Acids and Bases - MCAT Chemical and Physical Foundations of Biological Systems

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.

A buffer must contain either a weak base and its salt or a weak acid and its salt. A mixture of ammonia and ammonium chloride is an example of the first case, since ammonia, NH3, is a weak base and ammonium chloride, NH4Cl, contains its salt.

Though autoionization of water produces small amounts of H3O+ and OH-, each conjugate salts of H2O, they exist in such small amounts as to make any buffering effects negligible.

First, we need to determine from both the information presented in the paragraph and the chemical equation what HCO3- is doing. The paragraph tells us “HCO3- also \[serves\] to control the pH of the blood.” This is the definition of a buffer. A buffer is able to mitigate the addition of acid or base to a solution, or in this case, blood, by being deprotonated or protonated. For example, if the blood becomes too acidic, H+ will recombine with HCO3- to form H2CO3, thus stabilizing the pH. Alternatively, if the pH becomes too basic, H2CO3 will dissociate, releasing H+ into solution.

Whenever a concentration of the conjugate base is initially present in the solution, we say that the solution has been buffered. NaCN will completely dissolve in solution, meaning that there will be 1M of CN- ions in the solution initially (one mole of NaCN in 1L of soultion equals 1M). With this in mind, the new equilibrium equation for the acid's dissociation is as follows:

The initial concentrations of CN- and HCN are given in the question as 1M and 2M, respectively. The initial concentration of H+ is zero. As HCN begine to dissociate, H+ and CN- will each increase by X amount, while HCN will decrease by X amount.

Since 1M and 2M are much larger numbers than the value for X, the X values next to these numbers can be omitted from the equation.

By plugging this value into the equation , we determine that the pH of the solution is 8.9. (Remember that X is equal to the amount of increase in H+ ions).

This may seem wrong, because HCN is an acid, and we are used to acids resulting in solutions with pH values less than 7; however, HCN is a weak acid, and a large concentration of its conjugate base can result in a basic solution.

In a buffered solution, when the concentrations of acid and conjugate base are equal, we know the .

This is derived from the Henderson-Hasselbalch equation: .

When concentrations are equal, the log of is , and .

In our question, the pH is approximately one unit greater than the .

Because of this, we know that the log of the two concentrations must be equal to one.

The log of is , and therefore the conjugate base must be ten times greater than the acid; therefore the ratio of acid to base is approximately .

We can quickly narrow this question to two possible answers, since the buffered solution is more basic than the , and thus we would expect there to be a greater concentration of base than acid in the solution.

A buffer is best made of equal and copious amounts of an acid and its conjugate base. The acid must have a pKa as close as possible to the pH desired. For this question, the desired pH is 6. The best option will have a pKa of about 5.9.

Adding larger amounts of the acid than the base will skew the equilibrium in favor of the acid, increasing the overall proton concentration, and limiting the buffer ability to absorb more proton addition without changing pH significantly.

A good buffer solution consists of a weak acid with its conjugate base, or a weak base and its conjugate acid. A strong acid or base will never be a good buffer. We will go through each answer choice in detail.

Sulfuric acid is a strong acid and cannot be used in a buffer.

Neither sodium chloride, nor calcium chloride constitute a weak acid or weak base. Both of these are simple salts and will not be able to create a buffer solution.

Hydrochloric acid is a strong acid and cannot be used in a buffer.

In solution, these salts will produce sulfate ions and hydrogen sulfate ions. Hydrogen sulfate is a weak acid and sulfate is its conjugate base. These ions will form a buffer solution.

A buffer system is composed of one of the following scenarios:

1. A weak base with its conjugate acid.

2. A weak acid with its conjugate base.

The only option that does not qualify is the nitrate, nitrite pair. A difference of one oxygen does not make a buffer system.

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 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 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)

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.

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.

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.

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.

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.

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.

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.

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.