[cliverlong]

I feel a bit embarrassed asking this question - but I can't work it out.

If I add concentrated sulphuric acid to sodium chloride the text book says

NaCl(s) + H₂SO₄(aq) ----> NaHSO₄(s) + HCl(g)

Now no species changes its oxidation state - so I can't use redox / electrode potential to explain why the reaction occurs.

It is not adding acid to a base - as NaCl is a neutral salt - so I can't apply acid-base neutralisation.

So how can I explain the reaction?

However, I can explain the "second stage" of the sodium bromide and sulphuric acid reaction by redox potentials

KBr(s) + H₂SO₄(aq) ----> KHSO₄(s) + HBr(g)         displacement   << same problem as NaCl to explain
2HBr(g) + H₂SO₄(aq) ----> Br₂(aq) + 2H₂O + SO₂(g)   oxidation of HBr

since Br has ox state -1 in HBr and this is oxidized to 0 in Br₂ and oxidation state of sulphur changes from +6 to +4 - so I just find the appropriate half-cell / redox equation for that bit. Still can't explain the initial reaction though.


Ta

Clive

[Borek]

It is not just H₂SO₄(aq) - it must be concentrated!

[cliverlong]

It is not just H₂SO₄(aq) - it must be concentrated!

Yes, I agree. That condition is in the second line of my question.

My question still stands. WHY does the reaction occur? It's not redox, it's not acid-base. What is it?


Note, I did find this discussion

http://www.chemicalforums.com/index.php?topic=8137.msg36735

where (after a while) the two reactions are described but not explained.


Thanks again

Clive

[Borek]

HCl, while very strong acid, is not 100% dissociated, especialy in concentrated solutions of sulfuric acid, where there is abundance of H⁺ and almost no water. When you have HCl and no water, HCl gets airborne and flies away - and the reaction can proceed.

[cliverlong]

HCl, while very strong acid, is not 100% dissociated, especialy in concentrated solutions of sulfuric acid, where there is abundance of H⁺ and almost no water. When you have HCl and no water, HCl gets airborne and flies away - and the reaction can proceed.

Ahhh ... are you saying all the ions from NaCl and H₂SO₄are swishing around .. and one potential product is HCl ... Then because of the presence of vast amounts of H⁺ I need to argue in equilibrium terms ... and (say it quietly) use Le Chatelier to explain why the reaction proceeds from left to right because the HCl is generated as a gas and being constantly lost/driven out rather than staying in solution?


Ta,

Clive

[Borek]

Ahhh ... are you saying all the ions from NaCl and H₂SO₄are swishing around .. and one potential product is HCl ... Then because of the presence of vast amounts of H⁺ I need to argue in equilibrium terms ... and (say it quietly) use Le Chatelier to explain why the reaction proceeds from left to right because the HCl is generated as a gas and being constantly lost/driven out rather than staying in solution?

Looks OK to me. And you don't have to wisper Le Chatelier principle - it is very useful in such situations to help predict possible outcome or to help explain what is going on on qualitative level

[wilson]

It is not just H₂SO₄(aq) - it must be concentrated!

Does it matter if we write concentrated sulphuric acid as H₂SO_4(aq)? Must it be H₂SO_4(l)?

Ahhh ... are you saying all the ions from NaCl and H₂SO₄are swishing around .. and one potential product is HCl ... Then because of the presence of vast amounts of H⁺ I need to argue in equilibrium terms ... and (say it quietly) use Le Chatelier to explain why the reaction proceeds from left to right because the HCl is generated as a gas and being constantly lost/driven out rather than staying in solution?

Looks OK to me. And you don't have to wisper Le Chatelier principle - it is very useful in such situations to help predict possible outcome or to help explain what is going on on qualitative level

From what I am seeing so far, we are using chemical equilibrium to explain why the reaction moves foward. I just want to confirm whether this is this an acid base reaction, since it appears that sulphuric acid is acting as an acid. Chemguide confirms this: http://www.chemguide.co.uk/inorganic/group7/halideions.html#top

If I am not wrong, NaCl is not completely neutral. It is very slightly basic. The chloride ions can react with protons to form HCl, which is considered the conjugate acid. Also, sulphuric acid acts as the acid here and produces sodium hydrogen sulphate, which is the conjugate base (by looking at the anion, which had lost one hydrogen).

Thus, are we looking at an acid-base reaction here?

[cliverlong]

From what I am seeing so far, we are using chemical equilibrium to explain why the reaction moves foward. I just want to confirm whether this is this an acid base reaction, since it appears that sulphuric acid is acting as an acid. Chemguide confirms this: http://www.chemguide.co.uk/inorganic/group7/halideions.html#top

If I am not wrong, NaCl is not completely neutral. It is very slightly basic. The chloride ions can react with protons to form HCl, which is considered the conjugate acid. Also, sulphuric acid acts as the acid here and produces sodium hydrogen sulphate, which is the conjugate base (by looking at the anion, which had lost one hydrogen).

Thus, are we looking at an acid-base reaction here?

Good link! thanks !

So both my statements in original question were incorrect.

Now no species changes its oxidation state - so I can't use redox / electrode potential to explain why the reaction occurs.

Of course, looking closer. Cl has oxidation state -1 in NaCl and ox state 0 in Cl₂ - so a redox reaction does occur

It is not adding acid to a base - as NaCl is a neutral salt - so I can't apply acid-base neutralisation.

It appears from the link NaCl is very slightly basic - so we can use acid-base as an explanation.

Ta.

Clive

[Borek]

If I am not wrong, NaCl is not completely neutral. It is very slightly basic. The chloride ions can react with protons to form HCl, which is considered the conjugate acid.

Quite the opposite - if anything, it is slightly acidic

Na⁺ reacts more easily with OH^- (making solution slightly acidic) than Cl^- with H⁺ (making solution slightly basic).

But these effects are so small they can be safely ignored - pH of 0.1M solution of NaCl is around 6.98 or 6.99.

[wilson]

I always thought that NaOH is a stronger base than HCl is a strong acid. So NaCl will be slightly basic.

But if Chemguide says that H2SO4 acts as an acid in the reaction:
NaCl(s) + H2SO4(aq) ----> NaHSO4(s) + HCl(g)

Then NaCl should be the base right? At least, acting as the base.

If Chemguide is wrong (I don't know), then can we classify this reaction as a redox reaction?
But then again, I don't see any change in oxidation states.

Of course, looking closer. Cl has oxidation state -1 in NaCl and ox state 0 in Cl₂ - so a redox reaction does occur

Aren't we looking at HCl (not Cl2) here?

[wilson]

And one more thing: I think these should be the correct state symbols (can someone verify?):
NaCl(s) + H2SO4(l) ----> NaHSO4(s) + HCl(g)

H2SO4 is concentrated, so it is (l).
There is very little water, so NaHSO4 is not aqueous.

[cliverlong]

Of course, looking closer. Cl has oxidation state -1 in NaCl and ox state 0 in Cl₂ - so a redox reaction does occur

Aren't we looking at HCl (not Cl2) here?

I'll perform another U-turn  here.   

Yep, that was my original thinking. Cl oxidation state in NaCl and HCl is -1 - so no change in this reaction

So looks like a redox explanation is no good for this  reaction 

So acid-base, where NaCl is an acid , or a base, probably, maybe ...     

I'll revert to type and just "accept" it and not try to explain it    (pah, Chemistry    )

Clive

[Borek]

I would go for metathesis (double replacement). But I don't feel urge to classify each reaction

[wilson]

I would go for metathesis (double replacement). But I don't feel urge to classify each reaction

Any subcategory? Is acid-base reaction wrong?

[Borek]

If you go deep enough and you use all acid/base definitions, almost each reaction can be classified as acid/base. I won't call it acid base reaction, but treat it rather as humble opinion, not definitive statement.