Classify its composition and nature of the components.

Please indicate the range of pH values within which this system has a buffer capacity.

Write equations of reactions reflecting the mechanism of its action (ionic form).

Explain why the ammonia buffer system is not part of the blood

1.Composition and nature of components:

A) NH 4 OH (NH 3 x H 2 O) -ammonium hydroxide, weak electrolyte

B) NH 4 C1 - salt, ammonium chloride, strong electrolyte.

Ammonium hydroxide is a weak electrolyte; in solution it partially dissociates into ions:

NH 4 OH<=>NH 4 + + OH-

When ammonium chloride is added to the solution of ammonium hydroxide, salt, as a strong electrolyte, almost completely dissociates into ions:

NH 4 C1> NH 4 + + C1-

and suppresses the dissociation of the base, the equilibrium of which is shifted towards the reverse reaction.

The range of pH values within which the system under consideration has a buffer capacity is calculated by the formula:

where KB is the dissociation constant of NH 4 OH = 1.8 * 10 -5, C 0 is the base concentration, Cc is the salt concentration.

pH = 14-4.74 + log (C 0 / Cc) = 9.26 + log (C 0 / Cc). Depending on the C 0 / Cc ratio, the pH range is 8.26-10.26.

Ability ammonia buffer to maintain an almost constant pH value of the solution is based on the fact that the components included in them bind the H + and OH- ions introduced into the solution or formed as a result of the reaction proceeding in this solution. When a strong acid is added to the ammonia buffer mixture, H + ions will bind to molecules or ammonium hydroxide, and not increase the concentration of H + ions and decrease the pH of the solution:

NH 4 OH + H + = NH 4 + + H 2 O

When alkali is added, OH - ions will bind NH 4 + ions, thus forming a slightly dissociated compound, and not increase the pH of the solution:

NH 4 + + OH - = NH 4 OH

The ammonia buffer system is not included in the TOP RT composition of blood, since the range of pH values within which it will have a buffer capacity is in the alkaline region (pH greater than 8). Normal blood plasma pH is 7.40 ± 0.05, i.e. below the buffering area.

1. 3)Write the scheme of the reaction of the interaction of ethanal with methylamine.

Describe mechanism of this reaction.

Justify the role of the acid catalyst.

Explain the possibility of the reaction of hydrolysis of the resulting imine in an acidic and alkaline environment.

2. The mechanism of this reaction is nucleophilic addition followed by elimination of a water molecule

3. The role of the acid catalyst - protonation in stage a)

4. In the presence of dilute acids, imines are hydrolyzed with water to form carbonyl compounds and amines, this reaction is the reverse of the reaction for the synthesis of imines:

In the presence of alkali, hydrolysis does not take place

Ticket 4.

Thermodynamic system (TM) - this is any real object isolated from the environment in order to study the processes of exchange of energy and energy between its constituent parts, as well as between it and environment using thermodynamic methods

Classification of thermodynamic systems

3. Open exchange both matter and energy with the OS (an organism, an open vessel with boiling water)

4. Closed- exchanges with the OS only energy in the form of heat or work (gas in a closed cylinder)

5. Isolated- do not exchange anything or energy. There are no absolutely isolated people in nature.

By the presence of an interface inside the vehicle

1.Homogeneous- there is no interface, all components are in the water phase, all physical and chemical substances in any part of the volume are the same (mixture of gases)

2. Heterogeneous-contains the interface, the separating parts of the system (phases) are different according to the sv-you (blood)

Options- quantities that determine the state of the vehicle

Direct measurement possible

The main parameters are parameters that can be measured using the appropriate devices (m, V, C, density, volume)

State functions - internal energy E (U); enthalpy (H); entropy (S); Gibbs energy (G); free energy or Helmholtz energy

You can define the change in the values of the state function

∆X (X 2 -X 1), WHERE X-U, H, S, G, H

Thermodynamic state-the set of values of a certain number of physical. values characterizing all physical and hm sv-va systems

Types of states:

Non-equilibrium - parameters change spontaneously (glass with hot water)

Equilibrium parameters do not change without external influences

Stationary = constancy of parameters due to external parameters (inherent in living organisms)

Process-transition of the system from one state to another, accompanied by changes in thermodynamic parameters.

Classification-

by the constancy of parameters:

A) isochoric (v = const)

B) Isobaric (pressure- const)

C) isothermal (temperature = const)

By the sign of the thermal effect: exothermic and endothermic

By energy expenditure: spontaneous, non-spontaneous

By the nature of the course:-reversible - flow in the forward and backward directions through the same stages, without changes in the environment. environment.

Irreversible - all processes cannot proceed in the forward and reverse directions through one and the same stages.

Buffer solutions are solutions that keep the pH constant when diluted or a small amount of a strong acid or base is added.

Protolytic buffer solutions are electrolyte mixtures containing ions of the same name.

There are two types of protolytic buffers:

Acidic, consisting of a weak acid and an excess of a base conjugated with it (the salt formed by a strong base and the anion of this acid);
Basic, consisting of a weak base and an excess of conjugated acid (i.e., a salt formed by a strong acid and a cation of this base).

The buffer system equation is calculated using the Henderson-Hasselbach formula:

where pK = -ℓg K D.

C - molar or equivalent concentration of electrolyte (C = V N)

The mechanism of action of buffer solutions can be considered using the example of an acetate buffer: CH 3 COOH + CH 3 COONa.

The high concentration of acetate ions is due to the complete dissociation of a strong electrolyte - sodium acetate, and acetic acid in the presence of the anion of the same name exists in solution in almost non-ionized form.

1. When a small amount of hydrochloric acid is added, H + ions are bound with the conjugated base CH 3 COO - present in the solution, into a weak electrolyte CH 3 COOH.

CH 3 COO‾ + H + ↔ CH 3 COOH

It can be seen from the equation that the strong acid HCl is replaced by an equivalent amount of the weak acid CH 3 COOH. The amount of CH 3 COOH increases and according to V. Ostwald's law of dilution, the degree of dissociation decreases. As a result, the concentration of H + ions in the buffer increases, but very slightly, while the pH remains constant.

When acid is added to the buffer, the pH is determined by the formula:

2. When a small amount of alkali is added to the buffer, it reacts with CH 3 COOH. Acetic acid molecules will react with hydroxide ions to form H 2 O and CH 3 COO ‾:

CH 3 COOН + OH ‾ ↔ CH 3 COO‾ + H 2 O

As a result, the alkali is replaced with an equivalent amount of the weakly basic salt CH 3 COONa. The amount of CH 3 COOH decreases and according to V. Ostwald's law of dilution, the degree of dissociation increases due to the potential acidity of the remaining undissociated CH 3 COOH molecules. Consequently, the concentration of H + ions practically does not change, and the pH remains constant.

When alkali is added, the pH is determined by the formula:

3. When the buffer is diluted, the pH also does not change, because the dissociation constant and the ratio of the components remain unchanged.

Thus, the pH of the buffer depends on the dissociation constant and the concentration ratio of the components. The higher these values are, the higher the pH of the buffer. It should be noted that the pH of the buffer will be highest when the ratio of the components is equal to unity.

Buffer capacity is the ability of the buffer system to resist changes in the pH of the medium.

Buffer capacity (B) is expressed as the number of molar equivalents of a strong acid or alkali that must be added to one liter of buffer to shift the pH by one.

where B is the buffer capacity, n E is the amount of a mol-equivalent of a strong acid or alkali, pH H is the initial pH value (before adding an acid or alkali), pH K is the final pH value (after adding an acid or alkali), ΔрН is a change in pH ...

In practice, the buffer capacity is calculated using the formula:

where V is the volume of acid or alkali, N is the equivalent concentration of acid or alkali, V buf. is the volume of the buffer solution, Δ pH is the change in pH.

The buffer capacity depends on the concentration of electrolytes and the ratio of the components of the buffer. The greatest buffer capacity is possessed by solutions with a higher concentration of components and a ratio of components equal to unity.

The following buffer systems operate in the human body:

Bicarbonate buffer, which is the main buffer system of blood plasma; it is a system of rapid reaction, since the product of its interaction with acids CO 2 is quickly excreted through the lungs. In addition to plasma, this buffer system is contained in erythrocytes, interstitial fluid, and renal tissue.
The hemoglobin buffer is the main buffer system of erythrocytes, which accounts for about 75% of the total buffer capacity of the blood. The participation of hemoglobin in the regulation of blood pH is associated with its role in the transport of oxygen and CO2. The hemoglobin buffer system of the blood plays a significant role in several physiological processes at once: respiration, oxygen transport in tissues and in maintaining a constant pH inside erythrocytes, and ultimately in the blood.
Phosphate buffer is found both in the blood and in the cellular fluid of other tissues, especially the kidneys. In cells, it is represented by the salts K 2 HPO 4 and KH 2 PO 4, and in blood plasma and in the intercellular fluid, Na 2 HPO 4 and NaH 2 PO 4. It functions mainly in plasma and includes: dihydrogen phosphate ion Н 2 РО 4 - and hydrogen phosphate ion НРО 4 2-.
A protein buffer consists of a protein acid and its salt formed by a strong base.

Protein is an amphoteric electrolyte and therefore exerts its own buffering action. Interaction of buffer systems in the body by stages:

1. In the process of gas exchange in the lungs, oxygen enters the erythrocytes;

2. As the blood moves to the peripheral regions circulatory system oxygen is returned by the ionized form of HbO 2 -. In this case, the blood becomes venous from arterial. The oxygen released in the tissues is spent on the oxidation of various substrates, resulting in the formation of CO 2, most of which enters the erythrocytes.

3. In erythrocytes in the presence of carbonic anhydrase, the following reaction proceeds at a significant rate:

CO 2 + H 2 O ↔ H 2 CO 3 ↔ H + + HCO 3 -

4. The resulting excess of protons binds with hemoglobinate ions, while the binding of protons shifts the equilibrium of the reaction of stage (3) to the right, as a result of which the concentration of bicarbonate ions increases and they diffuse through the membrane into the plasma. As a result of the counter-diffusion of ions differing in acid-base properties (the chloride ion is protolytically inactive; the bicarbonate ion is the base under the conditions of the organism), a bicarbonate-chloride shift occurs. This explains the more acidic reaction of the medium in erythrocytes (pH = 7.25) compared to plasma (pH = 7.4).

5. Bicarbonate ions entering the plasma neutralize the excess of protons accumulating there, resulting from metabolic processes;

6. Formed CO 2 interacts with the components of the protein buffer system;

7. Excess protons are neutralized with phosphate buffer:

Н + + НРО 4 - ↔ Н 2 РО 4 -

8. After the blood enters the lungs again, the concentration of oxyhemoglobin increases in it (stage 1), which reacts with bicarbonate ions that have not diffused into the plasma. The resulting CO 2 is excreted through the lungs. As a result of a decrease in the concentration of HCO 3 - ions in this part of the bloodstream, their diffusion into erythrocytes and diffusion of chloride ions in the opposite direction are observed.

9. The kidneys also accumulate an excess of protons as a result of the reaction:

CO 2 + H 2 O ↔ H 2 CO 3 ↔ H + + HCO 3 -,

which is neutralized by hydrophosphate ions and ammonia (ammonia buffer):

H + + NH 3 ↔ NH 4 +

It should be noted that the maintenance of the constancy of the pH of various fluid systems of the body is influenced not so much by the buffer systems as by the functioning of a number of organs and systems: lungs, kidneys, intestines, skin, etc.

The pH of human blood averages 7.4, a change in this value even by one tenth of a unit leads to severe violations(acidosis or alkalosis). When the pH is outside the 6.8-7.8 range, it usually leads to death. The most important buffer system of the blood is carbon (HCO 3 - / H 2 CO 3), the second in importance is phosphate (HPO 2 -4 / H 2 PO -4), proteins also play a role in maintaining pH.

Buffer classification

Distinguish between natural and artificial buffer solutions. The natural buffer solution is blood containing bicarbonate, phosphate, protein, hemoglobin and acidic buffer systems. The artificial buffer can be an acetate buffer consisting of CH3COOH.

Buffer solutions can be acidic (pH< 7) или щелочную (рН > 7). .

Buffer systems can be of four types:

1) Weak acid and its anion:

For example: acetate buffer system

CH 3 COONa and CH 3 COOH, area of action pH = 3, 8 - 5, 8.

2) Weak base and its cation:

For example: ammonia buffer system

NH 3 and NH 4 Cl, range of action pH = 8, 2 - 10, 2.

3) Anions of acidic and medium salt:

For example: carbonate buffer system

Na 2 CO 3 and NaHCO 3, area of action pH = 9, 3 - 11.

4) A mixture of two acidic salts:

For example: phosphate buffer system

Na 2 HP0 4 and NaH 2 PO 4, area of action pH = 7.4 - 8.

The mechanism of action of buffers

Let's see what the properties of buffer solutions are based on, using the example of a buffer mixture of acetic acid and sodium acetate.

1) Dilution with water

Acetic acid is a weak acid, in addition, its dissociation is further reduced due to the presence of sodium acetate (the effect of the ion of the same name). buffer solution hydroxide tetraborate

Suppose that the solution in question is diluted with water 10 or 20 times. It would seem that due to a strong decrease in the concentration of acetic acid, the concentration of H + ions should decrease, but this does not happen, because the degree of dissociation of acetic acid increases with dilution, since the concentration of sodium acetate, which suppresses the dissociation of acetic acid in this solution, decreases. Therefore, when diluted with water, the pH practically does not change.

2) Addition of strong acid

When a small amount of a strong acid, such as hydrochloric acid, is added to the buffer mixture, the following reaction occurs:

CH 3 COONa + HCl = NaCl + CH 3 COOH.

H + ions entering the solution will bind to form acetic acid molecules with a low degree of dissociation. Thus, the concentration of H + ions will hardly increase and the pH of the solution will practically not change.

If the same amount of acid is added to pure water, all H + ions will remain in the solution, the concentration of hydrogen ions will increase many times and the pH of the solution will noticeably change. And hydrogen, as you know, is the most widespread chemical element.

3) Adding a small amount of alkali

The alkali added to the buffer mixture reacts with acetic acid:

CH 3 COOH + NaOH = CH 3 COONa + H 2 O.

OH - ions are bound by H + ions of acetic acid to form undissociated water molecules. However, the loss of these ions is replenished as a result of the dissociation of acetic acid molecules. Thus, the pH of the solution practically does not change after the addition of alkali.

If you add alkali to pure water, all OH ions will remain in solution. The concentration of OH - ions will increase sharply, the concentration of H + ions will accordingly decrease and the pH of the solution will change noticeably.

Similar phenomena are observed when small amounts of acids and alkalis are added to other buffers.

Chapter 6. PROTOLYTIC BUFFER SYSTEMS

A change in any factor that can affect the state of chemical equilibrium of a system of substances causes a reaction in it that tends to counteract the change being made.

A. Le Chatelier

6.1. BUFFER SYSTEMS. DEFINITION AND GENERAL PROVISIONS OF THE THEORY OF BUFFER SYSTEMS. CLASSIFICATION OF BUFFER SYSTEMS

Systems supporting protolytic homeostasis include not only physiological mechanisms (pulmonary and renal compensation), but also physicochemical buffering, ion exchange, and diffusion. Maintaining the acid-base balance at a given level is provided at the molecular level by the action of buffer systems.

Protolytic buffering systems are solutions that maintain a constant pH value both with the addition of acids and alkalis, and with dilution.

The ability of some solutions to keep the concentration of hydrogen ions unchanged is called buffering action, which is the main mechanism of protolytic homeostasis. Buffer solutions are mixtures of a weak base or weak acid and their salts. In buffer solutions, the main "active" components are the proton donor and acceptor, according to Brønsted's theory, or the electron pair donor and acceptor, according to Lewis's theory, which are an acid-base pair.

According to the belonging of the weak electrolyte of the buffer system to the class of acids or bases and according to the type of charged particle, they are divided into three types: acidic, basic and ampholytic. A solution containing one or more buffering systems is called a buffer solution. Buffer solutions can be prepared in two ways:

By partial neutralization of a weak electrolyte with a strong electrolyte:

By mixing solutions of weak electrolytes with their salts (or two salts): CH 3 COOH and CH 3 COONa; NH 3 and NH 4 Cl; NaH 2 PO 4

and Na 2 HPO 4.

The reason for the emergence of a new quality in solutions - buffering action - lies in the combination of several protolytic equilibria:

Conjugated acid-base pairs B / BH + and A - / HA are called buffer systems.

In accordance with Le Chatelier's principle, the addition of a weak acid HB + H 2 O ↔ H 3 O + + B - a strong acid or a salt containing anions B - to a solution, an ionization process takes place, shifting the equilibrium to the left (effect of the common ion) B - + H 2 O ↔ HB + OH -, and the addition of alkali (OH -) - to the right, since the concentration of hydronium ions will decrease due to the neutralization reaction.

When two isolated equilibria are combined (acid ionization and anion hydrolysis), it turns out that the processes that will proceed in them under the influence of the same external factors (addition of hydronium ions and hydroxide ions) are multidirectional. In addition, the concentration of one of the products of each of the combined reactions affects the equilibrium position of the other reaction.

The protolytic buffer system is a combined equilibrium of the processes of ionization and hydrolysis.

The equation of the buffer system expresses the dependence of the pH of the buffer solution on the composition of the buffer system:

An analysis of the equation shows that the pH value of the buffer solution depends on the nature of the substances that form the buffer system, the ratio of the concentration of components and temperature (since the value of pKa depends on it).

According to protolytic theory, acids, bases and ampholytes are protoliths.

6.2. TYPES OF BUFFER SYSTEMS

Acid type buffer systems

Acid buffer systems are a mixture of a weak acid HB (proton donor) and its salt B - (proton acceptor). They are usually acidic (pH<7).

Bicarbonate buffer system (buffer zone pH 5.4-7.4) - a mixture of weak carbonic acid H 2 CO 3 (proton donor) and its salt HCO 3 - (proton acceptor).

Hydrophosphate buffer system (buffer zone pH 6.2-8.2) - a mixture of a weak acid H 2 PO 4 - (proton donor) and its salt HPO 4 2- (proton acceptor).

The hemoglobin buffer system is represented by two weak acids (proton donors) - hemoglobin HHb and oxyhemoglobin HHbO 2 and conjugated weak bases (proton acceptors) - hemoglobinate - Hb - and oxyhemoglobinate anions HbO 2 -, respectively.

Buffer systems of the basic type

The main buffer systems are a mixture of a weak base (proton acceptor) and its salt (proton donor). They are usually alkaline (pH> 7).

Ammonia buffer system: a mixture of a weak base NH 3 H 2 O (proton acceptor) and its salt - a strong electrolyte NH 4 + (proton donor). Buffer zone at pH 8.2-10.2.

Buffer systems of the ampholytic type

Ampholytic buffer systems consist of a mixture of two salts or a salt of a weak acid and a weak base, for example CH 3 COONH 4, in which CH 3 COO - exhibits weak basic properties - a proton acceptor, and NH 4 + - a weak acid - a proton donor. A biologically significant buffer system of the ampholytic type is the protein buffer system - (NH 3 +) m -Prot- (CH 3 COO -) n.

Buffer systems can be viewed as a mixture of weak and strong electrolytes with the same ions (common ion effect). For example, in an acetate buffer solution, there are acetate ions, and in a bicarbonate solution, there are carbonate ions.

6.3. MECHANISM OF ACTION OF BUFFER SOLUTIONS AND DETERMINATION OF PH IN THESE SOLUTIONS. GENDERSON-HASSELBACH EQUATION

Let us consider the mechanism of action of acid-type buffer solutions using the example of the acetate buffer system CH 3 COO - / CH 3 COOH, which is based on the acid-base balance CH 3 COOH ↔ H + + CH 3 COO - (K И = 1.75 10 - 5). The main source of acetate ions is the strong electrolyte CH 3 COONa. When a strong acid is added, the conjugated base CH 3 COO - binds the added hydrogen cations, turning into a weak acid: CH 3 COO - + + H + ↔ CH 3 COOH (acid-base equilibrium shifts to the left). A decrease in the concentration of CH 3 COO - is balanced by an increase in the concentration of a weak acid and indicates a hydrolysis process. According to the Ostwald dilution law, an increase in the acid concentration somewhat lowers its degree of electrolytic dissociation and the acid practically does not ionize. Therefore, in the system: C to increases, C c and α decreases, - const, C to / C c increases, where C to is the acid concentration, C c is the salt concentration, α is the degree of electrolytic dissociation.

When alkali is added, the hydrogen cations of acetic acid are released and neutralized by the added OH - ions, binding to water molecules: CH 3 COOH + OH - → CH 3 COO - + H 2 O

(acid-base balance shifts to the right). Consequently, C to increases, C c and α decreases, - const, C to / C decreases.

The mechanism of action of the buffer systems of the basic and ampholytic types is similar. The buffer effect of the solution is due to a shift in the acid-base balance due to the binding of the added H + and OH - ions by the components of the buffer and the formation of low-dissociating substances.

Mechanism of action of protein buffer solution upon addition of acid: (NH 3 +) m -Prot- (COO -) n + nH + ↔ (NH 3 +) m -Prot- (COOH) n, with the addition of alkali - (NH 3 +) m -Prot- (COO -) n + mOH - ↔ (NH 2) m - Prot- (COO -) n + mH 2 O.

At high concentrations of H + and OH - (more than 0.1 mol / l), the ratio of the components of the buffer mixture changes significantly - C / C increases or decreases, and the pH can change. This is confirmed by Henderson-Hasselbach equation, which establishes the relationship [H +], K I, α and C to / C s. The equation

we deduce the example of an acid-type buffer system - a mixture of acetic acid and its salt CH 3 СОONа. The concentration of hydrogen ions in the buffer solution is determined by the ionization constant of acetic acid:

The equation shows that the concentration of hydrogen ions is in direct proportion to KI, α, the concentration of acid C to and inversely to C c and the ratio C to / C c. Taking the logarithm of both sides of the equation and taking the logarithm with a minus sign, we get the equation in logarithmic form:

The Henderson-Hasselbach equation for buffer systems of the basic and ampholytic types is derived using the example of deriving an equation for acid-type buffer systems.

For a buffer system of the basic type, for example, ammonia, the concentration of hydrogen cations in solution can be calculated based on the constant of acid-base equilibrium of conjugate acid

NH 4 + :

Henderson-Hasselbach equation for basic buffer systems:

This equation can be represented as:

For the phosphate buffer system HPO 4 2- / H 2 PO 4 - pH can be calculated using the equation:

where pK 2 is the dissociation constant of orthophosphoric acid at the second stage.

6.4. CAPACITY OF BUFFER SOLUTIONS AND ITS DETERMINING FACTORS

The ability of solutions to maintain a constant pH value is not unlimited. Buffers can be distinguished by their resistance to acids and bases added to the buffer.

The amount of acid or alkali that must be added to 1 liter of the buffer solution so that its pH value changes by one is called the buffer tank.

Thus, buffering capacity is a quantitative measure of the buffering action of a solution. The buffer solution has a maximum buffer capacity at pH = pK of an acid or base that forms a mixture at a ratio of its components, equal to one... The higher the initial concentration of the buffer mixture, the higher its buffer capacity. The buffering capacity depends on the composition of the buffer solution, the concentration and the ratio of the components.

You need to be able to choose the right buffer system. The choice is determined by the required pH range. The zone of buffer action is determined by the strength index of the acid (base) ± 1 unit.

When choosing a buffer mixture, it is necessary to take into account the chemical nature of its components, since the substances of the solution to which

the buffer system is formed, they can form insoluble compounds, interact with the components of the buffer system.

6.5. BLOOD BUFFER SYSTEMS

Blood contains 4 major buffering systems.

1.Hydrocarbonate. It accounts for 50% of the capacity. It works primarily in plasma and plays a central role in CO2 transport.

2.Protein. It accounts for 7% of the capacity.

3.Hemoglobin, it accounts for 35% of the capacity. It is represented by hemoglobin and oxyhemoglobin.

4. Hydrophosphate buffer system - 5% capacity. Hydrocarbonate and hemoglobin buffer systems perform

a central and extremely important role in CO2 transport and pH adjustment. In blood plasma pH 7.4. CO 2 is a product of cellular metabolism released into the blood. It diffuses through the membrane into erythrocytes, where it reacts with water to form H 2 CO 3. The ratio is set to 7 and the pH will be 7.25. Acidity rises, while reactions take place:

Formed HCO 3 - leaves through the membrane and is carried away by the blood stream. In the blood plasma at the same pH is 7.4. When venous blood enters the lungs again, hemoglobin reacts with oxygen to form oxyhemoglobin, which is a stronger acid: HHb + + O 2 ↔ HHbO 2. pH decreases, as a stronger acid is formed, the reaction occurs: HHbO 2 + HCO 3 - ↔ HbO 2 - + H 2 CO 3. Then CO 2 is released into the atmosphere. This is one of the mechanisms of transport of CO 2 and O 2.

Hydration and dehydration of CO 2 is catalyzed by the enzyme carbo-anhydrase, which is present in erythrocytes.

The bases also bind to the buffered blood solution and are excreted in the urine, mainly in the form of mono- and disubstituted phosphates.

In clinics, the reserve alkalinity of the blood is always determined.

6.6. QUESTIONS AND EXERCISES FOR SELF-CHECK PREPARATION FOR EXERCISES AND EXAMINATIONS

1.When combining what protolytic equilibria will the solutions have buffer properties?

2. To give the concept of buffer systems and buffering action. What is the buffering chemistry?

3. The main types of buffer solutions. The mechanism of their buffering action and the Henderson-Hasselbach equation, which determines the pH in buffer systems.

4. The main buffer systems of the body and their relationship. What does the pH of buffer systems depend on?

5. What is called the buffer capacity of the buffer system? Which of the blood buffer systems has the greatest capacity?

6. Methods for preparing buffer solutions.

7. The choice of buffer solutions for biomedical research.

8. Determine whether acidosis or alkalosis is observed in a patient if the concentration of hydrogen ions in the blood is 1.2.10 -7 mol / l?

6.7. TEST PROBLEMS

1. Which of the proposed systems is a buffer?

a) HCl and NaCl;

b) H 2 SO 4 and NaHSO 4;

c) H 2 CO 3 and NaHCO 3;

d) HNO 3 and NaNO 3;

e) HClO 4 and NaClO 4.

2. For which of the proposed buffer systems does the calculated formula pH = pK correspond?

a) 0.1 M solution NaH 2 PO 4 and 0.1 M solution Na 2 HPO 4;

b) 0.2 M solution of H 2 CO 3 and 0.3 M solution of NaHCO 3;

c) 0.4 M solution of NH 4 OH and 0.3 M solution of NH 4 Cl;

d) 0.5 M solution of CH 3 COOH and 0.8 M solution of CH 3 COONa;

e) 0.4 M solution NaHCO 3 and 0.2 M solution H 2 CO 3.

3. Which of the proposed buffer systems is a bicarbonate buffer system?

a) NH 4 OH and NH 4 Cl;

b) H 2 CO 3 and KHSO 3;

c) NaH 2 PO 4 and Na 2 HPO 4;

d) CH 3 COOH and CH 3 COOK;

e) K 2 HPO 4 and KH 2 PO 4.

4. Under what conditions is the pH of the buffer system equal to pK k?

a) when the concentration of the acid and its salt are equal;

b) when the concentration of the acid and its salt are not equal;

c) when the ratio of volumes of acid and its salt is equal to 0.5;

d) when the ratio of volumes of acid and its salt at the same concentrations is not equal;

e) when the acid concentration is 2 times higher than the salt concentration.

5. Which of the proposed formulas is suitable for calculating [H +], for the system CH 3 COOH and CH 3 SOOK?

6. Which of the following mixtures is part of the body's buffer system?

a) HCl and NaCl;

b) H 2 S and NaHS;

c) NH 4 OH and NH 4 Cl;

d) H 2 CO 3 and NaHCO 3;

e) Ba (OH) 2 and BaOHCl.

7. What type of acid-base buffer systems is a protein buffer?

a) weak acid and its anion;

c) anions 2 of acidic salts;

e) ions and molecules of ampholytes.

8. What type of acid-base buffer systems is ammonia buffer?

a) weak acid and its anion;

b) anions of acidic and medium salt;

c) anions 2 of acidic salts;

d) weak base and its cation;

e) ions and molecules of ampholytes.

9. What type of acid-base buffer systems is phosphate buffer?

a) weak acid and its anion;

b) anions of acidic and medium salt;

c) anions 2 of acidic salts;

d) weak base and its cation;

e) ions and molecules of ampholytes.

10. When is a protein buffering system not a buffer?

a) at the isoelectric point;

b) when adding alkali;

c) when adding acid;

d) in a neutral environment.

11. Which of the proposed formulas is suitable for calculating the [OH -] system: NH 4 OH and NH 4 Cl?

General chemistry: textbook / A. V. Zholnin; ed. V. A. Popkova, A. V. Zholnina. - 2012 .-- 400 p .: ill.

INTRODUCTION

BUFFER SOLUTIONS (buffer mixtures, buffers) - solutions containing buffer systems and therefore have the ability to maintain the pH at a constant level. They are usually prepared by dissolving in water taken in appropriate proportions of a weak acid and its salt formed by an alkali metal, partial neutralization of a weak acid with a strong alkali or a weak base with a strong acid, dissolving a mixture of salts of a polybasic acid. The pH value of the buffer solutions prepared in this way changes slightly with temperature. The range of pH values in which the buffer solution has stable buffering properties lies within pK ± 1 (pK is the negative decimal logarithm of the dissociation constant of the weak acid included in its composition). The most famous buffer solutions are: Sørensen glycine, Walpole acetate, Sørensen phosphate, Palic borate, Veronal Michaelis, Kolthof carbonate, Tris buffer, universal Veronal Michaelis, etc.

In laboratory practice, buffer solutions are used to maintain the active reaction of the medium at a certain constant level and to determine the pH value (pH) - as standard solutions with stable pH values, etc.

BUFFER MIXTURES

If water is added to a solution of any acid or alkali, then, of course, the concentration of hydrogen or hydroxyl ions decreases accordingly. But if you add a certain amount of water to a mixture of acetic acid and sodium acetate or to a mixture of ammonium hydroxide and ammonium chloride, then the concentration of hydrogen and hydroxyl ions in these solutions will not change.

The property of some solutions to keep the concentration of hydrogen ions unchanged when diluted, as well as when small amounts of strong acids or alkalis are added, is known as buffering.

Solutions containing at the same time any weak acid and its salt or any weak base and its salt and have a buffering effect are called buffer solutions. Buffer solutions can be considered as a mixture of electrolytes with ions of the same name. The presence in the solution of a weak acid or a weak base and their salts reduces the effect of dilution or the action of other acids and bases on the pH of the solution.

Such buffers are the following mixtures of CH 3 COOH + CH 3 C OON a, NH 4 OH + NH 4 Cl, Na 2 CO 3 + NaHCO 3, etc.

Buffer solutions, which are mixtures of weak acids and their salts, usually have an acidic reaction (pH<7). Например, буферная смесь 0,1М раствора СН 3 COOP + 0.1M CH solution 3 CO ONa has pH = 4.7.

Buffer solutions, which are mixtures of weak bases and their salts, are usually alkaline (pH> 7). For example, a buffer mixture of 0.1M solution N H 4 OH + 0.1 M solution of N H 4 C1 has a pH of 9.3.

Acid-base buffers

In a broad sense, buffer systems are systems that maintain a certain value of a parameter when the composition changes. Buffer solutions can be

- acid-base - maintain a constant pH value with the addition of small amounts of acid or base.

Redox - keep the potential of the system constant with the introduction of oxidizing agents or reducing agents.

known metallobuffer solutions that maintain a constant pH value.

In all cases, the buffer solution is a conjugated pair. In particular, acid-base buffers contain a conjugated acid-base pair. The buffering effect of these solutions is due to the presence of acid-base equilibrium general type:

HA ↔ H + + A -

acid conjugate

Base

В + Н + ↔ ВН +

O warp conjugate

Acid

Since in this section only acid-base buffer solutions are considered, we will call them buffer solutions, omitting "acid-base" in the name.

Buffer solutions are solutions that maintain a constant pH value by diluting and adding small amounts of acid or base.

Classification of buffer systems

1. mixtures of solutions of weak acids and their salts. For example, acetate buffer.

2. mixtures of solutions of weak bases and their salts. For example, ammonium buffer solution.

3. mixtures of solutions of salts of polybasic acids of various degrees of substitution. For example, phosphate buffered saline.

4. ions and molecules of ampholytes. These include, for example, amino acids and protein buffering systems. While in the isoelectric state, amino acids and proteins are not buffering. The buffering effect appears only when a certain amount of acid or alkali is added to them. In this case, a mixture of two forms of protein is formed: a) a weak "protein acid" + a salt of this weak acid; b) weak "protein base" + salt of this weak base. Thus, this type of buffer systems can be attributed to the buffer systems of the first or second type.

Calculation of pH buffer solutions

The calculation of the pH of buffer systems is based on the law of mass action for acid-base equilibrium. For a buffer system consisting of a weak acid and its salt, for example, acetate, the concentration of ions H + easy to calculate based on the equilibrium constant of acetic acid:

CH 3 COOH ↔ CH 3 COO - + H +

(1).

From (1) it follows that the concentration of hydrogen ions is

(2)

In the presence of CH 3 COONa acid-base balance of acetic acid is shifted to the left. Therefore, the concentration of undissociated acetic acid is practically equal to the concentration of the acid, i.e. [CH 3 COOH] = with acid.

The main source of acetate ions is a strong electrolyte CH 3 COONa:

CH 3 COONa → Na + + CH 3 COO -,

Therefore, we can assume that [ CH 3 COO -] = with salt ... Taking into account the assumptions made, equation (2) takes the form:

From this, the Henderson-Hasselbach equation for buffer systems consisting of a weak acid and its salt is obtained:

(3)

For a buffer system consisting of a weak base and its salt, for example, ammonia, the concentration of hydrogen ions in solution can be calculated from the dissociation constant of the weak base.

NH 3 × H 2 O = NH 4 OH ↔ NH 4 + + OH -

(4)

Let us express the concentration of ions OH - from the ionic product of water

(5)

and substitute in (4).

(6)

From (6) it follows that the concentration of hydrogen ions is

(7)

In the presence of NH 4 Cl acid-base balance shifted to the left. Therefore, the concentration of undissociated ammonia is practically equal to the concentration of ammonia, i.e. [ NH 4 OH] = with basic.

The main source of ammonium cations is a strong electrolyte NH 4 Cl:

NH 4 Cl → NH 4 + + Cl -,

Therefore, we can assume that [ NH 4 +] = with salt ... Taking into account the assumptions made, equation (7) takes the form:

(8)

This gives the Henderson-Hasselbach equation for buffer systems consisting of a weak base and its salt:

(9)

In a similar way, you can calculate the pH of a buffer system consisting of a mixture of solutions of salts of polybasic acids of various degrees of substitution, for example, phosphate, consisting of a mixture of solutions of hydrophosphate ( Na 2 HPO 4 ) and dihydrogen phosphate ( NaH 2 PO 4 ) sodium. Its action is based on acid-base balance:

H 2 PO 4 - ↔ Н + + HPO 4 2-

Weak acid conjugated base

(10)

Expressing the concentration of hydrogen ions from (10) and making the following assumptions:

[H 2 PO 4 -] = c (H 2 PO 4 -); [HPO 4 2-] = c (HPO 4 2-), we get:

(11).

Taking the logarithm of this expression and changing the signs to the opposite, we obtain the Henderson-Hasselbach equation for calculating the pH of the phosphate buffer system

(12),

Where pK b (H 2 PO 4 - ) - negative decimal logarithm of the dissociation constant

phosphoric acid in the second stage; with ( H 2 PO 4 -) and with (HPO 4 2- ), respectively, the concentration of acid and salt.

Buffer properties

The pH value of the buffer solutions remains unchanged upon dilution, which follows from the Henderson-Hasselbach equation. When the buffer solution is diluted with water, the concentrations of both components of the mixture decrease by the same number once. Consequently, the pH value should not change in this case. However, experience shows that some change in pH, although insignificant, still occurs. This is due to the fact that the Henderson-Hasselbach equation is approximate and does not take into account interionic interactions. When making accurate calculations, one should take into account the change in the activity coefficients of the conjugated acid and base.

Buffer solutions have little change in pH when small amounts of acid or base are added. The ability of buffer solutions to maintain a constant pH when small amounts of a strong acid or strong base are added to them is based on the fact that one component of the buffer solution can interact with H+ added acid, and the other with OH- added base. As a consequence, the buffer system can bind as H + and OH - and maintain a constant pH value up to a certain limit. Let us demonstrate this using the example of a formate buffer system, which is a conjugated acid-base pair HCOOH / HCOO - ... Equilibrium in a formate buffer solution can be represented by the equation:

HCOOH ↔ HCOO - + H +

When a strong acid is added, the conjugated base HCOO - binds added ions H + , turning into weak formic acid:

HCOO - + H + ↔ HCOOH

In accordance with Le Chatelier's principle, the balance shifts to the left.

When alkali is added, the protons of formic acid bind the added OH ions- into water molecules:

HCOOH + OH - → HCOO - + H 2 O

Acid-base balance according to Le Chatelier shifts to the right.

In both cases, there are small changes in the ratio HCOOH / HCOO - , but the logarithm of this ratio changes little. Consequently, the pH of the solution also changes slightly.

The essence of buffering action

The effect of buffer solutions is based on the fact that the individual components of the buffer mixtures bind the hydrogen or hydroxyl ions of the acids introduced into them and the base with the formation of weak electrolytes. For example, if a buffer solution containing a weak acid HA n and the salt of this acid Kt А n , add alkali, then the reaction of the formation of a weak electrolyte-water will occur:

H + + OH → H 2 O

Therefore, if alkali is added to the buffer solution containing the acid, then the hydrogen ions formed during the electrolytic dissociation of the HA acid n , bind with hydroxyl ions of the added alkali, forming a weak electrolyte - water. Instead of consumed hydrogen ions, due to the subsequent dissociation of the HA acid n , new hydrogen ions appear. As a result, the previous concentration of Н+ - ions in the buffer solution will be restored to their original value.

If a strong acid is added to the specified buffer mixture, then the reaction will occur:

H + + A n - → HA n

those. A n - - ions formed during the electrolytic dissociation of salt K t А n combining with hydrogen ions of the added acid, they form weak acid molecules. Therefore, the concentration of hydrogen ions from the added strong acid to the buffer mixture practically does not change. The effect of other buffers can be explained in a similar way.

PH value in buffer solutions

Changing the ratios and you can get buffer

solutions that differ in a smooth change in pH from their minimum possible values. In an aqueous solution of a weak acid

[H +] = √K HAn * C HAn

where

pH = - log [Н +] = - - log K HAn - - log C HAn

But since K HAn is a constant value, then it is best to represent it in the form pK HAn those. electrolytic dissociation constant index: pK Han = - lg K HAn.

Then we get that in an aqueous solution of a weak acid:

pH = - log [Н +] = - - pK HAn - - pC HAn

As a weak acid is added to an aqueous solution of its salt, the pH of the solution will change.

According to the equation, in a solution containing a mixture of a weak acid and its salt [H+] = K HAn

then

pH = - log [Н +] = - log K HAn - log C HAn + log C Kt А n.

Similarly, we derive the formula for weak bases:

[OH] = √K KtOH * C KtOH

pOH = - log [ОН] = - - log K KtOH - - log C KtOH

The concentration of hydrogen ions is also expressed by the following formula [H+] =, so

pH = pK w - (- pK KtOH - - log C KtOH)

According to the equation, in a solution containing a mixture of a weak base and its salt

[H +] =

T . e.

pH = - log [Н +] = - log K w + log K KtOH - logC Kt А n + log C KtOH.

There is no need to memorize the derived formula for the pH value, since they are very easily derived by taking the logarithm of simple formulas expressing the value [H+ ].

Buffer capacity

The ability of buffer solutions to maintain a constant pH value is not unlimited and depends on the qualitative composition of the buffer solution and the concentration of its components. When significant amounts of a strong acid or alkali are added to the buffer solution, a noticeable change in pH is observed. moreover, for different buffer mixtures, differing from each other in composition, differing from each other in composition, the buffering action is not the same. Consequently, buffer mixtures can be distinguished by the strength of their resistance to the action of acids and alkalis added to the buffer solution in equal amounts and at a certain concentration. The limiting amount of acid or alkali of a certain concentration (in mol / l or g-eq / l) that can be added to the buffer solution so that its pH value changes by only one unit is called the buffer capacity.

If the value [Н + ] of one buffer solution changes with the added strong acid less than the value of [Н+ ] another buffer solution when the same amount of acid is added, the first mixture has a larger buffer capacity. For the same buffer solution, the higher the concentration of its components, the larger the buffer capacity.

Buffering properties of solutions of strong acids and bases.

Solutions of strong acids and bases at a sufficiently high concentration also have a buffering effect. In this case, the conjugate systems are H 3 O + / H 2 O - for strong acids and OH- / H 2 O - for strong bases. Strong acids and bases are completely dissociated into aqueous solutions and therefore are characterized by a high concentration of hydronium ionsor hydroxyl - ions. The addition of small amounts of a strong acid or strong base to their solutions, therefore, has only a slight effect on the pH of the solution.

Preparation of buffer solutions

1. By diluting the appropriate fixed channels in a volumetric flask.

2. By mixing the amounts of suitable conjugated acid-base pairs calculated by the Henderson-Hasselbach equation.

3. Partial neutralization of a weak acid with a strong alkali or a weak base with a strong acid.

Since the buffering properties are very weak, if the concentration of one component is 10 times or more different from the concentration of the other, buffer solutions are often prepared by mixing solutions of equal concentration of both components or adding an appropriate amount of reagent to the solution of one component, leading to the formation of an equal concentration of the conjugated form. Reference literature contains detailed recipes preparation of buffer solutions for different pH values.

The use of buffer solutions in chemical analysis

Buffer solutions are widely used in chemical analysis in cases where, according to the experimental conditions chemical reaction must proceed with the observance of the exact pH value, which does not change when the solution is diluted or when other reagents are added to it. For example, during the oxidation-reduction reaction, during the precipitation of sulfides, hydroxides, carbonates, chromates, phosphates, etc.

Here are some cases of using them for analysis purposes:

Acetate buffer solution (CH3COOH + CH 3 COO Na ; pH = 5) is used for precipitation of non-precipitated in acidic or alkaline solutions. The harmful effects of acids inhibit sodium acetate, which reacts with strong acids. For example:

HC1 + CH 3 COO N a → CH 3 COOH + Na C1

or in ionic form

H + + CH 3 COO → CH 3 COOH.

Ammonium-ammonium buffer solution ( N H 4 OH + N H 4 C1; pH = 9) is used in the precipitation of carbonates of barium, strontium, calcium and their separation from magnesium ions; when precipitating sulfides of nickel, cobalt, zinc, manganese, iron; as well as in the isolation of hydroxides of aluminum, chromium, beryllium, titanium, zirconium, iron, etc.

Formate buffer solution (HCOOH + HCOO N a; pH = 2) is used in the separation of zinc ions, precipitated in the form ZnS in the presence of ions of cobalt, nickel, manganese, iron, aluminum and chromium.

Phosphate buffer solution ( N а 2 НРО 4 + N аН 2 RO; pH = 8) is used in many oxidation-reduction reactions.

To successfully use buffers for assay purposes, remember that not all buffers are assayable. The buffer mixture is chosen depending on its purpose. It must satisfy a certain qualitative composition, and its components must be present in the solution in certain quantities, since the effect of buffer mixtures depends on the ratio of the concentration of their components.

The above can be presented in the form of a table.

Buffer solutions used in the analysis

Buffer mix	The composition of the mixture (at a molar ratio of 1: 1)	NS
Formate	Formic acid and sodium formate
Benzoate	Benzoic acid and ammonium benzoate
Acetate	Acetic acid and sodium acetate
Phosphate	Mono and dibasic sodium phosphate
Ammonium	Ammonium hydroxide and ammonium chloride

Mixtures of acidic salts with different substitution of hydrogen by metal also have a buffering effect. For example, in a buffer mixture of dihydrogen phosphate and sodium hydrogen phosphate, the first salt plays the role of a weak acid, and the second, the role of its salt.

By varying the concentration of a weak acid and its salt, it is possible to obtain buffer solutions with specified pH values.

Complex buffer systems also operate in animals and plants that maintain constant pH in blood, lymph and other fluids. Buffer properties the soil also possesses, which tends to counteract external factors that change the pH of the soil solution, for example, when acids or bases are introduced into the soil.

CONCLUSION

So, buffer solutions are called solutions that supportconstant pH when diluted and small amounts of acid or base are added. An important property buffer solutions is their ability to maintain a constant pH value when diluting the solution. Solutions of acids and bases cannot be called buffer solutions, because when diluted with water, the pH of the solution changes. The most effective buffer solutions are prepared from solutions of a weak acid and its salt or a weak base and its salt

Buffer solutions can be considered as a mixture of electrolytes with ions of the same name. Buffer solutions play an important role in many technological processes. They are used, for example, in the electrochemical deposition of protective coatings, in the production of dyes, leather, photographic materials. Buffer solutions are widely used in chemical analysis and for calibrating pH meters.

Many biological fluids are buffer solutions. For example, the pH of the blood in the human body is maintained in the range from 7.35 to 7.45; gastric juice from 1.6 to 1.8; saliva from 6.35 to 6.85. The components of such solutions are carbonates, phosphates and proteins. In bacteriological studies, when growing bacteria, it is also necessary to use buffer solutions.

BIBLIOGRAPHIC LIST

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3. Kreshkov A.P., Yaroslavtsev A.A. Analytical chemistry course. Book 1. Qualitative Analysis - 2nd ed. Revised. - M.: Chemistry, 1964 - 432 p.

4. Chemistry: a reference book for high school students and those entering universities / Ed. Lydia R.A., Alikberova L.Yu. - M.: AST-PRESS SCHOOL, 2007. -512s.

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