I think Na2SO4 does not react with any of the compound you have mentioned because all the compounds (Na2SO4, NaOH, HCl) are strong electrolytes and their ions in solution do not interact with each other to cause a precipitation or other type of reaction. Acetic acid on the other hand is a weak acid and its interaction with Na+ ion produce a strong electrolyte which is not possible in solution. that is why acetic acid do not react with Na2SO4.

Depends on how broad you define "reaction" and which conditions you apply. In aqueous solution I would say no reaction would occur. A bit hydrolysis with HCl (aq.) (-> HSO4-).

Dear Andreas Leineweber, all of the reactions are at room conditions / room temperature. Is this the equation ? 2 HCl + Na2SO4 ----> 2 NaCl + H2SO4

And in diluted aqueous solution?! You can write this equation, but it is an equilibrium reaction, which is more something formal. Look at the pKa values.....

Hi Hor, I suggest you check the thermodynamic feasibility at room temperature.... using the delta G ....

As long as you are dealing with dilute solutions, you will have a variety of different dissolved ions along with some very small level of atomic pairs (non-dissociated molecules). As Andreas Leinweber points out, you could use a series of simultaneous equations using the pKas to figure out exactly what the concentration of each species will be. Where that all adds up depends upon the pH and the concentration of the added concentrations of each species (you said room T).

The big question is whether you really care about 10^-10 or lower concentrations on non-dissociated HCl or H2SO4 and is this really a "reaction"?

Thanks for your helping, ANdreas Leinweber, Ayodele GEORGE Olumide Bolarinwa and Stephen N Smith. Actually, I have 2 cases/experiments here.

1st, I will try to titrate my sample solution (consists of pyridine , acetic acid, my sample and imidazole as catalyst) with sodium hydroxide. I worry if my sample consist of Na2SO4 as impurity, it will influence my titration result (either react with sodium hydroxide or acetic acid).

2nd, I wil try to titrate my sample solution (consists of KOH, p-hydroxybenzoic acid and my sample) with HCl. I also worry if my sample consist of Na2SO4 as impurity, it will influence my titration result (react with HCl).

No it will not affect things in either case. During the titrations Na2SO4 will effectively act as spectator ions. In a large enough concentration the ionic activity will alter the pH slightly but negligibly. pKa2 of H2SO4 is unlikely to have any effect.

You are unlikely to be able to deconvolute useful data from complex mixtures of weak acids and bases. Why are you titrating them?

I think Na2SO4 does not react with any of the compound you have mentioned because all the compounds (Na2SO4, NaOH, HCl) are strong electrolytes and their ions in solution do not interact with each other to cause a precipitation or other type of reaction. Acetic acid on the other hand is a weak acid and its interaction with Na+ ion produce a strong electrolyte which is not possible in solution. that is why acetic acid do not react with Na2SO4.

No, there is no chemical reaction appear between these compounds. The sodium salt of all your mentioned compound completely ionized in aqueous solutions. Therefore by adding Na2SO4 to any of these solution you only change the ionic strength of solution.

Sayyed is right completely. In dilute solutions Na2SO4 will not disturb your titrations due to reasons given by above.  If you are doubt, you can add a known amount of Na2SO4 and you will see its effect. I would be surpsied if you could observe any.  It is easier than modelling with various calculations.

Sodium sulfate is salt of stong acid (sulfuric acid) and strong base (NaOH) and so no one of the weak acids (acetic and hydraulic acid, both are weak compared to sulfuric acid) can remove sulfuric acid from its salts (as sodium sulfate here). Also sodium sulfate and sodium hydroxide are both can not react together because both are salts of sodium and hydroxide can not produce acid as you expect.