7
$\begingroup$

My assumption is yes, as the $\mathrm{pH}$ levels of each will be similar and thus they have the same acidity. Is this correct or am I missing something?

$\endgroup$
25
$\begingroup$

The pH is the same, yes, but the weak acid has undissociated molecules "in reserve", so to speak, and thus can neutralize a lot more of any base.

Indeed, this concept of weak acids and their conjugate bases having molecules "in reserve", reacting only when called upon, is the principle behind acid-base buffers.

$\endgroup$
5
$\begingroup$

There are two main differences that can be observed, that Oscar's answer did not mention.

Firstly, when you add a strong base to the acid solution, the solution pH changes far more drastically when you had used a strong acid, simply because the amount of acid is very little to begin with.

Secondly, there is no buffer region for a strong acid solution. With a weak acid solution, at the rise in pH will slow down until a buffer region that lasts pretty much until nearly all the acid has been neutralized.

For instance a weak acid HA with disassociation constant $10^{-4.75}$ at concentration $0.1 \text{ mol} \text{ dm}^{-3}$ will have a pH of about $2.88$. A semi-strong acid with disassociation constant $10^{-2}$ will need to be at concentration about $0.0015 \text{ mol} \text{ dm}^{-3}$ to have the same pH. The pH of the solution as these two solutions are titrated against a strong base BOH can be easily proven to be the logarithm of the root of some cubic equation, which will have the following familiar shapes:

Titration pH (This graph is when the added base is at concentration $0.2 \text{ mol} \text{ dm}^{-3}$. The blue curve is for the weak acid and the orange curve is for the semi-strong acid.)

$\endgroup$
  • $\begingroup$ I didn't show the graph for a strong acid because it would be squashed too much to the left and there's really nothing to see. According to Wikipedia, what I called a semi-strong acid is still considered a weak acid, though on the strong side, but anyway it is strong enough for illustrative purposes. $\endgroup$ – user21820 Jun 10 '16 at 8:13
  • 1
    $\begingroup$ I am guessing the red line represents your "semi-strong" acid and the blue line represents the weak acid? It's a good idea to add it in imo $\endgroup$ – orthocresol Jun 10 '16 at 12:39
  • $\begingroup$ @orthocresol: Done. I had forgotten about that! $\endgroup$ – user21820 Jun 10 '16 at 13:07
0
$\begingroup$

The main difference between those two kinds of acid is the way that they dissociation occurs. A strong acid will rapidly dissociate in its ions with nearly a 100% of conversion p ex: HCl. Instead weak acids like acetic acid remain in equilibrium with its ions HAc <=> H+ + Ac- This means that in order to achieve a value of pH with a weak acid you need to consider that the acid will only partially increase the concentration of H+ ions so the final concentration will be given by the equilibrium constant (Ka) for this acid. Instead the pH given by a strong acid will be given simply by the concentration of the acid. So if you have HCl with a concentration of 10^-3 mol/L then the pH of that solution will be three. I hope this was helpful :)

$\endgroup$
  • $\begingroup$ Just fyi, acetic acid is HOAc, AcH actually represents acetaldehyde $\endgroup$ – orthocresol Jun 10 '16 at 12:36

Your Answer

By clicking “Post Your Answer”, you agree to our terms of service, privacy policy and cookie policy

Not the answer you're looking for? Browse other questions tagged or ask your own question.