Protolytic theory of acids and bases. Acids and bases in organizational chemistry. Conjugate acid and conjugate base. Acid-base equilibria, examples. The influence of substituents in a molecule on acidity and basicity Modern ideas about acid

Acids and bases exhibit their properties only in the presence of each other. Not a single substance will give up a proton if there is no proton acceptor in the system - a base, and vice versa. i.e. they form conjugate acid-base pair in which the stronger the acid, the weaker its conjugate base, and the stronger the base, the weaker its conjugate acid.

An acid donating a proton becomes a conjugate base, and a base accepting a proton becomes a conjugate acid. An acid is usually denoted AN and a base is B

For example: HC1- H + + C1 -, HC1 - strong acid; C1 - ion - conjugate weak base;

CH 3 COOH - CH 3 COO - + H +, CH 3 COOH is a weak acid, and CH 3 COO - is a conjugate strong base ion.

The general view can be represented as follows:

Н+¦ : A + B Н:В+ + А:-

set of fundamentals of resistance resistance

something basic

We have already said that the acidic properties of compounds are revealed only in the presence of a base, and the basic properties - in the presence of an acid, i.e. in compounds there is a certain acid-base equilibrium, for the study of which H 2 O is used as a solvent. In relation to H 2 O as an acid or as a base, the acid-base properties of compounds are determined.

For weak electrolytes, acidity is quantified TO Rav a reaction that involves the transfer of H + from an acid to H 2 O as a base.

CH 3 COOH + H 2 O - CH 3 COO - + H 3 O +

which is the basic acid

CH 3 COO - - acetate ion, conjugate base;

H 3 O + - hydronium ion, conjugate acid.

Using the value of the equilibrium constant of this reaction and taking into account that the concentration of H 2 O is almost constant, we can determine the product K? called acidity constant TO acidity (K A).

The higher the K a, the stronger the acid. For CH 3 COOH K a = 1.75 10 -5. such small values are inconvenient in practical work, therefore K a is expressed through rK A (pK = -?g K A). For CH 3 COOH pK a = 4.75. The lower the pKa value, the stronger the acid.

The strength of the bases is determined by the pK value of BH +.

Acidic properties of organic compounds with hydrogen-containing functional groups (alcohols, phenols, thiols, carboxylic acids, amines).

Organic acids

In organic compounds, depending on the nature of the element with which H + is associated, the following acids are distinguished:

HE- acids (carboxylic acids, phenols, alcohols)

CH - acids (hydrocarbons and their derivatives)

NH- acids (amines, amides, imides)

SH- acids (thiols).

An acid center is an element and its associated hydrogen atom.

The strength of the acid will depend on anion stability, those. from the conjugate base, which is formed when H + is separated from the molecule. The more stable the anion, the higher the acidity of the compound.

Anion stability depends on a number of factors that contribute to charge delocalization. The higher the charge delocalization, the more stable the anion, the stronger the acidic properties.

Factors influencing the degree of delocalization:

1. Nature of the heteroatom in the acid site
2. Electronic effects of atoms of hydrocarbon radicals and their substituents
3. The ability of anions to solvation.
1. Dependence of acidity on the heteroatom.

The nature of a heteroatom is understood as its electronegativity (E.O.) and polarizability. The larger the (E.O.), the easier the heterolytic cleavage in the molecule occurs. In periods from left to right, as the charge of the nucleus increases, (E.O) increases, i.e. the ability of elements to hold a negative charge. As a result of the shift in electron density, the bond between atoms is polarized. The more electrons and the larger the radius of the atom, the further the electrons of the outer energy level are located from the nucleus, the higher the polarizability and the higher the acidity.

Example: CH- NH- OH- SH-

increase in E.O. and acidity

C, N, O - elements of the same period. E.O. grows over the period, acidity increases. In this case, polarizability will not affect acidity.

The polarizability of atoms in a period changes slightly, so the main factor determining acidity is E.O.

Now consider OH-SH-

increased acidity

O, S - are in the same group, the radius in the group increases from top to bottom, therefore the polarizability of the atom also increases, which leads to an increase in acidity. S has a larger atomic radius than O, so thiols exhibit stronger acidic properties compared to alcohols.

Compare three compounds: ethanol, ethanethiol and aminoaethanol:

H 3 C - CH 2 - HE, H 3 C - CH 2 - SH and H 3 C - CH 2 - N.H. 2

1. Let's compare by radical - they are the same;
2. By the nature of the heteroatom in the functional group: S and O are in the same group, but S has a larger atomic radius and higher polarizability, therefore ethanethiol has stronger acidic properties
3. Now let’s compare O and N. O has a higher EO, therefore the acidity of alcohols will be higher.
2. The influence of the hydrocarbon radical and the substituents present in it

It is necessary to draw students' attention to the fact that the compounds being compared must have the same acid center and the same solvent.

Electron-withdrawing (E.A.) substituents contribute to the delocalization of electron density, which leads to the stability of the anion and, accordingly, an increase in acidity.

Electron-donating (E.D.) substituents on the contrary, they contribute to the concentration of electron density in the acid center, which leads to a decrease in acidity and an increase in basicity.

For example: monohydric alcohols exhibit weaker acidic properties compared to phenols.

Example: H 3 C > CH 2 > OH

1. The acid center is the same
2. The solvent is the same

In monohydric alcohols, the electron density shifts from the hydrocarbon radical to the OH group, i.e. the radical exhibits a +I effect, then a large amount of electron density is concentrated on the OH group, as a result of which H + is more tightly bound to O and breaking the O-H bond is difficult, therefore monohydric alcohols exhibit weak acidic properties.

In phenol, on the contrary, the benzene ring is E.A., and the OH group is E.D.

Due to the fact that the hydroxyl group enters into common conjugation with the benzene ring, delocalization of the electron density occurs in the phenol molecule and the acidity increases, because conjugation is always accompanied by an increase in acidic properties.

An increase in the hydrocarbon radical in monocarboxylic acids also affects the change in acidic properties, and when substituents are introduced into the hydrocarbon, a change in acidic properties occurs.

Example: In carboxylic acids, upon dissociation, carboxylate ions are formed - the most stable organic anions.

In the carboxylate ion, the negative charge due to p, p-conjugation is distributed equally between the two oxygen atoms, i.e. it is delocalized and, accordingly, less concentrated, therefore, in carboxylic acids the acid center is stronger than in alcohols and phenols.

With an increase in the hydrocarbon radical, which plays the role of E.D. the acidity of monocarboxylic acids decreases due to the decrease in d + on the carbon atom of the carboxyl group. Therefore, in the homologous series of acids, formic acid is the strongest.

When introducing E.A. substituent in a hydrocarbon radical, for example chlorine - the acidity of the compound increases, because due to the -I effect, the electron density is delocalized and d + on the C atom of the carboxyl group increases, therefore, in this example, trichloroacetic acid will be the strongest.

3. Effect of solvent.

The interaction of molecules or ions of a solute with a solvent is called a process solvation. The stability of an anion significantly depends on its solvation in solution: the more the ion is solvated, the more stable it is, and the greater the solvation, the smaller the size of the ion and the less delocalization of the negative charge in it.

The terms “acid” and “base” are used to refer to two groups of compounds that have a set of diametrically opposed properties. In 1923, I. Brønsted and T. Lowry proposed a general protolytic theory of acids and bases. According to this theory, the following definitions correspond to the concepts of acid and base.

An acid is a molecule or ion capable of donating a hydrogen cation (proton). Acid is a proton donor.

A base is a molecule or ion capable of attaching a hydrogen cation (proton). The base is a proton acceptor.

An acid, giving up a proton, turns into a particle tending to accept it, which is called conjugate base:

The base, adding a proton, turns into a particle that tends to give it away, which is called conjugate acid:

The combination of an acid and its conjugate base or a base and its conjugate acid is called conjugate acid-base pairs.

The strength of an acid is determined by its ability to donate a proton, i.e. a strong acid is an active proton donor. The strength of acids in aqueous solutions decreases in the following order:

The strength of a base is determined by its ability to accept a proton, i.e. a strong base is an active proton acceptor. The strength of bases in aqueous solutions, i.e. their affinity for protons, decreases in the series:

Strong acids, easily donating a proton, are converted into conjugate bases, which do not readily accept a proton. Therefore, the dissociation of these acids is almost irreversible:

Weak acids, having difficulty giving up a proton, are converted into conjugate bases that actively accept a proton, which makes the dissociation of weak acids a reversible process, and the equilibrium is shifted towards the undissociated form:

Strong and weak bases behave in a similar way, transforming as a result of the reaction into the corresponding conjugate acids, i.e. in these cases there are also conjugate acid-base pairs:

Some substances are capable of acting in some reactions as a proton donor, donating it to compounds that have a higher affinity for the proton, and in others - as a proton acceptor, taking it away from compounds with lower affinity for the proton. Such substances are called ampholytes.

Ampholytes are molecules or ions that can both donate and accept a proton, and therefore enter into reactions characteristic of both acids and bases. Ampholyte exhibits the properties of an acid or a base depending on what substances it interacts with. A typical ampholyte is water, since as a result of its electrolytic dissociation, both a strong acid and a strong base are formed:

In addition, water interacts with acids, acting as a base, and with bases, exhibiting the properties of an acid:

Ampholytes are hydroxides of some metals (Zn, Al, Pb, Sn, Cr):

Ampholytes are hydroanions of polybasic acids, for example HC0 3 -, HP0 4 2- and H2PO4-.

Ampholytes are also compounds whose molecules contain two different acid-base groups, for example, biologically important a-amino acids. As a result of the transfer of a proton from the carboxyl group to the amino group, the α-amino acid molecule transforms from a tautomer* that does not contain charged groups into a tautomer having a bipolar-ionic (zwitterionic) structure. Thus, a-amino acids are characterized by prototropic tautomerism(Section 21.2.1).

In the crystalline state and in aqueous solutions, this equilibrium for α-amino acids is almost completely shifted towards the tautomer with a bipolar structure. Thus, for glycine in an aqueous solution, the content of a tautomer with a bipolar ionic structure is 223,000 times greater than that of another tautomer.

Due to this structural feature, α-amino acid molecules exhibit acidic properties due to the ammonium group (NH 3 +), and basic ones due to the ionized carboxyl group (-COO-), acting as ampholytes:

Like all ampholytes, α-amino acids are weak electrolytes.

According to the protolytic theory, acids, bases and ampholytes are protoliths, and the process of transfer of a proton from an acid to a base is called protolysis and is explained by the fact that these two substances have different affinities for protons. An acid-base interaction always involves two conjugate acid-base pairs, and proton transfer always occurs towards the formation of weaker acids, including conjugate ones. If the propensity of the reactants to interact with a proton is commensurate, then protolytic equilibrium.

Protolytic, or acid-base, balance established as a result of competition for a proton(H+) between bases of interacting conjugate acid-base pairs(NA, A- And VN + , V). The protolytic equilibrium always shifts towards the formation of a weaker acid:

The protolytic equilibrium can be schematically represented by the following diagram:

Proton transfer always occurs from a strong acid To to the anion of a weak acid, which is accompanied by the displacement of a weak acid from its salt under the influence of a stronger acid.

Protolytic equilibrium is observed during the ionization of weak electrolytes in water (Section 7.2). Thus, the ionization of a weak acid in aqueous solutions is a consequence of competition for a proton between the anion of the weak acid and water, which acts as a base, i.e., a proton acceptor. This process is reversible and is characterized by an equilibrium constant K a:

When a weak base interacts with water, the latter, acting as a proton donor, promotes the ionization of this base, which is of an equilibrium nature:

for weak electrolyte acids and bases is characterized by the values of acidity constants K a and basicity K b accordingly (section 7.2). If these constants characterize the protolytic interaction of water with an acid or base of one conjugate pair HA, A or BH +, B, then the product of the acidity constants K a i basicity Kb, components of a given pair is always equal to the ionic product of water Kn 2 o = 1 * 10 -14 (at 22 °C):

These expressions allow us to replace the basicity constant in the case of aqueous solutions Q or basicity indicator pKb weak base to acidity constant K a or acidity level pK a conjugate acid of this base. In practice, to characterize the protolytic properties of a compound, the value is usually used rK a. Thus, the strength of ammonia in water as a base (pKb, = 4.76) can be characterized by the acidity index of the ammonium ion NH4+, i.e. the conjugate acid: pK a (NH4+) = 14 - 4.76 - 9.24. Therefore, in the case of aqueous solutions there is no need for a special table of constants or indicators! basicity, a single acidity scale presented in table is sufficient. 8.1, where the properties of bases are characterized by the constant K a or acidity indicator pK a their conjugate acids. The strongest acid in aqueous solutions is the hydrogen cation H + (more precisely H3O +), and the strongest base is the OH- anion. Magnitude pK a quantitatively characterizes the strength of weak electrolytes in aqueous solutions.

A weak acid is weaker, the higher its pKa value. A weak base is weaker, the lower the pK value of its conjugate acid.

Meaning pK a is equal to the pH value of an aqueous solution in which a given weak electrolyte is ionized by 50%: since in this case [A - ] = [HA], then K a= [H + ] and pK a= pH. Thus, for acetic acid in its aqueous solution with pH = pK a (CH 3 COOH) = = 4.76, the equality [CH 3 COO-] = [CH 3 COOH] takes place, and for an aqueous solution of ammonia the equality = will be observed in a solution with pH = pK and (NH4+) = 9.24.

In addition, the value pK a allows you to determine the pH value of aqueous solutions, where a given weak acid HA is found predominantly (99% or more) in the form of an anion (A") - this will be in solutions with a pH > pK a + 2; or in the form of molecules (NA) - in solutions with pH< pK a - 2. In the interval ArH = pK a ± 2 a weak electrolyte in aqueous solutions exists in both ionized and non-ionized forms in the ratio [A-]/[HA] from 100: 1 to 1: 100 respectively.

The given relations allow, knowing the value pK a biosubstrate, determine what form it will be in at a particular pH value in the body’s aqueous systems. In addition, knowledge of the magnitude pK a of a weak electrolyte allows you to calculate the pH of aqueous solutions of this electrolyte if its concentration is known.

According to the Lowry-Bronsted theory, acids are substances that can donate a proton, bases are substances that accept a proton:

If B is a strong base, then it is a weak acid. With the help, you can characterize the degree of dissociation of an acid or conjugate acid. Along with the acidity constant, there is also the concept of the basicity constant and its corresponding

According to Lewis's theory, acids are compounds that can accept and bases can donate a pair of electrons.

In a broad sense, acids are compounds that supply a cation, in a particular case a proton, or accept a pair of electrons with an atom or group of atoms, etc.).

Bases accept a cation, in a particular case a proton, or provide a pair of electrons with an atom or group of atoms

The acidity or basicity of a substance is manifested in the process of interaction with another substance, in particular with a solvent, and therefore is relative.

Many substances have amphoteric properties. For example, water, alcohols and acids are capable of donating a proton when interacting with bases, and accepting it with acids. In the absence of acids and bases, the dual nature of such compounds is manifested in autoprotolysis:

Dissociation of an acid in a solvent means the transfer of a proton to the solvent:

In this regard, the strength of the acid is expressed by the dissociation constant, which is characteristic only of a given solvent. Proton transfer occurs only in highly ionizing and solvating solvents, such as water.

The degree of acid dissociation during the transition from an aqueous to an organic medium decreases by 4-6 orders of magnitude.

Strongly solvating and ionizing solvents neutralize the strength of acids, while non-polar and low-iolar solvents, interacting with them at the level of hydrogen bonds, have a differentiating effect. In the latter case, the differences in acid strength become more significant.

In inert, non-polar solvents the probability of proton abstraction is very low, although due to internal electronic effects the bond can be highly polarized. Under such conditions, acidic properties manifest themselves in the self-association of HA molecules or in association with proton acceptors - bases. In the latter case, the measure of acidity is the association constant with any base chosen as a standard. For example, the association constant of benzoic acid and diphenylguanidine in benzene is

The protonizing power of an acid is also expressed through the acidity function, which characterizes the state of equilibrium during complex formation of acids and bases in organic solvents. Indicators that change color depending on the strength of the acid are most often used as bases, which makes it possible to study the system by spectroscopic methods. In this case, it is important that the bands of associated free bases be identified in the spectrum.

So, in an introductory medium, acids and bases form solvated ions, in an organic medium - ion pairs and their associates.

Close in meaning to the concept of association is the concept of complexation: due to donor-acceptor and dative interactions, electron-donor-acceptor complexes, also called charge transfer complexes, can be formed from ions and molecules

Types of electron donors: I) compounds with heteroatoms. containing lone pairs of electrons, ethers, amines, sulfides, iodides, etc. For example: diethyl ether from lampn. ldmethyl sulfide triphenylphosphine propyl iodide

2) compounds containing - bonds, ethylenes, acetylenes, benzene and its derivatives, and other aromatic systems;

3) compounds capable of transferring electrons - alkanes, cycloalkanes:

Types of electron acceptors: 1) metal compounds containing a vacant orbital (K-orbital): halides, etc., metal ions

2) compounds capable of accepting a pair of electrons per vacant antibonding halogen, mixed halogens

3) compounds with -bonds with strongly electronegative substituents, participating in complex formation due to antibonding tetracyanethylene trinitrobenzene

Thus, either the donor can interact with the vacant acceptor, forming a new MO with a decrease in the energy of the system:

In organic chemistry, -complexes are of greatest importance and -complexes are characterized by instability constants, which are essentially the constants of their dissociation.

The dissociation and association constants of acids and bases still do not describe their properties fully enough. An important role in understanding many chemical processes, and in particular the phenomenon of catalysis, was played by the concept of hard and soft acids and bases (the principle

ZHMKO). In accordance with this concept, related acids and bases interact most effectively: a soft acid with a soft base, a hard one with a hard one.

Signs of hard acids and bases (Table 8): 1) small size of the ion or molecule; 2) high electronegativity; 3) localized charge; 4) low polarizability; 5) lowest vacant orbitals (LVO) of acids have high energy; 6) the highest filled orbitals (HFO) of the bases have low energy.

According to Lewis, the acidic and basic properties of organic compounds are assessed by their ability to accept or provide an electron pair and subsequently form a bond. An atom that accepts an electron pair is an electron acceptor, and a compound containing such an atom should be classified as an acid. The atom that provides an electron pair is an electron donor, and the compound containing such an atom is a base.

Specifically, Lewis acids can be an atom, molecule or cation: a proton, halides of elements of the second and third groups of the Periodic Table, transition metal halides - BF3, ZnCl2, AlCl3, FeCl3, FeBr3, TiCl4, SnCl4, SbCl5, metal cations, sulfuric anhydride - SO3, carbocation. Lewis bases include amines (RNH2, R2NH, R3N), alcohols ROH, ethers ROR

According to Brønsted-Lowry, acids are substances that can donate a proton, and bases are substances that can accept a proton.

Conjugate acid and base:

HCN (acidic) and CN- (base)

NH3 (base) and NH4+ (acid)

Acid-base (or protolytic) equilibrium is an equilibrium in which a proton (H+) is involved.

HCOOH + H 2 O D H 3 O + + HCOO -

acid2 base1

H 2 O + NH 3 D NH 4 + + OH - .

acid1 base2 conjugate conjugate

acid2 base1

7. Types of isomerism in organic chemistry. Structural, spatial and optical isomerism. Chirality. Layout and configuration. R, S, Z, E – nomenclatures.

There are two types of isomerism: structural and spatial (stereoisomerism). Structural isomers differ from each other by the order of bonds of atoms in the molecule, stereo-isomers - by the arrangement of atoms in space with the same order of bonds between them.

Structural isomerism: carbon skeleton isomerism, positional isomerism, isomerism of various classes of organic compounds (interclass isomerism).

Structural isomerism

Isomerism of the carbon skeleton

Positional isomerism is due to different positions of the multiple bond, substituent, and functional group with the same carbon skeleton of the molecule:

Spatial isomerism

Spatial isomerism is divided into two types: geometric and optical.

Geometric isomerism is characteristic of compounds containing double bonds and cyclic compounds. Since free rotation of atoms around a double bond or in a ring is impossible, the substituents can be located either on the same side of the plane of the double bond or ring (cis position) or on opposite sides (trans position).

Optical isomerism occurs when a molecule is incompatible with its image in a mirror. This is possible when the carbon atom in the molecule has four different substituents. This atom is called asymmetric.

CHIRALITY, the property of an object to be incompatible with its image in an ideal plane mirror.

Various spatial structures that arise due to rotation around simple bonds without violating the integrity of the molecule (without breaking chemical bonds) are called CONFORMATIONS.

Structure of alkanes. Sp3 is the state of carbon. Characteristics of C-C and C-H bonds. The principle of free rotation. Conformation. Methods of representation and nomenclature. Physical properties of alkanes.

All carbon atoms in alkane molecules are in the state sp 3 -hybridization, the angle between the C-C bonds is 109°28", therefore the molecules of normal alkanes with a large number of carbon atoms have a zigzag structure (zigzag). The length of the C-C bond in saturated hydrocarbons is 0.154 nm

The C-C bond is covalent non-polar. The C-H bond is covalent and weakly polar, since C and H are close in electronegativity.

Physical properties

Under normal conditions, the first four members of the homologous series of alkanes are gases, C 5 -C 17 are liquids, and starting from C 18 are solids. The melting and boiling points of alkanes of their density increase with increasing molecular weight. All alkanes are lighter than water and are insoluble in it, but they are soluble in non-polar solvents (for example, benzene) and are themselves good solvents.

· Melting and boiling points decrease from less branched to more branched.

· Gaseous alkanes burn with a colorless or pale blue flame and release large amounts of heat.

Rotation of atoms around the s-bond will not lead to its breaking. As a result of intramolecular rotation along C–C s-bonds, alkane molecules, starting with ethane C 2 H 6, can take on different geometric shapes.
Various spatial forms of a molecule that transform into each other by rotating around C–C s-bonds are called conformations or rotary isomers(conformers).
Rotational isomers of a molecule are its energetically unequal states. Their interconversion occurs quickly and constantly as a result of thermal movement. Therefore, rotary isomers cannot be isolated in individual form, but their existence has been proven by physical methods.

alkanes .
methane, ethane, propane, butane –an

9. Hydrocarbons. Classification. Saturated hydrocarbons of the methane series. Homologous series. Nomenclature. Isomerism. Radicals. Natural springs. Fischer-Tropsch synthesis. Methods of preparation (from alkenes, carboxylic acids, halogen derivatives, by the Wurtz reaction)

The general (generic) name of saturated hydrocarbons is alkanes .
The names of the first four members of the methane homologous series are trivial: methane, ethane, propane, butane . Starting from the fifth, the names are derived from Greek numerals with the addition of a suffix –an

Radicals (hydrocarbon radicals) also have their own nomenclature. Monovalent radicals are called alkyls and is designated by the letter R or Alk.
Their general formula is C n H 2n+ 1 .
The names of the radicals are made up of the names of the corresponding hydrocarbons by replacing the suffix -an to suffix -il(methane - methyl, ethane - ethyl, propane - propyl, etc.).
Divalent radicals are named by replacing the suffix -an on -iliden(exception is the methylene radical = CH 2).
Trivalent radicals have the suffix -ilidin

Isomerism. Alkanes are characterized by structural isomerism. If an alkane molecule contains more than three carbon atoms, then the order of their connection may be different. One of the isomers of butane ( n-butane) contains an unbranched carbon chain, and the other, isobutane, contains a branched one (isostructure).

The most important source of alkanes in nature is natural gas, mineral hydrocarbon raw materials - oil and associated petroleum gases.

Alkanes can be prepared by the Wurtz reaction, which involves the action of metallic sodium on monohalogen derivatives of hydrocarbons.
2CH 3 –CH 2 Br (ethyl bromide) + 2Na ––> CH 3 –CH 2 –CH 2 –CH 3 (butane) + 2NaBr

· From alkenes

C n H 2n + H 2 → C n H 2n+2

Fischer-Tropsch synthesis

nCO + (2n+1)H 2 → C n H 2n+2 + nH 2 O

The table shows that these hydrocarbons differ from each other in the number of groups - CH2-. Such a series of similar structures, having similar chemical properties and differing from each other in the number of these groups is called a homologous series. And the substances that make it up are called homologues.

Formula	Name
CH 4	methane
C2H6	ethane
C 3 H 8	propane
C4H10	butane
C4H10	isobutane
C5H12	pentane
C5H12	isopentane
C5H12	neopentane
C6H14	hexane
C 7 H 16	heptane
C 10 H 22	dean

10. Saturated hydrocarbons (alkanes). Chemical and physical properties: radical substitution reactions. Halogenation, nitridation, sulfochlorination, sulfoxidation. The concept of chain reactions.

Physical properties

According to Brønsted-Lowry, acids are substances that can donate a proton, and bases are substances that can accept a proton.

Conjugate acid and base:

HCN (acidic) and CN- (base)

NH3 (base) and NH4+ (acid)

Acid-base (or protolytic) equilibrium is an equilibrium in which a proton (H+) is involved.

HCOOH + H 2 O D H 3 O + + HCOO -

acid2 base1

H 2 O + NH 3 D NH 4 + + OH - .

acid1 base2 conjugate conjugate

acid2 base1

7. Types of isomerism in organic chemistry. Structural, spatial and optical isomerism. Chirality. Layout and configuration. R, S, Z, E – nomenclatures.

Structural isomerism: carbon skeleton isomerism, positional isomerism, isomerism of various classes of organic compounds (interclass isomerism).

Structural isomerism

Isomerism of the carbon skeleton

Positional isomerism is due to different positions of the multiple bond, substituent, and functional group with the same carbon skeleton of the molecule:

Spatial isomerism

Spatial isomerism is divided into two types: geometric and optical.

CHIRALITY, the property of an object to be incompatible with its image in an ideal plane mirror.

Various spatial structures that arise due to rotation around simple bonds without violating the integrity of the molecule (without breaking chemical bonds) are called CONFORMATIONS.