[haz658]

          I have what I believe many here will consider a rather simple question, but one for which I have yet to find an answer while searching the Internet for the subject matter. It is a question which was on a test I took for undergraduate analytical chemistry years ago, I answered incorrectly, and to whose author I no longer have access.

                    "Hydroxybenzene is a monoprotic acid with a pK_a of 9.98. Under what conditions will this acid act as a strong acid? Under what conditions will its conjugate base act as a strong base? Provide a brief justification."

                 From what I have read, having a pK_a of this value would indicate that hydroxybenzene is a weak acid. This would mean that the reaction which the K_a describes does not proceed to products to an appreciable extent. I read somewhere that an increase in temperature causes endergonic reactions to increase formation of products, so I figure that this could potentially be a correct answer to this question. Another condition which would seem to make sense is if hydroxybenzene were added to a strongly basic solution. These are the only conditions I conceived. I didn't make any attempts at the second part of the problem, but it would seem that similarly conceived conditions would work for this scenario. Although, I believe the reverse reaction of the weak base being protonated to form phenol is favored, which would require a decrease in temperature to push the reaction forward(?)
               While I provided my thoughts above, any feedback or further ideas/suggestions would be greatly appreciated.
                                            Thanks in advance,
                                                                       Matt

[Borek]

This is not a very simple question, and to be honest with you I have no idea what was the intended answer. The only thing I can think of is to use protophilic solvent different then water - like DMF, piridine, ethylenediamine and so on.

Then, could be it is too early for me...

[Jorriss]

I haven't taken analytical chemistry so pardon me if this is done but, isn't the weak acid scale pretty fluid?

So, if the pKa is around 9, measured in water it's not that acidic, but if you add that acid to a solution that is basic, it will give up protons acting as a stronger acid. In other words, in a basic solvent, it will act like a strong acid?

It's like, nitric acid is a strong acid, but it acts as a base in a mixture of sulfuric and nitric acid.

[aeacfm]

from wikipedia
The acid dissociation constant for an acid is a direct consequence of the underlying thermodynamics of the dissociation reaction; the pKa value is directly proportional to the standard Gibbs energy change for the reaction. The value of the pKa changes with temperature and can be understood qualitatively based on Le Chatelier's principle: when the reaction is endothermic, the pKa decreases with increasing temperature; the opposite is true for exothermic reactions. The underlying structural factors that influence the magnitude of the acid dissociation constant include Pauling's rules for acidity constants, inductive effects, mesomeric effects, and hydrogen bonding.

[haz658]

Thank you all for your responses. I apologize for not replying sooner, but I have been sick and have not had the opportunity. It made me feel better- about not knowing the answer to this question- to see Borek state that he did not view this question as being simple. I considered the use of a protophilic solvent/basic solution, and the fact that several of you suggested these solutions makes me confident in their viability. Regarding aeacfm's post, I read that article from Wikipedia on the acid dissociation constant, and I mentioned the potential efficacy of increasing the temperature to increase the K_a. Reading it again, however, has caused me to wonder whether it meant endergonic instead of endothermic. I used the term endergonic in my original post, because the Wikipedia article states that pK_a is related to Gibbs free energy, which determines whether a reaction is endergonic or exergonic. This relationship is shown through the equation
           $$ \Delta G = -RTlnK /$$
So should the article replace endothermic with endergonic? Or were they basing this concept of increasing the temperature on the equation
                 $$ \Delta G = \Delta H - T \Delta S /$$
and thereby stating that the increase in temperature is impacting enthalpic aspect of the reaction?
      Either way, thank you all for the feedback regarding my question.

[Borek]

from wikipedia

How does the quote addresses the original question? Have you actually read it?

So, if the pKa is around 9, measured in water it's not that acidic, but if you add that acid to a solution that is basic, it will give up protons acting as a stronger acid. In other words, in a basic solvent, it will act like a strong acid?

As long as there is water present in quantities large enough, presence of other solvent is not that important. When water is present in small quantities, situation doesn't differ from the one I have described earlier.

Edit: sorry, I have started the post several hours ago but I had unexpected visitors, so it took way too much time. And I am not sure if I am able to think correctly right now

[haz658]

I finally found the answer which my professor desired. The text for that course was Daniel C. Harris's Sixth Edition Quantitative Chemical Analysis. While searching through the questions at the end of chapter 10, it asks the same question as the one I quoted, but with a more general skew. It directs you to the middle of the chapter, where he devotes very little space to the concept. To the point, it states in the text that the fraction of dissociation of an acid in solution is dependent upon its concentration. All weak electrolytes dissociate more as they are diluted (thus increasing their acidity with a decrease in their concentration). At some small concentration, these weak acids can dissociate nearly completely. In the other direction, the greater the concentration of a weak acid, the less it will dissociate (to some minimum, at which point it will decrease very little, if at all). This would mean that the conjugate (weak) base at this concentration will associate/protonate/hydrolyze completely, making it a strong base at high concentrations. Harris points out that at any given concentration a strong acid will dissociate more than a weak acid (and a strong base will associate/protonate/hydrolyze more than a weak base).