Electronegativity is the ability of atoms to displace electrons in their direction when forming a chemical bond. This concept was introduced by the American chemist L. Pauling (1932). Electronegativity characterizes the ability of an atom of a given element to attract a common electron pair in a molecule. Electronegativity values ​​determined by various methods differ from each other. In educational practice, they most often use relative rather than absolute values ​​of electronegativity. The most common is a scale in which the electronegativity of all elements is compared with the electronegativity of lithium, taken as one.

Among the elements of groups IA - VIIA:

electronegativity, as a rule, increases in periods (“from left to right”) with increasing atomic number, and decreases in groups (“from top to bottom”).

The patterns of changes in electronegativity among d-block elements are much more complex.

Elements with high electronegativity, the atoms of which have high electron affinity and high ionization energy, i.e., prone to the addition of an electron or the displacement of a pair of bonding electrons in their direction, are called nonmetals.

These include: hydrogen, carbon, nitrogen, phosphorus, oxygen, sulfur, selenium, fluorine, chlorine, bromine and iodine. According to a number of characteristics, a special group of noble gases (helium-radon) is also classified as nonmetals.

Metals include most of the elements of the Periodic Table.

Metals are characterized by low electronegativity, i.e., low ionization energy and electron affinity. Metal atoms either donate electrons to nonmetal atoms or mix pairs of bonding electrons from themselves. Metals have a characteristic luster, high electrical conductivity and good thermal conductivity. They are mostly durable and malleable.

This set of physical properties that distinguish metals from non-metals is explained by the special type of bond that exists in metals. All metals have a clearly defined crystal lattice. Along with atoms, its nodes contain metal cations, i.e. atoms that have lost their electrons. These electrons form a socialized electron cloud, the so-called electron gas. These electrons are in the force field of many nuclei. This bond is called metallic. The free migration of electrons throughout the volume of the crystal determines the special physical properties of metals.

Metals include all d and f elements. If from the Periodic Table you mentally select only blocks of s- and p-elements, i.e., elements of group A and draw a diagonal from the upper left corner to the lower right corner, then it turns out that non-metallic elements are located on the right side of this diagonal, and metallic ones - in the left. Adjacent to the diagonal are elements that cannot be unambiguously classified as either metals or non-metals. These intermediate elements include: boron, silicon, germanium, arsenic, antimony, selenium, polonium and astatine.

Ideas about covalent and ionic bonds played an important role in the development of ideas about the structure of matter, however, the creation of new physical and chemical methods for studying the fine structure of matter and their use showed that the phenomenon of chemical bonding is much more complex. It is currently believed that any heteroatomic bond is both covalent and ionic, but in different proportions. Thus, the concept of covalent and ionic components of a heteroatomic bond is introduced. The greater the difference in electronegativity of the bonding atoms, the greater the polarity of the bond. When the difference is more than two units, the ionic component is almost always predominant. Let's compare two oxides: sodium oxide Na 2 O and chlorine oxide (VII) Cl 2 O 7. In sodium oxide, the partial charge on the oxygen atom is -0.81, and in chlorine oxide -0.02. This effectively means that the Na-O bond is 81% ionic and 19% covalent. The ionic component of the Cl-O bond is only 2%.

List of used literature

1. Popkov V. A., Puzakov S. A. General chemistry: textbook. - M.: GEOTAR-Media, 2010. - 976 pp.: ISBN 978-5-9704-1570-2. [With. 35-37]
2. Volkov, A.I., Zharsky, I.M. Big chemical reference book / A.I. Volkov, I.M. Zharsky. - Mn.: Modern School, 2005. - 608 with ISBN 985-6751-04-7.

Of exceptionally great importance in biological systems is a special type of intermolecular interaction, the hydrogen bond, which occurs between hydrogen atoms chemically combined in one molecule and the electronegative atoms F, O, N, Cl, S belonging to another molecule. The concept of "hydrogen bond" was first introduced in 1920 by Latimer and Rodebush to explain the properties of water and other associated substances. Let's look at some examples of such a connection.

In paragraph 5.2 we talked about the pyridine molecule and it was noted that the nitrogen atom in it has two outer electrons with antiparallel spins that do not participate in the formation of a chemical bond. This "free" or "lone" pair of electrons will attract the proton and form a chemical bond with it. In this case, the pyridine molecule will go into an ionic state. If there are two pyridine molecules, they will compete to capture a proton, resulting in a compound

in which three dots indicate a new type of intermolecular interaction called hydrogen bonding. In this compound, the proton is closer to the left-handed nitrogen atom. With the same success, the proton may be closer to the right nitrogen atom. Therefore, the potential energy of a proton as a function of the distance to the right or left nitrogen atom at a fixed distance between them (approximately ) should be depicted by a curve with two minima. A quantum mechanical calculation of such a curve, carried out by Rhine and Harris, is shown in Fig. 4.

The quantum mechanical theory of the A-H...B hydrogen bond based on donor-acceptor interactions was one of the first to be developed by N. D. Sokolov. The reason for the bond is the redistribution of electron density between atoms A and B caused by the proton. Briefly, they say that a “lone pair” of electrons is shared. In fact, in

Rice. 4. Potential curve of proton energy as a function of the distance between the nitrogen atoms of two pyridine molecules.

Other electrons of molecules also participate in the formation of potential hydrogen bond curves, although to a lesser extent (see below).

Typical hydrogen bond energies range from 0.13 to 0.31 eV. It is an order of magnitude less than the energy of chemical covalent bonds, but an order of magnitude greater than the energy of van der Waals interactions.

The simplest intermolecular complex formed by hydrogen bonding is the complex. This complex has a linear structure. The distance between fluorine atoms is 2.79 A. The distance between atoms in a polar molecule is 0.92 A. When a complex is formed, an energy of about 0.26 eV is released.

With the help of hydrogen bonding, a water dimer is formed with a binding energy of about 0.2 eV. This energy is approximately one-twentieth the energy of the OH covalent bond. The distance between two oxygen atoms in the complex is approximately 2.76 A. It is less than the sum of the van der Waals radii of oxygen atoms, equal to 3.06 A. In Fig. Figure 5 shows the change in the electron density of water atoms calculated in the work during the formation of the complex. These calculations confirm that when a complex is formed, the distribution of electron density around all atoms of the reacting molecules changes.

The role of all atoms in the establishment of hydrogen bonds in the complex can also be judged by the mutual influence of two hydrogen bonds between the nitrogenous bases, thymine and adenine, that are part of the double helix of the DNA molecule. The location of the minima of the proton potential curves in two bonds reflects their mutual correlation (Fig. 6).

Along with the usual or weak hydrogen bond formed by hydrogen with an energy release of less than 1 eV, and characterized by a potential energy with two minima, hydrogen forms some complexes with a large energy release. For example, when creating a complex, an energy of 2.17 eV is released. This type of interaction is called strong

Rice. 5. Change in electron density around atoms in a complex formed by hydrogen bonds from two water molecules.

The charge of the electron is assumed to be equal to unity. In a free water molecule, the charge of 10 electrons is distributed so that near the oxygen atom there is a charge of 8.64, and at the hydrogen atoms

Rice. 6. Hydrogen bonds between nitrogenous bases: a - thymine (T) and adenip (A), which are part of the DNN molecules (arrows indicate the places of attachment of the bases to the chains of sugar and phosphoric acid molecules); - potential hydrogen bond curves; O - oxygen; - hydrogen; - carbon; - nitrogen.

hydrogen bond. When complexes with strong hydrogen bonds are formed, the configuration of the molecules changes significantly. The proton's potential energy has one relatively flat minimum located approximately at the center of the bond. Therefore, the proton is easily displaced. The easy displacement of the proton under the influence of an external field determines the high polarizability of the complex.

Strong hydrogen bonding does not occur in biological systems. As for the weak hydrogen bond, it is of decisive importance in all living organisms.

The exceptionally large role of hydrogen bonding in biological systems is due primarily to the fact that it determines the secondary structure of proteins, which is of fundamental importance for all life processes; with the help of hydrogen bonds, base pairs are held in DNA molecules and their stable structure in the form of double helices is ensured, and, finally, hydrogen bonds are responsible for the very unusual properties of water, which are important for the existence of living systems.

Water is one of the main components of all living things. Animal bodies are almost two-thirds water. The human embryo contains about 93% water during the first month. There would be no running water. Water serves as the main medium in which biochemical reactions occur in the cell. It forms the liquid part of blood and lymph. Water is necessary for digestion, since the breakdown of carbohydrates, proteins and fats occurs with the addition of water molecules. Water is released in the cell when proteins are built from amino acids. Physiological

Rice. 7. Ice structure. Each water molecule is connected by hydrogen bonds (three points) to four water molecules located at the vertices of the tetrahedron.

Rice. 8. Hydrogen bond in a dimer and “linear” hydrogen bond

the properties of biopolymers and many supramolecular structures (in particular, cell membranes) very significantly depend on their interaction with water.

Let's look at some properties of water. Each water molecule has a large electrical moment. Due to the high electronegativity of oxygen atoms, a water molecule can form hydrogen bonds with one, two, three, or four other water molecules. The result is relatively stable dimers and other polymer complexes. On average, each molecule in liquid water has four neighbors. The composition and structure of intermolecular complexes depend on water temperature.

Crystalline water (ice) has the most ordered structure at normal pressure and temperature below zero degrees Celsius. Its crystals have a hexagonal structure. The unit cell contains four water molecules. The cell structure is shown in Fig. 7. Around the central oxygen atom there are four other oxygen atoms located at the vertices of a regular tetrahedron at distances of 2.76 A. Each water molecule is connected to its neighbors by four hydrogen bonds. In this case, the angle between OH bonds in the molecule approaches the “tetrahedral” value of 109.1°. In a free molecule it is approximately 105°.

The structure of ice resembles that of diamond. However, in diamond there are chemical forces between the carbon atoms. A diamond crystal is a large molecule. Ice crystals are classified as molecular crystals. The molecules in a crystal retain essentially their individuality and hold each other together through hydrogen bonds.

Rice. 9. Experimental value of the shift in infrared vibration frequency in water during the formation of a hydrogen bond at an angle .

The ice lattice is very loose and contains many “voids”, since the number of nearest water molecules for each molecule (coordination number) is only four. When melting, the ice lattice is partially destroyed, at the same time some voids are filled and the density of water becomes greater than the density of ice. This is one of the main water anomalies. With further heating to 4° C, the compaction process continues. When heated above 4° C, the amplitude of anharmonic vibrations increases, the number of associated molecules in complexes (swarms) decreases, and the density of water decreases. According to rough estimates, the swarms at room temperature include about 240 molecules, at 37° C - about 150, at 45 and 100° C, 120 and 40, respectively.

The contribution of hydrogen bonding to the total energy of intermolecular interactions (11.6 kcal/mol) is about 69%. Due to hydrogen bonds, the melting points (0° C) and boiling points (100° C) of water differ significantly from the melting and boiling points of other molecular liquids, between the molecules of which only van der Waals forces act. For example, for methane these values ​​are respectively -186 and -161° C.

In liquid water, along with the remnants of the tetrahedral structure of ice, there are linear and cyclic dimers and other complexes containing 3, 4, 5, 6 or more molecules. It is important that the angle P formed between the OH bond and the hydrogen bond changes depending on the number of molecules in the cycle (Fig. 8). In a dimer this angle is 110°, in a five-membered ring it is 10°, and in a six-membered ring and hexagonal ice structure it is close to a bullet (“linear” hydrogen bond).

It turns out that the highest energy of one hydrogen bond corresponds to the angle. The energy of a hydrogen bond is proportional (Badger-Bauer rule) to the shift in the frequency of stretching infrared vibrations of the OH group in a water molecule compared to the vibration frequency of a free molecule. The maximum displacement is observed in the case of a “linear” hydrogen bond. In a water molecule in this case, the frequency decreases by , and the frequency decreases by . In Fig. Figure 9 shows a graph of the displacement ratio

frequency to maximum offset from angle . Consequently, this graph also characterizes the dependence of the hydrogen bond energy on the angle . This dependence is a manifestation of the cooperative nature of the hydrogen bond.

Multiple attempts have been made to theoretically calculate the structure and properties of water, taking into account hydrogen bonds and other intermolecular interactions. According to statistical physics, the thermodynamic properties of a system of interacting molecules located in volume V at constant pressure P in statistical equilibrium with a thermostat are determined through the partition function of states

Here V is the volume of the system; k - Boltzmann constant; T - absolute temperature; means that we need to take the trace of the statistical operator in curly brackets, where H is the quantum operator of the energy of the entire system. This operator is equal to the sum of the kinetic energy operators of the translational and rotational motions of molecules and the potential energy operator of the interaction of all molecules.

If all the eigenfunctions and the full energy spectrum E of the operator H are known, then (6.2) takes the form

Then the Gibbs free energy G of the system at pressure P and temperature T is determined by the simple expression

Knowing the Gibbs free energy, we find the total energy entropy volume.

Unfortunately, due to the complex nature of interactions between molecules in water (anisotropic dipole molecules, hydrogen bonds leading to complexes of variable composition, in which the energy of hydrogen bonds itself depends on the composition and structure of the complex, etc.), we cannot write the operator H in explicitly. Therefore, we have to resort to very large simplifications. Thus, Nameti and Scheraga calculated the partition function based on the fact that only five energy states of molecules in complexes can be taken into account, according to

with the number of hydrogen bonds they form (0, 1, 2, 3, 4) with neighboring molecules. Using this model, they even managed to show that the density of water is maximum at 4° C. However, later the authors themselves criticized the theory they developed, since it did not describe many experimental facts. Other attempts at theoretical calculations of the structure of water can be found in the review by Ben-Naim and Stillinger.

Due to the dipole nature of water molecules and the large role of hydrogen bonds, the interactions of water molecules with ions and neutral molecules in living organisms also play an extremely important role. Interactions leading to the hydration of ions and a special type of interaction called hydrophobic and hydrophilic will be discussed in the following sections of this chapter."

Speaking about the role of water in biological phenomena, it should be noted that all living organisms have very successfully adapted to a certain amount of hydrogen bonding between molecules. This is evidenced by the fact that the replacement of heavy water molecules has a very significant effect on biological systems. The solubility of polar molecules decreases, the speed of transmission of a nerve impulse decreases, the work of enzymes is disrupted, the growth of bacteria and fungi slows down, etc. Perhaps all these phenomena are due to the fact that the hydrogen interaction between molecules is stronger than the interaction between molecules. between the molecules of heavy water is indicated by its extremely high melting point (3.8 ° C) and high heat of fusion (1.51 kcal/mol). For ordinary water, the heat of fusion is 1.43 kcal/mol.

Link length - internuclear distance. The shorter this distance, the stronger the chemical bond. The length of a bond depends on the radii of the atoms forming it: the smaller the atoms, the shorter the bond between them. For example, the H-O bond length is shorter than the H-N bond length (due to less oxygen atom exchange).

An ionic bond is an extreme case of a polar covalent bond.

Metal connection.

The prerequisite for the formation of this type of connection is:

1) the presence of a relatively small number of electrons at the outer levels of atoms;

2) the presence of empty (vacant orbitals) on the outer levels of metal atoms

3) relatively low ionization energy.

Let's consider the formation of a metal bond using sodium as an example. The valence electron of sodium, which is located on the 3s sublevel, can relatively easily move through the empty orbitals of the outer layer: along 3p and 3d. When atoms come closer together as a result of the formation of a crystal lattice, the valence orbitals of neighboring atoms overlap, due to which electrons move freely from one orbital to another, establishing a bond between ALL atoms of the metal crystal.

At the nodes of the crystal lattice there are positively charged metal ions and atoms, and between them there are electrons that can move freely throughout the crystal lattice. These electrons become common to all atoms and ions of the metal and are called "electron gas". The bond between all positively charged metal ions and free electrons in the metal crystal lattice is called metal bond.

The presence of a metallic bond determines the physical properties of metals and alloys: hardness, electrical conductivity, thermal conductivity, malleability, ductility, metallic luster. Free electrons can carry heat and electricity, so they are the reason for the main physical properties that distinguish metals from non-metals - high electrical and thermal conductivity.

Hydrogen bond.

Hydrogen bond occurs between molecules that contain hydrogen and atoms with high EO (oxygen, fluorine, nitrogen). Covalent bonds H-O, H-F, H-N are highly polar, due to which an excess positive charge accumulates on the hydrogen atom, and an excess negative charge on the opposite poles. Between oppositely charged poles, forces of electrostatic attraction arise - hydrogen bonds.

Hydrogen bonds can be either intermolecular or intramolecular. The energy of a hydrogen bond is approximately ten times less than the energy of a conventional covalent bond, but nevertheless, hydrogen bonds play an important role in many physicochemical and biological processes. In particular, DNA molecules are double helices in which two chains of nucleotides are linked by hydrogen bonds. Intermolecular hydrogen bonds between water and hydrogen fluoride molecules can be depicted (by dots) as follows:

Substances with hydrogen bonds have molecular crystal lattices. The presence of a hydrogen bond leads to the formation of molecular associates and, as a consequence, to an increase in the melting and boiling points.

In addition to the listed main types of chemical bonds, there are also universal forces of interaction between any molecules that do not lead to the breaking or formation of new chemical bonds. These interactions are called van der Waals forces. They determine the attraction of molecules of a given substance (or various substances) to each other in liquid and solid states of aggregation.

Different types of chemical bonds determine the existence of different types of crystal lattices (table).

Substances consisting of molecules have molecular structure. These substances include all gases, liquids, as well as solids with a molecular crystal lattice, such as iodine. Solids with an atomic, ionic or metal lattice have non-molecular structure, they have no molecules.

Table

Feature of the crystal lattice	Lattice type
Molecular	Ionic	Nuclear	Metal
Particles at lattice nodes	Molecules	Cations and anions	Atoms	Metal cations and atoms
The nature of the connection between particles	Intermolecular interaction forces (including hydrogen bonds)	Ionic bonds	Covalent bonds	Metal connection
Bond strength	Weak	Durable	Very durable	Various strengths
Distinctive physical properties of substances	Low-melting or sublimating, low hardness, many soluble in water	Refractory, hard, brittle, many soluble in water. Solutions and melts conduct electric current	Very refractory, very hard, practically insoluble in water	High electrical and thermal conductivity, metallic luster, ductility.
Examples of substances	Simple substances - non-metals (in solid state): Cl 2, F 2, Br 2, O 2, O 3, P 4, sulfur, iodine (except silicon, diamond, graphite); complex substances consisting of non-metal atoms (except ammonium salts): water, dry ice, acids, non-metal halides: PCl 3, SiF 4, CBr 4, SF 6, organic substances: hydrocarbons, alcohols, phenols, aldehydes, etc.	Salts: sodium chloride, barium nitrate, etc.; alkalis: potassium hydroxide, calcium hydroxide, ammonium salts: NH 4 Cl, NH 4 NO 3, etc., metal oxides, nitrides, hydrides, etc. (compounds of metals with non-metals)	Diamond, graphite, silicon, boron, germanium, silicon oxide (IV) - silica, SiC (carborundum), black phosphorus (P).	Copper, potassium, zinc, iron and other metals
Comparison of substances by melting and boiling points.
	Due to the weak forces of intermolecular interaction, such substances have the lowest melting and boiling points. Moreover, the greater the molecular weight of the substance, the higher the t 0 pl. it has. Exceptions are substances whose molecules can form hydrogen bonds. For example, HF has a higher t0 pl. than HCl.	Substances have high t 0 pl., but lower than substances with an atomic lattice. The higher the charges of the ions that are located in the lattice sites and the shorter the distance between them, the higher the melting point of the substance. For example, t 0 pl. CaF 2 is higher than t 0 pl. KF.	They have the highest t 0 pl. The stronger the bond between the atoms in the lattice, the higher the t 0 pl. has substance. For example, Si has a lower t0 pl. than C.	Metals have different t0 pl.: from -37 0 C for mercury to 3360 0 C for tungsten.

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CONNECTION

Synonyms:

consistency, coherence, continuity, foldability, consistency, harmony, interaction, connection, articulation, concatenation, coupling, communication, means of communication, intercourse, communication, contact, association, relation, relationship, dependence, binding, ties, romance, connecting link, union, causation, public relations, tomba, intimate relationships, intrigue, correlation, duplex, umbilical cord, intercourse, bonding, religion, cohabitation, parataxis, connecting thread, continuity, adhesion, interconnectedness, correlation, conditionality, connection, kinship, putty, bond , cupids, affair, synapse, context, love, thread, mail, message, quadruplex. Ant. fragmentation

CONNECTION synonyms, what are they? CONNECTION, CONNECTION this is the meaning of the word CONNECTION, origin (etymology) CONNECTION, CONNECTION stress, word forms in other dictionaries

+ CONNECTION synonym - Dictionary of Russian synonyms 4

Ionic bond

(materials from the site http://www.hemi.nsu.ru/ucheb138.htm were used)

Ionic bonding occurs through electrostatic attraction between oppositely charged ions. These ions are formed as a result of the transfer of electrons from one atom to another. An ionic bond is formed between atoms that have large differences in electronegativity (usually greater than 1.7 on the Pauling scale), for example, between alkali metal and halogen atoms.

Let us consider the occurrence of an ionic bond using the example of the formation of NaCl.

From electronic formulas of atoms

Na 1s 2 2s 2 2p 6 3s 1 and

Cl 1s 2 2s 2 2p 6 3s 2 3p 5

It can be seen that to complete the outer level, it is easier for a sodium atom to give up one electron than to gain seven, and for a chlorine atom it is easier to gain one electron than to gain seven. In chemical reactions, the sodium atom gives up one electron, and the chlorine atom takes it. As a result, the electron shells of sodium and chlorine atoms are transformed into stable electron shells of noble gases (electronic configuration of the sodium cation

Na + 1s 2 2s 2 2p 6,

and the electronic configuration of the chlorine anion is

Cl – - 1s 2 2s 2 2p 6 3s 2 3p 6).

The electrostatic interaction of ions leads to the formation of a NaCl molecule.

The nature of the chemical bond is often reflected in the state of aggregation and physical properties of the substance. Ionic compounds such as sodium chloride NaCl are hard and refractory because there are powerful forces of electrostatic attraction between the charges of their “+” and “–” ions.

The negatively charged chlorine ion attracts not only “its” Na+ ion, but also other sodium ions around it. This leads to the fact that near any of the ions there is not one ion with the opposite sign, but several.

The structure of a crystal of sodium chloride NaCl.

In fact, there are 6 sodium ions around each chlorine ion, and 6 chlorine ions around each sodium ion. This ordered packing of ions is called an ionic crystal. If a single chlorine atom is isolated in a crystal, then among the sodium atoms surrounding it it is no longer possible to find the one with which the chlorine reacted.

Attracted to each other by electrostatic forces, the ions are extremely reluctant to change their location under the influence of external force or an increase in temperature. But if sodium chloride is melted and continued to be heated in a vacuum, it evaporates, forming diatomic NaCl molecules. This suggests that covalent bonding forces are never completely turned off.

Basic characteristics of ionic bonds and properties of ionic compounds

1. An ionic bond is a strong chemical bond. The energy of this bond is on the order of 300 – 700 kJ/mol.

2. Unlike a covalent bond, an ionic bond is non-directional because an ion can attract ions of the opposite sign to itself in any direction.

3. Unlike a covalent bond, an ionic bond is unsaturated, since the interaction of ions of the opposite sign does not lead to complete mutual compensation of their force fields.

4. During the formation of molecules with an ionic bond, complete transfer of electrons does not occur, therefore, one hundred percent ionic bonds do not exist in nature. In the NaCl molecule, the chemical bond is only 80% ionic.

5. Compounds with ionic bonds are crystalline solids that have high melting and boiling points.

6. Most ionic compounds are soluble in water. Solutions and melts of ionic compounds conduct electric current.

Metal connection

Metal crystals are structured differently. If you examine a piece of sodium metal, you will find that its appearance is very different from table salt. Sodium is a soft metal, easily cut with a knife, flattened with a hammer, it can be easily melted in a cup on an alcohol lamp (melting point 97.8 o C). In a sodium crystal, each atom is surrounded by eight other similar atoms.

Crystal structure of metallic Na.

The figure shows that the Na atom in the center of the cube has 8 nearest neighbors. But the same can be said about any other atom in a crystal, since they are all the same. The crystal consists of "infinitely" repeating fragments shown in this figure.

Metal atoms at the outer energy level contain a small number of valence electrons. Since the ionization energy of metal atoms is low, valence electrons are weakly retained in these atoms. As a result, positively charged ions and free electrons appear in the crystal lattice of metals. In this case, metal cations are located in the nodes of the crystal lattice, and electrons move freely in the field of positive centers, forming the so-called “electron gas”.

The presence of a negatively charged electron between two cations causes each cation to interact with this electron.

Thus, Metallic bonding is the bonding between positive ions in metal crystals that occurs through the attraction of electrons moving freely throughout the crystal.

Since the valence electrons in a metal are evenly distributed throughout the crystal, a metallic bond, like an ionic bond, is a non-directional bond. Unlike a covalent bond, a metallic bond is an unsaturated bond. A metal bond also differs from a covalent bond in strength. The energy of a metallic bond is approximately three to four times less than the energy of a covalent bond.

Due to the high mobility of the electron gas, metals are characterized by high electrical and thermal conductivity.

The metal crystal looks quite simple, but in fact its electronic structure is more complex than that of ionic salt crystals. There are not enough electrons in the outer electron shell of metal elements to form a full-fledged “octet” covalent or ionic bond. Therefore, in the gaseous state, most metals consist of monatomic molecules (i.e., individual atoms not connected to each other). A typical example is mercury vapor. Thus, the metallic bond between metal atoms occurs only in the liquid and solid state of aggregation.

A metallic bond can be described as follows: some of the metal atoms in the resulting crystal give up their valence electrons to the space between the atoms (for sodium this is...3s1), turning into ions. Since all the metal atoms in a crystal are the same, each has an equal chance of losing a valence electron.

In other words, the transfer of electrons between neutral and ionized metal atoms occurs without energy consumption. In this case, some electrons always end up in the space between the atoms in the form of “electron gas”.

These free electrons, firstly, hold the metal atoms at a certain equilibrium distance from each other.

Secondly, they give metals a characteristic “metallic shine” (free electrons can interact with light quanta).

Thirdly, free electrons provide metals with good electrical conductivity. The high thermal conductivity of metals is also explained by the presence of free electrons in the interatomic space - they easily “respond” to changes in energy and contribute to its rapid transfer in the crystal.

A simplified model of the electronic structure of a metal crystal.

******** Using the metal sodium as an example, let us consider the nature of the metallic bond from the point of view of ideas about atomic orbitals. The sodium atom, like many other metals, has a lack of valence electrons, but there are free valence orbitals. The single 3s electron of sodium is capable of moving to any of the free and close-in-energy neighboring orbitals. As atoms in a crystal come closer together, the outer orbitals of neighboring atoms overlap, allowing the electrons given up to move freely throughout the crystal.

However, the "electron gas" is not as disorderly as it might seem. Free electrons in a metal crystal are in overlapping orbitals and are to some extent shared, forming something like covalent bonds. Sodium, potassium, rubidium and other metallic s-elements simply have few shared electrons, so their crystals are fragile and fusible. As the number of valence electrons increases, the strength of metals generally increases.

Thus, metallic bonds tend to be formed by elements whose atoms have few valence electrons in their outer shells. These valence electrons, which carry out the metallic bond, are shared so much that they can move throughout the metal crystal and provide high electrical conductivity of the metal.

A NaCl crystal does not conduct electricity because there are no free electrons in the space between the ions. All electrons donated by sodium atoms are firmly held by chlorine ions. This is one of the significant differences between ionic crystals and metal ones.

What you now know about metallic bonding helps explain the high malleability (ductility) of most metals. Metal can be flattened into a thin sheet and drawn into wire. The fact is that individual layers of atoms in a metal crystal can slide one another relatively easily: the mobile “electron gas” constantly softens the movement of individual positive ions, shielding them from each other.

Of course, nothing like this can be done with table salt, although salt is also a crystalline substance. In ionic crystals, the valence electrons are tightly bound to the nucleus of the atom. The shift of one layer of ions relative to another brings ions of the same charge closer together and causes strong repulsion between them, resulting in the destruction of the crystal (NaCl is a fragile substance).

The shift of the layers of an ionic crystal causes the appearance of large repulsive forces between like ions and destruction of the crystal.

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