Is Mixing An Acid With A Base A Physical Change

Crit Care. 2000; 4(1): 6–xiv.

Determinants of blood pH in health and disease

John A Kellum

¹Academy of Pittsburgh Medical Center, Pittsburgh, Pennsylvania, USA.

Received 1999 November 29; Accepted 1999 Nov xxx.

Abstract

An advanced understanding of acid–base of operations physiology is as central to the exercise of disquisitional intendance medicine, as are an understanding of cardiac and pulmonary physiology. Intensivists spend much of their fourth dimension 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 claret plasma. This analysis has revealed three independent variables that regulate pH in blood plasma. These variables are carbon dioxide, relative electrolyte concentrations, and full weak acrid concentrations. All changes in blood pH, in health and in illness, occur through changes in these iii variables. Clinical implications for these findings are as well discussed.

Keywords: Acid–base of operations balance, acidosis, alkalosis, anion gap, arterial blood gases, stiff ion difference, strong ion gap

Introduction

Whereas well-nigh medical and surgical subspecialists business themselves with a specific organ (eg nephrology), region of the body (eg cardiothoracic surgery), or disease procedure (eg infectious disease), practitioners of critical intendance are more oftentimes concerned with the interaction of various organs and disease states. As such, our practice is oftentimes defined by certain syndromes (sepsis, multiorgan failure) and pathophysiologic atmospheric condition (shock) that do not confine themselves to the domains of a single subspecialty. Acrid–base of operations regulation is one of these 'areas' of medicine that crosses organ-specific boundaries, and the intensive care unit of measurement is often the place where severe derangements in this area exist. For these reasons, intensivists, and others called upon to intendance for critically ill patients in the intensive care unit, operating room, or emergency department, are expected to diagnose and manage complicated disorders of acrid–base of operations residual. This review provides a rather in-depth exam of chemistry and physiology of acid–base residuum in health and illness.

The concentration of H⁺ in blood plasma and diverse other body solutions is amongst the near tightly regulated variables in homo physiology. (Most of the principles discussed in this review are applicative to animal physiology also. A complete discussion of the differences between species, notwithstanding, particularly aquatic versus terrestrial species, is beyond the scope of the present review.) Acute changes in claret 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 every bit to what methods should exist used to sympathise these mechanisms [2]. Much of this controversy exists only considering the strict rules for causation (equally opposed to correlation) have not frequently been practical to the understanding of acrid–base residual, and methods that are useful clinically have frequently been used to empathise physiology without being subjected to appropriate scientific rigor. The use of various laboratory variables to diagnose an acid–base of operations disorder is analogous to the employ of the electrocardiogram to diagnose a myocardial infarction. Nevertheless, neither the changes in the electrocardiogram tracing nor the disturbances in electrical conduction that these changes reflect were ever considered to be the crusade of a myocardial infarction. In contrast, changes in HCO_three ^- (bicarbonate)concentration, for example, have been assumed to be responsible for metabolic acidosis or alkalosis. Failure to found causation has lead to numerous incorrect notions of acrid–base of operations physiology and has fueled years of, oft heated, contend [2,iii,4]. This review analyzes what is known about the causal relationships between acid–base variables and acrid–base of operations balance in health and disease.

Quantification, classification, and causation

In order to sympathise acid–base of operations physiology, we must kickoff agree on how to depict and measure out it. Since Sörensen (1868–1939) first introduced the pH notation, we have used the pH calibration to quantify acid–base residuum The pH scale has a tremendous reward considering it lends itself to colorimetric and electrometric techniques. At that place is likewise some physiologic relevance to the logarithmic pH calibration [5]. pH is a disruptive variable, however. It is a nonlinear transformation of H⁺ concentration – the logarithm of its reciprocal. Strictly speaking, pH tin only be thought of as a dimensionless representation of H⁺ concentration and is non, itself, a concentration. Indeed, pH is actually the logarithmic measure out of the volume required to contain ane mol/l of H⁺. In blood plasma at pH 7.4, this book is approximately 25 million liters [6].

Regardless of how we limited the concentration of H ⁺, either directly or as the pH, information technology is by and large accepted that changes in blood H ⁺concentration occur as the result of changes in volatile [partial carbon dioxide tension (pCO_two)] and nonvolatile acids (hydrochloric, sulfuric, lactic, etc). Clinically, we refer to changes in volatile acids equally 'respiratory' and changes in nonvolatile acids as 'metabolic'. There are three major methods of quantifying (describing) acid–base of operations disorders, and each differs only in cess of this latter, 'metabolic' component. These three methods quantify the metabolic component either by using HCO_iii - (in the context of pCO_two), the standard base excess (SBE),or the strong ion difference (SID). Although there has been significant argue regarding the accurateness 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 agreement of mechanisms) [9,x,xi].

Across Henderson and Hasselbalch

Since Hasselbalch adapted the Henderson equation to the pH note of Sörenson, we have used the post-obit equation to sympathise the human relationship betwixt respiratory and metabolic acrid–base of operations variables:

pH = pK × log [HCO₃ /(0.03 × pCO₂)] (1)

This is the Henderson–Hasselbalch equation, and it is of import to realize what this equation tells the states. Anincrease in pCO₂ volition issue in a decrease in the pH and anincrease in the HCO₃ ^- concentration. Thus, a patientfound to have a low claret pH, a condition known asacidemia, will either accept an increased pCO₂ or a pCO₂that is 'non increased'. In the one-time circumstance, we allocate the disorder as a 'respiratory acidosis'. We use the term 'acidosis' to depict the process resulting in acidemia and 'respiratory' because the apparent cause is an increased pCO₂. This is logical, because carbonic acrid results when CO₂ is added to h2o (or blood), and the resultant subtract in pH is entirely expected. In the latter condition pCO₂ is not increased, and thus in that location cannot be a respiratory acidosis. We therefore refer to this condition as 'metabolic' because some nonvolatile acid must be the crusade of the acidemia. We can opposite the to a higher place logic and easily allocate uncomplicated conditions of alkalemia equally either resulting from respiratory or metabolic alkaloses. Thus, equation 1 allows us to classify disorders according to the primary type of acrid existence increased or decreased. Over time physiology superimposes its effects on simple chemistry and the relationship betwixt pCO_two and HCO_three ^- is altered in guild to reduce the alterations in pH. By carefully examining the changes that occur in pCO₂ and HCO_iii ^- in human 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 instance i such rule is the convenient formula for predicting the expected pCO_ii in the setting of a metabolic acidosis [12]:

pCO_two = (1.5 × HCO_three ^-) + 8 ± 5 (two)

This rule tells us what the pCO₂ should exist secondary to the increase in alveolar ventilation that accompanies a metabolic acidosis. If pCO_two does not modify enough or changes too much, we classify the status as a 'mixed' disorder, with either a respiratory acidosis if the pCO_two is still besides high, or a respiratory alkalosis if the change is likewise swell. This rule, along with others (Tabular array 1) has been recently translated to SBE terminology [7]:

Table one

Observational acid–base patterns

Disorder	HCO₃ ^- (mmol/l)	pCO_two (mmHg)	SBE (mmol/I)
Metabolic acidosis	<22	= (1.v × HCO₃ ^-) + 8	<-five
		= forty + SBE
Metabolic alkalosis	>26	= (0.vii × HCO₃ ^-) + 21	>+v
		= 40 + (0.half dozen × SBE)
Astute respiratory acidosis	= [(pCO₂ – 40)/10] + 24	>45	= 0
Chronic respiratory acidosis	= [(pCO_two – xl)/3] + 24	>45	= 0.4 × (pCO₂ – 40)
Astute respiratory alkalosis	= [(40 – pCO₂)/5] + 24	<35	= 0
Chronic respiratory alkalosis	= [(40 – pCO₂)/ii] + 24	<35	= 0.four × (pCO_two – 40)

Compiled from various sources. The changes in SBE were taken from Schlichtig et al [seven]. pCO₂, fractional carbon dioxide tension; SBE, standard base of operations backlog.

pCO₂ = (forty + SBE) ± 5 (3)

For example, consider the post-obit arterial claret gas sample: pH7.31, pCO_two 31, HCO₃ ^- 15, SBE-nine.v. Equation ii tells us that the expected pCO_two =(1.five×xv)+viii ± v=30.5 ± five, and in Equation 3 the SBE added to twoscore also yields 30.v. The measured pCO₂ of 31 mmHg is consistent with a pure metabolic acidosis (ie no respiratory disorder).

It is also very of import to understand what the Henderson-Hasselbach equation does not tell us. First, it does non allow united states to discern the severity (quantity) of the metabolic derangement in a manner analogous to the respiratory component. For case, 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 HCO_three ^-. Second, Equation 1 does not tell us about any acids other than carbonic acrid. The human relationship between CO_ii 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 HCO_iii ^- concentration, is determined by the pCO_ii, however, which is in plow determined by the balance between alveolar ventilation and CO_two product. HCO₃ ^- cannot be regulated independent of pCO₂. The HCO₃ ^- concentration in the plasma will e'er increase equally the pCO₂ increases, but this is not an alkalosis. To empathize how the pH and HCO_iii ^- concentration are altered independent of pCO₂, we must look beyond Henderson and Hasselbach.

Base of operations excess

In order to address the commencement 'shortcoming' of the Henderson-Hasselbach equation – the inability to quantify the metabolic component – several methods have been devised. In 1948, Vocalizer and Hastings proposed the term 'buffer base' to define the sum of HCO₃ ^- plus the nonvolatile weak acid buffers (A^-) [thirteen]. A change in buffer base corresponds to a change in the metabolic component. The methods for calculating the alter in buffer base were later refined by investigators [xiv,xv] and refined further by others [16,17,xviii] to yield the base backlog methodology. Base backlog is the quantity of metabolic acidosis or alkalosis, defined every bit the amount of acid or base that must exist added to a sample of whole blood in vitro in guild to restore the pH of the sample to 7.40 while the pCO_two is held at 40 mmHg [fifteen]. Although this calculation is quite authentic in vitro, inaccuracy exists when practical in vivo in that base backlog changes with changes in pCO₂ [xix,20]. This issue is understood to be due to equilibration across the unabridged extracellular fluid space (whole blood + interstitial fluid). When the base backlog equation is modified to account for an 'average' content of hemoglobin across this entire infinite, a value of five g/dl is used instead, and this defines the SBE. Information technology 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 gauge that improves the accurateness of the base backlog. It tin be argued that the unabridged extracellular fluid space is involved in acid–base residue, because this fluid flows through claret vessles and lymphatics, mixing constantly [21]. Thus, the value of SBE is that information technology quantifies the change in metabolic acrid–base of operations status in vivo. It is of interest that base excess is just accurate in vivo when it assumes a constant hemoglobin concentration.

However, the base excess approach does not address the second trouble associated with using the Henderson-Hasselbach equation alone (ie it still does non tell usa about the mechanisms of metabolic acrid–base balance). For case the body does not 'regulate' the SBE. It is not a substance that can be excreted in the carrion or reabsorbed from the proximal tubule. Similarly, HCO₃^- is not a strong acid or base of operations and its improver to or removal from the plasma cannot be translated into changes in SBE. This not to say that changes in SBE and HCO_three ^- practise non correlate closely, considering they practice. However, correlation and causation are non the aforementioned affair. The difference has traditionally been ascribed to the effects of 'buffering', the argument being that stong acid (or base of operations), quantified past SBE, is 'buffered' by plasma proteins, hemoglobin, and finally past HCO_three ^-. 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 past Stewart [6,9], the fundamental concrete-chemic backdrop of biologic solutions dictate much of this and then-called 'buffering'.

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Changes in the relationship between partial carbon dioxide tension (pCO₂) and H⁺ concentration every bit function of changes in 'buffer' force. Individual curves are drawn for varying concentration of total nonvolatile buffers in mmol/l. Note that as the concentration of 'buffer' increases, the slope of the curve increases, making changes in H⁺ concentration more responsive to changes in CO_two.

Physical-chemical properties of biologic solutions

A concrete–chemical analysis of acid–base physiology requires the application of two basic principles. The beginning 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 2d is conservation of mass, which means that the corporeality of a substance remains constant unless information technology is added to or generated, or removed or destroyed. These principles may seem very basic indeed, but they are often overlooked in the assay of clinical acid–base physiology, leading to incorrect conclusions. For example, a hyperchloremic metabolic acidosis tin can only exist brought about in 2 ways. Kickoff, Cl^- ions can be added to the circulation, either via an exogenous source (eg HCl or saline) or via internal shifts (eg from the ruddy 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 ^- HCO_three or H⁺ concentration does non produce hyperchloremia, just rather hyperchloremia is a crusade of acidosis. This distinction is not merely semantics, whatever more than Copernicus' ascertainment that the Earth, rather than the sun moves [11,22].

In addition to these physical–chemical principles, near all solutions of biologic involvement share two important characteristics. First, nigh all are aqueous (composed of water), and 2nd, virtually are alkaline (OH^-concentration >H⁺ concentration). Because these characteristics are and so universal in human physiology, they are frequently ignored in reviews of physiology, especially for clinical medicine, but they are extremely of import. 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 h2o dissociation. Similarly, aqueous solutions that are alkaline conduct very differently compared with acidic solutions in terms of the extent to which changes in their composition influence changes in pH.

To illustrate this bespeak, first consider a i l solution of pure water. Pure h2o contains merely a small corporeality of H⁺ and OH^- ions and molecular H_iiO. Pure h2o is a neutral solution by definition, considering the H⁺ and OH^- concentrations are equal. The concentration of these ions is determined solely past the extent to which h2o dissociates and can be defined past a constant, Chiliad'_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/fifty of each Na⁺ and Cl^- to this ane lsolution of pure h2o, nosotros would have an aqueous solution that contains H⁺, OH^-, Na⁺ and Cl^- ions, and molecular water. Of note, the solution does non 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 so add together 10 mmol/fifty HCl, we will accept a solution that contains x mmol/l Na⁺ and 20 mmol/l Cl^-. This solution is acidic: OH^- =4.four × 10^-ix nmol/l, and H⁺ = ~ x mmol/l. Notation that in this acidic solution the H⁺ concentration increased past the amount of H⁺ added (ie 10 mmol/l). However, if we were to add ten mmol/l NaOH instead of HCl, the solution would instead contain 20 mmol/fifty Na⁺ and 10 mmol/l Cl^-, and would be alkaline: H⁺ = 4.iv × 10^-ix nmol/l and OH^- = ˜ ten mmol/fifty. If nosotros and then add 5 mmol/50 HCl to this alkaline solution, the resulting concentration of Na⁺ would exist 20 mmol/l and of Cl^- would exist 15 mmol/l. The final H⁺ concentration is now 8.8 × ten^-9 nmol/l and OH^- is approximately 5 mmol/l. Note that in this final case 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 exist farther 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 belongings of alkali metal solutions.

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 simply three contained variables that determine H⁺ concentration. Please notation 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 simply correlated. The distinction between independent and dependent, between causation and correlation, is as of import to acrid–base of operations physiology as to any other area of scientific discipline. Merely 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 (A_TOT).

Carbon dioxide

CO₂ is an independent determinant of pH and is produced by cellular metabolism or by the titration of HCO₃ - by metabolic acids. Unremarkably, 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 pCO_two 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, equally 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, pCO_ii will increment to a new steadystate that is adamant past the new relationship betwixt alveolar ventilation and CO₂ production. Acutely, this increase in pCO₂will increment both the H⁺ and the HCO_three ^- concentrations according to the Henderson-Hasselbach equation (Equation ane). 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. In that location is no change in the SBE. Tissue acidosis e'er occurs in respiratory acidosis, considering CO₂ diffuses into the tissues. If the pCO_two remains increased the body will endeavour to compensate by altering another independent determinant of pH, namely the SID.

Electrolytes (potent ions)

Blood plasma contains numerous ions. These ions can be classified both by charge, positive 'cations' and negative 'anions', as well equally by their tendency to dissociate in aqueous solutions. Some ions are completely dissociated in h2o, for example, Na⁺, K⁺, Ca²⁺, Mg²⁺, and Cl^-. These ions are called 'strong ions' to distinguish them from 'weak ions' (eg albumin, phosphate and HCO_iii ^-), which can be both as charged (dissociated) and uncharged forms. Sure ions such every bit lactate are so well-nigh completely dissociated that they may be considered strong ions nether physiologic weather. In a neutral common salt solution containing merely water and NaCl, the sum of strong cations (Na⁺) minus the sum of strong anions (Cl^-) is nil (ie Na⁺ = Cl^-). In blood plasma, all the same, strong cations (mainly Na⁺) outnumber strong anions (mainly Cl^-). The difference between the sum of all potent cations and all strong anions is known equally 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 guild to maintain electrical neutrality (Fig. two).

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Plot of pH versus strong ion divergence (SID). For this plot, total weak acid concentration (A_TOT) and partial carbon dioxide tension (pCO₂) were held constant at xviii mmol/fifty and forty mmHg, respectively. Assumes a h2o dissociation constant for claret of iv.4 × 10–fourteen (mol/fifty). Notation how steep the pH bend becomes at SID <20 mmol/50. Adapted from Kellum and Kellum [x].

In healthy humans, the plasma SID is betwixt 40 and 42 mmol/l, although it is oft quite different in critically ill patients. Co-ordinate to the principle of electrical neutrality, blood plasma cannot be charged, and so the remaining negative charges balancing the SID come up from CO_two and theweak acids (A^-) and, to very small extent, from OH^-. At physiologic pH, the contribution of OH^- is so pocket-size (nmol range) that it can be ignored. The total weak acrid concentration (mainly albumin and phosphate) tin can be considered together and for convenience is abbreviated A_TOT, where AH + A^- = A_TOT. The SID of a blood sample can exist 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], just it is actually 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 substantially 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 gauge of SID is as follows: (Na⁺ + K⁺ + Caⁱⁱ⁺+ Mg²⁺) – (Cl^- + lactate^-). This is oft referred to as the 'credible' SID with the understanding that some 'unmeasured' ions might also exist nowadays [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 credible SID an inaccurate estimate of SID. Similarly, these patients may have aberrant weak ions (eg proteins) that volition make the effective SID inaccurate. In healthly humans, nonetheless, the credible 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 status we have referred to every bit the strong ion gap (SIG), where apparent SID – constructive 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 apportionment of patients with a diverseness of diseases [25,26,27,28] and in animals under experimental weather condition [29].

The SIG is non the same as the anion gap (AG). Normally, the SIG is zippo, whereas the AG is 8–12mmol/50. The AG is an estimate of the sum of SIG + A^-. Thus, subtracting A^- from the AG approximates the SIG. A user-friendly and reasonably accurate way to approximate A^- is to apply the following formula [30]:

2 (albumin g/dl) + 0.five (phosphate mg/dl) (iv)

or for international units:

0.2 (albumin m/l) + ane.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 ane to 'calibrate' the AG for patients with abnormal albumin and/or phosphate concentrations.

Physiologic mechanisms

In order to alter the SID, the trunk must affect a change in the relative concentrations of strong cations and stiff anions. The kidney is the principal organ that affects this change. Still, the kidney tin only excrete a very small-scale amount of strong ion into the urine each minute, and several minutes to hours are therefore required to impact significantly on the SID. The handling of strong ions by the kidney is extremely of import because every Cl^- ion filtered only not reabsorbed increases the SID. Because almost of the homo diet contains similar ratios of stiff cations to strong anions, there is usually sufficient Cl^-available for this to exist the primary regulating mechanism. This is particularly apparent when we consider that renal Na⁺ and K⁺ handling are influenced by other priorities (eg intravascular volume and plasma K⁺ homeostasis). Accordingly, 'acid treatment' by the kidney is mostly mediated through Cl^- rest. How the kidney handles Cl^- is evidently very important. Traditional approaches to this problem have focused on H⁺ excretion and have emphasized the importance of NH_iii and its add-on cation NH_four ⁺. H⁺ excretion per se is irrelevant, notwithstanding, because h2o provides an essentially infinite source of free H⁺. Indeed the kidney does not excrete H⁺ any more than every bit NH₄+ than it does as H₂O. The purpose of renal ammoniagenesis is to let the excretion of Cl^- without excreting Na⁺ or K⁺. This is accomplished by supplying a weak cation (NH₄ ⁺) to excrete with Cl^-.

Thus, NH_iv ⁺ is important to systemic acid–base balance,not because of its carriage of H⁺ or because of its straight action in the plasma (normal plasma NH₄ ⁺ concentration is <0.01mmol/l), just considering of its 'coexcretion' with Cl^-. Of course, NH₄ ⁺ is not but produced in the kidney. Hepatic ammoniagenesis (and also glutaminogenesis) is important for systemic acid–base residual and, as expected, it is tightly controlled by mechanisms that are sensitive to plasma pH [31]. Indeed this reinterpretation of the part of NH₄ ⁺ in acid–base balance is supported by the evidence that hepatic glutaminogenesis is stimulated past acidosis [32]. Nitrogen metabolism by the liver can result in either urea, glutamine or NH_iv ⁺. Normally, the liver does not release more than a very minor amount NH₄ ⁺, only rather incorporates this nitrogen into either urea or glutamine. Hepatocytes have enzymes to enable them to produce either of these end products, and both allow for the regulation for plasma NH_iv ⁺ at suitably low levels. The product of urea or glutamine has significantly unlike effects at the level of the kidney. This is because glutamine is used by the kidney to generate NH₄ ⁺ and facilitatethe excretion of Cl^-. Thus, the production of glutamine can exist seen as having an alkalinizing effect on plasma pH because of the way in which the kidney utilizes it.

Farther support for this scenario comes from the recent discovery of an anatomic organization of hepatocytes co-ordinate to their enzymatic content [33]. Hepatocytes with a propensity to produce urea are positioned closer to the portal venule, and thus have the commencement take a chance at the NH4⁺delivered. Acidosis inhibits ureagenesis, however, and under these conditions more NH₄ ⁺ is bachelor for the downstream hepatocytes that are predisposed to produce glutamine. Thus, the leftover NH₄ ⁺ is 'packaged' equally glutamine for export to the kidney, where information technology is used to facilitate Cl^- excretion and hence increases the SID.

The alimentary canal also has important furnishings on the SID. Along its length, the gastrointestinal tract handles strong ions quite differently. In the breadbasket, Cl^- is pumped out of the plasma and into the lumen, reducing the SID of the gastric juice and thus reducing the pH. On the plasma side, SID is increased past the loss of Cl^- and the pH is increased, producing the and so-chosen 'alkaline tide' that occurs at the beginning of a meal when gastric acid secretion is maximal [34]. In the duodenum Cl^- is reabsorbed and the plasma pH is restored. Normally, only slight changes in plasma pH are evident considering Cl^- is returned to the circulation almost as before long equally it is existence removed. If gastric secretions are removed from the patient, nonetheless, either by suction catheter or vomiting, Cl^- will be progressively lost and the SID will steadily increase. It is of import to realize that it is the Cl^- loss, and not the H⁺ that is the determinant of plasma pH. Although H⁺ is 'lost' as HCl, it is also lost with every molecule of water removed from the body. When Cl^- (a potent anion) is lost without loss of a stiff cation, the SID is increased and therefore the plasma H⁺ concentration is decreased. When H⁺ is 'lost' as water (HOH) rather than HCl, at that place is no change in the SID and hence no change in the plasma H⁺ concentration.

In contrast to the stomach, the pancreas secretes fluid into the small intestine that has a SID much higher than that of plasma and is very low in Cl^-. Thus, the plasma perfusing the pancreas has its SID decreased, a miracle that peaks about 1 h afterwards a meal and helps annul the alkaline tide. If large amounts of pancreatic fluid are lost, for example from surgical drainage, an acidosis will result as a effect of the decreased plasma SID. In the large intestine, fluid also has a loftier SID because most of the Cl^- has been removed in the small intestine and the remaining electrolytes are more often than not Na⁺ and Thou⁺. The trunk unremarkably reabsorbs much of the h2o and electrolytes from this fluid, simply when astringent diarrhea exists large amounts of cations can be lost. If this loss is persistent, the plasma SID volition subtract and acidosis will result. Finally, whether the gastrointestinal tract is capable of regulating strong ion uptake in a compensatory manner has non been well studied. In that location is some show that the gut may modulate systemic acidosis in experimental endotoxemia by removing anions from the plasma [35]. The full capacity of this organ to affect acid–base of operations balance is unknown, however.

Pathophysiologic mechanisms

Metabolic acidoses and alkaloses are categorized according to the ions that are responsible. Thus we have lactic acidosis and chloride responsive alkalosis, etc. It is of import to recognize that metabolic acidosis is produced past a decrease in the SID, which produces an electrochemical strength that results in an increase in free H⁺ concentration. A decrease in SID may be brought about by the generation of organic anions (eg lactate, ketones), the loss of cations (eg diarrhea), the mishandling of ions (eg renal tubular acidosis), or the addition of exogenous anions (eg iatrogenic acidosis, poisonings). By contrast metabolic alkaloses occur as a upshot of an inappropriately big SID, although the SID need non exist greater than the 'normal' forty-42 mmol/50. This may be brought about by the loss of anions in excess of cations (eg vomiting, diuretics), or rarely by administration of potent cations in excess of strong anions (eg transfusion of large volumes of banked blood). Tables ii and 3 provide a useful means to differentiate the various causes of metabolic acidosis and alkalosis.

Table two

Differential diagnosis for metabolic acidosis (decreased SID)

Renal tubular acidosis:	Nonrenal:
urine SID (Na⁺ + Thou⁺ – Cl) > 0	urine SID (Na⁺ + M⁺ – Cl^-) < 0
Distal (blazon I): urine pH >five.v	Gastrointestinal: diarrhea, small
	bowel/pancreatic drainage
Proximal (type II): urine
pH <v.5/low serum M⁺	Iatrogenic: parenteral nutrition,
	saline, anion exchange resins
Aldosterone deficiency (blazon 4):
urine pH <5.5/loftier serum Yard⁺

SID, strong ion difference.

Tabular array 3

Differential diagnosis of a metabolic alkalosis (increased SID)

Chloride loss < sodium loss

Chloride responsive (urine Cl^- concentration <x mmol/l)

Gastrointestinal losses: airsickness, gastric drainage, chloride

wasting diarrhea (villous adenoma)

Postdiuretic employ

Posthypercapnea

Chloride unresponsive (urine Cl^- concentration >20 mmol/50)

Mineralocorticoid excess: primary hyperaldosteronism (Conn's

syndrome), secondary hyperaldosteronism, Cushing'southward

syndrome, Liddle's syndrome, Bartter'southward syndrome,

exogenous corticoids, excessive licorice intake

Ongoing diuretic use

Exogenous sodium load (>chloride)

Sodium salt administration (acetate, citrate): massive blood

transfusions, parenteral nutrition, plasma volume expanders,

sodium lactate (Ringer's solution)

Other

Severe deficiency of intracellular cations: Mg²⁺, K⁺

SID, potent ion difference.

In the intensive care unit acidosis is usually more of a problem than alkalosis, and in the critically ill the virtually common sources of metabolic acidosis are disorders of chloride homeostasis, lactate, and other anions. Hyperchloremic metabolic acidosis occurs either equally a consequence of chloride administration or secondary to abnormalities in chloride handling or related to movements of chloride from one compartment to another. The effect of chloride administration on the development of metabolic acidosis has been known for many years [36,37]. Recently, new attention has been paid to this expanse in light of ameliorate understanding of the mechanisms responsible for this event [38,39,forty]. It has now been shown in animal models of sepsis [38] and in patients undergoing surgery [39,40] that saline causes metabolic acidosis not by 'diluting' HCO₃ ^-, merely rather by its Cl^- content. From a concrete-chemical prospective, this is completely expected. HCO_three ^-is a dependent variable and cannot exist the crusade of the acidosis. Instead, Cl^- assistants decreases the SID (an independent variable) and produces an increase in water dissociation and hence H⁺ concentration. The reason why this occurs with saline assistants is that, although saline contains equal amounts of both Na⁺ and Cl^-, the plasma does non. When large amounts of salt are added, the Cl^- concentration increases much more the sodium concentration. For example, 0.nine% ('normal') saline contains 154 mmol/l of Na⁺ and Cl^-. Administration of big volumes of this fluid will accept a proportionally greater result on total body Cl^- than on total body Na⁺. Of note, it is the total body concentrations of these strong ions that must be considered and, although the true volume of distribution of Cl^- is less, similar Na⁺, the effective book of distribution (afterwards some time of equilibration) is equal to total torso h2o [38].

In that location are other of import causes of hyperchloremia (Table 2) and, in addition, this form of metabolic acidosis is common in critical illness, specially sepsis. Although saline resuscitation undoubtedly plays a function, there announced to exist unexplained sources of Cl^-, at least in animal models of sepsis [38]. We have hypothesized that this Cl^- comes from intracellular and interstitial compartments as a result of the partial loss of Donnan equilibrium due to albumin exiting the intravascular space. However, this hypothesis is yet untested.

In improver to Cl^-, several other strong ions may be present in the blood of critically sick patients. Lactate is perhaps the about of import of these, simply ketones, sulfates, and sure poisons (eg methanol, salicylate) are important in the advisable clinical weather. In addition, unexplained anions accept been shown to be present in the blood of many critically sick patients [25,26,27,28].

Weak acids

The third and last determinant of H⁺ concentration is A_TOT. The weak acids are mostly proteins (predominantlyalbumin) and phosphates, and they contribute the remaining charges to satisfy the principle of electroneutrality, such that SID–(CO₂ +A^-)=0. Yet, A^- is non an independent variable considering it changes with alterations in SID and pCO_ii. Rather, A_TOT (AH + A^-) is the independent variable, because its value is not determined past any other. The identification of A_TOT as the tertiary contained acrid–base variable has pb some authors to advise that a 3rd 'kind' of acrid–base of operations disorder exists [41,42]. Thus, along with respiratory and metabolic acidosis and alkalosis, nosotros would too have acidosis and alkalosis due to abnormalities in A_TOT. Yet, mathematical, and therefore chemic independence does not necessarily imply physiologic independence. Although the loss of weak acrid (A_TOT) from the plasma infinite is an alkalinizing process,in that location is no evidence that the body regulates A_TOT to maintain acrid–base of operations residue. Furthermore, there is no prove that nosotros as clinicians should treat hypoalbuminemia equally an acid–base disorder.

Critically ill patients oft have hypoalbuminemia andas such their A_TOT is reduced. These patients are not often alkalemic and their SID is besides reduced, however [43]. When these patients have a normal pH and a normal SBE and HCO₃ ^- concentration, it would seem most appropriate to consider this to be physiologic compensation for a decreased A_TOT [44], rather than classifying this condition every bit a circuitous acid–base disorder with a mixed metabolic acidosis/hypoalbuminemic alkalosis. Thus, information technology seems far more likely that this 'disorder' is in fact the normal physiologic response to a decreased A_TOT. Furthermore, because changes in A_TOT more often than not occur slowly, the evolution of alkalemia would require the kidney to continue to excrete Cl^- despite an evolving alkalosis. I would consider such a scenario to be renal-mediated hypochloremic metabolic alkalosis, the treatment for which would include fluids and/or chloride, depending on the clinical conditions. Stewart's designation of a 'normal' SID of approximately xl mmol/fifty was based on a 'normal' CO_ii and A_TOT [six,9]. The 'normal' SID for a patient with an albumin of 2g/dl would exist much lower (eg approximately 32 mmol/50).

Decision

Dissimilar many other areas in clinical medicine, the approach to acid–base of operations physiology has not frequently distinguished cause from effect. Although information technology is perfectly reasonable to describe an amending in acid–base of operations status by the observed changes in H⁺ and HCO_three ^-, this does not itself imply causation. The essence of the Stewart [vi,nine] approach is the agreement that only three variables are important in determining H⁺ concentration: pCO₂, SID and A_TOT. Neither H⁺ nor HCO₃ ^- can change unless one or more of these iii variables modify. The principle of conservation of mass makes this point more than than semantics. Strong ions cannot be created or destroyed to satisfy electroneutrality but H⁺ ions are generated or consumed by changes in water dissociation. Hence, in order to understand how the body regulates pH, we need only consider how it regulates these three independent variables. Other approaches to acid–base of operations physiology have ignored the stardom between independent and dependent variables, and although information technology is possible to describe an acid–base disorder in terms of H⁺ or HCO₃ ^- concentrations or SBE, it is incorrect to clarify the pathology or to plan treatment on the basis of altering these variables.

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