Sciencemadness Discussion Board
Not logged in [Login - Register]
Go To Bottom

Printable Version  
Author: Subject: Sodium peroxide 8-hydrate
Polverone
Now celebrating 15 years of madness
*********




Posts: 3144
Registered: 19-5-2002
Location: The Sunny Pacific Northwest
Member Is Offline

Mood: No Mood

[*] posted on 20-5-2002 at 18:14
Sodium peroxide 8-hydrate


This is from "Inorganic Syntheses Volume III"

SODIUM PEROXIDE 8-HYDRATE
Submitted by R.A. Penneman
Checked by A.D.F. Toy

Hydrates of sodium peroxide may be prepared by (1) the slow evaporation of a cold aqueous solution of sodium peroxide [footnote 1]; (2) the slow action of water vapor on solid sodium peroxide [footnote 2]; (3) the electrolysis of aqueous sodium hydroxide at temperatures between -10 and 0 degrees [footnote 3], and (4) precipitation from a cold solution of sodium hydroxide and hydrogen peroxide by means of alcohol [footnote 4].

The method outlined below is a modification of the last mention of these procedures and yields a chemically pure product. The success of this method depends on the use of an excess of sodium hydroxide, since no precipitate is obtained when hydrogen peroxide is present in excess. However, it was found that the amount of water in the hydrate depends on the temperature at which precipitation is carried out. The directions that follow give the 8-hydrate consistently, whereas a product analyzing almost exactly for Na2O2.11H2O is formed when hydrogen peroxide is added to a saturated sodium hydroxide solution at 0 degrees.

Procedure

Ten grams of carbonate-free sodium hydroxide is dissolved in 25 ml of water in a stoppered Erlenmeyer flask and cooled to 15 degrees. Ten grams of a 30 per cent solution of hydrogen peroxide (corresponding to a mol ratio of NaOH:H2O2 = 2.83:1) is added slowly with constant stirring at a rate such that the temperature does not rise above 18 degrees. Sixty milliliters of 95 per cent alcohol (cooled to 15 degrees) is added; the flask is then stoppered and shaken vigorously. The solution is allowed to stand about 1/2 hour, the supernatant liquid is decanted, and the washing is repeated with two 60 ml portions of cold alcohol. The white crystals are filtered with suction on a hardened filter paper and washed with ether. The compound is transferred quickly to a dessicator containing sulfuric acid (not in vacuo) and kept in a cold chest for 10 hours at a temperature not above 15 degrees. The yield is 18 g (92 per cent based on H2O2). The product may be preserved for a limited amount of time in the ice chest.

Analysis

Sodium (reported as Na2O) was determined by hydrolysis of a sample and titration with a standard acid. Peroxide oxygen was determined by dissolving a weighed sample in an excess of standard cerium(IV) nitrate solution and titrating the excess cerium(IV) ion with iron(II) sulfate using o-phenanthroline as indicator [footnote 5]. Analysis calculated for Na2O2.8H2O: Na2O, 27.88; O (peroxide) 7.2 Found: Na2O, 27.92; O, 7.11

Properties

The 8-hydrate is a white, crystalline powder which reacts readily with carbon dioxide, hence must be kept from contact with the atmosphere. It melts in its own water of crystallization at 30 degrees and decomposes to yield oxygen. If kept for a long period over sulfuric acid in a vacuum dessicator, the 8-hydrate loses 6 molecules of water to form the 2-hydrate, Na2O2.2H2O [footnote 2].

References

1. Harcourt, J. Chem. Soc., 14, 278 (1862).
2. Joubert: Compt. rend., 132, 86 (1901).
3. German patent 245531 (1911).
4. Fairly: J. Chem. Soc., 31, 125 (1877).
5. Waldent, Hammett, and Chapman: J. Am. Chem. Soc., 53, 3908 (1931).
View user's profile Visit user's homepage View All Posts By User
Coen
Harmless
*




Posts: 10
Registered: 20-5-2002
Location: Netherlands
Member Is Offline

Mood: No Mood

[*] posted on 20-5-2002 at 18:14


Nice, but not THAT interesting though.
If I remember I still have a bottle of that stuff somewhere.
View user's profile View All Posts By User
Polverone
Now celebrating 15 years of madness
*********




Posts: 3144
Registered: 19-5-2002
Location: The Sunny Pacific Northwest
Member Is Offline

Mood: No Mood

[*] posted on 20-5-2002 at 18:15


This was actually posted in response to a private conversation with another board member. We had discussed making metallic peroxides and he said that he didn't think sodium peroxide could exist in aqueous media. So I looked up that synthesis and posted it.
View user's profile Visit user's homepage View All Posts By User
Theoretic
International Hazard
*****




Posts: 756
Registered: 17-6-2003
Location: London, the Land of Sun, Summer and Snow
Member Is Offline

Mood: eating the souls of dust mites

[*] posted on 23-7-2003 at 09:15


Assuming the product is less soluble in water than NaOH, electrolysis could be a continuous method of production.
Also, electrolysis doesn't consume H2O2.
Instead, a warm or hot Na2O2 solution could be vacuum-distilled to GIVE H2O2 (due to hydrolysis).
:)
View user's profile View All Posts By User
Organikum
resurrected
*****




Posts: 2228
Registered: 12-10-2002
Location: Europe
Member Is Offline

Mood: fluffy!

[*] posted on 23-7-2003 at 14:27


Thats quite exactly what I was after, thanks Polverone! H2O2 is asscheap and easy here and I donĀ“t have sodium peroxide around in any form. So the described modification procedure is a present to me.



Restrict alcohol and Tobacco.
Legalize everything else.
Mandatory LSD for politicians and Franklyn.
View user's profile View All Posts By User
verode
Harmless
*




Posts: 32
Registered: 22-3-2005
Member Is Offline

Mood: No Mood

[*] posted on 8-10-2006 at 12:45


If you get NaOH drye +Methanol (drye) you add 30% H2O2 and you'll get better results with Na2O2.hydrate
Quote:
Originally posted by Polverone
This is from "Inorganic Syntheses Volume III"

SODIUM PEROXIDE 8-HYDRATE
Submitted by R.A. Penneman
Checked by A.D.F. Toy

Hydrates of sodium peroxide may be prepared by (1) the slow evaporation of a cold aqueous solution of sodium peroxide [footnote 1]; (2) the slow action of water vapor on solid sodium peroxide [footnote 2]; (3) the electrolysis of aqueous sodium hydroxide at temperatures between -10 and 0 degrees [footnote 3], and (4) precipitation from a cold solution of sodium hydroxide and hydrogen peroxide by means of alcohol [footnote 4].

The method outlined below is a modification of the last mention of these procedures and yields a chemically pure product. The success of this method depends on the use of an excess of sodium hydroxide, since no precipitate is obtained when hydrogen peroxide is present in excess. However, it was found that the amount of water in the hydrate depends on the temperature at which precipitation is carried out. The directions that follow give the 8-hydrate consistently, whereas a product analyzing almost exactly for Na2O2.11H2O is formed when hydrogen peroxide is added to a saturated sodium hydroxide solution at 0 degrees.

Procedure

Ten grams of carbonate-free sodium hydroxide is dissolved in 25 ml of water in a stoppered Erlenmeyer flask and cooled to 15 degrees. Ten grams of a 30 per cent solution of hydrogen peroxide (corresponding to a mol ratio of NaOH:H2O2 = 2.83:1) is added slowly with constant stirring at a rate such that the temperature does not rise above 18 degrees. Sixty milliliters of 95 per cent alcohol (cooled to 15 degrees) is added; the flask is then stoppered and shaken vigorously. The solution is allowed to stand about 1/2 hour, the supernatant liquid is decanted, and the washing is repeated with two 60 ml portions of cold alcohol. The white crystals are filtered with suction on a hardened filter paper and washed with ether. The compound is transferred quickly to a dessicator containing sulfuric acid (not in vacuo) and kept in a cold chest for 10 hours at a temperature not above 15 degrees. The yield is 18 g (92 per cent based on H2O2). The product may be preserved for a limited amount of time in the ice chest.

Analysis

Sodium (reported as Na2O) was determined by hydrolysis of a sample and titration with a standard acid. Peroxide oxygen was determined by dissolving a weighed sample in an excess of standard cerium(IV) nitrate solution and titrating the excess cerium(IV) ion with iron(II) sulfate using o-phenanthroline as indicator [footnote 5]. Analysis calculated for Na2O2.8H2O: Na2O, 27.88; O (peroxide) 7.2 Found: Na2O, 27.92; O, 7.11

Properties

The 8-hydrate is a white, crystalline powder which reacts readily with carbon dioxide, hence must be kept from contact with the atmosphere. It melts in its own water of crystallization at 30 degrees and decomposes to yield oxygen. If kept for a long period over sulfuric acid in a vacuum dessicator, the 8-hydrate loses 6 molecules of water to form the 2-hydrate, Na2O2.2H2O [footnote 2].

References

1. Harcourt, J. Chem. Soc., 14, 278 (1862).
2. Joubert: Compt. rend., 132, 86 (1901).
3. German patent 245531 (1911).
4. Fairly: J. Chem. Soc., 31, 125 (1877).
5. Waldent, Hammett, and Chapman: J. Am. Chem. Soc., 53, 3908 (1931).
View user's profile View All Posts By User
The_Davster
A pnictogen
*******




Posts: 2859
Registered: 18-11-2003
Member Is Offline

Mood: No Mood

[*] posted on 15-10-2006 at 20:49


I prepared several grams of sodium peroxide about a year ago, it seemed to be storing fine, perhaps loosing a bit of its powerfull oxidative power, but not much. I last looked at it a month ago and it seemed fine. Today I find the vial of it, and it is no longer crystaline...it is a foamy colorless slightly viscous liquid.:o
Can anyone think of how this would have happened, as in what reaction took place?
View user's profile View All Posts By User
guy
International Hazard
*****




Posts: 982
Registered: 14-4-2004
Location: California, USA
Member Is Offline

Mood: Catalytic!

[*] posted on 15-10-2006 at 21:44


Looks like hydrogen peroxide.



View user's profile View All Posts By User
not_important
International Hazard
*****




Posts: 3873
Registered: 21-7-2006
Member Is Offline

Mood: No Mood

[*] posted on 15-10-2006 at 21:44


Quote:
Originally posted by rogue chemist
I prepared several grams of sodium peroxide about a year ago, it seemed to be storing fine, perhaps loosing a bit of its powerfull oxidative power, but not much. I last looked at it a month ago and it seemed fine. Today I find the vial of it, and it is no longer crystaline...it is a foamy colorless slightly viscous liquid.:o
Can anyone think of how this would have happened, as in what reaction took place?


Sounds like either the cap started to leak a bit more, or the local humidity increased enough that the existing leak rate became important. When Na2O2 reacts with water it makes NaOH, which grabs onto additional water to make a concentrated solution, somewhat thick or viscous.

It's worth keeping bottles of moisture sensitive chemicals in a dessicator, even a plastic storage container with a good tight seal would work. Use CaSO4, MgSO4, silica gel, or molecular sieves for the dessicant, include a little indicating dessicant or anhydrous CuSO4 in a small vented container so you can tell when the dessicant needs to be regenerated. That gives some protection against cap seals giving out, or even not tightening the cap on the bottle enough.
View user's profile View All Posts By User
The_Davster
A pnictogen
*******




Posts: 2859
Registered: 18-11-2003
Member Is Offline

Mood: No Mood

[*] posted on 6-11-2006 at 21:58


Yup, turns out humidity had really changed a few months ago...So much for the notoriously dry weather of my province. Lost some nice things, small samples of finely divided rare earths, anhydrous silver perchlorate. Somehow the P2O5 stayed safe however. It would be a terrible loss loosing that. Now I have to teflon tape other bottles well, I hope it is not too late for the acetic anhydride.

And a open sack of CaCl2 goes into the cabinet.

:mad::(:mad::(




View user's profile View All Posts By User
not_important
International Hazard
*****




Posts: 3873
Registered: 21-7-2006
Member Is Offline

Mood: No Mood

[*] posted on 7-11-2006 at 00:25


As I said elsewhere, standard white teflon tape is fairly porous, it's been used to hold catalysts with 'fluor-tails'. It works to seal water pipe joints because it is not wetted by liquid water and any leak is trying to cross a long section of the tape - at least the width as opposed to the thinkness when you wrap it around a bottle cap. If you use the tape on the bottle threads - have it between the bottle and the cap - then it will do more good, but it's still not real effective against water vapour.

Red teflon tape, intended to be used with gas lines, is OK as it is much denser and less porous; ordinary PVC tape will do the job if all you are after is keeping water vapour out
View user's profile View All Posts By User
The_Davster
A pnictogen
*******




Posts: 2859
Registered: 18-11-2003
Member Is Offline

Mood: No Mood

[*] posted on 7-11-2006 at 21:02


Thanks not_important, I had many inflated beliefs of the power of teflon tape, now dispelled.

[Edited on 8-11-2006 by The_Davster]




View user's profile View All Posts By User
not_important
International Hazard
*****




Posts: 3873
Registered: 21-7-2006
Member Is Offline

Mood: No Mood

[*] posted on 7-11-2006 at 21:31


If you can find the red tape, it is pretty good. The other useful item is sheets of dense teflon that can be cut/punched for cap liners, or pre-made disks of the same. If you pick think enough ones, the cold flow of teflon gives a nice seal and protects the cap from being chewed up. Combine that with tape on the threads or a good tape on the outside and you've got decent long term storage. Of course properly sealed glass ampules are better, but less convenient.
View user's profile View All Posts By User
Rosco Bodine
International Hazard
*****




Posts: 6277
Registered: 29-9-2004
Member Is Offline

Mood: analytical

[*] posted on 7-11-2006 at 21:32


I wonder if acetone peroxide would react with the
NaOH dissolved in warm dry denatured alcohol .

This would eliminate the added water .

Such a reaction might also give sodium ethoxide , or of course nothing at all , but I would guess it will react ,
and the product may be dependant on the temperature
of the reaction , warmer temperature favoring the
ethoxide .
View user's profile View All Posts By User
Zinc
National Hazard
****




Posts: 472
Registered: 10-5-2006
Member Is Offline

Mood: No Mood

[*] posted on 14-1-2007 at 06:49


Can NaOH be replaced by some other soluble sodium salt (ZnO2 can be made with ZnSO4)?

[Edited on 14-1-2007 by Zinc]




View user's profile View All Posts By User
woelen
Super Administrator
*********




Posts: 6546
Registered: 20-8-2005
Location: Netherlands
Member Is Offline

Mood: interested

[*] posted on 14-1-2007 at 09:38


ZnO2 probably can be made from ZnO or Zn(OH)2, suspended in water. This also is possible with Ba(OH)2. You will not obtain it from ZnSO4.



The art of wondering makes life worth living...
Want to wonder? Look at http://www.oelen.net/science
View user's profile Visit user's homepage View All Posts By User
Zinc
National Hazard
****




Posts: 472
Registered: 10-5-2006
Member Is Offline

Mood: No Mood

[*] posted on 14-1-2007 at 09:53


Quote:
Originally posted by woelen
You will not obtain it from ZnSO4.


http://www.roguesci.org/theforum/showthread.php?t=276&hi...

ZnO2 can be made from ZnSO4.




View user's profile View All Posts By User
guy
International Hazard
*****




Posts: 982
Registered: 14-4-2004
Location: California, USA
Member Is Offline

Mood: Catalytic!

[*] posted on 14-1-2007 at 17:10


Quote:
Originally posted by Zinc
Quote:
Originally posted by woelen
You will not obtain it from ZnSO4.


http://www.roguesci.org/theforum/showthread.php?t=276&hi...

ZnO2 can be made from ZnSO4.


The procedure requires you precipitate Zn(OH)2 with ammonia.




View user's profile View All Posts By User
S.C. Wack
bibliomaster
*****




Posts: 1819
Registered: 7-5-2004
Location: Cornworld, Central USA
Member Is Offline

Mood: Enhanced

[*] posted on 14-1-2007 at 19:43


It is probable that any souble Zn+2 salt and any base capable of forming its hydroxide/zincate can be used. This says nothing of the likelihood of other Na salts being used for a similar production of its peroxide hydrate though, and it would be silly to try. Peroxo or peroxohydrate cpds are the likely result if they crystallize out at all..

According to US3305310, there are storage problems with the title cpd.
View user's profile Visit user's homepage View All Posts By User
Zinc
National Hazard
****




Posts: 472
Registered: 10-5-2006
Member Is Offline

Mood: No Mood

[*] posted on 25-1-2007 at 10:47


Today I made it following the synthesis posted by Polverone. I didn't wash it with ether because I don't have any.
Does anyone know if this will have any efect on the Na2O2?




View user's profile View All Posts By User
The_Davster
A pnictogen
*******




Posts: 2859
Registered: 18-11-2003
Member Is Offline

Mood: No Mood

[*] posted on 25-1-2007 at 18:15


No biggie, mine was not washed either, and lasted over a year.



View user's profile View All Posts By User
Theoretic
International Hazard
*****




Posts: 756
Registered: 17-6-2003
Location: London, the Land of Sun, Summer and Snow
Member Is Offline

Mood: eating the souls of dust mites

[*] posted on 26-1-2007 at 03:58


On the subject of other metals' peroxides, calcium peroxide forms from CaO and oxygen exothermically (its dissociation pressure at 55 degrees is 1 atm), and there have been experiments on preparing it directly.



View user's profile View All Posts By User
Ballermatz
Harmless
*




Posts: 19
Registered: 17-7-2007
Member Is Offline

Mood: No Mood

[*] posted on 4-11-2007 at 08:04


Doing some research now how to best synthesize BaO2 and I stumbled across this old thread. Just wanted to mention that I made ZnO2 two times myself, first time using the reaction posted here, second time I simply boiled a mixture of 30% H2O2 and ZnO. I got the same yellow poweder both times, with strong oxidizing capabilities. So I dont see the reason to use NaOH or NH3 to produce hydroxides/zincates first. The oxide seems to react rightaway with the O from decomposing H2O2.
View user's profile View All Posts By User

  Go To Top