There is no unified theory of chemical bonds; chemical bonds are conventionally divided into covalent (a universal type of bond), ionic (a special case of a covalent bond), metallic and hydrogen.
Covalent bond
The formation of a covalent bond is possible by three mechanisms: exchange, donor-acceptor and dative (Lewis).
According to metabolic mechanism The formation of a covalent bond occurs due to the sharing of common electron pairs. In this case, each atom tends to acquire a shell of an inert gas, i.e. obtain a completed external energy level. The formation of a chemical bond by exchange type is depicted using Lewis formulas, in which each valence electron of an atom is represented by dots (Fig. 1).
Rice. 1 Formation of a covalent bond in the HCl molecule by the exchange mechanism
With the development of the theory of atomic structure and quantum mechanics, the formation of a covalent bond is represented as the overlap of electronic orbitals (Fig. 2).
Rice. 2. Formation of a covalent bond due to the overlap of electron clouds
The greater the overlap of atomic orbitals, the stronger the bond, the shorter the bond length, and the greater the bond energy. A covalent bond can be formed by overlapping different orbitals. As a result of the overlap of s-s, s-p orbitals, as well as d-d, p-p, d-p orbitals with lateral lobes, the formation of bonds occurs. A bond is formed perpendicular to the line connecting the nuclei of 2 atoms. One and one bond are capable of forming a multiple (double) covalent bond, characteristic of organic substances of the class of alkenes, alkadienes, etc. One and two bonds form a multiple (triple) covalent bond, characteristic of organic substances of the class of alkynes (acetylenes).
Formation of a covalent bond by donor-acceptor mechanism Let's look at the example of the ammonium cation:
NH 3 + H + = NH 4 +
7 N 1s 2 2s 2 2p 3
The nitrogen atom has a free lone pair of electrons (electrons not involved in the formation of chemical bonds within the molecule), and the hydrogen cation has a free orbital, so they are an electron donor and acceptor, respectively.
Let us consider the dative mechanism of covalent bond formation using the example of a chlorine molecule.
17 Cl 1s 2 2s 2 2p 6 3s 2 3p 5
The chlorine atom has both a free lone pair of electrons and vacant orbitals, therefore, it can exhibit the properties of both a donor and an acceptor. Therefore, when a chlorine molecule is formed, one chlorine atom acts as a donor and the other as an acceptor.
Main characteristics of a covalent bond are: saturation (saturated bonds are formed when an atom attaches as many electrons to itself as its valence capabilities allow; unsaturated bonds are formed when the number of attached electrons is less than the valence capabilities of the atom); directionality (this value is related to the geometry of the molecule and the concept of “bond angle” - the angle between bonds).
Ionic bond
There are no compounds with a pure ionic bond, although this is understood as a chemically bonded state of atoms in which a stable electronic environment of the atom is created when the total electron density is completely transferred to the atom of a more electronegative element. Ionic bonding is possible only between atoms of electronegative and electropositive elements that are in the state of oppositely charged ions - cations and anions.
DEFINITION
Ion are electrically charged particles formed by the removal or addition of an electron to an atom.
When transferring an electron, metal and nonmetal atoms tend to form a stable electron shell configuration around their nucleus. A non-metal atom creates a shell of the subsequent inert gas around its core, and a metal atom creates a shell of the previous inert gas (Fig. 3).
Rice. 3. Formation of an ionic bond using the example of a sodium chloride molecule
Molecules in which ionic bonds exist in their pure form are found in the vapor state of the substance. The ionic bond is very strong, and therefore substances with this bond have a high melting point. Unlike covalent bonds, ionic bonds are not characterized by directionality and saturation, since the electric field created by ions acts equally on all ions due to spherical symmetry.
Metal connection
The metallic bond is realized only in metals - this is the interaction that holds metal atoms in a single lattice. Only the valence electrons of the metal atoms belonging to its entire volume participate in the formation of a bond. In metals, electrons are constantly stripped from atoms and move throughout the entire mass of the metal. Metal atoms, deprived of electrons, turn into positively charged ions, which tend to accept moving electrons. This continuous process forms the so-called “electron gas” inside the metal, which firmly binds all the metal atoms together (Fig. 4).
The metallic bond is strong, therefore metals are characterized by a high melting point, and the presence of “electron gas” gives metals malleability and ductility.
Hydrogen bond
A hydrogen bond is a specific intermolecular interaction, because its occurrence and strength depend on the chemical nature of the substance. It is formed between molecules in which a hydrogen atom is bonded to an atom with high electronegativity (O, N, S). The occurrence of a hydrogen bond depends on two reasons: firstly, the hydrogen atom associated with an electronegative atom does not have electrons and can easily be incorporated into the electron clouds of other atoms, and, secondly, having a valence s-orbital, the hydrogen atom is able to accept a lone pair electrons of an electronegative atom and form a bond with it through the donor-acceptor mechanism.
The most important characteristics of a bond include: length, polarity, dipole moment, saturation, directionality, strength, and multiplicity of the bond.
Link length– is the distance between the nuclei of atoms in a molecule. The bond length is determined by the size of the nuclei and the degree of overlap of the electron clouds.
The bond length in HF is 0.92∙10 -10, in HCl – 1.28∙10 -10 m. The shorter its length, the stronger the chemical bond.
Bond angle (Bond angle) call the angle between imaginary lines passing through the nuclei of chemically bonded atoms. ∟HOH=104 0 .5; ∟H 2 S=92.2 0; ∟H 2 S e =91 0 .0.
The most important characteristic of a chemical bond is energy, defining it strength.
The bond strength is quantitatively characterized by the energy expended to break it and is measured in kJ per 1 mole of substance.
Therefore, the bond strength is quantitatively characterized by the sublimation energy E subl. substances and energy of dissociation of a molecule into atoms E diss. . Sublimation energy refers to the energy expended to transition a substance from a solid to a gaseous state. For diatomic molecules, the binding energy is equal to the energy of dissociation of the molecule into two atoms.
For example, E diss. (and therefore E St.) in the H 2 molecule is 435 kJ/mol. In the F 2 molecule = 159 kJ/mol, in the N 2 molecule = 940 kJ/mol.
For not diatomic, but polyatomic molecules of type AB n, the average binding energy
by AB n =A+nB.
For example, the energy absorbed during the process
equal to 924 kJ/mol.
Communication energy
E OH = = = = 462 kJ/mol.
Conclusions about the structure of molecules and the structure of a substance are made based on the results obtained by different methods. In this case, the obtained information is used not only about bond lengths and energies, bond angles, but also other properties of the substance, such as magnetic, optical, electrical, thermal and others.
The set of experimentally obtained data on the structure of matter complements and generalizes the results of quantum chemical calculation methods that use the concept of the quantum mechanical theory of chemical bonding. Chemical bonding is believed to be primarily mediated by valence electrons. For s- and p-elements, the valence electrons are the electrons of the orbitals of the outer layer, and for d-elements, the electrons are the s-orbitals of the outer layer and the d-orbitals of the pre-outer layer.
The nature of the chemical bond.
A chemical bond is formed only if, as atoms approach each other, the total energy of the system (E kin. + E pot.) decreases.
Let's consider the nature of a chemical bond using the example of the molecular hydrogen ion H 2 +. (It is obtained by irradiating hydrogen molecules with H 2 electrons; in a gas discharge). For such a simple molecular system, the Schrödinger equation is most accurately solved.
In the hydrogen ion H 2 + one electron moves in the field of two nuclei - protons. The distance between the nuclei is 0.106 nm, the binding energy (dissociation into H atoms and H + ion) is 255.7 kJ/mol. That is, the particle is durable.
In the molecular ion H 2 + there are two types of electrostatic forces - the force of attraction of an electron to both nuclei and the force of repulsion between nuclei. The repulsive force manifests itself between the positively charged nuclei H A + and H A +, which can be represented in the form of the following figure. 3. The repulsive force tends to push the nuclei apart from each other.
Rice. 3. The force of repulsion (a) and attraction (b) between two nuclei, arising when they approach each other at distances on the order of the size of atoms.
Attractive forces act between the negatively charged electron e - and the positively charged nuclei H + and H +. A molecule is formed if the resultant of the forces of attraction and repulsion is zero, that is, the mutual repulsion of nuclei must be compensated by the attraction of the electron to the nuclei. Such compensation depends on the location of the electron e - relative to the nuclei (Fig. 3 b and c). What is meant here is not the position of the electron in space (which cannot be determined), but the probability of finding the electron in space. Location of electron density in space, corresponding to Fig. 3.b) promotes the convergence of nuclei, and the corresponding Fig. 3.c) – repulsion of nuclei, since in this case the attractive forces are directed in one direction and the repulsion of nuclei is not compensated. Thus, there is a binding region, when the electron density is distributed between the nuclei, and an antibonding or antibonding region, when the electron density is distributed behind the nuclei.
If an electron enters the bonding region, a chemical bond is formed. If the electron falls into the antibonding region, then a chemical bond is not formed.
Depending on the nature of the electron density distribution in the binding region, three main types of chemical bonds are distinguished: covalent, ionic and metallic. These bonds do not occur in their pure form, and usually a combination of these types of bonds is present in compounds.
Types of connections.
In chemistry, the following types of bonds are distinguished: covalent, ionic, metallic, hydrogen bond, van der Waals bond, donor-acceptor bond, dative bond.
Covalent bond
When a covalent bond is formed, atoms share electrons with each other. An example of a covalent bond is the chemical bond in the Cl 2 molecule. Lewis (1916) first proposed that in such a bond, each of the two chlorine atoms shares one of its outer electrons with the other chlorine atom. To overlap atomic orbitals, two atoms must come as close to each other as possible. A shared pair of electrons forms a covalent bond. These electrons occupy the same orbital, and their spins are directed in opposite directions.
Thus, a covalent bond is accomplished by sharing electrons from different atoms as a result of pairing of electrons with opposite spins.
Covalent bonding is a common type of bonding. Covalent bonds can occur not only in molecules, but also in crystals. It occurs between identical atoms (in molecules of H 2, Cl 2, diamond) and between different atoms (in molecules of H 2 O, NH 3 ...)
Mechanism of covalent bond formation
Let us consider the mechanism using the example of the formation of the H 2 molecule.
H+H=H 2, ∆H=-436 kJ/mol
The nucleus of a free hydrogen atom is surrounded by a spherically symmetric electron cloud formed by a 1s electron. When atoms approach a certain distance, their electron clouds (orbitals) partially overlap (Fig. 4).
Rice. 4. The mechanism of bond formation in a hydrogen molecule.
If the hydrogen atoms approaching before touching have a distance between the nuclei of 0.106 nm, then after the electron clouds overlap, this distance is 0.074 nm.
As a result, a molecular two-electron cloud appears between the centers of the nuclei, which has a maximum electron density in the space between the nuclei. An increase in the negative charge density between nuclei favors a strong increase in the attractive forces between nuclei, which leads to the release of energy. The greater the overlap of electron orbitals, the stronger the chemical bond. As a result of the formation of a chemical bond between two hydrogen atoms, each of them reaches the electronic configuration of a noble gas atom - helium.
There are two methods that explain from a quantum mechanical point of view the formation of the area of overlap of electron clouds, and the formation of a covalent bond, respectively. One of them is called the BC (valence bonds) method, the other MO (molecular orbitals).
The valence bond method considers the overlap of atomic orbitals of a selected pair of atoms. In the MO method, the molecule is considered as a whole and the distribution of electron density (from one electron) is spread over the entire molecule. From the position of MO 2H in H 2 are connected due to the attraction of nuclei to the electron cloud located between these nuclei.
Illustration of a covalent bond
Connections are depicted in different ways:
1). Using electrons as dots
In this case, the formation of a hydrogen molecule is shown by the diagram
N∙ + N∙ → N: N
2). Using square cells (orbitals), like placing two electrons with opposite spins in one molecular quantum cell
This diagram shows that the molecular energy level is lower than the original atomic levels, which means the molecular state of the substance is more stable than the atomic one.
3). A covalent bond is represented by a line
For example, H – N. This line symbolizes a pair of electrons.
If one covalent bond (one common electron pair) occurs between atoms, then it is called single, if more, then a multiple double(two common electron pairs), triple(three common electron pairs). A single bond is represented by one line, a double bond by two lines, and a triple bond by three lines.
The dash between the atoms shows that they have a generalized pair of electrons.
Classification of covalent bonds
Depending on the direction of overlap of electron clouds, σ-, π-, δ-bonds are distinguished. The σ bond occurs when electron clouds overlap along the axis connecting the nuclei of interacting atoms.
Examples of σ-bonds:
Rice. 5. Formation of a σ bond between s-, p-, d- electrons.
An example of the formation of a σ bond when s-s clouds overlap is observed in the hydrogen molecule.
The π bond occurs when the electron clouds on either side of the axis overlap, connecting the nuclei of atoms.
Rice. 6. Formation of π-bond between p-, d- electrons.
δ-coupling occurs when two d-electron clouds located in parallel planes overlap. The δ bond is less strong than the π bond, and the π bond is less strong than the σ bond.
Properties of covalent bonds
A). Polarity.
There are two types of covalent bonds: nonpolar and polar.
In the case of a nonpolar covalent bond, the electron cloud formed by a common pair of electrons is distributed in space symmetrically relative to the atomic nuclei. An example is diatomic molecules consisting of atoms of one element: H 2, Cl 2, O 2, N 2, F 2. Their electron pair belongs equally to both atoms.
In the case of a polar bond, the electron cloud forming the bond is shifted toward the atom with higher relative electronegativity.
Examples are the following molecules: HCl, H 2 O, H 2 S, N 2 S, NH 3, etc. Consider the formation of an HCl molecule, which can be represented by the following diagram
The electron pair is shifted to the chlorine atom, because the relative electronegativity of the chlorine atom (2.83) is greater than that of the hydrogen atom (2.1).
b). Saturability.
The ability of atoms to participate in the formation of a limited number of covalent bonds is called the saturation of a covalent bond. The saturation of covalent bonds is due to the fact that only electrons from external energy levels, that is, a limited number of electrons, participate in chemical interactions.
V) . Focus and covalent bond hybridization.
A covalent bond is characterized by directionality in space. This is explained by the fact that electron clouds have a certain shape and their maximum overlap is possible at a certain spatial orientation.
The direction of a covalent bond determines the geometric structure of molecules.
For example, for water it has a triangular shape.
Rice. 7. Spatial structure of a water molecule.
It has been experimentally established that in a water molecule H 2 O the distance between the hydrogen and oxygen nuclei is 0.096 nm (96 pm). The angle between the lines passing through the nuclei is 104.5 0. Thus, the water molecule has an angular shape and its structure can be expressed in the form of the presented figure.
Hybridization
As experimental and theoretical studies (Slater, Pauling) show, during the formation of some compounds, such as BeCl 2, BeF 2, BeBr 2, the state of the valence electrons of an atom in a molecule is described not by pure s-, p-, d- wave functions, but by their linear combinations . Such mixed structures are called hybrid orbitals, and the mixing process is called hybridization.
As quantum chemical calculations show, mixing the s- and p-orbitals of an atom is a process favorable for the formation of a molecule. In this case, more energy is released than in the formation of bonds involving pure s- and p-orbitals. Therefore, the hybridization of the electronic orbitals of an atom leads to a large decrease in the energy of the system and, accordingly, an increase in the stability of the molecule. The hybridized orbital is more elongated on one side of the nucleus than on the other. Therefore, the electron density in the region of overlap of the hybrid cloud will be greater than the electron density in the region of overlap of the s- and p-orbitals separately, as a result of which the bond formed by the electrons of the hybrid orbital is characterized by greater strength.
Several types of hybrid states occur. When s- and p-orbitals hybridize (called sp-hybridization), two hybrid orbitals arise, located at an angle of 180 0 relative to each other. In this case, a linear structure is formed. This configuration (structure) is known for most alkaline earth metal halides (for example, BeX 2, where X = Cl, F, Br), i.e. The bond angle is 180 0 C.
Rice. 8. sp hybridization
Another type of hybridization, called sp 2 hybridization (formed from one s and two p orbitals), leads to the formation of three hybrid orbitals, which are located at an angle of 120 0 to each other. In this case, a trigonal structure of the molecule (or a regular triangle) is formed in space. Such structures are known for compounds BX 3 (X=Cl, F, Br).
Rice. 9. sp 2 -hybridization.
No less common is sp 3 hybridization, which is formed from one s- and three p- orbitals. In this case, four hybrid orbitals are formed, oriented in space symmetrically to the four vertices of the tetrahedron, that is, they are located at an angle of 109 0 28 ". This spatial position is called tetrahedral. This structure is known for molecules NH 3, H 2 O and in general for elements of the II period. Schematically its appearance in space can be displayed in the following figure
Rice. 10. Spatial arrangement of bonds in the ammonia molecule,
projected onto a plane.
The formation of tetrahedral bonds due to sp 3 hybridization can be represented as follows (Fig. 11):
Rice. 11. Formation of tetrahedral bonds during sp 3 hybridization.
The formation of tetrahedral bonds during sp 3 hybridization using the example of a CCl 4 molecule is shown in Fig. 12.
Fig. 12. Formation of tetrahedral bonds during sp 3 - hybridization into CCl 4 molecules
Hybridization does not only concern s- and p-orbitals. To explain the stereochemical elements of III and subsequent periods, there is a need to construct hybrid orbitals simultaneously including s-, p-, d- orbitals.
Substances with covalent bonds include:
1. organic compounds;
2. solid and liquid substances in which bonds are formed between pairs of halogen atoms, as well as between pairs of hydrogen, nitrogen and oxygen atoms, for example, H2;
3. elements of group VI (for example, spiral chains of tellurium), elements of group V (for example, arsenic), elements of group IV (diamond, silicon, germanium);
4. compounds that obey the 8-N rule (such as InSb, CdS, GaAs, CdTe), when their constituent elements are located in II-VI, III-V groups in the periodic table.
In solids with covalent bonds, different crystal structures can be formed for the same substance, the binding energy of which is almost the same. For example, the structure of ZnS can be cubic (zincblende) or hexagonal (wurtzite). The arrangement of the nearest neighbors in zinc blende and wurtzite is the same, and the only and small difference in the energies of these two structures is determined by the arrangement of the atoms next to the nearest ones. This ability of some substances is called allotropy or polymorphism. Another example of allotropy is silicon carbide, which has a number of polytypes of different structures from purely cubic to hexagonal. These numerous crystalline modifications of ZnS, SiC exist at room temperature.
Ionic bond
Ionic bonding is an electrostatic force of attraction between ions with charges of opposite sign (i.e. + and −).
The idea of ionic bonding was formed on the basis of the ideas of V. Kossel. He suggested (1916) that when two atoms interact, one gives up and the other accepts electrons. Thus, an ionic bond is formed by the transfer of one or more electrons from one atom to another. For example, in sodium chloride, an ionic bond is formed by the transfer of an electron from a sodium atom to a chlorine atom. As a result of this transfer, a sodium ion with a charge of +1 and a chloride ion with a charge of -1 are formed. They are attracted to each other by electrostatic forces, forming a stable molecule. The electron transfer model proposed by Kossel allows one to explain the formation of such compounds as lithium fluoride, calcium oxide, and lithium oxide.
The most typical ionic compounds consist of metal cations belonging to groups I and II of the periodic system, and anions of non-metallic elements belonging to groups VI and VII.
The ease of formation of an ionic compound depends on the ease of formation of its constituent cations and anions. The ease of formation is higher, the lower the ionization energy of the atom donating electrons (electron donor), and the atom adding electrons (electron acceptor) has a higher affinity for the electron. Electron affinity is a measure of the ability of an atom to gain an electron. It is quantified as the change in energy that occurs when one mole of singly charged anions is formed from one mole of atoms. This is the so-called “first electron affinity” concept. The second electron affinity is the energy change that occurs when one mole of doubly charged anions is formed from one mole of singly charged anions. These concepts, that is, ionization energy and electron affinity, relate to gaseous substances and are characteristics of atoms and ions in the gaseous state. But it should be borne in mind that most ionic compounds are most stable in the solid state. This circumstance is explained by the existence of a crystal lattice in them in the solid state. The question arises. Why, after all, are ionic compounds more stable in the form of crystal lattices, and not in the gaseous state? The answer to this question is the calculation of the energy of the crystal lattice, based on the electrostatic model. In addition to this, this calculation is also a test of the theory of ionic bonding.
To calculate the energy of a crystal lattice, it is necessary to determine the work that needs to be spent on destroying the crystal lattice with the formation of gaseous ions. To carry out the calculation, the idea of the forces of attraction and repulsion is used. The expression for the potential energy of interaction of singly charged ions is obtained by summing the energy of attraction and energy of repulsion
E = E in + E out (1).
The energy of Coulomb attraction of ions of opposite signs is taken as Eat, for example, Na + and Cl - for the NaCl compound
E incoming = -e 2 /4πε 0 r (2),
since the distribution of electronic charge in a filled electron shell is spherically symmetrical. Due to the repulsion that occurs due to the Pauli principle when the filled shells of the anion and cation overlap, the distance to which the ions can approach is limited. The repulsive energy changes rapidly with internuclear distance and can be written as the following two approximate expressions:
E ott = A/r n (n≈12) (3)
E ott = B∙exp(-r/ρ) (4),
where A and B are constants, r is the distance between ions, ρ is a parameter (characteristic length).
It should be noted that none of these expressions correspond to the complex quantum mechanical process that leads to repulsion.
Despite the approximate nature of these formulas, they make it possible to quite accurately calculate and accordingly describe the chemical bond in the molecules of such ionic compounds as NaCl, KCl, CaO.
Since the electric field of an ion has spherical symmetry (Fig. 13), an ionic bond, unlike a covalent bond, has no directionality. The interaction of two oppositely charged ions is compensated by repulsive forces only in the direction connecting the centers of the ion nuclei; in other directions, compensation of the electric fields of the ions does not occur. Therefore, they are able to interact with other ions. Thus, the ionic bond is not saturable.
Rice. 13. Spherical symmetry of the electrostatic field
oppositely charged charges.
Due to the non-directionality and unsaturation of ionic bonds, it is energetically most favorable when each ion is surrounded by the maximum number of ions of the opposite sign. Due to this, the most preferred form of existence of an ionic compound is a crystal. For example, in a NaCl crystal, each cation has six anions as its nearest neighbors.
Only at high temperatures in the gaseous state do ionic compounds exist in the form of unassociated molecules.
In ionic compounds, the coordination number does not depend on the specific electronic structure of the atoms, as in covalent compounds, but is determined by the ratio of the sizes of the ions. With a ratio of ionic radii in the range of 0.41 - 0.73, octahedral coordination of ions is observed, with a ratio of 0.73-1.37 - cubic coordination, etc.
Thus, under normal conditions, ionic compounds are crystalline substances. The concept of two-ionic molecules, for example, NaCL, CsCl, does not apply to them. Each crystal consists of a large number of ions.
An ionic bond can be represented as a limiting polar bond, for which the effective charge of the atom is close to unity. For a purely covalent nonpolar bond, the effective charge of the atoms is zero. In real substances, purely ionic and purely covalent bonds are rare. Most compounds have a bond character intermediate between nonpolar covalent and polar ionic. That is, in these compounds the covalent bond is partially ionic in nature. The nature of ionic and covalent bonds in real substances is presented in Figure 14.
Rice. 14. Ionic and covalent nature of the bond.
The proportion of ionic character of a bond is called the degree of ionicity. It is characterized by the effective charges of atoms in a molecule. The degree of ionicity increases with increasing difference in electronegativity of the atoms forming it.
Metal connection
In metal atoms, the outer valence electrons are held much weaker than in non-metal atoms. This causes the loss of connection between electrons and individual atoms for a sufficiently long period of time and their socialization. A socialized ensemble of external electrons is formed. The existence of such an electronic system leads to the emergence of forces that keep positive metal ions in a close state, despite their charge of the same name. This bond is called metallic. Such a bond is characteristic only of metal and exists in the solid and liquid states of the substance. A metal bond is a type of chemical bond. It is based on the socialization of external electrons, which lose their connection with the atom and are therefore called free electrons (Fig. 15).
Rice. 15. Metal connection.
The existence of a metallic bond is confirmed by the following facts. All metals have high thermal conductivity and high electrical conductivity, which is ensured by the presence of free electrons. In addition, the same circumstance determines the good reflectivity of metals to light irradiation, their luster and opacity, high ductility, and positive temperature coefficient of electrical resistance.
The stability of the crystal lattice of metals cannot be explained by such types of bonds as ionic and covalent. Ionic bonding between metal atoms located at the sites of the crystal lattice is impossible, since they have the same charge. Covalent bonding between metal atoms is also unlikely, since each atom has 8 to 12 nearest neighbors, and the formation of covalent bonds with so many shared electron pairs is unknown.
Metal structures are characterized by the fact that they have a rather rare arrangement of atoms (internuclear distances are large) and a large number of nearest neighbors for each atom in the crystal lattice. Table 1 shows three typical metal structures.
Table 1
Characteristics of the structures of the three most common metals
We see that each atom participates in the formation of a large number of bonds (for example, with 8 atoms). Such a large number of bonds (with 8 or 12 atoms) cannot be simultaneously localized in space. The connection must be carried out due to the resonance of the vibrational motion of the external electrons of each atom, as a result of which the collectivization of all external electrons of the crystal occurs with the formation of an electron gas. In many metals, to form a metallic bond, it is enough to take one electron from each atom. This is exactly what is observed for lithium, which has only one electron in its outer shell. A lithium crystal is a lattice of Li + ions (spheres with a radius of 0.068 nm) surrounded by electron gas.
Rice. 16. Various types of crystalline packing: a-hexagonal close packing; b - face-centered cubic packing; c-body-centered cubic packing.
There are similarities between metallic and covalent bonds. It lies in the fact that both types of bonds are based on the sharing of valence electrons. However, a covalent bond only connects two adjacent atoms, and the shared electrons are in close proximity to the bonded atoms. In a metallic bond, several atoms participate in sharing valence electrons.
Thus, the concept of a metallic bond is inextricably linked with the idea of metals as a collection of positively charged ionic cores with large gaps between the ions filled with electron gas, while at the macroscopic level the system remains electrically neutral.
In addition to the types of chemical bonds discussed above, there are other types of bonds that are intermolecular: hydrogen bond, van der Waals interaction, donor-acceptor interaction.
Donor-acceptor interaction of molecules
The mechanism of formation of a covalent bond due to the two-electron cloud of one atom and the free orbital of another is called donor-acceptor. An atom or particle that provides a two-electron cloud for communication is called a donor. An atom or particle with a free orbital that accepts this electron pair is called an acceptor.
Main types of intermolecular interactions. Hydrogen bond
Between valence-saturated molecules, at distances exceeding the particle size, electrostatic forces of intermolecular attraction can appear. They are called van der Waals forces. The van der Waals interaction always exists between closely spaced atoms, but plays an important role only in the absence of stronger bonding mechanisms. This weak interaction with a characteristic energy of 0.2 eV/atom occurs between neutral atoms and between molecules. The name of the interaction is associated with the name of van der Waals, since it was he who first suggested that the equation of state, taking into account the weak interaction between gas molecules, describes the properties of real gases much better than the equation of state of an ideal gas. However, the nature of this attractive force was explained only in 1930 by London. Currently, the following three types of interactions are classified as van der Waals attraction: orientational, inductive, and dispersive (London effect). The energy of van der Waals attraction is determined by the sum of orientational, inductive and dispersion interactions.
E incoming = E or + E ind + E disp (5).
Orientation interaction (or dipole-dipole interaction) occurs between polar molecules, which, when approaching, turn (orient) toward each other with opposite poles so that the potential energy of the system of molecules becomes minimal. The greater the dipole moment of molecules μ and the smaller the distance l between them, the more significant the energy of orientation interaction is:
E or = -(μ 1 μ 2) 2 / (8π 2 ∙ε 0 ∙l 6) (6),
where ε 0 is the electrical constant.
Inductive interaction is associated with the processes of polarization of molecules by surrounding dipoles. It is more significant, the higher the polarizability α of a non-polar molecule and the greater the dipole moment μ of a polar molecule
E ind = -(αμ 2)/ (8π 2 ∙ε 0 ∙l 6) (7).
The polarizability α of a nonpolar molecule is called deformational, since it is associated with the deformation of the particle, while μ characterizes the displacement of the electron cloud and nuclei relative to their previous positions.
Dispersion interaction (London effect) occurs in any molecules, regardless of their structure and polarity. Due to the instantaneous mismatch of the centers of gravity of the charges of the electron cloud and nuclei, an instantaneous dipole is formed, which induces instantaneous dipoles in other particles. The movement of instantaneous dipoles becomes consistent. As a result, neighboring particles experience mutual attraction. The energy of dispersion interaction depends on the ionization energy E I and the polarizability of molecules α
E disp = - (E I 1 ∙E I 2)∙ α 1 α 2 /(E I 1 +E I 2) l 6 (8).
The hydrogen bond is intermediate between valence and intermolecular interactions. The hydrogen bond energy is low, 8–80 kJ/mol, but higher than the van der Waals interaction energy. Hydrogen bonding is characteristic of liquids such as water, alcohols, and acids and is caused by a positively polarized hydrogen atom. Small sizes and the absence of internal electrons allow a hydrogen atom present in a liquid in any compound to enter into additional interaction with a negatively polarized atom of another or the same molecule that is not covalently bonded to it
A δ- - H δ+…. A δ- - H δ+.
That is, an association of molecules occurs. The association of molecules leads to a decrease in volatility, an increase in the boiling point and heat of evaporation, and an increase in the viscosity and dielectric constant of liquids.
Water is a particularly suitable substance for hydrogen bonding because its molecule has two hydrogen atoms and two lone pairs on the oxygen atom. This determines the high dipole moment of the molecule (μ D = 1.86 D) and the ability to form four hydrogen bonds: two as a proton donor and two as a proton acceptor
(H 2 O….N – O…H 2 O) 2 times.
It is known from experiments that with a change in molecular weight in the series of hydrogen compounds of elements of the third and subsequent periods, the boiling point increases. If this pattern is applied to water, then its boiling point should not be 100 0 C, but 280 0 C. This contradiction confirms the existence of a hydrogen bond in water.
Experiments have shown that molecular associates are formed in liquid and especially in solid water. Ice has a tetrahedral crystal lattice. In the center of the tetrahedron there is an oxygen atom of one water molecule; at the four vertices there are oxygen atoms of neighboring molecules, which are connected by hydrogen bonds to their nearest neighbors. In liquid water, hydrogen bonds are partially destroyed, and in its structure there is a dynamic equilibrium between molecular associates and free molecules.
Valence bond method
The theory of valence bonds, or localized electron pairs, posits that each pair of atoms in a molecule is held together by one or more shared electron pairs. In the valence bond theory, a chemical bond is localized between two atoms, that is, it is two-center and two-electron.
The valence bond method is based on the following basic principles:
Each pair of atoms in a molecule is held together by one or more shared electron pairs;
A single covalent bond is formed by two electrons with antiparallel spins located on the valence orbitals of the bonding atoms;
When a bond is formed, the wave functions of electrons overlap, leading to an increase in the electron density between atoms and a decrease in the total energy of the system;
“Chemical bond” is the energy of destruction of the lattice into ions _Ekul = Uresh. Basic principles of the MO method. Types of overlap of atomic AOs. bonding and antibonding MOs with a combination of atomic orbitals s and s pz and pz px and px. H?C? C?H. ? - Repulsion coefficient. Qeff =. Ao. Basic theories of chemical bonding.
“Types of chemical bonds” - Substances with ionic bonds form an ionic crystal lattice. Atoms. Electronegativity. Municipal Educational Institution Lyceum No. 18 chemistry teacher Kalinina L.A. Ions. For example: Na1+ and Cl1-, Li1+ and F1- Na1+ + Cl1- = Na(:Cl:) . If e - are added, the ion becomes negatively charged. The atomic frame has high strength.
“The Life of Mendeleev” - July 18 D.I. Mendeleev graduated from the Tobolsk gymnasium. August 9, 1850 - June 20, 1855 while studying at the Main Pedagogical Institute. “If you do not know names, then the knowledge of things will die” K. Liney. Life and work of D.I. Mendeleev. Ivan Pavlovich Mendeleev (1783 - 1847), father of the scientist. Discovery of the periodic law.
“Types of chemical bonds” - H3N. Al2O3. The structure of matter." H2S. MgO. H2. Cu. Mg S.CS2. I. Write down the formulas of the substances: 1.c.N.S. 2.s K.P.S. 3. with I.S. K.N.S. NaF. C.K.P.S. Determine the type of chemical bond. Which of the molecules corresponds to the scheme: A A?
"Mendeleev" - Dobereiner's Triads of Elements. Gases. Work. Life and scientific feat. Periodic table of elements (long form). Newlands' "Law of Octaves" Scientific activity. Solutions. A new stage of life. The second version of Mendeleev's system of elements. Part of L. Meyer's table of elements. Discovery of the periodic law (1869).
“The Life and Work of Mendeleev” - Ivan Pavlovich Mendeleev (1783 - 1847), the scientist’s father. 1834, January 27 (February 6) - D.I. Mendeleev was born in the city of Tobolsk, in Siberia. 1907, January 20 (February 2) D.I. Mendeleev died of heart paralysis. DI. Menedeleev (South Kazakhstan region, Shymkent city). Industry. On July 18, 1849, D.I. Mendeleev graduated from the Tobolsk gymnasium.
Chemical bond.
Exercises.
1. Determine the type of chemical bond in the following substances:
Substance | Phosphorus chloride | Sulfuric acid | |||
Communication type | |||||
Substance | Barium oxide | ||||
Communication type |
2. Emphasize substances in which BETWEEN molecules exists hydrogen bond:
sulphur dioxide; ice; ozone; ethanol; ethylene; acetic acid; hydrogen fluoride.
3. How do they affect bond length, strength and polarity- atomic radii, their electronegativity, bond multiplicity?
A) The larger the radii atoms that form a bond, so link length _______
b) The higher the multiplicity (single, double or triple) bonds, so its strength ____________________
V) The greater the electronegativity difference between two atoms, the polarity of the bond ____________
4. Compare length, strength and polarity of bonds in molecules:
a) bond length: HCl ___HBr
b) bond strength PH3_______NH3
c) polarity of the CCl4 bond ______CH4
d) bond strength: N2 _______O2
e) bond length between carbon atoms in ethylene and acetylene: __________
f) polarity of bonds in NH3_________H2O
Tests. A4. Chemical bond.
1. The valence of an atom is
1) the number of chemical bonds formed by a given atom in a compound
2) oxidation state of the atom
3) the number of electrons given or received
4) the number of electrons missing to obtain the electron configuration of the nearest inert gas
A. When a chemical bond is formed, energy is always released
B. The energy of a double bond is less than that of a single bond.
1) only A is true 2) only B is true 3) both judgments are correct 4) both judgments are incorrect
3. In substances formed by combining identical atoms, chemical bond
1) ionic 2) covalent polar 3) hydrogen 4) covalent nonpolar
4. Compounds with a covalent polar and covalent nonpolar bond are respectively
1) water and hydrogen sulfide 2) potassium bromide and nitrogen
5. Due to the shared electron pair, a chemical bond is formed in the compound
1) KI 2) HBr 3) Li2O 4) NaBr
6. Select a pair of substances in which all bonds are covalent:
1) NaCl, HCl 2) CO2, BaO 3) CH3Cl, CH3Na 4) SO2, NO2
7. A substance with a polar covalent bond has the formula
1)KCl 2)HBr 3)P4 4)CaCl2
8. Compound with an ionic chemical bond
1) phosphorus chloride 2) potassium bromide 3) nitrogen oxide (II) 4) barium
9. In ammonia and barium chloride, the chemical bond is respectively
1) ionic and covalent polar 2) covalent nonpolar and ionic 3) covalent polar and ionic 4) covalent nonpolar and metallic
10. Substances with a covalent polar bond are
1) sulfur oxide (IV) 2) oxygen 3) calcium hydride 4) diamond
11. Which series lists substances with only polar covalent bonds:
1) CH4 H2 Cl2 2) NH3 HBr CO2 3) PCl3 KCl CCl4 4) H2S SO2 LiF
12. Which series lists substances with only ionic bonds:
1) F2O LiF SF4 2) PCl3 NaCl CO2 3) KF Li2O BaCl2 4) CaF2 CH4 CCl4
13. A compound with an ionic bond is formed when interacting
1) CH4 and O2 2) NH3 and HCl 3) C2H6 and HNO3 4) SO3 and H2O
14. In which substance are all chemical bonds covalent nonpolar?
1) Diamond 2) Carbon monoxide (IV) 3) Gold 4) Methane
15. The connection formed between elements with serial numbers 15 and 53
1) ionic 2) metal
3) covalent non-polar 4) covalent polar
16. Hydrogen bond is formed between molecules
1) ethane 2) benzene 3) hydrogen 4) ethanol
17. What substance contains hydrogen bonds?
1) Hydrogen sulfide 2) Ice 3) Hydrogen bromide 4) Benzene
18.Which substance contains both ionic and covalent chemical bonds?
1) Sodium chloride 2) Hydrogen chloride 3) Sodium sulfate 4) Phosphoric acid
19. The chemical bond in the molecule has a more pronounced ionic character
1) lithium bromide 2) copper chloride 3) calcium carbide 4) potassium fluoride
20. Three common electron pairs form a covalent bond in the molecule of 1) nitrogen 2) hydrogen sulfide 3) methane 4) chlorine
21. How many electrons are involved in the formation of chemical bonds in a water molecule?4) 18
22. The molecule contains four covalent bonds: 1) CO2 2) C2H4 3) P4 4) C3H4
23. The number of bonds in molecules increases in a series
1) CHCl3, CH4 2) CH4, SO3 3) CO2, CH4 4) SO2, NH3
24. In what compound is a covalent bond formed between atoms? by donor-acceptor mechanism? 1) KCl 2) CCl4 3) NH4Cl 4) CaCl2
25. Which of the following molecules requires the least amount of energy to decompose into atoms? 1) HI 2) H2 3) O2 4) CO
26. Indicate the molecule in which the binding energy is the highest:
1) N≡N 2) H-H 3) O=O 4) H-F
27. Indicate the molecule in which the chemical bond is the strongest:
1) HF 2) HCl 3) HBr 4) HI
28. Indicate a series characterized by an increase in the length of a chemical bond
1)O2, N2, F2, Cl2 2)N2, O2, F2, Cl2 3)F2, N2, O2, Cl2 4)N2, O2, Cl2, F2
29. The length of the E-O bond increases in the series
1) silicon oxide (IV), carbon oxide (IV)
2) sulfur(IV) oxide, tellurium(IV) oxide
3) strontium oxide, beryllium oxide
4) sulfur oxide(IV), carbon monoxide(IV)
30. In the series CH4 – SiH4 occurs increase
1) bond strength 2) oxidative properties
3) bond lengths 4) bond polarities
31. In what row are the molecules arranged in order of increasing polarity of bonds?
1) HF, HCl, HBr 2) H2Se, H2S, H2O 3) NH3, PH3, AsH3 4) CO2, CS2, CSe2
32. The most polar covalent bond in a molecule is:
1) CH4 2) CF4 3) CCl4 4) CBr4
33.Indicate the series in which the polarity increases:
1)AgF, F2, HF 2)Cl2, HCl, NaCl 3)CuO, CO, O2 4) KBr, NaCl, KF
Covalent chemical bond, its varieties and mechanisms of formation. Characteristics of covalent bonds (polarity and bond energy). Ionic bond. Metal connection. Hydrogen bond.
1. In ammonia and barium chloride, the chemical bond is respectively
1) ionic and covalent polar
2) covalent polar and ionic
3) covalent nonpolar and metallic
4) covalent nonpolar and ionic
2. Substances with only ionic bonds are listed in the following series:
1) F2, CCl4, KS1
2) NaBr, Na2O, KI
3. A compound with an ionic bond is formed by interaction
3) C2H6 and HNO3
4. In which series do all substances have a polar covalent bond?
1) HCl, NaCl. Cl2
4) NaBr. HBr. CO
5. In which series are the formulas of substances with only covalent polar
1) C12, NO2, HC1
6. Covalent nonpolar bond is characteristic of
1) C12 2) SO3 3) CO 4) SiO2
7. A substance with a polar covalent bond is
1) C12 2) NaBr 3) H2S 4) MgCl2
8. A substance with a covalent bond is
1) CaC12 2) MgS 3) H2S 4) NaBr
9. A substance with a covalent nonpolar bond has the formula
1) NH3 2) Cu 3) H2S 4) I2
10. Substances with non-polar covalent bonds are
1) water and diamond
2) hydrogen and chlorine
3) copper and nitrogen
4) bromine and methane
11. A chemical bond is formed between atoms with the same relative electronegativity
2) covalent polar
3) covalent nonpolar
4) hydrogen
12. Covalent polar bonds are characteristic of
1) KC1 2) HBr 3) P4 4) CaCl2
13. A chemical element in the atom of which the electrons are distributed among the layers as follows: 2, 8, 8, 2 forms a chemical bond with hydrogen
1) covalent polar
2) covalent nonpolar
4) metal
14. In the molecule of which substance does the bond between carbon atoms have the longest length?
1) acetylene 2) ethane 3) ethene 4) benzene
15. Three common electron pairs form a covalent bond in a molecule
2) hydrogen sulfide
16. Hydrogen bonds form between molecules
1) dimethyl ether
2) methanol
3) ethylene
4) ethyl acetate
17. Bond polarity is most pronounced in the molecule
1) HI 2) HC1 3) HF 4) NVg
18. Substances with non-polar covalent bonds are
1) water and diamond
2) hydrogen and chlorine
3) copper and nitrogen
4) bromine and methane
19. Hydrogen bonding is not typical for the substance
1) H2O 2) CH4 3) NH3 4) CH3OH
20. A covalent polar bond is characteristic of each of the two substances whose formulas are
2) CO2 and K2O
4) CS2 and RS15
21. The weakest chemical bond in a molecule
1) fluorine 2) chlorine 3) bromine 4) iodine
22. Which substance has the longest chemical bond in its molecule?
1) fluorine 2) chlorine 3) bromine 4) iodine
23. Each of the substances indicated in the series has covalent bonds:
1) C4H10, NO2, NaCl
2) CO, CuO, CH3Cl
4) C6H5NO2, F2, CC14
24. Each of the substances indicated in the series has a covalent bond:
1) CaO, C3H6, S8
2) Fe. NaNO3,CO
3) N2, CuCO3, K2S
4) C6H5N02, SO2, CHC13
25. Each of the substances indicated in the series has a covalent bond:
1) C3H4, NO, Na2O
2) CO, CH3C1, PBr3
3) Р2Оз, NaHSO4, Cu
4) C6H5NO2, NaF, CC14
26. Each of the substances indicated in the series has covalent bonds:
1) C3Ha, NO2, NaF
2) KS1, CH3Cl, C6H12O6
3) P2O5, NaHSO4, Ba
4) C2H5NH2, P4, CH3OH
27. Bond polarity is most pronounced in molecules
1) hydrogen sulfide
3) phosphine
4) hydrogen chloride
28. In the molecule of which substance are the chemical bonds the strongest?
29. Among the substances NH4Cl, CsCl, NaNO3, PH3, HNO3 - the number of compounds with ionic bonds is equal
30. Among the substances (NH4)2SO4, Na2SO4, CaI2, I2, CO2 - the number of compounds with a covalent bond is equal to
Answers: 1-2, 2-2, 3-4, 4-3, 5-4, 6-1, 7-3, 8-3, 9-4, 10-2, 11-3, 12-2, 13-3, 14-2, 15-1, 16-2, 17-3, 18-2, 19-2, 20-4, 21-4, 22-4, 23-4, 24-4, 25- 2, 26-4, 27-4, 28-1, 29-3, 30-4