BaFuxa
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Sulfuric Acid from Oxalic Acid and Copper Sulfate: does not work
Hi,
Today I tried making myself some H2SO4 via the oxalic acid + copper sulfate method. There are few videos on the internet claiming it a viable way to
get the sulfuric.
This is totally rubbish. What you get after distilling off the water is a transparent, very acidic solution ( PH around 1) that is lighter than water. Needless to say it did not react with iron nor aluminum.
I have no idea what the heck this is. Sulfurous acid, bisulfite ? Some of you guys had any success with this approach ?
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elementcollector1
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Have you perhaps tried boiling it to concentrate it?
Elements Collected:52/87
Latest Acquired: Cl
Next in Line: Nd
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hissingnoise
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Your stoichiometry was off, perhaps ─ did you, for instance, get the expected precipitate?
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clearly_not_atara
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I thought it was supposed to be magnesium sulfate?
Are you sure the product is lighter than water, and not just lighter than you expected it to be? This seems to suggest an impurity IMO as nothing in
the rxn mixture is lighter than water nor are the decomposition products (formic acid, carbonic acid, sulfurous acid).
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SWIM
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Quote: Originally posted by clearly_not_atara  | I thought it was supposed to be magnesium sulfate?
Are you sure the product is lighter than water, and not just lighter than you expected it to be? This seems to suggest an impurity IMO as nothing in
the rxn mixture is lighter than water nor are the decomposition products (formic acid, carbonic acid, sulfurous acid). |
Magnesium sulfate is definitely discussed for this purpose on some threads here and is supposed to work.
Maybe whatever you did somehow reduced the acid by oxidizing the copper?
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Melgar
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The "lighter than water" part is patently false. How did you come to the conclusion that it's lighter than water? OP seems far too sure of his own
correctness to actually be correct, in my opinion.
The first step in the process of learning something is admitting that you don't know it already.
I'm givin' the spam shields max power at full warp, but they just dinna have the power! We're gonna have to evacuate to new forum software!
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Metacelsus
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I think you mean oxidizing the oxalic acid (to CO2) while reducing the copper. Copper(II) is very unlikely to be oxidized further.
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Texium
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Thread Moved
29-10-2017 at 08:29 |
BaFuxa
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Ok ok I am just back from the shed.
First off, as for the weight ethanol is probably contaminating the solution as I distilled some before and just lightly washed my hardware.
Then you are correct, I did a stochiometry of 50/50 and upon further research I found that the correct ratio is 4/5 oxalic acid/ copper sulfate. So
yes it was off. I will try again with the prescribed amounts.
What was really interesting though is that I added some 6% hydrogen peroxide to the solution and it did attack steel nails. I tested the steel nails with the H2O2 alone and nothing happened.
It did nothing on aluminum though, probably not concentrated enough but I think this is sulfuric acid.
[Edited on 29-10-2017 by BaFuxa]
[Edited on 29-10-2017 by BaFuxa]
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SWIM
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| Quote: Originally posted by Metacelsus |
I think you mean oxidizing the oxalic acid (to CO2) while reducing the copper. Copper(II) is very unlikely to be oxidized further.
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No, actually I just had no idea what I was talking about.
Oxidation states of metals are not my strong point.
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AJKOER
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First, known that H2C2O4 is both an oxidizing and reducing agent!
The recommended preparation of ClO2 (which is an easily explosive gas unless diluted in an inert gas) is via chlorate + H2C2O4. The reaction proceeds:
2 KClO3 + H2C2O4 = 2 HClO3 + K2C2O4 (s)
and with more excess Oxalic acid:
2 HClO3 + H2C2O4 --> 2 H2O + 2 ClO2 (g) + 2 CO2(g)
so the net reaction is:
2 KClO3 + 2 H2C2O4 --> 2 H2O + 2 ClO2 (g) + 2 CO2 (g) + K2C2O4 (s) (source: see https://books.google.com/books?id=6wUmteTIc18C&pg=PA334&... )
The action of oxalic acid on HClO3 illustrates H2C2O4 acting as a reducing agent on chloric acid resulting in a safer mix of ClO2 diluted in CO2
gases.
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The reaction of H2C2O4 with CuSO4 proceeds as follows:
CuSO4 + H2C2O4 = CuC2O4 + H2SO4
Avoid an excess of oxalic acid as it may apparently, see below, be violently reduced by when the H2SO4 becomes concentrated.
Here are some important (safety) comments from a prior thread of mine (see http://www.sciencemadness.org/talk/post.php?action=reply&... ).
| Quote: Originally posted by AJKOER | | Quote: Originally posted by Formatik | .....
Another attempt of sulfuric acid from oxalic acid and CuSO4 (beware!):
I made another attempt with the copper sulfate and oxalic acid. But this time used larger amounts. This time I only filtered and siphoned the
filtrate, and did not evaporate and collect more solids. But this time I just boiled down the filtrate. Something very bad happened on boiling near
the end, all of the sulfuric acid and contents in the 600mL beaker ejected entierly! I think the sulfuric acid reacted violently with residual oxalate
(another crystallization would have been good) and the heating might have been too high.
The purity of acid made this way should be alright for some purposes. CuSO4 has a solubility of 0.19g in 100g of 92.70% H2SO4 at 25 C (Solubilities of
inorganic and organic compounds, 2nd ed. (1919) by A. Seidell). CuSO4 should be the end-product copper salt and the white solid that was seen earlier
in the brown acid.
On another similar note, aqueous copper sulfate yields no precipitate or any reaction of note when added slowly into an excess of aqueous citric acid.
The reactivity of citrates might be the reason why there is no reaction. Oxalic acid can be boiled with nitric acid and is able to partially resist
the attack.
..........
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Please do not attempt to push this reaction to the point where concentrated H2SO4 is formed. The violent ejection on heating, per recent research,
could be from the abrupt decomposition of unreacted Oxalic acid into H2O, CO and CO2. Source: Watts' dictionary of chemistry, Volume 3, by Henry
Watts, page 649 under 'Reactions'. To quote:
"-2. On heating with conc. H2SO4 or with P2O5 it is resolved into water, CO and CO2."
Link: http://books.google.com/books?pg=PA649&lpg=PA649&sig...
[Edited on 7-8-2012 by AJKOER] |
[Edited on 1-11-2017 by AJKOER]
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MrHomeScientist
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If you post more details on amounts, procedure, how you tested the products, etc. we can be much more helpful. It's hard to give good advice when the
report is "it didn't work". Assuming you followed an established procedure, I'd agree that low concentration may explain why it doesn't readily react
with metals. Aluminum isn't a great choice for a test because of its oxide layer, which can greatly slow reactions.
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