Can I combine an Acid dissociation with autoionization of H2O?

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The discussion focuses on calculating the final pH after adding KOH to HCH3COO, with the assumption that the reaction goes to completion due to a high K value. It is noted that the derived K value is essentially the reciprocal of the acetate's base dissociation constant, which indicates that acetate solutions can be slightly alkaline. While the approach using the Henderson-Hasselbalch equation is generally valid, caution is advised for very dilute solutions or when neutralizing more than 95% of the acid, as this can lead to inaccuracies due to the hydrolysis of the conjugate base. The conversation highlights the importance of validating assumptions in chemical equilibria, especially in edge cases. Overall, the method discussed is reasonable but should be applied with careful consideration of the specific conditions.
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Can I do this?

HCH3COO + H2O <<-> H3O+ + CH3COO- ; K= 1.8*10^-5

H3O+ + OH- <->>>>> 2H2O ; K= 10^14

Net Equation:
HCH3COO + OH- <->>> H2O + CH3COO-
K= (1.8*10^-5) * (10^14) = 1.8*10^9 = extremely large
The problem states that I'm adding a certain volume of a known [KOH] to a certain volume of a known [HCH3COO].

The goal is to calculate the final pH.

Since I don't know the K value of the rxn of HA w/ OH-, I set up 2 equations and combined their K values to derive the K value.

Since the new K value is extremely large, I assume the reaction goes fully to completely and that all OH- (limiting reagent) reacts with the HCH3COO to provide an equivalent amount of conjugate base and an excess HCH3COO.

Next, since I find the concentrations of the acid and the conjugate base. Since they are of appreciable amounts, I can assume that the initial volumes of both change negligibly.

Finally, I use the Henderson-Hasselbach equation to find the pH.*My question is: is this a legit way to solve this problem?

Thank you.
 
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Hammad Shahid said:
Since the new K value is extremely large

Actually it is not as large as you think.

I assume the reaction goes fully to completely and that all OH- (limiting reagent) reacts with the HCH3COO to provide an equivalent amount of conjugate base and an excess HCH3COO.

(...)

Finally, I use the Henderson-Hasselbach equation to find the pH.

That's a reasonable approach

Note however, that the K value you have derived is just a reciprocal of the base dissociation constant of acetate. While acetate is a weak base, it is still strong enough for the acetate solutions to be slightly alkaline. Assumption that the neutralization went to completion works OK for a typical buffer preparation, but will fail for more exotic cases (high dilution, or attempting to neutralize more than - say - 95% of the acid).
 
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Borek said:
Actually it is not as large as you think.
That's a reasonable approach

Note however, that the K value you have derived is just a reciprocal of the base dissociation constant of acetate. While acetate is a weak base, it is still strong enough for the acetate solutions to be slightly alkaline. Assumption that the neutralization went to completion works OK for a typical buffer preparation, but will fail for more exotic cases (high dilution, or attempting to neutralize more than - say - 95% of the acid).
In my class, we're told to assume the rxn goes to completion if K is 10^10 or greater, and since this is pretty close to that, I put in that assumption.

Oh wow, I just realized that the net reaction is the acetate with water rxn in reverse. Thanks, that'll save me quite some time on the actual tests.

Alright, so I understand why for very dilute solutions the approximation with Henderson-Hesselbach equation becomes less accurate, but why for attempting to neutralise more than 95%?
 
Hammad Shahid said:
why for attempting to neutralise more than 95%?

Don't treat this number too seriously, just a rule of thumb - if neutralization goes close to completion it is better to check if the assumption was valid.
 
Borek said:
Don't treat this number too seriously, just a rule of thumb - if neutralization goes close to completion it is better to check if the assumption was valid.
Ok I was actually concerned on why the assumption may not hold for neutralizing more acid?
 
The close we get to the 100% neutralization the more of the conjugate base gets hydrolized. If you assume 99% of the acid is neutralized and ignore the fact 1% of the conjugate base reacts with water, concentration of the undissociated acid estimated this way is wrong by 100%.
 
Borek said:
The close we get to the 100% neutralization the more of the conjugate base gets hydrolized. If you assume 99% of the acid is neutralized and ignore the fact 1% of the conjugate base reacts with water, concentration of the undissociated acid estimated this way is wrong by 100%.
I'm sorry, but I don't quite understand this. Is there an example of this that could be given.
 
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