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 valency. 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:
CCl, CO, CN, OH, CMg.
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 accessible 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. Fluorine acquires exactly the same electronic configuration by accepting one electron donated by sodium. As a result, forces of electrostatic attraction 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, and electrically conductive in aqueous solutions. 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 the 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 is important to note that H-N bonds 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).
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Figure 10 - Orientation interaction |
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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).
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Figure 11 - Inductive interaction |
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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).
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Figure 12 - Dispersion interaction |
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Chemical bonds are understood as various types of interactions that determine the stable existence of di- and polyatomic compounds: molecules, ions, crystalline substances. By its nature, a chemical bond is the interaction between positively charged nuclei and negatively charged electrons, as well as electrons with each other. The main types of chemical bonds: covalent, ionic, metallic. To describe covalent bonds, two methods are used - the valence bond method (VBC) and the molecular orbital method (MMO).
The BC method is based on the following provisions:
1. Only unpaired electrons of two atoms with oppositely directed spins (exchange mechanism of bond formation), or an electron pair of one atom, participate in the formation of a covalent bond
Donor and free orbital of another atom - acceptor (donor-acceptor mechanism).
2. A covalent chemical bond between two atoms arises as a result of overlapping atomic orbitals with the formation of electron pairs (sharing of two electrons).
According to the theory of valence bonds, a covalent bond is directed towards maximum overlap of atomic orbitals of interacting atoms.
The geometric (spatial) structure of a molecule consisting of more than two atoms is determined by the relative arrangement of atomic orbitals involved in the formation of chemical bonds. The AB 2 molecule can have a linear 
, or angular structure (a). The AB 3 molecule can have the shape of a regular triangle (b), a trigonal pyramid (c). Molecule AB 4 – tetrahedral shape (d).
A) 
b) 
V) 
G) 
The spatial structure of the molecule is determined by the type of hybridization of the valence orbitals of the central atom and the number of lone electron pairs contained in the valence electron layer.
Example 1. Describe the structure of molecules from the standpoint of the valence bond method: a) PH 3, b) BBr 3. Which atomic orbitals are involved in the formation of chemical bonds? Indicate the type of hybridization (if hybridization occurs). What is the spatial structure of these molecules?
Solution. a) Formation of a PH 3 molecule.
Let's write the electronic formulas of the atoms that form the PH 3 molecule in the ground (normal) state: 15 P 1s 2 2s 2 2p 6 3s 2 3p 3 ; 1 N 1s 1
Electron graphic formulas of the external level of these atoms:
In this molecule, the central atom is the phosphorus atom, in which only p orbitals, located on the same sublevel and having the same shape and the same energy, participate in the formation of three chemical bonds. Consequently, there is no hybridization in the PH 3 molecule.
To visualize the valence schemes, you can use the following method. Electrons located in the outer electron layer are indicated by dots located around the chemical symbol of the atom. The electrons shared by two atoms are shown by dots placed between their chemical symbols; a double or triple bond is indicated by two or three pairs of common points, respectively. Using these notations, the formation of the PH 3 molecule can be represented as follows:
H
ê
This scheme can be written differently: H – P – H, where each pair of electrons connecting two atoms corresponds to one line representing a covalent bond in the structural formulas.
Option 1
1. Determine the type of chemical bond in the compounds N₂, KF, HF, NH₃ and H₂S. Write the structural and electronic formulas of the compounds NH₃ and HF.
2. Draw the electronic formulas of a neutral lithium atom and ion. How do the structures of these particles differ?
Li: 1s2 2s1 – neutral lithium atom
Lithium cation (gave up one electron): Li+: 1s2 2s0
3. Determine the type of crystal lattice characteristic of each of the following substances: potassium chloride, graphite, sugar, iodine, diamond.
KCl is an ionic lattice, atomic, sugar is molecular, iodine is molecular, diamond is atomic.
Option 2
1. From the given formulas of substances, write down only the formulas of compounds with a covalent polar bond: CO₂, PH₃, H₂, OF₂, O₂, KF, NaCl.
CO2, PH3, OF2
2. Make up electronic formulas for the molecules of chlorine Cl₂, hydrogen sulfide H₂S and phosphine PH₃. 
3. Using specific examples, compare the physical properties of substances with molecular and crystal lattices. 
Option 3
1. Determine the type of chemical bond in the compounds SO₃, NCl₃, ClF₃, Br₂, H₂O and NaCl.
2. Make up electronic formulas for the molecules of iodine I₂, water and methane CH₄. 
3. Using specific examples, show how some physical properties of substances depend on the type of their crystal lattice. 
Option 4
1. From the given formulas of substances, write down only the formulas of compounds with a covalent nonpolar bond: I₂, HCl, O₂, NH₃, H₂O, N₂, Cl₂, PH₃, NaNO₃.
I2, O2, N2, Cl2
№2 . Which of these compounds are formed by covalent polar bond and which are covalent non-polar bond: F2, NO, NH3, H2O, O2, CO2, Cl2, NaCl, SO2.
№3 . Molecules of what substances are formed polar covalent bond: NH3, H2O, N2, HCl, SO3, Al, Cl.
№4. Determine the type of chemical bond (polar covalent, nonpolar covalent, ionic) in the substances: NO, HF, NaF, O2, CO2, Cl2, FeCl3, NaCl, KBr, CaF2, H2, CH4.
№5 . Make up chemical formulas and indicate the electron density shift in the compounds: 1) sodium with nitrogen, 2) calcium with chlorine, 3) hydrogen with fluorine.
Many thanks to everyone who will help! With:
1) Make diagrams of the formation of chemical bonds in H2 and NH3 molecules. Indicate the type of chemical bond and the valency of the atom of each element.2) From the list, write down the formulas of substances with a covalent nonpolar bond: H2O, H2, H2S, HCI.CI2.
Write their electronic structural formulas.
3) Write the electronic structural formulas of the OF2 and H2O molecules. In which molecule is the chemical bond the most polar and towards which atom does the shared electron pair shift?
a) fluorine F2; b) hydrogen fluoride HF; c) hydrogen sulfide H2S.
Write the electronic and structural formulas of these molecules. Indicate the type of chemical bond and the valency of the atom of each element.
Thank you very much in advance )))
aluminum ions4) diatomic aluminum molecules
2. Indicate which products are formed when iron reacts with sulfuric acid1) Fe2(SO4)3, SO2, H2O2) FeSO4, SO2, H2O3) Fe2(SO4)3, H2O4) FeSO4, H2
3. Indicate which metal does not interact with water under normal conditions1) Na 2) Ba 3) Cu 4) K
4. Which of the statements are true?1) the chemical bond in the crystal lattice of a metal is metallic2) in reactions with non-metals, metals are oxidizing agents3) metals relatively easily give up external electrons and turn into cations4) all metals interact with solutions of acids
5. Arrange the substances in order of increasing metallic properties1) Mg 2) Ba 3) Be 4) Ca 5) Sr
HELP ME PLEASE!!! VERY URGENT THANK YOU IN ADVANCE
a) Choose a substance with an ionic crystal lattice
1. table salt 2. resin
3. naphthalene 4. diamond
b) Metallic crystal lattice of a simple substance
1.Se 2.Fe 3.F2 4.Te
c) Substance with a molecular type of crystal lattice
1.ionic 2.molecular 3.atomic 4.metallic
d) Forged, plastic, electrically and thermally conductive substances with a lattice type
1.ionic 2.molecular 3.atomic 4.metal
e) Check the correct statement
1. metal crystal lattices have substances with metallic bonds
2.Atomic lattices are called crystal lattices at the nodes where ions are located
3.dry ice, naphthalene, sugar - substances with ionic crystal lattices
4. table salt, quartz - fusible substances