| Pages:
1
2
3 |
MR AZIDE
Hazard to Self
Posts: 64
Registered: 21-5-2012
Location: UNITED KINGDOM
Member Is Offline
Mood: Fizzing
|
|
Chevreul's Salt.
Found this interesting article in another old School Science review Sep 1986.
boiling solutions of Sodium metabisulphite Na2S2O5 ( sodium disulphate iv) added to boiling CuSO4 soln ,
and boiled for 4 mins gives Chevreuls salt, Cu3(SO3)2.2H2O) This is a red solid , which evolves a little
Sulphur Dioxide.
Copper (I) chloride will be obtained by dissolving chevreuls salt in dilute HCL, with some SO2 present.
3S2O52- + 3Cu2+ + %H2O ----->
Cu3(SO3)2.2H2O + 3So2 + 2H3O+ + SO42-
Says dissolve 6g of the metabisulphite in 30ml hot water. Dissolve 7g copper sulphate in another beaker in 100 ml water, asnd heat this to boiling.
With stirring as the metabisulphite is added, and boil for about 4 mins.
The precipitate if filtered off, and washed with several portions water.
This is the chevreuls salt.
Dssolving this in some dilute HCL will provide a white precip, within a greeny / blue solution. This white precip turns green slowly of exposure to
air . The CuCl.
Addition of excess HCL redissolved the precip.
Addition of ammonia gives the intense blue copper tetraammine (ii) complex.
[Edited on 13-6-2012 by MR AZIDE]
|
|
|
woelen
Super Administrator
Posts: 8027
Registered: 20-8-2005
Location: Netherlands
Member Is Offline
Mood: interested
|
|
Do you have this article online? This Chevreul's salt is quite special, it must have both copper(I) and copper(II) in its lattice. Cu3(SO3)2, means
that copper has charge +4 divided over 3 atoms. This means that two of the coppers must have oxidation state +1 and the last one has oxidation state
+2.
I never heard of it. A nice one to try at home 
|
|
|
turd
National Hazard
Posts: 800
Registered: 5-3-2006
Member Is Offline
Mood: No Mood
|
|
For crystal structure see Acta. Chem. Scand., 19, 2189-2199 (1965) (Free access)
|
|
|
reckless explosive
Harmless
Posts: 28
Registered: 1-12-2011
Member Is Offline
Mood: experimental
|
|
just did this reaction today went very smoothly and easily but since theres no info about it that i can find is there anything it is soluble in that
doesn't decompose it? because i thought some nice large crystals of this substance would be pretty beautiful looking. 
|
|
|
ScienceSquirrel
International Hazard
Posts: 1863
Registered: 18-6-2008
Location: Brittany
Member Is Offline
Mood: Dogs are pets but cats are little furry humans with four feet and self determination! 
|
|
Some interesting information on Chevreul's Salt;
http://www.scielo.br/scielo.php?script=sci_arttext&pid=S...
It looks like you can make it as red crystals, woohoo!
|
|
|
MR AZIDE
Hazard to Self
Posts: 64
Registered: 21-5-2012
Location: UNITED KINGDOM
Member Is Offline
Mood: Fizzing
|
|
I havent got this online, i dont have a image scanner, .
Further info on reaction with Dilute HCL:::
Cu3(SO4).2H2O + 4H3O+ + 2Cl- -----> 2CuCl + *H2O =
2SO2 + Cu2+
It says that the formation of Cu(II) ions favours the left side of eqilibrium.
2Cu+ <----> Cu2+ + Cu
So It can be seen that disproportionation happens.
I thought that the compound would be a precipitate, I havent done the reacion yet, I have got a small amount of the metabisulphite. So I might try
this.
Usefull way to make CuCl....
Interesting to hear that this forms red crystals..........any pictures,
monoclinic, orthorombic etc... etc.......
|
|
|
turd
National Hazard
Posts: 800
Registered: 5-3-2006
Member Is Offline
Mood: No Mood
|
|
Too lazy to read the publication I posted above?
|
|
|
AndersHoveland
Hazard to Other Members, due to repeated speculation and posting of untested highly dangerous procedures!
Posts: 1986
Registered: 2-3-2011
Member Is Offline
Mood: No Mood
|
|
So it seems that Cu+2 can oxidize sulfite to sulfate, if the reaction is driven by the formation of the Chevreul's Salt, which contains a portion of
its copper in the +1 oxidation state. These types of reactions are not extremely unusual, for example, Cu+2 can oxidize iodide to free iodine, the
reaction being driven by the formation of insoluble CuI.
Apparently Cu+1 can form strong complexes with iodide or certain sulfur-containing anions.
|
|
|
blogfast25
International Hazard
Posts: 10562
Registered: 3-2-2008
Location: Neverland
Member Is Offline
Mood: No Mood
|
|
Hmmm, a bit long winded, if you ask me.
Boil copper (II) sulphate solution with sodium chloride and copper wire:
Cu (II) + Cu (0) < === > 2 Cu+
Cu+ (aq) + Cl- (aq) === > CuCl (s)
White CuCl drops out.
[Edited on 15-6-2012 by blogfast25]
|
|
|
woelen
Super Administrator
Posts: 8027
Registered: 20-8-2005
Location: Netherlands
Member Is Offline
Mood: interested
|
|
I have experimented with copper sulfate and sulfite and bisulfite and I indeed can reproduce the results mentioned in this thread. I found, however,
that the reaction strongly depends on the actual pH of the solutions.
When a solution of Na2S2O5 is mixed with a solution of CuSO4, then a nice grass green solution is obtained. When this solution is heated to boidling,
then a brick red fine crystalline precipitate is formed and the liquid above the precipitate is (nearly) colorless.
I also made another solution of CuSO4 and excess of Na2S2O5 and put this aside at room temperature. After several hours, the bottom of the test tube
is covered by many nearly black crystals (very dark red) and the liquid has become lighter green.
With sodium sulfite (Na2SO3) the behavior is very different. This produces a dark green precipitate, which quickly turns dirty brown (mustard-like).
Probably first a basic copper(II)sulfite is formed, which then is converted to impure copper(I) oxide and free sulfate ion. This precipitate is slimy
and becomes flocculent on shaking, while the brick red Chevreul's salt is crystalline. The latter has a much nicer appearance and is much more compact
and hence can easily be separated.
I intend to do a more thorough investigation of this salt and the dependence on pH and there might follow a webpage about this interesting salt.
[Edited on 18-6-12 by woelen]
|
|
|
kavu
Hazard to Others
Posts: 207
Registered: 11-9-2011
Location: Scandinavia
Member Is Offline
Mood: To understand is to synthesize
|
|
I repeated the vague procedure available from Spectroscopic and magnetic properties of Chevreul’s salt, a mixed valence copper sulfite
Cu3(SO3)2·2H2O in Inorganica Chimica Acta Volume 295, Issue 1, 15 November 1999, Pages 125–127. It states that "Chevreul’s salt was synthesized
by heating a solution containing CuSO4 and NaHSO3 at a temperature between 60 and 70°C for 3 h. The concentration of CuSO4 was altered between 0.6
and 0.8 M, and the NaHSO3/CuSO4 mole ratio between 1.1 and 1.7." Unfortunately I don't have access to the original reference, which is the famous
Gmelin’s Handbook, Aufl. System No. 60, vol. 8, Verlag Chemie, Weinheim/Bergstrasse, 1958, p. 484.
In 1,1 ratio and 0,6 M solution in 10 ml scale the preparation worked like charm. A quick boil led to a deep red ppt. Heating the dried product led to
formation of black copper(II)oxide and volatile sulfur compounds (mostly SO2, white fumes can be seen on top of the solid) some of the sulfite also
turns to sulfate, though.
[Edited on 18-6-2012 by kavu]
[Edited on 18-6-2012 by kavu]
|
|
|
woelen
Super Administrator
Posts: 8027
Registered: 20-8-2005
Location: Netherlands
Member Is Offline
Mood: interested
|
|
I made a webpage about this interesting compound, as part of a somewhat more extensive set of experiments in which I investigated the redox reaction
between sulfite and copper(II) as function of pH. Quite a few interesting things were found. All of it is described in the following webpage:
http://woelen.homescience.net/science/chem/exps/copper_sulfi...
|
|
|
Poppy
Hazard to Others
Posts: 294
Registered: 3-11-2011
Member Is Offline
Mood: † chemical zombie
|
|
Bravo!
What about growing larger crystals now huh? Checking their opacity to UV and IR rays maybe 
[Edited on 6-25-2012 by Poppy]
|
|
|
MrHomeScientist
International Hazard
Posts: 1806
Registered: 24-10-2010
Location: Flerovium
Member Is Offline
Mood: No Mood
|
|
I just performed this synthesis tonight, and it went quite smoothly. Very simple, and the range of color changes is striking. Copper chemistry is just
amazing!
I simply dissolved a small spatula of sodium metabisulfite into a few mL of water in a test tube, and did the same with copper sulfate in another test
tube. This resulted in a clear solution, and a blue one. I combined the two to yield a beautiful emerald green liquid. Upon heating to boiling, the
red chevruel's salt precipitated out quickly. It's a nice, compact, crystalline precipitate that settles completely within seconds. My remaining
solution was still a light blue-green, indicating I had an excess of copper sulfate.
Now, I'd like to distinguish this from copper(I) oxide, which has a similar color can also be formed by addition of SO2 to solutions of
Cu2+ (according to the wiki article on cuprous oxide). This is to preempt any internet nay-sayers for when I post a video on my YouTube
page
I first took a few hundred mg of the salt and added concentrated ammonia solution, yielding the deep blue hexaaminecopper(II) complex. (proving that
the compound contains copper(II) and really isn't copper(I) oxide.
I wanted to check myself by testing copper(I) with ammonia, so following the suggestion in the OP I added a similar amount of the salt to diluted
hydrochloric acid. It formed a light green solution, and upon adding a bit more chevruel's salt I obtained a white precipitate of copper(I) chloride.
I pipetted off the liquid, added a small amount of water to wash the precipitate, pipetted off the water, and then added some concentrated ammonia. I
got the same deep blue solution! I believe this should have yielded a colorless copper(I) complex. Would slight contamination of copper(II) be enough
to give the same deep blue? The copper(I) complex is also claimed to be air-sensitive - would swirling the test tube to dissolve the CuCl be enough to
convert fully to the blue complex?
A very interesting and fun experiment overall.
|
|
|
deltaH
Dangerous source of unreferenced speculation
Posts: 1663
Registered: 30-9-2013
Location: South Africa
Member Is Offline
Mood: Heavily protonated
|
|
Very interesting inorganic chemistry, but I think one has to be careful that one indeed produces the mixed valent complex and not simply copper metal
powder precipitate (complete reduction). This particularly a concern as you're heating a reducing agent with an oxidant, it can be easy to overshoot
your mixed valent state and reduce fully to copper powder?
This red powder I see in some of the photos looks to me like copper powder? If I were a betting man, I would be placing my money on the @ woelen's
nearly black crystals that form on standing as being the actual complex. After all Cu(I) Cu (II) charge transfer should make it VERY strongly deeply
coloured.
|
|
|
woelen
Super Administrator
Posts: 8027
Registered: 20-8-2005
Location: Netherlands
Member Is Offline
Mood: interested
|
|
@MrHomeScientist: The colorless complex of copper(I) with ammonia is VERY air sensitive. The complex is colorless, but even a tiny amount of oxygen
from the air makes the solution deep blue. Simply dissolving Cu2O or CuCl in ammonia does not give a colorless solution. Any air, dissolve in the
ammonia or air around the ammonia causes the solution to turn blue. If you want to make a colorless solution of copper(I), then you need to have a
strong reductor in solution as well, e.g. Na2S2O4 (dithionite).
Testing the Chevreuls salt can be done by taking this, together with a very small pinch of sodium dithionite and dissolving this in ammonia. You then
get a blue solution. If you have a similar amount of Cu2O with the same small pinch of Na2S2O4, then the solution will be colorless and only turns
blue after somewhat longer shaking with fresh air around.
Another test may be to add the Chevreuls salt to 10% HCl. That should give a clearly present smell of SO2. With Cu2O you won't get any smell.
@deltaH: I tested my red precipitate by adding acid to it. It dissolves. Fine copper powder does not dissolve. When the precipitate was added to a
small amount of acid, then I could smell SO2.
|
|
|
deltaH
Dangerous source of unreferenced speculation
Posts: 1663
Registered: 30-9-2013
Location: South Africa
Member Is Offline
Mood: Heavily protonated
|
|
@woelen
Ok, you've convinced me it's not copper, but I'm not so convinced it's not Cu2O passivated with adsorbed surface sulfite that gives off a smell of SO2
with acid and causes the Cu2O to dissolve. This could be especially true in the case if the Cu2O particles are very small or porous, possibly even
nanosized aggregates capped off by surface sulfite species?
[Edited on 7-10-2013 by deltaH]
|
|
|
woelen
Super Administrator
Posts: 8027
Registered: 20-8-2005
Location: Netherlands
Member Is Offline
Mood: interested
|
|
Yes, I understand that sulfite can be coprecipitated with Cu2O. So, the test with the smell of SO2 is not 100% sure.
If you want to do a test which is fully conclusive, then most likely you need to do a quantitative test. If you look at the formula for Chevreul's
salt, then you see that there are two Cu(+) entities and two SO3(2-) entities. Fully oxidizing this to copper(II) and sulfate requires 6 electrons. If
you have Cu2O, then you need 6 electrons for 3 Cu2O.
The problem with such quantitative tests is that you need to work very accurately. One mole of Chevreul's salt is 387 grams, 3 moles of cuprous oxide
is 429 grams. This is fairly close. Impurities in either one of these, especially if precipitated from solution can easily be several percents (e.g.
humidity, coprecipitated sodium ions, coprecipitated sulfate ions) and then the difference between the two may be swamped in inaccuracy.
Do you see a practical method for distinguishing between the two? If it is within the reach of me as a amateur chemist, then I certainly am willing to
try it.
|
|
|
deltaH
Dangerous source of unreferenced speculation
Posts: 1663
Registered: 30-9-2013
Location: South Africa
Member Is Offline
Mood: Heavily protonated
|
|
@woelen
I hear you and very interesting challenge indeed... NOT easy!!!
Okay, I read that Cu2O is a semiconductor, so maybe that might open some avenues for testing? FYI Band gap = 2.17eV from Wiki's semiconductor data
table, so that means it should absorb STRONGLY all light with wavelength shorter than 571nm... basically all green light and higher, but importantly,
it should only scatter light below this. So red light is only scattered by Cu2O for example, not absorbed.
Now as for Chevreul's salt, that has broad adsorptions all over the place, see this paper on the electronic spectra of it: http://www.scielo.br/pdf/jbchs/v13n5/12815.pdf
Note that there is a broad absorbtion peak at 785nm (deep red) for Chevreul's salt, it's not the strongest, but it's there.
So I know this is not ideal, but maybe if you have can get a bright deep red LED and shine on suspect Cu2O and suspect Chevreul and compare the
difference on scattered light in a darkened room?
Would be very interesting if indeed the one sample Cu2O scattered the RED light but not so much absorb it, while the Chevreul's salt absorbs it more
strongly appearing darker.
What about then using ultrabright green LED, now the Cu2O should be STRONGLY absorbing, appearing black, compared to Chevreul's salt which should only
be partially absorbing green light.
Then ultrabright blue LED, BOTH Chevreul's salt AND Cu2O should be absorbing that STRONGLY according to the info above.
Photo's PLEASE if this works out  should look amazing!
[Edited on 7-10-2013 by deltaH]
|
|
|
MrHomeScientist
International Hazard
Posts: 1806
Registered: 24-10-2010
Location: Flerovium
Member Is Offline
Mood: No Mood
|
|
Woelen:
Thanks for the response. Unfortunately I don't have any dithionite, so I'm looking into alternatives. I think I'll try adding some of the salt to HCl
to produce SO2 gas, which I will then draw into a syringe and bubble through an acidified solution of potassium dichromate. If
SO2 is present, it reduces the solution and turns it from orange to green. This is the standard, specific test for sulfur dioxide. For
videos, I like to be able to show visual evidence whenever possible
deltaH:
It's fairly obviously not copper powder, when you see it in person. Copper powder is more of a brownish color, whereas this salt is a deep red just
like dark bricks. But as woelen said (and as I wrote in my original post), it is soluble in acid (and ammonia) which shows that it is indeed not
copper.
If the idea of using colored LEDs is plausible, I conveniently have access to some very bright, 1 amp (3 watt) red, green, and blue LEDs. Minor
obstacle is that I do not have any CuO to compare against, so I'd have to make it. I see there are multiple ways to make this compound - anyone have
experience with this? Any routes more recommended than others? This sentence from wikipedia seems like the simplest to me: "Copper(I) oxide can also
be prepared by reacting a copper-ammonia complex with hydrogen peroxide."
|
|
|
deltaH
Dangerous source of unreferenced speculation
Posts: 1663
Registered: 30-9-2013
Location: South Africa
Member Is Offline
Mood: Heavily protonated
|
|
@MrHomeScientrist
Yeah agreed, some 'standard' to compare to would be the way to go. Hope it gives nice results and pretty pictures!
Good luck
p.s. I'm jealous of you ultra brights and in all the colours too!
[Edited on 7-10-2013 by deltaH]
|
|
|
woelen
Super Administrator
Posts: 8027
Registered: 20-8-2005
Location: Netherlands
Member Is Offline
Mood: interested
|
|
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.
|
|
|
kmno4
International Hazard
Posts: 1503
Registered: 1-6-2005
Location: Silly, stupid country
Member Is Offline
Mood: No Mood
|
|
Some article about preparation of this salt (and not only).
Attachment: hhhh.7z (401kB)
This file has been downloaded 668 times
Слава Україні !
Героям слава !
|
|
|
deltaH
Dangerous source of unreferenced speculation
Posts: 1663
Registered: 30-9-2013
Location: South Africa
Member Is Offline
Mood: Heavily protonated
|
|
Thanks kmno4, I found this paragraph at the end of your last reference of particular importance:
"It appears from these results that the reduction of copper(II) sulfate solutions with sulfur dioxide at elevated temperatures yields Chevreul's salt,
copper metal and a mixture of anions containing sulfur, in the + IV and lower oxidation states. In contrast, the sulfur dioxide reduction of an
ammoniacal solution of copper(lI), i.e. a mixture of the ammine complexes of copper(IlL gives Chevreul's salt which is not contaminated with
copper metal or anionic species of sulfur in oxidation states less than + IV."
N.B. if correct!
|
|
|
MrHomeScientist
International Hazard
Posts: 1806
Registered: 24-10-2010
Location: Flerovium
Member Is Offline
Mood: No Mood
|
|
I completed a video and writeup of this experiment, which are now posted on their respective sites. If you google Chevreul's salt, my pages are the
second and third hits already! (this thread is the first)
Video: http://www.youtube.com/watch?v=CitNalVs01M
Writeup: http://thehomescientist.blogspot.com/2013/10/chevruels-salt....
I tested out deltaH's colored LED idea, and Chevreul's salt behaves just as predicted, which is pretty neat. Under red light it
appears red, of course. Under green light, it is indeed partially absorbing and it still appears reddish to me. Under blue light, is looks totally
black.
For doing a comparison, I tried making some CuO via a long chain of syntheses (cream of tartar -> Rochelle salt -> Fenton's reagent -> CuO),
but only got a tiny amount that won't scrape off my filter paper (it seems impregnated onto it). I didn't notice anything special about it under the
lights, but the paper could be interfering. I'll try making some more and this time pipette off the liquid instead of filtering, so I have a powder to
look at. Once I do that, I'll post photos of the results.
|
|
|
| Pages:
1
2
3 |
![[*]](images/xpblue/default_icon.gif){kind=icon}
