Examples of formulas for metallic bond substances. Metal bond: mechanism of formation

All metals have the following characteristics:

A small number of electrons at the outer energy level (except for some exceptions, which may have 6,7 and 8);

Large atomic radius;

Low ionization energy.

All this contributes to the easy separation of outer unpaired electrons from the nucleus. At the same time, the atom has a lot of free orbitals. The diagram of the formation of a metallic bond will precisely show the overlap of numerous orbital cells of different atoms with each other, which as a result form a common intracrystalline space. Electrons are fed into it from each atom, which begin to wander freely around different parts grates. Periodically, each of them attaches to an ion at a site in the crystal and turns it into an atom, then detaches again to form an ion.

Thus, A metallic bond is the bond between atoms, ions, and free electrons in a common metal crystal. An electron cloud moving freely within a structure is called an “electron gas.” It explains most of the physical properties of metals and their alloys.

How exactly does metal realize itself? chemical bond? Various examples can be given. Let's try to look at it on a piece of lithium. Even if you take it the size of a pea, there will be thousands of atoms. So let’s imagine that each of these thousands of atoms gives up its single valence electron to the common crystalline space. At the same time, knowing the electronic structure of this element, you can see the number of empty orbitals. Lithium will have 3 of them (p-orbitals of the second energy level). Three for each atom out of tens of thousands - this is the common space inside the crystal in which the “electron gas” moves freely.

A substance with a metal bond is always strong. After all, electron gas does not allow the crystal to collapse, but only displaces the layers and immediately restores them. It shines, has a certain density (usually high), fusibility, malleability and plasticity.

Where else is metal bonding sold? Examples of substances:

Metals in the form of simple structures;

All metals alloy with each other;

All metals and their alloys in liquid and solid states.

There are simply an incredible number of specific examples, since there are more than 80 metals in the periodic table!

The mechanism of education in general view is expressed next entry: Ме 0 - e - ↔ Ме n+ . From the diagram it is obvious what particles are present in the metal crystal.

Any metal can give up electrons, becoming a positively charged ion.

Using iron as an example: Fe 0 -2e - = Fe 2+

Where do the separated negatively charged particles - electrons - go? A minus is always attracted to a plus. The electrons are attracted to another (positively charged) iron ion in the crystal lattice: Fe 2+ +2e - = Fe 0

The ion becomes a neutral atom. And this process is repeated many times.

It turns out that the free electrons of iron are in constant movement throughout the entire volume of the crystal, breaking away and joining ions at lattice sites. Another name for this phenomenon is delocalized electron cloud. The term "delocalized" means free, not tied.

It is extremely rare that chemical substances consist of individual, unrelated atoms of chemical elements. Under normal conditions, only a small number of gases called noble gases have this structure: helium, neon, argon, krypton, xenon and radon. Most often, chemical substances do not consist of isolated atoms, but of their combinations into various groups. Such associations of atoms can number a few, hundreds, thousands, or even more atoms. The force that holds these atoms in such groups is called chemical bond.

In other words, we can say that a chemical bond is an interaction that provides the connection of individual atoms into more complex structures (molecules, ions, radicals, crystals, etc.).

The reason for the formation of a chemical bond is that the energy of more complex structures is less than the total energy of the individual atoms that form it.

So, in particular, if the interaction of atoms X and Y produces a molecule XY, this means that internal energy molecules of this substance are lower than the internal energy of the individual atoms from which it was formed:

E(XY)< E(X) + E(Y)

For this reason, when chemical bonds are formed between individual atoms, energy is released.

Electrons of the outer electron layer with the lowest binding energy with the nucleus, called valence. For example, in boron these are electrons of the 2nd energy level - 2 electrons per 2 s- orbitals and 1 by 2 p-orbitals:

When a chemical bond is formed, each atom tends to obtain the electronic configuration of noble gas atoms, i.e. so that there are 8 electrons in its outer electron layer (2 for elements of the first period). This phenomenon is called the octet rule.

It is possible for atoms to achieve the electron configuration of a noble gas if initially single atoms share some of their valence electrons with other atoms. In this case, common electron pairs are formed.

Depending on the degree of sharing of electrons, covalent, ionic and metallic bonds can be distinguished.

Covalent bond

Covalent bonds most often occur between atoms of nonmetal elements. If the nonmetal atoms forming a covalent bond belong to different chemical elements, such a bond is called a polar covalent bond. The reason for this name lies in the fact that atoms different elements They also have different abilities to attract a common electron pair. Obviously, this leads to a displacement of the common electron pair towards one of the atoms, as a result of which a partial negative charge is formed on it. In turn, a partial positive charge is formed on the other atom. For example, in a hydrogen chloride molecule the electron pair is shifted from the hydrogen atom to the chlorine atom:

Examples of substances with polar covalent bonds:

CCl 4, H 2 S, CO 2, NH 3, SiO 2, etc.

A covalent nonpolar bond is formed between nonmetal atoms of the same chemical element. Since the atoms are identical, their ability to attract shared electrons is also the same. In this regard, no displacement of the electron pair is observed:

The above mechanism for the formation of a covalent bond, when both atoms provide electrons to form common electron pairs, is called exchange.

There is also a donor-acceptor mechanism.

When a covalent bond is formed by the donor-acceptor mechanism, a shared electron pair is formed due to the filled orbital of one atom (with two electrons) and the empty orbital of another atom. An atom that provides a lone pair of electrons is called a donor, and an atom with a vacant orbital is called an acceptor. Atoms that have paired electrons, for example N, O, P, S, act as donors of electron pairs.

For example, according to the donor-acceptor mechanism, the formation of the fourth covalent N-H connections in the ammonium cation NH 4 +:

In addition to polarity, covalent bonds are also characterized by energy. Bond energy is the minimum energy required to break a bond between atoms.

The binding energy decreases with increasing radii of bonded atoms. Since we know that atomic radii increase down the subgroups, we can, for example, conclude that the strength of the halogen-hydrogen bond increases in the series:

HI< HBr < HCl < HF

Also, the bond energy depends on its multiplicity - the greater the bond multiplicity, the greater its energy. Bond multiplicity refers to the number of shared electron pairs between two atoms.

Ionic bond

An ionic bond can be considered as an extreme case of a polar covalent bond. If in a covalent-polar bond the common electron pair is partially shifted to one of the pair of atoms, then in an ionic bond it is almost completely “given” to one of the atoms. The atom that donates electron(s) acquires a positive charge and becomes cation, and the atom that has taken electrons from it acquires a negative charge and becomes anion.

Thus, ionic bond is a bond formed due to the electrostatic attraction of cations to anions.

The formation of this type of bond is typical during the interaction of atoms of typical metals and typical non-metals.

For example, potassium fluoride. The potassium cation is formed by the removal of one electron from a neutral atom, and the fluorine ion is formed by the addition of one electron to the fluorine atom:

An electrostatic force of attraction arises between the resulting ions, resulting in the formation of an ionic compound.

When a chemical bond was formed, electrons from the sodium atom passed to the chlorine atom and oppositely charged ions were formed, which have a completed external energy level.

It has been established that electrons from the metal atom are not completely detached, but are only shifted towards the chlorine atom, as in a covalent bond.

Most binary compounds that contain metal atoms are ionic. For example, oxides, halides, sulfides, nitrides.

Ionic bonding also occurs between simple cations and simple anions (F −, Cl −, S 2-), as well as between simple cations and complex anions (NO 3 −, SO 4 2-, PO 4 3-, OH −). Therefore, ionic compounds include salts and bases (Na 2 SO 4, Cu(NO 3) 2, (NH 4) 2 SO 4), Ca(OH) 2, NaOH)

Metal connection

This type of bond is formed in metals.

Atoms of all metals have electrons in their outer electron layer that have a low binding energy with the nucleus of the atom. For most metals, the process of losing outer electrons is energetically favorable.

Due to such a weak interaction with the nucleus, these electrons in metals are very mobile and the following process continuously occurs in each metal crystal:

М 0 — ne − = M n + ,

where M 0 is a neutral metal atom, and M n + a cation of the same metal. The figure below provides an illustration of the processes taking place.

That is, electrons “rush” across a metal crystal, detaching from one metal atom, forming a cation from it, joining another cation, forming a neutral atom. This phenomenon was called “electron wind,” and the collection of free electrons in a crystal of a nonmetal atom was called “electron gas.” This type of interaction between metal atoms is called a metallic bond.

Hydrogen bond

If a hydrogen atom in a substance is bonded to an element with high electronegativity (nitrogen, oxygen or fluorine), such a substance is characterized by a phenomenon such as hydrogen bond.

Since a hydrogen atom is bonded to an electronegative atom, a partial positive charge is formed on the hydrogen atom, and a partial negative charge is formed on the atom of the electronegative element. In this regard, electrostatic attraction becomes possible between a partially positively charged hydrogen atom of one molecule and an electronegative atom of another. For example, hydrogen bonding is observed for water molecules:

It is the hydrogen bond that explains the abnormally high melting point of water. In addition to water, strong hydrogen bonds are also formed in substances such as hydrogen fluoride, ammonia, oxygen-containing acids, phenols, alcohols, and amines.

163120 0

Each atom has a certain number of electrons.

Entering 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 opposite sign 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 of T and II groups periodic table 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 are insoluble organic compounds, not containing - OH groups, 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, characteristic property ionic bond - the complete movement of two 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 limiting cases of the set existing types chemical bonds, and in practice most bonds are intermediate.

Connections of two elements located at opposite ends of one or different periods Mendeleev's systems 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, halides and oxides of 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 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's 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 Brønsted-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 that 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 brief overview types of bonds, 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, ionic radii of metals and van der Waals radii of interacting molecules increase as their serial number 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

A metallic bond is a bond formed between atoms under conditions of strong delocalization (distribution of valence electrons over several chemical bonds in a compound) and electron deficiency in the atom (crystal). It is unsaturated and spatially non-directional.

Delocalization of valence electrons in metals is a consequence of the multicenter nature of the metal bond. The multicenter nature of the metal bond ensures high electrical conductivity and thermal conductivity of metals.

Saturability determined by the number of valence orbitals involved in the formation of a chemical. communications. Quantitative characteristic – valency. Valence is the number of bonds that one atom can form with others; - is determined by the number of valence orbitals involved in the formation of bonds according to the exchange and donor-acceptor mechanisms.

Focus – the connection is formed in the direction of maximum overlap of electron clouds; - determines the chemical and crystal chemical structure of a substance (how atoms are connected in a crystal lattice).

When a covalent bond is formed, the electron density is concentrated between the interacting atoms (drawing from notebook). In the case of a metallic bond, the electron density is delocalized throughout the crystal. (drawing from notebook)

(example from notebook)

Due to the unsaturated and non-directional nature of the metallic bond, metallic bodies (crystals) are highly symmetrical and highly coordinated. The vast majority of metal crystal structures correspond to 3 types of atomic packings in crystals:

1. GCC– grain-centered cubic close-packed structure. Packing density – 74.05%, coordination number = 12.

2. GPU– hexagonal close-packed structure, packing density = 74.05%, c.h. = 12.

3. BCC– volume is centered, packing density = 68.1%, c.h. = 8.

A metallic bond does not exclude some degree of covalency. Metal connection in its pure form is characteristic only of alkali and alkaline earth metals.

A pure metallic bond is characterized by an energy of the order of 100/150/200 kJ/mol, 4 times weaker than a covalent bond.

36. Chlorine and its properties. В=1(III, IV, V and VII) oxidation state=7, 6, 5, 4, 3, 1, −1

yellow-green gas with a pungent irritating odor. Chlorine occurs in nature only in the form of compounds. In nature in the form of potassium chloride, magnesium, nitrium, formed as a result of the evaporation of former seas and lakes. Receipt.prom:2NaCl+2H2O=2NaOH+H2+Cl2, by electrolysis of waters of chlorides Me.\2KMnO4+16HCl=2MnCl2+2KCl+8H2O+5Cl2/Chemically, chlorine is very active, directly combines with almost all Me, and with non-metals (except carbon, nitrogen, oxygen, inert gases), replaces hydrogen in hydrocarbons and joins unsaturated compounds, displaces bromine and iodine from their compounds. Phosphorus ignites in an atmosphere of chlorine PCl3, and with further chlorination - PCl5; sulfur with chlorine = S2Сl2, SСl2 and other SnClm. A mixture of chlorine and hydrogen burns. With oxygen, chlorine forms oxides: Cl2O, ClO2, Cl2O6, Cl2O7, Cl2O8, as well as hypochlorites (salts of hypochlorous acid), chlorites, chlorates and perchlorates. All oxygen compounds of chlorine form explosive mixtures with easily oxidized substances. Chlorine oxides are unstable and can spontaneously explode; hypochlorites slowly decompose during storage; chlorates and perchlorates can explode under the influence of initiators. in water - hypochlorous and salty: Cl2 + H2O = HClO + HCl. When aqueous solutions of alkalis are chlorinated in the cold, hypochlorites and chlorides are formed: 2NaOH + Cl2 = NaClO + NaCl + H2O, and when heated, chlorates are formed. When ammonia reacts with chlorine, nitrogen trichloride is formed. interhalogen compounds with other halogens. Fluorides ClF, ClF3, ClF5 are very reactive; for example, in a ClF3 atmosphere, glass wool spontaneously ignites. Known compounds of chlorine with oxygen and fluorine are chlorine oxyfluorides: ClO3F, ClO2F3, ClOF, ClOF3 and fluorine perchlorate FClO4. Application: production of chemical compounds, water purification, food synthesis, pharmaceutical industry, bactericide, antiseptic, bleaching of papers, fabrics, pyrotechnics, matches, destroys weeds in agriculture.

Biological role: biogenic, component of plant and animal tissues. 100g is the main osmotically active substance in blood plasma, lymph, cerebrospinal fluid and some tissues. Daily sodium chloride requirement = 6-9g - bread, meat and dairy products. Plays a role in water-salt metabolism, promoting tissue retention of water. Regulation of acid-base balance in tissues is carried out along with other processes by changing the distribution of chlorine between the blood and other tissues; chlorine is involved in energy metabolism in plants, activating both oxidative phosphorylation and photophosphorylation. Chlorine has a positive effect on the absorption of oxygen by roots, a component of iron sap.

37. Hydrogen, water. B = 1; st. oxide = + 1-1 The hydrogen ion is completely devoid of electron shells and can approach very close distances and penetrate into electron shells.

The most common element in the Universe. It makes up the bulk of the Sun, stars and other cosmic bodies. In a free state on Earth, it is found relatively rarely - it is found in oil and combustible gases, present in the form of inclusions in some minerals, and most of it in water. Receipt: 1. Laboratory Zn+2HCl=ZnCl2+H 2 ; 2.Si+2NaOH+H 2 O=Na 2 SiO 3 +2H 2; 3. Al+NaOH+H 2 O=Na(AlOH) 4 +H 2. 4. In industry: conversion, electrolysis: СH4+H2O=CO+3H2\CO+H2O=CO+ H2/Him St. In no.:H 2 +F 2 =2HF. When irradiated, illuminated, catalysts: H 2 + O 2 , S, N, P = H 2 O, H 2 S, NH 3 , Ca + H2 = CaH2\F2 + H2 = 2HF\N2 + 3H2 → 2NH3\Cl2 + H2 → 2HCl, 2NO+2H2=N2+2H2O,CuO+H2=Cu+H2O,CO+H2=CH3OH. Hydrogen forms hydrides: ionic, covalent and metallic. To ionic –NaH -& ,CaH 2 -& +H 2 O=Ca(OH) 2 ;NaH+H 2 O=NaOH+H 2 . Covalent –B 2 H 6 , AlH 3 , SiH 4 . Metal – with d-elements; variable composition: MeH ≤1, MeH ≤2 – are introduced into the voids between atoms. Conducts heat, current, solids. WATER.sp3-hybrid highly polar molecule at an angle of 104.5 , dipoles, the most common solvent. Water reacts at room temperature with active halogens (F, Cl) and interhalogen compounds with salts, weak forms and weak bases, causing their complete hydrolysis ; with carbonic and inorganic anhydrides and acid halides. acid; with active metal organ compounds; with carbides, nitrides, phosphides, silicides, hydrides of active Me; with many salts, forming hydrates; with boranes, silanes; with ketenes, carbon dioxide; with fluorides of noble gases. Water reacts when heated: with Fe, Mg, coal, methane; with some alkyl halides. Application:hydrogen -synthesis of ammonia, methanol, hydrogen chloride, TV fats, hydrogen flame - for welding, melting, in metallurgy for the reduction of Me from oxide, fuel for rockets, in pharmacy - water, peroxide-antiseptic, bactericide, washing, hair bleaching, sterilization.

Biological role: hydrogen-7kg, The main function of hydrogen is the structuring of biological space (water and hydrogen bonds) and the formation of a variety of organic molecules (included in the structure of proteins, carbohydrates, fats, enzymes) Thanks to hydrogen bonds,

copying a DNA molecule. Water takes part in a huge

number of biochemical reactions, in all physiological and biological

processes, ensures metabolism between the body and external environment, between

cells and within cells. Water is the structural basis of cells and is necessary for

maintaining their optimal volume, it determines the spatial structure and

functions of biomolecules.

Chemical bond

All interactions leading to the combination of chemical particles (atoms, molecules, ions, etc.) into substances are divided into chemical bonds and intermolecular bonds (intermolecular interactions).

Chemical bonds- bonds directly between atoms. There are ionic, covalent and metallic bonds.

Intermolecular bonds- connections between molecules. These are hydrogen bonds, ion-dipole bonds (due to the formation of this bond, for example, the formation of a hydration shell of ions occurs), dipole-dipole (due to the formation of this bond, molecules of polar substances are combined, for example, in liquid acetone), etc.

Ionic bond- a chemical bond formed due to the electrostatic attraction of oppositely charged ions. In binary compounds (compounds of two elements), it is formed when the sizes of the bonded atoms differ greatly from each other: some atoms are large, others are small - that is, some atoms easily give up electrons, while others tend to accept them (usually these are atoms of the elements that form typical metals and atoms of elements forming typical nonmetals); the electronegativity of such atoms is also very different.
Ionic bonding is non-directional and non-saturable.

Covalent bond- a chemical bond that occurs due to the formation of a common pair of electrons. A covalent bond is formed between small atoms with the same or similar radii. Prerequisite- the presence of unpaired electrons in both bonded atoms (exchange mechanism) or a lone pair in one atom and a free orbital in the other (donor-acceptor mechanism):

A)	H· + ·H H:H	H-H	H 2	(one shared pair of electrons; H is monovalent);
b)		NN	N 2	(three shared pairs of electrons; N is trivalent);
V)		H-F	HF	(one shared pair of electrons; H and F are monovalent);
G)			NH4+	(four shared pairs of electrons; N is tetravalent)

simple (single)- one pair of electrons,
double- two pairs of electrons,
triples- three pairs of electrons.

Double and triple bonds are called multiple bonds.

According to the distribution of electron density between the bonded atoms, a covalent bond is divided into non-polar And polar. A non-polar bond is formed between identical atoms, a polar one - between different ones.

Electronegativity- a measure of the ability of an atom in a substance to attract common electron pairs.
The electron pairs of polar bonds are shifted towards more electronegative elements. The displacement of electron pairs itself is called bond polarization. The partial (excess) charges formed during polarization are designated + and -, for example: .

Based on the nature of the overlap of electron clouds ("orbitals"), a covalent bond is divided into -bond and -bond.
-A bond is formed due to the direct overlap of electron clouds (along the straight line connecting the atomic nuclei), -a bond is formed due to lateral overlap (on both sides of the plane in which the atomic nuclei lie).

A covalent bond is directional and saturable, as well as polarizable.
The hybridization model is used to explain and predict the mutual direction of covalent bonds.

Hybridization of atomic orbitals and electron clouds- the supposed alignment of atomic orbitals in energy, and electron clouds in shape when an atom forms covalent bonds.
The three most common types of hybridization are: sp-, sp 2 and sp 3 -hybridization. For example:
sp-hybridization - in molecules C 2 H 2, BeH 2, CO 2 (linear structure);
sp 2-hybridization - in molecules C 2 H 4, C 6 H 6, BF 3 (flat triangular shape);
sp 3-hybridization - in molecules CCl 4, SiH 4, CH 4 (tetrahedral form); NH 3 (pyramidal shape); H 2 O (angular shape).

Metal connection- a chemical bond formed by sharing the valence electrons of all bonded atoms of a metal crystal. As a result, a single electron cloud of the crystal is formed, which easily moves under the influence of electrical voltage - hence the high electrical conductivity of metals.
A metallic bond is formed when the atoms being bonded are large and therefore tend to give up electrons. Simple substances with a metallic bond are metals (Na, Ba, Al, Cu, Au, etc.), complex substances are intermetallic compounds (AlCr 2, Ca 2 Cu, Cu 5 Zn 8, etc.).
The metal bond does not have directionality or saturation. It is also preserved in metal melts.

Hydrogen bond- an intermolecular bond formed due to the partial acceptance of a pair of electrons from a highly electronegative atom by a hydrogen atom with a large positive partial charge. It is formed in cases where one molecule contains an atom with a lone pair of electrons and high electronegativity (F, O, N), and the other contains a hydrogen atom bound by a highly polar bond to one of such atoms. Examples of intermolecular hydrogen bonds:

H—O—H OH 2 , H—O—H NH 3 , H—O—H F—H, H—F H—F.

Intramolecular hydrogen bonds exist in polypeptide molecules, nucleic acids, proteins, etc.

A measure of the strength of any bond is the bond energy.
Communication energy- the energy required to break a given chemical bond in 1 mole of a substance. The unit of measurement is 1 kJ/mol.

The energies of ionic and covalent bonds are of the same order of magnitude, the energy of hydrogen bonds is an order of magnitude lower.

The energy of a covalent bond depends on the size of the bonded atoms (bond length) and on the multiplicity of the bond. The smaller the atoms and the greater the bond multiplicity, the greater its energy.

The ionic bond energy depends on the size of the ions and their charges. The smaller the ions and the greater their charge, the greater the binding energy.

Structure of matter

According to the type of structure, all substances are divided into molecular And non-molecular. Among organic matter molecular substances predominate; among inorganic substances, non-molecular substances predominate.

Based on the type of chemical bond, substances are divided into substances with covalent bonds, substances with ionic bonds (ionic substances) and substances with metallic bonds (metals).

Substances with covalent bonds can be molecular or non-molecular. This significantly affects their physical properties.

Molecular substances consist of molecules connected to each other by weak intermolecular bonds, these include: H 2, O 2, N 2, Cl 2, Br 2, S 8, P 4 and others simple substances; CO 2, SO 2, N 2 O 5, H 2 O, HCl, HF, NH 3, CH 4, C 2 H 5 OH, organic polymers and many other substances. These substances do not have high strength, they have low temperatures melting and boiling, do not carry out electric current, some of them are soluble in water or other solvents.

Non-molecular substances with covalent bonds or atomic substances (diamond, graphite, Si, SiO 2, SiC and others) form very strong crystals (with the exception of layered graphite), they are insoluble in water and other solvents, have high temperatures melting and boiling, most of them do not conduct electric current (except for graphite, which is electrically conductive, and semiconductors - silicon, germanium, etc.)

All ionic substances are naturally non-molecular. These are solid, refractory substances, solutions and melts of which conduct electric current. Many of them are soluble in water. It should be noted that in ionic substances, the crystals of which consist of complex ions, there are also covalent bonds, for example: (Na +) 2 (SO 4 2-), (K +) 3 (PO 4 3-), (NH 4 + )(NO 3-), etc. The atoms that make up complex ions are connected by covalent bonds.

Metals (substances with metallic bonds) very diverse in their physical properties. Among them there are liquid (Hg), very soft (Na, K) and very hard metals(W, Nb).

Characteristic physical properties metals are their high electrical conductivity (unlike semiconductors, it decreases with increasing temperature), high heat capacity and ductility (for pure metals).

In the solid state, almost all substances are composed of crystals. According to the type of structure and type of chemical bond, crystals (" crystal lattices") divided by atomic(crystals are not molecular substances with a covalent bond), ionic(crystals of ionic substances), molecular(crystals of molecular substances with covalent bonds) and metal(crystals of substances with a metallic bond).

Tasks and tests on the topic "Topic 10. "Chemical bonding. Structure of matter."

Types of chemical bond - Structure of matter grade 8–9
Lessons: 2 Assignments: 9 Tests: 1
Assignments: 9 Tests: 1

Having worked through this topic, you should understand the following concepts: chemical bond, intermolecular bond, ionic bond, covalent bond, metallic bond, hydrogen bond, single bond, double bond, triple bond, multiple bonds, non-polar bond, polar bond, electronegativity, bond polarization, - and - bond, hybridization of atomic orbitals , binding energy.

You must know the classification of substances by type of structure, by type of chemical bond, the dependence of the properties of simple and complex substances on the type of chemical bond and the type of “crystal lattice”.

You must be able to: determine the type of chemical bond in a substance, the type of hybridization, draw up diagrams of bond formation, use the concept of electronegativity, a number of electronegativity; know how electronegativity changes in chemical elements of the same period and one group to determine the polarity of a covalent bond.

After making sure that everything you need has been learned, proceed to completing the tasks. We wish you success.

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