In analytical chemistry, buffer solutions are very often used. Buffer are called solutions, the pH of which practically does not change when small amounts of acids and bases are added to them or when they are diluted. Buffer solutions can be of four types.

1. Weak acid and her salt. For example, acetate buffer solution CH 3 COOH + CH 3 COONa.

2. Weak base and its salt. For example, ammonia buffer solution NH 4 OH + NH 4 C1.

3. A solution of two acidic salts. For example, phosphate buffered saline NaH 2 PO 4 + Na 2 HPO 4. In this case, the NaH 2 PO 4 salt plays the role of a weak acid.

4. Amino acid and protein buffers. pH and pOH buffer solutions depend on the value of the dissociation constant of the acid or base and on the ratio of the concentrations of the components. This dependence is expressed by the equations

pH = p K k- lg C (acid) (2.6)

pOH = pK 0- lg C (base),(2.7)

where pK k and pK 0- indicators of the dissociation constant of the corresponding acid and base; С (acid) - acid concentration; C (base) is the concentration of the base; С (salt) - salt concentration.

When preparing a buffer solution with the same concentration of acid (base) and salt, the pH or pOH of such a solution is numerically equal to pK k or pK 0, since C (acid) / C (salt) = 1 or C (base) / C (salt) = 1. By changing the ratio between the concentrations of acid (base) and salt, you can get a series of solutions with different concentrations of hydrogen ions, i.e. ... with different pH values.

Using the example of an acetate buffer solution, let us consider what is the basis of the property of buffer solutions to maintain a constant pH value. For an acetate buffer, the pH can be calculated using equation (2.6):

pH = pKsn 3 coon - lg C (CH 3 COOH) . (2.8)

When the acetate buffer solution is diluted with water, as can be seen from equation (2.8), the C (CH 3 COOH) / C (CH 3 CОONa) ratio does not change, since the concentrations of acid and salt decrease by the same number of times, and рКсн 3 СООН remains constant size. As a result, the pH of the buffer solution remains practically unchanged upon dilution.

Now suppose that 1 L of acetate buffer solution is prepared with the same concentration of both components equal to 0.1 M. For acetic acid pK= 4.76. Therefore, according to equation (2.8), the pH of such a buffer solution is equal to the following value:

pH = 4.76 - log0.1 / 0.1 = 4.76.

Add 10 millimoles of hydrochloric acid to this solution. As a result of the reaction

CH 3 COONa + HC1 → CH 3 COOH + NaCl

the concentration of the weak acid increases and the concentration of the salt decreases. The concentration of acetic acid will be equal to 0.1 M + 0.01M = 0.11M, and the concentration of the CH 3 COONa salt: 0.1M - 0.01M = 0.09M. Then the pH of the acetate buffer solution decreases by 0.08:

pH = 4.76 - log (0.11 / 0.09) = 4.76 - 0.079 = 4.68.

When the same amount of base is added instead of a strong acid, the latter reacts with acetic acid:

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

The acid concentration decreases (0.1M - 0.01M = 0.09M), but the salt concentration increases (0.1M + 0.01M = 0.11M). Then

pH = 4.76 - lg (0.09 / 0.11) = 4.76 - 0.09 = 4.67.

When an acid or base is added, the concentrations of the components of the buffer solution change insignificantly, and after equilibrium is established, the pH also changes insignificantly.

Adding 10 millimoles of HCl or NaOH to 1 liter of water creates a concentration of [H +] and [OH -] equal to 0.01M. In the first case, the pH will be equal to 2, in the second - 12, i.e. The pH will change by 5 units compared to the pH of pure water.

The ability of buffers to maintain pH substantially constant is limited. Any buffer solution practically maintains a constant pH only until a certain certain amount of acid or alkali is added. The ability of a buffer solution to resist a shift in pH is measured buffer capacity. This value is characterized by the number of moles of H + or OH, respectively, of a strong acid or alkali, which must be added to 1 liter of the buffer solution in order to shift the value of its pH by one unit.

Buffer solutions are widely used in qualitative and quantitative analysis to create and maintain a certain pH value of the medium during reactions. Thus, Ba 2+ ions are separated from Ca 2+ and Sr 2+ ions by precipitation with dichromate ions Cr 2 O 7 2- in the presence of an acetate buffer solution. In the determination of many metal cations using Trilon B by complexometry, an ammonia buffer solution (NH 4 OH + NH 4 Cl) is used.

Buffer solutions or buffer systems ensure a constant pH of biological fluids and tissues. The main buffer systems in the body are hydrocarbonate, hemoglobin, phosphate and protein. The action of all buffer systems in the body is interconnected. Hydrogen ions received from the outside or formed in the course of metabolism are bound into weakly dissociated compounds by one of the components of the buffer systems. However, in some diseases, a change in the pH value of the blood may occur. The shift of the blood pH value to the acidic region from the normal pH value of 7.4 is called acidosis to the alkaline region - alkalosis. Acidosis occurs in severe forms of diabetes mellitus, prolonged physical work and inflammatory processes. Alkalosis may occur if renal or hepatic impairment is severe, or if breathing is impaired.

QUESTIONS AND EXERCISES

1. What are buffers?

2. What are the main types of buffers? Give examples.

3. What does it depend on pH buffer solutions?

4. Why pH acetate buffer does not change significantly when small amounts are added to it nitric acid or potassium hydroxide?

5. Will it change pH phosphate buffer solution when diluted with water 10 times? Give an explanation.

6. Calculate: a) pH phosphate buffer solution, consisting of 16 ml of solution Na 2 HPO 4 with a concentration of 0.1 mol / l and 40 ml of solution NaH 2 PO 4 with a concentration of 0.04 mol / l, if pKH 2 PO - 4 = 6,8; b) how will it change pH this solution by adding 6 ml of solution to it HC1 with a concentration of 0.1 mol / l.

Answer: a) pH = 6.8; b) pH = 6,46; ∆рН = 0,34.

7. Give examples of the use of buffer solutions in analytical chemistry.

8. What is: a) acidosis; b) alkalosis?

The buffer solution is used to maintain a constant pH value. It consists of a mixture of a weak acid HA and a conjugated base A -. Equilibria coexist in a buffer solution:

HA + H 2 O ↔ H 3 O + + A -

A - + H 2 O ↔ HA + OH -

suppressing each other at sufficiently high C (HA) and C (A -); therefore, we can assume that [HA] = C (HA) and [A -] = C (A -). Using an expression for K a HA and neglecting the contribution of [H 3 O +] due to the dissociation of water, we obtain

The same expression can be obtained using the second equilibrium constant.

EXAMPLE 16. Calculate the pH of the buffer solution consisting of 0.10 M acetic acid and 0.10 M sodium acetate.

Solution. Here, all the conditions are met that make it possible to apply the formula (2-14) (acetic acid is a weak acid, the concentrations of the acid and conjugated base are quite high). So

EXAMPLE 17. Calculate the pH of the buffer solution consisting of 0.10 M ammonia and 0.20 M ammonium chloride.

Solution. By the formula (2-14) we find

An important characteristic of a buffer solution is its buffer capacity. Adding a strong base (acid) to a buffer solution, its pH can change with a change in the concentration of acid HA and conjugated base A -. Therefore, it is customary to represent the buffer capacity in the form

if added to the buffer solution strong foundation, and

if a strong acid is added to the buffered solution. Let us write down the material balance equation for a mixture of monobasic acid HA and conjugated base A -:

We express [ON] in terms of K a HA and substitute it into the material balance equation. Find [A -]:

(2-17)

Differentiating equation (2-17) with respect to dpH, taking into account that dc main =, we obtain

(2-18)

It is easy to see that at pH = pK and HA, i.e. - C (HA) = C (A -), the maximum buffer capacity is reached. It can be shown that

(2-19)

Formulas (2-18) and (2-19) follow from one another, if we remember that [HA] = a(HA) C (HA) and [A -] = a(A -) C (A -), as well as expressions for a(ON) and a(A -).

For highly diluted buffer solutions, the contribution of water dissociation should be taken into account. In this case, equation (2-19) becomes more complicated:

Here the first two terms describe the buffering action of water, the third - the buffering action of the acid and the conjugated base.

EXAMPLE 18. Calculate how the pH will change if 1.0 · 10 -3 mol of hydrochloric acid is added to 1.0 L of a buffer solution consisting of 0.010 M acetic acid and 0.010 M sodium acetate.

Solution. We calculate the pH of the buffer solution before adding hydrochloric acid:

The total concentration of the buffer solution is

For such a sufficiently concentrated buffer solution, the buffer capacity should be calculated using the formula (2-18):

Calculation using formula (2-19) gives the same result:

Calculate the change in pH

Thus, after adding hydrochloric acid, the pH of the buffer solution will be

pH = 4.75 - 0.087 = 4.66

This problem can be solved without resorting to calculating the buffer capacity, but by finding the quantities of the components of the buffer mixture before and after the addition of HC1. In the original solution

EXAMPLE 19... Derive an expression for the maximum buffer capacity of a solution with a total concentration of components with.

Solution. Let us find the conditions under which the buffer capacity is maximum. To do this, we differentiate expression (2-18) by pH and equate the derivative to zero

Hence [H +] = K and HA and, therefore, C (HA) = C (A -).

Using formulas (2-19) and (2-21), we obtain that

Calculation of the pH of mixtures of acids or bases. Let the solution contain two acids HA 1 and HA 2. If one acid is much stronger than the other, then almost always the presence of a weaker acid can be neglected, since its dissociation is suppressed. Otherwise, the dissociation of both acids must be taken into account.

If HA 1 and HA 2 are not too weak acids, then neglecting the autoprotolysis of water, the electroneutrality equation can be written as:

[H 3 O +] = [A 1 -] +

Let us find the equilibrium concentrations A 1 - and A 2 1 from the expressions for the dissociation constants HA 1 and HA 2:

Substitute the obtained expressions into the electroneutrality equation

After transformation we get

If the degree of dissociation of acids does not exceed 5%, then

For a mixture of P acids

Similarly for a mixture of monobases

(2-21)

where Section 1 and K a 2 - dissociation constants of conjugated acids. In practice, more often, perhaps, there are situations when one (one) of the acids (bases) present in the mixture suppresses the dissociation of others and therefore, to calculate the pH, the dissociation of only this acid (this base) can be taken into account, and the dissociation of the rest can be neglected. But other situations may also occur.

EXAMPLE 20. Calculate the pH of a mixture in which the total concentrations of benzoic and aminobenzoic acids are 0.200 and 0.020 M, respectively.

Solution. Although the values ​​of the dissociation constants of benzoic (K a= 1.62 · 10 -6, we denote K 1) and aminobenzoic (K a = 1.10 · 10 -5, we denote K 2) acids differ by almost two orders of magnitude, because of the rather large difference in acid concentrations, here it is necessary to take into account the dissociation of both acids. Therefore, by formula (2-20), we find

For the HA / A type I buffer system, the concentration of H + ions in solution can be easily calculated based on the dissociation constant of a weak acid (for simplicity of presentation, instead of ion activities in the expression for, we will use their concentrations):

HA ⇄ A - + H +;

where C (acid) and C (salt)- molar concentrations of acid and salt.

If equality (3) is logarithmized (take negative decimal logarithm left and right sides of the equation), then we get:

where the index "0" denotes the characteristics of the initial solutions of the acid and salt, by mixing which the required buffer mixture is obtained.

For a buffer system of type II B / BH +, for example, ammonium, hydroxide and hydrogen indices are calculated by the equations:

where is the index of the base dissociation constant.

V general view the equation for calculating the pH of buffer systems is as follows:

,

(7)

and is called the equation Henderson-Hasselbach.

It follows from the Henderson-Hasselbach equation that:

1. The pH value of buffer solutions depends on the dissociation constant of an acid or base and on the ratio of the amounts of components, but practically does not depend on dilution or concentration of solutions. Indeed, in these processes, the concentrations of the components of the buffer solution change proportionally; therefore, their ratio, which determines the pH of the buffer solution, remains unchanged.

If the concentrations of the components of the buffer solutions exceed 0.1 mol / L, then the calculations must take into account the activity coefficients of the ions of the system.

2. The indicator of the dissociation constant of a weak electrolyte determines the area of ​​the buffer action of the solution, i.e. that range of pH values ​​in which the buffer properties of the system are preserved. Since the buffering action continues until 90% of the component is consumed (i.e., its concentration has not decreased by an order of magnitude), the area (zone) of the buffering action differs from by 1 unit:

Ampholytes can have several zones of buffer action, each of which corresponds to the corresponding constant:

.

Thus, the maximum allowable ratio of the components of the solution, at which it exhibits a buffering effect, is 10: 1.

Example 1. Is it possible to prepare acetate buffer with pH = 6.5 if acetic acid is 4.74?

Since the buffer zone is defined as , for the acetate buffer, it is in the range of pH values ​​from 3.74 to 5.74. The pH value = 6.5 lies outside the zone of action of the acetate buffer; therefore, such a buffer cannot be prepared on the basis of the acetate buffer system.

Buffer capacity.

It is possible to add acid or alkali, without significantly changing the pH of the buffer solution, only in relatively small amounts, since the ability of buffer solutions to maintain a constant pH is limited.

The value that characterizes the ability of the buffer solution to resist the shift in the reaction of the medium when acids and alkalis are added is called buffer capacity (B). Distinguish the buffer capacity for acid () and alkali ().

Buffer capacity (B) is measured by the amount of acid or alkali (mol or mmol equivalent), the addition of which to 1 liter of buffer solution changes the pH by one.

In practice, the buffer capacity is determined by titration. For this, a certain volume of the buffer solution is titrated with a strong acid or alkali of a known concentration until the equivalence point is reached. Titration is carried out in the presence of acid-base indicators, with the right choice which fix the state when the component of the buffer system reacts completely. Based on the results obtained, the value of the buffer capacity (or) is calculated:

	(8)
	(9)

where WITH( to-you), WITH( click) - molar concentrations of acid and alkali equivalent (mol / l);

V (to-you), V (crack) - volumes of added solutions of acid or alkali (l; ml);

V (buffer) - the volume of the buffer solution (l; ml);

pH 0 and pH - the pH value of the buffer solution before and after titration with acid or alkali (the change in pH is taken as an absolute value).

The buffer capacity is expressed in [mol / L] or [mmol / L].

The buffer capacity depends on a number of factors:

1. The higher the absolute content of the components of the base / conjugate acid pair, the higher the buffer capacity of the buffer solution.

The buffer capacity depends on the ratio of the amounts of the components of the buffer solution, and, consequently, on the pH of the buffer. The buffer capacity is maximum at equal amounts of the components of the buffer system and decreases with a deviation from this ratio.

3. With different content of components, the buffer capacities of the solution for acid and alkali differ. So, in a buffer solution of type I, the higher the acid content, the greater the alkali buffer capacity, and the higher the salt content, the greater the acid buffer capacity. In a type II buffer solution, the higher the salt content, the greater the alkali buffer capacity, and the higher the base content, the greater the acid buffer capacity.

Chapter 6. PROTOLYTIC BUFFER SYSTEMS

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 action, ion exchange, 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 B - anions to a solution - leads to an ionization process that shifts the equilibrium to the left (common ion effect) 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 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

Acidic 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 their 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 considered as a mixture of weak and strong electrolytes with ions of the same name (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 I = 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 [Н +], К И, α and С to / С with. The equation

we deduce by 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 to inverse relationship from C s and the ratio C to / C s. 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 in order for its pH value to change by one is called the buffer capacity.

Thus, buffering capacity is a quantitative measure of the buffering action of a solution. The buffer solution has a maximum buffering capacity at pH = pK of an acid or base that forms a mixture when the ratio of its components is 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

central and extremely important role in the transport of CO 2 and the establishment of pH. 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 carboanhydrase, 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.

Buffer solution is a chemical reagent with a constantpH

Laboratory glassware, laboratory equipment, instruments and chemical substances these are the four main components of any modern laboratory, regardless of its specialization. Depending on the purpose, laboratory products - glassware, equipment, instruments are made of various materials: plastic, porcelain, quartz, borosilicate, laboratory glass, etc. It is only a matter of price and quality. Chemical reagents occupy a special place in the list of laboratory equipment - without them it is impossible to carry out even the simplest analysis, research, experiment.

In the practice of laboratory work employees often face such chemical solutions, which have or should have a certain indicator of the pH value. It is for these purposes that special buffer solutions are made.

What is this solution?

Buffer solutions - chemical reagents with a certain stable indicator of the concentration of hydrogen ions; mixture is weak concentrated acid and her salt. These solutions practically do not change their structure when concentrated, diluted with other chemical reagents, or when a small amount of highly concentrated alkalis or acids is added to it. To obtain a buffer solution with different pH values, it is necessary to change the concentration and ratio of the chemical solutions used.

This chemical reagent is able to maintain a certain pH value up to a certain level, depending on the specific amount of aggressive media, alkalis and acids. Each buffer has a specific buffer capacity - the equivalent ratio of alkali to acid elements.

Unfortunately, the acids and alkalis themselves cannot be attributed to buffer mixtures, since when they are diluted with water, the pH level of these aggressive media changes.

V laboratory practice the calibration buffer is also applicable. It is designed to adjust the accuracy of indicators of instruments that are used to determine the level of acid liquid substances- activity in various environments of hydrogen ions.

For work both in laboratory conditions and in private practice, it is recommended to use buffer mixtures of high stability prepared in specialized laboratories using laboratory glassware on special laboratory equipment and instruments. Self-preparation of this chemical reagent can be obtained with a large error.

What does the buffer solution consist of?

This chemical reagent contains water - a solvent and equally dissolved ions or molecules of substances that make up an acid-base or alkaline-acid buffer system. The buffer system is the interaction of a weakly concentrated acid with one of its salts.

Such chemical reagents, together with modern laboratory equipment and instruments, have found wide application in research in analytical chemistry, biology and microbiology, genetics, medicine, pharmaceuticals, research centers and other scientific fields.

The importance of buffered saline in humans

The natural buffer mixture is very important for the normal functioning of the body, as it maintains a constant pH level in biological fluids of tissues, organs, lymph and blood.

Storage conditions

Store this chemical reagent in a hermetically sealed container (glass or plastic bottles).

Where to buy laboratory equipment High Quality at an affordable price?

It is profitable to buy chemical reagents, devices, equipment, laboratory glassware in Moscow in a modern specialized store of chemical reagents Moscow retail and wholesale "Prime Chemicals Group". It is here that you will find a wide range of high quality goods from well-known brands at affordable prices. We also offer delivery both in the city and in the region.

“Prime Chemicals Group” - laboratory equipment from examination gloves to electronic laboratory scales with a quality mark.