Buffer capacity and factors influencing it. Buffer zone
The ability of solutions to maintain a constant pH value is not unlimited. A quantitative measure of the buffering effect of a solution is the buffer capacity (BC). Theoretically, BE is equal to the derivative of the amount of a strong acid or alkali substance added to 1 liter of buffer solution according to the change in pH:
Practically, BE is the number of equivalents of H+ or OH- ions that must be added to 1 liter of buffer solution in order to change its pH by one. When practically determining the buffer capacity for acid, the following formulas are used:
where CH (k-you) and V (k-you) are, respectively, the normal concentration and volume of the acid solution, when added to the volume of the buffer solution Vbuf.solution, the pH decreases by the amount ΔрН.
The buffer capacity for alkali is found according to the formula
where CH (basic) and V(basic) are, respectively, the normal concentration and volume of an alkali solution, when added to the volume of the buffer solution Vbuffer solution, the pH increases by the amount ΔрН
BE depends on the composition of the buffer solution, concentration and ratio of components.
· The greater the number of components of the base/conjugate acid base pair in solution, the higher the B.E. this solution.
· BE depends on the ratio of the concentrations of the components of the buffer solution, and, consequently, on the pH of the buffer solution.
· At pH = pKa, the ratio C (salt) / C (acid) = 1, i.e., the solution contains the same amount of salt and acid. With this ratio of concentrations, the pH of the solution changes to a lesser extent than with others, and, therefore, BE is maximum at equal concentrations of the components of the buffer system and decreases with deviation from this ratio. The BE of a solution increases as the concentration of its components increases and the ratio [HAn]/[KtAn] or [KtOH]/[KtAn] approaches unity.
The working area of the buffer system, i.e. the ability to counteract changes in pH when adding acids and alkalis, has an extent of approximately one pH unit on each side of the pH = pKa point. Outside this interval, the buffer capacity quickly drops to 0. pH interval = pKa ± 1
called the buffer zone
. A pronounced buffering effect is observed if the concentration of one of the components exceeds the concentration of the other by no more than 10 times. Accordingly, the boundaries of the buffer zone are:
Blood buffer systems: composition, distribution in plasma and erythrocytes, mechanism of action of hydrocarbonate, phosphate, protein buffer systems, normal blood pH, arterial and venous blood pH.
Blood contains 4 main buffer systems. 1. Hydrocarbonate. 2. Protein.3. Hemoglobin4. Phosphate buffer system.
Hydrocarbonate buffer
is represented by a mixture of substances H2CO3 and NaHCO3 in a ratio of 1: 20. This buffer is the main buffer system of blood plasma; it is a quick response system, since the product of its interaction with acids (CO2) is quickly eliminated through the lungs.
Mechanism of action
. If acids accumulate in the blood, the amount of HCO3- decreases and the reaction occurs: HCO3- + H+ ↔ H2CO3 ↔ H2O + CO2↑. The excess is removed by the lungs. However, the blood pH value remains constant, as the volume of pulmonary ventilation increases, which leads to a decrease in the volume of CO2. With an increase in blood alkalinity, the concentration of HCO3- increases: H2CO3 + OH- ↔ HCO3- + H2O. This leads to a slowdown in lung ventilation, so CO2 accumulates in the body and the buffer ratio remains unchanged.
Hemoglobin buffer-
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 CO2. 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. It is represented by two weak acids - hemoglobin and oxyhemoglobin and their conjugate bases - hemoglobinate and oxyhemoglobinate ions, respectively:
HHb ↔ H+ + Hb-
HHbO2 ↔ H+ +HbO2-
Oxyhemoglobin is a stronger acid (pKa = 6.95) than hemoglobin (pKa = 8.2). At pH = 7.25 (inside red blood cells), oxyhemoglobin is ionized by 65%, and hemoglobin by 10%, so the addition of oxygen to hemoglobin reduces the pH value of the blood, since a stronger acid is formed. On the other hand, as oxygen is released from oxyhemoglobin in the tissues, the pH value of the blood increases again.
The buffering properties of HHb are primarily due to the possibility of interaction of acid-reactive compounds with the potassium salt of hemoglobin with the formation of an equivalent amount of the corresponding potassium salt of the acid and free hemoglobin:
KHb + H2CO3 ↔ KHCO3 + HHb.
The resulting bicarbonate (KHCO3) balances the amount of incoming H2CO3, the pH is maintained, since the dissociation of potential H2CO3 molecules and the resulting hemoglobin acids occurs. This is how the blood pH is maintained within normal limits, despite the entry into the venous blood of a huge amount of CO2 and other acid-reacting metabolic products. In the capillaries of the lungs, hemoglobin (HHb) absorbs oxygen and turns into HHbO2, which leads to some acidification of the blood, displacement of some H2CO3 from bicarbonates and a decrease in the alkaline reserve of the blood, and in the tissues it releases it and absorbs CO2.
In the lungs: ННb + O2 ↔ HHbO2;
HHbO2 + HCO3- ↔ HbO2 + H2O + CO2 ↑
In tissues: HbO2 ↔ Hb- + O2; Hb- + Н2СО3 ↔ ННb + HCO3-
In addition, hemoglobin buffer is a complex protein and acts as a protein
buffer.
Phosphate buffer
is 5% of the buffer capacity. Contained both in the blood and in the cellular fluid of other tissues, especially the kidneys. In cells it is represented by the salts K2HPO4 and KH2PO4, and in the blood plasma and intercellular fluid Na2HPO4 and NaH2PO4. It functions mainly in plasma and includes: dihydrogen phosphate ion H2PO4- and hydrogen phosphate ion HPO42-. The ratio [HPO42- ]/[H2PO4-] in blood plasma (at pH = 7.4) is 4: 1. Therefore, this system has a buffer The capacity for acid is greater than for base. For example, when the concentration of H+ cations in the intracellular fluid increases, for example, as a result of processing meat food, they are neutralized by HPO42- ions:
H + + NPO42- ↔ N2PO41-
The resulting excess dihydrogen phosphate is excreted by the kidneys, which leads to a decrease in urine pH.
When the concentration of bases in the body increases, for example when eating plant foods, they are neutralized by H2PO41- ions:
OH ‾ + H2PO41- ↔ NPO42- + H2O
The resulting excess hydrogen phosphate is excreted by the kidneys, and the pH of the urine increases.
The excretion of certain components of the phosphate buffer system in the urine, depending on the food being processed, explains the wide range of urine pH values - from 4.8 to 7.5. The phosphate buffer system of the blood is characterized by a lower buffer capacity than the bicarbonate system due to the low concentration of blood components.
Protein buffer
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
When strong acids are formed in the body, they interact with the protein salt. This produces an equivalent amount of protein acid: HC1 + Pt-COONa ↔ Pt-COOH + NaCl. According to W. Ostwald's law of dilution, an increase in the concentration of a weak electrolyte reduces its dissociation, the pH remains virtually unchanged.
As alkaline foods increase, they interact with
Pt-COOH: NaOH + Pt-COOH ↔ Pt-COONa + H2O
The amount of acid decreases. However, the concentration of H+ ions increases due to the potential acidity of the protein-acid, so the pH practically does not change. Protein is an amphoteric electrolyte and therefore exhibits its own buffering effect.
Normally, the blood pH is 7.36, i.e. the reaction is weakly basic. Fluctuations in blood pH are extremely insignificant. Thus, under resting conditions, the pH of arterial blood corresponds to 7.4, and that of venous blood to 7.34. In cells and tissues, the pH reaches 7.2 and even 7.0, which depends on the formation of “acidic” metabolic products in them during the metabolic process. Under various physiological conditions, blood pH can change both acidic (up to 7.3) and alkaline (up to 7.5).
22. The concept of the acid-base state of the body: definition, significance for vital processes, alkaline blood reserve (%, mmol/l).
The acid-base state (ABS) refers to the ratio of the concentrations of hydrogen (H+) and hydroxyl (OH) ions in biological media. A necessary condition for the existence of a living organism is to maintain the constancy of this parameter of the internal environment. CBS is of paramount importance because:
· H+ ions are catalysts for many biochemical transformations;
· Enzymes and hormones exhibit biological activity at strictly defined pH values;
· The greatest changes in the concentration of H+ ions in the blood and interstitial fluid affect the value of their osmotic pressure.
· Deviation of blood pH (7.4) by 0.3 units. can lead to a coma, deviation by 0.4 units. may result in death. Saliva pH equal to 5 units. leads to the development of caries.
To the main indicators of WWTP
include blood pH, partial pressure of CO2, alkaline balance of blood. Normal blood pH is 7.4. A shift in pH towards an increase is called alkalosis, and towards a decrease - acidosis. The normal partial pressure of CO2 is 40 mm Hg. A decrease in this indicator is observed with respiratory alkalosis and metabolic acidosis. An increase in CO2 pressure is observed in respiratory acidosis and metabolic alkalosis.
Alkaline blood reserve
- an indicator of the functional capabilities of the blood buffer system, numerically coincides with the concentration of the bicarbonate anion (HCO3-) in the actual state of arterial blood plasma in the bloodstream.
Under physiological conditions it is 22-25 mmol/l. Another definition of blood alkaline reserve is the ability of circulating blood to bind CO2. It is calculated under conditions of equilibration of blood plasma at P(CO2) = 40 mm Hg: the total amount of CO2 is determined, from which the amount of physically dissolved CO2 in the blood serum under study is subtracted. The value is expressed in volume percent CO2 (in ml of CO2 per 100 ml of plasma), normally in humans it is 50-65 vol.% CO2. The concept of alkaline blood reserve
is closely related to the work of the body's hemoglobin buffer system, which helps maintain the pH level of circulating blood within physiological limits. A decrease in alkalinity indicates a decrease in the content of bicarbonates in the body, and an increase in it indicates an increase in them.
Buffer systems and their mechanism of action. Buffer capacity and its determining factors.
Buffer solutions are solutions whose pH changes little when small amounts of strong acids or alkalis are added to them, as well as when diluted. From the point of view of proton theory, the simplest buffer solution consists of a weak acid and its conjugate base or a weak base and its conjugate acid.
Classification of buffer systems 1. Acid. They consist of a weak acid and a salt of this acid. For example, an acetate buffer system (CH3COOH+ CH3COONa), a hydrocarbonate buffer system (H2CO3 + NaHCO3). 2. Basic. Consist of a weak base and its salt. For example, an ammonia buffer system (NH3H×2O + NH4Cl). 3. Salt. Consist of an acidic and a moderate salt or two acidic salts. For example, a carbonate buffer system (NaHCO3 + Na2CO3), a phosphate buffer system (KH2PO4 + K2HPO4). 4. Amino acid and protein. If the total charge of an amino acid or protein molecule is zero (isoelectric state), then solutions of these compounds are not buffers. Their buffering effect begins to appear when a certain amount of acid or alkali is added to them.
Mechanism of action of buffer systems:
1. Dilution. When diluted with water, the concentration of both components in the buffer system decreases to the same extent, so the value of their ratio will not change. pK(acids) and pK(bases) are constant at a given temperature and do not depend on dilution. Indeed, a simultaneous decrease in the concentrations of acid and salt in the acetate buffer system from 0.1 M to 0.001 M upon dilution with water changes the pH of the buffer solution from 4.63 to 4.73. Consequently, dilution ultimately changes the pH of buffer systems little.
2. Addition of acids and bases. When small amounts of strong acids or bases are added, the pH of buffer systems changes slightly. For example, consider an acetate buffer:
CH3COOH/ CH3COO–
When a small amount of HCl is added to the acetate buffer, H+ ions interact with the main component of the buffer solution: H+ + CH3COO– ⇄ CH3COOH. The degree of CH3COOH dissociation is small and the [H+] concentration remains virtually unchanged. The pH of the buffer solution will decrease, but only slightly. Thus, if X mol/l HCl is added to the acetate buffer, then the equation for calculating the pH of the buffer system takes the form: pH = pK(acids) + log
When a small amount of NaOH is added, the OH–– ions are neutralized by the acidic component of the buffer solution: OH–+ CH3COOH ⇄ CH3COO – + H2O.
As a result, the added strong base is replaced by an equivalent amount of a weak conjugate base (CH3COO–), which has less effect on the reaction of the medium. The pH of the buffer solution increases, but only slightly.
Thus, if Y mol/l NaOH is added to the acetate buffer, then the equation for calculating the pH of the buffer system takes the form: pH = pK(acid) + log
Buffer capacity (B) is the number of moles equivalent of a strong acid or alkali that must be added to 1 liter of a buffer solution to shift its pH by one.
The buffer capacity of the system is determined in relation to the added acid (Acid.) or base (alkali) (Rec.) and is calculated using the formulas: Acid.= Rec.=
39. Henderson-Hasselbach equation for calculating the pH of buffer systems (conclusion).
At high concentrations of H and OH, the ratio of the components of the buffer mixture changes significantly – Cc/Cc increases or decreases and the pH may change. This is confirmed by the Henderson-Hasselbach equation, which establishes the dependence of [H], Ki, α and Sk/Cc.
Using the example of an acid-type buffer system - a mixture of acetic acid and its salt CH3COONa. The concentration of hydrogen ions in the buffer solution is determined by the ionization constant of acetic acid:
CH3COOH↔CH3COO+H
Ki= =1.75*10-5
Let's transform it to:
=Ki or =Ki Sk/Ss.
Blood buffer systems.
Blood buffer systems are represented by blood plasma buffer systems and erythrocyte buffer systems. Plasma buffer systems are hydrocarbonate, protein and phosphate, the role of the latter is insignificant. They account for » 44% of the buffer capacity of the blood. Buffer systems of erythrocytes - hemoglobin, bicarbonate, organic phosphate system (phosphate). They account for 56% of the buffer capacity of the blood.
Since in the blood plasma the main role in the binding of H+ ions is played by bicarbonate anion, its concentration in the plasma determines the reserve alkalinity of the blood.
The hemoglobin buffer system is found only in red blood cells. The mechanism of its action is associated with the addition and release of oxygen. In this regard, hemoglobin (Hb) has oxidized HHbO2 and reduced HHb forms. HHb + O2 ⇄ HHbO2⇄ H+ + HbO2-
The hemoglobin buffer system in the body functions effectively only in combination with the bicarbonate system.
