Defining/Classifying Acids and Bases - MCAT Chemical and Physical Foundations of Biological Systems
Card 1 of 105
Which of the following is the strongest acid?
Which of the following is the strongest acid?
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Hydroiodic acid (HI) is the strongest acid listed. Charge density decreases as the atomic size gets larger, thus stabilizing the charge when a hydrogen is given off. Hydroflouric acid (HF) is a relatively weak acid because electronegative elements hold on to their valence electrons more tightly, and are thus less likely to dissociate.
Hydroiodic acid (HI) is the strongest acid listed. Charge density decreases as the atomic size gets larger, thus stabilizing the charge when a hydrogen is given off. Hydroflouric acid (HF) is a relatively weak acid because electronegative elements hold on to their valence electrons more tightly, and are thus less likely to dissociate.
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Acids and bases can be described in three principal ways. The Arrhenius definition is the most restrictive. It limits acids and bases to species that donate protons and hydroxide ions in solution, respectively. Examples of such acids include HCl and HBr, while KOH and NaOH are examples of bases. When in aqueous solution, these acids proceed to an equilibrium state through a dissociation reaction.
All of the bases proceed in a similar fashion.
The Brønsted-Lowry definition of an acid is a more inclusive approach. All Arrhenius acids and bases are also Brønsted-Lowry acids and bases, but the converse is not true. Brønsted-Lowry acids still reach equilibrium through the same dissociation reaction as Arrhenius acids, but the acid character is defined by different parameters. The Brønsted-Lowry definition considers bases to be hydroxide donors, like the Arrhenius definition, but also includes conjugate bases such as the A- in the above reaction. In the reverse reaction, A- accepts the proton to regenerate HA. The Brønsted-Lowry definition thus defines bases as proton acceptors, and acids as proton donors.
Reducing agents, such as 
, are used to donate hydrogen ions to chemical species. 
will dissociate in solution to form hydrogen ions, which then reduce the target compound.
Which of the following must be true of 
when it is used as a reducing agent?
Acids and bases can be described in three principal ways. The Arrhenius definition is the most restrictive. It limits acids and bases to species that donate protons and hydroxide ions in solution, respectively. Examples of such acids include HCl and HBr, while KOH and NaOH are examples of bases. When in aqueous solution, these acids proceed to an equilibrium state through a dissociation reaction.
All of the bases proceed in a similar fashion.
The Brønsted-Lowry definition of an acid is a more inclusive approach. All Arrhenius acids and bases are also Brønsted-Lowry acids and bases, but the converse is not true. Brønsted-Lowry acids still reach equilibrium through the same dissociation reaction as Arrhenius acids, but the acid character is defined by different parameters. The Brønsted-Lowry definition considers bases to be hydroxide donors, like the Arrhenius definition, but also includes conjugate bases such as the A- in the above reaction. In the reverse reaction, A- accepts the proton to regenerate HA. The Brønsted-Lowry definition thus defines bases as proton acceptors, and acids as proton donors.
Reducing agents, such as
Which of the following must be true of
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is not an acid when it is used as a reducing agent. This can confuse students because, 
is in fact donating a hydrogen ion to a compound that is being reduced; however, we are reducing something, and thus we are adding electrons. The only way that 
could be an effective reducing agent is if it carried electrons on its donated hydrogen.
As a result, we are donating hydride (
) ions, and not protons. Acids are exclusively those species that donate protons in solution, not hydrides.
As a result, we are donating hydride (
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Which of the following is considered a weak base?
Which of the following is considered a weak base?
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Strong bases are those that completely dissociate in water. All group I hydroxides are considered strong bases, including sodium hydroxide, potassium hydroxide, and lithium hydroxide. Some group II hydroxides are also strong bases, including calcium hydroxide, barium hydroxide, and strontium hydroxide.
Ammonia is considered a weak base and is frequently used as a buffer in solution with ammonium, its conjugate acid.
Strong bases are those that completely dissociate in water. All group I hydroxides are considered strong bases, including sodium hydroxide, potassium hydroxide, and lithium hydroxide. Some group II hydroxides are also strong bases, including calcium hydroxide, barium hydroxide, and strontium hydroxide.
Ammonia is considered a weak base and is frequently used as a buffer in solution with ammonium, its conjugate acid.
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None of the questions in this set require the use of a calculator. Math problems are intended to mimic the level of math intensity that you will see on the MCAT exam.
A Bronsted-Lowry acid will be able to .
None of the questions in this set require the use of a calculator. Math problems are intended to mimic the level of math intensity that you will see on the MCAT exam.
A Bronsted-Lowry acid will be able to .
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The Bronsted-Lowry definition of an acid is one of three acid definitions that can be used on the MCAT. The three definitions are as follows:
Arrhenius acid: increases H+ concentration in water (releases an H+)
Bronsted-Lowry acid: donates a proton to a Bronsted base (Bronsted bases accept protons)
Lewis acid: accepts a pair of electrons from a Lewis base (Lewis bases donate electrons)
An easy way to remember the difference between Bronsted-Lowry and Lewis is that the "e" in Lewis comes before the "e" in Bronsted; therefore the Lewis definition has to do with electrons.
The Bronsted-Lowry definition of an acid is one of three acid definitions that can be used on the MCAT. The three definitions are as follows:
Arrhenius acid: increases H+ concentration in water (releases an H+)
Bronsted-Lowry acid: donates a proton to a Bronsted base (Bronsted bases accept protons)
Lewis acid: accepts a pair of electrons from a Lewis base (Lewis bases donate electrons)
An easy way to remember the difference between Bronsted-Lowry and Lewis is that the "e" in Lewis comes before the "e" in Bronsted; therefore the Lewis definition has to do with electrons.
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Which of the following is not a Lewis acid?
Which of the following is not a Lewis acid?
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The definition of a Lewis acid is an electron acceptor. Try drawing the Lewis dot structures for these compounds, and you will find that NH4+, Na+, and Al3+ are missing electrons. They each have an empty orbital, allowing them to accept electrons.
BF3 is not an ion; however, we know that boron only has 3 valence electrons, which means that even when they are all bound to fluorine the molecule does not satisfy the octet rule. BF3 only has 6 valence electrons around boron, and can accept another electron pair to get to the octet state.
CaO is a Lewis base. In the Lewis dot structure, we can see that the oxygen molecule has two lone pairs that it can donate to other molecules, making it an electron donor.
The definition of a Lewis acid is an electron acceptor. Try drawing the Lewis dot structures for these compounds, and you will find that NH4+, Na+, and Al3+ are missing electrons. They each have an empty orbital, allowing them to accept electrons.
BF3 is not an ion; however, we know that boron only has 3 valence electrons, which means that even when they are all bound to fluorine the molecule does not satisfy the octet rule. BF3 only has 6 valence electrons around boron, and can accept another electron pair to get to the octet state.
CaO is a Lewis base. In the Lewis dot structure, we can see that the oxygen molecule has two lone pairs that it can donate to other molecules, making it an electron donor.
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Which of the following is a weak base?
Which of the following is a weak base?
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Group 1 hydroxides, group 1 oxides, Ba(OH)2, Sr(OH)2, Ca(OH)2, and metal amides are the only strong bases. Everything else is considered a weak base.
Group 1 hydroxides, group 1 oxides, Ba(OH)2, Sr(OH)2, Ca(OH)2, and metal amides are the only strong bases. Everything else is considered a weak base.
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Which of the following will dissolve completely in water?
Which of the following will dissolve completely in water?
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The question is asking which of the following is a strong acid. Of the following, the only strong acid provided is perchloric acid, HClO4. NH3 is a weak base; HCOOH, C6H5COOH, and HCN are all weak acids. A strong acid is one that dissociates completely in water.
The strong acids that you must know for the MCAT are:
Hydroiodic acid (HI)
Hydrobromic acid (HBr)
Hydrochloric acid (HCl)
Perchloric acid (HClO4 )
Sulfuric acid (H2SO4)
and Nitric acid (HNO3)
The question is asking which of the following is a strong acid. Of the following, the only strong acid provided is perchloric acid, HClO4. NH3 is a weak base; HCOOH, C6H5COOH, and HCN are all weak acids. A strong acid is one that dissociates completely in water.
The strong acids that you must know for the MCAT are:
Hydroiodic acid (HI)
Hydrobromic acid (HBr)
Hydrochloric acid (HCl)
Perchloric acid (HClO4 )
Sulfuric acid (H2SO4)
and Nitric acid (HNO3)
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Carbonic anhydrase is an important enzyme that allows CO2 and H2O to be converted into H2CO3. In addition to allowing CO2 to be dissolved into the blood and transported to the lungs for exhalation, the products of the carbonic anhydrase reaction, H2CO3 and a related compound HCO3-, also serve to control the pH of the blood to prevent acidosis or alkalosis. The carbonic anhydrase reaction and acid-base reaction are presented below.
CO2 + H2O 
H2CO3
H2CO3 
HCO3- + H+
HCO3- is the of H2CO3.
Carbonic anhydrase is an important enzyme that allows CO2 and H2O to be converted into H2CO3. In addition to allowing CO2 to be dissolved into the blood and transported to the lungs for exhalation, the products of the carbonic anhydrase reaction, H2CO3 and a related compound HCO3-, also serve to control the pH of the blood to prevent acidosis or alkalosis. The carbonic anhydrase reaction and acid-base reaction are presented below.
CO2 + H2O
H2CO3
HCO3- is the of H2CO3.
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First, we need to determine what type of reaction HCO3- is participating in. We can see that a H+ is being produced in the forward reaction, and consumed in the reserve reaction; thus, we are looking at an acid-base reaction. H2CO3 has lost a H+, making it a Arrhenius acid. HCO3- is the conjugate base of the acid H2CO3.
First, we need to determine what type of reaction HCO3- is participating in. We can see that a H+ is being produced in the forward reaction, and consumed in the reserve reaction; thus, we are looking at an acid-base reaction. H2CO3 has lost a H+, making it a Arrhenius acid. HCO3- is the conjugate base of the acid H2CO3.
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Which of the following describes a Brønsted-Lowry acid?
Which of the following describes a Brønsted-Lowry acid?
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There are three definitions for an acid that are commonly used: an Arrhenius acid, a Brønsted-Lowry acid, and a Lewis acid. Arrhenius acids are any substances that create hydrogen ions in solution, and Lewis acids are any substances that accept an electron pair. A Brønsted-Lowry acid is defined as any substance that donates a proton.
There are three definitions for an acid that are commonly used: an Arrhenius acid, a Brønsted-Lowry acid, and a Lewis acid. Arrhenius acids are any substances that create hydrogen ions in solution, and Lewis acids are any substances that accept an electron pair. A Brønsted-Lowry acid is defined as any substance that donates a proton.
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Which of the following acids has the weakest conjugate base?
Which of the following acids has the weakest conjugate base?
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The stronger the acid, the weaker its conjugate base. This is because a strong acid will dissociate almost entirely, meaning that its conjugate base will not accept protons frequently. The only strong acid as an option is nitric acid (HNO3). As a result, we conclude that nitric acid has the weakest conjugate base.
The stronger the acid, the weaker its conjugate base. This is because a strong acid will dissociate almost entirely, meaning that its conjugate base will not accept protons frequently. The only strong acid as an option is nitric acid (HNO3). As a result, we conclude that nitric acid has the weakest conjugate base.
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Which of the following salts would result in an acidic solution?
Which of the following salts would result in an acidic solution?
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Certain salts are able to affect the pH of a solution, because the dissociated ions can act as acids or bases. The conjugate bases of strong acids will not affect the pH, nor will the conjugate acids of strong bases.
NH4+ is the conjugate acid of ammonia (NH3), a weak base. As a result, it will donate protons to the solution and make the solution acidic.
Certain salts are able to affect the pH of a solution, because the dissociated ions can act as acids or bases. The conjugate bases of strong acids will not affect the pH, nor will the conjugate acids of strong bases.
NH4+ is the conjugate acid of ammonia (NH3), a weak base. As a result, it will donate protons to the solution and make the solution acidic.
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Acids and bases can be described in three principal ways. The Arrhenius definition is the most restrictive. It limits acids and bases to species that donate protons and hydroxide ions in solution, respectively. Examples of such acids include HCl and HBr, while KOH and NaOH are examples of bases. When in aqueous solution, these acids proceed to an equilibrium state through a dissociation reaction.
All of the bases proceed in a similar fashion.
The Brønsted-Lowry definition of an acid is a more inclusive approach. All Arrhenius acids and bases are also Brønsted-Lowry acids and bases, but the converse is not true. Brønsted-Lowry acids still reach equilibrium through the same dissociation reaction as Arrhenius acids, but the acid character is defined by different parameters. The Brønsted-Lowry definition considers bases to be hydroxide donors, like the Arrhenius definition, but also includes conjugate bases such as the A- in the above reaction. In the reverse reaction, A- accepts the proton to regenerate HA. The Brønsted-Lowry definition thus defines bases as proton acceptors, and acids as proton donors.
Instead of a monoprotic acid like 
, a scientist is studying a diprotic acid, like 
. Which of the following is true regarding 
?
I. It only lowers the pH in solution when both protons are released
II. It releases a much greater fraction of its first proton than its second
III. 
is its conjugate acid
Acids and bases can be described in three principal ways. The Arrhenius definition is the most restrictive. It limits acids and bases to species that donate protons and hydroxide ions in solution, respectively. Examples of such acids include HCl and HBr, while KOH and NaOH are examples of bases. When in aqueous solution, these acids proceed to an equilibrium state through a dissociation reaction.
All of the bases proceed in a similar fashion.
The Brønsted-Lowry definition of an acid is a more inclusive approach. All Arrhenius acids and bases are also Brønsted-Lowry acids and bases, but the converse is not true. Brønsted-Lowry acids still reach equilibrium through the same dissociation reaction as Arrhenius acids, but the acid character is defined by different parameters. The Brønsted-Lowry definition considers bases to be hydroxide donors, like the Arrhenius definition, but also includes conjugate bases such as the A- in the above reaction. In the reverse reaction, A- accepts the proton to regenerate HA. The Brønsted-Lowry definition thus defines bases as proton acceptors, and acids as proton donors.
Instead of a monoprotic acid like
I. It only lowers the pH in solution when both protons are released
II. It releases a much greater fraction of its first proton than its second
III.
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A diprotic acid will ionize fewer of its second protons than its first; that is, a smaller fraction of the 
ions will react to 
than 
molecules will react to 
.
The presence of protons in solution from the original ionization of 
in the first reaction drives the equilibrium toward 
(reactants) and away from 
(products) in the second reaction.
A diprotic acid will ionize fewer of its second protons than its first; that is, a smaller fraction of the
The presence of protons in solution from the original ionization of
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Acids and bases can be described in three principal ways. The Arrhenius definition is the most restrictive. It limits acids and bases to species that donate protons and hydroxide ions in solution, respectively. Examples of such acids include HCl and HBr, while KOH and NaOH are examples of bases. When in aqueous solution, these acids proceed to an equilibrium state through a dissociation reaction.
All of the bases proceed in a similar fashion.
The Brønsted-Lowry definition of an acid is a more inclusive approach. All Arrhenius acids and bases are also Brønsted-Lowry acids and bases, but the converse is not true. Brønsted-Lowry acids still reach equilibrium through the same dissociation reaction as Arrhenius acids, but the acid character is defined by different parameters. The Brønsted-Lowry definition considers bases to be hydroxide donors, like the Arrhenius definition, but also includes conjugate bases such as the A- in the above reaction. In the reverse reaction, A- accepts the proton to regenerate HA. The Brønsted-Lowry definition thus defines bases as proton acceptors, and acids as proton donors.
In the reverse reaction of 
, the proton is acting as a(n) , and is thus a .
Acids and bases can be described in three principal ways. The Arrhenius definition is the most restrictive. It limits acids and bases to species that donate protons and hydroxide ions in solution, respectively. Examples of such acids include HCl and HBr, while KOH and NaOH are examples of bases. When in aqueous solution, these acids proceed to an equilibrium state through a dissociation reaction.
All of the bases proceed in a similar fashion.
The Brønsted-Lowry definition of an acid is a more inclusive approach. All Arrhenius acids and bases are also Brønsted-Lowry acids and bases, but the converse is not true. Brønsted-Lowry acids still reach equilibrium through the same dissociation reaction as Arrhenius acids, but the acid character is defined by different parameters. The Brønsted-Lowry definition considers bases to be hydroxide donors, like the Arrhenius definition, but also includes conjugate bases such as the A- in the above reaction. In the reverse reaction, A- accepts the proton to regenerate HA. The Brønsted-Lowry definition thus defines bases as proton acceptors, and acids as proton donors.
In the reverse reaction of
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In terms of the passage, the lone proton can be considered a proton donor and would, therefore, be a Brønsted-Lowry acid. This is not an answer choice.
The third acid-base definition is the Lewis definition, which states that acids are electron acceptors and bases are electron donors. The negative charge on the 
signifies that it is a Lewis base with available electrons to donate. The proton is accepting these electrons from 
, and is thus acting as a Lewis acid.
In terms of the passage, the lone proton can be considered a proton donor and would, therefore, be a Brønsted-Lowry acid. This is not an answer choice.
The third acid-base definition is the Lewis definition, which states that acids are electron acceptors and bases are electron donors. The negative charge on the
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Which of the following is not a strong acid?
Which of the following is not a strong acid?
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Strong acids are those that fully dissociate in an aqueous environment. Most protonated halogens are considered strong acids, with the exception of hydrogen fluoride; hydrochloric, hydrobromic, and hydroiodic acid are considered strong acids.
Perchloric acid is perhaps the strongest acid, and is formed from a hydrogen atom and a perchlorate ion: 
. Nitric acid is formed from a hydrogen atom and a nitrate ion, making it highly soluble and making it easy to dissolve in solution: 
.
Acetic acid, and most other organic acids, are considered weak acids since they lack stability when deprotonated. Stronger acids will form ions or resonace structures that are highly stable, favoring the release of the bound proton, while weaker acids will bind the proton more tightly to maintain a more stable structure. Acetic acid has the structure 
.
Strong acids are those that fully dissociate in an aqueous environment. Most protonated halogens are considered strong acids, with the exception of hydrogen fluoride; hydrochloric, hydrobromic, and hydroiodic acid are considered strong acids.
Perchloric acid is perhaps the strongest acid, and is formed from a hydrogen atom and a perchlorate ion:
Acetic acid, and most other organic acids, are considered weak acids since they lack stability when deprotonated. Stronger acids will form ions or resonace structures that are highly stable, favoring the release of the bound proton, while weaker acids will bind the proton more tightly to maintain a more stable structure. Acetic acid has the structure
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The Haber-Bosch process, or simply the Haber process, is a common industrial reaction that generates ammonia from nitrogen and hydrogen gas. A worker in a company generates ammonia from the Haber process. He then dissociates the gaseous ammonia in water to produce an aqueous solution. Since ammonia is a base, it will accept a proton from water, generating 
and ammonium ion products. The two reactions involved are:
What happens when ammonium ions are dissociated in water?
The Haber-Bosch process, or simply the Haber process, is a common industrial reaction that generates ammonia from nitrogen and hydrogen gas. A worker in a company generates ammonia from the Haber process. He then dissociates the gaseous ammonia in water to produce an aqueous solution. Since ammonia is a base, it will accept a proton from water, generating
What happens when ammonium ions are dissociated in water?
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Ammonia is a base and produces ammonium ions in water; therefore, ammonium is the conjugate acid of ammonia. This means that ammonium ions in water will donate protons to water, regenerating the conjugate base (ammonia). Protonated water contains a molecular formula of 
, which are called hydronium ions.
Remember that hydronium ions are always produced during acid dissociation. The presence of more hydronium ions indicates that the solution is acidic and has a low pH.
Ammonia is a base and produces ammonium ions in water; therefore, ammonium is the conjugate acid of ammonia. This means that ammonium ions in water will donate protons to water, regenerating the conjugate base (ammonia). Protonated water contains a molecular formula of
Remember that hydronium ions are always produced during acid dissociation. The presence of more hydronium ions indicates that the solution is acidic and has a low pH.
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None of the questions in this set require the use of a calculator. Math problems are intended to mimic the level of math intensity that you will see on the MCAT exam.
A Bronsted-Lowry acid will be able to .
None of the questions in this set require the use of a calculator. Math problems are intended to mimic the level of math intensity that you will see on the MCAT exam.
A Bronsted-Lowry acid will be able to .
Tap to reveal answer
The Bronsted-Lowry definition of an acid is one of three acid definitions that can be used on the MCAT. The three definitions are as follows:
Arrhenius acid: increases H+ concentration in water (releases an H+)
Bronsted-Lowry acid: donates a proton to a Bronsted base (Bronsted bases accept protons)
Lewis acid: accepts a pair of electrons from a Lewis base (Lewis bases donate electrons)
An easy way to remember the difference between Bronsted-Lowry and Lewis is that the "e" in Lewis comes before the "e" in Bronsted; therefore the Lewis definition has to do with electrons.
The Bronsted-Lowry definition of an acid is one of three acid definitions that can be used on the MCAT. The three definitions are as follows:
Arrhenius acid: increases H+ concentration in water (releases an H+)
Bronsted-Lowry acid: donates a proton to a Bronsted base (Bronsted bases accept protons)
Lewis acid: accepts a pair of electrons from a Lewis base (Lewis bases donate electrons)
An easy way to remember the difference between Bronsted-Lowry and Lewis is that the "e" in Lewis comes before the "e" in Bronsted; therefore the Lewis definition has to do with electrons.
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Which of the following is not a Lewis acid?
Which of the following is not a Lewis acid?
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The definition of a Lewis acid is an electron acceptor. Try drawing the Lewis dot structures for these compounds, and you will find that NH4+, Na+, and Al3+ are missing electrons. They each have an empty orbital, allowing them to accept electrons.
BF3 is not an ion; however, we know that boron only has 3 valence electrons, which means that even when they are all bound to fluorine the molecule does not satisfy the octet rule. BF3 only has 6 valence electrons around boron, and can accept another electron pair to get to the octet state.
CaO is a Lewis base. In the Lewis dot structure, we can see that the oxygen molecule has two lone pairs that it can donate to other molecules, making it an electron donor.
The definition of a Lewis acid is an electron acceptor. Try drawing the Lewis dot structures for these compounds, and you will find that NH4+, Na+, and Al3+ are missing electrons. They each have an empty orbital, allowing them to accept electrons.
BF3 is not an ion; however, we know that boron only has 3 valence electrons, which means that even when they are all bound to fluorine the molecule does not satisfy the octet rule. BF3 only has 6 valence electrons around boron, and can accept another electron pair to get to the octet state.
CaO is a Lewis base. In the Lewis dot structure, we can see that the oxygen molecule has two lone pairs that it can donate to other molecules, making it an electron donor.
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Which of the following is a weak base?
Which of the following is a weak base?
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Group 1 hydroxides, group 1 oxides, Ba(OH)2, Sr(OH)2, Ca(OH)2, and metal amides are the only strong bases. Everything else is considered a weak base.
Group 1 hydroxides, group 1 oxides, Ba(OH)2, Sr(OH)2, Ca(OH)2, and metal amides are the only strong bases. Everything else is considered a weak base.
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Which of the following will dissolve completely in water?
Which of the following will dissolve completely in water?
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The question is asking which of the following is a strong acid. Of the following, the only strong acid provided is perchloric acid, HClO4. NH3 is a weak base; HCOOH, C6H5COOH, and HCN are all weak acids. A strong acid is one that dissociates completely in water.
The strong acids that you must know for the MCAT are:
Hydroiodic acid (HI)
Hydrobromic acid (HBr)
Hydrochloric acid (HCl)
Perchloric acid (HClO4 )
Sulfuric acid (H2SO4)
and Nitric acid (HNO3)
The question is asking which of the following is a strong acid. Of the following, the only strong acid provided is perchloric acid, HClO4. NH3 is a weak base; HCOOH, C6H5COOH, and HCN are all weak acids. A strong acid is one that dissociates completely in water.
The strong acids that you must know for the MCAT are:
Hydroiodic acid (HI)
Hydrobromic acid (HBr)
Hydrochloric acid (HCl)
Perchloric acid (HClO4 )
Sulfuric acid (H2SO4)
and Nitric acid (HNO3)
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Which of the following is the strongest acid?
Which of the following is the strongest acid?
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Hydroiodic acid (HI) is the strongest acid listed. Charge density decreases as the atomic size gets larger, thus stabilizing the charge when a hydrogen is given off. Hydroflouric acid (HF) is a relatively weak acid because electronegative elements hold on to their valence electrons more tightly, and are thus less likely to dissociate.
Hydroiodic acid (HI) is the strongest acid listed. Charge density decreases as the atomic size gets larger, thus stabilizing the charge when a hydrogen is given off. Hydroflouric acid (HF) is a relatively weak acid because electronegative elements hold on to their valence electrons more tightly, and are thus less likely to dissociate.
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