ScienceHideout
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Cleaning Sulfuric Acid- The Easy Way!
I have seen a couple threads where people have taken some dirty, brown Rooto brand H2SO4, and have cleaned it in various ways. Usually this involved
heating it to some ungodly temperature. Not only is this hard, but it is dangerous, too!
My local hardware store was closing, so they had a big clearance event. All drain cleaners 50% off... so I bought a half gallon of the rooto acid. I
had big plans. First thing I did is took an empty 2.5 L glass jug and sanded the surface with sandpaper, and then I sprayed onto it, a transparent
spray rubber. I got it nice and thick with several coats, then I sprayed some clear spray paint. I did this to a smaller bottle too. When I dropped
the smaller bottle full of water, it did not shatter, but held up rather well. I understand that H2SO4 may seep through the cracks and eat the rubber
coating, but hey, each second counts, and maybe the extra time it buys you would be enough to grab a 5 gallon bucket to put it in.
Anyways. I poured in this ugly brown H2SO4. I then added 30% H2O2, 1 mL for every 100 mL of H2SO4, so about 20 in total. I put a glass stir rod right
in the bottle and mixed it. The acid warmed up a bit, but it did not get too hot to touch, and there was no issue, even though I broke a rule about
adding things to acid. Anyways, nothing happened, I then put the cap on the bottle. After one hour, I loosened the cap, and heard a good 'hiss.' The
hiss was good because it meant the peroxide was decomposing, and I was not to be left with piranha solution (even though piranha solution is a 1:3
ratio, not a 1:100) for 3 days. It was crystal clear. I mean, it looked like I just got it from fisher. It was 100% colorless.
When you put it in your storage, I recommend leaving the cap loose for another 2 weeks or so just to ensure there is no pressure build up.
10 days later I tested it gravimetrically to find it as 94-95% H2SO4, so it is effectively 18 M. As far as the peroxide, I noticed no difference in
the properties of this than sulfuric acid. Really there was never enough H2O2 in it to be considered Piranha solution in the first place... But my
hypothesis is that when the H2O2 is mixed in and it sits, the solution oxidizes and slowly eats particles that give it a yellow color. The H2O2 then
undergoes self decomposition that leaves nothing but water, H2SO4 and the by-products of dissolving the mysterious brown stuff. 
hey, if you are reading this, I can't U2U, but you are always welcome to send me an email!
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Zyklon-A
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Good to know, thanks for the write-up. How concentrated was the acid originally, or don't you know? I have never bought drain cleaner, but instead buy
lab grade acids, perhaps I should start buying draincleaner if iI can find any. At Home Depot, I was informed that it would be illegal to sell
sulfuric acid in Texas, and I wouldn't be able to find it at any store, so I didn't look.
The black crap is probably only organic material, so once oxidized, there would be CO2, water, and because it's not likely to be just
hydrocarbons, whatever other oxidized elements may be present. Any ideas as to what else there might be?
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Texium
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Yeah, I think it must be a Texas thing, because I've never seen any acid drain cleaners for sale at any of the stores that I've gone to either. I also
have to order mine.
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ScienceHideout
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That is very true. I don't know exactly the concentration of the acid to start with, but I can make some assumptions. As you said, the by products of
the crap are CO2 and H2O. Let's assume the water that comes from dissolving the organics is negligible... because I am sure that not much water came
from that in the first place...
But I gravimetrically tested the finished stuff to 94-95%. I can't pinpoint exactly because neither the Lange's or CRC had an exact number for the
density, and it did not seem like a linear function. So let's assume 94.5%. I added to this, ~1% of the original volume of H2O2. I had about 2 liters,
so I added 20 mL of 30% H2O2. Density of H2O2 is about 1.1 g/cm3. So... that means 22 grams***. This was 30% H2O2, so that means I added the
equivalent of 6.6 g pure H2O2, and the remaining is 15.4 g of water. 6.6 g at 34g/mol is about .195 moles.
2 H2O2 ----> 2 H2O + O2
That means we are left with .195 moles of water once it has 'worn off.' .195 moles of water is 3.51 g, or 3.51 mL. So in total we added a net volume
of 19 mL of water. So, initially we had min 1892 mL x c = 1911 mL x .94 =.95 initial minimum
Minimum 1892 mL x c = 1911 mL x .95 =.96 initial maximum
Lol, all that math to tell you it was about 1% higher
EDIT: At *** this point I started singing that Taylor Swift song
[Edited on 25-8-2014 by ScienceHideout]
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Fantasma4500
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brilliant...
you could however try boiling the H2O2 further towards 80%, slowly and using a computer fan, upside down
you may be lucky to have a beaker where a fan will fit perfectly into, still surprises me every time that my computer fan fits like it was made for my
1L beaker
anyhow, really nice... i may actually try this with my H2SO4 as it seems to have a ever so slightly brownish tint, perhaps i may not, now that i get
to consider that i store it in HDPE.. perhaps another day where i get to have some neat glass container for it
potential improvement: H2SO4 past azeotropic level is known to be hygroscopic -- you could, if the hole is clean and has no acid on it use aluminium
foil, i think it would practically impossible to wrap it so perfect it would hold enough pressure to potentially shatter the bottle, perhaps just some
plastic wrap, which will again be the weakest spot, allowing pressure to escape
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ScienceHideout
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Nice input, antiswat!
Perhaps even something like parafilm would work? That has some pretty good acid resistance. Or perhaps some sort of contraption that would hold
desiccant!
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BobD1001
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Any Ace hardware should carry Rooto drain opener, which is basically pure sulfuric acid. There may be a few states with exceptions, but every one I
have ever been in here in PA has stocked it. Dirt cheap compared to sourcing online too. Mine has always been nearly crystal clear as purchased, no
need for any further purification or clarification with H2O2. I was quite shocked at how clear the two jugs I bought were, both had just the slightest
hit of yellow, barely noticeable. I've worked with lab grade sulfuric before that was much darker. Rooto gets my thumbs up!
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JefferyH
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Can anyone shed light on what is going on here? WHat is H2O2 doing and what exactly are the impurities in this drain cleaner?
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BromicAcid
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What's old is new again, using peroxide to clear up sulfuric acid was first done back in 2005 and there is plenty of information and speculation in
this thread:
http://www.sciencemadness.org/talk/viewthread.php?tid=3722
The general consensus is that peroxide + sulfuric acid = strong oxidizing agent. This solution attacks organic impurities, carbon sludge and such
that is residual from the production process (since they're not trying to make ACS grade acid here). There are also organic surfactants and
inhibitors that might be destroyed. Of course any inorganic ion contamination will stick around but may be oxidized to a higher state. In my
experience the acid became almost colorless which would tell me it is low in terms of iron and nickel at the least.
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ScienceHideout
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@Bromic Acid- That thread I saw a long time ago, and it is actually what gave me the idea to do this. However, it seems that there are two major flaws
with that procedure- Firstly, you add a large volume of H2O2 to a little bit of H2SO4, and therefore water, and you also must heat it.
Firstly, I am not one who just likes to have hot sulfuric acid in a beaker. Secondly, the method I used, yes, does take a bit longer, but it only
changes the acid's concentration by 1%! You don't waste peroxide, you don't heat, and best of all, you can do it directly in the jug that you wish to
store your acid in!
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Templar
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I have never had problems violating this "dont pour water in conc acids or bases" rule
Touch and go, take it slowly and if it doesnt do anything adverse, take it a bit faster
Although, a henry condensation was happening but I hadnt been watching the temperature of the solvent. It wasn't boiling as much as it should have
been and I foolishly only waited a few minutes before dropping in a stir bar for nucleation points.
Orange, orange everywhere. Thank god I had gloves/glasses on
Do not drop small objects into superheated fluids in flasks
I thought impurities in the rooto would be inorganic in nature and thus uneffected by piranha solution
He who fights with monsters should be careful lest he thereby become a monster. And if thou gaze long into an abyss, the abyss will also gaze into
thee.
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careysub
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The Caro's acid is unstable at room temperature, so it won't be sticking around for long.
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BromicAcid
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I had always thought it was the metal cations that caused the decomposition of Caro's acid at room temperature. A number of older literature sources
prepare the pure peroxymonosulfuric acid or the peroxydisulfuric acid at room temperature with no issues and even go so far as to purify them by
recrystalization (peroxydisulfuric acid melting at ca. 65C according to Brauer).
I did find that the amount of peroxide solution necessary to decolorize the sulfuric acid is much less than I needed in my initial trials. However it
is certainly advantageous to perform this decolorization at room temperature. But my method was still viable about 200C lower than my work distilling
the sulfuric acid 
http://www.sciencemadness.org/talk/viewthread.php?tid=23352#...
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careysub
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Quote: Originally posted by BromicAcid  | I had always thought it was the metal cations that caused the decomposition of Caro's acid at room temperature. A number of older literature sources
prepare the pure peroxymonosulfuric acid or the peroxydisulfuric acid at room temperature with no issues and even go so far as to purify them by
recrystalization (peroxydisulfuric acid melting at ca. 65C according to Brauer).
I did find that the amount of peroxide solution necessary to decolorize the sulfuric acid is much less than I needed in my initial trials. However it
is certainly advantageous to perform this decolorization at room temperature. But my method was still viable about 200C lower than my work distilling
the sulfuric acid
http://www.sciencemadness.org/talk/viewthread.php?tid=23352#... |
I'm going by what is reported here:
http://books.google.com/books?id=bUBFkv7PGr4C&pg=PA33&dq=caro%27s+acid+storage&hl=en&sa=X&ei=YQwAVLuoEs6uogT4jICgBQ&ved=0CDgQ6
AEwAA#v=onepage&q=caro's%20acid%20storage&f=false
It is a fairly complete, if short, account of Caro's acid.
In a typical solution it appears to hydrolyze slowly back to sulfuric acid, and apparently special care is needed to preserve the prepared reagent for
even a few days. It is pretty common that highly purified chemicals are more stable than otherwise.
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