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Test Prep MCAT Test Exam
Medical College Admission Test: Verbal Reasoning, Biological Sciences, Physical Sciences, Writing Sample (Page 3 )

Updated On: 19-Jan-2026
Page 3 of 164
PDF with 811 Questions

It is critical for the human body blood to maintain its pH at approximately 7.4. Decreased or increased blood pH are called acidosis and alkalosis respectively; both are serious metabolic problems that can cause death. The table below lists the major buffers found in the blood and/or kidneys.
Table 1
Buffer
pKa of a typical conjugate acid:*

Histidine side chains

Organic phosphates
N-terminal amino groups

6.1
6.3
6.8
7.0
8.0
9.2
*For buffers in many of these categories, there is a range of actual pKa values.
The relationship between blood pH and the pKa of any buffer can be described by the Henderson-Hasselbalch equation:
pH = pKa + log([conjugate base] / [conjugate acid])
Equation 1
Bicarbonate, the most important buffer in the plasma, enters the blood in the form of carbon dioxide, a byproduct of metabolism, and leaves in two forms: exhaled CO2 and excreted bicarbonate. Blood pH can be adjusted rapidly by changes in the rate of CO2 exhalation. The reaction given below, which is catalyzed by carbonic anhydrase in the erythrocytes, describes how bicarbonate and CO2 interact in the blood.

Reaction 1
The equilibrium as shown in Reaction 1 is most likely to proceed through which of the following intermediates?

  1. H2CO3
  2. 2H+ and CO32-
  3. CO2 and H3O2
  4. CO2 and H2

Answer(s): A

Explanation:

Carbonic acid, H2CO3, is the intermediate formed when carbon dioxide and water combine, making choice A the correct response. You should be able to recognize that when the products of Reaction 1 combine, carbonic acid will result. Choice B is incorrect because it is just not a logical intermediate: water and carbon dioxide would not react to form H+ and CO32-. Choice D and choice C can be eliminated because they both contain carbon dioxide, which is one of the two reactants.



It is critical for the human body blood to maintain its pH at approximately 7.4. Decreased or increased blood pH are called acidosis and alkalosis respectively; both are serious metabolic problems that can cause death. The table below lists the major buffers found in the blood and/or kidneys.
Table 1
Buffer
pKa of a typical conjugate acid:*

Histidine side chains

Organic phosphates
N-terminal amino groups

6.1
6.3
6.8
7.0
8.0
9.2
*For buffers in many of these categories, there is a range of actual pKa values.
The relationship between blood pH and the pKa of any buffer can be described by the Henderson-Hasselbalch equation:
pH = pKa + log([conjugate base] / [conjugate acid])
Equation 1
Bicarbonate, the most important buffer in the plasma, enters the blood in the form of carbon dioxide, a byproduct of metabolism, and leaves in two forms: exhaled CO2 and excreted bicarbonate. Blood pH can be adjusted rapidly by changes in the rate of CO2 exhalation. The reaction given below, which is catalyzed by carbonic anhydrase in the erythrocytes, describes how bicarbonate and CO2 interact in the blood.

Reaction 1
What would be the order of conjugate acid strength in the following buffers?

  1. Histidine side chains = organic phosphates > NH4+
  2. NH4+ > organic phosphates > histidine side chains
  3. Histidine side chains > organic phosphates > NH4+
  4. NH4+ > organic phosphates = histidine side chains

Answer(s): C

Explanation:

You should know that the smaller the pKa the stronger the acid. Looking at Table 1, it can be seen that the histidine side chains buffer has a stronger conjugate acid than both the organic phosphates buffer and the ammonium buffer. Choice C is therefore the correct response. Choice A is wrong because the pKa of the histidine side chains buffer is not equal to the organic phosphates buffer. Choice B is wrong because it is the reverse of what it should be. Choice D is wrong because, again, the pKa of the histidine side chains buffer is not equal to the organic phosphates buffer.



It is critical for the human body blood to maintain its pH at approximately 7.4. Decreased or increased blood pH are called acidosis and alkalosis respectively; both are serious metabolic problems that can cause death. The table below lists the major buffers found in the blood and/or kidneys.

Table 1
Buffer
pKa of a typical conjugate acid:*

Histidine side chains

Organic phosphates
N-terminal amino groups

6.1
6.3
6.8
7.0
8.0
9.2
*For buffers in many of these categories, there is a range of actual pKa values.
The relationship between blood pH and the pKa of any buffer can be described by the Henderson-Hasselbalch equation:
pH = pKa + log([conjugate base] / [conjugate acid])
Equation 1
Bicarbonate, the most important buffer in the plasma, enters the blood in the form of carbon dioxide, a byproduct of metabolism, and leaves in two forms: exhaled CO2 and excreted bicarbonate. Blood pH can be adjusted rapidly by changes in the rate of CO2 exhalation. The reaction given below, which is catalyzed by carbonic anhydrase in the erythrocytes, describes how bicarbonate and CO2 interact in the blood.

Reaction 1
The following graph shows the titration of 0.01 M H3PO4 with 10 M NaOH. Within which region of the titration curve will the concentration of H2PO4- become equal to that of HPO42-?

  1. II
  2. III
  3. IV
  4. V

Answer(s): B

Explanation:

In Region I the following reaction takes place: H3PO4(aq) + OH-(aq) H2PO4- (aq) + H2O.
As OH- is added, H2PO4- is formed and a buffer is realized. So, in this region, there is a point where the concentration of H3PO4 will equal that of H2PO4-. In Region II there is a point where exactly enough base has been added to react with all of the H3PO4; this point is called an equivalence point. In Region III the following reaction takes place: H2PO4-(aq) + OH-(aq) HPO42-(aq) + H2O.
Again, this system constitutes a buffer, and there is a point in this region where ­ after enough base has been added ­ the concentration of H2PO4- will equal that of HPO42-. Choice B is therefore the correct response. In Region IV there is another equivalence point when exactly enough base has been added to react with all of the H2PO4-. In Region V the following reaction takes place: HPO42-(aq) + OH-(aq) PO43-(aq) + H2O.
This system is, of course, a buffer as well.



It is critical for the human body blood to maintain its pH at approximately 7.4. Decreased or increased blood pH are called acidosis and alkalosis respectively; both are serious metabolic problems that can cause death. The table below lists the major buffers found in the blood and/or kidneys.
Table 1
Buffer
pKa of a typical conjugate acid:*

Histidine side chains

Organic phosphates

N-terminal amino groups

6.1
6.3
6.8
7.0
8.0
9.2
*For buffers in many of these categories, there is a range of actual pKa values.
The relationship between blood pH and the pKa of any buffer can be described by the Henderson-Hasselbalch equation:
pH = pKa + log([conjugate base] / [conjugate acid])
Equation 1
Bicarbonate, the most important buffer in the plasma, enters the blood in the form of carbon dioxide, a byproduct of metabolism, and leaves in two forms: exhaled CO2 and excreted bicarbonate. Blood pH can be adjusted rapidly by changes in the rate of CO2 exhalation. The reaction given below, which is catalyzed by carbonic anhydrase in the erythrocytes, describes how bicarbonate and CO2 interact in the blood.

Reaction 1
How does the titration of a weak monoprotic acid with a strong base differ from the titration of a strong monoprotic acid with a strong base?

  1. The equivalence point will occur at a higher pH.
  2. The equivalence point will occur at a lower pH.
  3. The equivalence point will occur at the same pH.
  4. Whether the equivalence point is higher or lower depends on the particular monoprotic acids used.

Answer(s): A

Explanation:

When a weak acid is reacted with a strong base, the equivalence point will be in the basic region. Consider the titration of equimolar solutions of acetic acid and NaOH. Before the equivalence point, the following reaction takes place:
HC2H3O2(aq) + OH-(aq) H2O + C2H3O2-(aq)
At the equivalence point, only C2H3O2- exists. When C2H3O2- undergoes hydrolysis (i.e., reacts with water), hydroxide ions are formed according to the following equilibrium:

The numerical value of the equilibrium constant along with the initial concentration of acetate is all that is needed to determine the hydroxide ion concentration. When equimolar solutions of a strong acid and a strong base are titrated, the equivalence point will be neutral. It is neutral because neither of the ions present at the equivalence point can undergo hydrolysis. Choice A is therefore the correct response. Choice B would be correct if a weak base was titrated with a strong acid.



It is critical for the human body blood to maintain its pH at approximately 7.4. Decreased or increased blood pH are called acidosis and alkalosis respectively; both are serious metabolic problems that can cause death. The table below lists the major buffers found in the blood and/or kidneys.
Table 1
Buffer
pKa of a typical conjugate acid:*

Histidine side chains

Organic phosphates
N-terminal amino groups

6.1
6.3
6.8
7.0
8.0
9.2
*For buffers in many of these categories, there is a range of actual pKa values.
The relationship between blood pH and the pKa of any buffer can be described by the Henderson-Hasselbalch equation:
pH = pKa + log([conjugate base] / [conjugate acid])
Equation 1
Bicarbonate, the most important buffer in the plasma, enters the blood in the form of carbon dioxide, a byproduct of metabolism, and leaves in two forms: exhaled CO2 and excreted bicarbonate. Blood pH can be adjusted rapidly by changes in the rate of CO2 exhalation. The reaction given below, which is catalyzed by carbonic anhydrase in the erythrocytes, describes how bicarbonate and CO2 interact in the blood.

Reaction 1
What would be the nature of the compensatory change that would take place in the respiratory system response to acidosis caused by organic acids?

  1. Breathing rate would increase and total blood CO2/HCO3- concentration would increase
  2. Breathing rate would increase and total blood CO2/HCO3- concentration would decrease
  3. Breathing rate would decrease and total blood CO2/HCO3- concentration would increase
  4. Breathing rate would decrease and total blood CO2/HCO3- concentration would decrease

Answer(s): B

Explanation:

As the passage states, the immediate buffering effect of bicarbonate is controlled by changes in the breathing rate. When acidosis occurs, the concentration of H+ is too high. As discussed earlier, Le Châtelier's principle applies: In order to decrease the concentration of H+, the concentration of carbon dioxide must decrease. If the breathing rate increases, more carbon dioxide is exhaled and its concentration in the blood decreases. Since the effect of this rapid breathing is to remove carbon dioxide from the body, the ultimate effect is to decrease the total CO2/HCO3- concentration in the blood. Therefore, choice B is correct. Choice A is wrong because the ratio of CO2/HCO3- would decrease, not increase. Choice C and choice D are wrong because the breathing rate would increase, not decrease.



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