Acid-Base Disturbances

Published on 13/02/2015 by admin

Last modified 22/04/2025

Print this page

This article have been viewed 2184 times

Chapter 31 Acid-Base Disturbances

In this chapter, the principles of normal acid-base homeostasis are briefly reviewed and the common disturbances encountered in cardiac surgery patients are discussed. In a short summary such as this, many details are touched on only briefly or are oversimplified. In particular, the Stewart principle ¹ for analysis of acid-base disturbance is mentioned only in passing (Box 31-1).

BOX 31-1 The Stewart Physiochemical Approach to Acid-Base Analysis

The approach to acid-base analysis described by Stewart ¹ is based on the fundamental principles of physical chemistry and describes the mechanism of acid-base disturbances. These principles are as follows:

1. Electroneutrality: In an aqueous solution the sum of all the positive ions must equal the sum of all the negative ions.

2. The law of mass action: The dissociation equilibria of all the incompletely dissociated substances must be satisfied.

3. The law of conservation of mass: The amount of a substance remains the same unless it is added, removed, generated, or destroyed.

From these principles, one can derive the three independent variables influencing pH:

1. Strong ion difference (SID).

2. Pco₂ (Partial pressure of CO₂).

3. The total weak acid concentration (A_tot).

SID

SID is the net charge balance of all the strong ions present, where a strong ion is one that is effectively completely dissociated (as its equilibrium constant (pK) is far from the pH). It is usually calculated as (Na⁺ + K⁺ + Ca²⁺ + Mg²⁺) − (Cl⁻ + lactate⁻) and is also known as the apparent strong ion difference. In healthy humans this value is 40 to 42 mEq/l.

P_aco₂

P_aco₂ is an independent variable if the system is open, that is, if ventilation is occurring and the tissue Pco₂ is in equilibrium with alveolar gas.

A_tot

A_tot is the total plasma concentration of nonvolatile buffers, which are mostly proteins (mainly albumin) and phosphates.

The main implication of the Stewart approach is that the concentration of hydrogen ions or bicarbonate cannot change unless one of the three independent variables changes. Similarly, changes caused by adding or removing hydrogen ions or bicarbonate in a compartment are accommodated by changes in the dissociation of water unless one or more of the independent variables changes. Acid-base disorders can therefore be classified in relation to the derangements of the independent variables. Metabolic acidoses are caused by a reduced SID or by an increased A_tot, whereas metabolic alkaloses are caused by an increased SID or a reduced A_tot. However, because no mechanisms appear to control A_tot, changes in A_tot are not considered to be acid-base disorders by some authors.¹²

ACID-BASE PHYSIOLOGY

Hydrogen Ions and pH

Small changes in the concentration of hydrogen ions can affect the ionic charge carried by enzymes, which can have a dramatic effect on their conformation and physiologic activity. Marked changes in hydrogen ion concentration can permanently denature enzymes and cause cell death. Thus, precise control of hydrogen ion concentration is essential.

By convention, pH (the negative log₁₀ of the hydrogen ion concentration), rather than the absolute concentration of hydrogen ions, is reported. The normal value for plasma pH is 7.4. An arterial pH below 7.35 constitutes an acidemia and a pH above 7.45 constitutes an alkalemia. Acidosis refers to a process whereby excess hydrogen ions are generated but, because of the body’s pH defense mechanisms, pH may not fall below 7.35. Alkalosis refers to a process whereby hydrogen ions are consumed but it does not necessarily result in a plasma arterial pH above 7.45.

Normal Regulation of pH

The body generates more than 200 liters of carbon dioxide and 80 mmol of nonvolatile acid each day as the result of oxidative metabolism. Despite this, the hydrogen ion concentration of extracellular fluid is kept relatively constant at approximately 0.000040 mmol/l. Thus, the pH of extracellular fluid is under constant threat from this large acid load. pH stability is achieved by three methods: (1) chemical buffers; (2) the lungs, which regulate the partial pressure of carbon dioxide; (3) the kidneys, which regulate the plasma bicarbonate concentration.

Buffers

A buffer is any compound that can reversibly bind with hydrogen ions and thereby attenuate the change in pH of a solution that arises from the addition of a hydrogen ion load.

All human buffers are weak acids. Weak acids are compounds that have a modest tendency to dissociate and release hydrogen ions (H⁺):

(31-1)

where HA is the weak acid and A⁻ is the conjugate base of that acid. The buffering solution contains both HA and A⁻ in equilibrium. If hydrogen ions are added to the solution, the equilibrium of this equation shifts to the left and some, but not all, of the hydrogen ions are removed from solution as HA. Thus, the fall in pH is attenuated. Conversely, if hydrogen ions are removed from solution, the equilibrium shifts to the right, and the rise in pH is minimized.

The ability of a buffer to attenuate a change in pH is determined by the tendency of the buffer to dissociate, which is characterized by the equilibrium constant (Ka) of the buffer:

(31-2)

Buffering ability is maximal when the concentration of the weak acid and its conjugate base are equal; that is, when [HA] = [A⁻], which occurs when the hydrogen ion concentration of the solution is equal to the dissociation constant (Ka) of the buffer; put another way, when the pH of the solution is equal to the pKa of the buffer (where pKa = the −log₁₀Ka). To function effectively, a buffer must have a pKa of within 1.0 unit of the pH of the solution being buffered.

Buffer systems may be open or closed. A closed buffer system is simply a fixed amount of buffer in solution. As hydrogen ions are added to the buffer system, [A⁻] falls, [HA] rises, and the ability of the buffer system to attenuate any further pH change diminishes. With an open buffer system, [HA] and [A⁻] can be independently controlled; for instance, HA can be eliminated from the solution and A⁻ regenerated. Thus, open buffer systems have a far greater capacity to deal with acid or alkali loads than do closed buffer systems.

Thus, the ability of a buffer to attenuate the change in pH of a solution in response to the addition or removal of hydrogen ions depends on: (1) the concentration of the buffer; (2) the starting pH of the solution; (3) the pKa of the buffer; and (4) whether the buffer system is open or closed.

Buffering Systems Within the Body.

There are four important buffers of extracellular fluid: (1) bicarbonate; (2) hemoglobin; (3) plasma proteins; (4) phosphate.

The bicarbonate buffer system is as follows:

(31-3)

Despite not having the same form as Equation 33-1, the bicarbonate buffer system reacts to changes in pH in a similar manner: the addition of H⁺ shifts the equation to the left, generating CO₂ and H₂O; the removal of H⁺ shifts the equation to the right, generating HCO₃⁻ (and H⁺). The pKa of the bicarbonate buffer system is 6.1, which at first glance seems too far removed from the pH of plasma to be very effective. However, because the concentration of bicarbonate in extracellular fluid is high (24 mmol/l) and because the bicarbonate buffer system is open, with the ability to independently control carbon dioxide and bicarbonate concentrations, it is quantitatively the most important buffer of the extracellular fluid.

The histidine amino acid found on hemoglobin and plasma proteins (pK_a 6.5) constitutes the majority of the nonbicarbonate buffering in the extracellular fluid. Although hemoglobin is an intracellular protein, the permeability of the red cell membrane is such that it can be considered an extracellular buffer.

Dihydrogen phosphate (pKa 6.8) is a relatively unimportant extracellular buffer because its concentration in plasma is low, approximately 1 mmol/l. However, dihydrogen phosphate plays an important role in buffering urine within the renal tubules, where it is concentrated. It is also an important buffer within cells.

Because all extracellular buffers are in solution with the same hydrogen ions, in effect, all buffers are in equilibrium with each other. This is known as the isohydric principle. Only the bicarbonate buffer system is open, so all of the extracellular buffers are regenerated by the elimination of carbon dioxide and the regeneration of bicarbonate. Furthermore, the pH of the entire system can be calculated from the Henderson-Hasselbalch equation as it applies to the bicarbonate buffer system:

(31-4)

where 6.1 is the pKa of the bicarbonate buffer; HCO₃⁻ is in mmol/l; and Pco₂ is in mmHg. The term 0.03 is the solubility coefficient of CO₂ in plasma and converts Pco₂ in mmHg into mmol/l of dissolved CO₂ or H₂CO₃.

In addition to extracellular buffers (and hemoglobin), some buffering occurs intracellularly. This is achieved by swapping hydrogen ions for potassium ions across cell membranes (to maintain electroneutrality). Intracellular buffering is quantitatively important in some forms of metabolic acidosis such as acute renal failure, and it contributes to the hyperkalemia seen in this condition.

Role of the Lungs: Control of P_aco₂

The carbon dioxide generated by oxidative metabolism is transported from the tissues to the lungs predominantly as bicarbonate (70%) or bound to hemoglobin (23%), with only a small proportion (7%) as dissolved gas. In the lungs, bicarbonate is reconverted to carbon dioxide and excreted in expired breath. Thus, the process of transporting carbon dioxide from the tissues to the lungs does not consume or generate bicarbonate buffer. However, the partial pressure of carbon dioxide in blood does influence pH.

The partial pressure of dissolved carbon dioxide in arterial blood (P_aco₂) is inversely proportional to alveolar ventilation and directly proportional to the rate of production of carbon dioxide (see Eq. 1-17). The carbon dioxide partial pressure in arterial blood is tightly controlled to a value close to 5.2 κPa (40 mmHg), corresponding to a concentration of dissolved gas of 1.2 mmol/l. A reduction in alveolar ventilation leads to an increase in P_aco₂, which drives the equilibrium of Equation 31-3 to the right, causing a fall in pH. This condition is termed respiratory acidosis. Conversely, an increase in alveolar ventilation drives the P_aco₂ down and the equilibrium of Equation 31-3 to the left, causing an increase in pH. This condition is termed respiratory alkalosis.

Role of the Kidneys: Control of Plasma HCO₃⁻ Concentration

The role of the kidneys in acid-base homeostasis is twofold: first, reabsorbing bicarbonate that is filtered at the glomerulus, and second, replacing bicarbonate that is consumed buffering metabolic acid.

On a normal diet, about 80 mmol/day of nonvolatile (non-carbon-dioxide or metabolic) acid is generated each day, mainly by the metabolism of proteins. This acid is buffered in the plasma by the bicarbonate buffer system and excreted as carbon dioxide gas, resulting in a net loss of bicarbonate. The cells of the distal convoluted tubule of the kidney regenerate this lost bicarbonate. In the process of regenerating bicarbonate, the distal tubular cells secrete an equimolar amount of hydrogen ions into the tubular fluid, resulting in acidification of the urine. With normal renal function, approximately 400 to 500 mmol of bicarbonate can be regenerated in this way (or, put another way, the kidneys can clear 400 to 500 mmol of H⁺ per day, providing that urinary buffering is ideal),^2,³ resulting in a urinary pH as low as 4.5. In this way, the kidneys regulate the plasma bicarbonate concentration (normal range 24 to 28 mmol/l) in the face of a changing metabolic acid load. The larger the metabolic acid load, the lower the urinary pH. This renal mechanism takes some hours to days to be fully effective.

From Equation 31-3 it is seen that an excess of hydrogen ions leads to a fall in plasma bicarbonate concentration, a condition known as metabolic acidosis. Conversely, the loss of hydrogen ions results in an increase in plasma bicarbonate concentration, a condition known as a metabolic alkalosis.

Assessment of Acid-Base Status

Three parameters are used to assess acid-base status: pH, P_aco₂, and HCO₃⁻. Both pH and P_aco₂ are measured directly by the blood-gas analyzer, whereas bicarbonate concentration is calculated (automatically) from the Henderson-Hasselbalch equation. Because bicarbonate concentration is influenced by changes in the partial pressure of carbon dioxide (due to shifts in the equilibrium shown in Equation 31-3), blood gas analyzers also report the standard bicarbonate or the base excess. The standard bicarbonate is what the actual bicarbonate concentration would be if the blood gas sample were equilibrated with carbon dioxide at 5.2 κPa in fully oxygenated blood at 37°C. The base excess is the amount (in mmol/l) of acid that would be required to titrate the blood sample to a pH of 7.4 at a P_aco₂ of 5.2 κPa at 37°C. A negative base excess indicates that there is excess acid in the blood; that is, a metabolic acidosis exists. Abnormalities of the standard bicarbonate and base excess provide a measure of the metabolic component of an acid-base disorder.

Primary, Compensated, and Mixed Acid-Base Disturbances

As already noted, there are four primary acid-base disturbances. Some disturbances may be compensatory and minimize the change in pH associated with the primary disturbances. Thus, a metabolic acidosis promotes a compensatory respiratory alkalosis (hyperventilation), whereas a respiratory acidosis promotes a compensatory metabolic alkalosis (renal bicarbonate production and retention). Respiratory compensation for a metabolic disturbance occurs over minutes to hours, whereas renal compensation for a respiratory process takes place over days. Compensation may return the pH close to normal but, with the possible exception of metabolic compensation for a respiratory alkalosis, the pH will not be restored to 7.40.

In addition to primary and compensated disturbances, two or more independent acid-base disturbances may occur simultaneously (e.g., respiratory acidosis due to respiratory failure and metabolic alkalosis due to potassium loss). These are termed mixed disturbances. To distinguish between compensated and mixed acid-base disturbances requires knowledge of the clinical context and of the expected changes in P_aco₂ and actual bicarbonate associated with compensation (Table 31-1). See stepwise evaluation of acid-base disturbances.

Table 31-1 Expected Responses to Primary Acid-Base Disorders

Metabolic acidosis	P_aco₂ should equal 1.5 · bicarbonate concentration + 8.
Metabolic alkalosis	P_aco₂ should increase by 0.5 to 1.0 mmHg for each mmol/l gain in plasma bicarbonate above 25.
Acute respiratory acidosis	For every 10 mmHg increase in P_aco₂, actual bicarbonate should increase by 1 mmol/l.
Chronic respiratory acidosis	For every 10 mmHg increase in P_aco₂, actual bicarbonate should increase by 4 mmol/l, up to a maximum of 12 mmol/l.
Acute respiratory alkalosis	For every 10 mm Hg decrease in P_aco₂, actual bicarbonate should decrease by 2 mmol/l.
Chronic respiratory alkalosis	For every 10 mmHg decrease in P_aco₂, actual bicarbonate should decrease by 5 mmol/l.

Note: P_aco₂ is indicated in mmHg. See text for details. The changes in actual bicarbonate that occur with acute respiratory acidosis and alkalosis are minor and result from shifts in the equilibrium demonstrated by Equation 31-3.

Consequences of Acidosis and Alkalosis

Acidemia

Respiratory acidosis causing acidemia is well tolerated down to a pH of about 7.1. Below a pH of 7.1, problems of impaired mentation, arrhythmias, and reduced myocardial contractility ⁴ may develop. Reduced contractility may be explained in part by reduced cellular sensitivity to catecholamines.⁵ Patients with elevated pulmonary vascular resistance and right ventricular dysfunction may develop acute right ventricular failure because modest respiratory acidosis increases pulmonary vascular resistance.⁶ Acidosis is usually a consequence of severe respiratory, central nervous system, metabolic, or cardio-vascular abnormalities; it is these problems that demand urgent attention, rather than the acidosis per se.

Alkalemia

Severe alkalemia (pH >7.6) causes cardiac arrhythmias and neurologic dysfunction. The latter is manifested by weakness and lethargy or, less commonly, by confusion, agitation, and seizures. Alkalemia, especially alkalemia due to respiratory alkalosis, promotes the binding of calcium to albumin and results in ionized hypocalcemia which, if severe, can cause hyperreflexia, tetany, and seizures. Alkalemia also shifts the oxygen-hemoglobin dissociation curve to the left, reducing oxygen unloading in the tissues. In addition, severe metabolic alkalosis can cause compensatory hypoventilation, which can contribute to hypoxemia.

Hypokalemia and hypovolemia commonly accompany metabolic alkalosis (see subsequent material) and generally are of greater concern than the alkalosis itself. Respiratory alkalosis may be indicative of severe hypoxemic respiratory failure (see subsequent material).

ACID-BASE DISTURBANCES IN THE CARDIOTHORACIC INTENSIVE CARE UNIT

Respiratory Acidosis

Equation 1-17 shows that P_aco₂ depends on both alveolar ventilation and carbon dioxide production. Thus, respiratory acidosis may result from increased carbon dioxide production or inadequate alveolar ventilation. Increased carbon dioxide production arises as a result of increased metabolic rate caused by fever, sepsis, hyperthyroidism, and overfeeding—particularly when calories derive predominantly from carbohydrates. Excess carbon dioxide production rarely leads to respiratory acidosis alone because alveolar ventilation can increase, but it can exacerbate respiratory acidosis resulting from other causes.

Inadequate alveolar ventilation can arise from hypoventilation due to oversedation, residual neuromuscular blockade, or neurologic dysfunction. In a patient receiving mechanical ventilation, hypoventilation may occur because of inappropriate ventilator settings. One particular situation in which respiratory acidosis can develop unexpectedly is when lung compliance suddenly decreases (e.g., due to mucus plugging) in a patient receiving pressure-control ventilation (see Chapter 29).

Alveolar hypoventilation also arises when dead-space ventilation increases and minute ventilation fails to increase proportionately (see Eq. 1-16). This occurs when there is a ventilation-perfusion (V/Q) mismatch (see Chapter 1); lung units with a high V/Q ratio represent alveolar dead space. V/Q mismatch occurs with most acute and chronic lung diseases and imposes the requirement for increased minute ventilation to achieve a normal P_aco₂. However, many patients with acute respiratory compromise actually have a low P_aco₂ (i.e., respiratory alkalosis). This is because the minute ventilation increases out of proportion to the increase in alveolar dead space (see subsequent discussion under Respiratory Alkalosis). The presence of respiratory acidosis in the context of acute hypoxemic respiratory failure implies either reduced ventilatory drive (e.g., because of reduced level of consciousness) or significantly increased work of breathing (e.g., due to poor lung compliance), and it may indicate impending respiratory arrest. In contrast to the unstable nature of acute hypercarbic respiratory failure, chronic respiratory acidosis due to lung disease (e.g., chronic bronchitis) may remain stable for years.

Acute respiratory acidosis results in marked acidemia because renal compensation takes several days to develop fully.

Treatment of Respiratory Acidosis

When considering treatment for respiratory acidosis, the following issues must be considered:

• Is the condition acute or chronic? In most circumstances, the important parameter is the pH, not the P_aco₂. Thus, a patient with compensated respiratory acidosis who is receiving mechanical ventilation should be ventilated to a normal pH, not a normal P_aco₂. Targeting the latter will render the patient alkalemic.

• Is the respiratory condition stable? The presence of a rising P_aco₂ in a spontaneously breathing patient may represent impending respiratory arrest, requiring acute intervention (see Chapter 27). Conversely, a patient with a chronic respiratory acidosis may be quite stable.

• Will correcting respiratory acidosis worsen acute lung injury? With acute lung injury, lung-protective ventilation with low tidal volumes (see Chapter 29) improves outcome but may result in respiratory acidosis. Attempts to normalize pH at the cost of high airway pressure or large tidal volumes are harmful. Even when pulmonary hypertension accompanies acute lung injury, modest hypercarbia is usually preferable to aggressive invasive ventilation.

Treatment of acute respiratory acidosis involves management of the underlying cause and ventilatory strategies that result in improved alveolar ventilation.

Respiratory Alkalosis

Respiratory alkalosis is caused by increased alveolar ventilation. In the intensive care unit (ICU), respiratory alkalosis may arise for the following reasons:

• Iatrogenic overventilation in patients receiving mechanical ventilation.

• As a response to hypoxic respiratory failure in spontaneously breathing patients. Patients with hypoxic respiratory failure hyperventilate for a number of reasons. First, severe hypoxemia (P_ao₂ <8 kPa) is a potent ventilatory stimulus. Disorders such as anemia and shock that result in tissue hypoxemia (but normal P_ao₂) also cause hyperventilation and respiratory alkalosis. Second, acute lung disease (e.g., pulmonary edema, pneumonia) can stimulate lung-irritating receptors and cause hyperventilation. Finally, anxiety, pain, and fear exacerbate hyperventilation. Thus, even though dead-space ventilation is increased by pulmonary pathology, this reduction is offset by increased minute ventilation.

• Direct stimulation of the medullary respiratory center occurs with fever, sepsis, and liver failure.

Acute respiratory alkalosis results in a significant disturbance in pH. For instance, an increase in alveolar ventilation to twice baseline raises the pH by 0.23. Treatment is directed at the underlying condition. Respiratory alkalosis that develops during weaning from mechanical ventilation is discussed in Chapter 29.

Metabolic Acidosis

Metabolic acidosis results from bicarbonate loss or from gain of hydrogen ions. These two mechanisms may be distinguished by calculating the anion gap.

Anion Gap

The anion gap is the difference between the measured cations and anions in plasma:

(31-5)

The normal range is 10 to 14 mmol/l. Sometimes potassium is included in the calculation, in which case the normal range is 14 to 18 mmol/l. Most unmeasured anions arise from negative charges carried on plasma proteins, particularly albumin. Thus, a low plasma concentration of albumin reduces the anion gap. For every 10 g/l decrease in plasma albumin concentration, the anion gap is reduced by 4 mmol/l.

Raised anion gap acidoses are caused by the presence of additional unmeasured anions. The two common causes of a raised anion gap acidosis in the cardiothoracic ICU are lactic acidosis and renal failure. Rarely, an anion gap acidosis occurs due to diabetic ketoacidosis or rhabdomyolysis.

Normal anion gap acidosis results from either bicarbonate dilution (secondary to fluid resuscitation with 0.9% sodium chloride) or abnormal bicarbonate loss from the gut or the kidneys (renal tubular acidosis). With a normal anion gap metabolic acidosis, the reduction in bicarbonate is mirrored by an increase in chloride concentration. For this reason, a normal anion gap acidosis is sometimes called a hyperchloremic acidosis. Urinary pH should be low (usually less than 5) with a metabolic acidosis due to increased renal secretion of hydrogen ions. A urinary pH above 5.3 is suggestive of impaired urinary acidification, and usually indicates renal tubular acidosis. The causes of metabolic acidosis in the cardiothoracic ICU are summarized in Table 31-2.

Table 31-2 Causes of Metabolic Acidosis

Increased anion gap	Ketoacidosis
	Renal failure
	Lactic acidosis (see Table 31-3)
	Citrate
	Methanol, salicylate, other toxins
Normal anion gap	Gastrointestinal bicarbonate loss
	Renal tubular acidosis
	Excessive chloride administration (large-volume resuscitation with 0.9% sodium chloride)
	Postchronic hypocarbia

Lactic Acidosis

In the presence of oxygen, glucose is metabolized into water and carbon dioxide within mitochondria. In the absence of oxygen, incomplete oxidation of glucose occurs, resulting in the production of lactate and hydrogen ions, which is known as lactic acidosis.⁷ (The incomplete oxidation of fatty acids results in the production of ketoacids.) Lactate production is a normal process that occurs in many tissues in the body, such as red blood cells, skin, brain, and muscle. In the liver, this lactate is metabolized to carbon dioxide and water or is used to produce glucose. Both processes consume hydrogen ions, and both generate bicarbonate (see Eq. 31-3). Thus, the lactate anion can be considered a precursor of bicarbonate. Lactate is normally present in low levels (<2 mmol/l) in plasma.

High plasma levels of lactate are pathologic and may cause acidemia. The causes of lactic acidosis (plasma lactate >5 mmol/l) are categorized as follows:

• Type A, caused by inadequate delivery of oxygen to tissues to meet metabolic demand. This is seen in cardiac arrest, circulatory shock, hypoxemia, profound anemia, and seizures.

• Type B, caused by increased lactate production or reduced lactate clearance. Increased lactate production occurs in the absence of tissue hypoxia in various disease states (particularly sepsis) and in association with a number of drugs (e.g., metformin). Reduced lactate clearance occurs with severe liver disease.

In the cardiothoracic ICU, lactic acidosis commonly occurs in patients with shock; a plasma lactate concentration of more than 5 mmol/l is a marker of increased likelihood of mortality.⁸

Lactic acidosis may also arise for other reasons, including:

• Infusion of β₂-agonists (see Chapter 3). Resolution of lactic acidosis occurs rapidly with discontinuation of drug treatment.

• Treatment with the biguanide antidiabetic agents metformin and phenformin.

• Limb ischemia (e.g., due to an arterial embolus or the effects of an intraaortic balloon pump).

• Hepatic ischemia or failure (see Chapter 34).

• Mesenteric ischemia or infarction (see Chapter 34).

Mesenteric infarction, hepatic ischemia, limb ischemia, and treatment with β₂-agonists commonly coexist with shock. The causes of lactic acidosis in the cardiothoracic ICU are summarized in Table 31-3. In patients with renal failure receiving continuous hemofiltration, moderate hyperlactatemia (5 to 10 mmol/l) in the absence of acidemia is a normal finding when sodium lactate replacement fluid is used.

Table 31-3 Mechanisms of Lactic Acidosis

Type A: Hypoxic
Global ischemia	Shock
	Severe anemia
	Severe hypoxemia
Focal ischemia	Mesenteric infarction
	Limb ischemia
Type B: Nonhypoxic
Delayed clearance of lactate	Hepatic or renal failure
Pyruvate dehydrogenase dysfunction	Sepsis, high catecholamine levels, thiamine deficiency
Uncoupling of oxidative phosphorylation	Protein breakdown, severe catabolism
Accelerated glycolysis	Sepsis, seizures, malignancies, parenteral nutrition
Lactate administration	Lactate-containing renal replacement fluid

Treatment of lactic acidosis should be directed to the underlying cause. If a type A lactic acidosis is present, the cause of shock must be investigated and treated urgently. Therapy should be directed to increasing oxygen delivery and reducing oxygen demand. For unremitting lactic acidemia (pH <7.1) despite appropriate cardiovascular support, even in the absence of renal failure, renal replacement therapy with bicarbonate-buffered replacement solution may be indicated.⁹ Treatment with intravenous sodium bicarbonate is generally not indicated because it generates excess carbon dioxide, causes hyperosmolarity, worsens intracellular acidosis, may cause rebound metabolic alkalosis, and does not improve outcome.^10,¹¹ However, with severe lactic acidosis (pH <7.1), 50 to 100 mmol of IV sodium bicarbonate may temporarily arrest the fall of pH while other treatments are commenced. Other buffers, such as Tham (tris-hydroxymethyl aminomethane) or carbicarb, generate less carbon dioxide than sodium bicarbonate but have not been subjected to randomized clinical trials. In patients who are mechanically ventilated, increasing minute ventilation attenuates the fall in pH. However, this should not be undertaken at the cost of high airway pressures or large tidal volumes.

Renal Dysfunction

Acute and chronic renal failure leads to an elevated anion gap metabolic acidosis due to the accumulation of sulfates, phosphates, urate, and other organic non-volatile acids. With chronic renal failure, acidosis develops when the glomerular filtration rate falls to about 25% of normal. With acute renal failure associated with tubular necrosis, acidosis is an early and consistent finding. Marked acidification of the urine is not possible, and metabolic compensation for respiratory acidosis (by de novo bicarbonate generation in the distal convoluted tubule) may be poor or nonexistent.

Normal Anion Gap Acidosis

The use of 0.9% sodium chloride as volume replacement can cause a hyperchloremic metabolic acidosis with a normal anion gap. Sodium chloride has a weak tendency to react with water-forming hydrochloric acid. Thus, the use of 0.9% sodium chloride to replace fluid losses, which effectively replaces bicarbonate with chloride, results in acidosis. Using the approach of Stewart (see Box 31-1), the strong ion difference is reduced, and the pH falls. If renal function is normal, this acidosis rapidly resolves as the kidneys regenerate bicarbonate and excrete excess chloride.

Normal anion gap acidosis also results from abnormal bicarbonate losses from the gut or the kidney. Bicarbonate loss from the gut occurs with severe diarrhea. When abnormal renal bicarbonate loss occurs, it is known as renal tubular acidosis. There are a wide variety of causes of renal tubular acidosis, including autoimmune disease, Addison disease, diabetes, and various drugs. Rarely, renal tubular acidosis follows acute tubular necrosis due to hypovolemic shock.

The cause of nonanion gap metabolic acidosis may be assessed by calculation of the urine anion gap:

(31-6)

The normal response to metabolic acidosis is to excrete more hydrogen ions into the urine, primarily as ammonium (NH₄⁺). Thus, with metabolic acidosis, the urinary anion gap should be negative because of the presence of high levels of ammonium, an unmeasured cation. A positive urinary anion gap indicates impaired urinary acidification such as that which occurs with renal tubular acidosis. Urinary pH is abnormally high (>5.3) with renal tubular acidosis.

In contrast to lactic acidosis, treatment with sodium bicarbonate is indicated for normal anion gap metabolic acidosis.

Metabolic Alkalosis

Metabolic alkalosis is generated primarily by an abnormal loss of hydrogen ions or by the administration of bicarbonate. Because the renal mechanisms that excrete excess bicarbonate are well developed, not only must there be a stimulus to generate a metabolic alkalosis, but there must also be a mechanism by which the alkalosis is maintained.

Generation of Metabolic Alkalosis

In the cardiothoracic ICU, metabolic alkalosis most commonly arises for the following reasons:

• Diuretic therapy. Loop and thiazide diuretics result in renal loss of both sodium and chloride. The loss of extracellular fluid, without concomitant loss of bicarbonate, leads to a “contraction alkalosis.”

• Gastric fluid losses. Gastric secretion of H⁺ causes an equimolar amount of HCO₃⁻ to be secreted into the plasma. Thus, excessive nasogastric fluid losses result in metabolic alkalosis.

• Administration of sodium bicarbonate or a bicarbonate precursor. Abnormal bicarbonate gain, either as sodium bicarbonate or as a bicarbonate precursor such as citrate (from blood products or renal replacement therapy anticoagulation), can cause metabolic alkalosis.

• Posthypercarbic metabolic alkalosis. Metabolic compensation for respiratory acidosis leads to increased plasma HCO₃⁻. Then, if minute ventilation increases (e.g., due to institution of invasive ventilation or reversal of sedation), resulting in a normal P_aco₂, metabolic alkalosis occurs. This underlines the need to target mechanical ventilation pH with, not P_aco₂.

The common causes of metabolic alkalosis are summarized in Table 31-4. Other causes, such as Conn or Cushing syndrome, occur rarely in a cardiac ICU.

Table 31-4 Causes of Metabolic Alkalosis

Alkali intake	Bicarbonate
	Citrate (massive blood transfusion, anticoagulation for CVVH)
	Carbonates (antacids)
	Acetate (TPN, intermittent hemodialysis)
Gastrointestinal acid loss	Vomiting
	Nasogastric tube losses with gastroparesis
	Chloride-losing diarrhea
Renal acid loss	Diuretics
	Posthypercarbic alkalosis
	Hypovolemia
	Nonreabsorbable anions (e.g., penicillin)

CVVH, Continuous venovenous hemofiltration; TPN, total parenteral nutrition.

Maintenance of Metabolic Alkalosis

The kidney has a great capacity to excrete excess bicarbonate, so maintenance of metabolic alkalosis requires excess renal bicarbonate retention. This occurs primarily in response to hypovolemia. Hypovolemia leads to activation of the renin-angiotensin-aldosterone system and a reduction in renal blood flow. In the distal nephron, aldosterone promotes sodium reabsorptions and excretion of either potassium or hydrogen ion (see Chapter 1). In hypokalemia, aldosterone augments the secretion of hydrogen ions (and therefore the production of plasma bicarbonate); in alkalosis, aldosterone promotes the loss of potassium. In the proximal nephron, hypovolemia results in enhanced sodium reabsorption via a sodium/hydrogen ion exchange mechanism. Thus, hypovolemia leads to abnormal hydrogen ion secretion (and therefore to bicarbonate reabsorption). To summarize: metabolic alkalosis is maintained by hypovolemia, particularly in the presence of hypokalemia. These mechanisms underlie an important physiologic principle: if necessary, circulating volume is defended in preference to acid-base and potassium homeostasis.