[yene]

Hi guys,

Consider a situation where a weak acid (H₃PO₄ in my case) is used to neutralize a diluted strong basic solution (NaOH in water).

So I know that the H₃PO₄ partially dissociates into H+ and H₂PO₄(knowing the molarity of H₃PO₄ and Ka of the reaction, I know that 14.7 M of H₃PO₄ dissociates into 0.3M of H+).

So here's my question: once the OH- ions are neutralized by those H+ ions, does the H₃PO₄ in the solution dissociate again in order to maintain equilibrium within that dissociation reaction? Since the new solution will have H₃PO₄ ions and no H+ ions, I'm wondering if there's a sort of 'Le Chatalier' effect that makes the H₃PO₄ 're-dissociate' into some more H+ ions.

This is a concern to me because I don't want to add enough H₃PO₄ to neutralize the OH- solution and then have it dissociate into even more H+ ions, thus giving me an overly acidic solution.

Any help is super appreciated.

Thanks guys.

[yene]

Perhaps the easier question is this:

If i have an excess of [OH-] ions in a solution, and use a weak acid to neutralize it, will the weak acid eventually completely dissociate?

My logic is the following:
At first, the weak acid would only partially dissociate, but all of the dissociated H+ ions would get totally consumed by neutralizing OH- ions. So to maintain equilibrium, the weak acid would partially dissociate again, and again all of the dissociated H+ ions would get totally consumed, and this cycle would continue until the weak acid is totally dissociated.

[chenbeier]

Phosphoric acid has 3 dissocation steps.

H₃PO₄ to H₂PO₄^- to HPO₄^2- to PO₄^3-

So first Phosporic acid will neutralized to the first step dihydrogen phosphate, if all H⁺ consumed it will  neutralize to hydrogen phosphate and finally to phosphate.
Regarding the pKs values will be phosphate in alkaline pH.

[Borek]

All forms of the dissociated acid are always present in the solution in equilibrium concentrations.

1M phosphoric acid contains approximately 0.92 M H₃PO₄, 0.081 M H₂PO₄^-, 6.3×10^-8 M HPO₄^2- and 3.5×10^-19 M PO₄^3-. Neutralizing the acid will change these values shifting the maximum fraction towards higher charge anions., around pH 13-14 the situation will reverse - solution will be dominated by PO₄^3- and H₃PO₄ will be present in 10^-19 M concentrations.

[yene]

Thanks for those replies, but in response:

The Ka for H2Po4 dissociation and HPO4 dissociation is way lower than that of the H3PO4 dissociation, therefore my understanding is that those subsequent reactions are negligible in this situation regarding Hydrogen production.

So just neglecting those subsequent reactions for now: in a solution with excess OH-, does the weak acid eventually fully dissociate?

[chenbeier]

In excess of NaOH you will have mostly PO₄^3-. Borek wrote the concentration already.

[yene]

Thanks for the reply. so is the answer "yes"? Does the weak acid eventually fully dissociate?

Also, Borek said that in the pH of 13-14 range we will see mostly phosphate ions.
I guess I should have clarified. In my situation, I have excess of sodium hydroxide, but not very much. It's a very diluted NaOH solution with pH of 10-11.

[Borek]

The Ka for H2Po4 dissociation and HPO4 dissociation is way lower than that of the H3PO4 dissociation, therefore my understanding is that those subsequent reactions are negligible in this situation regarding Hydrogen production.

So just neglecting those subsequent reactions for now: in a solution with excess OH-, does the weak acid eventually fully dissociate?

You can't neglect them. Concentrations of all forms depend on the pH, this in turn depends on the amount of base added.

To get exact concentrations you will need to solve full set of equilibrium equations (dissociation steps, mass and charge balance). For some acid/base combinations approximations can be used, but they are never trivial if you don't know how to solve whole system.

Alternatively, there is a software that allows such calculations: https://www.chembuddy.com/?left=BATE&right=pH-calculator

[yene]

For my system, I am not looking to bring the solution to an exact pH. I'm just looking to bring the solution to a pH somewhere within 5.5 and 9.5.

So for my application, do I really need to consider these subsequent reactions? Especially if the solution is never expected to get above pH of 12 or so?

And just for my education: if we had a weak monoprotic acid (i.e. that didn't have subsequent reactions), like nitrous acid, would all of the hydrogen be eventually dissociated and consumed in a neutralization with NaOH?

Thanks.

[Borek]

You will always have some equilibrium concentration of the protonated acid, no matter what the pH is. It can be so low it is negligible for most applications, but it will always be there. You don't need to worry abut it if all you need is a given pH range.

Just neutralize first proton completely and the second to 50%, you will get a buffer with pH=pKa2.