mewrox99
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Sodium chlorite storage, unstable?
How unstable is NaClO2 to store.
I currently have about 500g of NaClO2.
How labile is this stuff?
Abstract from Chemical Hazards in Industry 1993
An explosion occurred when recrystallised samples of sodium chlorite ignited. The ignition was believed to be caused by a combination of friction,
formation of chlorate and hypochlorite and high temperature. The chlorite burnt with a fierce red flame, which could only be slowed down by a dry
chemical fire extinguisher. When all sodium chlorite was consumed, the fire stopped. Chemists have been advised to store recrystallised sodium
chlorite in small amounts and to use appropriate protective clothing when handling the compound.
It is triple bagged in my HDPE oxidisers drum. So is void of light and moisture.
Since I'll never need 500g of this stuff.
I'm thinking that if is a liability, I should get rid of it. Would heavily diluting it in a bucket (ie 500g in 5L) of water and adding very dilute
stoichiometric Na Thiosulfate, outside in small portions be a safe way to get rid of it, should I ever need to
Also btw with the diluting, I am aware of explosions with strong (hypo)chlorite solutions so it will be diluted slowly following the NaOH/water rule
[Edited on 19-5-2011 by mewrox99]
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woelen
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Sodium chlorite as it is sold commercially is stable, because of the co-crystallized NaCl. I assume that your NaClO2 is not the pure substance, but
somewhere between 80% and 85%. The remainder is NaCl. This is the MMS formulation, although I would never use it for that purpose.
Pure NaClO2 slowly disproportionates to NaClO3 and NaCl. This process slows down more and more until the solid contains appr. 15% NaCl. The NaClO2 you
can buy from eBay sellers is intended for human consumption (MMS), and for this reason it must be absolutely free of NaClO3 (the latter is poisonous).
For this reason, the NaClO2 is crystallized together with NaCl in order to assure that no further disproportionation occurs and no chlorate is formed
in the solid phase.
Keep it in its container and store in a dry and dark place. If stored in that way, it will remain good for tens of years at least. I would not be
afraid of ignition of the material. Only if there is a strong external heat source, it can start decomposing and take part of a very fierce and hard
to extinguish fire.
Do NOT dissolve the material in water. Its solutions are not stable and slowly decompose, just like NaClO solutions are not stable. Over time these
solutions may even pressurize, due to release of oxygen gas. Liquid MMS solutions for this reason always must be stored in a fridge and these have
limited shelf life.
You now see that sellers of MMS start selling solid NaClO2 now, because of the limited shelf life of MMS solutions. Some examples of such sellers:
http://nhc.lefora.com/2008/06/10/site-2/
https://cart.mmsdr.com/index.php?main_page=product_info&...
So, please do not be afraid of having the NaClO2 around. As long as it is stored away from strong external heat sources, there is no risk at all of
spontaneous ignition.
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AJKOER
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Hi:
I generally agree with the comments.
However, there was a reported incidence on a cargo ship transporting NaClO/NaCl solution. The fire disaster occurred because the manufacturer left a
magnesium impurity in the NaClO that formed Magnesium Hypochlorite. And even though this was a trace impurity, that was sufficient!
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woelen
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Here we are talking about NaClO2, not NaClO. The latter is hypochlorite and it is less stable than chlorite. The stability of the oxo anions of
chlorine increases with increasing oxygen number:
ClO(-) < ClO2(-) < ClO3(-) < ClO4(-)
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AJKOER
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For those with pure NaClO2, I agree further oxidization may not be an issue. Although checking for Mg impurity in the form of hypochlorite, chlorite
or chlorate would make me sleep better.
To the general public, I would note that anytime you have a solution of NaClO, in time you also have some NaClO2 and NAClO3, a fact that bleach
manufacturers are not too forthcoming about (they just tell you to change your bleach, which is more of a legal defense than a disclosure). Buying
cheap bleach may also increase the likelihood of possible dangerous impurities like Magnesium (this is just my cautionary opinion, as the cargo ship
hypochlorite disaster was linked to a discount Chinese supplier). Light and raising the temperature do promote the decomposition (around 70C is good
for making chlorate and not boiling as is usually suggested). So don't store your bleach on top of your clothes dryer! Ph is also an important
variable especially after the small amount of NaOH in your bleach is neutralized.
The two step disproportionation is:
2 NaClO --> NaCl + NaClO2
NaClO + NaClO2 --> NaCl + NaClO3
Sleep well.
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woelen
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Don't stir up things more than is necessary. There is no need for mewrox99 (or whoever has NaClO2 in his house) to check for magnesium impurity in his
batch of NaClO2. There is NO magnesium impurity in his material.
Indeed, magnesium and NaClO2 make a perfect energetic and sensitive mix, which easily could be set off and which most likely will burn very
spectacularly. But why would someone put magnesium in this stuff? Unless this happens by accident, there is no reason to worry about such a
contamination.
I have NaClO3 as well and this forms very dangerous and explosive mixes with red phosphorus. Is it necessary for me now to check my NaClO3 for
presence of red phosphorus?
--------------------------------------------------------------------------------------------------
I have serious doubt about the disproportionation of hypochlorite going over NaClO2. I do not know the precise mechanism of this disproportionation
reaction, but it is a very complex and still only partially understood reaction. As far as I know the disproportionation of NaClO leads to NaClO3
through other intermediates.
But for the sake of experimenting, I'll try this reaction next weekend. I have NaClO2 and I will mix this with a solution of NaClO. Let's see if we
can make NaClO3 from this by simply mixing the two chemicals. That would be an interesting synth, because both chlorites and hypochlorites can be
obtained easily while chlorates more and more are frowned upon because of abuse by bombers and terrorists.
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AJKOER
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My last point, on your order of stability if one were to reverse the order and add perchlorates (Cl04), you now have the proper ranking of oxidizers
for ROCKET FUEL.
ClO4 > ClO3 > ClO2....
Perchlorates are the current item of choice for solid fuel rockets.
My actual point is that certain metals, perhaps only present as impurities, produce more dangerous salts that have acted as the primer to ignite their
otherwise more stable cousins.
So reader, be smart and be safe.
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AJKOER
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Sorry as was not apparently completely clear. I am NOT talking about the pure metal Magnesium. When one makes NaClO, for example, you may have also
made an impurity of Magnesium Hypochlorite. This is what actually is believed to have happen with the discount Chinese supplier. I say believed (per
the investigating authorities) as the entire cargo was consumed in a massive self-sustaining fire. Somewhat unbelievable for a fresh "stable"
NaClO/NaCl solution!
Now, when the ClO becomes ClO2, it occurs for all metal hypochlorites present, whether there are well behaved or not.
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AJKOER
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On my challenged disproportionation reactions, the reference: "Handbook of Detergents: Production Volume 142", pages 445 to 446 by Uri Zoller and Paul
Sosis. This is a Google book and these pages can be viewed online.
LINK:
http://books.google.com/books?id=dXn3aB1DKk4C&pg=PA247&a...
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AJKOER
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A further reference and a couple of minor corrections. The equation of interest on the disproportionation reaction starts on Page 444 (and not 445) in
Zoller's "Handbook of Detergents: Production Volume 142". For a reference on the alleged cargo ship disasters, see below which alludes, in this
incident, to a Japanese (and not Chinese) supplier and commercial hypochlorites (Calcium Hypochlorite) with a Magnesium oxide impurity in the lime
used to produce the hypochlorite:
"Bretherick's handbook of reactive chemical hazards" edited by P. G. Urben, page 1358.
Also a free online Google book.
Link:
http://books.google.com/books?id=_dW_2XPbo_oC&pg=PA1358&...
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mewrox99
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If at some stage, I need to get rid of it what would be the best method. I'm being more safety minded in my lab and that includes knowing how to
dispose of everything when I move in a few years.
Is this a safe way to get rid of NaClO2. Slowly dissolve it in a large (ca 5L amount) of water and slowly mix with Na Thiosulfate solution and pour
down drain?
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entropy51
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An old Aldrich catalog states:
Cautiously acidify a 3% solution to pH 2 with sulfuric acid. Cautiously add a 50% excess of aqueous sodium bisulfite with stirring at room
temperature. An increase in temperature indicates that a reaction is taking place.
I would first try this on a small scale (few grams) and would not run large batches. I would also recommend not buying larger amounts of chemicals
than you expect to use in a reasonable amount of time, no matter how cheap they happen to be. 500 grams is a rather large amount of a strong oxidizer
to have in a home lab.
Where I live the county has hazardous waste disposal days on which you can take household, garden, photographic and swimming pool chemicals down for
disposal, no questions asked. If available this would be your best option.
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woelen
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If you really want to get rid of the NaClO2 (which I would find a waste of money and resources), then the easiest way to do this is dissolve half of
it in 8 liters of water and mix this well, then add 1 liter of dilute hydrochloric acid (10% from hardware store) and put this aside OUTSIDE for one
day in bright daylight (preferrably with a lot of sunshine). The liquid will be yellow and contains 1% or so of ClO2, but the (sun)light slowly
decomposes the ClO2 formed and after one day you can flush all of it through the toilet (only do that if you have no septic tank and have a connection
to the municipal sanitary sewer!!). The 1% solution of ClO2 in water is not dangerous at all, there is no risk of explosion or otherwise violent
decomposition. It is smelly though and hence you have to keep it outside.
In two batches then all of it can be destroyed safely and at not too high cost.
------------------------------------------------------
Inspired by AJKOER's posts I also did some experiments with mixes of NaClO and NaClO2 and found some interesting results. A separate thread is made
for these in order to keep this thread on topic for disposal of NaClO2:
http://www.sciencemadness.org/talk/viewthread.php?tid=16378
[Edited on 22-5-11 by woelen]
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Formatik
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@mewrox99: In the hazard you cited, the NaClO2 was subjected to friction. Chlorites are usually shock sensitive. Even LiClO2, NaClO2, Ca(ClO2)2 and
Sr(ClO2)2 are known to explode on impact, for example. Impurities, certainly heavy metal compounds or combustibles (sulfur is especially dangerous)
tend to heighten shock sensitivity. Sensitivity ought to reduce with NaClO2 that has been recrystallized with NaCl. One could test a few milligrams by
subjecting it to hammer blows on a flat surface.
In Gmelin it's reported a NaClO2-solution doesn't change in 10 hours on a hot bath, and at room temperature it keeps for years. I would expect the
solid to keep for many years. Just keep it out of sunlight, heat, and away from incompatibles (check the incompatibility sheets).
Another way to destroy chlorite, is to mix it with H2O2 and then heat it: NaClO2 + 2 H2O2 = 2 H2O + NaCl + O2. I think it would be a waste of two good
reagents, though. Hypochlorites decompose similarly, but don't need the heating. E.g., aq. NaClO solutions decompose vigorously when mixed with H2O2
also forming H2O, NaCl, and O2.
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AJKOER
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For those interested in exploring the mechanics of the disproportionation reaction in hypochlorites forming chlorites and chlorates, I found a very
good educational reference:
"DICHLORINE MONOXIDE, HYPOCHLOROUS ACID, AND HYPOCHLORITES", Pages 553-554 and Page 559.
LINK:
http://www.scribd.com/doc/30121142/Dichlorine-Monoxide-Hypoc...
Rereading this chapter, I was impressed by the oxidizing ability of DiChlorine Mono-oxide, Cl2O, apparently even more powerful than HClO.
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Formatik
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A good reference from Kirk Othmer.
Below is attached a description of the stability of sodium chlorite.
Attachment: ie50367a007.pdf (708kB)
This file has been downloaded 1204 times
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woelen
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Thanks both of you for the very interesting reads. After having read both of them I see that chlorite is MUCH more stable than hypochlorite. Even
after 10 years of storage, the samples of chlorite only lost a few percents of activity at most. You must also keep in mind that these samples were
made with 1930's technology. I'm quite sure that our samples of chlorite are even more pure and the containers we have probably are much more tight
than the containers of the 1930's. Again, this is a good indication of its shelf life properties and it indicates that this material can be kept
around really safely. The only real risk is when there is a fire and the chlorite enhances the burning rate of nearby combustible materials.
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Formatik
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Indeed. From, those studies we can ascertain the mass loss of NaClO2 is on the order of 3 to 5% about per 10 years. Wether it decomposes at a faster
rate after that time period, I don't know, but doubt it. The study also apparently explains shock sensitivity of sodium chlorite (which also might
apply to other alkali and alkaline-earth chlorites):
| Quote: | | Levi (17) stated that sodium chlorite explodes upon percussion or impact. This has been investigated by placing a portion of sodium chlorite upon an
anvil and striking it with a hammer. Explosion will result unless the face of the hammer and anvil have been freed from the film of grease usually
present on such tools. When both the hammer and anvil are clean, there is no explosion when the chlorite is struck. The addition of a trace of oil or
grease will ensure explosion. Less adherent organic matter will probably, but not always, cause explosion when the mixture is struck.
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mewrox99
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Since there is no point keeping 500g of it and in the event of a fire having it around would be a very bad thing. I've decided to get rid of most of
it. I'm going to take it to the hazardous waste section of the town dump. Probably the easiest and safest way
[Edited on 29-5-2011 by mewrox99]
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turd
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Since I am very interested in solid solutions and clathrates, I would appreciate if you could shed any light on this co-crystallization. I have a hard
time seeing how Cl and ClO2 could have the same coordination environment and in a quick search I failed to find any stoichiometric Na/Cl/ClO2
compounds. Do you have a reference giving details on these co-crystals?
Edit: grammar. 
[Edited on 29-5-2011 by turd]
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woelen
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There is no real exact stoichiometric NaCl/NaClO2 compound. What is sold on eBay is 80% NaClO2 plus 20% NaCl, but it could equally well be 75% NaClO2
or 85% NaClO2. So, the cocrystallized solid is not a pure well-defined entity, but in fact a very intimate mix of both chemicals. It is sold as such,
because pure NaClO2 is not stable and slowly disproportionates, giving NaClO3 as one of the undesired toxic by-products, rendering the material
unsuitable for human consumption as MMS.
Maybe the term 'cocrystallization' is not the correct term. What is done is making the 80/20 mix of NaClO2 and NaCl not by simply mixing the solids,
because that would not prevent the decomposition of the NaClO2. The solid mix is made from solution, so that the solid material is a very intimate
mix, probably at the molecular level as kind of solid solution.
I used the term 'cocrystallized' because my sample has a crystalline appearance. The solid consists of small granules, but not round. The length of
the granules is appr. twice the diameter, hence the crystalline appearance.
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