The type of chemical bond formed between atoms is determined. chemical bond
The chemical bond, its types, properties, along with is one of the cornerstones of an interesting science called chemistry. In this article, we will analyze all aspects of chemical bonds, their significance in science, give examples and much more.
What is a chemical bond
In chemistry, a chemical bond is understood as the mutual adhesion of atoms in a molecule and, as a result of the force of attraction that exists between. It is through chemical bonds that the formation of various chemical compounds, this is the nature chemical bond.
Types of chemical bonds
The mechanism of formation of a chemical bond strongly depends on its type or type; in general, the following main types of chemical bond differ:
- Covalent chemical bond (which in turn can be polar or non-polar)
- Ionic bond
- connection
- chemical bond
similar people.
As for, a separate article is devoted to it on our website, and you can read in more detail at the link. Further, we will analyze in more detail all the other main types of chemical bonds.
Ionic chemical bond
The formation of an ionic chemical bond occurs when two ions with different charges are electrically attracted to each other. Ions usually with such chemical bonds are simple, consisting of one atom of the substance.
Diagram of an ionic chemical bond.
A characteristic feature of the ionic type of a chemical bond is its lack of saturation, and as a result, the most different amount oppositely charged ions. An example of an ionic chemical bond is the cesium fluoride compound CsF, in which the level of "ionicity" is almost 97%.
Hydrogen chemical bond
Long before the advent modern theory chemical bonds in it modern form scientists chemists have noticed that hydrogen compounds with non-metals have various amazing properties. Let's say the boiling point of water and together with hydrogen fluoride is much higher than it could be, here's a ready-made example of a hydrogen chemical bond.
The picture shows a diagram of the formation of a hydrogen chemical bond.
The nature and properties of the hydrogen chemical bond are due to the ability of the hydrogen atom H to form another chemical bond, hence the name of this bond. The reason for the formation of such a bond is the properties of electrostatic forces. For example, the general electron cloud in a hydrogen fluoride molecule is so shifted towards fluorine that the space around an atom of this substance is saturated with a negative electric field. Around the hydrogen atom, especially deprived of its only electron, everything is exactly the opposite, its electronic field is much weaker and, as a result, has a positive charge. And positive and negative charges, as you know, are attracted, in such a simple way, a hydrogen bond occurs.
Chemical bonding of metals
What chemical bond is typical for metals? These substances have their own type of chemical bond - the atoms of all metals are not arranged somehow, but in a certain way, the order of their arrangement is called the crystal lattice. The electrons of different atoms form a common electron cloud, while they weakly interact with each other.
This is what a metallic chemical bond looks like.
Any metal can serve as an example of a metallic chemical bond: sodium, iron, zinc, and so on.
How to determine the type of chemical bond
Depending on the substances taking part in it, if a metal and a non-metal, then the bond is ionic, if two metals, then it is metallic, if two non-metals, then it is covalent.
Properties of chemical bonds
In order to compare different chemical reactions different quantitative characteristics are used, such as:
- length,
- energy,
- polarity,
- the order of the links.
Let's analyze them in more detail.
The bond length is the equilibrium distance between the nuclei of atoms that are connected by a chemical bond. Usually measured experimentally.
The energy of a chemical bond determines its strength. In this case, energy refers to the force required to break a chemical bond and separate atoms.
The polarity of a chemical bond shows how much the electron density is shifted towards one of the atoms. The ability of atoms to shift their electron density towards themselves, or speaking plain language Pulling the blanket over you is called electronegativity in chemistry.
Crystals.
There are four types of chemical bonds: ionic, covalent, metallic and hydrogen.
Ionic chemical bond
Ionic chemical bond - this is a bond formed due to the electrostatic attraction of cations to anions.
As you know, the most stable is such an electronic configuration of atoms, in which there will be 8 electrons on the outer electronic level, like atoms of noble gases (or 2 for the first energy level). In chemical interactions, atoms tend to acquire just such a stable electronic configuration and often achieve this either as a result of the addition of valence electrons from other atoms (reduction process), or as a result of giving up their valence electrons (oxidation process). Atoms that have attached "foreign" electrons turn into negative ions, or anions. Atoms that donate their electrons turn into positive ions, or cations. It is clear that electrostatic attraction forces arise between anions and cations, which will keep them near each other, thereby carrying out an ionic chemical bond.
Since cations form mainly metal atoms, and anions - non-metal atoms, it is logical to conclude that this type of bond is typical for compounds of typical metals (elements of the main subgroups of groups I and II, except for magnesium and beryllium Be) with typical non-metals (elements of the main subgroup VII group). A classic example is the formation of alkali metal halides (fluorides, chlorides, etc.). For example, consider the scheme for the formation of an ionic bond in sodium chloride:
Two oppositely charged ions, bound by attractive forces, do not lose their ability to interact with oppositely charged ions, as a result of which compounds with an ionic crystal lattice are formed. Ionic compounds are solid, strong, refractory substances with high temperature melting.
Solutions and melts of most ionic compounds are electrolytes. This type of bond is characteristic of hydroxides of typical metals and many salts of oxygen-containing acids. However, when an ionic bond is formed, an ideal (complete) transition of electrons does not occur. An ionic bond is an extreme case of a covalent polar bond.
In an ionic compound, ions are presented as if in the form of electric charges with spherical symmetry of the electric field, which equally decreases with increasing distance from the charge center (ion) in any direction. Therefore, the interaction of ions does not depend on the direction, that is, the ionic bond, in contrast to the covalent bond, will be non-directional.
An ionic bond also exists in ammonium salts, where there are no metal atoms (their role is played by the ammonium cation).
covalent chemical bond
covalent bond(from the Latin "with" jointly and "vales" valid) is carried out by an electron pair belonging to both atoms. Formed between atoms of non-metals.
The electronegativity of non-metals is quite large, so that during the chemical interaction of two non-metal atoms, the complete transfer of electrons from one to the other (as in the case) is impossible. In this case, electron pooling is necessary to perform.
As an example, let's discuss the interaction of hydrogen and chlorine atoms:
H 1s 1 - one electron
Cl 1s 2 2s 2 2 p6 3 s2 3 p5 - seven electrons in the outer level
Each of the two atoms lacks one electron in order to have a complete outer electron shell. And each of the atoms allocates “for common use” one electron. Thus, the octet rule is satisfied. The best way to represent this is with the Lewis formulas:
Formation of a covalent bond
The shared electrons now belong to both atoms. The hydrogen atom has two electrons (its own and the shared electron of the chlorine atom), and the chlorine atom has eight electrons (its own plus the shared electron of the hydrogen atom). These two shared electrons form a covalent bond between the hydrogen and chlorine atoms. The particle formed when two atoms bond is called molecule.
Non-polar covalent bond
A covalent bond can form between two the same atoms. For example:
This diagram explains why hydrogen and chlorine exist as diatomic molecules. Thanks to the pairing and socialization of two electrons, it is possible to fulfill the octet rule for both atoms.
In addition to single bonds, a double or triple covalent bond can be formed, as, for example, in oxygen O 2 or nitrogen N 2 molecules. Nitrogen atoms each have five valence electrons, so three more electrons are required to complete the shell. This is achieved by sharing three pairs of electrons, as shown below:
Covalent compounds are usually gases, liquids, or relatively low-melting solids. One of the rare exceptions is diamond, which melts above 3,500°C. This is due to the structure of diamond, which is a continuous lattice of covalently bonded carbon atoms, and not a collection of individual molecules. In fact, any diamond crystal, regardless of its size, is one huge molecule.
A covalent bond occurs when the electrons of two nonmetal atoms join together. The resulting structure is called a molecule.
Polar covalent bond
In most cases, two covalently bound atom have different electronegativity and shared electrons do not belong equally to two atoms. Most time they are closer to one atom than to another. In a molecule of hydrogen chloride, for example, the electrons that form a covalent bond are located closer to the chlorine atom, since its electronegativity is higher than that of hydrogen. However, the difference in the ability to attract electrons is not so great that there is a complete transfer of an electron from a hydrogen atom to a chlorine atom. Therefore, the bond between hydrogen and chlorine atoms can be considered as something between an ionic bond (complete electron transfer) and a nonpolar one. covalent bond(symmetrical arrangement of a pair of electrons between two atoms). The partial charge on atoms is denoted by the Greek letter δ. Such a connection is called polar covalent bond, and the hydrogen chloride molecule is said to be polar, that is, it has a positively charged end (hydrogen atom) and a negatively charged end (chlorine atom).
The table below lists the main types of bonds and examples of substances:
Exchange and donor-acceptor mechanism of covalent bond formation
1) Exchange mechanism. Each atom gives one unpaired electron into a common electron pair.
2) Donor-acceptor mechanism. One atom (donor) provides an electron pair, and another atom (acceptor) provides an empty orbital for this pair.
Fig.1. Orbital radii of elements (r a) and length of one-electron chemical bond (d)
The simplest one-electron chemical bond is created by a single valence electron. It turns out that one electron is able to hold two positively charged ions in a single whole. In a one-electron bond, the Coulomb repulsive forces of positively charged particles are compensated by the Coulomb forces of attraction of these particles to a negatively charged electron. The valence electron becomes common to the two nuclei of the molecule.
Examples of such chemical compounds are molecular ions: H 2 + , Li 2 + , Na 2 + , K 2 + , Rb 2 + , Cs 2 + :
A polar covalent bond occurs in heteronuclear diatomic molecules (Fig. 3). The bonding electron pair in a polar chemical bond is close to the atom with a higher first ionization potential.
Characterizing the spatial structure of polar molecules, the distance d between atomic nuclei can be approximately considered as the sum of the covalent radii of the corresponding atoms.
Characterization of some polar substancesThe shift of the binding electron pair to one of the nuclei of the polar molecule leads to the appearance of an electric dipole (electrodynamics) (Fig. 4).
The distance between the centers of gravity of positive and negative charges is called the length of the dipole. The polarity of the molecule, as well as the polarity of the bond, is estimated by the value of the dipole moment μ, which is the product of the length of the dipole l by the value of the electronic charge:
Multiple covalent bonds
Multiple covalent bonds are represented by unsaturated organic compounds containing double and triple chemical bonds. To describe the nature of unsaturated compounds, L. Pauling introduces the concepts of sigma and π bonds, hybridization of atomic orbitals.
Pauling's hybridization for two S- and two p-electrons allowed the directionality of chemical bonds to be explained, in particular the tetrahedral configuration of methane. To explain the structure of ethylene, it is necessary to isolate one p-electron from four equivalent Sp 3 - electrons of the carbon atom to form an additional bond, called the π-bond. In this case, the three remaining Sp 2 -hybrid orbitals are located in the plane at an angle of 120° and form the main bonds, for example, a flat ethylene molecule (Fig. 5).
AT new theory Pauling, all binding electrons became equal and equidistant from the line connecting the nuclei of the molecule. Pauling's theory of a bent chemical bond took into account the statistical interpretation of the wave function by M. Born, the Coulomb electron correlation of electrons. A physical meaning appeared - the nature of the chemical bond is completely determined by the electrical interaction of nuclei and electrons. The more bonding electrons, the smaller the internuclear distance and the stronger the chemical bond between carbon atoms.
Three-center chemical bond
Further development of ideas about the chemical bond was given by the American physical chemist W. Lipscomb, who developed the theory of two-electron three-center bonds and a topological theory that allows predicting the structure of some more boron hydrides (borohydrides).
An electron pair in a three-center chemical bond becomes common to three atomic nuclei. In the simplest representative of a three-center chemical bond - the molecular hydrogen ion H 3 +, an electron pair holds three protons in a single whole (Fig. 6).
Fig. 7. Diboran
The existence of boranes with their two-electron three-center bonds with "bridge" hydrogen atoms violated the canonical doctrine of valency. The hydrogen atom, previously considered a standard univalent element, turned out to be bound by identical bonds with two boron atoms and became formally a divalent element. The work of W. Lipscomb on deciphering the structure of boranes expanded the understanding of the chemical bond. The Nobel Committee awarded the William Nunn Lipscomb Prize in Chemistry in 1976 with the wording "For his investigations into the structure of boranes (borohydrites) which elucidate the problems of chemical bonds".
Multicenter chemical bond
Fig. 8. Ferrocene molecule
Fig. 9. Dibenzenechromium
Fig. 10. Uranocene
All ten bonds (C-Fe) in the ferrocene molecule are equivalent, the Fe-c internuclear distance is 2.04 Å. All carbon atoms in a ferrocene molecule are structurally and chemically equivalent, the length of each C-C connections 1.40 - 1.41 Å (for comparison, in benzene the C-C bond length is 1.39 Å). A 36-electron shell appears around the iron atom.
Chemical bond dynamics
The chemical bond is quite dynamic. Thus, a metallic bond is transformed into a covalent bond during a phase transition during the evaporation of the metal. The transition of a metal from a solid to a vapor state requires the expenditure of large amounts of energy.
In vapors, these metals consist practically of homonuclear diatomic molecules and free atoms. When metal vapor condenses, the covalent bond turns into a metal one.
The evaporation of salts with a typical ionic bond, such as alkali metal fluorides, leads to the destruction of the ionic bond and the formation of heteronuclear diatomic molecules with a polar covalent bond. In this case, the formation of dimeric molecules with bridging bonds takes place.
Characterization of the chemical bond in the molecules of alkali metal fluorides and their dimers.
During the condensation of vapors of alkali metal fluorides, the polar covalent bond is transformed into an ionic one with the formation of the corresponding crystal lattice of the salt.
The mechanism of the transition of a covalent to a metallic bond
Fig.11. Relationship between the orbital radius of an electron pair r e and the length of a covalent chemical bond d
Fig.12. Orientation of the dipoles of diatomic molecules and the formation of a distorted octahedral cluster fragment during the condensation of alkali metal vapors
Fig. 13. Body-centered cubic arrangement of nuclei in alkali metal crystals and a link
Disperse attraction (London forces) causes interatomic interaction and the formation of homonuclear diatomic molecules from alkali metal atoms.
The formation of a metal-metal covalent bond is associated with the deformation of the electron shells of interacting atoms - valence electrons create a binding electron pair, the electron density of which is concentrated in the space between the atomic nuclei of the resulting molecule. A characteristic feature of homonuclear diatomic molecules of alkali metals is the long length of the covalent bond (3.6-5.8 times the bond length in the hydrogen molecule) and the low energy of its rupture.
The specified ratio between re and d determines the uneven distribution of electric charges in the molecule - in the middle part of the molecule, the negative electric charge of the binding electron pair is concentrated, and at the ends of the molecule - positive electric charges two atomic cores.
The uneven distribution of electric charges creates conditions for the interaction of molecules due to orientational forces (van der Waals forces). Molecules of alkali metals tend to orient themselves in such a way that opposite electric charges appear in the neighborhood. As a result, attractive forces act between the molecules. Due to the presence of the latter, alkali metal molecules approach each other and are more or less firmly drawn together. At the same time, some deformation of each of them occurs under the action of closer located poles of neighboring molecules (Fig. 12).
In fact, the binding electrons of the original diatomic molecule, falling into the electric field of four positively charged atomic cores of alkali metal molecules, break off from the orbital radius of the atom and become free.
In this case, the bonding electron pair becomes common even for a system with six cations. The construction of the crystal lattice of the metal begins at the cluster stage. In the crystal lattice of alkali metals, the structure of the connecting link is clearly expressed, having the shape of a distorted oblate octahedron - a square bipyramid, the height of which and the edges of the basis are equal to the value of the constant translational lattice a w (Fig. 13).
The value of the translational lattice constant a w of an alkali metal crystal significantly exceeds the length of the covalent bond of an alkali metal molecule, therefore it is generally accepted that the electrons in the metal are in a free state:
The mathematical construction associated with the properties of free electrons in a metal is usually identified with the "Fermi surface", which should be considered as a geometric place where electrons reside, providing the main property of the metal - to conduct electric current.
When comparing the process of condensation of alkali metal vapors with the process of condensation of gases, for example, hydrogen, salient feature in the properties of the metal. So, if weak intermolecular interactions appear during the condensation of hydrogen, then during the condensation of metal vapors, processes characteristic of chemical reactions occur. The condensation of metal vapor itself proceeds in several stages and can be described by the following procession: a free atom → a diatomic molecule with a covalent bond → a metal cluster → a compact metal with a metal bond.
The interaction of alkali metal halide molecules is accompanied by their dimerization. A dimeric molecule can be considered as an electric quadrupole (Fig. 15). At present, the main characteristics of alkali metal halide dimers (chemical bond lengths and bond angles) are known.
Chemical bond length and bond angles in dimers of alkali metal halides (E 2 X 2) (gas phase).
E 2 X 2 | X=F | X=Cl | X=Br | X=I | ||||
---|---|---|---|---|---|---|---|---|
d EF , Å | d ECl , Å | d EBr , Å | d EI , Å | |||||
Li 2 X 2 | 1,75 | 105 | 2,23 | 108 | 2,35 | 110 | 2,54 | 116 |
Na 2 X 2 | 2,08 | 95 | 2,54 | 105 | 2,69 | 108 | 2,91 | 111 |
K2X2 | 2,35 | 88 | 2,86 | 98 | 3,02 | 101 | 3,26 | 104 |
Cs 2 X 2 | 2,56 | 79 | 3,11 | 91 | 3,29 | 94 | 3,54 | 94 |
In the process of condensation, the action of orientational forces is enhanced, intermolecular interaction is accompanied by the formation of clusters, and then a solid. Alkali metal halides form crystals with a simple cubic and body-centered cubic lattice.
Lattice type and translational lattice constant for alkali metal halides.
In the process of crystallization, a further increase in the interatomic distance occurs, leading to the removal of an electron from the orbital radius of an alkali metal atom and the transfer of an electron to a halogen atom with the formation of the corresponding ions. Force fields of ions are evenly distributed in all directions in space. In this regard, in alkali metal crystals, the force field of each ion coordinates by no means one ion with the opposite sign, as it is customary to qualitatively represent the ionic bond (Na + Cl -).
In crystals of ionic compounds, the concept of simple two-ion molecules of the type Na + Cl - and Cs + Cl - loses its meaning, since the alkali metal ion is associated with six chloride ions (in a sodium chloride crystal) and eight chlorine ions (in a cesium chloride crystal. In this case, all interionic distances in crystals are equidistant.
Notes
- Handbook of inorganic chemistry. Constants of inorganic substances. - M .: "Chemistry", 1987. - S. 124. - 320 p.
- Lidin R.A., Andreeva L.L., Molochko V.A. Handbook of inorganic chemistry. Constants of inorganic substances. - M .: "Chemistry", 1987. - S. 132-136. - 320 s.
- Gankin V.Yu., Gankin Yu.V. How chemical bonds are formed and how chemical reactions proceed. - M .: publishing group "Border", 2007. - 320 p. - ISBN 978-5-94691296-9
- Nekrasov B.V. General chemistry course. - M .: Goshimizdat, 1962. - S. 88. - 976 p.
- Pauling L. The nature of the chemical bond / edited by Ya.K. Syrkin. - per. from English. M.E. Dyatkina. - M.-L.: Goshimizdat, 1947. - 440 p.
- Theoretical organic chemistry / ed. R.Kh. Freidlina. - per. from English. Yu.G. Bundel. - M .: Ed. foreign literature, 1963. - 365 p.
- Lemenovsky D.A., Levitsky M.M. Russian Chemical Journal (Journal of the Russian Chemical Society named after D.I. Mendeleev). - 2000. - T. XLIV, issue 6. - S. 63-86.
- Chemical Encyclopedic Dictionary / Ch. ed. I.L.Knunyants. - M .: Sov. Encyclopedia, 1983. - S. 607. - 792 p.
- Nekrasov B.V. General chemistry course. - M .: Goshimizdat, 1962. - S. 679. - 976 p.
- Lidin R.A., Andreeva L.L., Molochko V.A. Handbook of inorganic chemistry. Constants of inorganic substances. - M .: "Chemistry", 1987. - S. 155-161. - 320 s.
- Gillespie R. Geometry of molecules / per. from English. E.Z. Zasorina and V.S. Mastryukov, ed. Yu.A. Pentina. - M .: "Mir", 1975. - S. 49. - 278 p.
- Handbook of a chemist. - 2nd ed., revised. and additional - L.-M.: GNTI Chemical Literature, 1962. - T. 1. - S. 402-513. - 1072 p.
- Lidin R.A., Andreeva L.L., Molochko V.A. Handbook of inorganic chemistry. Constants of inorganic substances .. - M .: "Chemistry", 1987. - S. 132-136. - 320 s.
- Zieman J. Electrons in metals (introduction to the theory of Fermi surfaces). Advances in physical sciences .. - 1962. - T. 78, issue 2. - 291 p.
see also
- chemical bond- article from the Great Soviet Encyclopedia
- chemical bond- Chemport.ru
- chemical bond- Physical Encyclopedia
Covalent chemical bond, its varieties and formation mechanisms. Characteristics of a covalent bond (polarity and bond energy). Ionic bond. Metal connection. hydrogen bond
The doctrine of the chemical bond is the basis of all theoretical chemistry.
A chemical bond is such an interaction of atoms that binds them into molecules, ions, radicals, crystals.
There are four types of chemical bonds: ionic, covalent, metallic and hydrogen.
The division of chemical bonds into types is conditional, since all of them are characterized by a certain unity.
An ionic bond can be considered as the limiting case of a covalent polar bond.
A metallic bond combines the covalent interaction of atoms with the help of shared electrons and the electrostatic attraction between these electrons and metal ions.
In substances, there are often no limiting cases of chemical bonding (or pure chemical bonds).
For example, lithium fluoride $LiF$ is classified as an ionic compound. In fact, the bond in it is $80%$ ionic and $20%$ covalent. Therefore, it is obviously more correct to speak of the degree of polarity (ionicity) of a chemical bond.
In the hydrogen halide series $HF—HCl—HBr—HI—HAt$, the degree of bond polarity decreases, because the difference in the electronegativity values of the halogen and hydrogen atoms decreases, and in astatic hydrogen the bond becomes almost nonpolar $(EO(H) = 2.1; EO(At) = 2.2)$.
Different types of bonds can be contained in the same substances, for example:
- in bases: between the oxygen and hydrogen atoms in the hydroxo groups, the bond is polar covalent, and between the metal and the hydroxo group is ionic;
- in salts of oxygen-containing acids: between the non-metal atom and the oxygen of the acid residue - covalent polar, and between the metal and the acid residue - ionic;
- in salts of ammonium, methylammonium, etc.: between nitrogen and hydrogen atoms - covalent polar, and between ammonium or methylammonium ions and an acid residue - ionic;
- in metal peroxides (for example, $Na_2O_2$) the bond between oxygen atoms is covalent non-polar, and between the metal and oxygen it is ionic, and so on.
Different types of connections can pass one into another:
- during electrolytic dissociation in water of covalent compounds, a covalent polar bond passes into an ionic one;
- during the evaporation of metals, the metallic bond turns into a covalent non-polar, 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 an electron-nuclear interaction of atoms, accompanied by the release of energy.
Methods for the formation of a covalent bond. Characteristics of a covalent bond: bond length and energy
A covalent chemical bond is a bond that occurs between atoms due to the formation of common electron pairs.
The mechanism of formation of such a bond can be exchange and donor-acceptor.
I. exchange mechanism acts when atoms form common 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) $N_2$: three common electron pairs are formed between atoms in a nitrogen molecule:
II. Donor-acceptor mechanism Let us consider the formation of a covalent bond using 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 by the exchange mechanism, one - by the donor-acceptor mechanism.
Covalent bonds can be classified by the way in which the electron orbitals overlap, as well as by their displacement towards one of the bonded atoms.
Chemical bonds formed as a result of the overlap of electron orbitals along the 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 through lateral overlap:
Chemical bonds formed as a result of the "lateral" overlapping of electron orbitals outside the communication line, i.e. in two regions are called $π$ -bonds (pi-bonds).
By degree of bias common 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 ER - the property of pulling valence electrons towards themselves from other atoms. For example:
those. through a covalent non-polar bond, molecules of simple non-metal substances are formed. A covalent chemical bond between atoms of elements whose electronegativity differs is called polar.
The length and energy of a covalent bond.
characteristic covalent bond properties is its length and energy. Link length is the distance between the nuclei of atoms. A chemical bond is stronger the shorter its length. However, the measure of bond strength is binding energy, which is determined by the amount of energy required to break the 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 $0.074, 0.198$, and $0.109$ nm, respectively, and the binding energies are $436, 242$, and $946$ kJ/mol, respectively.
Ions. Ionic bond
Imagine that two atoms "meet": a metal atom of group I and a non-metal atom of group VII. A metal atom has a single electron in its outer energy level, while a non-metal atom lacks just one electron to complete its outer level.
The first atom will easily give up to the second its electron, which is far from the nucleus and weakly bound to it, and the second will give it a free place on its outer electronic level.
Then an 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 received electron. Such particles are called ions.
The chemical bond that occurs between ions is called ionic.
Consider the formation of this bond using the well-known sodium chloride compound (table salt) as an example:
The process of transformation of atoms into ions is shown in the diagram:
Such a transformation of atoms into ions always occurs during the interaction of atoms of typical metals and typical non-metals.
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 the numbers showing the number of atoms or ions in a molecule are called indexes.
metal connection
Let's get acquainted with how the atoms of metal elements interact with each other. Metals do not usually exist in the form of isolated atoms, but in the form of a piece, ingot, or metal product. What holds metal atoms together?
The atoms of most metals at the outer level contain a small number of electrons - $1, 2, 3$. These electrons are easily detached, and the atoms are converted into 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, and so on. Consequently, in the volume of a metal, atoms are continuously converted into ions and vice versa.
The bond in metals between ions by means of socialized electrons is called metallic.
The figure schematically shows the structure of a sodium metal fragment.
In this case, a small number of socialized electrons binds a large number of ions and atoms.
The metallic bond bears some resemblance to the covalent bond, since it is based on the sharing of outer electrons. However, in a covalent bond, the outer unpaired electrons of only two neighboring atoms are socialized, while in a metallic bond, all atoms take part in the socialization of these electrons. That is why crystals with a covalent bond are brittle, while those with a metal bond are, as a rule, plastic, electrically conductive, and have a metallic sheen.
The metallic bond is characteristic of both pure metals and mixtures of various metals - alloys that are in solid and liquid states.
hydrogen bond
A chemical bond between positively polarized hydrogen atoms of one molecule (or part of it) and negatively polarized atoms of strongly electronegative elements having unshared electron pairs ($F, O, N$ and less often $S$ and $Cl$), another molecule (or its parts) is called hydrogen.
The mechanism of hydrogen bond formation is partly electrostatic, partly donor-acceptor.
Examples of intermolecular hydrogen bonding:
In the presence of such a bond, even low molecular weight substances can under normal conditions be liquids (alcohol, water) or easily liquefying gases (ammonia, hydrogen fluoride).
Substances with a hydrogen bond have molecular crystal lattices.
Substances of molecular and non-molecular structure. Type of crystal lattice. The 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. A substance under given conditions 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. According to 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 molecules in such substances are very weak, much weaker than between atoms inside a molecule, and already at relatively low temperatures they break - the substance turns into a liquid and then into a gas (iodine sublimation). The melting and boiling points of substances consisting of molecules increase with increasing molecular weight.
Molecular substances include substances with an atomic structure ($C, Si, Li, Na, K, Cu, Fe, W$), among them there are metals and non-metals.
Consider the physical properties of 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.
The large sizes of atoms lead to a low density 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 an increase in the ordinal number of the element, because. the size of the atoms increases and the bonds weaken.
To substances non-molecular structures include ionic compounds. Most compounds of metals with non-metals 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
A substance, as is known, can exist in three states of aggregation: gaseous, liquid and solid.
Solids: amorphous and crystalline.
Consider how the features of chemical bonds affect 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 become fluid. In the amorphous state, for example, are plasticine and various resins.
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 frame is formed, called the 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 called crystal lattices, in the nodes of which there are ions. They are formed by substances with an ionic bond, which can bind both simple ions $Na^(+), Cl^(-)$, and complex $SO_4^(2−), OH^-$. Consequently, salts, some oxides and hydroxides of metals have ionic crystal lattices. For example, a sodium chloride crystal consists of alternating $Na^+$ positive ions and $Cl^-$ negative 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.
nuclear called crystal lattices, in the nodes of which there are individual atoms. In such lattices, the atoms are interconnected by very strong covalent bonds. An example of substances with this type of crystal lattice 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, practically insoluble.
Molecular crystal lattices.
Molecular called crystal lattices, at the nodes of which molecules are located. Chemical bonds in these molecules can be either polar ($HCl, H_2O$) or nonpolar ($N_2, O_2$). Despite the fact that the atoms within the molecules are bound by very strong covalent bonds, there are weak forces of intermolecular attraction between the molecules themselves. 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).
Metallic crystal lattices.
Substances with a metallic bond have metallic crystal lattices. At the nodes of such lattices there are atoms and ions (either atoms or ions, into which metal atoms easily turn, giving their outer electrons “for common use”). Such an internal structure of metals determines their characteristic physical properties: malleability, plasticity, electrical and thermal conductivity, and a characteristic metallic luster.