Chemical Buffer Systems

Chemical buffers such as bicarbonate and ammonia help keep blood pH in the narrow range compatible with life.

Learning Objective

Distinguish between buffer solutions, ventilation, and renal function as buffer systems to control acid-base balance

Key Points

The body's acid–base balance is normally tightly regulated, keeping the arterial blood pH between 7.38 and 7.42. Buffer solutions keep pH constant in a wide variety of chemical actions.
A buffer solution is a mixture of a weak acid and it conjugate base or a weak base and its conjugate acid.
The bicarbonate buffering system is key to maintaining optimal pH levels by regulating carbon dioxide concentration that in turn shifts acid-base imbalance.
Renal physiology has several powerful mechanisms to control pH by the excretion of excess acid or base.

Terms

pH
In chemistry, pH is a measure of the activity of the hydrogen ion concentration.
bicarbonate
Bicarbonate is alkaline, and a vital component of the pH buffering system of the human body (maintaining acid-base homeostasis).
buffer
A solution used to stabilize the pH (acidity) of a liquid.

Example

Anything that adversely affects an individual's bloodstream will have a negative impact on that individual's health and well being since the blood acts as a chemical buffer solution to keep all the body's cells and tissues properly balanced.

Full Text

Acid–base homeostasis is the part of human homeostasis concerning the proper balance between acids and bases, also called body pH. The body is very sensitive to its pH level, so strong mechanisms exist to maintain it. Outside the acceptable range of pH, proteins are denatured and digested, enzymes lose their ability to function, and death may occur.

A buffer solution is an aqueous solution consisting of a mixture of a weak acid and its conjugate base or a weak base and its conjugate acid. Its pH changes very little when a small amount of strong acid or base is added to it. Buffer solutions are used as a means of keeping pH at a nearly constant value in a wide variety of chemical applications.

Many life forms thrive only in a relatively small pH range so they utilize a buffer solution to maintain a constant pH. One example of a buffer solution found in nature is blood. The body's acid–base balance is normally tightly regulated, keeping the arterial blood pH between 7.38 and 7.42. Several buffering agents that reversibly bind hydrogen ions and impede any change in pH exist. Extracellular buffers include bicarbonate and ammonia, whereas proteins and phosphate act as intracellular buffers. The bicarbonate buffering system is especially key, as carbon dioxide (CO₂) can be shifted through carbonic acid (H₂CO₃) to hydrogen ions and bicarbonate (HCO^3-):

${ H }_{ 2 }O+{ CO }_{ 2 }\leftrightarrow H_{ 2 }{ CO }_{ 3 }\leftrightarrow { H }^{ + }+{ CO }_{ 3 }^{ - }$

Acid–base imbalances that overcome the buffer system can be compensated in the short term by changing the rate of ventilation. This alters the concentration of carbon dioxide in the blood, shifting the above reaction according to Le Chatelier's principle, which in turn alters the pH.

The kidneys are slower to compensate, but renal physiology has several powerful mechanisms to control pH by the excretion of excess acid or base. In response to acidosis, tubular cells reabsorb more bicarbonate from the tubular fluid, collecting duct cells secrete more hydrogen and generate more bicarbonate, and ammoniagenesis leads to increased formation of the NH₃ buffer. In responses to alkalosis, the kidneys may excrete more bicarbonate by decreasing hydrogen ion secretion from the tubular epithelial cells, and lowering rates of glutamine metabolism and ammonium excretion.

pH Range

Buffering agents keep blood pH between 7.38 and 7.42.

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