Henderson-Hasselbach equation - a mathematical expression that characterizes the capabilities of the buffer system. The equation shows how the acid-base balance of a buffer solution depends on the properties of the components of the acid-base buffer system and on the quantitative ratio of these components in the solution. An indicator of the acid-base balance in a solution is the hydrogen index, pH. The property of an acid (its ability to decompose into ions), as a component of a buffer system, is characterized by the value of the equilibrium constant, the dissociation constant of the acid, Ka. pK= – lgK D

The quantitative structure (composition) of the buffer system can be estimated as a salt/acid ratio. Given the above, the Henderson-Hasselbach equation looks like this:

pH = pK + log

The value of pH and pOH is affected by dissociation constant and the ratio of the concentrations of the components.

18. Buffer tank. Buffer zone.

Interval pH=pKa±1 called buffer zone .

Buffer capacity(V) expressed as the number of mole equivalents of a strong acid or base that must be added to one liter of buffer to shift the pH by one.

B - buffer capacity,

nE is the mole equivalent of a strong acid or alkali,

ΔpH is the change in pH.

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

V is the volume of acid or alkali,

N is the equivalent concentration of acid or alkali,

V buffer - the volume of the buffer solution,

Δ pH is the change in pH.

The buffer capacity depends on electrolyte concentrations and buffer ratios.

19. Quantitative determination of buffer capacity.

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

The higher initial concentration buffer mixture, the higher its buffer capacity

20. Buffer systems of blood: bicarbonate, phosphate, hemoglobin and protein

hemoglobin buffer It makes up 35% of the buffer capacity.

The main buffer system of erythrocytes, which accounts for about 75% of the total buffer capacity of the blood. The hemoglobin buffer system of the blood plays a significant role in: respiration, transport of oxygen to tissues and in maintaining a constant blood pH.

It is represented by two weak acids - hemoglobin and oxyhemoglobin and their conjugate bases - hemoglobinate and oxyhemoglobinate ions, respectively:

HHb ↔ H + + Hb -

HHbO 2 ↔ H + HbO 2 -

Phosphate buffer

It is found both in the blood and in the cell fluid of other tissues, especially the kidneys. In cells, it is represented by salts

K 2 NRO 4 and KN 2 RO 4, and in blood plasma and intercellular fluid

Na2HPO4 and NaH2PO4.

Functions primarily in plasma and includes: dihydrophosphate ion and hydrogen phosphate ion

H 2 RO 4 - and NRO 4 2-

This system plays a crucial role in biological environments − in the cell, in the juices of the digestive glands, in the urine.

bicarbonate buffer . It makes up 53% of the buffer capacity.

Presented:

H 2 CO 3 and NaHCO3

The bicarbonate buffer is the main buffer system in blood plasma; it is a rapid response system, since the product of its interaction with CO 2 acids is quickly excreted through the lungs.

Protein buffer It is 5% of the buffer capacity.

It consists of a protein-acid and its salt formed by a strong base.

Pt - COOH - protein-acid

Pt - COONa - protein-salt

1. When strong acids are formed in the body, they interact with the salt of the protein.

HC1 + Pt-COONa ↔ Pt-COOH + NaCl.

2. With an increase in alkaline products, they interact with Pt-COOH:

NaOH + Pt-COOH ↔ Pt-COONa + H 2 O

Protein is an amphoteric electrolyte and therefore exhibits its own buffering action.

Buffer solution is a chemical reagent with a constantpH

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In the practice of laboratory work, employees often encounter chemical solutions that have or should have a certain 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; a mixture of weakly concentrated acid and its salt. These solutions practically do not change their structure when concentrated, diluted with other chemical reagents, or when highly concentrated alkalis or acids are added to it in a small amount. To obtain a buffer solution with a different pH, it is necessary to change the concentration and ratio of the chemical solutions used.

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

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

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

For work both in laboratory conditions and in private practice, it is recommended to use high-stability buffer mixtures 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 is the buffer solution?

The composition of this chemical reagent includes water - a solvent and equally dissolved ions or molecules of substances that make up an acid-base or alkali-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, are widely used in research in analytical chemistry, biology and microbiology, genetics, medicine, pharmaceuticals, research centers and other scientific fields.

Importance of buffer solution for humans

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

Storage conditions

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

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

ON + 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] \u003d C (HA) and [A - ] \u003d C (A -). Using the expression for K a ON 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 a buffer solution consisting of 0.10 M acetic acid and 0.10 M sodium acetate.

Solution. Here, all conditions are met that allow applying the formula (2-14) (acetic acid is a weak acid, the concentrations of acid and conjugate base are quite high). That's why

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

Solution. According to 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 the acid HA and the conjugate base A - . Therefore, the buffer capacity is usually represented as

if a strong base is added to the buffer solution, and

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

Let us express [ON] in terms of K a ON and substitute into the material balance equation. Let's 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 a 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 one from the other, if we remember that [ON] = a(HA)C(HA) and [A - ] = a(A -) C (A -), as well as expressions for a(ON) and a(BUT -).

For highly dilute 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 describes the buffering action of an acid and a conjugate 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 but the 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 the calculation of the buffer capacity, but by finding the amounts 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 c.

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

Hence [H + ] = K a HA and, consequently, 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 the 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 +] \u003d [A 1 -] +

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

Let us 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 monobasic bases

(2-21)

where K a 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, one can take into account the dissociation of only this acid (this base), and neglect the dissociation of the rest. But there may be other situations.

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

Solution. Although the dissociation constants of benzoic (K a= 1.62 10 -6 , denote K 1) and aminobenzoic (K a = 1.10 10 -5 , denote K2) acids differ by almost two orders of magnitude; due to the rather large difference in acid concentrations, it is necessary to take into account the dissociation of both acids. Therefore, according to the formula (2-20) we find

A buffer solution or simply a buffer is a solution whose pH does not change significantly when small amounts of acid or base are added.

Buffer solutions can be classified into four types.

Buffer solutions containing a strong acid

Any strong acid such as nitric acid can be used as a low pH buffer. Strong acids are completely dissociated in aqueous solutions, and therefore their solutions are characterized by a high concentration of hydronium ions. The addition of a small amount of acid or base to a strong acid therefore has only a small effect on the pH of a solution of a strong acid.

For example, if 1 cm3 of hydrochloric acid with a concentration of 0.1 mol/dm3 is added to 100 cm3 of a nitric acid solution with a concentration of 0.01 mol/dm3, then it will decrease from 2.00 to 1.96. A change in pH of 0.04 can be considered negligible. To check the above pH values ​​of a solution before and after adding hydrochloric acid, use the equation

Let us now compare the indicated negligible decrease in pH with the result of adding a solution with a concentration of 0.1 mol/dm3 to 100 cm3 of pure water. In this case, the pH drops sharply from 7.00 to 4.00. Clearly, pure water does not act as a buffer solution because it does not keep the pH at approximately the same level. The concentrations of the buffer solutions correspond to the flat parts of the titration curves shown in Figs. 8.2. These parts of the titration curves are called buffer regions. In the buffer region, pH values ​​are insensitive to small changes in acid or base concentration.

Buffer solutions containing a strong base

Any strong base can be used as a buffer with a high value. Adding a small amount of acid or base to such a buffer has a negligible effect on. For example, when adding a solution of hydrochloric acid with a concentration of a solution with a concentration, a change from 12.00 to 11.96 occurs. The change in this case is only 0.04. This result can be verified using equation (6) and the relation

Buffer solutions containing a weak acid

Buffer solutions with stable values ​​ranging from 4 to 7 can be obtained using any weak acid and one of its salts. For this purpose, a mixture of acetic acid and sodium acetate is often used. Sodium acetate in aqueous solution is fully ionized

In contrast, acetic acid is only partially ionized.

When acid is added, this equilibrium shifts to the left, the content of added ions decreases and the original value is restored. The presence of sodium acetate in the buffer solution provides a large supply of ions that can compensate for the effect of added portions of acid.

When a base is added, it is neutralized by hydronium ions

The removal of ions as a result of this reaction causes equilibrium (7) to shift to the right. The concentration of ions and therefore the value of the solution remain constant. The presence of acetic acid in the buffer solution provides a large supply of non-dissociated molecules capable of dissociating and thus, if necessary, compensate for the addition of portions of the base.

The action of buffer solutions can be considered quantitatively on the basis of the law of mass action. As shown in the previous section, applying this law to the acetic acid dissociation equilibrium leads to the following expression for the acetic acid dissociation constant:

Taking the logarithm of this expression leads to the following result:

where are the total concentrations of the corresponding particles in the buffer solution. The dissociation constant of acetic acid is equal to the table. 8.1). This means that the acetic acid dissociation equilibrium described by

equation (7), is significantly shifted to the left. For this reason, the relative contribution of acetic acid to the total amount of ions in the buffer solution is small. The value in equation (8) is almost entirely due to the salt contribution, i.e. sodium acetate, which is completely dissociated into ions. Therefore,

Since acetic acid is little dissociated in the buffer solution, the acid concentration in the equilibrium mixture (7) approximately coincides with its initial concentration in the buffer solution. This allows you to write

Substituting the results obtained into equation (8), we obtain

The resulting relationship is called the Henderson equation for a buffer solution consisting of a weak acid and its salt. It can be used for various calculations, namely for calculating: buffer solution;

the amount of acid or salt required to obtain a buffer solution with the desired value

changes in the buffer solution when small portions of acid or base are added to it.

a) How much sodium acetate should be dissolved in acetic acid, having a concentration to obtain a buffer solution with

b) How will this buffer solution change if a solution having a concentration of

a) From equation (9) it is easy to find

By condition and

According to Table. 8.1.

Substituting all these values ​​into the resulting equation gives

Consequently,

This means that to obtain a buffer solution with one mole of sodium acetate should be dissolved in acetic acid.

Relative molar mass of sodium acetate:

Therefore, the mass of a mole of sodium acetate is

Thus, to obtain a buffer solution with 1.46 g of sodium acetate must be dissolved in acetic acid.

b) 1 cm3 of a solution having a concentration contains

0.001 mol It reacts with forming Therefore, the concentration will decrease by and the concentration will increase by 0.001 mol/dm3 (a slight increase in volume can be neglected). In this way,

So, when alkali is added to the buffer solution, a negligible change of 0.07 should occur.

When considering buffer solutions containing a weak acid, one special case arises. The Henderson equation shows that when the salt concentration is exactly equal to the acid concentration, the buffer solution is the same as that acid, i.e.

For example, if 100 cm3 of a 0.1 mol/dm3 solution is added to 100 cm3 of a 0.1 mol/dm3 solution, the resulting buffer should have a pH of 4.75 at 25°C.

Buffer solutions containing a weak base

Buffer solutions with stable values ​​ranging from 7 to 10 can be obtained by mixing a weak base with one of its salts. A typical buffer solution of this type is a solution of ammonia and ammonium chloride. In an aqueous solution, ammonium chloride completely dissociates

Ammonia only partially dissociates in water

When an acid is added to this buffer solution, it is neutralized by ions. As a result, equilibrium (10) shifts to the right. This shift maintains a constant ion concentration and hence a constant

When a base is added, equilibrium (10) shifts to the left, and the concentration of OH ions is maintained constant. The presence of ammonium chloride in the buffer solution provides a large supply of ions in it, which makes it possible to compensate for the effect of added portions of the base.

The Henderson equation for a buffer solution containing a weak base and one of its salts is

Buffer Applications

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, photographic materials and leather. In addition, buffer solutions are widely used in chemical analysis and for the calibration of pH meters (see Chapter 10).

Many biological and other systems depend on the buffer solutions they contain to maintain a constant pH. Normal pH values ​​for some of these systems are listed in Table. 8.6. For example, the pH of the blood in the human body is maintained in the range of 7.35 to 7.45, despite the fact that the content of carbon dioxide and, therefore, carbonic acid in the blood can vary widely. The buffer contained in the blood is a mixture of phosphate, bicarbonate and proteins. Protein buffers maintain the pH of tears at 7.4. In bacteriological studies, to maintain a constant pH of culture media used to grow bacteria, it is also necessary to use buffer solutions.

Table 8.6, pH values ​​for some biological systems and other solutions

INTRODUCTION

BUFFER SOLUTIONS (buffer mixtures, buffers) - solutions containing buffer systems and, as a result, having 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 the buffer solution has stable buffering properties lies within pK ± 1 (pK is the negative decimal logarithm of the dissociation constant of the weak acid included in its composition). The most well-known buffer solutions are: Serensen's glycine, Walpole's acetate, Serensen's phosphate, Palic's borate, Michaelis's veronal, Kolthoff's carbonate, Tris buffer, universal Michaelis's veronal, 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 (pH) - as standard solutions with stable pH values, etc.

BUFFER MIXTURES

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

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

Solutions that simultaneously contain some weak acid and its salt or some weak base and its salt and have a buffering effect are called buffer solutions. Buffer solutions can be considered as mixtures of electrolytes having ions of the same name. The presence of a weak acid or weak base and their salts in the 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 mixtures of CH 3 COOH + CH 3 C OON a, NH 4 OH + NH 4 Cl, Na 2 CO 3 + NaHCO 3, etc.

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

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

Acid-base buffer solutions

In a broad sense, buffer systems are called 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.

known metal buffer solutions that maintain a constant pH.

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

ON ↔ H + + A -

acid conjugate

Base

B + H + ↔ HH +

O conjugate warping

Acid

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

Buffer solutions are solutions that maintain a constant pH when diluted and small amounts of acid or base are added.

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 various 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 buffer. The buffering effect only appears when a certain amount of acid or alkali is added to them. In this case, a mixture of two forms of protein is formed: a) a weak "protein acid" + a salt of this weak acid; b) weak "protein base" + salt of this weak base. Thus, this type of buffer systems can be attributed to buffer systems of the first or second type.

pH Calculation of Buffer Solutions

The calculation of the pH of buffer systems is based on the law of mass action for acid-base balance. For a buffer system consisting of a weak acid and its salt, such as acetate, the ion concentration H+ easy to calculate from 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 CH 3 COONa acid-base balance of acetic acid is shifted to the left. Therefore, the concentration of undissociated acetic acid is practically equal to the concentration of acid, i.e. [CH 3 COOH] = with acid.

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

CH 3 COONa → Na + + CH 3 COO -,

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

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

(3)

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

NH 3 × H 2 O \u003d NH 4 OH ↔ NH 4 + + OH -

(4)

We express the concentration of ions oh- from the ionic product of water

(5)

and substitute into (4).

(6)

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

(7)

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

The main source of ammonium cations is a strong electrolyte NH4Cl:

NH 4 Cl → NH 4 + + Cl -,

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

(8)

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

(9)

Similarly, you can calculate the pH of a buffer system consisting of a mixture of solutions of salts of polybasic acids of various degrees of substitution, for example, phosphate, consisting of a mixture of solutions of hydrophosphate ( Na2HPO4 ) and dihydrophosphate ( 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 - ) is the 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 when diluted, as follows from the Henderson-Hasselbach equation. When the buffer solution is diluted with water, the concentrations of both components of the mixture decrease by the same number of times. Therefore, the pH value should not change. However, experience shows that some change in pH, although insignificant, does occur. This is explained by the fact that the Henderson-Hasselbach equation is approximate and does not take into account interionic interactions. Accurate calculations should take into account the change in the activity coefficients of conjugated acids and bases.

Buffer solutions change the 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 are added to them is based on the fact that one constituent of the buffer solution can interact with H+ added acid, and the other with OH- added base. As a result, the buffer system can bind both H + and OH - and up to a certain limit to maintain the constancy of the pH value. Let us demonstrate this using the example of a formate buffer system, which is a conjugated acid-base pair HCOOH/HCOO- . Equilibrium in a formate buffer solution can be represented by the equation:

HCOOH ↔ HCOO-+H+

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

HCOO - + H + ↔ HCOOH

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

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

HCOOH + OH - → HCOO - + H 2 O

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

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

The essence of the buffer action

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

H + + OH → H 2 O

Therefore, 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 spent hydrogen ions, due to the subsequent dissociation of acid HA n , new hydrogen ions appear. As a result, the former 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, then the following reaction will occur:

H + + A n - → ON n

those. And 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 added 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 and you can get buffer

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

[ H + ] = √K HAn * C HAn

where

pH = − lg [ Н + ] = − − 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 get that in an aqueous solution of a weak acid:

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

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

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

then

pH \u003d - lg [H + ] \u003d - lg K HAn - lg C HAn + lg C Kt A n.

Similarly, we derive the formula for weak bases:

[OH] = √KKtOH * CKtOH

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

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

pH = pK w − (− pK KtOH − − lg C KtOH )

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

[H+]=

t . e .

pH \u003d - log [H + ] \u003d - log K w + log K KtOH - logC Kt A n + log C KtOH.

There is no need to memorize the pH values ​​derived from the formula, 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 the buffer solution, a noticeable change in pH is observed. moreover, for different buffer mixtures, differing from each other in composition, differing from each other in composition, the buffer effect is not the same. Therefore, buffer mixtures can be distinguished by the strength of their resistance to the action of acids and alkalis introduced into the buffer solution in the same quantities and at a certain concentration. The limiting amount of acid or alkali of a certain concentration (in mol / l or g-eq / l), which can be added to a buffer solution so that its pH value changes by only one unit, is called the buffer capacity.

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

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- /N 2 O - for strong bases. Strong acids and bases are completely dissociated in aqueous solutions and therefore are characterized by a high concentration of hydronium ions.or hydroxyl ions. The addition of small amounts of a strong acid or strong base to their solutions therefore has only a negligible effect on the pH of the solution.

Preparation of buffer solutions

1. Dilution in a volumetric flask of the corresponding fixanals.

2. Mixing the amounts of suitable conjugated acid-base pairs calculated according to 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 differs by 10 times or more from the concentration of the other, buffer solutions are often prepared by mixing solutions of equal concentrations of both components or by adding an appropriate amount of a reagent to a solution of one component, leading to the formation of an equal concentration of the conjugated 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 those cases where, according to the conditions of the experiment, the chemical reaction must proceed while maintaining the exact pH value that does not change when the solution is diluted or when other reagents are added to it. For example, when carrying out an oxidation-reduction reaction, during the precipitation of sulfides, hydroxides, carbonates, chromates, phosphates, etc.

Here are some cases of their use for analysis purposes:

Acetate buffer solution (CH3COOH + CH 3 COO Na ; pH \u003d 5) is used for the precipitation of precipitates that are not precipitated in acidic or alkaline solutions. The harmful effect of acids is 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 in the precipitation of barium, strontium, calcium carbonates and their separation from magnesium ions; during precipitation of nickel, cobalt, zinc, manganese, and iron sulfides; as well as in the isolation of hydroxides of aluminum, chromium, beryllium, titanium, zirconium, iron, etc.

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

Phosphate buffer solution ( N a 2 HPO 4 + N aH 2 RO; pH = 8) is used in carrying out many redox reactions.

To successfully use buffer mixtures for analysis, it must be remembered 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 concentration of their components.

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

Buffer solutions used in the assay

buffer mixture	Composition of the mixture (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	One-substituted and disubstituted 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 role of its salt.

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

Complex buffer systems also operate in animal and plant organisms, maintaining a constant pH of blood, lymph and other fluids. The soil also has buffering properties, which tend to counteract external factors that change the pH of the soil solution, for example, when acids or bases are introduced into the soil.

CONCLUSION

So, buffer solutions are called solutions that supportconstant pH 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 the solution is diluted. Solutions of acids and bases cannot be called buffer solutions, because when diluted with water, the pH of the solution changes. The most effective buffer solutions are prepared from solutions of a weak acid and its salt or a weak base and its salt.

Buffer solutions can be considered as mixtures of electrolytes having ions of the same name. 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, photographic materials. Buffer solutions are widely used in chemical analysis and for the calibration of pH meters.

Many biological fluids are buffer solutions. For example, the pH of the 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, the cultivation of bacteria also requires the use of buffer solutions.

REFERENCES

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

2. Tsitovich I.K. Course of Analytical Chemistry: Textbook for High Schools. - 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 university applicants / Ed. Lidia 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, Goshimtekhizdat, 1933.