Disorders of Acid-Base Balance

Chapter Outline
ACID-BASE HOMEOSTASIS, 512
- Buffer Systems, 512
- Mechanisms of pH Buffering, 512
- Chemical Equilibria of Physicochemical Buffer Systems, 512
- Chemical Equilibria for the Carbon Dioxide–Bicarbonate System, 513
- Physiologic Advantage of an Open Buffer System, 513
- Regulation of Buffers, 513
- Integration of Regulatory Processes, 514
RENAL REGULATION, 514
SYSTEMIC RESPONSE TO CHANGES IN CARBON DIOXIDE TENSION, 515
- Acute Response: Generation of Respiratory Acidosis or Alkalosis, 515
- Chronic Response, 516
SYSTEMIC RESPONSE TO ADDITION OF NONVOLATILE ACIDS, 516
- Sources of Endogenous Acids, 516
- Hepatic and Renal Roles in Acid-Base Homeostasis, 517
- Neurorespiratory Response to Acidemia, 518
- Renal Excretion, 518
SYSTEMIC RESPONSE TO GAIN OF ALKALI, 519
- Distribution and Cellular Buffering, 519
- Respiratory Compensation, 519
- Renal Excretion, 519
STEPWISE APPROACH TO THE DIAGNOSIS OF ACID-BASE DISORDERS, 520
- Step 1: Measure Arterial Blood Gas and Electrolyte Values Simultaneously, 521
- Step 2: Verify Acid-Base Laboratory Values, 521
- Step 3: Define the Limits of Compensation to Distinguish Simple from Mixed Acid-Base Disorders, 521
- Step 4: Calculate the Anion Gap, 523
- Step 5 and 6: Recognize Conditions Causing Acid-Base Abnormalities with High or Normal Anion Gap, 524
RESPIRATORY DISORDERS, 525
- Respiratory Acidosis, 525
- Respiratory Alkalosis, 527
METABOLIC DISORDERS, 528
- Metabolic Acidosis, 528
- Metabolic Alkalosis, 550

The appropriate diagnosis and management of acid-base disorders requires accurate interpretation of the specific acid-base disorder, as well as consideration of the clinical setting in which these disorders occur as indicated by the patient’s history and physical examination. Accuracy requires simultaneous measurement of plasma electrolyte levels and arterial blood gas (ABG) values, as well as an appreciation by the clinician of the physiologic adaptations and compensatory responses that occur with specific acid-base disturbances. In most circumstances, these compensatory responses can be predicted through an analysis of the prevailing disorder in a stepwise manner. This chapter reviews acid-base homeostasis as a consequence of the integration of physiologic and compensatory responses. The approach in this chapter, sometimes referred to as the Boston method, uses measurements of arterial pH, CO ₂pressure (P co ₂), and [HCO ₃⁻] plus an analysis of the anion gap (corrected for a normal plasma albumin level of 4.5 g/dL), referring to the pathophysiologically established range of compensatory responses for simple disorders. This is the most widely used and generally accepted approach used clinically by nephrologists, and the easiest model to understand. Other methods, such as one requiring calculation and consideration of “base excess” (the Copenhagen method), or a physicochemical method (Stewart method), requiring calculation of the “strong ion difference,” the “strong ion gap,” and the total concentration of plasma weak acids, are not discussed here.

Acid-Base Homeostasis

Acid-base homeostasis operates to maintain systemic arterial pH within a narrow range. Although clinical laboratories consider the normal range to be between 7.35 and 7.45 pH units, pH in vivo in an individual is maintained within a much narrower range. This degree of tight regulation is accomplished through (1) chemical buffering in the extracellular fluid (ECF) and the intracellular fluid and (2) regulatory responses that are under the control of the respiratory and renal systems. Those chemical buffers, respiration, and renal processes efficiently dispose of the physiologic daily load of carbonic acid (as volatile CO ₂) and nonvolatile acids, mainly derived from dietary protein intake, and defend against the occasional addition of pathologic quantities of acid and alkali. Therefore chemical buffers within the extracellular and intracellular compartments serve to blunt changes in pH that would occur with retention of either acids or bases. In addition, the control of P co ₂by the central nervous system and respiratory system and the control of the plasma HCO ₃⁻by the kidneys constitute the regulatory processes that act in concert to stabilize the arterial pH.

The major buffer system in the body comprises a base (H ⁺acceptor), which is predominantly HCO ₃⁻, and an acid (H ⁺donor), which is predominantly carbonic acid (H ₂CO ₃):

H++HCO3−⇔H2CO3H++HCO3−⇔H2CO3
H + + HCO 3 − ⇔ H 2 CO 3

Extracellular H ⁺concentration ([H ⁺] _e) throughout the body is constant in the steady state. The HCO ₃⁻/H ₂CO ₃ratio is proportional to the ratio of all the other extracellular buffers such as PO ₄³⁻and plasma proteins:

[H+]e∝HCO3−H2CO3∝B−HB[H+]e∝HCO3−H2CO3∝B−HB
[ H + ] e ∝ HCO 3 − H 2 CO 3 ∝ B − HB

The intracellular H ⁺concentration ([H ⁺] _i), or pH _i, is also relatively stable. Both cellular ion exchange mechanisms and intracellular buffers (hemoglobin, tissue proteins, organophosphate complexes, and bone apatite) participate in the blunting of changes in both [H ⁺] _iand [H ⁺] _e. Extracellular and intracellular buffers provide the first line of defense against the addition of acid or base to the body (see “ Mechanisms of pH Buffering ” section later).

The second line of defense is the respiratory system. Pulmonary participation in acid-base homeostasis relies on the excretion of CO ₂by the lungs. The reaction is catalyzed by the enzyme carbonic anhydrase:

H++HCO3−↔H2CO3←→−−−anhydraseCarbonicH2O+CO2H++HCO3−↔H2CO3↔anhydraseCarbonicH2O+CO2
H + + HCO 3 − ↔ H 2 CO 3 ↔ anhydrase Carbonic H 2 O + CO 2

Large amounts of CO ₂(10 to 12 mol/day) accumulate as metabolic end products of tissue metabolism. This CO ₂load is transported in the blood to the lungs as hemoglobin-generated HCO ₃⁻and hemoglobin-bound carbamino groups :

Metabolism→CO2←→−−transportBloodLungsMetabolism→CO2↔transportBloodLungs
Metabolism → CO 2 ↔ transport Blood Lungs

Conventionally H ⁺concentration is expressed in two different ways, either directly as [H ⁺] or indirectly as pH. The relationship between these two factors can be written in mathematically equivalent forms:

pH=−log10[H+]pH=−log10[H+]
pH = − log 10 [ H + ]

[H+](Eq/L)=10−pH[H+](Eq/L)=10−pH
[ H + ] ( Eq / L ) = 10 − pH

When [H ⁺] is expressed (for numeric convenience) in nanomoles per liter (nmol/L) or nanomolar (nM), then:

[H+]=109−pH[H+]=109−pH
[ H + ] = 10 9 − pH

Buffer Systems

Acid-base chemistry deals with molecular interactions that involve the transfer of H ⁺. A large variety of molecules, both inorganic and organic, contain hydrogen atoms that can dissociate to yield H ⁺. The relationship between an undissociated acid (HA) and its conjugate, disassociated base (A ⁻) may be represented as follows:

HA⇔H++A−HA⇔H++A−
HA ⇔ H + + A −

In addition to the many inorganic and organic acid-base substances encountered in biologic systems, many protein molecules (e.g., hemoglobin) contain acidic groups that may dissociate, yielding a corresponding conjugate base.

Mechanisms of pH Buffering

Buffer systems are critical to the physiology and pathophysiology of acid-base homeostasis because they attenuate the pH change in a solution or tissue by reversibly combining with or releasing H ⁺. Thus the pH change of a solution during the addition of acid or base equivalents is smaller in the presence of a buffer system than would have occurred if no buffer systems were present. The acid or base load can be extrinsic, such as during systemic acid or base infusion, or intrinsic, resulting from net generation of new acid or base equivalents that are added to the extracellular or intracellular compartments.

Chemical Equilibria of Physicochemical Buffer Systems

As an example of a physicochemical buffer pair, consider a neutral weak acid (HA) and its conjugate weak base (A ⁻). Examples of such buffer pairs are acetic acid and acetate and the carboxyl groups on proteins. Another example of a physicochemical buffer pair is a neutral weak base (B) and its conjugate weak acid (BH ⁺):

BH+⇔B+H+BH+⇔B+H+
BH + ⇔ B + H +

Examples of such buffer pairs are NH ₃and NH ₄⁺and the imidazole group in proteins. A rigorous analysis of the kinetics of reversible reactions in solution yields the law of mass action, which states that at equilibrium (i.e., when the velocities of the forward and backward reactions are equal) the ratio of the concentration products of opposing reactions is a constant:

K′a=[H+][A−]HAK′a=[H+][A−]HA
K ′ a = [ H + ] [ A − ] HA

K′b=[H+][B−]BHK′b=[H+][B−]BH
K ′ b = [ H + ] [ B − ] BH

K′ _aand K′ _bare the equilibrium or dissociation constants for equations 10 and 11 , respectively.

Taking logarithms of both sides of equations 10 and 11 and defining pK′ _a= −log ₁₀(K′ _a) and pK′ _b= −log ₁₀(K′ _b) yields:

pH=pK′a+log10[A−][HA]pH=pK′a+log10[A−][HA]
pH = p K ′ a + log 10 [ A − ] [ HA ]

pH=pK′b+log10[B−][BH]pH=pK′b+log10[B−][BH]
pH = p K ′ b + log 10 [ B − ] [ BH ]

The dissociation constants K′ _aand K′ _bprovide an estimate of the strength of the acid and base, respectively. From equations 12 and 13 , it can be seen that the buffer pairs are half dissociated at pH = pK′. In other words, pK′ of a buffer pair is defined as the pH at which 50% of the buffer pair exists as the weak acid (HA) and 50% as the anion (A ⁻).

Chemical Equilibria for the Carbon Dioxide–Bicarbonate System

When CO ₂is dissolved in water, H ₂CO ₃is formed according to the reaction

CO2+H2O⇔H2CO3CO2+H2O⇔H2CO3
CO 2 + H 2 O ⇔ H 2 CO 3

The rate of this reaction, in the absence of the enzyme carbonic anhydrase, is slow, with a half-time of approximately 8 seconds at 37° C. The major portion of CO ₂remains as dissolved CO ₂; only approximately 1 part in 1000 forms H ₂CO ₃, a nonvolatile acid. Because H ₂CO ₃is a weak acid, it dissociates to yield H ⁺and HCO ₃:

H2CO3⇔H++HCO3−H2CO3⇔H++HCO3−
H 2 CO 3 ⇔ H + + HCO 3 −

The concentration of dissolved CO ₂is given by Henry’s law:

[CO2]dis=αCO2PCO2[CO2]dis=αCO2PCO2
[ CO 2 ] dis = α CO 2 PCO 2

where

is the physical solubility coefficient for CO ₂, which has a value of 0.0301 mmol/L in most body fluids, including plasma. Because the concentration of H ₂CO ₃is low and proportional to the concentration of dissolved CO ₂, equations 14 and 15 can be combined and treated as a single reaction:

CO2+H2O⇔H++HCO3−CO2+H2O⇔H++HCO3−
CO 2 + H 2 O ⇔ H + + HCO 3 −

The equilibrium constant for this reaction is given by:

K=[H+][HCO3−][CO2][H2O]K=[H+][HCO3−][CO2][H2O]
K = [ H + ] [ HCO 3 − ] [ CO 2 ] [ H 2 O ]

Defining K′ = K[H ₂O] as the apparent equilibrium constant and using equation 17 :

K′=[H+][HCO3−]αCO2PCO2K′=[H+][HCO3−]αCO2PCO2
K ′ = [ H + ] [ HCO 3 − ] α CO 2 PCO 2

Taking logarithms of both sides of equation 19 and recognizing that pK′ = log ₁₀(K′) allows the familiar Henderson-Hasselbalch equation to be derived:

pH=pK′+log10[HCO3−](αCO2PCO2)pH=pK′+log10[HCO3−](αCO2PCO2)
pH = p K ′ + log 10 [ HCO 3 − ] ( α CO 2 PCO 2 )

When pK′ = 6.1 is used in equation 20 , the Henderson equation is derived, which may be used in clinical interpretation of acid-base data:

[H+](nmol/L)=24PCO2(mm Hg)[HCO3−](mmol/L)[H+](nmol/L)=24PCO2(mm Hg)[HCO3−](mmol/L)
[ H + ] ( nmol / L ) = 24 PCO 2 ( mm Hg ) [ HCO 3 − ] ( mmol / L )

Physiologic Advantage of an Open Buffer System

The quantitative behavior of an open system buffer pair differs considerably from that of a buffer pair confined to a closed system. In an open system the buffer pair may be envisioned as occurring in two separate but communicating compartments (internal and external). The external compartment provides an effective infinite reservoir of the uncharged buffer pair component, to which the barrier between the internal and the external compartments (e.g., plasma cell membrane, vascular capillary endothelium) is freely permeable.

Physiologically, the most important open system buffer is the CO ₂-HCO ₃⁻system.

Adjustments in alveolar ventilation serve to maintain a constant arterial CO ₂pressure (Pa co ₂):

Acid(H+)(Expired gas)↓↑↑H++HCO3−→H2CO3→H2O+CO2Acid(H+)(Expired gas)↓↑↑H++HCO3−→H2CO3→H2O+CO2
Acid ( H + ) ( Expired gas ) ↓ ↑ ↑ H + + HCO 3 − → H 2 CO 3 → H 2 O + CO 2

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