Clinical Use of Henderson-Hasselbalch Equation

The H-H equation allows the pH, [HCO₃⁻], or Pco₂ to be computed if two of these three variables are known (shown as follows for Pco₂ and HCO₃⁻):

pH=6.1+log⁡[HCO3−]PaCO2×0.03

pH=6.1+log⁡[HCO3−]PaCO2×0.03

Blood gas analyzers measure pH and PCO₂ but compute [HCO₃⁻]. Assuming a normal arterial pH of 7.40 and a PaCO₂ of 40 mm Hg, arterial [HCO₃⁻] can be calculated as follows:

pH=6.1+log⁡([HCO3−]PCO2×0.03)

7.40=6.1+log⁡([HCO3−][40×0.03])

7.40=6.1+log⁡([HCO3−]1.2)

Solving for [HCO₃⁻]:

[HCO3−]=antilog (7.40−6.1)×1.2=antilog(1.3)×1.2=20×1.2=24 mEq/L

The H-H equation is useful for checking a clinical blood gas report to see if the pH, PCO₂, and [HCO₃⁻] values are compatible with one another. In this way, transcription errors and analyzer inaccuracies can be detected. It is also clinically useful to predict what effect changing one H-H equation component will have on the other components. For example, a clinician may want to know how low the arterial blood pH will fall for a given increase in PaCO₂.

Mini Clini

Applying the Henderson-Hasselbalch Equation in a Clinical Setting

 Problem

The respiratory therapist (RT) is caring for a mechanically ventilated patient. The patient has a tidal volume (V_T) of 800 ml and a breathing frequency of 10/min, yielding a minute ventilation () of 8 L/min. The patient’s PaCO₂ is 55 mm Hg, pH is 7.30, and bicarbonate is 26 mEq/L, and the therapist wishes to maintain a pH of 7.35. How much does the RT need to change the PaCO₂ to achieve this desired pH, and what change in the patient’s V_T does this require?

Solution

First, the therapist needs to calculate the PaCO₂ required to achieve a pH of 7.35 using the known values:

PaCO2=26 mEq/L0.03×antilog(7.35−6.1)

PaCO2=260.53

PaCO2=49 mm Hg

Next, the RT must calculate the required to produce a PaCO₂ of 49 mm Hg. Because is inversely proportional to PaCO₂, the following can be stated:

(V˙E)1×(PaCO2)1=(V˙E)2×(PaCO2)2

Where subscripts 1 and 2 represent current and future values. The RT then solves for ()₂ as follows:

(8L/min)×55 mm Hg=(V˙E)2×49 mm Hg

(8×55)49=(V˙E)2

8.98 L/min=(V˙E)2

Increasing the patient’s from 8 L/min to approximately 9 L/min yields a PaCO₂ of 49 mm Hg and a pH of approximately 7.35. Now the RT divides the new of 9 L/min by the respiratory frequency to calculate the new V_T required:

9 L/min/10=900 ml

A V_T of 900 mL at a rate of 10 breaths/min should produce an arterial pH of 7.35, according to the H-H equation.

Bicarbonate Buffer System

The bicarbonate buffer system is particularly effective in the body because it is an open system—that is, one of its components (CO₂) is continually removed through ventilation:

(Exhaled gas)←CO2+H2O←H2CO3←HCO3−+H+

In this way, HCO₃⁻ continues to buffer H⁺ as long as ventilation continues. Hypothetically, this buffering activity can continue until all body sources of HCO₃⁻ are used up in binding H⁺.

The bicarbonate buffer system can buffer only fixed acid. An increased fixed acid load in the body (e.g., lactic acid) reacts with HCO₃⁻ of the bicarbonate buffer system:

As shown, the process of buffering fixed acid produces CO₂, which is eliminated in exhaled gas. Large amounts of acid are buffered in this fashion. If the ability to ventilate is impaired, this type of buffering cannot occur.

The bicarbonate buffer system cannot buffer carbonic (volatile) acid, which accumulates in the blood whenever ventilation fails to eliminate CO₂ as fast as it is produced (hypoventilation). The resulting accumulation of CO₂ drives the hydration reaction in the direction that produces more carbonic acid, H⁺, and HCO₃⁻, as shown:

H⁺ produced by dissociating H₂CO₃ molecules cannot be buffered by the simultaneously produced HCO₃⁻ because hypoventilation prevents the reaction from reversing its direction. The closed nonbicarbonate buffer systems are the only buffers that can buffer carbonic acid.

Nonbicarbonate Buffer System

Table 13-1 lists the nonbicarbonate buffers in the blood. Of these, Hb is the most important because it is the most abundant. As mentioned, these buffers are the only ones available to buffer carbonic acid. However, they can buffer H⁺ produced by any acid, fixed or volatile. Because nonbicarbonate buffers (Buf⁻/HBuf) function in closed systems, the products of their buffering activity eventually accumulate, slowing or stopping further buffering activity:

H++Buf−↔HBuf

This slowing or stopping of buffering activity means that not all of the Buf⁻ is available for buffering activity. At equilibrium (denoted by the double arrow), Buf⁻ still exists in solution but cannot combine further with H⁺. In contrast, most of the HCO₃⁻ in the bicarbonate buffer system is available for buffering activity because it functions in an open system where equilibrium between reactants and products does not occur. Both open and closed systems function in a common fluid compartment (blood plasma) as illustrated in the following equation:

Most of the added fixed acid is buffered by HCO₃⁻ because ventilation continually pulls the reaction to the left. Smaller amounts of H⁺ react with Buf⁻ because equilibrium is approached, slowing the reaction.

Basic Kidney Function

To understand how the kidneys determine whether to excrete acidic or basic urine, some fundamental facts about renal function must be understood. The glomerulus is the component of the renal nephron responsible for filtering the blood. Hydrostatic blood pressure forces water, electrolytes, and other nonprotein substances through semipermeable glomerular capillary endothelium. The resulting filtrate is greatly modified in volume and composition as it flows through the nephron tubules. Excreted filtrate is called urine.

HCO₃⁻ is one of the electrolytes filtered from the blood at the glomerulus to become part of the tubular filtrate. In this way, base (HCO₃⁻) is removed from the blood. This loss of base is offset by the simultaneous secretion of the nephron tubular epithelium of H⁺ into the tubular lumen and into the filtrate. Under normal conditions, the rate of H⁺ secretion is almost the same as the rate of HCO₃⁻ filtration.⁴ In this way, the kidneys titrate H⁺ and HCO₃⁻ against each other to form CO₂ and H₂O.

H⁺ secretion begins with the diffusion of blood CO₂ into the tubule cell (Figure 13-2). Aided by the enzyme carbonic anhydrase, CO₂ reacts with H₂O to form carbonic acid, which forms HCO₃⁻ and H⁺. The tubule cell actively secretes H⁺ into the filtrate by means of countertransport, in which Na⁺ and H⁺ are simultaneously transported in opposite directions. That is, Na⁺ and H⁺ combine with opposite ends of a carrier protein in the luminal border of the tubule cell membrane. Sodium ions move into the cell down its high concentration gradient, providing the energy to secrete H⁺ into the tubular filtrate (see Figure 13-2).⁴

The rate of tubular H⁺ secretion increases if the concentration of H⁺ in the blood plasma increases. Conversely, the rate of H⁺ secretion decreases if blood plasma [H⁺] decreases (Figure 13-3). Any factor that increases PaCO₂, such as hypoventilation, increases H⁺ secretion, and any factor that decreases PaCO₂, such as hyperventilation, decreases H⁺ secretion.

HCO₃⁻ formed in the tubule cell from the reaction between CO₂ and H₂O (see Figure 13-2) diffuses back into the blood plasma because the luminal side of the tubule cell is relatively impermeable to HCO₃⁻. HCO₃⁻ and Na⁺ are reabsorbed whenever H⁺ is secreted into the tubular filtrate.

Reabsorption of Bicarbonate Ion

Because the luminal side of the renal tubule cell is relatively impermeable to HCO₃⁻, these ions are reabsorbed indirectly, as shown in Figure 13-2. The HCO₃⁻ in the filtrate reacts with the H⁺ secreted by the tubular cells. The resulting carbonic acid breaks down into CO₂ and H₂O. Because CO₂ is extremely diffusible through biologic membranes, it diffuses instantly into the tubule cell. There, CO₂ reacts rapidly with H₂O in the presence of carbonic anhydrase, rapidly forming HCO₃⁻ and H⁺. The HCO₃⁻ diffuses back through the nonluminal side of the tubule cell into the blood. The reabsorbed HCO₃⁻ ion is not the same HCO₃⁻ ion that existed in the tubular fluid. If the tubule cells secrete sufficient H⁺, all HCO₃⁻ in the tubular fluid is reabsorbed in this manner.

The net effect of secreting H⁺ (caused by high blood CO₂ or hypoventilation, as shown in Figure 13-2) is to reabsorb all filtrate HCO₃⁻, increasing the quantity of HCO₃⁻ in the blood. According to the H-H equation, this brings blood pH up toward the normal range.

If blood CO₂ is low, as is the case in a state of hyperventilation (see Figure 13-3), the ratio of HCO₃⁻ to dissolved CO₂ molecules increases, and the renal filtrate has more HCO₃⁻ than H⁺. Because HCO₃⁻ cannot be reabsorbed without first reacting with H⁺, the excess HCO₃⁻ is excreted in the urine, carrying positive ions such as Na⁺ or K⁺ in the filtrate. The net effect of secreting less H⁺ is to increase the quantity of HCO₃⁻ (base) lost in the urine. According to the H-H equation, this brings blood pH down toward the normal range. These renal responses to high and low blood PCO₂ are the mechanisms by which the kidneys compensate for respiratory acid-base disturbances.

Excess Hydrogen Ion Excretion and Role of Urinary Buffers

If no buffers existed in the filtrate to react with H⁺, the H⁺-secreting mechanism would soon cease to function because when the filtrate pH decreases to 4.5, H⁺ secretion stops.⁴ Buffers in the tubular fluid are essential for the secretion and elimination of excess H⁺ in acidotic states.

In Figure 13-2, more H⁺ than HCO₃⁻ is present in the filtrate. After all available HCO₃⁻ reacts with H⁺, the remaining H⁺ reacts with two other filtrate buffers, phosphate and ammonia, as illustrated in Figures 13-4 and 13-5. In Figure 13-4, phosphate and H⁺ react to form H₂PO₄⁻, which must be excreted with a positive ion to maintain tubular electroneutrality. Figure 13-5 shows that when urinary buffers are depleted, the resulting fall in filtrate pH stimulates the tubules to secrete ammonia. The NH₃ molecule buffers H⁺ by reacting with it to form the ammonium ion (NH₄⁺). To maintain electroneutrality, the kidney excretes a negatively charged ion to accompany NH₄⁺. This negative ion is chloride (Cl⁻), the most abundant filtrate anion.

When NH₄⁺ reacts with H⁺, HCO₃⁻ diffuses from the tubule cell into the blood (see Figure 13-5). The net effect of ammonia buffer activity is to cause more bicarbonate to be reabsorbed into the blood, counteracting the acidic state of the blood. Figure 13-5 shows that when Cl⁻ is excreted in combination with NH₄⁻, the blood gains HCO₃⁻. Blood [Cl⁻] and [HCO₃⁻] are reciprocally related (i.e., when one is high, the other is low). This relationship explains why people with chronically high blood PCO₂ tend to have low blood [Cl⁻] or hypochloremia. Activation of the ammonia buffer system enhances Cl⁻ loss and HCO₃⁻ gain.

Effect of Carbon Dioxide Hydration Reaction on [HCO₃⁻]

In the previous examples of pure (uncompensated) respiratory acidosis and alkalosis, it was assumed that [HCO₃⁻] did not change as the PaCO₂ level increased or decreased. Arterial [HCO₃⁻] does increase slightly as the PaCO₂ increases because the CO₂ hydration reaction generates HCO₃⁻. This reaction occurs primarily in the red blood cell because the catalytic enzyme, carbonic anhydrase, is present:

CO2+H2O−(carbonic anhydrase)→H2CO3→H++HCO3−

As H⁺ and HCO₃⁻ are rapidly produced, Hb immediately buffers H⁺, generating HCO₃⁻ in the process:

The amount of HCO₃⁻ increase depends on the amount of nonbicarbonate buffer that is available to accept the H⁺ produced by the hydration reaction. Generally, when the nonbicarbonate buffer concentration is normal, and the PCO₂ increase is acute, the hydration reaction increases the plasma [HCO₃⁻] approximately 1 mEq/L for every 10-mm Hg increase in PCO₂ higher than 40 mm Hg. Figure 13-6 illustrates this hydration reaction effect. Normal status is represented by point A: PaCO₂ of 40 mm Hg, pH of 7.40, and plasma HCO₃⁻ of 24 mEq/L. An acute increase in PaCO₂ from 40 mm Hg to 80 mm Hg proceeds from point A, moving to the left, up the normal blood buffer line (line BAC) to point D, where the buffer line intersects the PaCO₂ = 80 mm Hg isopleth. Point D indicates an HCO₃⁻ of approximately 28.5 mEq/L and a pH of approximately 7.18. This small change in [HCO₃⁻] should not be erroneously interpreted as early renal compensation.

Causes

Any process in which alveolar ventilation fails to eliminate CO₂ as rapidly as the body produces it causes respiratory acidosis. This acidosis could occur in different ways. A person’s ventilation may be decreased from a drug-induced central nervous system depression, or a person with limited ventilatory reserve may have a normal PaCO₂ at rest but cannot accommodate the increased CO₂ production associated with increased physical activity. Box 13-3 summarizes causes of respiratory acidosis.

If hypercapnia is uncompensated, respiratory acidosis occurs with decreased pH, increased PaCO₂, and normal or slightly increased [HCO₃⁻]. In this instance, the slightly increased [HCO₃⁻] is not a sign that the kidneys have started compensatory activity; it merely reflects the effect of CO₂ hydration reaction on [HCO₃⁻].

Compensation

Renal compensation for respiratory acidosis begins as soon as PaCO₂ increases. The kidney reabsorbs HCO₃⁻ from the renal tubular filtrate, bringing the arterial pH into the normal range (see Figure 13-2). However, this process cannot keep pace with an acutely increasing PaCO₂. Full compensation may take several days.

Partly compensated respiratory acidosis is characterized by increased PaCO₂, increased [HCO₃⁻], and an acid pH—still not quite up in the normal range. Fully compensated respiratory acidosis is characterized by a pH on the acidic side of the normal range (<7.40 but >7.35), increased PaCO₂, and increased [HCO₃⁻]. Increased [HCO₃⁻] in the presence of increased PaCO₂ is a sign that the PaCO₂ has been elevated for a considerable time (i.e., the kidneys have had sufficient time to compensate). The underlying pathologic process producing hypercapnia is still present; the kidneys simply mask the problem by maintaining a normal-range pH. Because of hypercapnia, the term acidosis is retained in classifying this condition (compensated respiratory acidosis). This term emphasizes that lung function is still abnormal, and, if unopposed by the renal compensatory mechanism, it would produce an acidosis.

Correction

The main goal in correcting respiratory acidosis is to improve alveolar ventilation. Various respiratory care modalities may be employed ranging from bronchial hygiene and lung expansion techniques to noninvasive positive pressure ventilation to endotracheal intubation and mechanical ventilation. If hypoventilation is chronic and compensation has restored pH within the normal range, corrective action aimed at decreasing PaCO₂ is inappropriate and possibly harmful. In this instance, rapidly decreasing PaCO₂ to normal would induce an alkalosis because of the compensatory [HCO₃⁻] retention by the kidneys (Table 13-6).

Causes

Any process in which ventilatory elimination of CO₂ exceeds production of CO₂ causes respiratory alkalosis. The most common cause of hyperventilation in patients with pulmonary disease is decreased PaO₂ (hypoxemia). Hypoxemia causes specialized neural structures to signal the brain, increasing ventilation (see Chapter 14). Anxiety, fever, stimulatory drugs, pain, and central nervous system injuries are possible causes of hyperventilation. Other possible causes include stimulation of irritant receptors in the lung parenchyma, which may occur in pneumonia or pulmonary edema.

Hyperventilation and respiratory alkalosis also may be iatrogenically induced (induced by medical treatment). Iatrogenic hyperventilation is most commonly associated with overly aggressive mechanical ventilation. It may also be associated with aggressive deep breathing and lung expansion respiratory care procedures. Decreased PaCO₂, increased pH, and normal-range [HCO₃⁻] characterize acute respiratory alkalosis. A slight decrease in [HCO₃⁻] is expected from the effect of the hydration reaction. Box 13-4 summarizes causes of respiratory alkalosis.

Clinical Signs

An early sign of respiratory alkalosis is paresthesia (numbness or a tingling sensation in the extremities). Severe hyperventilation is associated with hyperactive reflexes and possibly tetanic convulsions. The low PaCO₂ may constrict cerebral vessels sufficiently to impair cerebral circulation, causing light-headedness and dizziness.

Compensation

The kidneys compensate for respiratory alkalosis by excreting HCO₃⁻ in the urine (bicarbonate diuresis; see Figure 13-3). This activity brings arterial pH down into the normal range. As with respiratory acidosis, renal compensation is a slow process. Complete compensation may take days.

Partly compensated respiratory alkalosis is characterized by decreased PaCO₂, decreased [HCO₃⁻], and alkaline pH—still not quite down in the normal range. Fully compensated respiratory alkalosis is characterized by decreased PaCO₂, decreased [HCO₃⁻], and pH on the alkaline side of normal (pH > 7.40 but ≤ 7.45). Compensated respiratory alkalosis is sometimes called chronic respiratory alkalosis. The underlying hyperventilation and hypocapnia are still present. The term alkalosis is used in classifying this condition, although the pH is within the normal range.

Correction

Correcting respiratory alkalosis involves removing the stimulus causing the hyperventilation. If hypoxemia is the stimulus, oxygen (O₂) therapy is needed.

Alveolar Hyperventilation Superimposed on Compensated Respiratory Acidosis

Consider a patient with a compensated respiratory acidosis who has an arterial pH of 7.38, PaCO₂ of 58 mm Hg, and HCO₃⁻ of 33 mEq/L. If this patient becomes severely hypoxic, the hypoxia may stimulate increased alveolar ventilation if lung mechanics are not too severely deranged. This increased alveolar ventilation would acutely lower PaCO₂, possibly increasing the pH to the alkalotic side of normal. For example, the patient’s blood gas values might now be as follows: pH of 7.44, PaCO₂ of 50 mm Hg, and HCO₃⁻ of 33 mEq/L.

The novice might erroneously interpret these values as compensated metabolic alkalosis. This example shows that blood gas data alone are insufficient for rational acid-base assessment. Knowledge of the patient’s medical history and the nature of the current problem is essential to evaluate this problem accurately. The blood gas values in this example are described as acute hyperventilation (although PaCO₂ is >45 mm Hg) superimposed on compensated respiratory acidosis.

Causes

Metabolic acidosis can occur in one of the following two ways: (1) fixed (nonvolatile) acid accumulation in the blood or (2) an excessive loss of HCO₃⁻ from the body. An example of fixed acid accumulation is a state of low blood flow in which tissue hypoxia and anaerobic metabolism produce lactic acid. The resulting H⁺ accumulates and reacts with HCO₃⁻, reducing blood [HCO₃]. An example of bicarbonate loss is severe diarrhea, in which large stores of HCO₃⁻ are eliminated from the body, also producing a nonrespiratory acidosis.

Because these two kinds of metabolic acidosis are treated differently, it is important to identify the underlying cause. Analysis of the plasma electrolytes is helpful in distinguishing between these two types of metabolic acidosis. Specifically, measuring the anion gap is helpful in making this distinction.

Anion Gap

The law of electroneutrality states that the total number of positive charges must equal the total number of negative charges in the body fluids. Cations (positively charged ions) in the plasma produce a charge exactly balanced by plasma anions (negatively charged ions). Plasma electrolytes (cations and anions) routinely measured in clinical medicine are Na⁺, K⁺, Cl⁻, and HCO₃⁻. Normal plasma concentrations of these electrolytes are such that the cations (Na⁺ and K⁺) outnumber the anions (Cl⁻ and HCO₃⁻), leading to the so-called anion gap. Generally, K⁺ is ignored in calculating the anion gap:

Anion gap=[Na+]−([Cl−]+[HCO3−])

Figure 13-7, A shows that normal concentrations of these ions in the plasma are as follows: 140 mEq/L for Na⁺, 105 mEq/L for Cl⁻, and 24 mEq/L for HCO₃⁻, yielding an anion gap of 11 mEq/L (140 mEq/L − [105 mEq/L + 24 mEq/L] = 11 mEq/L). The normal anion gap range is 9 to 14 mEq/L.⁶

An increased anion gap (>14 mEq/L) is caused by metabolic acidosis in which fixed acids accumulate in the body. The H⁺ of these acids reacts with plasma HCO₃⁻, lowering its concentration; this leads to an increased anion gap (i.e., an increase in unmeasured anions) (see Figure 13-7, B). (When the H⁺ of fixed acids is buffered by HCO₃⁻, the anion portion of the fixed acid remains in the plasma, increasing unmeasured anion concentration.) A high anion gap indicates that fixed acid concentration in the body has increased.

Metabolic acidosis caused by HCO₃⁻ loss from the body does not cause an increased anion gap. Bicarbonate loss is accompanied by Cl⁻ gain, which keeps the anion gap within normal limits (see Figure 13-7, C). The law of electroneutrality helps explain the reciprocal nature of [HCO₃⁻] and [Cl⁻] in this instance. With a constant cation concentration, losing HCO₃⁻ means that another anion must be gained to maintain electroneutrality. In this case, the kidney increases its reabsorption of the most abundant anion in the tubular filtrate, the Cl⁻. The kind of metabolic acidosis in which HCO₃⁻ is lost from the body is sometimes called hyperchloremic acidosis because of the characteristic increase in plasma [Cl⁻]. Box 13-5 summarizes causes of anion gap and non–anion gap metabolic acidosis.

Compensation

Hyperventilation is the main compensatory mechanism for metabolic acidosis. The increased plasma [H⁺] of metabolic acidosis is buffered by plasma HCO₃⁻, reducing the plasma [HCO₃⁻] and the pH. A low pH activates sensitive receptors in the brain, signaling the respiratory muscles to increase ventilation. This increased ventilation lowers the blood’s volatile acid (H₂CO₃) and dissolved CO₂ levels, returning pH toward the normal range. Uncompensated metabolic acidosis suggests that a ventilatory defect is present. Metabolic acidosis accompanied by PaCO₂ of 40 mm Hg means that something prevents the lungs from responding appropriately to the brain’s stimulation. The defect may lie in nerve impulse transmission, the respiratory muscles, or the lungs themselves.

Symptoms

Respiratory compensation in metabolic acidosis may result in a great increase in minute ventilation, causing patients to complain of dyspnea. Hyperpnea (increased tidal volume depth) is a common finding during physical examination of patients with metabolic acidosis. In patients with severe diabetic ketoacidosis, a very deep, gasping type of breathing develops, called Kussmaul respiration. Neurologic symptoms of severe metabolic acidosis range from lethargy to coma.

Correction

The initial goal in severe acidemia is to increase the arterial pH greater than 7.20, a level below which serious cardiac arrhythmias are likely.⁷ If respiratory compensation maintains the pH at or above this level, immediate corrective action is usually not indicated. Treatment of the underlying cause of acid gain or base loss is the rational approach.

In cases of severe metabolic acidosis, intravenous infusion of NaHCO₃ may be indicated. If respiratory compensation is under way, only small amounts of NaHCO₃ are required to attain an arterial pH of 7.20. In any case, rapid correction of arterial pH greater than 7.20 by NaHCO₃ infusion is undesirable.

Causes

Metabolic alkalosis can occur in one of the following two ways: (1) loss of fixed acids or (2) gain of blood buffer base. Both processes increase plasma [HCO₃⁻]. To explain why losing fixed acid increases the plasma [HCO₃⁻], consider a situation in which vomiting removes gastric HCl from the body (Figure 13-8). In response to HCl loss, H⁺ diffuses out of the gastric cell into the gastric fluid, where Cl⁻ accompanies it; this forces the CO₂ hydration reaction in the gastric cell to the right, which generates HCO₃⁻. The HCO₃⁻ enters the blood in exchange for the Cl⁻. The plasma gains an HCO₃⁻ for each Cl⁻ (or H⁺) that is lost (see Figure 13-8).⁷

The causes of metabolic alkalosis are summarized in Box 13-6. Metabolic alkalosis is common in acutely ill patients and is probably the most complicated acid-base imbalance to treat because it involves fluid and electrolyte imbalances. Metabolic alkalosis is often iatrogenic, resulting from the use of diuretics, low-salt diets, and gastric drainage.

To understand how the loss of Cl⁻, K⁺, and fluid volume may cause alkalosis, one needs to understand how the kidney regulates Na⁺. Approximately 26,000 mEq of Na⁺ passes through the glomerular membrane daily, but the body’s daily Na⁺ intake averages only approximately 150 mEq.⁴ The kidney’s main job is to reabsorb Na⁺, not to excrete it. For this reason, and because Na⁺ has a major role in maintaining fluid balance, the kidney places a greater priority on reabsorbing Na⁺ than on maintaining Cl⁻, K⁺, or acid-base balance.

Normally, Na⁺ is reabsorbed through primary active transport (Figure 13-9), in which the sodium-potassium-adenosine triphosphatase (Na⁺,K⁺-ATPase) pump actively transports Na⁺ out of the renal tubule cell into the blood. This process causes Na⁺ to diffuse continually from the filtrate into the tubule cell. Cl⁻ must accompany Na⁺ to maintain electroneutrality in the filtrate. If blood Cl⁻ concentration is low (hypochloremia), less Cl⁻ is present in the filtrate, which means that the kidney relies more on other mechanisms to reabsorb Na⁺. These mechanisms, called secondary active secretion, require the kidney to secrete H⁺ or K⁺ into the filtrate in exchange for Na⁺. In this way, Na⁺ is reabsorbed, and filtrate electroneutrality is preserved. Figures 13-10 and 13-11 illustrate the secondary active secretion process for H⁺ and Na⁺, which may lead to depletion of blood H⁺ (alkalemia) and K⁺ (hypokalemia). Preexisting hypokalemia (e.g., from inadequate K⁺ intake) in the presence of hypochloremia places an even greater demand on the kidney to secrete H⁺ to reabsorb Na⁺; hypokalemia produces alkalosis. Dehydration (fluid volume depletion or hypovolemia) aggravates alkalosis and hypokalemia further because hypovolemia profoundly increases the kidney’s stimulus to reabsorb Na⁺.

Compensation

The expected compensatory response to metabolic alkalosis is hypoventilation (CO₂ retention). Traditionally, it was thought that the hypoxemia accompanying hypoventilation greatly limited respiratory compensation for metabolic alkalosis (i.e., hypoxemia stimulates ventilation and should prevent compensatory hypoventilation). However, more recent evidence does not support this theory.⁶ Metabolic alkalosis apparently blunts hypoxemic stimulation to ventilation. Individuals with PaO₂ levels of 50 mm Hg may still hypoventilate to PaCO₂ levels of 60 mm Hg to compensate for metabolic alkalosis. Nevertheless, significant CO₂ retention is not seen often in cases of metabolic alkalosis, probably because metabolic alkalosis commonly coexists with other conditions that may cause hyperventilation, such as anxiety, pain, infection, fever, or pulmonary edema.

Correction

Correction of metabolic alkalosis is aimed at restoring normal fluid volume and electrolyte concentrations, especially K⁺ and Cl⁻ levels. Inadequate fluid volume, especially if coupled with hypochloremia, causes excessive secretion and loss of H⁺ and K⁺ because of the great need to reabsorb Na⁺. In treating this type of alkalosis, it is important to supply adequate fluids containing Cl⁻. If hypokalemia is a primary factor, potassium chloride (KCl) is the preferred corrective agent. In cases of severe metabolic alkalosis, acidifying agents, such as dilute HCl or ammonium chloride may be infused directly into a large central vein.⁸

Base Excess

Base excess (BE) is determined by equilibrating a blood sample in the laboratory to a PCO₂ of 40 mm Hg (at 37° C) and recording the amount of acid or base needed to titrate 1 L of blood to a pH of 7.40. A normal BE is ±2 mEq/L. A “positive BE” (>+2 mEq/L) indicates a gain of base or loss of acid from nonrespiratory causes. A “negative BE” (<−2 mEq/L) indicates a loss of base or a gain of acid from nonrespiratory causes. The BE has the same limitation as the standard bicarbonate in that it is an in vitro, rather than in vivo, measurement. That is, in hypercapnia, the buffer base that diffused into the extravascular fluid in vivo cannot be recovered during laboratory in vitro titrations.

The reliance on BE to quantify metabolic acid-base abnormalities can be misleading. In cases of acute (uncompensated) respiratory acidosis, the BE commonly would be within the normal range, indicating correctly that the disturbance is purely respiratory in origin. However, when renal compensation has occurred to offset chronic hypercapnia, the BE measurement is elevated above the normal range because of the compensatory increase in plasma HCO₃⁻.

To illustrate, consider the Mini Clini in the left-hand column on p. 304 in which the patient has respiratory acidosis for which the body has compensated by renal retention of HCO₃⁻. If this patient’s blood were equilibrated in vitro to a PaCO₂ of 40 mm Hg, the HCO₃⁻ would decrease by only 2 to 3 mEq/L to 32 to 33 mEq/L, and the pH would increase to much greater than 7.45. This patient’s BE would be well above normal. The high BE may lead the clinician to conclude incorrectly that this patient has a primary metabolic alkalosis. However, in this instance, the high BE merely reflects the fact that renal compensation has occurred.