Keywords

Management of acute acid-base changes is a common part of the practice of anesthesiology. Historically, analysis of these changes focuses on the Henderson–Hasselbalch equation and the relationship among the three parameters—pH, PCO2, and HCO3 −. Respiratory changes are easily recognized and treated by changes in the partial pressure of CO2. Metabolic changes are easily recognized by bicarbonate or base excess (BE) abnormality; however, the treatment is not simple because the bicarbonate or BE is only a reflection of the actual problem. For example, a lactic acidosis will cause a change in bicarbonate, and if only looking at the bicarbonate, this lactic acidosis cannot be differentiated from a hyperchloremic metabolic acidosis. Unlike PCO2, HCO3 − is not an independent determinant of pH. In other words, we know that by changing the ventilator setting we can change the PCO2 and directly change the pH; however, a similar relationship between HCO3 − and pH does not exist.

The clinician can gain further information as to the cause of a metabolic problem through use of the anion gap. The anion gap is based on the concept of electroneutrality, that is, the sum of the cations must equal the sum of the anions in solution. The classic acid-base theory focuses on electroneutrality and the Henderson–Hasselbalch equation, but unfortunately, reliance on these two concepts fails to identify the independent determinants of metabolic acid-base change.

In this chapter a brief overview of classic acid-base theory will be discussed, and a new approach, the “physicochemical” acid-base approach will be introduced. This approach incorporates the Henderson–Hasselbalch equation and electroneutrality but rearranges the importance of each in acid-base analysis. This new approach identifies the independent determinants of pH, and by determining these variables, a better understanding is gained of the impact that fluid and electrolyte management have on acid-base status.

Classic Acid-Base Theory

Henderson–Hasselbalch Equation

Traditionally, assessment of [H+] or pH abnormalities has focused on the Henderson–Hasselbalch equation and its two primary components, PCO2 and HCO3 −.

$$ \begin{array}{l}\mathrm{p}\mathrm{H}=\mathrm{p}{K}_{\mathrm{a}}+ \log \left[{{\mathrm{HCO}}_3}^{-}\right]/\left[{\mathrm{CO}}_2\right]\\ {}\left({\mathrm{pK}}_{\mathrm{a}}\;\mathrm{is}\;\mathrm{the}\;\mathrm{negative}\kern0.24em \log\ \mathrm{of}\kern0.24em \mathrm{the}\kern0.24em \mathrm{acid}\kern0.24em \mathrm{dissociation}\kern0.24em \mathrm{constant}\right)\end{array} $$

From this equation, respiratory disorders are defined by changes in CO2 and metabolic disorders result from changes in the HCO3 −. Increases in CO2 cause a respiratory acidosis while decreases in CO2 cause a respiratory alkalosis. Similarly, an increase in HCO3 − causes a metabolic alkalosis and a decrease causes a metabolic acidosis. From this it would appear that by this equation, [H+] is determined by two variables, CO2 and HCO3 −.

By reviewing the derivation of the Henderson–Hasselbalch equation we discover that these two variables are interdependent and not independent as these definitions would suggest. The Henderson–Hasselbalch equation is derived from the carbonic acid equilibrium and its associated equilibrium equation.

$$ \begin{array}{l}{\mathrm{CO}}_2+{\mathrm{H}}_2\mathrm{O}={\mathrm{H}}_2{\mathrm{CO}}_3={{\mathrm{H}\mathrm{CO}}_3}^{-}+{\mathrm{H}}^{+}\\ {}\kern0.72em \left(\mathrm{carbonic}\kern0.24em \mathrm{acid}\kern0.24em \mathrm{equilibrium}\right)\end{array} $$
$$ \begin{array}{l} {\mathrm{K \, (equilibrium \, constant)}} = {[\mathrm {CO_2}][\mathrm {H_{2}O}]/[\mathrm {H^+}][\mathrm {HCO_3^-}]}\\ \qquad \qquad \kern0.72em \left(\mathrm{(equilibrium)}\kern0.24em \mathrm{equation}\right)\end{array} $$

From this equilibrium, it is seen that an increase in CO2 results in hydration of the CO2 and an increase in H2CO3. The H2CO3 will partially dissociate yielding equimolar quantities of H+ and HCO3 −. Both changes will be dictated by the equilibrium equation and its associated constant. As a result a change in CO2 must be matched by a change in HCO3 −. Thus, CO2 and HCO3 − are dependent on each other and not truly independent determinants of pH as is commonly implied.

Metabolic Indices

The CO2–HCO3 − relationship has resulted in the proposal of multiple metabolic indices. These indices are intended to circumvent the CO2–HCO3 − relationship. One of these indices is the standard bicarbonate concentration (SBC). The SBC attempts to correct for the interrelationship by standardization of the CO2. By exposing a blood sample to CO2 at a partial pressure of 40 mmHg, the sample will equilibrate to this standard CO2 partial pressure. From this standardization, any deviation of the HCO3 − from normal will be an indicator of a nonrespiratory problem. In 1960, Siggaard-Andersen proposed that the BE be the standard metabolic index.

Anion Gap

When using the SBC or the BE, the origin of a metabolic deviation is left unexplained. For instance, an abnormal BE would not tell a clinician whether a metabolic acidosis is a result of ketoacidosis, lactic acidosis, or hyperchloremia. For further understanding of a metabolic acidosis, the anion gap is utilized.

The anion gap is based on the concept of electroneutrality. The sum of the positive ions and the negative ions in a solution must be zero, (Σcations = Σanions). In other words, any body fluid will have no net charge. This charge balance means that the charge of Na+, K+, Mg2+, Ca2+, and H+ must be balanced by an equal and opposite charge of Cl−, SO4 2−, PO4 3−, CO3 2−, HCO3 −, OH−, lactate, and the charges on the proteins. The anion gap can be defined as

$$ \begin{array}{l}\mathrm{Anion}\kern0.24em \mathrm{gap}=\left[{\mathrm{Na}}^{+}\right]-\left(\left[{\mathrm{Cl}}^{-}\right]+\left[{{\mathrm{HCO}}_3}^{-}\right]\right)\\ {}=\mathrm{Unmeasured}\kern0.24em \mathrm{anions}-\mathrm{Unmeasured}\kern0.24em \mathrm{cations}\end{array} $$

By this definition K+, Ca2+, and Mg2+ have been relegated into a grouping of unmeasured cations. Likewise, SO4 2−, PO4 3−, lactate, and the proteins have been grouped into unmeasured anions. An increase in anion gap represents an increase in unmeasured anions or a decrease in unmeasured cations, and a decrease in anion gap can be caused by a decrease in unmeasured anions or an increase in unmeasured cations.

Traditional Approach of Arterial Blood gas Analysis

Arterial blood gases are routinely used to assess acid-base disturbances, which can be analyzed as follows (Fig. 45.1, Table 45.1):

Fig. 45.1
figure 1

Approach to determine acid-base disorder

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Table 45.1 Simplified approach to blood gas analysis (approximate equality)
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  1. 1.

    pH—The normal pH is 7.35–7.45. A blood pH less than 7.35 is termed as acidosis, while a pH higher than 7.45 is termed as alkalosis.

  2. 2.

    PaCO2—The normal PaCO2 is 35–45 mmHg. A PaCO2 less than 35 mmHg is termed as respiratory alkalosis, while a PaCO2 more than 45 mmHg is termed as respiratory acidosis. Adequacy of ventilation can be assessed by calculation of the dead space (V D) to tidal volume (V T) ratio, using the Bohr dead space equation:

    $$ \begin{array}{c}{V}_{\mathrm{D}}/{V}_{\mathrm{T}}=\left({\mathrm{P}}_{\mathrm{A}}{\mathrm{CO}}_2-{\mathrm{ETCO}}_2\right)/{\mathrm{P}\mathrm{ACO}}_2,\\ {}\mathrm{the}\kern0.24em \mathrm{normal}\kern0.24em \mathrm{ratio}\kern0.24em \mathrm{should}\kern0.24em \mathrm{be}\kern0.24em \mathrm{less}\kern0.24em \mathrm{than}\kern0.24em 0.3\end{array} $$
  3. 3.

    PaO2—Hypoxia is defined as a PaO2 < 60 mmHg. Adequacy of ventilation can be assessed by measuring the

    1. (a)

      Alveolar-arterial gradient of oxygen: PAO2 − PaO2

      $$ \begin{array}{c}{\mathrm{PAO}}_2=\left(\begin{array}{l}\mathrm{atmospheric}\kern0.24em \mathrm{pressure}-\\ {}\mathrm{water}\kern0.24em \mathrm{vapor}\kern0.24em \mathrm{pressure}\end{array}\right)\\ {}\times {\mathrm{FiO}}_2-{\mathrm{PaCO}}_2/0.8\end{array} $$

      where 0.8 is the respiratory quotient and is the ratio of CO2 produced to O2 consumed. Normally, about 80 % of CO2 is produced for 100 % O2 consumed (200 ml CO2:250 ml of O2)

    2. (b)

      Ratio of PaO2:FiO2, the P/F ratio. The lower the ratio, the worse the oxygenation. A P/F ratio less than 300 denotes acute lung injury, whereas a P/F ratio less than 200 denotes ARDS.

  4. 4.

    HCO3—The normal HCO3 is 22–26 mmol/L. A HCO3 less than 22 is termed as metabolic acidosis, while a HCO3 more than 26 is termed as metabolic alkalosis.

  5. 5.

    Assess compensatory changes.

Regulation of pH in the Body

Regulation of pH or the hydrogen ion concentration in the human body occurs mainly via three processes: the buffer systems, central and peripheral chemoreceptors, and the renal system. Causes and compensatory mechanisms for acid base disturbances are summarized in Tables 45.2 and 45.3. Adverse effects of acid-base disturbances are summarized in Table 45.4.

Table 45.2 Causes and compensation of respiratory acidosis and alkalosis
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Table 45.3 Causes and compensation of metabolic acidosis and alkalosis
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Table 45.4 Deleterious effects of respiratory and metabolic acidosis/alkalosis
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Buffer Systems

Buffers are chemicals/substances which tend to maintain the pH of the body fluids at 7.4. The main buffer systems in the body are the bicarbonate and the hemoglobin buffer systems. Additionally, some proteins and phosphate also have buffering capabilities.

Bicarbonate:

CO2 combines with water to form carbonic acid; the reaction accelerated by the enzyme carbonic anhydrase. The carbonic acid then dissociates into hydrogen and bicarbonate ions. The bicarbonate reaches the lung, where an opposite reaction occurs. Hydrogen ions are added to the bicarbonate to form carbonic acid, which dissociated into CO2 and water. The CO2 is then exhaled.

$$ \begin{array}{l}{\mathrm{CO}}_2+{\mathrm{H}}_2\mathrm{O}\overset{\mathrm{Carbonic}\kern0.24em \mathrm{anhydrase}}{\leftarrow }{\mathrm{H}}_2{\mathrm{CO}}_3\\ {}\kern0.84em \left(\mathrm{occurs}\;\mathrm{in}\;\mathrm{the}\;\mathrm{tissues}\right)\kern1.68em \end{array} $$
$$ \begin{array}{l}{{\mathrm{H}\mathrm{CO}}_3}^{-}+{\mathrm{H}}^{+}\to {\mathrm{H}}_2{\mathrm{CO}}_3\to {\mathrm{CO}}_2+{\mathrm{H}}_2\mathrm{O}\\ {}\kern1.56em \left(\mathrm{occurs}\kern0.24em \mathrm{in}\kern0.24em \mathrm{the}\kern0.24em \mathrm{lungs}\kern0.24em \mathrm{and}\kern0.24em \mathrm{kidneys}\right)\end{array} $$

Hemoglobin:

Hemoglobin also plays a role in buffering CO2. A similar reaction, as above, takes place in erythrocytes. CO2 diffuses freely into the erythrocytes, where it combines with water to form carbonic acid. The latter then dissociates into hydrogen and bicarbonate ions. The hydrogen ions are absorbed by the hemoglobin, and the bicarbonate ions are exchanged for chloride (Chloride shift) to maintain electroneutrality. A reverse reaction happens in the pulmonary capillaries, where bicarbonate combines with the hydrogen ions to form carbonic acid and ultimately CO2, which is exhaled. Also hemoglobin, especially deoxyhemoglobin, can directly combine with CO2 to form carbaminohemoglobin, which facilitates removal of CO2 from peripheral tissues.

Chemoreceptors

CO2 freely passes the plasma membrane of cells. In the brain it decreases the pH of CSF, thereby stimulating the central chemoreceptors causing an increase in minute ventilation. The increase in minute ventilation decreases the PaCO2 and maintains the pH. In addition, the peripheral chemoreceptors, which are present in the carotid bodies and the aortic arch, sense a decrease in blood pH or a decrease in PaO2 and stimulate the respiratory center in the brain to increase the minute ventilation.

Renal Buffering

Proximal tubule cells of the kidney absorb most of the bicarbonate from the glomerular filtrate and secrete hydrogen ions into the tubules. The kidneys thus regulate the pH by altering the absorption of bicarbonate and the secretion of hydrogen ions. Renal tubal acidosis results from wasting of bicarbonate ions due to a defect in the absorption of bicarbonate. The drug acetazolamide can cause a normal anion gap acidosis by inhibiting the reabsorption of bicarbonate ions in the renal proximal tubule.

The PhysicoChemical Approach

In 1981, Stewart proposed a change in the approach to acid-base problems. He recognized that multiple chemical interactions affect [H+] and that the carbonic acid equilibrium was just one of these interactions (Fig. 45.2). The focus of his approach is on the concept of electroneutrality, the basis of the anion gap. He incorporated the multiple chemical equilibria that affect [H+], including the carbonic acid equilibrium, into a single electroneutrality equation.

Fig. 45.2
figure 2

Relationship of SID change and H+ and OH− change (SID-strong ion difference)

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From this mathematical development, he found that [H+] is dependent on three independent variables: (1) the strong ion difference (SID) which is a modified anion gap, (2) the PCO2, and (3) the total weak acid concentration [Atot], which is primarily composed of protein and phosphate. For most purposes, the weak acid concentration does not change during a surgical procedure, so we can define the CO2 as the respiratory component which drives pH and the SID as the metabolic component which changes pH.

For simplicity, the SID = [Na+] + [K+] − [Cl−] − [lactate−]

Remembering the law of electroneutrality, we can think of H+ and OH− as charge buffers. As the relationship of the strong ions changes, so does the H+ and OH− change, as reflected in Fig. 45.2. For instance, an increase in the negatively charged chloride will result in a decrease in the SID and an increase in H+ to maintain electroneutrality, which results in acidosis. Because of the inverse relationship between H+ and OH−, it is sometimes easier to assess pH changes through changes in the basic OH−. Increased OH− leads to alkalosis, decreased OH− results in acidosis.

Specific Metabolic Abnormalities

From this general approach more specific metabolic problems can be addressed. There are three general mechanisms by which SID changes: changing the water content of plasma (contraction alkalosis and dilutional acidosis), changing the Cl− (hyperchloremic acidosis and hypochloremic alkalosis), and increasing the concentration of unidentified anions (organic acidosis).

SID Free Water Change

Dilutional Acidosis

Development of a dilutional acidosis is best illustrated by an example. If a liter of water contains 140 mEq/L of sodium and 110 mEq/L of chloride then the SID of that solution is 30 mEq. This represents a positive charge excess which needs to be balanced by a negative charge of 30 mEq. Hydroxyl ions (OH−) act as this charge equalizer. If we were to add another liter of water without adding any more electrolytes, the solution would contain 70 mEq/L of sodium and 55 mEq/L of chloride (Fig. 45.3). Now the SID is 15 mEq. Because we have decreased the positive charge contribution of the SID from 30 to 15 mEq, a fall in OH− would occur and a “dilutional” acidosis would be seen. In the operating room, dilutional acidosis can theoretically occur as part of the Trans Urethral Resection of the Prostate (TURP) syndrome.

Fig. 45.3
figure 3

Dilutional acidosis

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Contraction Alkalosis

Contraction alkalosis can be seen in the perioperative patient who has been fluid restricted or treated with diuretics. It can also be seen intraoperatively if evaporative loss of free water is not replaced. Similar to dilutional acidosis, this problem arises from free water and SID changes. If we return to the original volume of water containing 140 mEq/L of sodium and 110 mEq/L of chloride (as above), and boil off half of the water, it would result in a sodium concentration of 280 mEq/L and a chloride concentration of 220 mEq/L. Now the SID is 60 mEq, and the OH− “buffer” would increase so that the solution would remain electrically neutral.

Treatment of contraction alkalosis simply requires free water administration in the form of hypotonic solutions. Using the beaker model, treatment can be explained mechanistically. We would now add one liter of 0.45 % NaCl solution containing 77 mEq of Na+ and 77 mEq of Cl−. The final electrolyte concentration would contain 238 mEq of Na+ and 198 mEq of Cl−, and a SID of 40 mEq. By the use of this hypotonic fluid, we have changed the SID from 60 to 40 mEq resulting in a decrease in the OH− and a correction of the alkalosis.

SID Chloride Change

Hypochloremia

Chloride shifts occur in relation to gastrointestinal abnormality. If the hyperchloremic gastric contents are lost through vomiting or through gastric tube suction then a hypochloremia can result. Hypochloremia leads to an increase in SID. The positive charge increase associated with the SID must be balanced by an increase OH−. Treatment can be with normal saline administration. The treatment can be illustrated in the same fashion as free water changes. If we have a 1 L of water with 140 mEq/L of Na+ and a “hypochloremic” 95 mEq/L of Cl− then the SID is 45 mEq. If 1 L of normal saline is added, the beaker would then contain 147 mEq/L of Na+ and 125 mEq/L of Cl−, with the SID being 22 mEq/L. By shifting the SID, we have shifted the pH in the normal direction.

Hyperchloremia

Hyperchloremia results in an increase in H+. Hyperchloremia typically results from aggressive normal saline administration. Treatment of the elevated Cl− and decreased SID would be done by increasing the SID. This could be accomplished through sodium bicarbonate administration. Here, the Na+ is the effector agent and not the HCO3 −. The HCO3 − is a dependent variable and is rapidly excreted as CO2. Other ways of administering Na+ with a metabolizable anion are through the use of the sodium salts of lactate, gluconate, acetate, or citrate.

SID Unidentified Anions

SID can also be affected by the presence of organic acids such as lactate or ketoacids. Again, because these negatively charged molecules lower the SID, they result in an acidosis. Treatment is usually focused on stopping the development of acid. Resolution of the abnormal H+ can also be achieved by increasing the SID using NaHCO3.