weiming1998
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Aluminum as a substitute for acid in SO2 generators
As everyone knows, Al3+ is a Lewis acid that is capable of partial hydrolysis in water to create H+ ions in an equilibrium. Today, by chance, my
science teacher brought up the compound aluminum carbonate when we were learning about how to work out compound formulas (I'm only in year 8). I
immediately thought that this compound would either be unstable in water/nonexistant. I looked it up on Wikipedia (lame source, I know) http://en.wikipedia.org/wiki/Aluminium_carbonate
and found out that it hydrolyzes when it touches water. I then thought of the hydrated AlCl3 thermal decomposition reaction and had a new idea. What
if we were to combine a source of aqueous Al3+ ions with a sulfite/thiosulfate?
Here is my experiment: A water solution of AlCl3/Al2(SO4)3 (made by the reaction of CuSO4, NaCl and aluminum metal, so no acidic contaminants) were
poured into a dilute Na2S2O3 solution. A first, there were no reaction, but when I heated this in a beaker, a fluffy, white precipitate of Al(OH)3 and
some yellow S precipitate appeared. There were also a sharp smell of SO2 around the beaker.
After I made this discovery, I thought that perhaps you could utilize some waste lab Al waste in a SO2 generator with a sulfite/metabisulfite (if it
worked for thiosulfate, which liberated SO2 as well on acidification, why not sulfite?). You can cut acid costs, find a use for waste, and this might
even benefit the environment (the produced Al(OH)3 is insoluble and doesn't get absorbed into water as easily).
Theoretically, this reaction will work, because of the Lewis acidity of the Al3+. So the general equation: First:
2AlCl3+3Na2S2O3===>Al2(S2O3)3+6NaCl. Then, decomposition by hydrolysis: Al2(S2O3)3+3H2O==>2Al(OH)3+3SO2+3S.
But for some reason, I cannot find references or even mentions of this reaction anywhere on the internet, the reaction between Al3+ and
sulfite/thiosulfate.
Any comments, criticism, etc are welcome.
[Edited on 30-4-2012 by weiming1998]
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barley81
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The cheapest way to make SO2 is to burn sulfur. The second cheapest way, and probably the most economical for amateurs is the reaction of sulfuric
acid and sulfur at high temperatures. I think this reaction proceeds not by the formation of aluminum thiosulfate as an intermediate, but by
hydrolysis of the hydrated aluminum ion and protonation of the thiosulfate ion which leads to decomposition of the resulting thiosulfuric acid.
The pKa of aluminum ion is ~5, while the pKa of hydrogen thiosulfate is 1.74. The pKa of thiosulfuric acid itself is 0.6. So, depending on the
equilibrium constant for the decomposition of thiosulfuric acid, the reaction might not go all the way to completion.
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Pyridinium
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Quote: Originally posted by barley81  |
The pKa of aluminum ion is ~5, while the pKa of hydrogen thiosulfate is 1.74. The pKa of thiosulfuric acid itself is 0.6. So, depending on the
equilibrium constant for the decomposition of thiosulfuric acid, the reaction might not go all the way to completion. |
I think this is true, though maybe heat would help. I actually have some AlCl3 hexahydrate and some thiosulfate handy, so you have me wanting to try
this experiment. No time to do it at the moment, maybe tomorrow.
bisulfite / metabisulfite don't need much of an excuse to let go of SO2... although the pK1 of sulfurous acid is actually pretty low also (1.85).
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weiming1998
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| Quote: Originally posted by barley81 | The cheapest way to make SO2 is to burn sulfur. The second cheapest way, and probably the most economical for amateurs is the reaction of sulfuric
acid and sulfur at high temperatures. I think this reaction proceeds not by the formation of aluminum thiosulfate as an intermediate, but by
hydrolysis of the hydrated aluminum ion and protonation of the thiosulfate ion which leads to decomposition of the resulting thiosulfuric acid.
The pKa of aluminum ion is ~5, while the pKa of hydrogen thiosulfate is 1.74. The pKa of thiosulfuric acid itself is 0.6. So, depending on the
equilibrium constant for the decomposition of thiosulfuric acid, the reaction might not go all the way to completion. |
Sometimes, for salts that produce a gaseous product when acidified, the pKa isn't anywhere near as important as in a normal situation. It all depends
on equilibrium, so if you apply heat to the mixed aluminum/thiosulfate solution, it will shift this equilibrium to the right:
2Al(3+)+3S2O3-+6H2O<===>2Al(OH)3+3SO2+3S+3H2O
The removal of products from one side will shift it towards the side with the removed products, so if the SO2 exits the solution by heat, the reaction
will continue producing SO2 until one of the reactants are exhausted.
An example of this would be the drying of hydrated AlCl3. HCl has a pKa of about -6, but with heating of the concentrated hydrate solution to dryness,
even it can be driven off as a gas. Another example is the reaction of acetic acid and thiosulfate. Acetic acid has a lower pKa than thiosulfuric
acid, but the reaction still proceeds nonetheless.
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