Buffers - AP Chemistry
Card 0 of 168
You are in chemistry lab performing a titration. You were given 15 mL of an aqueous solution with an unknown concentration of acetic acid, 
to solve through titration with concentrated sodium hydroxide, 
. You know that the pKa of acetic acid is 4.75 and that your titrant is 0.1 M sodium hydroxide, 
.
The endpoint was determined at 10 mL of sodium hydroxide, 
. What is the pH after 5 mL of 
was added?
You are in chemistry lab performing a titration. You were given 15 mL of an aqueous solution with an unknown concentration of acetic acid,
The endpoint was determined at 10 mL of sodium hydroxide,
At the half end point, the 
. This can be determined by the Henderson-Hasselbalch equation if it is not clear.
Since the endpoint of the titration is that there are 10 mL of 0.1 M NaOH added, that means that there are 0.001 moles of acetic acid.
When 5 mL of NaOH is added, there are 0.0005 moles of acetic acid and 0.0005 moles of acetate formed.
Therefore pH= pKa
At the half end point, the
Since the endpoint of the titration is that there are 10 mL of 0.1 M NaOH added, that means that there are 0.001 moles of acetic acid.
When 5 mL of NaOH is added, there are 0.0005 moles of acetic acid and 0.0005 moles of acetate formed.
Therefore pH= pKa
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Determine which of these solution combinations form a buffer.
Determine which of these solution combinations form a buffer.
First to go through why the other ones are wrong:
Strong base + strong acid neutralizes and does not form a buffer solution
Strong base + weak base does not form a buffer - would need an acid
Strong base + weak acid = all weak acid converted to conjugate base
The correct answer is:
Strong base + weak acid = half converted to conjugate base with half leftover as weak acid, with all the components for a buffer
First to go through why the other ones are wrong:
Strong base + strong acid neutralizes and does not form a buffer solution
Strong base + weak base does not form a buffer - would need an acid
Strong base + weak acid = all weak acid converted to conjugate base
The correct answer is:
Strong base + weak acid = half converted to conjugate base with half leftover as weak acid, with all the components for a buffer
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Determine which combination of solutions would create a buffer solution.
Determine which combination of solutions would create a buffer solution.
For all the other options there is no ammonium leftover with which to serve as the weak acid in the buffer system, the ammonium is all used up and converted to ammonia. However in the correct answer choice, there is enough ammonium leftover after the reaction with the sodium hydroxide.
For all the other options there is no ammonium leftover with which to serve as the weak acid in the buffer system, the ammonium is all used up and converted to ammonia. However in the correct answer choice, there is enough ammonium leftover after the reaction with the sodium hydroxide.
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Which combination(s) would create a buffer solution?
I. Weak acid
II. Weak acid's conjugate base
III. Strong acid
IV. Strong base
V. Weak base
VI. Weak base's conjugate acid
Which combination(s) would create a buffer solution?
I. Weak acid
II. Weak acid's conjugate base
III. Strong acid
IV. Strong base
V. Weak base
VI. Weak base's conjugate acid
A buffer solution is formed from the equilibrium of a weak acid and its conjugate base, or from a weak base and its conjugate acid. It's ability to "buffer" the pH or keep it from changing in large amounts in from the switching between these two forms weak and its conjugate.
A buffer solution is formed from the equilibrium of a weak acid and its conjugate base, or from a weak base and its conjugate acid. It's ability to "buffer" the pH or keep it from changing in large amounts in from the switching between these two forms weak and its conjugate.
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Determine the pH of an aqueous solution of 0.01 M acetic acid, 
. The pKa of acetic acid is 4.75.
Determine the pH of an aqueous solution of 0.01 M acetic acid,
Since acetic acid is a weak acid, it has a Ka that is rather small, we have to do a RICE table to determine the equilibrium amount of hydronium, H3O+ to then determine the pH.
R 
I 0.1 M - 0 0
C -x +x +x
E 0.1 -x x x
So first we need to change our pKa to a Ka
where 
therefore
= 
= 
If we assume that x is very small compared to 0.1...
Where 
(note: when solving using the quadratic we come up with the same answer)
So if 
Since acetic acid is a weak acid, it has a Ka that is rather small, we have to do a RICE table to determine the equilibrium amount of hydronium, H3O+ to then determine the pH.
R
I 0.1 M - 0 0
C -x +x +x
E 0.1 -x x x
So first we need to change our pKa to a Ka
where
If we assume that x is very small compared to 0.1...
Where
(note: when solving using the quadratic we come up with the same answer)
So if
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Determine which solution(s) will yield a buffer solution.
I. 10 mL of 0.5 M HCl + 20 mL of 0.5 M acetate
II. 10 mL of 0.5 M HCl + 10 mL of 0.5 M acetate
III. 10 mL of 0.5 M HCl + 10 mL of 1.0 M acetate
IV. 10 mL of 0.5 M HCl + 10 mL of 1.5 M acetate
Determine which solution(s) will yield a buffer solution.
I. 10 mL of 0.5 M HCl + 20 mL of 0.5 M acetate
II. 10 mL of 0.5 M HCl + 10 mL of 0.5 M acetate
III. 10 mL of 0.5 M HCl + 10 mL of 1.0 M acetate
IV. 10 mL of 0.5 M HCl + 10 mL of 1.5 M acetate
These answers are correct because the two components needed to create a buffer solution are a weak acid and its conjugate base, or a weak base and its conjugate acid. In these cases, the first reaction to occur upon addition of the strong acid is the formation of the conjugate acid, acetic acid.
If the amount of initial 
is greater than HCl, then we will have some 
left over to act as a buffer with the created conjugate acid. This can be through a greater volume, or through a higher concentration as shown in the correct answers.
These answers are correct because the two components needed to create a buffer solution are a weak acid and its conjugate base, or a weak base and its conjugate acid. In these cases, the first reaction to occur upon addition of the strong acid is the formation of the conjugate acid, acetic acid.
If the amount of initial
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Which of the following acid and base pairs are capable of acting as a buffer?
Which of the following acid and base pairs are capable of acting as a buffer?
In this question, we're presented with a variety of acid/base pairs and we're asked to identify which one could act as a buffer.
Remember that a buffer is a pair of acid and its conjugate base that acts to resist substantial changes in pH. In order for a buffer to work, the acid base pair needs to exist in equilibrium. This way, when the pH of the solution changes, the equilibrium of the acid/base reaction will shift, such that the pH will not change drastically.
To have an acid/base pair in equilibrium, we'll need to look for a pair that contains a weak acid. Acids like 
and 
are so strong that they will dissociate completely. Of the answer choices shown, only the carbonic acid/bicarbonate system (
and 
) exists in equilibrium. Thus, this is the correct answer.
In this question, we're presented with a variety of acid/base pairs and we're asked to identify which one could act as a buffer.
Remember that a buffer is a pair of acid and its conjugate base that acts to resist substantial changes in pH. In order for a buffer to work, the acid base pair needs to exist in equilibrium. This way, when the pH of the solution changes, the equilibrium of the acid/base reaction will shift, such that the pH will not change drastically.
To have an acid/base pair in equilibrium, we'll need to look for a pair that contains a weak acid. Acids like
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Which of the following would best buffer a solution from a pH of 4 to 6?
Which of the following would best buffer a solution from a pH of 4 to 6?
A weak acid/base best buffers about 1 pH point above and below its pKa. The pKA closest to the middle of 4 and 6 (so want as close to 5) is acetic acid at 4.7.
A weak acid/base best buffers about 1 pH point above and below its pKa. The pKA closest to the middle of 4 and 6 (so want as close to 5) is acetic acid at 4.7.
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A buffer using acetic acid (pKa=4.76) is titrated with NaOH. What is the pH at half the equivalence point?
A buffer using acetic acid (pKa=4.76) is titrated with NaOH. What is the pH at half the equivalence point?
The pH at half the equivalence point is equal to the pKa of the acid.
The pH at half the equivalence point is equal to the pKa of the acid.
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Which of the following solutions has the greatest buffering capacity?
Which of the following solutions has the greatest buffering capacity?
Nitric Acid is a strong acid and can't buffer. Rubidium Hydroxide is a strong base and thus can't buffer. Of the remaining, both are weak acids, but the one with a greater concentration has a greater buffering capacity.
Nitric Acid is a strong acid and can't buffer. Rubidium Hydroxide is a strong base and thus can't buffer. Of the remaining, both are weak acids, but the one with a greater concentration has a greater buffering capacity.
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To create a buffer solution, you can use a weak acid and .
To create a buffer solution, you can use a weak acid and .
The definition of a buffer solution is that it contains a weak acid and its conjugate base, or a weak base and its conjugate acid. Since we are starting with a weak acid in this case, we need its conjugate base.
The definition of a buffer solution is that it contains a weak acid and its conjugate base, or a weak base and its conjugate acid. Since we are starting with a weak acid in this case, we need its conjugate base.
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Which of the following can be used in a buffer solution?
Which of the following can be used in a buffer solution?
For a buffer solution, you need a weak acid and its conjugate base, or a weak base and its conjugate acid. HCO3 from the NaHCO3 and CO3– from K2CO3 are this pair.
For a buffer solution, you need a weak acid and its conjugate base, or a weak base and its conjugate acid. HCO3 from the NaHCO3 and CO3– from K2CO3 are this pair.
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Which of the following will increase the pH of an 
buffer solution?
I. Removing carbonic acid
II. Adding sodium bicarbonate
Which of the following will increase the pH of an
I. Removing carbonic acid
II. Adding sodium bicarbonate
To answer this question we need to look at the reaction below:
An increase in the pH will result in a decrease in the concentration of hydrogen ions (
). Using Le Chatelier’s principle we can find out which answer choices will decrease 
.
Removing carbonic acid will decrease the concentration of 
. To maintain equilibrium, the reaction will shift to the left and make more reactants from products; therefore, there will be a decrease in the 
and an increase in pH.
Recall that salts like sodium bicarbonate, or 
, will dissociate in water and form ions. Sodium bicarbonate will form sodium (
) and bicarbonate (
) ions. This side reaction will result in an increase in the bicarbonate ion concentration. Le Chatelier’s principle will shift the equilibrium of the given reaction to the left and, therefore, decrease the 
. Adding sodium bicarbonate will increase the pH.
To answer this question we need to look at the reaction below:
An increase in the pH will result in a decrease in the concentration of hydrogen ions (
Removing carbonic acid will decrease the concentration of
Recall that salts like sodium bicarbonate, or
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Which of the following combinations cannot be used to produce a buffer solution?
Which of the following combinations cannot be used to produce a buffer solution?
Buffer solutions can be made via two methods. The first method involves adding equal amounts of a weak acid and a salt of its weak conjugate base (or vice versa). The second methods involves adding a weak acid and a half equivalent of a strong base (or vice versa).
is a weak acid and 
is a salt of its weak conjugate base; therefore, this can form a buffer.
is a weak base and 
is a salt of its weak conjugate acid; this can also form a buffer. Note that this is the converse of the first method (weak base with salt of weak acid), but it can still form a buffer solution.
is a strong acid and 
is a weak base; therefore, adding 
and a half equivalent of 
will create a buffer solution. This is the converse of the second method (adding a weak base to a half equivalent of strong acid).
and 
are both strong reagents (acid and base, respectively); therefore, they cannot form a buffer solution.
Buffer solutions can be made via two methods. The first method involves adding equal amounts of a weak acid and a salt of its weak conjugate base (or vice versa). The second methods involves adding a weak acid and a half equivalent of a strong base (or vice versa).
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Blood is a physiological buffer. The carbonic acid/bicarbonate system maintains blood’s pH at around 7.35. Carbon dioxide in blood undergoes a complex equilibrium reaction as follows:
Alterations to carbon dioxide levels can change the blood pH.
A patient has abnormally low levels of carbon dioxide in the blood. What can you conclude about this patient?
Blood is a physiological buffer. The carbonic acid/bicarbonate system maintains blood’s pH at around 7.35. Carbon dioxide in blood undergoes a complex equilibrium reaction as follows:
Alterations to carbon dioxide levels can change the blood pH.
A patient has abnormally low levels of carbon dioxide in the blood. What can you conclude about this patient?
The question states that the patient has low levels of carbon dioxide. If we look at the given reaction, we will notice that the reaction will compensate for this by shifting the reaction equilibrium to the left. This phenomenon is called Le Chatelier’s principle and occurs to maintain the equilibrium of the reaction; therefore, the reaction will create more carbon dioxide by utilizing bicarbonate and hydrogen ions in the blood. A decrease in hydrogen ion concentration in blood will increase the pH and cause alkalosis (basicity in the blood). Since carbon dioxide is the cause of alkalosis, this patient will experience respiratory alkalosis. If he experienced alkalosis due to a change in bicarbonate ion concentration, the patient will have metabolic alkalosis.
The ratio of carbonic acid to bicarbonate will stay the same because both will be used in equal amounts (1:1 ratio) to produce carbon dioxide. Increasing respiratory rate, or hyperventilation, will result in an increase in the amount of carbon dioxide expelled by the patient; this will decrease the carbon dioxide concentration in the blood and will worsen the respiratory alkalosis. Recall that we are utilizing the bicarbonate ion (in conjunction with hydrogen ions) to create carbonic acid. The carbonic acid will be further broken down to replenish the carbon dioxide. A decrease in the bicarbonate concentration will slow down this process.
The question states that the patient has low levels of carbon dioxide. If we look at the given reaction, we will notice that the reaction will compensate for this by shifting the reaction equilibrium to the left. This phenomenon is called Le Chatelier’s principle and occurs to maintain the equilibrium of the reaction; therefore, the reaction will create more carbon dioxide by utilizing bicarbonate and hydrogen ions in the blood. A decrease in hydrogen ion concentration in blood will increase the pH and cause alkalosis (basicity in the blood). Since carbon dioxide is the cause of alkalosis, this patient will experience respiratory alkalosis. If he experienced alkalosis due to a change in bicarbonate ion concentration, the patient will have metabolic alkalosis.
The ratio of carbonic acid to bicarbonate will stay the same because both will be used in equal amounts (1:1 ratio) to produce carbon dioxide. Increasing respiratory rate, or hyperventilation, will result in an increase in the amount of carbon dioxide expelled by the patient; this will decrease the carbon dioxide concentration in the blood and will worsen the respiratory alkalosis. Recall that we are utilizing the bicarbonate ion (in conjunction with hydrogen ions) to create carbonic acid. The carbonic acid will be further broken down to replenish the carbon dioxide. A decrease in the bicarbonate concentration will slow down this process.
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A researcher is trying to make a buffer solution from a weak acid and its weak conjugate base. The pKa of the acid is 5.9 and the desired pH of the buffer solution is 3.5. Which of the following is the best way to make this buffer solution?
A researcher is trying to make a buffer solution from a weak acid and its weak conjugate base. The pKa of the acid is 5.9 and the desired pH of the buffer solution is 3.5. Which of the following is the best way to make this buffer solution?
One way to make a buffer is by adding equal amounts of a weak acid to its weak conjugate base. For example, you can add 1M acetic acid to 1M acetate to create a buffer solution (note that both acetic acid and its conjugate base (acetate) are weak). However, when using this method you have to remember that the desired pH of the buffer solution has to equal the pKa of the weak acid. The question states that the pKa of the acid is 5.9 and the desired pH of the buffer is 3.5; therefore, it is not possible to make the buffer with the given acid. The researcher would have to find another acid that has a pKa near 3.5.
One way to make a buffer is by adding equal amounts of a weak acid to its weak conjugate base. For example, you can add 1M acetic acid to 1M acetate to create a buffer solution (note that both acetic acid and its conjugate base (acetate) are weak). However, when using this method you have to remember that the desired pH of the buffer solution has to equal the pKa of the weak acid. The question states that the pKa of the acid is 5.9 and the desired pH of the buffer is 3.5; therefore, it is not possible to make the buffer with the given acid. The researcher would have to find another acid that has a pKa near 3.5.
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There are many solution systems which can only function as desired when the pH of that solution stays within a narrow range. Maintaining a stable pH in an unstable environment is most often achieved by the use of a buffer system, which is composed of a conjugate acid-base pair. One physiologically important buffer system is the bicarbonate buffer system that resists changes in blood pH.
The acid dissociation constant of carbonic acid 
The normal blood pH is tightly regulated between 7.35 and 7.45
When blood pH falls below 7.35 a person is said to have acidosis. Depending upon how far the pH drops, this condition could lead to nervous system impairment, coma, and death.
What is the ratio of bicarbonate ion concentration to carbonic acid concentration at which an individual will be at the threshold of experiencing acidosis 
There are many solution systems which can only function as desired when the pH of that solution stays within a narrow range. Maintaining a stable pH in an unstable environment is most often achieved by the use of a buffer system, which is composed of a conjugate acid-base pair. One physiologically important buffer system is the bicarbonate buffer system that resists changes in blood pH.
The acid dissociation constant of carbonic acid
The normal blood pH is tightly regulated between 7.35 and 7.45
When blood pH falls below 7.35 a person is said to have acidosis. Depending upon how far the pH drops, this condition could lead to nervous system impairment, coma, and death.
What is the ratio of bicarbonate ion concentration to carbonic acid concentration at which an individual will be at the threshold of experiencing acidosis
When performing acid/base calculations, the Henderson-Hasselbalch equation is useful:
To use this equation we need to convert the 
to the 
, and can use the following definition to do so:
Since an individual will begin experiencing acidosis when the blood falls below 7.35 we can use 7.35 as the pH in the Henderson-Hasselbalch equation. Adding in the 
of 6.10 calculated from the 
gives the equation below:
Plug in known values and solve:
Correct Answer:
The answer is unitless because 
all units cancel out.
When performing acid/base calculations, the Henderson-Hasselbalch equation is useful:
To use this equation we need to convert the
Since an individual will begin experiencing acidosis when the blood falls below 7.35 we can use 7.35 as the pH in the Henderson-Hasselbalch equation. Adding in the
Plug in known values and solve:
Correct Answer:
The answer is unitless because
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Which of the following would best buffer a solution from a pH of 4 to 6?
Which of the following would best buffer a solution from a pH of 4 to 6?
A weak acid/base best buffers about 1 pH point above and below its pKa. The pKA closest to the middle of 4 and 6 (so want as close to 5) is acetic acid at 4.7.
A weak acid/base best buffers about 1 pH point above and below its pKa. The pKA closest to the middle of 4 and 6 (so want as close to 5) is acetic acid at 4.7.
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A buffer using acetic acid (pKa=4.76) is titrated with NaOH. What is the pH at half the equivalence point?
A buffer using acetic acid (pKa=4.76) is titrated with NaOH. What is the pH at half the equivalence point?
The pH at half the equivalence point is equal to the pKa of the acid.
The pH at half the equivalence point is equal to the pKa of the acid.
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Which of the following solutions has the greatest buffering capacity?
Which of the following solutions has the greatest buffering capacity?
Nitric Acid is a strong acid and can't buffer. Rubidium Hydroxide is a strong base and thus can't buffer. Of the remaining, both are weak acids, but the one with a greater concentration has a greater buffering capacity.
Nitric Acid is a strong acid and can't buffer. Rubidium Hydroxide is a strong base and thus can't buffer. Of the remaining, both are weak acids, but the one with a greater concentration has a greater buffering capacity.
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