Chemical Forums
Chemistry Forums for Students => High School Chemistry Forum => Topic started by: cliverlong on May 16, 2008, 08:15:29 AM
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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
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It is not just H2SO4(aq) - it must be concentrated!
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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
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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.
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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
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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 :)
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It is not just H2SO4(aq) - it must be concentrated!
Does it matter if we write concentrated sulphuric acid as H2SO4(aq)? Must it be H2SO4(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?
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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
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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.
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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?
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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.
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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 ... 8) ;D
I'll revert to type and just "accept" it and not try to explain it ??? (pah, Chemistry >:( )
Clive
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I would go for metathesis (double replacement). But I don't feel urge to classify each reaction :)
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I would go for metathesis (double replacement). But I don't feel urge to classify each reaction :)
Any subcategory? Is acid-base reaction wrong?
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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.
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If you go deep enough and you use all acid/base definitions, almost each reaction can be classified as acid/base.
Which supports my (not well thought out) opinion, that the Lewis definition of acids and bases is so broad that it doesn't really help classify or categorise or distinguish the interaction of various species.
Why categorise reactions? Because you can then give a general description of a type of reaction. e.g. acid as a proton donor reacting with metal carbonates. Then maybe (just maybe) predict what the result of adding chemical A to chemical B might produce. Rash thoughts, I realise.
Also reaction mechanisms are all the rage in EdExcel A-Level chemistry and this acid/salt reaction was another one to add to the collection.
I have an aversion to learning pages of reactions which may be due to laziness. I recognise that High School / A-Level chemistry probably focuses on elements that exhibit reasonably "regular" behaviour and that "most" chemistry may have to dealt with on a case-by-case basis. I'm just trying to contain the amount of "unconnected" learning I have to do.
Clive
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Classifying the reactions is not an explanation of why the reaction occurs. I agree with Clive that classification is a tool to help us predict what reactions will occur. In the case of these reactions I would argue prediction is far more difficult than rote memorization. Read on...
A question about the reactions of concentrated sulphuric acid with sodium halide salts cropped up in one of the CIE A2 exams. This question expected students to use the electrochemical series to predict what would happen in each reaction.
According to the mark scheme the reactions are as follows:
Eqtn1: NaCl + H2SO4 :rarrow: NaHSO4 + HCl
Eqtn2: 2NaBr + 3H2SO4 :rarrow: 2NaHSO4 + 2H2O + SO2 + Br2
Eqtn1 is not redox, but eqtn2 is because Br- (-1) ions are oxidized to Br2(g) (0) and sulphate ion (+6) are reduced to sulphur dioxide (+4)
Students are supposed to use the electrochemical series to work this out, however it doesnt work. The standard electrode potential for the reduction of Cl2 to Cl- = +1.36V and for the reduction of Br2 to Br- = +1.07V. Both of these values are more positive than the electrode potential for the reduction of SO42- to SO2 which = +0.17V. The electrode potentials would lead students to predict that neither Br- ions or Cl- ions present in the salts would be capable of reducing the sulphate ion! I puzzled over this until I realized that neither of the reactions are under standard conditions. The extremely high concentration of sulphate ions make them more subseptible to reduction- however, it is only the Br- ions that are capable of reducing them since Br2/Br- has the less positve standard electrode potential than Cl2/Cl-.
I don't understand how anyone could predict that the concentration is high enough to make the reduction using bromide ions thermodynamically favourable but not chloride ions. I suspect you just have to memorize the reactions then pretend that you used the electrochemical series!!! :'(