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. The pK for the phosphate buffer is 6. The phosphate buffer only plays a minor role in the blood, however, because H 3 PO 4 and H 2 PO 4 - are found in very low concentration in the blood.
Proper physiological functioning depends on a very tight balance between the concentrations of acids and bases in the blood. Acid-balance balance is measured using the pH scale, as shown in. A variety of buffering systems permits blood and other bodily fluids to maintain a narrow pH range, even in the face of perturbations.
A buffer is a chemical system that prevents a radical change in fluid pH by dampening the change in hydrogen ion concentrations in the case of excess acid or base. Most commonly, the substance that absorbs the ions is either a weak acid, which takes up hydroxyl ions, or a weak base, which takes up hydrogen ions.
The buffer systems in the human body are extremely efficient, and different systems work at different rates. It takes only seconds for the chemical buffers in the blood to make adjustments to pH.
The respiratory tract can adjust the blood pH upward in minutes by exhaling CO 2 from the body. The buffer systems functioning in blood plasma include plasma proteins, phosphate, and bicarbonate and carbonic acid buffers. The kidneys help control acid-base balance by excreting hydrogen ions and generating bicarbonate that helps maintain blood plasma pH within a normal range.
Protein buffer systems work predominantly inside cells. Nearly all proteins can function as buffers. Proteins are made up of amino acids, which contain positively charged amino groups and negatively charged carboxyl groups. The charged regions of these molecules can bind hydrogen and hydroxyl ions, and thus function as buffers. Buffering by proteins accounts for two-thirds of the buffering power of the blood and most of the buffering within cells. Hemoglobin is the principal protein inside of red blood cells and accounts for one-third of the mass of the cell.
During the conversion of CO 2 into bicarbonate, hydrogen ions liberated in the reaction are buffered by hemoglobin, which is reduced by the dissociation of oxygen. This buffering helps maintain normal pH. The process is reversed in the pulmonary capillaries to re-form CO 2 , which then can diffuse into the air sacs to be exhaled into the atmosphere. This process is discussed in detail in the chapter on the respiratory system.
Acids and bases are still present, but they hold onto the ions. The bicarbonate-carbonic acid buffer works in a fashion similar to phosphate buffers. The bicarbonate is regulated in the blood by sodium, as are the phosphate ions.
When carbonic acid comes into contact with a strong base, such as NaOH, bicarbonate and water are formed. 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. Photosynthesis, respiration and decomposition all contribute to pH fluctuations due to their influences on CO2 levels.
This influence is more measurable in bodies of water with high rates of respiration and decomposition. While carbon dioxide exists in water in a dissolved state like oxygen , it can also react with water to form carbonic acid:.
However, this equation can operate in both directions depending on the current pH level, working as its own buffering system. At a higher pH, this bicarbonate system will shift to the left, and CO3 2- will pick up a free hydrogen ion.
However, as CO2 levels increase around the world, the amount of dissolved CO2 also increases, and the equation will be carried out from left to right. This increases H2CO3, which decreases pH. The effect is becoming more evident in oceanic pH studies over time. The above equations also explain why rain has a pH of approximately 5. As raindrops fall through the air, they interact with carbon dioxide molecules in the atmosphere.
A pH level of 5. Natural, unpolluted rain or snow is expected to have pH levels near 5. Acid rain requires a pH below 5.
Carbonate materials and limestone are two elements that can buffer pH changes in water. When carbonate minerals are present in the soil, the buffering capacity alkalinity of water is increased, keeping the pH of water close to neutral even when acids or bases are added.
Additional carbonate materials beyond this can make neutral water slightly basic. As mentioned earlier, unpolluted rain is slightly acidic pH of 5. If rain falls on a poorly buffered water source, it can decrease the pH of nearby water through runoff. Anthropogenic causes of pH fluctuations are usually related to pollution. Acid rain is one of the best known examples of human influence on the pH of water. Any form of precipitation with a pH level less than 5.
This precipitation comes from the reaction of water with nitrogen oxides, sulfur oxides and other acidic compounds, lowering its already slightly acidic pH. These chemicals can come from agricultural runoff, wastewater discharge or industrial runoff. Wastewater discharge that contains detergents and soap-based products can cause a water source to become too basic.
Typical pH levels vary due to environmental influences, particularly alkalinity. The alkalinity of water varies due to the presence of dissolved salts and carbonates, as well as the mineral composition of the surrounding soil. The recommended pH range for most fish is between 6. Oceanic organisms like clownfish and coral require higher pH levels. Sensitive freshwater species such as salmon prefer pH levels between 7.
Natural precipitation, both rain and snow, has a pH near 5. This hormone increases gluconeogenesis and fat breakdown, as described above. Given that the pKa of pyruvic acid is 2. Thus the ratio of is equal to , or the ratio of is equal to. When this is true, the ratio of ionized to unionized form of the buffer is Thus the solution can best resist changes in pH, as hydrogen ions can be quenched or donated to solution to resist change.
Acetic acid would have almost the same buffering capacity since its pKa is almost as close to 4. Upon running lab tests, you determine that a patient has very low blood pH. Which of the following could have caused this low pH? Increased iron in blood. Increased red blood cells. Low blood pH suggests that the patient has high concentration of hydrogen ions.
To solve this question, we need to look at the following reaction, which represents the major blood buffer system:. One way the body controls the amount of hydrogen ions in the blood is by altering the amount of carbon dioxide.
Body controls carbon dioxide levels via breathing. Hyperventilation refers to increased breathing whereas hypoventilation refers to decreased breathing. During hyperventilation the person breathes out excess carbon dioxide decreasing the hydrogen ion concentration.
During hypoventilation, on the other hand, a person breathes slowly and retains carbon dioxide increasing the hydrogen ion concentration. The patient in this question has low blood pH high hydrogen ion concentration ; therefore, of the options, the patient must be hypoventilating.
What buffer system is most important for maintaining blood pH? A buffer system occurs when equal amounts of a weak acid and its conjugate base or vice versa is added together.
Maintaining blood pH is a very important aspect in maintaining homeostasis. The concentrations of the base and acid are altered accordingly to maintain a constant blood pH. Hydrochloric acid is found in the stomach to maintain an acidic pH, phosphoric acid does play a role in buffering the blood, but is not the major buffer. Acetic acid is found in vinegar and does not play a major role in regulating blood pH.
Which of the following molecule s will increase in response to high blood pH? Blood pH is maintained via the lungs and the kidneys. Lungs alter the amount of carbon dioxide expelled to maintain blood pH. Consider the reaction below. Carbon dioxide is decreased when pH is low high hydrogen ion concentration. Decreasing carbon dioxide will shift the reaction to the left and decrease the hydrogen ion concentration. Similarly, the body compensates for high blood pH by increasing carbon dioxide.
Kidneys alter blood pH by increasing or decreasing the excretion of bicarbonate ions. Using the reaction above, we can determine that increasing bicarbonate ion in blood will decrease hydrogen ion concentration whereas decreasing bicarbonate ion will increase hydrogen ion concentration.
To combat high blood pH low hydrogen ion concentration , the bicarbonate ion needs to increased in the blood. The kidneys do this by decreasing the excretion of the bicarbonate ions.
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