Sodium bicarbonate

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CHAPTER 95

Sodium bicarbonate

Anna E. Bartunek, MD and Wolfgang Schramm, MD

Sodium bicarbonate (NaHCO₃) is an inorganic salt that readily dissociates into Na⁺ and HCO₃⁻ to bind acids and strong bases; when reacting with acids, the sodium salt of the acid, H₂O, and CO₂ are the byproducts. NaHCO₃ has been used for centuries in foods, animal feeds, and industrial processes. In vertebrate animals, the HCO₃⁻ ion is the principal buffer in extracellular and interstitial fluid. Because of its ability to neutralize acid, NaHCO₃ has been frequently ingested by mouth as an antacid and administered intravenously to treat metabolic acidosis. Because of its effect on blood pH, NaHCO₃ is used to treat a variety of drug overdoses (chlorpropamide, phenobarbital, cocaine, and class Ia and Ic antiarrhythmic agents), to treat metabolic acidosis induced by ingestion of methanol (and to enhance formate elimination by the kidneys) and ethylene glycol poisoning, and to alkalinize urine (e.g., in patients with rhabdomyolysis). Although anesthesia providers commonly use NaHCO₃ to treat metabolic acidosis, the indications for doing so are not universally accepted.

Acid-base balance

The precise functions of enzymes and proteins within cells are dependent on pH; therefore, a number of mechanisms exist to maintain hydrogen ion concentration ([H⁺]) within a very narrow range in the face of CO₂, H⁺, and OH⁻ production during normal metabolism and despite physiologic and pathologic challenges. Intracellular and extracellular chemical buffering, transcellular diffusion or transport of electrically charged ions, and excretion or retention of acid or alkali by the kidneys and of CO₂ by the lungs maintain acid-base homeostasis. The most important buffer in the extracellular space is NaHCO₃, also referred to as the carbonic acid/bicarbonate buffer.

The carbonic acid/bicarbonate buffering system

Because CO₂ is not sufficiently soluble in blood, a number of mechanisms have had to evolve over time to facilitate the transport of CO₂ from the periphery to the lungs (Chapter 22). The most efficient of these mechanisms combines CO₂ with H₂O to form carbonic acid, which in turn dissociates into hydrogen ions and bicarbonate ions.

The reactions are reversible and subject to the law of mass action:

< ?xml:namespace prefix = "mml" />H2O + CO2 ↔ H2CO3 ↔ H+ + HCO3−

In peripheral tissues—where CO₂ is continuously produced—the reaction proceeds from the left to the right, whereas in the lung the reaction proceeds in the opposite direction. The H⁺ created by the reaction must be buffered in order to preserve physiologic pH.

The carbonic acid/bicarbonate buffer

A buffer is typically a weak acid in equilibrium with its conjugate base that minimizes changes in the pH of a solution when an acid or a base is added; bicarbonate in the previous reaction is the conjugate base of carbonic acid. The Henderson-Hasselbalch equation describes the effects of changes in carbonic acid and bicarbonate on pH. Because the blood concentration of H₂CO₃ is so low, it can be replaced by α × PCO₂, wherein α is the solubility coefficient for CO₂ in plasma:

pH=pK+logHCO3–α×PCO2

In human plasma, pK is 6.1, α = 0.03 mmol·L⁻¹ ·mm·Hg⁻¹ and pH is 7.4 at a body temperature of 37°C. This yields:

7.4=6.1+logHCO3–0.03×PCO2

The amount of hydrogen any chemical can buffer is highest when its pH equals pK. Yet, in human plasma, the pK of 6.1 of the carbonic acid [CO₂]/bicarbonate buffer is not within the optimal buffer range. Therefore, the buffering capacity is dependent on the HCO₃⁻/PCO₂ ratio, which is kept in a narrow range by means of neuroventilatory PCO₂ control. To preserve a pH of 7.4, the ratio of bicarbonate to partial pressure of CO₂ must be maintained at 20:1 because the log of 20 is 1.3 (1.30103). Indeed, this is the case because the normal HCO₃⁻ concentration is 24 mEq/dL and the normal PaCO₂ is 40 mm Hg. Substituting in the preceding equation yields

pH = 6.1 + log of [24/(0.03 × 40)] = 6.1 + log of [24/1.2]= 6.1 + log of 20 = 6.1 + 1.3= 7.4

Several causes of metabolic acidosis (Box 95-1) may occur simultaneously in patients with critical illness. Acidemia may impair cardiac contractile function, constrict pulmonary arteries, and reduce adrenergic receptor responsiveness to catecholamines; therefore, many clinicians often restore the HCO₃⁻/PCO₂ ratio to 20 by administering bicarbonate. Treatment of the underlying problem is then undertaken to ultimately restore and maintain a pH of 7.4.