How to distinguish bonds in chemistry. Types of chemical bonds: ionic, covalent, metallic

Covalent chemical bond, its varieties and mechanisms of formation. Characteristics of covalent bonds (polarity and bond energy). Ionic bond. Metal connection. Hydrogen bond

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

A chemical bond is understood as the interaction of atoms that binds them into molecules, ions, radicals, and crystals.

There are four types of chemical bonds: ionic, covalent, metallic and hydrogen.

The division of chemical bonds into types is conditional, since they are all characterized by a certain unity.

An ionic bond can be considered as an extreme case of a polar covalent bond.

A metallic bond combines the covalent interaction of atoms using shared electrons and the electrostatic attraction between these electrons and metal ions.

Substances often lack limiting cases of chemical bonding (or pure chemical bonding).

For example, lithium fluoride $LiF$ is classified as an ionic compound. In fact, the bond in it is $80%$ ionic and $20%$ covalent. It is therefore more correct, obviously, to talk about the degree of polarity (ionicity) of a chemical bond.

In the series of hydrogen halides $HF—HCl—HBr—HI—HAt$ the degree of bond polarity decreases, because the difference in the electronegativity values ​​of the halogen and hydrogen atoms decreases, and in astatine hydrogen the bond becomes almost nonpolar $(EO(H) = 2.1; EO(At) = 2.2)$.

Different types of bonds can be found in the same substances, for example:

1. in bases: between the oxygen and hydrogen atoms in hydroxo groups the bond is polar covalent, and between the metal and the hydroxo group it is ionic;
2. in salts of oxygen-containing acids: between the non-metal atom and the oxygen of the acidic residue - covalent polar, and between the metal and the acidic residue - ionic;
3. in ammonium, methylammonium salts, etc.: between nitrogen and hydrogen atoms - covalent polar, and between ammonium or methylammonium ions and the acid residue - ionic;
4. in metal peroxides (for example, $Na_2O_2$), the bond between oxygen atoms is covalent nonpolar, and between the metal and oxygen is ionic, etc.

Different types of connections can transform into one another:

— during electrolytic dissociation of covalent compounds in water, the covalent polar bond turns into an ionic bond;

- when metals evaporate, the metal bond turns into a nonpolar covalent bond, etc.

The reason for the unity of all types and types of chemical bonds is their identical chemical nature - electron-nuclear interaction. The formation of a chemical bond in any case is the result of electron-nuclear interaction of atoms, accompanied by the release of energy.

Methods for forming covalent bonds. Characteristics of a covalent bond: bond length and energy

A covalent chemical bond is a bond formed between atoms through the formation of shared electron pairs.

The mechanism of formation of such a bond can be exchange or donor-acceptor.

I. Exchange mechanism operates when atoms form shared electron pairs by combining unpaired electrons.

1) $H_2$ - hydrogen:

The bond arises due to the formation of a common electron pair by $s$-electrons of hydrogen atoms (overlapping $s$-orbitals):

2) $HCl$ - hydrogen chloride:

The bond arises due to the formation of a common electron pair of $s-$ and $p-$electrons (overlapping $s-p-$orbitals):

3) $Cl_2$: in a chlorine molecule, a covalent bond is formed due to unpaired $p-$electrons (overlapping $p-p-$orbitals):

4) $N_2$: in a nitrogen molecule three common electron pairs are formed between the atoms:

II. Donor-acceptor mechanism education covalent bond Let's consider the example of the ammonium ion $NH_4^+$.

The donor has an electron pair, the acceptor has an empty orbital that this pair can occupy. In the ammonium ion, all four bonds with hydrogen atoms are covalent: three were formed due to the creation of common electron pairs by the nitrogen atom and hydrogen atoms according to the exchange mechanism, one - through the donor-acceptor mechanism.

Covalent bonds can be classified by the way the electron orbitals overlap, as well as by their displacement towards one of the bonded atoms.

Chemical bonds formed as a result of overlapping electron orbitals along a bond line are called $σ$ -bonds (sigma bonds). The sigma bond is very strong.

$p-$orbitals can overlap in two regions, forming a covalent bond due to lateral overlap:

Chemical bonds formed as a result of “lateral” overlap of electron orbitals outside the communication line, i.e. in two areas are called $π$ -bonds (pi-bonds).

By degree of displacement shared electron pairs to one of the atoms they bond, a covalent bond can be polar And non-polar.

A covalent chemical bond formed between atoms with the same electronegativity is called non-polar. Electron pairs are not shifted to any of the atoms, because atoms have the same EO - the property of attracting valence electrons from other atoms. For example:

those. molecules of simple non-metal substances are formed through covalent non-polar bonds. A covalent chemical bond between atoms of elements whose electronegativity differs is called polar.

Length and energy of covalent bonds.

Characteristic properties of covalent bond- its length and energy. Link length is the distance between the nuclei of atoms. The shorter the length of a chemical bond, the stronger it is. However, a measure of the strength of the connection is binding energy, which is determined by the amount of energy required to break a bond. It is usually measured in kJ/mol. Thus, according to experimental data, the bond lengths of $H_2, Cl_2$ and $N_2$ molecules are respectively $0.074, 0.198$ and $0.109$ nm, and the bond energies are respectively $436, 242$ and $946$ kJ/mol.

Ions. Ionic bond

Let's imagine that two atoms “meet”: an atom of a group I metal and a non-metal atom of group VII. A metal atom has a single electron at its outer energy level, while a non-metal atom just lacks one electron for its outer level to be complete.

The first atom will easily give the second its electron, which is far from the nucleus and weakly bound to it, and the second will provide it with a free place on its outer electronic level.

Then the atom, deprived of one of its negative charges, will become a positively charged particle, and the second will turn into a negatively charged particle due to the resulting electron. Such particles are called ions.

The chemical bond that occurs between ions is called ionic.

Let's consider the formation of this bond using the example of the well-known compound sodium chloride (table salt):

The process of converting atoms into ions is depicted in the diagram:

This transformation of atoms into ions always occurs during the interaction of atoms of typical metals and typical non-metals.

Let's consider the algorithm (sequence) of reasoning when recording the formation of an ionic bond, for example, between calcium and chlorine atoms:

Numbers showing the number of atoms or molecules are called coefficients, and numbers showing the number of atoms or ions in a molecule are called indexes.

Metal connection

Let's get acquainted with how atoms of metal elements interact with each other. Metals usually do not exist as isolated atoms, but in the form of a piece, ingot, or metal product. What holds metal atoms in a single volume?

The atoms of most metals at the external level do not contain big number electrons - $1, 2, 3$. These electrons are easily stripped off and the atoms become positive ions. The detached electrons move from one ion to another, binding them into a single whole. Connecting with ions, these electrons temporarily form atoms, then break off again and combine with another ion, etc. Consequently, in the volume of the metal, atoms are continuously converted into ions and vice versa.

The bond in metals between ions through shared electrons is called metallic.

The figure schematically shows the structure of a sodium metal fragment.

In this case, a small number of shared electrons bind a large number of ions and atoms.

A metallic bond has some similarities with a covalent bond, since it is based on the sharing of external electrons. However, with a covalent bond, the outer unpaired electrons of only two neighboring atoms are shared, while with a metallic bond, all atoms take part in the sharing of these electrons. That is why crystals with a covalent bond are brittle, but with a metal bond, as a rule, they are ductile, electrically conductive and have a metallic luster.

Metallic bonding is characteristic of both pure metals and mixtures of various metals—alloys in solid and liquid states.

Hydrogen bond

A chemical bond between positively polarized hydrogen atoms of one molecule (or part thereof) and negatively polarized atoms of strongly electronegative elements having lone electron pairs ($F, O, N$ and less commonly $S$ and $Cl$) of another molecule (or its part) is called hydrogen.

The mechanism of hydrogen bond formation is partly electrostatic, partly donor-acceptor in nature.

Examples of intermolecular hydrogen bonding:

In the presence of such a connection, even low-molecular substances can, under normal conditions, be liquids (alcohol, water) or easily liquefied gases (ammonia, hydrogen fluoride).

Substances with hydrogen bonds have molecular crystal lattices.

Substances of molecular and non-molecular structure. Type of crystal lattice. Dependence of the properties of substances on their composition and structure

Molecular and non-molecular structure of substances

It is not individual atoms or molecules that enter into chemical interactions, but substances. Under given conditions, a substance can be in one of three states of aggregation: solid, liquid or gaseous. The properties of a substance also depend on the nature of the chemical bond between the particles that form it - molecules, atoms or ions. Based on the type of bond, substances of molecular and non-molecular structure are distinguished.

Substances made up of molecules are called molecular substances. The bonds between the molecules in such substances are very weak, much weaker than between the atoms inside the molecule, and even at relatively low temperatures they break - the substance turns into a liquid and then into a gas (sublimation of iodine). The melting and boiling points of substances consisting of molecules increase with increasing molecular weight.

TO molecular substances include substances with an atomic structure ($C, Si, Li, Na, K, Cu, Fe, W$), among them there are metals and non-metals.

Let's consider physical properties alkali metals. The relatively low bond strength between atoms causes low mechanical strength: alkali metals are soft and can be easily cut with a knife.

Large atomic sizes lead to low densities of alkali metals: lithium, sodium and potassium are even lighter than water. In the group of alkali metals, the boiling and melting points decrease with increasing atomic number of the element, because Atom sizes increase and bonds weaken.

To substances non-molecular structures include ionic compounds. Most compounds of metals with nonmetals have this structure: all salts ($NaCl, K_2SO_4$), some hydrides ($LiH$) and oxides ($CaO, MgO, FeO$), bases ($NaOH, KOH$). Ionic (non-molecular) substances have high melting and boiling points.

Crystal lattices

Matter, as is known, can exist in three states of aggregation: gaseous, liquid and solid.

Solids: amorphous and crystalline.

Let us consider how the characteristics of chemical bonds influence the properties of solids. Solids are divided into crystalline And amorphous.

Amorphous substances do not have a clear melting point; when heated, they gradually soften and turn into a fluid state. For example, plasticine and various resins are in an amorphous state.

Crystalline substances are characterized by the correct arrangement of the particles of which they are composed: atoms, molecules and ions - at strictly defined points in space. When these points are connected by straight lines, a spatial framework is formed, called a crystal lattice. The points at which crystal particles are located are called lattice nodes.

Depending on the type of particles located at the nodes of the crystal lattice and the nature of the connection between them, four types of crystal lattices are distinguished: ionic, atomic, molecular And metal.

Ionic crystal lattices.

Ionic are called crystal lattices, in the nodes of which there are ions. They are formed by substances with ionic bonds, which can bind both simple ions $Na^(+), Cl^(-)$, and complex $SO_4^(2−), OH^-$. Consequently, salts and some oxides and hydroxides of metals have ionic crystal lattices. For example, a sodium chloride crystal consists of alternating positive $Na^+$ and negative $Cl^-$ ions, forming a cube-shaped lattice. The bonds between ions in such a crystal are very stable. Therefore, substances with an ionic lattice are characterized by relatively high hardness and strength, they are refractory and non-volatile.

Atomic crystal lattices.

Atomic are called crystal lattices, in the nodes of which there are individual atoms. In such lattices, the atoms are connected to each other by very strong covalent bonds. An example of substances with this type of crystal lattices is diamond, one of the allotropic modifications of carbon.

Most substances with an atomic crystal lattice have very high melting points (for example, for diamond it is above $3500°C), they are strong and hard, and practically insoluble.

Molecular crystal lattices.

Molecular called crystal lattices, in the nodes of which molecules are located. Chemical bonds in these molecules can be both polar ($HCl, H_2O$) and nonpolar ($N_2, O_2$). Despite the fact that the atoms inside the molecules are connected by very strong covalent bonds, between the molecules themselves there are weak forces intermolecular attraction. Therefore, substances with molecular crystal lattices have low hardness, low melting points, and are volatile. Most solid organic compounds have molecular crystal lattices (naphthalene, glucose, sugar).

Metal crystal lattices.

Substances with metallic bonds have metallic crystal lattices. At the sites of such lattices there are atoms and ions (either atoms or ions, into which metal atoms easily transform, giving up their outer electrons “to common use"). This internal structure of metals determines their characteristic physical properties: malleability, ductility, electrical and thermal conductivity, characteristic metallic luster.

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You know that atoms can combine with each other to form both simple and complex substances. In this case, various types of chemical bonds are formed: ionic, covalent (non-polar and polar), metallic and hydrogen. One of the most essential properties of atoms of elements that determine what kind of bond is formed between them - ionic or covalent - This is electronegativity, i.e. the ability of atoms in a compound to attract electrons.

Conditional quantitative assessment electronegativity is given by the relative electronegativity scale.

In periods, there is a general tendency for the electronegativity of elements to increase, and in groups - for their decrease. Elements are arranged in a row according to their electronegativity, on the basis of which the electronegativity of elements located in different periods can be compared.

The type of chemical bond depends on how large the difference in electronegativity values ​​of the connecting atoms of elements is. The more the atoms of the elements forming the bond differ in electronegativity, the more polar the chemical bond. It is impossible to draw a sharp boundary between the types of chemical bonds. In most compounds, the type of chemical bond is intermediate; for example, a highly polar covalent chemical bond is close to an ionic bond. Depending on which of the limiting cases a chemical bond is closer in nature, it is classified as either an ionic or a covalent polar bond.

Ionic bond.

An ionic bond is formed by the interaction of atoms that differ sharply from each other in electronegativity. For example, the typical metals lithium (Li), sodium (Na), potassium (K), calcium (Ca), strontium (Sr), barium (Ba) form ionic bonds with typical non-metals, mainly halogens.

In addition to alkali metal halides, ionic bonds also form in compounds such as alkalis and salts. For example, in sodium hydroxide (NaOH) and sodium sulfate (Na 2 SO 4) ionic bonds exist only between sodium and oxygen atoms (the remaining bonds are polar covalent).

Covalent nonpolar bond.

When atoms with the same electronegativity interact, molecules with a covalent nonpolar bond are formed. Such a bond exists in the molecules of the following simple substances: H 2, F 2, Cl 2, O 2, N 2. Chemical bonds in these gases are formed through shared electron pairs, i.e. when the corresponding electron clouds overlap, due to the electron-nuclear interaction, which occurs when atoms approach each other.

When composing electronic formulas of substances, it should be remembered that each common electron pair is a conventional image of increased electron density resulting from the overlap of the corresponding electron clouds.

Covalent polar bond.

When atoms interact, the electronegativity values ​​of which differ, but not sharply, the common electron pair shifts to a more electronegative atom. This is the most common type of chemical bond, found in both inorganic and organic compounds.

Covalent bonds also fully include those bonds that are formed by a donor-acceptor mechanism, for example in hydronium and ammonium ions.

Metal connection.

The bond that is formed as a result of the interaction of relatively free electrons with metal ions is called a metallic bond. This type of bond is characteristic of simple substances - metals.

The essence of the process of metal bond formation is as follows: metal atoms easily give up valence electrons and turn into positively charged ions. Relatively free electrons detached from the atom move between positive metal ions. A metallic bond arises between them, i.e. Electrons, as it were, cement the positive ions of the crystal lattice of metals.

Hydrogen bond.

A bond that forms between the hydrogen atoms of one molecule and an atom of a strongly electronegative element(O,N,F) another molecule is called a hydrogen bond.

The question may arise: why does hydrogen form such a specific chemical bond?

This is explained by the fact that the atomic radius of hydrogen is very small. In addition, when displacing or completely donating its only electron, hydrogen acquires a relatively high positive charge, due to which the hydrogen of one molecule interacts with atoms of electronegative elements that have a partial negative charge that goes into the composition of other molecules (HF, H 2 O, NH 3) .

Let's look at some examples. Usually we depict the composition of water chemical formula H 2 O. However, this is not entirely accurate. It would be more correct to denote the composition of water by the formula (H 2 O)n, where n = 2,3,4, etc. This is explained by the fact that individual water molecules are connected to each other through hydrogen bonds.

Hydrogen bonds are usually denoted by dots. It is much weaker than ionic or covalent bonds, but stronger than ordinary intermolecular interactions.

The presence of hydrogen bonds explains the increase in water volume with decreasing temperature. This is due to the fact that as the temperature decreases, the molecules become stronger and therefore the density of their “packing” decreases.

When studying organic chemistry, the following question arose: why are the boiling points of alcohols much higher than the corresponding hydrocarbons? This is explained by the fact that hydrogen bonds also form between alcohol molecules.

An increase in the boiling point of alcohols also occurs due to the enlargement of their molecules.

Hydrogen bonding is also characteristic of many other organic compounds (phenols, carboxylic acids, etc.). From organic chemistry courses and general biology You know that the presence of a hydrogen bond explains the secondary structure of proteins, the structure of the double helix of DNA, i.e. the phenomenon of complementarity.

Any interaction between atoms is possible only if there is a chemical bond. Such a connection is the reason for the formation of a stable polyatomic system - a molecular ion, molecule, crystal lattice. A strong chemical bond requires a lot of energy to break, which is why it is the basic quantity for measuring bond strength.

Conditions for the formation of a chemical bond

The formation of a chemical bond is always accompanied by the release of energy. This process occurs due to a decrease potential energy systems of interacting particles - molecules, ions, atoms. The potential energy of the resulting system of interacting elements is always less than the energy of unbound outgoing particles. Thus, the basis for the emergence of a chemical bond in a system is the decrease in the potential energy of its elements.

Nature of chemical interaction

A chemical bond is a consequence of the interaction of electromagnetic fields that arise around the electrons and atomic nuclei of those substances that take part in the formation of a new molecule or crystal. After the discovery of the theory of atomic structure, the nature of this interaction became more accessible to study.

For the first time the idea of electrical nature chemical bonding originated with the English physicist G. Davy, who suggested that molecules are formed due to the electrical attraction of oppositely charged particles. This idea interested the Swedish chemist and natural scientist I.Ya. Bercellius, who developed the electrochemical theory of the occurrence of chemical bonds.

The first theory to explain the processes chemical interaction substances, was imperfect, and over time had to be abandoned.

Butlerov's theory

A more successful attempt to explain the nature of the chemical bond of substances was made by the Russian scientist A.M. Butlerov. This scientist based his theory on the following assumptions:

• Atoms in the bonded state are connected to each other in a certain order. A change in this order causes the formation of a new substance.
• Atoms bond with each other according to the laws of valence.
• The properties of a substance depend on the order of connection of atoms in the molecule of the substance. A different arrangement causes a change in the chemical properties of the substance.
• Atoms connected to each other most strongly influence each other.

Butlerov's theory explained the properties chemical substances not only by their composition, but also by the order of arrangement of atoms. This internal order of A.M. Butlerov called it “chemical structure”.

The theory of the Russian scientist made it possible to restore order in the classification of substances and provided the opportunity to determine the structure of molecules by their chemical properties. The theory also answered the question: why molecules containing the same number of atoms have different chemical properties.

Prerequisites for the creation of theories of chemical bonding

In his theory chemical structure Butlerov did not touch upon the question of what a chemical bond is. There was too little data for this then. internal structure substances. Only after the discovery of the planetary model of the atom, the American scientist Lewis began to develop the hypothesis that a chemical bond arises through the formation of an electron pair that simultaneously belongs to two atoms. Subsequently, this idea became the foundation for the development of the theory of covalent bonds.

Covalent chemical bond

Sustainable chemical compound can be formed when the electron clouds of two neighboring atoms overlap. The result of such mutual intersection is an increasing electron density in the internuclear space. The nuclei of atoms, as we know, are positively charged, and therefore try to be drawn as close as possible to the negatively charged electron cloud. This attraction is much stronger than the repulsive forces between two positively charged nuclei, so this connection is stable.

Chemical bond calculations were first performed by chemists Heitler and London. They examined the bond between two hydrogen atoms. The simplest visual representation of it might look like this:

As you can see, the electron pair occupies a quantum place in both hydrogen atoms. This two-center arrangement of electrons is called a “covalent chemical bond.” Covalent bonds are typical of molecules of simple substances and their non-metal compounds. Substances created by covalent bonds usually do not conduct electricity or are semiconductors.

Ionic bond

An ionic chemical bond occurs when two oppositely charged ions attract each other. Ions can be simple, consisting of one atom of a substance. In compounds of this type, simple ions are most often positively charged metal atoms of groups 1 and 2 that have lost their electron. The formation of negative ions is inherent in the atoms of typical nonmetals and their acid bases. Therefore, among the typical ionic compounds there are many alkali metal halides, such as CsF, NaCl, and others.

Unlike a covalent bond, an ion is not saturated: a varying number of oppositely charged ions can join an ion or group of ions. The number of attached particles is limited only by the linear dimensions of the interacting ions, as well as the condition under which the attractive forces of oppositely charged ions must be greater than the repulsive forces of equally charged particles participating in the ionic type compound.

Hydrogen bond

Even before the creation of the theory of chemical structure, it was experimentally noticed that hydrogen compounds with various non-metals have somewhat unusual properties. For example, the boiling points of hydrogen fluoride and water are much higher than might be expected.

These and other features of hydrogen compounds can be explained by the ability of the H + atom to form another chemical bond. This type of connection is called a “hydrogen bond.” The reasons for the occurrence of a hydrogen bond lie in the properties of electrostatic forces. For example, in a hydrogen fluoride molecule, the total electron cloud is so shifted towards fluorine that the space around an atom of this substance is saturated with negative electric field. Around a hydrogen atom, deprived of its only electron, the field is much weaker and has a positive charge. As a result, an additional relationship arises between the positive fields of electron clouds H + and negative F - .

Chemical bond of metals

The atoms of all metals are located in space in a certain way. The arrangement of metal atoms is called a crystal lattice. In this case, electrons of different atoms weakly interact with each other, forming a common electron cloud. This type of interaction between atoms and electrons is called a “metallic bond.”

It is the free movement of electrons in metals that can explain the physical properties of metallic substances: electrical conductivity, thermal conductivity, strength, fusibility and others.

3.3.1 Covalent bond is a two-center, two-electron bond formed due to the overlap of electron clouds carrying unpaired electrons with antiparallel spins. As a rule, it is formed between atoms of one chemical element.

It is quantitatively characterized by valence. Valency of the element - this is its ability to form a certain number of chemical bonds due to free electrons located in the atomic valence band.

A covalent bond is formed only by a pair of electrons located between atoms. It's called a split pair. The remaining pairs of electrons are called lone pairs. They fill the shells and do not take part in binding. The connection between atoms can be carried out not only by one, but also by two and even three divided pairs. Such connections are called double etc swarm - multiple connections.

3.3.1.1 Covalent nonpolar bond. A bond achieved through the formation of electron pairs that belong equally to both atoms is called covalent nonpolar. It occurs between atoms with practically equal electronegativity (0.4 > ΔEO > 0) and, therefore, a uniform distribution of electron density between the nuclei of atoms in homonuclear molecules. For example, H 2, O 2, N 2, Cl 2, etc. The dipole moment of such bonds is zero. The CH bond in saturated hydrocarbons (for example, in CH 4) is considered practically nonpolar, because ΔEO = 2.5 (C) - 2.1 (H) = 0.4.

3.3.1.2 Covalent polar bond. If a molecule is formed by two different atoms, then the overlap zone of electron clouds (orbitals) shifts towards one of the atoms, and such a bond is called polar . With such a bond, the probability of finding electrons near the nucleus of one of the atoms is higher. For example, HCl, H 2 S, PH 3.

Polar (unsymmetrical) covalent bond - bonding between atoms with different electronegativity (2 > ΔEO > 0.4) and asymmetric distribution of the common electron pair. Typically, it forms between two non-metals.

The electron density of such a bond is shifted towards a more electronegative atom, which leads to the appearance of a partial negative charge (delta minus) on it, and a partial positive charge (delta plus) on the less electronegative atom.

C ?  .

The direction of electron displacement is also indicated by an arrow:

CCl, CO, CN, OH, CMg.

The greater the difference in electronegativity of the bonded atoms, the higher the polarity of the bond and the greater its dipole moment. Additional attractive forces act between partial charges of opposite sign. Therefore, the more polar the bond, the stronger it is.

Except polarizability covalent bond has the property saturation – the ability of an atom to form as many covalent bonds as it has energetically available atomic orbitals. The third property of a covalent bond is its direction.

3.3.2 Ionic bonding. The driving force behind its formation is the same desire of atoms for the octet shell. But in some cases, such an “octet” shell can only arise when electrons are transferred from one atom to another. Therefore, as a rule, an ionic bond is formed between a metal and a non-metal.

Consider, as an example, the reaction between sodium (3s 1) and fluorine (2s 2 3s 5) atoms. Electronegativity difference in NaF compound

EO = 4.0 - 0.93 = 3.07

Sodium, having given its 3s 1 electron to fluorine, becomes a Na + ion and remains with a filled 2s 2 2p 6 shell, which corresponds to the electronic configuration of the neon atom. Exactly the same electronic configuration acquires fluorine by accepting one electron donated by sodium. As a result, electrostatic attractive forces arise between oppositely charged ions.

Ionic bond - an extreme case of polar covalent bonding, based on the electrostatic attraction of ions. Such a bond occurs when there is a large difference in the electronegativity of the bonded atoms (EO > 2), when a less electronegative atom almost completely gives up its valence electrons and turns into a cation, and another, more electronegative atom, attaches these electrons and becomes an anion. The interaction of ions of the opposite sign does not depend on the direction, and Coulomb forces do not have the property of saturation. Due to this ionic bond has no spatial focus And saturation , since each ion is associated with a certain number of counterions (ion coordination number). Therefore, ionic-bonded compounds do not have a molecular structure and are solid substances that form ionic crystal lattices, with high melting and boiling points, they are highly polar, often salt-like, in aqueous solutions electrically conductive. For example, MgS, NaCl, A 2 O 3. There are practically no compounds with purely ionic bonds, since a certain amount of covalency always remains due to the fact that a complete transfer of one electron to another atom is not observed; in the most “ionic” substances, the proportion of bond ionicity does not exceed 90%. For example, in NaF the bond polarization is about 80%.

In organic compounds, ionic bonds are quite rare, because A carbon atom tends neither to lose nor to gain electrons to form ions.

Valence elements in compounds with ionic bonds are very often characterized oxidation state , which, in turn, corresponds to the charge value of the element ion in a given compound.

Oxidation state - this is a conventional charge that an atom acquires as a result of the redistribution of electron density. Quantitatively, it is characterized by the number of electrons displaced from a less electronegative element to a more electronegative one. A positively charged ion is formed from the element that gave up its electrons, and a negative ion is formed from the element that accepted these electrons.

The element located in highest oxidation state (maximum positive), has already given up all of its valence electrons located in the AVZ. And since their number is determined by the number of the group in which the element is located, then highest oxidation state for most elements and will be equal group number . Concerning lowest oxidation state (maximum negative), then it appears during the formation of an eight-electron shell, that is, in the case when the AVZ is completely filled. For non-metals it is calculated by the formula Group number – 8 . For metals equal to zero , since they cannot accept electrons.

For example, the AVZ of sulfur has the form: 3s 2 3p 4. If an atom gives up all its electrons (six), it will acquire the highest oxidation state +6 , equal to the group number VI , if it takes the two necessary to complete the stable shell, it will acquire the lowest oxidation state –2 , equal to Group number – 8 = 6 – 8= –2.

3.3.3 Metal bond. Most metals have a number of properties that are general in nature and differ from the properties of other substances. Such properties are relatively high melting temperatures, the ability to reflect light, and high thermal and electrical conductivity. These features are explained by the existence of a special type of interaction in metals – metal connection.

In accordance with their position in the periodic table, metal atoms have a small number of valence electrons, which are rather weakly bound to their nuclei and can easily be detached from them. As a result, positively charged ions appear in the crystal lattice of the metal, localized in certain positions of the crystal lattice, and a large number of delocalized (free) electrons, moving relatively freely in the field of positive centers and communicating between all metal atoms due to electrostatic attraction.

This is an important difference between metallic bonds and covalent bonds, which have a strict orientation in space. Bonding forces in metals are not localized or directed, and free electrons forming an “electron gas” cause high thermal and electrical conductivity. Therefore, in this case it is impossible to talk about the direction of the bonds, since the valence electrons are distributed almost evenly throughout the crystal. This is what explains, for example, the plasticity of metals, i.e. the possibility of displacement of ions and atoms in any direction

3.3.4 Donor-acceptor bond. In addition to the mechanism of covalent bond formation, according to which a shared electron pair arises from the interaction of two electrons, there is also a special donor-acceptor mechanism . It lies in the fact that a covalent bond is formed as a result of the transition of an already existing (lone) electron pair donor (electron supplier) for the common use of the donor and acceptor (supplier of free atomic orbital).

Once formed, it is no different from covalent. The donor-acceptor mechanism is well illustrated by the scheme for the formation of an ammonium ion (Figure 9) (asterisks indicate the electrons of the outer level of the nitrogen atom):

Figure 9 - Scheme of formation of ammonium ion

The electronic formula of the ABZ of the nitrogen atom is 2s 2 2p 3, that is, it has three unpaired electrons that enter into a covalent bond with three hydrogen atoms (1s 1), each of which has one valence electron. In this case, an ammonia molecule NH 3 is formed, in which the lone electron pair of nitrogen is retained. If a hydrogen proton (1s 0), which has no electrons, approaches this molecule, then nitrogen will transfer its pair of electrons (donor) to this hydrogen atomic orbital (acceptor), resulting in the formation of an ammonium ion. In it, each hydrogen atom is connected to a nitrogen atom by a common electron pair, one of which is implemented via a donor-acceptor mechanism. It's important to note that H-N connections, formed by different mechanisms, do not have any differences in properties. This phenomenon is due to the fact that at the moment of bond formation, the orbitals of the 2s and 2p electrons of the nitrogen atom change their shape. As a result, four orbitals of exactly the same shape appear.

Donors are usually atoms with a large number of electrons, but with a small number of unpaired electrons. For elements of period II, in addition to the nitrogen atom, such a possibility is available for oxygen (two lone pairs) and fluorine (three lone pairs). For example, the hydrogen ion H + in aqueous solutions is never in a free state, since the hydronium ion H 3 O + is always formed from water molecules H 2 O and the H + ion. The hydronium ion is present in all aqueous solutions, although for ease of writing it is preserved symbol H+.

3.3.5 Hydrogen bond. A hydrogen atom associated with a strongly electronegative element (nitrogen, oxygen, fluorine, etc.), which “pulls” a common electron pair onto itself, experiences a lack of electrons and acquires an effective positive charge. Therefore, it is able to interact with the lone pair of electrons of another electronegative atom (which acquires an effective negative charge) of the same (intramolecular bond) or another molecule (intermolecular bond). As a result, there is hydrogen bond , which is graphically indicated by dots:

This bond is much weaker than other chemical bonds (the energy of its formation is 10 – 40 kJ/mol) and mainly has a partially electrostatic, partially donor-acceptor character.

The hydrogen bond plays an extremely important role in biological macromolecules, such inorganic compounds as H 2 O, H 2 F 2, NH 3. For example, O-H bonds in H2O are noticeably polar in nature, with an excess of negative charge – on the oxygen atom. The hydrogen atom, on the contrary, acquires a small positive charge  + and can interact with the lone pairs of electrons of the oxygen atom of a neighboring water molecule.

The interaction between water molecules turns out to be quite strong, such that even in water vapor there are dimers and trimers of the composition (H 2 O) 2, (H 2 O) 3, etc. In solutions, long chains of associates of this type can appear:

because the oxygen atom has two lone pairs of electrons.

The presence of hydrogen bonds explains the high boiling temperatures of water, alcohols, and carboxylic acids. Due to hydrogen bonds, water is characterized by such high melting and boiling temperatures compared to H 2 E (E = S, Se, Te). If there were no hydrogen bonds, then water would melt at –100 °C and boil at –80 °C. Typical cases of association are observed for alcohols and organic acids.

Hydrogen bonds can occur both between different molecules and within a molecule if this molecule contains groups with donor and acceptor abilities. For example, it is intramolecular hydrogen bonds that play the main role in the formation of peptide chains, which determine the structure of proteins. H-bonds affect the physical and chemical properties of a substance.

Atoms of other elements do not form hydrogen bonds , since the forces of electrostatic attraction of opposite ends of dipoles of polar bonds (O-H, N-H, etc.) are rather weak and act only at short distances. Hydrogen, having the smallest atomic radius, allows such dipoles to get so close that the attractive forces become noticeable. No other element with a large atomic radius is capable of forming such bonds.

3.3.6 Intermolecular interaction forces (van der Waals forces). In 1873, the Dutch scientist I. Van der Waals suggested that there are forces that cause attraction between molecules. These forces were later called van der Waals forces – the most universal type of intermolecular bond. The energy of the van der Waals bond is less than the hydrogen bond and amounts to 2–20 kJ/∙mol.

Depending on the method of occurrence, forces are divided into:

1) orientational (dipole-dipole or ion-dipole) - occur between polar molecules or between ions and polar molecules. As polar molecules approach each other, they orient themselves so that the positive side of one dipole is oriented toward the negative side of the other dipole (Figure 10).

Figure 10 - Orientation interaction

2) induction (dipole - induced dipole or ion - induced dipole) - arise between polar molecules or ions and non-polar molecules, but capable of polarization. Dipoles can affect non-polar molecules, turning them into indicated (induced) dipoles. (Figure 11).

Figure 11 - Inductive interaction

3) dispersive (induced dipole - induced dipole) - arise between non-polar molecules capable of polarization. In any molecule or atom of a noble gas, fluctuations in electrical density occur, resulting in the appearance of instantaneous dipoles, which in turn induce instantaneous dipoles in neighboring molecules. The movement of instantaneous dipoles becomes consistent, their appearance and decay occur synchronously. As a result of the interaction of instantaneous dipoles, the energy of the system decreases (Figure 12).

Figure 12 - Dispersion interaction

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 chloride 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: crystal cell is the skeleton of matter, 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 nonpolar 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.

As examples of covalent bonding chemical elements we can cite molecules of carbon dioxide (CO2) and carbon monoxide (CO), 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. 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.