Binding energy c h. Chemical bond breaking energy. Chemical bond length

Sexual Energy - Energy of Money Power is an aphrodisiac. Sex equals power. Michael Hutchinson Psychologist Carl Jung invented a psychological model for men and women, which he called anima and animus. He admitted that every man has an inner

Hybridization of atomic orbitals. 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 interacting atoms change, and above all the energy and occupancy of their outer orbitals.

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PAGE 13

Lebedev Yu.A. Lecture 0 2

Lecture number 0 2

Chemical bond. Chemical bond characteristics: energy, length, bond angle. Types of chemical bonds. Communication polarity. Quantum-mechanical concepts of the nature of covalent bonds. The concept of the method of valence bonds. Hybridization of atomic orbitals.- (c igma) and (pi) -connection. Geometric configuration of molecules. The electric moment of the dipole of the molecule. 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. Chemical bond characteristics: energy, length, bond angle.

We examined the structure of the electronic levels of isolated atoms. These are objects very rare 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 pieces of argon atoms in one breath.

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

The chemical bond is due to 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 on ours - valence electrons.

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

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

Typical chemical bond energies are hundreds of kJ / mol.

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

Typical chemical bond lengths are in 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.

By the nature of the interaction of valence electrons and the type of orbitals formed during the interaction,chemical bonds are classified 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 depicted by a valence line:

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

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

With the formation of a polar covalent bond, atoms acquire an additional charge - negative for an atom with a higher electronegativity and positive for an atom with a lower electronegativity:

H + Cl = HCl (
–
)

In the case when the difference in the electronegativities of the 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 donor-acceptor. The atom that provided the electron pair is called the donor, and the one that received it into the free orbital is called the acceptor.

The formation of donor-acceptor bonds is especially characteristic. d - metals with free or partially filled d -orbitals with the formation of complex compounds.

We'll talk about other types of communication later.

Quantum-mechanical concepts of the nature of covalent bonds.

From a modern point of view, a covalent bond occurs during the quantum mechanical interaction of all electrons of all interacting atoms. But, as we said in lecture No. 1, there is no exact solution to the Schrödinger equation describing the orbitals of many electrons in molecules. The problem of the quantum-mechanical description of a chemical bond is facilitated 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 applied programs that make it possible to perform calculations with high accuracy for a wide class of molecules and ions. 4

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

For a qualitative understanding of the structure of chemical compounds, two methods are used -valence bond method (MVS) and molecular orbital method (MO).

The concept of the method of valence bonds. Geometric configuration of molecules. The electric moment of the dipole of the 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 electrons forming a valence pair should be opposite (electrons with antiparallel spins form a bond).

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

(3) The bond is formed due to the overlap of the orbitals, which leads to an increase in the electron density between the nuclei of interacting atoms in the direction providing the maximum overlap.

Consider the formation of a chemical bond along the MVC in a water vapor molecule - H 2 O.

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 ... There are 6 electrons on the outer energy level. Sublevel 2 s is full. At sublevel 2 p on one of p -orbitals (put p y ,) there is an electron pair, and on the other two ( p x and p z ) - one unpaired electron each. It is they who will participate in the formation of chemical bonds.

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

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

Not a bad match for the simplest quality model without any calculations!

The electronegativity of oxygen according to Mulliken is 3.5, and that of hydrogen is 2.1. Therefore, 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.

A dipole is two equal charges located at a finite distance l apart. A dipole is characterized by a dipole moment

=

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

It is important to emphasize that the dipole moments of bonds are added vectorially and the total dipole moment depends on the geometry of the molecule. As you can 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 the experiment confirms this - the dipole moment of the water molecule is 1.84 Debye. (1 Debye equals 0.33 * 10-29 Cl * 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 three-dimensional bodies of various configurations (trigonal, tetragonal, hexagonal pyramids, bipyramids, rings composed of pyramids, etc.)

Read more about the relationship between the structure of chemical bonds and the geometry of molecules in the textbook on pages 119-128).

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

Let's go back to the overlapping of orbitals during bond formation. In our examplearea of ​​maximum overlap s and p -orbitals lies on the line connecting the centers of the atoms. This type of overlap is called-connection.

Consider another case - an oxygen molecule O 2 ... 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 just reviewed-connection. And the second one? It turns out that the second bond is formed due to another type of orbital overlap, which is called- communication.

Concept of and connections put forward F.Hund.

In education -bonds of the orbitals overlap in such a way that two overlap regions 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-connection is formed by smaller parts p -orbitals, in which the density of the "electron cloud" is greater, and therefore this bond is stronger-connection. 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 -connection and one -bond) and acetylene C 2 H 2 (C НС H - one -connection and two -bonds) their breaking energy is 247, 419, and 515 kJ / mol, respectively.

Now we can add to the list of MCS postulates:

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

Note that in the connections d - and f -metals, the formation of another type of bonds is possible --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 are formed, an important phenomenon can occur, which is calledorbital hybridization.

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

However, let's take a closer look at the electron diffraction diagram of the beryllium atom:

It can be seen from the diagram 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 the valueE ... But this energy is small and less than that which is released during the formation of a chemical bond. Therefore, the atom seeks to rearrange its orbitals in the course of interaction in order to achieve an energetically favorable final state. For such a rearrangement, the kinetic energy of the particles interacting with the given atom is used. We will talk in more detail about this source of energy when discussing questions of chemical kinetics. 5

Such a restructuring is called orbital hybridization, since in the course of this process a new one arises from the “two kinds” of orbitals.

In the language of wave functions, this is described by an equation connecting the hybrid wave function of the resulting orbitals with the original wave functions.

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

This process can be graphically depicted in the following diagram:

Note that the energy required for hybridization E hybrid less than the difference in energies of hybridizing orbitals E.

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

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

In this case, 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 by 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 tendency of the like-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 contain 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.

The modification of the MVS, taking into account the influence of such orbitals, is called the theory of repulsion of electron pairs of valence orbitals (OEPVO) and you can get acquainted with it from the textbook on pages 124 - 128.

Concept of the molecular orbital method.

We have considered the phenomenon of AO hybridization within the framework of the MFM. It turned out that the idea of ​​hybridization is fruitful even for deeper modeling of chemical bonds. It is the basis of the second method of describing them, which is considered in our course - the 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 "belong" not to any particular atom, but quantum-mechanically move throughout the entire molecular structure.

There are several varieties of the MO method, taking into account b O a greater or lesser number of factors and, accordingly, more or less complex mathematically. The simplest is the approximation that takes into account only the linear effects of electron interaction. 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 is the MO wave function,
is the wave function of the AO of the first atom,
is the wave function of the AO of the second atom, a and b - numerical coefficients characterizing the contribution of a given AO to the general structure of the MO.

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

It can be seen from the equation thatwhen two AOs interact, two MOs are obtained... One of them is called binding MO, and the other - loosening MO.

Why they got this name is clear from the figure, which shows the energy diagram of the orbitals in the molecule:

As can be seen from the figure, the binding MO has an energy lower than the energies of the initial AO, and the antibonding one is higher. (Respectively,). Naturally, in accordance with the principle of minimum energy, electrons in a molecule will first of all occupy the bonding orbital during bond formation.

In general, when interacting N AO turns out to be 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 symmetry of orbitals are preserved in the MO LCAO method.

This is what the outline looks like-binding (denoted asor) and -bonding (denoted as or) orbitals in the MO LCAO method:

And this is what the outlines look like- connecting ( ) and - loosening ( * ) orbitals by the MO LCAO 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 Mendeleev's Periodic Table is universal and has the form:

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

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

One of the serious advantages of the MO LCAO method in comparison with the VS 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 MVC, which we considered earlier. In accordance with this structure, all valence electrons andand -bonds in the molecule O 2 form electron pairs and the total spin of the molecule is zero.

The structure of the orbitals of this molecule by the MO LCAO method, obtained by filling the MO 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 antibonding
and
orbitals. Their magnetic moments add up and give the total magnetic moment of the molecule. Experiment shows that the magnetic moment of the oxygen molecule is 2.8(The intrinsic magnetic moment of the electron is 1). Considering that the total magnetic moment, in addition to its own electronic moment, includes the orbital one, the quantitative coincidence is very convincing evidence in favor of the validity of the MO method.

In the presence of a magnetic moment, the substance becomesparamagnetic -it 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 the magnetic properties, the analysis of the energy diagrams of the MO LCAO makes it possible to determinethe multiplicity (or order) of the chemical bond (CS or PS).

КС = ½ (N connection - N bit)

where N bond - the total number of electrons in the bonding orbitals; N bit Is the total number of electrons in antibonding orbitals).

We have considered various cases of manifestation and description of covalent chemical bonds. This is the main type of chemical bond, since the cause of its occurrence - the presence of valence electrons - is in the overwhelming majority of chemical elements.

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

When a chemical bond is formed, a redistribution of electron densities in space, which originally belonged to different atoms, occurs. Since the electrons of the outer level are 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 taking part 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 the penultimate d-electrons. From an energy point of view, the most stable is the atom, the outer level of which contains the maximum number of electrons (2 and 8 electrons). This level is called complete. The completed levels are highly durable and are characteristic of noble gas atoms, therefore, 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 by 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 seeks 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 chemical bonds. 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 appearance of a chemical bond always proceeds exothermically, since it leads to the appearance of new particles (molecules), which are more stable under normal conditions, and therefore, they are less energy than the original. 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 to itself. The electronegativity of atoms of elements changes gradually: in the periods of the periodic table, from left to right, its value increases, and in groups from top to bottom, it decreases.

The chemical bond, carried out due to the formation of common (bonding) electron pairs, is called covalent. 1) Let us examine 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 or a dash that symbolizes a pair of electrons: HH 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, the electronegativity of which differs slightly. A covalent bond formed by atoms with different electronegativity is called polar. With a covalent polar bond, the electron density from a common pair of electrons is shifted to an atom with greater electronegativity. Examples are H2O, NH3, H2S, CH3Cl molecules. The covalent (polar and non-polar) bond in our examples was formed due to the unpaired electrons of the bonding atoms. This mechanism for the formation of a covalent bond is called exchange. 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 the free orbital of another atom (acceptor). A well-known example is the formation of an ammonium ion: H ++: NH 3 -> [H: NH3 | +<=====>NH4 + acceptor donor electron ammonium ion. With the formation of an ammonium ion, 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 the same:

An ionic bond arises between atoms, the electronegativity of which is sharply different. Let us consider the method of formation using the example of sodium chloride NaCl. The electronic configuration of sodium and chlorine atoms can be represented: 11 Na ls2 2s2 2p 6 3s1; 17 Cl ls2 2p 6 Зs2 3p5 Like atoms with incomplete energy levels. Obviously, for their completion, it is easier for a sodium atom to donate one electron than to attach seven, and it is easier for a chlorine atom to attach one electron than to donate seven. In chemical interaction, the sodium atom completely donates one electron, and the chlorine atom accepts it. Schematically it can be written like this: Na. - l е -> 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. Forces of electrostatic attraction arise between the Na + and Cl- ions, as a result of which a compound is formed.

The chemical bond carried out by electrostatic attraction between ions is called ionic bond. Compounds formed by the attraction of ions are called ionic. Ionic compounds consist of individual molecules only in a vapor state. In the solid (crystalline) state, ionic compounds consist of regularly spaced positive and negative ions. Molecules are absent in this case. Ionic compounds form elements of the main subgroups I and II of groups and the main subgroups of VI and VII groups, sharply differing in magnitude of 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, salts of carboxylic acids, salts of amines 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 the bonding agent of two electron clouds between identical atoms. On the contrary, in the case of ionic bonding, the electron cloud that binds is almost entirely owned by one of the atoms. In most compounds, chemical bonds are intermediate between these types of bonds, that is, a polar covalent bond is carried out in them.

A metallic bond exists in metals in a solid in a liquid state. In accordance with the position in the periodic table, metal atoms have a small number of valence electrons (1-3 electrons) and a low ionization energy (electron detachment). Therefore, valence electrons are weakly retained in the atom, are easily torn off and have the ability to move throughout the crystal. In the nodes of the crystal lattice of metals there are free atoms, positively charged ions, and part of the valence electrons, freely moving in the volume of the crystal lattice, forms an "electron gas" that provides a bond between metal atoms. The bond, which is carried out by relatively free electrons between metal ions in the crystal lattice, is called a metal bond. The 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 carry out a covalent bond are mainly in the immediate vicinity of two connected atoms. In the case of a metal bond, the electrons that make the bond move around the entire piece of metal. This determines the general characteristics of metals: metallic luster, good conductivity of heat and electricity, malleability, ductility, etc. The general chemical property of metals is their relatively high reducibility.

Hydrogen bonds can form between a hydrogen atom bonded to an electronegative element and an electronegative element that has 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. O-H bonds have a pronounced polar character: The hydrogen bond is much weaker than ionic or covalent, but stronger than intermolecular interaction. Hydrogen bonds determine some of the physical properties of substances (for example, high boiling points). Hydrogen bonds are especially widespread in molecules of proteins, nucleic acids and other biologically important compounds, providing them with a certain spatial structure (organization).

Communication 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.

Link length (lw). The distance between the nuclei at the junction. The longer the bond length, the lower the bond energy.

Method of valence bonds.

• 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 exchange of electrons 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 interaction forces of the particles.
• C) in accordance with the Pauli principle, a chemical bond is formed only when electrons interact with different spins.
• D) the characteristics of the chemical bond (energy, length, polarity) are determined by the type of overlapping AO.

Method of valence bonds. The covalent bond is directed towards the 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.

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.

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

• 1. The number of MO is equal to the total number of AOs from which the MO is combined.
• 2. The energy of some MO turns out to be higher, others - lower than the energy of the initial AO. The average energy of MO obtained from a set of AOs approximately coincides with the average energy of these AOs.
• 3. Electrons fill MO, as well as AO, in the order of increasing energy, while the Pauli exclusion principle and Gund'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 overlapping of atomic orbitals.

Order and energy of communication. The order of communication is n = (Ncw-Np) / 2. Nw is the number of e on the bonding molecular orbitals, Np is the number of e on the antibonding molecular orbitals.

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

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

Tutorial

2. Astrakhan

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

Designed for students of engineering and technical non-chemical specialties.

Complies with state educational standards of higher professional education

Ill .: 15 figures, table: 1, bibliography: 6 titles, app.

Published by the decision of the Department of General, Inorganic and Analytical Chemistry (protocol No.__ dated _________ 2013)

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

© Ryabukhin Yu.I., 2013

© ASTU, 2013

INTRODUCTION

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

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

Chemical bond 1 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 arises 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 - this 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 resulting substance (E sv , kJ / mol) 2.

The higher the chemical bond energy, the stronger the bond. The energy of the chemical bond of a diatomic molecule is estimated by comparing it with the state prior to 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 energy of formation 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 overlapping of their electron shells.

For example, for hydrogen fluoride and hydrogen iodide:

l HF< l HI

Depending on the type of particles (atoms or molecules) to be joined, a distinction is made between intramolecular bonds, due to 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 - between the atoms there is 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 an increase in the repulsive forces between the nuclei.

Depending on how the common electron pair interacts with the nuclei of the atoms to be joined, three main types of chemical bonds are distinguished: ovalent, ionic and metallic, as well as hydrogen bonds.

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

Experimental data on the determination of 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

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 PEI of the bonded atoms.

Is it possible that in diatomic molecules (where there is more than one electron), the bond is formed by a different mechanism and there are additional forces not previously accounted for by us?

Before moving on to identifying these forces, let's try to explain this independence based on pre-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 table (4.2) into four groups:

Group A includes molecules consisting of identical atoms, in which the binding energy is 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 in comparison with the binding energy in the hydrogen molecule, which is 429 kJ / mol.

GroupWITH includes diatomic molecules consisting of different atoms, the binding energy in 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

Before we start explaining, let's clarify the questions that we need to cover.
First question:
Why is the binding energy between many-electron atoms much less or much more (table 4.2) than in a hydrogen molecule (H 2)?

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 potential energy value of the system atom + electron increases as a result of the bond of an electron with an atom. The data on the affinity of an atom for an electron, shown in Table 4.3, give us the numerical value of the energy gain when an electron is attached to an atom.

table 4.3

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 electron-electron repulsion increases due to the 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 attractive energy is greater than the loss of energy due to an increase in the repulsive energy.

Counting the change in energy when an electron is attached 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 exactly 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 attachment of an electron does not lead 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 ability to attach or not attach 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 it will not.

The data on the electron affinity of atoms given in Table 4.3 show that for atoms of the 1st, 2nd and 3rd periods, except 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 electron-electron repulsion. This conclusion is independently confirmed by 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 on the outer electron shells of atoms increases by one electron, and according to the calculation of the model of the hydrogen molecule H 2, the effective charges of the bonded atoms change. The effective charges of the bound nuclei change due to the attraction of the charged nuclei, and in connection with an increase in the number of electrons on the outer shells of the bound atoms.

In a hydrogen molecule, the approach of the nuclei leads to an increase in the force of attraction of the binding electrons to the 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 the data in 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 is reduced to 100 kJ / mol.
Recall that the outer electron shell of boron has 3 electrons, and the outer shell of beryllium contains 2 electrons. That is, when an electron is attached to 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 it is quite understandable a sharp decrease in the binding energy of atoms with a negative affinity of the atom for the electron, indicated in Table 4.3. However, though Ne,Be,Mg,Ar do not attach electrons, they create molecules, because the effective charge of the nuclei increases. The binding energy in these molecules (group A) is much lower than in other molecules.

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

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

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

Now consider the group WITH: Why is the binding energy in some of the diatomic molecules, made up of different atoms, significantly higher than in other molecules, which are made up of the same atoms?

Let's analyze the molecule NaCl... Sodium and chlorine atoms differ greatly in their electron affinity. Introduce bond formation as a two-step 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 Cl 2, should be more than during the formation of the molecule NaCl by the difference in their electron affinity.

When calculating a hydrogen molecule (Chapter 3), the binding energy (the energy required to separate the 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 bonding 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 the bonding pairs of electrons, as well as to estimate the energy of electron-electron repulsion, we must first find out the value of the effective charge of the bonded nuclei.

Ionization potential and binding energy in diatomic molecules