Type of chemical bond in metals definition. Metal connection. Metal crystal lattice and metal chemical bond. Covalent polar chemical bond
Atoms of most elements do not exist separately, as they can interact with each other. This interaction produces more complex particles.
The nature of a 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 on the outer electronic levels (valence electrons), being farthest from the nucleus, interact with it weakest, and therefore are able to break away from the nucleus. They are responsible for bonding atoms to each other.
Types of interactions in chemistry
Types of chemical bonds can be presented in the following table:
Characteristics of ionic bonding
Chemical reaction that occurs due to ion attraction having different charges is called ionic. This happens if the atoms being bonded have a significant difference in electronegativity (that is, the ability to attract electrons) and the electron pair goes to the more electronegative element. The result of this transfer of electrons from one atom to another is the formation of charged particles - ions. An attraction arises between them.
They have the lowest electronegativity indices typical metals, and the largest are typical non-metals. Ions are thus formed by the interaction between typical metals and typical nonmetals.
Metal atoms become positively charged ions (cations), donating electrons to their outer electron levels, and nonmetals 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 non-saturable, since the electrostatic interaction occurs in all directions; accordingly, the ion can attract ions of the opposite sign in all directions.
The arrangement of the ions is such that around each there 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 to form the corresponding ions:
Na 0 - 1 e = Na + (cation)
Cl 0 + 1 e = Cl - (anion)
In sodium chloride, there are six chlorine anions around the sodium cations, and six sodium ions around each chloride ion.
When interaction is formed between atoms in barium sulfide, the following processes occur:
Ba 0 - 2 e = Ba 2+
S 0 + 2 e = 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 are easily separated 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 separated from atoms.
The structure of the metal substance is as follows: the crystal lattice is the skeleton of the substance, and between its nodes electrons can move freely.
The following examples can be given:
Mg - 2е<->Mg 2+
Cs - e<->Cs+
Ca - 2e<->Ca2+
Fe-3e<->Fe 3+
Covalent: polar and non-polar
The most common type of chemical interaction is a covalent bond. The electronegativity values of the elements that interact do not differ sharply; therefore, only a shift of the common electron pair to a more electronegative atom occurs.
Covalent interactions can be formed by an exchange mechanism or a donor-acceptor mechanism.
The exchange mechanism is realized if each of the atoms has unpaired electrons on 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 on the outer electronic level, and the other has a free orbital, then when the atomic orbitals overlap, the electron pair is shared and interacts according to the donor-acceptor mechanism.
Covalent ones are divided by multiplicity into:
- simple or single;
- double;
- triples.
Double ones ensure the sharing of two pairs of electrons at once, and triple ones - three.
According to the distribution of electron density (polarity) between bonded atoms, a covalent bond is divided into:
- non-polar;
- polar.
A nonpolar bond is formed by identical atoms, and a polar bond is formed by different electronegativity.
The interaction of atoms with similar electronegativity is called a nonpolar bond. The common pair of electrons in such a molecule is not attracted to either atom, but belongs equally to both.
The interaction of elements differing in electronegativity leads to the formation of polar bonds. In this type of interaction, shared electron pairs are attracted to the more electronegative element, but are not completely transferred to it (that is, the formation of ions does not occur). As a result of this shift in electron density, partial charges appear on the atoms: the more electronegative one has a negative charge, and the less electronegative one has a positive charge.
Properties and characteristics of covalency
Main characteristics of a covalent bond:
- The length is determined by the distance between the nuclei of interacting atoms.
- Polarity is determined by the displacement of the electron cloud towards one of the atoms.
- Directionality is the property of forming bonds oriented in space and, accordingly, molecules having 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 determines its strength.
An example of a covalent non-polar interaction can be the molecules of hydrogen (H2), chlorine (Cl2), oxygen (O2), nitrogen (N2) and many others.
H· + ·H → H-H molecule has a single non-polar bond,
O: + :O → O=O molecule has a double nonpolar,
Ṅ: + Ṅ: → N≡N the molecule is triple nonpolar.
Examples of covalent bonds of chemical elements include molecules of carbon dioxide (CO2) and carbon monoxide (CO), hydrogen sulfide (H2S), 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. Oxygen has two unpaired electrons in its outer shell, while 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 metal substances form a metallic bond, metals with nonmetals form an ionic bond, simple nonmetal substances form a covalent nonpolar bond, and molecules consisting of different nonmetals form through a polar covalent bond.
A metallic bond is a chemical bond caused by the presence of relatively free electrons. Characteristic of both pure metals and their alloys and intermetallic compounds.
Metal link mechanism
Positive metal ions are located at all nodes of the crystal lattice. Between them, valence electrons move randomly, like gas molecules, detached from the atoms during the formation of ions. These electrons act as cement, holding the positive ions together; otherwise, the lattice would disintegrate under the influence of repulsive forces between the ions. At the same time, electrons are held by ions within the crystal lattice and cannot leave it. The coupling forces are not localized or directed.
Therefore, in most cases high coordination numbers (for example, 12 or 8) appear. When two metal atoms come close together, the orbitals in their outer shells overlap to form molecular orbitals. If a third atom approaches, its orbital overlaps with the orbitals of the first two atoms, giving another molecular orbital. When there are many atoms, a huge number of three-dimensional molecular orbitals arise, extending in all directions. Due to multiple overlapping orbitals, the valence electrons of each atom are influenced by many atoms.
Characteristic crystal lattices
Most metals form one of the following highly symmetrical lattices with close packing of atoms: body-centered cubic, face-centered cubic, and hexagonal.
In a body-centered cubic (bcc) lattice, the atoms are located at the vertices of the cube and one atom is at the center of the cube volume. Metals have a cubic body-centered lattice: Pb, K, Na, Li, β-Ti, β-Zr, Ta, W, V, α-Fe, Cr, Nb, Ba, etc.
In a face-centered cubic (fcc) lattice, the atoms are located at the vertices of the cube and at the center of each face. Metals of this type have a lattice: α-Ca, Ce, α-Sr, Pb, Ni, Ag, Au, Pd, Pt, Rh, γ-Fe, Cu, α-Co, etc.
In a hexagonal lattice, the atoms are located at the vertices and center of the hexagonal bases of the prism, and three atoms are located in the middle plane of the prism. Metals have this packing of atoms: Mg, α-Ti, Cd, Re, Os, Ru, Zn, β-Co, Be, β-Ca, etc.
Other properties
Freely moving electrons cause high electrical and thermal conductivity. Substances that have a metallic bond often combine strength with plasticity, since when atoms are displaced relative to each other, the bonds do not break. Another important property is metallic aromaticity.
Metals conduct heat and electricity well, they are strong enough, and can be deformed without destruction. Some metals are malleable (they can be forged), some are malleable (you can draw wire from them). These unique properties are explained by a special type of chemical bond that connects metal atoms to each other - a metallic bond.
Metals in the solid state exist in the form of crystals of positive ions, as if “floating” in a sea of electrons freely moving between them.
Metallic bond explains the properties of metals, in particular their strength. Under the influence of a deforming force, a metal lattice can change its shape without cracking, unlike ionic crystals.
The high thermal conductivity of metals is explained by the fact that if a piece of metal is heated on one side, the kinetic energy of the electrons will increase. This increase in energy will spread in a “sea of electrons” throughout the sample at high speed.
The electrical conductivity of metals also becomes clear. If a potential difference is applied to the ends of a metal sample, the cloud of delocalized electrons will shift in the direction of a positive potential: this flow of electrons moving in one direction represents the familiar electric current.
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Each atom has a certain number of electrons.
When entering into chemical reactions, atoms donate, gain, or share electrons, achieving the most stable electronic configuration. The configuration with the lowest energy (as in noble gas atoms) turns out to be the most stable. This pattern is called the “octet rule” (Fig. 1).
Rice. 1.
This rule applies to everyone types of connections. Electronic connections between atoms allow them to form stable structures, from the simplest crystals to complex biomolecules that ultimately form living systems. They differ from crystals in their continuous metabolism. At the same time, many chemical reactions proceed according to the mechanisms electronic transfer, which play a critical role in energy processes in the body.
A chemical bond is the force that holds together two or more atoms, ions, molecules, or any combination of these.
The nature of a 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 of 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, located in the highest energy orbitals. Accordingly, the outer shell of the atom containing these orbitals is called valence shell. Currently, 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 valence theory, 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 opposite signs, a chemical bond is formed, called by 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 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 of ionic compounds are unsaturated and non-directional, so they retain the possibility of electrostatic interaction with other ions. In Fig. Figures 2 and 3 show examples of ionic bonds corresponding to the Kossel model of electron transfer.
Rice. 2.
Rice. 3. Ionic bond in a molecule of table salt (NaCl)
Here it is appropriate to recall some properties that explain the behavior of substances in nature, in particular, consider the idea of acids And reasons.
Aqueous solutions of all these substances are electrolytes. They change color differently indicators. The mechanism of action of indicators was discovered by F.V. Ostwald. He showed that indicators are weak acids or bases, the color of which differs in the undissociated and dissociated states.
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 undergo 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 N 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 N+ , 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 accordance with proton According to the theory of Brønsted and Lowry, an acid is a substance containing molecules or ions that donate protons ( donors protons), and a base is a substance consisting of molecules or ions that accept protons ( acceptors protons). Note that in aqueous solutions, hydrogen ions exist in hydrated form, that is, in the form of hydronium ions H3O+ . This theory describes reactions not only with water and hydroxide ions, but also those carried out in the absence of a solvent or with a non-aqueous solvent.
For example, in the reaction between ammonia N.H. 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 conjugate pairs of acids and bases:
1)N.H. 4+ and N.H. 3
2) HCl And Cl ‑
Here, in each conjugate 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 Brønsted-Lowry theory helps 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 in reactions with aqueous solutions of ammonia, it is an acid.
1) CH 3 COOH + H2O ↔ H3O + + CH 3 COO- . Here, an acetic acid molecule donates a proton to a water molecule;
2) NH 3 + H2O ↔ NH 4 + + HE- . Here, an ammonia molecule accepts a proton from a water molecule.
Thus, water can form two conjugate pairs:
1) H2O(acid) and HE- (conjugate base)
2) H 3 O+ (acid) and H2O(conjugate base).
In the first case, water donates a proton, and in the second, it accepts it.
This property is called amphiprotonism. Substances that can react as both acids and bases are called amphoteric. Such substances are often found in living nature. For example, amino acids can form salts with both acids and bases. Therefore, peptides easily form coordination compounds with the metal ions present.
Thus, a characteristic property of an ionic bond is the complete movement of the bonding electrons to one of the nuclei. This means that between the ions there is a region 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 everyone atom. In this case, the shared bond electrons are distributed equally between the atoms. Examples of covalent bonds include homonuclear diatomic molecules H 2 , N 2 , F 2. The same type of connection is found in allotropes 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 WITH 2 N 5 HE, sulfur hexafluoride SF 6, acetylene WITH 2 N 2. All these molecules share the same electrons, and their bonds are saturated and directed in the same way (Fig. 4).
It is important for biologists that double and triple bonds have reduced covalent atomic radii compared to a single bond.
Rice. 4. Covalent bond in a Cl 2 molecule.
Ionic and covalent types of bonds are two extreme cases of the many existing types of chemical bonds, and in practice most bonds are intermediate.
Compounds of two elements located at opposite ends of the same or different periods of the periodic system predominantly form ionic bonds. As elements move closer together within a period, the ionic nature of their compounds decreases, and the covalent character increases. For example, the halides and oxides of elements on the left side of the periodic table form predominantly ionic bonds ( NaCl, AgBr, BaSO 4, CaCO 3, KNO 3, CaO, NaOH), and the same compounds of 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 one more modification.
In polyatomic ions and in complex biological molecules, both electrons can only come from one atom. It is called donor electron pair. An atom that shares 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 d-elements most important for metabolism is largely described by coordination bonds.
Fig. 5.
As a rule, in a complex compound the metal atom acts as an acceptor of an electron pair; 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. Lewis's 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 electron pair, which, by donating electrons, forms a covalent bond with Lewis acid.
That is, Lewis's 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 capable of accepting an electron pair.
Therefore, according to this theory, the cations are Lewis acids and the anions are Lewis bases. An example would be the following reactions:
It was noted above that the division of substances into ionic and covalent is relative, since complete electron transfer from metal atoms to acceptor atoms does not occur in covalent molecules. In compounds with ionic bonds, 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 by the electronic structure, charge and size of the ion; for anions it is higher than for cations. The highest polarizability among cations is for cations of higher charge and smaller size, for example, Hg 2+, Cd 2+, Pb 2+, Al 3+, Tl 3+. Has a strong polarizing effect N+ . Since the influence of ion polarization is two-way, it significantly changes the properties of the compounds they form.
The third type of connection isdipole-dipole connection
In addition to the listed types of communication, there are also dipole-dipole intermolecular interactions, also called 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 ( dispersive 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 Debaya(1D = 3.338 × 10‑30 coulomb meters - C × m).
In biochemistry, there is another type of connection - 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 similar electronegativity (such as chlorine and sulfur), the hydrogen bond is much weaker. The hydrogen atom is distinguished by one significant feature: when the bonding electrons are pulled away, its nucleus - the proton - is exposed and is no longer shielded 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 an important role in biochemistry, for example, to stabilize the structure of proteins in the form of an a-helix, or for the formation of a double helix of DNA (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. 1.
Table 1. Energy of intermolecular forces
Note: The degree of intermolecular interactions is reflected by the enthalpy of melting and evaporation (boiling). Ionic compounds require significantly more energy to separate ions than to separate molecules. The enthalpy of melting of ionic compounds is much higher than that of molecular compounds.
The fourth type of connection ismetal connection
Finally, there is another type of intermolecular bonds - metal: connection of positive ions of a metal lattice with free electrons. This type of connection does not occur in biological objects.
From a brief review of bond types, one detail becomes clear: an important parameter of a metal atom or ion - an electron donor, as well as an atom - an electron acceptor, is its 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 increases in groups of the periodic table. 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.
Of greatest importance for biologists and physicians are coordination(donor-acceptor) bonds considered by coordination chemistry.
Medical bioinorganics. G.K. Barashkov
As already indicated in paragraph 4.2.2.1, metal connection- electronic bonding of atomic nuclei with minimal localization of shared electrons both on individual (as opposed to an ionic bond) nuclei, and on individual (as opposed to a covalent bond) bonds. The result is an electron-deficient multicenter chemical bond in which shared electrons (in the form of “electron gas”) provide bonding to the maximum possible number of nuclei (cations) that form the structure of liquid or solid metallic substances. Therefore, the metallic bond as a whole is non-directional and saturated; it should be considered as limiting case of delocalization of a covalent bond. Let us recall that in pure metals the metallic bond appears primarily homonuclear, i.e. cannot have an ionic component. As a result, a typical picture of the electron density distribution in metals is spherically symmetrical cores (cations) in a uniformly distributed electron gas (Fig. 5.10).
Consequently, the final structure of compounds with a predominantly metallic type of bond is determined primarily by the steric factor and packing density in the crystal lattice of these cations (high CN). The BC method cannot interpret metallic bonds. According to MMO, a metallic bond is characterized by a deficiency of electrons compared to a covalent bond. Strict application of MMO to metallic bonds and connections leads to band theory(electronic model of a metal), according to which in the atoms included in the crystal lattice of a metal, there is an interaction of almost free valence electrons located in external electron orbits with the (electric) periodic field of the crystal lattice. As a result, the energy levels of electrons split and form a more or less wide band. According to Fermi statistics, the highest energy band is populated by free electrons up to complete filling, especially if the energy terms of an individual atom correspond to two electrons with antiparallel spins. However, it can be partially filled, which provides the opportunity for electrons to move to higher energy levels. Then
this zone is called the conduction zone. There are several basic types of relative arrangement of energy bands, corresponding to an insulator, a monovalent metal, a divalent metal, a semiconductor with intrinsic conductivity, a -type semiconductor and an impurity semiconductor/b-type. The ratio of energy bands also determines the type of conductivity of a solid.
However, this theory does not allow quantitative characterization of various metal compounds and has not led to a solution to the problem of the origin of real crystal structures of metal phases. The specific nature of chemical bonds in homonuclear metals, metal alloys and intermetallic heterocompounds is considered by N.V. Ageev)