Sciencemadness Discussion Board - The "WTF did I just make?" thread

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

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.

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

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

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.

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

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:

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 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 ...

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).

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.

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

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

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.

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

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

Brown particulate matter was clearly seen.

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.

[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

OK !

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.

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

aga - 6-11-2015 at 14:54

Seems to me that the Ammonia is just contributing OH- 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.

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 ....

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

MolecularWorld - 7-11-2015 at 08:03

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

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

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

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

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.

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.

The brown/black precipitate settled very quickly.

[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

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)2+ and SO₄2- ?)

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

[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 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

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

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:

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