Lactic acid and other organic acids are produced by the cells of the body and secreted into the blood. Despite the release of H+ by these acids, the arterial blood pH normally does not decrease but remains remarkably constant at pH 7.40 ± 0.05. This constancy is achieved, in part, by the buffering action of bicarbonate shown in the preceding equation. Bicarbonate serves as the major buffer of the blood.
Certain conditions could cause an opposite change in pH. For example, excessive vomiting that results in loss of gastric acid could cause the concentration of free H+ in the blood to fall and the blood pH to rise. In this case, the reaction previously described could be reversed:
The dissociation of carbonic acid yields free H+, which helps to prevent an increase in pH. Bicarbonate ions and carbonic acid thus act as a buffer pair to prevent either decreases or increases in pH, respectively. This buffering action normally maintains the blood pH within the narrow range of 7.35 to 7.45.
If the arterial blood pH falls below 7.35, the condition is called acidosis. A blood pH of 7.20, for example, represents significant acidosis. Notice that acidotic blood need not be acidic. An increase in blood pH above 7.45, conversely, is known as alkalo-sis. Acidosis and alkalosis are normally prevented by the action of the bicarbonate/carbonic acid buffer pair and by the functions of the lungs and kidneys. Regulation of blood pH is discussed in more detail in chapters 13, 16, and 17.
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This ebook provides an introductory explanation of the workings of the human body, with an effort to draw connections between the body systems and explain their interdependencies. A framework for the book is homeostasis and how the body maintains balance within each system. This is intended as a first introduction to physiology for a college-level course.