Regulation of Acid Base Balance

Dr Ameer Khalid  BHMSMD(Hom)

Regulation of acid-base balance, regulation of hydrogen ion concentration in the body fluids is actually meant. The hydrogen ion concentration in different solutions can vary from less than 10-14 Eq/litre to higher than 100 Eq/litre.

Only slight changes in hydrogen ion concentration from the normal value can cause marked alterations in the rates of chemical reactions in the cells, some being depressed and others accelerated. For this reason the regulation of hydrogen ion concentration is one of the most important aspects of homeostasis.

Acid and bases definition
An acid is defined as a molecule or an ion that can function as a proton donor. An acid is a molecule or an ion that can contribute a hydrogen ion to a solution.
A base is defined as a molecule or an ion that can function as a proton acceptor. A base is a molecule or an ion that will combine with hydrogen ions to remove them from a solution.

Relation between base and alkali:-alkali is combination of one of the alkaline metals –sodium, potassium or so forth with highly Basic ion such as hydroxyl ions.
Since the alkali is well known the word alkali is frequently used synonymously with the term base and for the similar reason the term alkalosis is used to mean the opposite of acidosis.

Strong and Weak Acids and Bases:– A strong acid is one that has a very strong tendency to dissociate into ions and therefore to discharge its hydrogen ion into the solution. A typical example is hydrochloric acid.
On the other hand, acids that release hydrogen ions with far less vigor are called weak acids. Examples are carbonic acid and sodium acid phosphate.

A strong base is one that reacts very powerfully with hydrogen ions and therefore removes these with extreme acidity from the solution. A typical example is the hydroxyl ion (OH-). A typical weak base is the bicarbonate ion (HCO3- ), for it binds much more weakly with hydrogen ions.
Most of the acids and bases that are concerned with normal regulation of acid-base balance in the body are weak acids and bases, the most important of which are carbonic acid and bicarbonate base.
Hydrogen Ion Concentration and pH of Normal Body Fluids and changes that occur in Acidosis and Alkalosis.
Expressing hydrogen ion concentration in terms of actual concentrations is a difficult procedure. Therefore, the symbol pH has come into usage for expressing the concentration; the term pH was introduced in 1909 by Sorensen who defined pH as negative log of the hydrogen ion concentration. pH is related to acid hydrogen ion concentration by the following formula.

A low pH corresponds to a high hydrogen ion concentration, which is called acidosis; and, conversely, a high pH corresponds to a low hydrogen ion concentration, which is called alkalosis.

pH and H+ concentration of body fluids

	H+ concentration mEq/L	pH
Extra cellular fluid arterial blood venous blood interstitial fluids	4.0 X 10-5 4.5 X 10-5 4.5 X 10-5	7.4 7.35 7.35
Intracellular fluid	1 X10-3 to 4.0 X 10-5	6.0 to 7.4
Urine	3 X 10-2 to 1 X 10-5	4.5 to 8.0
Gastric HCl	160	0.8

Since the normal pH of the arterial blood is 7.4, a person is considered to have acidosis whenever the pH is below this value and to have alkalosis when it rises above 7.4. The lower limit at which a person can live more than a few hours is about 6.8, and the upper limit approximately 8.0.
Defenses against changes in hydrogen ion concentration buffers, lungs and kidneys

There are three primary systems that regulate the hydrogen ion concentration in the body fluids
1. The chemical acid base buffer system of the body fluids, which immediately combine with acid or base to prevent excessive changes in hydrogen ion concentration.
2. The respiratory centre which regulates the removal of CO2 (and therefore H2CO3) from the extracelluar fluid.
3. The kidneys which can excrete either acid or alkaline urine, there by readjusting the extra cellular fluid hydrogen ion concentration toward normal during acidosis or alkalosis.

To prevent acidosis or alkalosis, several special control systems are available: (1) All the body fluids are supplied with acid-base buffer systems that immediately combine with any acid or base and thereby prevent excessive changes in hydrogen ion concentration. (2) If the hydrogen ion concentration does change measurably, the respiratory center is immediately stimulated to alter the rate of breathing. As a result, the rate of carbon dioxide removal from the body fluids is automatically changed; and, for reasons that will be presented later, this causes the hydrogen ion concentration to return toward normal. (3) When the hydrogen ion concentration changes from normal, the kidneys excrete either acid or alkaline urine, thereby also helping to readjust the hydrogen ion concentration of the body fluids back to normal.
The buffer systems can act within a fraction of a second to prevent excessive changes in hydrogen ion concentration. On the other hand, it takes 1 to 12 minutes for the respiratory system to make acute adjustments and another day or so to make still additional chronic adjustments. Finally, the kidneys, although providing the most powerful of all the acid-base regulatory systems, require many hours to several days to readjust the hydrogen ion concentration.
From this equation it can be seen that the strong hydrochloric acid is converted into the very weak carbonic acid. Therefore, addition of the HCl lowers the pH of the solution only slightly.
Now let us see what happens when a strong base, such as sodium hydroxide, is added to a buffer solution containing carbonic acid; the following reaction takes place:
NaOH + H2C03 NaHC03 + H20 (3)
This equation shows that the hydroxyl ion of the sodium hydroxide combines with a hydrogen ion from the carbonic acid to form water and that the other product formed is sodium bicarbonate. The net result is exchange of the strong base NaOH for the weak base NaHC03.

Acid base balance by buffers
Function of buffers
An acid-base buffer is a solution containing two or more chemical compounds that prevents marked changes in hydrogen ion concentration when either an acid or a base is added to the solution. As an example, if only a few drops of concentrated hydrochloric acid are added to a beaker of pure water, the pH of the water might immediately fall from a neutral 7.0 to as low as 1.0. However, if a satisfactory buffer system is present, the hydrochloric acid combines instantaneously with the buffer, and the pH falls only slightly. Perhaps the best way to explain the action of an acid base buffer is to consider an actual simple buffer system, such as the bicarbonate buffer, which is extremely important in regulation of acid-base balance in the body.

Buffering of hydrogen ions in the body fluids
A buffer is any substance that can reversibly bind H+. the general form of the buffering reaction is
Buffer + H+ H Buffer
In this example, a free H+ combines with the buffer to form a weak acid (HBuffer) that either remain as an unassociated molecule or dissociate back to buffer and H+. When the H+ ion concentration increases, the reaction is forced to right and more H+ bind to the buffer, as long as available buffer is present. Conversely when the H+ concentration decreases, the reaction shifts towards the left and H+ are released from the buffer. In this way changes H+ concentration are minimized.

I. The bicarbonate buffer system
A typical bicarbonate buffer system consists of a mixture of carbonic acid (H2C03) and sodium bicarbonate (NaHC03) in the same solution.
H2C03 is formed in the body by the reaction of CO2 with H2O
CO2+H2O H2CO3
Carbonic anhydrase
This reaction is slow unless enzyme carbonic anhydrase is present. This enzyme is especially abundant in the wall of lung alveoli; where CO2 is released; also epithelial cells of renal tubules where CO2 reacts with H2O to form HC03- .
H2C03 ionises weakly to form small amount of H+ and HC03-
H2CO3 H+ + HCO3-
The second component bicarbonate salt occurs predominantly as sodium bicarbonate NaHC03 in the extracellular fluid. NaHC03 ionises almost completely to form bicarbonate ions and sodium ions.
NaHC03 Na+ + HC03-
Now putting the entire system together,
CO2+H2O H2CO3 H+ + HCO3-
{ +Na+
Because of the weak dissociation of H2C03 the H+ concentration is extremely small.
When a strong acid such as HCl is added to bicarbonate buffer solution, the increased H+ released from acid.
HCl H+ + Cl- are buffered by HC03-
↑ H++ HC03- H2CO3 CO2+H2O
As a result, more H2C03 is formed, causing increased CO2 and H2O production. From this reaction, the hydrogen ions from the strong acid, HCl, react with HC03- to form CO2 and H2O. The excess CO2 greatly stimulates respiration, which eliminates the CO2 from the extracellular fluid.
The opposite reactions take place when a strong base, such as sodium hydroxide (Na OH), is added to the bicarbonate buffer solution.
Na OH + H2CO3 NaHCO3+H2O
In this case, the hydroxyl ion (OH-) from the NaOH combines with H2CO3 to form additional. Thus the weak base NaHCO3 replaces the strong base NaOH. At the same time, the concentration of H2CO3 decreases (because it reacts with NaOH), causing more CO2 to combines with H2O to replace the H2CO3.
CO2+H2O H2CO3 ↑ HCO3- + H+
+ +
NaOH Na
The net result is a tendency for the CO2 levels in the blood to decrease; but the decreased CO2 in the blood inhibits respiration and decreases the rate of CO2 expiration. The rise in blood HCO3- that occurs is compensated for by increased renal excretion of HCO3-.

Quantitative dynamics of the bicarbonate buffer system
All acids, including H2CO3 are ionized to some extent. From mass balance considerations, the concentrations of hydrogen ions and bicarbonate ions are proportional to the concentration of H2CO3.
H2CO3 H++ HCO3-
For any acid, the concentration of the acid relative to its dissociated ions is defined by the dissociation constant K’.
K’ = (1)

This equation indicates that in an H2CO3 solution, the amount of free hydrogen ions is equal to
H+= (2)
The concentration of undissociated H2CO3 cannot be measured in solution because it rapidly dissociates into CO2 and H2O or to H+ and HCO3-. However, the CO2 dissolved in the blood is directly proportional to the amount of undissociated H2CO3. There fore, equation (2) can be written as
H+= (3)
The dissociation constant (K) for Equation (3) is only about 1/400 of the dissociation constant (K’) of equation (2) because the proportionality ratio between H2CO3 and CO2 is 1to 400.
Equation (3) is written in terms of the total amount of CO2 dissolved in solution. However, most clinical laboratories measure the blood CO2 tension (Pco2) rather than the actual amount of CO2. Fortunately, the amount of CO2 in the blood is a linear function of times the solubility coefficient for CO2; under physiologic conditions the solubility coefficient for CO2 is 0.03 mmol/mmHg at body temperature .this means that 0.03mmmol of H2CO3 is present in the blood for each millimeter of mercury Pco2 measured. There fore equation (3) can be rewritten as
H+= (4)
Henderson-Hasselbalch Equation.
As discussed earlier, it is customary to express hydrogen ion concentration in pH units rather than in actual concentrations. Recall that pH is defined as
pH = -log H
The dissociation constant can be expressed in a similar manner pK = -log K
Therefore, we can express the hydrogen ion concentration in equation (4) in pH units by taking the negative logarithm of that equation, which yields
-log H+ = -log pK – log (5)
pH = pK — (6)
Rather than work with a negative logarithm, we can change the sign of the logarithm and invert the numerator and denominator in the last term, using the law of logarithms to yield
pH = pK + (7)
For the bicarbonate buffer system,the pK is 6.1, and equation 7 can be written as
pH = (8)
Equation 8 is the Henderson-Hasselbalch equation, and with it one can calculate the pH of a solution if the molar concentration of bicarbonate ion and the Pco2 are known.

From the Henderson-Hasselbalch, equation, it is apparent that an increase in bicarbonate ion concentration causes the pH to rise, shifting the acid-base balance toward alkalosis. And an increase in Pco2 causes the pH to decrease, shifting the acid base balance toward acidosis.

The Henderson-Hasselbalch equation, in addition to defining the determinants of normal pH regulation and acid-base balance in the extracellular fluid, provides insight into the physiologic control of acid and base composition of the extracellular fluid. The bicarbonate concentration is regulated mainly by the kidneys, whereas the Pco2 in extra cellular fluid is controlled by the rate of respiration. By increasing the rate of respiration, the lungs remove CO2 from the plasma, and by decreasing respiration, the lungs elevate Pco2. Normal physiologic acid-base homeostasis results from the coordinated efforts of both of these organs, the lungs and the kidneys, and acid-base disorders occur when one or both of these control mechanisms are impaired, thus altering either the bicarbonate concentration or the Pco2 of extracellular fluid.

When disturbances of acid-base balance result from a primary change in extracellular fluid bicarbonate concentration, they are referred to as metabolic acid-base disorders. Therefore, acidosis caused by a primary decrease in bicarbonate concentration is termed metabolic acidosis, whereas alkalosis caused by a primary increase in bicarbonate concentration is called metabolic alkalosis. Acidosis caused by an increase in Pco2 is called respiratory acidosis, whereas alkalosis caused by a decrease in Pco2 is termed respiratory alkalosis.

Bicarbonate Buffer System Titration Curve.
When the concentrations of HCO3- and CO2 are equal, the right-hand portion of equation 8 becomes the log of 1, which is equal to 0. Therefore, when the two components of the buffer system are equal, the pH of the solution is the same as the pK (6.1) of the bicarbonate buffer system. When base is added to the system, part of the dissolved CO2 is converted into HCO3- causing an increase in the ratio of HCO to CO, and increasing the pH, as is evident from the 1-lendersonHasselbalch equation. When acid is added, it is buffered by HCO,-, which is then converted into dissolved CO2 decreasing the ratio of HCO3- to CO2 and decreasing the pH of the extracellular fluid.

II. The phosphate buffer system
Although the phosphate buffer system is not important as an extracellular fluid buffer, it plays a major role in buffering renal tubular fluid and intracellular fluids.
The main elements of the phosphate buffer system are H2PO4- and HPO4- . When a strong acid such as HCl is added to a mixture of these two substances, the Hydrogen is accepted by the base HPO4- and converted to H2PO4-
HCl + Na2 HPO4 Na2HPO4+ NaCl
The result of this reaction is that the strong acid, HCl, is replaced by an additional amount of a weak acid, NaH2PO4, and the decrease in pH is minimized.
When a strong base, such as NaOH, is added to the buffer system, the OH- is buffered by the H2PO4 to form additional amounts of HPO4 + H20.
NaOH + NaH2PO4 Na2HPO4 + H20

In this case, a strong base NaOH, is traded for a weak base, Na2HP04, causing only a slight increase in the pH.

The phosphate buffer system has a pK of 6.8, which is not far from the normal pH of 7.4 in the body fluids; this allows the system to operate near its maximum buffering power. However, its concentration in the extracellular fluid is low, only about 8 per cent of the concentration of the bicarbonate buffer. Therefore, the total buffering power of the phosphate system in the extracellular fluid is much less than that of the bicarbonate buffering system.
In contrast to its rather insignificant role as an extracellular buffer, the phosphate buffer is especially important in the tubular fluids of the kidneys for two reasons: (1) phosphate usually becomes greatly concentrated in the tubules, thereby increasing the buffering power of the phosphate system, and (2) the tubular fluid usually has a considerably lower pH than extracellular fluid, bringing the operating range of the buffer closer to the pK (6.8) of the system.
The phosphate buffer system is also important in buffering intracellular fluids because the concentration of phosphate in these fluids is many times that in the extracellular fluids. Also, the pH of intracellular fluids is lower than that of extracellular fluid and therefore usually closer to the pK of the phosphate buffer system, compared with the extracellular fluid.

III. Proteins: important intracellular buffers
Proteins are among the most plentiful buffers in the body because of their high concentrations, especially within the cells.
The pH of the cells, although slightly lower than in the extracellular fluid, nevertheless changes approximately in proportion to extracellular fluid pH changes. There is a slight amount of diffusion of hydrogen and bicarbonate ions through the cell membrane, although these ions require several hours to come to equilibrium with the extracellular fluid, except for rapid equilibrium that occurs in the red blood cells. CO, however, can rapidly diffuse through all the cell membranes. This diffusion of the elements of the bicarbonate buffer system causes the pH in intracellular fluids to change when there are changes in extracellular pH. For this reason, the buffer systems within the cells help to prevent changes in pH of extracellular fluids but may take several hours to become maximally effective.

In the red blood cell, hemoglobin is an important buffer as follows:
H+ + Hb     HHb
Approximately 60 to 70 per cent of the total chemical buffering of the body fluids is inside the cells, and most of this results from the intracellular proteins. However, except for the red blood cells, the slowness of movement of hydrogen ions and bicarbonate ions through the cell membranes often delays for several hours the maximum ability of the intracellular proteins to buffer extracellular acid-base abnormalities.
Another factor besides the high concentration of proteins in the cells that contributes to their buffering power is the fact that the pKs of many of these protein systems are fairly close to 7.4.

References:
1. Text Book Of Medical Physiology-Guyton & Hall
2. Review Of Medical Physiology William .F. Ganong
3. Harpers Biochemistry -Robert K Murray, Daryl K Grunner, Peter A Mayes & Victor W Rodwell