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Author: Subject: The "WTF did I just make?" thread
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[*] posted on 25-2-2015 at 12:17


Quote: Originally posted by quantumcorespacealchemyst  
you don't think that an oxidizer KNO3 is too unreactive with an O2 source K2S2O8? the reactants are mixed and cold. the calcogenide is 0.998g Tm2Te3.

I have no idea what you mean by "an oxidizer KNO3 is too unreactive with an O2 source K2S2O8?"
It's spelled chalcogenide, and you still haven't told me what you're tying to do here. Are you trying to oxidize Thulium (VI) Telluride? Or Mercuy (I) sulfate?




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[*] posted on 25-2-2015 at 13:46


I am looking to get both Thulium and Mercury ions in solution and form salts, possibly double salts, with Te in their makeup and recover other (nitrate) salts of Tm and/or Hg.

I don't know if the premix is too reactive to heat up. I guess since they are both oxidizers, KNO3/K2S2O8, it may be fine as long as there are no compounds that can be oxidized too rapidly.

[Edited on 25-2-2015 by quantumcorespacealchemyst]
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[*] posted on 25-2-2015 at 14:42


@quantum
The irony of you posting in a thread titled "WTF did I just make?" has not escaped me.
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[*] posted on 25-2-2015 at 14:45


Quote: Originally posted by j_sum1  
@quantum
The irony of you posting in a thread titled "WTF did I just make?" has not escaped me.


I think that irony has escaped him, though.




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[*] posted on 25-2-2015 at 20:38


Quote: Originally posted by prof_genius  
I was making iodoform yesterday and I must have gotten the proportions wrong because the solution turned blue and released white fumes that I believe where HCl gas. Has this happened to anyone?


I am very interested in this
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[*] posted on 15-10-2015 at 23:16


it tuns out, the H2SO4 i was using was 2N, nowhere near the 98% o 95% i thought i had (wrote one o those somewhere). that explains my constant puzzlement about it's peculia properties and failed reactions.

Irony (from Ancient Greek εἰρωνεία (eirōneía), meaning "dissimulation, feigned ignorance"[1]), in its broadest sense, is a rhetorical device, literary technique, or event in which what appears, on the surface, to be the case, differs radically from what is actually the case. Irony may be divided into categories such as verbal, dramatic, and situational.

https://en.wikipedia.org/wiki/Irony

1. Liddell & Scott, A Greek-English Lexicon, v. sub εἰρωνεία.
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[*] posted on 31-10-2015 at 10:29


Quote: Originally posted by prof_genius  
I was making iodoform yesterday and I must have gotten the proportions wrong because the solution turned blue and released white fumes that I believe where HCl gas. Has this happened to anyone?


i still wonder about this. what is that?
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[*] posted on 5-11-2015 at 13:48
Unknown compound created by adding H2O2 to a solution of tetraamminecopper(ii) sulfate


From my previous post:

Quote:


No stoichiometric amounts were used:

1) I made a probably dilute solution of tetraaminecopper (ii) sulfate that looked similar to this, but far less(the greenish foam contains some of the product):
http://i.imgur.com/nNQoSmi.jpg

2) In the midst of adding H2O2 to the TACS solution:
https://youtu.be/aJE0ZPjMEas

3) Filtering some of the foam(you will see that this foam is not always created):
http://i.imgur.com/W3wfyo7.jpg

What was left in the container:
http://i.imgur.com/moHXrLw.jpg

4) After adding H2O2 to the mixture until it had stopped fizzing a lot, I was left with a dark-colored, minorly fizzing (as a cup of soda) solution:
https://youtu.be/SgRv2z9q0oU





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[*] posted on 5-11-2015 at 14:50


Despite the photos and videos, it's all a bit pointless.

You made dung by the Fizzy process.

Either Measure stuff or quit pretending it's Chemistry.

Any idiot can mix things together.

Sometimes the mixture will go Boom !

A Chemist (as i understand the idea) measures everything they can in order to glean some understanding from what they produced, including the Boom !

Mixing random quantities of things in a bucket isn't Science.




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[*] posted on 5-11-2015 at 14:59


Quote: Originally posted by aga  
Despite the photos and videos, it's all a bit pointless.

You made dung by the Fizzy process.

Either Measure stuff or quit pretending it's Chemistry.

Any idiot can mix things together.

Sometimes the mixture will go Boom !

A Chemist (as i understand the idea) measures everything they can in order to glean some understanding from what they produced, including the Boom !

Mixing random quantities of things in a bucket isn't Science.
Well aga, while you do have a point, experimentation does not necessarily have to be quantitative all the time. If I measured everything ultra precisely even when I was just performing a qualitative test, I know I would never get anything done. What Velzee has done consists of perfectly valid qualitative observations, though it is true that there is not much to learn from them. What he should do next is repeat it in a more controlled way now that he knows that something indeed will happen, taking care in preparing the complex, and then adding the peroxide dropwise until the effect is observed.

[Edited on 11-5-2015 by zts16]




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[*] posted on 5-11-2015 at 15:02


Exactly.

Edit:

If you ever see something happen when you 'just mix stuff up' then the true Chemist does it again and again until they have the Precise mix/procedure that gives that effect.

Without this approach, there would be no 'Body of Knowledge' to feed on.

No measuring, No Method = no Giants on whose shoulders we can stand.

Random mixing stuff up is Good.

Doing the Science afterwards is the Boring bit that wins you the Nobel.

[Edited on 5-11-2015 by aga]




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[*] posted on 5-11-2015 at 15:08


Quote: Originally posted by aga  
Exactly.
Yeah, so... maybe try to not sound like you're condemning the poor guy?



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[*] posted on 5-11-2015 at 16:35


I discovered the mystery substance when I misplaced a container of water with a container of H2O2 with other substances in it. Fizzing began, the substance became the color resembling dung(as aga described), and the container became very warm.

I then repeated the sequence of events many times, narrowing it down to a reaction of just two substances: 1)H2O2 and 2)TACS solution. NH4OH did not react with the H2O2; I don't believe that CuSO4 reacted with H2O2 either. I came across the thread that I had quoted my post from some time after my experiments, so I decided to repeat them.

The compound, when dry, seems to decompose, very slowly(some, I mean, very little of it became a lightish blue over the course of about a week or so), to copper hydroxide in air when dry. When heated, the compound seems to decompose to copper oxide.

When dry, it has a deep, almost black, green color to it. And, based on my observations, it crystallizes in a needle-like fashion. I haven't tested it further, and I have not found any other information on it since.

Compared to most on this forum, I am a very inexperienced chemist, therefore, please pardon my errors.

[Edited on 11/6/2015 by Velzee]




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[*] posted on 5-11-2015 at 17:25


I just replicated this. A spoonful of copper sulfate crystals were added to detergent-free clear household 3% ammonia solution, yielding the well-known dark purple solution of tetra-amine copper sulfate. Upon addition of 3% hydrogen peroxide, much gas is evolved, and the solution shifts to dark green.

However, as fizzing subsides, my solution returned to a coloration that was more purple than green. This suggests an unstable compound, easily decomposed by the heat of the gas-producing reaction.

Could it be a copper salt of peroxymonosulfuric acid? Such a compound, if it exists, would likely be unstable, decomposing to oxygen and light-blue copper sulfate. It would decompose on strong heating to copper oxide, but might first turn light blue (hydrous copper sulfate) or white (anhydrous copper sulfate).

It would be useful to know the composition of the evolved gasses. If it's oxygen, then it's simply the peroxide decomposing; the presence of nitrogen or nitrogenous gasses would suggest the ammonia is also decomposing.




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[*] posted on 5-11-2015 at 17:32


Addition of even more hydrogen peroxide yielded a solution that was dark green, as well as a lot of what appears to be suspended brown particles, likely copper oxide.

I can also confirm this green compound doesn't form on the addition of hydrogen peroxide to copper sulfate solution; ammonia is required.

{Edit}

Attempts to filter/settle the suspended particles produced a brown solid and a blue/purple solution. This leads me to suspect that either:
- the green mystery compound is highly unstable, and is decomposing to copper sulfate / TACS even as I attempt to filter it
or
- the green compound is, in fact, TACS or copper sulfate contaminated with suspended particles of copper oxide, making it appear green

{Edit}

The two photos below show the same aliquot of TACS & peroxide mixture in a small polystyrene cup with a flashlight under it. One shows the mixture after being left undisturbed for some time, allowing the suspended particles to settle and revealing a blue solution. The other shows the same solution, agitated lightly, making it appear dark green.

tacs peroxide light 1.jpg - 84kB tacs peroxide light 2.jpg - 85kB

This is enough to have me convinced that the green "compound" is indeed a blue compound contaminated with extremely fine particles of copper oxide. The exact reaction, and other products produced, remain unknown. I can say that the gasses evolved are practically odorless, and that the ammonia might be consumed in the reaction (the supernatant looked more like copper sulfate solution than tetraamminecopper(II) sulfate, and the copper sulfate was probably in excess; exact measurements are needed).

[Edited on 6-11-2015 by MolecularWorld]




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[*] posted on 5-11-2015 at 19:31


Quote: Originally posted by MolecularWorld  
Addition of even more hydrogen peroxide yielded a solution that was dark green, as well as a lot of what appears to be suspended brown particles, likely copper oxide.

I can also confirm this green compound doesn't form on the addition of hydrogen peroxide to copper sulfate solution; ammonia is required.

{Edit}

Attempts to filter/settle the suspended particles produced a brown solid and a blue/purple solution. This leads me to suspect that either:
- the green mystery compound is highly unstable, and is decomposing to copper sulfate / TACS even as I attempt to filter it
or
- the green compound is, in fact, TACS or copper sulfate contaminated with suspended particles of copper oxide, making it appear green

{Edit}

The two photos below show the same aliquot of TACS & peroxide mixture in a small polystyrene cup with a flashlight under it. One shows the mixture after being left undisturbed for some time, allowing the suspended particles to settle and revealing a blue solution. The other shows the same solution, agitated lightly, making it appear dark green.



This is enough to have me convinced that the green "compound" is indeed a blue compound contaminated with extremely fine particles of copper oxide. The exact reaction, and other products produced, remain unknown. I can say that the gasses evolved are practically odorless, and that the ammonia might be consumed in the reaction (the supernatant looked more like copper sulfate solution than tetraamminecopper(II) sulfate, and the copper sulfate was probably in excess; exact measurements are needed).

[Edited on 6-11-2015 by MolecularWorld]


Thank you for your attempt at this! I have a sample of this compound(if it is not copper oxide as you hypothesized) which remained its dark green color since I isolated it(weeks ago). I will try to post photos soon.

I would agree with your theory that this compound is simply a Cu oxide, had it not been for this:

Filtering the solution, after the reaction was complete, left an almost colorless solution(with the exception that some of the compound passed through the filter paper). If your theory's correct, where did the CuSO4 go?

If/When I obtain some more ammonia, I'll repeat this experiment, hopefully yeilding clearer results.

EDIT: Did you try adding enough H2O2 until no more visible bubbling occurs, and then filtering? If you haven't, that may explain why some reactants remained in the filtrate.



[Edited on 11/6/2015 by Velzee]




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[*] posted on 5-11-2015 at 20:00


Quote: Originally posted by Velzee  

Thank you for your attempt at this! I have a sample of this compound(if it is not copper oxide as you hypothesized) which remained its dark green color since I isolated it(weeks ago). I will try to post photos soon.

I would agree with your theory that this compound is simply a Cu oxide, had it not been for this:

Filtering the solution, after the reaction was complete, left an almost colorless solution(with the exception that some of the compound passed through the filter paper). If your theory's correct, where did the CuSO4 go?



I'm not sure I understand your question. Based on my results, I'd conclude that green solid you obtained is a mixture of copper oxide and [tetraammine] copper sulfate. Though I had thought you obtained the solid by evaporating the solution; are you saying that your filtration yielded a green crystalline solid and a colorless solution? That I can't explain.

I assumed my solution retained some blue color due to an excess of copper sulfate used in preparing the tetraamminecopper(II) sulfate. It's possible that reacting perfectly stoichiometric tetraamminecopper(II) sulfate with an excess of hydrogen peroxide will precipitate all of the copper as the oxide, an amorphous brown solid, leaving a colorless solution containing ammonium and/or sulfate ions. Since hydrogen peroxide only reacts with copper sulfate in the presence of ammonia, once the ammonia is 'used up' (assuming it is consumed in the reaction), any excess of copper sulfate used in preparing the tetraamminecopper(II) sulfate would remain in solution, coloring it light blue.


Quote: Originally posted by Velzee  
EDIT: Did you try adding enough H2O2 until no more visible bubbling occurs, and then filtering? If you haven't, that may explain why some reactants remained in the filtrate.


Nope. The bubbling continued until I ran out of peroxide. I also might conduct more tests once I replenish my supplies.

[Edited on 6-11-2015 by MolecularWorld]




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[*] posted on 6-11-2015 at 11:11


Quote: Originally posted by Velzee  
Compared to most on this forum, I am a very inexperienced chemist, therefore, please pardon my errors.

Many here are pretty inexperienced - i am for sure, so don't feel too lowly.

Appologies for being grumpy - i think i'm just jealous of the Fun you're having.

Errors are good : they are proof that something is happening.




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[*] posted on 6-11-2015 at 11:59


It's all good, aga! :)

I just took a photo of the mystery compound a few minutes ago, here it is:

tPxcnC9.jpg - 307kB

It is slightly more green than the camera allows.

[Edited on 11/6/2015 by Velzee]




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[*] posted on 6-11-2015 at 12:07


How, exactly, did you prepare the solid compound pictured?

As I said above, I was able to prepare a green liquid, but, on settling, it turned out to be a blue solution that looked green due to solid contaminants.




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[*] posted on 6-11-2015 at 12:12


OK.

Let's see if us inexperienced chemists can pull this apart in a logical way.

Firstly you make the teraamminecopper(II)sulphate complex

In solution we have the ions Cu2+, SO42-, NH4+, [Cu(NH3)4]2+ (had to google that) and water of course, so H3O+ and OH-.

Next you add H2O2, usually a powerful oxidiser.

What will it rip an electron from, and oxidise ?

Going from Blue (copper II) to Green (copper I) is interesting seeing as it's an Oxidiser being added.

Nope, i have no idea either !

So, what experiment would yield more information/clues ?

Reducing the pH of the starting solution to see if less OH- makes a difference ?

Dunno if tetraaminecopper(II)sulphate complex can exist in acidic solution - one way to find out ...




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[*] posted on 6-11-2015 at 12:34


***It should be noted that I used absolutely NO stoichiometric measurements when attempting to synthesize this compound***


1) I first prepared a saturated solution of CuSO4 by adding CuSO4 to a tall glass of water until no more would dissolved.

2) I then added enough NH4OH to convert the CuSO4 to tetraamminecopper(ii) sulfate, as seen here(ignoring the green on top): http://i.imgur.com/nNQoSmi.jpg

3) I then SLOWLY(to prevent foaming) added H2O2 until the solution stopped fizzing, leaving behind a fizzing solution, looking very similar to that of a cup of soda: https://youtu.be/SgRv2z9q0oU

4) I then attempted to filter the solution, but this proved difficult because some of the compound passed through the filter, yet the solution remained somewhat clear.

5)I allowed both the product on the filter paper to dry(this is the product that I have shown in my previous post), and the filtrate to evaporate, leaving very small and long needle-like crystals. I procrastinated collecting the product in the container(which took about a week to dry), and after two weeks, I disposed of it, keeping the product recovered from the filter paper.




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[*] posted on 6-11-2015 at 12:52


It's probably copper amine nitrites due to the oxidation of ammonia:

Y. Cudennec; et al. (1995). "Etude cinétique de l'oxydation de l'ammoniac en présence d'ions cuivriques". Comptes Rendus Académie Sciences Paris, série II,Méca; phys. chim. astron. 320 (6): 309–316.

Y. Cudennec; et al. (1993). "Synthesis and study of Cu(NO2)2(NH3)4 and Cu(NO2)2(NH3)2". European journal of solid state and inorganic chemistry. 30(1-2): 77–85.

[Edited on 6-11-2015 by deltaH]




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[*] posted on 6-11-2015 at 12:53


I suspect the overall reaction I had was something like:

[Cu(NH3)4]SO4 + 3 H2O2 → CuO + (NH4)2SO4 + 5 H2O + N2

More qualitative data:
- The amorphous brown precipitate (which I suspect is copper oxide) is quite voluminous; the precipitate formed from less than half a teaspoon of fine copper sulfate crystals, after settling for over 12 hours, had a volume of well over 100mL
- The washed brown precipitate dissolves readily in dilute sulfuric acid, to give a light sky-blue solution
- The washed brown precipitate dissolves very slowly in dilute ammonia, to give a dark purple/blue solution (probably tetraaminecopper hydroxide)
- Bubbling air through a dulute solution of tetraaminecopper(II) sulfate in excess ammonia, using an aquarium pump, for more than an hour at room temperature, did not produce any noticable reaction. It seems hydrogen peroxide (and/or heat) is required.

Based on all my experimental data, I suspect that the hydrogen peroxide reacts with the tetraaminecopper(II) sulfate to form: ammonium sulfate, copper oxide, water, and nitrogen gas.

I still need to confirm the identity of the gas(es); once I get more hydrogen peroxide, I'll test the gas with a burning/smoldering splint; the most likely candidates are nitrogen or oxygen.

I also need to attempt the reaction of perfectly stoichiometric tetraaminecopper(II) sulfate with hydrogen peroxide. If I'm right, a colorless solution of ammonium sulfate with a copper oxide precipitate should be produced. This would be odorless; addition of sodium hydroxide should produce a strong smell of ammonia.

I may also attempt this reaction with ice-cold solutions, combined slowly in an ice bath, to see if an unstable compound forms (that is otherwise being decomposed by the heat of the reaction).

Quote: Originally posted by aga  

Going from Blue (copper II) to Green (copper I) is interesting seeing as it's an Oxidiser being added.

I couldn't produce a green compound, only a blue solution that looked green due to fine solid contaminants. I'm now fairly certain the solids are an oxide of copper. The blue is probably unreacted copper sulfate.

Quote: Originally posted by aga  

Dunno if tetraaminecopper(II)sulphate complex can exist in acidic solution - one way to find out ...


I added some dilute sulfuric acid to tetraaminecopper(II) sulfate; it abruptly changed from dark purple-blue to light sky-blue. I've actually been considering using this color change to titrate my ammonia solution (lacking proper indicators).

[Edited on 6-11-2015 by MolecularWorld]




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[*] posted on 6-11-2015 at 12:59


Quote: Originally posted by Velzee  

5)I allowed both the product on the filter paper to dry(this is the product that I have shown in my previous post), and the filtrate to evaporate, leaving very small and long needle-like crystals. I procrastinated collecting the product in the container(which took about a week to dry), and after two weeks, I disposed of it, keeping the product recovered from the filter paper.


Based on my results, I strongly suspect that the substance in the filter was a mixture of copper oxide and undissolved copper sulfate, and the substance produced by evaporating the liquid is unreacted tetraamminecopper(II) sulfate also contaminated with copper oxide. In both cases, these are dark blue compounds, which might appear very dark green if contaminated.




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