I found a local source of sulfur from a farm supply store, problem is for some reason they mixed the sulfur with bentonite. The only property I saw
that might allow a simple separation is density, so I have been blending it with water in an old blender and separating the top and bottom layers.
This might be working, the top and bottom layers are very frothy, but appear to be a little different shade of yellow in color. But it requires
drying and repeating, and I'm not sure how to tell how pure the sulfur is. Any thoughts on how to purify this sulfur and bentonite mixture?Paddywhacker - 23-3-2014 at 21:40
I have never done this, but I have read on this site of dissolving out the sulphur with toluene, in which it is apparently slightly soluble.elementcollector1 - 23-3-2014 at 21:58
Xylene is also a contender, but the solubility is significantly less than in toluene.
I heard it had fairly good solubility in toluene, but I have yet to check the actual figure.Fantasma4500 - 23-3-2014 at 23:13
i once recrystallised sulfur from toluene, i used a decent amount although the toluene was impure, approx 60% toluene
it didnt have a very great solubility, perhaps heating would change this..?
you could however try melting the dry substance and pouring the sulfur off, i also have an idea of removing the intercrystalline H2SO4 making sulfur
unusable with KClO3 and KMnO4 by simply adding NH4OH to the molten sulfur -- supposing it will not make some interesting reaction insteadviolet sin - 23-3-2014 at 23:44
In my experience, with out heat, very little sulfur is dissolved in toluene. I once threw some in a pint mason jar with a decent hand full of sulfur
from a garden shop. in just under a year it has barely etched the surface of the pellets, and had tiny yellow crystals here and there on the jar.
looks like very little has been moved in all this time.
some random things I found, not all that helpful but a little interesting.
and https://www.google.com/patents/US2409408?dq=method+of+dissol...
"In fact, hot kerosene forms a highly satisfactory solvent for dissolving sulfur from its ores without the use of any additional reagent or catalyst.
The kerosene is heated to approximately 140 C. and the ore is either immersed therein or the hot kerosene is flowed through the ore"
Paddywhacker - 23-3-2014 at 23:52
Heating sulphur with saturated hydrocarbons is one method of producing H2S, so be careful.violet sin - 24-3-2014 at 00:29
thanks I was unaware of that and didn't see it mentioned while skimming over the patents. plante1999 - 24-3-2014 at 03:34
Drydistilling the sulphur could be the most efficient of the process since it is cheap in term of material.MrHomeScientist - 24-3-2014 at 06:29
That's what I was going to suggest - sublimation. Not sure what temperature this occurs at, but it's been claimed that S sublimes "easily," and I've
seen bags of sublimed garden sulfur before.WGTR - 24-3-2014 at 06:52
I've purified the garden store variety of sulfur by dissolving in hot toluene, decanting the clear
solution, then letting the sulfur crystallize out upon cooling.
Sulfur does sublime easily, but it's easy to go too far and melt the crystals into a tarry mess. When heating sulfur, I've observed that crystals will
be deposited further up on the walls of the tube, but the quantity is rather small. Using an improvised sublimator (something like a cold finger made
from an old test tube suspended inside a larger test tube) might be the best way to purify your sulfur. Be careful not to ignite the vapors; they're
smelly and toxic.
Subliming Sulfur (using vacuum?)
Xenon1898 - 24-3-2014 at 17:33
Thanks for those suggestions. I like the idea of sublimation for it's simplicity and effectiveness. For some reason I over looked the sublimation.
The melting point for monoclinic sulfur being 115 deg C and the melting point 444 C it shouldn't be too hard to avoid melting it with slow, controlled
heating.
Thanks for the hydrocarbon info, I hadn't considered those either. Since I can avoid the expense and slight carcinogenicity of an additional
hydrocarbon I will probably try the sublimation first.
I believe I will need to employ a vacuum for greatest effect, but I have to admit I don't have experience with the potential dangers of imploding
evacuated glassware. I will need to purchase a vacuum pump, probably a cheapish one (Harbor Freight, E-bay...) But should I just crank the vacuum
pump and not worry about it, or is there a guideline with how much vacuum is safe to expose the glassware to? I would avoid the larger, thin walled
glassware, but then my batch size is limited. What is the largest recommended glassware size, say using a heated round flask and condenser, that is
safe for a strong vacuum? Or should I try for a metal container and forget the breakable glass for vacuum sublimation?
[Edited on 25-3-2014 by Xenon1898]Oscilllator - 24-3-2014 at 19:51
While its true that subliming may be cheaper in the long run, if you have to buy a vacuum pump you are probably better of recrystallising it from
toluene. Because Sulfur is very soluble in hot toluene and much less soluble in the cold stuff, it should be relatively easy to do multiple
extractions quickly.
One other way could be to use a soxhlet extractor, although that would limit the batch size of your extraction and be quite slow.
Soxhlet
Xenon1898 - 24-3-2014 at 20:08
Yeah, but a Soxhlet is still a good idea. Maybe I should design some multi-unit leaching extractor tanks... ha just kidding. But seriously this
could be a perfect excuse to buy a vacuum pump, they seem useful for a well equipped chemistry lab. I could use it for purifying a lot of other
things as well via vacuum distillations too. But I remember seeing some safety film as a kid in elementary school where they applied a vacuum on a
bottle and it imploded, and shot glass shards all over the room. Does this ever happen to people using vacuum pumps using laboratory glassware?Oscilllator - 24-3-2014 at 21:07
Provided the glassware you use is of decent quality, and you don't use flat-bottomed equipment unless it is specifically vacuum rated you should be
fine. Things like RBF's, condensors etc are all designed to be able to take a vacuum.
One important point is that in terms of the amount of force that is put on the glassware, there is virtually NO DIFFERENCE between a cheapish vacuum
pump and interstellar space.
Provided the glassware you use is of decent quality, and you don't use flat-bottomed equipment unless it is specifically vacuum rated you should be
fine. .
Did you mean to say something like "unless it is specifically not vacuum rated..."?macckone - 25-3-2014 at 21:39
Provided the glassware you use is of decent quality, and you don't use flat-bottomed equipment unless it is specifically vacuum rated you should be
fine. .
Did you mean to say something like "unless it is specifically not vacuum rated..."?
Flat bottom flask should not be used unless they are vacuum rated.violet sin - 25-3-2014 at 23:07
found this looking for something else, has a list of solvents and sulfur solubility in it.
has a lot of good info on sulfur.
Sulfur is soluble in Halogenated Hydrocarbons such as Chloroform and Dichloromethane.annaandherdad - 26-3-2014 at 10:24
violet sin, I was interested in the high solubility of sulfur in liquid ammonia, according to the review article you posted. 38 gr / 100 gr solvent.
Must be under some pressure, since the temperature quoted is -20C. Dr.Bob - 26-3-2014 at 12:08
This is one case where a Soxhlet would be ideal, although I am pretty sure that you can buy nearly pure sulfur in many places, drugstores, farm and
garden supplies, online from pyrosuppliers. I have seen bags of 50-100 pounds for sale before, without any issues. Xenon1898 - 26-3-2014 at 20:02
Yes, sulfur is easily available in 99.9% purity. It is not so much that I can't buy it but rather the challenge of purifying a good sized bag of the
stuff I already have sitting around.
Method for Purifying Sulfur from CS2
Xenon1898 - 26-3-2014 at 20:20
This seems close enough to this thread to post here (hopefully) - see attachments for a method to produce very pure sulfur from CS2
Reference is George Brauer, ed., Handbook of Preparative Organic Chemistry, vol. 1, 1963, pg. 341-342
Very nice paper, thanks. It also has something I was looking for, vapor pressure curve data.
It might be interesting to try NH3 at room temp. [although it would be at 857 kPa at 20 deg C] Would be easy to remove the NH3 afterward too (I am
assuming). Too bad it doesn't list toluene.
[Edited on 28-3-2014 by Xenon1898]annaandherdad - 28-3-2014 at 08:06
I was wondering if you could make nice crystals of sulfur by dissolving it in liquid NH3. Of course you could try chlorinated hydrocarbons, too, but
the solubilities aren't as high as one would like. Besides, it would be something interesting to try with liquid NH3, which is pretty easy to make
with some dry ice (and a fume hood---all I have is a fan, but it's better than nothing).
Face Velocity
Xenon1898 - 28-3-2014 at 22:43
Yes a fan is better than nothing. If possible you could try to measure or estimate the velocity of the air moving where you need to work (where your
face is next to the lab bench). The industry standard for fume hoods is at least 100 feet per minute at the face of the fume hood. If you think all
the fumes are moving away to a safe place it is probably okay. Your health should be your top priority, then you can come back another day and do
more chemistry, a close second priority.... just my 2 cents.
