Buffer System: Meaning and Importance (With Diagram)

Let us learn about Buffer System. After reading this article you will learn about: 1. Meaning of Buffer System 2. Importance of Buffer System.

Meaning of Buffer System:

A buffer system has the property of resisting pH changes despite additions of acid or base. A buffer is a mixture of an acid that does not ionize completely in water and its corresponding base-for example, carbonic acid (H₂CO₃) and sodium bicarbonate (NaHCO₃).

If acid is added to this buffer, the added H⁺ ions combine with bicarbonate ions to produce more carbonic acid, using up some of the H⁺ ions (the Na⁺ ions do not participate in this reaction).

If base is added, some of the carbonic acid ionizes to produce bicarbonate ions and more H⁺, which counteracts some of the pH. In this way, the buffer minimizes the effects of the added acid or base on the pH. Buffers illustrate the reversibility of chemical reactions, the addition of acid drives the reaction in one direction, whereas addition of base drives the chemical reaction in the other direction.

An adequate buffer system can be obtained from a weak acid mixed with the salt from that acid and a strong base. For example,

where B⁺ = the cation of a strong base e.g., Na⁺; A⁻ = anion from week acid

This corresponds to a partially neutralized solution of weak acid (HA→ H⁺ +A⁻) with some reasonable proportion of the acid converted to the anion. Its pH will be higher than the pH of the weak acid alone.

The pH of a buffer solution is determined not on the absolute concentrations of buffer constituents but on their ratio, i.e., on the ratio of the amount of salt to weak acid as stated in the Henderson Hasselbalch equation:

pH = pK_a + log [salt]/[acid]

Where K_a is the dissociation constant of the acid and pK_a is the pH at which the acid is half neutralized and equal amounts of acid and ion (salt) are present. This equation is applicable to all the buffer systems. The buffer capacity, measured as the ability of the solution to minimize changes in pH due to addition of base, is strongest near the midpoint of a titration, when [A⁻] = [HA] and pH = pK_a.

The most important examples of biological buffer systems are as follows:

1. Bicarbonate Buffer:

The major buffer of extracellular fluid is the HCO₃/ H₂CO₃) system.

This results from a number of factors:

(i) There is considerably more HCO3 present in extracellular fluid than any other buffer component

(ii) There is an unlimited supply of CO₂

(iii) Physiological mechanisms maintain normal extracellular pH by controlling either the HCO3 or the CO₂ concentration of extracellular fluid

(iv) This buffer system operates in conjunction with Hb.

The Henderson- Hasselbalch equation for this buffer system may be written as:

Moreover, in true plasma total CO₂ content is equal to [H₂CO₃] + [HCO3]. Hence, if the total C0₂ content of true plasma is known and the pCO₂ of alveolar air, i.e., arterial pCO₂ is determined, the pH of the true plasma can be calculated. This method of determining blood pH (true plasma pH) was used before the advent of commercial pH meters.

The Henderson Hasselbalch equation is important, for it enables us to appreciate that the blood pH is dependent on the ratio of the concentration of free acid to the concentration of the buffer anion (bicarbonate). Therefore, acidaemia or aikalaemia could be caused by either a change in the nonvolatile component (HO₃⁻ concentration) or in the volatile component (H₂CO₃ concentration and hence pCO₂).

Since the [H₂CO₃] is fixed only by the alveolar CO₂ tension, if the gas tension is equivalent to that normally present in the blood, the HCO₃⁻ /H₂CO₃ system is more useful at pH 7.4 with a ratio of 20 then it would be at its pK 6.1 where the HCO₃⁻ would be exhausted by addition of 1.25 meq per litre of acid, and addition of this amount of alkali would result in a rise of 0.3 pH unit.

The buffer efficiency of HCO₃7H₂CO₃ is further enhanced by the presence of erythrocytes As CO₂ diffuses into cells, the H₂CO₃ reacts with Hb forming HCO₃⁻, which then enters the plasma in exchange for CP. This is not contingent upon deoxygenation of Hb, but is achieved more readily and with even less pH change when deoxygenation occurs simultaneously.

Conversely, lowering the CO₂ tension results in a reversal of this process with consequent diminution of plasma [HCO₃⁻]although the situation does not arise under physiological conditions, only in the presence of red cells does the total CO₂ content of plasma fall to zero at a dCO₂ of zero mm.

This is possible because there is sufficient Hb in whole blood to permit the following series of reactions:

H⁺ + HbO₂→O₂ +HHb⁺

HHb⁺ + HCO₃ ⇋Hb + H₂CO₃⇋H₂O + CO₂

The normal value of 20: 1 of the ratio of HCO₃ ⁻/H₂CO₃ is also maintained through respiratory regulation of the pH of extracellular fluid. In addition to the automatic self-adjustments made possible by intracorpuscular Hb (as described above), the body possesses two further safeguards (the respiratory apparatus and the kidneys) which by their control of plasma [H₂CO₃] and [HCO₃^–] respectively serve in auxiliary fashion to maintain constant pH of blood plasma.

Unlike the [HCO₃^–] the [H₂CO₃] is determined solely by the partial pressure of CO₂ in the alveolar air in equilbrium with extracellular fluid. This, in turn, is dependent upon the rate at which CO₂ leaving pulmonary blood is diluted with atmospheric air and hence upon the rate and depth of respiration.

The latter are regulated by the nervous system at the respiratory centre which is sensitive to pH and pCO₂ of extracellular fluid. When the pH falls below the normal because of diminished [HCO₃^–], respiration is stimulated, lowering alveolar pCO₂ and hence extracellular [H₂CO₃]. This tends to return the HOC₃⁻/H₂CO₃ ratio to its normal value of 20:1 and thus to restore pH toward 7.4. With a high plasma pH, the respiratory rate falls, alveolar pCO₂ and hence plasma [H₂CO₃] rise, and the pH moves toward 7.4.

But perfect compensation is not attained since the increased plasma [H₂CO₃] opposes the effect of elevated pH on the respiratory centre. If the respiratory rate falls sufficiently, the diminished pO₂, sensed by the carotid body chemoreceptors, could be a stimulus for increased respiratory activity.

The pH is dependent not on absolute concentrations but solely on the HOC₃^–/H₂CO₃ ratio. The buffer system of plasma can withstand the addition of 16 meq of acid or 29 meq of alkali per litre and still maintain the pH within the range compatible with life. With pulmonary compensation the normal pH range can be maintained despite addition of 23 meq of acid or 80 meq of alkali per litre of plasma.

2. Phosphate Buffer:

Although the contribution of phosphate buffer HPO²⁻₄/ H₂PO₄^– to the buffering power of the plasma is negligible because their plasma concentration is very low (3mg/100 ml), this buffer system is present (as is HCO₃^–/ H₂CO₃) in the tubular fluid of the kidney.

The Henderson-Hasselbalch equation for this buffer system may be written as:

Thus, phosphate equilibrium in plasma is such that 4 parts of the so-called basic phosphate (HPO²⁻₄) exists to 1 part of acidic phosphate (H₂PO⁻₄) at a pH of 7.4, for pK_phos is 6.8.

At a urinary pH of 5.8 , ten parts of phosphate are in the ‘acidic’ form and one part in the ‘basic’ form (5.8 = 6.8 + 1⁻), for each phosphate ion excreted as H₂PO₄^– one Na⁺ ion is saved and one H⁺ ion is excreted. The phosphates are by far the most important buffers in the urine, although in acid urine the formation of NH₄ from NH₃ and H⁺ helps to keep the [H⁺] lower than it would otherwise be.

Thus, phosphate buffer system contributes little to the total buffering capacity of blood because the blood phosphate concentration is very low compared with the amount of protein present. But the phosphate buffer pair is a major outlet for H⁺ via the urine which has a relatively high phosphate content. Hence, the pH of urine is provided by the excretion of H⁺ into the tubular fluid.

3. Oxyhaemoglobin Buffer:

The buffer pair formed by HbO₂ /H.HbO₂ i.e., oxyhaemoglolbin anion/ oxyhaemoglobin (acid) is also important. It is well known that uptake of C0₂ from the tissue depends upon the following reactions:

Thus, the uptake of CO₂ from the tissues is dependent upon the supply of hydrogen acceptors so that the reaction can proceed towards the right. The most important hydrogen acceptor in blood is the haemoglobin ion.

Most of the power of haemoglobin in mopping up hydrogen ions in the range pH 7.0 to 7.70 comes from the dissociation of its imidazole groups [haemoglobin contains 38 molecules of histidine (β- imidazole α-amino propionic acid) per molecule]. The dissociation tendency of the imidazole group like that of any weak acid is largely dependent upon the pH of the solution.

As the pH rises, the dissociation is reduced and as a result the H⁺ ions added are ‘mopped up’.

Oxyhaemoglobin dissociates more completely than does reduced haemoglobin. As a result, reduced haemoglobin produces less H⁺ at a given pH than does oxyhaemoglobin. Thus, when oxyhaemoglobin solution is reduced the pH of solution falls, i.e., reduced haemoglobin is a better H⁺ acceptor than oxyhaemoglobin.

This is of great importance in blood physiology because the entrance of CO₂ from the tissues to the capillary blood is accompainied by the simultaneous reduction of oxyhaemoglobin. As the uptake of CO₂ depends upon H⁺ acceptors, this increase in H⁺ acceptance by the reduced form facilitates the uptake and buffering of CO₂.

For each mMol CO₂ of oxyhaemoglobin reduced, abut 0.7 mMol of H⁺ can be taken up; as a result 0.7 mMol CO₂ can enter the blood without any change in pH.

As arterial blood enters the tissues, CO₂ diffuses into the erythrocytes, thus, potentially lowering the pH and the affinity of Hb for O₂ as shown in the following reactions:

On the other hand, in the lungs the loss of CO₂ which would potentially raise the pH, increases the affinity of the Hb for O₂, thus permitting the saturation of Hb with O₂ at a lower pO₂.

it should be noted that oxygenation of Hb results in a shift in the apparent pK_a of some acidic groups on the peptide chains from 7.71 to 6.17; thus HbO₂acts as a stronger acid than Hb. Approximately 0.7 mol of H⁺ is released as 1 equiv, of O₂ is bound.

Importance of Buffer System:

Buffer mixtures are very important in living organisms and the mineral world. An example of a natural buffer is the blood of mammals. It always contains free carbonic acid and sodium carbonate. Therefore, the pH of blood is always maintained at 7.4.

The buffering action of soils is very important in agriculture, because plants absorb artificial fertilizers from the soil to change the pH in solutions that they extract from the soil in an unfavourable direction. An imbalance in the buffering action of soil is detrimental to useful micro-organisms living in it.

Buffer solutions are very important in the treatment of domestic sewage, because the microorganisms which mineralize their organic matter thrive better in a neutral medium. A shift towards acidity or alkalinity inhibits the vital processes in the microbes, thus adversely affecting the working of sewage treatment plants.

Buffers play an important role in the chemical treatment of water to separate it from suspended matter of coagulation. The higher the buffer capacity of the treated water, the more efficient its purification with a hydrolyzing coagulant. The buffer capacity of neutral water accounts for its neutralising power.

Buffers are widely used in volumetric analysis. For example, an ammonium buffer (NH₄OH + NH₄CI) is used to determine Ca and Mg ions in water (triionometric method). Buffer mixtures are used to determine the pH of solutions colorimetrically and potentiometrically.

In cellular physiology, buffers are an important way for the cell to maintain constant or smoothly changing conditions. Since cellular metabolism is constantly producing and consuming protons, the pH of the cell in the absence of buffers would change rapidly between low and high levels.