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Author: Subject: H2O2 from sodium percarbonate
BASF
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[*] posted on 1-10-2003 at 06:51
H2O2 from sodium percarbonate


sodium percarbonate = Na2CO3*H2O2
Many commercial bleaching mixtures you can buy in any super market consist of up to 30% sodium percarbonate.
neutralize the carbonate with vinegar or NaHSO4 - draincleaner.
Some EDTA-containing soap added... voilà: a stabilized solution of dilute H2O2.

Not the high purity :D , but really OTC.

[Edited on 1-10-2003 by BASF]




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rikkitikkitavi
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[*] posted on 10-10-2003 at 09:19


how stable would H2O2 be in a Na2CO3 solution?

since H2O2 decomposition is catalyzed by OH- , would the solution have a very short shelf life? Almost faster than you can neutralize?

and if so, wouldnt it be possible to vacuum distill it instead?

EDTA for complexbinding heavy metals sounds like a good idea, even if detergents dont contain a lot of them. (discolouring the clothes)

/rickard
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[*] posted on 6-1-2004 at 16:26
Chemical Pseudonyms-PerCarbonate;HypoChlorate,etc


My understanding of Sodium Percarbonate is a Sodium Atom per Carbonate which is CO3.There is no H2O2 attachment or bond.
Just like sodium hypochlorite.Sodium with an excess of chlorine atoms.The excess of atoms comes from the oxygen bond to the sodium which allows chlorine molecules to bond to the Na-O bond




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[*] posted on 6-1-2004 at 23:19


the PER in Percabonate is for Peroxy which mean a Peroxo Bridge or -O-O- bond. Like in Peracid (R-COOOH)
<pre style="font-family: FixedSys, MonoSpace; color: black;">
O-O-H
|
R-C=O
</pre>

[Edited on 7-1-2004 by Blind Angel]




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[*] posted on 7-1-2004 at 02:17


Quote:

My understanding of Sodium Percarbonate is a Sodium Atom per Carbonate


LOL! ROFLMFAO!
And I suppose Bicarbonate is carbonate that isn't sure about its sexuality...

I would be a little more careful with my use of the word 'understanding' if I were you...
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[*] posted on 9-1-2004 at 14:07


Blind Angel:
<blockquote>quote:<hr>the PER in Percabonate is for Peroxy which mean a Peroxo Bridge or -O-O- bond. Like in Peracid (R-COOOH) <br />
<pre style="font-family: FixedSys, MonoSpace; color: black;">
O-O-H
|
R-C=O
</pre><hr></blockquote>it's not a salt of percarbonic acid. it's like perhydrates with hydrogen peroxide instead of water. sorry if I got your post wrong




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[*] posted on 27-1-2004 at 10:28


i know h2o2 feezes at -35 celisus, what happens if i put percarbonate +water in a freezer?
will i get h202 liquid form?
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[*] posted on 27-1-2004 at 15:37


No, Percarbonate has H2O2 bonded to it. Do I get water when I have Copper Sulfate above zero degrees. I don't think so.
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[*] posted on 6-8-2006 at 17:48


Sodium Percarbonate (Na2CO3)2(H2O2)3 a detergent used for cleaning wood
decks and concrete masonry is one third by weight Hydrogen Peroxide and
readily available at reasonable cost.
http://www.chemistrystore.com/sodium_percarbonate.htm
30% Muriatic acid and Urea are even more so.
If one neutralizes by titration HCl with Urea to obtain a solution of Urea Hydro-
chloride, then adds to this the Sodium Percarbonate, that should yield a salt
solution of NaCl with considerable fizziness from all the CO2, and a precipitate
after standing and cooling of the desired product Urea Peroxide addition compound.

Because the peroxohydrate is alkaline (pH 10-11 in solution) and hydrogen peroxide
is unstable in alkali, the hydrogen peroxide tends to decompose as it is liberated.
Countering this is the acidic nature of the Urea Hydrochloride. The only question
here really is how much yield could be expected.

.
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[*] posted on 6-8-2006 at 20:11


Quote:
Originally posted by franklyn
Sodium Percarbonate (Na2CO3)2(H2O2)3 a detergent used for cleaning wood
decks and concrete masonry is one third by weight Hydrogen Peroxide and
readily available at reasonable cost.
http://www.chemistrystore.com/sodium_percarbonate.htm
30% Muriatic acid and Urea are even more so.
If one neutralizes by titration HCl with Urea to obtain a solution of Urea Hydro-
chloride, then adds to this the Sodium Percarbonate, that should yield a salt
solution of NaCl with considerable fizziness from all the CO2, and a precipitate
after standing and cooling of the desired product Urea Peroxide addition compound.

Because the peroxohydrate is alkaline (pH 10-11 in solution) and hydrogen peroxide
is unstable in alkali, the hydrogen peroxide tends to decompose as it is liberated.
Countering this is the acidic nature of the Urea Hydrochloride. The only question
here really is how much yield could be expected.

.


I have done something similar. A long time ago I made a thread about "manganese tie dye". I know that if you add Sodium Percarbonate to an organic acid like acetic or oxalic acid you will have Hydrogen Peroxide in solution. I don't know how you would extract the peroxide though....

[Edited on 7-8-2006 by DeAdFX]
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[*] posted on 14-8-2006 at 22:08


"I have done something similar. A long time ago I made a thread about "manganese tie dye". I know that if you add Sodium Percarbonate to an organic acid like acetic or oxalic acid you will have Hydrogen Peroxide in solution. I don't know how you would extract the peroxide though...."

are you sure that peracetic acid will not form?




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[*] posted on 15-8-2006 at 16:14


Peracid formation from peroxide and a weak acid is in equilibrium. To form the peracid, a strong acid is usually also added.
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[*] posted on 24-8-2006 at 01:45


Sorry for bad english :(

There is a way to get pure H2O2 from commercial percarbonate. First neutralise carbonate part with acid, after that add soluble barium salt solution + NH3 solution, BaO2*8H2O will preticipate:

xH2O2*yNa2CO3 + yH2SO4 = xH2O2 + yNa2SO4 + CO2 + H2O
xH2O2 + xBaCl2 + 2xNH3 => xBaO2 + 2xNH4Cl

BaO2*8H2O is filtered off, washed with ice cold water (to remove traces of soluble salts) and heated until no more water fumes go off, resulting in anhydrous barium peroxide.

BaO2*8H2O => BaO2 + 8H2O

Anhydrous barium peroxide is mixed with sulfuric acid to release pure H2O2:

BaO2 + H2SO4 = BaSO4 + H2O2

In such way you can get H2O2 of any concentration, it depends only from concentration of sulfuric acid taken. If you'll get pure 100% H2SO4, result will be pure 100% H2O2. Remember to add some stabilizer to make H2O2 to be stable at storage, for example add some polyphoshate or any other reagent whitch forms stable complexes with heavy metall contaminants.
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[*] posted on 24-8-2006 at 03:55


Engager:

You seem to forget that bariumsulfate and carbonate are prone to form in the solution too. For the sulfate, it's solubility in water is low enough to be used as a contrast medium for x-ray analysis of the gastrointestinal tract, while most other barium salts are considered toxic or harmfull...

I've tried the synthesis of TCAP and HMTD with neutralized percarbonate solutions a long time ago already. It really becomes a mess though. I used 10% hydrochloric acid to neutralize, but a small spoonfull of this percarbonate goes a long way!

I got something like halve a gram of TCAP after 3 days of reaction, and from almost 50 grams of percarbonate in 300 ml HCL IIRC. Vast quantaties of NaCl precipitated from the cold during these 3 days and the resulting TCAP was heavily contaminated as was evident from the bright yellow light the TCAP emitted upon ignition...

I had no luck using nitric acid to neutralize the carbonate for the formation of any HMTD.


Like I expected, the presence of all these extra cations and anions from the neutralization interferes too much with the reaction mechanism to have any decent yields. Destiling the neutralized percarbonate solution is probably the only way to obtain H2O2 completely OTC...

[Edited on 24-8-2006 by nitro-genes]
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[*] posted on 24-8-2006 at 03:58


You are right, just use HCl or HNO3 for example they form soluble barium compounds :) And you can also use Sr/Ca salts instead of Ba.
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[*] posted on 24-8-2006 at 04:16


using barium for this is a total waste of time - if you wanted peroxide just oxidize the barium itself and use sulfuric (the old way of making peroxide) - but since the question is getting hydrogen peroxide from bicarbonate - the answer is no - not gonna happen - from what i understand h2o2 decomposes instantly in basic solutions - there is no way to isolate the peroxide ions from the hydroxyls - now perborates and persulfates are a totally different matter since they are not basic..
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[*] posted on 24-8-2006 at 04:48


Don't be so sure... If you use concentrated strong acids to neutralize the percarbonate solution there is quite some H2O2 present and the resulting solution can give the white coloration of the skin, just like 10%+ H2O2 solutions can do also!
I mean, if the H2O2 would instantly decompose as soon as the percarbonate goes into solution, there wouldn't be much use for it, would there? IIRC, The package says that the temperature for the percarbonate to become active is something like 40-50 deg C, so these solutions are actually quite stable...
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[*] posted on 24-8-2006 at 06:56


Nitro-genes is right. I have a pound of sodium peroxocarbonate and when I dissolve this in acid, then hardly any H2O2 in it decomposes. Even, when dissolved in plain water, the H2O2 remains present for quite some time, although in that situation the liquid never is completely clear. It is most turbid near the surface, while at the bottom it is almost clear. This is an indication of formation of VERY small bubbles of gas, which cannot be observed individually, but which cause the turbidity of the liquid.

For many aqueous chemistry experiments, requiring dilute H2O2 this stuff is perfectly suitable. Isolating the H2O2, however, is a different thing. You get a lot of additional cations and anions in such solutions and these may pose a problem in situations where more concentrated H2O2 is desired.




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[*] posted on 24-8-2006 at 07:50


maybe i should reword - trying to obtain concentrated peroxide from this would not work very well (he did state it would be diluted) - but 10 percent is better than the 3 percent from the drugstore - you can get 12 at some places if you look around - hmm would it be cheaper? - maybe there is a way to to refine the percarbonate up from 30 % - i always thought oxyclean was close to pure percarbonate

percarbonate can probably be made at home actually - i wonder what kind of concentration you could reach with the pure salt

[Edited on 24-8-2006 by jimmyboy]
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[*] posted on 24-8-2006 at 19:09


Quote:
Originally posted by jimmyboy
maybe i should reword - trying to obtain concentrated peroxide from this would not work very well (he did state it would be diluted) - but 10 percent is better than the 3 percent from the drugstore - you can get 12 at some places if you look around - hmm would it be cheaper? - maybe there is a way to to refine the percarbonate up from 30 % - i always thought oxyclean was close to pure percarbonate

[Edited on 24-8-2006 by jimmyboy]


Easier to concentrate the H2O2. Stablise it, store bought H2O2 already is, a bit of acetic acid plus polyphosphates or sodium silicate work if you make your the H2O2 from something else. Then distill off the water under somewhat reduced pressure; best if you use a short fractionating column; keep the temperature 50 C or less. Oh, and avoid dust getting into the solutions, clean that glassware well before using.

You can get concentrations of 50% or higher without too much effort, just the pain of seeing a liter of 3% shrink to about 100 ml of 30%.

Some hydrogen peroxide is stabilised with sodium stannate, I don't know if this might cause cloudiness on concentration.
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[*] posted on 17-3-2007 at 23:45


Hmm... well, I'm resuscitating this thread because I encountered an unexpected result which wasn't mentioned here.
I recently tried to obtain peroxide from Oxiclean by adding it a pinch at a time to concentrated (32%) HCl.
The reaction I expected was:
Na<sub>2</sub>CO<sub>3</sub> * xH<sub>2</sub>O<sub>2</sub> + HCl --> NaCl + CO<sub>2</sub> + H<sub>2</sub>O<sub>2</sub>

However, the strong chlorine gas smell I observed indicates that NaOCl was somehow formed and was reacting with the excess HCl.
Can anyone tell me whether or not this is reasonable? Will NaCl + H<sub>2</sub>O<sub>2</sub> + HCl make chlorine gas by in situ NaOCl? If so, then my peroxide is being used up to make hypochlorite which eliminates the possibility of isolating it from this method.




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[*] posted on 18-3-2007 at 09:43


You most like do obtain H2O2, but H2O2 reacts with concentrated HCl to form Cl2 and H2O:

H2O2 + 2HCl --> 2H2O + Cl2

You can also observe this reaction, when you add H2O2 to concentrated HCl.

If you used 30% H2SO4 instead of conc. HCl, then you would obtain H2O2 and CO2 gas would bubble out of solution.




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[*] posted on 18-3-2007 at 10:14


Thanks a bundle, woelen, unfortunately I don't have any sulfuric acid at the moment. Is it possible to produce it with HCl and MgSO<sub>4</sub>? Even if sulfuric acid can not be produced this way, would the presence of the sulfate ion allow me to obtain hydrogen peroxide from this method?



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[*] posted on 18-3-2007 at 10:15


No, it's the combination of oxidizer, acid and HCl.



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[*] posted on 21-3-2007 at 14:51


How much sodium percarbonate does oxyclean contain and how can you isolate it?
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