BIOLOGICAL BUFFERS
Many chemical reactions
are affected by the acidity of the solution in which they occur. In order for a
particular reaction to occur or to occur at an appropriate rate, the pH of the
reaction medium must be controlled. Such control is provided by buffer
solutions, which are solutions that maintain a particular pH. Biochemical
reactions are especially sensitive to pH. Most biological molecules contain
groups of atoms that may be charged or neutral depending on pH, and whether
these groups are charged or neutral has a significant effect on the biological
activity of the molecule.
In all multicellular
organisms, the fluid within the cell and the fluids surrounding the cells have a
characteristic and nearly constant pH. This pH is maintained in a number of
ways, and one of the most important is through buffer systems. Two important
biological buffer systems are the dihydrogen phosphate system and the carbonic
acid system.
The phosphate buffer
system operates in the internal fluid of all cells. This buffer system consists
of dihydrogen phosphate ions (H2PO4-) as
hydrogen-ion donor (acid) and hydrogen phosphate ions (HPO42-)
as hydrogen-ion acceptor (base). These two ions are in equilibrium with each
other as indicated by the chemical equation below.
H2PO4-(aq) H+(aq)
+ HPO42-(aq)
If additional hydrogen ions
enter the cellular fluid, they are consumed in the reaction with HPO42-,
and the equilibrium shifts to the left. If additional hydroxide ions enter the
cellular fluid, they react with H2PO4-,
producing HPO42-, and shifting the equilibrium to the
right. The equilibrium-constant expression for this equilibrium is
Ka =
|
[H +] [HPO42-]
|
|
[H2PO4-
|
The value of Ka for
this equilibrium is 6.23 × 10-8 at
25°C. From this equation, the relationship between the hydrogen-ion
concentration and the concentrations of the acid and base can be derived.
[H +]
|
= Ka
|
[H2PO4-]
|
|
[HPO42-]
|
Thus, when the concentrations
of H2PO4- and
HPO42- are the
same, the value of the molar concentration of hydrogen ions is equal to the
value of the equilibrium constant, and the pH is equal to the pKa (-log Ka),
namely 7.21. Buffer solutions are most effective at maintaining a pH near the
value of the pKa. In mammals, cellular fluid has a pH in the
range 6.9 to 7.4, and the phosphate buffer is effective in maintaining this pH
range.
Another biological fluid
in which a buffer plays an important role in maintaining pH is blood plasma. In
blood plasma, the carbonic acid and hydrogen carbonate ion equilibrium buffers
the pH. In this buffer, carbonic acid (H2CO3) is the
hydrogen-ion donor (acid) and hydrogen carbonate ion (HCO3-)
is the hydrogen-ion acceptor (base).
H2CO3(aq) H+(aq)
+ HCO3-(aq)
This buffer functions in
exactly the same way as the phosphate buffer. Additional H+ is
consumed by HCO3- and
additional OH- is consumed
by H2CO3. The value of Ka for
this equilibrium is 7.9 × 10-7, and the pKa is
6.1 at body temperature. In blood plasma, the concentration of hydrogen
carbonate ion is about twenty times the concentration of carbonic acid. The pH
of arterial blood plasma is 7.40. If the pH falls below this normal value, a
condition called acidosis is
produced. If the pH rises above the normal value, the condition is called alkalosis.
The concentrations of
hydrogen carbonate ions and of carbonic acid are controlled by two independent
physiological systems. Carbonic acid concentration is controlled by respiration,
that is through the lungs. Carbonic acid is in equilibrium with dissolved carbon
dioxide gas.
H2CO3(aq) CO2(aq)
+ H2O(l)
An enzyme called carbonic
anhydrase catalyzes the conversion of carbonic acid to dissolved carbon dioxide.
In the lungs, excess dissolved carbon dioxide is exhaled as carbon dioxide gas.
CO2(aq) CO2(g)
The concentration of hydrogen
carbonate ions is controlled through the kidneys. Excess hydrogen carbonate ions
are excreted in the urine.
The much higher
concentration of hydrogen carbonate ion over that of carbonic acid in blood
plasma allows the buffer to respond effectively to the most common materials
that are released into the blood. Normal metabolism releases mainly acidic
materials: carboxylic acids such as lactic acid (HLac). These acids react with
hydrogen carbonate ion and form carbonic acid.
HLac(aq) + HCO3-(aq) Lac-(aq)
+ H2CO3(aq)
The carbonic acid is converted
through the action of the enzyme carbonic anhydrase into aqueous carbon dioxide.
H2CO3(aq) CO2(aq)
+ H2O(l)
An increase in CO2(aq)
concentration stimulates increased breathing, and the excess carbon dioxide is
released into the air in the lungs.
The condition called respiratory
acidosis occurs when blood pH
falls as a result of decreased respiration. When respiration is restricted, the
concentration of dissolved carbon dioxide in the blood increases, making the
blood too acidic. Such a condition can be produced by asthma, pneumonia,
emphysema, or inhaling smoke.
Metabolic acidosis is
the decrease in blood pH that results when excessive amounts of acidic
substances are released into the blood. This can happen through prolonged
physical exertion, by diabetes, or restricted food intake. The normal body
response to this condition is increases breathing to reduce the amount of
dissolved carbon dioxide in the blood. This is why we breathe more heavily after
climbing several flights of stairs.
Respiratory alkalosis results
from excessive breathing that produces an increase in blood pH. Hyperventilation
causes too much dissolved carbon dioxide to be removed from the blood, which
decreases the carbonic acid concentration, which raises the blood pH. Often, the
body of a hyperventilating person will react by fainting, which slows the
breathing.
Metabolic alkalosis is
an increase in blood pH resulting from the release of alkaline materials into
the blood. This can result from the ingestion of alkaline materials, and through
overuse of diuretics. Again, the body usually responds to this condition by
slowing breathing, possibly through fainting.
The carbonic
acid-hydrogen carbonate ion buffer works throughout the body to maintain the pH
of blood plasma close to 7.40. The body maintains the buffer by eliminating
either the acid (carbonic acid) or the base (hydrogen carbonate ions). Changes
in carbonic acid concentration can be effected within seconds through increased
or decreased respiration. Changes in hydrogen carbonate ion concentration,
however, require hours through the relatively slow elimination through the
kidneys