Acid-Base Balance

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:

<SPAN role=presentation tabIndex=0 id=MathJax-Element-1-Frame class=MathJax style="POSITION: relative" data-mathml='pH=−log10[H+]’>pH=log10[?+]pH=−log10[H+]
pH = − log 10 [ H + ]
or more precisely,
<SPAN role=presentation tabIndex=0 id=MathJax-Element-2-Frame class=MathJax style="POSITION: relative" data-mathml='pH=−log10(aH+)’>pH=log10(??+)pH=−log10(aH+)
pH = − log 10 ( a H + )
where (a H+ ) designates the “activity,” of H + . (a H+ ) is determined with a hydrogen ion electrode and numerically approaches the concentration of H + in dilute solutions. Although the concentrations of H + in tissue fluids are typically very low compared to those of electrolytes, they can be responsible for important free energy differences across cellular membranes if the ratio between compartmental concentrations is large. For example, much of the energy stored in mitochondria is attributable to the ratio of H + ion concentrations which is maintained across the inner membrane of the mitochondria (see later). Because the concentration of H + is routinely divided by the thermodynamic standard state activity of a solution containing 1 mole/L of H + ions, no units are used for a H+ or pH.

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:

<SPAN role=presentation tabIndex=0 id=MathJax-Element-3-Frame class=MathJax style="POSITION: relative" data-mathml='CH3COOH+H2O⇌CH3COO−+H3O+’>CH3COOH+?2???3???+?3?+CH3COOH+H2O⇌CH3COO−+H3O+
CH 3 COOH + H 2 O ⇌ C H 3 C O O − + H 3 O +
where CH 3 COOH and CH 3 COO represent the “conjugate acid-base pair” for acetic acid. CH 3 COOH loses one H + when the reaction proceeds to the right. H 3 O + and H 2 O represent the conjugate acid-base pair of water. H 2 O accepts one H + as the reaction proceeds to the right.

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):

<SPAN role=presentation tabIndex=0 id=MathJax-Element-4-Frame class=MathJax style="POSITION: relative" data-mathml='HA+H2O⇌H3O++A−’>HA+?2??3?++?HA+H2O⇌H3O++A−
HA + H 2 O ⇌ H 3 O + + A −

<SPAN role=presentation tabIndex=0 id=MathJax-Element-5-Frame class=MathJax style="POSITION: relative" data-mathml='Ka=[H3O+][A−][HA]’>??=[?3?+][?][??]Ka=[H3O+][A−][HA]
K a = [ H 3 O + ] [ A − ] [ H A ]
where K a is the acid-dissociation constant. Rearranging this equation and taking the logarithms of both sides yields the generalized Henderson-Hasselbalch equation:
<SPAN role=presentation tabIndex=0 id=MathJax-Element-6-Frame class=MathJax style="POSITION: relative" data-mathml='pH=pKa+log[A−][HA]=pKa+log[base][acid]’>??=???+log[?][??]=???+log[????][????]pH=pKa+log[A−][HA]=pKa+log[base][acid]
p H = p K a + log [ A − ] [ H A ] = p K a + log [ b a s e ] [ a c i d ]
Buffering is maximal when pH = pK a . It is usually assumed that concentrations of H + released from water are negligible compared to those derived from the acid.

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 :

<SPAN role=presentation tabIndex=0 id=MathJax-Element-7-Frame class=MathJax style="POSITION: relative" data-mathml='CO2+OH−⇌CACAHCO3−’>CO2+OHCACAHCO3CO2+OH−⇌CACAHCO3−
CO 2 + OH − ⇌ CA CA HCO 3 −

<SPAN role=presentation tabIndex=0 id=MathJax-Element-8-Frame class=MathJax style="POSITION: relative" data-mathml='CO2+H2O⇌CACAH2CO3⇌H++HCO3−’>CO2+?2?CACA?2CO3?++HCO3CO2+H2O⇌CACAH2CO3⇌H++HCO3−
CO 2 + H 2 O ⇌ CA CA H 2 CO 3 ⇌ H + + HCO 3 −
reaction 6 provides the dominant pathway to 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:
<SPAN role=presentation tabIndex=0 id=MathJax-Element-9-Frame class=MathJax style="POSITION: relative" data-mathml='pH=pKa+log[HCO3−]αPCO2′>??=???+log[???3]????2pH=pKa+log[HCO3−]αPCO2
p H = p K a + log [ H C O 3 − ] α P C O 2

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Jul 21, 2019 | Posted by in CARDIOLOGY | Comments Off on Acid-Base Balance
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