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.

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. 

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.

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 pK_a as close as possible to the pH desired. For this question, the desired pH is 6. The best option will have a pK_a 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 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.

A buffer solution is used to keep the pH of the solution within a relatively narrow range. It must resist change in pH upon addition of acids or bases. A weak acid is used to lower the pH in response to addition of base or removal of acid, and its conjugate base is used to raise the pH in response to addition of acid or removal of base.

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 (). 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.

Since sodium hydroxide is a strong base, it will dissociate completely in water. This means that the concentration of the base will be equal to the concentration of hydroxide ions after the reaction runs to completion.

We can find the concentration of hydroxide ions via stoichiometry. One hydroxide ion is created from each molecule of sodium hydroxide that dissociates.

Since we have the concentration of hydroxide ions, we can solve for the pOH of the solution.

The question asks us to find the pH of the solution, so we will need to convert pOH to pH. To do so, we simply subtract the pOH from 14. 

The pH of the solution is 12.3. Because sodium hydroxide is a strong base, it makes sense that the pH is above 7.