ACID-BASE BALANCE IN NORMAL AND PATHOLOGICAL CONDITIONS
An important property of blood is its degree of acidity or alkalinity. Body acidity increases when the level of acidic compounds in the body rises (through increased intake or production, or decreased elimination) or when the level of basic (alkaline) compounds in the body falls (through decreased intake or production, or increased elimination). Body alkalinity increases with the reverse of these processes. The body's balance between acidity and alkalinity is referred to as acid-base balance.
The blood's acid-base balance is precisely controlled, because even a minor deviation from the normal range can severely affect many organs. The body uses different mechanisms to control the blood's acid-base balance.
One mechanism the body uses to control blood pH involves the release of carbon dioxide from the lungs. Carbon dioxide, which is mildly acidic, is a waste product of the metabolism of oxygen (which all cells need) and, as such, is constantly produced by cells. As with all waste products, carbon dioxide gets excreted into the blood. The blood carries carbon dioxide to the lungs, where it is exhaled. As carbon dioxide accumulates in the blood, the pH of the blood decreases. The brain regulates the amount of carbon dioxide that is exhaled by controlling the speed and depth of breathing. The amount of carbon dioxide exhaled, and consequently the pH of the blood, increases as breathing becomes faster and deeper. By adjusting the speed and depth of breathing, the brain and lungs are able to regulate the blood pH minute by minute.
The kidneys are also able to affect blood pH by excreting excess acids or bases. The kidneys have some ability to alter the amount of acid or base that is excreted, but because the kidneys make these adjustments more slowly than the lungs do, this compensation generally takes several days.
Yet another mechanism for controlling blood pH involves the use of buffer systems, which guard against sudden shifts in acidity and alkalinity. The pH buffer systems are combinations of a weak acid and weak base that exist in balance under normal pH conditions. The pH buffer systems work chemically to minimize changes in the pH of a solution by adjusting the proportion of acid and base. The most important pH buffer system in the blood involves carbonic acid (a weak acid formed from the carbon dioxide dissolved in blood) and bicarbonate ions (the corresponding weak base).
The hydrogen ion concentration of ECF is normally maintained within very close limits. To achieve this, each day the body must dispose of:
1. About 20 000 mmol of CO2 generated by tissue metabolism. CO2 itself is not an acid, but combines with water to form the weak acid, carbonic acid;
2. About 40-80 mmol of non-volatile acids, mainly sulphur-containing organic acids, which are excreted by the kidneys.
Transport of carbon dioxide
The CO2 produced in tissue cells diffuses freely down a concentration gradient across the cell membrane into the ECF and red cells. This gradient is maintained because red blood cell metabolism is anaerobic, so that no CO2 is produced there, and the concentration remains low. The following reactions then occur:
CO2 + H2O « H2CO3 (3.1)
H2CO3 « H+ + HCO3- (3.2)
Reaction 3.1, the hydration of CO2 to form carbonic acid (H2CO3), is slow, except in the presence of the catalyst carbonate dehydratase (also known as carbonic anhydrase). This limits its site in the blood mainly to erythrocytes, where carbonate dehydratase is located. Reaction 3.2, the ionisation of carbonic acid, then occurs rapidly and spontaneously. As a result, erythrocytes are the principal site of H+ and HCO3- formation in the blood. The H+ ions are mainly buffered inside the red cell by haemoglobin (Hb). Hb is a more effective buffer when deoxygenated, so its buffering capacity increases as it passes through the capillary beds and gives up oxygen to the tissues. Bicarbonate ions, meanwhile, pass from the erythrocytes down their concentration gradient into plasma, in exchange for chloride ions to maintain electrical neutrality.
In the lungs, the PCO2, in the alveoli is maintained at a low level by ventilation. The PCO2 in the blood of the pulmonary capillaries is therefore higher than the PC02 in the alveoli, so the PC02 gradient is reversed. CO2 then diffuses into the alveoli down its concentration gradient, and is excreted by the lungs. The above reaction sequence shifts to the left, carbonate dehydratase again catalysing reaction 3.1, but this time in the reverse direction.
Renal mechanisms for HCO3- reabsorption and H excretion
Glomerular filtrate contains the same concentration of HCO3- as plasma. At normal HCO3-, renal tubular mechanisms are responsible for reabsorbing virtually all this HCO3-. If this fails to occur, large amounts of HCO3- would be lost in the urine, resulting in an acidosis and reduction in the body's buffering capacity. In addition, the renal tubules are responsible for excreting 40-80 mmol of acid per day under normal circumstances. This will increase when there is an acidosis.
The mechanism of reabsorption of HCO3- is shown in Figure 3.1. HCO3- is not able to cross the luminal membrane of the renal tubular cells. H+ is pumped from the tubular cell into the lumen, in exchange for Na+. The H+ combines with HCO3- to form H2CO3 in the lumen. This dissociates to give water and CO2, which readily diffuses into the cell. In the cell, CO2 recombines with water under the influence of carbonate dehydratase to give H2CO3. This dissociates to H+ and HCO3- .
The HCO3- then passes across the basal membrane of the cell into the interstitial fluid. This mechanism results in the reabsorption of filtered HCO3- , but no net excretion of H+
The net excretion of H+ relies on the same renal tubular cell reactions as HCO3 reabsorption, but occurs after luminal HCO3- has been reabsorbed, and depends on the presence of other suitable buffers in the urine (Figure 3.2). The main urinary buffer is phosphate, most of which is present as HPO42, which can combine with H+ to form H2PO4-. Ammonia can also act as a urinary buffer, and is formed by the deamination of glutamine in renal tubular cells under the influence of the enzyme glutaminase. Ammonia readily diffuses across the cell membrane into the tubular lumen, where it combines with H+ to form NH4+. This does not pass across cell membranes, so passive reabsorption is prevented. Glutaminase is induced in chronic acidoses, stimulating increased ammonia production and therefore increased H+ excretion in the form of NH4+ ions.
Fig. 3.1 Reabsorption of bicarbonate in the renal tubule
Buffering of hydrogen ions
The lungs and the kidneys together maintain overall acid-base balance. However, the ECF needs to be protected against rapid changes in [H+]. This is achieved by various buffer systems. A buffer system consists of a weak (incompletely dissociated) acid in equilibrium with its conjugate base and H+. The capacity of a buffer for H+ is related to its concentration and the position of its equilibrium, being most effective at the [H+] at which the acid and conjugate base are present in
equal concentrations. Thus, Hb and plasma proteins act as efficient buffers in blood, since they are abundant and at a physiological [H+] of approximately 40 nmol/L have side groups that exist in an appropriate equilibrium. At this [H+], the bicarbonate buffer system has an equilibrium that is far removed from the ideal, with [HCO3-] being about 20 times greater than [H2CO3]. However, the effectiveness of the bicarbonate system is greatly enhanced in vivo by the fact that H2CO3 is readily produced or disposed of by interconversion with CO2. Furthermore, physiological control mechanisms act on this buffer system to maintain both PC02 and [HCO3- ] within limits, and hence to control [H+].
Any physiological buffer system could be used to investigate and define acid-base status, but the H2CO3/HCO3- buffer system has proved to be the most appropriate for this purpose, due to its physiological importance.
Investigating acid-base balance
The acid-base status of a patient can be fully characterised by measuring [H+] and PC02 in arterial or arterialised capillary blood specimens; [HCO3-] is then obtained by calculation:
[H+] = 180 × PCO2 / [HCO3-], so [HCO3-] = 180 × PCO2 / [H+]
Although standard bicarbonate, base excess and base deficit are still sometimes reported, these derived values are not necessary for the understanding of acid-base disturbances.
Collection and transport of specimens
Arterial blood specimens are the most appropriate for assessing acid-base status. However, unless an arterial cannula is in situ, these specimens may be difficult to obtain for repeated assessment of patients whose clinical condition is changing rapidly. Arterialised capillary blood specimens are also widely used, especially in infants and children. It is essential for the capillary blood to flow freely, and collection of satisfactory samples may be impossible if there is peripheral vasoconstriction or the blood flow is sluggish.
Patients must be relaxed, and their breathing pattern should have settled after any temporary disturbance (e.g. due to insertion of an arterial cannula), before specimens are collected. Some patients may hyperventilate temporarily because they are apprehensive.
Blood is collected in syringes or capillary tubes that contain sufficient heparin to act as an anticoagulant; excess heparin, which is acidic, must be avoided. If ionised Ca2+ is to be measured on the same specimen, as is possible with some instruments, calcium-balanced heparin must be used. Specimens must be free of air bubbles, since these will equilibrate with the sample causing a rise in PO2 and a fall in PCO2.
Acid-base measurements should be performed immediately after the sample has been obtained, or the specimen should be chilled until analysis.
Otherwise, glycolysis (with the production or lactic acid) occurs, and the acid-base composition of the blood alters rapidly. Specimens chilled in iced water can have their analysis delayed for as long as 4 h. However, the clinical reasons that gave rise to the need for full acid-base studies usually demand much more rapid answers.
Acid-base measurements are nearly always made at 37°C, but some patients may have body temperatures that are higher or lower than 37°C. Equations are available to relate [H+], PCO2 and PO2 , determined at 37°C, to 'equivalent' values that correspond to the patient's body temperature. However, reference ranges for acid-base data have only been established by most laboratories for measurements made at 37°C. Adjustment of analytical results to values that would have been obtained at the patient's temperature, according to these equations, may therefore be difficult to interpret. If treatment aimed at reducing an acid-base disturbance (e.g. NaHCO3 infusion) is given to a severely hypothermic patient, the effects of the treatment should be monitored frequently by repeating the acid-base measurements (at 37°C).
Disturbances of acid-base status
Acid-base disorders fall into two main categories, respiratory and metabolic.
1 Respiratory disorders A primary defect in ventilation affects the PCO2.
2 Metabolic disorders The primary defect may be the production of non-volatile acids, or ingestion of substances that give rise to them, in excess of the kidney's ability to excrete these substances. Alternatively, the primary defect may be the loss of H+ from the body, or it may be the loss or retention of HCO3-.
Acidosis is excessive blood acidity caused by an overabundance of acid in the blood or a loss of bicarbonate from the blood (metabolic acidosis), or by a buildup of carbon dioxide in the blood that results from poor lung function or slow breathing (respiratory acidosis).
Alkalosis is excessive blood alkalinity caused by an overabundance of bicarbonate in the blood or a loss of acid from the blood (metabolic alkalosis), or by a low level of carbon dioxide in the blood that results from rapid or deep breathing (respiratory alkalosis).
This is caused by CO2 retention due to hypoventilation (Table 3.2). It may accompany defects in the control of ventilation, or diseases affecting the nerve supply or muscles of the chest wall or diaphragm, or disorders affecting the ribcage or intrinsic lung disease.
In acute respiratory acidosis, a rise in PCO2 causes the equilibria in reactions 3.1 and 3.2 to shift to the right, as a result of which plasma [H+] and [HCO3-] both increase. Equilibration of H+ with body buffer systems limits the potential rise in [H+], and a new steady state is achieved within a few minutes.
Unless the cause of the acute episode of acidosis is resolved, or is treated quickly and successfully, renal compensation causes HCO3- retention and H+ excretion, thereby returning plasma [H+] towards normal while [HCO3-] increases. These compensatory changes can occur over a period of hours to days, by which time a new steady state is achieved and the daily renal H+ excretion and HCO3- retention return to normal. The patient then has the pattern of acid-base abnormalities of chronic respiratory acidosis.
This is due to hyperventilation (Table 3.3). The reduced PC02 that results causes the equilibrium positions of reactions 3.1 and 3.2 to move to the left. As a result, plasma [H+] and [HCO3-] both fall, although the relative change in [HCO3-] is small.
If conditions giving rise to a low P PC02 persist for more than a few hours, the kidneys increase HCO3- excretion and reduce H+ excretion. Plasma [H+] returns towards normal, whereas plasma [HCO3-] falls further. A new steady state will be achieved in hours to days, if the respiratory disorder persists. It is unusual for chronic respiratory alkalosis to be severe, and plasma [HCO3-] rarely falls below 12 mmol/L.
Table 3.3 Respiratory alkalosis
Increased production or decreased excretion of H+ leads to accumulation of H+ within the ECF (Table 3.4). The extra H+ ions combine with HCO3- to form H2CO3, disturbing the equilibrium in reaction 3.2, with a shift to the left. However, since there is no ventilatory abnormality, any increase in plasma [H2CO3] is only transient, as the related slight increase in dissolved CO2 is immediately excreted by the lungs. The net effect is that a new equilibrium rapidly establishes itself in which the product, [H+] x [HCO3-], remains unchanged, since [H2CO3] is unchanged. In consequence, the rise in plasma [H+] is limited, but at the expense of a fall in [HCO3-], which has been consumed in this process and may be very low. Its availability for further buffering becomes progressively more limited. Less often, metabolic acidosis arises from loss of HCO3- from the renal or GI tracts. Typically in these conditions, HCO3- does not fall to such a great extent, rarely being less than 15 mmol/L.
Table 3.4 Metabolic acidosis.
The rise in ECF [H+] stimulates the respiratory centre, causing compensatory hyperventilation. As a result, due to the fall in PCO2 , plasma [H+] returns towards normal, while plasma [HCO3-] falls even further. Plasma [H+] will not, however, become completely normal through this mechanism, since it is the low [H+] that drives the compensatory hyperventilation - as the [H+] falls, the hyperventilation becomes correspondingly reduced. In addition, if renal function is normal, H+ will be excreted by the kidney. It is quite common for patients with metabolic acidosis to have very low plasma [HCO3-], often below 10 mmol/L.
This is most often due to prolonged vomiting, but may be due to other causes (Table 3.5). The loss of H+ upsets the equilibrium in reaction 3.2, causing it to shift to the right as H2CO3 dissociates to form H+ (which is being lost) and HCO3-. However, because there is no primary disturbance of ventilation, plasma PCO2 remains constant, with the net effect that plasma [H+] falls and [HCO3-] rises. Respiratory compensation (i.e. hypoventilation) for the alkalosis is usually minimal, since any resulting rise in PCO2 or fall in PO2 will be a potent stimulator of ventilation. HCO3- is freely filtered at the glomerulus, and is therefore available for excretion in the urine, which would rapidly tend to restore the acid-base status towards normal. The continuing presence of an alkalosis means there is inappropriate reabsorption of filtered HCO3- from the distal nephron. This can be due to ECF volume depletion, potassium deficiency or mineralocorticoid excess.
Table 3.5 Metabolic alkalosis.
Other investigations in acid-base assessment
The full characterisation of acid-base status requires arterial or arterialised capillary blood samples, since venous blood PCO2 (even if 'arterialised') bears no constant relationship to alveolar PCO2. However, other investigations can provide some useful information.
Total CO2 (reference range 24-30 mmol/L)
This test, performed on venous plasma or serum, includes contributions from HCO3-H2CO3, dissolved CO2 and carbamino compounds. However, about 95 % of 'total CO2' is contributed by HCO3-. Total CO2 measurements have the advantages of ease of sample collection and suitability for measurement in large numbers, but they cannot define a patient's acid-base status, since plasma [H+] and PCO2 are both unknown. For example, an increased plasma [total CO2] may be due to either a respiratory acidosis or a metabolic alkalosis. However, when interpreted in the light of clinical findings, plasma [total CO2] can often give an adequate assessment of whether an acid-base disturbance is present and, if one is present, provide an indication of its severity. This is particularly true when there is a metabolic disturbance. However, patients with respiratory disturbances are much more likely to require full assessment of acid-base status, both for their definition and for monitoring and controlling their treatment.
Anion gap (reference range 10-20 mmol/L)
The anion gap (or ion difference) is obtained from plasma electrolyte results, as follows:
AG = ([Na+] + [K+]) - ([Cl] + [total CO2)
The difference between the cations and the anions represents the unmeasured anions or anion gap and includes proteins, phosphate, sulphate and lactate ions. The anion gap may be increased because of an increase in unmeasured anions.
This may be of help in narrowing the differential diagnosis in a patient with metabolic acidosis (table 3.6). In the presence of metabolic acidosis, raised anion gap points to the cause being exessive production of hydrogen ions or failure to
excrete them. As the acid accumulates in the ECF (e.g. in DKA), the HCO3- is titrated and replaced with unmeasured anions (e.g. acetoacetate) and the anion gap increases. In contrast, if the cause is a loss of HCO3- (e.g. renal tubular acidosis), there is a compensatory increase in Cl- and the anion gap remains unchanged (Table 3.4)
Plasma chloride (reference range 95-107 mmol/L)
The causes of metabolic acidosis are sometimes divided into those with an increased anion gap and those with a normal anion gap. In the latter group, the fall in plasma [total CO2], which accompanies the metabolic acidosis, is associated with an approximately equal rise in plasma [Cl-]. Patients with metabolic acidosis and a normal anion gap are sometimes described as having hyperchloraemic acidosis.
Increased plasma [Cl ], out of proportion to any accompanying increase in plasma [Na+], may occur in patients with chronic renal failure, ureteric transplants into the colon, renal tubular acidosis, or in patients treated with carbonate dehydratase inhibitors. Increased plasma [Cl-] may also occur in patients who develop respiratory alkalosis as a result of prolonged assisted ventilation. An iatrogenic cause of increased plasma [Cl-] is the IV administration of excessive amounts of isotonic or 'physiological' saline, which contains 155 mmol/L NaCl.
Patients who lose large volumes of gastric secretion (e.g. due to pyloric stenosis) often show a disproportionately marked fall in plasma [Cl-] compared with any hyponatraemia that may develop. They develop metabolic alkalosis, and are often dehydrated.
Oxygen delivery to tissues depends on the combination of their blood supply and the arterial O2 content. In turn the O2 content depends on the concentration of Hb and its saturation. Tissue hypoxia can therefore be caused not just by hypoxaemia, but also by impaired perfusion (e.g. because of reduced cardiac output or vasoconstric-tion), anaemia and the presence of abnormal Hb species. The full characterisation of the oxygen composition of a blood sample requires measurement of PO2, Hb concentration and percentage oxygen saturation. Hb measurements are widely available, and PO2 is one of the measurements automatically performed by most blood gas analysers as part of the full acid-base assessment of patients, and Hb saturation is measured using an oximeter.
Measurements of PO2 in arterial blood (reference range 12-15 kPa) are important, and are often valuable in assessing the efficiency of oxygen therapy, when high PO2 values may be found. Above a PO2 of 10.5 kPa, however, Hb is almost fully saturated with O2 (Figure 3.4), and further increases in PO2 do not result in greater O2 carriage. Conversely, as PO2 drops, initially there is little reduction in O2 carriage on Hb, but when it falls below about 8 kPa, saturation starts to fall rapidly. In addition, results of PO2 measurements may be misleading in conditions where the oxygen-carrying capacity of blood is grossly impaired, as in severe anaemia, carbon monoxide poisoning and when abnormal Hb derivatives (e.g. methaemoglobin) are present. Measurement of both the blood [Hb] and the percentage oxygen saturation are required in addition to PO2 under these circumstances.
Figure 3.4 The oxygen dissociation curve of Hb. It is important to note that, above a PO2 of approximately 9 kPa, Hb is over 95 % saturated with O2. Also shown in the figure is the value of the PO2 3.8 kPa, that corresponds to 50 % saturation with O2; this value is called the P50 value.
Indications for full blood acid-base and oxygen measurements
The main indications for full acid-base assessment, coupled with PO2 or oxygen saturation measurements, are in the investigation and management of patients with pulmonary disorders, severely ill patients in intensive care units and patients in the operative and peri-operative periods of major surgery who may often be on assisted ventilation. Other important applications include the investigation and management of patients with vascular abnormalities involving the shunting of blood.
Full acid-base assessment is less essential in patients with metabolic acidosis or alkalosis, for whom measurements of plasma [total CO2] on venous blood may give sufficient information.