Energy of connection with n. Basic types of chemical bonds. Characteristics of chemical bonding. Energy of communication. Link length. Chemical bond length

When a chemical bond is formed, a redistribution in space of electron densities that originally belonged to different atoms occurs. Since the electrons of the outer level are least tightly bound to the nucleus, these electrons play the main role in the formation of a chemical bond. The number of chemical bonds formed by a given atom in a compound is called valency. The electrons that take part in the formation of a chemical bond are called valence: for s- and p elements these are the outer electrons, for d-elements the outer (last) s-electrons and the penultimate d-electrons. From an energy point of view, the most stable atom is the one whose outer level contains the maximum number of electrons (2 and 8 electrons). This level is called completed. Completed levels are highly durable and are characteristic of noble gas atoms, so under normal conditions they are in the state of a chemically inert monatomic gas.

Atoms of other elements have incomplete external energy levels. In the process of a chemical reaction, the completion of external levels is achieved, which is achieved either by the addition or loss of electrons, as well as by the formation of common electron pairs. These methods lead to the formation of two main types of bonds: covalent and ionic. Thus, when a molecule is formed, each atom tends to acquire a stable outer electron shell: either two-electron (doublet) or eight-electron (octet). This pattern forms the basis of the theory of chemical bond formation. The formation of a chemical bond due to the completion of external levels in the atoms forming the bond is accompanied by the release of a large amount of energy, that is, the formation of a chemical bond always proceeds exothermically, since it leads to the appearance of new particles (molecules) that, under normal conditions, are more stable and, therefore, smaller energy than the original ones. One of the essential indicators that determine what kind of bond is formed between atoms is electronegativity, that is, the ability of an atom to attract electrons from other atoms. The electronegativity of atoms of elements changes gradually: in the periods of the periodic system from left to right its value increases and in groups from top to bottom it decreases.

A chemical bond carried out due to the formation of common (bonding) electron pairs is called covalent.1) Let us consider an example of the formation of a chemical bond between atoms with the same electronegativity, for example, the hydrogen molecule H2. The formation of a chemical bond in a hydrogen molecule can be represented as two points: H- + -H -> H: H or a dash that symbolizes a pair of electrons: H-H A covalent bond formed by atoms with the same electronegativity is called non-polar. Such a bond is formed by diatomic molecules consisting of atoms of one chemical element: H 2, Cl 2, etc. 2) Formation of a covalent bond between atoms whose electronegativity differs slightly. A covalent bond formed by atoms with different electronegativity is called polar. In a covalent polar bond, the electron density from a shared pair of electrons is shifted to the atom with higher electronegativity. Examples include the molecules H2O, NH3, H2S, CH3Cl. The covalent (polar and nonpolar) bond in our examples was formed due to the unpaired electrons of the bonding atoms. This mechanism of covalent bond formation is called exchange. Another mechanism for the formation of a covalent bond is donor-acceptor. In this case, the bond occurs due to two paired electrons of one atom (donor) and a free orbital of another atom (acceptor). A well-known example is the formation of ammonium ion: H++:NH 3 -> [H: NH3 | +<=====>NH4+ acceptor ammonium ion electron donor. When an ammonium ion is formed, the electron pair of nitrogen becomes common to the N and H atoms, that is, a fourth bond appears, which is no different from the other three. They are depicted the same way:

An ionic bond occurs between atoms whose electronegativity differs sharply. Let us consider the method of formation using sodium chloride NaCl as an example. The electronic configuration of sodium and chlorine atoms can be represented as: 11 Na ls2 2s2 2p 6 3s1; 17 Cl ls2 2p 6 Зs2 3р5 Like these are atoms with incomplete energy levels. Obviously, to complete them, it is easier for a sodium atom to give up one electron than to gain seven, and for a chlorine atom it is easier to gain one electron than to give up seven. During a chemical interaction, the sodium atom completely gives up one electron, and the chlorine atom accepts it. Schematically, this can be written as follows: Na. -- l e --> Na+ sodium ion, stable eight-electron 1s2 2s2 2p6 shell due to the second energy level. :Cl + 1е -->.Cl - chlorine ion, stable eight-electron shell. Electrostatic attraction forces arise between the Na+ and Cl- ions, resulting in the formation of a compound.

A chemical bond that occurs through electrostatic attraction between ions is called an ionic bond. Compounds formed by the attraction of ions are called ionic. Ionic compounds consist of individual molecules only in the vapor state. In the solid (crystalline) state, ionic compounds consist of regularly arranged positive and negative ions. In this case there are no molecules. Ionic compounds form elements of the main subgroups of groups I and II and the main subgroups of groups VI and VII that are sharply different in electronegativity. There are relatively few ionic compounds. For example, inorganic salts: NH4Cl (ammonium ion NH4 + and chlorine ion Cl-), as well as salt-like organic compounds: alcoholates, carboxylic acid salts, amine salts. Nonpolar covalent bonds and ionic bonds are two limiting cases of electron density distribution. A nonpolar bond corresponds to a uniform distribution of the connecting two electron clouds between identical atoms. On the contrary, in an ionic bond the connecting electron cloud belongs almost entirely to one of the atoms. In most compounds, chemical bonds are intermediate between these types of bonds, that is, they contain a polar covalent bond.

Metallic bonding exists in metals in the solid and liquid states. In accordance with their position in the periodic table, metal atoms have a small number of valence electrons (1-3 electrons) and low ionization energy (electron removal). Therefore, valence electrons are weakly held in the atom, easily detached and have the ability to move throughout the crystal. At the nodes of the crystal lattice of metals there are free atoms, positively charged electrons, and some of the valence electrons, moving freely in the volume of the crystal lattice, form an “electron gas” that ensures the connection between the metal atoms. The bond carried out by relatively free electrons between metal ions in a crystal lattice is called a metallic bond. A metallic bond arises due to the sharing of valence electrons by atoms. However, there is a significant difference between these types of communication. The electrons that perform a covalent bond generally reside in the immediate vicinity of the two bonded atoms. In the case of a metallic bond, the electrons that perform the bond move throughout the piece of metal. This determines the general characteristics of metals: metallic luster, good conductivity of heat and electricity, malleability, ductility, etc. A general chemical property of metals is their relatively high reducing ability.

Hydrogen bonds can form between a hydrogen atom bonded to an atom of an electronegative element and an electronegative element having a free pair of electrons (O,F,N). The hydrogen bond is caused by electrostatic attraction, which is facilitated by the small size of the hydrogen atom, and partly by donor-acceptor interaction. Hydrogen bonding can be intermolecular or intramolecular. 0-H bonds have a pronounced polar character: A hydrogen bond is much weaker than an ionic or covalent bond, but stronger than an intermolecular interaction. Hydrogen bonds determine some physical properties of substances (for example, high boiling points). Hydrogen bonds are especially common in molecules of proteins, nucleic acids and other biologically important compounds, providing them with a certain spatial structure (organization).

Binding energy (Eb). The amount of energy released during the formation of a chemical bond is called chemical bond energy [kJ/mol]. For polyatomic compounds, its average value is taken. The more Eb, the more stable the molecule.

Communication length (lсв). The distance between cores in a connection. The longer the bond length, the lower the bond energy.

Valence bond method.

A) a chemical bond between two atoms arises as a result of the overlap of AOs with the formation of electron pairs.
B) atoms entering into a chemical bond exchange electrons with each other, which form bonding pairs. The energy of electron exchange between atoms (the energy of attraction of atoms) makes the main contribution to the energy of a chemical bond. An additional contribution to the binding energy is made by the Coulomb forces of particle interaction.
C) in accordance with the Pauli principle, a chemical bond is formed only when electrons with different spins interact.
D) the characteristics of a chemical bond (energy, length, polarity) are determined by the type of overlapping AO.

Valence bond method. The covalent bond is directed towards maximum overlap of the AOs of the reacting atoms.

Valence. The ability of an atom to attach or replace a certain number of other atoms to form chemical bonds.

When transitioning to an excited state, one of the paired electrons moves into an empty orbital of the same shell.

Donor-acceptor mechanism: a common electron pair is formed due to the lone pair of electrons of one atom and the vacant orbital of another atom.

Molecular orbital method. The electrons in the molecule are distributed among MOs, which, like AOs, are characterized by a certain energy and shape. MOs span the entire molecule. The molecule is considered as a single system.

1. The number of MOs is equal to the total number of AOs from which the MOs are combined.
2. The energy of some MOs turns out to be higher, while others are lower than the energy of the original AOs. The average energy of MOs obtained from a set of AOs approximately coincides with the average energy of these AOs.
3. Electrons fill the MO, like the AO, in order of increasing energy, while the Pauli exclusion principle and Hund’s rule are observed.
4. AOs are most effectively combined with those AOs that are characterized by comparable energies and corresponding symmetry.
5. As in the BC method, the bond strength in the MO method is proportional to the degree of overlap of atomic orbitals.

Order and bond energy. The order of communication is n=(Nсв-Нр)/2. Nb is the number e in bonding molecular orbitals, Nр is the number e in antibonding molecular orbitals.

If Nсв = Nр, then n=0 and the molecule is not formed. As n increases, the binding energy in molecules of the same type increases. Unlike the AO method, the MO method allows that a bond can be formed by one electron.

Complex connections. Complex compounds that have covalent bonds formed according to the donor-acceptor mechanism

Communication energy is the energy that is released when a molecule is formed from single atoms. Binding energy is the energy that is absorbed when two atoms move an infinite distance away from each other. And the enthalpy of formation is the heat that is released when a substance is obtained from simple substances, that is, if we speak in the language of binding energies, first the atoms of simple substances are spread over an infinitely large distance (with the absorption of energy), then they combine to form the desired substance (energy is released ). The difference is the enthalpy of formation.

The binding energy differs from ΔH arr. The heat of formation is the energy that is released or absorbed during the formation of molecules from simple substances. So:

N 2 + O 2 → 2NO + 677.8 kJ/mol – ∆H arr.

N + O → NO - 89.96 kJ/mol – E St.

For polyatomic molecules with one type of bond, for example, for molecules AB n, the average bond energy is equal to 1/n part of the total energy of formation of a compound from atoms. Thus, the energy of formation of CH 4 = -1661.1 kJ/mol. Since there are four bonds in the CH 4 molecule, the energy of one C – H bond is 415.3 kJ/mol. An examination of the large body of currently known data on binding energies shows that the binding energy between a particular pair of atoms is often constant, provided that the rest of the molecule changes little. Thus, in saturated hydrocarbons Eb (C – H) = 415.3 kJ/mol, Eb (C – C) = 331.8 kJ/mol.

Bond energies in molecules consisting of identical atoms decrease in groups from top to bottom. Bond energies increase over the period. Electron affinity also increases in the same direction.

In the last paragraph we gave an example of calculating the thermal effect of a reaction:

C(tv) + 2 H 2 (g) = CH 4 (g) + 76 kJ/mol.

In this case, 76 kJ is not just the thermal effect of this chemical reaction, but also heat of formation of methane from elements .

ENTHALPY is the heat effect of a reaction, measured (or calculated) for the case when the reaction occurs in an open vessel (i.e. at constant pressure). Denoted as ΔH.

When the volume occupied by the reaction products is different from the volume occupied by the reactants, the chemical system can perform additional work PΔV (where P is the pressure and ΔV is the change in volume). Therefore, ΔH and ΔE are related to each other by the relationship:

ΔH = ΔE + PΔV

So, if the reaction is not carried out in a “bomb,” then ENTHALPY and THERMAL EFFECT coincide with each other. Enthalpy is also called "heat content". If we carry out the reaction to produce water in an open vessel, then 286 kJ/mol is the “heat” ΔH contained in hydrogen and oxygen for the case when we obtain water from them. Since the starting substances (hydrogen and oxygen) were in our experiment under standard conditions (25 o C and a pressure of 1 atm), and we also brought the reaction product (water) to standard conditions, we have the right to say that 286 kJ/mol is STANDARD HEAT OF FORMATION OF WATER or, what is the same - STANDARD ENTHALPY OF FORMATION OF WATER. If we obtain from the same elements not water, but hydrogen peroxide H 2 O 2, then the “heat content” of such a chemical system will be different (187.6 kJ/mol). During reactions that produce 1 mole of water or 1 mole of H 2 O 2, different amounts of energy are released, as would be expected. In what follows, we will more often refer to the standard heat of formation of substances as standard enthalpy of formation ΔH. To emphasize the validity of this value only for standard conditions, in the tables it is designated as follows: ΔН about 298

The small “zero” next to ΔH traditionally symbolizes a certain standard state, and the number 298 reminds us that the values are given for substances at 25 o C (or 298 K). Standard enthalpy not necessary must be the enthalpy of formation of the substance from elements. You can get the standard enthalpy value ΔH about 298 for any chemical reaction. But in our case, with the production of water from hydrogen and oxygen, we received exactly the standard enthalpy of formation of water. It is written like this: H 2 + 0.5 O 2 = H 2 O (ΔH o 298 = -286 kJ/mol)

Where does the minus sign in front of the thermal effect value come from? Here the author, with a sigh, must inform the reader about another feature of the representation of heat (and enthalpy) in thermodynamics. It's accepted here lost represent energy by any system with a minus sign. Consider, for example, the already familiar system of methane and oxygen molecules. As a result exothermic reactions occur between them allocation heat: CH 4 (g) + 2 O 2 (g) = CO 2 (g) + 2 H 2 O (l) + 890 kJ

This reaction can also be written by another equation, where the released ("lost") heat has a minus sign: CH 4 (g) + 2 O 2 (g) – 890 kJ = CO 2 (g) + 2 H 2 O (l )

According to tradition, the enthalpy of this and others exothermic reactions in thermodynamics are usually written with the sign "minus": ΔH o 298 = –890 kJ/mol (energy released).

On the contrary, if as a result endothermic reaction system absorbed energy, then the enthalpy of such an endothermic reaction is written with the sign "plus". For example, for the already familiar reaction of producing CO and hydrogen from coal and water (when heated): C(solid) + H 2 O(g) + 131.3 kJ = CO(g) + H 2 (g)

(ΔH o 298 = +131.3 kJ/mol)

You just need to get used to this feature of the thermodynamic language, although at first, confusion with signs can be quite annoying when solving problems.

Let's try to solve the same problem first in thermodynamic scale (where the heat released by the reaction has a minus sign), and then in thermochemical scale (which we used in the previous paragraph and where the energy released by the reaction has a plus sign).

So, here is an example of calculating the thermal effect of a reaction: Fe 2 O 3 (s) + 3 C (graphite) = 2 Fe (s) + 3 CO (g)

This reaction occurs in a blast furnace at a very high temperature (about 1500 o C). In reference books where it is used thermodynamic scale, you can find the standard heats of formation of Fe 2 O 3 (ΔH o 298 = –822.1 kJ/mol) and CO (ΔH o 298 = – 110.5 kJ/mol). The other two substances in this equation, carbon and iron, are elements, meaning their heat of formation is by definition zero. Therefore, the standard heat of the reaction under consideration is:

ΔH o 298 = 3× (-110.5) - (-822.1) = -331.5 + 822.1 = +490.6 kJ

So, the reduction reaction of iron(III) oxide carbon is endothermic(ΔH o 298 is positive!), and it would be necessary to spend 490.6 kJ to reduce one mole of Fe 2 O 3 with three moles of carbon if the starting substances before the start of the reaction and the products after the end of the reaction are under standard conditions (that is, at room temperature and atmospheric pressure). It doesn't matter that the starting materials had to be very heated for the reaction to occur. The value ΔH o 298 = +490.6 kJ reflects the “pure” thermal effect of an endothermic reaction, in which the reactants were first heated by an external heat source from 25 to 1500 o C, and at the end of the reaction the products cooled again to room temperature, releasing all the heat to the environment . In this case, the heat released will be less than what had to be spent on heating, because part of the heat was absorbed in the reaction.

Let's do the same calculation using thermochemical scale. Suppose the heats of combustion of carbon and iron in oxygen are known (at constant pressure):

1) C + 1/2 O 2 = CO + 110.5 kJ

2) 2 Fe + 3/2 O 2 = Fe 2 O 3 + 822.1 kJ

To get the thermal effect of the reaction we are interested in, we multiply the first equation by 3, and rewrite the second in reverse order:

1) 3 C + 3/2 O 2 = 3 CO + 331.5 kJ

2) Fe 2 O 3 + 822.1 kJ = 2 Fe + 3/2 O 2

Now let’s add both equations term by term: 3 C + 3/2 O 2 + Fe 2 O 3 + 822.1 kJ = 3 CO + 331.5 kJ + 2 Fe + 3/2 O 2

After reducing both sides of the oxygen equation (3/2 O 2) and transferring 822.1 kJ to the right side, we obtain: 3 C + Fe 2 O 3 = 3 CO + 2 Fe – 490.6 kJ

kinetics of chemical reactions- a branch of physical chemistry that studies the patterns of the occurrence of chemical reactions over time, the dependence of these patterns on external conditions, as well as the mechanisms of chemical transformations. Chemical kinetics is the science of the rates and patterns of the occurrence of chemical processes over time.

Chemical kinetics studies the mechanism of the process, i.e. those intermediate stages consisting of elementary acts through which the system passes from the initial state to the final state.

Chemical kinetics studies the rates of these steps and the factors that influence their rates.

The equation of a chemical reaction shows the initial state of the system (starting substances) and its final state (reaction products), but does not reflect the mechanism of the process.

Ticket number 10.
1.Characteristics of a chemical bond - energy, length, multiplicity, polarity.
The reason for the formation of a chemical bond.

A chemical bond is a set of interactions between atoms, leading to the formation of stable systems (molecules, complexes, crystals.). It occurs if, as a result of the overlapping of clouds of atoms, the total energy of the system decreases. A measure of strength is the bond energy, which is determined by the work required to break a given bond.
Types of chemicals bonds: covalent (polar, nonpolar, exchange and donor-acceptor), ionic, hydrogen and metallic.
Bond length is the distance between the centers of atoms in a molecule. The energy and length of bonds depend on the nature of the El distribution. density between atoms. The distribution of its density is influenced by the spatial orientation of the chemical. communications. If 2-atomic molecules are always linear, then the shapes of polyatomic molecules may be are different.
The angle between imaginary lines that can be drawn through the centers of bonded atoms is called valence. The distribution of e density also depends on the size of the at. and their eo. In homoatomic El. the density is evenly distributed. In heteroatomic ones, it is shifted in the direction that helps to reduce the energy of the system.
Binding energy is the energy that is released when a molecule is formed from single atoms. The binding energy differs from ΔHobr. The heat of formation is the energy that is released or absorbed during the formation of molecules from simple substances. So:

N2 + O2 → 2NO + 677.8 kJ/mol – ∆Hobr.

N + O → NO - 89.96 kJ/mol – E St.

The multiplicity of a bond is determined by the number of electron pairs involved in the bond between atoms. The chemical bond is caused by the overlap of electron clouds. If this overlap occurs along the line connecting the atomic nuclei, then such a bond is called a σ 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.
If a bond is formed by more than one pair of electrons, then it is called multiple.
A multiple bond is formed when there are too few electrons and bonding atoms for each bond-forming valence orbital of the central atom to overlap with any orbital of the surrounding atom.
Since p-orbitals are strictly oriented in space, they can overlap only if the p-orbitals of each atom perpendicular to the internuclear axis are parallel to each other. This means that in molecules with a multiple bond there is no rotation around the bond.

If a diatomic molecule consists of atoms of one element, such as molecules H2, N2, Cl2, etc., then each electron cloud formed by a common pair of electrons and carrying out a covalent bond is distributed in space symmetrically relative to the nuclei of both atoms. In such a case, the covalent bond is called nonpolar or homeopolar. If a diatomic molecule consists of atoms of different elements, then the overall electron cloud is shifted towards one of the atoms, so that an asymmetry in the charge distribution occurs. In such cases, the covalent bond is called polar or heteropolar.

To assess the ability of an atom of a given element to attract a common electron pair, the value of relative electronegativity is used. The greater the electronegativity of an atom, the more strongly it attracts the shared electron pair. In other words, when a covalent bond is formed between two atoms of different elements, the common electron cloud shifts to a more electronegative atom, and to a greater extent, the more the electronegativity of the interacting atoms differs. Electronegativity values of atoms of some elements relative to the electronegativity of fluorine, which is assumed to be 4.
Electronegativity naturally changes depending on the position of the element in the periodic table. At the beginning of each period there are elements with the lowest electronegativity - typical metals, at the end of the period (before the noble gases) - elements with the highest electronegativity, i.e. typical non-metals.

For elements of the same subgroup, electronegativity tends to decrease with increasing nuclear charge. Thus, the more typical a metal an element is, the lower its electronegativity; The more typical a nonmetal an element is, the higher its electronegativity.

The reason for the formation of a chemical bond. Atoms of most chemical elements do not exist in individual form, since they interact with each other, forming complex particles (molecules, ions and radicals). Electrostatic forces act between atoms, i.e. the force of interaction of electric charges, the carriers of which are electrons and atomic nuclei. Valence electrons play a major role in the formation of chemical bonds between atoms.
The reasons for the formation of a chemical bond between atoms can be sought in the electrostatic nature of the atom itself. Due to the presence of spatially separated regions of electrical charge in atoms, electrostatic interactions can occur between different atoms that can hold these atoms together.
When a chemical bond is formed, a redistribution in space of electron densities that initially belonged to different atoms occurs. Since the electrons of the outer level are least tightly bound to the nucleus, these electrons play the main role in the formation of a chemical bond. The number of chemical bonds formed by a given atom in a compound is called valency. For this reason, the outer level electrons are called valence electrons.

2.Characteristics of a chemical bond - energy, length, multiplicity, polarity.

Binding energy is the energy that is released when a molecule is formed from single atoms. The binding energy differs from ΔHobr. The heat of formation is the energy that is released or absorbed during the formation of molecules from simple substances. (Bond energies in molecules consisting of identical atoms decrease in groups from top to bottom)

For diatomic molecules, the bond energy is equal to the dissociation energy, taken with the opposite sign: for example, in the F2 molecule, the bond energy between F-F atoms is equal to - 150.6 kJ/mol. For polyatomic molecules with one type of bond, for example, for ABn molecules, the average bond energy is equal to 1/n part of the total energy of formation of a compound from atoms. Thus, the energy of formation of CH4 = -1661.1 kJ/mol.

If more than two different atoms are combined in a molecule, then the average binding energy does not coincide with the dissociation energy of the molecule. If a molecule contains different types of bonds, then each of them can be approximately assigned a certain value of E. This allows us to estimate the energy of formation of a molecule from atoms. For example, the energy of formation of a pentane molecule from carbon and hydrogen atoms can be calculated using the equation:

E = 4EC-C + 12EC-H.

Bond length is the distance between the nuclei of interacting atoms. The bond length can be roughly estimated based on atomic or ionic radii, or from the results of determining the size of molecules using Avogadro's number. So, the volume per one molecule of water: , o

The higher the order of the bond between atoms, the shorter it is.

Multiplicity: The multiplicity of a bond is determined by the number of electron pairs involved in the bond between atoms. The chemical bond is caused by the overlap of electron clouds. If this overlap occurs along the line connecting the atomic nuclei, then such a bond is called a σ 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.

If a bond is formed by more than one pair of electrons, then it is called multiple.

A multiple bond is formed when there are too few electrons and bonding atoms for each bond-forming valence orbital of the central atom to overlap with any orbital of the surrounding atom.

Since p-orbitals are strictly oriented in space, they can overlap only if the p-orbitals of each atom perpendicular to the internuclear axis are parallel to each other. This means that in molecules with a multiple bond there is no rotation around the bond.

Polarity: If a diatomic molecule consists of atoms of one element, such as molecules H2, N2, Cl2, etc., then each electron cloud formed by a common pair of electrons and carrying out a covalent bond is distributed in space symmetrically relative to the nuclei of both atoms. In such a case, the covalent bond is called nonpolar or homeopolar. If a diatomic molecule consists of atoms of different elements, then the overall electron cloud is shifted towards one of the atoms, so that an asymmetry in the charge distribution occurs. In such cases, the covalent bond is called polar or heteropolar.

The displacement of the total electron cloud during the formation of a polar covalent bond leads to the fact that the average density of negative electric charge is higher near a more electronegative atom and lower near a less electronegative one. As a result, the first atom acquires an excess negative charge, and the second - an excess positive charge; These charges are usually called the effective charges of the atoms in the molecule.

3. The reason for the formation of a chemical bond is the desire of metal and non-metal atoms, through interaction with other atoms, to achieve a more stable electronic structure, similar to the structure of inert gases. There are three main types of bonds: polar covalent, nonpolar covalent and ionic.

A covalent bond is called nonpolar if the shared electron pair is shared equally by both atoms. A covalent nonpolar bond occurs between atoms whose electronegativity is the same (between atoms of the same non-metal), i.e. in simple substances. For example, in molecules of oxygen, nitrogen, chlorine, bromine, the covalent bond is nonpolar.
A covalent bond is called polar if the shared electron pair is shifted to one of the elements. A polar covalent bond occurs between atoms whose electronegativity differs, but not greatly, i.e. in complex substances between non-metal atoms. For example, in molecules of water, hydrogen chloride, ammonia, and sulfuric acid, the bond is polar covalent.
An ionic bond is a bond between ions that occurs due to the attraction of oppositely charged ions. Ionic bonds occur between atoms of typical metals (the main subgroup of the first and second groups) and atoms of typical non-metals (the main subgroup of the seventh group and oxygen).
4. Chemical equilibrium. Equilibrium constant. Calculation of equilibrium concentrations.
Chemical equilibrium is a state of a chemical system in which one or more chemical reactions occur reversibly, and the rates in each forward-reverse reaction pair are equal. For a system in chemical equilibrium, the concentrations of reagents, temperature and other parameters of the system do not change over time.

A2 + B2 ⇄ 2AB

In a state of equilibrium, the rates of forward and reverse reactions become equal.

Equilibrium constant is a value that determines for a given chemical reaction the ratio between starting substances and products in a state of chemical equilibrium. Knowing the equilibrium constant of the reaction, it is possible to calculate the equilibrium composition of the reacting mixture, the maximum yield of products, and determine the direction of the reaction.

Ways to express the equilibrium constant:
For a reaction in a mixture of ideal gases, the equilibrium constant can be expressed in terms of the equilibrium partial pressures of the components pi according to the formula:

where νi is the stoichiometric coefficient (taken negative for starting substances, positive for products). Kp does not depend on the total pressure, on the initial amounts of substances, or on which reaction participants were taken as the initial ones, but depends on temperature.

For example, for the oxidation reaction of carbon monoxide:
2CO + O2 = 2CO2

The equilibrium constant can be calculated using the equation:

If the reaction occurs in an ideal solution and the concentration of the components is expressed in terms of molarity ci, the equilibrium constant takes the form:

For reactions in a mixture of real gases or in a real solution, instead of partial pressure and concentration, fugacity fi and activity ai are used, respectively:

In some cases (depending on the method of expression), the equilibrium constant can be a function of not only temperature, but also pressure. Thus, for a reaction in a mixture of ideal gases, the partial pressure of a component can be expressed according to Dalton’s law in terms of the total pressure and mole fraction of the component (), then it is easy to show that:

where Δn is the change in the number of moles of substances during the reaction. It can be seen that Kx depends on pressure. If the number of moles of reaction products is equal to the number of moles of starting substances (Δn = 0), then Kp = Kx.

Hybridization of atomic orbitals. The concept of the molecular orbital method. Energy diagrams of the formation of molecular orbitals for binary homonuclear molecules. When a chemical bond is formed, the properties of the interacting atoms and, above all, the energy and occupancy of their outer orbitals change.

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PAGE 13

Lebedev Yu.A. Lecture 0 2

Lecture No. 0 2

Chemical bond. Characteristics of a chemical bond: energy, length, bond angle. Types of chemical bonds. Communication polarity. Quantum mechanical ideas about the nature of covalent bonds. The concept of the valence bond method. Hybridization of atomic orbitals.- (c igma) and (pi)-connections. Geometric configuration of molecules. Electric dipole moment of a molecule. The concept of the molecular orbital method. Energy diagrams of the formation of molecular orbitals for binary homonuclear molecules. Sigma () and Pi( )-molecular orbitals. Dia- and paramagnetic molecules.

REMINDER

Schrödinger equation. - wave function.

E= f (n, l, m, s).

Chemical bond. Characteristics of a chemical bond: energy, length, bond angle.

We examined the structure of the electronic levels of isolated atoms. These are very rare objects in practice. The only exception is the inert gas argon with electronic formula 1 s 2 2 s 2 2 p 6 3 s 2 3 p 6 . And although it is “only” 0.93% vol in the atmosphere, each of you literally “swallows” about three hundred quintillion argon atoms in one breath.

All other substances and materials we deal with containchemically relatedatoms. The interaction of free atoms with each other leads to the formation of molecules, ions and crystals. These are “classical” chemical objects. However, recently objects such as nanostructures, surface compounds, berthollides and a number of other practically important “non-classical” chemical objects have acquired an important role.

A chemical bond is caused by the interaction of electrons in the outer electron shells of atoms.Those orbitals that take part in the formation of a chemical bond are calledvalence orbitals, and the electrons located on them are valence electrons.

When a chemical bond is formed, the properties of the interacting atoms and, above all, the energy and occupancy of their outer orbitals change.

When a chemical bond is formed, the total energy of electrons in valence orbitals is less than their energy in free atoms. This difference in energy is called chemical bond energy.

The typical value of chemical bond energy is hundreds of kJ/mol.

An important quantitative characteristic of a chemical bond is its length.Bond length is the distance between the nuclei of chemically bonded atoms in the stable state of the molecule.

The typical length of a chemical bond is tenths of a nanometer. 1

If two or more other atoms take part in the formation of a molecule when interacting with a given atom, then the question arises about its geometric structure or chemical structure. The foundations of the theory of the chemical structure of molecules were laid by A.M. Butlerov 2

One of the most important quantitative characteristics of the structure of complex molecules is bond angle - the angle formed by two directions of chemical bonds emanating from one atom.

Types of chemical bonds. Communication polarity.

According to the nature of the interaction of valence electrons and the type of orbitals formed during the interaction,chemical bonds are divided into the following main types:covalent (polar and non-polar), ionic, donor-acceptor, hydrogen and intermolecular (also called van der Waals).

Back in 1916, the American chemist G.N. Lewis 3 expressed the idea that a chemical bond is formed by an electron pair, which is graphically represented by a valence line:

F + F = F 2 (F-F).

If the electronegativity of the atoms is equal, then such a bond is called non-polar. If different polar.

When a polar covalent bond is formed, atoms acquire an additional charge: negative for an atom with higher electronegativity and positive for an atom with lower electronegativity:

H+Cl = HCl (

)

In the case when the difference in electronegativity of interacting atoms is large, the bond considered ionic:

Na + Cl = NaCl (Na + Cl - ).

If the electron pair forming a bond belonged to one of the atoms before the interaction, then such a bond is called a donor-acceptor bond. The atom that provided an electron pair is called a donor, and the atom that accepted it into a free orbital is called an acceptor.

The emergence of donor-acceptor bonds is especially characteristic d - metals that have free or partially filled d -orbitals to form complex compounds.

We will talk about other types of communication later.

Quantum mechanical ideas about the nature of covalent bonds.

From a modern point of view, a covalent bond arises from the quantum mechanical interaction of all electrons of all interacting atoms. But, as we already said in lecture No. 1, there is no exact solution to the Schrödinger equation, which describes the orbitals of many electrons in molecules. The task of quantum mechanical description of a chemical bond is simplified by the fact that during its formation the role of electrons located on the inner and outer electron shells is significantly different.

Therefore, it was possible to create various approximate methods for describing chemical bonds.

Quantum chemistry has a rich arsenal of application programs that allow calculations with high accuracy for a wide class of molecules and ions. 4

However, there is no universal and sufficiently accurate quantum chemical algorithm yet.

To qualitatively understand the structure of chemical compounds, two methods are usedvalence bond method (VBC) And molecular orbital (MO) method.

The concept of the valence bond method. Geometric configuration of molecules. Electric dipole moment of a molecule.

The main postulates of the valence bond method are:

1. A single covalent chemical bond is carried out by two valence electrons, which occupy two orbitals, one from each of the interacting atoms. In this case, the spins of the electrons forming the valence pair must be opposite (the bond is formed by electrons with antiparallel spins).

2. The original atomic orbitals (AO) retain their outline in the composition of the molecule.

3. A bond is formed due to the overlap of orbitals, leading to an increase in the electron density between the nuclei of interacting atoms in the direction that provides maximum overlap.

Let us consider the formation of a chemical bond along the MBC in a water vapor molecule H2O.

The molecule consists of one oxygen atom O and two hydrogen atoms H . Electronic formula of oxygen atom 1 s 2 2 s 2 2 p 4 . The outer energy level contains 6 electrons. Sublevel 2 s is filled. At sublevel 2 p on one of the p -orbitals (assuming p y ,) there is an electron pair, and on the other two ( p x and p z ) one unpaired electron. It is they who will participate in the formation of a chemical bond.

Electronic formula of hydrogen atom 1 s 1 . Hydrogen has one s -an electron whose orbital outline is a sphere, and it will participate in overlap with p -oxygen orbital, forming a chemical bond. Total of these sp -there will be two overlaps in a water molecule. And the structure of the molecule will look like this:

As can be seen from the figure, in a water molecule there are two covalent chemical bonds directed along the axes Z and X . Therefore, the bond angle in this model is 90 O . The experiment shows that this angle is 104.5 o.

Not a bad match for a simple high-quality model without any calculations!

The electronegativity of oxygen according to Mulliken is 3.5, and hydrogen 2.1. Consequently, each of the bonds will be polar, and the charge- will be on oxygen, and+ - on hydrogen, i.e. Three centers of electric charge are formed. Two electric dipoles are formed in the molecule.

Dipole is two equal charges located at a finite distance l from each other. A dipole is characterized by a dipole moment

A dipole is a vector directed from the negative pole to the positive pole. In a water molecule, two bond dipole moments are formed, which when added together give the total dipole moment of the molecule. The diagram of the dipole moments of a water molecule according to the MBC model has the form:

It is important to emphasize that the dipole moments of bonds add up vectorially and the total dipole moment depends on the geometry of the molecule. As we see, in this case, due to the fact that the bonds are directed at right angles to each other, the molecule as a whole turns out to be polar. And experiment confirms this: the dipole moment of a water molecule is 1.84 Debye. (1 Debye is equal to 0.33*10-29 Kl*m)

The geometric structure of bonds in molecules can be very diverse. Bonds can be located both on a plane and in space, forming molecules in the form of volumetric bodies of various configurations (trigonal, tetragonal, hexagonal pyramids, bipyramids, rings composed of pyramids, etc.)

You can read more about the relationship between the structure of chemical bonds and the geometry of molecules in the textbook on page 119 128).

- (c igma) and (pi)-connections.

Let's return to the overlap of orbitals during bond formation. In our examplemaximum overlap area s and p orbitals lies on the line connecting the centers of the atoms. This type of overlap is called- connections.

Let's consider another case - an oxygen molecule O2 . As we have already seen, the oxygen atom has two p -orbitals containing electrons capable of forming a chemical bond. Well known structural formula of oxygen O=O . There is a double bond in an oxygen molecule. One of them is the one just discussed-connection. And the second? It turns out that the second bond is formed due to another type of orbital overlap called- communication.

Concept of And connections was put forward by F. Hund.

During education -bonds of the orbitals overlap in such a way that two areas of overlap are formed, and they are located symmetrically relative to the plane on which the nuclei of interacting atoms lie.

Geometrically it looks like this:

Please note that- the bond is formed by smaller parts p -orbitals in which the density of the “electron cloud” is greater, and therefore this bond is stronger- connections. Indeed, experiment shows that in carbon compounds ethane C 2 H 6 (CH 3 - CH 3 one -bond), ethylene C 2 H 4 (CH 2 = CH 2 - one - communication and one -bond) and acetylene C 2 H 2 (C HC H - one - communication and two -bond) their breaking energy is respectively 247, 419 and 515 kJ/mol.

Now we can add to the list of MBC postulates:

4. If multiple (double and triple) bonds are formed in a molecule, then one of them will be-communication, and others -- connections).

Note that in connections d - and f -metals, it is possible to form another type of bond --bonds, when the overlap occurs in four spatial regions and the plane of symmetry is perpendicular to the line connecting the atomic nuclei.

Hybridization of atomic orbitals.

When chemical bonds form, an important phenomenon can occur calledorbital hybridization.

Consider the beryllium atom Be . Its electronic formula is 1 s 2 2 s 2 . Judging by the fact that all beryllium electrons are paired, such an atom should behave chemically like noble gases and not enter into chemical interactions.

However, let's look carefully at the electron diffraction diagram of the beryllium atom:

From the diagram it is clear that the beryllium atom has, in addition to the filled 2 s -orbitals three more free 2 p -orbitals! True, the energy of these orbitals is greater than the energy of 2 s -orbitals by magnitudeE . But this energy is small and less than that which is released during the formation of a chemical bond. Therefore, the atom tends to rearrange its orbitals during the interaction to achieve an energetically favorable final state. For such a restructuring, the kinetic energy of particles interacting with a given atom is used. We will talk in more detail about this energy source when discussing issues of chemical kinetics. 5

This rearrangement is called orbital hybridization, since during this process a new one arises from “two types” of orbitals.

In wavefunction language, this is described by an equation that relates the hybrid wavefunction of the resulting orbitals to the original wavefunctions.

The number of hybrid orbitals formed is equal to the number of orbitals that took part in the hybridization process.

Graphically this process can be represented by the following diagram:

Note that the energy required for hybridization E hybrid less than the energy difference between the hybridizing orbitals E.

The designation of hybrid orbitals retains the designation of the original orbitals. So, in this case (atom Be ), hybridize one s and one p -orbital, and both hybrid orbitals are denoted as sp -orbitals. The need for hybridization of only two orbitals is due to the fact that the beryllium atom has only two electrons at the outer energy level.

In other cases, when several identical orbitals are involved in hybridization, their numbers are indicated by an exponent. For example, when hybridizing one s and two p -there are three orbitals sp 2 -orbitals, and during hybridization of one s and three p -orbitals four sp 3 orbitals.

In the case under consideration, in accordance with Hund's rule, the beryllium atom receives two unpaired electrons and the ability to form two covalent chemical bonds.

Hybrid orbitals formed s, p and even d -orbitals differ little in shape and look like this (“asymmetrical dumbbell”):

Note that the number of hybrid orbitals is equal to the number of orbitals involved in their creationregardless of the number and type of hybridizing orbitals.

The location of hybrid orbitals in space is determined by their number.

Specifically, the beryllium atom has two hybrid sp -orbitals are located along one straight line (at an angle of 180 o ), which corresponds to the desire of the similarly charged electrons occupying them to move away from each other as much as possible:

More details You can read about the method of valence bonds and hybridization here:

http://center.fio.ru/method/resources/Alikberovalyu/2004/stroenie/gl_10.html#104

Molecules often have orbitals occupied by an electron pair (“lone electron pair”). Such orbitals do not take part in the formation of chemical bonds, but affect the geometric structure of the molecule.

A modification of the MBC that takes into account the influence of such orbitals is called the theory of repulsion of electron pairs of valence orbitals (EPVO) and you can get acquainted with it in the textbook on p. 124 128.

The concept of the molecular orbital method.

We examined the phenomenon of hybridization of joint stock companies within the framework of the MBC. It turned out that the idea of hybridization is also fruitful for deeper modeling of chemical bonds. It is the basis of the second method of their description, which is discussed in our course methodmolecular orbitals(MO).

The main postulate of this method is the statement that AOs of atoms interacting with each other lose their individuality and form generalized MOs, i.e. that electrons in molecules do not “belong” to any particular atom, but move quantum mechanically throughout the molecular structure.

There are several varieties of the ML method that take into account b O greater or lesser number of factors and, accordingly, more or less mathematically complex. The simplest approximation is one that takes into account only the linear effects of electron interactions. This approximation is called the MO LCAO (linear combination of atomic orbitals) method.

In the language of quantum mechanics, this statement for the simplest case of interaction of two orbitals is written as follows:

Where - MO wave function,
- wave function of the AO of the first atom,
- wave function of the second atom, a and b numerical coefficients characterizing the contribution of a given JSC to the overall structure of the MO.

Since a linear polynomial is written on the right side, this modification of the MO method is called LCAO.

From the equation it is clear thatwhen two AOs interact, two MOs are obtained. One of them is called binding MO, and the other loosening MO.

Why they received such a name is clear from the figure, which shows the energy diagram of orbitals in a molecule:

As can be seen from the figure, the binding MO has an energy lower than the energy of the original AO, and the antibonding MO has a higher energy. (Respectively,). Naturally, in accordance with the principle of minimum energy, the electrons in the molecule will first occupy the bonding orbital when forming a bond.

In general, when interacting N AO becomes N MO.

Sigma ( ) and pi( )-molecular orbitals.

As a result of quantitative calculations using the MO LCAO method, it turned out that the concepts ofAnd types of orbital symmetry are preserved in the LCAO MO method.

This is what outlines look like-binding (designated asor) and -antibonding (denoted as or) orbitals in the LCAO MO method:

And this is what the outlines look like- connecting ( ) And - loosening ( * ) orbitals using the LCAO MO method:

Energy diagrams of the formation of molecular orbitals for binary homonuclear molecules.

Calculating the energy of molecular orbitals for complex molecules that include the nuclei of various elements (heteronuclear molecules) is a complex computational task even for modern computers. Therefore, each calculation of individual molecules is a separate creative work.

Nevertheless, it turned out that the energy diagram for binary homonuclear molecules of elements of the second period of D.I. Mendeleev’s Periodic Table is universal and has the form:

Sometimes the literature provides different diagrams for elements B ,C,N and subsequent O, F, Ne , however, studies of the magnetic properties of the molecule B 2 at ultra-low temperatures do not clearly confirm the need to complicate the type of energy diagrams for B ,C,N.

Dia- and paramagnetic molecules. Multiplicity of connections according to MO LCAO.

One of the serious advantages of the MO LCAO method compared to the BC method is a more correct description of the magnetic properties of molecules and, in particular, an explanation of the paramagnetism of molecular oxygen. 6

Let us recall the structure of the oxygen molecule according to the MBC, which we examined earlier. According to this structure, all valence electrons andAnd -bonds in a molecule O2 form electron pairs and the total spin of the molecule is zero.

The structure of the orbitals of this molecule using the LCAO MO method, obtained by filling MOs with electrons in accordance with the above energy diagram, has the form:

As can be seen from this diagram, the oxygen molecule contains two unpaired electrons on the antibonding
And
orbitals. Their magnetic moments add up and give the total magnetic moment of the molecule. The experiment shows that the magnetic moment of the oxygen molecule is 2.8(Intrinsic magnetic moment of an electron 1). Considering that the total magnetic moment, in addition to its own electronic moment, also includes an orbital one, the quantitative coincidence very convincingly testifies in favor of the validity of the MO method.

In the presence of a magnetic moment, a substance becomesparamagneticit is “attracted by a magnet.” 7 In the absence of a magnetic moment, the substance diamagnetic it is “pushed out” by the magnetic field. 8

In addition to magnetic properties, analysis of the energy diagrams of MO LCAO makes it possible to determinemultiplicity (or order) of a chemical bond (CS or PS).

KS= ½(N connection N cut)

where N connections total number of electrons in bonding orbitals; N bit total number of electrons in antibonding orbitals).

We looked at various cases of manifestation and description of covalent chemical bonds. This is the main type of chemical bond, since the reason for its occurrence is the presence of valence electrons in the vast majority of chemical elements.

However, in some cases of interaction of atoms, special conditions arise that give rise to special types of bonds, which we will consider in the next lecture.

equal to the work that must be expended to divide a molecule into two parts (atoms, groups of atoms) and remove them from each other at an infinite distance. For example, if E. x. With. H 3 C-H in a methane molecule, then such particles are the methyl group CH 3 and the hydrogen atom H, if E. chemistry is considered. With. H-H in a hydrogen molecule, such particles are hydrogen atoms. E. x. With. - a special case of binding energy (See Bonding energy) , it is usually expressed in kJ/mol(kcal/mol); depending on the particles forming a chemical bond (See Chemical bond), the nature of the interaction between them (Covalent bond, Hydrogen bond and other types of chemical bonds), bond multiplicity (for example, double, triple bonds) E. x. With. has a value from 8-10 to 1000 kJ/mol. For a molecule containing two (or more) identical bonds, E. ch. With. each bond (bond breaking energy) and the average bond energy equal to the average value of the breaking energy of these bonds. Thus, the energy of breaking the HO-H bond in a water molecule, i.e., the thermal effect of the reaction H 2 O = HO + H is 495 kJ/mol, energy of breaking the H-O bond in the hydroxyl group - 435 kJ/mol, average E. x. With. equal to 465 kJ/mol. The difference between the values of the rupture energies and the average E. ch. With. due to the fact that during partial dissociation (See Dissociation) of a molecule (breaking one bond), the electronic configuration and relative arrangement of the atoms remaining in the molecule change, as a result of which their interaction energy changes. The value of E. x. With. depends on the initial energy of the molecule; this fact is sometimes referred to as the dependence of E. x. With. on temperature. Usually E. x. With. are considered for cases when the molecules are in the standard state (See Standard states) or at 0 K. It is these values of E. x. With. are usually given in reference books. E. x. With. - an important characteristic that determines reactivity (See Reactivity) substances and used in thermodynamic and kinetic calculations of chemical reactions (See Chemical reactions). E. x. With. can be indirectly determined from calorimetric measurements (see Thermochemistry) , by calculation (see Quantum chemistry) , and also using mass spectroscopy (See Mass spectroscopy) and spectral analysis (See Spectral analysis).

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