testimento
Hazard to Others
Posts: 351
Registered: 10-6-2013
Member Is Offline
Mood: No Mood
|
|
Easiest way to produce sulfur?
What might be the easiest way to produce elemental sulfur from sulfates, mainly?
|
|
|
ScienceSquirrel
International Hazard
Posts: 1863
Registered: 18-6-2008
Location: Brittany
Member Is Offline
Mood: Dogs are pets but cats are little furry humans with four feet and self determination! 
|
|
As a product from a modified Le Blanc soda process?
http://en.wikipedia.org/wiki/Leblanc_process
Large scale yucky smelly chemistrythat needs a big garden with distant neighbours.
|
|
|
WGTR
National Hazard
Posts: 971
Registered: 29-9-2013
Location: Online
Member Is Offline
Mood: Outline
|
|
Probably via sulfides as an intermediary, as ScienceSquirrel mentioned.
From here:
If you can access this preview page here, it lists a eutectic of Na2SO4 and Na2S at 740*C.
Na2SO4 melts at 884*C, and Na2S melts at 1176*C, but a partial reduction of Na2SO4 with charcoal
should allow a process that stays molten under 900*C.
Going from there, a partial oxidation of H2S with oxygen in the Claus Process.
Of course, as was already mentioned, this whole process just stinks. Not only that, but the gasses
are toxic, although you probably already know that. Maybe plante1999 has tried this one.
[Edited on 1-12-2014 by WGTR]
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
From H2S, simply shake with a chlorine Bleach (NaOCl) solution (or even that other oxygen based bleach). This is a recognized scrubbing method for the
removal of Hydrogen sulfide gas. No high temperature synthesis required, just a sulfide/acid to generate the H2S.
The product is a fine Sulfur suspension, which would be very reactive. However, I would expect filtering to be an issue. Reaction:
NaOCl (aq) + H2S (g) --> NaCl (aq) + S (s) + H2O
---------------------------------------------------------------------------
From SO2, the direct reduction with activated carbon at 650 C in an oxygen free atmosphere for over 7 hours of heating is not rationale as not all of
the Sulfur cleanly separates from the Carbon!
Reference: http://www.ems.psu.edu/~radovic/PLW/1968_6_917_Stacy_Carbon....
Based on Table 4 observed products, my take on the implied reaction at 650 C in an O2 free environment:
9 SO2 + 10 C --Heat to 65O C--> 9 S + 2 CO + 8 CO2
[Edited on 13-1-2014 by AJKOER]
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
There is very little NaOH added to Bleach. But, if a concern add a little vinegar, which would also address any added Na2CO3. The created HOCl will
also react with H2S as follows:
H2S + HOCl --> S (s) + H2O + HCl
But do avoid an excess of the HOCl, as some of the Sulfur could be oxidized with time.
Now, while I have outlined a cheap and simple method, working with poisonous H2S is not only unpleasant, it can be dangerous. Once you can no longer
smell the gas, it is time to exit. Unbelievably, Hydrogen sulfide, which can also be absorbed through your skin, is not too far from HCN in its
toxicity! So try to perform all reactions in a sealed vessel outdoors (and away from your neighbors).
[EDIT] What I find unsettling about H2S is that if you do manage to be alive at the end of the preparation, it could still be your last! One can amass
a lethal dose without being fully aware of your impending death in the coming hours, in essence, a 'walking dead' scenario. Not cool, IMHO.
[Edited on 13-1-2014 by AJKOER]
|
|
|
bismuthate
National Hazard
Posts: 803
Registered: 28-9-2013
Location: the island of stability
Member Is Offline
Mood: self reacting
|
|
Adding HCl to sodium thiosulfate (which is very cheap) will percipitate sulfur which could then be filtered off.
Although why don't you just buy sulfur? It's insanley cheap.
[Edited on 13-1-2014 by bismuthate]
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
Here is a cheap safe path. Add FeS2 (iron pyrite so called fool's gold) to a solution of Bleach and vinegar (Acetic Acid or HAc). Form a second
solution of H2SO4 and NaCl (or, just use HCl if available). Combine solutions in a closed vessel (capable of handling any pressure build-up, perhaps a
compressed plastic bottle) and shake. Pressure issues can also be addressed by testing increasingly coarse forms of FeS2 and adjusting acid strength.
Some of the expected reactions:
HAc + NaOCl --> NaAc + HOCl
H2SO4 + 2 NaCl --> Na2SO4 + 2 HCl
FeS2 + 2 HCl --> FeCl2 + H2S (g) + S (s)
H2S + HOCl --> H2O + S (s) + HCl
Net: HAc + H2SO4 + NaOCl + 2 NaCl + FeS2 = NaAc + Na2SO4 + FeCl2 + 2 S(s) + HCl + H2O
Whatever Hydrogen sulfide formed is rapidly changed to Sulfur. Again, avoid having an excess of NaOCl (or HOCl) as some of the Sulfur may be oxidized.
Note, the chlorine bleach is treated with vinegar to resolve any NaOH or Na2CO3 additives, but if deemed not an issue, one can forget the vinegar. Be
mindful that some bleaches have significant amounts of NaCl.
[EDIT] This source (http://water.me.vccs.edu/concepts/chlorchemistry.html ) suggests the use of a little extra NaOCl, but not an excess of HOCl as the Sulfur could be
oxidized all the way to H2SO4. To quote:
"Let's consider one example, in which chlorine reacts with hydrogen sulfide in water. Two different reactions can occur:
Hydrogen Sulfide + Chlorine + Oxygen Ion → Elemental Sulfur + Water + Chloride Ions
H2S + Cl2 + O2- → S + H2O + 2Cl-
Hydrogen Sulfide + Chlorine + Water → Sulfuric Acid + Hydrochloric Acid
H2S + 4 Cl2 + 4 H2O → H2SO4 + 8 HCl "
[Edited on 13-1-2014 by AJKOER]
|
|
|
ScienceSquirrel
International Hazard
Posts: 1863
Registered: 18-6-2008
Location: Brittany
Member Is Offline
Mood: Dogs are pets but cats are little furry humans with four feet and self determination!
|
|
Iron pyrites is iron disulphide, not iron sulphide.
It does not react with hydrochloric or sulphuric acids.
It can be dissolved in concentrated nitric acid to give a solution of iron III nitrate with the evolution of sulphur dioxide.
I very much doubt that it will react with a mixture of vinegar and bleach.
[Edited on 13-1-2014 by ScienceSquirrel]
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
Actual ScienceSquirrel, thanks for differentiating between FeS2 and FeS, but I use the formula and the common name of Fool's gold, so I doubt if there
is any confusion.
I am not sure on your source for the reactivity of FeS2, but here are a few quotes from Atomistry.com (link: http://iron.atomistry.com/iron_disulphide.html ), to quote:
"Aerated waters charged with calcium carbonate appear to decompose pyrites very slowly at the ordinarv temperature, yielding limonite. Thus: -
4FeS2 + 15O2 + 3H2O + 8CaCO3 = 2Fe2O3.3H2O + 8CaSO4 + 8CO2;
Distilled water, in the presence of air, slowly oxidises pyrites to ferrous sulphate and sulphuric acid."
Wow, explain that last chemical equation with some simple concepts! And, also interestingly, there is more:
"A convenient chemical method has been devised, however, which enables a discrimination to be made with certainty. It consists in boiling the mineral
with iron alum, containing 1 gram of ferric iron and 16 c.c. of 25 per cent, sulphuric acid per litre. The proportion of sulphur oxidised in the case
of pyrites is 60.4 per cent, of the total sulphur contained in the mineral; in the case of mareasite it is only 18 per cent. The reaction may be
considered as taking place in two stages, namely: -
FeS2 + Fe2(SO4)3 = 3FeSO4 + 2S,
6Fe2(SO4)3 + 2S + 8H2O = 12FeSO4 + 8H2SO4. "
So apparently boiling FeS2 with an acid Iron salt, Iron and some H2SO4 came determine what is real gold (and differentiates marcasite from pyrite in
the quantity of free Sulfur formed).
So, in the cold (no boiling), FeS2 may have a reactivity issue with H2SO4. Upon boiling, still no direct confirmation of the formation of H2S from
FeS2, per ScienceSquirrel comment. More work (testing and research) needed to confirm my suggested path. This source (http://www.sciencedirect.com/science/article/pii/S0008443398... ) does note, interestingly, to quote:
"Various aspects of electrochemical phenomena involved during MnO2–FeS2 dissolution in dilute HCl have been examined"
However, I am amazed on how vinegar (and not necessarily pure Acetic acid) plus regular impure Bleach (NaOCl) is able to attack chunks of Iron (from a
magnet) in the cold (and these solutions are dilute). I suspect the reaction is actually, in part, electrochemical in nature, and complex as there is
an absence of general discussion in the literature (what I observed when someone designated as a pundit doesn't want to get it wrong). So, I for one,
would not be surprised if fine powdered FeS2 does react, to an obviously small extent, in the cold vinegar/bleach while H2SO4 on FeS2 struggles at
best. Note, per the FeS2/O2/CaCO3 equation above, O2 in the cold (weaker than HOCl) slowly attacks FeS2.
--------------------------------------------------------------------
[EDIT] Why do I claim the reaction of Fe and HOCl is electrochemical and otherwise chemically complex, here is some of my proposed reactions:
Preliminary chemical reactions:
Fe + HOCl --> FeO + HCl
HOCl --FeO--> HCl + 1/2 O2 (See http://books.google.com/books?id=mGVbIoW2lNAC&pg=PT462&a... probably via the formation and decomposition of Iron hypochlorite, or I suspect,
via the presence of Nickel (and then NiO) in the Iron)
Electrochemical reactions (see, for example, for the Iron-Air Battery: http://patentscope.wipo.int/search/en/detail.jsf;jsessionid=... )
At cathode: HOCl + H+ + e- --> 1/2 Cl2(g) + H2O (similar to the Bleach battery)
2nd half-reaction: 1/4 O2 + 1/2 H2O + e- --> OH- (Iron-air battery)
At anode: Fe + 2 OH- ⇒ Fe(OH)2 + 2 e- (Iron-air battery)
Total net electrochemical reaction:
HOCl + Fe + 1/4 O2 + 3/2 H2O --> Fe(OH)2 + 1/2 Cl2 (g) + H2O
Or: 2 HOCl + 2 Fe + 1/2 O2 + H2O --> 2 Fe(OH)2 + Cl2 (g)
Subsequent possible reactions:
Fe + 2H20→ Fe(OH)2 + H2 (Iron-air battery)
Fe(OH)2 + 2 HCl --> FeCl2 + 2 H2O
2 FeCl2(aq) + Cl2(g) → 2 FeCl3(aq) (see Wiki http://en.wikipedia.org/wiki/FeCl3 )
and as Sodium acetate has been cited as a catalyst, the possible formation of acetate ligands to add to the complexity. The reactions I have detailed
account for the observed formation of Cl2, FeCl3 and H2.
[Edited on 14-1-2014 by AJKOER]
|
|
|
WGTR
National Hazard
Posts: 971
Registered: 29-9-2013
Location: Online
Member Is Offline
Mood: Outline
|
|
However you make it, one useful bit of info to know is the moderate solubility of sulfur in hot toluene:
https://www.sciencemadness.org/whisper/viewthread.php?tid=39...
This property can be used to separate sulfur from a dried mess of salts and other crud.
[Edited on 1-13-2014 by WGTR]
|
|
|
plante1999
International Hazard
Posts: 1936
Registered: 27-12-2010
Member Is Offline
Mood: Mad as a hatter
|
|
I indeed made sulphur in the past using various means.
Sulphur dioxide and hydrogen sulphide plus moisture works really well. And hot charcoal reduction of sulphur dioxide works too.
The best source for sulphur from sulphate is from the sulphide produced in the Leblanc process, descriped in one of the PDF I made avaible in the
prepublications. Acidifing such sulphide make a good amount of hydrogen sulphide, which can be burned to sulphur dioxide.
I never asked for this.
|
|
|
ScienceSquirrel
International Hazard
Posts: 1863
Registered: 18-6-2008
Location: Brittany
Member Is Offline
Mood: Dogs are pets but cats are little furry humans with four feet and self determination!
|
|
I have tried treating pyrites with concentrated acid, fusion with alkali. etc and it is not a reactive material under less than harsh conditions.
AJKOER should buy some pyrites, put the pipe and port away and try some experiments.
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
Quote: Originally posted by ScienceSquirrel  | I have tried treating pyrites with concentrated acid, fusion with alkali. etc and it is not a reactive material under less than harsh conditions.
AJKOER should buy some pyrites, put the pipe and port away and try some experiments. |
As I noted above, one source, "Electrochemical phenomena in MnO2–FeS2 leaching in dilute HCl. part 2 : studies on polarization measurements", by
R.K. Paramguru1 at the Regional Research Laboratory (Council of Scientific and Industrial Research) Bhubaneswar, India, link: http://www.sciencedirect.com/science/article/pii/S0008443398... claimed that there was an electrochemical phenomena involving MnO2–FeS2
dissolution in dilute HCl.
I find this interesting given this equation:
MnCl2 + HOCl + H2O = MnO2 (s) + 3 HCl (see page 4 at http://www.ecs.umass.edu/cee/reckhow/courses/371/371l22/371l... )
as upon adding HCl to both sides, noting that HCl + HOCl = Cl2 + H2O:
MnCl2 + Cl2 + 2 H2O = MnO2 (s) + 4 HCl
which is, of course, the reaction for the generation of Chlorine written in reverse. My point, the reaction is reversible.
Now, instead of adding HCl to each side, lets add FeS2 to each side:
MnCl2 + HOCl + FeS2 + H2O = MnO2 (s) + FeS2 + 3 HCl
Now, the author cited a dissolution for the right side of this equation. But, this implies that FeS2 dissolves in a solution of MnCl2 and HOCl.
Now, at this point, I would normally have to agree with ScienceSquirrel and perform the reaction. However, I just discover a source that claims it is
in fact the case. See page 141 at http://books.google.com/books?id=QeDTeUalP0IC&pg=PA141&a... which cites a commercial process for the recovery of gold contained in Iron pyrite
via HOCl in an aqueous chloride solution. So my new method is relatively new, but apparently, not original.
|
|
|
Zyklon-A
International Hazard
Posts: 1547
Registered: 26-11-2013
Member Is Offline
Mood: Fluorine radical
|
|
Just don't use your sulfur for pyro, as it will contain a lot of acid, from most of these methods.
|
|
|
bfesser
Resident Wikipedian
Posts: 2114
Registered: 29-1-2008
Member Is Offline
Mood: No Mood
|
|
AJKOER, take heed of this (facetious?) advice. Your constant spamming of poorly referenced theoretical reactions
is getting out of hand.
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
Bfesser:
If you read my sources, apparently the electrochemical phenomena involving MnO2–FeS2 dissolution in dilute HCl has been well researched/documented.
This was actually stated, per my recollection, in my cited references (note the Journals, Electrochemistry and Metallury). Now, the published year was
2000 in the cited mining application, which was, unfortunately, after I (and perhaps many others) graduated.
In my opinion, chemistry is a large and growing science, but the only thing 'poor' here, speaking for myself, is perhaps my personal knowledge of
these important commercially significant disciplines, and, with respect to the gold mining application, apparently well funded research at that.
I think it is also important to explain how I originally arrived at my speculation (since confirmed) of the action of HOCl in the presence of a
chloride on FeS2, by experimentation reported on ScienceMadness.org no less. I refer anyone interested to my numerous prior thread references and
experimentation results on the so-called Bleach battery. Even prior to this, I attempted to explain the reaction of Fe in a NaOCl/Acetic acid solution
based on my personal observations on ScienceMadness and elsewhere. As such, based on my experimentation and continuing research as is publicly
documented, I formulated my speculation presented to ScienceSquirrel, and those are the facts, for the record to the best of my recollection.
[EDIT] Some references not previously cited:
"Kinetics and mechanism of MnO2 dissolution in H2SO4 in the presence of pyrite" in Journal of Applied Electrochemistry, February 1999, Volume 29,
Issue 2, pp 191-200.
"Oxidation of FeS2 and FeS by MnO2 in Marine Sediments" by Axel Schippers & Bo Barker Jørgensen at Max Planck Institute, for Marine
Microbiology, Celsiusstrasse, Germany.
"Electrochemical phenomena in MnO2–FeS2 leaching in dilute HCl. part 3: manganese dissolution from indian ocean nodules", published in Canadian
metallurgical quarterly, 1998.
"Biogeochemistry of pyrite and iron sulfide oxidation in marine sediments", published in Geochimica et Cosmochimica Acta, Vol. 66, No. 1, pp.
85–92, 2002.
"Electro-generative mechanism for simultaneous leaching of pyrite and MnO2 in presence of A. ferrooxidans", link: http://www.ysxbcn.com/down/upfile/soft/200856/20085621187488...
"Galvanic interaction between sulfide minerals and pyrolusite", published in Journal of Solid State Electrochemistry, March 2000, Volume 4, Issue 4,
pp 189-198
......
[Edited on 14-1-2014 by AJKOER]
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
A simple path to Sulfur via FeS2 may be to just heat it in an oxygen free environment per one old text with a partial release of Sulfur.
A most recent source (see page 2 at http://link.springer.com/article/10.1023/A:1010135425009#pag... ) notes upon heating between 683 K to 789 K, the following reaction:
FeS2 --> FeS + S
The FeS can then be used to form H2S for processing as previously indicated.
Heating in O2 results in the loss of the Sulfur (to SO2 and SO3).
--------------------------------------------------------------------------------------
Another old text "Watts' Dictionary of Chemistry, Revised and Entirely Rewritten, Volume 3, page 64 (link: http://books.google.com/books?id=st7nAAAAMAAJ&pg=PA64&am... ) states that concentrated HCl decomposes FeS2 with the formation of H2S and S. My
take on such a reaction would be:
FeS2 + 2 HCl (Conc) --> FeCl2 + S + H2S (g)
One source of such HCl could be by the action of concentrated H2SO4 on NaCl.
--------------------------------------------------------------------
Bfesser:
The Wikipedia article on FeS2 does not have a single reaction or any mention of recent work. Not even a simple thermal decomposition reaction. Other
topics, for example, active awards (see https://www.google.com/url?sa=t&rct=j&q=&esrc=s&... ) "Modeling and Analytical Surface Analysis of Iron Pyrite (FeS2) for Thin-Film
Photovoltaics", $1,602,797.
[Edited on 15-1-2014 by AJKOER]
|
|
|
sasan
Hazard to Self
Posts: 92
Registered: 22-2-2014
Location: TEHRAN / IRAN
Member Is Offline
Mood: Radiative
|
|
try reacting thiosulfates with acids.sulfur would be form in crystaline form.filter out the sulfur.i think it is the easiest the way to sulfur
Na2S2O3 + 2 HCl → 2 NaCl + S + SO2 + H2O
if you dont have thiosulfate then see the sodium thiosulfate paje in wiki
any ideas?
|
|
|
sasan
Hazard to Self
Posts: 92
Registered: 22-2-2014
Location: TEHRAN / IRAN
Member Is Offline
Mood: Radiative
|
|
another path to sulfur>
wiki:Sodium thiosulfate is also used in analytical chemistry. It can, when heated with a sample containing aluminium cations, produce a white
precipitate:
2 Al3+ + 3 S2O32− + 3 H2O → 3 SO2 + 3 S + 2 Al(OH)3
reacting with HCl(aq) is better
|
|
|
phlogiston
International Hazard
Posts: 1375
Registered: 26-4-2008
Location: Neon Thorium Erbium Lanthanum Neodymium Sulphur
Member Is Offline
Mood: pyrophoric
|
|
A fast and easy method to obtain hydrogen sulfide from calcium sulfate (plaster of paris): mix with aluminium powder and ignite. An extremely
exothermic reaction (thermite-like) will take place. Break up the residue and add concentrated acid to release hydrogen sulfide.
It has been mentioned already, but just to make sure: H<sub>2</sub> is extremely poisonous. Although it is notorious for how bad it
smells, you can not rely on smell to detect the gas, because it quickly deadens your sense of smell. Therefore, it is very easy to create dangerous
levels of this gas unknowingly. Please be extremely careful when you pursue these experiments!
-----
"If a rocket goes up, who cares where it comes down, that's not my concern said Wernher von Braun" - Tom Lehrer |
|
|
Theoretic
National Hazard
Posts: 776
Registered: 17-6-2003
Location: London, the Land of Sun, Summer and Snow
Member Is Offline
Mood: eating the souls of dust mites
|
|
| Quote: Originally posted by ScienceSquirrel | I have tried treating pyrites with concentrated acid, fusion with alkali. etc and it is not a reactive material under less than harsh conditions.
AJKOER should buy some pyrites, put the pipe and port away and try some experiments. |
Are you sure about this!? that might not be pyrite, or the percentage of pyrite in your sample is actually quite small, leading to the appearance of
unreactivity.
because the original 'acid test' is to react it with (sulfuric) acid, and if it is pyrite, it will bubble vigorously, emitting clouds of H2S (gold
will not react)
pyrite will spontaneously oxidize in air also, this is a major source of mine acid runoff and spontaneous underground ignition.
|
|
|
Zyklon-A
International Hazard
Posts: 1547
Registered: 26-11-2013
Member Is Offline
Mood: Fluorine radical
|
|
From Wikipedia:
| Quote: |
Sulfur may be prepared by bubbling hydrogen sulfide steam through an aqueous solution of Iodine, forming hydroiodic acid (which is distilled) and
elemental sulfur (this is filtered). |
I may use this method to make some hydroiodic acid. (Not for cookin'.)
[Edit] Does anyone what an "aqueous solution of Iodine" means?
[Edited on 6-3-2014 by Zyklonb]
|
|
|
![[*]](images/xpblue/default_icon.gif){kind=icon}
