Acids, Bases, Buffers: The Traditional Understanding

An acid is strong when it dissociates its hydrogen ion easily because the corresponding base has a low affinity for it (high Ka–dissociation constant value). Weak acids only partially dissociate because their corresponding base has high affinity for hydrogen ions (low K_a value). When the latter is the case, both the hydrogen ion and the base form of the parental molecule are present in the resulting solution in equimolar proportions. As a consequence, this solution has the ability to resist changes in acidity (or pH) after the addition of a strong acid or base. This solution, constituted by an acid-base pair, is known as a buffer. The dissociation constant K_a, as in the case of pH, can be expressed as log₁₀K_a, is termed pK_a, and represents a measure of the tendency of the acid-base pair to ionize. Strong acids have a low pK_a (high K_a) and weak acids have high pK_a (low K_a). A weak acid-base pair is more effective as a buffer in living systems when the pK_a is close to physiologic pH.^6,7,10

Physiologic buffering systems may be classified into two general categories: the bicarbonate/carbonic acid (HCO₃⁻/H₂CO₃) buffering system, that acts in both the extracellular space and inside the cells, and the nonbicarbonate buffers, which include hemoglobin and oxyhemoglobin, organic phosphates, inorganic phosphates, and plasma proteins. Extracellular buffers (HCO₃⁻, plasma proteins) represent the body’s first and immediate line of defense against any alteration of pH of the body fluids. After extracellular buffering occurs, a second intracellular phase takes place, as long as intracellular buffers (intracellular proteins, dibasic phosphates, hemoglobin-oxyhemoglobin, and carbonate in bones) reach buffering capacity over the next several hours.¹⁰ The buffering systems cannot truly eliminate hydrogen ions from the body, but they tamper sudden changes in [H⁺] (or pH) and “buy time” until a new balance can be achieved. Altogether, extracellular and intracellular buffers provide a volume of distribution close to that of the total body water to any exogenous acid load (or deficit). This represents a formidable capacity for resisting changes in pH.

Acid Production in the Body

Acid production in the body occurs in two major ways. The first one involves CO₂ production during oxidative metabolism of carbohydrates, fat, and amino acids. The CO₂ is then hydrated by the cytoplasmic enzyme carbonic anhydrase to produce carbonic acid (H₂CO₃), a weak acid that dissociates in hydrogen ion and bicarbonate.

(3)

The second pathway involves nonvolatile metabolic acids that are produced by the normal daily catabolic load (oxidation of sulfur-containing amino acids, hydrolysis of pyrophosphate and orthophosphate esters) or during situations with incomplete catabolism of fat and carbohydrates, such as lactic acidosis or diabetic ketoacidosis (DKA). The free hydrogen ions are neutralized by extracellular and intracellular buffers, but because these acids are not in equilibrium in the normal plasma, they must be metabolized, mainly in the liver, and then excreted by the kidneys.¹⁰

The major acid-base buffering system in the blood is the bicarbonate/carbonic acid system, which has the tremendous advantage of interconversion of CO₂ with H₂CO₃. Any increase in [H⁺] (or drop in pH) will shift the former reaction (Equation 3) to the left through an increase in both alveolar ventilation and elimination rate of CO₂. This respiratory response begins within minutes, but may not reach a steady state for 12 to 24 hours. This method of achieving a rapid decrease of the actual acid component is a unique characteristic of the bicarbonate/carbonic acid buffering system. For the other buffers, the addition or removal of hydrogen ions has the corresponding opposite effect on the buffer components, limiting the maximum buffering capacity. On the contrary, the capacity of the bicarbonate-CO₂ system is greatly increased because the lungs can eliminate a vast amount of CO₂ per day. Similarly, the kidneys can eliminate or regenerate bicarbonate as needed, although this response is slower than that of the lungs.

The Classic Paradigm: H⁺ Depends on CO₂, H₂CO₃, and HCO₃⁻

The quantitative importance of the bicarbonate/carbonic acid buffering system was clearly noticed by Henderson when he developed his famous equation, which was modified by Hasselbalch shortly after. The Henderson-Hasselbalch equation^4,5 expresses the relationship of the bicarbonate/carbonic acid buffering system to pH:

(4)

The pH is equal to a constant (pK_a, the log₁₀ of K_a, the first dissociation constant for H₂CO₃, which has a value of 6.1 for human plasma in physiological conditions) plus the log of the ratio of HCO₃⁻ (proton acceptor) to H₂CO₃ (proton donor). The modified Henderson-Hasselbalch equation takes into consideration that H₂CO₃ is in equilibrium with dissolved CO₂ and is rewritten as⁹:

(5)

What the Henderson-Hasselbalch equation tells us is how pH is affected by the change in the ratio of the concentration of nondissociated acid HA (in this case carbonic acid or PCO₂) to the concentration of the conjugated base or anion A⁻ (in this case HCO3⁻). When pH is altered as the result of changes in the volatile component (increases or decreases of PCO₂), clinically speaking, the change is referred to as respiratory.⁹ When pH is modified by changes in nonvolatile acids (e.g., lactic acid, keto acids), or by changes in serum cations and/or anions (e.g., hyperchloremia/hypochloremia), it is referred to as metabolic in origin.^5,8,9 The term acidosis is used to describe the process that tends to produce an increase in H⁺, whether or not there is a change in pH. Alkalosis is the opposite; that is, the process that tends to produce a decrease in H⁺, with or without changes in pH. Acidemia and alkalemia are the corresponding terms for those situations in which blood pH actually changes.⁹ Thus this equation allows disorders to be classified according to the primary type of acid being increased or decreased. For example, if a patient’s pH is low (acidemia), then the patient may have either an increased PCO₂ or a PCO₂ within normal or decreased values. If the PCO₂ is increased (alveolar hypoventilation), the condition is classified as respiratory acidosis. If it is not increased, there cannot be a respiratory acidosis. Therefore some nonvolatile acid or anion must be the cause of the acidemia, and this is then referred to as metabolic acidosis. If these examples are reversed, then alkalemia can also be classified as resulting from either respiratory or metabolic alkaloses.⁹

The Henderson-Hasselbalch equation is useful but has intrinsic limitations.^5,8 The first one is that it does not quantify the severity of the metabolic derangement as it does for the respiratory component. In a respiratory acidosis, the increase in the PCO₂ quantifies the derangement even when there are mixed disorders. The metabolic component can only be approximated, however, by the change in HCO₃⁻. Although the relationship between PCO₂ and HCO₃⁻ provides a useful clinical guide for uncovering a metabolic origin of a given derangement, this is dampened because PCO₂, H₂CO₃, and HCO₃⁻ are all interlinked (Equations 3 and 5), so the bicarbonate will also increase if PCO₂ increases.^6,8 The second important characteristic of the Henderson-Hasselbalch equation is that it does not provide information about any acids other than carbonic acid.

Using the Brönsted-Lowry definition of acid as a “proton donor,”^4,6,7,11 many physiologists dismissed defining chloride as an acid and sodium as a base, arguing that there was an insufficient link between these and other electrolytes and the subsequent changes in hydrogen ions. Thus, by accepting only Brönsted-Lowry acids while looking for factors controlling the nonrespiratory component of acid-base balance, many physiologists focused on the plasma bicarbonate concentration and the Henderson-Hasselbalch equation.⁵ This was the beginning of the still-dominant concept that plasma bicarbonate is not only the best indicator of acid-base status, but also a main determinant of it. Unfortunately, this way of thinking overlooked the fact that all weak acids in a given aqueous solution such as plasma can be inserted into a Henderson-Hasselbalch type of equation to calculate pH, as Lawrence Henderson himself demonstrated in 1908.⁴ The reason is easy to understand: for a single solution containing several weak acids (as human plasma), all the weak acids are equilibrated with a single pool of hydrogen ions. This is called the isohydric principle.⁵ Consider two buffer systems: bicarbonate/carbonic acid and phosphate/phosphoric acid:

(6)

(7)

Thus according to the isohydric principle, the ratio of any pair of conjugate base or anion and its nondissociated acid will be able to describe the H⁺ or pH.^5,12 This means that although the ratio of PCO₂ to bicarbonate can describe what the pH and acid-base status is, the system bicarbonate/carbonic acid is not necessarily the primary or underlying mechanism for explaining changes in pH.

Correlation and causation are not the same.^5,8 Immersed in the bicarbonate-centered approach, investigators yielded basically three solutions for analyzing the “metabolic side” of the equation: the base excess method, the six bicarbonate “rules of thumb,” and the use of anion gap (AG).^{4,5,9,13–15}

Base Excess and Standard Base Excess

The change in pH of a given buffered solution is dependent on both the amount of strong acid (or strong base) that is added, and on the buffering capacity of the system. As acid is added, for every H⁺ that is buffered, one molecule of conjugate base of the buffer is consumed. Therefore, quantifying the changes in the concentration of the conjugate base of the buffer is more useful than the degree of change in the pH in estimating the amount of nonvolatile acids present.

Siggaard-Andersen, from Copenhagen, developed the “base excess” method in the late 1950s.^4,14,15 Base excess (BE) is defined as the amount of strong acid (or strong base), in moles per liter (mol/L), that must be added to a whole blood sample to return the pH of the sample to 7.40, while the PCO₂ is maintained at 40 mm Hg.^5,15,16 Therefore if the blood sample is normal (i.e., its pH is 7.4 and its PCO₂ is about 40 mm Hg), the BE will be 0 mmol/L. Positive values mean literally an excess of metabolic bases; negative values mean an excess of metabolic acids.^9,15,16 To apply the BE to the clinical setting, a nomogram was developed that was later mathematically transcribed to allow BE calculation by blood gas analyzers.^17,18 However, several flaws appeared.^19–21 For clinical accuracy, other assumptions had to be incorporated (correction factors, adjusted formulas, nomogram modifications),^5,18,22 and, remarkably, an empiric estimate of hemoglobin concentration throughout the entire extracellular fluid space (whole blood plus interstitial fluid) had to be established, in order to consider the net effect of the intracellular buffers, hemoglobin the main one.⁵ A hemoglobin concentration of 50 g/L (5 g/dL)^5,22 was chosen to calculate standard base excess (SBE), a parameter reported nowadays by modern gas analyzers and used in the classical approach to acid-base balance (Table 68-1). Although SBE has good correlation with bicarbonate levels and quantifies the change in metabolic acid-base status in vivo, its accuracy depends on the 5 g/dL of hemoglobin assumption, and it still does not provide information about the origin or mechanisms of the metabolic acid-base derangement because SBE is not a substance that can be regulated, absorbed, or excreted by the body.⁸ Perhaps the main pitfall is that the SBE value represents the net effect of all metabolic acid-base abnormalities. Therefore the effects of coexisting metabolic acidoses and alkaloses may cancel each other, and the normal figure of the SBE will mistakenly suggest that no acid-base derangement exists. Actually, 15% to 18% of critically ill adults with acid-base disorders have a normal SBE.^13,23

Table 68–1 Classical Acid-Base Parameters

Bicarbonate Rules

In a “great transatlantic debate,”^18,21 strong criticism of the SBE from Schwartz and Relman from Boston yielded the six bicarbonate “rules of thumb” (Table 68-2).¹⁸ The Henderson-Hasselbalch equation easily disclosed derangements into “respiratory” and “metabolic,” the so-called simple derangements.⁹ Because the body always seeks to tightly control H⁺, however, several physiologic responses become active over time, and then the relationship between PCO₂ and HCO₃⁻ is modified so pH changes can be minimized. This minimization gives origin to more complex or “mixed” conditions.^9,10,24 Despite this, through careful examination of the changes that occur in PCO₂ and HCO₃⁻ in relation to each other, it was possible to find patterns and to derive rules (the so-called six rules of thumb) for uncovering mixed disorders and to differentiate acute from chronic respiratory unbalances. These rules describe the physiological compensation to acid-base changes to optimize acid-base homeostasis. With the expected physiologic compensation allowed for, residual changes in CO₂ or bicarbonate are then seen as the mechanisms for changes in acid-base status (Tables 68-2 and 68-3). This approach is still the most practical acid-base balance diagnostic tool for the busy clinician.⁹

Table 68–2 Classical Acid-Base Approach: Observational Acid-Base Patterns

Table 68–3 Classical Acid-Base Approach: Additional Clues

Anion Gap and Corrected Anion Gap

The AG was introduced as a complementary diagnostic tool for either SBE or bicarbonate rules-of-thumb approaches to metabolic disturbances.²⁵ It is based on the principle of electroneutrality, which states that there is no electrical charge in plasma. Accordingly, serum positive-charged cations must equal serum negative-charged anions⁹ (see Table 68-1):

(8)

The plasma concentrations of SO₄²⁻, OH⁻, CO₃²⁻, and H⁺ are quite small and can be neglected. Concentrations of Na⁺, K⁺, Cl⁻, and HCO₃⁻ (in the form of total CO₂) are reported in a standard chemistry panel. The sum of the remaining anion species, proteins⁻, PO₄³⁻, and XA⁻, is defined as the unmeasured anions (UA), whereas the sum of the remaining cation species, Ca²⁺ and Mg²⁺, is defined as the unmeasured cations (UC). Thus AG is commonly defined as⁹:

(9A)

Sometimes [K+] and its effect are disregarded²⁶:

(9B)

Typically, normal values are 16 mEq/L (if K⁺ is included) or 12 mEq/L (without K⁺), with variations of ±2 to 4 mEq/L. These values may be influenced by the way the values of some of the parameters—particularly chloride levels—are measured, so it is always better to consult each institution’s own expected “normal AG.”^26,27

It must be realized that calculated AG is affected by any change in the concentrations of either the UA or UC. Plasma proteins and phosphate levels can be significantly decreased in the critical care setting. Reduced protein and phosphate levels result in a decreased UA and hence in a decreased AG. Therefore, an increase in AG from organic acids may be masked by a decrease in AG from low protein and phosphate levels and thus it is wise to make the proper corrections.²⁸ This is particularly true in the pediatric critical care setting, where hypoalbuminemia is nearly ubiquitous.^26,29,30 Phosphate influence is rather small, and so is the effect of globulins, since their pK_a is much greater than plasma pH (they do not easily dissociate) and hence they do not have a significant H⁺ contribution. Figge et al. showed that in most patients, the AG could be corrected (AG_CORR) as follows²⁶:

(10)

The albumin-corrected AG (AG_CORR) can unmask an organic acidosis previously undetected in the setting of hypoalbuminemia, and adds sensitivity for detecting unmeasured anions, including lactate, although the specificity for hyperlactatemia is poor.^28,31–33 Thus lactate must be directly determined.

Despite the fact that it is not useful for quantifying the metabolic derangement, AG is a powerful tool for the clinician in the categorization of metabolic acidoses, as not all metabolic acidoses result in an elevated AG.³⁴ As an hypothetical example, the reader might consider that acid in the form of HCl is added to the circulation. Upon dissociation, Cl⁻ remains as the conjugate base, while the plasmatic bicarbonate buffer is “consumed” by the hydrogen ion. The result is that an increase in Cl⁻ is balanced by a decrease in HCO₃⁻, that is, there is metabolic acidosis without a change in the AG. Thus through the AG_CORR it is possible to differentiate hyperchloremic acidosis (normal AG) from “gap acidoses” (increased AG).^26,34 The last category includes all conditions of metabolic simple acidosis caused by increased concentrations of nonvolatile anions other than chloride, usually unmeasured (or unmeasurable), such as lactate, keto acids, phosphate, sulfates (and other anions in renal failure setting), salicylate, some β-lactam antibiotics, organic acids from congenital errors of metabolism, and acetate from parenteral nutrition (Table 68-4).^31,34 Non-AG hyperchloremic acidoses include excessive infusion of chloride salts (saline or parenteral nutrition) and increased renal or gastrointestinal bicarbonate losses. On the contrary, metabolic alkalosis may stem from hypochloremia from chloride loss or bicarbonate gain, and from hypoalbuminemia.³¹ AG performance is not so good in mixed acid-base physiology, despite several attempts to improve it. This represents an important limitation to its use in critical care patients, in whom single metabolic acid-base disorders are more the exception than the rule. In spite of several attempts, the “corrections” and adaptations for the AG that have been proposed in order to detect mixed disorders have not fulfilled the expectations.^24,31

Table 68–4 Causes of Metabolic Acidosis in Critically Ill Patients

∗ Total parenteral nutrition may cause both anion gap and non-anion gap acidosis. In the first case, an excessive amount of exogenous acids is the cause, mainly if liver or renal functions are impaired; in the latter, an unbalanced, excessive chloride content in the formulation provokes the derangement, which is easily produced in the setting of renal failure.

† This class corresponds to the gastrointestinal loss of bicarbonate type of acidosis, according to the classical approach.

‡ This class corresponds to the renal loss of bicarbonate type of acidosis, according to the classical approach.

Water as the Main Source of Hydrogen Ions: The Stewart Approach

It is remarkable that not one of the classic acid-base approaches examined the role of water as a virtually inexhaustible source of hydrogen ion, as Peter Stewart later did.^5,8,35 Although pure water dissociates only slightly into H⁺ and OH⁻ in plasma, the presence of electrolytes, CO₂, and other weak acids produces powerful electrochemical forces influencing water dissociation.⁸ In the late 1970s and early 1980s, Peter Stewart proposed that Arrhenius’ more general definition of an acid, along with Naunyn’s ideas from 1900, was more useful to acid-base physiology than the Brönsted-Lowry definition.^36,37 Applying several basic principles of physical chemistry, particularly electroneutrality, conservation of mass, and dissociation equilibrium of partially dissociated substances (electrolytes and others),^5,8,13,38 Stewart developed a mathematical model of acid-base balance. He defined the system as an aqueous solution that contains strong ions that are completely dissociated at physiologic pH, weak acids that are partially dissociated at physiologic pH, and carbon dioxide that is in equilibrium with an external partial pressure of carbon dioxide. The main concepts of this approach are summarized in Table 68-5, and a classification of primary acid-base disturbances based in this system is presented in Table 68-6. Thus according to Stewart, there are only three independent controlling variables of H+ concentration^5,8,38–40:

Table 68–5 Concepts and Parameters Used in the Quantitative Physical-Chemical Approach to Acid-Base Balance: A Summary

Table 68–6 Classification of Primary Acid-Base Disturbances: Independent Variables-Oriented Approach

“Independent variables” mean that PCO₂, SID, and A_TOT are causally related to the hydrogen ion, rather than being merely correlated.^38,39 Normal acid-base status occurs when the independent variables have normal values. Abnormality of one or more of the independent variables underlies all acid-base disturbances. Adjustment of the independent variables is the essence of all therapeutic interventions, because none of the dependent variables (e.g., pH, BE, HCO₃⁻) can be changed primarily or individually; the dependent variables change all of them simultaneously if, and only if, one or more of the independent variables changes.^8,13,39 Since PCO₂ is regulated by respiration, its changes result in respiratory acid-base disorders. Metabolic acid-base disorders stem both from changes in either SID or A_TOT.

CO₂ and Bicarbonate in Stewart’s Approach

In Stewart’s approach, CO₂ represents a primary independent determinant of pH, just as in the bicarbonate-BE-centered approaches, that is, there are no changes in the understanding of the respiratory acid-base derangements. However, the role of bicarbonate is reduced from being causative to a mere indicator, as hydrogen ion and bicarbonate concentrations are totally dependent on the three independent controlling variables. Hence, according to Stewart, the major use for the bicarbonate rules of thumb and SBE are to determine the extent of the clinical acid-base disorder, rather than the mechanism.^5,8,39

The rise of the PCO₂, according to the Henderson-Hasselbalch equation (Equation 5), will increase both H⁺ and HCO₃⁻ concentrations. Therefore the change in HCO₃⁻ concentration is mediated by chemical equilibrium (Equation 3) and not by any systemic adaptive response.⁸ The total CO₂ concentration (and hence the [HCO₃⁻]) is determined by the PCO₂, which is in turn determined by the balance between alveolar ventilation and CO₂ production at tissue level. Therefore HCO₃⁻ cannot be regulated independently of PCO₂. Plasma bicarbonate concentration will always increase if PCO₂ increases, but this is not an alkalosis. Therefore the increased bicarbonate concentration is not “buffering” the rise of [H⁺], and there will be no change in the SBE. As CO₂ easily diffuses through membranes, when PCO₂ increases there is always tissue acidosis. If PCO₂ remains high, the body will attempt to compensate by altering another independent determinant of pH. The kidneys are the main organs involved in this compensation.^8,39

The Strong Ion Difference

Strong ion difference (SID) is defined as the charge difference between the sum of all strong cations and the sum of all strong anions. When SID is calculated from direct measurement of serum strong ions, it is termed apparent SID (SID_APP),^8,31,42,43 with the understanding that more unmeasured (or unmeasurable) ions might also be present (see Table 68-5):

(11A)

(11B)

From actual measurements in healthy volunteers, the normal range of SID_APP has been estimated from 38 to 42 mEq/L.^5,8,13,42,44 In critical illnesses, SID_APP may be substantially reduced, even when there is no evidence (by the traditional approach) of a metabolic acid-base derangement.^42,45 Gunnerson et al.⁴² found in stable critical care patients a SID_APP value of 33 ± 5.6 mEq/L, but others have found different values. Thus, in spite of the lack of consensus about what value of SID_APP one should expect in an ICU population, it is clear that their SID_APP values are different from healthy volunteers’ figures.^23,44,45 The tendency of critical care patients to have a lower SID_APP is not surprising, given the fact that the positive charge of the SID_APP is balanced by the negative charges of the weak acids, [A⁻] (albumin and phosphate), and total CO₂. Because PCO₂ is usually modified by the body for other reasons (e.g., hyperventilation in a patient that tends to be hypoxemic) and because hypoalbuminemia is the rule rather than the exception in this setting, [A⁻] tends to be reduced, and this decrease leads to a reduction in SID_APP to keep the normal pH. Thus a typical critical care patient might have a SID_APP around 30 to 35 mEq/L, rather than the “usual” range of 38 to 42 mEq/L. If this patient then has metabolic acidosis, the SID_APP will decrease further. The lower the baseline of the SID_APP, the greater the susceptibility to a subsequent acid load.^42,44,45

Additional strong anions may become significative (e.g., hyperlactatemia, sulfate, keto acids, anions from Krebs cycle) or appear de novo in disease states (several organic acids from inborn errors of metabolism), or enter the body as medications or toxic substantes (salicylates, certain β-lactam antibiotics, etc.). They may eventually have an influence on the value of SID_APP.^44–47 Because it is usually not possible to have a direct measure of [XA⁻] at bedside (with the exception of lactate), it is not always convenient to rely only on SID_APP to determine the strong ion status. The more ill the patient, the more accurate is this statement. Then, in critically ill patients, it is better to calculate SID to account for electrical neutrality. According to the principle of electroneutrality (Equation 8), blood plasma cannot be charged, so there should be remaining negative charges balancing the “excess” of plasma strong cations in relation to the strong anions, which is the origin of SID_APP. These balancing negative charges come from total CO₂ (~[HCO₃⁻]) and from two substances that act as nonvolatile weak acids, that have concentrations in plasma large enough so that changes in them can produce significant acid disturbances: the weak acids (A⁻) albumin and inorganic phosphate. Thus SID can be derived, accounting for electrical neutrality, as the sum of HCO₃⁻ plus the negative electric charges contributed by albumin (Alb⁻) and by inorganic phosphate (Pi⁻).¹³ This calculation of SID is known as effective SID (SID_EFF)^13,48:

(12A)

A less-accurate, simplified version of this formula, which assumes a stable pH of 7.4, appears in some publications³¹:

(12B)

where factors 0.28 and 1.8 are, respectively, the negative electric charges displayed by 1 g albumin and 1 mmol of [Phosphate] in plasma at pH 7.4 (0.6 instead of 1.8 if [Phosphate] is in mg/dL).⁴⁸

Human serum albumin must be treated as a polyproteic macromolecule, exhibiting multiple apparent equilibrium dissociation constants corresponding to different classes of amino acids side chains. Accordingly human albumin charge should be expressed as a linear function of the pH. Thus Eqation 12A from Figge et al.⁴⁸ is more useful in sicker patients. Actually, SID was derived from a mathematical model with much more complex nonlinear equations, which might be encountered by the clinician while reviewing certain books and articles.^49–51 Constable and coworkers^42,52 have now developed some adaptations for clinical application, which are claimed to be more accurate:

(12C)

SID has a powerful electrochemical effect on water dissociation, and hence on H⁺. Both H⁺ and OH⁻ behave as a weak cation and anion, respectively. Because practically all the cations in plasma are strong ions, except for H⁺, only H⁺ concentration can vary in response to changes in anions. On the other hand, there are several anions that are weak ions, and thus there are several “options” for anion molecules that can change their charges if cations change.^8,53 As the SID_APP increases (strong cations in excess over strong anions) plasma becomes positively charged. Hence H⁺, a weak cation, decreases, and pH increases, to maintain electrical neutrality, and alkalosis is produced. On the other hand, if the SID_APP decreases (strong anions in relative excess over strong cations), plasma becomes negatively charged, more water is dissociated producing more H⁺ for maintaining electroneutrality, and acidosis is produced.^8,40,51

Changes in SID result from either a relative or an absolute change in strong ion concentrations. A change in free water content produces an alteration in the relative concentration of the strong ions, and hence a change on SID. If, as a result, Na⁺, K⁺, and Cl⁻ decrease, the absolute value for SID will decrease proportionately. Thus an excess of free water lowers SID and results in a tendency towards metabolic acidosis. In a similar manner, a free water deficit causes a metabolic alkalosis by increasing the SID through a relative increase in concentration of all strong ions. Relative and absolute changes in sodium concentration primarily result from osmoregulation, and thus they are reflective of changes in free water. The remaining strong cations, potassium, magnesium, and calcium, as they are tightly regulated by the body for other functions (coagulation, membrane excitability, neuromuscular plates, muscle contraction, etc.) do not vary significantly enough to directly cause alterations in acid-base balance.¹³ On the other hand, changes in strong anion concentrations are significant to acid-base status. Traditionally, hypochloremia (from gastrointestinal or renal losses) is associated with metabolic alkalosis, and hyperchloremia (e.g., from resuscitation with saline infusion) is associated with metabolic acidosis. Because chloride concentration can be affected by free water content just as is sodium concentration, according to some authors,³¹ it is important to take in account alterations in free water content that might exist, and thus Cl⁻ would need to be corrected based on the sodium concentration:

(13)

As the SID requires an actual change in the relative concentrations of strong cations or anions, the kidneys are the primary organs regulating SID. As the kidneys can excrete only small amounts of strong ion into the urine each minute, several minutes to hours are required to achieve a significant change in the SID. The control of the kidney is important and precise because every chloride ion that is filtered and not reabsorbed increases the SID. Acid handling of the kidney has traditionally focused on H⁺ excretion and the role of ammonia (NH₃) and ammonium (NH₄⁺) as hydrogen ion carriers.^9,10,54 Because water provides an essentially infinite source of H⁺, its excretion, per se, is possibly not as relevant as formerly taught. In fact, the net H⁺ excretion by the kidney as water molecules is larger than the excretion as NH₄⁺. Therefore the purpose of renal NH₄⁺ production is to allow the excretion of Cl⁻ without the excretion of Na⁺ or K⁺. Thus NH₄⁺ is important to systemic acid-base balance, not for its role as an H⁺ carrier or for its direct action in the plasma (normal [NH₄⁺] is ~0.01 mmol/L), but because it allows a “safe” excretion of Cl⁻.^8,55

A urinary anion gap (uGap) was described some time ago.⁹ This urinary gap has also been designated by some authors⁸ as “urine strong ion difference” (uSID):

(14)

Certainly, sodium, potassium, and chloride are all strong ions, so the name is, at least, functionally correct. What is the value in measuring the uSID? As its serum analogue, the uSID estimates the unmeasured urinary ions. In terms of its amount, the most important unmeasured urinary strong anion (Ur⁻) is SO₄²⁻ (derived from the metabolism of sulfur-containing amino acids), whereas the most important unmeasured cation (Ur⁺) is ammonium (NH₄⁺). In physiologically “normal” conditions, uSID must equal the plasmatic SID_APP, that is, 38 to 42 mEq/L. When a strong ion enters the plasma (e.g., lactate, or a chloride load), plasmatic SID will obviously decrease. Normal kidneys will react by increasing their excretion of chloride, thereby decreasing the plasma chloride concentration, while Na⁺ and K⁺ must be maintained within normal ranges. This is accomplished by increasing the excretion rate of NH₄⁺, which is the more efficient way to increase chloride elimination without losing sodium or potassium. The increased excretion of chloride will decrease the uSID.⁵⁵

In addition to the kidney, both the liver and the gastrointestinal tract may have an influence on SID. In the stomach, Cl⁻ is pumped out of the plasma and into the lumen, reducing the SID (and pH) of the gastric fluids, but increasing the SID (and pH) in the plasma side (alkaline tide), because of the exit of Cl⁻ to the stomach. However, only a slight change in plasma pH becomes evident, because Cl⁻ is reabsorbed in the duodenum just as fast as it was pumped out. However, if gastric secretions are removed, either by vomiting or suction catheter, the SID will progressively increase, as will the pH, as a result of Cl⁻ loss. Although H⁺ is excreted as HCl, it is also lost with every water molecule that exits the body. When the body loses Cl⁻, a strong ion, without also losing a strong cation, SID is increased; therefore H⁺ is decreased and alkalosis ensues. When H⁺ is excreted as H₂O rather than as HCl, there is no change in SID, and consequently there is no change in H⁺. At present it is not known whether the gastrointestinal tract, the liver, and the pancreas are capable of compensatory actions to regulate strong ion uptake.^9,39