Practice: Imaging tissue structures using muon tomography. Practice: The effect of contrast agents on MRI signals. Practice: Analyzing electrocardiogram voltage signals. Practice: The mechanics of standing balance. Practice: What conditions influence lyophilisation? Practice: Your heart does work: A relationship of pressure and volume. Practice: Heat engines and work. Practice: Tension in the muscles. Practice: Patients in a wheelchair. Practice: Biological adaptations in response to physical constraints: the case of Atlantic salmon.
Practice: Purification of caffeine. Practice: Ideal gases in medicine. Practice: What causes red blood cells to sickle? Practice: Qualitative analysis of contaminated water supply. Practice: Synthesis of anti-tumor drug Combretastatin and its derivatives. Practice: Testing new suture material. Practice: Swine flu in Finland.
Practice: Ancient Greek catapults and medical packs. Practice: Forces on a kidney stone. Practice: Knee injuries in athletes. Practice: The speed of a neural impulse. Practice: Blood flow in the arteries. Practice: Concussions in professional athletes. Practice: The world's fastest mammal. Practice: Force of tension - Passage 1. Practice: Frictional forces on mobility walkers. Practice: An elevator in a hospital. Practice: A ramp in an administrative office.
However, the relationship shown in Equation 11 is frequently referred to as the Henderson-Hasselbach equation for the buffer in physiological applications. In Equation 11, pK is equal to the negative log of the equilibrium constant, K, for the buffer Equation We may begin by defining the equilibrium constant, K 1 , for the left-hand reaction in Equation 10, using the Law of Mass Action:. K a see Equation 9, above is the equilibrium constant for the acid-base reaction that is the reverse of the left-hand reaction in Equation It follows that the formula for K a is.
The equilibrium constant, K 2 , for the right-hand reaction in Equation 10 is also defined by the Law of Mass Action:. Because the two equilibrium reactions in Equation 10 occur simultaneously, Equations 14 and 15 can be treated as two simultaneous equations. Solving for the equilibrium concentration of carbonic acid gives.
Rearranging Equation 16 allows us to solve for the equilibrium proton concentration in terms of the two equilibrium constants and the concentrations of the other species:. Because we are interested in the pH of the blood, we take the negative log of both sides of Equation Recalling the definitions of pH and pK Equations 2 and 12, above , Equation 18 can be rewritten using more conventional notation, to give the relation shown in Equation 11, which is reproduced below:. As shown in Equation 11, the pH of the buffered solution i.
This optimal buffering occurs when the pH is within approximately 1 pH unit from the pK value for the buffering system, i. However, the normal blood pH of 7. The lungs remove excess CO 2 from the blood helping to raise the pH via shifts in the equilibria in Equation 10 , and the kidneys remove excess HCO 3 - from the body helping to lower the pH. The lungs' removal of CO 2 from the blood is somewhat impeded during exercise when the heart rate is very rapid; the blood is pumped through the capillaries very quickly, and so there is little time in the lungs for carbon dioxide to be exchanged for oxygen.
The ways in which these three organs help to control the blood pH through the bicarbonate buffer system are highlighted in Figure 3, below.
This figure shows the major organs that help control the blood concentrations of CO 2 and HCO 3 - , and thus help control the pH of the blood. Removing CO 2 from the blood helps increase the pH. Removing HCO 3 - from the blood helps lower the pH. Why is the buffering capacity of the carbonic-acid-bicarbonate buffer highest when the pH is close to the pK value, but lower at normal blood pH?
The answer to this question lies in the shape of the titration curve for the buffer, which is shown in Figure 4, below. It is possible to plot a titration curve for this buffer system, just as you did for your solution in the acid-base-equilibria experiment.
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 carbonic acid or 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 This is the titration curve for the carbonic-acid-bicarbonate buffer. Note that the pH of the blood 7. Note: The percent buffer in the form of HCO 3 - is given by the formula:. The slope of the curve is flattest where the pH is equal to the pK value 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. Here, a shift in the relative concentrations of bicarbonate and carbon dioxide produces a large change in the pH of the solution. Hence, at the physiological blood pH of 7. Other buffers perform a more minor role than the carbonic-acid-bicarbonate buffer in regulating the pH of the blood.
As with the phosphate buffer, a weak acid or weak base captures the free ions, and a significant change in pH is prevented. Bicarbonate ions and carbonic acid are present in the blood in a ratio if the blood pH is within the normal range. With 20 times more bicarbonate than carbonic acid, this capture system is most efficient at buffering changes that would make the blood more acidic.
Carbonic acid levels in the blood are controlled by the expiration of CO 2 through the lungs. In red blood cells, carbonic anhydrase forces the dissociation of the acid, rendering the blood less acidic.
Because of this acid dissociation, CO 2 is exhaled see equations above. The level of bicarbonate in the blood is controlled through the renal system, where bicarbonate ions in the renal filtrate are conserved and passed back into the blood. However, the bicarbonate buffer is the primary buffering system of the IF surrounding the cells in tissues throughout the body. The respiratory system contributes to the balance of acids and bases in the body by regulating the blood levels of carbonic acid Figure CO 2 in the blood readily reacts with water to form carbonic acid, and the levels of CO 2 and carbonic acid in the blood are in equilibrium.
When the CO 2 level in the blood rises as it does when you hold your breath , the excess CO 2 reacts with water to form additional carbonic acid, lowering blood pH. The loss of CO 2 from the body reduces blood levels of carbonic acid and thereby adjusts the pH upward, toward normal levels.
As you might have surmised, this process also works in the opposite direction. Excessive deep and rapid breathing as in hyperventilation rids the blood of CO 2 and reduces the level of carbonic acid, making the blood too alkaline.
This brief alkalosis can be remedied by rebreathing air that has been exhaled into a paper bag. Rebreathing exhaled air will rapidly bring blood pH down toward normal. Minor adjustments in breathing are usually sufficient to adjust the pH of the blood by changing how much CO 2 is exhaled. This situation is common if you are exercising strenuously over a period of time.
To keep up the necessary energy production, you would produce excess CO 2 and lactic acid if exercising beyond your aerobic threshold. In order to balance the increased acid production, the respiration rate goes up to remove the CO 2. This helps to keep you from developing acidosis. The body regulates the respiratory rate by the use of chemoreceptors, which primarily use CO 2 as a signal. Peripheral blood sensors are found in the walls of the aorta and carotid arteries.
These sensors signal the brain to provide immediate adjustments to the respiratory rate if CO 2 levels rise or fall. Yet other sensors are found in the brain itself. Changes in vitamin D metabolism, phosphate metabolism and secondary hyperparathyroidism are more important than the acidosis in causing loss of bone mineral in uraemic patients.
The loss of bone mineral due to these other factors releases substantial amounts of buffer. They were hidden on buffers and so these hydrogen ions were hidden from view. Henderson-Hasselbalch Equation. The bicarbonate buffer system is an effective buffer system despite having a low pKa because the body also controls pCO 2.
The phosphate buffer system is NOT an important blood buffer as its concentration is too low The concentration of phosphate in the blood is so low that it is quantitatively unimportant. Note The pKa2 value is actually 7. Isohydric Principle All buffer systems which participate in defence of acid-base changes are in equilibrium with each other.
Buffering in different sites Respiratory disorders are predominantly buffered in the intracellular compartment. Deoxyhaemoglobin is a better buffer than oxyhaemoglobin Another factor which makes haemoglobin an important buffer is the phenomemon of isohydric exchange.
Because of this ease of movement, CO 2 is not important in causing differences in pH on the two sides of the cell membrane. How does bone act as a buffer? Release of calcium carbonate from bone is the most important buffering mechanism involved in chronic metabolic acidosis. Loss of bone crystal in uraemic acidosis is multifactorial and acidosis is only a minor factor BOTH the acidosis and the vitamin D3 changes are responsible for the osteomalacia that occurs with renal tubular acidosis.
References Worthley LI. Hydrogen ion metabolism. Anaesth Intensive Care Nov; 5 4 PubMed Pitts RF. Mechanisms for stabilizing the alkaline reserves of the body. Harvey Lect ; 48 Regional Refresher Courses in Anesthesiology.
Acidosis and bone. Miner Electrolyte Metab ; 20 The Major Body Buffer Systems. Buffer System. For metabolic acids.
0コメント