rstar
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A new sulphate salt
Hi geeks,
I added some sodium sulfate to equal amount of sulfur powder, and then heated it. Some amount of sulfur reacted, some melt and vaporized. However the
product was like whitish powder with a reddish hint. The picture is given:
It shows reducing properties, as i added it to a dilute KmnO4 solution(looking pink), it changed the color of the solution to yellowish, probably due
to MnO2 formation.
Can anyone say what is this salt ?? 
I think it can be any one of these: Na2SO3, Na2S2O3, Na2S2O5, or any other salt or a mixture of salts. I suggest if anyone has those ingredients can
also perform this experiment.
Thanks in advance for any help 
[Edited on 2-10-2011 by rstar]
"A tidy laboratory means a lazy chemist "
- Jöns Jacob Berzelius
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peach
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Try adding a drop of hydrochloric acid to it.
If it stinks of rotting swamps and sewers, you've made sodium sulphide.
[Edited on 2-10-2011 by peach]
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Jor
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How can you get sodium sulfide (oxidation state -2) by heating sulfur (oxidation state 0) and sulfate (6+) ? 
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rstar
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Quote: Originally posted by Jor  | | How can you get sodium sulfide (oxidation state -2) by heating sulfur (oxidation state 0) and sulfate (6+) ? |
Thats right,
The substance(s) formed must contain sulfur, in some oxidation state between +1 to +5.
Oxidation state of S in different compounds are:
+2 = Thiosulfate (S<sub>2</sub>O<sub>3</sub><sup>2-</sup>
+3 = Dithionite (S<sub>2</sub>O<sub>4</sub><sup>2-</sup>
+4 = Metabisulfite (S<sub>2</sub>O<sub>5</sub><sup>2-</sup> and Sulfite (SO<sub>3</sub><sup>2-</sup>
+5 = Dithionate (S<sub>2</sub>O<sub>6</sub><sup>2-</sup>
+<sup>5</sup>/<sub>2</sub> = Tetrathionate (S<sub>4</sub>O<sub>6</sub><sup>2-</sup>
[Edited on 2-10-2011 by rstar]
"A tidy laboratory means a lazy chemist "
- Jöns Jacob Berzelius
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AndersHoveland
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This is rather interesting. I do not know what salt you have formed, but it is probably a combination of thiosulfite and metabisulfite, or possibly
even dithionate.
(2)Na2SO4 + (2) S --> Na2S2O3 + Na2S2O6
Na2SO4 + S --> Na2S2O4
I'm not saying let's go kill all the stupid people...I'm just saying lets remove all the warning labels and let the problem sort itself out.
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rstar
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| Quote: Originally posted by AndersHoveland | I do not know what salt you have formed, but it is probably a combination of thiosulfite and metabisulfite, or possibly even dithionate.
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Can you tell me some properties of the salts, that might be formed so as to test them and find the actual products.
The salt cannot be Sodium Dithionate because, it is stated in Wikipedia that :
"Sodium dithionate is a very stable compound which is not oxidized by permanganate, dichromate or bromine."
When I added this salt mix to a CuSO4 solution, a brown precipitate was formed.
"A tidy laboratory means a lazy chemist "
- Jöns Jacob Berzelius
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barley81
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Take some and acidify it, then it will be easier to tell what it is.
This should also be in the 'Beginnings' section.
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peach
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Whoops!
| Quote: | | I suggest if anyone has those ingredients can also perform this experiment. |
Righto
One empty container
I've added some sodium sulphate, cooked it with a blow torch for a while to make sure it's dry and then reweighed the container
I shook the two together
Melted the sulphur, which then began to boil, and then ignited.
Reweighing it.
Masses:
Container 18.22068g
Container + sulphate 22.54209g
Container + something after heating 22.54591g
Mass of sulphate in container, 22.54209 - 18.22068 = 4.32141g
Mass of whatever is left in there = 22.54591 - 18.22068 = 4.32523g
Mass change = +0.00382g
Number of moles of sulphate going in, 4.32141 / 142.04 = 0.03042
Thiosulphate
4.80970
Dithionite
5.29633
Metabisulphite
5.78305
Sulphite
3.83422
Dithionate
5.29633
It seems, in my test of it, nothing has happened. The powder left over is actually a lot lighter than it looks in the photo. It is slightly grey, but
that could just be traces of organic stuff.
[Edited on 2-10-2011 by peach]
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AndersHoveland
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Then this is indeed quite a mystery. The mass of the sulfate was virtually unchanged by heating it with sulfur, yet this new salt can reduce
permanganate. Could the similar masses be an unlikely coincidence?
The molar mass of Sulfur is 32.065 g/mol, that for oxygen is almost exactly half that value at 15.9994 g/mol. It could be possible that some of the
oxygen in the sulfate is leaving in the form of SO2. From your measurements, the reaction would seem to be:
Na2SO4 + (2)S --> Na2S2O2 + SO2
where the final product would have the net composition Na2S2O2, but is not necessarily a single substance. The substance may be sodium sulfide mixed
with an equimolar ratio with sodium sulfate.
(2)Na2SO4 + (4)S --> Na2SO4 + Na2S + (2)SO2
But if this is so, there is the question of why the resulting product has an exact 1 to 1 ratio of sulfide to unreacted sulfate. Could it be possible
that a double salt may have formed, with the constituent sulfate being less reactive with sulfur than it was originally?
More tests are obviously needed.
[Edited on 2-10-2011 by AndersHoveland]
I'm not saying let's go kill all the stupid people...I'm just saying lets remove all the warning labels and let the problem sort itself out.
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Polverone
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Are your sodium sulfate and sulfur powder perfectly pure and neutral? Small amounts of basic impurities (carbonates, hydroxides) could react with the
sulfur to form sulfide, polysulfide, and/or thiosulfate.
PGP Key and corresponding e-mail address
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peach
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[I am no longer receiving reply emails from the site. I've never knowingly blocked SM and am on googlemail. Has anyone else had the problem recently?
It's always worked fine before.]
The sulphur should be reasonably pure. It's a rod I had to powder up myself.
The sulphate is eBay grade drying agent in a bottle, but it's been repacked.
Anders is also correct that the balance alone is not a particularly solid piece of evidence. I'm sure I got the sulphur out, as it's easy to see when
it's still there (the sample changes colour when warmed - which mine no longer does). And I was doing the weighing as soon as the things had come back
to RT. But alone, it's not really enough.
So I had another look this morning.
-------------------------------------------------------------
Mine is a similar colour and form to rstars, like sand. It is not pink or red however. I'd say that's just some organic traces that have been burnt
with the intense heat, but I would expect them to be closer to brown, black or grey, where as this is beige. It could still be just some form of trace
contaminant however.
<a href="http://img513.imageshack.us/i/img0979rx.jpg/" target="_blank"><img src="http://img513.imageshack.us/img513/989/img0979rx.jpg"
alt="Free Image Hosting at www.ImageShack.us" border="0"/></a><br>
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I add some potassium permanganate to a test tube and attempt to get the colour to show up on the camera properly.
<a href="http://img651.imageshack.us/i/img0981wn.jpg/" target="_blank"><img src="http://img651.imageshack.us/img651/4756/img0981wn.jpg"
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alt="Free Image Hosting at www.ImageShack.us" border="0"/></a><br>
And then spoon some of element unknown in. It readily dissolves, but there is no colour change.
<a href="http://img411.imageshack.us/i/img0983it.jpg/" target="_blank"><img src="http://img411.imageshack.us/img411/6185/img0983it.jpg"
alt="Free Image Hosting at www.ImageShack.us" border="0"/></a><br>
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I empty it out, rinse the tube a few times, add some more of the sand and a squirt of hydrochloric. There is no real smell, but there is a vanishingly
faint smell of egg. It is so faint, I'm not sure if I can even smell it. I ask two other people to give it a smell (they have no idea what it is).
They say "sulphur" and "egg". The sample it's self lacks the smell.
<a href="http://img43.imageshack.us/i/img0985xv.jpg/" target="_blank"><img src="http://img43.imageshack.us/img43/9185/img0985xv.jpg"
alt="Free Image Hosting at www.ImageShack.us" border="0"/></a><br>
I go to the shop and exchange time credits for a pack of ?Duralons? (hmmmm). I was hoping they'd have some water bombs, but no luck.
<a href="http://img59.imageshack.us/i/img0990m.jpg/" target="_blank"><img src="http://img59.imageshack.us/img59/3693/img0990m.jpg" alt="Free
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I empty the rest of the powder into a tiny flask. I'm weighing it because, if any significant amount of gas comes off the sample, I can then measure
the volume and go back to the mass to compare the two.
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I fit a tap, greased up with vacuum grease, to the flask and zip tie a flat Duralon to the exit. I lift the stopper just clear, squirt 2.5ml of
concentrated hydrochloric down the gap and immediately close it, leaving to attend to my cheese and ketchup sammich for 15 minutes.
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On returning, the balloon does not seem to have inflated by any tangible amount. I close the tap, remove the top, then open the tap and have a sniff.
I have purposefully made hydrogen sulphide before and can smell a very faint trace of it there. But, given that the detection threshold for the stuff
is right down at an eye popping 0.47 parts per billion, I'd have expected a decent stench from even a fraction of a gram of sulphide.
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| Quote: | | When I added this salt mix to a CuSO4 solution, a brown precipitate was formed. |
Annoying suggestion, but have you tried adding it to water? I'm wondering if the brown you have noted for the permanganate and CuSO4 is actually a
response and not something already present.
-------------------------------------------------------------
| Quote: | | Are your sodium sulfate and sulfur powder perfectly pure and neutral? |
Perfectly pure, I can't really test at home.
Neutral though, will do!
One pH probe + beaker. I squirt some water in from the wash bottle.
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That's novel. 
<a href="http://img405.imageshack.us/i/img0997gm.jpg/" target="_blank"><img src="http://img405.imageshack.us/img405/6927/img0997gm.jpg"
alt="Free Image Hosting at www.ImageShack.us" border="0"/></a><br>
I've run out of distilled water and need to buy some, so I'm using filtered tap water for this. I assume the probe or beaker has traces of acid left
on it, so rinse both. No. Then squirt water directly at the probe's bulb. No. Then put it under the tap and get something closer to neutral. That wash
bottle is new (QuickFit gave me it for free in some promotional thing), so where's this acidic pH coming from? Ye shall see. I'll just use tap water
and check the incremental change instead.
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I add about a spoonful of the sodium sulphate and stir it in.
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The pH drops by about .1 units, so there can't be anything basic in there.
<a href="http://img214.imageshack.us/i/img1001h.jpg/" target="_blank"><img src="http://img214.imageshack.us/img214/668/img1001h.jpg"
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Mean whilst (things are about to take a detour), somewhere in the kitchen, I inspect the water filter. Acidic down below.
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Tap water pH up top. I think I'm onto the source right m'ehe. There's only one other logical thing left to try.
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alt="Free Image Hosting at www.ImageShack.us" border="0"/></a>
Tap water in the jug alone. And it's tap water pH. The change is being effected by the Brita filter cartridge.
| Quote: | How does a Brita water filter work?
BRITA cartridges contain a combination of ion exchange resin and activated carbon. The carbon absorbs chlorine, pesticides and organic pollutants,
improves taste, and eliminates odours and discoloration. It also contains an inhibitor that prevents bacterial growth. The ion exchange resin removes
the temporary hardness, which causes limescale; it also significantly reduces levels of metals such as copper and lead. |
- Brita.net
Ion exchange resin leaking out? Or perhaps it is selectively removing basic impurities better than acidic?
Place your bets, I shall contact Brita.
Finally, I wonder how it compares to that stuff that rots my teeth. Ahhh, not quite as bad, but there's science fair project material in there! "Does
filtering our tap water accelerate tooth decay?" This is big, big stuff kids! Don't mess around with that standard (Cola), impress your school with
cola versus tap water versus filtered.  I'm serious, and I'd like to see the
results. If anyone has kids, or you're in school yourself, give it a try and post up the results.
[Edited on 3-10-2011 by peach]
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rstar
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| Quote: Originally posted by peach | | I empty it out, rinse the tube a few times, add some more of the sand and a squirt of hydrochloric. There is no real smell, but there is a vanishingly
faint smell of egg. |
I added that stuff into Dilute HNO3 and got a similar egg-like smell
"A tidy laboratory means a lazy chemist "
- Jöns Jacob Berzelius
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peach
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The egg smell will be hydrogen sulphide, but hydrogen sulphide has such an immense stench that people around my house have been able to smell it out
in the street when I've had literally a few grams of aluminium sulphide sitting around; it's even managed to get round closed double glazed doors
(through the gaps in the frame).
If there was any appreciable amount of sulphide in there, it'd be an intense smell.
With it being that faint, even when collected for 15 minutes from a gram sample, there can only be minute amounts of sulphide in there.
Sulphur alone smells similar to hydrogen sulphide, but there are differences. H2S smells more like rotting. Sulphur dioxide is quite different to
both.
[Edited on 3-10-2011 by peach]
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ScienceSquirrel
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If there were traces of base in the sodium sulphate they would not be detectable by the pH meter but they would be sufficent to form traces of
sulphide.
I would put this down as no reaction as the very slight reaction might be down to traces of impurities.
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blogfast25
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Peach:
I'm betting your 'sand' is sodium sulphate and that in short nothing much happened during th heating of Na2SO4 with sulphur.
Try recrystallising the product with as little water as possible: 100 g @ 100 C for 42.5 g product, then cool and ice.
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peach
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I think yourself and Mr Squirrel are correct.
I would suggest the next test would be for rstar to try adding some to water and see if it goes brown and how easily it dissolves.
If it goes straight into solution and the water is brown, it's dirty sodium sulphate.
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rstar
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When i added a spatula of it to a few ml of water, some of it dissolved on simple shaking, and a little more dissolved in heating, but there was still
some grey stuff remaining, which did not dissolve, even after shaking hard and heating strong.
"A tidy laboratory means a lazy chemist "
- Jöns Jacob Berzelius
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unionised
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| Quote: Originally posted by Jor | | How can you get sodium sulfide (oxidation state -2) by heating sulfur (oxidation state 0) and sulfate (6+) ? |
quite easily, as long as you accept a very poor yield.
Na2 SO4 <---> Na2O + SO3
(the eqm will be massively over to the left but some SO3 will form)
S + 2 SO3 --> 3 SO2
That reaction happens quite well. If you like you can consider it as the thermolysis of SO3 to SO2 and O2 and then burning the S in the O2. Both those
reactions are known to take place at high temps.
The overall effect is to remover SO3 and swap it for SO2.
Of course, SO2 + Na2O --> Na2SO3, but that eqm is less favoured than the one for the SO3 which is a stronger acid
So, there's a plausible mechanism for at least a little Na2O to form.
The next bit expplains where the unexpected oxidation state of sulphur comes from; the reducing agent is sulphur.
That sounds odd, but here's an analogy.
Cl2 +2 NaOH ---> NaOCl + NaCl
One atom of chlorine oxidises another while being reduced.
With sulphur, in practice the equations and products are complicated, but to illustrate the idea
3 S + 3 Na2O --> 2 Na2S + Na2SO3
Adding acid to this gives H2S which you can smell, even if there are only traces. Both products will reduce permanganate (and so will sulphur itself
if it's finely divided)
Another issue is that any organic matter present will reduce sulphate to sulphide if you get it hot enough.
The reduction of NaSO4 by coal used to be a commercial process.
A bit of paper dust from a filter would be enough to get a smell of H2S.
I think the reason the weight doesn't change much is simply that nothing much happens.
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