Sciencemadness Discussion Board

Chevreul's Salt.

MR AZIDE - 13-6-2012 at 10:56


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 - 13-6-2012 at 22:20

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 :P

turd - 14-6-2012 at 02:59

For crystal structure see Acta. Chem. Scand., 19, 2189-2199 (1965) (Free access)

reckless explosive - 14-6-2012 at 07:16

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 - 14-6-2012 at 07:52

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 - 14-6-2012 at 11:03

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 - 15-6-2012 at 01:59

Quote: Originally posted by MR AZIDE  

monoclinic, orthorombic etc... etc.......

Too lazy to read the publication I posted above?

AndersHoveland - 15-6-2012 at 02:09

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 - 15-6-2012 at 05:05

Quote: Originally posted by MR AZIDE  

Usefull way to make CuCl....



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 - 17-6-2012 at 23:08

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 - 18-6-2012 at 04:06

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 - 24-6-2012 at 10:10

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 - 24-6-2012 at 20:59

Bravo!
What about growing larger crystals now huh? Checking their opacity to UV and IR rays maybe :o

[Edited on 6-25-2012 by Poppy]

MrHomeScientist - 6-10-2013 at 17:32

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 - 6-10-2013 at 22:27

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 - 6-10-2013 at 23:22

@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 - 7-10-2013 at 01:12

@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 - 7-10-2013 at 03:38

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 - 7-10-2013 at 05:06

@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 :D should look amazing!

[Edited on 7-10-2013 by deltaH]

MrHomeScientist - 7-10-2013 at 06:19

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 - 7-10-2013 at 06:58

@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 - 7-10-2013 at 12:23

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 - 7-10-2013 at 23:01

Some article about preparation of this salt (and not only).

Attachment: hhhh.7z (401kB)
This file has been downloaded 668 times

deltaH - 8-10-2013 at 01:30

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 - 15-10-2013 at 06:23

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.

deltaH - 15-10-2013 at 08:37

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

woelen - 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.

blogfast25 - 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...

bfesser - 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.

MrHomeScientist - 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]

deltaH - 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!!!

Volanschemia - 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.

Eddygp - 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?

Amos - 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.

woelen - 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.

Eddygp - 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.

woelen - 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

Eddygp - 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!?

woelen - 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.

vmelkon - 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+)

vmelkon - 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.

vmelkon - 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]

DalisAndy - 4-9-2015 at 07:25

So is this salt similar to Iron(IV) oxide? In its bonding behavior

Texium - 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.

Chevreul's Salt synthesis

DalisAndy - 30-11-2015 at 15:53

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.

LargeV - 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?

DalisAndy - 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

MrHomeScientist - 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]

DalisAndy - 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

Chemist_Cup_Noodles - 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.

DraconicAcid - 14-4-2016 at 08:21

Quote: Originally posted by Chemist_Cup_Noodles  
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.


It's a salt, and you think it will dissolve in toluene? Good luck.

Acetonitrile *might* work, but you may not get the original salt back.

crystal grower - 14-4-2016 at 08:43

http://chem.pieceofscience.com/?p=988
This is my article about chevruls salt, if I would ever succesfully recrystalize it, I'll add it to the article.

The Volatile Chemist - 15-4-2016 at 15:20

Just curious, has DMSO ever been tried? Sorry, I don't know my solvents too well...

Fidelmios - 11-2-2017 at 14:24

Just to add on a little bit here.. I recently preformed the experiment outlined by Mrhomescientist (http://thehomescientist.blogspot.com/2013/10/chevruels-salt....) but did so with 3x the reagents. After mixing the two solutions I decided to heat a large amount of it, and finish the experiment outlined above, but a small amount of it I decided to try and produce the salt with a slow evaporation. (I also wanted to see if it affected the crystal size)




While the slow evaporation DID work, it took quite a long time, and many impurities remained in the evaporation dish (most of which were unreacted reagents). The crystal size didn't increase any significant amount.

MrHomeScientist - 13-2-2017 at 07:02

Interesting! Glad the experiment worked for you. Woelen was able to make slightly larger crystals on his site: http://woelen.homescience.net/science/chem/exps/copper_sulfi...