6.1 INTRODUCTION
Clinical problems of pH are all related to pH of the plasma of whole blood. pH in extracellular fluid is always close to that of blood. pH inside cells differs from that of blood but it is not recognised as being an important clinical problem apart from blood pH changes.
In the clinical situation if the actual pH of the blood is lowered, one can usually assume that the primary disturbance has been the addition to the blood of acid or the removal of base and vice versa.
6.2 CHEMICAL CLASSIFICATION OF CAUSES OF CHANGES IN BLOOD pH
Blood pH may be changed if acids or bases are added to or removed from the blood. Secretion of an acid (e.g. gastric juice) implies that the acid involved (HCl in this case) is removed from the blood.
The acids which can cause changes in blood pH are:
A. Metabolic or non-respiratory acids :-
1. Organic
1.a. Lactic Acid
1.b. Keto Acids
1.a.i. Hypoxia
1.a.ii. Drug Induced
1.a.iii. Idiopathic
1.b.i. Diabetes
1.b.ii. Starvation
1.c.Might be called "true" metabolic acidoses
2.a. Inorganic
2.a.i. Sulphuric Acid
2.a.ii. Phosphoric Acid
2.a.iii. Hydrochloric Acid
2.c. Inorganic acids, increase in renal failure
B. Respiratory acids >> Organic only>> Carbonic Acid.
A rise in concentration of any of these acids in the blood causes a fall in the pH of the blood. Loss of acid from the blood (e.g. into gastric juice) causes a rise in the pH. Only HCl and H2CO3 can be lost from the blood in appreciable quantities.
The bases which can cause changes in blood pH are:
•NaHC03
•KHC03
Administration of base by mouth or parenterally may cause blood pH to rise if rate of excretion does not match rate of administration. Loss of alkaline fluid from bowel (diarrhoea, intestinal obstruction or intestinal fistulae), or urine (after acetoazolamide) will cause blood pH to fall.
6.3 CLINICAL CLASSIFICATION OF CAUSES OF CHANGES IN BLOOD pH
Clinical states of pH disturbence (acid-base inbalance) can conveniently be divided into two groups, i.e. (a)respiratory and (b)metabolic or non-respiratory. The reasons for this division into respiratory and non-respiratory are that:
i) the compensatory mechanisms (Section 3.5.1) and treatments (Section 7) of the two types are different.;
ii) the recognition of non-respiratory disturbances is masked by compensatory alterations in PCO2 and the recognition of changes in pH caused by PCO2 changes are masked by renal compensation.
6.3.1 RESPIRATORY ACIDOSIS. This is synonymous with CO2 retention and is usually a sign of hypoventilation. Compensation is renal. There is renal loss HCl in the form of buffer or as NH4Cl. During recvovery chloride has to be supplied and retained.
6.3.1.1 Causes of hypoventilation:
•Central nervous system
•Peripheral nervous system
•Neuromuscular transmission
•Muscle disorders
Chest wall abnormalities
Lung and airway disorders.
6.3.1.2 Inhalational of CO2 This is another cause of respiratory acidosis, but it is only likely to occur under situations of re-breathing, e.g. under anaesthesia or during resuscitation with a Water's cannister circuit without the cannister, i.e. ward resuscitators or Type C anaesthetic systems. (Mapleson, 1954).
6.3.1.3 Increased production of CO2. This very rarely causes a high PaCO2. In thyrotoxicosis and fever CO2 production is raised but the rise is well within the capacity of a normal respiratory system. In malignant hyperpyrexia high PaCO2s have been recorded due undoubtedly to increased production.
NaHCO2 therapy in non-respiratory acidosis causes a rise in PaCO2 until the CO2 generated by the neutralization of HCO3- has been excreted (Singer et al, 1956). This effect may be prolonged if CO2 excretion is impaired (Ostrea and Odell, 1972).
Acute respiratory acidosis is diagnosed by high PaCO2 and low actual pH associated with near normal non-respiratory pH, (standard HCO3-and base excess). (See Appendix 4.2).
Chronic respiratory acidosis is compensated by loss of H+Cl- by the kidney and generation (not reabsorption) of (Na+) HCO3-. (See Appendix 3.4). Compensation causes slightly lowered actual pH with a high non-respiratory pH (positive base excess and high standard HCO3-). Steroids and diuretics (except acetazolamide (Diamox) ) may produce a non-respiratory alkalosis which may cause a respiratory acidosis to appear to be "over compensated". There will then be a high (actual) pH.
In chronic respiratory failure ventilation is maintained by hypoxic drive and also by the acid pH of the blood. If the pH stimulus is removed by the actual pH being raised, some of the respiratory drive will be removed causing a diminution of respiratory effort and further CO2 retention.
When a chronic respiratory acidosis (with compensation, i.e. high non-respiratory pH, Base excess, and Standard HCO3-.) is improved by increasing alveolar ventilation the actual pH may rise above 7.4. The resulting chemical picture of high pH, high non-respiratory pH (positive base excess), and high PaCO2 is indistinguishable chemically from a primary metabolic alkalosis with compensatory CO2 retention. Clinical information is required to distinguish them. (Plot this and other examples on a standard Siggaard-Andersen nomogram as an exercise). ( See Footnote to table 4.2.2.)
N.B. The compensatory renal or respiratory changes which partially correct pH disturbances are chemically true pH changes in the opposite direction to the original disturbances. It is a semantic problem whether the compensatory renal change induced for example by primary chronic CO2 retention is called a metabolic alkalosis or not. This semantic muddle can be avoided if the term alkalosis is used only in a non-precise term.
6.3.2. RESPIRATORY ALKALOSIS. This is associated with hyperventilation. Usually these are acute so there is no time for renal compensation, but if prolonged, such as in acclimatization to high altitudes, there would probably be renal compensation.
6.3.2.1. Deliberate induced hyperventilation during anaesthesia.
6.3.2.2. Some causes of hypoxia associated with hyperventilation. This would include "air hunger" of hypovolaemia, severe ventilation-perfusion abnormalities and acclimatization to high altitudes. Critically ill patients without arterial hypoxaemia can have severe hypocapnia (Mazzara et al,1974). This may be explained by low cardiac output.
In asthma attacks, unless severe, the PaCO2 is usually lowered, i.e. the alveolar ventilation is increased (McFadden and Lyons, 1968). This is at least sometimes associated with hypoxaemia due to mismatching of ventilation and perfusion.
6.3.2.3. Fever. This may cause hypocapnia due presumably to central stimulation of the respiratory control mechanisms (Chapot et al, 1974).
6.3.2.4. Some types of C.N.S. damage.
6.3.2.5. Hysterical hyperventilation.
6.3.3. METABOLIC ACIDOSIS.
This (non-respiratory acidosis) is due to increase in acids (i.e. H+ donating substances) other than H2CO2 or decrease in base (i.e. H+ acceptors) in the blood. Compensation is by hyperventilation. This lowers the PaCO2 thus deducing the any pH change. The causes of non-respiratory acidosis are:
Increased alimentary or parenteral intake of acid or alimentary loss of base. (Addition and subtraction).
Increased production of acid. (Accumulation).
Failure of excretion of acid or loss of base by the renal system.
Increased Intake of Acid or Loss of Base
6.3.3.1 Increased alimentary or parenteral intake of acid or alimentary loss of base.
6.3.3.1.1. Adding acid. The acid content of blood may be raised by ingestion or injection of ammonium chloride or dilute hydrochloric acid. The hydrochloric acid directly increases [H+]. The ammonium chloride produces hydrochloric acid by the NH3 being split off and converted to urea. Adding ammonium chloride directly to blood would not change the pH greatly until the NH4+ has been metabolised.
6.3.3.1.2 Alimentary loss if Base. Loss of intestinal contents by diarrhoea, low small bowel obstruction or intestinal fistulae causes loss of fluid of high pH, i.e. containing an excess of base (Na+HCO3- or K+HCO3-). This results in the fluid left in the body having a lower base content than normal. The removal of base causes the blood pH to fall (Leading Article, 1966). A similar disturbance has been reported from loss of lymph. (Siegler et al, 1978).
6.3.3.1.3 Intravenous infusions. These can cause an acidosis.
6.3.3.1.3.1 Stored blood for transfusion. This has la ow pH. The anticoagulant contains citric acid. When mixed with the blood the pH drops. PCO2 rises because of the action of acid on bicarbonate ions. Most of the pH drop is due to this CO2 which does not escape from the stored blood. The non-respiratory pH of the stored blood is not as low as the actual pH. The pH of stored blood does not fall progressively if stored for up to three weeks at 4�C (Gaudry, Joseph and Duffy, 1974).
The acidic salt of sodium citrate and "dilutional" acidosis are the main contributers to the non-respiratory pH of stored blood. There is a variable amount of lactic acid acid in stored blood This usually contributes less to the non-respiratory pH change. If stored concentrated red cells are washed there will be a considerable and even an increased non-respiratory acidosis in the product. If 27meq NaHCO2 is added to each litre of washing fluid this would not occur, but would introduce other problems.
Stored blood transfusion rarely causes a non-respiratory acidosis if the circulation and temperature are maintained normal. It is possible that if the circulation and temperature are not maintained, that metabolic acidosis could occur with massive transfusion. In a situation where a massive transfusion is given it will not usually be possible to distinguish the low pH due to blood, from that due to the condition for which the blood is being given. From personal observations the low base excess observed during liver transplantation appears to be mainly a combination of lactic acidosis from the transfused blood and released or generated in the new liver and dilutional acidosis. The sodium citrate in the anti-coagulant solution can cause a metabolic alkalosis after the citrate has been metabolised and replaced by bicarbonate ion.
Howlands et al (1965) have advocated routine use of sodium bicarbonate during massive blood transfusion to counteract the acidosis of the infused blood. This is usually unnecessary (Bookallil and Joseph, 1968), and will probably only aggravate post transfusion alkalosis (Miller et al, 1971).
If the temperature and circulation are not maintained as is common in when more than say 5 litres in 2 hours are transfused in trauma the composition of the circulating blood approaches that of the transfused blood. In liver transplantation 3 times the blood volume may be replaced in an hour and up to 30 times the blood volume during the operation. This can result in pH<6 .9="" acidosis.="" alkali="" and="" appars="" base="" be="" cause="" circulatory="" dilutional="" effects="" ersonal="" excess="" font="" is="" lactic="" litre.="" little="" meq="" observation="" occurs="" recovery="" spontaneous="" the="" therapy="" there="" to="" without="">
6.3.3.1.3.2 "Dilution Acidosis" (Shires et al, 1948; Garella et al, 1973) is an acidosis due to the intravenous infusion of a neutral non-buffer solution.
Intravenous solutions which contain only non-acids or non-bases, e.g. sodium chloride, glucose and other carbohydrates usually have a pH slightly less than 7. This is usually due to pharmaceutical details in the preparation. In these non-buffer solutions this low pH represents a very small amount of acid. One would only have to add a fraction of a miliequivalent of strong base to a litre to make the solution alkaline.
If such a non-buffer solution (e.g. Saline or 5% glucose) is equilibrated with a gas mixture containing 40mmHg CO2 it has a pH of 4.9 (Gaudry et al, 1972). 24meq Na0H (or NaHC03) would have to be added to it to give a pH of 7.4, therefore a neutral solution plus 24meq NaHC03 will on infusion cause no change in pH in the blood. The original solution of saline or glucose will act in a similar fashion to an injection of 24meq HCl for each litre infused. This will only be importasnt if large volumes of intravenous fluid are given in a short time,e.g. in burns or cholera (see Diuretic Alkalosis, 6.3.4.2).
6.3.3.1.3.3 Metabolic acidosis associated with intravenous alimentation. Intravenous feeding with amino acid solution in the form of hydrochlorides can cause a metabolic acidosis associated with a high serum Cl level (Chan, Ghadimi, Kaminski and Heird et al, 1973). Lactic acidosis may be associated with intravenous feeding regimes which contain fructose, xylotol or sorbitol (Alberti and Nattrass, 1977).
In some cases hypertonic glucose also has been associated with lactic acidosis (Ames et al, 1975).
6.3.3.4 Ingestion or injection of oganic acids. This would not ordinarily produce change in pH because the liver would metabolise them. In liver disease and during liver transplantation organic acids introduced to the body gain access to the systemic circulation.
6.3.3.5 Accumulation of Acid. Excess acid may accumulate in the blood from processes of metabolism and cause a fall in the pH. There are two main mechanisms for this:
6.3.3.5.1 Hypoxia. Hypoxia from any cause prevents the products of anaerobic glycolysis being further metabolised. Excess lactic acid appears in the blood and lowers the pH.
The causes of hypoxia are:
Low oxygen inspired air.
Lung disorders.
Hypoventilation.
Low cardiac output (including shock states and myocardial infarction (Neaverson, 1966, Kirby et al, 1966).
Blood defect; hypovolaemia, anaemia or CO poisoning.
Tissue toxins, e.g. cyanide. Cyanide is important in sodium nitroprusside toxicity. This may be important in combined respiratory and circulatory failure if sodium nitroprusside is used to lower peripheral resistance.
In cardiac arrest there have been several studies which document low blood pH (Ledingham et al, 1962; Chazan et al, 1968 and Fillmore et al, 1970). All cardiac arrests do not have non-respiratory acidosis. Respiratory acidosis (high PaCO2) is a feature of some. High PaCO2 for practical purposes means inadequate alveolar ventilation, which of course, might be present. Fillmore et al. imply that aspiration may be a cause of high PaCO2. This is only the case if it associated with hypoventilation. If ventilation is inadequate not only will the PaCO2 be high but what is more important the oxygenation will be inadequate unless oxygen is given. High levels of alveolar oxygen will not be present if mouth-to-mouth ventilation is being used. A transient rise in PaCO2 will occur for a few minutes after NaHC03 administration, due to increased CO2 release (Bishop et al, 1976).
Circulatory occlusion of any large area causes accumulation of organic acids in the area supplied. On restoration of the circulation these acids are distributed systemically. This is a possible cause of acidosis, but in practice if the general circulation and temperature are maintained in a satisfactory state only transient acidosis results as the acids are quickly metabolised (Bookallil and Joseph, 1968).
Non-hypoxic lactic acidosis is a syndrome being more commonly discussed. It is a serious syndrome of unknown aetiology (Oliva: Leading Article, 1970; Leading Article, 1973; Harken, 1976: Alberti et al, 1977). It probably includes syndromes of varying aetiology, e.g. phenformin associated acidosis in diabetes (Gale et al, 1976), and the acidosis associated with parenteral nutrition with some carbohydrates (See 6.3.3.1.3.3 and Appendix 6.3.3.2.1).
6.3.3.5.2 Diabetes and Starvation. In both these states excess acetone and keto-acids are produced. There will be some attempt at renal excretion (ketonuria) but if this is inadequate the quantity of acid in the blood will rise and the pH wil fall (Peters and Van Slyke, 1931). All acidoses in diabetes are not due to keto-acid. In some instances lactic acid is the cause. Patients clinically diagnosed as diabetic keto-acidosis sometimes have alkalosis (Lim and Walsh, 1976).
6.3.3.5.3 Others. There are some other disturbances of intermediary metabolism which may cause blood pH to fall from accumulation of organic acids. One of these which is mentioned in 6.3.3.1.3.3 occurs from intravenous administration of some sugars.Another is iso-valeric acidaemia - a rare inborn error of metabolism due to an enzyme deficiency (Cohn et al, 1978).
6.3.3.3 Failure of excretion of acid. This may lead to acidosis. Normally during metabolism some inorganic acids are produced, i.e. sulphuric and phosphoric acids. The anions have to be excreted by the kidney covered either by Na+ , K+ , H+ (small amount) or NH4+ (produced in the kidney). Normally the quantity of acid involved is not great (40-60meq/day), but over a long period, if there is failure of excretion, accumulation will occur with a fall in pH. This is renal acidosis (Relman, 1968). Renal acidosis may also be due to loss of base induced by acetazolamide (Diamox). Renal acidosis is usually a manifestation of generalised renal failure but there is also a syndrome of renal tubular acidosis where metabolic acidosis of renal origin is an isolated disorder.
6.3.4 METABOLIC ALKALOSIS (non-respiratory alkalosis). This is due to loss of HCl from the ECF or addition of alkali. Metabolic alkalosis is compensated by respiratory depression which causes CO2 retention (Tuller and Mehdi, 1971; Shear et al, 1973; Aquino et al, 1973) but may also cause hypoxia. The pH is usually raised but may be high normal if there is much CO2 retention. Metabolic alkalosis is due to:
6.3.4.1. Loss of gastric juice containing HCl. Patients with pyloric obstruction lose some K+ and Na+ as well as HCl. The loss of K+ is mainly through the kidney (Kassirer et al, 1966). At first the urine is alkaline but after stable conditions are established the urine becomes acidic as the normal inorganic acid load from protein breakdown still has to be excreted (Schwartz et al, 1978). If it were not, it would correct the alkalosis. The acid urine used to be thought to be paradoxical (Van Slyke and Evans, 1947), and was attributed to K+ deficiency.
The situation of a chronic metabolic alkalosis with acid urine is probably the best clinical example where balance or status as destinct from input, output or turnover (section 2.1) should be distinguished. The blood pH status is stable and alkaline. For the non-respiratory pH to remain high the normal acid output must continue, i.e. acid excretion in the urine will be normal and the urine pH will be low.
6.3.4.2 Diuretic alkalosis. Thiazide diuretics, frusemide and ethacrynic acid can produce a metabolic alkalosis. HCl, its equivalent NH4Cl or HCl having acted on phosphate buffer is lost in the urine. The central role of Cl in the production of diuretic alkalosis has been established (Kassirer et al, 1965).
6.3.4.3 Ingestion or injection of excess base, e.g. Na+HCO3- or Na+OH-. Post transfusional or post-cardiac surgery metabolic alkalosis is usually due to the administration and metabolism of sodium citrate (Kappogoda et al, 1973; Barcenas et al, 1976). If NaHCO2 is given to "correct" an acidosis it may result in a high serum Na and osmolality and later an alkalosis.
6.3.4.4 Steroid alkalosis (Kassirer et al, 1970; Schambelan et al, 1971).
6.4 CLINICAL DIAGNOSIS OF ACID-BASE OF pH DISTURBANCES
Any of the clinical conditions mentioned can be associated with pH or acid-base status disturbances and one can usually suspect pH changes by the clinical picture, e.g. hypoventilating, acute on chronic obstructive airway disease, probably has CO2 accumulation; intestinal obstruction or severe diarrhoea probably has acidosis due to loss of base Na+ and/or K+ + HCO3- ); a diabetic who is drowsy and hyperventilating with urinary glucose and ketones probably has keto-acidosis; "shock" with poor tissue perfusion may have lactic acidosis.
The signs of acid-base disturbances are not diagnostic but in association with the clinical picture, a suspicion of variable certainty can be entertained. Definite daignosis is only possible by measurements on blood which should usually be arterial but can be capillary, venous or mixed venous. See tables 4.2.1 and 4.2.2.
6.5 PHYSIOLOGICAL EFFECTS OF pH DISTURBANCES
It is often stated that blood pH below some arbitrary level, usually 6.8, is incompatible with survival (Andreoli 1988). This is not so. Low pHs' have been recorded in association with CO2 retention (Schultz, 1960, pH 6.71, PaCO2 234mmHg. Prys-Roberts et al, 1967, pH 6.86, PaCO2 248mmHg) and severe exercise (Osnes and Hermansen, 1972, Hermansen and Osnes, 1972) with little physiological disturbance attributed to the low pH, and with complete recovery.
Two cases of strychnine poisoning who recovered after low pHs' were recorded, have been reported, (Goldstein, 1975, pH 6.57 PaCO2 18mmHg; Loughhead et al, 1978, pH 6.59 PaCO2 59.5mmHg).
If the blood pH is abnormal, the cause of the disturbance is usually severe, and compensatory or correcting mechanisms have not occurred or have been insufficient. It is in practice impossible to distinguish the physiololgical effects of the cause from those of the pH change itself.
There are many effects of pH changes which are reversible when the cause and/or the pH change are removed. As far as the whole organism is concerned the cardiovascular effects of pH change are the most important. A low pH is said to aggravate cardiovascular depression. This then further lowers pH due to lowered perfusion, thus setting up a vicious circle.
There is evidence that correcting a low non-respiratory pH in circulatory insufficiency improves the prognosis (Manger, 1962, Ledingham et al, 1962). Low non-respiratory pH can worsen arrhythmias (Anderson, 1968). Artificially induced pH changes due to infusion of acid in normal animals does not reduce cardiovascular activity with moderate falls in pH (Anderson, 1967, 1968, Caress et al, 1968). Maybe a low pH has a much greater effect in an already impaired circulatory system than in the normal. There is controversy (Stackpool,1986, Narins & Cohen 1987) about use of alkali in states of of low pH.
The cardiovascular effects of pH changes are probably different in the various chemical causes, e.g. hypoxic lactic acidosis produces more disturbance than a similar pH change due to CO2 (Schultz, 1960, Prys-Roberts et al, 1967).
Alteration of enzyme action when the pH is abnormal is usually given as the mechanism of the deleterious effects of pH changes. This, of course, is the basic mechanism in all disturbances but it does not indicate in which system the effect is critical. It is cardiovascular effects of pH changes which set the limit.
High pH levels also produce physiological disturbances. These are less well documented and of less clinical importance. The main clinically important effects are CNS effects (tetany and impaired consciousness) hypoventilation (if the high pH is not due to hyperventilation and a low PCO2) and possible cardiovascular effects (Streisand et al, 1971; Lawson et al, 1973; Galmarini et al, 1973; Lock et al, 1975).
The serum potassium concentration falls when the PaCO2 is acutely lowered. Return of the potassium concentration may be slower than the return of the pH when the PCO2 recovers (Edwards et al, 1977; Sanchez et al, 1978; Finsterer et al, 1978).
Appendices
A.6.3.3.2.1. Lactic Acidosis
A review article "Lactic Acidosis" by Alberti and Nattrass (1977) gives a rational and useful classification of lactic acidosis drawing on the book by Cohen and Woods, "Clinical and Biochemical Aspects of Lactic Acidosis" (1976).
This classification is:
Type A Shock syndrome and hypoxia.
Type B1 Common disease states, e.g. diabetes, liver disease, infection.
Type B2 Drugs, e.g. phenoformin, sorbitol, fructose.
Type B3 Hereditary metabolic disorders.
They state that as these categories are better defined, the reporting of so-called iodiopathic lactic acidosis decreases.
It is sensible to treat the cause if this is known. This may be all that is required in Type A. Restored circulation and oxygenation will permit the lactic acid to be metabolised. In Type B more active measures are suggested to correct the pH abnormality as a low pH may cause the liver to produce further lactic acid. Large doses of NaHCO2 (up to 2,500mMol) have been given. This quantity would usually produce problems of its own which would necessitate dialysis. Such situations are very rare or else are unrecognised.
In the preface of Cohen and Woods' Monograph which contains an exhaustive literature review, it is stated, "We strongly suspect that the comparatively small number of cases (of the syndrome of severe lactic acidosis without shock) recorded in the literature reflect merely the tip of the iceberg". One might as logically say that "I strongly suspect that there is no iceberg"!!!!
3. CONTROL OF THE HYDROGEN ION
ACTIVITY (pH) IN THE BODY
3.1. 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 (Van Slyke 1966).
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. (See Section 8).
There are three mechanisms which diminish pH changes in body fluid: buffers; respiratory; renal.
3.2. 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 as defined in Section 2.4. (See also section 3.3.2 and Appendix A3.2).
3.2.1. 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 (Rune et al 1968). (See Appendix 3.2.1).
3.2.2. 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.
3.2.3. 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.
3.3 RESPIRATORY CONTROL
3.3.1 Normal. In the normal state (at 37�C) 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.
3.3.2 Effect of Control of PCO2 in Minimising pH Changes due to Non-Respiratory Acids or to Bases. Maintenance of the PaCO2 level (even without compensation, see 3.3.3), 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. (See Appendix 3.2).
3.3.3 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.
In summary, sections 3.3.2 and 3.3.3 state that:
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. (See Appendix 3.2).
3.4 RENAL CONTROL
3.4.1 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).
3.4.2 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.
3.4.3 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.
3.4.4 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- (see 3.4.5) (Polak et al 1961). 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 (Peters and Van Slyke, 1931).
A high pH due to low PaCO2 is probably compensated by renal excretion of base, i.e. NaHCO2 or KHC03.
See Appendix 3.4 for extra discussion of renal control of pH.
3.4.5 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.
3.5 COMPENSATION AND CORRECTION
3.5.1 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.
3.5.2 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.
Appendixes
A.3.1 Control of Plasma pH
The mechanism of control of plasma pH at 7.4 is unknown. The receptor which senses changes in pH is unknown, as is the reason why 7.4 is the appropriate pH. Teleologically, the pH at any particular site is that pH which is optimal for enzyme action.
As with most other physiological parameters in the body controlled at particular levels (e.g. haemoglobin level, 14 g%; blood volume, 5 litres), the mechanism of control is unknown. Discussions of how these parameters are controlled, are discussions of how the values are returned to normal if disturbed (e.g. how bone marrow reacts to lowering of haemoglobin level).
The concept of the "normal" pH value is an example of the philosophic problem of induction, i.e. why do all reasonable medical scientists expect and believe that all or almost all stable human organisms will have a blood pH of a particular value?
A classical explanation (Hume) is that it is 'habit or custom' which convinces. This leads to complete scepticism. Bertrand Russell discusses the problem of induction in "The Problems of Philosophy", Chapter 6, 1912 and "The History of Western Philosophy", Chapter 17, Hume, 1946. Karl Popper gives a satisfying and so far not refuted theoretical justification of our expecting further repetition of past events. This he does in among other places in "Objective Knowledge", Chapter 1, 1972.
The concept of normal values and their ranges has been challenged by Schwartz et al, 1973. It is erroneous to assume that if the normal range is the mean � 2 standard deviations that any value outside this range is importantly abnormal, i.e. that it leads to a decision being made. The probability of a particular value being importantly abnormal varies if the expectation of it being abnormal varies, i.e. pH values of 7.4 or 7.2 have different meaning in say a patient with chronic obstructive airways disease and in a patient in whom no pH abnormality is suspected.
A.3.2 Examination of "Buffering" Properties of HCO3-:H2CO2 System
Most texts state that the HCO3- : H2CO2 system is an efficient physiological buffer because the components of the pair are controlled separately (Pitts, 1974). As it is not a chemical buffer of any reasonable efficiency at the blood pH use of the term "buffer" in respect to HCO3- : H2CO2 action introduces considerable confusion. This is illustrated in the following example.
Plasma has a [HCO3-] of approximately 24meq/l and [H2CO2] of 1.2meq/l, hence:
If 1.2meq HCl is added to 1 litre of a solution of 24meq NaHCO2 in water, 1.2meq HCO3- will be converted to H2CO3, so:
If the H2CO3 is held constant then 1.2meq HCl causes the [HCO3- ] to diminish by 1.2meq but the H2CO3 level remains constant at 1.2meq/l, therefore:
7.38 is a trivial drop in pH whereas 7.08 is a large change.
Acting as a buffer the HCO3- : H2CO2 system would have allowed a marked change in pH. Chemically the main process which actually diminished the pH change when HCl was added was the removal of another weaker acid (H2CO3) which was formed when the H+ of the HCl was neutralised by HCO3-.
If base is added to the HCO3- : H2CO3 system at pH 7.4 the change in pH is even greater than when an acid is added. If 1.2meq NaOH is added, the [HCO3-] rises 1.2meq/l or the following equation moves to the right:
H2CO3 ↔ H+ + HCO3- + OH- ↔ H2O + HCO3-
therefore [H2CO3] falls to a very low level so that [HCO3-]/ [H2CO3] rises greatly, and therefore pH rises.
In vivo H2CO3 is prevented from falling so the ratio [HCO3-] / [H2CO3] changes only slightly when [HCO3-] rises following the addition of Na0H, therefore the pH changes only slightly. In other words, if a base is added to the blood the main reason the pH changes very slightly is that a weak acid (H2CO3) is added to keep the PaCO2 constant as the CO2 is used up in partly neutralizing the added base:
Na+ + OH- + CO2 → Na+ + HCO3-
Inspection of the Siggaard-Andersen nomogram shows greater changes in pH if the H2CO3 (PCO2) level in blood is altered, than if similar quantities of strong acid or base are added while the PCO2 is held at 40mmHg. Altering the level of H2CO3 by, say 1.2meq in blood, is not the same as adding 1.2meq H2CO3. In blood in vitro raising the level of H2CO3 1.2meq (80mmHg PCO2) causes the [HCO3-] to rise by 6meq/l (see Appendix 2.4.2 ) therefore in fact 7.2meq H2CO3 is added to blood if the PCO2 is raised from 40 to 80mmHg and the level of H2CO3 raised 1.2meq/l. Adding 7.2meq H2CO3 to blood in vitro therefore gives a pH of 7.08, whereas adding 7.2meq HCl when PaCO2 is 40mmHg gives a pH of 7.29.
A.3.2.1. Other Factors
Factors other than buffer action which diminish the pH change due to transfer of CO2 as blood passes through capilliaries
Except in the case of CO2, only a small quantity of acid or base is transferred during any one passage of blood through a capillary. When CO2 is transferred two mechanisms other than buffer action reduce pH fluctuations.
a) Reduced haemoglobin anion is a stronger base than oxygenated haemoglobin. Less haemoglobin exists in the ionised form when it is deoxygenated. During the process of deoxygenation CO2 is added. The resulting fall in pH is diminished by the reduced haemoglobin taking up more H+ ions than the oxygenated haemoglobin.
b) Some of the CO2 is carried directly combined with haemoglobin as carbamino-haemoglobin. This is more readily formed by reduced than by oxygenated haemoglobin.
A.3.4 Extra Discussion of Renal Control of pH
A knowledge of the conventional postulated mechanism of renal control of pH is assumed in this section (see Relman, 1968 and Pitts, 1974). Text book descriptions of renal control of pH (Pitts, 1974) are often not consistent with the overall ("black-box") function of the kidney in transferring acid or base from the blood to the urine. Briefly, Pitts and others state that pH of blood is controlled by the components of the H2CO3 : HCO3- buffer systems being under separate control, i.e. the respiratory system controls the [H2CO3] by controlling PaCO2 and the kidneys control the HCO3- level by controlling the renal threshold of HCO3-. The ratio [HCO3-] : [H2CO3] is, therefore, controlled and, thus, according to the Henderson-Hasselbalch equation pH is controlled. As the pH is controlled so are the ratios of all the other buffer pairs. In this section I attempt to show [HCO3-] is not a controlled variable but rather a dependent variable. It appears to me that pH itself is a controlled variable and [HCO3-] is dependent on pH and PCO2.
The conventional description of control of non-respiratory disturbances emphasizes the control of [HCO3-] (Pitts, 1974) rather than the excretion of acid (which must include an anion). It is axiomatic that if H+ is taken from blood and excreted, an anion must also be excreted. Exchange of H+ for Na+ cannot be the full explanation of excretion of H+ in the urine. If the blood pH rises by Na+ for H+ exchange the blood Na+ + HCO3- level would rise. This Na+ rise does not occur. Sodium would have to come from an external source. Excretion of H+ = excretion of acid, therefore acid level falls in the blood, and the level of anion of the particular acid must fall.
The extracellular fluid (E.C.F.) is the source of urinary acid. It consists of a mixture of equal numbers of anions and cations in solution. When the E.C.F. becomes more alkaline the total quantity of basic anion or base (HCO3-, OH- etc.) must increase. (This does not apply to the basic anion of a weak acid if that acid's removal is the cause of the rise in pH, e.g. in acute hypocapnia [HCO3-] falls although pH rises). When the pH rises due to loss of acid other than carbonic, H2CO2 dissociates to H+ and HCO3-. The H+ partially replaces the H+ of the acid which has been removed.
The HCO3- level in the blood cannot rise without changes in levels of other anions or cations in the blood. To do so without effects on Na+, Cl- or other ions would defy electrical neutrality, i.e. the blood would become negatively charged and the urine positive. If the HCO3- level in the blood rises either the cation level (Na+) must rise or the non-basic anion (Cl-) level fall. The Na+ level can be raised only by retention of exogenous Na+ if the E.C.F. volume does not fall. The non-basic anion level could fall by the kidneys excreting Cl- with the H+. This is the usual way the kidneys extract and excrete acid from the blood.
For some years the medical literature stated that Na+ and Cl- were not important in the control of blood pH because they are neither acids nor bases (Smith, 1951). It is now recognised that Cl- is vital in control of pH under some circumstances (i.e. when it is the anion of hydrochloric acid) (Schwartz et al, 1968). This should have been obvious from considerations of electricaly neutrality.
The kidney cannot increase the [HCO3-] in the blood only by reabsorption of HCO3- from the glomerular filtrate. Such action would leave the blood level unchanged if total reabsorption of HCO3- occurred. HCO3- must be generated and [Na+] or [Cl-] changes must occur as well.
The kidney is said to control pH by controlling the renal threshold of HCO3- (Pitts, 1974). In this explanation threshold of HCO3- rises if PaCO2 rises and vice versa. If this were the correct explanation the HCO3- level should fall during the recovery phase of chronic CO2 retention, because the HCO3- threshold should now fall. In fact the [HCO3- ] and the pH remain high until Cl- in some form is given (Polak et al, 1961).
If a K+ deficiency is induced in an animal without an accompanying acid-base disturbance and then K2S04 administered, the pH and [HCO3-] of blood rise (Bleich et al, 1966). This is impossible to explain if [HCO3-] is a controlled variable. The explanation is that S042- is not retained by the kidney. When it is excreted it must go with a cation. The kidney has Na+ , K+ or H+ to choose between. As the levels of Na+ and K+ are controlled, the excretion has to be as 2H+ + S042- or 2(NH4)+ + S042-. As the urine becomes acid the blood must become alkaline and the [HCO3-] in blood must rise.
The persisting high [HCO3-] due to loss of HCl in pyloric stenosis has been explained as a high threshold for HCO3- (Kassirer et al, (a) 1966). It would occur if HCl (i.e. gastric juice) was removed in a nephrectomised patient. The blood leaving the stomach would have a high pH while acid was being secreted. Therefore, the systemic blood pH would rise, and with it the [HCO3-] (Le Quesne, 1961). Administration of HCl would correct the high [HCO3-] without intervention by the kidneys (Bradham, 1968). Unless chloride ion in some form is given the pH and [HCO3-] cannot be corrected (Schwartz et al, 1968).
Conclusion
a) If it can be shown that the kidneys directly control all the electrolytes in the E.C.F. except HCO3- , then HCO3- level must be uncontrolled.
b) The kidneys generally can correct only states of excess. Deficiency states have to be corrected by exogenous supplies of raw materials. The kidney is able to correct high levels of Na+ , K+ ,Cl- or H+ but not HCO3- unless some acid is added or control of Na+ is lost (acetazolamide, Diamox). Most causes of non-respiratory alkalosis are acid deficiency states so the high [HCO3-] is incidental to the low [H+] state. The low acid state can be corrected only by giving acid or a salt from which acid can be generated. If the acid used has a non-resorbable anion, e.g. H2S04 or HN03, the correction of the alkalosis will be temporary (Tanner, Schwartz and Bleich, 1966).
Note: Since this section was written (1972) Knud Engel and Paul Kildeberg "Physiological Viewpoint on Clinical Acid-Base Diagnosis", (1977) have written along similar lines. They point out the absence of value in the concepts of bicarbonate reabsorption and renal bicarbonate threshold etc. I think this view is correct. They emphasize the distinction of metabolisable and non-metabolisable acids. I think it is more useful to think about organic and inorganic acids, although all organic acids may not be metabolisable.
The method of describing inputs and outputs which they recommend although consistent with their and my approach is unnecessarily complex in the clinical situation for which the method is advocated.
Schwartz and Cohen (1978) have reviewed several of the paradoxes inherent in the conventional description of renal acidification mechanisms. The approach used in this book is consistent with their hypotheses, i.e. that variation of acid secretion by the kidney occurs in a direction which would appear to be appropriate for pH homeostasis only if this is consistent with other homeostatic mechanisms, i.e. Na & K control. They think, and I agree, that pH control has a low priority in homeostatic mechanisms.
In contrast the conventional approach postulates a H+ excretion mechanism which is controlled by deviations of blood pH and that K+ is involved in the H+ excretion mechanism.
7. TREATMENT
7.1 RESPIRATORY ACIDOSIS
Respiratory Acidosis is corrected by increasing alveolar ventilation and/or treating the cause. Acute respiratory acidosis can probably be corrected rapidly (Prys-Roberts et al, 1967) but chronic respiratory acidosis should be corrected slowly. Rapid lowering of a high PaCO2 has been associated with fits and cardiovascular collapse.
If the patient has a compensatory disturbance (high non-respiratory pH and positive base excess) which has raised the pH towards normal, the kidneys have to correct this disturbance which is equivalent to a metabolic alkalosis. This compensatory metabolic alkalosis does not require treatment if the patient is given a mixed diet with adequate Cl- (Polak et al, 1961), and provided there are no complicating factors in its genesis or in the cardiovascular or renal systems. If the compensatory non-respiratory alkalosis is preventing the patient from lowering his PaCO2 the alkalosis may need to be treated by modifying the steroid or diuretic therapy. Acetazolamide (Diamox) can be used to correct the metabolic alkalosis. It differs from other diuretics in producing an acidosis by the loss of Na+ HCO3- in the urine. When this drug causes loss of base in the urine the sodium content of the E.C.F. must fall. Although acetazolamide will produce a lessening of alkalosis by loss of base, rather than the physiological mechanism of acid retention, the coincidental loss of Na+ may be a benefit if the patient has any cardiac failure.
7.2 RESPIRATORY ALKALOSIS
7.2.1. Treat the cause
e.g. correct hypoxia or shock if they are causing hyperventilation.
7.2.2. Elevate the PaCO2.
This can be corrected by administering CO2, increasing the dead space or lowering the minute ventilation. These measures will rarely be thought to be necessary.
7.3 METABOLIC ACIDOSIS
7.3.1 Treat the cause.
Stop alimentary loss of base; correct hypoxia; reduce renal acid load by diet; drain abscess in diabetic ketosis and give insulin (see 7.3.2.2.3, ketoacidosis) ; treat shock with intra-venous fluids and stop haemorrhage etc (see 7.3.2.2.2.2, shock) .
7.3.2. Correction of Acidosis
7.3.2.1. Administration of NaCl.
If the acidosis is (a) not affecting the cardiac action and (b) renal function is adequate, the acidosis may be corrected by giving sufficient NaCl (Na+ + Cl-) solution for the kidney to (i) correct the acidosis by excreting HCl (H+ + Cl-) or NH4Cl (NH4+ + Cl-) and (ii) repair any deficit in E.C.F. volume. This approach applies in alimentary causes of metabolic acidosis where the kidneys are usually able to correct the defects if enough saline is given (Hesse et al, 1966). Correction may be more rapid if Hartmann's solution rather than 0.9% NaCl solution is given to correct the pH disturbance as there is less for kidney to do. The lactate ion has to be converted to HCO3- and some H+ + Cl- will have to be excreted but not as much as with NaCl solution.
7.3.2.2. Administration of Base.
Indications for direct correction of acidosis by giving base:
7.3.2.2.1. The cause cannot be corrected. e.g. renal acidosis, where the kidneys fail to excrete inorganic acid (an end product of protein metabolism). If this defect is the sole manifestation of renal impairment (i.e. renal tubulcar acidosis), it is rational to neutralise the acid with NaHCO2 which can be given by mouth. In most instances renal failure is not manifest solely by acidosis. Usually dialysis or transplantation is necessary to correct the multiple effects of renal failure which include acidosis.
7.3.2.2.2. Where the acidosis is depressing the circulation (i.e. to break the viscious circle of myocardial depression which aggravates acidosis). This is the indication in cardiac arrest (Chazan et al, 1968) or in shock.
7.3.2.2.2.1. Cardiac Arrest. In cardiac arrest, acute lactic acidosis, it is said, may prevent the circulation restarting. NaHCO2 is often given in an empirical and probably excessive dose of "1 bottle" (i.e. 500mls 4.2% NaHC03 = 250meq). PaCO2 rises as some HCO3- is converted to H2CO2. Serum [Na+] will rise and circulatory overload may be caused by the Na+ load. After recovery a metabolic alkalosis will occur until the Na+ + HCO3-are excreted.
There is controversy surrounding the use of base in cardiac arrest (Stackpool 1986, Narins and Cohen 1987), as in the emergency situation it is impractical to obtain biochemical evidence before treatment (Leading Article, Lancet, 1976). Non-respiratory acidosis does not occur in all cases of cardiac arrest (Stewart, 1964, Chazan et al, 1968). Inadequate ventilation causing a high PCO2 is as frequent a cause of low pH in cardiac arrest as is metabolic acidosis. There is no reason for thinking that the low pH due to CO2 has much cardiovascular depressant effect (Schultz et al, 1960; Prys-Roberts et al, 1967; Gerst et al, 1964).
If the circulation is restored rapidly after resuscitation is commenced (i.e. pulse and/or consciousness) than any NaHCO2 infusion which may have been started should be stopped until the situation is assessed biochemically (Rackwitz et al, 1976).
Mattar et al, 1974, investigated 12 patients in whom NaHCO3 was administered during cardiac arrest resuscitation. Full measurements were not available in all cases. The doses of NaHCO3 were between 45 and 270meq. The mean changes were:
before NaHCO3
after NaHCO3
Osmolality ( mOsmol/litre)
302
377
pH
7.34
7.54
HCO2 (mEq/litre)
21
50
Na (mEq/litre
138
170
It has also been shown (Fillmore et al, 1970) that even if it is possible to effectively adjust pH abnormalities during resuscitation by use of alkali, restarting heart action still may not occur. Presumably this lack of restarting was be due to overwhelming myocardial damage or inadequate circulation provided by the cardiac massage. This was suggested by the rising lactic acid level which was probably an effect rather than the cause of continuing inadequate heart action.
It is still common practice to administer NaHCO3 during cardiac arrest. In some institutions NaHCO3 infusion is left set up on each cardiac arrest trolley to save time in starting its administration.
The dose required to correct a low non-respiratory pH is arbitrary in the individual patient. I would advise that a base-line specimen of blood should be taken before the NaHCO3 administration, if this is possible. Venous blood from a central vein if one (e.g. the internal jugular) is being used to administer drugs is satisfactory, but arterial might be better. Follow up blood measurements would then be done as resuscitation continues and is completed, to examine pH, osmolality, Na, K , PCO2, PO2 and lactate levels.
In summary I would conclude that correction of non-respiratory pH during cardiac resuscitation is not as important as once was thought, and that monitoring of biochemical changes before and after such correction should be routine.
7.3.2.2.2.2. Shock. In shock accompanied by acidosis it is rational to administer base together with other haemodynamic management, i.e. raising the C.V.P., giving inotropic agents, oxygen, etc. (Manger et al; 1962, Ledingham, 1962; MacKenzie, 1965). When base is given in shock it is not rational to give it over some hours after having decided on a dose. The acidosis should be corrected as quickly as possible in two or three steps controlled by non-respiratory pH, base excess or standard bicarbonate measurements, e.g. first dose (7.4 - non-respiratory pH) x body weight (kg) x 7meq or - (0.1 x base excess x body weight) meq. Subsequent doses are estimated after the effect of the first dose has been observed. Dosage of NaHCO2 (or any other intravenously administered electrolytes) cannot be determined by formulae. Response to initial dosage will suggest magnitude of subsequent doses. Factors which would have to be taken into account if dosage was to be predicted would have to include cause of acidosis, the circulatory state and the magnitude of the acidosis. The degree of acidosis alters the requirements of NaHCO2 by a function which is not a direct proportion (Garella et al, 1973).
7.3.2.2.3. Diabetic Keto-Acidosis. The use of HCO3- in diabetic keto-acidosis is also controversial. Acidosis has been claimed to be a cause of insulin resistance in diabetic acidosis (Walker et al, 1963). This study was not controlled. Regimes including HC03 have been advocated without conclusive evidence of benefit (Solar et al, 1973; Solar et al, 1974). Administering NaHC03 in diabetic keto-acidosis may exaggerate changes in serum K+ particularly if this is changing rapidly following high doses of insulin.
It now obvious that many of the problems of managing diabetic acidosis (i.e. ketotic and non-ketotic) were iatrogenic, i.e. due to large intermittent I.V. and I.M. dosage of insulin, rapid and erratic alterations in blood pH, K+ and glucose levels. Diabetic acidosis can be corrected over a period of about 6 hours with low dose I.V. infusions of insulin. There is steady biochemical and clinical improvement without swings in blood glucose or K+ levels. Only small doses of K+ are necessary and in most instances HCO3- is not given, although when it has been, it has apparently not produced adverse effects (Alberti et al, 1973; Page et al, 1974; Kidson et al, 1974; Semple et al, 1974; King et al, 1974 and Shaw et al, 1974). If NaHC03 is to be used in diabetic acidosis its use should probably be limited to patients with severe acidosis (pH<7 100meq="" 70kg="" a="" and="" be="" dose="" font="" greater="" in="" initial="" not="" patient.="" probably="" should="" than="" the="">
The low PaCO2 which is present in the initial phases of diabetic acidosis may persist after the blood pH has returned to normal. This is probably not due to delay in return of the CSF pH to normal (King et al, 1974).
7.3.2.2.4 Neonatal resuscitation. NaHCO2 is used in neonatal resuscitation (Clark et al, 1968). The use is analogous to its use in cardiac arrest. The need or efficacy of the treatment as a routine has not been established. When used, pH electrolyte studies should be done, at least retrospectively. Hypernatraemia and intracranial haemorrhage have been associated with administration of NaHCO2 in the neonatal period (Simmon et al, 1974; Volpe, 1974).
7.3.2.2.5 Neonatal Respiratory Distress Syndrome (RDS). NaHCO2 has been used to correct the "chronic" acidosis of this syndrome (Usher, 1963). The acidosis is presumably due to hypoxia and correction of this (Daily et al, 1971; Smith and Daily, 1971) if possible, would be more rational (Dell and Winters, 1972). In other cases CO2 retention may be the cause of low pH. In these cases increased ventilation would be the rational treatment. (Ostrea et al, 1976). When it has been given in the patients with RDS to correct acidosis, NaHCO2 gives an acute rise in PaCO2. Tham causes an acute fall in PaCO2. These changes in PaCO2 last for some minutes (Baum et al, 1975).
7.3.2.3 Complications of NaHCO2 Therapy
Metabolic alkalosis which may then cause respiratory depression.
Hypernatraemia and hyperosmolality (Mattar et al, 1974; Bishop et al,1976).
Changes in other electrolytes especially lowering of serum [K+].
Fluid retention in patients who have disorders which will lead to fluid retention if excess Na+ is given.
Acute rise in PaCO2 due to neutralization of HCO3- (Singer et al, 1956; Ostrea and Odell, 1972; Baum et al, 1975).
Intracranial haemorrhage (Simmons et al, 1974; Volpe, 1974; Papil et al, 1978).
7.3.2.4 Tham
Tham (Nahas, 1961) is an organic base (often referred to as a buffer) used to correct acidosis. It has no obvious advantages over sodium bicarbonate. Its claimed advantages include ability to correct intracellular acidosis. A separate syndrome of intracellular acidosis is not clinically recognised. Its main value might be in situations in which Na+ load may be undesirable, e.g. cardiogenic shock. Although, as Tham is an osmotically active agent it may have similar effects on extracellular fluid volume as Na+. It might be of use if one wished to correct the low pH of respiratory acidosis directly (Manfredi et al, 1960), as some have claimed (Mithoefer et al, 1965 and 1968) that bronchospasm is relieved by direct correction of low pH. Its main disadvantages are inconvenience of preparation (it is supplied in a powder) and that it causes respiratory depression. Respiratory depression would presumably be caused by any substance which raised blood pH including Na0H or NaHCO2. Such depression could be managed by intermittent positive pressure ventilation so is not in itself an absolute reason for not using Tham.
7.4 METABOLIC ALKALOSIS
7.4.1 Remove the cause (e.g. relieve pyloric obstruction or modify diuretic regime).
7.4.2.1 Administration of NaCl. Ingestion or injection of sufficient sodium chloride solution for the kidney to correct the alkalosis by excretion of Na+ + HCO3- .
7.4.2.2 Administration of Acid. Direct correction of alkalosis with ammonium chloride or hydrochloridc acid solution, infusion or ingestion (Bradham, 1968, Leading Article (a) 1974, Sanderson, 1974; Pain et al, 1974; Abouna, 1974; Harken et al, 1975 and Worthley, 1977). This is indicated only if the alkalosis is very severe or renal or cardiac function are poor. Usually there is an associated reduction of extracellular volume so some Na has to be given in the form of NaCl. Compensation for non-respiratory alkalosis is CO2 retention achieved by hypoventilation. The hypoventilation may result in hypoxaemia which may necessitate oxygen therapy.
Clinical problems of pH are all related to pH of the plasma of whole blood. pH in extracellular fluid is always close to that of blood. pH inside cells differs from that of blood but it is not recognised as being an important clinical problem apart from blood pH changes.
In the clinical situation if the actual pH of the blood is lowered, one can usually assume that the primary disturbance has been the addition to the blood of acid or the removal of base and vice versa.
6.2 CHEMICAL CLASSIFICATION OF CAUSES OF CHANGES IN BLOOD pH
Blood pH may be changed if acids or bases are added to or removed from the blood. Secretion of an acid (e.g. gastric juice) implies that the acid involved (HCl in this case) is removed from the blood.
The acids which can cause changes in blood pH are:
A. Metabolic or non-respiratory acids :-
1. Organic
1.a. Lactic Acid
1.b. Keto Acids
1.a.i. Hypoxia
1.a.ii. Drug Induced
1.a.iii. Idiopathic
1.b.i. Diabetes
1.b.ii. Starvation
1.c.Might be called "true" metabolic acidoses
2.a. Inorganic
2.a.i. Sulphuric Acid
2.a.ii. Phosphoric Acid
2.a.iii. Hydrochloric Acid
2.c. Inorganic acids, increase in renal failure
B. Respiratory acids >> Organic only>> Carbonic Acid.
A rise in concentration of any of these acids in the blood causes a fall in the pH of the blood. Loss of acid from the blood (e.g. into gastric juice) causes a rise in the pH. Only HCl and H2CO3 can be lost from the blood in appreciable quantities.
The bases which can cause changes in blood pH are:
•NaHC03
•KHC03
Administration of base by mouth or parenterally may cause blood pH to rise if rate of excretion does not match rate of administration. Loss of alkaline fluid from bowel (diarrhoea, intestinal obstruction or intestinal fistulae), or urine (after acetoazolamide) will cause blood pH to fall.
6.3 CLINICAL CLASSIFICATION OF CAUSES OF CHANGES IN BLOOD pH
Clinical states of pH disturbence (acid-base inbalance) can conveniently be divided into two groups, i.e. (a)respiratory and (b)metabolic or non-respiratory. The reasons for this division into respiratory and non-respiratory are that:
i) the compensatory mechanisms (Section 3.5.1) and treatments (Section 7) of the two types are different.;
ii) the recognition of non-respiratory disturbances is masked by compensatory alterations in PCO2 and the recognition of changes in pH caused by PCO2 changes are masked by renal compensation.
6.3.1 RESPIRATORY ACIDOSIS. This is synonymous with CO2 retention and is usually a sign of hypoventilation. Compensation is renal. There is renal loss HCl in the form of buffer or as NH4Cl. During recvovery chloride has to be supplied and retained.
6.3.1.1 Causes of hypoventilation:
•Central nervous system
•Peripheral nervous system
•Neuromuscular transmission
•Muscle disorders
Chest wall abnormalities
Lung and airway disorders.
6.3.1.2 Inhalational of CO2 This is another cause of respiratory acidosis, but it is only likely to occur under situations of re-breathing, e.g. under anaesthesia or during resuscitation with a Water's cannister circuit without the cannister, i.e. ward resuscitators or Type C anaesthetic systems. (Mapleson, 1954).
6.3.1.3 Increased production of CO2. This very rarely causes a high PaCO2. In thyrotoxicosis and fever CO2 production is raised but the rise is well within the capacity of a normal respiratory system. In malignant hyperpyrexia high PaCO2s have been recorded due undoubtedly to increased production.
NaHCO2 therapy in non-respiratory acidosis causes a rise in PaCO2 until the CO2 generated by the neutralization of HCO3- has been excreted (Singer et al, 1956). This effect may be prolonged if CO2 excretion is impaired (Ostrea and Odell, 1972).
Acute respiratory acidosis is diagnosed by high PaCO2 and low actual pH associated with near normal non-respiratory pH, (standard HCO3-and base excess). (See Appendix 4.2).
Chronic respiratory acidosis is compensated by loss of H+Cl- by the kidney and generation (not reabsorption) of (Na+) HCO3-. (See Appendix 3.4). Compensation causes slightly lowered actual pH with a high non-respiratory pH (positive base excess and high standard HCO3-). Steroids and diuretics (except acetazolamide (Diamox) ) may produce a non-respiratory alkalosis which may cause a respiratory acidosis to appear to be "over compensated". There will then be a high (actual) pH.
In chronic respiratory failure ventilation is maintained by hypoxic drive and also by the acid pH of the blood. If the pH stimulus is removed by the actual pH being raised, some of the respiratory drive will be removed causing a diminution of respiratory effort and further CO2 retention.
When a chronic respiratory acidosis (with compensation, i.e. high non-respiratory pH, Base excess, and Standard HCO3-.) is improved by increasing alveolar ventilation the actual pH may rise above 7.4. The resulting chemical picture of high pH, high non-respiratory pH (positive base excess), and high PaCO2 is indistinguishable chemically from a primary metabolic alkalosis with compensatory CO2 retention. Clinical information is required to distinguish them. (Plot this and other examples on a standard Siggaard-Andersen nomogram as an exercise). ( See Footnote to table 4.2.2.)
N.B. The compensatory renal or respiratory changes which partially correct pH disturbances are chemically true pH changes in the opposite direction to the original disturbances. It is a semantic problem whether the compensatory renal change induced for example by primary chronic CO2 retention is called a metabolic alkalosis or not. This semantic muddle can be avoided if the term alkalosis is used only in a non-precise term.
6.3.2. RESPIRATORY ALKALOSIS. This is associated with hyperventilation. Usually these are acute so there is no time for renal compensation, but if prolonged, such as in acclimatization to high altitudes, there would probably be renal compensation.
6.3.2.1. Deliberate induced hyperventilation during anaesthesia.
6.3.2.2. Some causes of hypoxia associated with hyperventilation. This would include "air hunger" of hypovolaemia, severe ventilation-perfusion abnormalities and acclimatization to high altitudes. Critically ill patients without arterial hypoxaemia can have severe hypocapnia (Mazzara et al,1974). This may be explained by low cardiac output.
In asthma attacks, unless severe, the PaCO2 is usually lowered, i.e. the alveolar ventilation is increased (McFadden and Lyons, 1968). This is at least sometimes associated with hypoxaemia due to mismatching of ventilation and perfusion.
6.3.2.3. Fever. This may cause hypocapnia due presumably to central stimulation of the respiratory control mechanisms (Chapot et al, 1974).
6.3.2.4. Some types of C.N.S. damage.
6.3.2.5. Hysterical hyperventilation.
6.3.3. METABOLIC ACIDOSIS.
This (non-respiratory acidosis) is due to increase in acids (i.e. H+ donating substances) other than H2CO2 or decrease in base (i.e. H+ acceptors) in the blood. Compensation is by hyperventilation. This lowers the PaCO2 thus deducing the any pH change. The causes of non-respiratory acidosis are:
Increased alimentary or parenteral intake of acid or alimentary loss of base. (Addition and subtraction).
Increased production of acid. (Accumulation).
Failure of excretion of acid or loss of base by the renal system.
Increased Intake of Acid or Loss of Base
6.3.3.1 Increased alimentary or parenteral intake of acid or alimentary loss of base.
6.3.3.1.1. Adding acid. The acid content of blood may be raised by ingestion or injection of ammonium chloride or dilute hydrochloric acid. The hydrochloric acid directly increases [H+]. The ammonium chloride produces hydrochloric acid by the NH3 being split off and converted to urea. Adding ammonium chloride directly to blood would not change the pH greatly until the NH4+ has been metabolised.
6.3.3.1.2 Alimentary loss if Base. Loss of intestinal contents by diarrhoea, low small bowel obstruction or intestinal fistulae causes loss of fluid of high pH, i.e. containing an excess of base (Na+HCO3- or K+HCO3-). This results in the fluid left in the body having a lower base content than normal. The removal of base causes the blood pH to fall (Leading Article, 1966). A similar disturbance has been reported from loss of lymph. (Siegler et al, 1978).
6.3.3.1.3 Intravenous infusions. These can cause an acidosis.
6.3.3.1.3.1 Stored blood for transfusion. This has la ow pH. The anticoagulant contains citric acid. When mixed with the blood the pH drops. PCO2 rises because of the action of acid on bicarbonate ions. Most of the pH drop is due to this CO2 which does not escape from the stored blood. The non-respiratory pH of the stored blood is not as low as the actual pH. The pH of stored blood does not fall progressively if stored for up to three weeks at 4�C (Gaudry, Joseph and Duffy, 1974).
The acidic salt of sodium citrate and "dilutional" acidosis are the main contributers to the non-respiratory pH of stored blood. There is a variable amount of lactic acid acid in stored blood This usually contributes less to the non-respiratory pH change. If stored concentrated red cells are washed there will be a considerable and even an increased non-respiratory acidosis in the product. If 27meq NaHCO2 is added to each litre of washing fluid this would not occur, but would introduce other problems.
Stored blood transfusion rarely causes a non-respiratory acidosis if the circulation and temperature are maintained normal. It is possible that if the circulation and temperature are not maintained, that metabolic acidosis could occur with massive transfusion. In a situation where a massive transfusion is given it will not usually be possible to distinguish the low pH due to blood, from that due to the condition for which the blood is being given. From personal observations the low base excess observed during liver transplantation appears to be mainly a combination of lactic acidosis from the transfused blood and released or generated in the new liver and dilutional acidosis. The sodium citrate in the anti-coagulant solution can cause a metabolic alkalosis after the citrate has been metabolised and replaced by bicarbonate ion.
Howlands et al (1965) have advocated routine use of sodium bicarbonate during massive blood transfusion to counteract the acidosis of the infused blood. This is usually unnecessary (Bookallil and Joseph, 1968), and will probably only aggravate post transfusion alkalosis (Miller et al, 1971).
If the temperature and circulation are not maintained as is common in when more than say 5 litres in 2 hours are transfused in trauma the composition of the circulating blood approaches that of the transfused blood. In liver transplantation 3 times the blood volume may be replaced in an hour and up to 30 times the blood volume during the operation. This can result in pH<6 .9="" acidosis.="" alkali="" and="" appars="" base="" be="" cause="" circulatory="" dilutional="" effects="" ersonal="" excess="" font="" is="" lactic="" litre.="" little="" meq="" observation="" occurs="" recovery="" spontaneous="" the="" therapy="" there="" to="" without="">
6.3.3.1.3.2 "Dilution Acidosis" (Shires et al, 1948; Garella et al, 1973) is an acidosis due to the intravenous infusion of a neutral non-buffer solution.
Intravenous solutions which contain only non-acids or non-bases, e.g. sodium chloride, glucose and other carbohydrates usually have a pH slightly less than 7. This is usually due to pharmaceutical details in the preparation. In these non-buffer solutions this low pH represents a very small amount of acid. One would only have to add a fraction of a miliequivalent of strong base to a litre to make the solution alkaline.
If such a non-buffer solution (e.g. Saline or 5% glucose) is equilibrated with a gas mixture containing 40mmHg CO2 it has a pH of 4.9 (Gaudry et al, 1972). 24meq Na0H (or NaHC03) would have to be added to it to give a pH of 7.4, therefore a neutral solution plus 24meq NaHC03 will on infusion cause no change in pH in the blood. The original solution of saline or glucose will act in a similar fashion to an injection of 24meq HCl for each litre infused. This will only be importasnt if large volumes of intravenous fluid are given in a short time,e.g. in burns or cholera (see Diuretic Alkalosis, 6.3.4.2).
6.3.3.1.3.3 Metabolic acidosis associated with intravenous alimentation. Intravenous feeding with amino acid solution in the form of hydrochlorides can cause a metabolic acidosis associated with a high serum Cl level (Chan, Ghadimi, Kaminski and Heird et al, 1973). Lactic acidosis may be associated with intravenous feeding regimes which contain fructose, xylotol or sorbitol (Alberti and Nattrass, 1977).
In some cases hypertonic glucose also has been associated with lactic acidosis (Ames et al, 1975).
6.3.3.4 Ingestion or injection of oganic acids. This would not ordinarily produce change in pH because the liver would metabolise them. In liver disease and during liver transplantation organic acids introduced to the body gain access to the systemic circulation.
6.3.3.5 Accumulation of Acid. Excess acid may accumulate in the blood from processes of metabolism and cause a fall in the pH. There are two main mechanisms for this:
6.3.3.5.1 Hypoxia. Hypoxia from any cause prevents the products of anaerobic glycolysis being further metabolised. Excess lactic acid appears in the blood and lowers the pH.
The causes of hypoxia are:
Low oxygen inspired air.
Lung disorders.
Hypoventilation.
Low cardiac output (including shock states and myocardial infarction (Neaverson, 1966, Kirby et al, 1966).
Blood defect; hypovolaemia, anaemia or CO poisoning.
Tissue toxins, e.g. cyanide. Cyanide is important in sodium nitroprusside toxicity. This may be important in combined respiratory and circulatory failure if sodium nitroprusside is used to lower peripheral resistance.
In cardiac arrest there have been several studies which document low blood pH (Ledingham et al, 1962; Chazan et al, 1968 and Fillmore et al, 1970). All cardiac arrests do not have non-respiratory acidosis. Respiratory acidosis (high PaCO2) is a feature of some. High PaCO2 for practical purposes means inadequate alveolar ventilation, which of course, might be present. Fillmore et al. imply that aspiration may be a cause of high PaCO2. This is only the case if it associated with hypoventilation. If ventilation is inadequate not only will the PaCO2 be high but what is more important the oxygenation will be inadequate unless oxygen is given. High levels of alveolar oxygen will not be present if mouth-to-mouth ventilation is being used. A transient rise in PaCO2 will occur for a few minutes after NaHC03 administration, due to increased CO2 release (Bishop et al, 1976).
Circulatory occlusion of any large area causes accumulation of organic acids in the area supplied. On restoration of the circulation these acids are distributed systemically. This is a possible cause of acidosis, but in practice if the general circulation and temperature are maintained in a satisfactory state only transient acidosis results as the acids are quickly metabolised (Bookallil and Joseph, 1968).
Non-hypoxic lactic acidosis is a syndrome being more commonly discussed. It is a serious syndrome of unknown aetiology (Oliva: Leading Article, 1970; Leading Article, 1973; Harken, 1976: Alberti et al, 1977). It probably includes syndromes of varying aetiology, e.g. phenformin associated acidosis in diabetes (Gale et al, 1976), and the acidosis associated with parenteral nutrition with some carbohydrates (See 6.3.3.1.3.3 and Appendix 6.3.3.2.1).
6.3.3.5.2 Diabetes and Starvation. In both these states excess acetone and keto-acids are produced. There will be some attempt at renal excretion (ketonuria) but if this is inadequate the quantity of acid in the blood will rise and the pH wil fall (Peters and Van Slyke, 1931). All acidoses in diabetes are not due to keto-acid. In some instances lactic acid is the cause. Patients clinically diagnosed as diabetic keto-acidosis sometimes have alkalosis (Lim and Walsh, 1976).
6.3.3.5.3 Others. There are some other disturbances of intermediary metabolism which may cause blood pH to fall from accumulation of organic acids. One of these which is mentioned in 6.3.3.1.3.3 occurs from intravenous administration of some sugars.Another is iso-valeric acidaemia - a rare inborn error of metabolism due to an enzyme deficiency (Cohn et al, 1978).
6.3.3.3 Failure of excretion of acid. This may lead to acidosis. Normally during metabolism some inorganic acids are produced, i.e. sulphuric and phosphoric acids. The anions have to be excreted by the kidney covered either by Na+ , K+ , H+ (small amount) or NH4+ (produced in the kidney). Normally the quantity of acid involved is not great (40-60meq/day), but over a long period, if there is failure of excretion, accumulation will occur with a fall in pH. This is renal acidosis (Relman, 1968). Renal acidosis may also be due to loss of base induced by acetazolamide (Diamox). Renal acidosis is usually a manifestation of generalised renal failure but there is also a syndrome of renal tubular acidosis where metabolic acidosis of renal origin is an isolated disorder.
6.3.4 METABOLIC ALKALOSIS (non-respiratory alkalosis). This is due to loss of HCl from the ECF or addition of alkali. Metabolic alkalosis is compensated by respiratory depression which causes CO2 retention (Tuller and Mehdi, 1971; Shear et al, 1973; Aquino et al, 1973) but may also cause hypoxia. The pH is usually raised but may be high normal if there is much CO2 retention. Metabolic alkalosis is due to:
6.3.4.1. Loss of gastric juice containing HCl. Patients with pyloric obstruction lose some K+ and Na+ as well as HCl. The loss of K+ is mainly through the kidney (Kassirer et al, 1966). At first the urine is alkaline but after stable conditions are established the urine becomes acidic as the normal inorganic acid load from protein breakdown still has to be excreted (Schwartz et al, 1978). If it were not, it would correct the alkalosis. The acid urine used to be thought to be paradoxical (Van Slyke and Evans, 1947), and was attributed to K+ deficiency.
The situation of a chronic metabolic alkalosis with acid urine is probably the best clinical example where balance or status as destinct from input, output or turnover (section 2.1) should be distinguished. The blood pH status is stable and alkaline. For the non-respiratory pH to remain high the normal acid output must continue, i.e. acid excretion in the urine will be normal and the urine pH will be low.
6.3.4.2 Diuretic alkalosis. Thiazide diuretics, frusemide and ethacrynic acid can produce a metabolic alkalosis. HCl, its equivalent NH4Cl or HCl having acted on phosphate buffer is lost in the urine. The central role of Cl in the production of diuretic alkalosis has been established (Kassirer et al, 1965).
6.3.4.3 Ingestion or injection of excess base, e.g. Na+HCO3- or Na+OH-. Post transfusional or post-cardiac surgery metabolic alkalosis is usually due to the administration and metabolism of sodium citrate (Kappogoda et al, 1973; Barcenas et al, 1976). If NaHCO2 is given to "correct" an acidosis it may result in a high serum Na and osmolality and later an alkalosis.
6.3.4.4 Steroid alkalosis (Kassirer et al, 1970; Schambelan et al, 1971).
6.4 CLINICAL DIAGNOSIS OF ACID-BASE OF pH DISTURBANCES
Any of the clinical conditions mentioned can be associated with pH or acid-base status disturbances and one can usually suspect pH changes by the clinical picture, e.g. hypoventilating, acute on chronic obstructive airway disease, probably has CO2 accumulation; intestinal obstruction or severe diarrhoea probably has acidosis due to loss of base Na+ and/or K+ + HCO3- ); a diabetic who is drowsy and hyperventilating with urinary glucose and ketones probably has keto-acidosis; "shock" with poor tissue perfusion may have lactic acidosis.
The signs of acid-base disturbances are not diagnostic but in association with the clinical picture, a suspicion of variable certainty can be entertained. Definite daignosis is only possible by measurements on blood which should usually be arterial but can be capillary, venous or mixed venous. See tables 4.2.1 and 4.2.2.
6.5 PHYSIOLOGICAL EFFECTS OF pH DISTURBANCES
It is often stated that blood pH below some arbitrary level, usually 6.8, is incompatible with survival (Andreoli 1988). This is not so. Low pHs' have been recorded in association with CO2 retention (Schultz, 1960, pH 6.71, PaCO2 234mmHg. Prys-Roberts et al, 1967, pH 6.86, PaCO2 248mmHg) and severe exercise (Osnes and Hermansen, 1972, Hermansen and Osnes, 1972) with little physiological disturbance attributed to the low pH, and with complete recovery.
Two cases of strychnine poisoning who recovered after low pHs' were recorded, have been reported, (Goldstein, 1975, pH 6.57 PaCO2 18mmHg; Loughhead et al, 1978, pH 6.59 PaCO2 59.5mmHg).
If the blood pH is abnormal, the cause of the disturbance is usually severe, and compensatory or correcting mechanisms have not occurred or have been insufficient. It is in practice impossible to distinguish the physiololgical effects of the cause from those of the pH change itself.
There are many effects of pH changes which are reversible when the cause and/or the pH change are removed. As far as the whole organism is concerned the cardiovascular effects of pH change are the most important. A low pH is said to aggravate cardiovascular depression. This then further lowers pH due to lowered perfusion, thus setting up a vicious circle.
There is evidence that correcting a low non-respiratory pH in circulatory insufficiency improves the prognosis (Manger, 1962, Ledingham et al, 1962). Low non-respiratory pH can worsen arrhythmias (Anderson, 1968). Artificially induced pH changes due to infusion of acid in normal animals does not reduce cardiovascular activity with moderate falls in pH (Anderson, 1967, 1968, Caress et al, 1968). Maybe a low pH has a much greater effect in an already impaired circulatory system than in the normal. There is controversy (Stackpool,1986, Narins & Cohen 1987) about use of alkali in states of of low pH.
The cardiovascular effects of pH changes are probably different in the various chemical causes, e.g. hypoxic lactic acidosis produces more disturbance than a similar pH change due to CO2 (Schultz, 1960, Prys-Roberts et al, 1967).
Alteration of enzyme action when the pH is abnormal is usually given as the mechanism of the deleterious effects of pH changes. This, of course, is the basic mechanism in all disturbances but it does not indicate in which system the effect is critical. It is cardiovascular effects of pH changes which set the limit.
High pH levels also produce physiological disturbances. These are less well documented and of less clinical importance. The main clinically important effects are CNS effects (tetany and impaired consciousness) hypoventilation (if the high pH is not due to hyperventilation and a low PCO2) and possible cardiovascular effects (Streisand et al, 1971; Lawson et al, 1973; Galmarini et al, 1973; Lock et al, 1975).
The serum potassium concentration falls when the PaCO2 is acutely lowered. Return of the potassium concentration may be slower than the return of the pH when the PCO2 recovers (Edwards et al, 1977; Sanchez et al, 1978; Finsterer et al, 1978).
Appendices
A.6.3.3.2.1. Lactic Acidosis
A review article "Lactic Acidosis" by Alberti and Nattrass (1977) gives a rational and useful classification of lactic acidosis drawing on the book by Cohen and Woods, "Clinical and Biochemical Aspects of Lactic Acidosis" (1976).
This classification is:
Type A Shock syndrome and hypoxia.
Type B1 Common disease states, e.g. diabetes, liver disease, infection.
Type B2 Drugs, e.g. phenoformin, sorbitol, fructose.
Type B3 Hereditary metabolic disorders.
They state that as these categories are better defined, the reporting of so-called iodiopathic lactic acidosis decreases.
It is sensible to treat the cause if this is known. This may be all that is required in Type A. Restored circulation and oxygenation will permit the lactic acid to be metabolised. In Type B more active measures are suggested to correct the pH abnormality as a low pH may cause the liver to produce further lactic acid. Large doses of NaHCO2 (up to 2,500mMol) have been given. This quantity would usually produce problems of its own which would necessitate dialysis. Such situations are very rare or else are unrecognised.
In the preface of Cohen and Woods' Monograph which contains an exhaustive literature review, it is stated, "We strongly suspect that the comparatively small number of cases (of the syndrome of severe lactic acidosis without shock) recorded in the literature reflect merely the tip of the iceberg". One might as logically say that "I strongly suspect that there is no iceberg"!!!!
3. CONTROL OF THE HYDROGEN ION
ACTIVITY (pH) IN THE BODY
3.1. 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 (Van Slyke 1966).
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. (See Section 8).
There are three mechanisms which diminish pH changes in body fluid: buffers; respiratory; renal.
3.2. 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 as defined in Section 2.4. (See also section 3.3.2 and Appendix A3.2).
3.2.1. 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 (Rune et al 1968). (See Appendix 3.2.1).
3.2.2. 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.
3.2.3. 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.
3.3 RESPIRATORY CONTROL
3.3.1 Normal. In the normal state (at 37�C) 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.
3.3.2 Effect of Control of PCO2 in Minimising pH Changes due to Non-Respiratory Acids or to Bases. Maintenance of the PaCO2 level (even without compensation, see 3.3.3), 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. (See Appendix 3.2).
3.3.3 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.
In summary, sections 3.3.2 and 3.3.3 state that:
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. (See Appendix 3.2).
3.4 RENAL CONTROL
3.4.1 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).
3.4.2 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.
3.4.3 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.
3.4.4 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- (see 3.4.5) (Polak et al 1961). 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 (Peters and Van Slyke, 1931).
A high pH due to low PaCO2 is probably compensated by renal excretion of base, i.e. NaHCO2 or KHC03.
See Appendix 3.4 for extra discussion of renal control of pH.
3.4.5 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.
3.5 COMPENSATION AND CORRECTION
3.5.1 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.
3.5.2 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.
Appendixes
A.3.1 Control of Plasma pH
The mechanism of control of plasma pH at 7.4 is unknown. The receptor which senses changes in pH is unknown, as is the reason why 7.4 is the appropriate pH. Teleologically, the pH at any particular site is that pH which is optimal for enzyme action.
As with most other physiological parameters in the body controlled at particular levels (e.g. haemoglobin level, 14 g%; blood volume, 5 litres), the mechanism of control is unknown. Discussions of how these parameters are controlled, are discussions of how the values are returned to normal if disturbed (e.g. how bone marrow reacts to lowering of haemoglobin level).
The concept of the "normal" pH value is an example of the philosophic problem of induction, i.e. why do all reasonable medical scientists expect and believe that all or almost all stable human organisms will have a blood pH of a particular value?
A classical explanation (Hume) is that it is 'habit or custom' which convinces. This leads to complete scepticism. Bertrand Russell discusses the problem of induction in "The Problems of Philosophy", Chapter 6, 1912 and "The History of Western Philosophy", Chapter 17, Hume, 1946. Karl Popper gives a satisfying and so far not refuted theoretical justification of our expecting further repetition of past events. This he does in among other places in "Objective Knowledge", Chapter 1, 1972.
The concept of normal values and their ranges has been challenged by Schwartz et al, 1973. It is erroneous to assume that if the normal range is the mean � 2 standard deviations that any value outside this range is importantly abnormal, i.e. that it leads to a decision being made. The probability of a particular value being importantly abnormal varies if the expectation of it being abnormal varies, i.e. pH values of 7.4 or 7.2 have different meaning in say a patient with chronic obstructive airways disease and in a patient in whom no pH abnormality is suspected.
A.3.2 Examination of "Buffering" Properties of HCO3-:H2CO2 System
Most texts state that the HCO3- : H2CO2 system is an efficient physiological buffer because the components of the pair are controlled separately (Pitts, 1974). As it is not a chemical buffer of any reasonable efficiency at the blood pH use of the term "buffer" in respect to HCO3- : H2CO2 action introduces considerable confusion. This is illustrated in the following example.
Plasma has a [HCO3-] of approximately 24meq/l and [H2CO2] of 1.2meq/l, hence:
If 1.2meq HCl is added to 1 litre of a solution of 24meq NaHCO2 in water, 1.2meq HCO3- will be converted to H2CO3, so:
If the H2CO3 is held constant then 1.2meq HCl causes the [HCO3- ] to diminish by 1.2meq but the H2CO3 level remains constant at 1.2meq/l, therefore:
7.38 is a trivial drop in pH whereas 7.08 is a large change.
Acting as a buffer the HCO3- : H2CO2 system would have allowed a marked change in pH. Chemically the main process which actually diminished the pH change when HCl was added was the removal of another weaker acid (H2CO3) which was formed when the H+ of the HCl was neutralised by HCO3-.
If base is added to the HCO3- : H2CO3 system at pH 7.4 the change in pH is even greater than when an acid is added. If 1.2meq NaOH is added, the [HCO3-] rises 1.2meq/l or the following equation moves to the right:
H2CO3 ↔ H+ + HCO3- + OH- ↔ H2O + HCO3-
therefore [H2CO3] falls to a very low level so that [HCO3-]/ [H2CO3] rises greatly, and therefore pH rises.
In vivo H2CO3 is prevented from falling so the ratio [HCO3-] / [H2CO3] changes only slightly when [HCO3-] rises following the addition of Na0H, therefore the pH changes only slightly. In other words, if a base is added to the blood the main reason the pH changes very slightly is that a weak acid (H2CO3) is added to keep the PaCO2 constant as the CO2 is used up in partly neutralizing the added base:
Na+ + OH- + CO2 → Na+ + HCO3-
Inspection of the Siggaard-Andersen nomogram shows greater changes in pH if the H2CO3 (PCO2) level in blood is altered, than if similar quantities of strong acid or base are added while the PCO2 is held at 40mmHg. Altering the level of H2CO3 by, say 1.2meq in blood, is not the same as adding 1.2meq H2CO3. In blood in vitro raising the level of H2CO3 1.2meq (80mmHg PCO2) causes the [HCO3-] to rise by 6meq/l (see Appendix 2.4.2 ) therefore in fact 7.2meq H2CO3 is added to blood if the PCO2 is raised from 40 to 80mmHg and the level of H2CO3 raised 1.2meq/l. Adding 7.2meq H2CO3 to blood in vitro therefore gives a pH of 7.08, whereas adding 7.2meq HCl when PaCO2 is 40mmHg gives a pH of 7.29.
A.3.2.1. Other Factors
Factors other than buffer action which diminish the pH change due to transfer of CO2 as blood passes through capilliaries
Except in the case of CO2, only a small quantity of acid or base is transferred during any one passage of blood through a capillary. When CO2 is transferred two mechanisms other than buffer action reduce pH fluctuations.
a) Reduced haemoglobin anion is a stronger base than oxygenated haemoglobin. Less haemoglobin exists in the ionised form when it is deoxygenated. During the process of deoxygenation CO2 is added. The resulting fall in pH is diminished by the reduced haemoglobin taking up more H+ ions than the oxygenated haemoglobin.
b) Some of the CO2 is carried directly combined with haemoglobin as carbamino-haemoglobin. This is more readily formed by reduced than by oxygenated haemoglobin.
A.3.4 Extra Discussion of Renal Control of pH
A knowledge of the conventional postulated mechanism of renal control of pH is assumed in this section (see Relman, 1968 and Pitts, 1974). Text book descriptions of renal control of pH (Pitts, 1974) are often not consistent with the overall ("black-box") function of the kidney in transferring acid or base from the blood to the urine. Briefly, Pitts and others state that pH of blood is controlled by the components of the H2CO3 : HCO3- buffer systems being under separate control, i.e. the respiratory system controls the [H2CO3] by controlling PaCO2 and the kidneys control the HCO3- level by controlling the renal threshold of HCO3-. The ratio [HCO3-] : [H2CO3] is, therefore, controlled and, thus, according to the Henderson-Hasselbalch equation pH is controlled. As the pH is controlled so are the ratios of all the other buffer pairs. In this section I attempt to show [HCO3-] is not a controlled variable but rather a dependent variable. It appears to me that pH itself is a controlled variable and [HCO3-] is dependent on pH and PCO2.
The conventional description of control of non-respiratory disturbances emphasizes the control of [HCO3-] (Pitts, 1974) rather than the excretion of acid (which must include an anion). It is axiomatic that if H+ is taken from blood and excreted, an anion must also be excreted. Exchange of H+ for Na+ cannot be the full explanation of excretion of H+ in the urine. If the blood pH rises by Na+ for H+ exchange the blood Na+ + HCO3- level would rise. This Na+ rise does not occur. Sodium would have to come from an external source. Excretion of H+ = excretion of acid, therefore acid level falls in the blood, and the level of anion of the particular acid must fall.
The extracellular fluid (E.C.F.) is the source of urinary acid. It consists of a mixture of equal numbers of anions and cations in solution. When the E.C.F. becomes more alkaline the total quantity of basic anion or base (HCO3-, OH- etc.) must increase. (This does not apply to the basic anion of a weak acid if that acid's removal is the cause of the rise in pH, e.g. in acute hypocapnia [HCO3-] falls although pH rises). When the pH rises due to loss of acid other than carbonic, H2CO2 dissociates to H+ and HCO3-. The H+ partially replaces the H+ of the acid which has been removed.
The HCO3- level in the blood cannot rise without changes in levels of other anions or cations in the blood. To do so without effects on Na+, Cl- or other ions would defy electrical neutrality, i.e. the blood would become negatively charged and the urine positive. If the HCO3- level in the blood rises either the cation level (Na+) must rise or the non-basic anion (Cl-) level fall. The Na+ level can be raised only by retention of exogenous Na+ if the E.C.F. volume does not fall. The non-basic anion level could fall by the kidneys excreting Cl- with the H+. This is the usual way the kidneys extract and excrete acid from the blood.
For some years the medical literature stated that Na+ and Cl- were not important in the control of blood pH because they are neither acids nor bases (Smith, 1951). It is now recognised that Cl- is vital in control of pH under some circumstances (i.e. when it is the anion of hydrochloric acid) (Schwartz et al, 1968). This should have been obvious from considerations of electricaly neutrality.
The kidney cannot increase the [HCO3-] in the blood only by reabsorption of HCO3- from the glomerular filtrate. Such action would leave the blood level unchanged if total reabsorption of HCO3- occurred. HCO3- must be generated and [Na+] or [Cl-] changes must occur as well.
The kidney is said to control pH by controlling the renal threshold of HCO3- (Pitts, 1974). In this explanation threshold of HCO3- rises if PaCO2 rises and vice versa. If this were the correct explanation the HCO3- level should fall during the recovery phase of chronic CO2 retention, because the HCO3- threshold should now fall. In fact the [HCO3- ] and the pH remain high until Cl- in some form is given (Polak et al, 1961).
If a K+ deficiency is induced in an animal without an accompanying acid-base disturbance and then K2S04 administered, the pH and [HCO3-] of blood rise (Bleich et al, 1966). This is impossible to explain if [HCO3-] is a controlled variable. The explanation is that S042- is not retained by the kidney. When it is excreted it must go with a cation. The kidney has Na+ , K+ or H+ to choose between. As the levels of Na+ and K+ are controlled, the excretion has to be as 2H+ + S042- or 2(NH4)+ + S042-. As the urine becomes acid the blood must become alkaline and the [HCO3-] in blood must rise.
The persisting high [HCO3-] due to loss of HCl in pyloric stenosis has been explained as a high threshold for HCO3- (Kassirer et al, (a) 1966). It would occur if HCl (i.e. gastric juice) was removed in a nephrectomised patient. The blood leaving the stomach would have a high pH while acid was being secreted. Therefore, the systemic blood pH would rise, and with it the [HCO3-] (Le Quesne, 1961). Administration of HCl would correct the high [HCO3-] without intervention by the kidneys (Bradham, 1968). Unless chloride ion in some form is given the pH and [HCO3-] cannot be corrected (Schwartz et al, 1968).
Conclusion
a) If it can be shown that the kidneys directly control all the electrolytes in the E.C.F. except HCO3- , then HCO3- level must be uncontrolled.
b) The kidneys generally can correct only states of excess. Deficiency states have to be corrected by exogenous supplies of raw materials. The kidney is able to correct high levels of Na+ , K+ ,Cl- or H+ but not HCO3- unless some acid is added or control of Na+ is lost (acetazolamide, Diamox). Most causes of non-respiratory alkalosis are acid deficiency states so the high [HCO3-] is incidental to the low [H+] state. The low acid state can be corrected only by giving acid or a salt from which acid can be generated. If the acid used has a non-resorbable anion, e.g. H2S04 or HN03, the correction of the alkalosis will be temporary (Tanner, Schwartz and Bleich, 1966).
Note: Since this section was written (1972) Knud Engel and Paul Kildeberg "Physiological Viewpoint on Clinical Acid-Base Diagnosis", (1977) have written along similar lines. They point out the absence of value in the concepts of bicarbonate reabsorption and renal bicarbonate threshold etc. I think this view is correct. They emphasize the distinction of metabolisable and non-metabolisable acids. I think it is more useful to think about organic and inorganic acids, although all organic acids may not be metabolisable.
The method of describing inputs and outputs which they recommend although consistent with their and my approach is unnecessarily complex in the clinical situation for which the method is advocated.
Schwartz and Cohen (1978) have reviewed several of the paradoxes inherent in the conventional description of renal acidification mechanisms. The approach used in this book is consistent with their hypotheses, i.e. that variation of acid secretion by the kidney occurs in a direction which would appear to be appropriate for pH homeostasis only if this is consistent with other homeostatic mechanisms, i.e. Na & K control. They think, and I agree, that pH control has a low priority in homeostatic mechanisms.
In contrast the conventional approach postulates a H+ excretion mechanism which is controlled by deviations of blood pH and that K+ is involved in the H+ excretion mechanism.
7. TREATMENT
7.1 RESPIRATORY ACIDOSIS
Respiratory Acidosis is corrected by increasing alveolar ventilation and/or treating the cause. Acute respiratory acidosis can probably be corrected rapidly (Prys-Roberts et al, 1967) but chronic respiratory acidosis should be corrected slowly. Rapid lowering of a high PaCO2 has been associated with fits and cardiovascular collapse.
If the patient has a compensatory disturbance (high non-respiratory pH and positive base excess) which has raised the pH towards normal, the kidneys have to correct this disturbance which is equivalent to a metabolic alkalosis. This compensatory metabolic alkalosis does not require treatment if the patient is given a mixed diet with adequate Cl- (Polak et al, 1961), and provided there are no complicating factors in its genesis or in the cardiovascular or renal systems. If the compensatory non-respiratory alkalosis is preventing the patient from lowering his PaCO2 the alkalosis may need to be treated by modifying the steroid or diuretic therapy. Acetazolamide (Diamox) can be used to correct the metabolic alkalosis. It differs from other diuretics in producing an acidosis by the loss of Na+ HCO3- in the urine. When this drug causes loss of base in the urine the sodium content of the E.C.F. must fall. Although acetazolamide will produce a lessening of alkalosis by loss of base, rather than the physiological mechanism of acid retention, the coincidental loss of Na+ may be a benefit if the patient has any cardiac failure.
7.2 RESPIRATORY ALKALOSIS
7.2.1. Treat the cause
e.g. correct hypoxia or shock if they are causing hyperventilation.
7.2.2. Elevate the PaCO2.
This can be corrected by administering CO2, increasing the dead space or lowering the minute ventilation. These measures will rarely be thought to be necessary.
7.3 METABOLIC ACIDOSIS
7.3.1 Treat the cause.
Stop alimentary loss of base; correct hypoxia; reduce renal acid load by diet; drain abscess in diabetic ketosis and give insulin (see 7.3.2.2.3, ketoacidosis) ; treat shock with intra-venous fluids and stop haemorrhage etc (see 7.3.2.2.2.2, shock) .
7.3.2. Correction of Acidosis
7.3.2.1. Administration of NaCl.
If the acidosis is (a) not affecting the cardiac action and (b) renal function is adequate, the acidosis may be corrected by giving sufficient NaCl (Na+ + Cl-) solution for the kidney to (i) correct the acidosis by excreting HCl (H+ + Cl-) or NH4Cl (NH4+ + Cl-) and (ii) repair any deficit in E.C.F. volume. This approach applies in alimentary causes of metabolic acidosis where the kidneys are usually able to correct the defects if enough saline is given (Hesse et al, 1966). Correction may be more rapid if Hartmann's solution rather than 0.9% NaCl solution is given to correct the pH disturbance as there is less for kidney to do. The lactate ion has to be converted to HCO3- and some H+ + Cl- will have to be excreted but not as much as with NaCl solution.
7.3.2.2. Administration of Base.
Indications for direct correction of acidosis by giving base:
7.3.2.2.1. The cause cannot be corrected. e.g. renal acidosis, where the kidneys fail to excrete inorganic acid (an end product of protein metabolism). If this defect is the sole manifestation of renal impairment (i.e. renal tubulcar acidosis), it is rational to neutralise the acid with NaHCO2 which can be given by mouth. In most instances renal failure is not manifest solely by acidosis. Usually dialysis or transplantation is necessary to correct the multiple effects of renal failure which include acidosis.
7.3.2.2.2. Where the acidosis is depressing the circulation (i.e. to break the viscious circle of myocardial depression which aggravates acidosis). This is the indication in cardiac arrest (Chazan et al, 1968) or in shock.
7.3.2.2.2.1. Cardiac Arrest. In cardiac arrest, acute lactic acidosis, it is said, may prevent the circulation restarting. NaHCO2 is often given in an empirical and probably excessive dose of "1 bottle" (i.e. 500mls 4.2% NaHC03 = 250meq). PaCO2 rises as some HCO3- is converted to H2CO2. Serum [Na+] will rise and circulatory overload may be caused by the Na+ load. After recovery a metabolic alkalosis will occur until the Na+ + HCO3-are excreted.
There is controversy surrounding the use of base in cardiac arrest (Stackpool 1986, Narins and Cohen 1987), as in the emergency situation it is impractical to obtain biochemical evidence before treatment (Leading Article, Lancet, 1976). Non-respiratory acidosis does not occur in all cases of cardiac arrest (Stewart, 1964, Chazan et al, 1968). Inadequate ventilation causing a high PCO2 is as frequent a cause of low pH in cardiac arrest as is metabolic acidosis. There is no reason for thinking that the low pH due to CO2 has much cardiovascular depressant effect (Schultz et al, 1960; Prys-Roberts et al, 1967; Gerst et al, 1964).
If the circulation is restored rapidly after resuscitation is commenced (i.e. pulse and/or consciousness) than any NaHCO2 infusion which may have been started should be stopped until the situation is assessed biochemically (Rackwitz et al, 1976).
Mattar et al, 1974, investigated 12 patients in whom NaHCO3 was administered during cardiac arrest resuscitation. Full measurements were not available in all cases. The doses of NaHCO3 were between 45 and 270meq. The mean changes were:
before NaHCO3
after NaHCO3
Osmolality ( mOsmol/litre)
302
377
pH
7.34
7.54
HCO2 (mEq/litre)
21
50
Na (mEq/litre
138
170
It has also been shown (Fillmore et al, 1970) that even if it is possible to effectively adjust pH abnormalities during resuscitation by use of alkali, restarting heart action still may not occur. Presumably this lack of restarting was be due to overwhelming myocardial damage or inadequate circulation provided by the cardiac massage. This was suggested by the rising lactic acid level which was probably an effect rather than the cause of continuing inadequate heart action.
It is still common practice to administer NaHCO3 during cardiac arrest. In some institutions NaHCO3 infusion is left set up on each cardiac arrest trolley to save time in starting its administration.
The dose required to correct a low non-respiratory pH is arbitrary in the individual patient. I would advise that a base-line specimen of blood should be taken before the NaHCO3 administration, if this is possible. Venous blood from a central vein if one (e.g. the internal jugular) is being used to administer drugs is satisfactory, but arterial might be better. Follow up blood measurements would then be done as resuscitation continues and is completed, to examine pH, osmolality, Na, K , PCO2, PO2 and lactate levels.
In summary I would conclude that correction of non-respiratory pH during cardiac resuscitation is not as important as once was thought, and that monitoring of biochemical changes before and after such correction should be routine.
7.3.2.2.2.2. Shock. In shock accompanied by acidosis it is rational to administer base together with other haemodynamic management, i.e. raising the C.V.P., giving inotropic agents, oxygen, etc. (Manger et al; 1962, Ledingham, 1962; MacKenzie, 1965). When base is given in shock it is not rational to give it over some hours after having decided on a dose. The acidosis should be corrected as quickly as possible in two or three steps controlled by non-respiratory pH, base excess or standard bicarbonate measurements, e.g. first dose (7.4 - non-respiratory pH) x body weight (kg) x 7meq or - (0.1 x base excess x body weight) meq. Subsequent doses are estimated after the effect of the first dose has been observed. Dosage of NaHCO2 (or any other intravenously administered electrolytes) cannot be determined by formulae. Response to initial dosage will suggest magnitude of subsequent doses. Factors which would have to be taken into account if dosage was to be predicted would have to include cause of acidosis, the circulatory state and the magnitude of the acidosis. The degree of acidosis alters the requirements of NaHCO2 by a function which is not a direct proportion (Garella et al, 1973).
7.3.2.2.3. Diabetic Keto-Acidosis. The use of HCO3- in diabetic keto-acidosis is also controversial. Acidosis has been claimed to be a cause of insulin resistance in diabetic acidosis (Walker et al, 1963). This study was not controlled. Regimes including HC03 have been advocated without conclusive evidence of benefit (Solar et al, 1973; Solar et al, 1974). Administering NaHC03 in diabetic keto-acidosis may exaggerate changes in serum K+ particularly if this is changing rapidly following high doses of insulin.
It now obvious that many of the problems of managing diabetic acidosis (i.e. ketotic and non-ketotic) were iatrogenic, i.e. due to large intermittent I.V. and I.M. dosage of insulin, rapid and erratic alterations in blood pH, K+ and glucose levels. Diabetic acidosis can be corrected over a period of about 6 hours with low dose I.V. infusions of insulin. There is steady biochemical and clinical improvement without swings in blood glucose or K+ levels. Only small doses of K+ are necessary and in most instances HCO3- is not given, although when it has been, it has apparently not produced adverse effects (Alberti et al, 1973; Page et al, 1974; Kidson et al, 1974; Semple et al, 1974; King et al, 1974 and Shaw et al, 1974). If NaHC03 is to be used in diabetic acidosis its use should probably be limited to patients with severe acidosis (pH<7 100meq="" 70kg="" a="" and="" be="" dose="" font="" greater="" in="" initial="" not="" patient.="" probably="" should="" than="" the="">
The low PaCO2 which is present in the initial phases of diabetic acidosis may persist after the blood pH has returned to normal. This is probably not due to delay in return of the CSF pH to normal (King et al, 1974).
7.3.2.2.4 Neonatal resuscitation. NaHCO2 is used in neonatal resuscitation (Clark et al, 1968). The use is analogous to its use in cardiac arrest. The need or efficacy of the treatment as a routine has not been established. When used, pH electrolyte studies should be done, at least retrospectively. Hypernatraemia and intracranial haemorrhage have been associated with administration of NaHCO2 in the neonatal period (Simmon et al, 1974; Volpe, 1974).
7.3.2.2.5 Neonatal Respiratory Distress Syndrome (RDS). NaHCO2 has been used to correct the "chronic" acidosis of this syndrome (Usher, 1963). The acidosis is presumably due to hypoxia and correction of this (Daily et al, 1971; Smith and Daily, 1971) if possible, would be more rational (Dell and Winters, 1972). In other cases CO2 retention may be the cause of low pH. In these cases increased ventilation would be the rational treatment. (Ostrea et al, 1976). When it has been given in the patients with RDS to correct acidosis, NaHCO2 gives an acute rise in PaCO2. Tham causes an acute fall in PaCO2. These changes in PaCO2 last for some minutes (Baum et al, 1975).
7.3.2.3 Complications of NaHCO2 Therapy
Metabolic alkalosis which may then cause respiratory depression.
Hypernatraemia and hyperosmolality (Mattar et al, 1974; Bishop et al,1976).
Changes in other electrolytes especially lowering of serum [K+].
Fluid retention in patients who have disorders which will lead to fluid retention if excess Na+ is given.
Acute rise in PaCO2 due to neutralization of HCO3- (Singer et al, 1956; Ostrea and Odell, 1972; Baum et al, 1975).
Intracranial haemorrhage (Simmons et al, 1974; Volpe, 1974; Papil et al, 1978).
7.3.2.4 Tham
Tham (Nahas, 1961) is an organic base (often referred to as a buffer) used to correct acidosis. It has no obvious advantages over sodium bicarbonate. Its claimed advantages include ability to correct intracellular acidosis. A separate syndrome of intracellular acidosis is not clinically recognised. Its main value might be in situations in which Na+ load may be undesirable, e.g. cardiogenic shock. Although, as Tham is an osmotically active agent it may have similar effects on extracellular fluid volume as Na+. It might be of use if one wished to correct the low pH of respiratory acidosis directly (Manfredi et al, 1960), as some have claimed (Mithoefer et al, 1965 and 1968) that bronchospasm is relieved by direct correction of low pH. Its main disadvantages are inconvenience of preparation (it is supplied in a powder) and that it causes respiratory depression. Respiratory depression would presumably be caused by any substance which raised blood pH including Na0H or NaHCO2. Such depression could be managed by intermittent positive pressure ventilation so is not in itself an absolute reason for not using Tham.
7.4 METABOLIC ALKALOSIS
7.4.1 Remove the cause (e.g. relieve pyloric obstruction or modify diuretic regime).
7.4.2.1 Administration of NaCl. Ingestion or injection of sufficient sodium chloride solution for the kidney to correct the alkalosis by excretion of Na+ + HCO3- .
7.4.2.2 Administration of Acid. Direct correction of alkalosis with ammonium chloride or hydrochloridc acid solution, infusion or ingestion (Bradham, 1968, Leading Article (a) 1974, Sanderson, 1974; Pain et al, 1974; Abouna, 1974; Harken et al, 1975 and Worthley, 1977). This is indicated only if the alkalosis is very severe or renal or cardiac function are poor. Usually there is an associated reduction of extracellular volume so some Na has to be given in the form of NaCl. Compensation for non-respiratory alkalosis is CO2 retention achieved by hypoventilation. The hypoventilation may result in hypoxaemia which may necessitate oxygen therapy.
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