dex
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Distilling H2SO4 from NaHSO4
I spent a couple of days searching the forum for a cheap and effective way of producing concentrated H2SO4 and came across an interesting post by
JJay, in which he stated that you could heat NaHSO4 to obtain H2SO4 vapors, quoting the book "Small-Scale Synthesis of Laboratory Reagents" by Leonid
Lerner.
I have seen this reaction being discussed here for making oleum but not as a cheap way to make H2SO4 for the amateur chemist. In particular, the book
states:
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Oleum can be obtained by the pyrolysis of NaHSO4, which in the fluid state is equiv-
alent to an equimolar mixture of H2SO4 and Na2SO4:
2NaHSO4 ≡ H2SO4 + Na2SO4. (equation 20.5)
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The book then further states
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Equation 20.5 turns out to be a realizable reaction in practice. The present experi-
ment shows that H2SO4 can be distilled from NaHSO4 with almost 100% efficiency.
While collecting the distillate as a single fraction gives 100% H2SO4, unlike the SO3/
H2O system, introducing a cut in the distillate fraction yields oleum, with the lower
bp fraction being therefore a correspondingly weaker acid (because the combined
fractions must give 100% H2SO4).
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If I can read this correctly, then why isn't it suggested more often as a common way of producing sulfuric acid? I may be missing something here which
is why I'm making this post before attempting anything too dangerous. Honestly, reading about it in the book makes it sound so easy, but then
absolutely no one mentions it in the usual H2SO4 talks, nor has anyone posted this reaction on youtube, so that makes me a little anxious about
attempting it by myself right now.
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Fluorite
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I think it's more about when NaHSO₄ is heated and decomposes into H₂SO₄ and Na₂SO₄, the resultant solid sodium sulfate can form a hard block
inside the glassware that would be very hard to remove. This block could also expand as it forms, potentially putting stress on the glass and causing
it to crack or break.
I was wondering if there might be alternative setups to mitigate this. For example, would it be possible to heat the bisulfate on a glass plate and
use vacuum suction connected to an inverted funnel placed above it to collect the H₂SO₄ vapors? That might avoid the risk of damaging enclosed
glassware.
Alternatively, maybe using a quartz tube reactor with a dry inert gas flow (like nitrogen or argon) to sweep the H₂SO₄ vapors could work? Quartz
can handle higher temperatures and might be less likely to suffer from the thermal stress.
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dex
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Thank you for your insight. Is the Na2SO4 product that hard to remove? Looking at Thy Labs' video on sodium sulfate it doesn't look that bad but I
don't know.
The author of the book I mentioned also wrote a lab report. I can't post the whole book here because of copyright but if you can find it I invite you
to read the entire section on SO3, that is very interesting. In particular he does use quartz glassware for heating, but he uses a simple distillation
arm, into a 2-neck RBF immersed in a cold bath with a condenser on the 2nd neck.
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bnull
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Mood: Dazed and confused.
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Lerner writes on page 181 of the same book that a "result of practical significance in the present experiment is that (...)
Na2SO4, which is solid at the maximum reaction temperature, does not attack quartz or expand on cooling", so expansion is not an
issue.
For me it is both the reaction temperature and the concentration. Unless you're in the UK, you can buy a bottle of ~35% sulfuric acid for batteries,
which is more than twice the concentration obtained from decomposition of bisulfate in the range 260 to 420 °C. In either case, bisulfate or battery,
distillation will be necessary to remove water. And dealing with sulfur trioxide is not exactly a pleasure.
By the way, could you please share the title or DOI of the report? Lerner was probably too modest to reference his own publication in his book.
Quod scripsi, scripsi.
B. N. Ull
P.S.: Did you know that we have a Library?
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dex
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| Quote: | | By the way, could you please share the title or DOI of the report? Lerner was probably too modest to reference his own publication in his book.
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Sorry for the confusion, I meant the report that is in the book. :')
Thank you for clearly laying out the downsides. I am not in the UK but I don't think you can buy battery acid in the EU anymore. It's such an
important reagent and I just want a reliable way to get as much as I want, and NaHSO4 is so cheap, I might just take the time to replicate his entire
setup with the box oven. If I do not take cuts that should avoid having to deal with the SO3 right? I'm not too sure about that, or even how to handle
SO3...
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Keras
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Quote: Originally posted by dex  |
If I can read this correctly, then why isn't it suggested more often as a common way of producing sulfuric acid? I may be missing something here which
is why I'm making this post before attempting anything too dangerous. Honestly, reading about it in the book makes it sound so easy, but then
absolutely no one mentions it in the usual H2SO4 talks, nor has anyone posted this reaction on youtube, so that makes me a little anxious about
attempting it by myself right now. |
This reaction needs blowtorches and a quartz RBF as well as some special glass pieces (a 65° bend, an air condenser…) to be conducted effectively.
There are also a lot of SO₃ fumes escaping, which makes it impractical unless you operate in a fume hood or outside.
If you want to make concentrated sulphuric acid you first have to preheat the RBF at 350/400 °C for a while in order to evaporate the major part of
the water that the reaction produces, which needs a heating mantle that can reach this high.
Really, the best solution to make sulphuric acid (diluted) is to do what we devised with NurdRage: make a concentrated solution of copper sulphate and
add the stoichiometric amount of oxalic acid. You’ll precipitate copper oxalate which is totally insoluble, and be left after filtration with
diluted sulphuric acid. Painless, harmless and odorless.
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woelen
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I tried the reaction with NaHSO4.H2O some years ago, without success.
The solid "melts" very easily, the solid NaHSO4.H2O melts in its own water of crystallization at well be low 100 C.
On further heating, the water can be boiled away, leaving behind anhydrous NaHSO4.
The next step occurs at below 300 C. More water is lost. What remains behind is solid Na2S2O7. And here the fun stops.
In order to get free SO3 from the Na2S2O7 you need insane heating. In glass, this is not possible. Maybe in quartz, but I do not have that. I heated
until my test tube became soft, but still no SO3.
I also tried with Na2S2O8. You can easily convert this to Na2S2O7 by heating, but again, at that point the fun stops.
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Keras
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I’m surprised. A single blowtorch suffices to get plenty of sulphur trioxide fumes, as shown in the attached picture (the test tube is made of
quartz).
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dex
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Thanks everyone. Great picture. I'll try to get my hands on quartz glassware. Copper sulfate + oxalic acid are more than twice as expensive as NaHSO4
from what I can see. It'll take some time but I'll report back on my results if I do it.
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Keras
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Please do. We’ll be pleased to help
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RU_KLO
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Check Nurd Rage videos regarding oleum.
https://www.youtube.com/watch?v=hUyJ6CibhSg&t=5s&ab_...
https://www.youtube.com/watch?v=wB2zzm8VP9Y&ab_channel=N...
Go SAFE, because stupidity and bad Luck exist.
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