Fundamental Concepts
As pulmonary physicians have become increasingly involved in the care of critically ill patients, a thorough understanding of acid-base metabolism has become indispensable. Regulation of arterial pH is a critical factor in maintaining stable extracellular fluid and intracellular acid-base homeostasis. Arterial pH is kept under tight control by both pulmonary and renal mechanisms, each of which must also regulate other processes such as gas exchange in the lungs and fluid and electrolyte balance by the kidneys. Although arterial pH is usually well guarded, it can be preempted by other priorities. For example, hypoxia stimulates the carotid bodies, resulting in hyperventilation and respiratory alkalosis. Furthermore, metabolic alkalosis is often perpetuated by the renal response to contraction of the extracellular volume in patients who have had severe vomiting.
Understanding how respiratory and metabolic mechanisms interact to govern pH has been complicated by the introduction of a bewildering assortment of conflicting acid-base approaches. The relative merits of each must be judged in terms of carefully selected chemical definitions and concepts, which are briefly considered in this chapter. This is followed by a review of some of the more important disorders of acid-base balance.
Acid-Base Chemistry
pH Versus H +
In aqueous solutions, “free” hydrogen ions (protons) are associated with clusters of water molecules, but for convenience these are designated as H + or H 3 O + . Rather than expressing acidity in terms of H + concentration ([H + ], normally 35 to 45 nanomoles/L in plasma), the logarithmic function (“pH,” normally 7.35 to 7.45 in plasma) is generally preferred both for convenience of representing a broad range of concentrations and because the free energy associated with changes in hydrogen concentration is related to the ratio rather than the difference between these concentrations:
pH = − log 10 [ H + ]
pH = − log 10 ( a H + )
Conjugate Acids and Bases
The Brønsted-Lowry (BL) concept has largely supplanted the Arrhenius and earlier approaches for describing acid-base reactions in chemical, physiologic, and clinical studies. By the BL criteria, an acid is a proton (H + ) donor, whereas a base is an H + acceptor. For example, BL acids are designated in roman font and BL bases in italics in the following reaction:
CH 3 COOH + H 2 O ⇌ C H 3 C O O − + H 3 O +
Strong Versus Weak Ions
Ions that are completely ionized in water (e.g., Na + , K + , and Cl − ) are not considered as BL bases or acids because they neither accept nor donate H + ions, and they are sometimes referred to as “spectator ions.” Because the pH of the extracellular fluid is normally approximately 7.4, some interpretations of acid-base chemistry categorize many organic acids (with dissociation constants below ≈ 4.0, see later) as “strong” acids because less than 0.1% of these acids remains undissociated in the extracellular milieu. Detection of excess concentrations of relatively strong anions such as lactate may indicate excessive intake, production, or retention of lactic acid. However, from the BL perspective, the corresponding lactate anion behaves as a weak base rather than an acid because lactate anions can accept H + ions. Furthermore, the presence of lactate anions may actually reflect infusions of solutions that promote alkalosis rather than acidosis. For example, an infusion of Ringer lactate initially dilutes the plasma, which tends to cause a dilutional acidosis (see later). However, subsequent metabolism of lactate to 
results in alkalinization.
Buffer Systems.
Conjugate acid-base pairs can be used to minimize changes in pH when strong acids or bases (e.g., HCl and NaOH) are added to aqueous solutions. If the constituents of a solution are neither created nor destroyed and do not exchange with the environment, the system is referred to as “closed.” The effectiveness of a closed buffer system is maximal when the concentrations of the conjugate acids and bases are similar and exceed those of the strong acids or bases that are added, in other words, when the H 3 O + of the solution is close to the dissociation constant (K a ) of the buffer pair. For a weak acid (HA):
HA + H 2 O ⇌ H 3 O + + A −
K a = [ H 3 O + ] [ A − ] [ H A ]
p H = p K a + log [ A − ] [ H A ] = p K a + log [ b a s e ] [ a c i d ]
Arterial pH is maintained at approximately 7.4, well above the pKa of 
/P co 2 buffer pair (6.1). This reflects the fact that this buffer pair is volatile and bicarbonate concentrations are kept at concentrations 20 times greater than those of dissolved carbon dioxide. Relatively high concentrations of 
relative to those of carbon dioxide reflect in part a steady state relationship between the lungs and kidney that presumably consumes more energy than buffers maintained at equilibrium , but that allows this acid-base pair to efficiently neutralize nonvolatile acids produced in the body.
Carbon Dioxide and Bicarbonate
By the early 20th century the importance of reactions of carbon dioxide and bicarbonate ion ( 
) with H + and OH − in acid-base balance was well established :
CO 2 + OH − ⇌ CA CA HCO 3 −
CO 2 + H 2 O ⇌ CA CA H 2 CO 3 ⇌ H + + HCO 3 −
from carbon dioxide, but rates of formation of either H 2 CO 3 (carbonic acid) in reaction 6 or 
directly in reaction 5 are relatively slow in the absence of a catalyst. These rates are normally accelerated by carbonic anhydrase (CA) present in erythrocytes, vascular endothelium, alveolar epithelium, and in most other organs, including the kidney. Under physiologic conditions, [H 2 CO 3 ] is much less concentrated than [carbon dioxide], and the relative amounts of 
and carbon dioxide can be calculated from the conventional Henderson-Hasselbalch equation:p H = p K a + log [ H C O 3 − ] α P C O 2
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