and large charges, so they form covalent-type compounds. The greater the charge, the greater the polarizing bond. 2. The polarizing effect of cations increases with the transition from ions with an s-electron cloud to an incomplete one and to an 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 of compounds with covalent bonds will be lower than those with ionic bonds. The higher 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 perspective of the theory of polarization, the color of compounds can be explained: white reflects everything; black – absorbs; transparent – lets through. This is related: if the shell is deformed, then the quantum levels of the electrons move closer together, reducing the energy barrier, so low 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 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 greatest polarizing effect is found in d-elements. The greater the charge, the greater the polarizing effect. 18. Ionic bond as a limiting case of a 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 determined by 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 p-elements of group 7, due to their high electron affinity. The electrical charges of ions cause their attraction and repulsion. Ions can be thought of as charged balls, the force fields of which 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 characterized by non-directionality. The interaction of ions of opposite signs 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. 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 a bond in which the electrons of each individual atom belong to all the atoms in contact. The energy difference between “molecular” orbitals in such a bond is small, so electrons can easily move from one “molecular” orbital to another and, therefore, move in the volume of the metal. Metals differ from other substances in their high electrical 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. easily move in the crystal. The ability of electrons to move around a 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, in metals there is a highly delocalized chemical bond. The metallic bond has a certain similarity to the covalent bond, since it is based on the sharing of valence electrons. However, in the formation of a covalent bond, the valence electrons of only two interacting atoms participate, while in the formation of a metallic bond, all atoms take part in the sharing of electrons. That is why the metallic bond does not have spatial directionality 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. Mechanism of education. Features of the physical properties of substances with hydrogen bonds. Examples. -
A 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). Hydrogen bonding is very common and plays an important role in the association of molecules, in the processes of crystallization, dissolution, formation of crystal 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 the hydrogen bond. Its peculiarity is that a 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, hydrogen bonding was interpreted as an electrostatic interaction. It has now been concluded that the donor-acceptor interaction plays a major role in hydrogen bonding. Hydrogen bonds are formed not only between molecules of different substances, but also in 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 within molecules is known, especially in organic compounds. Its formation is facilitated by the presence in the molecule of an acceptor group A-H and a donor group B-R. In the A-H molecule, A is the most electronegative element. The formation of hydrogen bonding in polymers, such as peptides, results in a helical structure. DNA, deoxyribonucleic acid, the keeper of the code of heredity, has similar structures. Hydrogen bonds are not strong. They easily form and break at ordinary temperatures, 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 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: 1. The presence of a permanent dipole in the molecule. In this case, as a result of the 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 be formed, which arises as a result of the polymerization of molecules by the 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. London introduced this concept. For symmetric atoms, the only forces acting 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 consists only of 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 - 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- This is a state in which a substance has its own volume and its own shape. In solids, the interaction forces between particles are very strong. Almost all substances exist in the form of several solids. The reactivity and other properties of these bodies are usually different. The ideal solid state corresponds to a hypothetical ideal crystal. Liquid state- This is a state in which a substance has its own volume, but does not have its own shape. The liquid has a certain structure. In structure, the liquid state is intermediate between a solid state with a strictly defined periodic structure and a 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 shape. The continuous movement of particles in a liquid determines strongly pronounced self-diffusion and its fluidity. The structure and physical properties of a liquid depend on the chemical identity of the particles that form it. 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. Intermolecular interaction forces occur when molecules come close to each other. Weak intermolecular interaction determines the 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 impede this desire. Due to the weak intermolecular interaction at low pressure and high temperatures, all typical gases behave approximately the same, but already 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 ambient conditions. if its temperature is 0 degrees and the pressure is 1* 10 Pa. -
Crystalline 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, resulting 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 influence of forces on the crystal. Particles in crystals undergo thermal vibrations around the equilibrium position or around 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 relative arrangement of particles. The bonds between structural units are not equivalent, therefore amorphous bodies do not have a specific melting point - during the heating process 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 signs alternating at the sites. In atomic lattices, atoms are connected by covalent or metallic bonds. Example: diamond (atomic-covalent lattice), metals and their alloys (atomic-metal lattice). The nodes of a molecular crystal lattice are formed by molecules. In crystals, molecules are connected through 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, and 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 interactions 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 the inner sphere. There is an outer sphere. The inner and outer spheres are interconnected by an ionic bond. Ideas about the mechanism of chemical bond formation based on the example of the hydrogen molecule extend to other molecules. The theory of chemical bonding, created on this basis, is called the valence bond method. (MVS). Key points: 1) a covalent bond is formed as a result of the overlap of two electron clouds with oppositely directed spins, and the resulting common electron cloud belongs to two atoms; 2) the stronger the covalent bond, the more the interacting electron clouds overlap. The degree to which electron clouds overlap 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 external energy level and d-electrons of the pre-external energy level take part in the formation of the bond. In a chlorine molecule, each of its atoms has a complete outer level of eight electrons s 2 p 6, and two of them (electron pair) belong equally 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 molecules of chlorine Cl 2 (a) and hydrogen chloride HCl (b) A chemical bond for which the line connecting the atomic nuclei is the axis of symmetry of the connecting electron cloud is called sigma (σ)-bond. It occurs when atomic orbitals overlap head-on. Bonds when s-s orbitals overlap in the H 2 molecule; p-p-orbitals in the Cl 2 molecule and s-p-orbitals in the HCl molecule are sigma bonds. “lateral” overlap of atomic orbitals is possible. When overlapping p-electron clouds oriented perpendicular to the bond axis, i.e. along the y- and z-axis, two overlap regions are formed, located on either side of this axis. This covalent bond is called pi (p)-bond. There is less overlap of electron clouds during π bond formation. In addition, the overlap regions lie further from the nuclei than during the formation of a σ bond. Due to these reasons, the π bond has less strength compared to 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-bond, 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. Sigma and pi bonds (σ- and π-bonds)
covalent chemical bonds characterized by a specific, but different spatial symmetry of the electron density distribution. As is known, a covalent bond is formed as a result of the sharing of electrons of interacting atoms. The resulting electron cloud of the σ bond is symmetrical with respect to the bond line, that is, the line connecting the nuclei of interacting atoms. Simple bonds in chemical compounds are usually (t-bonds (see Simple communication). The electron cloud of the π bond is symmetrical relative to the plane passing through the communication line ( rice. 1
, b), and in this plane (called the nodal plane) the electron density is zero. The use of Greek letters σ and π is associated with their correspondence to Latin letters s And r in the designation of electrons of the atom, with the participation of which for the first time it becomes possible to form σ- and π-bonds, respectively. Because clouds of atomic r-orbitals ( p x, r y, p z) are symmetrical about the corresponding axes of Cartesian coordinates ( X, at, z), then if one r-orbital, for example p z, takes part in the formation of the σ bond (axis z- communication line), the remaining two r-orbitals ( p x, p y) can take part in the formation of two π-bonds (their nodal planes will be yz And xz respectively; cm. rice. 2
). Can also take part in the formation of σ and π bonds d- (cm. rice. 1
) And f-electrons of the atom. Lit.: Pimentel G., Spratly R., How quantum mechanics explains chemical bonding, trans. from English, M., 1973; Shustorovich E. M., Chemical communication, M., 1973. E. M. Shustorovich. Rice. 1. Schematic representation of the spatial orientation of orbitals during the formation of a σ bond as a result of s - s-, s - p σ-, p σ - p σ -interactions (a) and π-bond as a result of p π -, p π -, d π - d π - interactions (b). Rice. 2. Schematic representation of clouds of p x -, p y -, p z - electrons. The axes of Cartesian coordinates and the nodal planes of p x - and p y -orbitals are shown. Great Soviet Encyclopedia. - M.: Soviet Encyclopedia.
1969-1978
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- (models) models of field theory, in which m scalar fields (i=1, ..., m) can be considered as defining a mapping of d-dimensional time space (of an arbitrary signature) into a certain manifold M of dimension with a metric... Physical encyclopedia Fig. 1. Sigma connection ... Wikipedia Greek alphabet Αα Alpha Νν Nu ... Wikipedia Sigma (σ) - and pi (π) - connections- a covalent chemical property characterized by a certain, but different spatial symmetry of the electron density distribution. The resulting electron cloud σ communication is symmetrical relative to the communication line,... ... Encyclopedic Dictionary of Metallurgy - (from Latin cumulo collect, accumulate) a system of bonds in which at least one atom is connected by double bonds to two neighboring atoms. K. s. in the Sigma and Pi connection group)). σ bonds are formed by two atomic orbitals of the C atom in... ... Covalent bond using the example of a methane molecule: complete outer energy level: hydrogen (H) has 2 electrons, and carbon (C) has 8 electrons. Covalent bond is a bond formed by directed valence electron clouds. Neutral... ... Wikipedia delta-sigma modulator- Modification of a delta modulator, at the input of which an integrator is turned on, and upon reception the reverse operation is performed, i.e. differentiation of the demodulator output signal. From an engineering point of view, the implementation of a delta-sigma modulator is no more difficult than... ... Technical Translator's Guide project Sigma- A project separated in 1976 from the secret American project Aquarius. The goal of the project is to establish communication with aliens and is probably carried out at one of the Air Force bases in the state. New Mexico. E. Project Sigma D. Project Sigma ... Explanatory ufological dictionary with equivalents in English and German Type Open share... Wikipedia One of the most important types of intramolecular mutual influence of atoms and bonds in organic compounds; is caused by the interaction of electronic systems of atoms (primarily valence electrons, see Valence). The main sign... ... Great Soviet EncyclopediaSee what “Sigma and pi bonds” are in other dictionaries:
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