Common Ion effect for Acid/Base systems

[SN2reaction]

Hey guys, here's a question from the PR science workbook that I'm kinda confused about.

It states: "Addition of sodium acetate to a solution of acetic acid ill cause the pH to:

A) remain constant because sodium acetate is a buffer.
B) remain constant because sodium acetate is neither acidic or basic.
C) decrease due to the common ion effect.
D) increase due to the common ion effect."

The answer is D, that pH will increase, because adding a base to such a system would increase the basicity.

My question is, why doesn't LeChatelier's principle affect the pH?

Can't you write the equilibrium for aqueous acetic acid as follows:


CH3COOH + H20 <--> H30+ + CH3COO-

Now if you add sodium acetate to this solution, wouldn't the law of mass action imply that the reactants will be favored, and so the excess sodium acetate would be converted to acetic acid, thereby decreasing the pH?

Any help would be GREATLY appreciated!

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[Rabolisk]

Your reasoning is correct. However, pH is only determined by [H+], nothing else. Le Chatelier's principle would favor acetic acid and decrease H+ and acetate ion.

 

[SN2reaction]

of course! thanks!!

 

[TieuBachHo]

The molar solubility of acetic acid does not change eventhough the amount of acetic acid appreciates per Le Chatelier's principle. The common ions (acetate ions) increases as being added. Thus, the solution will be more basic than acidic as you would think. So pH would increase (NOT decrease).

 

[Rabolisk]

Molar solubility has no meaning with regards to a solution of acetic acid, since we are not talking about a salt dissolving in water, but whether an aqueous solution of an acid dissociates or not.

 

[TieuBachHo]

Molar solubility has no meaning with regards to a solution of acetic acid, since we are not talking about a salt dissolving in water, but whether an aqueous solution of an acid dissociates or not.

Then how would you justify the changes in pH if the solution reaches its equalibrium? If acid dissociation constant does not change, then pH would stay the same, am I right? However, additional sodiums in the solution would accomodate the dissociated negative-acetate ions right? So pH would increase, is this reasoning wrong?

 

[Rabolisk]

First, molar solubility has do with whether a substance dissolves in a particular solvent. Acetic acid is dissolved without regards to its dissociation. All aqueous weak acids work the same way. That's why I said that molar solubility has no meaning in this reaction. Of course, molar solubility of the salt sodium acetate does matter, and it is considered to be infinite. But you said molar solubility of acetic acid, which is a meaningless concept.

Second, the acid dissociation constant does not change. Equilibrium constants only change as a result of temperature change. What changes is the concentration of the particles. Acetate ion concentration increases when sodium acetate is added. To counteract this increase, H+ and acetate anion react to form acetic acid, thus increasing the concentration of acetic acid, and decreasing [H+], and increasing pH.

I'm not entirely sure what you mean by additional sodiums accommodating acetate ions, but I think what you mean and what I say are the same thing. Your reasoning is correct; I was just correcting a couple of wrong things you said.

 

[TieuBachHo]

OK ... I see what you are saying now. Nevermind. I just look at the O.P. question about a decrease in pH which was wrong.

 

[Rabolisk]

I said [H+] decreases, not pH decreases. The OP also said that pH would increase, which I agreed with. Dissociation constant is different from acid/base concentration. Ka = [H+][CH3COO-]/[CH3COOH], which is always constant at constant temperature. Hendersen-Hasselbach equation says that pH = pKa + log[base]/[acid]. Obviously [base]/[acid] can change without changing Ka, which is how that equation works. If change in [base]/[acid] produced a change in Ka, and thus pKa, that equation would not be useful. Don't confuse acid dissociation constant with % dissociated.

 

[TieuBachHo]

Here is the OP question and I quote, "Now if you add sodium acetate to this solution, wouldn't the law of mass action imply that the reactants will be favored, and so the excess sodium acetate would be converted to acetic acid, thereby decreasing the pH?"

PS: I apologize for stealing your thread LOL.

 

[Rabolisk]

I posted that before I read your edited post before that. I definitely see how you could have seen that I said the opposite. Nevertheless, I think this was helpful to the OP as well as anyone wondering about common ion effect and acidity.