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

Figure 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. Electronic cloud, formed by it and the electronegative atom is strongly shifted towards the latter. As a result, the hydrogen nucleus remains weakly shielded.

Oxygen atoms of hydroxyl groups of two molecules carboxylic acids, alcohols or phenols may converge closely 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 crystal 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 the acidity of the donor and the basicity of the proton acceptor.

Importance of intermolecular hydrogen bond

Hydrogen bond plays a significant role in the manifestations of physical - chemical properties connections.

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 their difference physical properties from the relevant 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.

Chemical bonds in molecules are usually very strong, their energy is in the range of 100-150 kJ/mol. In addition, there are so-called hydrogen bonds, the strength of which is 10-40 kJ/mol. The length of these bonds is respectively 270-230 pm. A hydrogen bond between atoms Ea and Ev is the interaction carried out by a hydrogen atom connected to Ea or Ev by a chemical bond.

The image of a hydrogen bond in the general case has the form: Ea-H...Ev. It is obvious that the hydrogen bond is three-center, since three atoms take part in its formation. For such a bond to occur, it is necessary that the atoms of Ea and Ev have high electronegativity. These are atoms of the most negative elements: nitrogen (OEO = 3.0), oxygen (OEO = 3.5), fluorine (OEO = 4.0) and chlorine (OEO = 3.0). A hydrogen bond is formed by the combination of an ls-AO hydrogen and two 2r-AO atoms Ea and Ev. 2p orbitals are oriented along one straight line. Therefore, the hydrogen bond is linear. A hydrogen bond is called: 1) intramolecular, if the atoms Ea and Ev connected by this bond belong to the same molecule; 2) intermolecular, if the atoms Ea and Ev are in different molecules. Intramolecular hydrogen bonds play a vital role biological role, since they determine, for example, the helical structure polymer molecules proteins. In proteins, these are N-H...0 bonds between amino acid residues. Intermolecular hydrogen bonds are no less important. They are used to connect circuits nucleic acids, forming a double helix. There are two types of bonds between the nucleic bases N-H...N and N-H...0. Average kinetic energy thermal movement molecules has a value of the order of 3/2 RT. At a human body temperature of 37 °C (310 K) this is about 4 kJ/mol. The strength of hydrogen bonds is in the range of 10-40 kJ/mol. Therefore, they are strong enough to withstand constant impacts from surrounding molecules and ensure the shape of polymers remains constant. biological structures. At the same time, when active molecules strike, hydrogen bonds are periodically broken, then restored again, ensuring the occurrence of various life processes. The considered examples clearly illustrate the wider range of applications of the MO LCAO method than the BC method. Nevertheless, the BC method can be successfully used to predict the properties and structure of many substances, including complex compounds.

Question 37. Modern content of the concept of “complex compounds” (CS). Structure of the CS: central atom, ligands, complex ion, inner and outer sphere, coordination number of the central atom, dentation of ligands.

Complex connections- the most extensive and diverse class of compounds. Living organisms contain complex compounds of biogenic metals with proteins, amino acids, porphyrins, nucleic acids, carbohydrates, and macrocyclic compounds. The most important life processes occur with the participation of complex compounds. Some of them (hemoglobin, chlorophyll, hemocyanin, vitamin B12, etc.) play significant role in biochemical processes. Many drugs contain metal complexes. For example, insulin (zinc complex), vitamin B12 (cobalt complex), platinol (platinum complex), etc. Complex connections are called connections that exist both in crystalline state, and in solution, a feature of which is the presence of a central atom surrounded by ligands. Complex compounds can be thought of as complex compounds higher order, consisting of simple molecules capable of independent existence in solution. The structure of complex compounds, or simply complexes, was discovered by the Swiss scientist A. Werner in 1893. Many provisions of his theory formed the basis modern ideas about the structure of complexes. In the molecules of complex compounds, there is a central atom or ion M and n-molecules (or ions) L directly associated with it, called ligands. The central atom with the surrounding ligands form inner sphere MLn complex. Depending on the ratio of the total charge of the ligands and the complexing agent, the inner sphere may have a positive charge, for example, 3+, or a negative charge, for example, 3-, or a zero charge, for example, as for 0. In addition to the ligands, the complex may contain m other particles X , not directly associated with the central atom. Particles X form outer sphere complex, they neutralize the charge of the inner sphere, but are not covalently bound to the complexing agent. The general formula of a complex compound has the form: Xm, where M is the central atom; L - ligand; X - outer-sphere particle (molecule or ion); particles of the inner sphere are enclosed in square brackets. Complex compounds are often called coordination compounds. The number n of ligands is accordingly called the coordination number, and the inner sphere is called the coordination number. Central atom(complexing agent) - an atom or ion that occupies central position in a complex connection. The central atom coordinates the ligands, geometrically correctly positioning them in space. The role of a complexing agent is most often performed by particles that have free orbitals and a sufficiently large positive nuclear charge, and therefore can be electron acceptors. These are cations of transition elements. The most powerful complexing agents are elements of groups IB and VIIIB. Rarely, neutral atoms of d-elements and atoms of non-metals in varying degrees of oxidation act as complexing agents. The number of free atomic orbitals provided by the complexing agent determines its coordination number. The value of the coordination number depends on many factors, but is usually equal to twice the charge of the complexing ion. The strongest complexes are formed by d-elements. Complex compounds of Mn, Fe, Co, Cu, Zn, and Mo are especially important for human life. Amphoteric p-elements Al, Sn, Pb also form various complexes. Biogenic s-elements Na, K, Ca, Mg can form fragile complex compounds with ligands of a certain structure. Most often, the complexing agent is an atom of an element in a positive oxidation state. Negative conditional ions (i.e., atoms in a negative oxidation state) play the role of complexing agents relatively rarely. This is, for example, the nitrogen atom (-III) in the ammonium + cation, etc. The complexing atom may have a zero oxidation state. Thus, carbonyl complexes of nickel and iron, having the composition and , contain nickel(0) and iron(0) atoms. In a complex ion or neutral complex, ions, atoms, or simple molecules (L) are coordinated around the complexing agent. All these particles (ions or molecules) that have chemical bonds with the complexing agent are called ligands(ligands are donors of electron pairs). IN complex ions The 2- and 4- ligands are Cl- and CN- ions, and in the neutral complex the ligands are NH3 molecules and NCS- ions. Ligands, as a rule, are not bound to each other, and repulsive forces act between them. In some cases, intermolecular interaction of ligands with the formation of hydrogen bonds is observed. Ligands can be various inorganic and organic ions and molecules. The most important ligands are ions CN-, F-, Cl-, Br-, I-, NO2-, OH-, SO3S2-, C2O42-, CO32-, molecules H2O, NH3, CO, urea (NH2)2CO. The most important characteristic complexing agent is the number of chemical bonds it forms with ligands, or coordination number(CC). This characteristic of a complexing agent is determined mainly by the structure of its electronic shell and is determined by the valence capabilities of the central atom or conventional complexing ion. When the complexing agent coordinates monodentate ligands, the coordination number is equal to the number of attached ligands. And the number of polydentate ligands attached to the complexing agent is always less than value coordination number. The value of the coordination number of a complexing agent depends on its nature, degree of oxidation, the nature of the ligands and the conditions (temperature, nature of the solvent, concentration of the complexing agent and ligands, etc.) under which the complexation reaction occurs. The CN value can vary in various complex compounds from 2 to 8 and even higher. The most common coordination numbers are 4 and 6. Complexing elements with the oxidation state +II (ZnII, PtII, PdII, CuII, etc.) often form complexes in which they exhibit a coordination number of 4, such as 2+, 2-, 0. In aqua complexes, the coordination number of the complexing agent in the +II oxidation state is most often 6: 2+. Complexing elements with oxidation states +III and +IV (PtIV, AlIII, CoIII, CrIII, FeIII) in complexes, as a rule, have an CN of 6. For example, 3+, 3-. Complexing agents are known that have an almost constant coordination number in complexes of different types. These are cobalt(III), chromium(III) or platinum(IV) with CN 6 and boron(III), platinum(II), palladium(II), gold(III) with CN 4. However, most complexing agents have a variable coordination number. For example, for aluminum(III) CN 4 and CN 6 are possible in complexes - and -. Most often, the ligand is connected to the complexing agent through one of its atoms by a single two-center chemical bond. These types of ligands are called monodentate. Monodentate ligands include all halide ions, cyanide ions, ammonia, water, and others. Some common ligands such as water molecules H2O, hydroxide ion OH-, thiocyanate ion NCS-, amide ion NH2-, carbon monoxide CO in complexes are predominantly monodentate, although in some cases (in bridging structures) they become bidentate. There are a number of ligands that are almost always bidentate in complexes. These are ethylenediamine, carbonate ion, oxalate ion, etc. Each molecule or ion of a bidentate ligand forms two chemical bonds with the complexing agent in accordance with the features of its structure:

3 What chemical bond is called a hydrogen bond? What are the features of hydrogen bonding? What can be said about the strength of hydrogen bonds compared to covalent and ionic ones? What is the significance of hydrogen bonding in chemistry and biology?

A hydrogen bond is a chemical bond between hydrogen atoms and atoms of strongly electronegative elements (fluorine, oxygen, nitrogen). A hydrogen bond is usually formed between two neighboring molecules. For example, it is formed between molecules of water, alcohols, hydrogen fluoride, and ammonia.

This is a very weak bond - about 15-20 times weaker than a covalent bond. Thanks to it, some low-molecular substances form associates, which leads to an increase in the melting and boiling points of substances.

Abnormally high melting and boiling points are characteristic of water (if we consider group VI hydrogen compounds). All hydrogen compounds of group VI, except water, are gases.

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 connected covalent bond 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, which are not electronegative, do not carry a partial positive charge and therefore cannot 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 temperature boiling (+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 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 characteristic features hydrogen bonds is that they have the greatest strength in cases where the mutual orientation of interconnected molecules provides 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 large number intramolecular hydrogen bonds (Chapters 7, 8 and 27).

1)orientation(polar molecules, due to the electrostatic interaction of opposite ends of dipoles, are oriented in space so that the negative ends of the dipoles of some molecules are turned to the positive ends of the dipoles of other molecules)

2)induction(also observed in substances with polar molecules, but it is usually much weaker than the orientational one. A polar molecule can increase the polarity of a neighboring molecule. In other words, under the influence of the dipole of one molecule, the dipole of another molecule can increase, and a non-polar molecule can become polar)

3)dispersive(these forces interact between any atoms and molecules, regardless of their structure. They are caused by instantaneous dipole moments that occur in concert in a large group of atoms)

35. Hydrogen bond, its biological role.

36. Complex compounds. Werner's theory. Role in a living organism.

37. Dissociation of complex compounds. Instability constant of complex ions.

38. 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.

39. Redox reactions. Types of redox reactions.

Types of redox reactions:

1) Intermolecular- reactions in which oxidizing and reducing atoms are found in molecules 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

40. The most important oxidizing agents and reducing agents. Redox duality.