Can I combine an Acid dissociation with autoionization of H2O?

In summary: Well, for example, let's say you have a buffer with acetic acid and ammonium chloride at a concentration of 1 M. If you add 0.1 M of ammonium chloride, the buffer will become increasingly alkaline. The ammonium chloride will react with the acetic acid, and the resulting solution will be more alkaline than the original buffer.In summary, assuming the reaction goes to completion gives a more accurate K value than just using the base dissociation constant of acetate. However, if the reaction is close to completion the approximation with the Henderson-Hesselbach equation may not be accurate.
TL;DR Summary
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.

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).

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%?

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.

1. What is the difference between acid dissociation and autoionization of H2O?

Acid dissociation refers to the process in which an acid molecule breaks apart into ions when dissolved in water. On the other hand, autoionization of H2O is the process in which water molecules spontaneously dissociate into hydrogen ions (H+) and hydroxide ions (OH-).

2. Can acid dissociation and autoionization of H2O occur simultaneously?

Yes, they can occur simultaneously. When an acid is dissolved in water, both acid dissociation and autoionization of H2O will occur at the same time.

3. How does the concentration of H+ and OH- ions change during the combination of acid dissociation and autoionization of H2O?

The concentration of H+ ions will increase due to the acid dissociation, while the concentration of OH- ions will decrease due to the formation of H+ ions through autoionization of H2O.

4. What is the role of the equilibrium constant in the combination of acid dissociation and autoionization of H2O?

The equilibrium constant (K) is a measure of the extent to which the acid dissociation and autoionization of H2O reactions occur. It is used to determine the concentration of H+ and OH- ions in a solution at equilibrium.

5. How does temperature affect the combination of acid dissociation and autoionization of H2O?

Temperature can affect the rates of both acid dissociation and autoionization of H2O, which in turn can affect the equilibrium concentrations of H+ and OH- ions. Generally, an increase in temperature will lead to an increase in the concentration of H+ ions and a decrease in the concentration of OH- ions.

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