What are buffer solutions needed for? Calculation of pH of buffer solutions. Buffer solutions can be divided into four types

For a type I buffer system HA/A, the concentration of H + ions in the 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):

NA ⇄ A - + H + ;

Where C(acid) And C(salt)– molar concentrations of acid and salt.

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

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

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

where is the index of the base dissociation constant.

In general, the equation for calculating the pH of buffer systems is as follows:

(7)

and is called the equation Henderson-Hasselbach.

From the Henderson-Hasselbach equation it follows that:

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

If the concentrations of the components of buffer solutions exceed 0.1 mol/l, then the activity coefficients of the system ions must be taken into account in the calculations.

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 permissible ratio of solution components at which it exhibits a buffering effect is 10:1.

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

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

Buffer capacity.

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

The value characterizing the ability of a buffer solution to counteract the displacement of the reaction of the medium when adding acids and alkalis is called buffer capacity (B). Buffer capacity is distinguished by acid () and alkali ().

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

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

	(8)
	(9)

Where WITH( whoa), WITH( slot) - molar concentrations of acid and alkali equivalent (mol/l);

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

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

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

Buffer capacity is expressed in [mol/l] or [mmol/l].

Buffer capacity depends on a number of factors:

1. The greater 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 components of the buffer solution, and therefore on the pH of the buffer. The buffer capacity is maximum with equal quantities of components of the buffer system and decreases with deviation from this ratio.

3. With different contents of components, the buffer capacities of the solution for acid and alkali are different. Thus, in a type I buffer solution, 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 greater the salt content, the greater the alkali buffer capacity, and the greater the base content, the greater the acid buffer capacity.

Where C(acid) And C(salt)– molar concentrations of acid and salt.

If equality (3) is taken logarithmically (take the negative decimal logarithm of the left and right sides of the equation), we obtain:

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

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

where is the index of the base dissociation constant.

In general, the equation for calculating the pH of buffer systems is as follows:

(7)

and is called the equation Henderson-Hasselbach.

From the Henderson-Hasselbach equation it follows that:

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

If the concentrations of the components of buffer solutions exceed 0.1 mol/l, then the activity coefficients of the system ions must be taken into account in the calculations.

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

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

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

Solution.

Example 2. Calculate the pH of a buffer solution, 100 ml of which contains 1.2 g of acetic acid and 5.88 g of potassium acetate, if for acetic acid = 4.74.

Solution.

The molar concentrations of acid and salt in the buffer solution are:

Substituting these values into equation (7), we obtain:

Solution.

Since the molar concentrations of acid and salt are equal, when calculating pH using formula (5), only the volume ratio of the components can be used:

Example 4. Calculate the pH value of the buffer solution obtained by pouring 20 ml of ammonia water solution with C(NH 3 H 2 O) = 0.02 mol/l and 10 ml of ammonium chloride solution with C(NH 4 Cl) = 0.01 mol/ l. (NH 3 H 2 O) = 1.8 10 −5. Find the pH of the buffer diluted 5 times.

Solution.

In the case of a type II buffer system, the pH of the solution is calculated using equation (6¢):

Substituting the corresponding values, we get:

When diluted, the pH of buffer solutions does not change. Therefore, the pH of a buffer solution diluted 5 times will be 9.86.

Example 5. The buffer solution was obtained by pouring 100 ml of a CH 3 COOH solution with C(CH 3 COOH) = 0.02 mol/l and 50 ml of a CH 3 COONa solution with C(CH 3 COONa) = 0.01 mol/l. (CH 3 COOH) = 1.8×10 -5. Calculate:

a) pH of the resulting buffer;

b) change in pH of the buffer when adding 5 ml of HCl solution with C(HCl) = 0.01 mol/l.

c) buffer capacity of the solution for alkali.

Solution.

To calculate the pH of the resulting buffer, we use formula (5):

When an acid is added, the following reaction occurs:

CH 3 COONa + HCl CH 3 COOH + NaCl,

as a result of which the quantities of components of the buffer system change.

Taking into account the relation n(x) = C(x)×V(x), equation (7) can be presented as:

Since the amounts of reacted and formed substances are equal, the change in the amounts of acid and salt in the buffer solution will be the same value x:

In the initial buffer mixture the quantities of components are:

Let's find the value x:

Thus, the difference in pH values will be , i.e. the change in pH is negligible.

Buffer capacity.

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

In practice, the buffer capacity is determined by titration. To do this, a certain volume of the buffer solution is titrated with a strong acid or alkali of known concentration until the equivalence point is reached. Titration is carried out in the presence of acid-base indicators, with the correct choice of which the state is recorded 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( whoa), WITH( slot) - molar concentrations of acid and alkali equivalent (mol/l);

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

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

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

Buffer capacity is expressed in [mol/l] or [mmol/l].

Buffer capacity depends on a number of factors:

1. The greater 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 components of the buffer solution, and therefore on the pH of the buffer. The buffer capacity is maximum for equal quantities of buffer system components and decreases with deviation from this ratio.

Example 2. To prepare acetate buffer mixtures, solutions of acid and salt of the same molar concentration were mixed in the following volume ratios:

Composition of the buffer system	Volume ratios of buffer system components
solution I	solution II	solution III
CH3COOH
CH 3 COONa

Without resorting to calculations, determine in which of the three buffer solutions the following will be observed:

a) the highest pH value;

b) maximum buffer capacity;

c) the largest buffer capacity for acid.

Solution.

In the case of equal concentrations of components, equation (5) takes the form:

Since it is the same in all three solutions, the pH value of the buffer will be determined by the ratio. Hence, highest value pH solution I will have ():

Solution II is characterized by the maximum buffer capacity, since the ratio of the components in it is 1:1.

The acid buffer capacity for an acetate buffer is determined by the content of the conjugate base, i.e. salts: the higher it is, the greater the acid buffer capacity of the solution. That's why:

Thus, solution I will have the greatest acid capacity.

The student must be able to:

1. Calculate the pH of buffer systems.

2. Calculate the buffer capacity of the solution.

Solutions whose pH remains almost unchanged by the addition of small volumes of strong acids and alkalis, as well as by dilution, are called

buffer.

Most often, mixtures of solutions of weak acids and their salts, or mixtures of solutions of weak bases and their salts, or, finally, mixtures of solutions of salts of polybasic acids of varying degrees of substitution are used as buffer solutions.

For example: UNDC

formate, pH = 3.8

CH3 COOH

acetate, pH = 4.7

CH3 COONa

NaH2PO4

phosphate, pH = 6.6

Na2HPO4

	NH4OH		ammonia, pH = 9.25
	NH4CI

	Let's consider the mechanism of action of buffer systems:
	1. When an acid is added to a solution, its hydrogen ions bind into
	weak acid:
	CH3 COOH		CH3 COOH


	CH3 COONa		CH3 COOH
	2. When a base is added to a solution, the hydroxide ion binds to
	weak electrolyte (H2 O):
	CH3 COOH		CH3 COONa


	CH3 COONa		CH3 COONa

The formation of weak electrolytes when an acid or base is added to a buffer solution determines pH stability.

Calculation of pH of buffer solutions

1. Buffer solutions formed	pH = pKacid -	With acid
weak acid and its salt
		With salt
		With salt
	pK – strength indicator of acid:
	рК = – log Kacids

2. Buffer solutions formed	pOH = pCobas.		From the base
weak bases and their salts.
			With salt
			With salt
	knowing that pH + pH = 14, hence
	pH = 14 - pHKn.		From the base

			With salt
			With salt

The ability of buffer systems to maintain a constant pH is determined by its buffer capacity. It is measured by the number of mole equivalents of a strong acid or strong base that must be added to 1 liter

buffer system of the solution to change the pH by one.

We calculate the capacity of the buffer mixture using the formulas:

where B is the buffer capacity;

CA, CB – concentrations of substances in the buffer mixture.

The higher the concentration of the mixture components, the greater the buffer capacity. In order for the action of the buffer mixture to be sufficiently effective, that is, so that the buffer capacity of the solution does not change too much,

the concentration of one component should not exceed the concentration of another component by more than 10 times.

EXAMPLES OF SOLVING TYPICAL PROBLEMS

Calculation of pH of buffer solutions formed

weak acid and its salt

Example 1. Calculate the pH of a mixture of 0.03 N solution of acetic acid CH3 COOH with

0.1 N solution of CH3 COONa, if the strength index of the acid pK = 4.8.

pK(CH3 COOH) = 4.8 C(f(CH3 COOH) =

0.03 mol/l C(f(CH3 COONa) =

Since M(f) = M for CH3 COOH and CH3 COONa, then for these substances C = C(f)

pH pKacid. - lg Sour. Ssoli

pH 4.8 - lg 0.03 4.8 lg 0.3 4.8 - (-0.52) 5.32 0.1

Answer: pH = 5.32

Example 2. Calculate the pH of a solution obtained by mixing 20 ml

0.05m solution nitrous acid HNO2 and 30 ml of 1.5 m sodium nitrite solution

NaNO2.


V(HNO2) = 20 ml	1. Find the volume of the solution after mixing
C(HNO2) = 0.05 mol/l	acids HNO2 and salts NaNO2 and their concentrations
V(HNO2) = 30 ml	in the resulting mixture:
C(HNO2) = 1.5 mol/l	V = 20 + 30 = 50 ml

	C(HNO2)	0.02 mol/l

2. From the table we find that pK HNO 2 = 3,29.

3. Calculate pH:

C(NaNO2) 1.5 30 0.9 mol/l Answer: pH = 4.94 50

Example 3. How much 0.5 m solution of sodium acetate CH3 COONa must be added to 100 ml of 2 m solution of acetic acid CH3 COOH to obtain a buffer solution with pH = 4?


C(CH3 COONa) = 0.5 mol/l

		With salt
	With sour

	With salt

Therefore, the ratio of acid concentration to salt concentration

should be equal to 5.754:1.

2. Find the acid concentration in the buffer system:

4. Find the amount of 0.5 m solution of sodium acetate CH3 COONa containing

Example 4. In what molar ratios should solutions of salts of the composition NaH2 PO4 and Na2 HPO4 be taken to obtain a buffer system with pH = 6?

1. According to the conditions of the problem, we only know the pH value. Therefore, according to

Using the pH value, we find the concentration of hydrogen ions:

pH = - log = 6 or log = –6. Hence = 10-6 mol/l.

2. In this buffer system, the H2 PO4 ion acts as an acid

NaH2 PO4  Na+ + H2 PO4 ¯ K2 (H3 PO4) = 6.2 10 -8.

3. Knowing the concentration of hydrogen ions and the value of the constant

acid dissociation, we calculate the ratio of acid concentration to salt concentration in a given buffer system:

		C acid.
K2(H3PO4)			or = K2 (H3 PO4 )
K2(H3PO4)			or = K2 (H3 PO4 )
		With salt

				1 10 - 6

	K2(H3PO4)
	K2(H3PO4)

Calculation of pH of buffer systems formed

weak bases and their salts

Example 5. Calculate the pH of a buffer solution containing 0.1 mol/l NH4 OH

and 0.1 mol/l NH4 Cl, if the dissociation constant of NH4 OH is 1.79 10-5.

С(NH4 OH) = 0.1 mol/l

С(NH4 Cl) = 0.1 mol/l

КNH4OH = 1.79 10–5

1. pK NH 4 OH - log 1.79 10 -5 - (0.25- 5) 4.75

2.pH 14 - pKbas. lg Pine

With salt

14 - 4.75 lg 0.1 9.25 0.1

Answer: pH=9.25.

Example 6. Calculate the pH of an ammonia buffer system containing 0.5 m

ammonium hydroxide and ammonium chloride. How will the pH change when added to

1 liter of this mixture is 0.1 m HCI and when adding 0.1 m NaOH to 1 liter of the mixture and diluting the solution with water 10 times, if pK NH4 OH = 4.75?

C(NH4 OH)= 0.5 mol/l

С(NH4 Cl) = 0.5 mol/l

С(HCl) = 0.1 mol/l

С(NaOH) = 0.1 mol/l

p KNH 4 OH = 4.75

1. pH before dilution - ?

2. pH after adding HCI - ?

3. pH after adding NaOH - ?

4. pH after dilution with water - ?

pH 14 - pK lg C basic.

With salt

1. pH 14 - 4.75 lg 0.5 0.5 9.25

2. When adding 0.1 m HCl to a buffer solution, the concentration of NH 4OH

will decrease by 0.1 m and become equal

0.4 m, and the concentration of NH4 CI increases to 0.6 m. Therefore:

pH 14 - 4.75 lg 0.4 0.6 9.074

3. When adding 0.1 m NaOH to 1 liter of this mixture, the concentration of NH4 OH

will increase to 0.6 m, and the concentration of NH4 Cl will decrease to 0.4 m. As a result, we obtain: pH 14 - 4.75 lg 0.6 0.4 9.426

4. When diluting the buffer solution with water 10 times, we will have: pH 14 - 4.75 lg 0.05 0.05 9.25

Example 7. Calculate pOH and pH of a solution containing 8.5 g of ammonia in 1 liter and

107 g ammonium chloride.


m(NH3) = 8.5 g	1. Find molar concentrations
m(NH4Cl) = 107 g	ammonia and ammonium chloride:

rON -? pH - ?	C(NH3)

C(NH4CI)

2. Calculate pOH and pH:

C base

C salt

4,75 (0,6) 5,35 ;

Answer: pH = 8.65, pH = 5.35

Calculation of buffer capacity

buffer

mixture, if it is obtained by

mixing 0.1 m CH3 COOH and 0.1 m CH3 COONa?

C(CH3 COOH) = 0.1 mol/l

Because C(CH3 COOH) = C(CH3 COONa) = 0.1 m, then

С(CH3 COONa) =

we use the formula:

0.1mol/l

C A C B

0,12

0.115 mol/l

C(CH3 COONa) =

because

= K C

KCH 3 COOH = 18 10 –5 C = 1 mol/l

In order to lower the pH by one, you need to add the following to the solution:

number of moles of acid at which Acid 10

Therefore, we can create the equation:

1. SELF-CONTROL TASKS

What is the pH of a mixture consisting of 100 ml of 23N HCOOH and 30 ml of 15N

2. HCOOK solution. How will the pH of a buffer solution composed of 0.01 m Na change?

2 HPO4 and

3. 0.01 m NaH2 PO4, if you add 10–4 mol HCl to it. Calculate the pH of a solution containing 0.05 mol/l NH

4 OH and 0.05 mol/l

4. NH4 Cl (КNH4 OH = 1.8 10-5 ). Calculate the buffer capacity of a solution containing 0.4 mol Na in 1 liter

2 HPO4

The pH of buffer solutions is calculated using the Henderson–Hasselbach equation:

– for an acid buffer the equation has the form

– for the main buffer

The equations show that the pH of a buffer solution of a given composition is determined by the ratio of the concentrations of acid and salt or base and salt, and therefore does not depend on dilution. When the volume of the solution changes, the concentration of each component changes by the same number of times.

Buffer capacity

The ability of buffer solutions to maintain a constant pH is limited. Those. adding acid or alkali without significantly changing the pH of the buffer solution is possible only in limited quantities.

The value characterizing the ability of a buffer solution to counteract the displacement of the reaction of the medium when acids and alkalis are added is called the buffer capacity of the solution (B).

Buffer capacity is measured by the number of moles of equivalents of a strong acid or alkali, the addition of which to 1 liter of a buffer solution changes the pH by one.

Mathematically, the buffer capacity is defined as follows:

B by acid (mol/l or mmol/l):

where n(1/z HA) is the number of moles of acid equivalents, pH 0 and pH is the pH of the buffer solution before and after adding the acid, V B is the volume of the buffer solution.

In alkali (mol/l or mmol/l):

where n (1/z BOH) is the number of moles of alkali equivalents, the rest of the designations are the same.

Buffer capacity depends on a number of factors:

1. From the nature of the added substances and components of the buffer solution. Because Some substances can form insoluble compounds or complexes or give other undesirable reactions with components of the buffer system, then the concept of buffer capacity loses its meaning.

2. From the initial concentration of the components of the buffer system.

The greater the number of components of an acid-base pair in a solution, the greater the buffer capacity of this solution.

The limit of the ratio of the concentrations of the components of the buffer solution at which the system still retains its properties. The pH interval = pK ± 1 is called the buffer zone of the system. This corresponds to the range of the salt / salt ratio from 1/10 to 10/1.

In k (blood) = 0.05 mol/l; V to (plasma) = 0.03 mol/l; V to (serum blood) = 0.025 mol/l

Blood buffer systems

Especially great importance Buffer systems have a role in maintaining the acid-base balance of organisms. The pH value of most intracellular fluids is in the range from 6.8 to 7.8.

Acid-base balance in human blood is ensured by hydrocarbonate, phosphate, protein and hemoglobin buffer systems. The normal pH value of blood plasma is 7.40 ± 0.05.

The hemoglobin buffer system provides 35% buffer capacity of the blood: . Oxyhemoglobin is a stronger acid than reduced hemoglobin. Oxyhemoglobin usually comes in the form of a potassium salt.

Carbonate buffer system : It ranks first in terms of its power. She is presented carbonic acid(H 2 CO 3) and sodium or potassium bicarbonate (NaHCO 3, KHCO 3) in a proportion of 1/20. Bicarbonate buffer is widely used to correct disorders of the acid-base state of the body.

Phosphate buffer system . Dihydrogen phosphate has the properties of a weak acid and interacts with alkaline products entering the blood. Hydrogen phosphate has the properties of a weak alkali and reacts with stronger acids.

The protein buffer system plays the role of neutralizing acids and alkalis due to amphoteric properties: in an acidic environment, plasma proteins behave like bases, in a basic environment - like acids:

Buffer systems are also present in tissues, which help maintain tissue pH at a relatively constant level. The main tissue buffers are proteins and phosphates. pH is also maintained by the lungs and kidneys. Excess carbon dioxide is removed through the lungs. Kidneys with acidosis excrete more monobasic sodium phosphate, and with alkalosis - more alkaline salts: dibasic sodium phosphate and sodium bicarbonate.

Examples of problem solving

Solution:

We calculate the pH of an acidic buffer solution using the formula, then

Answer: 5,76

Solution:

We calculate the buffer capacity using the formula:

Answer: 0.021 mol/l

Example 3.

The buffer solution consists of 100 ml of 0.1 mol/l acetic acid and 200 ml of 0.2 mol/l sodium acetate. How will the pH of this solution change if 30 ml of 0.2 mol/l sodium hydroxide solution is added to it?

Solution:

We calculate the pH of the buffer solution using the formula:

When NaOH is added to a buffer solution, the amount of salt increases and the amount of acid in the buffer solution decreases:

0,006 0,006 0,006

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

We calculate n (NaOH) = 0.03 l · 0.2 mol/l = 0.006 mol, therefore in the buffer solution the amount of acid decreases by 0.006 mol, and the amount of salt increases by 0.006 mol.

We calculate the pH of the solution using the formula:

Hence: pH 2 – pH 1 = 5.82 – 5.3 = 0.52

Answer: change in pH of the buffer solution = 0.52.

Tasks for independent decision

4. To titrate 2 ml of blood to change the pH from the initial value (7.36) to the final value (7.0), it was necessary to add 1.6 ml of 0.01 M HCl solution. Calculate the acid buffer capacity.

5. How many moles of sodium acetate must be added to 300 ml of acetic acid to reduce the concentration of hydrogen ions by 300 times (K dis (CH 3 coon) = 1.85.10 -5).

6. When biochemical research use phosphate buffer with pH = 7.4. In what ratio should solutions of sodium hydrogen phosphate and sodium dihydrogen phosphate with a concentration of 0.1 mol/l each be mixed to obtain such a buffer solution (pK(H 2 PO 4 -) = 7.4).

7. What violations of the CBS are observed with the following indicators: blood pH = 7.20, Pco 2 = 38 mm Hg. Art., BO = 30 mmol/l, SBO = -4 mmol/l. How to eliminate this violation of the CBS?

Test tasks

Chapter 6. PROTOLYTIC BUFFER SYSTEMS

A change in any factor that can influence the state of chemical equilibrium of a system of substances causes a reaction in it that seeks 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 that maintain protolytic homeostasis include not only physiological mechanisms (pulmonary and renal compensation), but also physicochemical buffering effects, ion exchange, diffusion. Maintaining acid-base balance at a given level is ensured at the molecular level by the action of buffer systems.

Protolytic buffer systems are solutions that maintain a constant pH value both when adding acids and alkalis, and during dilution.

The ability of some solutions to maintain a constant concentration of hydrogen ions is called buffer action, which is the main mechanism of protolytic homeostasis. Buffers are mixtures of a weak base or weak acid and their salt. In buffer solutions, the main “active” components are a proton donor and acceptor, according to Brønsted’s theory, or an electron pair donor and acceptor, according to Lewis’s theory, representing an acid-base pair.

Based on whether the weak electrolyte of the buffer system belongs 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 buffer systems is called a buffer solution. Buffer solutions can be prepared in two ways:

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; NaH2PO4

and Na 2 HPO 4 .

The reason for the emergence of a new quality in solutions - buffering action - is 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, adding a weak acid HB + H 2 O ↔ H 3 O + + B - a strong acid or salt containing B - anions to a solution, an ionization process occurs, shifting the equilibrium to the left (effect common ion) B - + H 2 O ↔ HB + OH - , and the addition of alkali (OH -) - to the right, since due to the neutralization reaction the concentration of hydronium ions will decrease.

When combining two isolated equilibria (acid ionization and anion hydrolysis), it turns out that the processes that will occur in them under the influence of the same external factors(adding hydronium and hydroxide ions), 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 buffer system equation expresses the dependence of the pH of the buffer solution on the composition of the buffer system:

Analysis of the equation shows that the pH value of the buffer solution depends on the nature of the substances forming the buffer system, the ratio of the concentrations of the components and temperature (since the pKa value depends on it).

According to the protolytic theory, acids, bases and ampholytes are protolytes.

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 tend to have an acidic environment (pH<7).

Hydrocarbonate 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).

Hydrogen phosphate buffer system (buffer zone pH 6.2-8.2) - a mixture of 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 conjugate weak bases (proton acceptors) - hemoglobinate - Hb - and oxyhemoglobinate anions HbO 2 -, respectively.

Basic type buffer systems

Basic buffer systems are a mixture of a weak base (proton acceptor) and its salt (proton donor). They usually have an alkaline environment (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.

Ampholyte type buffer systems

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 - proton donor. A biologically significant buffer system of the ampholyte 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 having ions of the same name (common ion effect). For example, in an acetate buffer solution there are acetate ions, and in a hydrocarbonate 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, the action of which is based on the acid-base equilibrium 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 conjugate base CH 3 COO - binds the added hydrogen cations, turning into a weak acid: CH 3 COO - + + H + ↔ CH 3 COOH (the 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 the process of hydrolysis. According to Ostwald's law of dilution, an increase in the concentration of an acid slightly reduces its degree of electrolytic dissociation and the acid practically does not ionize. Consequently, in the system: C to increases, C to and α decreases, - const, C to /C to increases, where C to is the acid concentration, 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 into water molecules: CH 3 COOH + OH - → CH 3 COO - + H 2 O

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

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

The mechanism of action of a protein buffer solution when adding acid: (NH 3 +) m -Prot-(COO -) n + nH+ ↔ (NH 3 +) m -Prot-(COOH) n, when adding 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 to / C increases or decreases and the pH may change. This is confirmed by Henderson-Hasselbalch equation, which establishes the dependence of [H + ], K I, α and C to /C s. The equation

We derive this using the example of an acid-type buffer system - a mixture of acetic acid and its salt CH 3 COONa. 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 directly dependent on KI, α, acid concentration Ck and inversely dependent on Cc and the ratio C to /Cc. By 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 the equation for buffer systems of the acid type.

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

N.H. 4 + :

Henderson-Hasselbach equation for basic type buffer systems:

This equation can be represented as:

For a 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 in the second step.

6.4. CAPACITY OF BUFFER SOLUTIONS AND FACTORS DETERMINING ITS

The ability of solutions to maintain a constant pH value is not unlimited. Buffer mixtures can be distinguished by the strength of their resistance to the action of acids and bases introduced into the buffer solution.

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

Thus, the buffer capacity is a quantitative measure of the buffering effect of a solution. A buffer solution has a maximum buffer capacity at pH = pK of the acid or base forming a mixture with a ratio of its components equal to unity. The higher the initial concentration of the buffer mixture, the higher its buffer capacity. The buffer capacity depends on the composition of the buffer solution, concentration and ratio of components.

You need to be able to choose the right buffer system. The choice is determined by the required pH range. The buffer action zone is determined by the strength 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 are added

buffer system, can form insoluble compounds and interact with the components of the buffer system.

6.5. BLOOD BUFFERING SYSTEMS

Blood contains 4 main buffer systems.

1. Hydrocarbonate.

It accounts for 50% of the capacity. It operates primarily in plasma and plays a central role in CO 2 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 the transport of CO 2 and the establishment of pH. Blood plasma pH is 7.4. CO 2 is a product of cellular metabolism released into the blood. Diffuses through the membrane into red blood cells, 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 increases, and the following reactions take place: The resulting HCO 3 - exits through the membrane and is carried away by the blood stream. In blood plasma, the pH is 7.4. When venous blood returns to the lungs, hemoglobin reacts with oxygen to form oxyhemoglobin, which is a stronger acid: HHb + + O 2 ↔ HHbO 2. The pH decreases as more is formed, the reaction occurs: HHbO 2 + HCO 3 - ↔ HbO 2 - + H 2 CO 3. CO 2 is then released into the atmosphere. This is one of the mechanisms for the transport of CO 2 and O 2.

Hydration and dehydration of CO 2 is catalyzed by the enzyme carbonic anhydrase, which is found in red blood cells.

Bases are also bound by the blood buffer and excreted in the urine, mainly in the form of mono- and dibasic phosphates.

In clinics, reserve blood alkalinity is always determined.

6.6. QUESTIONS AND EXERCISES TO SELF-TEST YOUR PREPARATION FOR CLASSES AND EXAMINATIONS

1. When combining which protolytic equilibria will the solutions have buffering properties?

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

3. Main types of buffer solutions. The mechanism of their buffering action and the Henderson-Hasselbach equation that 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 the buffer capacity of a buffer system called? Which blood buffer system has the greatest capacity?

6. Methods for obtaining buffer solutions.

7. Selection of buffer solutions for medical and biological 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 TASKS

1. Which of the proposed systems is a buffer system?

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 calculation 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 NH 4 OH and 0.3 M solution NH 4 Cl;

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

e)0.4 M NaHCO solution 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 KNSO 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 KN 2 PO 4.

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

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

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

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

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

e) when the acid concentration is 2 times greater 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 system is a protein buffer?

a) a weak acid and its anion;

c) anions of 2 acid salts;

e) ions and molecules of ampholytes.

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

a) a weak acid and its anion;

b) anions of acidic and medium salts;

c) anions of 2 acid salts;

d) weak base and its cation;

e) ions and molecules of ampholytes.

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

a) a weak acid and its anion;

b) anions of acidic and medium salts;

c) anions of 2 acid salts;

d) weak base and its cation;

e) ions and molecules of ampholytes.

10. When is a protein buffer 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; edited by V. A. Popkova, A. V. Zholnina. - 2012. - 400 pp.: ill.