Sciencemadness Discussion Board

Precipitated copper powder

Mephisto - 8-10-2004 at 02:01

In many organic synthesis 'British Drug Houses' precipitated copper powder is used as catalyst. I always thought, that's common copper powder, which can be made from copper sulphate solution and zinc.

But now I saw a direct contest between this freshly prepared copper powder from CuSO4 + Zn (=A) and the 'British Drug Houses' precipitated copper powder (=B). After several runs A gives a maximum yield of 83% of a desired organic compound. Under similar conditions B gives a yield of 98.9%.

So what's the mystery of the 'British Drug Houses' precipitated copper powder? 'A' is already very fine and has a big surface, so this point can't be the reason for B's superiority. Has B some additives? Why is it soo much better??

Since I can't find any useful information neither in Roempp nor in Ullmann's and even not in the CA: Has anyone information about 'British Drug Houses' precipitated copper powder?

mick - 8-10-2004 at 11:29

As far as I am aware the British Drug House became BDH, which became part of Merk, which is now part of VWR.
I think the the stuff you are talking about is because they have been making it for years and make it better than anyone else.

The spec. on BDH stuff was fairly high and their containers were very good.
edit mick

[Edited on 8-10-2004 by mick]

Self-made Copper powder (fine precipitation)

chemoleo - 10-10-2004 at 08:45

You can of course make copper powder yourself, by reducing it straight from solution. Not Zn powder, as the resulting Cu powder is at least the size of the Zn powder. Whereas precipitation straight from a copper II+ solution should yield nanosized Cu particles.

Here was my attempt at this, I boilded a CuSO4 solution (10g) and added 5 g of Vit C (ascorbic acid). The solution turned dark green (from blue), and soon after, copper powder was precipitated.

This is several grams worth. I tried to upscale this, but no success - I am not quite sure how the stoichiometry of the reaction would work out. Also, gas is devoloped during the course of the reaction.
The copper powder is so fine it coated a teflon-coated stirbar with it.

Anyway, this is something you may want to try.

PS what sort of reactions would you use the Cu powder for? Examples?

PS it is definitely Cu, as it is totally unreactive to HCl (Cu2O reacts with it), plus it is very heavy (hence easily washed and isolated).

[Edited on 10-10-2004 by chemoleo]

Quantum - 10-10-2004 at 09:39

Interesting Chemleo!

Could acetic acid be used instead?

[Edited on 10-10-2004 by Quantum]

garage chemist - 10-10-2004 at 09:50

Chemoleo, are you sure that this is actually metallic copper?

I have a feeling that this is could be cuprous oxide (Cu2O). It has exactly the same colour as copper and can be mistaken as copper powder very easily.

Very fine copper powder can be produced by heating black copper oxide in a stream of hydrogen. The resulting powder is so fine that it produces a perfect shiny metallic surface on pressing it with the blade of a knife.

The black CuO is made from CuSO4 and NaOH and boiling for some time. (combine the CuSO4 and NaOH as two solutions, NOT as solids! Otherwise the Cu(OH)2 that forms first coats the particles and prevents further reaction)

[Edited on 10-10-2004 by garage chemist]

Mephisto - 10-10-2004 at 10:00

chemoleo: Thanks, I will try out your reaction, since it's a nice OTC-able one. It's good, that you have made a test for Cu2O, because your reaction remembers to the Fehling's reaction...

The Cu powder should serve as a catalyst in the reaction of 5-iodovanillin with sodium methoxide to syringaldehyde (I've got only a theoretical interest in this way). For more information see Can. J. Chem. 31, 476-483 (1953). You can find the contest of both Cu powders (A and B) on page 5.

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[Edited on 7/22/13 by bfesser]

chemoleo - 10-10-2004 at 10:16

Garage chemist, yes it is copper.
I have played with commerical Cu2O before, and it is 1) a darker colour than Cu and more importantly 2) soluble in HCl, forming first insoluble WHITE CuCl, which then is oxidised by air to green CuCl2.
There is definitely not even a hint of a reaction.

But let me do this test with conc HCl again, I used only a 5 molar solution.
Anyway, the heaviness is also indicative of Cu, as it really sinks immediately. QUite unlike Cu2O, which forms more of a suspension and takes a few minutes to settle.

PS Quantum, no you cant use HAc. You have to use something that is strongly reducing, and ascorbic acid is strongly reducing.

Edit: you can get fine and shiny copper powder from arts supplies, too, as it is used as a filler in polyester resins to produce the coppery look! i.e. tiranti for starters.

[Edited on 10-10-2004 by chemoleo]

garage chemist - 10-10-2004 at 13:42

OK, when it doesn't react with HCl then it's surely copper.

You may want to thoroughly wash and dry the powder and then test the electrical conductivity. This would be the final proof that it's copper.

Where did you hear about the reaction of CuSO4 with Ascorbic acid? This is very interesting!
I never heard of a "wet" reaction that produces metallic copper before, hence my thoughts about Cu2O.

I'd still produce copper powder from CuO and H2 though. The copper powder comes out perfectly dry and I hate drying wet precipitates.

chemoleo - 10-10-2004 at 14:06

Hmm, I had a look at it now, and the 5M HCl which contained the copper powder has become very slightly green. I guess that is to be expected - i.e. oxygen aids in dissolving small amounts. Washing this, and adding 35% HCl didnt change things really, except that a slightly green colouration was noticed, which didn't become more.
Compared to this, when I dissolved commercial Cu2O, it would dissolve immediately to form white CuCl.

What is more peculiar that my copper powder, the batch that was left to dry, turned brown coloured - not all, only what was under the surface. So now I have a powder consisting of dark grains together with the copper coloured grains. I suspect the fineness of the Cu was enough to be oxidised by air, in the presence of trace acid and water? A bit strange nonetheless. But then, who has played with nanosized Cu before? When I get back, I will measure whether any current is conducted. I wish I had done that yesterday, as now it contains this brown/black stuff (it's very hard to say, as the particles are so small). Actually I will take half of that batch and try to redissolve it in HCl, I should remove the oxide that way presumably..

Anyway, that method is something I tried way back, when I had restricted amounts of chemicals. I was hoping to make Cu2O then, to turn this into CuCl. But as you can see I never got it to work then.

Reducing CuO with H2? Hmm... and the hassle of heat, hydrogen gas and so on? I am not quite so sure about that! Definitely more hassle than just boiling a solution. But then... H2 reduction is certainly more economical and quantitative!

Edit/Followup: Ah ok, no problem about the dark brown powder. It seems only coated at the surface. When adding 5M HCl, it turned immediately back to the rosy copper colour :) - and the solution above turned ever so slightly green - indicating a very thin oxide layer! I guess that's to be expected with finely divided copper like that; even copper sheets tarnish after a while.

[Edited on 10-10-2004 by chemoleo]

chemoleo - 10-10-2004 at 17:08

Yes, it definitely conducts, meaning it is truly copper.
Because the powder so rapidly oxidises in air, I cleaned it with HCl, and then measured the relative conductencies of the supernatant (the liquid (water) containing no copper powder, and the precipitate (copper powder and water). The supernatant had about 500 kOhm, while the precipitate in the same liquid went down to 10-40 kOhm.
Not ultimately a proof, but taken together with the chemical & observational data, I now have no doubt it truly is copper precipitated straight from a Cu2+ solution.

Apart from that... I rubbed the powder onto a surface, and it actually became shiney, as one would expect from a metal powder. But most convincing to me, anyway, is still how quickly the powder settles. It doesn't take 2 seconds. Only a high density compound would behave that way.

Good bye, Cu2O :)

PS to keep the powder reduced, I'd suggest to dry it in a vaccum dessicator, or over nitrogen/Ar/CO2 etc, and to keep it dry and away from air at all times. To me, the fact that it oxidises so happily actually shows it has a decent surface area (i.e. being very finely divided).

The_Davster - 10-10-2004 at 17:12

After the copper powder was precipitated and filtered, what color was your filtrate? Mine is still bright green. What is this due to? Some copper ascorbate compound?
Obviously I just tried this but the ammount of copper powder I got seemed like under a gram. The ascorbic acid I was using was commercial vitamin C but they were by some hippie brand so I had to filter out some insoluabe particles. The solutions of copper (II) sulfate and ascorbic acid were combined while hot and were left for the copper powder to precipitate. I recieved no gas emmited during any stage of the reaction.

I am going to try boiling the green filtrate for a while, to see if any more copper will precipitate.

chemoleo - 10-10-2004 at 17:20

Yes it was still green, dark in my case (I guess I made sure to use an excess ;)). Copper ascorbate - I'm not sure. It would imply that ascorbic acid is stronger than H2SO4, which surely it is not. Yet the colour change seems to suggest that. It baffled me too.
As to the amount of copper - weigh it. It has a rel. high density, making you misjudge the amounts.
As to gas - it evolved slowly, while I continued to boil it. I can't really say whether this had an effect on the yield.
In fact, I was baffled by the yield alltogether, as I added more CuSO4/Vit C to the existing reaction, but no more Cu seemed to precipitate.
What might help is actually to construct a mechanism of the reaction.
Also, if you continue boiling it down, the colour will go very dark, and a distinct smell is produced - to my nose it smells like decomposing sugars.

I have a feeling that only a certain part of the reaction produces the copper powder, and thereafter, whatever you do, you won't get more. Maybe the product (copper, vitamin C's breakdown products) inhibiting the reaction. I just don't know. Looks like the subject for a PhD thesis :D

PS good to see someone else is doing some experimenting!

[Edited on 11-10-2004 by chemoleo]

The_Davster - 10-10-2004 at 18:53

Chemoleo, were you using lab grade ascorbic acid or OTC vit C? After disolving my vit C in water the aqueous layer was yellow so I am assuming some other impurities are in vit C. After filtering the hot dissolved VitC the residue left in the filter paper smelled strongly like citrus and unlike the vitC tablets prior to purification. Hmm, possibly because I was using citrus flavored vitC:D.

I attempted this reaction again moments ago and I appear to have gotten much more copper powder than before. Last time as I was filtering the citrus impurities out of the Vit C I just had the stem of the funnel leading into the copper sulfate solution. I assume this lead to some cooling of both reactants. However this time I allowed the filtrate from purifing the vit C to go into a 250 mL beaker. The solutions of copper sulfate and vit C were separatly heated to nearly boiling then mixed quickly and boiled again. I recieved much more copper powder than before.:cool:

Also, Chemoleo, in your pic of the copper powder, did it look like that upon precipitation and filtering? Or is that pic after you rinsed it with HCl? Mine is much darker after precipitation and filtering.

A more exact way for chemoleo's excellent method

Mephisto - 11-10-2004 at 00:15

There's already a method for making nano-sized copper powder with ascorbic acid in the Chemical Abstracts. The use of PVP is just necessary, if even a finer product is desired. The mol-ratio is the most usefull thing in the following text.


Copper nano-powder was synthesized with bluestone (CuSO4×5H2O) as starting material, Vitamin C (Vc) as reducer and polyvinylpyrrolidone as protectant and dispersant. After Vitamin C (Vc) was mixed with polyvinylpyrrolidone (PVP), the mix was added to soln. of bluestone. When the ratio of n(Vc) and n(CuSO4×5H2O) was 2.45:1, the particle size increased while the increase of reagent concn. In this way, the yield of copper powder was 95%. Structure of products was 20-40 nm confirmed by XRD and TEM.

*[Synthesis of nanometer-copper powder by reduction. Xiao, Han; Wang, Rui; Yu, Lei; Zou, Guitian. Department of Chemistry, Guizhou Nationalities College, Guiyang, Guizhou Province, Peop. Rep. China. Guizhou Shifan Daxue Xuebao, Ziran Kexueban (2003), 21(1), 4-6.]

[Edited on 11-10-2004 by Mephisto]

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[Edited on 7/22/13 by bfesser]

chemoleo - 11-10-2004 at 07:30

Wow, that's actually quite recent! Should have published it back in 92 or so when I first tried it - teenage production of Cu powder :D

Rogue - I used Vit C powder, 100%, straight from a drug store. It might indeed help to get rid of citrus smell :P
The copper powder you see on the picture is the water rinsed powder straight after the precipitation reaction- no HCl involved. I only used HCl to remove the brown coating after I allowed the Cu powder to dry on air. If yours is darker - I wouldnt know why. As long as it does look like copper, and sinks very rapidly to the bottom of the tube/beaker, you should be fine.

kevinlimse - 12-10-2004 at 01:39

Is there a method for Aluminium powder?

I am a fish - 12-10-2004 at 11:20

Originally posted by kevinlimse
Is there a method for Aluminium powder?

Elemental aluminium is too reactive to be formed in water.

However, it may be possible to form powders of other unreactive metals, such as bismuth, lead and tin.

Oxydro - 12-10-2004 at 13:00

Now what would be impressive would be someone figuring out how to easily make aluminum powder by a similar reaction, but in a solvent other than water, that aluminum will not react with enough to spoil the reaction. Have to be careful drying the precipitate, that fine aluminum would be pyrophoric, I imagine?

guy - 13-10-2004 at 17:20

So this reaction
CuSO4 + Ascorbic acid, does the remaining sulfate combine with the H+ to form sulfuric acid?

chemoleo - 13-10-2004 at 18:26

Possibly it would. It would be interesting to follow the pH during the course of the reaction.
Possibly the pH will stop the reaction lateron. I think I mentioned above that I had the feeling that not much more precipitation happened once I added more ascorbic acid/CuSO4. Didn't make much sense to me at the time. This may actually explain it. I guess someone (me?) ought to measure the pH during the course of the reaction :)

The_Davster - 13-10-2004 at 19:07

Well, since you and I are the only ones who have done this reaction(so far), and I currently only possess broad scale pH paper...;)

Anyway, I weighed my yields from the two tries at this experiment. The first trial yielded 0.2g of a dark red precipitate, the second trial yielded 0.6g of identically colored precipitate.

The precipitated powder was added to dilute HCl to attempt to remove any oxide coating(and I wanted it shiny:D). This was unsuccesful, the precipitate retained its origional dark color.

I believe I can attribute such low yields to the hippie brand of vitamin C.
Quote from the bottle
"Each tablet contains;
Vitamin c.........500mg
(ascorbic acid, sodium ascorbate)
Prepared in a specially formulated nutrient base of natural fruit pulp extracts, acerola, rosehips and the complete vitamin C complex of rutin, hesperidin and citrus bioflavonoids. Sweetened with Crystallose(tm)( a complex of phenylalanine, glucose, and cellulose)."
So as I said before, a hippie brand. The tablets do taste good, in fact, having the bottle in front of me while I quoted caused me to eat a couple:P. I believe the fact that this vit C is ascorbic acid and sodium ascorbate may have reduced my yields due to not all 500mg being ascorbic acid which is necesary for this reaction.
Tomorow I will try to obtain a bottle of (relatively)pure ascorbic acid and try that instead of this brand.

PS. I currently am evaporating some of the green filtrate at room temperature in the hopes that I obtain some sort of green crystalls. I realize it will still be contaminated with the other crud in the vit C tablets, unfortunatly.

[Edited on 14-10-2004 by rogue chemist]

chemoleo - 13-10-2004 at 19:15

I guess I should have said that when I tried Na-ascorbate (analytical) the other day, it didn't work. I haven't thought about it until this moment... in fact the picture above is the attempt with ascorbic acid rather than the ascorbate.
Now I wonder why. It formed the green solution nonetheless, but yet no precipitate. To a proper organic chemist, this should sort out the mechanism I guess?

PS I guess that pretty much negates what I said above, i.e. that the pH might prevent the precipitation of further copper. Because then the salt form of ascorbic acid should facilitate the reaction?

[Edited on 14-10-2004 by chemoleo]

Mephisto - 14-10-2004 at 03:05

L-Ascorbic acid is a reducing agent due to its enediol structure (two OH-groups with a C=C bond in between). In form of the Na-ascorbate there is a Na<sup>+</sup> instead of the H-atom in one OH-group. The resulting negative charge on the O-atom is delocalized in a p-bonding system with the keto-group (R'-O<sup>-</sup> and R"=O 'mesomers' to R'=O and R"-O<sup>-</sup> ==> delocalized electron). By this, the enediol structure may lose its reducing effect and Na-ascorbate won't reduce Cu-ions.

That's just a guess and I have got no references for this.

Theoretic - 15-10-2004 at 06:27

"Very fine copper powder can be produced by heating black copper oxide in a stream of hydrogen. The resulting powder is so fine that it produces a perfect shiny metallic surface on pressing it with the blade of a knife."

Maybe it actually kind of welded to the surface of the knife?

I was doing a reduction of CuSO4 with Al foil some months ago and found the copper produced was dark red (even if it has never hit air). Maybe it's because it was extremely porous?
It might be interesting sintering the powder - could compact copper be made of it? Might be a useful way of preparing small copper equipment parts.

vulture - 15-10-2004 at 06:59

Pure copper is dark red. It resembles terracotta somewhat.

chemoleo - 15-10-2004 at 07:20

And precipitated copper is somewhat brighter ... at least as long as one uses pure reagents! (i.e. no citrus flavours etc :P)

Tacho - 15-10-2004 at 09:38

Originally posted by chemoleo
Apart from that... I rubbed the powder onto a surface, and it actually became shiney, <snip>

Chemoleo, does the surface become conductive?

chemoleo - 15-10-2004 at 09:49

Well, as I said above, I noticed a massive increase in conductivity, if the Cu powder was present. So yes, it is. However, the oxide layer that forms after a few minutes/hours on air prevents conductivity. That is why the powder that I left on air to dry didn't conduct. Only the freshly HCl-washed powder conducted (which was still slightly wet with distilled H2O, while the very same distilled H2O had a conductivity up to 500 kOhm).
That's why I recommended to keep the powder wet (i.e. in EtOH) until before use, and then to dry it rapidly.
Or, use the PVP from the chemical abstract paper above to protect it (it's mentioned to be employed as a protectant).

IPN - 15-10-2004 at 11:05

As I had never made copper powder in any way I thought I'd give it a try.

I took ~80ml of strong CuCl2 solution (I don't know how strong it was as I was evaporating it down to get the CuCl2) and added few pieces of Al-foil. It boiled vigorously and some dark red powder witch was mostly very fluffy formed on the bottom. I then took a piece of pure Al-bar ( ~40cm x 5cm x 0.4cm) and supported it so that it would sink into the solution as it dissolves.
This method as I noted does not make very fine powder. But it's ok until I get some ascorbic acid as I think I’ll clean this batch with some HCl and close it into a glass vial so it will remain nice and shiny in my shelf.

Reverend Necroticus Rex - 16-10-2004 at 06:19

Would formic acid perhaps prove sufficient for reduction? just for the reason that would entail no extra expense to purchase ascorbid acid, as I have a bottle ofr 40% formic lying around much untouched:)

Theoretic - 16-10-2004 at 10:53

Yes, formic acid will do.

chemoleo - 16-10-2004 at 17:34

Now, theoretic, where did you get that idea from ? Although formic acid has the HC=O group, it doesn't mean it has as strong a reducing character as ascorbic acid.
You may know the Fehling test, where aldehyde sugars are employed to produce Cu2O. Note, Cu2O. So why would formic acid suddenly produce Cu?
Obviously the reduction with ascorbic acid isn't really obvious, else these fellows wouldn't have deemed it worth publishing in 2003. Not any reducing group does. It looks like ascorbic acid is quite unique in this ability to reduce Cu2+ to Cu!

My guess is nothing will happen at all, if you are lucky you might get Cu-formiate.

S.C. Wack - 16-10-2004 at 17:52

Rex- Not with CuSO4. This anhydrous salt can be used to dry HCOOH. chemoleo is of course correct.

Tacho - 17-10-2004 at 02:50


Formic acid does NOT work.
Hydroquinone does NOT work.
Sodium Bissulfite produces a precipitate that's NOT Cu.

Ascorbic acid WORKS. It takes a bit long though. Some copper deposits on the glass surface (possible mirror?) after 1/2 hour. Overnight a deposit forms.

A thought: I have doubts if this is indeed much more fine than the powder created by Zn powder in CuSO4. The vit C powder precipitates quickly, as chemoleo pointed out. The copper powders I have made from Zn sometimes take long to precipitate. Maybe the crystals growing on Zn are more "fractal" than the ones reduced by vit C.

But I haven't done any tests to back me up on this, just an impression.

Esplosivo - 17-10-2004 at 03:01

I tried the precipitation of Cu using Ascorbic acid and CuSO4. The Cu ppt formed is very fine, but nearly as fine as that which is precipitated by the rxn between Zn ad CuSO4. Could the size of the particles of Cu depend on the concentrations of the solutions being used? I used a somewhat dilute soln. of Ascorbic acid.

The_Davster - 17-10-2004 at 16:20

Well, I went out and bought the purest vit C I could find OTC. The contaminants are methylcellulose, microcrystalline cellulose, talc, magnesium stearate.

The yield is much higher, and shinier, than with the aforementioned hippie brand I was previously using but....

The copper was bound by one or more of the contaminants, so all the copper appears to be plasticised. It looks, and feels like chewing gum.

Sorry, for the crappy pics, my digital camera is very bad. In the pictures the copper is actually brighter than it appears.

EDIT: I am just going to order some ascorbic acid instead of dealing with these contaminants:mad:

copper plasticised1.JPG - 236kB

The_Davster - 17-10-2004 at 16:21

Sorry about the double post, just wanted to show another picture.

copper plasticised 2.JPG - 231kB

chemoleo - 17-10-2004 at 16:50

Rogue - why don't you use lemon juice? It might be cleaner than the stuff you are working with!!:D

Anyway - on the degree of fineness- tacho/esplosivo you may have noticed that there is very fine stuff, and the more grainy variant. The more grainy variant can, when dry, be crushed, to be very fine. This is when I got a piece of paper shiny with copper dust :)
Also - I think the overall surface area of the precipitated powder has to be greater by definition than Zn-reduced Cu, simply because reduction is not localised in space to the reductant particles (i.e. Zn). Instead, reduction occurs everwhere at once, i.e. in solution.
What I rather suspect is that the Cu conglommerates, aggregates or whatever, in a loose kind of fashion. This, as it seems from the Chemical Abstracts, above, can be prevented with various agents, here a polymer. Might be a nice idea to do this whole thing with polyethyleneglycol, or even glycerol, to see whether the coarsity of the Cu changes. However, here the hydroxyls might interfere.
Hell, I might give it a go tomorrow :)

Edit: Rogue, the picture is above 100 kb, so it might be deleted. But I don't know whether that's true for pictures in the upload directory only.

[Edited on 18-10-2004 by chemoleo]

FrankRizzo - 17-10-2004 at 18:22

What about using dissolved dextrin to create a smaller precipitate? It works for lead azide, why not here?

Tacho - 18-10-2004 at 03:08

Originally posted by chemoleo
Also - I think the overall surface area of the precipitated powder has to be greater by definition than Zn-reduced Cu, simply because reduction is not localised in space to the reductant particles (i.e. Zn). Instead, reduction occurs everwhere at once, i.e. in solution.


Your work is great, as the document published only in 2003 proves. It may be a very specific property of vit C to reduce copper like that.

However, just for the sake of discussion, I disagree with your statement above. It's about crystal growth. They may grow on the Zn surface or on nucleating points (!?). One doesn't have to be larger by definition.

vulture - 18-10-2004 at 07:57

A solid dispersion can NEVER achieve the "surface area" and reactivity of a dissolved agent.

With vitamine C you are working with single molecules of vitamine C, whilst the Zn powder consists of clusters of more than thousand atoms with relatively little surface.

Tacho - 18-10-2004 at 08:30

The fact is that the precipitated copper is not in a colloidal state. Its a powder, so we are talking of crystal growth on a nucleating point.

Wait!... what am I doing? BlahBlahBlahing! I'm sorry Descartes!

I'll experiment tonight with Zn powder and VitC and measure time of precipitation.

mick - 18-10-2004 at 13:01

I thought the realy active stuff was made by amalgamating the metal with aluminium and then treating it with alkali.
I have used Raney nickel but there is also Raney cobalt and Raney copper

Edite typo

[Edited on 18-10-2004 by mick]

chemoleo - 18-10-2004 at 15:16

Tacho, I should think that we are dealing with amorpheus aggregates of Cu here, rather than 'crysallised' Cu. I.e. one Cu atom binds randomly to another, in an unordered fashion. This explains why the powder can be crushed easily. While, when it is derived from Zn powder, the particles will at least have the size of the Zn powder, whereby the reaction allows for better alignment of the Cu atoms - so if anything - Zn-derived Cu is more crystalline than Cu precipitated straight from solution.
Anyway, I fully agree with vulture on this.

Nice idea about the dextrin btw. Shame I haven't got any here :(

Mick, of course one can make Cu that way too - but this is bound to be not quite as simple as with Zn/Vit C. But ultimately that should yield extremely fine metal powder.

axehandle - 18-10-2004 at 16:49


I have used Raney nickel but there is also Raney cobalt and Raney copper

No worry mate, we'll through another shrimp on the barbie!

FrankRizzo - 18-10-2004 at 17:11


Do you have any corn starch in your kitchen? You can turn it into dextrin fairly easily by baking in an oven @425F until golden brown.

[Edited on 19-10-2004 by FrankRizzo]

axehandle - 18-10-2004 at 17:16

Indeed. I turned 500g of corn starch into dextring by keeping it at 200 degrees C for 2 hours

EDIT: Or I was drunk when I did that: I feel... uncertain.

[Edited on 2004-10-19 by axehandle]

Tacho - 19-10-2004 at 03:47

The vit C copper is much better looking than tha Zn copper. No matter how much you wash the Zn copper with HCl or H2SO4. BTW, concentrated HCl (fuming) made both coppers disappear (without fizzing->no zinc)! This surprised me.

The vit C copper looks like copper. The Zn copper looks like grey-redish dirt.

The vit C copper seems to have coarser grain and finer grain, I mean it's "graduated" from fine to coarse. The finer grain is quite finer than the one obtained by Zn, it remains "floating" long after the Zn copper has settled. The coarser grains seems to be still finer than Zn grain, but not much.

Chemoleo and Vulture:

My reasoning was based on the silver crystals that grow on a piece of copper in silver nitrate solution (the classic "silver tree" demonstration). They are quite smaller than the copper piece. A similar mechanism could produce smaller copper crystals than the Zn grain. That doen't seem to be the case though. Maybe, as chemoleo pointed out, they are amorphous.

Anyway, I'm very interested in this. This copper powder seems quite different from the ones I have seen (from Zn and Al). I hope I can find a cheap source of rather pure ascorbic acid to do some further testing. It looks so conductive and reactive... You should try to make some, axehandle, I'm sure it will make your brain start ticking.

Chemoleo, what was the reasoning behind you choosing ascorbic acid as a reducing agent? Maybe we could find another one.

Dextrin huh?

[Edited on 19-10-2004 by Tacho]

Organikum - 19-10-2004 at 05:57

Copper precipitated on steelwool, the kind I posted a pic in the benzene-thread looks very nice, copperlike. It forms these "trees" which can easily brushed up and yield finely divided copper of the color it should look like.
Dont forget to add some NaCl to get this going.

Copperoxide can be reduced to elemental copper by hot ammonia vapors .

neutrino - 19-10-2004 at 13:16

Originally posted by Tacho
concentrated HCl (fuming) made both coppers disappear (without fizzing->no zinc)! This surprised me.

Copper shouldn't be attacked by HCl, as copper is lower on the reactivity series than hydrogen. Could there have been an oxidizer in your acid?

On the topic of acquisioton of vit C, I've found it in a vitamin store, but I can't seem to firgure out its purity. 8oz has a strength of 1065g??? But 8oz = 226.8g. Does anyone know what to make of this?

edit: The link function doesn't seem to eb working right, so here's the address:;$sessio...

[Edited on 19-10-2004 by neutrino]

Ium - 19-10-2004 at 17:12

What you are after is pharmaceutical grade Non-Flavoured Ascorbic Acid. Avoid buffered Calcium or Sodium derivatives. Check on the container or package for a reference to purity or active constituents.
While searching for the same particular brand you mention, I came across that errror. "Strength: 1065 g" is supposed to mean the serving size in mg. They also mention 1/4 teaspoon as the suggested serving size for a more general measurement. Both measurements are approximately the same amount. The refernce to "8 oz." is the weight of the full amount of Ascorbic Acid.

[Edited on 20-10-2004 by Ium]

[Edited on 20-10-2004 by Ium]

neutrino - 19-10-2004 at 19:14

I can't believe I didn't see that, thanks for pointing it out. I don't see any mention of anything other than vit C, so it should be reasonable pure unless they 'cut' it with some inert stuff.

Mr. Wizard - 19-10-2004 at 20:27

"Ascorbic acid WORKS. It takes a bit long though. Some copper deposits on the glass surface (possible mirror?) after 1/2 hour. Overnight a deposit forms." quote from Tacho

I just repeated the Copper Sulfate / Ascorbic Acid experiment and got a very fine reddish powder that acts just like copper, a strong magnet slowed it's fall down the side of the test tube, showing it conducts.

It would be very nice to make a copper mirror. Google showed a site that gave a copper mirror with hydrazine hydrate and copper acetate. It was done in a test tube that was heated while the copper was reduced by the hydrazine. Maybe if the glass was cleaned as the glass on a telescope mirror, with nitric acid, after being degreased, it would form a mirror? Has anyone seen a copper mirror form?

[Edited on 20-10-2004 by Mr. Wizard]

[Edited on 20-10-2004 by Mr. Wizard]

Theoretic - 20-10-2004 at 00:06

neutrino: probably the amount of air dissolved in the acid worked as an oxidizer.

Tacho - 20-10-2004 at 04:28

Copper mirror recipes

Tacho - 24-10-2004 at 15:33

I got some reagent grade ascorbic acid and did some tests. I noticed a few things: 1) heat strongly accelerates the reaction but, using heat, the precipitate has that “plasticized” look that rogue chemist displayed in his pictures. Since there is only pure reactants, I assume the copper agglomerates by itself. 2) I could get no mirror, even treating the glass with Sn ions. 3) Just for the record: The green solution left after the copper precipitates makes a good plating solution, I got a beautiful copper plating on graphite on my very first try. Maybe I was lucky, but next time I want to do copper plating, I’ll try this solution (I saved it).

I’m a bit disappointed, because I was expecting quantitative yields using pure reagents, but the reaction is… my English fails… Temperamental?

Any other use for 40g of ascorbic acid? The label says "laboratory use" and mentions heavy metal residue, so I don't think I should use it to cure my next cold...

chemoleo - 24-10-2004 at 16:15

Yes I agree the reaction is temperamental. Boiling does the job very quickly, that's how I always did it. I don't know why you get some 'plasticed' look, all I got with reagent grade Vit C/CuSO4 was fine grains as in the pic above. These don't stick to each other, and can be easily crushed to dust.
What confuses me is the quantitative yield, too. How can they, in that paper, claim 95% of the theoretical yield if you still got enough left to do some plating? By the way, I tried the reduction in the presence of 30% polyethyleneglycol (6000 average MW), but the powder was visibly not finer or coarser- but then I boiled it too. What was interesting is that the PEG is nicely soluble on its own, but as soon as I added the CuSO4 to this, the CuSO4 seemed to form an emulsion, i.e. as a dark blue liquid! I have the impression that a variety of inorganic salts could be purified/concentrated that way, i.e. add the salt in question to a conc. solution of PEG, shake/mix for a while, let it settle, and decant off the PEG/water, to be left with a concentrated solution of the salt!

Anyway, i will try the experiment with PVP at some point, as I managed to obtain some.

What to do with the remaining Vit C? Comon, the obvious! Test it with Ni, Co, Mn, etc etc salts! Who knows, maybe this works on other metals too!

PS interesting page on the copper mirror - shame only it requires some very expensive metal salts.

[Edited on 25-10-2004 by chemoleo]

Tacho - 24-10-2004 at 16:39

About quantitative yields: I didn't do any carefull measurement, but the ammount of copper precipitated is clearly not proportional to vit C added.

About the use for leftover vit C: Yes! Of course! Obvious. How embarassing...

Condutive paint powder...

Tacho - 26-10-2004 at 03:43

I tested the vit C in 3 metal salt solutions: Circuit etching ferric chloride, nickel chloride and silver nitrate. Nothing happened to nickel; ferric went from yellow-brown to green and, with extra vitamin, to a clear colorless solution (nice way to remove those ferric chloride stains). No precipitates.

Silver nitrate, OTOH, gave very interesting results: a fine gray sludge of metallic silver powder. Unfortunately, this powder did more than precipitate, it coagulated (sorry, but this word is perfect to describe what happens) in a blob in the same way that the “silver tree” crystals do.

Even so, this powder is very conductive, the best conductive powder I’ve seen and it’s certainly the thing to make homebrew conductive paint. If rubbed on a gypsum surface with my finger, it made a nice, very conductive surface (0 ohms) even on my fingertip.

I tried different solvents for the reaction, mixes of isopropyl, methyl, ethyl alcohols, acetic acid and ethylene glycol. No improvement.

I also tried different bonding glues: PVA, sodium silicate, acrylic latex and some alcohol soluble varnish. So far, PVA glue gave the best results, but far from a perfect paint.

I presume that, to obtain a good homemade conductive paint, one of this three routes should be used:

1- Carefully pick the coagulum and put it on filter paper to dry. After it’s dry, crunch it to a fine powder and mix with some bonding agent. I think this is the best route, silver does not oxidize and AgS is condutive, so the powder can be handled and stored safely.

2- Mix something with the reactants that will keep the silver in emulsion; an emulsifying agent or just a different solvent.

3- Do the reaction in a rotating flask with some coarse sand in it, so that the silver never has the chance to coagulate.

This is a very interesting result. Thanks chemoleo. I knew this vit C thing would give me something new to my list of voodoo tricks.

Mephisto - 27-10-2004 at 15:36

Ascorbic acid works for all reactions with a higher redox potential than +0.166 V, since ascorbic acid itself has this redox potential (according to the Merck Index XIII).

Cu<sup>2+</sup> + 2e<sup>-</sup> → Cu EΘ = +0.34 V

Ag<sup>+</sup> + e<sup>-</sup> → Ag EΘ = +0.80 V

Fe<sup>3+</sup> + 3e<sup>-</sup> → Fe EΘ = -0.04 V

Fe<sup>3+</sup> + e<sup>-</sup> → Fe<sup>2+</sup> (yellow to green) EΘ = +0.77 V

Ni<sup>2+</sup> + 2e<sup>-</sup> → Ni EΘ = -0.23 V

For example Hg and Au from their salts will work, too.

Many thanks to Tacho, for verifying this in such scientific manner!

Tacho - 28-10-2004 at 04:51

What books don’t say:

Electroplating with some quality is difficult. To obtain solid metallic deposits demand know-how and problematic chemicals.

I could never obtain a good NICKEL deposit using nickel chloride as electrolyte... until I tried the leftover solution of nickel chloride plus ascorbic acid.

I obtained a beautiful solid metal deposit on graphite and on copper cathodes. Almost shiny and easily polishable. I tried different concentrations of the chloride with equal results.

I used very low voltage. I didn’t check voltages or currents, and won’t be able to check them today or in the next few days, due to personal problems. Further tests will have to wait a while.

I believe THIS is news.

Who was looking for a nickel crucible?

axehandle - 28-10-2004 at 05:33

If any of you one wants some food grade pure ascorbic acid to do great stuff with and isn't scared of customs taking an initial interest in the white powder in that strange overseas letter, PM me. I've got close to 4kg left, of which I need only a small part.

Ium - 28-10-2004 at 05:48

Ascorbic acid has found use in electroplating solutions for various metals for some time now. However I am unable to find much information on the exact role of the acid in plating baths. Most articles seem to suggest that it is used because of it's strong reducing effect on metallic salts and so aiding the plating process by creating salts more easily freed of the metal.

It may help for other members to search for additional references to ascorbic acid's use in electroplating, as so far I have had little success.

Edit: Patents have some juicy info.


In this technique, ascorbic acid is used in the presence of an electrolytic aid. The purpose of adding ascorbic acid is to (a) control the pH of electrolyzed water of cathode side and (b) remove the free chlorine in electrolyzed water of anode side

[Edited on 28-10-2004 by Ium]

Low yield

Mephisto - 28-10-2004 at 21:13

A solution of 43.2 g (≈ 245 mmol) ascorbic acid in 140 ml 50 °C warm, distilled water was added to a second solution of 25 g (≈ 100 mmol) copper(II)-sulphate pentahydrate in 80 ml 50 °C warm, distilled water.

Before mixing.

After 30 minutes reaction time (the mix was stirred in this time).

The formed copper powder was filtered off and washed two times with distilled water and once with spirit to give 1.8 g (dry) copper powder. This is 28% of the theoretically possible yield. (Note: scale inaccuracy ±0.2 g)

Maybe higher concentrated solutions of CuSO4 and ascorbic acid were yielding more copper powder. Since I need only some grams Cu-powder to catalyze a reaction, the yield isn't very important for me.

I found easily a Vit C product that consists of pure ascorbic acid, without any contaminants. 100 g costs 1,49 € but it could be somewhere cheaper, since I bought it in the first drugstore, which was on my way.

BTW: German version of this experiment on

[Edited on 29-10-2004 by Mephisto]

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[Edited on 7/22/13 by bfesser]

garage chemist - 30-10-2004 at 13:11

LambdaSyn is a great site! I just read some of the other syntheses and they're GREAT! There are some syntheses there which are completely new to me.
I especially like the Ammonium chlorate synthesis from KClO3, (NH4)2SO4 and Ethanol
and the Ethyl nitrate synthesis.

Mephisto - 30-10-2004 at 13:40

Nice to hear that. After a period of stagnation I will work more on the website. Actually I got 102 synthesis-texts in the Rich Text Format on my HDD (many of them with photos), within the next weeks I try to convert a large part of them into HTML an put them online. So, I think, it's worth to visit the website once a week to check out the updates.

TheBear - 14-11-2004 at 12:52

Just a thought about the conductivity tests: Cu2O is a semiconductor with photoelectric properties.

chloric1 - 6-1-2005 at 15:24

Originally posted by kevinlimse
Is there a method for Aluminium powder?

If a am not mistaken aluminum could be isolated from an electrolyte based on DMSO as a solvent. I am not sure what aluminum compound though. Besides, by the time you work it out, the immense investment in reagents would far outweigh what puny yeilds you could possibly obtain. :(

neutrino - 6-1-2005 at 15:44

The point isn’t obtaining large amounts of superpure powder; it’s knowing how to make anything potentially useful at all.

Theoretic - 7-1-2005 at 05:40

"Just a thought about the conductivity tests: Cu2O is a semiconductor with photoelectric properties."
Maybe a Cu2O mirror can be created on a suitable surface with slow depositing (low concentration), then playing around with the photosensitive element thus made? Integrate it into an electric circuit with an alarm in it, so the first rays of the sun at sunrise wake you up and you start your day nice and early? :D

Archimede - 21-5-2005 at 18:46

As an electrician I can find fair amounts of scrap clean copper. Can I just dissolve it in solution using acetic acid/NaCl for then use it with electrolisis to plate stuff?
Same apply for Zinc scrap pieces.

Is there a way to process the vinegar to get a higher concentration of acetic acid?

I am new to this subject so throw me a bone pls.

neutrino - 21-5-2005 at 19:44

Not to be rude, but questions like this are better asked in the 'beginnings' section. Ask there and it shall be answered.

CD-ROM-LAUFWERK - 22-5-2005 at 02:32

a methode of producing very very very fine Cu powder is to form coppernitrate, leting the solution react whit NaOH-pellets and reducing the CuO whit glucose!
this power is much finer as that on the foto from chemoleo!
u cant see a singl particle!
it takes hours to siting down in water!
that stuff is the finest powder i'v ever seen, no Al power is finer!

chemoleo - 22-5-2005 at 05:43

And you are sure it isn't Cu2O? Try dissolving it in HCl.
It'd surprise me if the reduction potential of glucose is as high as that of ascorbic acid.

CD-ROM-LAUFWERK - 22-5-2005 at 15:28

i fergot to say that u have 2 cook the mixture of NaOH, CuO and glucose for a few minutes (~5min), i did that in the microwave, dont know if that mater...
i will test it later, no problem!

The_Davster - 4-9-2005 at 12:49

I have been doing some more experiments with the reaction between ascorbic acid and CuSO4 recently. I have pure acsorbic acid now, so no more problems with purity:P.
Using 5g vitC and 10g CuSO4, a dark precipitate of copper powder is produced(much more than in my previous experiments). However it takes a long time to settle. The green solution left over is what I have mainly been experimenting with.
1. Green solution +NaOH results in an orange/yellow precipitate.
2. Green solution + HNO3 results in a solution the normal Cu2+ colour(light blue)
3. Solution produced in 2 +NaOH results in the same precipitate as in 1
4. Green solution + Zn results in a black precipitate and bubbling (the most confusing result so far)
5. Green solution +Na2CO3 results in a green precipitate

[Edited on 4-9-2005 by rogue chemist]

12AX7 - 4-9-2005 at 18:10

1. Cu2O.
2. NO3- oxidizes to Cu++.
3. This suggests it isn't Cu(I) leftover in 1. Maybe the ppt is a complex?
4. How acidic is the solution, and pure the Zn? Most of the time when I'm dissolving Zn it leaves a black colloid, probably copper and other insoluble constituents of the 90% alloy.
5. Malachite?


The_Davster - 4-9-2005 at 18:23

It is pure zinc powder and I have never gotten any black matter while dissolving zinc. Also the hydroxide ppt looks nothing like Cu2O, way too yellow.

12AX7 - 4-9-2005 at 20:14

You sure? That's what I thought first time I made Cu2O. It starts almost yellow then runs to orange then red as the crystals grow, if they continue growing.

Ok, pure zinc? No idea on that!


woelen - 5-9-2005 at 01:02

Using 5g vitC and 10g CuSO4, a dark precipitate of copper powder is produced(much more than in my previous experiments). However it takes a long time to settle.

Are you sure it is copper and not some copper (I) compound (e.g. impure Cu2O)? Add some of this dark stuff to dilute HCl. If it is copper it will not dissolve, otherwise it may become white and eventually will dissolve.

The green solution left over is what I have mainly been experimenting with.
1. Green solution +NaOH results in an orange/yellow precipitate.
2. Green solution + HNO3 results in a solution the normal Cu2+ colour(light blue)
3. Solution produced in 2 +NaOH results in the same precipitate as in 1
4. Green solution + Zn results in a black precipitate and bubbling (the most confusing result so far)
5. Green solution +Na2CO3 results in a green precipitate

1. Indeed Cu2O, possibly in a hydrous form.

2. The blue color is due to Cu(2+), but certainly not all copper is oxidized to the +2 oxidation state. I doubt whether the HNO3 is oxidizing it, unless it is very concentrated. Usually oxygen from the air oxidizes the stuff. Copper (I) is colorless and copper (II) is blue. In nitric acid and sulphuric acid, the mix also is blue. In hydrochloric acid, however, you'll get deep brown, almost black complexes of mixed valency (cations, with multiple copper ions in it, and with an oxidation state both of +1 and +2 in a single cation). I have done quite some research on copper chemistry and put all my findings on my website, see this link
Another observation I have made is that copper (I) complexes seem to reverse react in strongly acidic environments, which are not coordinating to copper (I), such as nitric acid or sulphuric acid solutions. Copper (II) ions and ascorbic acid can coexist pefectly at very low pH.

3. This can easily be explained. The blue solution contains copper (II) and free ascorbic acid. On addition of excess NaOH, this reacts again to yellow cuprous oxide. The addition of alkali strongly shifts the redox potential for copper (II) to copper (I). This is why at low pH the blue copper (II) remains in solution, while at high pH the copper (II) is reduced by the ascorbate, formed in the alkaline solution.

4. Not easy to explain. The bubbling most likely is due to some acidity, but the black stuff I cannot easily explain. My guess would be that this is very impure copper metal, formed at the zinc metal surface. Zinc may also reduce other stuff in solution and this may give rise to the black stuff.

5. What kind of green color? Is it a bright cyan green color, or a darker moss-green or brown-green color. The cyan compound is basic copper carbonate, if the precipitate is dark green like some moss plants, then it is copper carbonate, mixed with copper (I) compounds. It is remarkable what the combination of copper (I) and copper (II) in a single compound does with its color. It impairs a very dark green/brown color to such compounds.

I also have done quite some experimenting with copper (I) oxide. I found a relation between the speed of formation and the color. Fast formation --> bright yellow, slow formation --> brick red. All colors from bright yellow through several oranges to brick red can be made. See the following two pictures I made of a copper (I) solution and the two extremes of the copper (I) oxides.

Copper (I) solution in dilute ammonia, which is oxidized by oxygen from air, forming the deep blue tetrammine complex.

Yellow Cu2O:

Red Cu2O:

What I found is that making pure copper metal from solutions is not as straightforward as I liked. One process, which gives very fine pure copper metal is preparing such a solution in ammonia, as shown in the first picture and then adding a solution of sodium dithionite. Then pure copper metal is formed. First the liquid becomes blue when viewed through transmitted light and red when viewed with reflected light. Later it becomes turbid and metallic copper settles at the bottom as a very fine brown powder.

Another way is to dissolve copper (II) sulfate or copper (II) chloride in dilute hydrochloric acid and adding sodium sulfite and heating a little. This gives rise to larger crystals of copper metal, but this procedure seems to be somewhat unreliable. I have tried multiple times and not always does it produce copper metal.

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[Edited on 7/22/13 by bfesser]

neutrino - 5-9-2005 at 03:40

2. I’m guessing that one of the byproducts of the initial reaction is some kind of ligand that complexes unreacted Cu<sup>2+</sup> ions, forming that green color. Adding a strong acid causes the free H<sup>+</sup> ions to displace the copper ions from their complex, leaving blue Cu<sup>2+</sup> ions floating around.

The_Davster - 5-9-2005 at 05:13

1. After looking at your pic wolean, the result is indeed Cu2O.
2. It was 70% HNO3 added to to the green solution
4. I will do more tests later today
5. Olive green

It indeed is copper powder produced as I washed it with 31% HCl without any dissolving. We discussed this a few pages back I believe, and Chemoleo found it conductive as well.

Bright metallic concentrate

bluebriant - 29-11-2008 at 07:47

I have had some success using the CUSO4 + Ammonia + ascobic acid method to precipetate copper powder. However, the powder ends up looking black. I would like to precipitate a bright metallic powder. Is this possible?

Jor - 29-11-2008 at 17:50

I have very fine lab-grade copper powder. It's BLACK as well!
It actually stains my fingers when I touch it.

Is this even more fine than the red superfine copper powder, you guys are talking about?

bluebriant - 30-11-2008 at 04:33

Yes I think it is copper, but is so fine so as not to be able to reflect the wavlength of the copper powder. Perhaps the precipitate needs to be produce more slowly...?

[Edited on 30-11-2008 by bluebriant]

Nerro - 30-11-2008 at 05:43

An experiment I did years go involved dissolving Mg shavings in ascorbic acid. The resulting substance dissolved completely in water and was much more reactive than just ascorbic acid or Mg shavings. Addition of this solution to a concentrated solution of CuSO4 in water reacted immediately to form a yellow-brown precipitate which then slowly converted into copper precipitate. This copper was exceedingly fine and once dry it got into absolutely everything!

bluebriant - 30-11-2008 at 06:21

Interesting...Unfortunately I think I need to produce copper in a more cystalline form (rather than amorphous). This will probably need some sort of cyrstalline promotor of sorts, e.g. a sodium silicate matrix?

12AX7 - 1-12-2008 at 03:36

You'll get metallic crystals from cementation (e.g. powdered zinc). Good old fashioned sintering of the colloidial powder (under hydrogen at maybe >600°C) will of course do something, although it may just stick together as a solid lump. :P

I've got fine copper before, but I can't quite remember why. Maybe it was Cu2O I dissolved in H2SO4 (which disproportionates to CuSO4(aq) and Cu(s)).


Ketone - 4-12-2008 at 14:35

That's pretty cool. Nanozized Cu.. Nice. Never even thought of reducing copper(II) salts with ascorbic acid..
(Seems like one often overlooks those conceptually simple but useful reactions for some reason..)
This could be quite useful and seems v. easy if done right and pure ascorbic acid is used..
Luckily I can get it undiluted it at the grocery-store for a few dimes here. :D

Though for those that can't...
Shouldn't Oxalic acid work too? I think it's also more widely available as a pure chemical than ascorbic acid, as it's sold as a rust remover.

Amos - 24-6-2014 at 20:04

I'd really like to revive this thread in order to explore what was in that lovely green solution. Incidentally I made copper powder in this same fashion after hearing about it somewhere else, and now I have loads of the solution left over. Did Davster ever finish his tests? Whether he did or didn't, it would be nice if we at least had an equation relating the reduction of copper sulfate by this method; it's a bit beyond me. I'm quite curious as to where those sulfate ions went.

Praxichys - 25-6-2014 at 10:34

Dry nano-scale copper powder is easily made from the thermal decomposition of copper oxalate as well. This might prove to be more convenient.

Amos - 25-6-2014 at 10:56

Oh yes, the oxalate method is great for metal powders, they're even so fine they're pyrophoric sometimes, as in the case of iron. I'm probably going to make a new thread so that this question of the green leftover solution from the Ascorbic acid method gets some proper attention.

opendna - 5-7-2014 at 14:47

I'm going to hypothesize a formula which seems to be consistent with my observations. Open to correction.

CuSO4*5H20 + C6H8O6 = Cu+ + H2SO4 + C6H6O6 + H2O
copper sulfate + ascorbic acid = copper + sulfuric acid + dehydroascorbic acid + water

(cross-posted here)