jokoron
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Making H2SO4 using sodium bisulfite
I'm quite new to Chemistry so don't facepalm too hard.
This video covers how to make Sulfuric acid using sodium metabisulfite
https://www.youtube.com/watch?v=okvvD3-DF9U
Will this method work if you replace metabisulfite with bisulfite?
Also how does the cold temperature of the oxidizer effect the concentration?
Why was he using sodium metabisulfite instead of the much more common bisulfite?
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Amos
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Sodium metabisulfite dissolves in water to give essentially the same solution as bisulfite would.
Na2S2O5 + H2O = 2 NaHSO3
Sodium metabisulfite is usually much easier to obtain as a consumer product (in the U.S. it's sold as both a root killer and a sanitizing agent for
homebrewing)
Cold temperature is generally beneficial when dissolving any gas in solution due to the higher solubility and lower volatility of gases at lower
temperatures.
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AJKOER
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A home chemist will much more likely have cheap chlorine bleach (NaOCl) rather than expensive 30% H2O2.
As detailed in the YouTube, mix HCl and aqueous sodium metabisulfite to generate SO2. The latter gas reacts with H2O to form H2SO3. More directly, one
can also burn sulfur in the presence of air or pure oxygen (from the action of dilute H2O2 on NaOCl) to form SO2 gas and dissolve in water.
Next, mix HCl and NaOCl to generate Cl2, a corrosive toxic gas (perform outdoors). It poorly dissolves in water in an equilibrium reaction as follows:
Cl2 + H2O = HCl + HOCl
However, in the presence of SO2/H2SO3, the dissolution of the chlorine is much improved with the creation of sulfuric acid and HCl as the SO2/H2SO3
effectively removes the hypochlorous acid, HOCl, thereby moving the above equilibrium reaction to the right. The final reaction is given by:
HOCl + H2SO3 = HCl + H2SO4 (reference: https://www.google.com/url?sa=t&source=web&rct=j&... )
The volatile HCl can be boiled off (along with any unreacted SO2) leaving H2SO4 and recovered to generate more SO2 or Cl2.
Note, one can avoid working with chlorine gas (as I have done successfully on many occasions) by adding CaCl2 to aqueous NaOCl and treating the
solution with CO2. Let the CaCO3 precipitate settle out, decant and quickly deploy the dilute HOCl (which is, in the current context of creating
sulfuric acid, would be treating the aqueous hypochlorous acid with SO2).
[Edited on 3-12-2016 by AJKOER]
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macckone
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Ajkoer, with the method you describe, there will be substantial sodium compounds in solution, requiring distillation of the resultant sulfuric acid.
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AJKOER
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Macckone you are correct, the chlorine gas free path here is more difficult. One way to remove the NaCl from the HOCl (before treating it with SO2 as
a path to H2SO4) is to distill half of the volume of the HOCl/NaCl and discard the remaining half. Per Watts Dictionary of Chemistry (please see
comments and links in a prior SM thread at https://www.sciencemadness.org/whisper/viewthread.php?tid=17... ), the more volatile than water HOCl/Cl2O is driven largely off first, so the
concentration of the hypochlorous acid is nearly doubled and the NaCl is removed. Note, concentrated HOCl is increasingly unstable and should be
cooled, free from strong light and used promptly.
Also, the distillation of the NaCl/H2SO4 would likely drive off the volatile HCl, leaving just Na2SO4.
No easy fixes. For example, try adding dry MgSO4 and freezing out the Na2SO4(H2O)9.H2O, where the last water molecular is impure containing the
solution water, H2SO4/MgCl2. This may concentrate the acid, and raise its activity coefficent (as aqueous MgCl2, which is divalent, is more ionic than
NaCl), but still leaves a chloride presence.
The easier and purer path to H2SO4 is, as I first described, via chlorine.
[Edited on 5-12-2016 by AJKOER]
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AJKOER
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There may be another easy(?) path to oxidizing H2SO3 to H2SO4. Just pass air through aqueous SO2 in the presence of a small amount of, say, MnCl2.
Source: "Kinetics and Mechanism of the Oxidation of HSO3- by O2. 2. The Manganese(II)-Catalyzed Reaction", by Robert E. Connick and Yi-Xue Zhang,
Inorg. Chem., 1996, 35 (16), pp 4613–4621,DOI: 10.1021/ic951141i. To quote from the synopsis:
"The manganous ion is a powerful catalyst for the chain reaction of bisulfite ion with O2: 2HSO3- + O2 → 2SO4(2-) + 2H+. At relatively high
manganous ion concentrations (∼10-4 M, pH 4.5), the rate does not depend on oxygen or bisulfite ion concentration but is proportional to the square
power of the manganous ion concentration. In the mechanism for the uncatalyzed reaction a simple replacement of the propagation reactions b and c by
reaction 3 accounts for the [Mn2+]2 dependence in the three-term rate law. This term combined with the term for the uncatalyzed reaction leads to the
quantitative prediction of the third term which contains [HSO3-][Mn2+]. It is proposed that SO5•- oxidizes the manganous ion to Mn(III) and Mn(III)
is rapidly reduced back to Mn2+ by HSO3- with the formation of the chain carrier SO3•-. The ratio of the bimolecular rate constant of (3) to the sum
of the rate constants of (b) and (c) is 124. The addition of the initiator catalyst S2O8(2-) to the manganous ion catalyzed reaction confirms that
manganese enters the reaction mechanism through a propagation step."
Link: http://pubs.acs.org/doi/abs/10.1021/ic951141i?src=recsys&...
The potential problem with the method is sourcing the manganous ion by converting battery MnO2 into MnCl2. To that end, here is a collection of likely
important chemical reactions relating to processing Manganese salts, in particular, MnO, Mn2O3, and MnO2 that may be present in old batteries:
2 HCl (dilute) + MnO → MnCl2 + H2O (see http://chemiday.com/en/reaction/3-1-0-1402 )
4 HCl (conc) + MnO2 → Cl2 + MnCl2 + 2H2O (see http://chemiday.com/en/reaction/3-1-0-177 )
6 HCl (conc, heated) + Mn2O3→ 2 MnCl2 + Cl2 + 3 H2O (see http://chemiday.com/en/reaction/3-1-0-7473 )
See also this thread https://www.sciencemadness.org/whisper/viewthread.php?tid=67... .
So this method requires concentrated HCl to eventually make dilute H2SO4 with a manganese impurity.
[Edit] More recent literature suggests the more effective employment of cobalt salts in place of manganese, where the cobalt salt in lower valent
states is claimed to be active in even trace amounts. See, for example, "Cobalt-catalyzed sulfate radical-based advanced oxidation: A review on
heterogeneous catalysts and applications", by
Peidong Hu abd Mingce Long, link: http://www.sciencedirect.com/science/article/pii/S0926337315... .
Recently, a member commented on preparing cobalt salts from pottery grade cobalt carbonate (see https://www.sciencemadness.org/whisper/viewthread.php?tid=71... ).
Note, cobalt salts are toxic (see https://www.atsdr.cdc.gov/phs/phs.asp?id=371&tid=64 ) and present an environmental hazard.
[Edited on 7-12-2016 by AJKOER]
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