dale_wilson69
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remove H2SO4 from H20?
If I mixed H2SO4 with H20. Are there anyway to separate the two again?
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12AX7
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H - two - O, not H-twenty.
Sure are ways, but you're kind of fighting thermodynamics here. Notice how it got hot when you mixed them? Kinda likes to be that way.
Really funny question. I mean really, it's not often we get questions like, "I pulled the pin on my grenade, any way to put it back in?" The easy
solution is not to do it, but I suppose such solutions are too obvious for today's instant-gratification society. Nobody wants to take the time and
sit down, UTFSE and learn, from others mistakes rather than their own.
(Best read in an Anthony Bourdain voice.)
Tim
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BromicAcid
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The boiling points are very different, some H<sub>2</sub>SO<sub>4</sub> will co-distill with your water, but the percentage
will be very small. Greater than 2% of your water though will remain with your H<sub>2</sub>SO<sub>4</sub> as it will form
the azeotrope and cling to the sulfuric acid, additional water may also stay without excessive heating. There are other chemical ways too, adding
barium hydroxide will precipitate out the barium sulfate leaving behind water, thereby separating the two, although you will alter the amount of water
and your sulfuric acid will be a different chemical form. There are lots of other ways too but distillation is the most straight forward.
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microcosmicus
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Distillation, recrystallization, and precipitation of the sulfate ion
come to mind.
Thermodynamics does not stop you from unmixing them, just
demands that there will be a cost paid in energy for undoing
the disorder created by making a solution.
For more information, see your favorite introductory chemistry
and/or thermodynamics text.
[Edited on 27-1-2008 by microcosmicus]
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vulture
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| Quote: |
Nobody wants to take the time and sit down, UTFSE and learn, from others mistakes rather than their own. |
Maybe that's because they have to waste time wading through your smartypants drivel in nearly every thread lately? Remove a letter from UTFSE, switch
some around and you have what you should've done because you haven't got anything useful to say.
Leave the rest to the moderators, thank you very much.
There are valid reasons why people would want to remove trace amounts of H2SO4 from water.
I'd go for precipitation with Ba and distill the water leaving some behind.
[Edited on 27-1-2008 by vulture]
One shouldn't accept or resort to the mutilation of science to appease the mentally impaired.
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YT2095
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| Quote: | Originally posted by vulture
I'd go for precipitation with Ba and distill the water leaving some behind. |
just an idea, but if you use the Hydroxide or Carbonate of Barium (not the pottery grade stuff), I was thinking that these would be Ideal salts, as
the Chloride or nitrate with make the corresponding acid, and the Oxide with make H2O2 (also not great).
[Edited on 27-1-2008 by YT2095]
\"In a world full of wonders mankind has managed to invent boredom\" - Death
Twinkies don\'t have a shelf life. They have a half-life! -Caine (a friend of mine)
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kilowatt
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Personally I'd just keep the dilute H2SO4 and get more concentrated H2SO4. If I really wanted to separate them though I would react it with copper
oxide or chloride to form copper sulfate. You can thermally decompose this to get SO3, the anhydride of H2SO4. You don't just add this *VERY*
hydrogscopic volatile anhydride into water to get H2SO4, however; it is too exothermic. Normally you add it to concentrated H2SO4 to form a partial
anhydride known as oleum. This is then carefully added to water in stochiometric quantities to form H2SO4. This way you can get 100% H2SO4 and
higher (partial anhydride), which is not really possible with distillation. The SO3 will inevitably contain impurities from corrosion in the metal
retort necessary to do this, but it can be distilled at just 45°C to get excellent purity.
Just a distillation to get to 98% H2SO4 involves boiling at over 330°C anyway, so while admittedly this is a more difficult/hardcore approach, if you
don't have a glass reflux rig that can handle the heat and you want very dry pure H2SO4 with excellent yield, and you can rig up a simple retort from
stainless steel or so, it will be the best option. Be sure you are familiar with the properties and hazards of SO3 and oleum before considering such
a method and realize you will need ~700°C temperatures for the decomposition. SO3 vapor is very noxious and the liquid will burn you with more
immediate severity than almost any other substance (besides maybe pirhana solution). It absorbs moisture with tremendous fury.
Whether or not this is remotely worth it depends on what you want, and what you can get.
The mind cannot decide the truth; it can only find the truth.
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