Keras
National Hazard
Posts: 912
Registered: 20-8-2018
Location: (48, 2)
Member Is Offline
|
|
Notes on the new NurdRage video: making sulphuric acid from sulphates and oxalic acid
The newest NurdRage's video shows an interesting way to produce sulphuric acid, namely by reacting a sulphate salt with oxalic acid. Since oxalic acid
is way weaker than sulphuric acid, one expects only a minute amount of sulphuric acid to be made, but the fact that most metal oxalates are insoluble
displaces the equilibrium towards the right and, in fine, leads to an almost quantitative reaction (IMHO).
NurdRage tested the reaction with iron sulphate, which might not have been the best idea: someone on the Patreon thread remarked that iron (II)
sulphate is easily oxidised into iron (III) sulphate, whose oxalate is much more soluble, and that impacts the yield negatively. I therefore suggested
to use copper (II) sulphate instead, which is widely available, does not risk over-oxidation (This reaction is even quoted verbatim on Wikipedia).
Furthermore, copper (II) oxalate is given as ‘insoluble’ in water, even more so than iron (II) oxalate.
NurdRage agreed, and, as you can see in his video, while the reaction did happen as expected, he was unable to filter out the resulting copper (II)
oxalate which seems to crystallise in very fine powder, and went all through his fritted funnel.
So I decided to retry his experiment, changing a few things.
To 20 mL of water in a beaker is added 6 g (24 mmol) of copper (II) sulphate. This is well under the solubility at 25 °C, but I wanted to keep
leeway. When the copper (II) sulphate had dissolved, I directly added 3.8 g of oxalic acid dihydrate (30 mmol, a slight excess). The solution
immediately turned from a medium blue to a light blue – copper (II) oxalate has a somewhat fascinating blue colour, both light and saturated, almost
fluorescent. (See picture.)
[NurdRage dissolves the oxalic acid in water and mix both solutions, but that leads to pointless dilution of the result.]
After five minutes of stirring at room temperature I set up the vacuum filtration by using a fritted funnel into which I put a small filtering disk
made of glass fibres, and atop of it a not-too-thick layer of diatomaceous earth (Celite). (See picture). I then cautiously tipped the contents of the
beaker into the funnel so as to avoid to disturb the layer of Celite, and turned the vacuum pump on. To my delight, all the copper oxalate remained
stuck over the Celite, while I was able to collect 20 mL of a limpid, transparent liquid in my filtering flask. (See pictures). I will have to titrate
the resulting solution tomorrow. If the reaction is quantitative, it should result in 24 mmol of sulphuric acid into 20 mL of water, or a 1.2 M
solution.
In any case, this process seems to work fine, provided extra mesures are taken to filter out the resulting mixture.
   
EDIT: PS. Since boiling water can dissolve around 2 kg of copper (II) sulphate per litre, it is theoretically possible to produce a 8 M solution of
sulphuric acid, which is slightly under 50% w/w.
But thinking about it, this process could even produce much more concentrated sulphuric acid. Since the reactants are consumed, water is liberated to
dissolve more, so, in theory, there could be more copper (II) sulphate and oxalic acid put in than a given amount of water can dissolve. Assuming
that:
1. Copper oxalate is insoluble in sulphuric acid, whatever its concentration.
2. There is no reverse reaction.
But that seems somehow highly suspicious.
P.PS: If some copper oxalate makes it to your fritted disc, you can remove it using 23% HCl. Copper oxalate will react to some extent with
concentrated HCl to form yellowish soluble CuCl₂ that you can then wash out.
[Edited on 23-6-2024 by Keras]
|
|
|
Sulaiman
International Hazard
Posts: 3704
Registered: 8-2-2015
Location: 3rd rock from the sun
Member Is Offline
|
|
Nice experiment. Worth trying myself.
Would a coagulant/flocculent such as potassium alum help?
CAUTION : Hobby Chemist, not Professional or even Amateur
|
|
|
Keras
National Hazard
Posts: 912
Registered: 20-8-2018
Location: (48, 2)
Member Is Offline
|
|
Nurd tried a flocculent, IIRC, but it didn't help in the slightest.
By the way, the video is here.
[Edited on 23-6-2024 by Keras]
|
|
|
Alkoholvergiftung
Hazard to Others
Posts: 180
Registered: 12-7-2018
Member Is Offline
|
|
I ve made nitric acid with Calciumnitrate and oxalic acid. I ve disolve the oxalic acid (300g in 300ml H2O) max concentration. I ve found when the
calciumnitate solution is at roometemperature the mud can be much easier filtered if both a hot there goes a lot throught the filter paper. Maybe the
same goes for Copper sulfate and oxalic acid.
[Edited on 23-6-2024 by Alkoholvergiftung]
|
|
|
Keras
National Hazard
Posts: 912
Registered: 20-8-2018
Location: (48, 2)
Member Is Offline
|
|
The problem with filter paper is that it will get destroyed by the acid if it is too concentrated.
|
|
|
digga
Harmless
Posts: 43
Registered: 11-6-2018
Member Is Offline
|
|
If the acid is too strong for the filter paper, perhaps centrifuging the mixture will suffice. I have this mental image of swing a roped bottle like
a slingshot. Better have a really good knot.
[Edited on 24-6-2024 by digga]
|
|
|
Twospoons
International Hazard
Posts: 1325
Registered: 26-7-2004
Location: Middle Earth
Member Is Offline
Mood: A trace of hope...
|
|
If the two solutions were poured in carefully as two layers (denser one on the bottom), then allowed to diffuse together, would that make larger
particles of copper oxalate? It would be slower, of course, but worth the wait if separation is easier.
Helicopter: "helico" -> spiral, "pter" -> with wings
|
|
|
Precipitates
Hazard to Others
Posts: 134
Registered: 4-12-2023
Location: SE Asia
Member Is Offline
Mood: Acid hungry
|
|
Quote: Originally posted by Keras  |
The problem with filter paper is that it will get destroyed by the acid if it is too concentrated. |
Filter paper just may not be fine enough to adequately filter the solution. I had problems with silver phosphide before - eventually
achieving separation by allowing it to settle out in the fridge for over a week! I'm not sure if the same would occur for copper oxalate.
Interestingly, I later got around this problem by performing a hot precipitation - the particles clumping together into a large airy mass.
Perhaps performing the reaction while hot, or some other alterations in conditions may help here.
I would try it, but all my copper sulphate has been converted to anhydrous and will be used to dry formic acid.
|
|
|
Keras
National Hazard
Posts: 912
Registered: 20-8-2018
Location: (48, 2)
Member Is Offline
|
|
So I did the titration I talked about in my first message. I was not aiming for a five decimals figure, but rather just a ballpark.
I measured 5 mL of the solution obtained after filtration, added a few tiny flakes of phenolphthalein prepared 20 mL of 1 M sodium hydroxide solution
and… I used around 11 mL of sodium hydroxide to neutralise the solution. If I had had 100% 1.2 M sulphuric acid, I would've used 12. So, if I’m
not mistaken (I’m going to redo the measure), the yield is around 11/12 or 91%. Probably a bit lower since there was an excess of oxalic acid. I
would say 85%+, which more or less jibes well with the visual absence of blue colour in the filtrate, indicating almost all the copper ions have been
chelated.
If this figure is true, it is definitely much higher yielding than the iron sulphate way.
[Edited on 24-6-2024 by Keras]
|
|
|
clearly_not_atara
International Hazard
Posts: 2793
Registered: 3-11-2013
Member Is Offline
Mood: Big
|
|
What was wrong with magnesium sulfate? Shouldn't that be massively cheaper than copper, and with no disposal issues?
|
|
|
Keras
National Hazard
Posts: 912
Registered: 20-8-2018
Location: (48, 2)
Member Is Offline
|
|
Nurd couldn’t make it work either. He got < 5% yield.
|
|
|
![[*]](images/xpblue/default_icon.gif){kind=icon}
