Conjugated pairs of acids and bases. Protolytic theory of acids and bases

Acids and bases exhibit their properties only in the presence of each other. Not a single substance will donate a proton if there is no proton acceptor - a base in the system, and vice versa. they form conjugated 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 donates a proton to become a conjugate base, and a base accepts a proton to become a conjugate acid. Acid is usually denoted as AN and base as B.

For example: HC1- H + + C1 -, HC1 is a strong acid; C1 - ion - conjugated weak base;

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

The general view can be represented as follows:

H+¦: A + B H:B+ + A:-

to-ta bases resist. resist.

to-ta basics

We have already said that the acidic properties of compounds are found 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 balance, for the study of which H 2 O is used as a solvent. With respect to H 2 O as an acid or as a base, the acid-base properties of the compounds are determined.

For weak electrolytes, acidity is quantified To equal a reaction that consists in 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 +

to-that basic base 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 practically constant, it is possible to determine the product K? called the acidity constant To acidity (K a).

The more K a, the stronger the acid. For CH 3 COOH K a \u003d 1.75 10 -5. such small values are inconvenient in practical work, therefore K a is expressed through RK a (рК = -?g К a). For CH 3 COOH pKa = 4.75. The smaller the pKa value, the stronger the acid.

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

Acid 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 detached from the molecule. The more stable the anion, the higher the acidity of the compound.

The stability of the anion 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 affecting the degree of delocalization:

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

The nature of a heteroatom is understood as its electronegativity (E.O.) and polarizability. The more (E.O.) the easier the heterolytic gap in the molecule is carried out. In periods from left to right, with an increase in the charge of the nucleus, (E.O) increases, i.e. the ability of elements to hold a negative charge. As a result of the displacement of the electron density, the bond between the 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. increases over time, acidity increases. In this case, the polarizability will not affect the acidity.

The polarizability of atoms in the period varies slightly, therefore, 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 than alcohols.

Compare three compounds: ethanol, ethanethiol and aminoethanol:

H 3 C - CH 2 - HE, H 3 C - CH 2 - SH and H 3 C - CH 2 - NH 2

1. 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, higher polarizability, therefore ethanethiol has stronger acidic properties
3. Now let's compare O and N. O has a higher EO, hence the acidity of alcohols will be higher.
2. Influence of the hydrocarbon radical and its substituents

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

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

Electron donating (ED) 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 + I effect, then a large amount of electron density is concentrated on the OH group, as a result of which H + is more firmly bonded to O and the breaking of 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 is included in the common p-p conjugation with the benzene ring, electron density delocalization occurs in the phenol molecule and acidity increases, tk. conjugation is always accompanied by an increase in acidic properties.

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

Example: in carboxylic acids, during 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 two oxygen atoms, i.e. it is delocalized and, accordingly, less concentrated; therefore, the acid center in carboxylic acids 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 a decrease in q + on the carbon atom of the carboxyl group. Therefore, in the homologous series of acids, formic acid is the strongest.

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

3. Influence of the solvent.

The interaction of molecules or ions of a solute with a solvent is called a process solvation. The stability of an anion essentially 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 in relation to two groups of compounds that have a set of diametrically opposed properties. In 1923, I. Bronsted and T. Lowry proposed a general protolytic theory of acids and bases. According to this theory, the concepts of acid and base correspond to the following definitions.

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

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

An acid, donating a proton, turns into a particle that seeks to accept it, which is called conjugate base:

The base, by attaching 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 conjugated 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 series:

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

Strong acids, easily donating a proton, turn into conjugate bases, which poorly attach a proton. Therefore, the dissociation of these acids proceeds almost irreversibly:

Weak acids, hard donating a proton, turn into conjugated 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, turning into the corresponding conjugate acids as a result of the reaction, i.e. in these cases there are also conjugated acid-base pairs:

Some substances are able to act in some reactions as a proton donor, donating it to compounds that have a higher proton affinity, and in others, as a proton acceptor, taking it away from compounds with a lower proton affinity. Such substances are called ampholytes.

Ampholytes are molecules or ions capable of both donating and accepting 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. Water is a typical ampholyte, 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, showing 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. The a-amino acid molecule, as a result of the transfer of a proton from the carboxyl group to the amino group, transforms from a tautomer* that does not contain charged groups into a tautomer with a bipolar ionic (zwitterionic) structure. Thus, a-amino acids are characterized prototropic tautomerism(Section 21.2.1).

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

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

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

According to the protolytic theory, acids, bases and ampholytes are protoliths The process by which a proton changes from an acid to a base is called protolysis and is explained by the fact that these two substances have different affinities for the proton. An acid-base interaction always involves two conjugated acid-base pairs, and the transition of a proton always occurs in the direction of the formation of weaker acids, including conjugated ones. If the propensity to interact with the proton of the reagents is commensurate, then there is protolytic balance.

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

Schematically, the protolytic equilibrium can be represented by the following scheme:

A proton transfer always comes from a strong acid to anion of a weak acid, which is accompanied by the displacement of the weak acid from its salt by the action of a stronger acid.

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

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

for weak electrolytes, the strength of acids and bases is characterized by the acidity constants K a and basicity K b respectively (Section 7.2). If these constants characterize the protolytic interaction of water with an acid or base of one conjugated pair HA, A or BH +, B, then the product of the acidity constants K a i basicity kb, components of this pair is always equal to the ionic product of water Kn 2 o \u003d 1 * 10 -14 (at 22 ° C):

These expressions make it possible to replace the basicity constant in the case of aqueous solutions ky or basicity index pK weak base on acidity constant K a or for acidity RK a conjugate acid of this base. In practice, to characterize the protolytic properties of a compound, the value is usually used RK a. So, 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. conjugate acid: pK a (NH4 +) \u003d 14 - 4.76 - 9.24. Therefore, in the case of aqueous solutions, there is no need for a special table of constants or index! basicity, a single scale of acidity, presented in Table. 8.1, where the properties of the bases are characterized by the constant K a or an indicator of acidity RK 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. Value RK a quantitatively characterizes the strength of weak electrolytes in aqueous solutions.

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

Meaning RK a is equal to the pH value of an aqueous solution in which this weak electrolyte is ionized by 50%: since in the atom case [A - ] \u003d [NA], then K a= [H + ] and RK a= pH. So, 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 be observed in solution with pH = pKa (NH4+) = 9.24.

In addition, the value RK a allows you to determine the pH value "of aqueous solutions, where this weak acid HA is predominantly (99% or more) in the form of an anion (A") - this will be in solutions with pH> pK a + 2; or in the form of molecules (HA) - in solutions with pH< pK a - 2. In the interval ApH = 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 above relations allow, knowing the value RK a biosubstrate, to determine in what form it will be at one or another pH value in the water systems of the body. In addition, knowledge of the value RK a of a weak electrolyte makes it possible to calculate the pH of aqueous solutions of this electrolyte, if its concentration is known.

Lecture #4

ORGANIC ACIDS AND BASES

Proton theory of acids and bases of Bronsted.
Classification of acids and bases according to Bronsted.
Influence of structural factors on acidity and basicity.
Lewis acids and bases. The theory of hard and soft acids and bases.

There are currently two main
theories of acids and bases: the Bronsted theory and the Lewis theory.

Proton theory of acids and
grounds of Bronsted

Bronsted acids - uh then the connections
capable of donating a proton (proton donors).

Founding of Bronsted - are compounds that can accept a proton
(proton acceptors). To interact with a proton, the base must have
a free pair of electrons or p-bond electrons.

Acids and bases form conjugates
acid-base pairs, for example:

In general :

The strength of the acid HA will depend on the strength of the base
:AT. Therefore, to create a unified scale, the strength of acids and bases of Bronsted
determined relative to water, which is an amphoteric compound and can
exhibit both acidic and basic properties.

The strength of acids is determined by the equilibrium constant
their interactions with water as a base, for example:

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

Since in dilute solutions
=const, then it can be added to
equilibrium constant, which is called the acidity constant:

In practice, the values are often used
pK a = - lg K a . How
less pK value a, the stronger
acid.

The strength of the bases is determined by the constant
equilibrium of their interaction with water as an acid:

RNH 2 + H 2 O  RNH 3 + + OH -

—
basicity constant.

For conjugate acids and bases
K a K b =K W . Thus, in
conjugated acid-base pair, the stronger the acid, the weaker the base and
vice versa. The strength of the base is often expressed not by the basicity constant, but by the constant
acidity of the conjugate acid.
For example, for the base RNH 2 magnitude is
conjugate acid acidity constant:

RNH 3 + + H 2 O  RNH 2 + H 3 O +

In practice, the value is often used . The larger the value, the
stronger base.

Organic classification
acids and bases

Bronsted acids and bases are classified according to
the nature of the atom at an acidic or basic center.

Depending on the nature of the element with which
proton bound, there are four main types of organic acids
Bronsted:

O-H - acids- carboxylic acids,
alcohols, phenols;
S-H - acids- thiols;
N-H - acids- amines, amides,
imides;
C-H - acids— hydrocarbons and their
derivatives.

Depending on the
the nature of the atom, to the lone pair of electrons of which a proton is attached,
Bronsted foundations are divided into three main types:

ammonium bases- amines,
nitriles, nitrogen-containing heterocyclic compounds;
oxonium bases- alcohols,
ethers, aldehydes, ketones, carboxylic acids and their functional
derivatives;
sulfonium bases- thiols,
sulfides.

special type
Bronsted's grounds represent p - bases in which the center of basicity is
electrons p - communications
(alkenes, arenes).

Influence of structural factors on
relative strength of acids and bases

The strength of an acid or base is determined
the equilibrium position of the acid-base interaction and depends on the difference
free energies of initial and final compounds. Therefore, the factors that
stabilize the conjugate base to a greater extent than the acid, increase
acidity and reduce basicity. Factors stabilizing predominantly
an acid compared to a base act in the opposite direction.
Since conjugate bases usually carry a negative charge, then
anion stabilizing factors contribute to an increase in acidity.

The effect of structure on the strength of acids and
grounds.

Bronsted acids.

The strength of an acid depends on the nature of the atom at
acid center and its structural environment.

For assessing the relative strength of acids, the following are important:
characteristics of an atom at an acidic center as its electronegativity and
polarizability.

Other things being equal, for elements of the same
period with an increase in the electronegativity of the atom, the acidity of the compounds
increases, since the high electronegativity of the atom at the acid center
stabilizes the anion formed during the elimination of a proton. Yes, acidity.
decreases in the series:

OH-acids> NH-acids>
CH-acids

	CH 3 O-H	CH 3 NH-H	CH 3 CH 2 -H
pK a	16	30	40

The electronegativity of an atom depends not only on
from its nature, but also from the type of hybridization and increases with increasing
s-character of hybrid orbitals. At the same time, acidity increases
connections:

The increase in the acidity of the compounds, despite
a decrease in the electronegativity of atoms in a subgroup is associated with an increase in their
polarizability as the radius of the atom increases. Large polarizability of the atom
contributes to better delocalization of the negative charge and increased stability
conjugate base.

With the same nature of the atom with acid
in the center, the strength of an acid is determined by its structural environment. Strength increase
acid promotes delocalization of the negative charge in the conjugate base
(anion) and its dispersion over more atoms.

So, carboxylic acids are one of the strongest
organic acids. Their strength is due to the stabilization of the carboxylate anion for
delocalization of the negative charge in the conjugated system. As a result
the negative charge in the carboxylate anion is dispersed between two atoms
oxygen, and both C-O bonds are absolutely equivalent:

Phenols are stronger acids than
alcohols, due to resonance stabilization of the phenolate anion, negative charge
which is delocalized along the aromatic ring:

As a result, the strength of organic OH-acids
may be placed in the following order:

	ROH		H2O		ArOH		RCOOH
pK a	16-17		15,7		8-11		4-5

The introduction of a substituent in the acid-bound
the center of the hydrocarbon radical affects the strength of the acid. electron-withdrawing
substituents increase, and electron-donating - reduce acidity. Influence
electron-withdrawing substituents is related to their ability to delocalize
negative charge and most
stabilize the conjugate base (anion). Effect of electron donor
substituents, on the contrary, leads to destabilization of the anion.

Electron-withdrawing substituents increase
strength of aliphatic and aromatic carboxylic acids, electron donor
substituents act in the opposite direction:

	Cl-CH 2 -COOH	H-COOH	CH 3 -COOH
pK a	2,8	3,7	4,7


	+M > -I		-M and -I
pK a	4,47	4,20	3,43

Substituents have a similar effect on
acidity of alcohols and phenols.

Founding of Bronsted.

With the same structural environment for
elements of the same period with an increase in the electronegativity of the atom at the main
the center of the basicity of the compounds decreases:

ammonium bases > oxonium bases I

	ROH	RNH 2
	~2	~10

The decrease in basicity is due to the fact that more
an electronegative atom holds the lone pair of electrons more firmly,
which he must give to form a bond with the proton.

Increases in the s-character of hybrid orbitals
leads to a decrease in basicity:

For elements of one subgroup with increasing
core charge basicity decreases:

oxonium bases > sulfonium
grounds

Introduction of electron-donating substituents
increases, and the introduction of electroacceptor - lowers the basicity. So,
electron-donating substituents increase the basicity of aliphatic and
aromatic amines, increasing the propensity of the nitrogen electron pair to attack
proton. Electron-withdrawing substituents, on the contrary, reduce the electron density
lone pair of nitrogen electrons and make it less susceptible to attack
proton:


		9,2	10,6	10,7

If a free pair of nitrogen electrons is in
conjugation with a double bond or aromatic ring, the basicity is reduced.
So, in aniline, a free pair of nitrogen electrons is conjugated with an aromatic
ring.

The protonation of aniline leads to a violation
conjugation and is energetically less favorable than the protonation of aliphatic
amines.


	10,6	4,6	0,9

Amides of carboxylic acids are very weak
bases due to the conjugation of a pair of nitrogen electrons with a carbonyl group. AT
As a result, the nitrogen atom acquires a partial positive, and the oxygen atom -
partial negative charge, and the protonation of amides occurs, as a rule,
by oxygen atom.

Basicity of nitrogen-containing heterocyclic
compounds is also determined by the availability of a pair of nitrogen electrons to attack
proton. Saturated nitrogen-containing heterocycles have a high basicity, in
in which the nitrogen atom is in the sp state 3 -hybridization. Basicity of the pyridinium nitrogen atom
(sp 2 hybridization) below. Finally,
the pyrrole nitrogen atom is practically devoid of basic properties, since its
protonation means the destruction of an aromatic heterocyclic
systems:


pK a	11,27	5,2	— 0.3

Acids and bases
Lewis

J. Lewis proposed a more general theory
acids and bases.

Lewis foundations they are the couple's donors
electrons (alcohols, alcoholate anions, ethers, amines, etc.)

Lewis acids - these are the acceptors of the pair
electrons, those. compounds that have
vacant orbital (hydrogen ion and metal cations: H + ,
Ag + , Na + , Fe 2+ ;
halides of elements of the second and third periods BF 3 ,
AlCl 3 , FeCl 3 , ZnCl 2 ; halogens; tin and sulfur compounds:
SnCl 4, SO3).

Thus, the foundations of Bronsted and Lewis are −
they are the same particles. However, according to Bronsted, basicity is the ability
attach only a proton, while Lewis basicity is a more
wide and means the ability to interact with any particle having
low-lying free orbital.

Lewis acid-base interaction is
donor-acceptor interaction and any heterolytic reaction can be
represent as the interaction of a Lewis acid and a Lewis base:

A single scale for comparing the strength of acids and
Lewis bases do not exist, since their relative strength will depend on
which substance is taken as the standard (for acids and bases of Bronsted such
water is the standard). To assess the ease of flow of acid-base
interaction according to Lewis R. Pearson proposed a qualitative theory
"hard" and "soft" acids and bases.

Rigid bases have a high
electronegativity and low polarizability. They are difficult to oxidize. Them
the highest occupied molecular orbitals (HOMO) have low energy.

Soft grounds have low
electronegativity and high polarizability. They oxidize easily. Their higher
occupied molecular orbitals (HOMO) have high energy.

Hard acids have a high
electronegativity and low polarizability. They are difficult to recover. Them
the lowest free molecular orbitals (LUMO) have low energy.

Soft acids have a low
electronegativity and high polarizability. They are easy to recover.
Their lowest free molecular orbitals (LUMOs) are high energy.

The hardest acid
H + , the softest
CH 3 Hg + . Most
rigid bases - F- and
oh- , the softest
I- and N - .

Table 5. Hard and soft acids
and foundations.

Rigid	Intermediate	Soft
acids
H + , Na + , K + , Mg 2+ , Ca 2+ , Al 3+ , Fe 3+ , BF 3 , AlCl 3 , RC + = O	Cu 2+, Fe 2+, Zn 2+ , R 3 C +	Ag + , Hg 2+ , I 2
Foundations
H 2 O, OH - , F - , ROH, RO -, R 2 O, NH 3, RNH 2	ArNH 2, Br -, C 5 H 5 N	R 2 S, RSH, RS - , I - , H - , C 2 H 4 , C 6 H 6

Principle of hard and soft acids and bases
Pearson (GIC principle):

Hard acids predominantly
react with hard bases and soft acids with soft
grounds.

This is expressed in high reaction rates and in
the formation of more stable compounds, since the interaction between close
orbitals are more energy efficient than the interaction between orbitals,
significantly different in energy.

The GMLC principle is used to determine
predominant direction of competing processes (reactions of elimination and
nucleophilic substitution, reactions involving ambident nucleophiles); for
targeted creation of detoxifiers and medicines.

Lecture #4

ORGANIC ACIDS AND BASES

Proton theory of acids and bases of Bronsted.
Classification of acids and bases according to Bronsted.
Influence of structural factors on acidity and basicity.
Lewis acids and bases. The theory of hard and soft acids and bases.

There are currently two main
theories of acids and bases: the Bronsted theory and the Lewis theory.

Proton theory of acids and
grounds of Bronsted

Bronsted acids - uh then the connections
capable of donating a proton (proton donors).

Founding of Bronsted - are compounds that can accept a proton
(proton acceptors). To interact with a proton, the base must have
a free pair of electrons or p-bond electrons.

Acids and bases form conjugates
acid-base pairs, for example:

In general :

The strength of acids is determined by the equilibrium constant
their interactions with water as a base, for example:

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

Since in dilute solutions
=const, then it can be added to
equilibrium constant, which is called the acidity constant:

In practice, the values are often used
pK a = - lg K a . How
less pK value a, the stronger
acid.

The strength of the bases is determined by the constant
equilibrium of their interaction with water as an acid:

RNH 2 + H 2 O  RNH 3 + + OH -

—
basicity constant.

RNH 3 + + H 2 O  RNH 2 + H 3 O +

In practice, the value is often used . The larger the value, the
stronger base.

Organic classification
acids and bases

Bronsted acids and bases are classified according to
the nature of the atom at an acidic or basic center.

Depending on the nature of the element with which
proton bound, there are four main types of organic acids
Bronsted:

O-H - acids- carboxylic acids,
alcohols, phenols;
S-H - acids- thiols;
N-H - acids- amines, amides,
imides;
C-H - acids— hydrocarbons and their
derivatives.

Depending on the
the nature of the atom, to the lone pair of electrons of which a proton is attached,
Bronsted foundations are divided into three main types:

ammonium bases- amines,
nitriles, nitrogen-containing heterocyclic compounds;
oxonium bases- alcohols,
ethers, aldehydes, ketones, carboxylic acids and their functional
derivatives;
sulfonium bases- thiols,
sulfides.

special type
Bronsted's grounds represent p - bases in which the center of basicity is
electrons p - communications
(alkenes, arenes).

Influence of structural factors on
relative strength of acids and bases

The effect of structure on the strength of acids and
grounds.

Bronsted acids.

The strength of an acid depends on the nature of the atom at
acid center and its structural environment.

For assessing the relative strength of acids, the following are important:
characteristics of an atom at an acidic center as its electronegativity and
polarizability.

OH-acids> NH-acids>
CH-acids

	CH 3 O-H	CH 3 NH-H	CH 3 CH 2 -H
pK a	16	30	40

Phenols are stronger acids than
alcohols, due to resonance stabilization of the phenolate anion, negative charge
which is delocalized along the aromatic ring:

As a result, the strength of organic OH-acids
may be placed in the following order:

	ROH		H2O		ArOH		RCOOH
pK a	16-17		15,7		8-11		4-5

Electron-withdrawing substituents increase
strength of aliphatic and aromatic carboxylic acids, electron donor
substituents act in the opposite direction:

	Cl-CH 2 -COOH	H-COOH	CH 3 -COOH
pK a	2,8	3,7	4,7


	+M > -I		-M and -I
pK a	4,47	4,20	3,43

Substituents have a similar effect on
acidity of alcohols and phenols.

Founding of Bronsted.

With the same structural environment for
elements of the same period with an increase in the electronegativity of the atom at the main
the center of the basicity of the compounds decreases:

ammonium bases > oxonium bases I

	ROH	RNH 2
	~2	~10

The decrease in basicity is due to the fact that more
an electronegative atom holds the lone pair of electrons more firmly,
which he must give to form a bond with the proton.

Increases in the s-character of hybrid orbitals
leads to a decrease in basicity:

For elements of one subgroup with increasing
core charge basicity decreases:

oxonium bases > sulfonium
grounds


		9,2	10,6	10,7

The protonation of aniline leads to a violation
conjugation and is energetically less favorable than the protonation of aliphatic
amines.


	10,6	4,6	0,9


pK a	11,27	5,2	— 0.3

Acids and bases
Lewis

J. Lewis proposed a more general theory
acids and bases.

Lewis foundations they are the couple's donors
electrons (alcohols, alcoholate anions, ethers, amines, etc.)

Lewis acid-base interaction is
donor-acceptor interaction and any heterolytic reaction can be
represent as the interaction of a Lewis acid and a Lewis base:

Rigid bases have a high
electronegativity and low polarizability. They are difficult to oxidize. Them
the highest occupied molecular orbitals (HOMO) have low energy.

Soft grounds have low
electronegativity and high polarizability. They oxidize easily. Their higher
occupied molecular orbitals (HOMO) have high energy.

Hard acids have a high
electronegativity and low polarizability. They are difficult to recover. Them
the lowest free molecular orbitals (LUMO) have low energy.

Soft acids have a low
electronegativity and high polarizability. They are easy to recover.
Their lowest free molecular orbitals (LUMOs) are high energy.

The hardest acid
H + , the softest
CH 3 Hg + . Most
rigid bases - F- and
oh- , the softest
I- and N - .

Table 5. Hard and soft acids
and foundations.

Rigid	Intermediate	Soft
acids
H + , Na + , K + , Mg 2+ , Ca 2+ , Al 3+ , Fe 3+ , BF 3 , AlCl 3 , RC + = O	Cu 2+, Fe 2+, Zn 2+ , R 3 C +	Ag + , Hg 2+ , I 2
Foundations
H 2 O, OH - , F - , ROH, RO -, R 2 O, NH 3, RNH 2	ArNH 2, Br -, C 5 H 5 N	R 2 S, RSH, RS - , I - , H - , C 2 H 4 , C 6 H 6

Principle of hard and soft acids and bases
Pearson (GIC principle):

Hard acids predominantly
react with hard bases and soft acids with soft
grounds.

According to Lewis, the acidic and basic properties of organic compounds are measured by the ability to accept or donate an electron pair, followed by the formation of 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. An atom that provides an electron pair is an electron donor, and a compound containing such an atom is a base.

Specifically, Lewis acids can be an atom, molecule or cation: proton, halides of elements of the second and third groups of the Periodic system, 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 Bronsted-Lowry, acids are substances capable of donating a proton, and bases are substances that accept a proton.

Conjugate acid and base:

HCN (acid) and CN- (base)

NH3 (base) and NH4+ (acid)

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

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

acid 2 base 1

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

acid 1 base 2 conjugate conjugate

acid 2 base 1

7. Types of isomerism in organic chemistry. Structural, spatial and optical isomerism. Chirality. Compatibility and configuration. R,S, Z,E - nomenclature.

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

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

Structural isomerism

Isomerism of the carbon skeleton

Position isomerism is due to the different position of the multiple bond, substituent, 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 cycle is impossible, substituents can be located either on one side of the plane of the double bond or cycle (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, property of an object to be incompatible with its reflection in an ideal flat 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.

The structure of alkanes. Sp3 is the state of carbon. Characterization of C-C and C-H bonds. The principle of free rotation. conformation. Methods of representation and nomenclature. Physical properties of alkanes.

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