Sciencemadness Discussion Board - Sulfur as an oxidizing agent

Just tried elemental sulfur in 100mL of a 1M oxalic acid solution, no reaction evident. This is actually surprising, since the table of reduction potentials suggests that sulfur is the most favorable to be reduced in the hypothesized reaction, leaving oxalate to be oxidized to carbon dioxide.

[Edited on 28-9-2015 by Upsilon]

I'm happy you tried it and I'm not that surprised. The problem might be the insolubility of sulfur and ionic nature of oxalic acid in water.

You might have success by melting oxalic acid with sulfur in a test tube. Again, be careful please... use very small amounts.

In short, the reaction is thermodynamically favourable, BUT that says nothing about kinetic barriers. However, kinetic barriers can be overcome, heat is the usual method, failing that or when there are selectivity issues, one can try to employ a catalyst to help speed things up.

Good luck and thanks for trying it out!

[Edited on 28-9-2015 by deltaH]

Praxichys - 28-9-2015 at 11:49

Use the action of acid on Iron (II) sulfide. I assume your aversion to metal sulfides is caused by their unavailability and/or violence of manufacture.

Heat a well-mixed, stoicheometric (1:1 molar) pile of iron filings and powdered sulfur with a torch. The reaction to FeS2 is nice and gentle; it looks as if the pile is slowly smoldering. When cool, grind into a fine power and remove any unreacted iron with a magnet. Store in a stoppered bottle. If you do this in a small tin can, you can manufacture >100g each time, enough for over 1mol H2S.

Exposure of the sulfide to acids will release H2S. Strong mineral acids are preferred but it will work with many organic acids as well. Be careful. H2S kills your sense of smell at ca. 100ppm, causes unconsciousness at ca. 300ppm, and is rapidly fatal at 1000ppm. (Keep in mind 1000ppm is 0.1wt%)

[Edited on 28-9-2015 by Praxichys]

Upsilon - 28-9-2015 at 12:18

Looks like I'll be doing that then. Oxalic acid and sulfur have conveniently similar melting points so it shouldn't be too much of a hassle. I'll try it out sometime this week.

Praxichys, that may very well be a better method, but I'm going to try this out anyway and see if it's more efficient. If not, then oh well, that's what experimentation is all about. That, and I don't have any iron filings on hand

deltaH - 28-9-2015 at 12:33

They may have similar melting points, but not sure if they will be miscible when molten. That said and with no water, oxalic acid can't ionise easily and in its unionised form it is symmetrical and so non-polar, so maybe it works. Experimentation tells all. At the very least, the higher temperatures should help a lot!

I look forward to your results

[Edited on 28-9-2015 by deltaH]

Brom - 28-9-2015 at 14:41

Do you not have powdered aluminum? If so the Al2S3 route is an extremely simple path to H2S

Upsilon - 28-9-2015 at 14:50

Quote: Originally posted by Brom

Do you not have powdered aluminum? If so the Al2S3 route is an extremely simple path to H2S

I considered that as well, but aluminum powder isn't exactly cheap. I would argue that even requiring an acid to free the H2S, the iron sulfide route is still cheaper than using aluminum powder.

Brom - 29-9-2015 at 07:17

Ok I understand the quote button now. I was having trouble with the edit feature. After I would edit I could not post it after it was fixed. Thanks for the help. Sorry for not being so good at the computer thing. I guess I'm getting old

Brom - 29-9-2015 at 07:37

I figured it all out now. Thanks again for the help

mayko - 29-9-2015 at 08:52

Doesn't oxalic acid decompose on heating to release carbon monoxide?

deltaH - 29-9-2015 at 09:35

Quote: Originally posted by mayko

Doesn't oxalic acid decompose on heating to release carbon monoxide?

That's an interesting point, it does decompose to carbon monoxide, carbon dioxide and water vapour between 150°C - 200°C (see the thermogravimetric plot below from http://www.azom.com/article.aspx?ArticleID=3024) , so again, please be careful. Even if unsuccessful, an odourless toxic gas will [nevertheless] be evolved.

I guess success will depend on which has the faster kinetics at those temperatures: oxidation or the [self] decomposition of oxalic acid. Reminds me of the behaviour of hydrogen peroxide where there is a similar competition, albeit in a reversed role as oxidising agent (oxalic acid being a reducing agent).

[Edited on 29-9-2015 by deltaH]

AJKOER2 - 29-9-2015 at 13:55

Here is another aqueous approach, to quote Atomistry on H2S, which contains dated historical accounts (link: http://sulphur.atomistry.com/hydrogen_sulphide.html ):

"(b) Sulphur can also be reduced to hydrogen sulphide at the ordinary temperature if "nascent" hydrogen is used; thus, powdered sulphur yields some hydrogen sulphide if treated with aluminium, tin, iron or zinc and hydrochloric acid, the result being improved by the additional presence of acetic acid or alcohol, which will increase the solubility of the sulphur. The reduction can also be effected electrolytically by having powdered sulphur in contact with a platinum cathode immersed in dilute acid, e.g. hydrochloric acid."

My updated rendition of the pathway is electro-chemical based on the so called Aluminum Sulfur battery (see, for example, "Aluminum and sulfur electrochemical batteries and cells", US patent 5,413,881, link: http://www.google.com/patents/US5413881 ). The cell reaction on which I am suggesting a modification is given generally as:

2 Al + 3S + 3OH- + 3H2 O → 2Al(OH)3 + 3HS- ; Ecell =1.8 V

Now, if one adds acid (say HCl) to both sides of the equation, with sufficient acid, one could have, for example:

HCl + NaHS → NaCl + H2S (g)
--------------------------------------

[Edit] A side comment is that all of the cited metals (aluminium, tin, iron and zinc) associated historically with "nascent hydrogen" formation above have a possibly interestingly corresponding electro-chemical basis as I argued above. Specifically, for example, existing metal-air batteries include aluminium-air, tin-air, iron-air and zinc-air.

[Edited on 30-9-2015 by AJKOER2]

aga - 29-9-2015 at 14:09

Why is it AJKOER2 and not good old AJKOER ?

Your link-finding abilities are legendary : all fascinating and almost 100% relevant.

Upsilon - 29-9-2015 at 14:13

I believe I have a reaction! To prevent decomposition of the oxalic acid, I created a stable temperature environment. I dissolved about 32g of potassium hydroxide in 30 mL of water to create a solution that boils at ~120C. I then added 0.5g of sulfur and 1.97g of oxalic acid dihydrate to a test tube and mixed thoroughly. The test tube is then submerged into the potassium hydroxide solution, which is carefully monitored for temperature so that it does not boil (it wouldn't matter if it did, I just don't want to be splashing it everywhere).

With this setup I have successfully melted the mixture, and I am observing bubbling. It could possibly be water vapor, but I highly doubt it. It has been bubbling for a decent amount of time now; I would think all water would have been driven off by now. To be sure, I will let it run to completion; theoretically nothing should be left after all is done according to the equation:
H2C2O4 + S -> H2S + 2CO2

AJKOER2, looks interesting. I'll look more into it once I have finished my experiment.

Metacelsus - 29-9-2015 at 14:28

To be more sure, you could also try passing the gas through a copper sulfate solution, and watch for a CuS precipitate.

blogfast25 - 29-9-2015 at 14:34

With this setup I have successfully melted the mixture, and I am observing bubbling. It could possibly be water vapor, but I highly doubt it. It has been bubbling for a decent amount of time now; I would think all water would have been driven off by now. To be sure, I will let it run to completion; theoretically nothing should be left after all is done according to the equation:
H2C2O4 + S -> H2S + 2CO2

I think you should have smelled H2S by now. Its smell threshold is only 0.47 ppb.

Upsilon - 29-9-2015 at 14:35

Quote: Originally posted by Cheddite Cheese

To be more sure, you could also try passing the gas through a copper sulfate solution, and watch for a CuS precipitate.

I'll have to do that next time then. I don't have any test tube stoppers with holes. Even if I did I wouldn't really want to try and rig something up while the reaction is taking place.

Upsilon - 29-9-2015 at 14:50

I think you should have smelled H2S by now. Its smell threshold is only 0.47 ppb.

If I were sitting here over it in stagnant air, yes. But I know it is a dangerous gas, so I have taken some precautions. Those precautions being outdoors with an enormous box fan blowing over the setup. Low tech but surprisingly effective.

Either that or it's because I have a headcold right now and dont have much of a sense of smell currently. Probably not the ideal time to do experiments with hydrogen sulfide, but I've been out here with it for over an hour already, and I'm still very much alive. Like I said I am upwind from the big box fan so I doubt very many H2S molecules are making their way over to me.

[Edited on 29-9-2015 by Upsilon]

aga - 29-9-2015 at 14:52

I think you should have smelled H2S by now. Its smell threshold is only 0.47 ppb.

=470 parts per Trillion ?

Jeez.

blogfast25 - 29-9-2015 at 15:36

Yes. Probably an evolutionary early warning system: it's also very toxic.

Upsilon - 29-9-2015 at 17:27

Just finished with the experiment. Not everything was reacted; curiously, the contents of the tube formed a solid column. About half of this column stayed at the bottom of the tube, and the upper half migrated up the tube somehow and did not melt. There is also a little bit of foamy unreacted sulfur left at the bottom of the tube. I weighed the contents and it came out to 1.93g, which is actually greater than the mass of the reactants (excluding the weight of the water in the hydrated oxalic acid). It has been raining here all day and is quite humid, so maybe a lot of water was absorbed. I can't say that these are conclusive results, but the molten portion of the mixture did bubble quite steadily for multiple hours until all of the molten mixture was gone.

blogfast25 - 29-9-2015 at 18:06

@Upsilon:

As I wrote above: had the reaction proceeded even at a small conversion rate you would have smelled rotten eggs. S8 is a rather sluggish reactant in many situations.

[Edited on 30-9-2015 by blogfast25]

Upsilon - 29-9-2015 at 18:16

@Upsilon:

As I wrote above: had the reaction proceeded even at a small conversion rate you would have smelled rotten eggs. S8 is a rather sluggish reactant in many situations.

[Edited on 30-9-2015 by blogfast25]

Not sure if you saw my post earlier or not, where I said that I had a giant box fan blowing over the setup which I was sitting upwind from, as well as having very poor sense of smell currently due to a head cold.

blogfast25 - 29-9-2015 at 18:46

Hmmm... even in those conditions 0.47 ppb is easily reached. Next time have some filter paper soaked in CuSO4 placed 'down wind'? It seems a little odd to set up a nice experiment that has no means of detecting a reaction product that's so easily... detected!

[Edited on 30-9-2015 by blogfast25]

Upsilon - 29-9-2015 at 19:21

[Edited on 30-9-2015 by blogfast25]

I know; next time I'm going to try actually bubbling the gas into a copper sulfate solution to see if I get a precipitate. I was hoping that I would be able to run the whole reaction to completion. Having nothing left in the tube would be pretty clear evidence of the hypothesized reaction, I think.

Another oddity to note is that my potassium hydroxide temperature buffer solution got way hotter than the boiling point elevation calculation predicted. I only wanted to elevate the boiling point by 20C; the calculation called for about 32g of KOH in 30mL of water. But it got darn close to 150C on a few occasions and showed no signs of boiling. A lot of water did evaporate off, though. Maybe the evaporating water caused the solution to supersaturate, thus allowing a very high BP elevation? When I let it cool, I got these really cool KOH crystals forming.

deltaH - 29-9-2015 at 20:55

Using KOH is a big mistake IMHO. If you did form H2S, it would have most surely formed K2S which is completely odourless at very high pH. You see H2S is a weak acid. Similarly, if you formed CO2, then you would, in fact, have formed K2CO3.

Also, strong electrolyte solutions will MAKE SURE no sulfur can dissolve because the sulfur in non-polar and the solution highly ionic. This then means that your reaction becomes purely a surface heterogeneous one, which with sluggish kinetics won't do much.

Finally, oxalic acid can form and be sold as the dihydrate. This wouldn't be a good version to use for this reaction because the melt would then be strongly ionic since oxalic acid is moderately a strong acid. Anhydrous oxalic acid, on the other hand, is a non-polar molecule because it is symmetrical and so might even dissolve a little sulfur at elevated temperature. Do you know for sure whether you have anhydrous oxalic acid or the dihydrate?

I strongly suggest you just take a TINY amount of sulfur and oxalic acid (preferably anhydrous) neat and gently melt them in a test tube immersed in some hot oil if desired, to test this. You might even use excess oxalic acid. The sulfur should be reduced faster than the oxalic acid decomposes with any luck.

Glad you're taking the appropriate precautions.

Thanks for posting your results and good luck!

[Edited on 30-9-2015 by deltaH]

madcedar - 30-9-2015 at 05:44

Lime sulphur controls a range of fungal diseases and pests on fruit trees, tomatoes, grapes, roses etc. It's available from your nearest horticultural store and it's basically a solution of calcium polysulphides. Just add HCl and Bob's your uncle.

How to kill yourself using the "detergent method" of suicide

blogfast25 - 30-9-2015 at 06:21

Using KOH is a big mistake IMHO. If you did form H2S, it would have most surely formed K2S which is completely odourless at very high pH. You see H2S is a weak acid. Similarly, if you formed CO2, then you would, in fact, have formed K2CO3.

The KOH solution was only there as an isothermal heating medium, surrounding the reactor. Small amounts of any H2S formed would have been absorbed but unlikely all of it. The use of KOH solution as a heating medium is a strange choice but it doesn't invalidate the experiment, IMHO.

[Edited on 30-9-2015 by blogfast25]

deltaH - 30-9-2015 at 08:05

Sorry, brain freeze, I completely missed that immersed in a bath of it
I should stop skimming threads early in the morning before rushing off lol

Okay, perfectly fine then and it's not looking good for the method.

Perhaps oxalic acid dihydrate is being employed and this is forming a strongly ionic melt and so sluggish to the kinetics? If Upsilon can confirm he is using anhydrous oxalic acid, then I think this thing is officially dead down to bad kinetics, but again, I'd try heating it purposefully to decomposition to see if there's any hope.

****************

Lime-sulfur can easily be prepared from lime and sulfur and I think that should react cleanly with oxalic acid to form H2S, CO2 and calcium oxalate.

Lime-sulfur consists of ionic polysulfides, mainly S5(2-) which ought to be reduced in solution by oxalic acid much more easily than the unactivated non-polar S8 molecule. In some ways, the polysulfides can be viewed as an 'activated' form of sulfur.

Lime-sulfur recipe from wikipedia:

Quote:

A New York State Agricultural Experiment Station recipe for the concentrate is 80 lb. of sulfur, 36 lb. of quicklime, and 50 gal. of water (192 g sulfur, 86 g quicklime of high purity per 1000 g (1 liter) of water). Ca. 2.2:1 is the ratio (by weight) for compounding sulfur and quicklime; this makes the highest proportion of calcium pentasulfide. If calcium hydroxide is used, an increase by 1/3 or more (to 115 g/L or more) might be used with the 192 g/L of sulfur. If the quicklime is 85%, 90%, or 95% pure, use 101 g/L, 96 g/L, or 91 g/L; if impure hydrated lime is used, similarly increase its quantity. Avoid using lime that is less than 90% pure. Boil for an hour, stirring and adding small amounts of hot water to compensate for evaporation.

Lime-sulfur may react from solution with oxalic acid to form H2S, just another suggestion to get around the sluggish kinetics of sulfur.

[Edited on 30-9-2015 by deltaH]

blogfast25 - 30-9-2015 at 08:19

http://www.sciencemadness.org/talk/viewthread.php?tid=63790&...

He's using the dihydrate. It's easy to dehydrate: just heat to about 200 C, off the top of my head.

deltaH - 30-9-2015 at 08:54

Quote:

He's using the dihydrate. It's easy to dehydrate: just heat to about 200 C, off the top of my head.

Isn't 200C too high, the TGA plot I attached above suggests that it decomposes fully by 200C?

If it dehydrates easily, then it should dry if molten and held say at 120C for some time, where it won't decompose. This may be the bubbling upsilon observed, i.e. just steam.

Anyway, as I said, the observation posted thus far doesn't bode well for the idea, I suggest moving on to preparing lime-sulfur, which is VERY easily done and then try reacting oxalic acid with that. Should solve the kinetic problem.

deltaH - 30-9-2015 at 09:06

Apologies for its morbidity, but I just read that a large number of people in Japan commit suicide by making H2S by reacting lime-sulfur containing bath salts with acid. See this bulletin on the disturbing trend https://info.publicintelligence.net/MAchemicalsuicide.pdf

So that lends more weight to the lime-sulfur + oxalic acid route. Note, while any acid mixed with lime sulfur yields some H2S... enough to even kill you, I think one would get much better yields of H2S on a sulfur basis if using oxalic acid since you still need a reducing agent to convert 4/5 of the sulfur present theoretically.

Again, WARNING anyone reading this about the extreme toxicity of H2S gas!

Upsilon - 30-9-2015 at 09:31

It just seems strange to me that it would continue bubbling steadily for hours if it was just steam. The sulfur and oxalic acid formed a clear honey-colored melt which slowly vanished over time. Studying the test tube closely, the bubbles would appear at random places in the melt and rapidly grow until taking off upwards.

deltaH - 30-9-2015 at 09:40

Quote:

The sulfur and oxalic acid formed a clear honey-colored melt which slowly vanished over time.

hmm, this is indeed an interesting observation and more encouraging that something happened than what I initially understood things to be, particularly the observation that it vanished. It could be that the sulfur did form H2S or it formed something else. Not sure what 'else' could have formed... I can speculate carbon oxysulfide (COS), water and CO2 perhaps? Although COS is also described to have a bad smell.

The honey colour seems indicative of dissolved sulfur, so I'd say you probably successfully dehydrated the oxalic acid.

Speculative alternate hypothesis equation:

HOOCCOOH(l) + S(l) + heat => H2O(g) + COS(g) +CO2(g)

blogfast, is there a simple filter-paper-type test for COS?

If things indeed went this route, then at least there's an interesting and simple way to make COS

There's always the lime-sulfur route to H2S...

[Edited on 30-9-2015 by deltaH]

Upsilon - 30-9-2015 at 17:51

I'm going to attempt the experiment again likely tomorrow; this time I'll cover the tube with filter paper soaked in CuSO4 solution. If it's really H2S gas then it should slowly turn black. Of course some basic copper carbonate would form as well (because of the CO2 gas), but that shouldn't matter since it's not even close to black.

Upsilon - 1-10-2015 at 15:03

Just finished with my second attempt. Much more conclusive results this time. I was able to melt all of the mixture this time, and almost all of it reacted. I even took some videos and photos so you all can see what happened (videos are in GIF form because I'm too lazy to upload a real video

). Details:

0.5g of sulfur powder and 1.97g of oxalic acid dihydrate were thoroughly mixed together in a test tube. A piece of filter paper was fixed over the mouth of the test tube with a rubber band. The test tube was then submerged in hot canola oil, which never got hotter than 150C. The filter paper on the test tube opening was doused with copper(II) sulfate solution via pipette, but not so much was added as to have some run down into the reaction mixture. The reaction was allowed to run for approximately 6 hours. Within a couple of hours the copper(II) sulfate-soaked filter paper began to turn noticeably black. By the time the heat was turned off, only a very small amount of melt was left, and reaction rate had drastically slowed. Strange crystal structures were found above the leftover melt.

Weight of reactants (excluding the water in the oxalic acid dihydrate): 1.91g
Weight of tube contents at end: 0.98g

Reaction soon after the contents melted (color is not as obvious as before since the canola oil is colored as well)

Reaction about 4 hours in

Reaction about 5 and a half hours in

Reaction immediately after turning off heat, about 6 hours in (strange crystal formation is visible partially submerged in the melt; strange crack-like formations on test tube visible, though they are not actually cracks in the glass)

Tube recovered from the oil after cooling (crystal is hard to see but present, crack-like formations again visible)

Top of tube, illustrating the effect on the CuSO4-doused filter paper

Inside of the recovered filter paper, clearly showing the probable CuS precipitate

And that's all! I would say this is pretty conclusive that the hypothesized reaction took place.

blogfast25 - 1-10-2015 at 16:48

@Upsilon:

Nice experiment and write-up (we really need a lot more of the latter on SM, so keep up the effort!)

What all this means (probably) is that reaction rate is very small and what little reaction does happen is driven by removal of two of the reaction products (CO2 and H2S) from the mix - Le Chatelier principle.

It doesn't appear to be a practical way to produce H2S though...

[Edited on 2-10-2015 by blogfast25]

Upsilon - 1-10-2015 at 17:03