Reaction between hydrogen halide and sulfuric acid

[yucheng]

I am studying the reaction

$$\ce{NaX(s) + H2SO4(conc.) -> NaHSO4 + HX \uparrow} \tag{1}$$

with the conditions similar to this video below

Solid halides with concentrated sulfuric acid


What reaction is this? Precipitation reaction? In other words, will sodium bisulfate be precipitated? I believe it would in the case of concentrated sulfuric acid

Does it make a difference if sulfuric acid is in its molecular form as opposed to dissociated form?
This is related to the question: what if concentrated acid was replaced with dilute acid? Apparently this works for bromides (see Handbook of Preparative Inorganic Chemistry see page 285) But isn't sodium bisulfate soluble? I'm thinking of putting table salt in dilute sulfuric acid etc. Will we get hydrochloric acid solution?

Why would this reaction even occur in the first place for both dilute and concentrated sulfuric acid? (See my post below). I hope there is a better reason than "thermodynamics".....

Thanks in advance!

 

[yucheng]

My first thought was that (1) appears to be a Bronsted-Lowry acid-base reaction, right? Why would the halides be protonated? Is this related to displacement of weak acid by strong acids? But aren't HCl, HBr, HI stronger than H2SO4?

 

[Borek]

I would classify it as a double displacement.

The driving factor here is the fact that produced hydrides of halogens are gaseous and - if you use a strong acid - they have no solution to dissolve in. In effect they quickly disperse and run away from the mixture. Thermodynamically it is the entropic factor -TΔS.

This can be easily confirmed by using less concentrated acid: there will be no reaction, just a simple dissolution.

 

[yucheng]

This can be easily confirmed by using less concentrated acid: there will be no reaction, just a simple dissolution.

I was reading this post:
https://chemistry.stackexchange.com...oncentrated-sulfuric-acid-and-sodium-bromide/
"This reaction occurs at all conditions of temperature and concentrations, even in dilute solutions, where the reagents and products are transformed into ions."

Is the above statement true, regarding reaction of NaBr and H2SO4? If it does, that this is a simple double displacement reaction is rather doubtful......

Also, it appears that reaction with dilute sulfuric acid has indeed been done:
PREPARATION of CONSTANTBOILING HYDROBROMIC ACID
https://pubs.acs.org/doi/pdf/10.1021/ed014p187
(see the second page of this paper for details, or the Handbook I cited above)

 

[Borek]

This stackexchange statement is wrong for me - writing the reaction taking place in a solution using compounds and not as a net ionic one is a nonsense. In the solution there is no reaction - you deal with ions only:

H⁺ + HSO₄^- + Na⁺ + Br^- → Na⁺ + HSO₄^- + H⁺ + Br^-

You can group these ions any way you like, but it doesn't mean anything has changed/reacted.

But: heating the solution changes the situation. Even in relatively diluted solutions of halide acids you can expect halide hydrogen to be volatile enough to evaporate. And any solution containing both H⁺ and X^- ions can be considered to be an "acid solution".

With heating we are basically back to thermodynamics and -TΔS, there is no other explanation.

That's all ignoring the oxidation of halide ions to elemental halides.

 

[yucheng]

This stackexchange statement is wrong for me - writing the reaction taking place in a solution using compounds and not as a net ionic one is a nonsense. In the solution there is no reaction - you deal with ions only:

H+ + HSO4- + Na+ + Br- → Na+ + HSO4- + H+ + Br-

You can group these ions any way you like, but it doesn't mean anything has changed/reacted.

I believe that at high enough concentration, sodium hydrogen sulfate might be precipitated out of the solution... perhaps comparing lattice enthalpies with NaBr?

What about this:
https://chemistry.stackexchange.com...ized-by-h2so4-whereas-nabr-isnt-oxidised?rq=1

The author of the answer argues that oxidation of hydrogen iodide is spontaneous, whereas that of hydrogen bromide is not spontaneous. However, in the video I've shown above, the reaction proceeds without heating. Perhaps the displacement of the halide from its salt is highly exothermic such that the temperature is just right for oxidation of hydrogen bromide to happen?

In relation to the video, as well, isn't the textbook mentioned in the question (the OP) wrong?

 

[snorkack]

What about this:
https://chemistry.stackexchange.com...ized-by-h2so4-whereas-nabr-isnt-oxidised?rq=1

The author of the answer argues that oxidation of hydrogen iodide is spontaneous, whereas that of hydrogen bromide is not spontaneous. However, in the video I've shown above, the reaction proceeds without heating. Perhaps the displacement of the halide from its salt is highly exothermic such that the temperature is just right for oxidation of hydrogen bromide to happen?

In relation to the video, as well, isn't the textbook mentioned in the question (the OP) wrong?

Well, the oxidation behaviour of sulphuric acid depends a lot on concentration, and somewhat on temperature. And while the reaction potential gives the thermodynamic limits on possible reactions, not all possible reactions will be kinetically spontaneous.
Note that the reaction of HI or HBr with H₂SO₄ will compete against the reaction with alcohols. When there is no alcohol to react with, and no alcohol or product water to dilute the acid, HBr does get oxidized (as in the experiments). But HI is so much more reactive with sulphuric acid that it gets oxidized even in presence of alcohol.