Basic Concepts
Intracellular and extracellular buffers are the most immediate mechanism of defense against changes in systemic pH. Bone and proteins constitute a substantial proportion of these buffers. However, the most important buffer system is the HCO3−/CO2 buffer system. The Henderson-Hasselbach equation (Equation 1) describes the relationship of pH, bicarbonate (HCO3−), and PCO2:
where HCO3− is in milliequivalents per liter and PCO2 is in millimeters of mercury. Equation 2 represents the reaction (water [H2O]):
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This buffer system is physiologically most important because of its quantitative capacity to buffer acid or alkali loads and because of the capacity for independent regulation of HCO3− and PCO2 by the kidneys and lungs, respectively. In fact, this latter aspect of independent regulation is the most powerful aspect of this system. Although the lungs and kidneys can compensate for disorders of the other, normal homeostasis requires that both CO2 and HCO3− be normal. Disorders of CO2 are usually referred to as respiratory disorders, and disorders of HCO3− or fixed acids are referred to as metabolic disorders.
Arterial CO2 is predominantly regulated by alveolar ventilation after production in peripheral tissues; CO2 is often referred to as a gaseous acid, because its addition to aqueous solutions produces carbonic acid, which then releases H+ and HCO3− (Equation 2 driven to the right). Plasma HCO3− is predominantly regulated by renal acid-base handling, and it will be discussed extensively below. The kidneys reabsorb, produce, and in some circumstances, excrete HCO3−. Plasma HCO3− is normally consumed daily by dietary acids and metabolic acids. As expressed by Equation 1, raising HCO3− or lowering PCO2 will raise systemic pH, and lowering HCO3− or raising PCO2 will lower pH.
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In physiologic systems, the addition of an acid and the loss of alkali are essentially equivalent; for instance, as obvious in Equation 2, loss of HCO3− will pull the equation to the right, producing more H+. A frequent example of this is the loss of HCO3− with diarrhea or proximal renal tubular acidosis. Conversely, the physiologic addition of alkali and the loss of acid are essentially equivalent. Therefore, the excretion of acid by the kidneys is equivalent to the production of base or HCO3−; this point becomes important in considering below how the kidneys produce more HCO3−.
Typical high-animal protein Western diets and endogenous metabolism produce acid, typically on the order of 1 mEq/kg body wt per day or approximately 70 mEq/d for a 70-kg person. Phosphoric acid and sulfuric acid are significant products of this normal metabolism of dietary nutrients, such as proteins and phospholipids. To maintain acid-base homeostasis, these nonvolatile acids must be excreted by the kidney. Other nonvolatile acids, such as ketoacids and lactic acids, are produced in pathologic conditions. Nonvolatile acid loads (or loss of HCO3−, which is an equivalent process) in excess of the excretory capacity of the kidneys cause metabolic acidosis. Of note, vegetarian diets with high fruit and vegetable content are not acid producing and may produce a net alkali load (5); this may be an important consideration in the progression or treatment of CKD (6). Although the kidneys normally control plasma HCO3−, a few other factors have been considered. Endogenous acid production may be regulated, at least under certain circumstances (7); for instance, lactic acid and ketoacid production are decreased by a low pH. Also, hepatic production of HCO3− in the metabolism of proteins and amino acids is altered by systemic acid-base balance. Therefore, a role for hepatic contribution to the control of plasma HCO3− has been hypothesized (8).
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