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deltaH
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Very nice indeed, very well done MrHomeScientist... pity no pics of the LED's 
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woelen
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This indeed is a nice write-up with good pictures. Another thing which you may find interesting is repeating the experiment, but now without heating.
This may lead to formation of nice large crystals, especially if you do the experiment at a scale of a few grams.
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blogfast25
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Nice write up and photos. Good demonstration of the dual oxidation state of the copper. Might have a bash at this myself...
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bfesser
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Wonderful work, <strong>MrHomeScientist</strong>. I second the idea of taking some photos of the compound under monochromatic
illumination.
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MrHomeScientist
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Thanks for all the high praise, especially from such illustrious members as yourselves!  Seeing that people enjoy the things I do is great motivation for me. I'm very glad to contribute to the scientific hobbyist community.
I'll have another attempt at making Cu2O and see if I can get a sample worthy of the illumination test. If that doesn't pan out for
whatever reason, I'll at least post some photos of Chevreul's salt under illumination by itself.
[edited to fix formula]
[Edited on 10-17-2013 by MrHomeScientist]
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deltaH
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Yes the Cu2O is important because the only difference in the two is that Cu2O should appear black under green light while Chevreul's is only partially
dark!
Well done again for your most excellent work!!!
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Volanschemia
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I've just made Chevreul's salt for the first time very successfully, using the CuSO4 + Na2S2O5 method.
A truly fascinating synthesis. I was wondering however, what is actually happening in the reaction. Obviously a copper compound is formed when the
CuSO4 and Na2S2O5 are combined, but not sure what. Would it be Copper(II) Sulphite? And then this
compound either decomposes into Chevreul's Salt, or another reaction takes place.
This may be mentioned in one of the references posted in this thread, but I have not read through them all. If someone knows the answer that would be
great!
Thanks.
"The chemists are a strange class of mortals, impelled by an almost insane impulse to seek their pleasures amid smoke and
vapor, soot and flame, poisons and poverty; yet among all these evils I seem to live so sweetly that may I die if I were to change places with the
Persian king" - Johann Joachim Becher, 1635 to 1682.
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Eddygp
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How does this salt react with ammonia? Does it form a complex, like the reaction of ammonia and copper(II) sulfate?
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Amos
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Quote: Originally posted by Eddygp  | | How does this salt react with ammonia? Does it form a complex, like the reaction of ammonia and copper(II) sulfate? |
One would be very inclined to say so, as the color of the resulting solution looks identical to the tetraammine complex that results from the
dissolution of other copper(II) compounds in aqueous ammonia. Now where the copper(I) ends of going is beyond me.
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woelen
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The copper(I) also goes into solution, it forms a diammine complex [Cu(NH3)2](+). This complex is colorless.
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Eddygp
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| Quote: Originally posted by woelen | | The copper(I) also goes into solution, it forms a diammine complex [Cu(NH3)2](+). This complex is colorless. |
This is what I suspected... I am thinking of the possible uses of that complex. If it is isolated, it would probably respond to oxidation (when H2O2
is added, dichromate, etc.) and probably under reductive conditions. Either case, it might be interesting as a soluble copper indicator of
oxidative/reductive stress.
Really, I am just speculating about the possible outcomes of these circumstances. I am not too experienced in how complexes differ from standard metal
cations.
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woelen
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The copper(I) diammine complex is very easily oxidized. It is impossible to have a colorless solution of this without a strong reductor, also
dissolved.
I made the colorless solution from a copper(II) salt, dissolved in ammonia to which also some sodium dithionite is added. The latter is the reductor.
At the surface, such solutions turn blue, but when you shake it, it becomes colorless again.
http://woelen.homescience.net/science/chem/solutions/cu.html
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Eddygp
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| Quote: Originally posted by woelen | The copper(I) diammine complex is very easily oxidized. It is impossible to have a colorless solution of this without a strong reductor, also
dissolved.
I made the colorless solution from a copper(II) salt, dissolved in ammonia to which also some sodium dithionite is added. The latter is the reductor.
At the surface, such solutions turn blue, but when you shake it, it becomes colorless again.
http://woelen.homescience.net/science/chem/solutions/cu.html |
Hmm what a shame. Can it be reduced to copper metal? Or a copper(0) compound!?
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woelen
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It certainly can be reduced to copper metal, but this requires a strong reductor. Hydrazine is a suitable reductor, borohydride also does the job.
With hydrazine you can make copper mirrors from such a colorless copper(I)-solution.
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vmelkon
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What is the full equation for this reaction?
I am coming up with
3 CuSO4 + 4 Na2S2O5 → Cu2SO3.CuSO3 + 5 SO2 + 4 Na2SO4
What happens to the S2O5?
Looks like it releases 2 electrons.
4 S2O5(2-) → 2 SO3(2-) + 5 SO2 + SO4(2-) + 2 e-
The 2 electrons are captured by the 2 copper atoms
3 Cu(2+) → 2 Cu+ + Cu(2+)
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vmelkon
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| Quote: Originally posted by woelen | A good method of making Cu2O is to use Fehling's reagent (use Google for more info on this) and add this to a solution of glucose. All chemicals
needed for this are very easy to obtain.
Another method is dissolving copper sulfate in a solution of citric acid and excess sodium hydroxide or sodium carbonate. The citric acid forms a
complex with the copper(II) ions and this complex reacts with glucose at high pH and in this process Cu2O is formed as red precipitate, use excess
glucose. Allow standing in a warm place for a day or so (a good way is to take a bucket of hot water and place the vessel with the solution of the
copper/citrate complex and glucose in this. |
Here is another one:
From
https://www.youtube.com/watch?v=kop1sWzTK-I
The chemist makes a sodium potassium tartrate. Heats it to 50 °C. Adds 3% H2O2 and adds 1 mL of CuSO4 solution. Stirring is used and the solution
bubbles oxygen and CO2.
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vmelkon
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Hi, it's me again.
I did some tests with my chevreuil's salt.
1. I added a ammonia solution --- yes, you get that dark blue color from the aminocopper complex.
2. I added a dilute solution of HCl to another test tube with some chevreuil's salt. A white solid is visible.
I added the dark blue solution (#1) to the #2 and a brownish or redish precipitate was formed. Not sure what this was. Maybe Cu2O. Perhaps
the Cu2+ was reduced by the presence of SO32-.
Shaking it more and it disappears. Did it forms CuCl and it dissolved in the water?
Video:
https://youtu.be/1aFpGjW-hG0
[Edited on 28-6-2015 by vmelkon]
Signature ==== Is this my youtube page? https://www.youtube.com/watch?v=tA5PYtul5aU
We must attach the electrodes of knowledge to the nipples of ignorance and give a few good jolts.
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DalisAndy
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So is this salt similar to Iron(IV) oxide? In its bonding behavior
Elements Collected: 19/81 (Excluding all radioactive, using placecard for those)
Any tips or good sources are welcome.
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Texium
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I think you mean iron(II,III) oxide
(magnetite), there isn't an iron(IV) oxide to my knowledge. But yes, it is also a mixed valance compound: copper(I,II) sulfite would be the systematic
name for it.
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DalisAndy
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Chevreul's Salt synthesis
I most recently made a small batch of Chevreul's Salt, following the procedure on mrhomescientist's blog. I noticed when I had been heating the
mixture for a bit, like 1 minute maybe. I noticed a naturally strong smell, very very strong and acidic. I also accumulated a yellow solid on the
upper part of my beaker, where there was no fluid. Is this normal? or did i do something wrong. Also any tips on drying the stuff? When I tried to
heat it to drive off water, it made a large pop sound.
Elements Collected: 19/81 (Excluding all radioactive, using placecard for those)
Any tips or good sources are welcome.
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LargeV
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The smell is sulfur dioxide and I first thought the yellow was sulfur, but I don't think sulfur explodes?
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DalisAndy
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| Quote: Originally posted by LargeV | | The smell is sulfur dioxide and I first thought the yellow was sulfur, but I don't think sulfur explodes? |
The Pop was after I had separated out the liquid and transferred the salt to a test tube for storage. I heated the slat in the tube
Elements Collected: 19/81 (Excluding all radioactive, using placecard for those)
Any tips or good sources are welcome.
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MrHomeScientist
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Trapped water boiling off, perhaps? I never encountered any yellow solids. Is it crystalline or amorphous (sandy or gelatinous)? The latter would
point to sulfur, though I don't know how it would form.
The sharp acidic smell is indeed sulfur dioxide. Sodium metabisulfite releases it when dissolved in water. I didn't make a point of that in my post
but perhaps I should have - it's not too friendly to the lungs! I'm glad the experiment worked for you.
[Edited on 12-1-2015 by MrHomeScientist]
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DalisAndy
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| Quote: Originally posted by MrHomeScientist | Trapped water boiling off, perhaps? I never encountered any yellow solids. Is it crystalline or amorphous (sandy or gelatinous)? The latter would
point to sulfur, though I don't know how it would form.
The sharp acidic smell is indeed sulfur dioxide. Sodium metabisulfite releases it when dissolved in water. I didn't make a point of that in my post
but perhaps I should have - it's not too friendly to the lungs! I'm glad the experiment worked for you
[Edited on 12-1-2015 by MrHomeScientist] |
Thank you and yes it was very sand like. But the test tube had be heating for less than a few seconds before it popped. I do know I was not using
distilled water. I was getting impatient so I used a faucet near by. I know that it's has a high level of phosphates in it, enough that it can
sustain algae grow. My father has a fish tank and had to stop using the tap water for that reason
Elements Collected: 19/81 (Excluding all radioactive, using placecard for those)
Any tips or good sources are welcome.
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Texium
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Threads Merged
22-12-2015 at 08:09 |
Chemist_Cup_Noodles
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I think it would be beneficial to perform a recrystallization of this fascinating salt. I have already thrown some pretty good polar solvents at it,
and I wasn't very hopeful they would work at all but it was all that I have. I'm sure maybe toluene or xylene would do the trick. I haven't seen any
data at all on this salt's solubility so maybe it's up to us to generate it.
I'll be honest-- We're throwing science at the wall here to see what sticks. No idea what it'll do.
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