Sciencemadness Discussion Board - Preparation of sodium chlorite

Preparation of sodium chlorite

garage chemist - 1-2-2005 at 08:51

As I found out that my planned project of making HClO4 requires vacuum and because our groundwater pump isn't connected in winter (tap water is expensive here), I decided to do something different instead.

Sodium chlorite is a bleaching agent often used in industry because it doesn't attack the fibers (as does hypochlorite) and produces no chlorinated organics. It is also used to produce in- situ chlorine dioxide, which is a disinfectant much more environmentally friendly than chlorine.

NaClO2 can decompose exothermically when heated, and the silver and mercury salts of chlorous acid are violently explosive, being interesting candidates for primary explosives.

This link outlines the production of sodium chlorite, basically, ClO2 is first produced from NaClO3, NaCl and H2SO4, this is then purified and reacted with aqueous NaOH to which a reducing agent has been added (to avoid the formation of chlorate which would otherwise occur).

I don't like this process because of the explosion hazard involved in handling the ClO2. The ClO2 is diluted with air in the industrial process, but this is difficult in a home setup.

In a textbook, there was a remark that sodium chlorite could be produced by reducing sodium chlorate with oxalic acid. This would of course be the best method, but nothing was said about reaction conditions.

Oxalic acid is always oxidised to CO2, so when there's CO2 evolution on mixing (and heating) solutions of oxalic acid and NaClO3, it would be a success.

I'll try this out soon.

Chlorine Dioxide

chloric1 - 1-2-2005 at 09:55

I posted,about a year ago, that I successfully and safely produced Chlorine Dioxide in quantity by placing potassium chlorate, oxalic acid, and sulfuring acid in a test tube then submerging the tube in hot water. Tthe CO2 from the oxalic acid safely dilutes the chlorine dioxide which I may add, has a seductive oxidizer smell

That is not even the exciting part! The reaction is ENTIRELY temperature dependent. So if you need to stop the gas flow, put your apparatus in ice water. The reaction will resume when its gets warm again.

According to one of my Inorganic Synthesis books, it is plausable to make sodium peroxide octahydrate from sodium hydroxide solution and hydrogen peroxide at subfreezing temperatures. This mixture would definately reduce chlorine dioxide to the chlorite.

I plan to redo my chlorine dioxide preparations soon. This time, I wish to reduce acidic chlorate with urea. I saw this in a patent but which one I cannot recall.

Oh, just thought I would mention one important key factor. I believe the temperature response to my previous reaction was largely based on temperature/solubility curves for both oxalic acid and potassium chlorate. Sodium Chlorate mixture may not be as easy to control.

[Edited on 2/1/2005 by chloric1]

garage chemist - 1-2-2005 at 13:29

Why do you want ClO2? Also for making chlorites?
I am trying to make NaClO2 without using ClO2, maybe it was a bit unclear in my message?

Theoretic - 1-2-2005 at 13:41

When reacted with chlorine, NaNO3 gives NaClO2:
NaNO3 + Cl2 => NaClO2 + NOCl.

garage chemist - 1-2-2005 at 13:52

Is anhydrous NaNO3 used ?

In aqueous solution, the NOCl would hydrolyse, the resultant HCl would liberate ClO2 from the NaClO2 and nitrogen oxides would also form.

Esplosivo - 1-2-2005 at 14:26

I think it must be in the anhydrous state. Theoretic, can KNO3 be used instead of the NaNO3?

The_Davster - 1-2-2005 at 15:01

I have seen camping water purifier tablets for sale which are 6% sodium chlorite. An expensive sourse, but noteworthy.

BromicAcid - 1-2-2005 at 15:24

Theoretic, I've seen a similar reaction to make NOCl:

2NO2 + KCl -----> KNO3 + NOCl

So, then you could just add chlorine gas and do your equation:

NaNO3 + Cl2 ---> NaClO2 + NOCl

Anyway, the reactoin that I posted needs 2.4% moisture according to the prep I pulled it from. ClO2 is fairly soluble in water, slow addition of HCl to chlorate could result in a solution of ClO2 in water with NaCl, then add some base to get your chlorite, as long as the temp is kept down and acid additoin slow you may get away without ClO2 gas being freely evolved.

BromicAcid - 2-2-2005 at 07:34

I'm at my library right now and decided to look up preparations of sodium chlorite, these are abstracted from The Encyclopedia of Chemical Reactions, Jacobson, copywrite 1948, Fifth Edition 1961, Vol II:

Quote:

An aqueous mixture of a base and sulfur or an inorganic sulfurous compound reacts with chlorine dioxide to form the chlorite of the base metal. The yield is high at room temperature.
6ClO2 + S + 8NaOH ---> 6NaClO2 + Na2SO4 + 4H2O
2ClO2 + Na2SO3 + 2NaOH ---> 2NaClO2 + Na2SO4 + H2O

G.P. Vincent, U.S. Pat. 2,092,944, Sept. 14, 1937

An aqueous mixture of a base and a carbonaceous material such as carbon reduces chlorine dioxide to the chlorite of the base metal in high yields at 20-50C.
4ClO2 + C + 6NaOH ---> 4NaClO2 + Na2CO3 + 3H2O

G.P. Vincent, U.S. Pat., 2,092,945, Sept 14, 1937
Ref., M. C. Taylor, J. F. White, G.P. Vincent, G.L. Cunningham, Ind. Eng. Chem., 32, 899 (1940)

Chlorine dioxide is absorbed in a solution of hydrogen peroxide and sodium bicarbonate to form sodium chlorite at room temperature in high yield. Also works with KHCO3 to form potassium chlorite. The mol ratio of peroxide to bicarbonate is 1 to 2.
2ClO2 + 2NaHCO3 + H2O2 ---> 2NaClO2 + 2CO2 + O2 + 2H2O
2ClO2 + 2KHCO3 + H2O2 ---> 2KClO2 + 2CO2 + O2 + 2H2O

E.C. Soule, U.S. Pat. 2,332,180, Oct. 19, 1943

Of course these being better then just absorbing in straight NaOH due to that reaction producing almost an equimolar mixture of chlorate and chlorite if done at medium temperature.

Theoretic - 2-2-2005 at 08:39

I don't actually know if KNO3 can be used instead of NaNO3, but there seems no reason why not. I don't know much about this reaction, I only saw the equation mentioned and that's it. If you do the reaction in acidic aqueous solution, then I think NOCl won't hydrolyse and will just fizz out (I assume its solubility depends directly on pH), and ClO2 will evolve as well (the problem lies in separating the two

). Or, a nonaqueous solvent could be used.

chloric1 - 2-2-2005 at 09:57

Quote:

Originally posted by BromicAcid
I'm at my library right now and decided to look up preparations of sodium chlorite, these are abstracted from The Encyclopedia of Chemical Reactions, Jacobson, copywrite 1948, Fifth Edition 1961, Vol II:

Quote:

Of course these being better then just absorbing in straight NaOH due to that reaction producing almost an equimolar mixture of chlorate and chlorite if done at medium temperature.

Yeh all these seem decent and applicable. I will read the patents you stated in my leisure and sometime I will compare these methods with my own method I stated previously

chloric1 - 2-2-2005 at 09:59

Quote:

Or, a nonaqueous solvent could be used.

Uh Yeh, I am not sure about that. A non-aqueous solvent commonly available would not stand up to these reaction conditions. I guess if you where suicidal you could liquify NO2 to N2O4 for a solvent

[Edited on 2/2/2005 by chloric1]

Esplosivo - 2-2-2005 at 11:40

Quoted from Inorganic and Theoretical Chemistry by F. Sherwood Taylor - Ninth edition [Pg.715-716].

Quote:

...Sodium chlorite is a powerful but stable oxidising agent, and is much used in bleaching cellulose materials. The chlorites of some of the heavy metals have been prepared by precipitation.

The part of interest is the last sentence of the quote. Which heavy metal which in the presence of NaOCl would remain unaffected and would precipitate as the chlorite?

[Edited on 2-2-2005 by Esplosivo]

S.C. Wack - 2-2-2005 at 14:14

Where does hypochlorite come in to this? Don't they mean precipitation from a soluble chlorite and the heavy metal nitrate? Or what?

Esplosivo - 3-2-2005 at 08:20

Quote:

Originally posted by S.C. Wack
Where does hypochlorite come in to this? Don't they mean precipitation from a soluble chlorite and the heavy metal nitrate? Or what?

When using the method Theoretic describe, which is the following reaction, NaOCl is produced.
NaNO3 + Cl2 --> NaClO2 + NOCl
I am interested in isolating the pure chlorite, isolating it from the mixture.

In the reference I found it states that the chlorite of a soluble salt, such as in this case the NaClO2, can be precipitated, say with BaCl2 solution to form Ba(ClO2)2 which is most probably insoluble. What I wanted to know is if the heavy metal would be affected by the NaOCl.

At the library today I found that the chlorites in solutions are thermally stable, to a certain extent, surely more than hypochlorites (no values were given). Maybe the hypochlorite can be disproportionated to give the chlorate and the chloride, and the chlorite ppted by some metal salts which does not ppt with the chlorate and chloride.

[Edited on 3-2-2005 by Esplosivo]

garage chemist - 4-2-2005 at 11:44

Esplosivo, NOCl (Nitrosyl chloride) is a gas! The Method you quoted gives pure Chlorite.

Yesterday I did a quick test: I mixed a spatula of NaClO3 with the same amount of oxalic acid, dissolved the mix in water and heated it, nothing happened, no gas evolution.
Even on addind a few drops of HCl, still nothing happened. So this doesn't work... thats bad, I will have to use ClO2.

Esplosivo - 4-2-2005 at 11:53

Quote:

Originally posted by garage chemist
Esplosivo, NOCl (Nitrosyl chloride) is a gas! The Method you quoted gives pure Chlorite.

Yesterday I did a quick test: I mixed a spatula of NaClO3 with the same amount of oxalic acid, dissolved the mix in water and heated it, nothing happened, no gas evolution.
Even on addind a few drops of HCl, still nothing happened. So this doesn't work... thats bad, I will have to use ClO2.

You're right, I'm sorry. I assumed that the reaction is carried out in aqueous conditions where NOCl would hydrolyse. Anyways. To liberate the ClO2/CO2/CO mixture concentrated sulfuric acid is needed, which reacts with both the solid oxalic acid and the solid chlorate - (The CO produced will react with some of the ClO2 being oxidised to CO2). This reaction is mentioned in reference books as the production of ClO2 to reduce possible explosion hazards of the gas. Use solid chemicals, not solutions. Like this it should work. Hope this helps.

[Edited on 4-2-2005 by Esplosivo]

Axt - 30-3-2005 at 19:19

Any further forays in chlorites Garage Chemist?

A quote from PATR is interesting, "Lead chlorite mixed with sugar detonates violently on percussion, and such a mixt has been used in detonators."

I searched for more info, since the PATR reference seemed to suggest a way of getting to KClO2 without ClO2, but most reference back to an Italian journal, I could only get its abstract as published in the London journal <a href="http://pulse.altlist.com/images/chlorites-abstract.pdf">here</a>, which doesn't give what I was after.

I think the methods already posted are going to be the easiest;

2 KClO3 + 2 H2SO4 + C2H2O4 -> 2 ClO2 + 2 CO2 + 2 H2O + 2 KHSO4

2 ClO2 + 2 NaHCO3 + H2O2 -> 2 NaClO2 + 2 CO2 + O2 + 2 H2O

<a href="http://www.sciencemadness.org/scipics/axt/chlorites.zip">This 4.3mb zip</a> contains the following from JACS online:

Chemistry of Chlorites
Production of Sodium Chlorite
Silver Chlorite
Small Scale Generation of ClO2
Sodium Chlorite - Properties and Production

[Edited on 9-12-2005 by Axt]

garage chemist - 31-3-2005 at 05:11

I didn't experiment further with chlorites, because I didn't want to risk destroying my glassware with a ClO2 explosion.

The articles you posted helped a lot. ClO2, when passed into NaOH/H2O2 mixture yields pure chlorite. H2O2 is too expensive for industrial use, but perfect for us.
I will use air as a carrier gas to dilute the ClO2 (instead of oxalic acid as a CO2 producer).
Air will be pumped throught a flask where conc. H2SO4 is reacted with NaClO3.
The resulting air/ClO2 mixture will be passed through a washing bottle containing NaOH/H2O2 solution.

JohnWW - 31-3-2005 at 05:37

With regard to Explosivo's post:
Except possibly at very low temperatures, heavy metal cations would be more likely to catalyze the decomposition of the chlorite(III) anion, rather than form a solid precipitate of the chlorites - which, if they could be obtained, are likely to be highly explosive.