Question

Describe the different mechanism of pH maintenance at the cellular and blood level.

Answer 1

mechanism of pH maintenance at the cellular level

Intracellular pH is typically lower than extracellular pH due to lower concentrations of HCO3−. A rise of extracellular (e.g., serum) partial pressure of carbon dioxide (pCO2) above 45 mmHg leads to formation of carbonic acid, which causes a decrease of pHi as it dissociates:

H2O + CO2 ⇌ H2CO3 ⇌ H+ + HCO3–

Since biological cells contain fluid that can act as a buffer, pHi can be maintained fairly well within a certain range.Cells adjust their pHi accordingly upon an increase in acidity or basicity, usually with the help of CO2 or HCO3– sensors present in the membrane of the cell.[3] These sensors can permit H+ to pass through the cell membrane accordingly, allowing for pHi to be interrelated with extracellular pH in this respect.

Major intracellular buffer systems include those involving proteins or phosphates. Since the proteins have acidic and basic regions, they can serve as both proton donors or acceptors in order to maintain a relatively stable intracellular pH. In the case of a phosphate buffer, substantial quantities of weak acid and conjugate weak base (H2PO4– and HPO42–) can accept or donate protons accordingly in order to conserve intracellular pH:

OH– + H2PO4– ⇌ H2O + HPO42–

H+ + HPO42– ⇌ H2PO4–

In organelles

Approximate pHs of various organelles within a cell.

The pH within a particular organelle is tailored for its specific function.

For example, lysosomes have a relatively low pH of 4.5.Additionally, fluorescence microscopy techniques have indicated that phagocytes also have a relatively low internal pH.Since these are both degradative organelles that engulf and break down other substances, they require high internal acidity in order to successfully perform their intended function.

In contrast to the relatively low pH inside lysosomes and phagocytes, the mitochondrial matrix has an internal pH of around 8.0, which is approximately 0.9 pH units higher than that of inside intermembrane space. Since oxidative phosphorylation must occur inside the mitochondria, this pH discrepancy is necessary to create a gradient across the membrane. This membrane potential is ultimately what allows for the mitochondria to generate large quantities of ATP.

MECHANISM OF pH CONTROL OF BLOOD

NORMAL pH

There is a normal pH value in each body compartment (i.e. extracellular fluid, plasma, intracellular fluid etc). Intracellular pH is difficult to measure and may vary in different types of cells and in different parts of cells.

pH of the plasma (i.e. pH of the plasma of whole blood = conventional "blood" pH) is controlled at 7.4 (7.35 - 7.45). This section discusses the processes which restore the blood pH to normal if it is displaced.

Changes in plasma pH reflect pH changes in other compartments. When the source of pH change is intracellular the plasma pH change will be in the same direction as the intracellular pH change but of lesser magnitude. When the primary change is in the extracellular fluid the magnitude of any intracellular change will be less than the extracellular change.

Theoretically, opposite pH changes could occur from shifts of acid or base from one point of the body to another. Proving that such a change has occurred is generally impossible.

K+ shifts are said to do this but the evidence is nebulous and the conclusions conflicting.

There are three mechanisms which diminish pH changes in body fluid: buffers; respiratory; renal.

THE BUFFER SYSTEMS OF THE BODY

(a) Proteins are the most important buffers in the body. They are mainly intracellular and include haemoglobin. The plasma proteins are buffers but the absolute amount is small compared to intracellular protein. Protein molecules possess basic and acidic groups which act as H+ acceptors or donors respectively if H+ is added or removed.

(b) Phosphate buffer (H2PO4- : HP042-) is mainly intracellular. The pK of this sytem is 6.8 so that it is moderately efficient at physiological pH's. The concentration of phosphate is low in the extracellular fluid but the phosphate buffer system is an important urinary buffer.

(c) H2CO2 : HCO3- is not an important true buffer system because normal blood pH (7.4) is so far from its pK (6.1). H2CO3 and HCO3- are involved in pH control but they are not acting as a buffer system

Normal Acid Load. In dealing with the normal acid load from diet and metabolism (carbonic and other acids) the buffers are only involved in diminishing pH changes in the blood as it passes through capillaries (e.g. when CO2 is added in tissue capillaries or removed in pulmonary capillaries). The normal pH of the intracellular and interstitial fluids is maintained not by buffer action but because acids are removed at the same rate as they are added. As there is no change in the quantity of acid in the interstitial fluid or cells with time there is no change in pH. The buffers diminish change in pH due to short term minor physiological disturbances eg. breath-holding (addition of CO2), severe exercise (lactic acid) or during secretion of gastric acid

Abnormal Acid Balance. If there is an abnormality in the acid balance, i.e. acid is added faster than it is removed, resulting in a raised level of acid, the change in pH is less than would have occurred if the same imbalance had occurred in a non-buffer solution.

Low Buffer States. In theory a low protein level, (i.e. hypoproteinaemia or anaemia) may make a patient more sensitive to a positive acid balance. This is not recognised clinically. Abnormalities of the buffering system as such do not produce appreciable abnormalities in either the pH status or the acid-base balance because a situation where there were little or no buffers (phosphate and protein) would be incompatible with life.

RESPIRATORY CONTROL

Normal. In the normal state (at 37C) when there is no non-respiratory disturbance in pH the carbonic acid level is kept constant in the blood at 1.2meq/l or PaCO2 of 40mmHg (5.3kPa). (PaCO2 x 0.03) = H2CO3 meq/litre.

Effect of Control of PCO2 in Minimising pH Changes due to Non-Respiratory Acids or to Bases. Maintenance of the PaCO2 level, is very important in diminishing pH changes when non-respiratory changes occur. In a closed system where only true buffering could occur, large changes in PCO2 would occur when the levels of acids other than carbonic or bases are changed. For example, if strong acid were added part of all of the bases in the buffer pairs would be neutralised. Some HCO3- would therefore become H2CO3 causing the PCO2 to rise. In the body this rise in PCO2 would stimulate the respiratory centre causing a period of hyperventilation which would lower the PCO2 to normal. If strong base were added some of the OH- would combine with CO2 to give HCO3-.

OH- + CO2 → HCO3-

The PCO2 would fall. In the body CO2 would be retained to keep the PCO2 from falling.

The control of PCO2 level necessitates either excretion or retention of CO2 by the lungs. This process greatly diminishes the pH change induced by the non-respiratory acids or bases. In effect carbonic acid is being added or taken away to diminish changes that would have been caused by the non-respiratory base or acid respectively.

Compensation of Non-Respiratory Disturbances. The respiratory system can also produce rapid compensation for changes in pH by altering the level of PaCO2. The change in pH alters respiratory control. This causes the alveolar ventilation to alter such that the PaCO2 moves in a direction to cause the pH to return towards normal, i.e. the PaCO2 moves away from normal (40mmHg) in a direction which returns pH towards normal.

1. By preventing the PaCO2 from changing when other acids or bases are altered, pH changes are much less than would occur if the system were acting as a true buffer system.

2. Respiratory compensation by moving the PaCO2 away from normal in the opposite direction to that which would have occurred in vitro still further reduces pH changes.

3 RENAL CONTROL

Normal. The renal system controls the volume and composition of extra-cellular fluid (E.C.F.). It manipulates the E.C.F. electrolytes to maintain the pH at 7.4. In contrast to the rates of change in PCO2 which the respiratory system can produce (minutes) when compensating for pH changes, the renal compensation is slow (days).

Excess Acid or Base (Non-respiratory). If excess acid other than carbonic, or base is added to the internal environemtn, the kidneys excrete them, thus restoring the composition and pH of extracellular fluid to normal. Until the kidney clears the blood of abnormal constituents the pH (assuming the PCO2 is normal) will remain abnormal.

Deficiency of Acid or Base (Non-respiratory). If a disturbance is due to loss of acid other than carbonic the kidneys are unable to restore normality unless the acid deficiency is restored, e.g. in the alkalosis of pyloric obstruction correction depends on replacement of lost HCl. Correction of an acidosis due to loss of base (Na+ or K+ + HCO3- ), e.g. diarrhoea, requires the administration of Na and/or K salts from which Na+ or K++ HCO3- can be formed.

NOTE: The kidney can correct states of excess but not states of deficiency.

Changes in CO2. If the pH is low because of a high PaCO2 (acute respiratory acidosis) the kidneys raise the blood pH towards normal by excreting acid H+ + Cl-, NH4+ + Cl- or 2Na++ Cl- + H2PO4-. Since the urine has a lower pH than the blood entering the kidney, the renal venous blood must have a higher non-respiratory pH than the renal arterial blood. The renal venous blood then mixes into the systemic circulation and raises the pH of the systemic blood towards normal. In the body, the H2CO2 supply is practically inexhaustible, therefore the following equation moves to the right as the [H+] concentration is lowered.

H2CO3 → H+ + HCO3-

Therefore the [HCO3-] in blood rises.

A high pH due to low PaCO2 is probably compensated by renal excretion of base, i.e. NaHCO2 or KHC03.

Limitation of Acidity of Urine. Limitation of acidity of urine. The kidneys cannot produce a urine pH of much less than 4.4. Strong acids can be removed from the blood and excreted in the urine by:

(a) Reacting with the basic salt of phosphoric acid in the urine without producing a great fall in pH of urine;

H+ + HP042- → H2P04- or more fully;

H+ + anion- + 2Na+ + HP042- → H2P04- + 2Na+ + anion-

(b) A more slowly developing process is the addition of NH3 (a base) to the urine.

NH3 + H+ → NH4-

Therefore pH of the urine falls less when a given quantity of acid is removed from the blood and added to the urine. NH3 is generated from glutamine and amino acids leaving organic acids which are metabolised to CO2 and water, therefore the generation of NH3 does not result in any permanent change in blood pH because the acids produced are easily metabolised and excreted as CO2 by the lung.

COMPENSATION AND CORRECTION

Compensation. When a respiratory (PCO2) change occurs in response to a non-respiratory pH disturbance and vice versa, the resulting secondary disturbance in pH is referred to as a compensatory change in response to the primary disturbance.

Compensation of pH disturbances does not completely restore the pH to normal. The pH abnormality which remains after compensation is in the same direction as the primary disturbance unless some complicating factor has intervened.

Some authors refer to the compensatory response as an acidosis or alkalosis whereas others refer only to the primary disturbance as an acidosis or alkalosis. Inserting the terms "primary" or "compensatory" before acidosis or alkalosis should eliminate the communication problem. A primary acidosis (or alkalosis) can be chemically indistinguishable from a compensatory acidosis (or alkalosis). Compensation is not a normal state.

Correction of a acid-base disturbances implies reversal of the chemical cause of the disturbance (e.g. hydrochloric acid, sulphuric acid, lactic acid, CO2). This is necessary before full correction of the disturbance can be said to have occurred.