INTRODUCTION

BUFFER SOLUTIONS (buffer mixtures, buffers) - solutions containing buffer systems and, as a result, have the ability to maintain pH at a constant level. They are usually prepared by dissolving a weak acid and its alkali metal salt taken in appropriate proportions in water, partially neutralizing a weak acid with a strong alkali or a weak base with a strong acid, and dissolving a mixture of salts of a polybasic acid. The pH value of buffer solutions prepared in this way varies slightly with temperature. The range of pH values in which a 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 well-known buffer solutions are: glycine Serensen, acetate Walpole, phosphate Serensen, borate Palich, veronal Michaelis, carbonate Colthoff, 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 - 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, the concentration of hydrogen and hydroxyl ions in these solutions will not change.

The ability of some solutions to maintain a constant concentration of hydrogen ions when diluted, as well as when small amounts of strong acids or alkalis are added, is known as buffering action.

Solutions that simultaneously contain a weak acid and its salt or a weak base and its salt and have a buffering effect are called buffer solutions. Buffering solutions can be considered as mixtures of electrolytes having the same ions. The presence of a weak acid or weak base and their salts in a solution reduces the effect of dilution or the action of other acids and bases on the pH of the solution.

Such buffer solutions are the following CH mixtures 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 + 0.1 M CH solution 3 CO ONa has pH = 4.7.

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

Acid-base buffer solutions

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 by adding small amounts of acid or base.

Redox keep the potential of the system constant when oxidizing or reducing agents are introduced.

metal buffer solutions are known that maintain a constant pH value.

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

NA ↔ N + + A -

conjugate acid

Base

B + N + ↔ VN +

ABOUT warp conjugate

Acid

Since this section discusses only acid-base buffer solutions, we will call them buffer solutions, omitting “acid-base” in the name.

Buffer solutions are solutions that maintain a constant pH value by dilution and the addition of small amounts of acid or base.

Classification of buffer systems

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

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 varying degrees of substitution. For example, phosphate buffer solution.

4. ions and molecules of ampholytes. These include, for example, amino acids and protein buffer systems. Being in an isoelectric state, amino acids and proteins are not buffers. The buffering effect occurs 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 classified as buffer systems of the first or second type.

Calculation of pH of buffer solutions

The basis for calculating the pH of buffer systems is 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 ion concentration 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 equal to

(2)

In the presence of CH3 COONa The acid-base equilibrium of acetic acid is shifted to the left. Therefore, the concentration of undissociated acetic acid is almost equal to the concentration of the acid, i.e. [SN 3 COOH] = acidic

Main source of acetate ions strong electrolyte CH3COONa:

CH 3 COONa → Na + + CH 3 COO - ,

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

From this we obtain the Henderson-Hasselbach equation for buffer systems consisting of a weak acid and its salt:

(3)

For a buffer system consisting of a weak base and its salt, for example, ammonia, the concentration of hydrogen ions in the solution can be calculated based on 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 it into (4).

(6)

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

(7)

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

Main source of ammonium cations strong electrolyte NH4Cl:

NH 4 Cl → NH 4 + + Cl - ,

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

(8)

From this we obtain 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 varying degrees of substitution, for example, phosphate, consisting of a mixture of solutions of hydrogen phosphate ( Na2HPO4 ) and dihydrogen phosphate ( NaH2PO4 ) sodium. Its action is based on acid-base balance:

H 2 PO 4 - ↔ H + + HPO 4 2-

Weak acid conjugate base

(10)

Expressing from (10) the concentration of hydrogen ions 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 reversing the signs, 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.

Properties of buffer solutions

The pH value of buffer solutions remains unchanged upon dilution, as follows from the Henderson-Hasselbalch equation. When the buffer solution is diluted with water, the concentrations of both components of the mixture decrease by the same number of times. Therefore, the pH value should not change. However, experience shows that some change in pH, although minor, does occur. This is explained by the fact that the Henderson-Hasselbalch 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 conjugate acid and base.

Buffers change pH little 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 is added to them is based on the fact that one component of the buffer solution can react with H+ added acid, and the other with OH- added base. As a result, the buffer system can bind both 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 conjugate 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 conjugate base HCOO- binds added ions H+ , turning into weak formic acid:

HCOO - + H + ↔ HCOOH

According to Le Chatelier's principle, the equilibrium shifts to the left.

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

HCOOH + OH - → HCOO - + H 2 O

According to Le Chatelier, the acid-base equilibrium 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 changes slightly.

The essence of buffer action

The action of buffer solutions is based on the fact that the individual components of buffer mixtures bind hydrogen or hydroxyl ions of the acids and bases introduced into them to form weak electrolytes. For example, if a buffer solution containing the weak acid HA n and a 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

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

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

N + + A n - → NA n

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

pH value in buffer solutions

By changing the ratios you can get buffer

solutions characterized by 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 [ H + ] = − − log K HAn − − log C HAn

But since K HAn is a constant value, it is best to represent it in the form pK HAn those. indicator of the electrolytic dissociation constant: pK Han = − log K HAn .

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

рН = − log [Н + ] = − − pK HAn − − pC HAn

As a weak acid salt is added to an aqueous solution, 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+ ] = KHAn

That

рН = − log [Н + ] = − log K HAn − log C HAn + log C Kt А n .

We similarly derive the formula for weak bases:

[OH] = √K KtOH * C KtOH

pOH = − log [ OH ] = − − log K KtOH − − log C KtOH

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

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 [ H + ] = − log K w + log K KtOH − logC Kt A n + log C KtOH.

There is no need to memorize the derived formula for pH values, since they are very easily derived by taking the logarithm of simple formulas expressing the value of [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 a buffer solution, a noticeable change in pH is observed. Moreover, for different buffer mixtures that differ from each other in composition, differing from each other in composition, the buffer effect is not the same. Consequently, buffer mixtures can be distinguished by the strength of their resistance to the action of acids and alkalis introduced into the buffer solution in equal quantities and 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 a buffer solution so that its pH value changes by only one unit is called a buffer capacity.

If the value [H + ] of one buffer solution changes with the addition of a strong acid less than the value of [H+ ] of another buffer solution when adding the same amount of acid, then the first mixture has a greater buffer capacity. For the same buffer solution, the higher the concentration of its components, the greater the buffer capacity.

Buffer properties of solutions of strong acids and bases.

Solutions of strong acids and bases at sufficiently high concentrations also have a buffering effect. The conjugate systems in this case 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 in aqueous solutions and are therefore characterized by a high concentration of hydronium ionsor hydroxyl ions. Adding small amounts of a strong acid or a strong base to their solutions therefore has only a minor effect on the pH of the solution.

Preparation of buffer solutions

1. By diluting the appropriate fixatives in a volumetric flask.

2. By mixing the quantities of suitable conjugate acid-base pairs calculated using the Henderson-Hasselbach equation.

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

Since 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 concentrations of both components or by adding to a solution of one component the appropriate amount of reagent, leading to the formation of an equal concentration of the conjugate form. The reference literature contains detailed recipes for preparing buffer solutions for various pH values.

Application of buffer solutions in chemical analysis

Buffer solutions are widely used in chemical analysis in cases where, according to the experimental conditions, a chemical reaction must occur while maintaining an exact pH value that does not change when the solution is diluted or when other reagents are added to it. For example, during an 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 the precipitation of sediments that cannot be precipitated in acidic or alkaline solutions. The harmful effects of acids are suppressed by sodium acetate, which reacts with a strong acid. For example:

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

or in ionic form

H + + CH 3 COO → CH 3 COOH.

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

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

Phosphate buffer solution ( N a 2 NPO 4 + N aH 2 RO; pH = 8) is used in many oxidation-reduction reactions.

To successfully use buffer mixtures for analytical purposes, it is necessary to remember that not every buffer mixture is suitable for analysis. The buffer mixture is selected 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 concentrations of their components.

The above can be presented in table form.

Buffer solutions used in analysis

Buffer mixture	Mixture composition (at a molar ratio of 1:1)	pH
Formate	Formic acid and sodium formate
Benzoate	Benzoic acid and ammonium benzoate
Acetate	Acetic acid and sodium acetate
Phosphate	Mono- and di-sodium phosphate
Ammonium	Ammonium hydroxide and ammonium chloride

Mixtures of acid 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 plays 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.

Animal and plant organisms also have complex buffer systems that maintain a constant pH of blood, lymph and other fluids. Soil also has buffering properties, 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 solutions that supporta constant pH value when diluted and small amounts of acid or base are added. An important property of 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 they are 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

Buffering solutions can be considered as mixtures of electrolytes having the same ions. Buffer solutions play an important role in many technological processes. They are used, for example, in the electrochemical application of protective coatings, in the production of dyes, leather, and 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 blood in the human body is maintained between 7.35 and 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.

BIBLIOGRAPHICAL LIST

1. Kreshkov A.P. Fundamentals of analytical chemistry. Book 1. - M: Chemistry, 1965. -498 pp.

2. Tsitovich I.K. Analytical chemistry course: Textbook for universities. - St. Petersburg: “Lan”, 2007 - 496 p.

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 applicants to universities / Ed. Lydia R.A., Alikberova L.Yu. - M.:AST-PRESS SCHOOL, 2007. -512s.

5. Osipov Yu.S., Great Russian Encyclopedia: in 30 volumes. T.4.- M.: Great Russian Encyclopedia 2006. - 751 p.

6. Mikhailenko Ya.I., Introduction to chemical analysis, Goskhimtekhizdat, 1933.

One of the main properties of living organisms is maintaining acid-base homeostasis at a certain level. Protolytic homeostasis– constancy of pH of biological fluids, tissues and organs. This is expressed in fairly constant pH values of biological media (blood, saliva, gastric juice, etc.) and the body’s ability to restore normal pH values when exposed to protoliths. System supporting protolytic homeostasis, includes not only physiological mechanisms (pulmonary and renal compensation), but also physicochemical ones: buffering action, ion exchange and diffusion.

Buffer solutions are called solutions that maintain the same pH value when diluted or added with a small amount of a strong acid or base. Protolytic buffer solutions are mixtures of electrolytes containing ions of the same name.

There are mainly two types of protolytic buffer solutions:

Acidic i.e. consisting of a weak acid and an excess of its conjugate base (a salt formed by a strong base and an anion of this acid). For example: CH 3 COOH and CH 3 COONa - acetate buffer

CH 3 COOH + H 2 O ↔ H 3 O + + CH 3 COO - excess conjugated

grounds

CH 3 COONa → Na + + CH 3 COO -

Basic ones, i.e. consisting of a weak base and an excess of its conjugate acid (i.e., a salt formed by a strong acid and a cation of this base). For example: NH 4 OH and NH 4 Cl – ammonia buffer.

NH 3 + H 2 O ↔ OH - + NH 4 + excess

Base

conjugate

NH 4 Cl → Cl - + NH 4 + acids

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

pH = pK + ℓg, pOH = pK + ℓg
,

where pK = -ℓg K D.

C – molar or equivalent electrolyte concentration (C = V N)

Mechanism of action of buffer solutions

Let's consider it 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 the strong electrolyte - sodium acetate, and acetic acid, in the presence of the anion of the same name, exists in solution in an almost non-ionized form.

When a small amount of hydrochloric acid is added, H + ions bind to the conjugate base CH 3 COO - present in the solution into the weak electrolyte CH 3 COOH.

CH 3 COO ‾ +H + ↔ CH 3 COOH (1)

From equation (1) it is clear that the strong acid HC1 is replaced by an equivalent amount of the weak acid CH 3 COOH. The amount of CH 3 COOH increases and, according to W. Ostwald’s dilution law, the degree of dissociation decreases. As a result, the concentration of H + ions in the buffer increases, but very slightly. The pH remains constant.

When adding an acid to a buffer, the pH is determined by the formula:

pH = pK + ℓg

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 COOH + OH ‾ ↔ CH 3 COO ‾ + H 2 O (2)

As a result, the alkali is replaced by an equivalent amount of the weakly basic salt CH 3 COONa. The amount of CH 3 COOH decreases and, according to W. Ostwald’s dilution law, the degree of dissociation increases due to the potential acidity of the remaining undissociated CH 3 COOH molecules. Consequently, the concentration of H + ions remains virtually unchanged. The pH remains constant.

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

pH = pK + ℓg

When diluting the buffer, the pH also does not change, because the dissociation constant and the ratio of 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. The pH of the buffer will be greatest when the component ratio is equal to one.

To quantitatively characterize the buffer, the concept is introduced buffer capacity.

Buffers are solutions that keep their pH constant when diluted or added with a small amount of a strong acid or base.

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

There are two types of protolytic buffer solutions:

Acidic, consisting of a weak acid and an excess of its conjugate base (a salt formed by a strong base and an anion of this acid);
Basic, consisting of a weak base and an excess of its conjugate 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 electrolyte concentration (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.

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

CH 3 COO‾ + H + ↔ CH 3 COOH

From the equation it can be seen that the strong acid HC1 is replaced by an equivalent amount of the weak acid CH 3 COOH. The amount of CH 3 COOH increases and, according to W. Ostwald’s dilution law, 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 adding an acid to a 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 COOH + OH ‾ ↔ CH 3 COO‾ + H 2 O

As a result, the alkali is replaced by an equivalent amount of the weakly basic salt CH 3 COONa. The amount of CH 3 COOH decreases and, according to W. Ostwald’s dilution law, 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 adding alkali, the pH is determined by the formula:

3. When diluting the buffer, the pH also does not change, because the dissociation constant and the ratio of 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 is worth noting that the pH of the buffer will be greatest when the component ratio is equal to one.

Buffer capacity is the ability of the buffer system to counteract changes in the pH of the environment.

Buffer capacity (B) is expressed as the number of mole equivalents of 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 mole equivalent of a strong acid or alkali, pH H is the initial pH value (before adding acid or alkali), pH K is the final pH value (after adding acid or alkali), ΔpH is the 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 buffer. - volume of buffer solution, Δ pH - change in pH.

The buffer capacity depends on the concentration of electrolytes and the ratio of buffer components. Solutions with a higher concentration of components and a component ratio equal to unity have the greatest buffer capacity.

The following buffer systems operate in the human body:

Bicarbonate buffer, which is the main buffer system of blood plasma; it is a quick response system, since the product of its interaction with CO 2 acids is quickly eliminated through the lungs. In addition to plasma, this buffer system is found in red blood cells, interstitial fluid, and renal tissue.
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 CO 2. 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 red blood cells, 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 the blood plasma and intercellular fluid Na 2 HPO 4 and NaH 2 PO 4. It functions mainly in plasma and includes: dihydrogen phosphate ion H 2 PO 4 - and hydrogen phosphate ion HPO 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 exhibits its own buffering effect. Interaction of buffer systems in the body by stages:

1. During the process of gas exchange in the lungs, oxygen enters the red blood cells;

2. As blood moves to the peripheral parts of the circulatory system, oxygen is released in the ionized form of HbO 2 -. In this case, the blood changes from arterial to venous. 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 red blood cells.

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

СО 2 + Н 2 О ↔ Н 2 СО 3 ↔ Н + + НСО 3 -

4. The resulting excess protons bind to 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 that differ in acid-base properties (the chloride ion is protolytically inactive; the bicarbonate ion is a base under body conditions), a hydrocarbonate-chloride shift occurs. This explains the more acidic reaction of the environment in erythrocytes (pH = 7.25) compared to plasma (pH = 7.4).

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

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

7. Excess protons are neutralized by phosphate buffer:

N + + NPO 4 - ↔ N 2 PO 4 -

8. After the blood returns to the lungs, the concentration of oxyhemoglobin increases (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. Excess protons also accumulate in the kidneys as a result of the reaction:

СО 2 + Н 2 О ↔ Н 2 СО 3 ↔ Н + + НСО 3 - ,

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

H + + NH 3 ↔ NH 4 +

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

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

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 support protolytic homeostasis include not only physiological mechanisms (pulmonary and renal compensation), but also physicochemical buffering effects, ion exchange, and 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 a salt containing anions B - to a solution, an ionization process occurs, shifting the equilibrium to the left (common ion effect) 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) are differently directed. 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 a stronger acid 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.

Classification of buffer solutions

There are natural and artificial buffer solutions. A natural buffer solution is blood, which contains bicarbonate, phosphate, protein, hemoglobin and acid buffer systems. An artificial buffer solution can be an acetate buffer consisting of CH3COOH.

Buffer solutions may have an acidic reaction (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, range 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 acid and medium salt:

For example: carbonate buffer system

Na 2 CO 3 and NaHCO 3, range of action pH = 9.3 - 11.

4) A mixture of two acid salts:

For example: phosphate buffer system

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

Mechanism of action of buffer solutions

Let us understand 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 influence of the ion of the same name). buffer solution hydroxide tetraborate

Let's assume 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 with dilution the degree of dissociation of acetic acid increases, since the concentration of sodium acetate, which suppresses the dissociation of acetic acid in this solution, decreases. Therefore, when diluted with water, the pH will remain virtually unchanged.

2) Adding strong acid

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

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

H + ions entering the solution will bind into 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 solution, the concentration of hydrogen ions will increase many times and the pH of the solution will change noticeably. And hydrogen, as you know, is the most common chemical element.

3) Adding a small amount of alkali

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 into 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 will remain virtually unchanged after adding alkali.

If you add alkali to clean water, all OH - ions will remain in solution. The concentration of OH - ions will increase sharply, the concentration of H + ions will correspondingly 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 buffer mixtures.