Sciencemadness Discussion Board - What is this Chemical?

It is pale yellow, highly water soluble solid.

Lead (II) oxide is insoluble in water.

TheMrbunGee - 19-7-2016 at 08:14

Quote: Originally posted by NEMO-Chemistry

Maybe Lead (II) oxide?

does Pb even form blue salts?

clearly_not_atara - 19-7-2016 at 14:36

If you have enough of the substance and a heating element, you can try carbothermal reduction as a way to possibly extract any metals in it. This won't work for all metals -- and it can produce toxic carbon monoxide -- but it gives us an opportunity to narrow the possibilities down a lot. Heat the salt with powdered carbon at 1000 C or so or until it melts; see if any metallic residue is produced, and if so, you can begin checking for solubility.

You can also go the other way by passing ozone through an aqueous solution of this salt. This will take most metals to their highest oxidation state, which at least aids colorimetry.

NEMO-Chemistry - 19-7-2016 at 17:43

Quote: Originally posted by NEMO-Chemistry

Maybe Lead (II) oxide?

It is pale yellow, highly water soluble solid.

Lead (II) oxide is insoluble in water.

Then i guess its not!

TheMrbunGee - 20-7-2016 at 08:52

Quote: Originally posted by clearly_not_atara

If you have enough of the substance and a heating element, you can try carbothermal reduction as a way to possibly extract any metals in it. This won't work for all metals -- and it can produce toxic carbon monoxide -- but it gives us an opportunity to narrow the possibilities down a lot. Heat the salt with powdered carbon at 1000 C or so or until it melts; see if any metallic residue is produced, and if so, you can begin checking for solubility.

You can also go the other way by passing ozone through an aqueous solution of this salt. This will take most metals to their highest oxidation state, which at least aids colorimetry.

I think it is most likely chromium salt.

Solutions mixed with:
HCl makes clear Blue solution;
HNO3 makes in the beginning greenish solution, wich turns brown and unclear;
Ammonia - nothing changes.

DraconicAcid - 20-7-2016 at 09:02

A chromium salt with that pale yellow colour?

If it is a chromium salt, then heating it with hydroxide and peroxide will give chromate, which can be easily tested for.

Sulaiman - 20-7-2016 at 09:26

before using up / damaging your sample,
you could re-crystalise it from water,
maybe it will be more recognisable than the mass of crystals crashed out,
and you wil have a purer sample of compound-x

EDIT: there may be the original compound and/or decomposition products such as oxide/hydroxide/carbonate... that would confuse the isimple dentification of the original compound.

[Edited on 20-7-2016 by Sulaiman]

TheMrbunGee - 20-7-2016 at 21:57

Quote:

If it is a chromium salt, then heating it with hydroxide and peroxide will give chromate, which can be easily tested for.

It did not gave chromate, but H2O2 decomposed (@ moderate rate).

Quote: Originally posted by Sulaiman

before using up / damaging your sample,
you could re-crystalise it from water,
maybe it will be more recognisable than the mass of crystals crashed out,
and you wil have a purer sample of compound-x

EDIT: there may be the original compound and/or decomposition products such as oxide/hydroxide/carbonate... that would confuse the isimple dentification of the original compound.

[Edited on 20-7-2016 by Sulaiman]

I started recrystallizing the compound, managed to dissolve 100g of compound in 170 ml of boiling water (58-59g/100ml 100C), hot filtered out and slowly lowering the temperature to maybe find out the crystal shape.

Quote:

A chromium salt with that pale yellow colour?

saturated solutions are quite dark orange..(picture)

13833403_1130072447034084_494398725_o.jpg - 169kB

[Edited on 21-7-2016 by TheMrbunGee]

Sulaiman - 21-7-2016 at 00:08

IMO that is a suspension, not a solution which would be clear,
filter the suspension and crystalise the filtrate

How embarrassing that the combined brains of SM can't help identify an inorganic compound

TheMrbunGee - 21-7-2016 at 03:39

Quote: Originally posted by Sulaiman

IMO that is a suspension, not a solution which would be clear,
filter the suspension and crystalise the filtrate

How embarrassing that the combined brains of SM can't help identify an inorganic compound

Yes, there was some insoluble particles, that went through filter. I did filter it.

Solution is now @ room temperature and some crystals have formed. Pictures below!

Macros of crystals for crystal fanats!

[Edited on 21-7-2016 by TheMrbunGee]

Sulaiman - 21-7-2016 at 10:14

my first guess https://en.wikipedia.org/wiki/Sodium_thiosulfate
the yellow suspension is sulphur.

I've not got enough knowledge to know if this is compatible with the earlier experimental results.

Edited : no gas with H₂SO₄ is confusing,
plus the colours are more for chromium salts of which I have no experience
.. any more would be guesswork based on more reading..

I thought I recognised the crystal shape ... wrong ... not sodium thiosulfate.

[Edited on 21-7-2016 by Sulaiman]

[Edited on 21-7-2016 by Sulaiman]

Neme - 21-7-2016 at 10:31

These crystals look like potassium ferrocyanide, but you said there were not gases produced in reaction with sulfuric acid, right? Also the other tests are not making this guess likely.

gdflp - 21-7-2016 at 11:16

That is almost certainly potassium ferrocyanide, at the very least it is part of a mixture. The solubility behavior is exactly what you would expect, and in the reaction with nitric acid, the excess acid will oxidize the all of the iron to the +3 oxidation state creating that brown solution. Those crystals are quite distinctive as well. For final confirmation, I would try reacting it a ferric salt of some kind to see how it reacts, a dark blue solution should be the result. I have a feeling that your HCl might have some iron contamination, hence the blue color when the two are mixed. You don't have to worry about HCN too much when mixing the ferrocyanide with acids so long as the mixture isn't heated, but take care in any event.

[Edited on 7-21-2016 by gdflp]

TheMrbunGee - 21-7-2016 at 20:30

These crystals look like potassium ferrocyanide, but you said there were not gases produced in reaction with sulfuric acid, right? Also the other tests are not making this guess likely.

Might be because I added acid to weak solution of it!

BUT YOU ARE RIGHT! Reaction with FeCl2 solution confirmed, that it is Potassium ferrocyanide! THank you!

OK, I will now prepare next mythical substance!

Neme - 22-7-2016 at 09:01

Nice!

Please post another one soon, it's really funny for me

Sulaiman - 22-7-2016 at 12:06

yes, another one, i need more embarassment

The Volatile Chemist - 22-7-2016 at 12:42

This is nice, better than the thread where someone just picked a compound out of the blue. And it's like a whodunit. I guessed it was K₄Fe(CN)₆ before that was figured out. Also, the granule size in the original picture is a giveaway once some of the chemical properties were shown. Potassium ferrocyanide is almost always sold in that form.

What is this Chemical? [2]

TheMrbunGee - 22-7-2016 at 22:23

Welcome to the part 2 of the game - What is this Chemical?

This time we have dark powder with a greenish accent, it is quite heavy, and kind of inert, only thing I got it to react with is aluminum powder, why I am guessing it is a metal oxide. Reaction was fast and I could not collect any metal particles.

Doesn't react with:
HCl;
H2SO4;
HNO3;
NaOH;

Could not melt it with jet lighter.

Can't be oxidized with KNO3.

[Edited on 23-7-2016 by TheMrbunGee]

Amos - 22-7-2016 at 22:27

Which acids and alkalis was it reacted with?

Texium - 22-7-2016 at 22:33

I would try fusing it with molten potassium hydroxide and potassium nitrate and then dissolving the melt (once it cools) in a sodium hydroxide solution (20% or so). I think it may be low quality manganese dioxide. If it is this should yield a dark blue-green solution of manganate(VI). It could also be a low quality copper(II) oxide that has been calcined to the point of being unreactive. If it is then that same treatment should yield a dark blue cuprate solution.

kavu - 23-7-2016 at 00:49

Looks a lot like chromium(III)oxide and the reactivity profile would match as well. You could check if it catalyses the oxidation of ammonia (a classic demo experiment http://www.rsc.org/learn-chemistry/resource/res00000988/chro...)

TheMrbunGee - 23-7-2016 at 01:57

So the catalytic the oxidation of ammonia did not work. chromium (III) oxide is also kind of green, more green than this powder.

Fusing it with molten potassium hydroxide and potassium nitrate also did not change anything. :/

It is remarkably heavy, may be it is lead compound? It is really almost as heavy as lead powder.

EDIT: I was able to collect a bunch of metal particles after few thermite reactions. this will make the task much more easy, because the particles are magnetic!

Pictures of particles below!

So - nickel or cobalt oxide?

[Edited on 23-7-2016 by TheMrbunGee]

DraconicAcid - 23-7-2016 at 06:37

Collect the metal particles and dissolve them in hydrochloric acid. The colour of the solution (or the residue, if allowed to go to dryness) will tell us a lot.

TheMrbunGee - 23-7-2016 at 07:46

Quote: Originally posted by DraconicAcid

Collect the metal particles and dissolve them in hydrochloric acid. The colour of the solution (or the residue, if allowed to go to dryness) will tell us a lot.

Extracted the metal and added nitric acid to it, green solution tells that it is nickel, and the chemical I began with is nickel (III) oxide!

This was kind of easy one!

Neme - 23-7-2016 at 10:07

Shame I got home too late to have a chance, another one please!

TheMrbunGee - 24-7-2016 at 09:56

Shame I got home too late to have a chance, another one please!

Next one will be out tomorrow!

Herr Haber - 24-7-2016 at 11:46

Pokemon Go? PAH !

Now that's a treasure hunt I like

TheMrbunGee - 25-7-2016 at 02:44

A wild Chemical appears!

So this one is colorless flaky crystals. Inorganic. Label was mostly ripped off, but I saw Na, so it is sodium salt.

Solubility is around 80g/100 ml (95C) water.

It kind off acts as a flux, because i was melting it on piece of Al foil and it spread on the foil when melted, but I have Na2B4O7 and it looks different and acts different when melted.

Right now I am cooling the saturated solution, to get pure crystals, and it looks like nothing has precipitated at around 30C .

The crystals have a tiny bit of faint yellow color, but I am not sure if that is just an impurity or what.

I have some things to do now, but When I’ll be back I will do more tests.

The next day:

I had divided the solution I tried to crystalize and put one in the freezer (-15) and other left outside (about +18 at night)

The one in the freezer precipitated super tiny crystals that sat in suspension and the outside one precipitated larger flaky crystals, just like the solid ones in the first picture..

So I did a few tests :
Nothing visible happened with :
1)HCl(conc.)
2)H2SO4 (I dropped small amount of crystals in conc. acid and they turned a tiny bit more yellow, but that’s all. )
3)HNO3 (conc.)
4)KOH (solution)
5)NaOH (solution)
6)NH4OH (solution)
7)Na2Cr4O7 (solution)

I forgot to mention that it is crystal hydrate.

And finally i poured CuSO4 solution into mystery crystal solution and got reaction!

Video

[Edited on 26-7-2016 by TheMrbunGee]

Neme - 27-7-2016 at 07:13

Well I think we need a clue. Could you please try reaction with Ca2+?

Praxichys - 27-7-2016 at 07:37

That looks a lot like sodium thiosulfate pentahydrate.

The reaction with copper sulfate should have been more yellow if that was the case, but it fits all the other (non)reactions. Hmmm.

Neme - 27-7-2016 at 07:56

Quote: Originally posted by Praxichys

That looks a lot like sodium thiosulfate pentahydrate.

The reaction with copper sulfate should have been more yellow if that was the case, but it fits all the other (non)reactions. Hmmm.

I don't agree, there should be reaction with hydrochloric acid producing elemental sulfur. Your theory would agree with the sulfuric acid test tho (yellow from sulfur).

mayko - 27-7-2016 at 09:44

Have you tried adding a bit to some tincture of iodine? That would pretty quickly confirm/disconfirm thiosulfate.

TheMrbunGee - 27-7-2016 at 21:00

Quote: Originally posted by mayko

Have you tried adding a bit to some tincture of iodine? That would pretty quickly confirm/disconfirm thiosulfate.

Test was negative. it is not Na2S2O3

RocksInHead - 27-7-2016 at 21:17

This is kind of a long shot (considering it is crystaline) but it might be sodium carbonate maybe? Seeing how you added the CuSO4 and got a green insoluble salt it may have been Na2CO3, considering that Cu3(OH)2(CO3)2 is greenish blue.

TheMrbunGee - 27-7-2016 at 22:11

Quote: Originally posted by RocksInHead

This is kind of a long shot (considering it is crystaline) but it might be sodium carbonate maybe? Seeing how you added the CuSO4 and got a green insoluble salt it may have been Na2CO3, considering that Cu3(OH)2(CO3)2 is greenish blue.

Wouldn't sodium carbonate react with acids?

Neme - 27-7-2016 at 23:44

Quote: Originally posted by RocksInHead

Wouldn't sodium carbonate react with acids?

Yep, it would...
Could you please do test with Ca2+?

TheMrbunGee - 28-7-2016 at 01:49

Could you please do test with Ca2+?

I added mystery solution to CaCl2 solution and after a second - white insoluble precipitate appeared.

[Edited on 28-7-2016 by TheMrbunGee]

Neme - 28-7-2016 at 08:11

I have had suspicion that it could be phosphate group, this would totally agree. Are you able to perform any other test for phosphates?

PS: Maybe test for arsenate would be an option if the mysterious solution reveals itself as phosphateless.

RocksInHead - 28-7-2016 at 08:14

Quote: Originally posted by RocksInHead

Wouldn't sodium carbonate react with acids?

OH MY GOD I AM AN IDIOT.

Neme - 28-7-2016 at 08:51

I have had suspicion that it could be phosphate group, this would totally agree. Are you able to perform any other test for phosphates?

PS: Maybe test for arsenate would be an option if the mysterious solution reveals itself as phosphateless.

Ummm, I'm wondering if you really could have access to sodium arsenate. Maybe just ignore that as an option.

DraconicAcid - 28-7-2016 at 09:14

Ummm, I'm wondering if you really could have access to sodium arsenate. Maybe just ignore that as an option.

But doing the Marsh test for arsenic would be so much fun!

TheMrbunGee - 28-7-2016 at 12:22

Test with (NH4)6Mo7O24 showed that it is not phosphate nor arsenate. :/

DraconicAcid - 28-7-2016 at 12:40

Might be oxalate.

TheMrbunGee - 28-7-2016 at 12:51

Quote: Originally posted by DraconicAcid

Might be oxalate.

It does not decompose @ 260C. :?

Neme - 28-7-2016 at 23:54

Test with (NH4)6Mo7O24 showed that it is not phosphate nor arsenate. :/

Wow, I'm getting out of ideas, could you try some tests with cationts? e.g. Fe2+, Fe3+, or even Ag+ if you have some silver nitrate you could spare.

TheMrbunGee - 29-7-2016 at 04:18

Wow, I'm getting out of ideas, could you try some tests with cationts? e.g. Fe2+, Fe3+, or even Ag+ if you have some silver nitrate you could spare.

With FeCl3 /FeCl2 (Oxidized FeCl2).

I will now try to isolate silver nitrate from the solution I got from dissolving piece of silver in nitric acid, but it had copper in it, so it won’t work for this test. That is coming soon.

[Edited on 30-7-2016 by TheMrbunGee]

Neme - 31-7-2016 at 23:51

I'm sorry, even after lot of searching the net I don't really have a clue.

Metallus - 1-8-2016 at 02:05

"snip"

Is it soluble in Ethanol?

[Edited on 1-8-2016 by Metallus]

Neme - 1-8-2016 at 03:30

I think he said it's inorganic.

Metallus - 1-8-2016 at 04:57

I think he said it's inorganic.

You are right. I'll give it another shot then, even though there are very few alternatives, considering the most common anions.

You tested it with molybdate to see if it was phosphate but the test was negative. I've only worked with molybdates in a couple of occasions, but what if this was the molybdate? Highly soluble in water, not much big of reaction in those environments... What if you add your salt to an acidified solution of sugar or some other reducing agent?

TheMrbunGee - 1-8-2016 at 21:44

I think he said it's inorganic.

You might be on to something! I added sodium phosphate and nitric acid to the salt and got nice yellow color of Phosphomolybdate. (picture)

sugar, nitric acid, water and salt x did not do anything, did I had to add water?

wiki says that Na molybdate would react explosively with molten Mg, I mixed the salt with Mg powder and heated the mixture - got no reaction at all.

Also after heating the xsalt, sugar water and nitric acid mixture something yellowish precipitated and blue spots apeared on the walls of test tube.

[Edited on 2-8-2016 by TheMrbunGee]

[Edited on 2-8-2016 by TheMrbunGee]

Metallus - 1-8-2016 at 23:11

Quote:

Have you heated the molybdate/sugar solution? Alternatively you can also try to add a reducing agent to the yellow suspension you've obtained when you added phosphate (for phosphate tests, stannous chloride is often employed; it'll be like doing an inverse phosphate test).

[Edited on 2-8-2016 by Metallus]

Metallus - 2-8-2016 at 08:33

Reduction of molybdates should yield molybdenum blue, but I can't understand well the yellow part. Also, about the texture of your solution, how much sugar did you add? I fear that the yellow tint might be due to HNO3 oxidizing sugar. You could try to repeat this test with a non-oxidizing acid like HCl or keep the HNO3 medium and use a different reductant.

Moreover, as you said before, molybdates should react if you heat it with magnesium or zinc but you said that nothing happened. On the other hand, tungstates have similar reactions to molybdates but I only worked with WO3 once, so my experience with tungstates is close to nil. I've just read a few tests to perform on tungstates that require thiocyanate and SnCl2. Do you have access to them? I've checked all the other reactions that you listed, and they are all compatible with both molybdates and tungstates to a good extent (even the precipitation of copper).

Oh ye, is the salt... heavy?

Nice pictures btw.

TheMrbunGee - 2-8-2016 at 09:31

Reduction of molybdates should yield molybdenum blue, but I can't understand well the yellow part. Also, about the texture of your solution, how much sugar did you add? I fear that the yellow tint might be due to HNO3 oxidizing sugar. You could try to repeat this test with a non-oxidizing acid like HCl or keep the HNO3 medium and use a different reductant.

Moreover, as you said before, molybdates should react if you heat it with magnesium or zinc but you said that nothing happened. On the other hand, tungstates have similar reactions to molybdates but I only worked with WO3 once, so my experience with tungstates is close to nil. I've just read a few tests to perform on tungstates that require thiocyanate and SnCl2. Do you have access to them? I've checked all the other reactions that you listed, and they are all compatible with both molybdates and tungstates to a good extent (even the precipitation of copper).

Oh ye, is the salt... heavy?

Nice pictures btw.

Thanks!

The salt is not extraordinary heavy. :?

My HCl is yellow, because it is old! :/ I have Tin, so I could make SnCl2!

I am now making tin chloride and potassium thiocyanate! Do some of them have to be solids or really pure?

DraconicAcid - 2-8-2016 at 09:50

If it's not soluble in organic solvents, you could determine its density.

Metallus - 2-8-2016 at 10:06

Thanks!

It would have been better if the HCl were a clear solution. Is it as yellow as the product of HNO3/sugar/salt mix? If it's just faintly yellow due to iron impurities but clearer, I would attempt a quick HCl/sugar/salt mix (HCl 4-5M should suffice. Perhaps first completely dissolve sugar in diluted HCl, then dump a small quantity of the salt).

The thing is, SnCl2 requires HCl solution in order to not hydrolize and precipitate. If it is too yellow and molybden blue forms, I'd expect a green product.

Potassium thiocyanate + SnCl2 in HCl is used to extract tungsten compounds because it should form a yellow/greenish complex. Again, if we start already from a yellow coloration, it might be hard to accurately distinguish the two.

PS: is the pH of the salt solution basic?

TheMrbunGee - 2-8-2016 at 22:01

At the moment I don’t have anything to test pH. :/

My HCl is even more yellow, then that mixture before. Interesting thing - I added acetic acid to the xsalt and it turned a bit yellow too! :?

Ok, I made some Potassium thiocyanate yesterday and I think some tin has dissolved in HCl, Should I just add them all together?

I added SnCl to the xsolution. Black precipitate formed. And there is some dark yellow/orange solution forming (2nd and 3rd frame) . Potassium thiocyanate does not change anything. (Picture)

I also dissolved xsalt in my dirty HCl and then added HCl/SnCl2 mixture. Looked like much more orange solution formed. But there was parts of test tube sides where deep blue color had formed!

vlcsnap-2016-08-03-08h51m17s590.png - 1MB

BTW, when I mix SnCl2 with yellow HCl it becomes colorless!

Last test I just did:

I took HCl and added Potassium thiocyanate to it. (Yellow goes to pink/red) then I added SnCl2 solution (everything becomes clear) and then I added xsolution (pictures) and I have a orange/red/brown solution.

Edit:
I Added much more xsolution to last test result and I got the deep blue molybden blue.

So it is confirmed for molybden. but which one is it.. :?

[Edited on 3-8-2016 by TheMrbunGee]

Metallus - 3-8-2016 at 00:24

So, to sum up:

- Sodium molybdate:
1) 84g/100ml solubility at 100°C (yours was approx 80g/100ml at 95°C)
2) Melts at 687°C (yours melts before that, if you were able to melt it on aluminium foil which supposedly melts at T < 660°C, which leads me to think it is not pure)
3) reacts with phosphate to form canary yellow complex (yours formed that yellow complex with phosphate)
4) gets reduced to molybdenum blue when reacted with an acid solution of a reductant (yours formed the blue molybden)
5) in the presence of thiocyanate and a reductant, it forms red blood solution (Your solution of HCl and thiocyanate is pink/red at start most likely because of Fe(III) impurities that form FeSCN^2- (red); this disappears when you reduce it with SnCl₂. Adding the molybdate then yields the red/brown color)
6) it is reported in here http://molybdenum.atomistry.com/detection_of_molybdenum.html that reaction with acetic acid yields a yellow coloration (you said that yours turned yellow) and it also confirms many other reactions that you've carried out (I dont' know how much truthful this page is, but I found most of those reactions in literature as well).

TheMrbunGee - 3-8-2016 at 00:40

So, to sum up:

- Sodium molybdate:
1) 84g/100ml solubility at 100°C (yours was approx 80g/100ml at 95°C)
2) Melts at 687°C (yours melts before that, if you were able to melt it on aluminium foil which supposedly melts at T < 660°C, which leads me to think it is not pure)
3) reacts with phosphate to form canary yellow complex (yours formed that yellow complex with phosphate)
4) gets reduced to molybdenum blue when reacted with an acid solution of a reductant (yours formed the blue molybden)
5) in the presence of thiocyanate and a reductant, it forms red blood solution (Your solution of HCl and thiocyanate is pink/red at start most likely because of Fe(III) impurities that form FeSCN^2- (red); this disappears when you reduce it with SnCl₂. Adding the molybdate then yields the red/brown color)
6) it is reported in here http://molybdenum.atomistry.com/detection_of_molybdenum.html that reaction with acetic acid yields a yellow coloration (you said that yours turned yellow) and it also confirms many other reactions that you've carried out (I dont' know how much truthful this page is, but I found most of those reactions in literature as well).

Sodium molybdate it is. I was thinking of it, but no reaction with Mg was strange! :? I will try it some other way!

Ok cool, we got another one!

Next one coming soon!

So Next I have pack of three!

All those 3 chemicals came in one box, so I guess the have some use all 3 together.

I cut out what was left of description on the box and I can read most of it.

It translates to :

Sodium (something)sulfate or (something)sulfite cryst.
sodium metabisulfite
ammonium chloride?

May be someone can confirm those and knows why they were used for and sold together?

[Edited on 3-8-2016 by TheMrbunGee]

Metallus - 3-8-2016 at 05:15

Probably they were used as food additives, as all of them (sodium sulfate, bisulfate, metabisulfite, ammonium chloride) find an application as food preservatives.

Does the box hint at anything else? I assume the compound names were the first 3 lines, since I could spot a Natriu (sodium) and Ammoni (ammonium).

Can you translate what the description says in the last 2 lines?

Also, a quick way to identify whether the first is sodium sulphate or bisulphate, is to simply add it to a solution of a carbonate. The bisulphate is acid and will react with the carbonate to produce CO₂. If you see bubbles, you have your answer.

[Edited on 3-8-2016 by Metallus]

TheMrbunGee - 3-8-2016 at 06:55

Probably they were used as food additives, as all of them (sodium sulfate, bisulfate, metabisulfite, ammonium chloride) find an application as food preservatives.

Does the box hint at anything else? I assume the compound names were the first 3 lines, since I could spot a Natriu (sodium) and Ammoni (ammonium).

Can you translate what the description says in the last 2 lines?

Also, a quick way to identify whether the first is sodium sulphate or bisulphate, is to simply add it to a solution of a carbonate. The bisulphate is acid and will react with the carbonate to produce CO₂. If you see bubbles, you have your answer.

[Edited on 3-8-2016 by Metallus]

Yes, those was the compound names, (see the translation, bold in last post. first one is more than sulfate I can read ...osulfate I am not that educated in Russian chemistry, so I am not sure what it starts with.

That piece was all that gave any info. Nothing more.. :/

Last two lines are company name "biochemreaktive" and like factory number or something..

One more problem is that I don’t know which bag is which compound..

I will add carbonate to all of them and we will see..

Also, wouldn't it be strange to sell boxes of not that common food additives like that?

So I added sodium carbonate to all the solutions (I will number them from 1st - the biggest bag and 3rd the smallest), 1st and 2nd did nothing and 3rd did the bubbling.

Did some Russian research and got that sodium bisulfite is "Гидросульфит натрия" and there is that "o" still visible.

Ok, so we have:
Sodium bisulfite
metabisulfite (smallest bag)
Ammonium chloride.

I will figure out which is which!

This one was also easy.

too easy!

Next one coming soon!

[Edited on 3-8-2016 by TheMrbunGee]

Texium - 3-8-2016 at 07:13

Ok, so if it's "osulfate" then you probably have sodium thiosulfate. Here's what I would do:

-Take a bit of each of the substances, and mix them as powders with solid sodium hydroxide.
-Add a few drops of water to each of the piles. The one that is ammonium chloride should react and produce ammonia gas, while the others do nothing.
-Dissolve a bit of each of the remaining two in water in separate test tubes. Add some dilute HCl. Both should react to produce sulfur dioxide, but there will also be a fine yellow precipitate of elemental sulfur in the one that is thiosulfate.

Metallus - 3-8-2016 at 07:41

You can rule out bisulphate by reacting it with a solution of sodium carbonate; the evolution of CO₂ as bubbles will confirm the acidic properties of the bisulfate.

You can rule out the metabisulfite by reacting your powders with oxidants like
- Permanganate: starting from an acidified solution of your X salt, add small drops of a diluted solution of KMNO₄. If the purple drop gets discoloured, it means the permanganate got reduced, indicating the presence of the sulfite.
- Iodine/iodide: partially oxidize potassium iodide to iodine with few drops of hydrogen peroxide (you will get a brown solution which is I₃^-) then add a few drops of the iodine/iodide mix to a solution of your could-be-sulfite and see if it gets discoloured from brown to pale yellow, indicating the reduction of I₂ to I^-.
- Potassium cromate/dicromate: starting from an acidified solution (H₂SO₄) of cromate/dicromate, add a few drops of your could-be-sulfite solution; the reduction of orange/red Cr (VI) to green Cr (III) will indicate the presence of sulfite.
These reactions will not interfere with the other 2 compounds.

You can rule out the ammonium chloride by reacting your powders with a solution of NaOH. The cringe smell of ammonia should be enough. Alternatively, if you have access to phenol, you could try the colorimetric method (phenol + sodium hypocloride) which will yield a blue solution.
Otherwise, if you can confirm the first 2, then by exclusion the 3rd will be the chloride (which you can further confirm by precipitating the Cl^- with AgNO₃).

If the first is not bisulphate but neutral sulphate, you can precipitate it with barium chloride (even though this might interfere with the metabilsulfite, if you haven't ruled it out yet)

These methods should provide instant results.

EDIT: I'm at work so I was too slow to reply.

If it's sodium bisulfite, it will get oxidized the same way the metabisulfite does. Since you said that the 3rd bag (which is metabisulfite) reacted with the carbonate, I suppose I overlooked the acidic properties of the metabisulfite (even though the oxidation tests would have ruled it out from bisulphate anyways). See if one other of the first 2 bags gets oxidized and see if the remaining one precipitates AgCl.

[Edited on 3-8-2016 by Metallus]

[Edited on 3-8-2016 by Metallus]

TheMrbunGee - 3-8-2016 at 08:20

Quote: Originally posted by zts16

Ok, so if it's "osulfate" then you probably have sodium thiosulfate. Here's what I would do:

You can rule out bisulphate by reacting it with a solution of sodium carbonate; the evolution of CO₂ as bubbles will confirm the acidic properties of the bisulfate.

I did the tests and:

Big bag - sodium thiosulfate - sulfur precipitated with HCl

Middle bag - ammonium chloride - ammonia released with NaOH

Small bag - Sodium metabisulfite (acidic nature)

Thank You guys!

Still - interesting combo to sell in one box.. :?

Next chemical looks like to be CaCO3, could not dissolve anything in 400ml 100C water. Also reacts with acids and releases CO2. So I think I got it and now start to prepare next one!

Texium - 3-8-2016 at 08:48

Next chemical looks like to be CaCO3, could not dissolve anything in 400ml 100C water. Also reacts with acids and releases CO2. So I think I got it and now start to prepare next one!

You're probably right, but it could be another alkali earth carbonate. To be sure, add some sulfuric acid or a soluble sulfate salt to the solution dissolved in HCl. If you don't get a precipitate, it's probably actually magnesium carbonate. If you do, you're probably correct about calcium carbonate, but to verify it though, you could do a flame test with some of the HCl solution on a wood splint. If it's orangish, it's probably calcium, but if it's bright red then it's really strontium carbonate, and green would indicate barium carbonate.

[Edited on 8-3-2016 by zts16]

TheMrbunGee - 3-8-2016 at 09:31

Quote: Originally posted by zts16

Next chemical looks like to be CaCO3, could not dissolve anything in 400ml 100C water. Also reacts with acids and releases CO2. So I think I got it and now start to prepare next one!

I heated it a lot (to more than 1000 C) so it is not MgCO2. But you are right! I dissolved it in HCl and flame test gave a red color to the flame.. Could it also be Li2CO3 :?

And why wiki says that SrCO3 is hydroscopic? is it? If it is then mine is Li2CO3. Because my powder has been standing outdoors for years and it has not lumped up

[Edited on 3-8-2016 by TheMrbunGee]

Texium - 3-8-2016 at 09:53

Well, is the sulfate soluble in water, or not? Lithium sulfate is soluble, strontium sulfate is not.

Compounds can be hygroscopic without being deliquescent, as anhydrous hygroscopic solids can absorb water from the air, without necessarily liquifying. Think of magnesium sulfate or sodium carbonate.

TheMrbunGee - 3-8-2016 at 10:14

Quote: Originally posted by zts16

Well, is the sulfate soluble in water, or not? Lithium sulfate is soluble, strontium sulfate is not.

Compounds can be hygroscopic without being deliquescent, as anhydrous hygroscopic solids can absorb water from the air, without necessarily liquifying. Think of magnesium sulfate or sodium carbonate.

Then it is SrCO3.

Yes, but MgSO4 clumps up, just like NaCO3.. This one is nice freeflowing powder, not even small cumps.. strange..

ok, soon I will prepare next one!

So the next one is kind of ready!

These are colorless crystals.
There is a huge endothermic solubility in water. Insoluble in acetone. Melting point is 100C-200C (low) and liquid boils leaving smoke (no significant odor) until nothing is left.
Solutions and crystals was reacted with:
HCl nothing
H2SO4 dissolves good and fast, gets really hot, but nothing changes. (Picture below)
NaOH nothing
CuSO4 nothing
SnCl nothing
Al nothing (mixed and heated strongly)
KNO3 nothing (mixed and heated strongly)
Sugar nothing (mixed and heated strongly)
Iodine nothing (mixed and heated strongly)
Phosphate test negative
HNO3 small amount of xsalt and ~1ml of conc. HNO3 was mixed. White paste formed (something less soluble formed) when I added more water - everything dissolved and I was left with clear solution.

Nothing much, but something!

Xsalt itself!

H2SO4 and xsalt after cooling. (Plastic container melted from the heat)

[Edited on 4-8-2016 by TheMrbunGee]

Metallus - 4-8-2016 at 03:19

I would have said NH₄NO₃ but some non-reactions leave me wondering.

I'll go have lunch now, I'll edit later.

- - - - -

Ok I'm back.
NH₄NO₃ melts at 150-200°C and decomposes into water and nitrogen. Its endothermic reaction with water is used in the instant-ice bags (which usually contain NH₄NO₃ and a few ml of water contained in a thin bag that breaks when you bend it).

You won't get any appreciable reaction with acids, but when you mix it with NaOH I would expect to smell some ammonia. Can you smell anything when you mix the two?

The low solubility in conc HNO₃ is common to many other nitrates.

When you mix it with H₂SO₄ I would expect the formation of HNO₃ and perhaps some NO_x. Can you appreciate any faint yellow tint? I remember that when I mixed H₂SO₄ with KNO₃, the solution would release some heat, but not enough to melt a plastic cup.

If it is indeed a nitrate, you could try to add it to an acid solution of iodide or Fe(II) or sugar and see if something happens. Preferably use HCl/CH₃COOH to acidify, as other oxidizing acids will interfere.

However that appearance... AN is white, yours are transparent crystals with a yellow tint.

You could perhaps try adding a little amount of AgNO₃; if nothing precipitates, then it might indeed be a nitrate. If something precipitates, then it could be the chloride (but you would be able to smell the HCl and ammonia released on heat, so I doubt it).

[Edited on 4-8-2016 by Metallus]

Sulaiman - 4-8-2016 at 04:12

does the mystery compound have an odour ?

TheMrbunGee - 4-8-2016 at 05:00

NH₄NO₃ melts at 150-200°C and decomposes into water and nitrogen. Its endothermic reaction with water is used in the instant-ice bags (which usually contain NH₄NO₃ and a few ml of water contained in a thin bag that breaks when you bend it).

You won't get any appreciable reaction with acids, but when you mix it with NaOH I would expect to smell some ammonia. Can you smell anything when you mix the two?

The low solubility in conc HNO₃ is common to many other nitrates.

When you mix it with H₂SO₄ I would expect the formation of HNO₃ and perhaps some NO_x. Can you appreciate any faint yellow tint? I remember that when I mixed H₂SO₄ with KNO₃, the solution would release some heat, but not enough to melt a plastic cup.

If it is indeed a nitrate, you could try to add it to an acid solution of iodide or Fe(II) or sugar and see if something happens. Preferably use HCl/CH₃COOH to acidify, as other oxidizing acids will interfere.

However that appearance... AN is white, yours are transparent crystals with a yellow tint. 100% sure all this stuff is inorganic?

[Edited on 4-8-2016 by Metallus]

It may be organic..

With NaOH cant smell ammonia and with H2SO4 it doesnt smell like HNO3. I don't think it is nitrate, because it is not oxidant. no reaction with sugar nor with Al.

Metallus - 4-8-2016 at 05:09

It may be organic..

With NaOH cant smell ammonia and with H2SO4 it doesnt smell like HNO3. I don't think it is nitrate, because it is not oxidant. no reaction with sugar nor with Al.

If it is organic, dump it inside 2M H₂SO₄ + dichromate mix. That shit will oxidize almost anything. If you have permanganate, all the better.

A more economic version can also be H₂O₂ in NaOH.

If it readily oxidizes or burns, we can rule out some common inorganic alternatives.

[Edited on 4-8-2016 by Metallus]

[Edited on 4-8-2016 by Metallus]

TheMrbunGee - 4-8-2016 at 05:50

when melted reacts with magnesium powder and emits ammonia gas!

I mixes xsalt with water and KMnO4 and slowly added conc. H2SO4 after a bunch of drops some gas started to evolve. I put a burning match to the bubbles and it went off. (N2, CO2)

Metallus - 4-8-2016 at 06:10

It really sounds like ammonium nitrate to me. In acid KMnO₄ it decomposes to release water and ammonia. Does the reaction yield a brown/black precipitate (MnO_x) ?

Please, try to perform one of the reductions I listed above.

Quote:

If it is indeed a nitrate, you could try to add it to an acid solution of iodide or Fe(II) or sugar and see if something happens. Preferably use HCl/CH3COOH to acidify, as other oxidizing acids will interfere.

I remind you that an acid medium is required to help the nitrate oxidize. Just dumping Al or sugar to a neutral solution of nitrate won't do a thing. Try dissolving your salt in HCl and then perform again those reactions. Heat if necessary.

Alternatively, add your salt to a 1:1 diluted H₂SO₄ and then drop some copper in it. HNO₃ readily oxidizes copper to copper nitrate. If you see red fumes and a blue solution, then that's it.

Again, another quick test is adding a little amount of AgNO₃ (expensive). Most anions will precipitate but NO₃^-.

Sorry if I'm fixated with AN, but I just can't think of another inorganic salt with those properties (the low melting point and evolution of odorless gases coupled with the highly endothermic reaction with water really scream AN to me).

[Edited on 4-8-2016 by Metallus]

TheMrbunGee - 4-8-2016 at 23:52

It really sounds like ammonium nitrate to me. In acid KMnO₄ it decomposes to release water and ammonia. Does the reaction yield a brown/black precipitate (MnO_x) ?

Please, try to perform one of the reductions I listed above.

Quote:

It is not ammonium nitrate. 100% not. I have worked with AN, and this does not do things AN does.

To be extra sure I added H2SO4 to the salt and added copper. Nothing happens, mixture doesn’t smell like nitric acid or anything and does not react with Cu. not at all! Not when concentrated and not diluted.

I did some more tests and it is urea.

Metallus - 5-8-2016 at 00:21

Wow, I was oblivious of urea's endothermic reaction with water.

I have only studied it from theory, never actually got to use it in lab, didn't know it had these properties as well. I am now reading that certain cold packs also contain urea instead of AN.

One more thing learnt.

Keep them coming, I'm having fun.

TheMrbunGee - 5-8-2016 at 00:51

Wow, I was oblivious of urea's endothermic reaction with water.

I have only studied it from theory, never actually got to use it in lab, didn't know it had these properties as well. I am now reading that certain cold packs also contain urea instead of AN.

One more thing learnt.

Keep them coming, I'm having fun.

Good to know!

OK, a trick questeon - What the f... can I do with about 8 kg of MgSO4 * 7H2O??!?

I had these 6 jars as unidentified compounds, they all looked similiar and I did quick tests with NaOH and they all are magnesium sulfate. that spoiled my excitment. :/

I think I will dispose them and leave me with about 2 kg max.

Metallus - 5-8-2016 at 00:58

That's Epsom salt!

If you are "clogged" down there, eat a tablespoon of it and you'll shit demons.

Btw, have you performed other tests to be sure it's MgSO4 and not some other compound of a polyvalent cation? Redox reactions are usually a nice way to rule out common salts (carbonate, sulphate) of alkaline metals/earths.

[Edited on 5-8-2016 by Metallus]

Neme - 5-8-2016 at 04:30

You can grow huge polycrystal

Nah, you can use it as fertilizer I think and I heard someone is using it to prevent animals from nibbling trees. I don't think it has many uses since it's already hydrated.

Sulaiman - 5-8-2016 at 06:02

the jars are more valuable than their contents, but a little as fertiliser, some as bath-salts?
better to use than dispose.
re-crystalise some for stock, air-drying gravity filtered is tedious, definitely use vacuum if available.
Anhydrous MgSO4 seems to store quite well in a jam jar, maybe roast a little?

TheMrbunGee - 5-8-2016 at 06:57

That's Epsom salt!

If you are "clogged" down there, eat a tablespoon of it and you'll shit demons.

Btw, have you performed other tests to be sure it's MgSO4 and not some other compound of a polyvalent cation? Redox reactions are usually a nice way to rule out common salts (carbonate, sulphate) of alkaline metals/earths.

[Edited on 5-8-2016 by Metallus]

Yes, it is MgSO4!

Weak solution of salt.

Precipitate from reaction with NaOH solution.

This formed when concentrated CrXsolution was added to concentrated solution of NaOH

Edit:
Cr2(SO4)3*15(H2O) is green!

Solutions react with Mg powder.

[Edited on 5-8-2016 by TheMrbunGee]

DraconicAcid - 5-8-2016 at 07:37