Acids from salts

[chem1a]

Just a quick q

 

is there any reason all the posts here (bar one) refer to the use of H2SO4 when making acids (eg nitric and ethanoic) from salts rather than any other strong acids.

 

Is HCL less effective ? or is H2SO4 cheaper /more available where most of you are (H2SO4 is at least a 40 min drive from where i am for 5mol acid compared to HCL which is in every hardware and pool store)

 

HCL appeared to be very effective in the production of boric acid from Borax...

 

dissolved 100g borax in massive excess of HCL (200 ml) & heated to 60c (not all dissolved).

 

Diluted acid borax mix with a futher 100ml H20 & kept at 60c then filtered.

 

Refluxed resultant filtrate with ethanol & burnt to accquire a green/orange/blue flame.

 

to confirm that the solid produced was indeed boric acid plain borax was refluxed with ethanol and burnt - only a blue/orange flame was obtained confirming suspicions that Triethyl borate had been created in the first refluxing step thereby confirming that Boric acid existed.

[John Cuthber]

Nitric acid reacts with HCl which pretty much rules it out.

Often the scheme is to add an acid to a salt then distill out the acid. There needs to be a reason why the reaction goes in the right direction ie the way you want. With sulphuric acid (boils about 300C) you can remove the acid you want eg HCl, acetic, nitric by distilling it. If you were to mix HCl with sodium acetate you would certainly get acetic acid produced. However if you were to try to distill out the acetic acid at least some of the (volatile) HCl would also come over. With H2SO4 that's not going to happen.

H2SO4 is also cheap (industrially, rather than from the shop up the road) and a strong acid.

Phosphoric acid might work for some of these reactions too- it's available and cheap. It's not as strong an acid but that doesn't matter so much if you can remove the product by distillation.

Unfortunately hot H3PO4 kills glassware.

[raivo]

Phosphoric acid might work for some of these reactions too- it's available and cheap. It's not as strong an acid but that doesn't matter so much if you can remove the product by distillation.

Unfortunately hot H3PO4 kills glassware.

 

As far as i know phosphoric acid will work for all common volatile acid synths but as you said, it will ruin glassware. I read from old book that oxalic acid can be used instead of sulfuric acid for such reactions. It is nonvolatile and reacts very slowly (if at all) with boiling HCL or HNO3. Yields will probably be low as oxalic acid is rather weak acid. I have not tested this myself so maybe it does not work at all.

[chem1a]

Ahhh .. that makes so much sense ! lol i cant believe i didnt think past the actual reaction to the isolation of the produced acids....

 

good to know H2SO4 doesnt boil away until higher temps - that makes conc H2SO4 much more accessible (i will make that 40 min drive now !)

 

I will pull out my jar of NaCH3COO & distill away ! ... ive been looking forward to seeing solid GAA

 

thanks all for your input

[woelen]

For preparation of some acids, H3PO4 is the preferred acid. HBr and HI cannot be made in a pure state from bromide/iodide salts and H2SO4. With bromide, some Br2 and SO2 come over as well, but the majority still is HBr. With iodide, the method of adding H2SO4 is totally useless. The concentrated acid oxidizes the iodide to iodine, and the acid itself is converted to a dirty mix of SO2, S and even H2S.

 

With H3PO4, which is non-oxidizing, even at high concentration, one can obtain quite pure HBr and even HI.

[budullewraagh]

I'll vouch for the reaction of KI and H2SO4 yielding H2S. That was a botched experiment on my part, but hey, I was really young and naive.

 

As for acid production from salts and other acids, the rule is to use acids that are stronger than the conjugate acid of the salt used. H2SO4 is both cheap and the strongest common acid, so we can make any organic or mineral acid from the reaction of it and a salt. As Woelen said, however, there are some exceptions concerning the oxidizing power of H2SO4.