Ambiguity as to source of high pH

[neilparker62]

Re ambiguity regarding the source of high pH in a solution following titration of NaOH (strong base) against a solution of acetic acid (weak acid) as described in attached pdf. Can we resolve the ambiguity ?

 

[mjc123]

What do you think about the suggestion in the last sentence?

 

[Borek]

A bit moot for me. Basically if you see the solid "final" color you never know if you overshoot or not (doesn't even matter if it is an acid/base titration or something else). That's why it is important to titrate very slowly close to the end so that you can see color change while it occurs.

 

[Mayhem]

Consider the appearance of a weak acid titrated strong base titration curve (right)

If you overshoot the equivalence point by just 0.05-0.1 ml your pH goes from 8.72 to to 11-12 (approx). Your source of high pH is (unsurprisingly) due to excess strong base.

 

[neilparker62]

Consider the appearance of a weak acid titrated strong base titration curve (right)

View attachment 360681
If you overshoot the equivalence point by just 0.05-0.1 ml your pH goes from 8.72 to to 11-12 (approx). Your source of high pH is (unsurprisingly) due to excess strong base.

So you won't know really - whether you've hit the neutralisation point or the equivalence point ?

 

[neilparker62]

What do you think about the suggestion in the last sentence?

Well it's fine but I wonder how easy it is to do practically ? Because the pH is changing so fast once it gets close to 7.

 

[Mayhem]

So you won't know really - whether you've hit the neutralisation point or the equivalence point ?

Phenolpthalein changes color at pH ≥ 8.3. It would therefore be colorless a pH 7. It is a misconception that phenolpthalein indicates the equivalence point, but it changes color at the endpoint, i.e. after all of the weak acid has been depleted and the solution is no longer buffered.

Colorimetric titration (such as phenolpthalein) is seldom useful in real life and potentiometric (pH meter in solution) is much more reliable.

 

[symbolipoint]

Colorimetric titration (such as phenolpthalein) is seldom useful in real life and potentiometric (pH meter in solution) is much more reliable.

I at least, fully believe you! But you will find people who will argue against us.

Also about those graphs in post #4, the one on the right hand side looks like titration of a diprotic weak acid. First "equivalence" point appears to be about pH 4.

 

[Mayhem]

at least, fully believe you! But you will find people who will argue against us.

Also about those graphs in post #4, the one on the right hand side looks like titration of a diprotic weak acid. First "equivalence" point appears to be about pH 4.

Not quite. The region before the inflection of the equivalence point is governed by the Henderson-Hasselbalch equation.

The region around the equivalence point must be modeled differently, as it exists in the limit of [HA] -> 0, where the equation approaches infinity and afterwards there is no acidic medium and only strong base, and the HH equation is not applicable. Instead pH = -log([H⁺]).

Of course for di- and multiprotic acids, the region after the equivalence points (except the last) are also modelled after the HH equation, which is why we can determine multiple pKa values.

Some non-idealities contribute to deviations from the fitting of the log function.

 

[Borek]

I at least, fully believe you! But you will find people who will argue against us.

Also about those graphs in post #4, the one on the right hand side looks like titration of a diprotic weak acid. First "equivalence" point appears to be about pH 4.

Nope, looks perfectly OK to me - 0.1 M acetic acid plus 0.1 M NaOH.

Diprotic looks either differently - oxalic acid (pKa1 = 1.252, pKa2 = 4.266)

or if the difference between pKa1 and pKa2 gets lower, very similar to just the weak acid, with longer initial slope, and there is no first endpoint visible (here - pKas of 3 and 4.266)

 

[symbolipoint]

Borek what I say is that the graph on the right looks like it is showing a pH change, not too strong, but a pH change at about the 5 ml. mark, pH 4. That looks to me like a first-proton neutralization. The second proton neutralization at pH 8.

 

[Mayhem]

Borek what I say is that the graph on the right looks like it is showing a pH change, not too strong, but a pH change at about the 5 ml. mark, pH 4. That looks to me like a first-proton neutralization. The second proton neutralization at pH 8.

See my post. The first region is governed by the Henderson-Hasselbalch equation. It is indeed called the buffer equation in some texts.

I know the "turn" doesnt look truly logarithmic but this could simply be an experimental error, such as not allowing the system to properly reach equilibrium between measurements.

 

[Borek]

Borek what I say is that the graph on the right looks like it is showing a pH change, not too strong, but a pH change at about the 5 ml. mark, pH 4. That looks to me like a first-proton neutralization. The second proton neutralization at pH 8.

I understand what you are saying, by my experience tells me otherwise - compare the curves I calculated (just as examples of what to expect from monoprotic and diprotic acids) with the curve in question.

Besides, if the acid is diprotic, how come first equivalence is at 5 mL, and the second at 25? If anything, it should be at 10 mL.

 

[Borek]

To make things even funnier (or more confusing, if you prefer ), here is a plot for a strong acid (0.1 M HCl) titrated with a strong, but diluted base (0.001 M NaOH). Initially pH changes are dominated by dilution effects, and that makes the curve look very similar to that of the titration of a weak acid.