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Author: Subject: iron disulfide and bleach?
symboom
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shocked.gif posted on 3-11-2011 at 19:33
iron disulfide and bleach?


it that even possible well tore apart a lithium battery that iron disulfide stinks so bad so i was thinking the h2s is being givin off ill just add bleach. thinking it should form sulfur yellowish brown powder precipitate evolved after awhile ive heared after a while then where did the
did it form iron sulfate? which should be soluble in water. so i was thinking its iron hydroxide which would be more likely. but as soon as i added the bleach i smelt no h2s nothing. so if it is where did the sulfur compound go
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[*] posted on 3-11-2011 at 20:01


Oxidized to sulfate I would imagine.



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Ozone
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[*] posted on 3-11-2011 at 20:11


Bleach contains a large amount of NaOH. You probably got colloidal Fe(OH)3, as a red-brown hydrate. It might oxidize slowly, over time, but the solubility of Fe3+ is very poor at pH > 4 (bleach is > 11).

The H2S was likely Na2S, which was probably (quite rapidly) oxidized to Na2SO4.

I suppose,

O3




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[*] posted on 3-11-2011 at 20:46


Quote: Originally posted by Ozone  
Bleach contains a large amount of NaOH. You probably got colloidal Fe(OH)3, as a red-brown hydrate. It might oxidize slowly, over time, but the solubility of Fe3+ is very poor at pH > 4 (bleach is > 11).

The H2S was likely Na2S, which was probably (quite rapidly) oxidized to Na2SO4.

I suppose,

O3


but i started with iron disulfide

FeS2 + NaClO --> first forming iron sulfate and sodium chloride
which then that should form sodium sulfate and iron chloride?

Fe(OH)3 + Na2SO4 + NaCl unbalanced

one things for sure this is a multistep reaction just not sure what reacted
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AndersHoveland
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[*] posted on 3-11-2011 at 22:39


I will make this easy for you:

(2)FeS2 + (15)NaOCl + (8)NaOH --> (2)Fe(OH)3 + (4)Na2SO4 + (15)NaCl + H2O

[Edited on 4-11-2011 by AndersHoveland]
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blogfast25
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[*] posted on 4-11-2011 at 06:52


Bleach doesn't oxidide sulphide to sulphate, rather it is oxidised to elemental sulphur. I know this from practical experience with bleach based H2S scrubbers: you end up with a bottle full of elemental sulphur!

Also the slag from thermite reactions boosted with CaSO4 contains CaS. To get rid of this smelly substance (CaS + 2 H2O === > Ca(OH)2 + H2S!) I treat it with thin bleach: again elemental sulphur forms.

Hypochlorite (bleach) is a mighty oxidiser but it isn't capable of oxidising sulphides to sulphates or sulphur to sulphate. If it did have that capability then every Dick, Tom and Harry amateur would making sulphuric acid from sulphur and bleach!

(2)FeS2 + (15)NaOCl + (8)NaOH --> (2)Fe(OH)3 + (4)Na2SO4 + (15)NaCl + H2O

... is a pretty equation: it just doesn't happen in the real world, Anders.

To oxidise FeS2, roasting in hot air is used. That oxidises the sulphur to SO2 or sulphate:

2 FeS2 + 11/2 O2 --> Fe2O3 + 4 SO2 ΔH = - 1666 kJ/mol (1)
3 FeS2 + 8 O2 --> Fe3O4 + 6 SO2 ΔH = - 2381 kJ/mol (2)
FeS2 + 3 O2 --> FeSO4 + SO2 ΔH = - 1054 kJ/mol (3)

(Source: http://www.saimm.co.za/Conferences/Sulphur2009/101-110_Runke...)



[Edited on 4-11-2011 by blogfast25]




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[*] posted on 4-11-2011 at 09:37


Quote: Originally posted by blogfast25  
To oxidise FeS2, roasting in hot air is used. That oxidises the sulphur to SO2 or sulphate
The roasting of sulfide ores for SO2 production is an important minority source of SO2 gas for commercial H2SO4 production. If I remember my market share numbers right, it's somewhere around 25% - 40% depending on region. As I recall, it's used to preprocess certain ores for primary metal extraction, prior to the smelting operation.
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[*] posted on 4-11-2011 at 09:53


I dug around on the net and found the attached paper. Because I am at work, I don't have time to write a lengthy post.

Cheers,

O3

Attachment: Bunyakan et al 2005 rxn of H2S with NaOCl.pdf (319kB)
This file has been downloaded 5163 times





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[*] posted on 4-11-2011 at 10:52


Quote: Originally posted by blogfast25  
Bleach doesn't oxidide sulphide to sulphate, rather it is oxidised to elemental sulphur. I know this from practical experience with bleach based H2S scrubbers: you end up with a bottle full of elemental sulphur

Sulfur should be just an intermediate in the oxidation. At such basic pH's, as the sodium hypochorite solutions have, the colloidal sulfur rapidly disproportionates to sulfite and sulfide which are both easily oxidized further. Once the pH drops to less basic values, due to the conversion of the hypochlorite, you get an accumulation of elemental sulfur which is now unable to disproportionate fast enough. This might explain your observations.




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[*] posted on 4-11-2011 at 11:09


Yes, that would explain it.



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AndersHoveland
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[*] posted on 4-11-2011 at 14:57


Quote: Originally posted by blogfast25  
Bleach doesn't oxidide sulphide to sulphate, rather it is oxidised to elemental sulphur.


"Hydrogen sulfide produces an odor which smells like rotten eggs. It reacts with chlorine to form sulfuric acid and elemental sulfur (depending on the ratio or reactants)."

page 161 of "Prudent practices in the laboratory: handling and disposal of chemicals", clearly states that hyochlorite oxidizes Na2S to Na2SO4.
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[*] posted on 13-11-2011 at 13:18


Quote: Originally posted by blogfast25  
"Bleach doesn't oxidide sulphide to sulphate, rather it is oxidised to elemental sulphur. I know this from practical experience with bleach based H2S scrubbers: you end up with a bottle full of elemental sulphur!

Also the slag from thermite reactions boosted with CaSO4 contains CaS. To get rid of this smelly substance (CaS + 2 H2O === > Ca(OH)2 + H2S!) I treat it with thin bleach: again elemental sulphur forms.

Hypochlorite (bleach) is a mighty oxidiser but it isn't capable of oxidising sulphides to sulphates or sulphur to sulphate. If it did have that capability then every Dick, Tom and Harry amateur would making sulphuric acid from sulphur and bleach!"

While this comment is predominately correct, is it completely? Can excess H2S create some sulphates and, if so, how?

Per Watt's Dictionary Chemistry, HClO can oxidize S to H2SO4 (but I would not expect anything but a very dilute solution). However, we must first form the required HClO from the basic NaClO/NaOH solution by reacting with a weak acid. Perhaps, the very dilute basic NaOH in the bleach can have its pH slightly lowered by an excess of H2S. Also, besides H2S dissolved in water, any dissolved CO2 for that matter, could now slowly liberate HClO.

NaClO + H2CO3 --> NaHCO3 + HClO

Granted, however, gaseous H2S on contact with NaClO would be primarily reduced as previously noted:

H2S (g) + NaClO + H2O --> 2 H2O + S (s) + NaCl

But, assuming some HClO is around, my conjecture as to possible reaction paths to explain Watt's statement:

2 HClO + S --> SO2 + 2 HCl

SO2 + H2O <---> H2SO3

H2SO3 + HClO --> H2SO4 + HCl

and, as such, perhaps some sulfates would indeed be formed.
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[*] posted on 13-11-2011 at 15:07


Quote: Originally posted by AJKOER  


any dissolved CO2 for that matter, could now slowly liberate HClO.

NaClO + H2CO3 --> NaHCO3 + HClO



~0.066 M HClO ? that's some pretty strong stuff!!!! (assuming complete dissasociation of carbonic acid into the acid and carbonate salt)

[Edited on 13-11-2011 by sternman318]
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[*] posted on 13-11-2011 at 15:40


sterman318:

Yes, carbonic acid is stronger than HClO, and hence the reaction proceeds.

I have done this (and in agreement with the observation of another source), the reaction unfortunately proceeds very slowly as compared to using say vinegar on NaClO to yield HClO.

Note, my comment only conditionally supports Nicodem, quoted below:

Quote: Originally posted by Nicodem  
Quote: Originally posted by blogfast25  
Bleach doesn't oxidide sulphide to sulphate, rather it is oxidised to elemental sulphur. I know this from practical experience with bleach based H2S scrubbers: you end up with a bottle full of elemental sulphur

Sulfur should be just an intermediate in the oxidation. At such basic pH's, as the sodium hypochorite solutions have, the colloidal sulfur rapidly disproportionates to sulfite and sulfide which are both easily oxidized further. Once the pH drops to less basic values, due to the conversion of the hypochlorite, you get an accumulation of elemental sulfur which is now unable to disproportionate fast enough. This might explain your observations.


My conditions are relating to how the solutions becomes more acidic (or not), and how long the process may take (slow I suspect, but could be wrong on this point having not performed the reaction myself).
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