Sciencemadness Discussion Board
Not logged in [Login ]
Go To Bottom

Printable Version  
 Pages:  1  2  
Author: Subject: Chevreul's Salt.
deltaH
Dangerous source of unreferenced speculation
*****




Posts: 1663
Registered: 30-9-2013
Location: South Africa
Member Is Offline

Mood: Heavily protonated

[*] posted on 15-10-2013 at 08:37


Very nice indeed, very well done MrHomeScientist... pity no pics of the LED's :(



Mind your step or step your mind. Website: www.ideashack.org
View user's profile Visit user's homepage View All Posts By User
woelen
Super Administrator
*********




Posts: 7733
Registered: 20-8-2005
Location: Netherlands
Member Is Offline

Mood: interested

[*] posted on 15-10-2013 at 11:33


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.



The art of wondering makes life worth living...
Want to wonder? Look at https://woelen.homescience.net
View user's profile Visit user's homepage View All Posts By User
blogfast25
Thought-provoking Teacher
*****




Posts: 10340
Registered: 3-2-2008
Location: Old Blighty
Member Is Offline

Mood: No Mood

[*] posted on 15-10-2013 at 12:11


Nice write up and photos. Good demonstration of the dual oxidation state of the copper. Might have a bash at this myself...



View user's profile Visit user's homepage View All Posts By User
bfesser
Resident Wikipedian
*****




Posts: 2114
Registered: 29-1-2008
Member Is Offline

Mood: No Mood

[*] posted on 16-10-2013 at 07:46


Wonderful work, <strong>MrHomeScientist</strong>. I second the idea of taking some photos of the compound under monochromatic illumination.



View user's profile View All Posts By User
MrHomeScientist
International Hazard
*****




Posts: 1806
Registered: 24-10-2010
Location: Flerovium
Member Is Offline

Mood: No Mood

[*] posted on 17-10-2013 at 05:56


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]
View user's profile Visit user's homepage View All Posts By User
deltaH
Dangerous source of unreferenced speculation
*****




Posts: 1663
Registered: 30-9-2013
Location: South Africa
Member Is Offline

Mood: Heavily protonated

[*] posted on 17-10-2013 at 06:13


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!!!




Mind your step or step your mind. Website: www.ideashack.org
View user's profile Visit user's homepage View All Posts By User
Volanschemia
National Hazard
****




Posts: 340
Registered: 16-1-2015
Location: Victoria, Australia
Member Is Offline

Mood: Pretty much all of them!

[*] posted on 17-3-2015 at 22:44


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.
View user's profile View All Posts By User
Eddygp
International Hazard
*****




Posts: 858
Registered: 31-3-2012
Location: University of York, UK
Member Is Offline

Mood: Organometallic

[*] posted on 19-3-2015 at 04:56


How does this salt react with ammonia? Does it form a complex, like the reaction of ammonia and copper(II) sulfate?



there may be bugs in gfind

[ˌɛdidʒiˈpiː] IPA pronunciation for my Username
View user's profile View All Posts By User
Amos
International Hazard
*****




Posts: 1390
Registered: 25-3-2014
Location: Yes
Member Is Offline

Mood: No

[*] posted on 19-3-2015 at 05:47


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.




View user's profile View All Posts By User
woelen
Super Administrator
*********




Posts: 7733
Registered: 20-8-2005
Location: Netherlands
Member Is Offline

Mood: interested

[*] posted on 19-3-2015 at 11:03


The copper(I) also goes into solution, it forms a diammine complex [Cu(NH3)2](+). This complex is colorless.



The art of wondering makes life worth living...
Want to wonder? Look at https://woelen.homescience.net
View user's profile Visit user's homepage View All Posts By User
Eddygp
International Hazard
*****




Posts: 858
Registered: 31-3-2012
Location: University of York, UK
Member Is Offline

Mood: Organometallic

[*] posted on 19-3-2015 at 14:00


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.




there may be bugs in gfind

[ˌɛdidʒiˈpiː] IPA pronunciation for my Username
View user's profile View All Posts By User
woelen
Super Administrator
*********




Posts: 7733
Registered: 20-8-2005
Location: Netherlands
Member Is Offline

Mood: interested

[*] posted on 20-3-2015 at 00:14


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




The art of wondering makes life worth living...
Want to wonder? Look at https://woelen.homescience.net
View user's profile Visit user's homepage View All Posts By User
Eddygp
International Hazard
*****




Posts: 858
Registered: 31-3-2012
Location: University of York, UK
Member Is Offline

Mood: Organometallic

[*] posted on 20-3-2015 at 02:01


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!?




there may be bugs in gfind

[ˌɛdidʒiˈpiː] IPA pronunciation for my Username
View user's profile View All Posts By User
woelen
Super Administrator
*********




Posts: 7733
Registered: 20-8-2005
Location: Netherlands
Member Is Offline

Mood: interested

[*] posted on 20-3-2015 at 09:18


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.



The art of wondering makes life worth living...
Want to wonder? Look at https://woelen.homescience.net
View user's profile Visit user's homepage View All Posts By User
vmelkon
International Hazard
*****




Posts: 669
Registered: 25-11-2011
Location: Canada
Member Is Offline

Mood: autoerotic asphyxiation

[*] posted on 8-6-2015 at 18:32


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+)




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.
Yes my evolutionary friends. We are all homos here.
View user's profile View All Posts By User
vmelkon
International Hazard
*****




Posts: 669
Registered: 25-11-2011
Location: Canada
Member Is Offline

Mood: autoerotic asphyxiation

[*] posted on 19-6-2015 at 20:56


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.




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.
Yes my evolutionary friends. We are all homos here.
View user's profile View All Posts By User
vmelkon
International Hazard
*****




Posts: 669
Registered: 25-11-2011
Location: Canada
Member Is Offline

Mood: autoerotic asphyxiation

[*] posted on 28-6-2015 at 15:12


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.
Yes my evolutionary friends. We are all homos here.
View user's profile View All Posts By User
DalisAndy
Hazard to Others
***




Posts: 129
Registered: 8-5-2015
Member Is Offline

Mood: No Mood

[*] posted on 4-9-2015 at 07:25


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.
View user's profile View All Posts By User
Texium
Administrator
********




Posts: 3844
Registered: 11-1-2014
Location: Salt Lake City
Member Is Offline

Mood: Triturated

[*] posted on 4-9-2015 at 08:20


Quote: Originally posted by DalisAndy  
So is this salt similar to Iron(IV) oxide? In its bonding behavior
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.



Come check out the Official Sciencemadness Wiki
They're not really active right now, but here's my YouTube channel and my blog.
View user's profile Visit user's homepage View All Posts By User
DalisAndy
Hazard to Others
***




Posts: 129
Registered: 8-5-2015
Member Is Offline

Mood: No Mood

[*] posted on 30-11-2015 at 15:53
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.
View user's profile View All Posts By User
LargeV
Harmless
*




Posts: 12
Registered: 28-11-2015
Location: Connecticut
Member Is Offline

Mood: No Mood

[*] posted on 30-11-2015 at 16:43


The smell is sulfur dioxide and I first thought the yellow was sulfur, but I don't think sulfur explodes?
View user's profile View All Posts By User
DalisAndy
Hazard to Others
***




Posts: 129
Registered: 8-5-2015
Member Is Offline

Mood: No Mood

[*] posted on 30-11-2015 at 17:11


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.
View user's profile View All Posts By User
MrHomeScientist
International Hazard
*****




Posts: 1806
Registered: 24-10-2010
Location: Flerovium
Member Is Offline

Mood: No Mood

[*] posted on 1-12-2015 at 06:43


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]
View user's profile Visit user's homepage View All Posts By User
DalisAndy
Hazard to Others
***




Posts: 129
Registered: 8-5-2015
Member Is Offline

Mood: No Mood

[*] posted on 1-12-2015 at 08:10


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.
View user's profile View All Posts By User
Texium
Administrator
Threads Merged
22-12-2015 at 08:09
Chemist_Cup_Noodles
Harmless
*




Posts: 46
Registered: 27-5-2015
Location: Northern VA
Member Is Offline

Mood: Anticipative

[*] posted on 14-4-2016 at 04:50


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.
-Cave Johnson, Portal 2
Add yourself to this map of SciMad members! https://zeemaps.com/map?group=388676&add=1
View user's profile Visit user's homepage View All Posts By User
 Pages:  1  2  

  Go To Top