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:

1. 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);
2. 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.

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.

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:

1. 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.
2. 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.
3. 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-.
4. 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.

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.

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.

Mechanism of buffer action (using the example of ammonia buffer)

Let's consider the mechanism of action of the buffer system using the example of an ammonia buffer system: NH 4 OH (NH 3 x H 2 O) + NH 4 C1.

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

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

When ammonium chloride is added to a solution of ammonium hydroxide, the salt, as a strong electrolyte, almost completely dissociates into ions NH 4 C1 > NH 4 + + C1 - and suppresses the dissociation of the base, the equilibrium of which shifts towards the reverse reaction. Therefore, C (NH 4 OH)? C (base); and C (NH 4 +) ? C (salt).

If in a buffer solution C (NH 4 OH) = C (NH 4 C1), then pH = 14 - pKosn. = 14 + log 1.8.10-5 = 9.25.

The ability of buffer mixtures to maintain an almost constant pH value of a solution is based on the fact that their components bind H+ and OH- ions introduced into the solution or formed as a result of the reaction occurring in this solution. When a strong acid is added to an ammonia buffer mixture, H+ ions will bind to ammonia or ammonium hydroxide molecules rather than increasing the concentration of H+ ions and decreasing the pH of the solution.

When adding alkali, OH - ions will bind NH 4 + ions, forming a slightly dissociated compound, rather than increasing the pH of the solution.

The buffering effect ceases as soon as one of the components of the buffer solution (conjugate base or conjugate acid) is completely consumed.

To quantitatively characterize the ability of a buffer solution to resist the influence of strong acids and bases, a value called buffer capacity is used. As the concentration of a buffer solution increases, its ability to resist changes in pH when acids or alkalis are added increases.

The property of solutions to maintain the pH value within certain limits when small amounts of acid or alkali are added is called buffering action. Solutions that have a buffering effect are called buffer mixtures.

For the titration case: oxalic acid and potassium hydroxide, draw the titration curve, indicate the titration case, titration jump, equivalence point, indicators used

Titration jump: pH = 4-10. The maximum error in% is less than 0.4.

Indicators - thymolphthalein, phenolphthalein.

Reducing agent, which elements of the periodic table of elements can be reducing agents and why?

A reducing agent is a substance that gives up electrons during a reaction, i.e. oxidizes.

Reducing agents can be neutral atoms, negatively charged non-metal ions, positively charged metal ions in a lower oxidation state, complex ions and molecules containing atoms in an intermediate oxidation state.

Neutral atoms. Typical reducing agents are atoms with 1 to 3 electrons in their outer energy level. This group of reducing agents includes metals, i.e. s-, d- and f-elements. Non-metals, such as hydrogen and carbon, also exhibit reducing properties. In chemical reactions they give up electrons.

Strong reducing agents are atoms with low ionization potential. These include atoms of elements of the first two main subgroups of the periodic system of elements D.I. Mendeleev (alkali and alkaline earth metals), as well as Al, Fe, etc.

In the main subgroups of the periodic system, the reducing ability of neutral atoms increases with increasing radius of the atoms. So, for example, in the series Li - Fr, the weaker reducing agent will be Li, and the strong reducing agent will be Fr, which is generally the strongest reducing agent of all the elements of the periodic table.

Negatively charged nonmetal ions. Negatively charged ions are formed by adding one or more electrons to a neutral nonmetal atom:

So, for example, neutral atoms of sulfur and iodine, which have 6 and 7 electrons in their outer levels, can add 2 and 1 electron, respectively, and turn into negatively charged ions.

Negatively charged ions are strong reducing agents, since under appropriate conditions they can give up not only weakly held excess electrons, but also electrons from their outer level. Moreover, the more active a nonmetal is as an oxidizing agent, the weaker its reducing ability in the state of a negative ion. And vice versa, the less active a nonmetal is as an oxidizing agent, the more active it is in the negative ion state as a reducing agent.

The reducing ability of negatively charged ions with the same charge increases with increasing atomic radius. Therefore, for example, in the group of halogens, the iodine ion has a greater reducing ability than bromine and chlorine ions, while fluorine does not exhibit reducing properties at all.

Positively charged metal ions in the lowest oxidation state. Metal ions in the lowest oxidation state are formed from neutral atoms as a result of the loss of only part of the electrons from the outer shell. For example, atoms of tin, chromium, iron, copper and cerium, when interacting with other substances, can initially give up a minimum number of electrons.

Metal ions in a lower oxidation state can exhibit reducing properties if states with a higher oxidation state are possible for them.

In the OVR equation, arrange the coefficients using the electronic balance method. Specify the oxidizing agent and the reducing agent.

K 2 Cr 2 O 7 + 6FeSO 4 + 7H 2 SO 4 = K 2 SO 4 + Cr 2 (SO 4) 3 + 3Fe 2 (SO 4) 3 + 7H 2 O

1 Cr 2 +6 +3e x 2 Cr 2 +3 oxidizing agent

6 Fe +2 - 1е Fe +3 reducing agent

2KMnO 4 + 5H 2 S + 3H 2 SO 4 = K 2 SO 4 + 2MnSO4 + 5S + 8H 2 O

2 Mn +7 + 5е Mn +2 oxidizing agent

5 S -2 - 2е S 0 reducing agent

Buffer systems(buffers) are solutions that have the property of sufficiently, persistently, and maintaining a constant concentration of hydrogen ions both when adding acids or alkalis, and during dilution.

Buffer systems (mixtures or solutions) are of two main types in composition:

a) from a weak acid and its salt formed by a strong base;

b) from a weak base and its salt formed by a strong acid.

In practice, the following buffer mixtures are often used: acetate buffer CH 3 COOH + CH 3 COONa, bicarbonate buffer H 2 CO 3 + NaHCO 3, ammonia buffer NH 4 OH + NH 4 Cl, protein buffer protein acid + protein salt, phosphate buffer NaH 2 PO 4 + Na 2 HPO 4

A phosphate buffer mixture consists of two salts, one of which is a monometallic salt and the other a dimetallic salt of phosphoric acid.

Acetate buffer.

Let's consider buffering mechanism. When hydrochloric acid is added to the acetate buffer, interaction occurs with one of the components of the mixture (CH3COOH); From equation (a), the strong acid is replaced by an equivalent amount of the weak acid (in this case, HCl is replaced by CH 3 COOH). In accordance with Ostwald's dilution law, an increase in the concentration of acetic acid reduces the degree of its dissociation, and as a result, the concentration of H + ions in the buffer increases slightly. When alkali is added to the buffer solution, the concentration of hydrogen ions and pH also changes slightly. The alkali will react with another component of the buffer, (CH 3 COOH) through a neutralization reaction. As a result of this, the added alkali is replaced by an equivalent amount of a weakly basic salt, which affects the reaction of the medium to a lesser extent. The CH3COO~ anions formed during the dissociation of this salt will have some inhibitory effect on the dissociation of acetic acid.

Buffer solutions, depending on their composition, are divided into 2 main types: acidic and basic.

An example of an acidic buffer is an acetate buffer solution containing a mixture of acetic acid and sodium acetate (CH3COOH + CH3COONa). When an acid is added to such a solution, it interacts with the salt and displaces an equivalent amount of a weak acid: CH3COONa + HCl ó CH 3 COOH + NaCl. In the solution, instead of a strong acid, a weak one is formed, and therefore the pH value decreases slightly. If an alkali is added to this buffer solution, it is neutralized by a weak acid, and an equivalent amount of salt is formed in the solution: CH3COOH + NaOH ó CH3COONa + H 2 O. As a result, the pH almost does not increase. To calculate the pH in a buffer solution, using an acetate buffer as an example, we will consider the processes occurring in it and their influence on each other. Sodium acetate almost completely dissociates into ions, the acetate ion undergoes hydrolysis, like an ion of a weak acid: CH3COONa -> Na + + CH 3 COO ~ CH3COO - + NOH ó CH3COON + OH - . Acetic acid, also included in the buffer, dissociates only to a small extent: CH3COOH ó CH 3 COO + H -- Weak dissociation of CH3COOH is even more suppressed in the presence of CH3COON, therefore the concentration of undissociated acetic acid is taken to be almost equal to its initial concentration: [CH3COOH] = c r . On the other hand, hydrolysis of the salt is also suppressed by the presence of acid in the solution. Therefore, we can assume that the concentration of acetate ions in the buffer mixture is practically equal to the initial salt concentration without taking into account the concentration of acetate ions formed as a result of acid dissociation: [СН3СОО] = с с . This equation is called the buffer solution equation (Henderson Hasselbach equation ). His analysis for a buffer solution formed by a weak acid and its salt shows that the concentration of hydrogen ions in the buffer solution is determined by the dissociation constant of the weak acid and the ratio of the concentrations of the acid and salt. Henderson-Hasselbach equation for basic type buffer systems:

31. Capacity of buffer solutions and factors determining it. Blood buffer systems. Hydrogen carbonate buffer. Phosphate buffer.

Buffer capacity(B) is the amount of strong acid or strong base that must be added to one liter of a buffer solution to change its pH by one. It is expressed in mol/l or more often in mmol/l and is determined by the formula: B = (c V) / d pH Vb, where B is the buffer capacity; c is the concentration of a strong acid or base (mol/l); V is the volume of added strong electrolyte (l); V b - volume of buffer solution (l); d pH - change in pH.

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.