Ganong's Review of Medical Physiology, 23rd Edition

614 SECTION VIIRespiratory Physiology

Fruits are the main dietary source of alkali. They contain Na+
and K+ salts of weak organic acids, and the anions of these salts
are metabolized to CO 2 , leaving NaHCO 3 and KHCO 3 in the
body. Such ingestion contributes little to changes in pH and a
more common cause of alkalosis is loss of acid from the body
as a result of vomiting of gastric juice rich in HCl. This is, of
course, equivalent to adding alkali to the body.

BUFFERING IN THE BLOOD

Acid and base shifts in the blood are largely controlled by
three main buffers in blood: (1) proteins, (2) hemoglobin, and
(3) the carbonic acid–bicarbonate system. Plasma proteins
are effective buffers because both their free carboxyl and their
free amino groups dissociate:

The second buffer system is provided by the dissociation of
the imidazole groups of the histidine residues in hemoglobin:

In the pH 7.0–7.7 range, the free carboxyl and amino
groups of hemoglobin contribute relatively little to its buffer-
ing capacity. However, the hemoglobin molecule contains 38
histidine residues, and on this basis—plus the fact that hemo-
globin is present in large amounts—the hemoglobin in blood
has six times the buffering capacity of the plasma proteins. In
addition, the action of hemoglobin is unique because the imi-
dazole groups of deoxyhemoglobin (Hb) dissociate less than
those of oxyhemoglobin (HbO 2 ), making Hb a weaker acid
and therefore a better buffer than HbO 2. Titration curves for
Hb and HbO 2 are shown in Figure 36–8.
The third and major buffer system in blood is the carbonic
acid–bicarbonate system:

The Henderson–Hasselbalch equation for this system is

The pK for this system in an ideal solution is low (about 3),

and the amount of H 2 CO 3 is small and hard to measure accu-

rately. However, in the body, H 2 CO 3 is in equilibrium with

CO2:

If the pK is changed to pK' (apparent ionization constant;

distinguished from the true pK due to less than ideal condi-

tions for the solution) and [CO 2 ] is substituted for [H 2 CO 3 ],

the pK' is 6.1:

The clinically relevant form of this equation is:

since the amount of dissolved CO 2 is proportional to the par-

tial pressure of CO 2 and the solubility coefficient of CO 2 in

mmol/L/mm Hg is 0.0301. [HCO 3 – ] cannot be measured

directly, but pH and PCO 2 can be measured with suitable

accuracy with pH and PCO 2 glass electrodes, and [HCO 3 – ]

can then be calculated.

The pK' of this system is still low relative to the pH of the

blood, but the system is one of the most effective buffer systems

in the body because the amount of dissolved CO 2 is controlled

by respiration. Additional control of the plasma concentration

of HCO 3 – is provided by the kidneys. When H+ is added to the

blood, HCO 3 – declines as more H 2 CO 3 is formed. If the extra

H 2 CO 3 were not converted to CO 2 and H 2 O and the CO 2

excreted in the lungs, the H 2 CO 3 concentration would rise.

When enough H+ has been added to halve the plasma HCO 3 – ,

the pH would have dropped from 7.4 to 6.0. However, not only

is all the extra H 2 CO 3 that is formed removed, but also the H+

rise stimulates respiration and therefore produces a drop in

RCOOH→←RCOO−+H+

pH=pK ́ [RCOO−]

RCOOH+log[RCOOH]

RNH 3 +→←RNH 2 +H+

[RNH 2 ]

pH=pK ́RNH 3 +log

[RNH 3 +]

H H

C C

HN NH+ HN N

←→ + H+

R R

CH C HC C

H 2 CO 3 ←→H++HCO 3 −

pH=pK+log

[HCO 3 −]

[H 2 CO 3 ]

FIGURE 36–8 Titration curves for hemoglobin. Individual

titration curves for deoxygenated hemoglobin (Hb) and oxygenated

hemoglobin (HbO 2 ) are shown. The arrow from a to c indicates the

number of millimoles of H that can be added without pH shift. The

arrow from a to b indicates the pH shift on deoxygenation.

+1.0

+0.5

−0.5

0

7.30 7.40 7.50 7.60 7.70

pH

c

a b

Hb

mmol

HbO 2

mmol of H+

added to

1 mmol of

HbO 2 or Hb

mmol of H+

removed

from 1

mmol of

HbO 2 or Hb

H 2 CO 3 ←→CO 2 +H 2 O

pH=6.10+log

[HCO 3 −]

[CO2]

[HCO 3 −]

pH=6.10+log

0.0301 PCO 2