Single connection. Covalent bond Which compounds contain a double covalent bond
simple bond (ordinary bond, single bond)- a chemical covalent bond carried out by a pair of electrons moving in the field of two atomic nuclei. In chemical compounds in which the average number of electrons binding each pair of atomic nuclei is not equal to two, ... ...
Connection- : See also: chemical bond metallic bond ionic bond covalent bond ... encyclopedic Dictionary in metallurgy
Mutual attraction of atoms, leading to the formation of molecules and crystals. It is customary to say that in a molecule or in a crystal between neighboring atoms there are ch. The valence of an atom (which is discussed in more detail below) indicates the number of bonds ...
metallic bond- interatomic bond, characteristic of metals with a uniform density of the electron gas. The metallic bond is due to the interaction of a negatively charged electron gas and positively charged ionic cores, ... ... Encyclopedic Dictionary of Metallurgy
covalent bond- interatomic bond due to the collectivization of the outer electrons of the interacting atoms. Covalent bonds are characterized by saturation and directionality. Saturation is manifested in the fact that such a covalent bond enters ... ... Encyclopedic Dictionary of Metallurgy
ionic bond- electro, heterovalent bond one of the types chemical bond, which is based on the electrostatic interaction between oppositely charged ions. Such bonds in a relatively pure form are formed in halides ... ... Encyclopedic Dictionary of Metallurgy
chemical bond- mutual attraction of atoms, leading to the formation of molecules and crystals. The valence of an atom shows the number of bonds formed by a given atom with neighboring ones. The term "chemical structure" was introduced by Academician A. M. Butlerov in ... ... Encyclopedic Dictionary of Metallurgy
Ordinary bond, single bond, chemical covalent bond carried out by a pair of electrons (with antiparallel spin orientation) moving in the field of 2 atomic nuclei. For example, in the molecules of H2, Cl2 and HCl there is one covalent ... ... Great Soviet Encyclopedia
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Each atom has a certain number of electrons.
Entering into chemical reactions, atoms donate, acquire, or socialize electrons, reaching the most stable electronic configuration. The configuration with the lowest energy is the most stable (as in noble gas atoms). This pattern is called the "octet rule" (Fig. 1).
Rice. one.
This rule applies to all connection types. Electronic bonds between atoms allow them to form stable structures, from the simplest crystals to complex biomolecules that eventually form living systems. They differ from crystals in their continuous metabolism. However, many chemical reactions proceed according to the mechanisms electronic transfer, which play an important role in the energy processes in the body.
A chemical bond is a force that holds together two or more atoms, ions, molecules, or any combination of them..
The nature of the chemical bond is universal: it is an electrostatic force of attraction between negatively charged electrons and positively charged nuclei, determined by the configuration of the electrons in the outer shell of atoms. The ability of an atom to form chemical bonds is called valency, or oxidation state. The concept of valence electrons- electrons that form chemical bonds, that is, those located in the most high-energy orbitals. Accordingly, the outer shell of an atom containing these orbitals is called valence shell. At present, it is not enough to indicate the presence of a chemical bond, but it is necessary to clarify its type: ionic, covalent, dipole-dipole, metallic.
The first type of connection isionic connection
According to Lewis and Kossel's electronic theory of valency, atoms can achieve a stable electronic configuration in two ways: first, by losing electrons, becoming cations, secondly, acquiring them, turning into anions. As a result of electron transfer, due to the electrostatic force of attraction between ions with charges of the opposite sign, a chemical bond is formed, called Kossel " electrovalent(now called ionic).
In this case, anions and cations form a stable electronic configuration with a filled outer electron shell. Typical ionic bonds are formed from cations of T and II groups of the periodic system and anions of non-metallic elements of groups VI and VII (16 and 17 subgroups - respectively, chalcogens and halogens). The bonds in ionic compounds are unsaturated and non-directional, so they retain the possibility of electrostatic interaction with other ions. On fig. 2 and 3 show examples of ionic bonds corresponding to the Kossel electron transfer model.
Rice. 2.
Rice. 3. Ionic bond in the sodium chloride (NaCl) molecule
Here it is appropriate to recall some of the properties that explain the behavior of substances in nature, in particular, to consider the concept of acids and grounds.
Aqueous solutions of all these substances are electrolytes. They change color in different ways. indicators. The mechanism of action of indicators was discovered by F.V. Ostwald. He showed that the indicators are weak acids or bases, the color of which in the undissociated and dissociated states is different.
Bases can neutralize acids. Not all bases are soluble in water (for example, some organic compounds that do not contain -OH groups are insoluble, in particular, triethylamine N (C 2 H 5) 3); soluble bases are called alkalis.
Aqueous solutions of acids enter into characteristic reactions:
a) with metal oxides - with the formation of salt and water;
b) with metals - with the formation of salt and hydrogen;
c) with carbonates - with the formation of salt, CO 2 and H 2 O.
The properties of acids and bases are described by several theories. In accordance with the theory of S.A. Arrhenius, an acid is a substance that dissociates to form ions H+ , while the base forms ions HE- . This theory does not take into account the existence of organic bases that do not have hydroxyl groups.
In line with proton Bronsted and Lowry's theory, an acid is a substance containing molecules or ions that donate protons ( donors protons), and the base is a substance consisting of molecules or ions that accept protons ( acceptors protons). Note that in aqueous solutions, hydrogen ions exist in a hydrated form, that is, in the form of hydronium ions H3O+ . This theory describes reactions not only with water and hydroxide ions, but also carried out in the absence of a solvent or with a non-aqueous solvent.
For example, in the reaction between ammonia NH 3 (weak base) and hydrogen chloride in the gas phase, solid ammonium chloride is formed, and in an equilibrium mixture of two substances there are always 4 particles, two of which are acids, and the other two are bases:
This equilibrium mixture consists of two conjugated pairs of acids and bases:
1)NH 4+ and NH 3
2) HCl and Cl ‑
Here, in each conjugated pair, the acid and base differ by one proton. Every acid has a conjugate base. A strong acid has a weak conjugate base, and a weak acid has a strong conjugate base.
The Bronsted-Lowry theory makes it possible to explain the unique role of water for the life of the biosphere. Water, depending on the substance interacting with it, can exhibit the properties of either an acid or a base. For example, in reactions with aqueous solutions of acetic acid, water is a base, and with aqueous solutions of ammonia, it is an acid.
1) CH 3 COOH + H 2 O ↔ H 3 O + + CH 3 SOO- . Here the acetic acid molecule donates a proton to the water molecule;
2) NH3 + H 2 O ↔ NH4 + + HE- . Here the ammonia molecule accepts a proton from the water molecule.
Thus, water can form two conjugated pairs:
1) H 2 O(acid) and HE- (conjugate base)
2) H 3 O+ (acid) and H 2 O(conjugate base).
In the first case, water donates a proton, and in the second, it accepts it.
Such a property is called amphiprotonity. Substances that can react as both acids and bases are called amphoteric. Such substances are often found in nature. For example, amino acids can form salts with both acids and bases. Therefore, peptides readily form coordination compounds with the metal ions present.
In this way, characteristic property ionic bond - the complete movement of a bunch of binding electrons to one of the nuclei. This means that there is a region between the ions where the electron density is almost zero.
The second type of connection iscovalent connection
Atoms can form stable electronic configurations by sharing electrons.
Such a bond is formed when a pair of electrons is shared one at a time. from each atom. In this case, the socialized bond electrons are distributed equally among the atoms. An example of a covalent bond is homonuclear diatomic H molecules 2 , N 2 , F 2. Allotropes have the same type of bond. O 2 and ozone O 3 and for a polyatomic molecule S 8 and also heteronuclear molecules hydrogen chloride Hcl, carbon dioxide CO 2, methane CH 4, ethanol FROM 2 H 5 HE, sulfur hexafluoride SF 6, acetylene FROM 2 H 2. All these molecules have the same common electrons, and their bonds are saturated and directed in the same way (Fig. 4).
For biologists, it is important that the covalent radii of atoms in double and triple bonds are reduced compared to a single bond.
Rice. four. Covalent bond in the Cl 2 molecule.
Ionic and covalent types of bonds are two limiting cases of many existing types of chemical bonds, and in practice most of the bonds are intermediate.
Connections of two elements located at opposite ends of one or different periods systems of Mendeleev, predominantly form ionic bonds. As the elements approach each other within a period, the ionic nature of their compounds decreases, while the covalent character increases. For example, the halides and oxides of the elements on the left side periodic table form predominantly ionic bonds ( NaCl, AgBr, BaSO 4 , CaCO 3 , KNO 3 , CaO, NaOH), and the same compounds of the elements on the right side of the table are covalent ( H 2 O, CO 2, NH 3, NO 2, CH 4, phenol C6H5OH, glucose C 6 H 12 O 6, ethanol C 2 H 5 OH).
The covalent bond, in turn, has another modification.
For polyatomic ions and in complex biological molecules both electrons can only come from one atom. It is called donor electron pair. An atom that socializes this pair of electrons with a donor is called acceptor electron pair. This type of covalent bond is called coordination (donor-acceptor, ordative) communication(Fig. 5). This type of bond is most important for biology and medicine, since the chemistry of the most important for the metabolism of d-elements in to a large extent described by coordination links.
Pic. 5.
As a rule, in a complex compound, a metal atom acts as an electron pair acceptor; on the contrary, in ionic and covalent bonds, the metal atom is an electron donor.
The essence of the covalent bond and its variety - the coordination bond - can be clarified with the help of another theory of acids and bases, proposed by GN. Lewis. He somewhat expanded the semantic concept of the terms "acid" and "base" according to the Bronsted-Lowry theory. The Lewis theory explains the nature of the formation of complex ions and the participation of substances in nucleophilic substitution reactions, that is, in the formation of CS.
According to Lewis, an acid is a substance capable of forming a covalent bond by accepting an electron pair from a base. A Lewis base is a substance that has a lone pair of electrons, which, by donating electrons, forms a covalent bond with Lewis acid.
That is, the Lewis theory expands the range of acid-base reactions also to reactions in which protons do not participate at all. Moreover, the proton itself, according to this theory, is also an acid, since it is able to accept an electron pair.
Therefore, according to this theory, cations are Lewis acids and anions are Lewis bases. The following reactions are examples:
It was noted above that the subdivision of substances into ionic and covalent ones is relative, since there is no complete transfer of an electron from metal atoms to acceptor atoms in covalent molecules. In compounds with an ionic bond, each ion is in the electric field of ions of the opposite sign, so they are mutually polarized, and their shells are deformed.
Polarizability determined electronic structure, charge and size of the ion; it is higher for anions than for cations. The highest polarizability among cations is for cations more charge and smaller, such as Hg 2+ , Cd 2+ , Pb 2+ , Al 3+ , Tl 3+. Has a strong polarizing effect H+ . Since the effect of ion polarization is two-sided, it significantly changes the properties of the compounds they form.
The third type of connection -dipole-dipole connection
In addition to the listed types of communication, there are also dipole-dipole intermolecular interactions, also known as van der Waals .
The strength of these interactions depends on the nature of the molecules.
There are three types of interactions: permanent dipole - permanent dipole ( dipole-dipole attraction); permanent dipole - induced dipole ( induction attraction); instantaneous dipole - induced dipole ( dispersion attraction, or London forces; rice. 6).
Rice. 6.
Only molecules with polar covalent bonds have a dipole-dipole moment ( HCl, NH 3, SO 2, H 2 O, C 6 H 5 Cl), and the bond strength is 1-2 debye(1D \u003d 3.338 × 10 -30 coulomb meters - C × m).
In biochemistry, another type of bond is distinguished - hydrogen connection, which is a limiting case dipole-dipole attraction. This bond is formed by the attraction between a hydrogen atom and a small electronegative atom, most often oxygen, fluorine and nitrogen. With large atoms that have a similar electronegativity (for example, with chlorine and sulfur), the hydrogen bond is much weaker. The hydrogen atom is distinguished by one essential feature: when the binding electrons are pulled away, its nucleus - the proton - is exposed and ceases to be screened by electrons.
Therefore, the atom turns into a large dipole.
A hydrogen bond, unlike a van der Waals bond, is formed not only during intermolecular interactions, but also within one molecule - intramolecular hydrogen bond. Hydrogen bonds play in biochemistry important role, for example, to stabilize the structure of proteins in the form of an a-helix, or to form a DNA double helix (Fig. 7).
Fig.7.
Hydrogen and van der Waals bonds are much weaker than ionic, covalent, and coordination bonds. The energy of intermolecular bonds is indicated in Table. one.
Table 1. Energy of intermolecular forces
Note: The degree of intermolecular interactions reflect the enthalpy of melting and evaporation (boiling). Ionic compounds require much more energy to separate ions than to separate molecules. The melting enthalpies of ionic compounds are much higher than those of molecular compounds.
The fourth type of connection -metallic bond
Finally, there is another type of intermolecular bonds - metal: connection of positive ions of the lattice of metals with free electrons. This type of connection does not occur in biological objects.
From overview types of bonds, one detail is clarified: an important parameter of an atom or ion of a metal - an electron donor, as well as an atom - an electron acceptor is its the size.
Without going into details, we note that the covalent radii of atoms, the ionic radii of metals, and the van der Waals radii of interacting molecules increase as their atomic number in the groups of the periodic system increases. In this case, the values of the ion radii are the smallest, and the van der Waals radii are the largest. As a rule, when moving down the group, the radii of all elements increase, both covalent and van der Waals.
The most important for biologists and physicians are coordination(donor-acceptor) bonds considered by coordination chemistry.
Medical bioinorganics. G.K. Barashkov
Atoms of most elements do not exist separately, as they can interact with each other. In this interaction, more complex particles are formed.
The nature of the chemical bond is the action of electrostatic forces, which are the forces of interaction between electric charges. Electrons and atomic nuclei have such charges.
Electrons located at the outer electronic levels (valence electrons), being farthest from the nucleus, interact with it the weakest, and therefore are able to break away from the nucleus. They are responsible for the binding of atoms to each other.
Types of interaction in chemistry
The types of chemical bond can be represented as the following table:
Ionic bond characteristic
The chemical interaction that is formed due to ion attraction having different charges is called ionic. This happens if the bonded atoms have a significant difference in electronegativity (that is, the ability to attract electrons) and the electron pair goes to a more electronegative element. The result of such a transition of electrons from one atom to another is the formation of charged particles - ions. There is an attraction between them.
have the lowest electronegativity typical metals, and the largest are typical non-metals. Ions are thus formed by interactions between typical metals and typical non-metals.
Metal atoms become positively charged ions (cations), donating electrons to external electronic levels, and non-metals accept electrons, thus turning into negatively charged ions (anions).
Atoms move into a more stable energy state, completing their electronic configurations.
The ionic bond is non-directional and not saturable, since the electrostatic interaction occurs in all directions, respectively, the ion can attract ions of the opposite sign in all directions.
The arrangement of ions is such that around each is a certain number of oppositely charged ions. The concept of "molecule" for ionic compounds doesn't make sense.
Examples of Education
The formation of a bond in sodium chloride (nacl) is due to the transfer of an electron from the Na atom to the Cl atom with the formation of the corresponding ions:
Na 0 - 1 e \u003d Na + (cation)
Cl 0 + 1 e \u003d Cl - (anion)
In sodium chloride, there are six chloride anions around the sodium cations, and six sodium ions around each chloride ion.
When an interaction is formed between atoms in barium sulfide, the following processes occur:
Ba 0 - 2 e \u003d Ba 2+
S 0 + 2 e \u003d S 2-
Ba donates its two electrons to sulfur, resulting in the formation of sulfur anions S 2- and barium cations Ba 2+ .
metal chemical bond
The number of electrons in the outer energy levels of metals is small; they easily break away from the nucleus. As a result of this detachment, metal ions and free electrons are formed. These electrons are called "electron gas". Electrons move freely throughout the volume of the metal and are constantly bound and detached from atoms.
The structure of the metal substance is as follows: the crystal lattice is the backbone of the substance, and electrons can move freely between its nodes.
The following examples can be given:
Mg - 2e<->Mg2+
Cs-e<->Cs+
Ca-2e<->Ca2+
Fe-3e<->Fe3+
Covalent: polar and non-polar
The most common type of chemical interaction is a covalent bond. The electronegativity values of the interacting elements do not differ sharply, in connection with this, only a shift of the common electron pair to a more electronegative atom occurs.
Covalent interaction can be formed by the exchange mechanism or by the donor-acceptor mechanism.
The exchange mechanism is realized if each of the atoms has unpaired electrons in the outer electronic levels and the overlap of atomic orbitals leads to the appearance of a pair of electrons that already belongs to both atoms. When one of the atoms has a pair of electrons at the outer electronic level, and the other has a free orbital, then when the atomic orbitals overlap, the electron pair is socialized and the interaction occurs according to the donor-acceptor mechanism.
Covalent are divided by multiplicity into:
- simple or single;
- double;
- triple.
Doubles provide the socialization of two pairs of electrons at once, and triples - three.
According to the distribution of electron density (polarity) between the bonded atoms, the covalent bond is divided into:
- non-polar;
- polar.
A non-polar bond is formed by the same atoms, and a polar bond is formed by electronegativity different.
The interaction of atoms with similar electronegativity is called a non-polar bond. The common pair of electrons in such a molecule is not attracted to any of the atoms, but belongs equally to both.
The interaction of elements differing in electronegativity leads to the formation of polar bonds. Common electron pairs with this type of interaction are attracted by a more electronegative element, but do not completely transfer to it (that is, the formation of ions does not occur). As a result of such a shift in the electron density, partial charges appear on atoms: on a more electronegative one, a negative charge, and on a less electronegative one, a positive one.
Properties and characteristics of covalence
The main characteristics of a covalent bond:
- The length is determined by the distance between the nuclei of the interacting atoms.
- Polarity is determined by the displacement of the electron cloud to one of the atoms.
- Orientation - the property to form space-oriented bonds and, accordingly, molecules that have certain geometric shapes.
- Saturation is determined by the ability to form a limited number of bonds.
- Polarizability is determined by the ability to change polarity under the influence of an external electric field.
- The energy required to break a bond, which determines its strength.
Molecules of hydrogen (H2), chlorine (Cl2), oxygen (O2), nitrogen (N2) and many others can be an example of a covalent non-polar interaction.
H+ + H → H-H molecule has a single non-polar bond,
O: + :O → O=O the molecule has a double nonpolar,
Ṅ: + Ṅ: → N≡N the molecule has a triple non-polar.
As an example of a covalent bond chemical elements you can bring molecules of carbon dioxide (CO2) and carbon monoxide (CO) gas, hydrogen sulfide (H2S), of hydrochloric acid(HCL), water (H2O), methane (CH4), sulfur oxide (SO2) and many others.
In the CO2 molecule, the relationship between carbon and oxygen atoms is covalent polar, since the more electronegative hydrogen attracts electron density to itself. oxygen has two unpaired electron at the outer level, and carbon can provide four valence electrons to form the interaction. As a result, double bonds are formed and the molecule looks like this: O=C=O.
In order to determine the type of bond in a particular molecule, it is enough to consider its constituent atoms. Simple substances metals form metallic, metals with non-metals - ionic, simple substances non-metals - covalent non-polar, and molecules consisting of different non-metals are formed through a covalent polar bond.
chemical bond- these are the interactions of electrons and the atomic nucleus of one particle (atom, ion, molecule, etc.) with electrons and the atomic nucleus of another particle, holding these particles in a stable or metastable chemical compound. The modern description of the chemical bond is carried out on the basis of quantum mechanics. The main characteristics of a chemical bond are strength, length, polarity.
Communication types
- Single electron chemical bond
- metal connection
- covalent bond
- Ionic bond
- Van der Waals connection
- hydrogen bond
- Two-electron three-center chemical bond
The simplest one-electron covalent chemical bond
The simplest one-electron chemical bond is created by a single valence electron. It turns out that one electron is able to hold two positively charged ions in a single whole. In a one-electron bond, the Coulomb repulsive forces of positively charged particles are compensated by the Coulomb forces of attraction of these particles to a negatively charged electron. The valence electron becomes common to the two nuclei of the molecule.
Examples such chemical compounds are molecular ions: H 2+, Li 2+, Na 2+, K 2+, Rb 2+, Cs 2+
Single covalent bond
A single covalent chemical bond is created by a bonding electron pair. In all existing theories(the theory of valence bonds, the theory of molecular orbitals, the theory of repulsion of valence electron pairs, the Bohr model of chemical bonding) the bonding electron pair is located in the space between the atoms of the molecule. Distinguish between polar and non-polar covalent bonds.
A nonpolar covalent bond takes place in homonuclear diatomic molecules in which the bonding electronI pair is equidistant from both nuclei of the molecular system.
Distance d between atomic nuclei can be viewed as the sum of the covalent radii of the corresponding atoms.
The distance between atomic nuclei in a single two-electron covalent bond is shorter than the same distance in the simplest one-electron chemical bond.
Multiple covalent bonds
Multiple covalent bonds are represented by unsaturated organic compounds containing double and triple chemical bonds. To describe the nature of unsaturated compounds, L. Pauling introduces the concepts of sigma- and π-bonds, hybridization of atomic orbitals.
Pauling's hybridization for two S- and two p-electrons made it possible to explain the directionality of chemical bonds, in particular, the tetrahedral configuration of methane. To explain the structure of ethylene, it is necessary to isolate one p-electron from four equivalent Sp3 electrons of the carbon atom to form an additional bond, called the π-bond. In this case, the three remaining Sp2 hybrid orbitals are located in the plane at an angle of 120° and form the main bonds, for example, a planar ethylene molecule.
In the case of the acetylene molecule, only one S- and one p-orbitals take part in hybridization (according to Pauling), and two Sp-orbitals are formed, located at an angle of 180 ° and directed to opposite sides. Two "clean" p-orbitals of carbon atoms overlap in pairs in mutually perpendicular planes, forming two π-bonds of a linear acetylene molecule.
The views of L. Pauling are reflected in his book “The nature of the chemical bond, which for many years has become table book chemist. In 1954, L. Pauling was awarded Nobel Prize in Chemistry with the wording "For the study of the nature of the chemical bond and its application to the determination of the structure of complex compounds."
However, the physical meaning of the selective hybridization of atomic orbitals remained unclear; hybridization was an algebraic transformation to which physical reality could not be attributed.
Linus Pauling made an attempt to improve the description of the chemical bond by eliminating the selectivity of hybridization of orbitals in the molecules of unsaturated compounds and creating the theory of a bent chemical bond. In his report at a symposium on theoretical organic chemistry dedicated to the memory of Kekule (London, September 1958), L. Pauling proposed a new way to describe a double bond as a combination of two identical bent chemical bonds, and a triple bond - three bent chemical bonds. On this
Symposium L. Pauling categorically stated:
There may be chemists who think that the most important innovation... was the description of the σ,π description for the double or triple bond and conjugated systems instead of the description by means of bent bonds. I maintain that the σ,π description is less satisfactory than the curved link description, that this innovation is only transitory and will soon die out.
AT new theory Pauling, all binding electrons became equal and equidistant from the line connecting the nuclei of the molecule. Pauling's bent chemical bond theory allowed for a statistical interpretation wave function M. Born, Coulomb electron correlation of electrons. A physical meaning appeared - the nature of the chemical bond is completely determined by the electrical interaction of nuclei and electrons. The more bonding electrons, the smaller the internuclear distance and the stronger the chemical bond between carbon atoms.
Three-center chemical bond
Further development of ideas about the chemical bond was given by the American physical chemist W. Lipscomb, who developed the theory of two-electron three-center bonds and a topological theory that makes it possible to predict the structure of some more boron hydrides (borohydrides).
An electron pair in a three-center chemical bond becomes common to three atomic nuclei. In the simplest representative of a three-center chemical bond - the molecular hydrogen ion H3 +, an electron pair holds three protons in a single whole.
There are four single covalent bonds in the diborane molecule. B-H connections and two two-electron three-center bonds. The internuclear distance in a single covalent B-H bond is 1.19 Å, while the similar distance in a three-center B-H-B bond is 1.31 Å. The angle of the three-center bond B-H-B (φ) is 830. The combination of two three-center bonds in the diborane molecule makes it possible to keep the nuclei of boron atoms at a distance dB-B = 2 1.31 sin φ/2 = 1.736 Å. The nuclei of the binding hydrogen atoms are located at a distance h = 1.31 · cos φ/2 = 0.981 Å from the plane in which four single B-H covalent bonds are located.
Three-center bonds can be realized not only in a triangle of two boron atoms and one hydrogen atom, but also between three boron atoms, for example, in framework borohydrides (pentaborane - B 5 H 9, decaborane - B 10 H 4, etc.). These structures contain ordinary (terminal) and three-center bond (bridge) hydrogen atoms and triangles of boron atoms.
The existence of boranes with their two-electron three-center bonds with "bridge" hydrogen atoms violated the canonical doctrine of valency. The hydrogen atom, previously considered a standard univalent element, turned out to be bound by identical bonds with two boron atoms and became formally a divalent element. The work of W. Lipscomb on deciphering the structure of boranes expanded the understanding of the chemical bond. The Nobel Committee awarded the William Nunn Lipscomb Prize in Chemistry in 1976 with the wording "For his studies of the structure of boranes (borohydrites) which elucidate the problems of chemical bonds".
Multicenter chemical bond
In 1951, T. Keely and P. Pawson unexpectedly obtained a completely new organo-iron compound during the synthesis of dicyclopentadienyl. The preparation of a previously unknown, extremely stable yellow-orange crystalline iron compound immediately attracted attention.
E. Fisher and D. Wilkinson independently established the structure of the new compound - two cyclopentadienyl rings are arranged in parallel, in layers, or in the form of a "sandwich" with an iron atom located between them in the center (Fig. 8). The name "ferrocene" was proposed by R. Woodward (or rather, an employee of his group, D. Whiting). It reflects the presence in the compound of an iron atom and ten carbon atoms (zehn - ten).
All ten bonds (C-Fe) in the ferrocene molecule are equivalent, the Fe-c internuclear distance is 2.04 Å. All carbon atoms in a ferrocene molecule are structurally and chemically equivalent, the length of each C-C connections 1.40 - 1.41 Å (for comparison, in benzene the C-C bond length is 1.39 Å). A 36-electron shell appears around the iron atom.
In 1973, Ernst Otto Fischer and Jeffrey Wilkinson were awarded the Nobel Prize in Chemistry for their pioneering work done independently in the field of organometallic, so-called sandwich compounds. Indvar Lindqvist, a member of the Royal Swedish Academy of Sciences, in his speech at the presentation of the laureates, stated that "the discovery and proof of new principles of bonds and structures found in sandwich compounds is a significant achievement, the practical significance of which at present cannot yet be predicted."
At present, dicyclopentadienyl derivatives of many metals have been obtained. Transition metal derivatives have the same structure and the same bond nature as ferrocene. The lanthanides do not form a sandwich structure, but a structure resembling a three-beam star [The atoms of La, Ce, Pr, Nd, therefore, create a fifteen-center chemical bond.
Soon after ferrocene, dibenzenechromium was obtained. Dibenzene-molybdenum and dibenzene-vanadium were prepared according to the same scheme. In all compounds of this class, the metal atoms hold together two six-membered rings. All 12 metal-carbon bonds in these compounds are identical.
Uranocene [bis(cyclooctatetraene)uranium] has also been synthesized, in which the uranium atom holds two eight-membered rings. All 16 uranium-carbon bonds in the uranocene are identical. Uranocene is obtained by reacting UCl 4 with a mixture of cyclooctatetraene and potassium in tetrahydrofuran at minus 300 C.
Fig.1. Orbital radii of elements (r a) and length of one-electron chemical bond (d)
The simplest one-electron chemical bond is created by a single valence electron. It turns out that one electron is able to hold two positively charged ions in a single whole. In a one-electron bond, the Coulomb repulsive forces of positively charged particles are compensated by the Coulomb forces of attraction of these particles to a negatively charged electron. The valence electron becomes common to the two nuclei of the molecule.
Examples of such chemical compounds are molecular ions: H 2 + , Li 2 + , Na 2 + , K 2 + , Rb 2 + , Cs 2 + :
A polar covalent bond occurs in heteronuclear diatomic molecules (Fig. 3). The bonding electron pair in a polar chemical bond is close to the atom with a higher first ionization potential.
The distance d between atomic nuclei, which characterizes the spatial structure of polar molecules, can be approximately considered as the sum of the covalent radii of the corresponding atoms.
Characterization of some polar substancesThe shift of the binding electron pair to one of the nuclei of the polar molecule leads to the appearance of an electric dipole (electrodynamics) (Fig. 4).
The distance between the centers of gravity of positive and negative charges is called the length of the dipole. The polarity of the molecule, as well as the polarity of the bond, is estimated by the value of the dipole moment μ, which is the product of the length of the dipole l and the value of the electronic charge:
Multiple covalent bonds
Multiple covalent bonds are represented by unsaturated organic compounds containing double and triple chemical bonds. To describe the nature of unsaturated compounds, L. Pauling introduces the concepts of sigma and π bonds, hybridization of atomic orbitals.
Pauling's hybridization for two S- and two p-electrons allowed the directionality of chemical bonds to be explained, in particular the tetrahedral configuration of methane. To explain the structure of ethylene, it is necessary to isolate one p-electron from four equivalent Sp 3 electrons of the carbon atom to form an additional bond, called the π-bond. In this case, the three remaining Sp 2 -hybrid orbitals are located in the plane at an angle of 120° and form the main bonds, for example, a flat ethylene molecule (Fig. 5).
In Pauling's new theory, all binding electrons became equal and equidistant from the line connecting the nuclei of the molecule. Pauling's theory of a bent chemical bond took into account the statistical interpretation of the wave function by M. Born, the Coulomb electron correlation of electrons. A physical meaning appeared - the nature of the chemical bond is completely determined by the electrical interaction of nuclei and electrons. The more bonding electrons, the smaller the internuclear distance and the stronger the chemical bond between carbon atoms.
Three-center chemical bond
Further development of ideas about the chemical bond was given by the American physical chemist W. Lipscomb, who developed the theory of two-electron three-center bonds and a topological theory that makes it possible to predict the structure of some more boron hydrides (borohydrides).
An electron pair in a three-center chemical bond becomes common to three atomic nuclei. In the simplest representative of a three-center chemical bond - the molecular hydrogen ion H 3 +, an electron pair holds three protons in a single whole (Fig. 6).
Fig. 7. Diboran
The existence of boranes with their two-electron three-center bonds with "bridge" hydrogen atoms violated the canonical doctrine of valency. The hydrogen atom, previously considered a standard univalent element, turned out to be bound by identical bonds with two boron atoms and became formally a divalent element. The work of W. Lipscomb on deciphering the structure of boranes expanded the understanding of the chemical bond. The Nobel Committee awarded the William Nunn Lipscomb Prize in Chemistry in 1976 with the wording "For his studies of the structure of boranes (borohydrites) which elucidate the problems of chemical bonds".
Multicenter chemical bond
Fig. 8. Ferrocene molecule
Fig. 9. Dibenzenechromium
Fig. 10. Uranocene
All ten bonds (C-Fe) in the ferrocene molecule are equivalent, the Fe-c internuclear distance is 2.04 Å. All carbon atoms in the ferrocene molecule are structurally and chemically equivalent, the length of each C-C bond is 1.40 - 1.41 Å (for comparison, in benzene the C-C bond length is 1.39 Å). A 36-electron shell appears around the iron atom.
Chemical bond dynamics
The chemical bond is quite dynamic. Thus, a metallic bond is transformed into a covalent bond during a phase transition during the evaporation of the metal. The transition of a metal from a solid to a vapor state requires the expenditure of large amounts of energy.
In vapors, these metals consist practically of homonuclear diatomic molecules and free atoms. When metal vapor condenses, the covalent bond turns into a metal one.
The evaporation of salts with a typical ionic bond, such as alkali metal fluorides, leads to the destruction of the ionic bond and the formation of heteronuclear diatomic molecules with a polar covalent bond. In this case, the formation of dimeric molecules with bridging bonds takes place.
Characterization of the chemical bond in the molecules of alkali metal fluorides and their dimers.
During the condensation of vapors of alkali metal fluorides, the polar covalent bond is transformed into an ionic one with the formation of the corresponding crystal lattice of the salt.
The mechanism of the transition of a covalent to a metallic bond
Fig.11. Relationship between the orbital radius of an electron pair r e and the length of a covalent chemical bond d
Fig.12. Orientation of the dipoles of diatomic molecules and the formation of a distorted octahedral cluster fragment during the condensation of alkali metal vapors
Fig. 13. Body-centered cubic arrangement of nuclei in alkali metal crystals and a link
Disperse attraction (London forces) causes interatomic interaction and the formation of homonuclear diatomic molecules from alkali metal atoms.
The formation of a metal-metal covalent bond is associated with the deformation of the electron shells of the interacting atoms - valence electrons create a binding electron pair, the electron density of which is concentrated in the space between the atomic nuclei of the resulting molecule. A characteristic feature of homonuclear diatomic molecules of alkali metals is the long length of the covalent bond (3.6-5.8 times the bond length in the hydrogen molecule) and the low energy of its rupture.
The specified ratio between re and d determines the uneven distribution of electric charges in the molecule - in the middle part of the molecule, the negative electric charge of the binding electron pair is concentrated, and at the ends of the molecule - positive electric charges two atomic cores.
The uneven distribution of electric charges creates conditions for the interaction of molecules due to orientational forces (van der Waals forces). Molecules of alkali metals tend to orient themselves in such a way that opposite electric charges appear in the neighborhood. As a result, attractive forces act between the molecules. Due to the presence of the latter, alkali metal molecules approach each other and are more or less firmly drawn together. At the same time, some deformation of each of them occurs under the action of closer located poles of neighboring molecules (Fig. 12).
In fact, the binding electrons of the original diatomic molecule, falling into the electric field of four positively charged atomic cores of alkali metal molecules, break off from the orbital radius of the atom and become free.
In this case, the bonding electron pair becomes common even for a system with six cations. The construction of the crystal lattice of the metal begins at the cluster stage. In the crystal lattice of alkali metals, the structure of the connecting link is clearly expressed, having the shape of a distorted oblate octahedron - a square bipyramid, the height of which and the edges of the basis are equal to the value of the constant translational lattice a w (Fig. 13).
The value of the translational lattice constant a w of an alkali metal crystal significantly exceeds the length of the covalent bond of an alkali metal molecule, therefore it is generally accepted that the electrons in the metal are in a free state:
The mathematical construction associated with the properties of free electrons in a metal is usually identified with the "Fermi surface", which should be considered as a geometric place where electrons reside, providing the main property of the metal - to conduct electric current.
When comparing the process of condensation of alkali metal vapors with the process of condensation of gases, for example, hydrogen, salient feature in the properties of the metal. So, if weak intermolecular interactions appear during the condensation of hydrogen, then during the condensation of metal vapors, processes characteristic of chemical reactions. The condensation of metal vapor itself proceeds in several stages and can be described by the following procession: a free atom → a diatomic molecule with a covalent bond → a metal cluster → a compact metal with a metal bond.
The interaction of alkali metal halide molecules is accompanied by their dimerization. A dimeric molecule can be considered as an electric quadrupole (Fig. 15). At present, the main characteristics of alkali metal halide dimers (chemical bond lengths and bond angles) are known.
Chemical bond length and bond angles in dimers of alkali metal halides (E 2 X 2) (gas phase).
E 2 X 2 | X=F | X=Cl | X=Br | X=I | ||||
---|---|---|---|---|---|---|---|---|
d EF , Å | d ECl , Å | d EBr , Å | d EI , Å | |||||
Li 2 X 2 | 1,75 | 105 | 2,23 | 108 | 2,35 | 110 | 2,54 | 116 |
Na 2 X 2 | 2,08 | 95 | 2,54 | 105 | 2,69 | 108 | 2,91 | 111 |
K2X2 | 2,35 | 88 | 2,86 | 98 | 3,02 | 101 | 3,26 | 104 |
Cs 2 X 2 | 2,56 | 79 | 3,11 | 91 | 3,29 | 94 | 3,54 | 94 |
In the process of condensation, the action of orientational forces is enhanced, intermolecular interaction is accompanied by the formation of clusters, and then a solid. Alkali metal halides form crystals with a simple cubic and body-centered cubic lattice.
Lattice type and translational lattice constant for alkali metal halides.
In the process of crystallization, a further increase in the interatomic distance occurs, leading to the removal of an electron from the orbital radius of an alkali metal atom and the transfer of an electron to a halogen atom with the formation of the corresponding ions. Force fields of ions are evenly distributed in all directions in space. In this regard, in alkali metal crystals, the force field of each ion coordinates by no means one ion with the opposite sign, as it is customary to qualitatively represent the ionic bond (Na + Cl -).
In crystals of ionic compounds, the concept of simple two-ion molecules such as Na + Cl - and Cs + Cl - loses its meaning, since the alkali metal ion is associated with six chloride ions (in a sodium chloride crystal) and eight chlorine ions (in a cesium chloride crystal. In this case, all interionic distances in crystals are equidistant.
Notes
- Handbook of inorganic chemistry. Constants of inorganic substances. - M .: "Chemistry", 1987. - S. 124. - 320 p.
- Lidin R.A., Andreeva L.L., Molochko V.A. Handbook of inorganic chemistry. Constants of inorganic substances. - M .: "Chemistry", 1987. - S. 132-136. - 320 s.
- Gankin V.Yu., Gankin Yu.V. How chemical bonds are formed and how chemical reactions proceed. - M .: publishing group "Border", 2007. - 320 p. - ISBN 978-5-94691296-9
- Nekrasov B.V. General chemistry course. - M .: Goshimizdat, 1962. - S. 88. - 976 p.
- Pauling L. The nature of the chemical bond / edited by Ya.K. Syrkin. - per. from English. M.E. Dyatkina. - M.-L.: Goshimizdat, 1947. - 440 p.
- Theoretical organic chemistry / ed. R.Kh. Freidlina. - per. from English. Yu.G. Bundel. - M .: Ed. foreign literature, 1963. - 365 p.
- Lemenovsky D.A., Levitsky M.M. Russian Chemical Journal (Journal of the Russian Chemical Society named after D.I. Mendeleev). - 2000. - T. XLIV, issue 6. - S. 63-86.
- Chemical Encyclopedic Dictionary / Ch. ed. I.L.Knunyants. - M .: Sov. Encyclopedia, 1983. - S. 607. - 792 p.
- Nekrasov B.V. General chemistry course. - M .: Goshimizdat, 1962. - S. 679. - 976 p.
- Lidin R.A., Andreeva L.L., Molochko V.A. Handbook of inorganic chemistry. Constants of inorganic substances. - M .: "Chemistry", 1987. - S. 155-161. - 320 s.
- Gillespie R. Geometry of molecules / per. from English. E.Z. Zasorina and V.S. Mastryukov, ed. Yu.A. Pentina. - M .: "Mir", 1975. - S. 49. - 278 p.
- Handbook of a chemist. - 2nd ed., revised. and additional - L.-M.: GNTI Chemical Literature, 1962. - T. 1. - S. 402-513. - 1072 p.
- Lidin R.A., Andreeva L.L., Molochko V.A. Handbook of inorganic chemistry. Constants of inorganic substances .. - M .: "Chemistry", 1987. - S. 132-136. - 320 s.
- Zieman J. Electrons in metals (introduction to the theory of Fermi surfaces). Advances in physical sciences .. - 1962. - T. 78, issue 2. - 291 p.
see also
- chemical bond- article from the Great Soviet Encyclopedia
- chemical bond- Chemport.ru
- chemical bond- Physical Encyclopedia
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