Hey,

Absolutely no worries! Acid Base can get confusing if you try to really understand it fully, so I would argue this question is more complex than it seems.

We are safe to treat the weak acid as 100% HA to start for 2 reasons.

1. If HA is a weak acid, the dissociation is quite small (as you have alluded to).

But more importantly

2. When HA dissociates it forms A- and H+ (or H3O+ more accurately). This will actually react with the base before the HA will as H3O+ has a lower pKa than that of HA.

Therefore for all intents and purposes, it looks like you are completely titrating 100 % HA, even if some of it is in actuality dissociated.

I hope this makes things clearer!

For number 2: What will react with the base? That makes sense that H3O+ has a lower pKa because it is a stronger acid. Would you consider H3O+ a strong acid? It's not on our list of SA's...

Both HA and H3O+ formed as a result of dissociation of HA can react with the base. Technically, H3O+ is considered a strong acid according to it's pKa, but we don't really talk about it as a strong acid on the MCAT.

My main point is is that H3O+ still needs to be neutralized in a titration, meaning that acid dissociation can't lessen the amount of base you need for a titration.

So if the dissociation of HA is so small, where does all the H3O+ come from then? I take it that it looks like there is 100% HA because the dissociation is small at a lower, and then once you get to the buffer point, you start to see HA become dissociated, but I don't really understand how we have a low pH to begin with if there is such little dissociation.

pH is interesting - depending on the initial concentration of your acid, pH 2 is a pretty reasonable pH, maybe a bit on the low side. For reference, 1M solutions of HCl are pH 0 (10^0 = 1).

Just as a quick proof that this pH of 2 is reasonable, I'll work out what the pH of a 1 M solution of acetic acid is:

1. Acetic acid = pKa of 5 (Ka = 10^-5)

Assuming a 1 M solution of Acetic Acid:

2. Ka = [H+][A-]/[HA]

10^-5 = x^2/ 1
10^2.5 = x = [H+]

Therefore pH = 2.5, and therefore this is your starting pH at the titration. This seems unreasonable, but recall pH 2 is only about 10^-2 M H+, which is not an exceedingly large amount of dissociation.

Additionally, totally different question. When we speak about the auto-ionization of water, do we say that both pH and pOH increase? Likewise with heating water, pH decreases. But see here, if pH decreases, doesn't that mean that pOH must also increase because you are getting H+ from H20, and if you are making H+ you are also making OH-?

Yep! When you heat water, Kw increases, and necessarily, H+ and OH+ concentrations must increase at the same proportion - the idea here is a pH and pOH range of 0-14 is meaningful at certain conditions only - you can have this scale extend to -1 to 15 etc depending on your temperature. For that particular test question, it was tricky because they only asked about the pH, not the pOH.

Let me know if any of this needs clarification!

HA is not dissociating that much, so we assume it is at 100% start. While this isn't the truth, it doesn't really matter. Because even for a weak acid like acetic acid, it only dissociated around [10^-2.5 H+'s], which is quite small.

Yeah essentially, you might even argue that you can't tell the difference because when you create H3O+ you have to neutralize that in the titration too.

Then, when we add the base, we notice that the concentration of HA is not changing because H3O+ is the thing that is reacting with the strong base (because pKa of H3O+ is lower than pKa HA). Then we approach the buffer where we get 50% HA and 50% A-, which is at the pH = pKA of our initial acid solute. Then, we keep adding base and we hit the equivalence point which means that we have 100% A- (conjugate base of our initial acid in solution). This means, that at this point we equal moles of titrant and conjugate base in solution, or, the moles of our titrant added is equal to the moles of our initial acid.

This is a good summary, I might pick that we notice that the pH isn't changing a ton rather than saying the concentration of HA doesn't change(because some of the HA will get pulled to dissociate further as H3O+ is consumed via le Chatelier's).

The buffer statement is a good summary, as is the description of reaching equivalence point. The last thing I might add is that at the equivalence point, it will be slightly above pH 7 as A- is a weak base.

And finally, just confirm, pI of an amino acid is the point where the amino acid has a net neutral charge? So the carboxy group is -COO-, and the amino group is -NH3+. Not when the carboxy group is -COOH AND when the amino group is -NH2? Because this will never happen? The COO- picks up a hydrogen at a low pH, and the NH3+ picks up a hydrogen at a high pH?

Yep, a pH is the pH where you have a net zero charge. For most amino acids this is where the COO- and NH3+ coexist, for those with acidic or basic R groups you have to make sure the R group is neutral as well.

Yeah, I can't think of a pH where there will be COOH and NH2 on an amino acid in solution, as COOH requires low pH to be protonated, while NH2 would be protonated as NH3+ at low pH.


Thanks for all the help!

No problem! I'm impressed by your dedication to really trying to get these concepts down fully.