The steps involved in supplying bicarbonate ions to the system are seen in Figure It is also possible that salts in the filtrate, such as sulfates, phosphates, or ammonia, will capture hydrogen ions. If this occurs, the hydrogen ions will not be available to combine with bicarbonate ions and produce CO 2. In such cases, bicarbonate ions are not conserved from the filtrate to the blood, which will also contribute to a pH imbalance and acidosis.
The hydrogen ions also compete with potassium to exchange with sodium in the renal tubules. If more potassium is present than normal, potassium, rather than the hydrogen ions, will be exchanged, and increased potassium enters the filtrate. When this occurs, fewer hydrogen ions in the filtrate participate in the conversion of bicarbonate into CO 2 and less bicarbonate is conserved.
If there is less potassium, more hydrogen ions enter the filtrate to be exchanged with sodium and more bicarbonate is conserved. Chloride ions are important in neutralizing positive ion charges in the body.
If chloride is lost, the body uses bicarbonate ions in place of the lost chloride ions. Thus, lost chloride results in an increased reabsorption of bicarbonate by the renal system.
Acid-Base Balance: KetoacidosisDiabetic acidosis, or ketoacidosis, occurs most frequently in people with poorly controlled diabetes mellitus. When certain tissues in the body cannot get adequate amounts of glucose, they depend on the breakdown of fatty acids for energy. When acetyl groups break off the fatty acid chains, the acetyl groups then non-enzymatically combine to form ketone bodies, acetoacetic acid, beta-hydroxybutyric acid, and acetone, all of which increase the acidity of the blood.
Ketoacidosis can be severe and, if not detected and treated properly, can lead to diabetic coma, which can be fatal. A common early symptom of ketoacidosis is deep, rapid breathing as the body attempts to drive off CO 2 and compensate for the acidosis. Another common symptom is fruity-smelling breath, due to the exhalation of acetone. Other symptoms include dry skin and mouth, a flushed face, nausea, vomiting, and stomach pain.
Treatment for diabetic coma is ingestion or injection of sugar; its prevention is the proper daily administration of insulin.
A person who is diabetic and uses insulin can initiate ketoacidosis if a dose of insulin is missed. Among people with type 2 diabetes, those of Hispanic and African-American descent are more likely to go into ketoacidosis than those of other ethnic backgrounds, although the reason for this is unknown. A variety of buffering systems exist in the body that helps maintain the pH of the blood and other fluids within a narrow range—between pH 7. A buffer is a substance that prevents a radical change in fluid pH by absorbing excess hydrogen or hydroxyl ions.
Several substances serve as buffers in the body, including cell and plasma proteins, hemoglobin, phosphates, bicarbonate ions, and carbonic acid.
The bicarbonate buffer is the primary buffering system of the IF surrounding the cells in tissues throughout the body. The respiratory and renal systems also play major roles in acid-base homeostasis by removing CO 2 and hydrogen ions, respectively, from the body. Skip to content Learning Objectives By the end of this section, you will be able to: Identify the most powerful buffer system in the body Identify the most rapid buffer system in the body Describe the protein buffer systems. Explain the way in which the respiratory system affects blood pH Describe how the kidney affects acid-base balance.
Chapter Review A variety of buffering systems exist in the body that helps maintain the pH of the blood and other fluids within a narrow range—between pH 7.
Review Questions. Critical Thinking Questions 1. Describe the conservation of bicarbonate ions in the renal system.
Describe the control of blood carbonic acid levels through the respiratory system. Glossary hypercapnia abnormally elevated blood levels of CO 2 hypocapnia abnormally low blood levels of CO 2. We can derive a convenient equation for calculating the pH of a buffer solution. Taking the logarithm of each side of Equation 5 gives:. Note that [HA] and [A - ] are the concentrations of the weak acid and its conjugate base at equilibrium.
Since the dissociation of a weak acid is typically very small, we can assume that [HA] and [A - ] are similar to the initial concentrations, [HA] i and [A - ] i.
Equation 8 is known as the Henderson-Hasselbalch Equation. This equation shows that the pH of a buffer solution is very close to the pK a value of the weak acid making the buffer. This is because the logarithm term will be small unless the concentrations of A - and HA differ by several orders of magnitude. For a basic buffer consisting of a weak base and its conjugate acid, one can begin with Eq.
By far the most important buffer for maintaining acid-base balance in the blood is the carbonic acid-bicarbonate buffer. The dissolved carbon dioxide and bicarbonate ion are at equilibrium Eq. The equilibrium on the right is an acid-base reaction where carbonic acid is the acid and water is the base. The equilibrium on the left is the association of the dissolved carbon dioxide with a water molecule to form carbonic acid. This equilibrium favors the CO 2 side; hence, the concentration of H 2 CO 3 in solution is very small.
Since both reactions are at equilibrium, we can simplify Eq. Remember, we are doing all these simplifications to obtain an equation that tells us how the concentrations of carbon dioxide and bicarbonate affect the pH of our blood.
Starting from Equation 12 , and following the steps of how Henderson-Hasselbalch Equation is derived from Eq. At normal body temperature, the value of pK is 6. Notice that Equation 13 is in a similar form to the Henderson-Hasselbalch Equation. However, Equation 13 does not meet the strict definition of a Henderson-Hasselbalch equation because this equation takes into account a non-acid-base reaction i.
Nonetheless, the relationship shown in Equation 11 is frequently referred to as the Henderson-Hasselbalch Equation for the buffer in physiological applications.
Equation 13 shows that the pH of the buffered solution the blood is dependent only on the ratio of the amount of CO 2 to the amount of HCO 3 - at a given temperature, so that pK remains constant. The buffering capacity of a buffer is highest when the pK a value of the buffer is closest to the desired pH value.
This can be explained by the Henderson-Hasselbach equation. According to Eq. The slope of the curve is flattest thus the change in pH is smallest where the pH is equal to the pK value of 6. Here, the buffering capacity is greatest because a shift in the relative concentrations of bicarbonate and carbon dioxide produces only a small change in the pH of the solution. However, at pH values higher than 7. In this case, a shift in the relative concentrations of bicarbonate and carbon dioxide produces a large change in the pH of the solution.
In this plot, the vertical axis shows the pH of the buffered solution in this case, the blood. The horizontal axis shows the composition of the buffer. On the left-hand side of the plot, most of the buffer is in the form of dissolved carbon dioxide, and on the right-hand side of the plot, most of the buffer is in the form of bicarbonate ion.
Conversely, as base is added, the pH increases and the buffer shifts toward greater HCO 3 - concentration Equation Hence, the physiological blood pH of 7.
The lungs remove excess CO 2 from the blood helping to raise the pH as equilibria in Eq. When the pH of the body is excessively high a condition known as alkalosis , the kidneys remove bicarbonate ion HCO 3 - from the blood helping to lower the pH as equilibria in Eq. Pathway Medicine. Search form Search.
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