Buffer systems maintian a constant pH in blood
The body maintains the pH of blood at around 7.4. If the pH level
changes just a few tenths of a pH unit, serious health consequences
can result. A decrease in blood ph is called
acidosis, an increase is called
Three different buffer systems exist in blood, the bicarbonate buffer
and the phosphate buffer are composed of "simple" chemicals. In
addition the carbonyl groups (-COOH) and the amide group (-NH2)
present on proteins allow some of these to act as buffers.
The bicarbonate buffer and the phospate buffer can be described by the
- What is the optimal pH for the bicarbonate buffer?
- What is the optimal pH for the phosphate buffer?
- In each buffer, which species react with added acid?
- In each buffer, which species react with added base?
The pH for the bicarbonate buffer seems to be outside of its
Buffer capacity is usually defined as +/- 1 pH unit of the pKa.
Notice that the pH of blood is is one unit away from the pKa of
carbonic acid. Calculate the ratio of bicarbonate to carbonic
acid implied by this. [Answer]
The ratio of bicarbonate to carbonic acid seems to be quite large (and
in general this system would not be considered ideal for maintaining a
pH of 7.4). However, physiologic conditions make this buffer ideal
In addition, the phosphate buffer as well as the buffering ability of
proteins in plasma are also available to maintain blood pH.
- excess acid is produced by the body as a byproduct of exercise
(lactic acid) making the higher concentration of the conjugate base
(bicarbonate) an advantage
- the body has the ability to obtain more carbonic acid by
reabsorbing carbon dioxide from the lungs. Recall that carbonic acid
is an aqueous solution of carbon dioxide.
Return to Chem 104 home page