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Author: Subject: Sodium sulphite spontaneous oxidation
Boffis
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[*] posted on 21-9-2020 at 02:14
Sodium sulphite spontaneous oxidation


I recently carried out the preparation of sodium dithionate (Na2S2O6) using the method in a standard college level chemistry book. The process involves the preparation of silver sulphite and then its thermal decomposition.

The reaction is almost quantitative:
2Na2SO3 + 2AgNO3 = 2Ag + 2NaNO3 + Na2S2O6

The preparation uses exactly these proportions. When I carried out the preparation I found the the precipitation of silver was low and I had to add about 10% excess of sulphite to ppt all of the silver. The sodium sulphite I used was Fison's Lab grade reagent but its fairly old (late 1980s vintage). It is nominally the heptahydrate but has effloresced a little and so the water content is a little lower than theory (hence I should have added a slight excess of sulphite to start with!).

When I began to investigate the sulphite I discovered that it contains nearly 10% sodium sulphate after drying in the oven. I used the classical method of assay by BaCl2 in sufficient hot dilute HCl to drive off the sulphite as SO2 first then added the BaCl2.

The label claims a max 0.2% sulphate as sodium salt. This raises the question as to whether sodium sulphite slowly oxidises on long exposure to air. Has anyone else ever come across such an occurrence or can add any comment.

Next question is can the sulphite be purified by recrystallisation?
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[*] posted on 21-9-2020 at 03:58


Quote: Originally posted by Boffis  
This raises the question as to whether sodium sulphite slowly oxidises on long exposure to air. Has anyone else ever come across such an occurrence or can add any comment.

Next question is can the sulphite be purified by recrystallisation?

Yes it oxidises in air.
Probably rather faster in an oven.

"A heptahydrate is also known but it is less useful because of its greater susceptibility toward oxidation by air.[1]"
From wiki


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Boffis
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[*] posted on 21-9-2020 at 09:15


Thanks Unionised, Hmmm, thats interesting. I wonder how easy it is to recrystallise?
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Boffis
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[*] posted on 26-9-2020 at 14:23


Quote: Originally posted by Boffis  
Thanks Unionised, Hmmm, thats interesting. I wonder how easy it is to recrystallise?


Difficult!
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[*] posted on 26-9-2020 at 18:27


I like how you answered your own question :)
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[*] posted on 26-9-2020 at 23:57


The heptahydrate is the form, which was sold 40 years ago. As a teenage boy I also used the heptahydrate, big crystals, hexagonal. Nowadays, the sodium sulfite you can buy always is the anhydrous form in the form of small crystals (fine powder). The anhydrous form is much more resistant against aerial oxidation. My sample of hydrated sulfite had become a white powder over several years and most of that white powder was sodium sulfate.



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[*] posted on 27-9-2020 at 01:03


Quote: Originally posted by woelen  
The heptahydrate is the form, which was sold 40 years ago. As a teenage boy I also used the heptahydrate, big crystals, hexagonal. Nowadays, the sodium sulfite you can buy always is the anhydrous form in the form of small crystals (fine powder). The anhydrous form is much more resistant against aerial oxidation. My sample of hydrated sulfite had become a white powder over several years and most of that white powder was sodium sulfate.


I also have some heptahydrate. Really old sample, but still have reducing properties. I don't know how much sulfate it contains, but still have quite high sulfite content.




If you are interested in aqueous inorganic chemistry look at https://colourchem.wordpress.com/main-page/

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Boffis
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[*] posted on 27-9-2020 at 09:22


I tried to recrystallise it by dissolving 300g of crude salt in 300ml of warm water. It dissolve rapidly and easily before the solution had reach even 60 C. I vacuum filtered it and allowed it to cool for a couple of days but no crystals formed so I took a few of the more transparent crystals from the jar and rinsed them in water to remove the white powder and then seeded the solution with them. By the following morning large six-side plates of sodium sulphite heptahydrate had formed. I was busy that day so I left them for another day by which time masses of slender colourless prisms of unmistakeable sodium sulphate decahydrate had formed on top; bollocks I though and warmed the solution to try and dissolve them again but they would not dissolve completely. Instead a fine granular white ppt formed that would not dissolve even when heating almost to boiling! I left the stuff to cool and masses of the platy heptahydrate formed but comtaminated with the white granules (anhydrous sodium sulphate or sulphite?). I'll have another go shortly.
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[*] posted on 27-9-2020 at 11:55


Quote: Originally posted by Boffis  
but they would not dissolve completely. Instead a fine granular white ppt formed that would not dissolve even when heating almost to boiling!

That could easily be sodium sulphate; it has a strange solubility curve with a maximum at about 32C (though that will probably be different in the presence of the sulphite.
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[*] posted on 27-9-2020 at 13:53


@unionised; I think you are right. I have just reheated the beaker slowly and when the crystal mass began to break up I put a stir bar in and everything dissolve quickly before the temperature reached 45 C!! It appears that this recrystallisation is going to require a very narrow temperature window, about 15 C at the lower end and about 40-45 at the top end.
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[*] posted on 2-10-2020 at 07:47


Update. From just under 300g of crude, 40 year old (like woelen's), heptahydrate I recovered roughly 130g a good material that gives only a slight ppt with BaCl2 in dilute HCl. I have used most of it immediately so that I don't have this problem again. Evaporating down the filtrate gave a mixture of sulphate and sulphite crystals so I discarded it.
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