schleerblitz
Harmless
Posts: 6
Registered: 5-1-2019
Member Is Offline
|
|
Hydrogen Peroxide from Sodium Percarbonate
Hi everyone! I have recently found an old bag of sodium percarbonate at home, and when looking up the formula, I thought, I should try extracting
hydrogen peroxide from it. I have already tried adding it to cc sufuric acid, creating just a solution of it with water, and even dissolving out the
Hydrogen peroxide using ethanol, adn acetone, but I could not succeed. Based on research online, I found that the only thing that could succeed, would
be dissolving out the hydrogen peroxide with some organic solvent, that does not dissolve the sodium carbonate itself, but I could not find which
organic solvent to use. I would apprechiate any help.
|
|
|
Ubya
International Hazard
Posts: 1247
Registered: 23-11-2017
Location: Rome-Italy
Member Is Offline
Mood: I'm a maddo scientisto!!!
|
|
https://pubs.acs.org/doi/pdf/10.1021/ja02260a013
ethyl acetate and diethyl ether are good options
---------------------------------------------------------------------
feel free to correct my grammar, or any mistakes i make
---------------------------------------------------------------------
|
|
|
violet sin
International Hazard
Posts: 1480
Registered: 2-9-2012
Location: Daydreaming of uraninite...
Member Is Offline
Mood: Good
|
|
THE PARTITION COEFFICIENTS OF HYDROGEN PEROXIDE BETWEEN WATER AND CERTAIN ORGANIC SOLVENTS
Attachment: walton1916.pdf (390kB)
This file has been downloaded 720 times
|
|
|
schleerblitz
Harmless
Posts: 6
Registered: 5-1-2019
Member Is Offline
|
|
Thanks to everyone, for responding, I will try ethyl acetate, once I can make it, but I don't think I will be able to get any diethyl ether.
|
|
|
Sulaiman
International Hazard
Posts: 3692
Registered: 8-2-2015
Location: 3rd rock from the sun
Member Is Offline
|
|
although I appreciate doing things purely for educational purposes,
concentrating OTC hydrogen peroxide by evaporation would save money, chemicals, time and effort.
Save the percarbonate for situations where it is usable as-is.
(if nothing else ... cleaning glassware)
P.S. hydrogen peroxide in diethyl ether sounds rather hazardous,
https://en.wikipedia.org/wiki/Diethyl_ether_peroxide
I'd be wary of any solvent plus hydrogen peroxide.
https://www.sigmaaldrich.com/chemistry/solvents/learning-cen...
CAUTION : Hobby Chemist, not Professional or even Amateur
|
|
|
12thealchemist
Hazard to Others
Posts: 181
Registered: 1-1-2014
Location: The Isle of Albion
Member Is Offline
Mood: Rare and Earthy
|
|
I tried this some years back. I tried acetone, and let a few grams of sodium percarbonate sit over a few tens of mls of acetone in a tall jar for
probably a couple of weeks. The solution turned brown, then black, then did absolutely nothing for a while. One evening I came back home, and found
that it had exploded, sending shards of sticky glass across my bench. I was lucky I wasn't around when it went off - I certainly wasn't expecting it.
|
|
|
j_sum1
Administrator
Posts: 6320
Registered: 4-10-2014
Location: At home
Member Is Online
Mood: Most of the ducks are in a row
|
|
Thanks for that, 12thealchemist. This is on my don't-do list.
|
|
|
Fery
International Hazard
Posts: 1015
Registered: 27-8-2019
Location: Czechoslovakia
Member Is Offline
|
|
Hi 12thealchemist, thx for sharing this sad experience, I hope it will forever distract everyone from producing acetone peroxide, which is good only
to commit random suicide.
Ether peroxides are dangerous close to the end of ether distillation. Nobody wants to experience exploding sharp pieces of glass apparatus blown
apart. Before distilling ether, it is necessary to remove peroxides from ether by e.g. FeSO4.
Perhaps the ethylacetate suggested by Ubya is the safest choice... but personally I wouldn't perform such extractions.
schleerblitz - do not synthesize ethylacetate, you can buy it cheaply instead (if not from chem supplier, then as a nonacetone nail polish remover)
Question to all - even we find safe and powerful solvent to extract H2O2, how will we later manage to move it back into water phase in substantial
amount? The only way to use it would be using the organic solvent instead of solution of H2O2 in H2O, right?
|
|
|
Ubya
International Hazard
Posts: 1247
Registered: 23-11-2017
Location: Rome-Italy
Member Is Offline
Mood: I'm a maddo scientisto!!!
|
|
Quote: Originally posted by Fery  |
Question to all - even we find safe and powerful solvent to extract H2O2, how will we later manage to move it back into water phase in substantial
amount? The only way to use it would be using the organic solvent instead of solution of H2O2 in H2O, right? |
I was thinking about it. Back extracting the organic layer with water the will give a low percentage of H2O2 in solution, but some organic solvent
will dissolve in the aqueous solution. Distilling in a vacuum the organic solvent is counter productive, one could add water to the percarbonate and
distill the water/H2O2 under vacuum skipping the organic extraction. Probably if an aqueous solution is needed, the organic contamination is
inevitable
---------------------------------------------------------------------
feel free to correct my grammar, or any mistakes i make
---------------------------------------------------------------------
|
|
|
Bedlasky
International Hazard
Posts: 1239
Registered: 15-4-2019
Location: Period 5, group 6
Member Is Offline
Mood: Volatile
|
|
Let react sodium percarbonate with HCl to make H2O2 + NaCl solution. Heat it to get out CO2 and cool it. Than add NaOH and precipitate out SrO2 or
BaO2 using some soluble Sr or Ba salt. Filtred out peroxide, wash it well and then dissolve it in H2SO4. Insoluble sulfate is filtred out. Simplier is
bought 10% H2O2 and concentrated by this way.
|
|
|
unionised
International Hazard
Posts: 5126
Registered: 1-11-2003
Location: UK
Member Is Offline
Mood: No Mood
|
|
Is it only me who finds the name of the second author here amusingly apt?
"Consequently the values obtained in this
investigation are appreciably lower than those of Osipoff and Popoff. "
Anyway, Any flammable solvent mixed with hydrogen peroxide is a bad idea.
They don't need to form any specific compound- a mixture of two liquids can explode .
It's a pity the partition coefficient for chloroform is so poor since its the least flammable of the common solvents.
[Edited on 8-6-20 by unionised]
|
|
|
Ubya
International Hazard
Posts: 1247
Registered: 23-11-2017
Location: Rome-Italy
Member Is Offline
Mood: I'm a maddo scientisto!!!
|
|
| Quote: Originally posted by Bedlasky | | Let react sodium percarbonate with HCl to make H2O2 + NaCl solution. Heat it to get out CO2 and cool it. Than add NaOH and precipitate out SrO2 or
BaO2 using some soluble Sr or Ba salt. Filtred out peroxide, wash it well and then dissolve it in H2SO4. Insoluble sulfate is filtred out. Simplier is
bought 10% H2O2 and concentrated by this way. |
Uranyl Peroxide 
---------------------------------------------------------------------
feel free to correct my grammar, or any mistakes i make
---------------------------------------------------------------------
|
|
|
clearly_not_atara
International Hazard
Posts: 2787
Registered: 3-11-2013
Member Is Offline
Mood: Big
|
|
Had the thought that you could react sodium percarbonate with metaboric acid to give the perborate. While perborate is generally available, boric acid
and percarbonate are both much cheaper.
However, this reaction would ideally be carried out by dry grinding, which makes temperature control very difficult. Is there a way around this? Would
it even heat up in the first place?
|
|
|
macckone
Dispenser of practical lab wisdom
Posts: 2168
Registered: 1-3-2013
Location: Over a mile high
Member Is Offline
Mood: Electrical
|
|
There is no easy way to remove the peroxide from the percarbonate.
With sulfuric acid you form persulfate.
With hydrochloric acid you will form hypochlorite and chlorate and even chlorine.
With boric acid you will form perborate.
|
|
|
symboom
International Hazard
Posts: 1143
Registered: 11-11-2010
Location: Wrongplanet
Member Is Offline
Mood: Doing science while it is still legal since 2010
|
|
Sodium percarbonate is an abduct of sodium carbonate and hydrogen peroxide. sodium carbonate is insoluble in alcohol but hydrogen peroxide is soluble
in alcohol couldn't anhydrous ethanol or isopropanol be used due to isopropanol properties. It can only oxidize alcohol with a catalyst btw. So it
would be
Sodium carbonate perciptate with hydrogen peroxide/isopropanol
I do like the idea of hydrochloric acid and sodium percarbonate due to sodium chloride being insoluble in hydrochloric acid also hydrogen chloride
gas could be used to reduce the amount of water used
Also for isopropanol hydrogen chloride is double and sodium chloride is not even without hydrogen chloric acid.
[Edited on 9-10-2020 by symboom]
|
|
|
clearly_not_atara
International Hazard
Posts: 2787
Registered: 3-11-2013
Member Is Offline
Mood: Big
|
|
Well, yes, that is exactly the goal, but the goal is also to get a decent yield of the perborate and no thermal runaway decomposition.
|
|
|
Fyndium
International Hazard
Posts: 1192
Registered: 12-7-2020
Location: Not in USA
Member Is Offline
|
|
Is it possible to use detergent based sodium percarbonate solution to make sodium perborate from sodium metaborate solution? The percarbonate should
yield a mixture of hydrogen peroxide and sodium carbonate, and the peroxide is supposed to react with the metaborate, forming 2g/100mL soluble
perborate, which is expected to crash out? The perborate would then be washed with good amount of water to remove detergent.
This is because hydrogen peroxide is more valuable agent and if it could be substituted with otc bulk percarbonate, it would free the agent for better
use.
[Edited on 14-10-2020 by Fyndium]
|
|
|
macckone
Dispenser of practical lab wisdom
Posts: 2168
Registered: 1-3-2013
Location: Over a mile high
Member Is Offline
Mood: Electrical
|
|
boric acid is a stronger acid than carbonic acid so it displaces it on the percarbonate. Mixing percarbonate with sodium borate won't achieve the
same reaction but some perborate could form.
When percarbonate is displaced it dissolves into H2CO3 and H2O2. Carbon dioxide is lost and eventually you would get straight perborate. Or you
could simply add 3% peroxide to sodium metaborate and evaporate the water.
Or you could just buy 25lbs for $150 off of ebay.
|
|
|
symboom
International Hazard
Posts: 1143
Registered: 11-11-2010
Location: Wrongplanet
Member Is Offline
Mood: Doing science while it is still legal since 2010
|
|
Boric acid is quoted as being "an acid by name only" and is even too weak to react with solids such as copper carbonate. It does not produce protons
in water.
Used in solvents that cannot dissolve the carbonate but can leach the H2O2 out of it
.
What if the solution of sodium carbonate is chilled with CO2 bubbled in to convert it to sodium bicarbonate.
Sodium bicarbonate seems to not for a compound with hydrogen peroxide
Sodium bicarbonate soubility
69 g/L (0 °C)
96 g/L (20 °C)
165 g/L (60 °C)
Damn never mind
Sodium carbonate
G/100 mL
16.4 (15 °C)
34.07 (27.8 °C)
48.69 (34.8 °C)
48.1 (41.9 °C)
45.62 (60 °C)
43.6 (100 °C
Or converting it to to magnesium carbonate
By adding magnesium sulfate converting it to sodium sulfate which can be frozen out.
Which seems to not adhere to Hydrogen peroxide
Interesting wood bleaching with
Hydrogen peroxide and ammonium bicarbonate solution.
https://patents.google.com/patent/US3034851A/en
[Edited on 28-10-2020 by symboom]
[Edited on 28-10-2020 by symboom]
|
|
|
macckone
Dispenser of practical lab wisdom
Posts: 2168
Registered: 1-3-2013
Location: Over a mile high
Member Is Offline
Mood: Electrical
|
|
symboom,
The only reason borate replaces the carbonate is due to loss of CO2 from solution.
This is a slow reaction.
Additionally sodium perborate is a compound while sodium percarbonate is an abduct.
In a saturated solution sodium perborate will precipitate and CO2 will escape driving the reaction forward despite otherwise unfavorable kinetics.
|
|
|
Fyndium
International Hazard
Posts: 1192
Registered: 12-7-2020
Location: Not in USA
Member Is Offline
|
|
Whatever I made, it behaved like sodium perborate. I dissolved sodium percarbonate in water and mixed in solution that is made from borax and NaOH.
This causes precipitate to fall, which is decanted, filtered, washed and dried. No need to use expensive and restricted H2O2 to generate perborate.
|
|
|
macckone
Dispenser of practical lab wisdom
Posts: 2168
Registered: 1-3-2013
Location: Over a mile high
Member Is Offline
Mood: Electrical
|
|
Fyndium since perborate is the least soluble, that would make sense.
However, I would not expect a good yield but since the reactance are very very cheap, that may not matter.
|
|
|
Fyndium
International Hazard
Posts: 1192
Registered: 12-7-2020
Location: Not in USA
Member Is Offline
|
|
I described the process in another thread so I don't go into details here, but it is a viable method to produce it even in kg scale if necessary
without H2O2. Biggest losses are induced by the dissolved perborate in the reaction liquor because of the sheer volume of it, even when cooled down.
Na2CO3 begins to precipitate at about 7C in that concentration, so it is pretty much the limit for extracting pure perborate.
|
|
|
Fyndium
International Hazard
Posts: 1192
Registered: 12-7-2020
Location: Not in USA
Member Is Offline
|
|
Btw, is it necessary to dissolve all sodium percarbonate in order to leach the H2O2 into water? Just a wild thought, could it be slurried with 0C
water to extract water-miscible H2O2 and leave most of the sodium carbonate behind? Other, partly as crazy idea was to vacuum distill H2O2-water
directly off the percarb, possibly by constantly adding more water into agitated slurry and then do a final concentration on the collected dilute
condensate? Or just a vacuum dry distillation, crudely separating anhydrous H2O2 from sodium carbonate..  This sounds intriguing enough that I might wanna actually try it out in a microscale to see what kind of explosion it
ends up. The decomposition starts at 50C according to SM wiki, and official decomposition temp is 100-120C. This is apparently survivable to hydrogen
peroxide if there are no decomposing catalysts present, so conceptually it should be doable in a vacuum. Presuming the product itself does not contain
accelerants, the glassware would be rinsed with piranha and distilled water to remove any impurities, as almost pure H2O2 would result.
There was also a thread ( https://www.sciencemadness.org/whisper/viewthread.php?tid=30... ) which suggested adding calcium chloride to precipitate calcium carbonate and
leaving sodium chloride. I doubt this would be anyway beneficial if the product is intended to be distilled?
I'm interested in this, because percarb is extremely cheap and otc, while hydrogen peroxide is generally not available except as a business, and when
it is, it's like 20€ per 1L of 3%, so any process even with high losses is very beneficial. Although, using sulfuric acid to neutralize the
carbonate turns things upside down, on the other hand.
(Can H2O2 be concentrated by drying with peroxide-inert desiccant?)
[Edited on 16-1-2021 by Fyndium]
[Edited on 16-1-2021 by Fyndium]
|
|
|
![[*]](images/xpblue/default_icon.gif){kind=icon}
