Free H⁺combines with the anionic base (A^–) to form a weak acid (HA), which can either remain in this state or dissociate back to A^– and H⁺ depending on the [H⁺]. A particular buffer is most effective when the local pH is within 1 pH unit of its pK_a (Table 7.1).

Chemical buffer systems are located throughout the body in a range of fluids and tissues as shown in Table 7.1. The three most important chemical buffers are the bicarbonate, phosphate, and protein buffering systems.

Bicarbonate buffer system

The bicarbonate system is responsible for approximately 80% of extracellular fluid (ECF) buffering, and contributes between 10% and 15% to total chemical buffering capacity. The system consists of a water solution containing H2CO₃ (a weak acid) and its conjugate base, HCO₃⁻(usually as the sodium salt NaHCO₃). The entire system may be represented as follows:

TBW, total body water.

^a T he imidazole ring of histidine has a p Ka of 6.4–7.0 while the α-amino group has a pKa of 7.4–7.9.

The formation of H₂CO₃ from CO₂ and H₂O is catalyzed by the enzyme carbonic anhydrase, which is abundantly located in alveolar and renal tubular epithelial cells and in red blood cells. The H₂CO₃ then ionizes weakly to form very small amounts of H⁺ and HCO₃^–. The second component of the system, NaHCO₃, ionizes almost completely, forming large amounts of HCO₃ ⁻ and Na⁺.

Based on the low pK_a value of 6.1 and the limited concentrations of the two principle elements of the system (H₂CO₃ and HCO₃^–), it would be reasonable to assume that the bicarbonate system lacks significant buffering power. However, because the system is “open at both ends“ (i.e., both pCO₂ and [HCO₃^–] can be adjusted via respiratory and renal mechanisms, respectively), the effectiveness of the bicarbonate buffer system is greatly increased.

Phosphate buffer system

The phosphate buffer system is not important in the ECF but does play a major role in buffering renal tubular fluid and intracellular fluid (ICF). The main components of the system are H₃PO₄ (phosphoric acid, a triprotic weak acid) and its conjugate bases, H₂PO₄^–, HPO₄^2– _{and PO4^3– (dihydrogen phosphate, hydrogen phosphate, and phosphate, respectively, usually in the form of sodium salts). At physiological pH values, the prevailing two species are H2PO4^– and HPO4^2–.}

While the pK_a for the phosphate buffer system is 6.8, making it a potentially significant buffer in the ECF, its concentration here is only a fraction of the concentration of the bicarbonate buffer, so its contribution to ECF buffering is negligible.

In contrast, this system plays a key role in urine because phosphate is greatly concentrated in renal tubules and tubular fluid has a lower pH than ECF, making it an ideal fluid compartment for this buffer. The phosphate system is also important in the ICF for similar reasons (i.e., high concentrations are available and ICF pH is very close to 6.8).

Protein buffer system

Proteins constitute the most plentiful buffers in body fluids. While plasma proteins, notably albumin, make a modest contribution to ECF buffering, it is the contribution of intracellular proteins, which is by far the most important. Approximately 60–70% of the body‘s total chemical buffering capacity occurs in the ICF and most of this is attributed to proteins. This is a function of the high concentration of protein found in cells and the fact that most protein buffers have pK_a values fairly close to that of ICF. The most important protein-dissociable group is the imidazole ring of histidine (pK_a 6.4 to 7.0), while α-amino groups (pK_a 7.4–7.9) also play an important secondary role.

Hemoglobin is a key intracellular protein buffer. Because it is confined to red blood cells, it is also often classified as a blood buffer along with plasma proteins. Hemoglobin is responsible for more than 80% of the nonbicarbonate buffering capacity of whole blood while albumin makes up most of the other 20%. This is because hemoglobin is present in about twice the concentration and contains about three times the number of histidine residues per molecule compared to albumin.

Deoxyhemoglobin is a more effective buffer than oxyhemoglobin and this accounts for about 30% of the Haldane effect (the greater ability of deoxyhemoglobin to form carbamino compounds accounts for the other 70%). This means that oxygen unloading actually improves buffering capacity at exactly the location where additional H⁺ is being produced by dissociation of H₂CO₃ to facilitate CO₂ transport as HCO₃^–.

Other factors impacting intracellular buffering

While the previous discussion has referred to the ECF and ICF as two totally distinct compartments, it is important to recognize that changes in [H⁺] are effectively communicated between them. Transfer of key elements across cell membranes has a significant impact on buffering and is facilitated by three major processes.

First, due to its high lipid solubility, CO₂ crosses cell membranes easily where it dissolves in H₂O to form H₂CO₃ ^–, which then dissociates to H⁺ and HCO₃^–. Due to the bicarbonate system’s ineffectiveness at buffering CO₂ (i.e., volatile acid) in the ECF, almost all buffering of respiratory acidosis and alkalosis occurs intracellularly.

Second, cell membranes contain various proton–cation antiporters, which facilitate entry of H⁺ into cells in exchange for other cations, such as Na⁺ and K⁺, to maintain electroneutrality. The delivered H⁺ are then buffered by the phosphate and protein buffer systems. More than half of buffering for metabolic acidosis and about one-third of buffering for metabolic alkalosis occurs intracellularly.

Finally, the presence of the HCO₃^–/Cl ^{– antiporter (or chloride pump) facilitates the movement of HCO₃ – _{produced from the dissociation of H2CO3 out of the cell in exchange for Cl–. This is referred to as the chloride shift.}}

The isohydric principle

Chemical buffer systems do not operate independently but work together in equilibrium with each other. Because there is only one [H⁺] in a given compartment, anything that changes the balance of one buffer system will affect the balance of all the others. This is known as the isohydric principle. Clinically, this means that measuring the concentrations of any one acid–base pair will provide a picture of overall acid–base balance in the body. Conventionally, the components of the bicarbonate system (i.e., H₂CO₃ measured as pCO₂, and HCO₃^–) are measured because they are accessible and easily determined. Blood gas machines measure pH and pCO₂ directly and the [HCO₃^–] is then calculated using the Henderson–Hasselbalch equation.

Respiratory regulation of acid–base balance

The second line of defense against acid–base abnormalities involves the respiratory system, which is able to excrete volatile acid (CO₂) and control [H⁺] by initiating changes in ventilation. Ventilatory changes occur rapidly and have profound effects on [H⁺]. Changes in alveolar ventilation are inversely proportional to changes in arterial pCO₂ (P_aCO₂) and directly proportional to total body CO₂ production. Changes in P_aCO₂, in turn, are inversely proportional to pH according to the Henderson–Hasselbach equation. Therefore, a decrease in alveolar ventilation will cause an increase in P_aCO₂ and a decrease in pH, while an increase in alveolar ventilation will cause a decrease in P_aCO₂ and an increase in pH.

Not only does alveolar ventilation impact pH, but the reverse also occurs. A decrease in pH will stimulate ventilatory drive while an increase in pH will inhibit it. Thus, the respiratory system operates as a typical negative feedback loop to maintain P_aCO₂ and therefore [H⁺] within a narrow range. Peripheral and central chemoreceptors sense changes in P_aCO₂, the respiratory center of the medulla integrates this information, and the muscles of respiration respond accordingly. This type of acid–base regulation is referred to as physiological buffering because it is an open system, as opposed to chemical buffering, which is a closed system. In general, the overall buffering capacity of the respiratory system is one to two times greater than the buffering power of all ECF chemical buffers combined.

Renal regulation of acid–base balance

While the lungs are responsible for elimination of volatile acid (CO₂), the kidneys serve as the primary means by which fixed, metabolic acids are eliminated from the body. Despite the fact that the amount of fixed acid produced daily is far, far smaller than the amount of CO₂ produced, the renal system is critical as there is no other mechanism for excretion of these acids. Quantitatively, even more important than excretion of fixed acids is the kidneys’ ability to reabsorb bicarbonate, thereby conserving the primary buffer system of the ECF. The kidneys accomplish these two goals through three fundamental processes: (1) active secretion of H⁺ into the tubular lumen, (2) indirect reabsorption of filtered HCO₃^{– into tubular cells and ultimately the blood, and (3) production of new HCO₃– in tubular cells involving the phosphate and ammonia buffer systems.}

The first two mechanisms are interrelated in that the process of active H⁺ secretion into the renal tubular lumen ultimately results in reabsorption of HCO₃^–. The cellular mechanisms of H⁺ secretion depend on the location within the kidney. In the proximal tubules, the thick ascending loop of Henle and the early distal tubule, H⁺ is secreted via secondary active transport. In the intercalated cells, H⁺ is secreted by primary active transport.

In both cases secreted H⁺ combines with filtered HCO₃^–, forming H₂CO₃, which dissociates into CO₂ and H₂O. The CO₂ diffuses easily into the tubular cell where it recombines with H₂O under the influence of carbonic anhydrase to generate a new H₂CO₃ molecule, which in turn dissociates into HCO₃^{– and H+. The HCO₃– is then transported across the basolateral membrane by either Na+–HCO₃– cotransport or Cl––HCO₃– exchange. Under normal circumstances, each time an H+ is formed in a tubular cell an HCO₃– is also formed and ultimately released back into the blood. Therefore, it can be said that HCO₃– and H+ “titrate“ each other in the tubules.}

In acidosis, there is excess H⁺ relative to HCO₃^– but only a small part of this can be excreted in the ionic form, as urine pH has a lower limit of approximately 4.5. Consequently, excretion of large amounts of urine H⁺ is accomplished by combining it with phosphate and ammonia buffers in the tubular lumen. This buffering actually results in the formation of “new“ (i.e., not previously filtered) HCO₃^–, which ultimately enters the blood. The ammonia system is quantitatively more important than the urinary phosphate system. In the proximal tubular cells, glutamine is metabolized to form ammonium and HCO₃^–. The ammonium is secreted into the tubular lumen while the new HCO₃^{– is transported back across the basolateral membrane. In the collecting tubules, H+ is secreted into the tubular lumen as described previously, but now it combines with ammonia, which is readily able to diffuse across the luminal membrane. The result is ammonium, which is much less diffusable and becomes trapped in the urine, thereby providing a means of H+ excretion. For each ammonium excreted, a new HCO₃– is formed and added back to the blood.}

Clinically, it is possible to quantify net renal acid excretion by determining HCO₃^{– excretion, the amount of new HCO₃– added to the blood (which is calculated by measuring urinary ammonium concentration), and urinary titratable acid. Titratable acid refers to the rest of the non-HCO₃–, nonammonium buffer excreted in the urine (i.e., mainly phosphate).}

As secretion of H⁺ into the tubular lumen is a key step in all aspects of renal acid–base balance, it is not surprising that this process is tightly regulated. Key factors that may increase or decrease H⁺ secretion include changes in the following: pCO₂, [H⁺], ECF fluid volume, angiotensin II levels, aldosterone levels, and potassium levels.

Temperature effects on acid–base balance

Evaluations of acid–base balance in a clinical setting typically involve measurement of ECF components simply because blood is an easily accessible fluid. However, it has been suggested that the intracellular environment is, in fact, the key compartment due to the phenomenon of ion trapping. There are two competing theories that seek to explain the regulation of intracellular pH across variations in temperature, the alpha-stat hypothesis, and the pH-stat hypothesis. While an in-depth discussion of these theories is not possible here, a brief definition of each is provided as they are relevant to the clinical interpretation of blood gas values.

The alpha-stat hypothesis states that the degree of ionization (referred to as “alpha“) of the imidazole groups of intracellular proteins remains constant despite changes in temperature. In other words, the ideal intracellular pH is the pH of neutrality (i.e., the state where [H⁺] = [OH^–]) and this pH will vary with changes in temperature. If the alpha-stat theory is correct, then blood gas values should not be corrected for the patient‘s body temperature.

The pH-stat theory states that the intracellular pH should be kept constant despite changes in temperature. If the pH-stat theory is correct, then temperature-corrected blood gas values should be used for interpretation. It should be noted that all blood gas analyzers are set to 37 °C and all measurements are performed at this temperature. When other patient temperatures are entered the computer in the machine uses various correction formulae to calculate the so-called corrected values.

While both approaches continue to be used in clinical practice, the reality is, with significant variations in patient temperature, we simply do not fully understand the complexity of effects on cellular metabolism, vascular function, and respiration. Therefore, until definitive evidence is produced, both corrected and uncorrected blood gas values are of questionable utility in patients with significant deviations from 37°C.

The traditional approach to acid–base analysis

Table 7.2. Causes of acid–base abnormalities in dogs and cats classified according to the traditional approach

Simple acid–base disorders

A simple acid–base disorder is one where there is a single primary etiologic disturbance. According to the traditional approach, there are four simple acid–base disorders to consider: (1) respiratory acidosis, (2) metabolic acidosis, (3) respiratory alkalosis, and (4) metabolic alkalosis. Respiratory disorders are caused by abnormal processes that alter pCO₂ levels, while metabolic disorders are caused by abnormal processes that alter [HCO₃^–]. A list of common causes of the four primary acid–base disorders in dogs and cats classified according to this approach can be found in Table 7.2.

Secondary or compensatory acid–base changes usually occur in response to most primary acid–base abnormalities in an effort to restore [H⁺] to normal. In simple acid base disorders, these compensatory responses are predictable and a stepwise analysis is conducted according to Figure 7.1. With primary metabolic abnormalities, respiratory compensation involves either hyper- or hypoventilation. In most cases, the respiratory response occurs rapidly and is complete within hours, assuming the metabolic disturbance remains stable. With primary respiratory abnormalities, the compensatory response is biphasic. Acutely, the metabolic response is limited to chemical buffering, which may have a very modest impact on [HCO₃^–]. With more chronic respiratory disorders, the kidneys respond and are able to excrete or reabsorb HCO₃^{– as appropriate. This renal response is slow to be initiated and takes 2–5 days to reach its maximum potential.}

The degree of expected compensation for acid–base disorders varies among species. In clinical canine patients, guidelines derived from healthy experimental dogs are commonly used and are presented in Table 7.3. There is limited data available for evaluating compensatory responses in cats but experimental evidence based on a small number of animals suggests cats do not appear to initiate a ventilatory response to metabolic acidosis. Consequently, extrapolation of canine values to feline patients may be misleading and cannot be recommended.

Figure 7.1. Analysis of simple acid–base disorders according to the traditional approach.

Table 7.3. Expected compensatory changes to primary acid–base disorders in dogs according to the traditional approach

Mixed acid–base disorders

A mixed acid–base disorder is present when there are two or more primary etiologic disturbances present simultaneously. Mixed disorders may be additive (such as respiratory and metabolic acidosis) or offsetting (such as respiratory alkalosis and metabolic acidosis) with regard to their effects on pH. A triple disorder is an example of a complex acid–base disorder where a respiratory disturbance further complicates concurrent metabolic acidosis and metabolic alkalosis. With the traditional approach, differentiating a mixed acid–base disorder from a simple disorder and its compensatory response is accomplished by evaluating the magnitude of the compensation as discussed above and shown in Table 7.3. If the compensatory response is not within the range expected for the primary disorder then a mixed disorder is diagnosed. Caution must be exercised when evaluating the appropriateness of metabolic compensation to a primary respiratory disorder as it will depend on the chronicity of the respiratory abnormality, which may be difficult to ascertain clinically.

In general, the possibility of a mixed acid–base disorder should be suspected if one or more of the following are observed: (1) the presence of a normal pH with abnormal PCO₂ and/or [HCO₃^–], (2) a pH change in a direction opposite that predicted for the known primary disorder, and (3) pCO₂ and [HCO₃^–] changing in opposite directions.

Mixed acid–base disorders are more difficult to decipher than simple disorders and a diagnosis must be made in the context of the patient’s clinical signs and other diagnostic information. In most cases, the clinical significance of the acid–base perturbations will be a function of the underlying cause(s), which should become the focus of treatment.

Limitations of the traditional approach

While intuitively appealing and relatively simple in nature, the traditional bicarbonate-based approach remains essentially descriptive. In the mid-1960s, a group of researchers expanded on this concept in an attempt to better quantify and predict the nature of various acid–base disturbances. This approach, using acid–base maps or nomograms, became known as the Boston approach and these nomograms are still found in many textbooks. Unfortunately, the Boston/traditional approach has significant limitations. It fails to take into account the influence of nonbicarbonate buffer systems (i.e., plasma proteins and phosphate) as well as the effect of serum electrolytes (notably Na⁺, K⁺. and Cl^–). Also, it incorrectly assumes that both pCO₂ and HCO₃^– are independent variables when in fact only pCO2 is independent.

Consequently, the traditional approach is unable to provide insight into complex mixed acid–base disorders and can actually be misleading in certain circumstances. Metabolic acidosis, in particular, poses significant problems for the traditional approach. This abnormality has been reported to be the most common acid–base disturbance in dogs and cats and accompanies such diverse conditions as sepsis, trauma, renal failure, and endocrine emergencies. In these settings, the traditional approach is not helpful. While it still retains a place in clinical acid–base analysis, its use should be reserved for patients with simple acid–base disorders and normal serum albumin and phosphate concentrations.

Alternative approaches to acid–base analysis

The base excess approach to acid–base analysis

The anion gap approach to acid–base analysis

In the case of organic acidosis, fixed acids such as lactic acid or ketoacids are added to plasma where they rapidly dissociate into their anions (which are not measured as part of a typical biochemical profile) and H⁺ (which is buffered by HCO₃^–). The net effect is a decrease in measured anions (i.e., HCO₃^–) with no change in Cl^–, leading to an increased AG. These types of acidosis are therefore classified as increased AG or normochloremic metabolic acidosis. In contrast, the loss of HCO₃^– that may accompany gastrointestinal fluid loss or renal tubular acidosis is associated with Cl^– retention so the AG doesn’3 change. Consequently, these types of acidosis are classified as normal AG or hyperchloremic metabolic acidosis. Causes of metabolic acidosis classified according to the AG approach are shown in Table 7.4.

Table 7.4. Causes of metabolic acidosis in dogs and cats classified according to the AG approach

A significant factor that confounds interpretation of the AG is hypoalbuminemia. Since albumin is the major unmeasured anion in the normal state, the presence of hypoalbuminemia will decrease the AG and obscure the presence of anions, which would otherwise increase it. In other words, a patient with concurrent lactic acidosis and hypoalbuminemia may have a normal calculated AG and this would be clinically misleading. In dogs, the AG can be adjusted for changes in protein concentration by using the following formula:

While not as substantial as the effect of albumin, phosphate can also impact the AG. In dogs, the AG can be adjusted for both albumin and phosphate by using the following formula:

While the AG is a useful tool in certain situations it remains limited in its ability to elucidate complex mixed acid–base disorders. Attempts to optimize its usefulness have been explored in humans and have led to the concept of the delta ratio or the delta gap. Both of these parameters attempt to relate the change in AG to the change in [HCO₃^–], each from their respective reference values. Neither the delta ratio nor the delta gap has been studied in dogs or cats.

The strong ion approach to acid–base analysis

In 1983, Stewart published his now widely accepted model of acid–base physiology, which has been referred to by various names including the Stewart approach, the quantitative approach, the physicochemical approach, and the strong ion approach. According to Stewart there are three independent determinants of acid–base balance: (1) pCO₂, (2) the difference between strong cations and strong anions (strong ion difference or SID), and (3) the total concentration of weak acids (A_TOT). As is the case with the traditional approach, the pCO₂ variable is the primary determinant of the respiratory component of plasma pH. The SID and A_TOT variables provide independent measures of the metabolic component of plasma pH. Therefore, according to the strong ion approach, there are four possible primary metabolic acid–base disorders instead of only two with the traditional approach. Metabolic acidosis is caused by a decrease in SID or an increase in A_TOT, while metabolic alkalosis is caused by an increase in SID or a decrease in A_TOT. Please refer to Table 7.5 for common metabolic acid–base disorders classified according to the strong ion approach.

SID

Strong ions (also known as nonbuffer ions) are defined as any ion that is completely dissociated at physiological pH and include Na⁺, K⁺, Ca²⁺, Mg²⁺. and Cl⁺. as well as unmeasured ions such as lactate, β-hydroxybutyrate, acetoacetate, and sulfate. The SID is a term used to describe the difference between the sum of strong cations and the sum of strong anions. It is usually represented by the following equation (where [A^–] refers to all unmeasured strong anions):

Table 7.5. Causes of metabolic acid–base abnormalities classified according to the strong ion approach

SID acidosis (↓ SID)

Dilution acidosis (↓ Na+)

• With hypervolemia

– Severe liver disease

– Congestive heart failure

– Nephrotic syndrome

• With normovolemia

– Psychogenic polydipsia

– Hypotonic fluid infusion

• With hypovolemia

– Vomiting

– Diarrhea

– Hypoadrenocorticism

– Third space loss

– Diuretic administration

Hyperchloremic acidosis (↑ Cl–corr)

• Loss of Na+ relative to Cl–

– Diarrhea

• Gain of Cl– relative to Na+

– Fluid therapy (0.9% NaCl, 7.2% NaCl, KCl-supplemented fluids

• Cl– retention

– Renal failure

– Hypoadrenocorticism

Organic acidosis (↑ unmeasured strong ions)

• Uremic, keto- or lactic acidosis

• Toxicities

– Ethylene glycol

– Salicylate

SID alkalosis (↑ SID)

Concentration alkalosis (↑ Na+)

• Pure water loss

– Water deprivation

– Diabetes insipidus

• Hypotonic fluid loss

– Vomiting

– Nonoliguric renal failure

– Postobstructive diuresis

Hypochloremic alkalosis (↓ Cl– corr)

• Gain of Na+ relative to Cl–

– 1sotonic or hypertonic

NaHCO3– administration

• Loss of Cl– relative to Na+

– Vomiting of stomach contents

– Thiazide or loop diuretics

ATOT acidosis (↑ ATOT )

Hyperalbuminemia

• Water deprivation

Hyperphosphatemia

• Translocation

– Tumor cell lysis

– Tissue trauma/rhabdomyolysis

• 1ncreased intake

– Phosphate-containing enemas

– 1ntravenous phosphate

• Decreased loss

– Renal failure

– Urethral obstruction

– Uroabdomen

ATOT alkalosis (↓ATOT)

Hypoalbuminemia

• Decreased production

– Chronic liver disease

– Acute phase response to inflammation

– Malnutrition/starvation

• Extracorporeal loss

– Protein-losing nephropathy

– Protein-losing enteropathy

• Sequestration

– 1nflammatory effusions

– Vasculitis

Under physiologically normal conditions, the SID of plasma is approximately 40–44mEq/L and this difference is balanced by the negative charge on HCO₃^– and A_TOT. Under pathological conditions, additional strong anions accumulate and [A^–] becomes clinically significant, indicating organic acidosis. The SID may be calculated by measuring as many strong ions as possible and summing their charges. This original definition has now come to be known as the apparent SID (SID_app), which distinguishes it from the effective SID (to be discussed later). With the capability to routinely measure lactate concentrations in the clinical setting, some recommend the addition of [lactate] to the SID_app calculation as follows:

In fact, because Na⁺ and Cl^– are quantitatively the most important strong ions, SID_app is often further simplified in clinical situations as follows:

Concentration of weak acids (A_TOT)

Stewart used the term A_TOT to reflect the sum of all weak acids in both the dissociated and undissociated states (i.e., A^– plus HA) according to the law of conservation of mass. The major contributors to A_TOT are albumin, globulins, and inorganic phosphate, and these weak acids are also sometimes referred to as buffer ions. Changes in the concentrations of any of these, either directly or through changes to plasma free water content, will affect pH.

Simplifications of the strong ion approach

The strong ion approach provides a sound mechanistic framework to explain why plasma pH changes in any setting and successfully integrates electrolyte and acid–base physiology. Unfortunately, Stewart’s equations are mathematically complex and thus not easily applicable to the clinical setting. Consequently, a number of simplified modifications have been developed to facilitate clinical use of Stewart’s strong ion approach.

The effective SID modification

Some 10 years after Stewart published his strong ion theory, Figge et al. proposed a modification to Stewart’s approach, which attempts to quantify the effects of unmeasured strong anions using a parameter that he referred to as the strong ion gap (SIG). The term effective SID (SID_eff) was proposed to represent the contribution of weak acids. Therefore, unmeasured strong anions (i.e., the SIG) can be estimated by subtracting the SID_eff (calculated using formulas involving the weak acids HCO₃^–, albumin, and phosphate) from the SID_app (calculated according to the equation presented previously). Despite its potential usefulness in humans, Figge’s formulas for calculation of SID_eff are complex and have not been evaluated in dogs and cats, so this approach is not currently feasible in veterinary medicine. The general concept of the SIG, however, has been revisited by others and will be discussed again below.

The base excess modification

In 1993, Fencl et al. published an additional modification of the Stewart approach that resurrected the concept of SBE and combined this with principles of the strong ion theory. The so-called Fencl–Stewart approach uses a set of equations to estimate the magnitude of effect of changes in SID and a_tot on the SBE value reported by blood gas analyzers. Equations for four distinct parameters have been proposed that can be calculated based on routine laboratory measurements: (1) the albumin effect or SBE_Alb, (2) the phosphate effect or SBE_Phos, (3) the free water effect or SBE_FW (based on sodium), and (4) the chloride effect or SBE_Cl. With the availability of lactate analyzers, a fifth parameter, the lactate effect or SBE_Lact, can also be calculated. A sixth effect, the unmeasured strong ion effect or SBE_XA, can then be calculated by subtracting the sum of the individual SBE effects from the reported SBE value or SBE_reported (see Table 7.6). Some authors use different terminology and refer to the SBE_XA as the base excess gap (BEG). Values of SBE_XA less than –5 mEq/L are suggestive of an increase in unmeasured strong anions.

Table 7.6. Formulas for calculating contributions of key parameters according to the base excess modification of the strong ion approach

SBE effect (in mEq/L)	Formulas
Effects relating to ATOT:
(1) SBEAlb	3.7([AIbnormaI ref] – [Albmeasured]) in g/dL 0.37([AIbnormaI ref] – [Albmeasured]) in g/L])
(2) SBEphos	1.8([PhosnormaI ref] – [Phosmeasured]) in mmol/L 0.58([PhosnormaI ref] – [Phosmeasured]) in mg/dL
Effects relating to SID:
(3) sbefw	Dogs 0.25([Na+measured] – [Na+normal ref]) in mEq/L Cats 0.22([Na+measured] – [Na+normal ref]) in mEq/L
(4) sbecl	[Cl–normal ref] – [Cl–corr] in mEq/L, where Cl– corr = [Cl– measured] x ([Na+normal ref]/[Na+measured])
(5) SBELact	–1[Lactmeasured] in mmol/L
(6) sbexa (also known as BEG)	SBEreported – (SBEAlb + SBEPhos + sbefw + sbeci + SBELact)

See text for additional definitions.

Alb_{normal ref}, midpoint of the albumin reference range; Alb_measured, patient’ s albumin concentration; Phos_{normal ref}, midpoint of the phosphorous reference range; Phos_measured, patient’s phosphorous concentration; Na⁺_{normal ref}, midpoint of the sodium reference range; Na⁺_measured, patient’ s sodium concentration; Cl^–_{normal ref}, midpoint of chloride reference range; Cl^–_corr, corrected chloride; Cl^–_measured, patient’s chloride concentration; Lact_measured, patient’s lactate concentration.

Table 7.7. Summary of metabolic acid–base influences identified by the base excess modification of the strong ion approach

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