Critical Care volume 4, Article number: 6 [2000] Cite this article
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Abstract
An advanced understanding of acid–base physiology is as central to the practice of critical care medicine, as are an understanding of cardiac and pulmonary physiology. Intensivists spend much of their time managing problems related to fluids, electrolytes, and blood pH. Recent advances in the understanding of acid–base physiology have occurred as the result of the application of basic physical-chemical principles of aqueous solutions to blood plasma. This analysis has revealed three independent variables that regulate pH in blood plasma. These variables are carbon dioxide, relative electrolyte concentrations, and total weak acid concentrations. All changes in blood pH, in health and in disease, occur through changes in these three variables. Clinical implications for these findings are also discussed.
Introduction
Whereas most medical and surgical subspecialists concern themselves with a specific organ [eg nephrology], region of the body [eg cardiothoracic surgery], or disease process [eg infectious disease], practitioners of critical care are more often concerned with the interaction of various organs and disease states. As such, our practice is often defined by certain syndromes [sepsis, multiorgan failure] and pathophysiologic conditions [shock] that do not confine themselves to the domains of a single subspecialty. Acid–base regulation is one of these 'areas' of medicine that crosses organ-specific boundaries, and the intensive care unit is often the place where severe derangements in this area exist. For these reasons, intensivists, and others called upon to care for critically ill patients in the intensive care unit, operating room, or emergency department, are expected to diagnose and manage complicated disorders of acid–base balance. This review provides a rather in-depth examination of chemistry and physiology of acid–base balance in health and disease.
The concentration of H+ in blood plasma and various other body solutions is among the most tightly regulated variables in human physiology. [Most of the principles discussed in this review are applicable to animal physiology as well. A complete discussion of the differences between species, however, particularly aquatic versus terrestrial species, is beyond the scope of the present review.] Acute changes in blood pH induce powerful regulatory effects at the level of the cell, organ, and organism [1]. The mechanisms responsible for local, regional, and systemic acid–base balance are incompletely understood though, and controversy exists in the literature as to what methods should be used to understand these mechanisms [2]. Much of this controversy exists only because the strict rules for causation [as opposed to correlation] have not often been applied to the understanding of acid–base balance, and methods that are useful clinically have often been used to understand physiology without being subjected to appropriate scientific rigor. The use of various laboratory variables to diagnose an acid–base disorder is analogous to the use of the electrocardiogram to diagnose a myocardial infarction. However, neither the changes in the electrocardiogram tracing nor the disturbances in electrical conduction that these changes reflect were ever considered to be the cause of a myocardial infarction. In contrast, changes in HCO3- [bicarbonate]concentration, for example, have been assumed to be responsible for metabolic acidosis or alkalosis. Failure to establish causation has lead to numerous incorrect notions of acid–base physiology and has fueled years of, often heated, debate [2,3,4]. This review analyzes what is known about the causal relationships between acid–base variables and acid–base balance in health and disease.
Quantification, classification, and causation
In order to understand acid–base physiology, we must first agree on how to describe and measure it. Since Sörensen [1868–1939] first introduced the pH notation, we have used the pH scale to quantify acid–base balance The pH scale has a tremendous advantage because it lends itself to colorimetric and electrometric techniques. There is also some physiologic relevance to the logarithmic pH scale [5]. pH is a confusing variable, however. It is a nonlinear transformation of H+ concentration – the logarithm of its reciprocal. Strictly speaking, pH can only be thought of as a dimensionless representation of H+ concentration and is not, itself, a concentration. Indeed, pH is actually the logarithmic measure of the volume required to contain 1 mol/l of H+. In blood plasma at pH 7.4, this volume is approximately 25 million liters [6].
Regardless of how we express the concentration of H +, either directly or as the pH, it is generally accepted that changes in blood H +concentration occur as the result of changes in volatile [partial carbon dioxide tension [pCO2]] and nonvolatile acids [hydrochloric, sulfuric, lactic, etc]. Clinically, we refer to changes in volatile acids as 'respiratory' and changes in nonvolatile acids as 'metabolic'. There are three major methods of quantifying [describing] acid–base disorders, and each differs only in assessment of this latter, 'metabolic' component. These three methods quantify the metabolic component either by using HCO3 - [in the context of pCO2], the standard base excess [SBE],or the strong ion difference [SID]. Although there has been significant debate regarding the accuracy and utility of each method compared with the others, all three yield virtually identical results when used to quantify the acid–base status of a given blood sample [7,8]. The only differences between these three approaches are conceptual [ie in how they approach the understanding of mechanisms] [9,10,11].
Beyond Henderson and Hasselbalch
Since Hasselbalch adapted the Henderson equation to the pH notation of Sörenson, we have used the following equation to understand the relationship between respiratory and metabolic acid–base variables:
pH = pK × log [HCO3 /[0.03 × pCO2]] [1]
This is the Henderson–Hasselbalch equation, and it is important to realize what this equation tells us. Anincrease in pCO2 will result in a decrease in the pH and anincrease in the HCO3- concentration. Thus, a patientfound to have a low blood pH, a condition known asacidemia, will either have an increased pCO2 or a pCO2that is 'not increased'. In the former circumstance, we classify the disorder as a 'respiratory acidosis'. We use the term 'acidosis' to describe the process resulting in acidemia and 'respiratory' because the apparent cause is an increased pCO2. This is logical, because carbonic acid results when CO2 is added to water [or blood], and the resultant decrease in pH is entirely expected. In the latter condition pCO2 is not increased, and thus there cannot be a respiratory acidosis. We therefore refer to this condition as 'metabolic' because some nonvolatile acid must be the cause of the acidemia. We can reverse the above logic and easily classify simple conditions of alkalemia as either resulting from respiratory or metabolic alkaloses. Thus, equation 1 allows us to classify disorders according to the primary type of acid being increased or decreased. Over time physiology superimposes its effects on simple chemistry and the relationship between pCO2 and HCO3- is altered in order to reduce the alterations in pH. By carefully examining the changes that occur in pCO2 and HCO3- in relationship to each, however, one can discernhighly conserved patterns. In this way rules can be established to allow one to discover mixed disorders and to separate chronic from acute respiratory derangements. For example one such rule is the convenient formula for predicting the expected pCO2 in the setting of a metabolic acidosis [12]:
pCO2 = [1.5 × HCO3-] + 8 ± 5 [2]
This rule tells us what the pCO2 should be secondary to the increase in alveolar ventilation that accompanies a metabolic acidosis. If pCO2 does not change enough or changes too much, we classify the condition as a 'mixed' disorder, with either a respiratory acidosis if the pCO2 is still too high, or a respiratory alkalosis if the change is too great. This rule, along with others [Table 1] has been recently translated to SBE terminology [7]:
pCO2 = [40 + SBE] ± 5 [3]
For example, consider the following arterial blood gas sample: pH7.31, pCO2 31, HCO3- 15, SBE-9.5. Equation 2 tells us that the expected pCO2 =[1.5×15]+8 ± 5=30.5 ± 5, and in Equation 3 the SBE added to 40 also yields 30.5. The measured pCO2 of 31 mmHg is consistent with a pure metabolic acidosis [ie no respiratory disorder].
It is also very important to understand what the Henderson-Hasselbach equation does not tell us. First, it does not allow us to discern the severity [quantity] of the metabolic derangement in a manner analogous to the respiratory component. For example, when there is a respiratoryacidosis, the increase in the pCO2 quantifies the derangement even when there is a mixed disorder. However, the metabolic component can only be approximated by the change in HCO3-. Second, Equation 1 does not tell us about any acids other than carbonic acid. The relationship between CO2 and HCO3- provides a useful clinical 'roadmap' to guide the clinician in uncovering the etiology of an acid–base disorder as described above. The total CO2concentration, and hence the HCO3- concentration, is determined by the pCO2, however, which is in turn determined by the balance between alveolar ventilation and CO2 production. HCO3- cannot be regulated independent of pCO2. The HCO3- concentration in the plasma will always increase as the pCO2 increases, but this is not an alkalosis. To understand how the pH and HCO3- concentration are altered independent of pCO2, we must look beyond Henderson and Hasselbach.
Base excess
In order to address the first 'shortcoming' of the Henderson-Hasselbach equation – the inability to quantify the metabolic component – several methods have been devised. In 1948, Singer and Hastings proposed the term 'buffer base' to define the sum of HCO3- plus the nonvolatile weak acid buffers [A-] [13]. A change in buffer base corresponds to a change in the metabolic component. The methods for calculating the change in buffer base were later refined by investigators [14,15] and refined further by others [16,17,18] to yield the base excess methodology. Base excess is the quantity of metabolic acidosis or alkalosis, defined as the amount of acid or base that must be added to a sample of whole blood in vitro in order to restore the pH of the sample to 7.40 while the pCO2 is held at 40 mmHg [15]. Although this calculation is quite accurate in vitro, inaccuracy exists when applied in vivo in that base excess changes with changes in pCO2 [19,20]. This effect is understood to be due to equilibration across the entire extracellular fluid space [whole blood + interstitial fluid]. When the base excess equation is modified to account for an 'average' content of hemoglobin across this entire space, a value of 5 g/dl is used instead, and this defines the SBE. It should be pointed out that this value does not represent the true content of hemoglobin suspended in the volume of whole blood together with interstitial fluid, but rather is an empiric estimate that improves the accuracy of the base excess. It can be argued that the entire extracellular fluid space is involved in acid–base balance, because this fluid flows through blood vessles and lymphatics, mixing constantly [21]. Thus, the value of SBE is that it quantifies the change in metabolic acid–base status in vivo. It is of interest that base excess is only accurate in vivo when it assumes a constant hemoglobin concentration.
However, the base excess approach does not address the second problem associated with using the Henderson-Hasselbach equation alone [ie it still does not tell us about the mechanisms of metabolic acid–base balance]. For example the body does not 'regulate' the SBE. It is not a substance that can be excreted in the feces or reabsorbed from the proximal tubule. Similarly, HCO3 - is not a strong acid or base and its addition to or removal from the plasma cannot be translated into changes in SBE. This not to say that changes in SBE and HCO3- do not correlate closely, because they do. However, correlation and causation are not the same thing. The difference has traditionally been ascribed to the effects of 'buffering', the argument being that stong acid [or base], quantified by SBE, is 'buffered' by plasma proteins, hemoglobin, and finally by HCO3-. The resulting changes in HCO3-and pH are then a result of this buffering process. These buffers are actually weak acids, however, and their addition to plamsa both decreases the pH and increases the responsiveness to pCO2 [Fig. 1]. Furthermore, as explained by Stewart [6,9], the fundamental physical-chemical properties of biologic solutions dictate much of this so-called 'buffering'.
Table 1 Observational acid–base patterns
Full size table
Physical-chemical properties of biologic solutions
A physical–chemical analysis of acid–base physiology requires the application of two basic principles. The first is electroneutrality, which dictates that, in aqueous solutions, the sum of all positively charged ions must equal the sum of all negatively charged ions. The second is conservation of mass, which means that the amount of a substance remains constant unless it is added to or generated, or removed or destroyed. These principles may seem very basic indeed, but they are often overlooked in the analysis of clinical acid–base physiology, leading to incorrect conclusions. For example, a hyperchloremic metabolic acidosis can only be brought about in two ways. First, Cl- ions can be added to the circulation, either via an exogenous source [eg HCl or saline] or via internal shifts [eg from the red cell]. Second, Cl- ions can be retained or reabsorbed, whereas water and other ions [ie Na+] are excreted so that the relative concentration of Cl- increases. A decrease in - HCO3 or H+ concentration does not produce hyperchloremia, but rather hyperchloremia is a cause of acidosis. This distinction is not merely semantics, any more than Copernicus' observation that the Earth, rather than the sun moves [11,22].
In addition to these physical–chemical principles, almost all solutions of biologic interest share two important characteristics. First, virtually all are aqueous [composed of water], and second, most are alkaline [OH-concentration >H+ concentration]. Because these characteristics are so universal in human physiology, they are often ignored in reviews of physiology, especially for clinical medicine, but they are extremely important. Aqueous solutions contain a virtually inexhaustible source of H+. Although pure water dissociates only slightly into H+ and OH-, electrolytes and CO2 produce powerful electrochemical forces that influence water dissociation. Similarly, aqueous solutions that are alkaline behave very differently compared with acidic solutions in terms of the extent to which changes in their composition influence changes in pH.
To illustrate this point, first consider a 1 l solution of pure water. Pure water contains only a small amount of H+ and OH- ions and molecular H2O. Pure water is a neutral solution by definition, because the H+ and OH- concentrations are equal. The concentration of these ions is determined solely by the extent to which water dissociates and can be defined by a constant, K'w. Water dissociation is temperature sensitive because K'w is, but, at all times, the concentrations of H+ and OH- must be equal, and H+× OH- = K'w. If we add 10 mmol/l of each Na+ and Cl- to this 1 lsolution of pure water, we would have an aqueous solution that contains H+, OH-, Na+ and Cl- ions, and molecular water. Of note, the solution does not contain any molecules of NaOH, HCl, or NaCl, because both Na+ and Cl- are strong ions and as such are completely dissociated in water. The solution we now have is still a neutral solution by definition, and at 25ºC the concentrations of both H+ and OH- are approximately 100 nmol/l, or pH7.0. If we then add 10 mmol/l HCl, we will have a solution that contains 10 mmol/l Na+ and 20 mmol/l Cl-. This solution is acidic: OH- =4.4 × 10-9 nmol/l, and H+ = ~ 10mmol/l. Note that in this acidic solution the H+ concentration increased by the amount of H+ added [ie 10 mmol/l]. However, if we were to add 10 mmol/l NaOH instead of HCl, the solution would instead contain 20 mmol/l Na+ and 10 mmol/l Cl-, and would be alkaline: H+ = 4.4 × 10-9 nmol/l and OH- = ~ 10 mmol/l. If we then add 5 mmol/l HCl to this alkaline solution, the resulting concentration of Na+ would be 20 mmol/l and of Cl- would be 15 mmol/l. The final H+ concentration is now 8.8 × 10-9 nmol/l and OH- is approximately 5 mmol/l. Note that in this final example 5 mmol/l of H+ were added to the solution, yet the final concentration of free H+ changed by less than billionth of this amount. It should be further noted that the solution I have described contains no 'buffers', and thus what is often attributed to the power of buffering systems is merely a physical–chemical property of alkaline solutions.
Figure 1
Changes in the relationship between partial carbon dioxide tension[pCO2] and H+ concentration as function of changes in'buffer' strength. Individual curves are drawn for varyingconcentration of total nonvolatile buffers in mmol/l. Note that as theconcentration of 'buffer' increases, the slope of the curveincreases, making changes in H+ concentration more responsive tochanges in CO2.
Full size image
Determinants of hydrogen concentration
From the preceding discussion it is apparent that, for aqueous solutions, water is the primary source of H+, and the determinants of H+concentration are the determinants of water dissociation. Fortunately, even for an aqueous solution as complex as blood plasma, there are but three independent variables that determine H+ concentration. Please note that I use the term 'determine' rather than 'describe' because, as shown by Stewart [6,9], these three variables are mathematically independent determinants of the H+ concentration. Thus, these variables are causally related to the H+ concentration rather than being merely correlated. The distinction between independent and dependent, between causation and correlation, is as important to acid–base physiology as to any other area of science. Only by the careful analysis of causal variables can mechanisms be determined. For blood plasma, these three variables are pCO2, SID, and the total weak acid concentration [ATOT].
Carbon dioxide
CO2 is an independent determinant of pH and is produced by cellular metabolism or by the titration of HCO3 - by metabolic acids. Normally, alveolar ventilation is adjusted to maintain the arterial pCO2 between 35 and 45 mmHg. When alveolar ventilation is increased or decreased out of proportion to pCO2 production, a respiratory acid–base disorder exists. CO2 production by the body [at 220 ml/min] is equal to 15000 mmol/day of carbonic acid [23]. This compares with less than 500 mmol/day for all nonrespiratory acids. The respiratory center, in response to signals from pCO2, pH, and partial oxygen tension, as well as some from exercise, anxiety, wakefulness, and others, controls alveolar ventilation. A precise match of alveolar ventilation to metabolic CO2 production attains the normal arterial pCO2 of 40 mmHg. Arterial pCO2 is adjusted by the respiratory center in response to altered arterial pH produced by metabolic acidosis or alkalosis in predictable ways.
When CO2 elimination is inadequate relative to the rate of tissue production, pCO2 will increase to a new steadystate that is determined by the new relationship between alveolar ventilation and CO2 production. Acutely, this increase in pCO2will increase both the H+ and the HCO3- concentrations according to the Henderson-Hasselbach equation [Equation 1]. Thus, this change in HCO3- concentration is mediated by chemical equilibrium, and not by any systemic adaptation. Similarly, this increased HCO3- concentration does not 'buffer' H+ concentration. There is no change in the SBE. Tissue acidosis always occurs in respiratory acidosis, because CO2 diffuses into the tissues. If the pCO2 remains increased the body will attempt to compensate by altering another independent determinant of pH, namely the SID.
Electrolytes [strong ions]
Blood plasma contains numerous ions. These ions can be classified both by charge, positive 'cations' and negative 'anions', as well as by their tendency to dissociate in aqueous solutions. Some ions are completely dissociated in water, for example, Na+, K+, Ca2+, Mg2+, and Cl-. These ions are called 'strong ions' to distinguish them from 'weak ions' [eg albumin, phosphate and HCO3-], which can exist both as charged [dissociated] and uncharged forms. Certain ions such as lactate are so nearly completely dissociated that they may be considered strong ions under physiologic conditions. In a neutral salt solution containing only water and NaCl, the sum of strong cations [Na+] minus the sum of strong anions [Cl-] is zero [ie Na+ = Cl-]. In blood plasma, however, strong cations [mainly Na+] outnumber strong anions [mainly Cl-]. The difference between the sum of all strong cations and all strong anions is known as the SID. SID has a powerful electrochemical effect on water dissociation, and hence on H+ concentration. As SID becomes more positive, H+, a 'weak' cation, decreases [and pH increases] in order to maintain electrical neutrality [Fig. 2].
In healthy humans, the plasma SID is between 40 and 42 mmol/l, although it is often quite different in critically ill patients. According to the principle of electrical neutrality, blood plasma cannot be charged, so the remaining negative charges balancing the SID come from CO2 and theweak acids [A-] and, to very small extent, from OH-. At physiologic pH, the contribution of OH- is so small [nmol range] that it can be ignored. The total weak acid concentration [mainly albumin and phosphate] can be considered together and for convenience is abbreviated ATOT, where AH + A- = ATOT. The SID of a blood sample can be estimated from the value of the remaining negative charge, because SID-[CO2 +A -]=0. This estimate of SID has been termed the 'effective' SID [24], but it is really no different from buffer base described over half a century ago [13]. Thus SID and buffer base are mirror images of each other. Recall that SBE is essentially the change in buffer base in vivo, and hence SBE defines the change in SID from the equilibrium point where pH=7.4 and pCO2 = 40 mmHg [8].
An alternative estimate of SID is as follows: [Na+ + K+ + Ca2++ Mg2+] – [Cl- + lactate-]. This is often referred to as the 'apparent' SID with the understanding that some 'unmeasured' ions might also be present [24]. Neither effective SID nor the apparent SID are perfect estimates of the true SID. Blood samples from patients may contain unmeasured strong ions [eg sulfate, ketones] making the apparent SID an inaccurate estimate of SID. Similarly, these patients may have abnormal weak ions [eg proteins] that will make the effective SID inaccurate. In healthly humans, however, the apparent SID and the effective SID are nearly identical, and are thus valid estimates of SID [24]. Furthermore, when the apparent SID and the effective SID are not equal, a condition we have referred to as the strong ion gap [SIG], where apparent SID – effective SID = SIG, abnormal strong and/or weak ions must be present [25]. The SIG is positive when unmeasured anions exceed unmeasured cations, and negative when unmeasured cations exceed unmeasured anions. Unexplained anions, and in some cases cations, have been found in the circulation of patients with a variety of diseases [25,26,27,28] and in animals under experimental conditions [29].
The SIG is not the same as the anion gap [AG]. Normally, the SIG is zero, whereas the AG is 8–12mmol/l. The AG is an estimate of the sum of SIG + A-. Thus, subtracting A- from the AG approximates the SIG. A convenient and reasonably accurate way to estimate A- is to use the following formula [30]:
2 [albumin g/dl] + 0.5 [phosphate mg/dl] [4]
or for international units:
0.2 [albumin g/l] + 1.5 [phosphate mmol/l] [5]
Note that the 'normal' AG for a person with no unmeasured anions or cations in their plasma is equal to A-, such that AG - A- = SIG = 0. This technique allows one to 'calibrate' the AG for patients with abnormal albumin and/or phosphate concentrations.
Figure 2
Plot of pH versus strong ion difference [SID]. For this plot,total weak acid concentration [ATOT] and partial carbon dioxidetension [pCO2] were held constant at 18 mmol/l and 40 mmHg,respectively. Assumes a water dissociation constant for blood of 4.4 ×10–14 [mol/l]. Note how steep the pH curve becomes at SID