Sciencemadness Discussion Board - What Was this Substance?

What Was this Substance?

hodges - 17-12-2003 at 17:21

Many years ago when I was in high school, I messed around with chemistry. I messed with such things as making my own black powder, making a mixture that would ignite when I hit it with a broomstick (had red phosphorous and KClO4), various reactions with chlorine (I found for example that fine steel wool would ignite spontaneously in chlorine gas).

Anyway, I learned that sulfur would dissolve in a solution of sodium hydroxide. After I tried this, I didn't really have anything I could do with the resulting solution. So I decided to drop in a ball of aluminum foil, since I knew the NaOH would react with the aluminum. Anyway, it did react. The mass became very thick as it boiled, and it was a dark orange color. As the mixture cooled, it changed from dark orange to a very dark green. So dark, in fact, that it looked black, but a small amount in water would make the water a dark green. Later I added HCL to the original compound. It immediately bubbled vigorously, and from the smell it was definitely hydrogen sulfide. I was outside but close to the house and some of the gas apparently came in an open window, causing my mom to run out and ask what the *#! I was doing to make all that smell.

I started using this reaction as a source of H2S, since I didn't have any ZnS FeS to play with. I noticed sometimes when I would add the HCl not only would I get H2S but also I would see what appeared to be finely divided sulfur particles in the gas coming off the reaction. Don't know if the sulfur was formed in the air or whether it was just in the solution and brought into the air by the bubbles. Another thing I noticed about this substance was that if I put a small amount of it on a popsicle stick and set this in the sun, it would appear to "melt".

Does anyone know what this substance might have been? If anyone is interested, I'm sure it can be reproduced, since I did it many times.

Hodges

guaguanco - 17-12-2003 at 19:01

Quote:

Originally posted by hodges
...
Another thing I noticed about this substance was that if I put a small amount of it on a popsicle stick and set this in the sun, it would appear to "melt".

Does anyone know what this substance might have been? If anyone is interested, I'm sure it can be reproduced, since I did it many times.

Hodges

If you're talking about the substance that appeared to be finely divided sulfur, I'd bet that it was finely divided sulfur.

Sulfur (as I vaguely recall) dissolves in alkali to form various polysulfides in solution. The color changes with aluminum are interesting; some of the polysulfides are orange in color, and I would bet there's some equilibrium with free sulfur. The green color might be something weird between the Al(+3) and the polysulfides, but I'm just guessing wildly.

Saerynide - 17-12-2003 at 22:08

This sounds interesting. Can you post some pictures?

chemoleo - 18-12-2003 at 10:00

I agree with Guaguanco.
First you get sodiumpolysulphides (NaSx), which forms hydrate salts and is orange red in colour. The stuff smells bad even when dry, got 500g of it. Then, when you add the Al, you may get some weird double salt or NaxAlySz (but never heard of either of them) (Besides, the remaining NaOH might have dissolved the Al, whereby the resulting Al(OH)4- ion could form some type of complex with the Sx2-).
What strikes me is the green/black colour change. In any case, the fact that you get H2S and finely divided sulphur argues for the presence of polysulphides even after Al addition. H2Sx will of course decay in its free state (after the addition of HCl), resulting in finely divided sulphur precipitation and H2S liberation.
If my theory is correct, dissolve S in NaOH, until it is saturated and add HCl. You should then get H2S and S (precip)....

unionised - 18-12-2003 at 12:07

Al foil is not pure Al. There are other metals in it, notably copper and iron the last time I checked. These might account for part of the colour.

chemoleo - 18-12-2003 at 12:48

Doubt it. Many times I dissolved Al foil in NaOH, and the colour never went green (due to [Cu(OH)4]2-]. Hence certainly not large amounts of transition metals.
Hodges, did you ever try to evaporate the water from your green/black solution? What happens then? Do you get crystals? Different colours? Or does it all become a dark sludge? If it's crystals, it gotta be a single molecular product, so no polysulphides (with varying amounts of S, which prevent crystal lattice molecular alignments). That would be cool if it was. I would bet though, the latter is the case, where you dont get crystals.
Another thing, did you try to precipitate it by some other means, i.e. with ethanol/acetone? If it precipitates, you could analyse this, i.e. check whether Al is present in the precipitate. or sulphide. etc. Also, maybe try and add other metals than Al, add iron, or Cu, or lead...etc and their respective salts. They should all produce tell-tale products... maybe you can figure it out that way.

hodges - 18-12-2003 at 18:22

Quote:

Originally posted by chemoleo
Hodges, did you ever try to evaporate the water from your green/black solution? What happens then? Do you get crystals? Different colours? Or does it all become a dark sludge? If it's crystals, it gotta be a single molecular product, so no polysulphides (with varying amounts of S, which prevent crystal lattice molecular alignments). That would be cool if it was. I would bet though, the latter is the case, where you dont get crystals.
Another thing, did you try to precipitate it by some other means, i.e. with ethanol/acetone?

I recall having the substance around for several weeks after preparation sometimes, and it didn't appear to be drying. It was sort of a paste - a dark green to black paste - which appeared to melt (get thinner) when heated by sunlight. I'm sure there was excess NaOH present, which would explain it not drying. I never did try to precipitate it with any organic compounds.

I did notice a similar green color, although not nearly as strong, with another reaction. I used to make "fire sticks" using what amounted basically black powder rolled up in aluminum foil and then twisted tightly. If I let these sit out on the sidewalk overnight after burning, and the air was moist, in the morning they had a similar dark green color. They also gave off a faint odor similar to, but not exactly like H2S (and resulted in my father yelling for me to come outside to clean up the "dog turds" I left on the sidewalk!). I'm thinking the sulfur in the black powder may have reacted with the aluminum to form some type of sulfide. However, I never saw the orange color that I would get when doing the NaOH + S + Al reaction.

Hodges

chemoleo - 18-12-2003 at 19:37

ok, clearly what you have to do then is to dissolve S in NaOH until saturation, i.e. until no more S is being dissolved at that point. Filter the remaining S off.
Check the pH at that point. It should be alkaline (as H2Sx is a weak acid compared to the strong base NaOH). Then, neutralise that solution to pH 7 and add the Al (you may get sulphur preciptiation there, upon neutralisation). If you get bubbles at that pH (possibly H2, unless it smells of H2S), it would be odd, as the Al only dissolves in acidic or basic media. That would be interesting indeed

I guess at that point we would have to rethink everything.
Anyway, if the Al does dissolve at neutral pH 7, try and evaporate the H2O. It is possible that you will get crystals then, as the low pH will (by it's sheer hygroscopicity) imped crystal formation.
Next, when you dissolve the Al, do you get bubbles? maybe collect the bubbles, and burn it to see whether it's hydrogen. Otherwise, smell will tell the H2S evolution.

If Al DOESN'T dissolve at neutral pH, then you will at least know that the low pH is causing dissolution of the Al (hence OH- are a requirement). Then you also know that the [Al(OH)4]- ion has formed (in excess OH-).
What then remains is to figure out what the [Al(OH)4]- does in the presence of sulphur. In this case I would make two separate solutions, i.e. 1) make NaSx by dissolving S in NaOH until saturation, and 2) make [Al(OH)4]- by dissolving Al in NaOH until no solid particles remain.
Then mix the two. If the colour goes green, you KNOW at least that the reaction occurs by combining [Al(OH)4]- and NaSx. If that happens, post again, and we have a thought about this

Edit: Anyway, first things first. Dissolve S in NaOH as much as you can. then add HCl. If you get S precipitation, and H2S, then all is good

[Edited on 19-12-2003 by chemoleo]

blip - 19-12-2003 at 10:49

Al2S3 hydrolyses to H2S and Al(OH)3 in both water and moist air. It is yellow, so maybe there was some carbon left after the ignition. I guess the Al2S3 was made from the reaction between the "container" and some sulfur in the composition.

unionised - 19-12-2003 at 16:00

Chemoleo,
Sulphur is an oxidising agent. Polysulphide is another.
The reaction mixture containing Al and NaOH is a strong reducing agent. Can you see why copper would not dissolve if it were present in the Al foil. Did you (on any of the many occasions you dissolved Al in NaOH) ever notice the black stuff left over? Did you analyse it? Do you actually have any valid evidence for gainsaying my observation? Did you notice the bit where I said I had actually checked this?

Pure Al2S3 is not yellow BTW.

hodges - 19-12-2003 at 16:02

Unfortunately, I probably can't do much experimenting with this given that I presently live in an apartment. I just stumbled across this forum a few days ago, but it was almost 25 years ago when I did those experiments.

I have always been fascinated by just about any kind of science. I spend much of my free time working with electronics and microcontrollers. I like to revisit past interests from time to time. New developments occur and/or new information becomes available over time. But nowdays to keep from annoying my neighbors or being suspected of some type of evil activities, I stick mostly to things I can do inside. That means avoiding the energetic stuff and reactions that produce gases. Perhaps at some point I will talk a friend who lives in a house in the country into letting me experiment on his property. Though I will probably avoid the high energy stuf - I'm getting a bit too old to make explosives I think LOL.

Hodges

blip - 19-12-2003 at 16:42

Hmm, <a href="http://www.espi-metals.com/msds's/aluminumsulfide.pdf" target="_blank">this MSDS</a> said aluminum sulfide was yellow. While composing my post, I lost the formatting and ended up having to redo some things, like subscripts. I wasn't about to retype much, so I didn't include my original reference to it. Sorry.

chemoleo - 20-12-2003 at 10:41

Unionised, sorry, I don't doubt there are transisiton metals in Al foil. What I do doubt that they interfere in the reaction, producing the green/black colour. I did indeed notice the black carbon like stuff. But after filtration of that solution, it is clear, hence I should think that there isn't a large amount of soluble transition metal ions in the [Al(OH)4]- solution.
I fully agree that those conditions (i.e. high alkalinity) would NOT dissolve these metals.
Anyway, are you saying that finely divided Cu/Fe as a left over from Al dissolution would then somehow react with the NaSx to give this green/black colour?? If not, then this colour is certainly not related to the transition metals, if you see what I am saying

Best test would be to have those two solutions separate, i.e. *filtered* NaSx and [Al(OH)4]-, the former yellow/red/brown and the latter clear. Then, if you mix the two, and still get the green colour, I should not think that it is related to transition metals.

Well, I am off for holidays soon, and then I get the chance to test this...with chemically pure Al

unionised - 20-12-2003 at 14:44

If you look back at the first post in this thread you will see that the Al foil was added to a solution of sodium polysulphide. The oxidant was there first. Once the Al had dissolved there would still (probably) be some polysulphide left to oxidise the other trash.
(Add enough foil, and you might manage to reduce most of the other bits back out of solution again)

chemoleo - 20-12-2003 at 15:35

Hmm, I can see what you are trying to say. But have we thought this thru?

1. doesn't Al, when dissolved in NaOH, produce nascent hydrogen, just like with Zn and HCl? in this case, would this react with the polysulphide to form H2S or something? which in turn would be absorbed by left-over NaOH?
2. let's assume that solid precipitated copper and iron are released, upon dissolution of Al foil in NaOH/NaSx. This you say gets oxidised by the polysulphide (if 1. holds, not much of the polysulphide is left). In other words, this is similar of a reaction of copper/iron and sulphur, in their isolated states. This reaction requires quite a bit of activation energy to get started (you will know if you ever tried this). Isn't it thus unlikely that this occurs in solution? In addition, *even* if Cu/Fe reacts with an alkaline NaSx solution, you get Copper sulphide (I doubt the polysulphides of copper are stable - the only I know of are double salts, such as NH4CuS4). CuS/Cu2S is insoluble and precipitates out of solution, the former being black.

The question remains then, where does the green colour come from, if no Cu/Fe salts are *in solution*??

unionised - 20-12-2003 at 16:10

No, we haven't thought this through. I have though.
Copper polysulphide is (slightly) soluble in water. (Vogel's qualitative inorgnic analysis) I can't easily find data on all the other possible impurities in cooking foil.
Granted, the Al will reduce a lot of the polysulphides, air will tend to oxidise them back.
Given that polysulphides are yellow and that Al and Na compounds (when pure) are generally colourless (Unless the counter ion is coloured) what other explanation for the green colour springs to mind? My money is still on an impurity.
Adding a solution of a pure Al salt (a whole lot easier to get than pure Al foil) to the NaOH + S solution might prove this.

chemoleo - 20-12-2003 at 16:38

Quote:

No, we haven't thought this through. I have though.

You aren't exactly Mr.Humble ey?

Read this

Quote:

Originally posted by chemoleo
Best test would be to have those two solutions separate, i.e. *filtered* NaSx and [Al(OH)4]-, the former yellow/red/brown and the latter clear. Then, if you mix the two, and still get the green colour, I should not think that it is related to transition metals.

Well, I am off for holidays soon, and then I get the chance to test this...with chemically pure Al

What you say is precisely what I said already many posts ago. No need repeating it. (so much for thoroughly reading my posts :mad:

)

I accepted from the start that the colour might be a result of impurities (well as soon as you mentioned it), all I wanted to do since was to figure a mechanistic way. However, unlike you I don't *yet* exclude other possibilities. I dont see how the Sx2- ion should dissolve elemental copper etc in solution.
Interestingly, you said this ...

Quote:

The reaction mixture containing Al and NaOH is a strong reducing agent. Can you see why copper would not dissolve if it were present in the Al foil. Did you (on any of the many occasions you dissolved Al in NaOH) ever notice the black stuff left over?

Now, dont you say it doesn't dissolve? Ok probably you mean it does in the presence of NaSx. Now, see and this is what I struggle to explain as yet. Maybe I should mix some finely divided copper with NaSx, and see whether I get this colour...

But before we go any further, I will be testing this soon with chemically pure Al (from Merck!). Should be interesting! If this remains colourless, then I will test NaSx with copper. If it turns green, then all is good. End of discussion.

How about we start a poll?

Is it impurity or something genuine

- I am indecisive. You can start betting now

. I accept Paypal. As I am the person doing the actual experiment, I should get, say, 30% of the winning bets

, regardless who wins

[Edited on 21-12-2003 by chemoleo]

my 1.3 cents

Polverone - 20-12-2003 at 21:52

I can't contribute much to this thread, but I too have seen the blue/green compound. I saw it just a few days ago, actually. I was hoping to react sodium sulfide and aluminum... but photo grade sodium sulfide isn't anhydrous. How disappointing. In any case, I got a bright orange mess that became a blue/green mess (well, at least part went dark blue/green) when added to water.

I've seen the same blue/green color before, but only transiently, when I was trying to make sodium polysulfides in acetone solution. The blue/green appeared and faded within a few seconds as I added sulfur powder to a mix of acetone and NaOH, and I only saw it one time (later, similar attempts did not produce the interesting color).

unionised - 21-12-2003 at 11:53

I did the experiment.
I disolved sulphur in aqueous NaOH and added alum as a source of relatively pure Al. No green colour (just the yellowish white of the mixed S and Al2O3).

Disolving impure Al foil in NaOH may give a solution containing other materials. Some transition metals form carbonate complexes as well as hydroxy complexes for example. Even under the reducing conditions I don't see how you could rule out coloidal metals. That's why I said to use a salt (rather than the solution prepared from a metal foil, known to be impure, you see, I really had thought this through; sorry for not mentioning it)

I'm not exactly Mr humble. OTOH I didn't respond to a post that said "perhaps it's down to impurities by saying "I doubt it, because I did an experiment under different conditions and got a different answer."
When I pointed out that the presence of an oxidant might disolve the transition metals I was told I hadn't thought it through.
And, by the way, no, it's not a reaction of Cu or Fe with S in the isolated states. Its the reaction of Cu or Fe with a solution of polysulphide.
Mix some citric acid with some sodium bicarbonate and the reaction is very slow. Solids don't generally react very quickly. Add some water and you find out how much better the reaction goes if you have (at least) one of the reactants in solution.
No big suprise, the materials are in much better contact.
I don't have a problem with the idea that copper might not disolve in the absense of an oxidant but it does in the presence of one. Try Copper and HCl; not a lot happens unless you add an oxidising agent like peroxide.
Thats why I think it might disolve with polysulphide present. Polysulphide is both an oxidant and a complexant.
Since you say you still don't see how copper could disolve in a polysulphide.
Cu + Na2 Sx --> CuS + Na2 S(x-1)
Followed by the dissolution of the sulphide as a polysulphide complex.

Fortunately for you, it is unlawfull to bet on certainties so I can't take up the bet having done the experiment.
Now, what was it you wanted me to be humble about?

Edit.
I forgot to ask, if disloving Al (or nascent H, if you like) reduces all the polysulphide then, granted there is none left for me to argue that it oxidises the other metals, OTOH there is none left to make green stuff with Al, so what's the green stuff?

[Edited on 21-12-2003 by unionised]

chemoleo - 21-12-2003 at 12:21

lol... just very quickly:

1. yes it could be something colloidal, it was never mentioned before, so how could I rule it out?
2. I am aware of the complexant properties of Sx2-, and that's why I thought maybe the whole thing isn't so easy.
3. I *did* say I was going to use pure Al, not foil.
4. I said, *we* haven't thought it through, which includes everyone in this thread/forum. No personal attack here.
5. At least you were being challenged lol - although I am sure you will say, challenge, what? are you kidding? me mr. oxford grad??

6. bet on certainties? lol... I wouldn't yet go so far

, I reserve my judgement

Nice experiment anyhow. But, doesn't the 'I did an experiment under different conditions and got a different answer' statement hold, to a small extent? You appear to be leaving out that bit with 'nascent hydrogen evolution', right?
Well of course, this we don't take into account, because you have thought it through, while humbleness isn't even an issue due to certainties ?

PS lol online fights are funny.... shall we just stop? or maybe U2U, so we (which is myself and you) won't bore the rest of the forum lol

Edit: response to your own edit- this is exactly what I have been thinking, and what's confusing. This is why I reserve judgement on this. pz.

Edit 2: I apologise, Mr.Humble is not the right word. The opposite of Mr. Modest might be more descriptive

Oxford with capital 'O'? Lol, how about we start splitting hairs?

[Edited on 21-12-2003 by chemoleo]

unionised - 21-12-2003 at 13:02

OK, Outbreak of peace.

I forgot to mention that the other expt I did was to get some rather crude Al sulphate (shett Al + H2SO4) and add that to the polysulphide. I got a green colour. I don't know what the impurities were but it was a while since it saw any nascent hydrogen. (BTW, Oxford, gets a capital letter

)

Finally the experiments

chemoleo - 26-1-2004 at 18:07

Ok, over holidays I finally followed this up.

Have a look at this:
[img]http://www.sciencemadness.org/scipics/NaOH+S+Al.jpg[/img]

Reaction 1 is when I dissolved powdered sulphur in strong hot near -saturated NaOH, until it seemingly wouldnt dissolve anymore. It was odd that I was not able to dissolve much. I regret now that I didnt weigh out amounts etc, but at the time this would have been hampered anyhow by the lack of a decent scale. Anyway, it's a nice red solution that doesnt smell.

Reaction 2 was the product after I added either pure analytical grade aluminium, or aluminium foil . The reaction did not behave differently in either case, but it was interesting that the solution turned from clear red to a very dark red tinge (not clear, see pic), the viscosity inreased significantly (the resulting foam made trouble as it started to flow out of the beaker)

REaction 3 is after the Al dissolution subsided, and after a bit of water was added. The addition of water immediately changed colour to a deep brown/black.
Note everything was done twice, once with Al foil and pure Al.

REaction 4 is the same as 3, except diluted more. It started to get a slightly greenish tinge. Note, on the plate, tiny particulates could be seen.

Reaction 5 - left is pure Al, right is impure foil. This is what happened after I diluted the initially dark red gloopy solution with destilled water, - it progressively changed colour, and when I diluted it lots, and let it settle, I obtained a BLACK precipitate in either case! NOw what the hell is this?? I haven't got a clue. Also, note that the colour in the supernatant solution is clear!! No red-ness whatsoever.

Reaction 6: The left side contains the NaOH/S, the right side pure AL dissolved in NaOH conc until no more gas bubbles would evolve, and the middle is a mixture of left and right. Notice that no reaction occurs, the milkyness of the Al(OH3) and the colour of the NaOH/S - showing that the use of alum is no good.

It's too late to go thru it in detail, suffice to say that I lack an explanation. At least we have established it is not impurities. There is more behind it ....

PS sadly I didnt get the chance to have a look at the precipitate, to figure out what it really is

[Edited on 27-1-2004 by chemoleo]

The_Davster - 26-1-2004 at 18:55

I got interested in this black precipate, so I randomly typed in chemical names(aluminum sulfide, and then sodium sulfide) into google searching for their MSDS. I believe that the black precipate might be sodium sulfide.
http://chinatrona.com/brocher/english/msds/sodium%20sulphide...

Note that under stability it says it darkens under exposure to light and that it "Decomposes under the influence of moisture, water and acids, forming toxic and combustible gas (hydrogen sulfide)."

[Edited on 27-1-2004 by rogue chemist]

chemoleo - 27-1-2004 at 12:44

Huh? according to your description it is everything *but* sodium sulphide.
My sodium sulphide is red-brown (Na2Sx9H2O) (from a photography supplier, so not homemade), and in fact it didnt react with Al - unlike the NaOH/S + Al. That is presumably because the pure Na2S is not alkaline enough. Also, sodium sulphide doesnt turn black, and it doesnt precipitate once you *dilute* it!!!! It should be the way round.

Anyway - let's try and analyse this.

1. Conclusion No 1 is that it is not impurity in the Al that is causing this.
2. Conclusion No 2 is that an additional reaction occurs by dilution with H2O dest. *after* the Al has fully reacted with the NaOH/S
3. Conclusion No 3 is that Al by itself does not react with Na2S, at least the solution needs to be strongly alkaline with NaOH.
4. Pure Al(OH)3 is not sufficient for the reaction. It needs to be the Al[OH]4 ion, or a derivative thereof. I now wish that I hadnt dissolved Al in NaOH up to saturation (which yielded Al[OH]3 only, as I would then have been able to mix the clear un-precipitated Al[OH]4- solution with NaOH/S. Would have been interesting to see whether nascent hydrogen is necessary for the reaction.

I am baffled by the precipitation that occurs only after the fully reacted mix is diluted with H2O. The colour/texture changes from a gloopy dark red to liquid brown black to black greenish - just by the addition of H2O.

Could maybe someone else repeat this experiment, and do a few things on the black precipitate (I am confident it is reproducible)??
For instance, add HCl and see whetehr it develops H2S (smell). See whether it dissolves (reacts), and observe any possible colour changes.
Additionally try to add it to conc. NaOH. if nothing happens, heat it up. Observe again what happens.
Once we know this, we could take it from there....

Also, it should be interesting to evaluate what the initial red reaction product is (picture 2). For instance, what would happen if conc. H2SO4 was added? What would happen, if solid NaOH would be added? How does the colour from dark red proceed to black - titration... How stable is a drop of that red stuff in air? after a while? does it change?

Lots could be done... for someone who has the time & stuff & motivation

- and of course it seems you could use Al foil, or Al parts from bikes etc.

(I cant right now as I lack some of the chemicals here)

hodges - 27-1-2004 at 18:10

As I recall, just dissolving sulfur in NaOH yielded a yellow solution. It did not turn red until I added the aluminum. But that could have been due to the heat of the reaction. I know sulfur itself turns a reddish-brown color at some point after melting - the color could be for the same reason. I always saw the green color once the solution cooled. I think the green resulted not from dilution but instead from the fact that the solution cooled.

It would be interesting to note the effects of temperature on the solution. A good experiment would be to dissolve sulfur in cold NaOH solution and note the color. Then heat it without adding any Al to see if the red color develops. Then cool it to see if the red color goes away. I seem to recall adding HCl once before adding any Al, and the only thing that I remember happening was the sulfur precipitating again. No H2S gas. But that would be a good thing to retry - also testing the effect of temperature on the solution after the Al has dissolved.

I can't really try this reaction on a very large scale now. I might be able to do it inside on a small (test-tube size) scale some Saturday when my closest neighbor is away.

Hodges

Reaction Notes

hodges - 14-2-2004 at 12:40

I tried this today, on a small scale.

Here is a test tube of NaOH with some sulfur dissolved in it:

After heating this solution, it looked like this:

When allowed to cool after heating, it didn't look much different (I removed a small amount of the solution for another test):

After adding a small ball of aluminum foil and allowing the reaction to complete, here was the result (note that some of the solution was consumed/vaproized):

When adding dilute HCl to the NaOH/S solution that had been previously heated,
sulfur precipitated and a (very small) amount of H2S was noted. When adding dilute HCl to the final product, vigorous bubbling of H2S was observed.

I'm thinking the dark green / black product is Al2S3, in NaOH solution. Adding HCl neutralizes the NaOH (at least locally), allowing the Al2S3 to react with water to produce H2S. But is Al2S3 green?

Theoretic - 15-2-2004 at 07:18

Al2S3 reacts with any water there is, high pH is not a hindrance.

chemoleo - 15-2-2004 at 07:55

Once you added the Al foil, did you dilute the solution with dH2O? If not, try that. I bet you get the same thing as I did...
and still I am not any closer to the mystery of what this black precipitating stuff is!

Edit: Make sure you add Al until saturation! You can in fact add quite a lot!

[Edited on 15-2-2004 by chemoleo]

chemoleo - 23-6-2004 at 16:51

Well, remember this experiment?
NaOH + S + heat --> (Na)ySx ?
Then (Na)ySx + Al --> ???

Well, have a look at picture No 5 from my last post. You wil see a black precipitate. Strangely enough, now, after 5 months, it has turned orange-brown, and it looks deceptively like some iron oxido hydrate. Rusty colour. I honestly dont know what to think of it. The Aluminium I used was pure, delivered from Merck, of analytical quality. The NaOH was pure too. The sulphur - well there may have been some antimony, or whatever... but I doubt it would be more than one percent.
What do you think?
Unionised? (you may have to re-read the thread to afford a judgement

)
Seriously, this whole thing baffles me.
I dissolve sulphur in hot NaOH (aq), which forms a red solution. To this I add Al (pure, not foil). This yields initially a dark green solution/suspension (upon dissolution with destilled water), and then, a black preicipate. After a few months, this preciptitate takes on a rusty orange-brown-red colour.
All I can say is - these were clean reagents to my knowledge. What happened???
Why has it taken on this rusty colour? I'd think oxidation of course. Iron would be a candiate. But where would it come from, if these are analytical grade compounds?
Could it be some kind of polysulphide?
Chemistry is indeed a mystery at times! I am not sure whether to trust my reagents anymore

Theoretic - 24-6-2004 at 06:37

Maybe the thioaluminate ion AlS2- is formed? Or something else with Al-S bonds?

chemoleo - 25-6-2004 at 07:26

And you wouldnt have more info on this? thioaluminates... but why the colour change? Or this precipitating behaviour upon dilution with H2O?

unionised - 25-6-2004 at 13:57

Have you the equipment/ reagents to test the rusty lookig stuff t see if it's iron based?
Even the pink colour with salicylate would be a strong hint (just in case you don't have ICP/MS kit lying around in your garage)

the mysterious "green from sulfur" revisited

Polverone - 9-7-2004 at 17:07

I made some sodium polysulfides by mixing a small amount of water, photo-grade Na2S, and agricultural sulfur together in a narrow-necked glass bottle. Since I wanted to see how heat/light/air degraded the mixture, I plugged the top of the bottle with a cotton ball (to keep insects/debris out) and left it outside where it would be exposed to sun. It took about a week for most of the sulfur to dissolve. There was always some cloudiness left that I think may have been from non-sulfur components of the agricultural dusting sulfur. After more than a month of heat, light, and air, the solution had changed from the rich, deep polysulfide color to a weak yellow/green (mostly yellow). After several days of yellow/green color, I decided it was time to filter and examine.

I filtered through a plug of cotton. The filtrate was still quite alkaline, though considerably less than when I started. The solution, though it was still warm from sitting in the sun, had no scent. On addition of phosphoric acid to a portion of the filtrate, the temperature rose, bubbling commenced, and a precipitate of sulfur formed. I first smelled H2S, followed and overpowered by SO2. So it appears that a considerable amount of thiosulfate was formed, as I expected.

The now-clear filtrate still has a yellow color and is going back in the cotton-plugged bottle to see if it changes any further.

But - the green stuff! The plug of cotton that I used to filter my liquid was colored a dark green at the top from material it had trapped. The color struck me as very similar to that produced by reacting aluminum with sulfides. Coincidence?

The green color was destroyed (changed to white/yellow) by the addition of acid.

Nerro - 9-4-2005 at 02:44

I made the green stuff too a while back. I dissolved NaOH untill no more would dissolve and then I added 99.95% pure sulfer powder to the (boiling hot) solution. It boiled even more vigorously and reacted to form a deepred solution of what presume to have been a NaSx solution. I then added small pieces of Al (from a pure aluminium pipe) to the solution. They all dissolved so I then put in a 1ft section of the pipe. This partially dissolved within 6 hours. I removed the bit that remained. The solution had become near-black. But upon dilution and carefull closer inspection it was indeed (as in the aformentioned cases) green.

Upon addition of H2SO4 (40% solution) to small amount of the green solution a lot of H2S was evolved and a white precipitate of what I presume to be Al2(SO4)3 or Al(HSO4)3 (probably the former considering the fact that the H2S is removed from the equation as a gas.) was formed.

Also actual green crystals were formed in the solution which indicates a certain amount of stability. The crystals did not seem to react with the air when they were removed from the solution. The were needle shaped and darkgreen.

NaAlS2 maybe?

[Edited on 9/4/2005 by Nerro]

mick - 9-4-2005 at 07:13

Back to the original post.

The stuff that melted in the sun could be an emulsion of sulphur and conc NaOH, plastic sulphur, the NaOH stops it from dryng out.
Addition of Al should convert NaSS... SSNa to NaSS..SH.
In a concentrated set up the SH end could pick up some water and adding acid could give H2S and some SO2. I think H2S and SO2 could precipitate sulphur in the gaseous phase in damp air.
An old container of flaked sodium hydrosulphide has every colour from yellow to pink to green and darker on the surface.

I used to do a lot of poly and hydrosulphide reductions when I was young (sodium and ammonium) but they smell to much now and the waste would be a problem. They reduced organic disulphide bond well.

Curiosity, has anyone tried using the black residue from dissolving impure Al in NaOH for catalytic hydrogenation. It should have a large surface area and be fairly active unless it contains a load of carbon, silicon and stuff.

mick