Control Of Acid Base Balance
The pH of arterial blood is
normally approximately 7.4 ([H+] 40 nmol/L). Regulation of acid base status so
that blood pH remains between 7.35 and 7.45 (45-35 nmol/L) is vital for the
correct functioning of the body. Carriage of CO2 in blood and its
removal in the lungs (Chapter 9) have an important influenc on acid base
status, as about 100 times more acid equivalents are expired per day in the
form of CO2/carbonic acid than are excreted as fied acids by the
kidneys. Nevertheless, renal mechanisms are crucially important for regulation
of acid-base balance, and for compensating respiratory disorders (see below).
Buffers bind or release H+
according to the pH; this limits the change
in pH that occurs when acid is added. The relationship between the amount of
acid equivalent added to a solution containing a buffer and the resultant
change in pH is known as the buffer curve. Buffers are most effective
when pH is close to their pKA (log of dissociation constant, KA;
see Fig. 10a). The most important buffers in blood are bicarbonate (HCO3−)
and haemoglobin. CO2 combines with water to form carbonic acid
(H2CO3), which dissociates to HCO3− and H+
(Chapter 9). The relationship between pH, Pco2 and [HCO3−]
is de- scribed by the Henderson–Hasselbalch equation (Fig. 10a), where
pKA is 6.1 and [CO2] can be calculated as Pco2
CO2 solubility, which is 0.03 mmol L/mmHg (0.23 mmol L/kPa). In
normal blood, [HCO3−] is 24 mmol and Pco2 40 mmHg
(5.3 kPa), and pH calculates as 7.4. Whatever their actual values, the
important points to remember are that if the ratio [HCO3−]/[CO2]
remains constant at 20, then pH will remain at 7.4, and:
Although the pKA of
the bicarbonate system (6.1) is further away from blood pH (7.4) than would
seem ideal for a buffer, the fact that Pco2 and HCO3−
can be independently controlled by ventilation (Chapter 9) and the kidneys,
respectively, means that in practice it makes an effective buffer system.
Haemoglobin is an
important buffer, particularly when deoxygenated (Chapter 9), and significantl
improves the buffering capacity of whole blood compared with plasma (Fig. 10b;
the steeper the line, the better the buffering). All other blood proteins combined
have slightly more than 20% of the buffering capacity of haemoglobin.
Acidosis, alkalosis and compensation
The relationship between pH, HCO3−
and Pco2 can be portrayed using a Davenport diagram (Fig.
10b). HCO3− is plotted against pH for given values of Pco2.
The line marked BAC is the buffer line for whole blood; in the absence
of other changes (e.g. anaemia and polycythaemia), changes in Pco2
alter HCO3− and pH along this line. Point A represents normal
conditions (pH 7.4, HCO3− 24 mmol, Pco2 40
mmHg/5.3 kPa). An acute rise in Pco2 (hypercapnia) due to
hypoventilation (e.g. acute respiratory failure) will decrease the [HCO3−]:
Pco2 ratio and consequently pH (see above). This respiratory
acidosis is represented by a move from A to B (Fig. 10c); points A to C
represent a respiratory alkalosis (e.g. hyperventilation). A sustained
respiratory acidosis caused by chronic respiratory failure (Chapter 23) can over days be partially compensated
by excretion of H+ (as phosphate and ammonium)
and reabsorption of HCO3− in the kidneys. The [HCO3−]/Pco2
ratio is thus largely restored and pH returns towards
normal. This renal compensation is described by the arrow between B and
D (Fig. 10c). Conversely, a respiratory alkalosis may be compensated by
increased renal excretion of HCO3− (C to E).
The term metabolic acidosis (or
alkalosis) is used when acid-base status is disturbed by changes in HCO3−
rather than CO2 - as a result, for example, of renal disease or increased
H+ production (table). A metabolic acidosis (Fig. 10c, G) may be
partially compensated by increased ventilation and a reduction in Pco2
(G to E), initiated by detection of acid pH by the chemoreceptors (Chapter 11).
There can be
little respiratory
compensation for metabolic alkalosis (F), as this may require
unsustainable falls in ventilation.
Base excess
Measurement of pH alone therefore
gives little indication of acid-base status (Fig. 10d); although pH may be
normal, Pco2 and [HCO3−] may not be (D, E). Measurements of blood pH, Pco2
and Po2 are always taken clinically. Base excess (or base
deļ¬cit negative base excess) is a calculated value representing
the amount of acid that would be needed to titrate the blood back to a pH of
7.4 at a Pco2 of 5.3 kPa. For example, in Fig. 10c an
increase in Pco2 to 60 mmHg results in a respiratory acidosis
(B). Full compensation back to pH 7.4 (D) requires [HCO3−] to be
increased to approximately 35 mmol/L; thus following renal compensation (D),
the base excess is approximately 11 mmol/L (the difference between A and D). In
a pure metabolic acidosis, the base excess is negative and greater than the
difference between the actual and normal HCO3−, as haemoglobin and
buffers must also be titrated. Base excess is normally calculated from the pH
and Pco2 automatically by clinical blood gas analysers, and
corrected for haemoglobin concentration. Together with the Pco2,
base excess may be useful for diagnosis of the cause of an acid-base
disturbance, but should be used with caution as a basis for treatment as the
whole body buffer line may differ significantl from that of blood in vitro,
due to contributions from interstitial fluids (Fig. 10b).
Metabolic and respiratory acid-base
disorders may often be combined, making diagnosis diff cult. A common example
is respiratory failure (Chapter 23), where concomitant hypoxia can cause
metabolic acidosis in addition to the primary respiratory acidosis. A useful diagnostic
aid is the Flenley nomogram (Fig. 10d). Only one type of disturbance is
likely if the patient's arterial pH and Pco2 fall within a
band (95% confidenc limits).
Common causes of acid–base disorders
