Reactions - AP Chemistry
Card 0 of 1366
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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When balanced, what is the value of 
in the following chemical equation?
When balanced, what is the value of
Recall that a balanced chemical equation has the same number of each element on one side.
Start by counting the number of each element on each side.
There are the following numbers of moles of each reactant:
There are the following numbers of moles of each product:
Add coefficients in front of the molecular compounds in the equation until there are the same numbers of sulfur, oxygen, lithium, and selenium on each side.
The balanced chemical equation is the following:
Both products and reactants now have the following number of moles:
Since the coefficient in front of 
is 
, 
must equal to 
.
Recall that a balanced chemical equation has the same number of each element on one side.
Start by counting the number of each element on each side.
There are the following numbers of moles of each reactant:
There are the following numbers of moles of each product:
Add coefficients in front of the molecular compounds in the equation until there are the same numbers of sulfur, oxygen, lithium, and selenium on each side.
The balanced chemical equation is the following:
Both products and reactants now have the following number of moles:
Since the coefficient in front of
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Balance the following chemical equation:
Balance the following chemical equation:
A balanced chemical equation will have the same number of each atom on both sides of the reaction.
Start by balancing the number of nitrate.
Since we have 
nitrate on the left, we must also have 
nitrate on the right.
This equation now gives 
potassium on the left, and 
potassium on the right. Balance the potassium.
This equation is balanced because there are equal numbers of each atom on both sides.
A balanced chemical equation will have the same number of each atom on both sides of the reaction.
Start by balancing the number of nitrate.
Since we have
This equation now gives
This equation is balanced because there are equal numbers of each atom on both sides.
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According to Le Chatelier's principle, which of the following occurs when you compress a system containing at least one gas species?
According to Le Chatelier's principle, which of the following occurs when you compress a system containing at least one gas species?
According to Le Chatelier's principle, when you compress a system, its volume decreases, so partial pressure of the all the gases in the system increases. The system will act to try to decrease the pressure by decreasing the moles of gas.
According to Le Chatelier's principle, when you compress a system, its volume decreases, so partial pressure of the all the gases in the system increases. The system will act to try to decrease the pressure by decreasing the moles of gas.
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If heat is added to an endothermic reaction, in which direction will the equilibrium shift according to Le Chatelier's principle?
If heat is added to an endothermic reaction, in which direction will the equilibrium shift according to Le Chatelier's principle?
In an endothermic reaction, heat can be treated as a reactant. Thus, if you add more reactant (heat), the system will shift to get rid of the extra reactant and shift to the right to form more products.
In an endothermic reaction, heat can be treated as a reactant. Thus, if you add more reactant (heat), the system will shift to get rid of the extra reactant and shift to the right to form more products.
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If heat is added to an exothermic reaction, in which direction will the equilibrium shift according to Le Chatelier's principle?
If heat is added to an exothermic reaction, in which direction will the equilibrium shift according to Le Chatelier's principle?
In an exothermic reaction, heat can be treated as a product. Thus, if you add more product (heat), the reaction will shift to the left to form more reactants.
In an exothermic reaction, heat can be treated as a product. Thus, if you add more product (heat), the reaction will shift to the left to form more reactants.
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Figure 1: Ammonia gas formation and equilibrium
Experimental data shows that the reaction shifts to the left at very cold temperatures. Using this information, what type of reaction is shown in Figure 1?
Figure 1: Ammonia gas formation and equilibrium
Experimental data shows that the reaction shifts to the left at very cold temperatures. Using this information, what type of reaction is shown in Figure 1?
This is an application of Le Chatlier's Principle. When you take away heat from the reaction, the reaction shifts toward the left in order to compensate from the heat loss. The reaction may then be rewritten to include energy as a reactant.
Since the energy is on the reactant side, the reaction is endothermic.
This is an application of Le Chatlier's Principle. When you take away heat from the reaction, the reaction shifts toward the left in order to compensate from the heat loss. The reaction may then be rewritten to include energy as a reactant.
Since the energy is on the reactant side, the reaction is endothermic.
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Which of the following reactions will be favored when the pressure in a system is increased?
I. 
II. 
III. 
Which of the following reactions will be favored when the pressure in a system is increased?
I.
II.
III.
With increased pressure, each reaction will favor the side with the least amount of moles of gas. In this problem we are looking for the reactions that favor the products in this scenario. I will favor reactants, II will favor products, III will favor reactants.
With increased pressure, each reaction will favor the side with the least amount of moles of gas. In this problem we are looking for the reactions that favor the products in this scenario. I will favor reactants, II will favor products, III will favor reactants.
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Which of the following will cause an equilibrium shift in an exothermic reaction towards the products?
I. Decreasing the temperature
II. Evaporating the product
III. Adding a catalyst
Which of the following will cause an equilibrium shift in an exothermic reaction towards the products?
I. Decreasing the temperature
II. Evaporating the product
III. Adding a catalyst
I) Decreasing the temperature would take away heat from the system (a product), driving the reaction towards the products. II) Evaporating product would take a product away from the system, driving the reaction towards the products. III) Adding a catalyst only affects the rate of the reaction and does not effect equilibrium.
I) Decreasing the temperature would take away heat from the system (a product), driving the reaction towards the products. II) Evaporating product would take a product away from the system, driving the reaction towards the products. III) Adding a catalyst only affects the rate of the reaction and does not effect equilibrium.
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Which of the following stresses to a system at equilibrium cause an increase in the production of CH3OH ? CO(g)+ 2H2(g) → CH3OH(g)
(a) H2 is added. (b) The volume is increased. (c) Argon is added. (d) Removing CO.
Which of the following stresses to a system at equilibrium cause an increase in the production of CH3OH ? CO(g)+ 2H2(g) → CH3OH(g)
(a) H2 is added. (b) The volume is increased. (c) Argon is added. (d) Removing CO.
LeChatelier’s principle states that if a stress is applied to a rxn mixture at equilibrium,
reaction occurs in the direction that relieves the stress. Therefore, adding H2 will produce
more methanol. By increasing the volume and decreasing the pressure, there will be a net
reaction in the direction that increases the number of moles of gas. So since there are 3
moles of gas on the products side and only 1 mole of gas on the reactants side, if the volume is increased less methanol will be produced. Since Ar is an inert, it will have no effect on the amount of methanol produced. Removing CO will cause a shift in the reaction from right to left causing less methanol to be produced.
LeChatelier’s principle states that if a stress is applied to a rxn mixture at equilibrium,
reaction occurs in the direction that relieves the stress. Therefore, adding H2 will produce
more methanol. By increasing the volume and decreasing the pressure, there will be a net
reaction in the direction that increases the number of moles of gas. So since there are 3
moles of gas on the products side and only 1 mole of gas on the reactants side, if the volume is increased less methanol will be produced. Since Ar is an inert, it will have no effect on the amount of methanol produced. Removing CO will cause a shift in the reaction from right to left causing less methanol to be produced.
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Which of the following does not affect the equilbrium of a reaction?
Which of the following does not affect the equilbrium of a reaction?
Le Chatelier's principle states that the concentration of reactants/products, the addition/subtraction of heat, and changing the volume of a reaction would all be factors that affect equilibrium. A catalyst alters the reaction rate without changing equilibrium.
Le Chatelier's principle states that the concentration of reactants/products, the addition/subtraction of heat, and changing the volume of a reaction would all be factors that affect equilibrium. A catalyst alters the reaction rate without changing equilibrium.
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