Sciencemadness Discussion Board

The "WTF did I just make?" thread

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Texium - 31-8-2014 at 17:02

Thought I'd make a thread for when you are trying to do one thing, and then something totally unexpected happens instead. (Not sure if this phenomenon happens to everybody. It certainly happens to me a lot!)

To start it off, I was experimenting with various manganates. One that I made is what I think is calcium manganate (Hard to tell though, since there is almost no information about it out there. It isn't even in the CRC handbook). Anyway, I made it from solutions of K2MnO4 and CaCl2. It formed a purple-mauve precipitate, and the solution turned purple. The weird thing though was that after I vacuum filtered it, the purple solution turned red and turbid, and the CaMnO4 precipitate turned beige. I think what caused this was me not cleaning my büchner funnel well enough. (It smelled a little bit like chlorine still from MnO2+HCl)

I still have absolutely no idea what compounds were formed, but the red color is fascinating and totally unexpected!

Brain&Force - 31-8-2014 at 21:27

I think it disproportionated into manganese dioxide and permanganate as the solution became less basic.

bismuthate - 1-9-2014 at 04:51

I will be using this thread so much. :)
When i mixed solutions of Iron (III) chloride (with some ferrous in it possibly) and a small amount of sodium chromate I got a grey precipitate. What's this? could it be iron chromium oxide formed by ferrous ions?

Texium - 1-9-2014 at 08:37

Quote: Originally posted by Brain&Force  
I think it disproportionated into manganese dioxide and permanganate as the solution became less basic.
That's what I thought at first too, but this stuff was really red! I'd seen the disproportionation happen a couple of times before and this looked completely different. I guess it will remain a mystery though. Since it was most likely caused by contaminants on the funnel, it would be quite difficult to repeat.

Amos - 1-9-2014 at 08:43

Maybe see if it is repeatable? You might be headed towards a whole new class of colored solutions for me to collect:D

gdflp - 1-9-2014 at 12:01

As for WTF reactions, this one was quite a surprise for me http://www.sciencemadness.org/talk/viewthread.php?tid=32902

Antiswat - 1-9-2014 at 13:17

i once tried making manganese sulfate, reacting MnO2 with ascorbic acid (which apparently is very useful??)
it turned brownish, and overall seemed to be Mn3O4
then i attempted to react this new weird substance with H2SO4 and zinc metal powder, turned reddish
at some point i estimated out of others guesses that what i may have formed might be Mn2(SO4)3
i recall that it doesnt want to react very well with carbonates or anything alike

just thought i found a very neat method of making manganese salts that can be used, but apparently i just have some reddish clusterfuck now of some sort

http://puu.sh/bhsvh.jpg
ill just leave out trying to put it up here as a picture as its 100x800 and will fill entire page
as seen it has a nice colour, and its alot more clear than i recall it was
could do some reactions if anyone are interested in it etc

Brain&Force - 1-9-2014 at 20:29

I wonder if manganese(II) permanganate is stable.

Happy 1001st post Antiswat.

bismuthate - 2-9-2014 at 03:23

Well no, I don't think it could be because Mn(II) is oxidized by permanganate.

TheChemiKid - 2-9-2014 at 04:40

What about Mn(IV) Permanganate?

bismuthate - 2-9-2014 at 08:54

I would think Mn(III) permanganate would be more likely.

Texium - 2-9-2014 at 13:39

So you mean something like Mn(MnO4)3?
Err... I think not.

bismuthate - 2-9-2014 at 13:55

Well yeah it would probably just make Mn2O3.

Texium - 2-9-2014 at 15:02

I think that what I made was probably MnCl3. I saw a picture of Mn(III) solution on woelen's site, and it looked quite similar. Also, it has since separated into a brownish solid and a clear solution, quite similar looking to his second picture of Mn(III).

Antiswat - 3-9-2014 at 09:12

why dont you ask to have this thread stickied somewhere?? i think either other chemists does too simple reactions or are too proud to confess that they are confused over what the hell they have witnessed taking place infront of their eyes, perhaps im just one of the chosen few to get confused over what may once in a while happen in chemical reactions..
sticky thread!! you have my vote

prof_genius - 3-9-2014 at 09:35

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?

bismuthate - 19-1-2015 at 11:25

I performed this experiment:
I dissolved 5g of sodium molybdate in around 75mL of 3% H2O2.
I then mixed it with 50mL of a solution of what should be 2.5% copper acetate.
It formed a brown precipitate which settles rather slowly.
I doubt this is copper peroxomolybdate, but I can't think of another explanation for it. Does anybody know what this could be?

[Edited on 19-1-2015 by bismuthate]

j_sum1 - 19-1-2015 at 22:02

Quote: Originally posted by zts16  
To start it off, I was experimenting with various manganates.
[snip]
The weird thing though was that after I vacuum filtered it, the purple solution turned red and turbid, and the CaMnO4 precipitate turned beige. I think what caused this was me not cleaning my büchner funnel well enough. (It smelled a little bit like chlorine still from MnO2+HCl)

I still have absolutely no idea what compounds were formed, but the red color is fascinating and totally unexpected!


On a similar note, I have always wondered about the red at the end of this demonstration. I have never been able to duplicate it. I read woelen's description on manganates fully (linked upthread by zts16. So much to learn.) He never mentions this red either.

[edit faulty link]

[Edited on 20-1-2015 by j_sum1]

bismuthate - 20-1-2015 at 04:21

It's most likely Mn+3, but I can't conclude anything beyond that.

Eddygp - 20-1-2015 at 12:04

While cleaning some cheap glassware with ethanol, I noticed some small black specks from some previous experiment that I have conducted. Curiously enough, the specks dissolved in the ethanol yielding a vivid purple colour. I have been working with ammonium phosphates, iodides, sodium hydroxide and hexane. Any ideas? I don't think that it is iodine.

Molecular Manipulations - 20-1-2015 at 12:24

Why don't you think it's iodine? You had iodide present. Iodide can easily be oxidized to iodine. Where there any oxidizers?
Small amounts of iodine can make for a very bright purple color in ethanol.

kecskesajt - 13-2-2015 at 11:17

I tried to make 2-propyl nitrite with homemade KNO2 wich is made using KNO3 and lead metal.So I mixed 2-propanol with the nitrite, cooled down and added H2SO4 in small portions.At first it turned yellow than into a black-gooey.Smells like a car motor but 20 times more concentrated.Can somebody help?

Hg, Hg sulfates (I, II), KNO3, H2SO4, K2S2O8

quantumcorespacealchemyst - 25-2-2015 at 11:39

I have small unknown amount of Hg and Hg sulfates with 4.314gKNO3, 3.8ml 98% H2SO4, 5ml H2O and 6.531g Potassium persulfate (K2S2O8). It was combined at around room temperature and kept colder. No reaction yet.

I am unsure if heating and or refluxing/boiling will cause an accelerated reaction/runaway reaction.

Please advise me on the safest approach.

I have 0.998g of a calcogenide I want to add, and moniter it's reaction rate at around boiling. Before doing this, I am concerned with the nature of getting the reaction mix I have described, safely to the boiling point/reflux and get a good idea of it's own reactivity (thermodynamic behaviour?), prior to monitering another reagent addition.


thanks

[Edited on 25-2-2015 by quantumcorespacealchemyst]

Molecular Manipulations - 25-2-2015 at 12:08

A chalcogenide? Could you be any more vague about the nature of this chemical?
Quote:

I am unsure if heating and or refluxing/boiling will cause an accelerated reaction/runaway reaction.

I'm unsure, of what reaction you want to happen. But I'm sure it won't be a "runaway". Everything here is as oxidized as it could be (except half of the mercury is in the 1st oxidation sate), I can't think of a single reaction that could happen at all.
Quote:

Please advise me on the safest approach.

Find the nearest toxic waist processing facility and throw away your entire <s>lab</s> reaction.


[Edited on 25-2-2015 by Molecular Manipulations]

quantumcorespacealchemyst - 25-2-2015 at 12:12

Do you 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 don't believe I understand the action of K2S2O8 well.


[Edited on 25-2-2015 by quantumcorespacealchemyst]

Molecular Manipulations - 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?

quantumcorespacealchemyst - 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]

j_sum1 - 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.

blogfast25 - 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.

quantumcorespacealchemyst - 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

quantumcorespacealchemyst - 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.

quantumcorespacealchemyst - 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?

Unknown compound created by adding H2O2 to a solution of tetraamminecopper(ii) sulfate

Velzee - 5-11-2015 at 13:48

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


aga - 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.

Texium - 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]

aga - 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]

Texium - 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?

Velzee - 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]

MolecularWorld - 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.

MolecularWorld - 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]

Velzee - 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]

MolecularWorld - 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]

aga - 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.

Velzee - 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]

MolecularWorld - 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.

aga - 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 Cu<sup>2+</sup>, SO4<sup>2-</sup>, NH4<sup>+</sup>, [Cu(NH3)4]<sup>2+</sup> (had to google that) and water of course, so H3O<sup>+</sup> and OH<sup>-</sup>.

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<sup>-</sup> makes a difference ?

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

Velzee - 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.

deltaH - 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]

MolecularWorld - 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]

MolecularWorld - 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.

aga - 6-11-2015 at 13:54

Starting with 3.00g CuSO4 (5H2O ?) dissolved in 40ml DIW

start.JPG - 165kB

2ml of 33% Ammonia was added.

TACS is destroyed in acidic conditons.
1 ml of conc H2SO4 was added to the Right hand test tube.

TACS in acid.JPG - 185kB

Re-basifying the right-hand sample with more ammonia gave a two-layer system in which the TACS colour re-appeared.

make basic.JPG - 156kB

Adding 3% H2O2 caused fizzing/mild foaming with a brown tint to the bubbles.

Brown particulate matter was clearly seen.

H2O2 added.JPG - 148kB

Adding H2O2 to just the CuSO4 solution gave no reaction.

Adding a few grains of NaOH to the left hand H2O2 + CuSO4 solution caused fizzing and a lot of brown particulate matter to precipitate.

+NaOH.JPG - 162kB


[Edited on 6-11-2015 by aga]

deltaH - 6-11-2015 at 13:58

Copper sulfate is dissolved in excess ammonia, results in the formation of a solution of Schweizer's reagent (tetraaaminecopper(II) hydroxide) and ammonium sulfate.

Now you add hydrogen peroxide. This oxidises some of the ammonia and maybe forms copper(I) oxide which precipitates as the brown stuff, but the latter is surprising.

NOTE: Please stop speaing about "copper oxide", copper has two oxides, black CuO or cupric oxide or copper(II) oxide and the red-brown cuprous oxide, Cu2O or copper(I) oxide.

It is possible that the ammonia doesn't get oxidised at all and this can be confirmed by doing the same reaction, but using sodium hydroxide to precipitate copper(II) hydroxide and then add the peroxide. If it forms the same brown stuff, then the ammonia is simply acting as a base and not really doing anything else. If not, ammonia is getting oxidised and plays a key role here.

[Edited on 6-11-2015 by deltaH]

aga - 6-11-2015 at 14:31

Quote: Originally posted by deltaH  
It is possible that the ammonia doesn't get oxidised at all and this can be confirmed by doing the same reaction, but using sodium hydroxide to precipitate copper(II) hydroxide and then add the peroxide.

OK !

Copper Sulphate solution with some random amount of NaOH added.

Before.JPG - 142kB

Now add the 3% H2O2

After.JPG - 162kB

Fizzed like crazy and some foam left the test tube at moderate speed.

It is now a Green solution.

Will leave it overnight to see if it has any precipitate.

@Velzee

Big thanks !

I was being grumpy and jealous - this is great fun !

[Edited on 6-11-2015 by aga]

deltaH - 6-11-2015 at 14:46

Very interesting, I didn't expect the copper hydroxide to go green upon addition of the peroxide in the case of the sodium hydroxide addition :o

aga - 6-11-2015 at 14:54

Seems to me that the Ammonia is just contributing OH<sup>-</sup> to the mix.

Strange.

After NaOH, KOH and Ammonia, i struggle to lay hands on another strong base.

Got plenty of acids, which means i'm acid-biased.

DOH !

MolecularWorld - 6-11-2015 at 15:12

Quote: Originally posted by deltaH  
Copper sulfate is dissolved in excess ammonia, results in the formation of a solution of Schweizer's reagent (tetraaaminecopper(II) hydroxide) and ammonium sulfate.

No, it forms <a href="http://en.wikipedia.org/wiki/Tetraamminecopper(II) sulfate" target="_blank">tetraamminecopper(II) sulfate</a> <img src="../scipics/_wiki.png" />.
Quote:

Now you add hydrogen peroxide. This oxidises some of the ammonia and maybe forms copper(I) oxide which precipitates as the brown stuff, but the latter is surprising.

Yep.
Quote:

NOTE: Please stop speaing about "copper oxide", copper has two oxides, black CuO or cupric oxide or copper(II) oxide and the red-brown cuprous oxide, Cu2O or copper(I) oxide.

I was aware of this, but I didn't know which of the oxides was being produced, and since they behave similarly under these conditions, it really doesn't matter. I will henceforth use the term "an oxide of copper" when referring to a compound of copper and oxygen at an unknown ratio.
Quote:

It is possible that the ammonia doesn't get oxidised at all and this can be confirmed by doing the same reaction, but using sodium hydroxide to precipitate copper(II) hydroxide and then add the peroxide. If it forms the same brown stuff, then the ammonia is simply acting as a base and not really doing anything else. If not, ammonia is getting oxidised and plays a key role here.

There is no hydroxide in tetraamminecopper(II) sulfate. It's not like making Tollens' reagent, at no point in the addition of ammonia to copper sulfate did a precipitate appear. As such, reacting hydrogen peroxide with copper hydroxide is a completely different and largely unrelated reaction.
Quote: Originally posted by aga  

Copper Sulphate solution with some random amount of NaOH added.
Now add the 3% H2O2
Fizzed like crazy and some foam left the test tube at moderate speed.
It is now a Green solution.

That's surprising. The liquid appears to be a much brighter green than any my reactions produced. Gotta get some more peroxide and try it myself!

[Edited on 7-11-2015 by MolecularWorld]

aga - 6-11-2015 at 15:31

Doing is believing.

Edit

Or at least verifying ...

[Edited on 6-11-2015 by aga]

MolecularWorld - 6-11-2015 at 15:38

Here's an interesting reference:
Quote:

"At 0°C neutral hydrogen peroxide converts an aqueous suspension of cupric hydroxide into the brown, crystalline peroxide."
"The product is always more or less impure. When moist, it decomposes rapidly with evolution of oxygen and formation of cupric oxide, but the decomposition of the dry substance is slow."
"The compounds obtained by the action of hydrogen peroxide on cupric hydroxide are regarded by Aldridge and Applebey as probably consisting of a mixture in varying proportions of cupric oxide and hydroxide with a yellow, gelatinous peroxide of the formula CuO2."

A mixture of yellow peroxide and blue hydroxide could produce the green coloration. If the above is applicable, I would expect your green liquid to turn brown.

Velzee - 6-11-2015 at 15:49

This is a lot more interesting than I first imagined! xD I want more supplies!

aga - 6-11-2015 at 15:53

That's the way it goes - something looks Boring and can turn out to be fascinating !

We shall see if my Finest Green of MostGreenyNess goes Brown.

I'll dump it into a petri dish so it can oxidise faster (would have thought the H2O2 would have done that already ...)

aga - 7-11-2015 at 07:50

Well, i was about to dump the Green liquid into a petri dish and saw this ....

rack.JPG - 152kB

The colours are hard to make out.

There is a dark green layer on top with a blue layer on the bottom.

Here's a close-up in sunlight

closer.JPG - 313kB

MolecularWorld - 7-11-2015 at 08:03

Quote: Originally posted by aga  
Well, i was about to dump the Green liquid into a petri dish and saw this ....
The colours are hard to make out.
There is a dark green layer on top with a blue layer on the bottom.
Here's a close-up in sunlight


If you were using 3% peroxide, there was probably an excess of copper hydroxide, which settled out. Some hydroxide stays in solution/suspension, and, if there's a "yellow, gelatinous peroxide", as my reference above states, the mixture in suspension would look green.

If the liquid is a mixture of copper hydroxide and copper peroxide, it should decompose on gentle heating to brown/black cupric oxide with the production of some oxygen. The hydroxide by itself would also decompose, but without producing gasses.

Which reaction produced the brown precipitate and colorless solution in the tube at the left?

aga - 7-11-2015 at 08:51

Erm (blush) i cannot remember.

This is a big problem with random 'mixing stuff up'.

Certainly two or more of copper sulphate/ammonia/NaOH/hydrogen peroxide.

I do recall that one went brown the instant the reactants were mixed, with strong fizzing.


[Edited on 7-11-2015 by aga]

MolecularWorld - 7-11-2015 at 09:42

Quote: Originally posted by MolecularWorld  

Which reaction produced the brown precipitate and colorless solution in the tube at the left?

Quote: Originally posted by aga  
Erm (blush) i cannot remember.
This is a big problem with random 'mixing stuff up'.
Certainly two or more of copper sulphate/ammonia/NaOH/hydrogen peroxide.
I do recall that one went brown the instant the reactants were mixed, with strong fizzing.

Based on your posts above, it's either: TACS + peroxide; or copper sulfate + peroxide + NaOH (in that order). You might be able to differentiate between the two by adding some hydroxide to the colorless liquid, only the TACS + peroxide mixture would smell of ammonia (assuming the ammonia isn't completely destroyed by the reaction).

Re-reading your posts, I realized something: the only difference between
Quote: Originally posted by aga  

Adding H2O2 to just the CuSO4 solution gave no reaction.
Adding a few grains of NaOH to the left hand H2O2 + CuSO4 solution caused fizzing and a lot of brown particulate matter to precipitate.

and
Quote: Originally posted by aga  

Copper Sulphate solution with some random amount of NaOH added.
Now add the 3% H2O2
Fizzed like crazy and some foam left the test tube at moderate speed.
It is now a Green solution.

...is that, in the first case, any copper peroxide formed would be subject to the localized heating around the dissolving NaOH prills, causing it to immediately decompose to cupric oxide and oxygen. In the second case, there might not have been enough heat to decompose the copper peroxide, which it why it persisted to give a "green" liquid (actually a suspension of brown/yellow and blue particles). Again, if the green liquid is a mixture of suspended of copper peroxide and copper hydroxide, gentle heating should decompose it to cupric oxide and oxygen. This will also happen slowly at room temperature (though the gas evolution may be too little and too slow to notice). Alternately, due to the ejection of foam, it's possible very little of the peroxide reacted with the copper hydroxide, but all of the copper peroxide thus formed decomposed, creating a only a small amount of cupric oxide in suspension with the rest of the copper hydroxide. In which case, the green liquid would decompose on heating, but without evolution of any gasses.

aga - 7-11-2015 at 11:42

@Velzee

This what proper chemists can do without even knowing the precise quantities !

(I mean MolecularWorld, not me.)

OK. MW, will go heat the green/blue stuff and post another photo.

aga - 7-11-2015 at 12:05

The Green part is actually looking Blacker than earlier, however it is Night here. Some Blue is still visible at the bottom.

start.JPG - 135kB

Removing some of the water and shaking gives a very dark green colour, almost black.

Gentle heating for ~1 minute has converted all the Colour to brown/black with no bubbles evolved.

agitated.JPG - 138kB

The brown/black precipitate settled very quickly.

precip.JPG - 154kB

[Edited on 7-11-2015 by aga]

MolecularWorld - 7-11-2015 at 12:25

@aga: I'm far from "proper chemist", but I have read a lot and do enjoy the occasional "random mixing stuff up". :)

Based on the all the experimental data presented here by aga, Velzee, and myself, as well as this Wikipedia article that appeared mysteriously a few hours ago, I can say with a high degree of confidence that our reactions of hydrogen peroxide with tetraamminecopper(II) sulfate solution and copper hydroxide suspension are both producing <a href="http://en.wikipedia.org/wiki/Copper_peroxide" target="_blank">copper peroxide</a> <img src="../scipics/_wiki.png" />, an unstable olive-green compound. This decomposes rapidly, to give oxygen gas and mixtures containing varying amounts of brown/black cupric oxide and excess blue-colored reactants, which look green in combination.

I'm still curious as to what the byproducts of the reaction between tetraamminecopper(II) sulfate solution and hydrogen peroxide are. I'll be doing some experiments later this week in an attempt to determine the nature of complete reaction.

aga - 7-11-2015 at 12:33

Please 'fess up when you're done (report your findings !).

aga - 7-11-2015 at 12:57

Quote: Originally posted by MolecularWorld  

No, it forms <a href="http://en.wikipedia.org/wiki/Tetraamminecopper(II) sulfate" target="_blank">tetraamminecopper(II) sulfate</a> <img src="../scipics/_wiki.png" />.

As a very inexperienced amateur chemist, i feel that i must challenge that assertion.

Tetraaminecopper(II) sulphate should really only exist as a Solid.

In solution it's a mix of Ions (tetraaminecopper(II)<sup>2+</sup> and SO4<sup>2-</sup> ?)

It's a bit like saying NaCl in solution is NaCl, wheras in aqueous solution it is Na<sup>+</sup> and Cl<sup>-</sup>

[Edited on 7-11-2015 by aga]

MolecularWorld - 7-11-2015 at 13:06

Remember, that assertion was in response to:
Quote: Originally posted by deltaH  
Copper sulfate is dissolved in excess ammonia, results in the formation of a solution of Schweizer's reagent (tetraaaminecopper(II) hydroxide) and ammonium sulfate.

You're right that there's an equilibrium between a variety of ionic species. But in perfectly stoichiometric tetraamminecopper(II) sulfate, the ratio of ammonia:copper:sulfate is 4:1:1. If it were instead a mixture of tetraamminecopper(II) hydroxide and ammonium sulfate, the ratio would be different; eg, for there to be a 1:1 copper:sulfate ratio, there would be 6 ammine/ammonium ions.

[Edited on 7-11-2015 by MolecularWorld]

aga - 7-11-2015 at 13:27

Wow ! (at least) 6 species in an ionic soup.

No wonder this was a tricky one to work out.

Unlikely to be in 'correct' stoichiometric amounts either in the OP or my own random mix-ups, so i guess anything could happen, including the ionic versions of what deltaH said.

pH seems to be a Key part in all of this.

Excellent that your heating suggestion turned out as expected.

JJay - 8-11-2015 at 19:20

Woot! I just got a near-quantitative yield using OTC reagents in a synth that had pretty crappy yields in the literature.

MolecularWorld - 8-11-2015 at 19:53

Quote: Originally posted by JJay  
Woot! I just got a near-quantitative yield using OTC reagents in a synth that had pretty crappy yields in the literature.


All right, I'll play along. Per the topic of this thread: WTF did you just make?

JJay - 9-11-2015 at 02:11

Quote: Originally posted by MolecularWorld  
Quote: Originally posted by JJay  
Woot! I just got a near-quantitative yield using OTC reagents in a synth that had pretty crappy yields in the literature.


All right, I'll play along. Per the topic of this thread: WTF did you just make?


I'm actually trying to figure that out... there seems to be some controversy over what the product of this reaction is... right now I am distilling some DCM off of an extraction from a reaction with the product to see if I get anything. The reaction did not run as I expected, but it could be that one of the side products simply didn't decompose as fast as I expected.

EDIT: The synth was too good to be true.... I definitely did not get the product I wanted; I got some isomer in near-quantitative yield. I don't know what it is, but I don't really care either.


[Edited on 9-11-2015 by JJay]

Back to tetraamminecopper(II) sulfate and ammonia...

MolecularWorld - 13-11-2015 at 11:24

I currently lack the means to determine the exact concentration of my ammonia solution, but the MSDS states it's at least 2%. Based on that, I prepared a solution of tetraamminecopper(II) sulfate with an excess of ammonia by adding 10g of copper sulfate pentahydrate to 1L of this ammonia solution. This was diluted again by half for use in the video below.

When 3% hydrogen peroxide is added to this solution, much gas is evolved, and the liquid turns a dingy green. This may be due to the formation of copper peroxide, or due to suspended particles of cupric oxide in the blue solution.

<iframe sandbox width="280" height="156" src="//www.youtube.com/embed/P_nXcoSV0ks?rel=0" frameborder="0" allowfullscreen></iframe>

The gasses reignite a smoldering splint, though not as rapidly as I would expect from pure oxygen. The volume of gas is too large to be explained simply from the decomposition of the peroxide. I suspect the gasses produced at this point are a mixture of nitrogen and oxygen, with a greater concentration of oxygen than air.

More peroxide is added, until rate of gas production drops significantly. This indicates the primary reaction is complete, and the peroxide is simply being decomposed catalytically by the cupric oxide particles. This leaves a dark brown suspension of extremely fine cupric oxide particles in a colorless solution. There was not enough gas produced at this point for me to confirm it's identity with a smoldering splint.

The solution was heated gently and allowed to settle, resulting in a brown precipitate and a light blue solution.

Blue solution with brown precipitate.jpg - 82kB

My tests earlier in this thread show that the fine cupric oxide will dissolve slowly in ammonia, and as my tetraamminecopper(II) sulfate had excess ammonia, this gives the slight blue coloration. With perfectly stoichiometric tetraamminecopper(II) sulfate, I expect a colorless solution. I lack the means to determine the exact concentration of my ammonia, so I'm unable to prepare a 'pure' tetraamminecopper(II) sulfate solution at this time.

The colorless (slightly blue) solution was decanted. A small aliquot was added to sodium hydroxide: a stronger ammonia odor results, and the solution is nearly decolorized. This suggests the presence of ammonium, and very little copper in solution. Another aliquot was added to calcium chloride solution: the solution turns cloudy, and a small amount of white precipitate eventually settles. This suggests the presence of sulfate.

Thus:
On addition of hydrogen peroxide to tetraamminecopper(II) sulfate solution:
- Copper peroxide is formed.
- Some of the ammonia is oxidized to nitrogen (and water)
- The copper peroxide is decomposed by the heat of the oxidizing ammonia, to give cupric oxide and oxygen
- The solution remaining after the cupric oxide precipitates is ammonium sulfate

Overall reaction:
2 [Cu(NH3)4]SO4 + 6 H2O2 → 2 CuO + 2 (NH4)2SO4 + 12 H2O + 2 N2 + O2

I may repeat these tests once I can prepare stoichiometric tetraamminecopper(II) sulfate solution.

[Edited on 13-11-2015 by MolecularWorld]

LargeV - 28-11-2015 at 10:30

I dissolved 1.5g sodium bisulfate in 20 ml of water and added 10 ml of 3% hydrogen peroxide and then added a 1987 US penny.

I left it out for 6 hours and when I put in a few pellets sodium hydroxide, it made this weird olive yellow-green substance that went away when the solution was swished.

I pippeted 1 ml of the solution into a solution with half a gram of sodium hydroxide, and it made a green precipitate/colloid.

This turned brown-black when heated, so I presume I accidentally made copper peroxide.

I then left it out for a day, and I come back to see that the penny has a huge black oxide layer and bubbles quite violently when I put another gram of sodium bisulfate in.

Then, I put in another few pellets of sodium hydroxide. The sodium hydroxide pellets sit inside the solution for a second and start to grow a huge white insoluble snake out the top of the pellets.

I thought this was weird, so I pipetted 1.5 ml in the sodium hydroxide solution with the precipitate of copper peroxide I left out yesterday, and it dissolved the copper peroxide, which I am pretty sure is insoluble.

What even?

http://i.imgur.com/fZre1F1.png http://i.imgur.com/4y7NNt2.png



[Edited on 11/28/2015 by LargeV]

Texium - 28-11-2015 at 11:56

Starting to wish that I never started this thread... It's turned into a string of unscientific and poorly executed experiments that all need to be redone in a more controlled way.

aga - 28-11-2015 at 16:28

Cheer up !

LargeV just made it better.

What's in a 1987 US penny ?

Copper mostly i guess.

[Edited on 29-11-2015 by aga]

Texium - 28-11-2015 at 18:40

US pennies made after 1983 have been made of zinc plated with copper. So it's quite likely that there is some zinc contamination present. I'd recommend using copper wire as a source for copper as it needs to be quite high purity to function properly. Copper pipe is decent, but not quite as good. Pennies, even the pre-1983 ones, will contain zinc. Though often cited as being pure copper, pre-1983 pennies also contain 5% zinc.

Perhaps try again with pure copper wire.

LargeV - 28-11-2015 at 18:42

I will try to find some! Thank you for the feedback :D

LargeV - 28-11-2015 at 19:40

I have a suggestion of what this mystery snake actually is. Because of the zinc contamination, it might have formed zinc hydroxide, which is white and insoluble like the snake! I will report back on my results with pure copper.

EDIT: I have my results! It was, in fact, the zinc! I tried with a piece of pure copper metal, it dissolved to form copper sulfate which I did not see in the first trial.

I then tried it with a 1972 penny which also did form copper sulfate without the zinc contamination too, as evidenced by only copper peroxide forming on addition of the hydroxide!

[Edited on 11/29/2015 by LargeV]

ave369 - 29-11-2015 at 01:04

I prepared fresh Cr2O3 by thermally decomposing ammonium dichromate. This Cr2O3, I tried to dissolve in 25% HCl, wanting to prepare CrCl3. The solution came out yellow. It does not change its color when I add strong acids (H2SO4 conc) or strong bases (NaOH). It, however, does change color to very pale green when I add a reducer such as sodium thiosulfate.

What is the yellow thing? Is it some complex of CrCl3, or hexavalent chromium?

Addendum: adding very strong NaOH resulted in the solution changing color to pale green and a wispy precipitation forming.

Addendum 2: I added potassium permanganate to the yellow liquid, and it turned deep brown, almost opaque, but with no precititate. This brown liquid was acidic and very bleaching, it discolored the paper strip after turning it red, and smelled of chlorine. But no massive evolution of chlorine.

Addendum 3. Re-did the experiment with a mixture of sulfuric and hydrochloric acids dissolving Cr2O3. Situation normal, a green solution which, I think, is a mixture of chromium chloride and sulfate. Addition of KMnO4 results in massive evolution of chlorine.

[Edited on 29-11-2015 by ave369]

MolecularWorld - 29-11-2015 at 01:44

@ave369: Did the Cr2O3 dissolve completely in the HCl? If not, then my guess is there was still some ammonium chromate/dichromate left in your chromium(III) oxide, which simply leached out in the HCl.

I just tested a dilute, yellowish, potassium dichromate solution (I have no ammonium dichromate): concentrated sulfuric acid turned it slightly more orange (though still yellow), while a large amount of sodium hydroxide turned it greenish.

[This was posted after Addendum 2 & before Addendum 3.]

[Edited on 29-11-2015 by MolecularWorld]

ave369 - 29-11-2015 at 01:47

I'll repeat the experiment with a specimen of well leached Cr2O3 I have.

LargeV - 4-12-2015 at 13:42

I tried making copper peroxide in dilute solutions, and it made this orange precipitate with no signs of copper hydroxide.

0.5g CuSO4 + 3ml H2O2 + 10ml water + 1.5ml 0.5M NaOH

[Edited on 12/4/2015 by LargeV]

MolecularWorld - 4-12-2015 at 13:56

That's similar to a reaction aga documented upthread.
Copper peroxide is only stable below 6*C, above which it decomposes to other copper oxides and oxygen; the orange is probably a finely divided CuO/Cu2O precipitate.

[Edited on 4-12-2015 by MolecularWorld]

(If you all don't mind)Back to my previous experiments:

Velzee - 21-12-2015 at 17:53

I've obtained more NH4OH, so I may repeat some of my previous experiments, and perhaps make some more accurate observations.

Meanwhile, I discovered something while reading the Wikipedia article of NH4OH:


Quote:


When ammonium hydroxide is mixed with dilute hydrogen peroxide in the presence of a metal ion, such as Cu2+, the peroxide will undergo rapid decomposition.


Which explains partially of what I observed, and corroborates what @MolecularWorld was hypothesizing. Hm, I should have done some more research, I admit.

Aurium - 27-12-2015 at 11:48

Hi guys,

So I tried to make some nitrocellulose. I had made it many times before with success.
This time however I tried using cotton string, expecting to end up with string form NC.
Regular H2SO4 + KNO3 + HNO3 nitration bath. I cannot remember the exact amounts I used right now. I made some calculations conserving pH and [NO3], anyway,
This happened:



nitro 1.jpg - 423kB nitro 2.jpg - 363kB nitro 3.jpg - 438kB


So I added the string unrolled. Immediately after hitting the acid it curled up into a big ball.
I finished the nitration as best I could and washed it neutral. I then broke it into smaller pieces.

Now, this thing, plastic like nitrocellulose:

- Not ignitable at all. Just decomposes over flame.
- Not soluble in Acetone. At all, not soluble in dry boiling acetone whatsoever.
- Is quite rigid.

I am aware that NC was used a common plastic in older days, but I find it strange that this isn't soluble in acetone.


Some insight as to why my NC turned out so different?

Velzee - 13-5-2016 at 20:32

It's not what I made; it's what did he make:

https://m.youtube.com/watch?v=DJGVZNth68k

j_sum1 - 13-5-2016 at 21:31

Quote: Originally posted by Velzee  
It's not what I made; it's what did he make:

https://m.youtube.com/watch?v=DJGVZNth68k

There is a thread on this from about a year ago. Look up "transparent sodium". Also look at some related YT vids by thunderf00t. He does the same with potassium.

The transparent bead is not sodium. Best guesses are that it is Na2O or NaOH or some combinations thereof. But there are a whole bunch of unanswered questions:

What is the blue/black coating that appears between the flame and the transparent bead?
What exactly is the paper doing to cause this effect? (I have done some experimenting and found that with Na, some contact with the paper is necessary -- it is not merely a matter of limited availability to water. K behaves differently and will give the effect without paper.)
Why does the bead crackle and explode at the end?
If the bead is solid NaOH or Na2O, why does it not sink in the water.

There is a lot going on. Given that the actual mechanism for Na explosion in water was not determined properly until about 18 months ago (Coulombic explosion), I think it will take some time for all of these questions to be completely resolved. Nevertheless, if you have some Na then it is worth having a play with it. I for one would prefer to see a bunch of YT clips investigating this than repeated clips of hoons chucking sodium or potassium into ponds.

symboom - 23-8-2017 at 19:49

The tetramine copper sulfate reaction with hydrogen peroxide forms copper peroxide
Check my link in my signature I have pictures of what is described

[Edited on 24-8-2017 by symboom]

[Edited on 24-8-2017 by symboom]

Melgar - 23-8-2017 at 21:46

I've noticed that when I dissolve nickel or copper using HCl and peroxide, that it eventually starts to smell different, and actually a bit... pleasant. This isn't really a word I'm used to using to describe compounds made exclusively of strong oxidizers. My best guess is that it's perchloric acid, since it doesn't smell like chlorine or a strong acid. Although this isn't among the various methods that are described to synthesize it, I'd guess that the metal being dissolved works as a catalyst, because nickel and copper catalyze so many other reactions.

Melgar - 23-8-2017 at 21:50

Quote: Originally posted by LargeV  
I have a suggestion of what this mystery snake actually is. Because of the zinc contamination, it might have formed zinc hydroxide, which is white and insoluble like the snake! I will report back on my results with pure copper.

EDIT: I have my results! It was, in fact, the zinc! I tried with a piece of pure copper metal, it dissolved to form copper sulfate which I did not see in the first trial.

I then tried it with a 1972 penny which also did form copper sulfate without the zinc contamination too, as evidenced by only copper peroxide forming on addition of the hydroxide!

[Edited on 11/29/2015 by LargeV]


Close! You formed sodium zincate.

myristicinaldehyde - 1-9-2017 at 10:39

I have something mildly interesting.

A solution of copper sulfate and dilute sulfuric acid, when left with tin pellets, plates the pellets with copper. However, the supernatant, after dropping a precipitate of tin oxide/hydroxide/oxysulfate is a clear green, like the tetrachlorocuprate complex, but maybe slightly more blue.

Why isn't the solution the gentian blue of hexaaquacopper(ii)? (My best guess is the presence of cu(i) ions, but they are often white and insoluble, which doesn't explain a clear green solution)

j_sum1 - 24-10-2017 at 18:41

I came up with a weird one today.

Some background -- Following a textbook procedure to determine the empirical formula of copper oxide -- teaching students about gravimetric analysis.

Procedure called for CuO to be dissolved in sulfuric acid. Then zinc added to precipitate out the copper quantitatively. Then washing, drying and weighing to determine the mass of Cu in the original sample.

We did not follow procedure exactly. In an effort to accelerate the dissolving process we used a decent excess of concentrated sulfuric acid (the "add a splash" method.)
This led to an overly acidic solution and hence the zinc granules added reacted producing H2.
Feeling a need for excess zinc I finished off the bottle. It was reagent grade but who knows its history.

The result was a yellowish solution and a grey black sediment.

The sediment showed traces of copper in a flame test but not much. It resisted dissolving in either HCl or H2SO4.
The yellow solution showed no indication of copper present. It had trace amounts of Fe3+.

Curious as to what may have caused the yellow tinge one of the students suggested a lead contaminant. We added some potassium iodide to check this but did not get the instantaneous vivid yellow that one would expect if lead was present.

What we did get was a slowly-forming orange precipitate (photographed). The colour was reminiscent of cadmium compounds and we discussed the likelihood of this. It was a nice idea but I think can be dismissed since the colour does not match CdI2 or CdSO4. CuI2 is not a match either.

So, I am curious. What the hell did we just make?

2017-10-25 12.13.58.jpg - 242kB

[edit]
The colour is brighter than the photo indicates. Similar I guess to the colour of potassium dichromate.

[Edited on 25-10-2017 by j_sum1]

ninhydric1 - 24-10-2017 at 20:24

Could it be precipitated Cu2O (copper(I) oxide)? The color looks a bit too fluorescent to be Cu2O in my opinion but the colors are pretty close.

j_sum1 - 24-10-2017 at 20:51

Quote: Originally posted by ninhydric1  
Could it be precipitated Cu2O (copper(I) oxide)? The color looks a bit too fluorescent to be Cu2O in my opinion but the colors are pretty close.

I confess to not being too familiar with copper (I). But that idea had occurred to me. It (my mystery compound) is very orange and bright. I am not sure why the oxide would precipitate after the addition of iodide. And I would have thought that either of the two copper oxides would be incompatible with conditions that are still highly acidic.

FWIW, a spritz of the solution through a bunsen flame did not reveal any green or blue coloration. And there was no visible reaction with concentrated ammonia. I guess that only rules out Cu(II).
I might add some peroxide to the yellow solution and see if any blueness appears.



[Edit] Clarification and addendum.


And no visible reaction attempting to oxidise any Cu(I) to Cu(II). There did seem to be a temperature rise though.

[Edited on 25-10-2017 by j_sum1]

Melgar - 25-10-2017 at 03:34

I bet it's a sulfide of some sort. Zinc will reduce sulfuric acid to zinc sulfide, especially in a concentrated solution, then hydrolyze to ZnO and H2S. H2S would react quickly with any copper salts, precipitating CuS, which is black. Elemental sulfur could be responsible for some of the yellow color, or there could be any of a number of mixed sulfides.

Texium - 25-10-2017 at 05:01

Quote: Originally posted by Melgar  
I bet it's a sulfide of some sort. Zinc will reduce sulfuric acid to zinc sulfide, especially in a concentrated solution, then hydrolyze to ZnO and H2S. H2S would react quickly with any copper salts, precipitating CuS, which is black. Elemental sulfur could be responsible for some of the yellow color, or there could be any of a number of mixed sulfides.
That sounds like a bit of a stretch. CuS is very black indeed, just as dark as CuO, and if it was in there, you'd be looking at a muddy brown color, not a bright orange.

Edit: Also interesting to look back at when I started this thread in my early days. My original puzzle of a red manganese solution was most certainly just manganate disproportionating into permanganate and very finely divided manganese dioxide. I couldn't believe at the time though that a solution could contain a precipitate while still appearing completely transparent, and it wasn't until later experiments with Prussian blue that I came to realize that.

[Edited on 10-25-2017 by zts16]

Melgar - 25-10-2017 at 05:22

Quote: Originally posted by zts16  
That sounds like a bit of a stretch. CuS is very black indeed, just as dark as CuO, and if it was in there, you'd be looking at a muddy brown color, not a bright orange.

The precipitate was described as being black, although the orange color is almost certainly from iodine. I missed that part the first time around. Add sulfuric acid to a potassium iodide solution, and see what color it turns. No metals necessary.

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