Literature review. Hydrogen bonding The most important oxidizing and reducing agents

Complex connections. Werner's theory. Role in a living organism.

Dissociation of complex compounds. Instability constant of complex ions.

Chemical bonding in complex compounds (examples).

In crystalline complex compounds with charged complexes, the connection between the complex and outer-sphere ions ionic, connections between the remaining particles of the outer sphere – intermolecular(including hydrogen ones). In most complex particles there are bonds between the central atom and the ligands covalent. All of them or part of them are formed according to the donor-acceptor mechanism (as a consequence - with a change in formal charges). In the least stable complexes (for example, in aqua complexes of alkali and alkaline earth elements, as well as ammonium), the ligands are held by electrostatic attraction. Bonding in complex particles is often called donor-acceptor or coordination bonding.

Redox reactions. Types of redox reactions.

Types of redox reactions:

1) Intermolecular- reactions in which oxidizing and reducing atoms are found in molecules of different substances, for example:

H 2 S + Cl 2 → S + 2HCl

2) Intramolecular- reactions in which oxidizing and reducing atoms are found in molecules of the same substance, for example:

2H 2 O → 2H 2 + O 2

3) Disproportionation(auto-oxidation-self-healing) - reactions in which the same element acts both as an oxidizing agent and as a reducing agent, for example:

Cl 2 + H 2 O → HClO + HCl

4)Reproportionation- reactions in which one oxidation state is obtained from two different oxidation states of the same element, for example:

NH 4 NO 3 → N 2 O + 2H 2 O

The most important oxidizing and reducing agents. Redox duality.

Restorers	Oxidizing agents
Metals	Halogens
Hydrogen	Potassium permanganate (KMnO 4)
Coal	Potassium manganate (K 2 MnO 4)
Carbon(II) monoxide (CO)	Manganese (IV) oxide (MnO 2)
Hydrogen sulfide (H 2 S)	Potassium dichromate (K 2 Cr 2 O 7)
Sulfur(IV) oxide (SO2)	Potassium chromate (K 2 CrO 4)
Sulfurous acid H 2 SO 3 and its salts	Nitric acid (HNO 3)
Hydrohalic acids and their salts	Sulfuric acid (H 2 SO 4) conc.
Metal cations in lower oxidation states: SnCl 2, FeCl 2, MnSO 4, Cr 2 (SO 4) 3	Copper(II) oxide (CuO)
Nitrous acid HNO 2	Lead(IV) oxide (PbO2)
Ammonia NH 3	Silver oxide (Ag 2 O)
Hydrazine NH 2 NH 2	Hydrogen peroxide (H 2 O 2)
Nitric oxide (II) (NO)	Iron(III) chloride (FeCl 3)
Cathode during electrolysis	Berthollet salt (KClO 3)
Metals	Anode during electrolysis

Hydrogen bonds are not unique to water. They form readily between any electronegative atom (usually oxygen or nitrogen) and a hydrogen atom covalently bonded to another electronegative atom in the same or another molecule (Figure 4-3). Hydrogen atoms covalently bonded to highly electronegative atoms such as oxygen always carry partial positive charges and are therefore capable of forming hydrogen bonds, whereas hydrogen atoms covalently bonded to carbon atoms that are not electronegative do not carry partial positive charges and, therefore, are unable to form hydrogen bonds. It is this difference that is the reason that butyl alcohol in the molecule of which one of the hydrogen atoms is bonded to oxygen and can thus form a hydrogen bond with another molecule of butyl alcohol has a relatively high boiling point (+117 ° C). On the contrary, butane, which is not capable of forming intermolecular hydrogen bonds, since all the hydrogen atoms in its molecules are bonded to carbon, has a low boiling point (- 0.5 ° C).

Some examples of biologically important hydrogen bonds are shown in Fig. 4-4.

Rice. 4-3. Hydrogen bonds. In this type of bond, the hydrogen atom is unevenly distributed between two electronegative atoms. And to which hydrogen is bonded covalently serves as a hydrogen donor, and the electronegative atom of another molecule serves as an acceptor. In biological systems, the electronegative atoms involved in the formation of hydrogen bonds are oxygen and nitrogen; carbon atoms take part in the formation of hydrogen bonds only in rare cases. The distance between two electronegative agoms connected by a hydrogen bond varies from 0.26 to 0.31 nm. Common types of hydrogen bonds are shown below.

One of the characteristic features of hydrogen bonds is that they are strongest in cases where the mutual orientation of the molecules connected to each other provides the maximum energy of electrostatic interaction (Fig. 4-5). In other words, a hydrogen bond is characterized by a certain orientation and, as a result, is capable of holding both molecules or groups associated with it in a certain mutual orientation. Below we will see that it is precisely this property of hydrogen bonds that contributes to the stabilization of strictly defined spatial structures characteristic of protein molecules and nucleic acids containing a large number of intramolecular hydrogen bonds (Chapters 7, 8 and 27).

Concept of hydrogen bond

A hydrogen atom bonded to a strongly electronegative atom (oxygen, fluorine, chlorine, nitrogen) can interact with the lone electron pair of another strongly electronegative atom of this or another molecule to form a weak additional bond - a hydrogen bond. In this case, a balance can be established

Picture 1.

The appearance of a hydrogen bond is predetermined by the exclusivity of the hydrogen atom. The hydrogen atom is much smaller than other atoms. The electron cloud formed by it and the electronegative atom is strongly shifted towards the latter. As a result, the hydrogen nucleus remains weakly shielded.

The oxygen atoms of the hydroxyl groups of two molecules of carboxylic acids, alcohols or phenols can come close together due to the formation of hydrogen bonds.

The positive charge on the nucleus of a hydrogen atom and the negative charge on another electronegative atom attract each other. The energy of their interaction is comparable to the energy of the previous bond, so the proton is bound to two atoms at once. The bond to a second electronegative atom may be stronger than the original bond.

A proton can move from one electronegative atom to another. The energy barrier for such a transition is insignificant.

Hydrogen bonds are among the chemical bonds of medium strength, but if there are many such bonds, then they contribute to the formation of strong dimeric or polymeric structures.

Example 1

Formation of a hydrogen bond in the $\alpha $-helical structure of deoxyribonucleic acid, diamond-like structure of crystalline ice, etc.

The positive end of the dipole in the hydroxyl group is at the hydrogen atom, so a bond can be formed through the hydrogen to anions or electronegative atoms containing lone pairs of electrons.

In almost all other polar groups, the positive end of the dipole is located inside the molecule and is therefore difficult to access for binding. In carboxylic acids $(R=RCO)$, alcohols $(R=Alk)$, phenols $(R=Ar)$, the positive end of the dipole $OH$ is located outside the molecule:

Examples of finding the positive end of the $C-O, S-O, P-O$ dipole inside a molecule:

Figure 2. Acetone, dimethyl sulfoxide (DMSO), hexamethylphosphortriamide (HMPTA)

Since there are no steric hindrances, hydrogen bonding is easy to form. Its strength is mainly determined by the fact that it is predominantly covalent in nature.

Typically, the presence of a hydrogen bond is indicated by a dotted line between the donor and acceptor, for example, in alcohols

Figure 3.

Typically, the distance between two oxygen atoms and a hydrogen bond is less than the sum of the van der Waals radii of the oxygen atoms. There must be mutual repulsion of the electron shells of oxygen atoms. However, the repulsive forces are overcome by the force of the hydrogen bond.

Nature of hydrogen bond

The nature of the hydrogen bond is electrostatic and donor-acceptor in nature. The main role in the formation of hydrogen bond energy is played by electrostatic interaction. Three atoms take part in the formation of an intermolecular hydrogen bond, which are located almost on the same straight line, but the distances between them are different. (the exception is the $F-H\cdots F-$ connection).

Example 2

For intermolecular hydrogen bonds in ice, $-O-H\cdots OH_2$, the $O-H$ distance is $0.097$ nm, and the $H\cdots O$ distance is $0.179$ nm.

The energy of most hydrogen bonds lies in the range of $10-40$ kJ/mol, and this is much less than the energy of a covalent or ionic bond. It can often be observed that the strength of hydrogen bonds increases with increasing acidity of the donor and basicity of the proton acceptor.

Importance of intermolecular hydrogen bond

Hydrogen bonding plays a significant role in the manifestations of the physical and chemical properties of a compound.

Hydrogen bonds have the following effects on compounds:

Intramolecular hydrogen bonds

In cases where closure of a six-membered or five-membered ring is possible, intramolecular hydrogen bonds are formed.

The presence of intramolecular hydrogen bonds in salicylic aldehyde and o-nitrophenol is the reason for the difference in their physical properties from the corresponding meta- And pair- isomers.

$o$-Hydroxybenzaldehyde or salicylic aldehyde $(A)$ and $o$-nitrophenol (B) do not form intermolecular associates, therefore they have lower boiling points. They are poorly soluble in water, since they do not participate in the formation of intermolecular hydrogen bonds with water.

Figure 5.

$o$-Nitrophenol is the only one of the three isomeric representatives of nitrophenols that is capable of steam distillation. This property is the basis for its isolation from a mixture of nitrophenol isomers, which is formed as a result of the nitration of phenols.

Hydrogen bonds are a specific bond that is created by the H atom, which is found in the groups OH, NH, FH, ClH and sometimes SH, and H bonds these groups with the valence-saturated atoms N2, O2 and F.

Hydrogen bonds determine the structure and properties of water, as the most important and basic solvent in biological systems. Hydrogen bonds are involved in the formation of macromolecules, biopolymers, as well as bonds with small molecules.

Uwater = 4-29 kJ/mol

The main contribution to hydrogen bonds comes from electrostatic interactions, but they are not limited to them. A proton moves along a straight line connecting electronegative atoms and experiences different influences from these atoms.

This graph is a special case, the relationship between N-H...N and N...H-N. R is the distance between interacting particles. 2 free energy minima are located near the first or second interacting N atom.

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  Hydrogen communications–specific connection, which is created by the H atom, which is located in the groups OH, NH, FH, ClH and sometimes SH, and H binds these groups to the valence-saturated atoms N2, O2 and F.


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  Hydrogen connection And her role V biological systems. Hydrogen communications–specific connection, which is created by the H atom that is in the group.


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  Hydrogen connection And her role V biological systems.
  She built in the form of a network of protein fibrillar molecules, among which a significant role plays alpha actinin.


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  Hydrogen connection And her role V biological systems. Hydrogen communications–specific connection


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  Hydrogen connection And her role V biological systems. Hydrogen communications–specific connection, which is created by the H atom, which is in the OH groups, ... more ».


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  Hydrogen connection And her role V biological systems. Hydrogen communications–specific connection, which is created by the H atom, which is in the OH groups, ... more ».


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  Role V biological systems.
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  2) intermolecular, if the EA and EV atoms are in different molecules. Intramolecular hydrogen communications play the most important biological role, since they determine, for example, the helical structure of polymeric protein molecules.


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  Shuttle transfer mechanisms hydrogen. home role The TCA cycle is the formation of a large amount of ATP.
  In this transport system hydrogen from cytoplasmic NAD is transferred to mitochondrial NAD, therefore 3 ATP molecules are formed in mitochondria and...


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  Role diffusion in substance transfer processes V biological systems.
  Intermolecular and intramolecular hydrogen connection Chemical communications in molecules there is usually a very pro... more ».

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The content of the article

HYDROGEN BONDING(H-bond) is a special type of interaction between reactive groups, with one of the groups containing a hydrogen atom prone to such interaction. Hydrogen bonding is a global phenomenon that spans all of chemistry. Unlike ordinary chemical bonds, the H-bond does not appear as a result of targeted synthesis, but arises under suitable conditions itself and manifests itself in the form of intermolecular or intramolecular interactions.

Features of hydrogen bonding.

A distinctive feature of the hydrogen bond is its relatively low strength, its energy is 5–10 times lower than the energy of a chemical bond. In terms of energy, it occupies an intermediate position between chemical bonds and van der Waals interactions, those that hold molecules in the solid or liquid phase.

In the formation of an H-bond, the electronegativity of the atoms participating in the bond plays a decisive role - the ability to attract electrons of a chemical bond from the partner atom participating in this bond. As a result, a partial negative charge d- appears on atom A with increased electronegativity, and a positive charge d+ appears on the partner atom, and the chemical bond is polarized: A d- –H d+.

The resulting partial positive charge on the hydrogen atom allows it to attract another molecule, also containing an electronegative element, thus, electrostatic interactions make the main contribution to the formation of the H-bond.

The formation of an H-bond involves three atoms, two electronegative (A and B) and the hydrogen atom H located between them; the structure of such a bond can be represented as follows: B···H d+ –A d- (a hydrogen bond is usually indicated by a dotted line ). Atom A, chemically bonded to H, is called a proton donor (Latin donare - give, donate), and B is its acceptor (Latin acceptor - receiver). More often than not, there is no true “donation” and H remains chemically bound to A.

There are not many donor atoms A that supply H for the formation of H-bonds, practically only three: N, O and F, while at the same time the set of acceptor atoms B is very wide.

The very concept and term “hydrogen bond” was introduced by W. Latimer and R. Rodebush in 1920 in order to explain the high boiling points of water, alcohols, liquid HF and some other compounds. Comparing the boiling temperatures of related compounds H 2 O, H 2 S, H 2 Se, and H 2 Te, they noticed that the first member of this series - water - boils much higher than it followed from the pattern formed by the rest members of the series. From this pattern it followed that water should boil 200 ° C lower than the observed true value.

Exactly the same deviation is observed for ammonia in a series of related compounds: NH 3, H 3 P, H 3 As, H 3 Sb. Its true boiling point (–33°C) is 80°C higher than expected.

When a liquid boils, only van der Waals interactions are destroyed, those that hold the molecules in the liquid phase. If the boiling temperatures are unexpectedly high, then, consequently, the molecules are additionally bound by some other forces. In this case, these are hydrogen bonds.

Similarly, the increased boiling point of alcohols (compared to compounds that do not contain an -OH group) is the result of the formation of hydrogen bonds.

Currently, spectral methods (most often infrared spectroscopy) provide a reliable way to detect H-bonds. The spectral characteristics of AN groups connected by hydrogen bonds differ markedly from those cases when such a bond is absent. In addition, if structural studies show that the distance between the B – H atoms is less than the sum of the van der Waals radii, then the presence of an H bond is considered to be established.

In addition to the increased boiling point, hydrogen bonds also manifest themselves during the formation of the crystalline structure of a substance, increasing its melting point. In the crystal structure of ice, H-bonds form a three-dimensional network, with water molecules arranged in such a way that the hydrogen atoms of one molecule are directed towards the oxygen atoms of neighboring molecules:

Boric acid B(OH) 3 has a layered crystal structure, each molecule is connected by hydrogen bonds to three other molecules. The packing of molecules in a layer forms a parquet pattern assembled from hexagons:

Most organic substances are insoluble in water; when this rule is violated, it is most often the result of the interference of hydrogen bonds.

Oxygen and nitrogen are the main donors of protons; they take on the function of atom A in the triad B···H d+ –A d- discussed earlier. They, most often, act as acceptors (atom B). Thanks to this, some organic substances containing O and N as atom B can dissolve in water (the role of atom A is played by oxygen in water). Hydrogen bonds between organic matter and water help to “pull apart” the molecules of the organic matter, transferring it into an aqueous solution.

There is a rule of thumb: if an organic substance contains no more than three carbon atoms per oxygen atom, then it is easily soluble in water:

Benzene is very slightly soluble in water, but if we replace one CH group with N, we get pyridine C 5 H 5 N, which is miscible with water in any ratio.

Hydrogen bonds can also manifest themselves in non-aqueous solutions, when a partial positive charge appears on hydrogen, and nearby there is a molecule containing a “good” acceptor, usually oxygen. For example, chloroform HCCl 3 dissolves fatty acids, and acetylene HCєCH is soluble in acetone:

This fact has found important technical application; acetylene under pressure is very sensitive to slight shocks and explodes easily, and its solution in acetone under pressure is safe to handle.

Hydrogen bonds play an important role in polymers and biopolymers. In cellulose, the main component of wood, hydroxyl groups are located in the form of side groups of a polymer chain assembled from cyclic fragments. Despite the relatively weak energy of each individual H-bond, their interaction throughout the polymer molecule leads to such powerful intermolecular interaction that the dissolution of cellulose becomes possible only when using an exotic highly polar solvent - Schweitzer's reagent (ammonia complex of copper hydroxide).

In polyamides (nylon, nylon) H-bonds arise between carbonyl and amino groups >C=O···H–N

This leads to the formation of crystalline regions in the polymer structure and an increase in its mechanical strength.

The same thing happens in polyurethanes, which have a structure close to polyamides:

NH-C(O)O-(CH 2) 4 -OC(O)-NH-(CH 2) n -NH-C(O)O-

The formation of crystalline regions and subsequent strengthening of the polymer occurs due to the formation of H-bonds between carbonyl and amino groups >C=O···H–N<.>

In a similar way, parallelly laid polymer chains in proteins are united, but H-bonds also provide protein molecules with a different way of packing - in the form of a spiral, while the turns of the helix are secured by the same hydrogen bonds that arise between the carbonyl and amino groups:

The DNA molecule contains all the information about a specific living organism in the form of alternating cyclic fragments containing carbonyl and amino groups. There are four types of such fragments: adenine, thymine, cytosine and guanine. They are located in the form of lateral pendants along the entire DNA polymer molecule. The order of alternation of these fragments determines the individuality of each living creature. When paired, the interaction of carbonyl C=O and amino groups of NH, as well as amino groups of NH and nitrogen atoms not containing hydrogen, creates H-bonds; it is they that hold two DNA molecules in the form of the well-known double spirals:

Complexes of some transition metals are prone to forming H-bonds (as proton acceptors); Complexes of metals of groups VI–VIII are most likely to participate in H-bonding. In order for such a bond to arise in some cases, the participation of a powerful proton donor, for example, trifluoroacetic acid, is necessary. At the first stage (see figure below), an H-bond occurs with the participation of the iridium metal atom (complex I), which plays the role of acceptor B.

Then, when the temperature decreases (from room temperature to –50° C), the proton passes to the metal and the usual M–H bond appears. All transformations are reversible; depending on the temperature, the proton can move either to the metal or to its donor - the acid anion.

In the second stage, the metal (complex II) accepts a proton, and with it a positive charge, and becomes a cation. A common ionic compound is formed (like NaCl). However, having passed to the metal, the proton retains its constant attraction to various acceptors, in this case to the acid anion. As a result, an H-bond appears (marked with asterisks), further tightening the ion pair:

The hydrogen atom can participate in the role of atom B, that is, a proton acceptor in the case when a negative charge is concentrated on it, this is realized in metal hydrides: M d+ –H d-, compounds containing a metal – hydrogen bond. If a metal hydride reacts with a proton donor of moderate strength (for example, fluorinated rubs-butanol), then an unusual dihydrogen bridge arises, where hydrogen forms an H-bond with itself: M d+ –H d- ···H d+ –A d- :

In the complex shown, wedge-shaped lines with solid filling or cross hatching indicate chemical bonds directed to the vertices of the octahedron.

Mikhail Levitsky