Chapter 7

Acid­base balance

OUTLINE

• Hydrogen ion homeostasis
• Concept of pH
• Buffer systems
• Henderson­Hasselbach equation
• Does the patient have an acid­base disorder
• Calculated vs. measured HCO₃^-
• Acid­base nomogram
• Acidemia and alkalemia
• Acidosis and alkalosis
• Anion gap
• Primary vs. compensatory processes
• Acid­base map
• In vivo titration curve for carbon dioxide
• Base excess
• Acute vs. chronic respiratory disorders
• Acute vs. chronic metabolic disorders
• Mixed acid­base disorders
• Clinical approach to acid­base diagnosis
  

HYDROGEN ION HOMEOSTASIS

"Life is a struggle, not against sin, not against the Money Power, not against malicious animal magnetism, but against hydrogen ions."
H.L. MENCKEN

Mencken was neither a physician nor physiologist, but he knew the importance of hydrogen ions. Enzyme systems operate at an optimal hydrogen ion concentration ([H⁺]), and variation from this optimal can markedly affect enzyme activity. For the blood plasma, optimal [H⁺] is 40 nanomoles/L. As shown in Table 7­1, the importance of H⁺ is out of proportion to its minuscule concentration.
Strictly speaking, hydrogen ions are protons and do not exist in the naked state in body fluids; instead they react with water (H20) to form hydronium ions, such as H30⁺ and H5O₂⁺. For clinical purposes H⁺ can be used to represent these hydrated protons. Because [H⁺] is so critical to enzyme function yet the absolute concentration is small and difficult to manipulate, the concept of pH was developed and is now universally used to represent [H⁺].^*

CONCEPT OF pH

The pH is the negative logarithm of the hydrogen ion concentration ([H⁺]):
(Eqn 7-1)

A pH of 7.4 represents a [H⁺] of 40 nmoles/L, or 4 x 10­5 moles/L (for univalent ions, mmoles/ L equal mEq/L). By definition pH does not have units.

Since pH is the negative log of [H⁺], the lower the pH, the greater the [H⁺] and hence the greater the acidity; the higher the pH, the lower the [H⁺] and the greater the alkalinity (or the less the acidity). Use of a logarithmic expression also means that a pH change of one whole unit, e.g., from 7.0 to 8.0, represents a tenfold change in [H⁺] .

Table 7­2 shows the relationship between pH and the relative acidity of the blood. A pH change from 7.40 to 7.30 represents a 25% increase in blood [H⁺]. A similar numerical change of conventional measurement, such as an increase in serum uric acid from 7.3 mg% to 7.4 mg%, represents only a 1.4% increase.
The range of normal arterial pH (7.36 to 7.44) encompasses approximately two standard deviations of the normal population; anything outside this range is considered abnormal. Clinically, the "safe" range for pH is approximately 7.30 to 7.52; within this range, pH per se is not usually life­threatening. A pH outside this range is potentially life­threatening because of altered enzymatic activity and enhanced myocardial irritability, and direct steps should be taken to return the pH to normal. Although 7.30 to 7.52 may at first seem a narrow range, it represents a [H⁺] ranging from 50 to 30 nmoles/L or a change from the normal 40 nmoles/L of plus or minus 25%. A similar range for serum sodium is 175 to 105 mEq/L!

BUFFER SYSTEMS

A buffer system counteracts the effects of adding acid or alkali to the blood. The resulting pH change is less than if the buffer were not present. Blood contains two basic buffer systems: bicarbonate and nonbicarbonate. Each consists of a weak acid or acids and their conjugate base or bases.
The bicarbonate system buffers the effects of fixed acids and alkalies that are added to the blood; the acid component is H2CO3 and the base is HCO₃^-. The nonbicarbonate system consists mainly of proteins and phosphates and serves to buffer changes in carbon dioxide. These two systems are represented by the equations in Fig. 7­1. Since the nonbicarbonate system is a heterogeneous group of compounds, the acid component is represented by HBuf and the base by Buf. Note that carbon dioxide is part of an open system, since any buildup in plasma (aqueous or dissolved CO₂) can be excreted by healthy lungs.

Fig. 7­1. Bicarbonate and nonbicarbonate buffer system. The two systems are in equilibrium with each other.

The bicarbonate and nonbicarbonate buffer systems are in equilibrium with each other. Measuring the components of either system will give the hydrogen ion concentration ([H⁺]) or the pH of the blood. However, since the nonbicarbonate system is a heterogeneous group of molecules, it is easier to measure the bicarbonate buffer components in order to determine pH.
An extremely small quantity of H2C03 is present in the blood compared with dissolved CO₂ (approximately 1 to 400). Since H2CO3 is in equilibrium with dissolved CO₂, the latter (measured as PaCO₂) can be used as the acid component in calculating pH. Therefore measurement of HCO₃^- and PaCO₂ will provide the pH.

HENDERSON­HASSELBALCH EQUATION

The Henderson­Hasselbalch equation relates blood pH to the components of the bicarbonate buffer system, as shown in Equation 2.

(Eqn 7-2)

where pK is the negative log of the dissociation constant of carbonic acid and has the value 6.1. The pH of the blood is equal to the pK of the bicarbonate buffer system plus the logarithm of the following ratio bicarbonate concentration ([HCO₃^-]) over 0.03 times the arterial partial pressure of carbon dioxide (PaCO₂). The constant 0.03 converts PaCO₂ from mm Hg to mmoles/L. Inserting normal values gives 7.4, the normal blood pH.

(Eqn 7-3)


(Eqn 7-4)


(Eqn 7-5)

It is not necessary to memorize the full Henderson­Hasselbalch equation to intelligently manage acid­base disorders. It is important to understand that pH reflects a ratio of HCO₃^- to PaCO₂.
The bicarbonate buffer system is the most important of the body's buffer systems for several reasons. This system provides the major way to buffer the additions of fixed acid and alkali to the blood. Since one of its components is carbon dioxide, the system is open, i.e., the respiratory system allows for excretion of huge amounts of carbon dioxide. Also, since carbon dioxide is readily diffusible across all cell membranes, the results of buffering can be reflected quickly in intracellular compartments.
Since there are three variables in the bicarbonate buffer system (Equation 2), measurement of any two will define the third. The body preferentially wants to maintain normal pH and does so by altering the numerator (HCO₃^-) or denominator (PaCO₂) of the Henderson­Hasselbalch equation as necessary.

DOES THE PATIENT HAVE AN ACID­BASE DISORDER?

It is important to recognize when a patient has an acid­base disorder since that recognition is the first step toward diagnosis and therapy. If any of the three variables in the Henderson­Hasselbalch equation are abnormal, the answer to this question is yes. Any acid­base derangement will be reflected in one or more components of the bicarbonate system: pH, PaCO₂, HCO₃^- (see the box on p. 20 for the range of normal values).
A single abnormal component, even without knowledge of the other two, always indicates an acid­base disorder. This is particularly important since an abnormal HCO₃^- is often found in venous blood (as part of the serum electrolytes measurement) without a concomitant blood gas measurement. An abnormal HCO₃^- value alone cannot define or diagnose an acid­base disorder but nonetheless points to its presence. For example, an elevated HCO₃^- suggests either metabolic alkalosis or respiratory acidosis.

CALCULATED VS. MEASURED HCO₃^-

Incorrect therapeutic decisions can occur if blood gas values are accepted at face value. They should always be examined for physiologic correctness, particularly when considering acid­base disorders, which seem prone to misdiagnosis. For example, a PaCO₂ of 49 mm Hg, pH of 7.35, and HCO₃^- of 16 mEq/L may be interpreted as a metabolic acidosis (low pH and low HCO₃^-) when in fact there is a transcription error: the HCO₃^- should be 26 and cannot possibly be 16 if the pH is 7.35 and the PaCO₂ is 49 mm Hg.
Such errors can be avoided if it is remembered that HCO₃^-, PaCO₂, and pH must satisfy the Henderson­Hasselbalch equation. If PaCO₂ and pH have been measured, arterial HCO₃^- can be calculated and does not have to be measured. The HCO₃^- is routinely measured as one of the serum electrolytes (on venous blood), and this measurement can pose a problem when a comparison is made with the blood gas HCO₃^-. Often, the measured venous HCO₃^- does not agree with the arterial HCO₃- that has been calculated from the Henderson­Hasselbalch equation. When this happens there are several possible reasons as shown in the box below.

Note that the pK of 6.1, on which the calculated HCO₃^- is based, may vary among patients. The significance of such variation is somewhat controversial (Hood and Campbell, 1981). At most, the variation is slight (+ 0.012 for extreme conditions) and would not account for the wide discrepancy often found between measured venous HCO₃^- and calculated arterial HCO₃^-.

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