Alternative Measures of AcidBase Imbalance

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CLINICAL CORRELATION 25.9 Salicylate Poisoning
Salicylates are the most common cause of poisoning in children. A typical pathway of salicylate intoxication is plotted in the accompanying figure. The first effect of salicylate overdose is stimulation of the respiratory center, resulting in respiratory alkalosis. Renal compensation occurs, lowering the
A typical pathway of salicylate intoxication. Data replotted from Singer, R. B. Medicine (Baltimore) 33:1, 1954.
plasma [HCO3–]. A second, delayed effect of salicylate may then appear, metabolic acidosis. Since [HCO3–] had been lowered by the previous compensatory process, the victim is at a particular disadvantage in dealing with the metabolic acidosis. In addition, but not shown in the graph, respiratory stimulation sometimes persists after the acidosis has run its course. Rational management of salicylate intoxication requires knowledge of the plasma pH and the plasma [HCO3–] or its equivalent throughout the course of the condition.
in compensated respiratory alkalosis, with their plasma pH within the normal range. For the other types of acid–base imbalance, the exact degree of compensation expected of a patient with a given clinical picture is well worked out, but a detailed discussion is beyond the scope of this chapter. Suffice it to say that if a patient is compensating, but not as well as expected, this is taken to mean that the patient cannot compensate appropriately and must therefore have a mixed acid–base disturbance.
25.13— Alternative Measures of Acid–Base Imbalance
Modern clinical laboratories generally report plasma bicarbonate concentration, and the value is used by physicians just as we have used it here. Some laboratories, however, report total plasma CO2, that is, the sum of bicarbonate and dissolved CO2, and this is always slightly higher than [HCO3–]. At pH 7.4, for example, the ratio of [HCO3–] to [CO2] is 20 : 1 (dissolved CO2 is only 1 : 21 of the total CO2); if [HCO3–] is 24 meq L–1, [CO2] is 1.2 meq L–1 and total CO2 is 25.2 meq L–1. At pH 7.1, HCO3– is still 10 times as concentrated as dissolved CO2. Because the major contributor to total CO2 is HCO3–, total CO2 is often used in the same manner as bicarbonate to make clinical judgments. Strictly speaking, total CO2 also includes that in carbamino proteins, but current clinical laboratory practice is to ignore this when making a blood gas and pH report. If it were included in a total CO2 measurement, it would not change the interpretation of the measurement, since the CO2 in carbamino proteins, like dissolved CO2, represents only a small fraction of the total CO2.
The clinical importance of bicarbonate as a gauge of the whole body's ability to buffer further loads of metabolic acid (see Clin. Corr. 25.9) has led to several ways of expressing what the [HCO3–] would be if there were no respiratory component or respiratory compensation involved in a patient's condition. Base excess is one of these expressions. It is defined as the amount of acid that would have to be added to blood to titrate it to pH 7.4 at a , only the metabolic contribution to acid–
base imbalance (primary metabolic imbalance and nonrespiratory compensatory processes) would be measured. If a blood sample were acidic under the conditions of the titration, alkali would have to be added instead of acid, and the base excess would be negative.
The concept and the quantitation of base excess are most easily understood from the pH–bicarbonate diagram. In our discussion of the blood buffer line we saw how increasing the in blood, where other buffers are present, would result in a rise in [HCO3–] and a virtually identical decrease in the concentration of other buffer bases. This was because equivalent amounts of the other buffer bases were consumed as they buffered carbonic acid. Since virtually all the carbonic acid formed was buffered, for every HCO3– formed one conjugate base of some other system was consumed. In this situation the total base in the blood is not measurably changed; only the distribution of HCO3– and nonbicarbonate buffer conjugate base is changed. Thus, as long as one remains on the blood buffer line, [HCO3–] can change but total base will not. There will be no positive or negative base excess.
If, however, renal activity, diet, or some metabolic process adds or removes HCO3–, then a positive or negative base excess will occur. The patient's status will no longer be described by a point on the buffer line, and the base excess will be the difference between the observed plasma [HCO3–] and the [HCO3–] on the buffer line at the same pH (Figure 25.26). To calculate this difference, the position of the buffer line, which can be determined from knowledge of the slope and the point representing the normal state, must be known. In the