Chemical buffers, such as bicarbonate and ammonia, help keep the blood’s pH in the narrow range that is compatible with life.
Distinguish between buffer solutions, ventilation, and renal function as buffer systems to control acid–base balance
- The body’s acid– base balance is tightly regulated to keep the arterial blood pH between 7.38 and 7.42. Buffer solutions keep the pH constant in a wide variety of chemical actions.
- A buffer solution is a mixture of a weak acid and its conjugate base, or a weak base and its conjugate acid.
- The bicarbonate buffering system maintains optimal pH levels and regulates the carbon dioxide concentration that, in turn, shifts any acid–base imbalance.
- Renal physiology controls pH levels through several powerful mechanisms that excrete excess acid or base.
- bicarbonate: An alkaline, vital component of the pH buffering system of the human body that maintains acid–base homeostasis.
- buffer: A solution used to stabilize the pH (acidity) of a liquid.
- pH: In chemistry, a measure of the activity of the hydrogen ion concentration.
Anything that adversely affects an individual’s bloodstream will have a negative impact on that individual’s health since the blood acts as a chemical buffer solution to keep all the body’s cells and tissues properly balanced.
Acid–base homeostasis concerns the proper balance between acids and bases; it is also called body pH. The body is very sensitive to its pH level, so strong mechanisms exist to maintain it. Outside an 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 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 phosphates 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 (HCO3−):
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 and shifts 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, the tubular cells reabsorb more bicarbonate from the tubular fluid, and the collecting duct cells secrete more hydrogen and generate more bicarbonate, and ammoniagenesis leads to an increase of the NH3 buffer.
In its responses to alkalosis, the kidneys may excrete more bicarbonate by decreasing hydrogen ion secretion from the tubular epithelial cells, and lower the rates of glutamine metabolism and ammonium excretion.
pH range: Buffering agents keep blood pH between 7.38 and 7.42.