home Botany Bond breaking energy h s. Ionization potential and binding energy in diatomic molecules. Chemical bond energy

Bond breaking energy h s. Ionization potential and binding energy in diatomic molecules. Chemical bond energy

Bond energy is an important concept in chemistry. It determines the amount of energy required to break a covalent bond between two gas atoms. This concept not applicable to ionic bonds. When two atoms combine to form a molecule, one can determine how strong the bond between them is - it is enough to find the energy that must be expended to break this bond. Remember that a single atom does not have a binding energy, this energy characterizes the strength of the bond between two atoms in a molecule. To calculate the binding energy for any chemical reaction, simply determine the total number of broken bonds and subtract the number of bonds formed from it.

Steps

Part 1

Identify broken and formed bonds

Write down an equation to calculate the binding energy. By definition, bond energy is the sum of broken bonds minus the sum of bonds formed: ΔH = ∑H (broken bonds) - ∑H (bonds formed) . ΔH denotes the change in binding energy, which is also called the enthalpy of binding, and ∑H corresponds to the sum of the binding energies for both sides of the chemical reaction equation.

Write down the chemical equation and label all the bonds between the individual elements. If the reaction equation is given in the form of chemical symbols and numbers, it is useful to rewrite it and indicate all the bonds between atoms. Such a visual record will allow you to easily count the bonds that are broken and formed during this reaction.

Learn the rules for counting broken and formed bonds. In most cases, the average values of the binding energy are used in the calculations. The same bond can have slightly different energies, depending on the particular molecule, so average bond energies are usually used. .

Breaks of a single, double and triple chemical bond are considered as one broken bond. Although these bonds have different energies, in each case one bond is considered to be broken.
The same applies to the formation of a single, double or triple bond. Each such case is considered as the formation of one new connection.
In our example, all bonds are single.

Determine which links are broken on the left side of the equation. The left side of a chemical equation contains the reactants, and it represents all the bonds that break as a result of the reaction. This is an endothermic process, that is, it takes a certain amount of energy to break chemical bonds.
- To our example left side reaction equation contains one H-H connection and one Br-Br bond.
Count the number of bonds formed on the right side of the equation. The reaction products are shown on the right. This part of the equation represents all the bonds that form as a result of a chemical reaction. This is an exothermic process and it releases energy (usually in the form of heat).
- In our example, the right side of the equation contains two H-Br bonds.
Part 2
Calculate the binding energy
1. Find the required binding energies. There are many tables that list the binding energies for a wide variety of compounds. Such tables can be found on the Internet or a reference book on chemistry. It should be remembered that the binding energies are always given for molecules in the gaseous state.
2. Multiply the bond energies by the number of bonds broken. In a number of reactions, one bond can be broken several times. For example, if a molecule consists of 4 hydrogen atoms, then the binding energy of hydrogen should be taken into account 4 times, that is, multiplied by 4.
  - In our example, each molecule has one bond, so the bond energies are simply multiplied by 1.
  - H-H = 436 x 1 = 436 kJ/mol
  - Br-Br \u003d 193 x 1 \u003d 193 kJ / mol
3. Add up all the energies of broken bonds. After you multiply the binding energies by the corresponding number of bonds on the left side of the equation, you need to find the total.
  - Let's find the total energy of broken bonds for our example: H-H + Br-Br = 436 + 193 = 629 kJ/mol.

Tutorial

Astrakhan

Chemical bond: Tutorial/ Ryabukhin Yu. I. - Astrakhan: Astrakhan. state tech. un-t, 2013. - 40 p.

Designed for students of engineering and non-chemical specialties.

Corresponds to the state educational standards of higher professional education

Illustr.: 15 pic., Table: 1, bibliography: 6 titles, app.

Printed by decision of the department "General, inorganic and analytical chemistry”(Minutes No. __ dated _________ 2013)

Reviewer: Cand. chem. Sciences, Associate Professor Lebedeva A.P.

INTRODUCTION

In nature, chemical elements in the form of free atoms (with the exception of noble gases - elements of group VIIIA) are practically not found. Usually, the atoms of a chemical element interact either with each other or with atoms of other elements, forming chemical bonds with the appearance of simple or complex substances, respectively. At the same time, molecules of different substances interact with each other.

The doctrine of the chemical bond is the basis of all theoretical chemistry.

Chemical bond 1 - this is a set of forces that bind atoms to each other into more stable structures - molecules or crystals.

The formation of molecules and crystals is mainly due to the Coulomb attraction between electrons and atomic nuclei.

The nature of the chemical bond was understood only after the discovery of the laws of quantum (wave) mechanics that govern the microworld. Modern theory answers the questions why a chemical bond occurs and what is the nature of its forces.

The formation of chemical bonds is a spontaneous process; otherwise, neither simple nor complex substances would exist. From a thermodynamic point of view, the reason for the formation of a chemical bond is a decrease in the energy of the system.

The formation of a chemical bond is accompanied by the release of energy, and its breaking requires the expenditure of energy.

The characteristics of a chemical bond are its energy and length.

Chemical bond energy is the energy released in the process of its formation and characterizing its strength; the binding energy is expressed in kJ per mole of the formed substance (E St. , kJ/mol) 2 .

The greater the energy of a chemical bond, the stronger the bond. The energy of the chemical bond of a diatomic molecule is estimated by comparing it with the state preceding its formation. For polyatomic molecules with the same type of bond, the average chemical bond energy is calculated (for example, for H 2 O or CH 4).

Average chemical bond energy is determined by dividing the formation energy of a molecule by the number of its bonds.

Chemical bond length called the distance between the nuclei of atoms in a molecule.

The bond length is determined by the size of the bonding atoms and the degree of overlap of their electron shells.

For example, for hydrogen fluoride and hydrogen iodide:

l HF< l HI

Depending on the type of connected particles (atoms or molecules), there are intramolecular bonds through which molecules are formed, and intermolecular bonds, leading to the formation of associates from molecules or to the binding of atoms of individual functional groups in a molecule. These types of bonds differ sharply in energy: for intramolecular bonds, the energy is 100–1000 kJ/mol 1, and for intermolecular bonds, it usually does not exceed 40 kJ/mol.

Consider Education intramolecular chemical bond on the example of the interaction of hydrogen atoms.

When two hydrogen atoms approach each other, a strong exchange interaction occurs between their electrons with antiparallel spins, leading to the appearance of a common electron pair. This increases the electron density in the internuclear space, which contributes to the attraction of nuclei, interacting atoms. As a result, the energy of the system decreases and the system becomes more stable - chemical bond(Fig. 1).

Rice. 1. Energy diagram of the formation of a chemical bond between hydrogen atoms

The system has a minimum of energy at a certain distance between the nuclei of atoms; with further approach of the atoms, the energy increases due to the increase in the repulsive forces between the nuclei.

Depending on how the common electron pair interacts with the nuclei of the atoms being joined, there are three main types of chemical bond: oval, ionic, metallic, and hydrogen bonds.

Comparison of data on the number of electrons in the outer shell with the number of chemical bonds that a given atom can form showed that the principles of chemical bond formation, revealed in the study of the hydrogen molecule, are also valid for other atoms. This is because the connection is electrical nature and is formed by two electrons (one from each atom). Therefore, a correlation should be expected between the first ionization energy (PEI) of atoms (having an electrostatic origin) and their binding energy in diatomic molecules.

Experimental data on determining the binding energy for a number of diatomic molecules (in the gas phase) formed from atoms of the 2nd and 3rd periods are shown in Table 4.2 and in Fig. 4.2.1.

Table 4.2

Molecule A 2	Bond energy (kJ/mol)	Molecule	Bond energy (kJ/mol)

Rice. 4.2-1 Binding energy in molecules from elements of the second and third periods depending on the PEI of the element

These data (see Table 4.2, Fig. 4.2-1) show that the binding energy between atoms is practically independent of the SEI of the bonded atoms.

Is it possible that in diatomic molecules (where there is more than one electron) the bond is formed according to a different mechanism and there are additional forces previously ignored by us?

Before proceeding to the identification of these forces, let us try to explain this independence based on existing interactions.
Let's start by examining additional factors that explain the lack of expected correlation and independence experimental data on the measurement of PEI from the binding energy in diatomic molecules.
We divide the table (4.2) into four groups:

Group A includes molecules consisting of identical atoms with a binding energy below 40 kJ/mol. In the gas phase, these molecules break down into atoms.

Group B includes diatomic molecules consisting of identical atoms, the binding energy in which ranges from 400 kJ/mol to 1000 kJ/mol. Indeed, the binding energy in these molecules differs significantly upwards compared to the binding energy in the hydrogen molecule, which is 429 kJ/mol.

GroupFROM includes diatomic molecules consisting of different atoms, the binding energy of which varies from 340 kJ/mol to 550 kJ/mol.

GroupD includes diatomic molecules with identical atoms, the binding energy of which is 50-350 kJ/mol.

TABLE 4.4
COMMUNICATION ENERGYIN MOLECULES

Binding energy (kJ/mol) in a series of diatomic molecules

group A		group B
molecule	binding energy	molecule	binding energy
Be 2	30	C2	602
Ne 2	4	N 2	941
	7.6	O2	493
Ar 2	7	P2	477
		S2	421
group C		group D
molecule	energy	molecule	energy
LiF	572	B2	274
NaF	447	Br2	190
LiCl	480	Cl2	239
NaCl	439	F2	139
		Li 2	110
		Na 2	72

Before we begin the explanation, let's clarify the issues we need to cover.
The first question:
Why is the binding energy between multielectron atoms much less or much more (Table 4.2) than in a hydrogen molecule (H2)?

To explain the significant deviation of the binding energy in polyatomic molecules from the binding energy in the hydrogen molecule, it is necessary to deepen our understanding of the reason why the number of electrons in the outer shell is limited.
The attachment of an electron to an atom occurs when there is a gain in energy, or, in other words, if absolute the value of the potential energy of the system atom + electron increases as a result of the bond between the electron and the atom. The data on the affinity of an atom for an electron, indicated in Table 4.3, give us the numerical value of the gain in energy when an electron is attached to an atom.

Table 4.3

First ionization energy (PEI) and electron affinity for elements of the 1st, 2nd and 3rd periods in the table of elements (kJ/mol)




Affinity




Affinity

When an electron is attached to an atom, the total energy of attraction of electrons to the nucleus increases due to an increase in the number of electrons attracted to the nucleus. On the other hand, the energy of interelectron repulsion increases due to an increase in the number of electrons. That is, the attachment of an electron to an atom occurs if, as a result of this connection, the gain in energy of attraction is greater than the loss of energy due to an increase in the energy of repulsion.

Calculating the change in energy when an electron is added to an atom hydrogen gives an energy gain of 3.4 eV. That is, the hydrogen atom must have a positive electron affinity. This is what is observed in the experiment.

A similar calculation of the change in potential energy when an electron is attached to an atom helium shows that the addition of an electron leads not to an increase in potential energy, but to its decrease. Indeed, the affinity of the helium atom, according to experiment, is less than zero.

Therefore, the possibility of attaching or not attaching an electron to an atom is determined by differences in the change in the absolute values of the potential energy of attraction of all electrons to the nucleus and mutual interelectronic repulsion. If this difference is greater than zero, then the electron will join, and if it is less than zero, then no.

The data on the affinity of atoms for the electron given in Table 4.3 show that for atoms of the 1st, 2nd and 3rd periods, in addition to be,mg,Ne,Ar the increase in the energy of attraction during the attachment of electrons to the nucleus is greater than the increase in the energy of repulsion.
In the case of atoms be,mg,Ne,Ar, the increase in the energy of attraction during the attachment of electrons to the nucleus is lower than the increase in the energy of interelectron repulsion. An independent confirmation of this conclusion is the information on PEI for atoms of the 2nd and 3rd periods given in Table 4.2 (group A).

When a chemical bond is formed, the number of electrons in the outer electron shells of atoms increases by one electron, and according to the calculation of the hydrogen molecule model H 2, the effective charges of the bound atoms change. The effective charges of the bound nuclei change due to the attraction of charged nuclei, and due to an increase in the number of electrons in the outer shells of the bound atoms.

In a hydrogen molecule, the approach of nuclei leads to an increase in the force of attraction of binding electrons to nuclei by 50%, which is equal to an increase in the effective charge of the bound nuclei by 0.5 proton units (see Chapter 3).

In terms of energy gain, bond formation is something like an intermediate process between the attachment of an electron to a neutral atom (measured electron affinity) and the attachment of an electron to an atom whose nuclear charge increases by 1 unit.

According to Table 4.3, when going from lithium (PEI - 519 kJ/mol) to beryllium (PEI - 900 kJ/mol), PEI increases by 400 kJ/mol, and when going from beryllium to boron (PEI - 799 kJ/mol ) the energy gain decreases to 100 kJ/mol.
Recall that the outer electron shell of boron has 3 electrons, and the outer shell of beryllium has 2 electrons. That is, when an electron joins beryllium with a simultaneous increase in the nuclear charge by one proton unit, the bound electron enters the outer shell of beryllium, while the energy gain will be 100 kJ/mol less than when an electron enters the outer shell of lithium (during the transition from lithium to beryllium).

Now, the sharp decrease in the binding energy of atoms with a negative atom-to-electron affinity, indicated in Table 4.3, is quite understandable. However, even though Ne,be,mg,Ar do not attach electrons, they create molecules, because the effective nuclear charge increases. The binding energy in these molecules (group BUT) is much lower than in other molecules.

Now let's answer second question: Why is the binding energy in the group B diatomic molecules shown in Table 4.2. 1.5-2 times greater than the binding energy in the hydrogen molecule?

On the outer shells of carbon atoms (C) nitrogen (N) and oxygen (o) are, respectively, 4, 5 and 6 electrons. The number of bonds that these atoms form is limited by the number of extra electrons that can enter the outer shell when the bond is formed. So carbon atoms (C) nitrogen (N) and oxygen (O) can form 4, 3 and 2 chemical bonds, respectively. Accordingly, between the two atoms shown in Table 4.4, not one, but several chemical bonds can be formed, which implies a much greater gain in energy, compared with the formation of 1 bond in a diatomic molecule, where the bonded atoms have 1 electron in the outer shell

If the atoms are connected by one chemical bond, then such a bond is called a single bond. chemical bond or common chemical bond. When atoms are linked by several chemical bonds (double or triple), such bonds are called multiple bonds. Multiple bonds, for example, in nitrogen molecules (N 2) and oxygen (O2) are described structural formulas: N ≡ N and O=O.

Now consider the group FROM: Why is the binding energy in some of the diatomic molecules consisting of different atoms much greater than in other molecules that are composed of the same atoms?

Let's disassemble the molecule NaCl. Sodium and chlorine atoms are very different in electron affinity. We present bond formation as a two-stage process. At the first stage, the gain in energy is obtained due to the affinity of atoms for electrons. That is, from this point of view, the gain in energy during the formation of a molecule Cl2, must be greater than when forming a molecule NaCl by the difference in their electron affinity.

When calculating the hydrogen molecule (Chapter 3), the binding energy (the energy required to separate molecules into atoms) was the sum of two components:

the difference between the electronic energy of a hydrogen molecule and two hydrogen atoms;

additional energy spent on heating unseparated molecules.

Calculating the first component, we calculate the energy of the molecule, which is equal to the difference between the energy of attraction of the nuclei of hydrogen atoms to the binding pair of electrons and the sum of the repulsive energy of the interelectronic and internuclear forces.

To estimate the energy of attraction of nuclei to binding pairs of electrons, as well as to estimate the energy of interelectron repulsion, we must first find out the value of the effective charge of the bound nuclei.

Ionization potential and binding energy in diatomic molecules

When a chemical bond is formed, there is a redistribution in space of the electron densities that originally belonged to different atoms. Since the electrons of the outer level are the least strongly 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 valence. The electrons involved in the formation of a chemical bond are called valence: for s- and p elements, these are external electrons, for d-elements, external (last) s-electrons and 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 complete. Completed levels are highly durable and 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 carried out, which is achieved either by the addition or release of electrons, as well as the formation of common electron pairs. These methods lead to the formation of two main types of bonds: covalent and ionic. Thus, during the formation of a molecule, each atom tends to acquire a stable outer electron shell: either two-electron (doublet) or eight-electron (octet). This regularity is the basis of the theory of the formation of a chemical bond. The formation of a chemical bond due to the completion of external levels in the atoms forming the bond is accompanied by the release a large number energy, that is, the occurrence 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 they have less energy than the original ones. One of the essential indicators that determine what bond is formed between atoms is electronegativity, that is, the ability of an atom to attract electrons from other atoms. The electronegativity of the 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.

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, a 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) The formation of a covalent bond between atoms, the electronegativity of which differs slightly. A covalent bond formed by atoms with different electronegativity is called a polar bond. With a covalent polar bond, the electron density from a common pair of electrons is shifted to an atom with a higher electronegativity. Molecules H2O, NH3, H2S, CH3Cl can serve as examples. The covalent (polar and nonpolar) bond in our examples was formed due to the unpaired electrons of the bonding atoms. Such a mechanism for the formation of a covalent bond is called an exchange mechanism. Another mechanism for the formation of a covalent bond is donor-acceptor. In this case, the bond arises 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 the ammonium ion: H++:NH 3 -> [H: NH3 | +<=====>NH4+ is an ammonium electron donor ion acceptor. When the ammonium ion is formed, the electron pair of nitrogen becomes common for the N and H atoms, that is, a fourth bond appears, which does not differ from the other three. They are portrayed in the same way:

An ionic bond occurs between atoms whose electronegativity differs sharply. Consider the method of formation using the example of sodium chloride NaCl. The electronic configuration of sodium and chlorine atoms can be represented as: 11 Na ls2 2s2 2p 6 3s1; 17 Cl ls2 2p 6 Zs2 3p5 How are these atoms with incomplete energy levels. Obviously, to complete them, it is easier for a sodium atom to give up one electron than to add seven, and it is easier for a chlorine atom to add one electron than to give up seven. In a chemical interaction, the sodium atom completely gives up one electron, and the chlorine atom accepts it. Schematically, this can be written as: Na. -- l e --> Na+ sodium ion, stable eight-electron 1s2 2s2 2p6 shell due to the second energy level. :Cl + 1e -->.Cl - chlorine ion, stable eight-electron shell. Electrostatic attraction forces arise between the Na+ and Cl- ions, as a result of which a compound is formed.

A chemical bond carried out by 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. There are no molecules in this case. Ionic compounds form elements of the main subgroups of groups I and II and the main subgroups of groups VI and VII, which are sharply different in terms of electronegativity. There are relatively few ionic compounds. For example, inorganic salts: NH4Cl (ammonium ion NH4 + and chloride ion Cl-), as well as salt-like organic compounds: salt alcoholates carboxylic acids, amine salts Non-polar covalent bond and ionic bond are two limiting cases of electron density distribution. A non-polar bond corresponds to a uniform distribution of a binding two electron cloud between identical atoms. On the contrary, with ionic bonding, the binding electron cloud almost entirely belongs to one of the atoms. In most compounds, chemical bonds are intermediate between these types of bonds, that is, they carry out a polar covalent bond.

The metallic bond exists in metals in the solid and liquid state. In accordance with the position in the periodic system, metal atoms have a small number of valence electrons (1-3 electrons) and low ionization energy (electron detachment). Therefore, valence electrons are weakly retained 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 horses, and part of the valence electrons, freely moving in the volume of the crystal lattice, forms an "electron gas" that provides a bond between the metal atoms. The connection that is made with respect to free electrons between metal ions in a crystal lattice is called a metallic bond. A metallic bond arises due to the socialization of valence electrons by atoms. However, there is a significant difference between these types of communication. The electrons that carry out a covalent bond are mainly in the immediate vicinity of the two connected atoms. In the case of a metallic bond, the bonding electrons travel throughout the piece of metal. This determines common features metals: metallic luster, good conductivity of heat and electricity, malleability, ductility, etc. General chemical property 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 due to electrostatic attraction, which is facilitated by the small size of the hydrogen atom, and in part, by donor-acceptor interaction. The hydrogen bond can be intermolecular and intramolecular. The 0-H bonds have a pronounced polar character: The hydrogen bond is much weaker than the ionic or covalent bond, but stronger than the intermolecular interaction. Hydrogen bonds are responsible for some physical properties substances (eg. high temperatures boiling). Hydrogen bonds are especially common in the molecules of proteins, nucleic acids and other biologically important compounds, providing them with a certain spatial structure (organization).

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

Bond length (lsv). Distance between cores in a compound. The longer the bond length, the lower the bond energy.

Method valence bonds.

A) a chemical bond between two atoms arises as a result of the overlap of AO 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 chemical bond energy. An additional contribution to the binding energy comes from 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.

The method of valence bonds. The covalent bond is directed towards the maximum overlap of the AO of the reacting atoms.

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

Upon transition to an excited state, one of the paired electrons passes into a free 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.

Method of molecular orbitals. Electrons in a molecule are distributed over 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 MO is total number AO, of which MO is combined.
2. The energy of some MO turns out to be higher, others - lower than the energy of the original AO. The average MO energy obtained from a set of AOs approximately coincides with the average energy of these AOs.
3. Electrons fill MO, as well as AO, in ascending order of 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 VS method, the bond strength in the MO method is proportional to the degree of overlap of atomic orbitals.

The order and energy of communication. Communication order n=(Nsv-Nr)/2. Nsv - the number e on the binding molecular orbitals, Nr is the number of e on the antibonding molecular orbitals.

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

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

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

Chemical bond - a set of interactions of atoms, leading to the formation of stable systems (molecules, complexes, crystals.). It arises if, as a result of the overlapping of e clouds of atoms, the total energy of the system decreases. The measure of strength is the bond energy, which is determined by the work required to break this bond.
Types of chem. bonds: covalent (polar, non-polar, exchange and donor-acceptor), ionic, hydrogen and metallic.
The 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 distribution El. density between atoms. The distribution of e density is affected by the spatial orientation of the chemical. connections. If 2-atomic molecules are always linear, then the shapes of polyatomic molecules can be different.
The angle between imaginary lines that can be drawn through the centers of bonded atoms is called the valence angle. The density distribution e also depends on the size of a. and their eo. In homoatomic El. density is evenly distributed. In heteroatomic it is shifted in the direction that contributes to a decrease in the energy of the system.
The binding energy is the energy that is released during the formation of a molecule from single atoms. The binding energy differs from ΔHrev. 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 – ∆Harr.

N + O → NO - 89.96 kJ / mol - E St.

The bond multiplicity is determined by the number of electron pairs involved in the bond between atoms. The chemical bond is due to the overlap of electron clouds. If this overlap occurs along the line connecting the nuclei of atoms, 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 the bond is formed by more than one pair of electrons, then it is called a multiple.
A multiple bond is formed when there are too few electrons and bonding atoms for each bondable valence orbital of the central atom to overlap with any orbital of the surrounding atom.
Since the 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 the 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 with respect to the nuclei of both atoms. In this case, the covalent bond is called non-polar or homeopolar. If a diatomic molecule consists of atoms various elements, then the common electron cloud is shifted towards one of the atoms, so that there is an asymmetry in the charge distribution. In such cases, the covalent bond is called polar or heteropolar.

To assess the ability of an atom of a given element to pull a common electron pair towards itself, the value of relative electronegativity is used. The greater the electronegativity of an atom, the stronger it attracts a common 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. The values of the electronegativity of atoms of some elements in relation to the electronegativity of fluorine, which is taken equal to 4.
Electronegativity naturally changes depending on the position of the element in the periodic system. At the beginning of each period there are elements with the lowest electronegativity - typical metals, at the end of the period (before 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 an element is a metal, the lower its electronegativity; the more typical a non-metal an element is, the higher its electronegativity.

The reason for the formation of a chemical bond. Majority atoms chemical elements in an individual form does not exist, 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 nuclei of atoms. Valence electrons play the main role in the formation of a chemical bond 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 in atoms of spatially separated regions with an electric charge, electrostatic interactions can occur between different atoms that can hold these atoms together.
When a chemical bond is formed, there is a redistribution in space of electron densities that originally belonged to different atoms. Since the electrons of the outer level are the least strongly bound to the nucleus, it is precisely these electrons that 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 valence. For this reason, the outer level electrons are called valence electrons.

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

The binding energy is the energy that is released during the formation of a molecule from single atoms. The binding energy differs from ΔHrev. The heat of formation is the energy that is released or absorbed during the formation of molecules from simple substances. (The 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 atoms F-F equal to - 150.6 kJ / mol. For polyatomic molecules with one type of bond, for example, for ABn molecules, the average binding energy is equal to 1/n of the total energy of formation of a compound from atoms. So, the energy of formation of CH4 = -1661.1 kJ / mol.

If more than two different atoms combine in a molecule, then the average binding energy does not coincide with the value of the dissociation energy of the molecule. If different types of bonds are present in the molecule, then each of them can be approximately assigned certain value 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 by the equation:

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

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

The higher the bond order 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 due to the overlap of electron clouds. If this overlap occurs along the line connecting the nuclei of atoms, 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 the bond is formed by more than one pair of electrons, then it is called a multiple.

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

Since the 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 the 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 with respect to the nuclei of both atoms. In this case, the covalent bond is called non-polar or homeopolar. If a diatomic molecule consists of atoms of different elements, then the common electron cloud is shifted towards one of the atoms, so that there is an asymmetry in the charge distribution. In such cases, the covalent bond is called polar or heteropolar.

The displacement of the common electron cloud during the formation of a polar covalent bond leads to the fact that the average density of the negative electric charge turns out to be higher near a more electronegative atom and lower near a less electronegative one. As a result, the first atom acquires an excess negative, 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 the atoms of metals and non-metals, 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: covalent polar, covalent non-polar and ionic.

A covalent bond is called non-polar if the shared electron pair equally belongs to both atoms. A covalent non-polar bond occurs between atoms whose electronegativity is the same (between atoms of the same non-metal), i.e. in simple substances. For example, in the molecules of oxygen, nitrogen, chlorine, bromine, the bond is covalent non-polar.
A covalent bond is called polar if the shared electron pair is shifted towards one of the elements. A covalent polar bond occurs between atoms whose electronegativity differs, but not much, i.e. in complex substances between non-metal atoms. For example, in the molecules of water, hydrogen chloride, ammonia, sulfuric acid, the bond is covalent polar.
An ionic bond is a bond between ions, carried out due to the attraction of oppositely charged ions. An ionic bond occurs 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 balance. Equilibrium constant. Calculation of equilibrium concentrations.
Chemical equilibrium - state chemical system, in which one or more chemical reactions reversibly proceed, and the rates in each pair of forward-reverse reactions are equal to each other. For a system in chemical equilibrium, the concentrations of reagents, temperature, and other parameters of the system do not change with time.

A2 + B2 ⇄ 2AB

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

Equilibrium constant - a value that determines for a given chemical reaction the ratio between starting materials 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 limiting yield of products, and determine the direction of the reaction.

Ways of expressing the equilibrium constant:
For a reaction in a mixture ideal gases the equilibrium constant can be expressed in terms of the equilibrium partial pressures of the components pi by the formula:

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

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

The equilibrium constant can be calculated from the equation:

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

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

In some cases (depending on the way of expression), the equilibrium constant can be a function not only of temperature, but also of pressure. So, for a reaction in a mixture of ideal gases partial pressure component can be expressed according to Dalton's law through the total pressure and the 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 the pressure. If the number of moles of reaction products is equal to the number of moles of starting materials (Δn = 0), then Kp = Kx.